A
MANUAL
OF
INORGANIC CHEMISTRY,
ARRANGED TO FACILITATE THE
EXPERIMENTAL DENONSTRATION
OF THE
FACTS AND  PRINCIPLES OF THE SCIENCE.
BY
CHARLES W. —EIOT,           FRANK H. STORER,
PROFESSOR OF ANALYTICAL CHEMISTRY AND PROFESSOR OF GENERAL AND INDUSTRIEt
METALLURGY,                 CHEMISTRY,
IN THE MASSACHUSETTS INSTITUTE OF TECHNOLOGY.
FIFTH THOUSAND.
NEW YORK:
IVISON, BLAKEMAN, TAYLOR, & COMPANY,
138 & 140 GRAND STREET.
CHICAGO: 133 & 135 STATE STREET.
1876.




Entered according to Act of Congress in the year 1866, by
F. H. STORER AND C. W. ELIOT,
In the Clerk's Office of the District Court of the Distriot of Massachusetts.




THE AUTHORS
INSCRTBR THIS BOOK TO THEIR TEACIIEI IN CHEMIITRY,
PROF. JOSIAH P. COOKE,
OF IIARVARD COLLEGE,
IN TOKEN OF
GRATITUDE AND FRIENDSHIP.








PREFACE TO THE FIRST EDITION.
IN preparing this manual, it has been the authors' object to
facilitate the teaching of chemistry by the experimental and
inductive method. The book will enable the careful student to
acquaint himself with the main facts and principles of chemistry,
through the attentive use of his own perceptive faculties, by a
process not unlike that by which these facts and principles were
first established. The authors believe that the study of a science
of observation ought to develop and discipline the observing
faculties, and that such a study fails of its true end if it become
a mere exercise of the memory.
The minute instructions, given in the descriptions of experiments and printed in the smaller type, are intended to enable the
student to see, smell, and touch for himself; these detailed
descriptions are meant for laboratory use. In order to mark as
clearly as possible the distinction between chemistry and chemical
manipulation, the necessary instructions on the latter subject
have been put in an Appendix. In cases in which it is impossible
for every student to experiment for himself, the authors hope that
this manual will make it easy for the teacher, even if he be not a
professional chemist, to exhibit to his class, in a familiar and
inexpensive manner, experiments enough to supply ocular demonstration of the leading facts and generalizations of the science.
Judging from their own experience, the authors venture to hope
that even professional chemists may find it convenient to have




Vi                       PREFACE.
always at hand the details of several hundred experiments,
covering the ground of an extensive course of chemical lectures.
The student of this manual is supposed to be already acquainted
with the rudiments of physics. The chemist must often depend
upon physical properties for his means of characterizing the
numerous substances with which he deals, and he is nearly concerned with the physical properties of gases and vapors; but
chemistry has now so wide a scope and so great an importance as
to deserve to be studied by itself, and not merely as an appendix
to the subject of molecular physics.
Like all elementary text-books, this manual is a mere compilation; it embodies in a somewhat new form previously existing
statements of well-recognized facts and principles which have
become the common stock of the science. There is little original
in the book, except its arrangement and method, in part, and its
general tone. The authors have, of course, drawn largely from
the invaluable compilations made by Gmelin, Otto, and Watts,
and they have also availed themselves freely of the text-books of
Stoeckhardt and Miller and the writings of Hofmann.
The book is not written in the interest of any particular theory
or system of notation, but is intended to exhibit, so far as is possible within the limits proper to an elementary manual, the present
state of the science.
The authors will feel very grateful to any one who will communicate to them errors, detected in using the book, or suggestions
for its improvement.
BOsTON, June 1867.




PREFACE TO THE SECOND EDITION.
THE authors have thoroughly revised this second edition of
their manual. Practical experience in using the book with two
classes in the laboratory, the questions of students, and the
suggestions of friends have enabled them to improve some of the
detailed directions for experiments, and to make a few other
changes and additions calculated to smooth difficulties or supply
defects. Special pains have been taken to make the printing of
this edition as accurate as possible.
BOSTON, December 1867,








TABLE OF CONTENTS.
PAGE
Introduction. —    Subject matter of Chemistry. Chemical change. Analys:s and synthesis. Fact and theory.1-4
Chap. I.-Air.  Atmospheric pressure.  Properties.  Analysis. Air a
mixture. Composition of air.410
Chap. II.-Oxygen. Preparation and properties of oxygen. Oxygen supports combustion. Oxides. Oxidation. Wide diffilsion of oxygen...  11-15
Chap. III.-Nitrogen. Preparation and properties of nitrogen...  16-18
Chap. IV.-Water. Properties of water. The gramme. Specific gravity.
Specific heat. Ice. Steam. Analysis, electrolysis, and synthesis of water. Hydrogen. Atoms and molecules. Molecular hypothesis. Atomic weights. Chemical combination. The chemical force. Water in nature. Distillation. Preparation of pure water. Solution. Solution and chemical combination compared  18-39
Chap. V.-Hydrogen. Preparation and properties of hydrogen. SymbolL.
Diffusion ol gases. Diffusive power and inflammability of hydrogen. Heat
from burning hydrogen. Unit of heat. Oxyhydrngcn blowpipe. Form of gasflames. Explosive mixtures of hydrogen and oxygen. Oxygen burns in hydrogen
as well as hydrogein in oxygen.                                            39-0
Chap. VI. —Compounds of oxygen, hydrogen, and nitrogen. Peroxide of
hydrogen. Nitric acid. One of the meanings of the terms acid and alkaline.
Neutralization. Nitrous oxide. Composition of nitrous oxide. Atomic weight
of nitrogen. Properties of nitrous oxide. Nitric oxide. A test for free oxygen.
The terminations ous and ic. Hyponitric acid. Nitrous acid. Analysis of nitric
acid. Synopsis of the oxides of nitrogen. Law of multiple proportions. Definite
and obscure chemical action. Air a mixture. Anhydrous and hydrated nitric
acid. Oxidizing and reducing agents. Atomic weights and combining weights.
Molecular formulae. Combining weight of nitric acid. Nitric acid reactions.
Nitrogen and hydrogen. Ammonia. Analysis and synthesis of ammonia.
Nascent state. Composition of ammonia. Ammonium. Salts of ammonium.
Sources of ammonia. Ammonia-water. Empirical and rational formula. Dualistic formulae. Typical formulae. Uses of symbolic formula....  50-91
Chap. VI —-Chlorhydric acid. Properties, analysis, and composition of
chlorhydric acid. Atomic weight of chlorine. Synthesis of chlorhydric acid.
Mlanufacture of chlorhydric acid. Practical application of chemical equations.
Chemical affinity. Preparation and uses of chlorhydric acid. Aqua regia.. 91-105




X                                 CONTENTS.
Chap. VIII.-Chlorine. Preparation and properties of chlorine. Chlorine-water. Metals burn in chlorine. Chlorine burns in hydrogen, and hydrogen in chlorine. Combustion in chlorine. Uses of chlorine. How chlorine acts
in bleaching and disinfecting. Action of chlorine on ammonia. Chloride of
nitrogen. Oxides of chlorine. Hypochlorous acid. Chlorous acid. Hypochloric acid. Chloric acid. Perchloric acid.105-119
Chap. IX. —Bromine.  Bromhydric acid. Bromic acid.  Hypobromous
acid. Chloride of bromine. Bromide of nitrogen.119-123
Chap. X. —Iodine. Extraction and properties of iodine. Tests for iodine.
Iodohydric acid. Iodic acid. Periodic acid. Iodide of nitrogen. Chlorides of
iodine. Bromides of iodine. The chlorine group of elements... 123-135
Chap. XI. —Fluorine. Fluorhydric acid. Etching glass. 135-139
Chap. XII. —Ozone and Antozone. Allotropism. Prepz ration and properties of ozone. Tests for ozone. Ozone a disinfectant. C0'one in the atmosphere. Preparation and properties of antozone. The ant ),one cloud. Antozone formed in all processes of combustion. Antozone oxidizes water. Differences between ozone and antozone....139-154
Chap. XII. —Sulphur. Crystallization of sulphu-r. Crystalline structure.
The six systems of crystallization. Sulphur a dimorphous element. Change of
prismatic into octahedral sulphur. Methods of obtaining crystals. Soft sulphur.
Milk of sulphur. Metals burn in sulphur as in oxygen.  Uses of sulphur. Sulphydric acid. Preparation, analysis, and properties of sulphydric acid. Sulphuretted-hydrogen water. Sulphydric acid as a reagent. Ready decomposition
of sulphydric acid. Persulphide of hydrogen. Greek and Latin numeral prefixes.
Oxides of sulphur. Sulphurous acid. Preparation, properties, and composition
of sulphurous acid. Liquid sulphurous acid. Bleaching-power of sulphurous
acid. Oxidation of sulphurous acid. Sllphites. Sulphuric acid.!Manufacture
and properties of sulphuric acid. Its behavior towards water and ice. Hydrates
of sulphuric acid. Sulphates. The terms monobasic, bibasic, &c. Fuming sulphuric acid. Anhydrous sulphuric acid. Preparation and properties of anhydrous sulphuric acid.  Hyposulphurous acid.  Hyposulphites.  Chlorides of
sulphur.154-198
Chap. XIV.,-Selenium and Tellurium. Properties of selenium. Isomorphism. Atomicvolume. Tellurium. Compounds of tellurium. The sulphur
group of elements.....198-203
Chap. XV.-Combination by volume. Synopsis of the gaseous compounds
previously studied. Condensation ratios. Combining-weight and volume-weight.
Double or product volume.' Molecular condition of elementary gases. Molecular formulase.....203-209
Chap. XVI.-Phosphorus. Allotropic modifications of phosphorus. Friction matches. Phosphorescence. Solutions of phosphorus. Poisonous properties of phosphorus. Manufacture of phosphorus. Red phosphorus. Amorphous and crystallized red phosphorus.  Safety-matches.  Phosphorus with
oxidizing agents. Phosphuretted hydrogen. Preparation and analysis of phos



CONTENTS.                                  xi
phuretted hydrogen gas. Composition of phosphuretted hydrogen as compared
with that of ammonia. Unit-volume weight of phosphorus twice its atomic
weight. Liquid and solid phosphuretted hydrogen. Compounds of phosphorus
and oxygen. Red oxide of phosphorus. Hypophosphorous acid. Phosphorous
acid. Spontaneous combustion. Phosphoric acid. Hydrates of phosphoric
acid. Meta- and pyro-phosphoric acids. Chlorides of phosphorus. Dissociation.
Bromides, iodides, and sulphides of phosphorus..209-241
Chap. XVII.-Arsenic. Sources and properties of arsenic. Arseniuretted hydrogen. Oxides of arsenic. Manufacture of arsenious acid. Isomerism. Properties and uses of arsenious acid. Reduction of arsenious acid.
Solubility of arsenious acid. Arsenic acid. Salts of arsenic acid; their analogy
to the phosphates. Detection of arsenic in cases of poisoning. Diffusion of
liquids. Dialysis. Crystalloids and colloids. Destroying the organic matters.
Precipitation and reduction of sulphide of arsenic. Marsh's test. Chloride of
arsenic. Sulphides of arsenic.  Sulphur-salts. Sulpharsenites and Sulpharseniates..                                                             241-265
Chap. XVIII.-Antimony. Sources, properties, and alloys of antimony.
Antimoniuretted hydrogen.  Testing for antimony.  Distinguishing between
arsenic and antimony. Teroxide of antimony. Antimoniate of antimony. Antimonic acid. Metantimonic acid. Chlorides of antimony. Chloridizing agents.
Sulphides of antimony. Sulphur-salts..266-281
Chap. XIX.-Bismuth.  Sources and properties of bismuth.  Fusible
metal. Teroxide of bismuth. Bismuthic acid. Chloride of bismuth. Sulphbde
of bismuth. The nitrogen group of elements.. 281-287
Chap. XX.-Carbon.  Allotropic modifications of carbon.  Diamond.
Graphite or plumbago. Graphitic acid. Gas-carbon. Coke. Anthracite.
Charcoal. Preparation of charcoal. Distillation of wood and of coal. Illuminating gas. Lampblaclk. Properties of charcoal. Reducing-power of charcoal.
Deflagration. Stability of charcoal. Charcoal absorbs gases. Induces combinations. Disinfects. Decolorizes. Compounds of carbon and hydrogen. Organic chemistry. Homologous series. Mlarsh-gas or light carburetted hlydrogen.
Atomic weight of carbon. Typical hydrogen compounds. Composition of illuminating gas. Carbonic acid. Preparation and properties of carbonic acid.
Ventilation of wells. Diffusion of gases. Solubility of carbonic acid. Carbonic
acid produced in the processes of fermentation, respiration, decay, and combustion. Liquid and solid carbonic acid. Decomposition, analysis, and synthesis of
carbonic acid. Carbonates. Carbonic oxide. Preparation and properties of carbonic oxide.  Carbonic oxide a deoxidizing agent.  Heat evolved in the combustion of carbonic oxide. Dissociation of carbonic oxide. Composition of carbonic oxide. Combustion. Luminosity of flames. The Bunsen burner. Quantity
and intensity flames. All flames gas-flames. Form of luminous flames. Blastlamps and blowpipes. Oxidizing and reducing flames. Chimneys. Indestructibility of matter. Kindling-temperature. Flames and fires cxtinguished by wire
gauze and other good conductors. Safety-lamps. Flaming fires. Loss of heat
from incomplete combustion. Chlorides of carbon. Compounds of carbon and
nitrogen. Bisulphide of carbon..287-304
Chap. XXI.-Boron. Source s of boron compounds. Allotropism of boron.
Boracic acid. Chloride of boron. Fluoridc of boron. Fluoboric acid. Sulhliide
and nitride of boron.... 365-372




X11                                CONTENTS.
Chap. XXII.-Silicon. Abundance of silicon. Preparation of silicon.
Modifications of silicon. Silicon and hydrogen. Oxide of silicon. Silicic acid.
Soluble silicic acid. Varieties of silicic acid. Silica in natural waters. Silicates.
Formulie of silicates. Decomposition of silicates. Composition and formula of
silicic acid. Chloride of silicon. Composition of chloride of silicon. Compound
of silicon, hydrogen, and chlorine. Fluoride of silicon. Fluosilicie acid. Fluosilicates. Sulphide of silicon. The carbon group..372-391
Chap. XXIII.-Sodium. Sources of sodium salts. Chloride of sodium.
Salt-works. Properties and uses of salt. Bromide and iodide of sodium. Sulphate of sodium. Glauber's salt. Supersaturated solutions. Bisulphate of sodium. Carbonate of sodium.  Leblauc's process.  Soda-ash.  Soda-crystals.
Bicarbonate of sodium. Yeast powders. Sulphides of sodium. Metallic sodium.
Hydrate of sodium. Caustic soda. Bases and acids. Replacement. Direct combination. Phosphates of sodium. Borax. Borax as a blowpipe test. Silicates
of sodium. Waterglass. Glass. Devitrified and colored glass. Hyposulphite
of sodium.... 391-414
Chap. XXIV.-Potassium. Sources of potassium. Carbonate of potassium. Bicarbonate of potassium. Hydrate of potassium. Caustic potash.
Strength of bases. Alkalimetry. Volumetric analysis. Valuation of potash and
soda-ash. Metallic potassium. Cream of tartar. Chloride of potassium. Bromide of potassium. Iodide of potassium. Cyanide of potassium. Sulphides of
potassium. Liver of sulphur. Sulphates of potassium. Nitrate of potassium.
Refining of saltpetre. Saltpetre not explosive. Gunpowder. Chlorate of potassium.414-436
Chap. XXV.-Ammonium-salts.  The hypothetical metal ammonium.
Hydrate of ammonium. Chloride of ammonium. Sulphate of ammonium. Nitrate of ammonium. Carbonates of ammonium. Sulphides of ammonium.   437-443
Chap. XXVI. —Lithium. Spectrum analysis. Rubidium and Caesium.
Thallium...............443-448
Chap. XXVII.-Silver. Occurrence of silver. Extraction of silver.
Metallic silver. The term metal. Acids and bases. Silver coin. Nitrate of
silver. Indelible ink. Oxides of silver. Fulminating silver. Photography.
The daguerreotype. Photography on glass. Photography on paper.  Chloride
of silver. Atomicweight of silver. Bromide of silver. Iodide of silver. Cyanide
of silver. Sulphide of silver. Sulphate of silver. The alkali group. Quantivalence. Atomicity........ 448-466
Chap. XXVIII.-Calcium-Strontium-Barium-Lead. Calcium. Carbonate of calcium. Calcareous petrifactions. Calc-spar and arragonite. Oxide
of calcium. Hydrate of calcium. Milk of lime. Slaked lime. Mortar. Caustic
lime. Uses of lime. Sulphate of calcium. Plaster of Paris, or calcined gypsum.
Plaster casts. Incrustation of steam-boilers. Hard water. Testing water. Phosphates of calcium. Chloride of calcium. Hypochlorite of calcium.  Bleachingpowder. Preparation of chlorine and oxygen from bleaching-powder. Chlorimetry. Nitrate of calcium. Sulphydrate of calcium. Suiphite of calcium. Metallic calcium.-Strontium  and barium. Peroxide of barium. Making oxygen
with peroxide of barium. Strontium salts. Barium salts. Uses of strontium




CONTENTS.                Xin
and barium salts. The calcium group. The alkaline earths.-Lead. Extraction
of lead. Separation of silver from lead. Oxides of lead. Corrosion of lead by
water. Leaden water-pipes. Suboxide of lead. Protoxide of lead or litharge.
Cupellation. Peroxide of lead. Red lead or minium. Sesquioxide of lead. Sulphides of lead. Chloride of lead. Sugar of lead. White lead. Silicate of lead. 467-500
Chap. XXIX. - Magnesium-Zinc-Cadmium.   Metallic magnesium.
Oxide of magnesium. Calcined magnesia. Hydraulic magnesia. Chloride of
magnesium. Sulphate of magnesium or Epsom salts. Carbonate of magnesium.
Citrate of magnesium. Phosphate of magnesium and ammonium.-Zinc. Granulated zinc. Galvanized iron. The galvanic current. Correlation of forces. Replacement of one metal by another. Oxide of zinc. Chloride of zinc. Sulphate
ot zinc. Alloys of zinc.-Cadmium. Sources and properties of cadmium. Unitvolume weight of cadmium half its atomic weight. The magnesium group.. 501-511
Chap. XXX. -Aluminum-Glucinum-Chromium-Manganese-TronCobalt-Nickel-Uranium. Metallic aluminum. Alloys of aluminum. Aluminum bronze. Atomic weight of aluminum. Oxide of aluminum. Hydrate of
aluminum. Aluminates. Mordants in dyeing. Lakes. Chloride of aluminum.
Sulphate of aluminum. Alum. Acetate of aluminum. Silicates of aluminum.
Clay. Earthenware, bricks, and pottery. Glazes. Hydraulic cement. Concrete.Glucinum.-Chromium. Oxides of chromium. Chlorides of chromium. Sulphate of chromium. Chromic acid. Chromates. Photolithography.-Manganese.
Protoxide of manganese. Sesquioxide of manganese. Alums. Binoxide of
manganese. Manganic acid. Manganates Chameleon mineral. Permanganio
acid. Permanganate of potassium.-Iron. Ores of iron. Extraction of iron.
The bloomary process. The blast furnace. Fluxes. Cast iron. Impurities of
iron. Wrought iron. The puddling process. Steel. The Bessemer process.
Protoxide of iron or ferrous oxide. Sesquioxide of iron or ferric oxide. Oxidation
by iron-rust. Ferric hydrate. Sulphide of iron. Iron pyrites. Chlorides of
iron. Ferrous sulphate or copperas. Ink. Dyeing. Ferric sulphate. Nitrates
of iron. Silicates of iron. Cyanides of iron. Prussian blue. Oxidizing ferrous
and reducing ferric salts.-Cobalt and nickel.  The terminations ous and ic.Uranium. Salts of the sesquioxides. The sesquioxide group. Atomic volume of
the alums. Rare elements allied to aluminum and iron.5. 12-561
Chap. XXXI.-Copper-Mercury. Extraction of copper from its ores.
Properties of copper. Alloys of copper. Dioxide of copper. Protoxide of copper.
Hydrate of copper.  Sulphides of copper. Chlorides of copper. Sulphate of
copper. Nitrate of copper. Verdigris.-Mercury. Its extraction and properties. Unit-volume weight of mercury half its atomic weight. Oxides of
mercury. Red oxide of mercury as a source of oxygen. Sulphides of mercury.
Vermilion. Calomel. Corrosive sublimate. Mercuric iodide. Sulphates and
nitrates of mercury. Amalgams. Detecting mercury... 562-577
Chap. XXXII.-Titanium. Tin. Its ore. Its extraction and propertiea
Tinning. Protoxide of tin. Binoxide of tin. Stannicacid. Metastannicacid. Stannates. Sulphides of tin. Protochloride of tin. Its reducing-power. Bichloride
of tin. Alloys of tin. The tin group.5... 577-585
Chap. XXXIII.-Molybdenum. Bisulphide of molybdenum. Oxides
of molybdenum.-Vanadium. Its occurrence. Its oxides.-Tungsten. Its occurrence and properties. Oxides of tungsten. Wolfram. Tungstate of sodium  585-587




X1V                               CONTENTS.
Chap. XXXIV.-Gold. Its widediffusion. Itsphysicalproperties. Indestructibility of gold. Refining of gold. Chloride of gold. Gilding. Alloys
of gold.-Platinum. Chlorides of platinum. Platinum sponge. Platinum black.
The platinum group..587-596
Chap. XXXV.-Atomic weights and symbols of the elements. Classification. Natural groups. Atomic heats of the elements. Atomic heats of compounds.
Value ol specific heats. Two sets of atomic weights in use. Barred symbols.
The student's view of discussions concerning theories.. 597-605
Appendix.-Glass tubing. Cutting and cracking glass. Bending, drawing,
and closing glass tubes. Blowing bulbs. Lamps.  Blast-lamps and blowers.
Caoutchouc stoppers, tubing, and sheets. Corks. Cork-cutters. Putting tubes
through corks. Supports for vessels-iron stand, sand-bath, wire gauze. Pneumatic trough. Collecting gases. Safety-tubes. Gas-holders. Defiagratingspoons. Platinum foil and wire. Filtering. Filters. Drying of gases. Dryingtubes. Chloride-of-calcium tubes Spring clip. Screw compressor. Waterbath. Iron retort. Self-regulatinggas-generator. Glass retorts. Flasks. Beakers.
Test-tubes. Test-glasses. Measuring-glasses. Burettes. Reading-off Burettes.
Pipettes. Wash-bottle. Porcelain dishes and crucibles. Rings to support roundbottomed vessels. Crucibles. Tongs. Pincers. Mortars. Spatulae. Thermometers. Furnaces. The metrical system of weights and measures. Table for the
conversion of degrees of the centigrade thermometer into degrees of Fahrenheit's
thermometer.  Table for the conversion of grammes into grains, and centimetres
into inches.                                                              i..  ~,  ~ ~     *   ~   i




EMENDATIONS TO BE MADE IN THE TEXT.
Page 2, line 25. Strike out the words "both copper and sulphur
have disappeared," and insert the words, "it suddenly glows vividly,
at the instant when the copper and sulphur combine."
Page 16, line 11. Insert after the word " air," the words " See Appendix, Fig. XVII."
Page 61, line 3 from foot. Add the words" See Appendix, ~ 17."
Page 173, Exp. 94. Instead of the first sentence (three and a half
lines) substitute the following: " Mix 75 or 100 grms. of flowers of
sulphur and an equal weight of slaked lime. Heat half a litre of
water in a capacious flask, until it boils vigorously. Pour the mixture
of lime and sulphur into the water, little by little, slowly enough not
check the boiling. The movement of ebullition prevents the solid
mazter from becoming impacted upon the flask."
Page 255, line 23. Strike out the words " consists of" and insert
" may be made by folding a circular piece of parchment paper into the
form of a cone, like the filter represented in Figr. XX. of the Appendix,
and cementing the upper edges of the paper with a drop of a thick
solution of shellac in alcohol. When nearly filled with the liquid to be
dialyzed and placed in a tumbler of water, the cone floats upright
and presents a large surface for the diffusion of the liquids. A more
elaborate form of dialyzer may be made with"
Page 26 7, line 6 from foot. Between the words " the " and " globule," insert the words " space about."
Page 267, line 12 fiom foot. Strike out the words " Chapter XX,"
and insert " ~ 431 and Exp. 220."
Page 299, line 1. After the word " lamp," insert " see Exp. 193."
Page 315, line 17. For 1" potassium," read " sodium."
Page 397, line 4. Insert after " sodium," the words "' in bleaching."
Page 498, lines 12 and 13 from  foot. Instead of " SbO" and
SbS5 " read " Sb206" and" Sb2 Sr."
Page 518, line 20. Strike out the words "is still," and insert "was
until lately."
Page 518, line 21. Strike out the words "in Europe."
Page 518, line 22.  Strike out the word "there."




Xvi                      EMENDATIONS.
Page 518, line 25. Insert after " clay," the words " or by passing
ammonia gas into a mixture of dilute sulphuric acid and roasted
shale."
Page 554, line 17 from foot. After " iron," insert the words "' containing a small percentage of sulphur."
Page 554, line 16 from foot. Insert after "formulated," the word
"approximately."
Page 590, line 16.  Instead  of "AuO" and  "AuOs,"  read
"Au2 0" and " Au2O~."
Appendix, page xiii., line 10.  After the word "motion," insert
"(6) The tube should always be thrust completely through the cork,
so that it may project a centimetre beyond the lower end of the cork."
ADDITIONAL EMENDATIONS.
Page 221, line 3 of Exp. 116. After "No. 2 " insert "about 25 C. M.
long."
Page 291, line 2 of Exp. 147. After "powder" read " Hold a pane of
window-glass before the eyes and stir the mixed powders into 14 grms.,"
etc.
Page 361, line 9 from foot. For "needed" read "theoretically necessary."
Page 369, line 1. Before "as" insert " something."
Appendix, page x., lines 16 and 17. Strike out the words " the tube g i
is closed by inserting a glass rod."




MANUAL
OF
INORGANIC CHEMISTRY.
INTJRODUCTION.
1. THE various objects which constitute external Nature present to the observing eye an infinite variety of quality and circumstance.., Some bodies are hard, others soft; some are brittle,
others tough or elastic; some natural objects are endowed with
life,-they grow; others are lifeless,-they may be moved, but
never move themselves; some bodies are in a state of incessant
change, while others are so immovable and unchangeable that
they seem everlasting. In the midst of this infinite diversity of
external objects, where lies the domain of Chemistry?  What is
the subject-matter of this science?
When air moves in wind, when water moves in tides, or in the
fall of rain or snow, the air and water remain air and water still;
their constitution is not changed by the motion, however frequent
or however great. A bit of granite, thrown off from the ledge
by frost, is still a bit of granite, and no new or altered thing. If
a solid piece of iron be reduced to filings, each finest morsel is
metallic iron still, of the same substance as the original piece, as
will appear to the eye if a morsel be sufficiently magnified under
the microscope. The melted, fluid lead in the hot crucible, and
the solid lead of the cold bullet cast from* it, are the same in
substance, only differing in respect. to temperature.  In all these




2                    CHEMICAL, CHANGE.
cases the changes are external and non-essential, not intimate
and constitunional; they are calledphysical changes.
2. When iron is exposed to the weather, it becomes cowered
with a brownish, earthy coating which bears no outward resemblance to the original iron; and, if exposed long enough, the
metal completely disappears, being wholly changed into this very
different substance, rust. A piece of coal burns in the grate,
and soon it vanishes, leaving nothing behind but a little ashes.
Dead vegetable or animal matters, buried in the ground, soon
putrefy, decay, and disappear. So, too, the fragment of granite
which frost has broken from the ledge, exposed for centuries to
the action of air and rain, becomes changed; it " weathers," and
after a time could no longer be recognized as granite. All these
changes involve alterations in the intimate constitution of the
bodies which undergo them; they are called chemical changes.
Experiment 1. —Mix thoroughly 3 grammes (for Tables of the
Metrical System of Weights and Mleasures, see Appendix) of coarsely
powdered sulphur with 8 grammes of copper filings, or fine turnings.
Put the mixture into a tube of hard glass, No. 3,  Fig. 1.
about 12 centimetres long, and closed at one
end. (For the manipulation of tubing, see
Appendix, ~~ 1-4.)  Hold the tube by the
open end, with the wooden nippers, as in
Fig. 1, and heat the mixture over the gas-lamp i
(Appendix, ~ 5), until both copper and sulphur
have disappeared.
Before heat was applied, the mixture of the          I
two substances was simply mechanical, and the
copper might have been completely separated
from the sulphur, by due care and patience;
but during the ignition the copper and sulphur
have united chemically, and there has been
formed a substance which, while containing both, has no external
resemblance to either. In the new body, the eye can detect neither
copper nor sulphur.
If 10 grammes of metallic iron be allowed to rust away completely
in moist air, the pile of rust which remains when the metal has disappeared, will weigh much more than 10 grammes. The iron has combined with two of the constituents of the atmosphere, the gas called
oxygen and.the vapor of water. The weight of the rust is the sum




ANALYSIS AND SYNTHESIS.                 3
of the weights of the metallic iron, and of the water and oxygen with
which the iron has combined.
Processes by which the whole character and appearance of the
bodies concerned are changed, as in these experiments, so that
essentially new bodies are formed from the old, are chemical
processes. It is the function of the chemist, on the one hand, to
investigate the action of each substance upon every other, to take
note of the phenomena which attend this action, and to study the
properties of the combinations which result from it; and on the
other, to contrive the resolution of compound bodies into their
simpler constituents.  He further seeks the general laws by
which the intimate combinations of matter are controlled. With
these ends in view, the chemist endeavors to pull to pieces, or in
technical language, to analyze, every natural substance on which
he lays hands.  Having thus found out the composition of the
substance, he seeks to put it together again, or to recompose it
out of its constituent parts. By one or both of these two processes, analysis (unloosing) and synthesis (putting together), the
chemist studies all substances.
3. There are two questions which the chemist asks himself
concerning every natural substance. First, Of what is it comnposed? With the means at his disposal the chemist may, or may
not, be able to resolve the substance into simpler constituents.
If he succeeds in decomposing it, he obtains the answer to this
first question; if the body cannot be decomposed by any known
method of analysis, the substance must be regarded as being
already at its simplest.  Such simple bodies are called elements.
Secondly, the chemist asks, How does this new substance comport itself when brought into contact with other substances already
familitr? There are sixty-five substances which are, at present,
admitted to be simple, primary substances, or elements. Of compound bodies, formed by the union of these elements with each
other, we find a series, numbering many thousands, in the inorganic kingdom. of Nature, comprising all the diversified mineral
constituents of the earth's crust; while another series, far more
complex in composition, and almost innumerable in multitude,
exists in the vegetable and animal world.  The task of the
chemist in thoroughly answering his second question would
B2




4                    FACT AND THEORY.
clearly be endless, were it not for the existence of general properties common to extensive groups of both elementary and
compound bodies, and of general laws which chemical processes
invariably obey.
While, therefore, the chemist seeks the answers to the two
fundamental questions above stated, he is at the same time
inquiring what relations exist between the properties of a body
and its composition, and he is also studying that natural and
invariable sequence of chemical phenomena, which, when fully
known, will constitute the perfect science of chemistry.
4. Generalizations from observed facts, so long as they are
uncertain and incomplete, are called hypotheses and theories;
when tolerably complete and reasonably certain, they are called
laws. The attention of the student should be constantly directed
to the keen discrimination between facts and the speculations
founded upon those facts-between the actual evidence of our
trained senses, brought intelligently to bear upon chemical phenomena, and the reasonings and abstract conclusions based upon
this evidence-between, in short, that which we may know, and
that which we may believe.
CHAPTER I.
A I R.
5. WE are everywhere surrounded by an atmosphere of invisible
gas, called air. All objects upon the earth's surface are bathed
with it. In motion, it is wind, and we recognize its existence
by its lowerful effects; but in the stillest places it exists as
well.
The presence of air in any ordinary bottle, flask, or other
hollow vessel which is empty, in the sense in which this word is
ordinarily applied, can be shown very simply by attempting to
put some other substance into the vessel, under such conditions
that the air cannot pass out from it. Or the air can actually be




ATMOSPHERIC PRESSURE.                    5
pumped out of the bottle; and it can be removed by other means,
both mechanical and chemical.
Exp. 2.-Wrap around the throat of a funnel, with narrow outlet,
a strip of moistened cloth or paper, so that the funnel shall
fit tightly into the neck of a bottle. After the funnel has  Fig. 2.
been fitted to the bottle, as is shown in Fig,. 2, fill it with
water, and observe that this water does not run into the
bottle. The bottle which we have called empty is in reality
filled with air, and it is this air which prevents the water
from entering the bottle. If, now, the funnel be lifted slightlyl
so that the mouth of the bottle shall no longer be completely
closed by it, the air within the bottle will pass out, and the
water in the funnel will instantly flow down.
6. In order to pump the air out of any hollow vessel, an apparatus known as the air-pump is commonly employed. Descriptions of this machine may be found in the text-books on physics.
For purposes of illustration, a portion of the air can always
be removed by suction, or by mechanically displacing the air by
some other substance, as when the finger thrusts the air out of a
thimble.
BEltp. 3.-Fit to any small flask or phial a perforated cork (for the
manipulation of corks, see Appendix, ~ 8), to which has been adapted
a short piece of glass tubing, No. 7. Tie upon this glass tube a short
piece of caoutchouc tubing. Suck part of the air out of the flask, and
then nip the caoutchouc tube with thumb and finger, so that no air
shall reenter. Immerse the neck of the flask in a basin of water, and
release the caoutchouc tube. Water will instantly rise into the flask.
to take the place of the air which has been sucked out.
7. The water, in this experiment, is forced into the flask by
the pressure of the superincumbent atmosphere. Air has weight.
It is subject to the law of gravitation, and is attracted towards
the earth's centre in the same way as other ponderable matter.
It has been found by experiment that a litre of dry air, at the
temperature of 00, and under a pressure of 760 millimetres,
weighs 1'2932 gramme.  It has also been determined by the
physicists that the force with which the air is attracted to the
earth is on an average equal to a weight of 1-033 kilogramme
to the square centimetre of surface. That is to say, the ocean
of air above us presses down upon every square centimetre of
the earth's surface with a force equal to that which would be




6                    PROPERTIES OF AIR.
exerted by a bar of metal, or other substance, a centimetre square
in section, and long enough to weigh 1'033 kilogramme.
If such a bar were constructed of iron, it would be 1'3 metre
long; if of water (and a bar of this substance can readily be
made by enclosing the water in a tube), it would be 10'33 metres
long; if of mercury, 76 centimetres long.  A clear conception of
the atmospheric pressure is important to the chemist, because of
its bearing upon the mechanical collection, manipulation, and
measurement of gases.
8. In addition to the qualities already mentioned, we find air
to be tasteless, and odorless, colorless when in small depths, but
exhibiting a blue tint when seen in large masses, as when in the
absence of clouds we look at the sky, or at a distant mountain.
Several of its properties are not peculiar, but are common to all
gases. Dry air obeys precisely the law of Mariotte up to a pressure of several atmospheres; that is to say, its volume diminishes
or increases in proportion to the pressure to which it is subjected; but it has never yet been condensed to a liquid.  Upon
being heated one degree centigrade, it expands    or in other
terms 0'003665, its volume at 0~. At the temperature of 0~, air
is,773 times lighter than water at 4~; that is, than water at its
maximum density.
These physical properties of air have been enumerated in order
to a distinct acquaintance with this gas, and that we may more
clearly comprehend its chemical properties as we now proceed to
study them.
9. Of what is air composed?  When a bar of iron is heated
in the air, as at a blacksmith's forge, it becomes covered with a
coating, which flies off in scales when the iron is beaten upon the
an il. If a bit of tin be melted in a shallow crucible, its surface
soon becomes covered with a white earth, or ashes. At a high
temperature both iron and tin slowly burn in the air, and are
converted into earth or ashes. With the exception of gold, silver,
platinum, and a few other exceedingly rare metals, all the metals
burn, or rust, when heated in the air.  If no air be present, this
rust or ashes will not be formed, however long or intensely the
metal may be heated.  But in what manner is the rust formed?
Is something driven out of the metal into the air, or does some



ANALYSIS OF AIR.                     7
thing come out of the air and unite with the metal?  Experiment shall answer.
Exp. 4.-Place 20 grammes of tin-foil in a porcelain dish, 4 or
5 c.m. in diameter. Counterbalance the dish, with its contents, upon
scales which will turn promptly with 0'5 grm. when thus loaded.
Heat the dish, placed upon the wire gauze of the iron-stand (see Appendix, ~ 9), over the gas-lamp moderately and continuously for two
or three hours. or until a large part of the metal has been converted
into ashes.  To facilitate the change, move the dish frequently, so
that a new surface of metal shall be often exposed to the air. Replace
the dish, when perfectly cool, upon the pan of the balance, and observe
that the weight of the dish and its contents has very decidedly increased.
A more rapid demonstration may be obtained by substituting for
the tin-foil a piece of magnesium wire or ribbon. This metal takes
fire when touched with a lighted match.
10. It is possible that daring the heating the metal may have
lost something, but it is certain that it has gained more. This
additional matter has not come from any alteration in the dish,
for'it is made of materials expressly adapted to resist such treatment, and a little cleaning will restore it to exactly its original
condition.  We have, then, taken something out of the air,
which, gaseous in the air, has become solid in the white ashes of
the tin.
If this something with which the tin has united can be separated again from the metal, and restored to the gaseous condition,
it. will be easy to compare it with common air, and so learn
whether the matter which combined with the heated tin is air
itself, or only a part of the air. It is quite possible to recover
the gas which enters into the composition of the rust of iron, or
copper, or tin; but the processes required are too circuitous for
the present purpose. From the rust of other common metals,
such as lead, manganese, or mercury, the absorbed gas can be
very easily expelled.  The rust of mercury is the most easily
decomposed.  Mercury-rust may be prepared by ihe long heating
of the metal in the air, in a manner strictly analogous to the
method already applied to the preparation of tin-rust.
Exp. 5.-Put into a tube of hard glass, No. 3, about 12 c.m.
long, and closed at one end as in Exp. 1, 10 grammes of the rust of




8                      ANALYSIS OF AIR.
mercury, a substance which is sold under the name of red oxide
of mercury. Such tubes of hard glass will be hereafter designated
as " ignition-tubes."  Attach to this ignition-tube, by means of aperforated cork, or caoutchouc stopper, a
delivery-tube of glass, No. 8, of such       Fig. 3.
shape and length that it shall reach
beneath the inverted saucer in the
pan of water, as represented in Fig. 3.
The point of departure, for the construction of the apparatus, is the top
of the gas-lamp placed upon the foot
of the iron-stand; the end of the ignition-tube should be about 4 c.m. -U
above the top of the lamp.  The
other details of the apparatus may be
learned from the figure.
Upon lighting the lamp, the air within the ignition-tube will expand, and a portion of it will pass out through the delivery-tube. This
air should be collected in a small bottle by itself, and thrown away.
The volume of air thus thrown away should of course not be much
greater than that of the tubes. (For a description of the pneumatic
trough, see Appendix, ~ 10.)
As the ignition-tube becomes hot, gas will be freely given off from
the red oxide of mercury contained in it. It is necessary to avoid
heating intensely a single small spot of the ignition-tube, lest the glass
soften, and, yielding to the pressure from within, blow outward, and
so spoil the tube and arrest the experiment. The gas-flame should be
so placed and regulated as to heat 3 or 4 c.m. of the tube at once.
Collect the escaping gas in bottles of 100 to 150 c. c. capacity.
As soon as the disengagement of gas slackens, lift the iron-stand up,
and take the delivery-tube out of the water, taking care that no water
shall remain in the end of the tube. Then, and not till then, extinguish
the lamp. (See Appendix, ~ 10.) In the upper part of the ignitiontube, and sometimes in the delivery-tube also, metallic mercury will be
found condensed in minute globules. The liquid metal is volatile at
the temperature to which it has been subjected, and has distilled away
from the hot part of the tube, and condensed upon the cooler part.
Thus is recovered the metallic mercury from which the red
mercury-rust was originally prepared by long heating in contact
with air. Is the gas, which the mercury originally took from the
atmosphere, air itself, or something different?
E.2. 6.-Introduce a lighted splinter of soft wood into a bottle of




AIR A MIXTURE.                       9
the gas collected in the last experiment. It will burn with much
greater brilliancy than in the air. Attach a bit of wax taper to a
piece of wire; light the taper, blow it out, and while the wick still
glows, introduce it into a second bottle of the gas of Exp. 5. The
glowing wick will burst into flame, and the taper will burn with extraordinary brilliancy.
11. It is very obvious, from these experiments, that the gas
which enters into the composition of mercury-rust is not air
itself. But since it came originally from the air, if it is not the
whole of air, it must be a part or constituent of air. This gas,
which causes combustible substances to burn with such intensity,
is indeed a constant constituent of the air, and a thorough study of
all its properties will hereafter convince us that it is a chemical
element of very various powers and great importance. It is called
oxygen, and under this name it will form the subject of the next
chapter.
12. But if oxygen be not air itself, but only a constituent of air,
it follows that air must have other constituents, or at least one
other constituent. If mercury be long heated in contact with a
certain confined portion of air, it will abstract from this air, as we
have seen, one of its ingredients, namely, oxygen, and there will be
left behind whatever of air is not oxygen. This experiment, the
original one by which the illustrious chemist Lavoisier demonstrated that air is a constant mixture of two different gases,
deserves careful study, both for its philosophical value and its
historical importance. The actual experiment lasts several days,
and is therefore unsuitable for repetition by the student.  A
description of it will suffice.
Into a flask, provided with a long neck, some metallic mercury was
introduced; the neck of the flask was then bent, as shown in Fig 4,
the flask placed upon a furnace, and the end of the neck plunged into
a basin of mercury; a jar was then placed over the end of the tube,
and a portion of the air within the jar was sucked out by means of a
bent tube; the mercury thereupon rose in the jar, and the point at
which it stood was accurately noted. The thermometer and barometer
were also simultaneously observed. Fire was then lighted in the
ftrnace, and the heat maintained for twelve days at a point just below
that required to make the mercury boil. The mercury became gradually covered with red scales, and the air in the jar, which at first
expanded fiom the action of the heat, slowly decreased in bulk until




10                   COMPOSITION OF Am.
fresh scales were no longer formed. From these red scales Lavoisier
obtained, by the method already exhibited (Exp. 5), the element oxygen. The residual air in the jar proved, on examination, to be unfit
for the support of combustion and
of animal life; a candle was instantlv extinguished by it, as if
plunged in water, and small aniinals, thrust into the gas, died in
a few seconds.  The gas is, in
reality, a second elementary substance, distinguished by marked
chemical and physical peculiarities.  It is called nitrogen, and   
under this name will be more
completely studied in another chapter.
13. The experiment of Lavoisier not only affords the means of
separating the two different gases of which air is composed, but
also determines the proportions in which they are mixed in air.
If the diminution in bulk. which the air in the jar undergoes
during the whole progress of the experiment be accurately
measured, it will be found that the bulk of the residual gas, the
nitrogen, is only four-fifths of the original volume of air. The
lost fifth is the oxygen which has combined with the mercury.
The air, then, is not an element, but is compound, and its two
principal ingredients are the  elementary bodies oxygen and
nitrogen, mixed in the proportion of four measures of nitrogen
to one of oxygen. It is quite possible to prove by synthesis
what analysis has thus taught. On putting together four measures of nitrogen and one measure of oxygen, a mixture is obtained, which, except by very refined experiments, is not to be
distinguished from pure air. Aqueous vapor is another normal
constituent of the actual atmosphere, and small traces of other
gases than nitrogen and oxygen are always present in it, as will
be set forth hereafter.




OXYGEN.                         11
CHAPTER IT.
OXYGE N.
14. Oxygen gas may be prepared by heating red oxide of
mercury, as described in Exp. 5; or, far more conveniently, by
heating a mixture of chlorate of potassium and black oxide of
manganese. For the present we have to regard these substances
merely as materials suitable for the preparation of oxygen. Their
constitution will be studied hereafter.
Exp. 7.-Mix intimately 5 grmins. of chlorate of potassium with
5 grms. of black oxide of manganese, which has been previously well
dried. Plhce the mixture in a tube of hard glass, No. 1, 12 or 15 c.m.
in length. Attach to this ignition-tube, by means of a perforated
cork or caoutchouc stopper, a delivery-tube of glass, No. 7, as represented in Fig 3, and described upon page 8. Heat the mixture in the
ignition-tube and collect the gas which will be given offl in bottles or
jars of the capacity of about 250 c. c.  The first 100 c. c. or so of gas
should be rejected, since it will be contaminated with the air originally contained in the apparatus.
It is easy to determine the moment at which the evolution of oxygen
commences, by noting the increased size of the bubbles of this gas as
contrasted with those of the expanded air, and the greater rapidity
with which the bubbles of oxygen come over. For every grm. of
chlorate of potassium taken, about 230 c. c. of oxygen gas should be
obtained.
Besides the general precautions described under Exp. 5, the following should here be observed. 1. Both the chlorate of potassium and
the oxide of manganese should be perfectly dry and pure-that is, free
from moisture, dust, or particles of organic matter. 2. So soon as the
oxygen begins to be delivered, the heat beneath the ignition-tube
should be diminished, if need be, and so regulated that the evolution
of gas shall be tranquil and uniform. 3. The uppermost portions of
the mixture should be heated before the lower. 4. An ignition-tube
should never be filled to more than one-thilid its total capacity, lest
portions of its contents be projected into the delivery-tube. The
chief danger to be guarded against in using ignition-tubes is the stoppage of their outlets with solid matter. 6. The ignition-tube should




12                  PROPERTIES OF OXYGEN.
always be inclined as represented in Fig. 3, and never placed upright
in the flame.
The oxygen thus obtained is to be employed in the experiments
shortly to be described. In case large quantities of oxygen are needed,
a similar mixture of equal weights of chlorate of potassium and black
oxide of manganese is heated in a retort of iron or copper, and the gas
is collected in large metallic vessels, called gas-holders, such as are
described in ~ 11 of the Appendix.
Besides the methods above described, there are many other
ways of preparing oxygen. Several of these methods will be
described hereafter, when the materials employed can be intelligently studied.
15. Oxygen is a transparent and colorless gas, not to be distinguished by its aspect from atmospheric air. Like air, it has
neither taste nor smell. It is, however, somewhat heavier than
air. If the weight of a measure of air be taken as 1,.then the
weight of the same measure of oxygen is found to be 1'1056.
At 00, and a pressure of 760 m.m. of mercury I litre of oxygen
gas weighs 1'4298 grm.
Since oxygen is thus heavier than air, it is not absolutely necessary,
in collecting it, or transferring it from one vessel to another, that we
should operate over water, as has been directed. When a gas is much
heavier, or much lighter, than atmospheric air, it may often be conveniently collected by displacement. A bottle can readily be filled
with oxygen from the gas-holder by carrying the delivery-tube to the
bottom of the upright bottle, and allowing the gas to flow in slowly,
as if it were water. In a short time the air will be wholly displaced,
and the bottle filled with oxygen. The progress of the operation can
be followed by testing the contents of the upper part of the bottle from
time to time, with a glowing match; when this bursts sharply into
flame the gas may be assumed to be pure enough for all ordinary
purposes.
16. Oxygen has never yet been reduced to the liquid condition.
Of all known substances it exerts the smallest refracting power
upon the rays of light. Compared with that of atmospheric air,
its refractive power is as 0'830 to 1'000. The specific heat of
oxygen, compared with that of an equal weight of water taken as
unity, is 0'2182; it has a lower capacity than other gases for
absorbing and radiating heat. Water dissolves it in small proportion; 100 volumes of water dissolve, at the ordinary tempera



OXYGEN SUPPORTS COMBUSTION.                13
ture, about 3 volumes of oxygen. It exhibits decided magnetic
properties. A glass tube filled with oxygen and suspended by a
fibre of silk between the ends of a strong horse-shoe magnet soon
comes to rest in such a position that its long axis is parallel to a
line drawn between the two poles. As with iron, its susceptibility to magnetization is diminished, or even temporarily suspended,
by a sufficient elevation of temperature. Its magnetic force, cor,
pared with that of the air, is as 17'5 to 3-4, that of a vacuum being
taken as 0.
17. A striking characteristic of oxygen is its power of supporting combustion. This has already been illustrated in Exp. 6,
and may be further exhibited by a great variety of experiments.
Exp. 8.-Place in a deflagrating spoon (see Appendix,
~ 12), a bit of sulphur as large as a pea. Light the sul-  Fig. 5
phur, and thrust it into a bottle of oxygen. It will burn
with a beautiful blue flame, and much gnore brilliantly
than in air. An acid, suffocating gas is produced..Exp. 9.-Place a piece of charcoal, that of bark is
best, in a deflagrating spoon. Kindle the charcoal by
holding it in the flame of a lamp, and then introduce it
into a bottle of oxygen. It will burn vividly, throwing
off brilliant sparks if bark charcoal has been employed.
In this experiment, as in the preceding, the products of the combustion are obviously gaseous, no solid substance being formed.
Exp. 10.-A piece of phosphorus, the size of a small pea, having
been well dried between pieces of blotting-paper, is placed in a deflagrating spoon, touched with a hot wire or a lighted match, and then
thrust into a jar of oxygen. It will burn with intense brilliancy and
the formation of a dense white smoke.
It should be observed that phosphorus is a substance which inflames
very readily in the air, when subjected to friction or any slight elevation of temperature. It is hence so dangerous that it must always be
kept under water. It should also be cut under water.
18. Many substances commonly called incombustible, because
they do not burn readily in ordinary air, burn vigorously in
oxygen. Of these, metallic iron may be taken as an example.
Exp. 11.-Convert two-thirds of a piece of fine piano-wire 20 or
30 c.m. long into a spiral, by winding it tightly around a glass rod,
No. 6, or lead pencil, and then withdrawing the rod or pencil. Thrust
the straight end of this wire into a cork or piece of thin board large




14                         OXIDES.
enough to cover one of the bottles of oxygen obtained in Exp. 7.
Upon the lower end of the wire coil tie a small piece of waxed
thread, or touch with a bit of melted sulphur the end of the spiral, so
that a small bead of sulphur shall adhere to the wire. Kindle the
thread or sulphur, and quickly place the wire in a bottle of oxygen, at
the bottom of which has been spread a layer of sand.
The burning sulphur heats the iron to redness, which then burns
brilliantly, with scintillation. From time to time, glowing balls of
molten matter fall off from the wire, and bury themselves in the sand
at the bottom of the bottle, or even melt into the glass.
If an abundance of oxygen be at hand, this experiment had better
be performed in a jar of 2 or 3 litres capacity,-a watch-spring which
has been rendered flexible by igniting it, and then allowing it to cool slowly, being the best material with which to  Fig. 6.
form the spiral coil. In this case it is well to tie a bit of
tinder upon the end of the coil as the kindling material, or
to attach a piece of twine to the wire, and soak this in
sulphur. But the experiment succeeds well even in very
small bottles of oxygen, provided the wire be fine, and
that the quantity of sulphur, or other matter, employed _
for kindling be not too large.
19. This experiment clearly proves what has been already
stated, that iron, when red-hot, combines with oxygen. It is
the burnt or oxidized iron which falls in globules to the bottom
of the bottle. This substance is called oxide of iron. The compounds which are formed by the union of oxygen with other
elements are called oxides.  The substances which have been
heretofore mentioned under the more familiar name of rust, like
iron-rust, tin-rust, mercury-rust, are called in chemistry oxides
-as the oxide of iron, oxide of tin, and oxide of mercury.
20. It will have been observed that the combinations obtained
in the foregoing experiments are of very various quality. Some
of these compounds are solid, others gaseous; some are acid and
caustic, while others are tasteless and innocuous. They agree
only in this, that they all contain oxygen. All these bodies will
be studied in detail hereafter. It concerns us now more particularly to realize the number and variety of the bodies into which
oxygen enters as an essential ingredient. In fact, the most
imnportant quality of oxygen is that, with a single exception, it
unites with all the other elements to form compounds.




OXIDATION.                       15
This act of combination~is often accompanied by development
of light and heat, as in the foregoing experiments; in the
affairs of common life we daily witness similar effects. All the
ordinary phenomena of fire and light depend upon the union of
the body burned with the oxygen of the air.  Indeed, the term
combustion may for all ordinary purposes be regarded as synonymous with oxidation.
Combustion is less vivid in air than in pure oxygen, because
of the nitrogen with which the oxygen of the air is diluted. If
a substance combines slowly with oxygen, it may often happen
that no evolution of heat or light can be detected.  Thus, when
a small piece of iron rusts at the ordinary temperature of the air,
there is perceived neither light nor heat. although a combination
of iron and oxygen has been formed, as in Exp. 11. Heat is
really disengaged in the slow rusting of iron, as in every act of
chemical combination, but it is taken up and carried away by the
circumambient air at the moment of its formation, so that it cannot usually be perceived. Slow oxidation, such as this, is often
spoken of as slow combustion.
21. Many of the compounds of oxygen are very familiar
bodies. Indeed, oxygen is the most widely diffused and the most
abundant of all known substances.  Not only does it occur in
the air, of which it constitutes about one-fifth the volume, as has
been already remarked, but at least one-third of the solid crust
of the globe is composed of it. It is the chief ingredient of
water, as will appear in a subsequent chapter. It enters largely
into the composition of plants and animals. Silica, in all its
varieties of sand, flint, quartz, rock-crystal, &c., contains about
half its weight of oxygen, and the same is true of the various
kinds of clay, and of chalk, limestone, and marble. Oxygen is
absolutely essential to the maintenance of animal and vegetable
life. The chemistry of the respiration of animals depends upon
the absorption of oxygen from the air respired. In the absence
of oxygen suffocation ensues.




16                      WNITROGOE.
CHAPTER IIIT.
NITROGEN.
22. The common modes of obtaining nitrogen depend upon
the removal of oxygen from the air.
Exp. 12.-Fill a tube of hard glass, No. 2, about 35 centimetres
long, with bright, not too coarse, copper turnings; place this tube
upon a semicylindrical trough of sheet iron, and support it upon a
ring of the iron stand, as shown in Fig. 7. It is well to interpose a
thin layer of asbestos between the tube and the iron trough, in order
to prevent the glass from adhering to the metal when it becomes soft
by heat. By means of corks, attach to one end of this tube a deliverytube leading to the water-pan, as shown in the figure; and connect the
Fig. 7.
--!
other end with a gas-holder filled with air. Light the lamps (for a
description of these, see Appendix, ~ 5) beneath the tube which contains the copper turnings, and wait until the copper has become redhot; then allow air to flow slowly out of the gas-holder over the hot
metal. The heated copper will combine with the oxygen of this air,
and retain the whole of it, so that nothing but nitrogen will be delivered
at the water-pan. This nitrogen may be collected in small bottles and
tested' with lighted splinters of wood, which should be instantly extinguished on being immersed in it.
23. A still simpler method of preparing nitrogen is to burn
out the oxygen from a confined portion of air, by phosphorus, or
by a jet of hydrogen.




PROPERTIES OF NITROGEN.                 17
Exp. 13.-Float a small porcelain capsule upon the surface of the
water-pan; a large cork must be placed beneath the capsule if the latter
will not float of itself. In the capsule put about a cubic centimetre of
phosphorus, and set it on fire. Invert over the whole a wide-mouthed
bottle, of the capacity of a litre or more, and hold this bottle so that
its mouth shall dip beneath the surface of the water. During the first
moments of the combustion, the heat developed thereby will cause the
air within the bottle to expand to such an extent that a few bubbles
of the air will be expelled; but after several seconds water will rise
into the bottle to take the place of the oxygen which has united with
the phosphorus.
The dense white cloud, which fills the bottle at first, is a compound
of phosphorus and oxygen which is soluble in water. It will, therefore, soon be absorbed by the water in the pan, and will disappear, so
that at the close of the experiment nearly pure nitrogen will be left
in the bottle. But, as the phosphorus ceases to burn before the last
traces of oxygen are exhausted (compare ~ 180), the nitrogen obtained
by this method is never absolutely pure.
As soon as the phosphorus goes out, the bottle should be shaken in
such a way that the porcelain capsule may be upset, and sunk in the
water-pan. The properties of the nitrogen may now be examined.
24. Nitrogen is a transparent, colorless, tasteless, odorless,
ijcondensable gas.  It is not quite so heavy as air.  If a measure of air weigh 1 gramme, then an equal measure of nitrogen
will weigh 0-9714 gramme. At 0~, and a pressure of 760 millimetres of mercury, 1 litre of nitrogen wveighs 1-256 gramme.
The specific heat of the gas is 0-244, that of an equal weight of
water being 1'000. A litre of water at 00 dissolves only 20 cubic
centimetres of nitrogen. Its refractive power in regard to light
is to that of atmospheric air as 1'034 to 1'000.
25. In its chemical deportment towards other substances, nitrogen is remarkably different from  oxygen.  Whilst oxygenr is
active and, as it were, aggressive, nitrogen, at least when in. the
condition in which it exists in air, is remarkably inert and
indifferent as regards entering into combination with other
bodies.  It is marked rather by the absence of salient characteristics than by any active properties of its own.  Mlany of the
metals, sulphur, phosphorus, and numerous other substances, may
be kept unchanged for any length of time in vessels filled with
nitrogen. A burning candle will instantly be extinguished when
c




18                        WATER.
thrust into a jar of nitrogen gas, for with the nitrogen the constituents of the candle have no tendency to combine.
As it extinguishes flame, so it destroys life. Animals cannot
live in an atmosphere of pure nitrogen. It may, indeed, be
breathed for a short time with impunity, but it does not support
respiration. It is not poisonous; if it were, it could not be
breathed in such large quantities as it is in air. An animal immersed in it dies, simply from want of oxygen.
26. As a diluent of the oxygen in the air, nitrogen is essential
to the existing balance and order of Nature.  All animal and
vegetable life-most inanimate matter, even-stands in harmonious relations with the chemical composition of the atmosphere.
The presence of so large a proportion of nitrogen in the air preven+s the too rapid action, as regards combustion and respiration,
that would take place in an atmosphere of unmixed oxygen.
Nitrogen is widely diffused in nature. Besides occurring in
the air, it is a constituent part of all animals and vegetables, and
of many of the products resulting from their decomposition.
Notwithstanding the indisposition of nitrogen in the free state to
enter into combination, a very large number of interesting and
important compounds can be formed by resorting to indirect
methods of effecting its union with other elements.
CHAPTER IV.
WATE R.
27. Another natural substance, quite as common as air, is
water. Three-fourths of the earth's surface is covered with it.
It is diffused through the atmosphere in the form of vapor, and
is a constituent of all animal and vegetable substances, and of
many minerals. We take up next this familiar substance, in
order that we may gain, through the study of it, a deeper insight
into chemical principles, and enlarge our experience by making
acquaintance with a new element. Let us first define vzith precision the external and physical properties of water, and then




PROPERTIES OF WATER.                  19
apply the two chemical methods, of analysis and synthesis, to the
closer investigation of its essential nature.
28. At the ordinary temperature of the air, pure water is a
transparent liquid, devoid of taste or smell. In thin layers it
appears to be colorless, but large masses of it are distinctly blue,
as seen in mid-ocean, in many deep lakes of pure water, and in
masses of ice, such as icebergs and some glaciers where it is possible to look through the ice from below.
This color can be seen upon the small scale by looking down through
a column of pure water, 2 metres long, upon pieces of white porcelain.
The water may be held in a glass tube, 5 c.m. wide, which has been
blackened internally with lamp-black and wax to within 1'25 c.m. of
the end, which is closed by a cork. Fill the tube with chemically
pure water, throw into it a few pieces of porcelain, and place it in a
vertical position? on a white plate. On now looking through the column of water at the bits of porcelain, which can only be illumined by
light reflected from the white plate through the rim of clear glass, it
will be observed that they exhibit a pure blue tint, the intensity of
which will diminish in proportion as the column of water is shortened.
The blue coloration may also be recognized when a white object is
illuminated through the column of water, by sunlight, and seen at
the bottom of the tube through a small lateral oaening in the black
coating.
29. At 40, the temperature at which it is densest, water is
773 times heavier than air at 00; at 150 it is 819 times heavier
than air of the same temperature. A cubic centimetre of water
at its greatest density, that is, at 40, weighed in a vacuum, is our
unit of weight-a gramme. One litre of water, which measures
1000 cubic centimetres, therefore, weighs a kilogramme. Water
is compressible and elastic; by the pressure of one atmosphere it
can be reduced to the extent of about 47-millionths of its original
volume, and this is true for every added atmosphere of pressure
so far as experiment has extended. Water expands upon being
heated, though at a less rate than other liquids; the rate of expansion increases with the temperature. Notwithstanding the
fact that water expands when cooled below 40, as well as when
warmed above that temperature, its refractive power on light
continues to increase regularly below 4~, as though it contracted.
The refractive index increases continuously between +5.2~0
c2




20                  PROPERTTES OF WATER.
and -13~0, the direction of the variation not changing in the
passage through the point of maximum density. At 00 the index
is 1-333.
30. Pure water at 0~, a temperature always to be obtained by
melting ice, is taken as a standard to which the weights of equal
bulks of other substances, liquid or solid, a 9 referred. In other
words, the specific gravity of water is tak -l as 1; and in terms
of this unit the specific gravities of all ott;r liquid and solid substances are expressed. The specific gra'ity of gold, for example,
is 19'3; that is to say, the weights of equal bulks of water and
of gold are to one another as 1 to 19~-.
31. Water is also the standard f specific heat. By specific
heats are meant the relative capacities for heat of the same
weights of different substances, at the same temperature.  For
example, to raise 1 kilogramme of mercury from 0~ to 1~ requires
only one-thirtieth of the quantity of heat necessary to raise 1 kilogramme of water from  0~ to 10. Water having been made the
standard of specific heat, its capacity for heat is denoted by 1,
and that of mercury will accordingly be 0'033. At the same
temperature, and for equal weights, water has a greater capacity
for heat than any solid or liquid known. Hence it results that
the specific heats of all solid and liquid substances are expressed
by fractions.
Water conducts heat very slowly; it may be boiled many
minutes at the top of the test-tube, which is held all the while
by the lower end, in the fingers, without inconvenience.
32. When exposed to a certain 0rgree of cold, water crystallizes, with formation of ice, or snow, according to circumstances;
and upon being heated sufficiently it is transformed into an
invisible gas, called steaim. Both these changes, however, are
purely physical, and therefore do not fi.ll within the province of
this manual. The chemical composition of the water is the
same, whether it be liquid, solid, or gaseous. The temperature
at which ice melts is one of the fixed points of the centigrade
thermometer, numbered 0~, and the temperature at which water
boils, under a pressure of 76 c.m. of mercury, is the other fixed
point, numbered 1000. Water evaporates at all temperatures,
and is therefore a constant ingredient of the atmosphere. Even




STEAM.                          21
ice slowly evaporates, at temperatures far below 0~, without first
passing into the liquid condition.
In crystallizing, that is to say, in assuming the solid form,
water increases in volume. The specific gravity of ice is only
0-916, which is equivalent to saying that, in the act of freezing,
916 c. c. (cubic centimetres) of water will be changed into a litre
of ice. From this fact result many familiar phenomena, such as
the floating of ice, the upheaving and disintegrating action of
frost, and the bursting of pipes and other hollow vessels when
water is frozen in them. The crystals of ice belong to the socalled hexagonal system; they are six-sided prisms, with regular
faces; by agglomeration they produce stellar and fern-like forms
of infinite variety and great beauty. Ice is a slow conductor of
heat, and a non- conductor of electricity. It becomes electric by
friction.
Steam is a colorless, transparent gas, as invisible as atmospheric air. It is lighter than air, the weight of any given
volume of steam, at the ordinary temperature, being to that of
the same volume of air as 0'622 to 1, a ratio deduced by calculation from the composition of steam. At 1000, the boiling-point
of water, the ratio of the weights of equal volumes of steam and
air is 0'455 to 1, and one volume of water furnishes about 1700
volumes of steam of 100~. When steam is heated by itself, without the presence of any liquid water, it is called super7teated
steam; but when there is water present, so that no excess of
heat can accumulate in the steam, above the quantity needed for
its formation under the pressure at which it exists, the steam is
called saturated, meaning saturated with water. When steam
escapes into the air, there is formed a multitude of little bubbles
or vesicles, composed of a film of water filled with air, precisely
similar to the vesicles seen in clouds and fogs. This steam-cloud
is sometimes improperly spoken of as steam or vapor, an error
against which the student should be upon his guard. Similar
fogs of air-filled vesicles are formed whenever the atmrosphere is
cooled to a temperature so low that the aqueous vapor contained
in it can no longer exist in the gaseous state.
33. Let us pass now to the analysis of water. Of what is
water composed?  We can determine this point by methods




22                    ANALYSIS OF WATER.
similar to those which were adopted in the examination of air.
There are several metals which, upon being brought into contact
with water, will abstract from it one of its ingredients, in the
same way that we have seen them abstract oxygen from the air.
Some metals can abstract this ingredient even at the ordinary
temperature. Thus the metal called sodium, on being brought
into contact with water, decomposes it, and, uniting with one of its
constituents, sets free another as a gas. This new gas is called
Hydrogen.
Exp. 14.-Make a small cvlinder of wire gauze by rolling a piece of
fine gauze, about 6 c.m. square, around a thick piece of No. 6 glass
tubing. Twist fine wire around the cylinder in order to preserve its
form, then slip the cylinder off the glass, and close one end of it by
pressure with a stout pair of pincers. Within this cylinder of wire
gauze place a piece of metallic sodium as large as a small pea, and
then close the upper end of the cylinder by pressure with the pincers,
as before. Quickly place the wire gauze cylinder under the mouth of
a small bottle of 100 or 200 c. c. capacity, which has previously been
filled with water and left inverted in the water-pan.
As soon as water comes in contact with the sodium, bubbles of gas
will issue from the wire gauze cage, and, rising through the water,
will collect at the top of the inverted bottle. If no gas is generated
during the first moment after the wire gauze has been placed in the
water, move the bottle to and fro, in such manner that the gauze
cylinder may be shaken about, and water forced through its interstices.
As soon as the evolution of gas has ceased, close the mouth of the
bottle with a small plate of glass, turn it mouth uppermost, remove the
plate, and touch a lighted match to the gas. The gas will take fire
with a sudden flash.
34. At a low red heat water can be decomposed by several of
the metals, such as iron, tin, zinc, and magnesium. Iron is well
adapted for this experiment..Exp. 15.-Provide a piece of iron gas-pipe, about 35 c.m. long, and
1 c.m. or more in internal diameter; fill it with small, bright ironturaings, and support it upon a ring of the iron stand over one or two
wire-gauze gas-lamps. By means of perforated corks, connect with
the iron tube, on the one hand, a glass delivery-tube leading to the
water-pan, as shown in the figure, and, upon the other, a delivery-tube
coming from a thin-bottomed glass flask, half full of water, and supported upon a ring of a second iron stand. Light the lamps beneath




ELECTROLYSIS OF WATER.                 23
the iron tube, and wait until its contents have become ied-hot; then
heat the water in the flask until it boils slowly. As the aqueous
Fig. 8.
vapor passes over the hot iron-turnings it will be decomposed, one of
its constituents will unite with the iron, and hydrogen will pass off
through the delivery-tube and may be collected in bottles at the
water-pan, so soon as the air originally contained in the tubes and
flask has all been expelled.
If, at the close of this experiment, and after the tube has become
cold, the iron be removed from the tube, it will be found to be covered with a black coating similar to that which forms on iron heated
in the air.
35. By these experiments it has been proved that one of, the
components of water is a gas called hydrogen. But with the
exception of the coating upon the iron of Exp. 15, we have as
yet nothing to indicate what other ingredients the water may
contain. Such evidence can, however, be readily obtained by
resorting to another method of analysis. If a galvanic current is
caused to flow through water, the force by which the constituents
of the water are held together will be overcome, and the water
will be resolved into the elements of which it is composed. On
immersing the platinum poles of a galvanic battery in water, to
which a little sulphuric acid has been added for the purpose of
increasing its conducting-power, minute bubbles of gas will immediately be given off from these poles, and will be seen rising
through the liquid. We have here abundant proof of the powerful action exerted by the battery upon the water. But the experiment will be much more satisfactory if it be made in a
vessel so arranged that the evolved gases may be collected for
examination.




24                 ELECTROLYSIS OF WATER.
For this purpose the apparatus shown in Fi g. 9 can be   Fig. 9.
conveniently employed. The test-glass, nearly full of
water which has been mixed with fiom  It to ~ of sulphuric acid, carries two platinum wires cemented with
shellac into the glass. These wires terminate above in
thin plates of platinum; over each of these plates there
is inverted a long, narrow test-tube full of water, acidulated in the same way as that in the test-glass. Upon 
connecting  the wires with a galvanic battery,- two
Bunsen's cells of medium size will be sufficient,-the  t!?Ijl
water will be decomposed, and the resulting gases, as
they are given off at the platinum plates, will rise, transparent and colorless, into the test-tubes. On comparing
the bulks of the two gases, it will be found that twice as x/
much gas has collected in the one tube as in the other.
If the test-tube containing the larger volume of gas be
now closed with the thumb, turned mouth uppermost,
and the gas within touched with a lighted match, it will take fire
and burn with the characteristic flame of hydrogen. It is, in fact,
hydrogen.
If the smaller volume of gas in the other tube be examined in the
same way, it will not inflame, although it gives intense brilliancy to
the combustion of the match. If a splinter of wood, retaining but a
single ignited spark, be immersed in the gas, it instantly exhibits a
vivid incandescence, and in a moment bursts into flame. This gas is
oxygen. It is thus proved that out of water may be unloosed two
volumes of hydrogen and one volume of oxygen.
If now the platinum plates be pressed so near together that a single
large test-tube, full of acidulated water, can be placed over both, the
gas obtained by passing the galvanic current will exhibit properties
differing from those of either hydrogen or oxygen. It is in fact a
mechanical mixture of these gases in the proportions in which they
would unite chemically to form water. The mixture is precisely
similar to that which would have been obtained if the two volumes of
hydrogen and one volume of oxyven, previously collected in two separate tubes, had been mixed in one. On touching a lighted match to
the mixed gas it instantly explodes with great violence, the hydrogen
burning suddenly, so that for a moment a flash of flame fills the whole
interior of the tube. Incited by the burning match, the hydrogen and
oxygen have combined chemically to form water, a portion of which
is deposited as dew upon the inner walls of the tube.
At the temperature of the air, and under ordinary circum



SYNTHESIS OF WATER.                 25
stances, oxygen and hydrogen do not combine chemically. Upon
being brought together they simply mix with one another mechanically in conformity with the physical laws which govern the
diffusion of gases. But under the influence of heat, of electricity,
and of certain other agents, the two gases will unite chemically,
and will thus again be converted into the water from which they
were derived.
36. It remains to be investigated whether hydrogen and oxygen, during their conversion into water, undergo any change of
volume, or merely combine to produce a measure of gaseous
water exactly equal to the sum of the measures of the constituents. To determine this point it is necessary to compare the
joint volumes of the constituents of the water with the volume of
the product formed, at a temperature high enough to maintain
the latter in the purely gaseous condition known as dry steam.
Through the closed end of a U-tube (Fig. 10) two platinum wires are
passed, and welded tightly to the glass walls of the tube. The outer
ends of these wires are formed into loops for the attachment of approFig. 11.
Fig. 10.
C      ~       C
priate battery wires; their inner ends are separated by a distance of
two millimetres. The general arrangement of the apparatus to be
employed is shown in Fig. 11. The U-tube is first completely filled
with mercury, and then the screw-compressor (Appendix, ~ 16) at a
is opened so as to afford a gradual exit to the metal in the open limb.
By means of a delivery-tube reaching down the open limb to the bend




26                   SYNTHESIS OF WATER,
of the tube, we introduce from a gas-holder (see Appendix, 5 11) a
quantity of a mixture of oxygen and hydrogen, made in the proportions in which they form water, —namely, two-thirds hydrogen, and
one-third oxygen, —in such a manner that the gas shall bubble up
through the mercury into the sealed limb, from which, of course, the
mercury escapes as the gas enters. A column of gas 2-5 or 30 c.m.
high is thus admitted.
It must be borne in mind that this mixture of hydrogen and oxygen
is very explosive; fire should be carefully kept away from the vicinity
of the gas-holder which contains it, and any remnant of the mixture
which is not used should be thrown away at the end of the experiment.
The gas-filled limb of the U-tube is next surrounded by a high glass
cylinder, b c, the ends of which are fitted with corks; through the
lower cork pass the U-tube, and a small glass tube, which is' connected
with a condensing worm, d, kept cool with water; through the upper
cork pass the wires which are to carry the electric current to the platinum points at the top of the U-tube, and a bent glass tube coming
from the flask, e. The top of the cylinder b c rises about 5 c.m. above
the sealed extremity of the U-tube. In the flask e, fusel oil, a liquid
which boils at 1320, a point much higher than the temperature at
which water becomes a gas, is kept in constant ebullition. The vapor
rising from the flask penetrates the annular space between the U-tube
and the enclosing cylinder, and quickly raises the tube to its own temperature. These strong-smelling vapors are not allowed to escape into
the atmosphere, but are carried out from the bottom of the cylinder
b c, into the condenser d.
When thus heated, the column of mixed oxygen and hydrogen in
the tube expands, and its height is marked by a caoutchouc ring, previously slipped over the cylinder b c. Care must be taken, before doing
this, to brincg the mercury to the same level in both limbs of the Utube, by adding or withdrawing mercury as may be required. A few
centimetres of mercury are next poured into the open limb, which is
then closed with a good cork. Between this cork and the mercury
intervenes a column of air, 8 or 10 c.m. in length, which will act as a
spring, and break the shock caused by the explosion of the mixed
gases. This mixture is now to be inflamed by causing an electric
spark to pass between the platinum points within the tube. This
spark may be obtained from a Ruhmkorff coil, or from an electrical
machine. The gases instantly rush into combination, with an intense
energy which produces the phenomena called explosive, and at the
high temperature which exists within the tube (1320) the water
formed retains the gaseous condition. On removing the cork, and




SYNTHESIS OF WATER.                   27
allowing the mercury to flow through the screw-compressor until it
is level in both limbs of the U-tube, it becomes obvious that the original volume of the gaseous mixture is diminished by one-third; the
residuary two-thirds are dry steam. If the U-tube is allowed to cool,
this steam will condense into liquid water.
Figs. 12 and 13 represent another form of apparatus for performing
this important expelin. en. In Fig. 12 the U-tube of Fig. 11 is replaced
Fig. 12.
Fig. 14.
L~ ~.Fig. 13.
by an iron U, such as is used in connecting parallel steam-pipes, into
which two straight glass tubes are fitted by means of caoutchouc
stoppers. One of these tubes is closed at one end, carries two platinum
wires, and may be graduated with advantage; the other is a plain
tube open at both ends. The iron U is fastened to a solid iron foot,
and into its side a small iron tube is inserted, to which is attached a
caoutchouc connector with a screw compressor, as in the other appa



28                  SYNTHESIS OF WATER.
ratus. The jacket-tube (Fig. 13) is a large tube, open at the bottom
and closed at the top with a caoutchouc stopper carrying the deliverytube for the hot vapor; near the bottom an exit-tube for the vapor is
welded into the glass. This jacket-tube, when in position, slips
tightly over the same caoutchouc stopper through which passes the
graduated tube. The advantages of this apparatus are that its parts
are not rigid, that it may be taken apart for cleaning, and that it may
be made with greater ease and fewer resources than the other apparatus. It is, moreover, capable of very general application in the
analysis of gases. Fig. 14 shows the ease with which the closed limb
of the U-tube in this apparatus may be charged. The closed limb and
the iron U being full of mercury, and the apparatus placed in an iron
pan suited to catch the overflow of mercury, the long open tube is
taken away, and the bent gas-delivery tube is inserted beneath the
closed limb through the iron U. The iron U acts simply like the
water-pan or pneumatic trough. When the required amount of gas
has been introduced, the open glass tube is replaced, and the level of
mercury in the two tubes is adjusted at will. For the experiment
now under consideration, it is well to employ the actual mixture ot
gases obtained by the electrolysis of water; and Fig. 14 represents a
simple bottle provided with two platinum plates and a single deliverytube for this purpose. The gas evolved is dried by passing over an
absorbent called chloride of calcium in a "dryingtube" (Appendix, ~ 15). The only noticeable feature   Fig. 15.
in this apparatus for electrolysis is the manner in which
the wires from the battery are connected with the platinum plates. This is more clearly shown in Fig. 15.
There passes through the caoutchouc stopper a short
piece of glass tubing, open at the top and drawn to a
point at the lower end, so as to enclose and hold tightly
a piece of platinum wire, as large as a common knitting-needle, previously placed within it. With the wire
thus welded to the glass is connected a thin plate of
platinum, which hangs in the liquid to be decomposed; this plate may be folded or rolled up. A little
mercury is poured into the glass tubes, and the battery-wires are
simply placed in the mercury when the operator desires to start the
decomposition.
This experiment demonstrates that two volumes of hydrogen
and one volume of oxygen are compacted, when chemically
united, into two volumes of steam.
37. We have thus established the composition of water by




ATOMS AND MOLECULES.                 29
analysis, having, through the agency of the electric current, resolved water into two gaseous constituents, hydrogen and oxygen,
and we have also demonstrated, by the converse or synthetical
method, that hydrogen and oxygen are its only constituents,
since we have ieproduced water by effecting the chemical union
of these two elementary materials mixed in due proportion.
If equal volumes of hydrogen and oxygen be represented
by equal squares, having the initials of the elements inscribed
therein, the composition of water by volume, and the condensation which occurs when the chemical union of the elements takes
place, may be thus expressed to the eye:
Each smallest possible or
greatest conceivable volume of
steam will invariably yield, on  I     [          H0
decomposition, its own volume                    1
of hydrogen, and half its vo-   1   3
lume of oxygen.
38. It has been agreed to call by the name " atom" the smallest
quantity. of an element which can be conceived to exist in combination; this technical term is applied only to the chemical
elements, and to certain chemical knots, or groups, of elements,
which, under conditions hereafter to be studied, play the part of
an element.
It has further been agreed among chemists to call by the
name "molecule," the least quantity of a compound, or of an
element, which can exist by itself uncombined. or take part in
any chemical process; a molecule always contains more. than
one atom, but these atoms may be either of one, two, or of
several kinds.
39. Physical experiments upon the expansion and contraction
of numerous gases, simple and compound, have proved that all
gases comport themselves in sensibly the same manner under
like variations of temperature and pressure; whence it has been
inferred that the intimate mechanical structure of all gases, compound as well as simple, is the same. This theoretical conception is expressed in the following propositions, of which the
second is the more general and includes the first:



30                MOLECULAR HYPOTHESIS.
The elementary gases contain, under like conditions of temperature and pressure, equal numbers of atoms in equal volumes.
Equal volumes of all gases, whether simple or compound, contain under like conditions, the same number of molecules.
The idea of an atom is complete and independent in itself; the
idea of a molecule is partly a consequent of the idea of an atom,
and partly of the physical facts which the definition helps to
formulate.
These definitions and hypotheses have found acceptance, partly
on the strength of experimental evidence, partly because of their
adaptation to the mathematical mode of investigating physical
problems which border on the domain of chemistry, but chiefly
on account of the clearness and formal consistency which they
have imparted to chemical language and modes of thought.
Chemical symbolization and nomenclature are mainly based on
the above definitions and hypotheses, which therefore justly demand the student's closest attention. Let us apply them to the
chemical compound, water.
40. The molecule of water, or least quantity of water which is
conceived to exist by itself, must yield, like any other quantity
when resolved into its elements, twice as large a volume of hydrogen as of oxygen. In accordance with the physical hypothesis
above explained, the molecule must consequently contain twice
as many atoms of hydrogen as of oxygen. The bulk and weight
of the molecule and atom are not absolute quantities, on account
of their assumed infinitesimal character. None but relative facts
can be known touching these hypothetical quantities, which are
both less than any assignable quantity, although one must be
smaller than the other. We shall express in the simplest terms
all our actual knowledge of the matter, and shall at the same time
conform to our definitions, in saying that a molecule of water
contains two atoms of hydrogen and one atom of oxygen. The
symbol 1H20 which we have already used to indicate the volumetric composition of water (~ 37) will now receive an added meaning; the I will represent for us an atom of hydrogen, and the O
an atom of oxygen.




ATOMIC WEIGHTS.                   31
When the proportions in which two bodies combine by volume,
and their specific gravities, or equal-volume weights, are known,
it is a matter of easy calculation to determine the proportions in
which they combine by weight. The specific gravity of oxygen,
or its density compared with that of air, has already been given,
namely, 1'1056. The specific gravity of hydrogen likewise referred
to air as the term of comparison, has been found by the most exact
experiments yet made to be 0'06926. Oxygep is therefore 16
times heavier than hydrogen. If hydrogen be made the standard
of specific gravity for gases, its specific gravity will be denoted
by 1, and that of oxygen will be 16. Now any measure of dry
steam is, as we have seen, resolvable into its own measure of
hydrogen and half that measure of oxygen; the weights of equal
measures of hydrogen and oxygen are as 1 to 16: but there is
twice as much hydrogen as oxygen in bulk, therefore the weight
of the hydrogen generated from any quantity of water, small or
great, is to the weight of the oxygen simultaneously produced, as
2 to 16. In 18 parts by weight of ste'am, water, or ice, there are
then 2 parts by weight of hydrogen and 16 of oxygen: and it
matters not what the absolute weight of these parts may be; the
proposition is as true of kilogrammes as of grammes, of the milligramme as of the millionth of the milligramme of water, in either
of its physical states.
Applying these facts of observation to our abstract definitions
of molecule and atom, it will appear that the molecule of water,
the least proportional weight in which it is conceived to exist
uncombined, must be composed, like any other mass of water, of
2 parts by weight of hydrogen, and 16 parts by weight of oxygen;
but in conformity with our definitions and hypotheses we conceive
of the molecule as consisting of two atoms of hydrogen and one
of oxygen; one proportional part by weight of hydrogen is then,
in chemical language, synonymous with one atom of hydrogen, and
16 of the same parts by weight is the relative quantity of the
atom of oxygen. As for volume, so for weight, absolute quantities
are entirely unattainable; the numbers express proportions only.
The numbers 1 and 16 are called the atomic weights of hydrogen
and oxygen respectively; they express the proportions by weight




32                 CHEMICAL COMBINATION.
in which these two elements enter into combination. If these
numbers be borne in mind, the symbol of water, H20, will always
remind us that water consists of 1 part by weight of hydrogen
and 8 parts of oxygen.
That any given weight of water, as for example one gramme,
is one-ninth hydrogen and eight-ninths oxygen, is a fact capable
of experimental demonstration. It is not difficult to decompose
a convenient weight of water, and to actually weigh separately
the hydrogen and the oxygen which are produced; the weights
of the two gases will invariably be to each other as 1 to 8 or as
2 to 16. The great value of the symbols used in chemistry may
be well illustrated by the amount of information condensed into
the concise expression H120: we learn from it the number and
names of the elements entering into the composition of water, and
the ratios in which the elements are united by volume and by
weight.
41. This discussion of the constitution of water rests upon two
solid facts of observation, namely the composition of water by
volume and its composition by weight; all else is plausible hypothesis and convenient theory. The strong chemical compound,
water, admirably illustrates the essential changes which the elements undergo, when they are joined together by that peculiar
force whose play it is the object of chemistry to study. Nothing
can be more striking than the contrast between the properties of
hydrogen and oxygen gases, or of a mechanical mixture of these
elements, and those of the liquid water which is produced by their
chemical union; even in dry steam, a prominent property of
hydrogen (inflammability) and a marked characteristic of oxygen
(power of supporting combustion) entirely disappear. In mechanical mixtures the constituents may be mingled in any proportions; in chemical compounds the elements are forcibly united
in definite volumetric and ponderal proportions, and the individuality of the elements is lost in the formation of a new substance with new properties. The CHEMICAL FORcE is. a peculiar
power, distinct from, though akin to, the forces of Light, Heat,
and Electricity; it is the province of chemistry to investigate the
conditions, modes, and effects of its action.




WATER IN NATURE.                      33
42. Having thus succeeded in determining the constituents of
air and water, we are naturally led to inquire whether it be not
possible to resolve oxygen, nitrogen, and hydrogen themselves
into simpler forms of matter. To this question but one answer
can be made-the result of the accumulated experience of many
philosophers of this and former generations-namely, that oxygen,
nitrogen, and hydrogen are incapable of decomposition by any
means as yet at our disposal. They resist the most powerful influences of electricity and heat, and they issue unchanged from every
variety and form of chemical reaction hitherto devised in the
hope of resolving them into simpler forms of matter. We are,
therefore, justified in regarding these gases as simple bodies, or
elements, in contradistinction to decomposable bodies, such as air
and water.
43. The water which occurs in nature is never absolutely pure.
In the form of ice, and as it falls from the clouds as rain or snow,
it is, indeed, tolerably free from  foreign sub,tances; but after
having once soaked into the ground, it becomes ch'arged with a
variety of mineral and other substances, which, being soluble in
water, are dissolved by it as it trickles through the earth.
Where the proportion of soluble matter contained in the water
is unusually large, and particularly if it possesses marked medicinal properties. the water is called mineral water, and the springs
from which it issues are known as mineral springs. Sea-water
may be regarded as a variety of mineral water.
44. For the conduct of chemical investigations it is often necessary to purify natural water. This is done by a process called
distillation. As a general rule distilled water is employed in all
delicate chemical operations.
Exp. 16.-Into a retort of 500 c. c. capacity, put 200 or 300 c. c. of
well-water. Thrust the neck of the retort into a half-litre receiver
placed in a pan of cold water. Cover the receiver with a cloth, or with
coarse paper, and upon this pour cold water from time to time, or pile
upon it fragments of ice. Place the retort upon wire gauze, on a ring
of the iron lamp-stand, and adjust the distance of the retort from the
lamp as described in Exp. 5, Fig. 3. Light the lamp beneath the retort, and bring the water to boiling. As fast as the water in the retort is converted into steam, this vapor will pass over into the cold
D




34                       DISTILLATION,
receiver, and will there be condensed again to the liquid condition.
Continue to boil until about three-quarters of the water in the.retort
has evaporated.
The earthy and saline ingredients of well-water are for the most
part not volatile; very few of them  are
capable of accompanying the water as it
goes off in vapor; hence the greater part of
the original impurity of the water will remain behind in the retort.
Besides the non-volatile impurities, there  
are often contained in well-water certain
volatile substances, such as ammoniacal salts
and organic matter, which pass over into
the receiver with the aqueous vapor; but
since it has been found that most of these
volatile matters go over with the first portions of the steam, it is only necessary to throw away that portion of
the distillate which is first condensed, in order to obtain thereafter
water of a high degree of purity.
This experiment must not only be so regulated that the retort shall
not boil over, but care must be taken that vapor alone shall pass off.
The ebullition should be so moderate that none of the particles of
water which are thrown up mechanically from the surface of the liquid
can be projected into the neck of the retort, or carried thither by the
current of steam.
45. In the operation of distillation the substance to be distilled
must in the first place be converted into the condition of vapor,
this vapor must next be transferred to another vessel, and there,
by refrigeration, be again condensed to the liquid state. As will
appear from the foregoing experiment, the vaporization is effected
in the retort or still, and the refrigeration in the condenser.  In
the experiment above given the receiver acts at once as receiver
and condenser; but a far more efficient apparatus can be constructed by interposing a long tube between the retort and the
receiver. This tube may be wrapped with cloths upon which
bits of ice are laid, or water is poured; or better, the tube may
be enclosed in a larger tube, or a metallic pipe, through which a
current of cold water is made to circulate. The water, which
may be iced if need be, is poured in through the funnel at the
lower end ofithe tube,.and passes out at the top (Fig. 17).  This




DISTILLATION.                      35
exceedingly convenient and valuable form of condenser was invented by Weigel; it is, however, commonly called Liebig's.
Exp. 17.-Place a few drops of the distilled water obtained in the
preceding experiment upon a piece of platinum foil (Appendix, ~ 13).
Hold the foil with iron pincers above the gas-flame in such a manner
that the liquid may slowly
evaporate without boiling or             Fig. 17.
spirting. After the water has
disappeared, no residue will
be found upon the foil. Take
out of the retort the same
number of drops of water, and
evaporate them upon the foil
as - before. A very decided
residue of earthy matter will
be left upon the foil.
46. In this country, where ice can be had in abundance at all
times, it may often be employed as a convenient substitute for
distilled water. In freezing, that is, in crystallizing, water rejects a great part of the foreign substances which were dissolved
in it. Hence, by collecting ice and remelting it, there can be
obtained water which is nearly pure.
Rain-water, also, especially that which has been collected in
the open country, is often pure enough to be used for chemical
purposes.
47. But even after all the mineral and all the organic matters
have been removed, the water is not yet absolutely pure. It still
contains oxygen and nitrogen in solution. Both of these gases
are soluble in water to a certain extent; and since the water
upon the surface of the earth is all the while in contact with air,
it must necessarily become charged with the constituents of'the
air. A method of collecting these gases for examination will be
described in a subsequent chapter; we are here more particularly
concerned with their removal. This may be effected by longcontinued boiling.
Exp. 18.-In a common long-necked medicine-phial of thin " blown"
glass, and of the capacity of about half a litre, place 300 or 400 c. c. of
recently distilled water. Draw out and bend the neck of the phial, as
D2




36                   SOLUTION ~i WATER.
shown in Fig. 18, and tie upon its point a short piece of caoutchouc
tubing. Boil the water steadily during half
an hour. Finally, nip the open end of the       Fig. 18.
caoutchouc tube, at the same moment remove
the lamp from beneath the flask, and instantly
seal the neck of the phial by directing the
flame of a blowpipe against the narrow spot.
This water can now be preserved for an indefinite length of time, without undergoing
change. Upon inverting the phial, the water,
which has been thus thoroughly freed from
air, will fall about as if it were a solid body,
and will strike against the glass with a sudden
shock. The apparatus is in fact nothing else
than the so-called water-hammer of the physicists.
It follows, then, that whenever absolutely pure water is needed
for chemical investigations, -natural water must first be distilled,
with the precautions above indicated, and this distilled water
must subsequently be thoroughly boiled, in order to expel the
gases which it holds in solution. Water so purified, though
necessary for chemical purposes, is unfit to support the life of
fishes or other animals which breathe in water, and is not suitable
for drinking. It is not only insipid and unpalatable, but is not
refreshing like ordinary water. Even if only a part of the dissolved gases have been removed, as is the case with water which
has been recently distilled, the taste of the water is still flat, and
repugnant. Hence, on board vessels where fresh water is prepared by distilling sea-water, the distillate should be left for
some time in contact with air, in order that by absorbing the
constituents of the air it may become fit for drinking.
48. As might be inferred from the foregoing, water has the
property of dissolving many substances, solid, liquid, and gaseous.
Sugar, for example, dissolves readily in water; but sand is insoluble therein.
A substance is said to be soluble when it is capable of being
divided in, and dispersed through water so intimately and completely that its particles become invisible, and can no longer be
separated by filtration; the result of this coalescence, or the




SOLUTION.                        37
solution as it is termed, is a transparent liquid, as a general rule
scarcely less mobile than the water itself.
Of the various substances soluble in water, some dissolve in
far larger proportion than others. With some liquids (as alcohol,
for example) water can be mixed in any proportion; but of ether
it dissolves but little, and of oil none. The proportion of any
substance that can be dissolved in a given quantity of water is
usually limited, and under fixed conditions is definite and peculiar
for each substance. When a given quantity of water has dissolved as much of a substance as it is capable of dissolving at the
temperature and pressure to which it happens to be exposed, the
solution is said to be saturated.  Generally speaking, solid substances dissolve in far larger quantity in hot than in cold water,
though with gases and some exceptional solids the contrary
obtains. From the saturated hot solution of any saline substance, crystals are usually deposited during the process of cooling. But so long as a solution is neither exposed to variations
of temperature, nor changed by the addition of another substance
or by the abstraction of either of its parts, it will usually deposit
nothing, and will remain unaltered during an indefinite period
of time.
During the act of solution, the first portions of the solid dissolve
with comparative rapidity, the subsequent portions dissolving
more and more slowly, until complete saturation is attained. In
preparing a solution of any solid, at the ordinary temperature of
the air, it is therefore inadvisable to add a large quantity of water
all at once; a much more satisfactory result will usually be obtained if the substance be rubbed in a mortar with repeated
small portions of water, the several portions of the solution being
poured off into a common receptacle as fast as the water becomes
nearly saturated.
There are many other liquids besides water which are commonly used as solvents; but as water is the commonest solvent of
all, and the most universally applicable, some of the general principles of solution may here be appropriately set forth.
49. Solution, though in many cases closely allied to chemical
action, is usually treated of as a distinct process. From the bestmarked chemical action it differs in several particulars. In true




38                        SOLUTION.
chemical combination the union of the several ingredients is so
close and intimate, that their properties are merged and lost in
those of the compound; while in the solvents proper, such as
water, alcohol, and benzine, the particles of the dissolved matter
appear to be merely mechanically divided and diffused through
the liquid.  The chemical properties of the dissolved matter
undergo no essential change during the act of solution, but remain unimpaired.  When common salt is dissolved in water, the
brine retains the peculiar taste of the salt, and behaves like salt
itself towards most chemical agents; moreover the salt can
readily be recovered unchanged by evaporating the water. But
if a piece of chalk be placed in muriatic acid, chemical decomposition and combination will immediately occur, the first signalized by a violent effervescence, the second resulting in the
formation of a liquid which contains neither chalk nor muriatic
acid, if the materials have been mixed in due proportion, but
which yields, on evaporation, a solid chemical compound, containing one of the constituents of each.
Chemical combination, as usually defined, occurs in fixed proportions only, whereas solution takes place in indefinite proportions; not only may many substances, as alcohol and glycerine,
be mixed with water in every proportion, but where the solubility
of a substance is limited in one direction, as that of common
salt, of which only about 0'355 part is dissolved by one part of
water at 150, the substance can nevertheless be dissolved in every
possible proportion below this maximum. Chemical action is
most marked between substances of unlike character; but with
solution the rule is different. In general, solution occurs most
readily when the solvent is not far removed in composition and
properties from the body dissolved.
Extreme cases of chemical action upon the one hand and of
solution on the other, are readily distinguishable.  But there is
a wide range between these extremes, and it is well nigh impossible to find a point at which the line of demarcation shall be
drawn.  Many cases which at first sight seem to be examples of
simple solution can readily be shown to depend in part upon
chemical force.
The majority of chemists are now inclined to regard most




HYDROGEN.                        39
instances of solution as feeble exhibitions of the chemical force,
or at all events as intermediate between purely chemical and
merely mechanical action. Solution facilitates chemical action
between heterogeneous materials, both by overcoming the force
of cohesion by which the particles of homogeneous solids are held
together, and also by bringing the particles of the unlike bodies
into intimate contact with one another through the vehicle of
the common solvent. Cohesion resists chemical action as it does
gravity; but solution overcomes cohesion, frees the particl s from
the bonds which held them, and, as we may imagine, leaves them
free to enter into other combinations.
CHAPTER V.
H YDROGE 0   N.
50. The commonest method of preparing hydrogen is by treating
zinc or iron with dilute sulphuric or muriatic acid. Unless very
large quantities of the gas are needed, this method is cheaper and
more convenient than either of those heretofore mentioned.
Exp. 19.-To a bottle 18 or 20 c.m. high, and of 500 or 600 c. c.
capacity, the mouth of which has an internal diameter of 2'5 to 3 c.m.,
fit a caoutchouc stopper or a
sound cork, furnished with a              Fig. 19.
thistle-tube (Fig. 19) and a gas
delivery-tube, of No. 6 glass.
Within the bottle put 15 or 20
grms. of granulated zinc, or small
scraps of the sheet metal, and as
much water as will cover the zinc
and close the end of the thistletube. Replace the cork in the
bottle, taking care to press it in
tightly, and gradually pour in
common muriatic acid through
the thistle-tube.  The thistle



40                PREPARATION OF HYDROGEN.
tube must reach nearly to the bottom of the bottle, so that its point
may dip beneath the water; and the muriatic acid must be added by
small successive portions, not more than a large thimbleful at a time.
On the addition of the first portions of the acid, chemical action
will ensue, the contents of the bottle will become warm, and gas will
be seen to escape from the liquid. This gas is hydrogen.
After all the air has been expelled from the bottle the hydrogen may
be collected over the water-pan, in inverted bottles filled with water,
or it may be passed into a gas-holder (Appendix, ~ 11). The rapidity
with which the gas shawl be evolved is easily controlled by regulating
the supply of acid; and the moment'at which the hydrogen ceases to
be contaminated with air can be determined by collecting small portions of the escaping gas, in wide-mouthed bottles of about 50 c. c.
capacity, and testing its quality by means of a lighted match. In
doing this the small bottle filled with gas must not be turned over, but
should be carefully lifted from the water without changing its vertical
position, and the lighted match should then be applied to the mouth
of the bottle. If the hydrogen be pure, it will burn tranquilly at the
mouth of and within the bottle; but in case the gas is still mixed
with much air, a sharp explosion will occur at the moment when the
match is applied to it.  In order to avoid these explosions, which
would be exceedingly dangerous if the volume of mixed gases were
large, it is indispensably necessary, in preparing hydrogen, to take care
that none of the gas shall be admitted into the gas-holder until all the
atmospheric air has been expelled from the bottle in which the gas is
generated. So, too, in experimenting with hydrogen, no light should
ever be brought into contact with the contents of the bottle into which it
is generated, or with any large quantity of the gas, until the purity of
the sample, or rather its non-explosive character, has been demonstrated by applying to a very small volume of the gas the test above
described.
This experiment, which has here been executed with zinc, can be
equally well performed with iron-filings, and with several other of the
less common metals.
Muriatic acid, or, in chemical nomenclature, chlorhydric acid,
is a compound of hydrogen and another element, called chlorine,
which will shortly be described.  The chemical composition of
this substance can be represented by the symbol H C1, in which
H represents, as before, the least proportional weight of hydrogen which exists in combination, and C1 the least proportional
weight of chlorine. We may likewise abbreviate the word zino




PROPERTIES OF HYDROGEN.               41
to the symbol Zn; and the chemical process, or reaction, by
which the hydrogen is liberated may then be symbolized by the
equation
2 H C+ Zn=Zn C12+2 H.
Since hydrogen is a gas, it escapes as such, and there remains
dissolved in the water within the bottle a compound of the elements chlorine and zinc, called chloride of zinc (Zn C12). The
zinc which was free, enters into combination, and the hydrogen
which was in combination, is set free; in other words, the zinc
has been substituted for, or has replaced, the hydrogen. It may
here be stated that chemists of all nations have agreed to represent each of the elements by a symbol, which consists either of
the initial letter of the Latin name of the element, or, when the
names of two or more elements begin with the same letter, of
the initial letter together with the first of the succeeding letters
in the Latin name which is distinctive. Thus Fe (Ferrum) is the
symbol of iron, W (Wolframium) of tungsten, C of carbon, Ca of
calcium, Cd of cadmium, Cl of chlorine, and Cr of chromium.
51. Hydrogen is a transparent, colorless, and tasteless gas,
odorless when pure. It is not poisonous, though animals die
from suffocation when immersed in it, as they do in an atmosphere
of nitrogen. It has never been condensed to a liquid. It is the
lightest substance known, being about 141 times as light as air.
If 1 volume of air weighs I gramme, an equal volume of hydrogen
will weigh only 0X0693 grm. It is 11,160 times as light as
water, 151,700 times as light as quicksilver, and 236,000 times
as light as platinum. 1 litre of hydrogen at 00, and a pressure
of 76 c.m. mercury, weighs 0'089578 grm.  Hydrogen is the
most suitable standard of specific gravity for gases, as water is for
liquids and solids; when thus used as the standard, its specific
gravity is, of course, unity. The student should, however, be informed that air is used by many writers as the standard of specific
gravity for gases.
52. The exceeding lightness of hydrogen can be illustrated in
various ways. From an inverted bottle, even though it be open
below, hydrogen will escape but slowly; but if a bottle of hydrogen be opened in the air, with the mouth upward, the gas will




42                 LIGHTNESS OF HYDROGEN.
quickly escape. Hence it can readily be poured or decanted upwards from one vessel to another.
Exp. 20.-Tlift from the water-pan a thick, strong, wide-mouthed
bottle, of 200 to 300 c. c. capacity, full of hydrogen, taking care to
hold it in a perpendicular position, with the mouth downward. With
the other hand place another bottle of equal size and strength, but
containing only air, beside the hydrogen-bottle, so that the mouths of
the bottles shall touch at one end. Gradually turn down the hvdrogen bottle, and at the same time push its mouth beneath that of the
air-bottle in such manner that the bottle which originally contained
the hydrogen shall at last stand upright beneath the inverted bottle.
During this operation, the lighter hydrogen flows up into the upper
bottle, while the heavier air sinks into the lower. If a burning match
be now thrust into the upper bottle, the hydrogen within it will take
fire; but upon applying the match to the lower bottle, originally full
of hydrogen, there will be found in it nothing but air.
In like manner, hydrogen may be collected by displacement-an upright delivery-tube being carried from the bottle in which the gas is
generated, to the top of an inverted recipient. The student will do
well to remember as a general rule, that in manipulating with hydrogen we must operate in a manner precisely opposite to that which
would be adopted if we were at work with water. Where water would
flow down, hydrogen will flow up.
Owing to its lightness, hydrogen is well adapted for filling
balloons.; and it is still sometimes employed for this purpose in
military operations, being prepared by means of hot iron, as in
Exp. 15. For purposes of illustration, soap-bubbles filled with
hydrogen will serve as well as balloons of more costly construction.
Exp. 21.-By means of a caoutchouc tube, attach an ordinary tobacco pipe to a gas-holder containing hydrogen. (See Appendix ~ 11,
Fig. xvii.) Dip the pipe in a solution of soap for a moment, then turn
it mouth uppermost, and slowly open the stopcock of the gas-holder so
that hydrogen may flow out and inflate the film of water upon the
mouth of the pipe.  The bubble will soon break away from the pipe
and rise rapidly through the air. If a burning match be applied to
the bubble the hydrogen within it will, of course, burst into flame.
53. There is another noticeable peculiarity of hydrogen which
is directly connected with its extreme lightness. It possesses in a




DIFFUSION OF GASES.                  43
high degree the power of diffusion. This diffusive power is a
physical property common to all gases and vapors; in the
case of hydrogen it is only the intensity of the diffusive
power which is remarkable.
When different gases, which have no chemical action,
upon each other, are brought into contact, they will not
remain separate, but will commingle. This tendency of
the gases to intermingle is so st ong that it will not only
overcome the greatest differences of specific gravity, but
even cause the spread of gases directly against powerful
currents of air or vapor. If a bottle of oxygen, standing
upright, be connected with an inverted bottle full of hydrogen, by means of a tube a metre in length, and no more
than 8 or 10 m.m. in diameter, both the bottles will be
found to be filled with a uniform mixture of the two gases
after the lapse of a very few hours.  Upon now touching
a lighted match to the open mouth of either bottle, the
gaseous mixture will explode. As a precautionary measure
it is best in this experiment to employ the thick, strong
bottles in which soda-water is kept-or, in lack of these, strong
wide-mouthed bottles enveloped in thick towels.
The velocities with which gases diffuse are in the inverse ratio
of the square rootj of their specific gravities. Hence it happens
that hydrogen being the most attenuated of all gases, diffuses with
the greatest rapidity. Compared with that of oxygen, its rate
of diffusion is as 4 to 1; that is to say, the relative
rates of diffusion of the two gases are inversely as  Fig. 21
the square roots of the numbers 1 and 16, which
represent the specific gravities of hydrogen and oxygen respectively.
Exp. 22.-A glass tube, 3 or 4 c.m. in diameter, and 30
or 40 c.m. long, is closed at one end with a plug of plaster
of Paris I or 2 c.m. thick. The tube is then set aside for
a day or two, in order that the plaster may become dry.
When the plug is dry, fill the tube with hydrogen by
displacement, and set it upright in a glass of water. Water
will rise rapidly in the tube, since hydrogen escapes
through the plaster more rapidly than air can enter the tube through




44                HYDROGEN INFLAMMABLE.
this porous plug. That. some air does enter, however, can be shown
by exploding the contents of the tube by applying a lighted match,
after the lapse of some time. Of course, if the tube be left to itself, air
will slowly enter through the plaster, so that the water within the
tube will in due time sink to the level of the outside liquid.
On account of its high diffusive power, hydrogen can be kept
only in perfectly tight vessels. It has been found that it will
leak rapidly under a pressure of 27 or 28 atmospheres through
stopcocks that are perfectly tight for nitrogen at a pressure of
even 50 or 60 atmospheres; from the same cause it cannot be
kept for any length of time in bladders or rubber bags. If a
sheet of paper be held a short distance in front of the opening of
a gas-holder from which hydrogen is escaping, the current of gas
will pass directly through the paper, and can be inflamed upon the
other side of the sheet. The high diffusive power of hydrogen,
which is to some extent shared by its compounds also, is an
obstacle to be overcome before ballooning can be made practicable.
Sound is propagated in hydrogen but little better than in a
vacuum.  The specific heat of hydrogen is 3'4046, that of an
equal weight of water being 1'000; it is 0'2356, that of an
equal volume of air being 0'2377. It refracts light very powerfully.
54. Hydrogen is exceedingly inflammable, as has been already
seen; that is to say, the temperature at which it takes fire is
comparatively low. But, as a matter of course, it extinguishes
any burning body which is immersed in it, since oxygen is necessary for the support of combustion.
Exp. 23.-Carefully lift from the water-pan a bottle of 200 or 300
c. c. capacity, completely full of hydrogen, slowly carry the bottle, the
mouth of which is of course held downward, to a burning candle or
splinter of wood, and depress the bottle over this flame. The hydrogen will take fire and burn, below, at the mouth of the bottle where it
is in contact with the oxygen of the atmosphere; but the flame of the
candle will be extinguished the moment it becomes completely enveloped by the hydrogen. The candle can easily be relit by slowly
lifting the bottle until the wick is brought into contact with the air
and the burning hydrogen.
E'xp. 24.-Fill a bottle of the capacity of 400 or 500 c. c. with hydrogen, close the mouth with a cork or a plate of glass, stand the bottle




UNIT OF HEAT.                     45
upon the table with the mouth upward, remove the stopper, inflame
the hydrogen, and immediately pour out of a pitcher into the bottle a
large quantity of water. The flame will instantly be increased, since
the water will force the gas out of the bottle into the air. Within the
bottle the hydrogen can only burn gradually, since it takes time for the
outside air to enter; but if the gas be pushed out of the bottle into
the air, it will burn at once.
In the familiar instances where water extinguishes fire, it does so by
reducing the temperature of the combustible-that is to say, by cooling
it to below the temperature at which it will take fire. It would act
thus in this case, were it not for the lightness and mobility of the
hydrogen, by virtue of which this gas immediately escapes froml contact with the water.
55. It has been seen that the hydrogen-flame affords only an
exceedingly feeble light; buit it would be a grave error to infer
that but little heat is developed by the combustion.
The temperature of the hydrogen-flame is in reality very
high. Indeed it has been found that when a given weight of
hydrogen enters into chemical union with oxygen, more heat is
developed than in the burning of the same weight of any other
substance.
In order to determine the amount of heat which is developed
in any act of combination, this heat can be transferred to water,
and there estimated either by the quantity of water heated, or
the amount of steam produced.  A unit of heat is that amount
of heat which will raise 1 gramme of water from 0~ centigrade
to 1~.
The amount of heat evolved during the combustion of a body
is as constant and unvarying as any other direct consequent of
its properties, and the quantity of heat evolved is absolutely the
same, no matter whether the combustion occurs in air or pure
oxygen, or whether it be slow or rapid. The actual amount of
heat developed during the most vivid combustion is no greater
than when the same combustible combines with oxygen, by gradual oxidation, without visibly burning.  From  1 gramme of
hydrogen, as it unites with oxygen, there are evolved 34462 units
of heat.
Although the same amount of heat is developed, in the aggregate, when a litre of hydrogen burns in the air, as when it burns




46                OXY-HYDROGEN BLOWPIPE.
in pure oxygen, it is none the less true that a far hotter flame
is obtained by burning the hydrogen in oxygen than in air. If
the combustion be complicated by the presence of nitrogen from
the air, a great deal of heat will expand its energy in expanding
this useless nitrogen. Moreover, since the nitrogen occupies much
room, it will, as it were, keep asunder the particles of hydrogen
and oxygen; the flame will thus be made longer and more dispersed, or, in other words, the heat evolved by the union of the
hydrogen and oxygen will be spread over more space than if the
nitrogen were absent.  Where nothing but hydrogen and oxygen
are present, and in the precise proportions in which these gases
unite to form water, the flame produced by their union will be to
all intents and purposes solid, and the heat will be concentrated
in the smallest possible space.
Exp. 25.-Provide two gas-holders (see Appendix, ~ 11), one full
of hydrogen, the other full of oxygen, also a metallic jet so constructed
that the tube wh!~h carries the oxygen shall pass through the
centre of the hydrogen tube as shown in Fig. 22. Screw the jet on
to the oxygen gas-holder, and connect the other opening with the
hydrogen gas-holder by means of a caoutchouc tube. Open the cock
of the hydrogen gas-holder, and inflame the gas at the point of the jet,
then slowly open the cock of the oxygen gas-holder until the flame of
the burning hydrogen has been reduced to a fine pencil. This apparatus is known as the compound, or oxy-hydrogen blowpipe. In order
to insure a steady flame, care must be taken that a constant and sufficient pressure be maintained upon the contents of the gas-holders.
Fig. 22.
Exp. 26.-In the flame of the oxy-hydrogen blowpipe, described in
the foregoing paragraph, hold the end of a piece of platinum wire,
about 10 c.m. long and less than one m.m. in diameter. The platinum
will melt and fall down in drops.
The intense heat of the oxy-hydrogen flame is thus admirably illustrated; for platinum is an exceedingly infusible metal, which can




FORM OF GAS-FLAMES,                   47
scarcely be softened in the hottest furnace. The drops of platinum
should be caught upon sand, or in water. In case the melted platinum
falls into water, a portion of the water will be decomposed into hydrogen and oxygen, bubbles of which will be seen issuing from the water.
Some of this gas can easily be collected, by an assistant, in an inverted
ignition-tube filled with water, and its explosive character demonstrated by applying a lighted match.
Exp. 27.-If a piece of chalk or lime, scraped to a fine point, be held
in the flame of the oxy-hydrogen blowpipe, it will quickly become
white-hot and evolve light of great brilliancy, almost comparable with
that of the sun. If the bit of lime be long exposed to the intense heat,
it will undergo incipient fusion, and afford less light than at first.
Where a constant light is desired, cylinders or plates of chalk are
kept continually moving before the flame by mechanical power, so that
fresh portions shall continually be brought into the flame. This is the
so-called Drummond or Calcium light, often employed for night signals
and optical experiments.
56. No matter in what way hydrogen is burned, whether in
the pure state or in combination with other materials, whether in
pure oxygen or in the air, the product of the combustion is always
water. At the high temperature of the flame, this water must of
course remain in the condition of a gas, but it can readily be
brought to the liquid state by reducing the temperature.
Exp. 28. —Over a jet of burning hydrogen, best obtained from a
gas-holder, hold a dry, cold bottle. The glass soon becomes covered
with a film of dew, as the water generated by the union of hydrogen
and oxygen condenses in droplets upon the cold sides of the bottle.
57. As the burning jet of hydrogen is the simplest instance of
combustion with flame, some exact knowledge of the form and
quality of flames may here be gained.
As hydrogen, or any other gas, issues from a small orifice into
the air by force of pressure from behind, the escaping gas
assumes a certain definite shape in accordance with the physical
conditions to which it is exposed, just as a fountain of water
takes form in accordance with the size and shape of the orifice
from which the water is expelled, the pressure by which it is expelled, the gravity of the water, the resistance of the air, and the
force and direction of the wind, and so forth.
If a lighted match be brought into contact with the column or




48       EXPLOSIVE MIXTURE OF HYDROGEN AND OXYGEN.
fountain of gas, the gas will take fire and burn; that is to say,
the hydrogen will enter into combination with oxygen as fast as
the latter can be furnished from the air.  But, in any event, the
column of gas will burn only upon the outside, for there alone
can the oxygen of the air come into contact with the hydrogen.
Neither the interior of the flame nor the contents of the reservoir
from which the gas is flowing can burn; for they consist of pure
hydrogen, which, as has been shown in Exp. 23, will by itself
immediately extinguish combustion.
All ordinary gas-flames are, of necessity, hollow. They are
visible to us through the evolution of light which is an accompaniment of the chemical action.  But even if no combustion
were going on upon its surface, the escaping column of gas could
still be made visible by causing it to pass through a quantity of
dust or other fine powder before coming into the air. A portion
of the solid matter would be transported by the current of gas,
and the form assumed by the latter would be made manifest.'-he shape of the unignited gas-column would of course be
somewhat different from that of the burning flame; in the latter,
not only is the outer edge of the column sharply defined by the
zone of combustion, but the actual form of the column itself is
modified through the expansion of the gas as it becomes heated by
the enveloping fire.
58. If, instead ol' burning pure hydrogen as it flows into the
air, as in the foregoing experiments, the gas be first mixed with
air or oxygen and then ignited, a very different result will be
obtained. The hydrogen being now in contact with oxygen at all
points, the entire mass of gas will burn with a violent explosion
the instant a light touches it.
Exp. 29.-Introduce 2 volumes of hydrogen and 5 volumes of air
into a strong round-bottomed bottle such as is used for soda-water.
Close the mouth of the bottle with a cork, and shake violently in order
that the gases shall be mixed. A small quantity of water should be
left in the bottle to act as a stirrer. Grasp the bottle firmly in one
hand, remove the cork with the other, and apply the open mouth of
the bottle to a lighted candle. An explosion will immediately ensue.
Exp. 30.-Into a gas-holder, bladder, or rubber bag, introduce a
mixture of 2 vols. hydrogen and 1 vol. oxygen. Connect therewith,




EXPLOSIVE MIXTURE OF HYDROGEN AND OXYGEN.        49
by means of caoutchouc tubing, a tobacco-pipe, or bit of glass tube.
Press the gas through the pipe into a dish of soap-suds, in such manner
that there shall be formed upon the surface of the suds a mass of foam
as large as an egg. Close the gas-holder and remove it from the vicinity ol the suds. On now touching the foam with a long, lighted
stick an exceedingly violent explosion will occur.
It is well to avoid the olrrmation of a large quantity of foam in this
experiment, since the concussion is in any event deafening. In this
experiment the explosive mixture is purposely confined in an exceedingly flimsy envelope, in order that no harm may be done by the
fragments.
Care should be taken to throw away any remnant of the mixture of
hydrogen and oxygen which may have been left in the gas-holder at
the close of the experiment, and upon no account should fire ever be
brout ht into its vicinity.
59. The cause of these loud explosions is twofold.  By the
act of combination water is formed, and at the same time intense
heat is emitted; the water, or rather steam, is thereby enormously
expanded, so that for a moment there is violent motion outwards
in all directions. This outward motion would scatter about in a
most dangerous manner the fragments of any vessel in which,
through carelessness, a mixture of hydrogen and oxygen might
be ignited-unless, indeed, the vessel were very strong, small,
and of large aperture. But in the next instant, as the steam
condenses, there is an even more violent motion inwards.
The original mixture of hydrogen and oxygen occupies about
2000 times as much space as the liquid water which results from
the combination of these gases.  Hence a partial vacuum  is
formed, into which air rushes from all sides; and it is the heavy
and sudden undulations thus communicated to the air which
occasion the noise. The outward and inward shocks follow oie
another so quickly that the ear cannot distinguish between them.
Mixtures of hydrogen and air produce less violent explosions
than mixtures of hydrogen and oxygen, because of the inert
nitrogen in the air, which acts as an elastic pad, or cushion, to
break the force of the shock.
60. Since air is everywhere about us, and since all ordinary
conmbustions occur in it, it has become customary to speak of it
and of oxygen as supporters of combustion, in contradistinction to
E




50              OXYGEN BURNS IN HYDROGEN.
the so-called cowz'.stibles, such as hydrogen. These terms are
often convenient, but, as will appear from the following experiment, they have only a relative, and no absolute significance.
Expl. 31.-Provide a tube of thin glass, the neck of a broken retort,
for example, 30 or 40 c.m. long, and 3 or 4 c.m. in d ameter; fix it in
a vertical position, so that the lower opening shall be 20 or 30 c.m.
above the table, and connect the upper opening with a gas-holder filled
with hvdrogen.
To a second gas-holder, containing oxygen, attach a caoutchouc
tube, and to the end of this fit a piece of glass tubing, No. 7, 25 or 30
c.m. long, bent at a right angle at 5 or 10 c.m. from the end which is
attached to the caoutchouc tube, and drawn out to a fine open point
at the other. The caoutchouc tube must be lone enough to reach as
far as the lower mouth of the wide vertical glass tube above
mentioned.
Open now the stopcock of the hydrogen gas-holder so that the
ver'ical tube shall be filled with the gas, then apply a lighted match
to the mouth of this tube, and regulate the flow of gas so that the
latter may continue to burn slowly at the lower edge of the tube.
Finally, open the stopcock of the oxygen gas-holder, so that a current of this gas shall flow through the pointed delivery-tube, and thrust
this tube up into the middle of the wide vertical tube which has been
filled with hydrogen.
As the stream of oxygen passes through the burning hydrogen at
the bottom of the vertical tube, it takes fire, and afterwards continues
to burn in the atmosphere of hydrogen within the tube.
CHAPTER VI.
OTHER COMPOUNDS OF THE ELEMENTS ALREADY STUDIED.
61. Oxygen and hydrogen do not unite directly in any other
proportion than that in which they form water; but by indirect
means, too,complex for profitable study at this stage, a molecule
of water can be made to combine with an atom of oxygen, forming'
a new substance, known by the name of peroxide of hydrogen.
Its formula is,.in accordance with this statement, t2,O, and it




PEROXIDE OF HYDROGEN.                 51
would yield, if completely decomposed into its elements, equal
volumes of hydrogen and oxygen; its composition by weight
must be 2 parts of hydrogen to 32 of oxygen; it contains just
twice as large a proportion of oxygen as water. The best method
of preparing this substance does not yield the pure thing itself,
but only a concentrated solution of it in water; the specific gravity of this solution is 1-45: it is colorless, transparent, and somewhat syrupy, has a metallic taste, corrodes the skin, and bleaches
vegetable colors.
How different a substance this is from water, appears at once
from this enumeration of some of its properties. It is very unstable, being readily decomposed by heat, and by contact with
various substances at the ordinary temperature, into oxygen and
water. Its instability and the intense chemical activity of which
it is capable, emphatically distinguish this, as yet obscure, body
from the neutral, stable, inactive compound of hydrogen and
oxygen, common water.
62. But though the peroxide of hydrogen is not water, it is
nevertheless a true chemical compound of the same elements which
are united in water; it is definite and constant in composition,
and its properties are as unlike those of its elementary constituents as are those of water. A new fact of great significance
here comes plainly into view. Two of the elements are evidently.
capable of combining in two definite proportions to form two
chemical compounds, each differing from the other and from its
primary constituents. The study of the compounds of nitrogen
and oxygen will bring into clearer view the general principle of
which this fact is a single illustration. These compounds form a
series of five members, all derived, more or less directly, from
common nitric acid.
63. Nitric Acid.-Two abundant sources of this material are
found in nature, and are familiar as articles of commerce.  Saltpetre or nitre, a whitish, saline, crystallized substance, now
mainly brought from India, is one of these sources; a similar
substance, known as Chili-saltpetre, or soda-nitre, is collected on
a desert tract in Chili and Peru, and forms a valuable article of
export from those countries. These two substances only differ
from each other in this, —that the first contains the metal potasE2




52                PREPARATION OF NITRIC ACID.
sium, the second the very similar metal sodium, in either case
combined with definite proportions of the elements nitrogen and
oxygen.  By the reaction of sulphuric acid (oil of vitriol) on
either of these two substances, nitric acid is obtained. On a
small scale, for laboratory purposes, saltpetre, or, as it is called
in chemistry, nitrate of potassium, is generally employed; on a
manufacturing scale, soda-saltpetre or nitrate of sodium is used,
because this salt costs less than nitrate of potassium, and also
contains a larger proportion of nitric acid, which it yields up with
greater facility.
Exp. 32. —Into a tubulated, glass-stoppered retort of 250 e. c. capacity put 40 grammes of powdered nitrate of potassium, or, better,
34 grammes of powdered nitrate of sodium if it can be obtained, and
through the tubulature pour 50 grammes of strong sulphuric acid,
which has been weighed out in a bottle previously counterpoised upon
the balance with shot or coarse sand. Imbed the bottom of the retort
in sand, contained in a small iron pan placed over the gas-lamp on a
ring of the iron-stand. Thrust the neck of the retort into the receiver
with two tubulatures; the retort-neck should fit the tubulature of the
receiver with tolerable accuracy. The second tubulature of the receiver should be left open, or loosely covered with a bit of glass, in
order to avoid the possibility of any pressure being created within the
retort during the operation. Place the receiver in a pan of cold water,
and cover it with cloth or bibulous paper, which must be kept consta.ntly wet during the distillation. (See Fig. 16, p. 34.) Apply a moderate heat to the sand-bath; reddish vapors will appear for a moment,
then disappear, and a yellowish fuming liquid will begin to condense
in the receiver. Towards the end of the operation the red vapors
reappear; after this has happened, and the saline matter in the retort
has attained a state of tranquil fusion, while very little liquid passes
over into the receiver, the lamp may be put out, for the process is
finished.
The very acid, corrosive, and poisonous liquid in the receiver is
nitric acid; its faint color is not its own, but is due to the presence of
a compound of nitrogen and oxygen shortly to be described. Transfer
the liquid to a glass-stoppered bottle, and keep it for future use. In
all manipulations with nitric acid it is necessary to avoid getting it
upon the skin, since it produces rather permanent yellow stains.
As the retort cools, the residue solidifies into a white, saline mass,
which must be dissolved out of the vessel by heating it with water after
the apparatus has become thoroughly cold. It will be observed that




ACID AND ALKALINE REACTIONS.                53
the liquid sulphuric acid which was used has disappeared, though the
saline residue is still intensely acid.
64. Pure nitric acid is colorless, and is about half as heavy
again as water. It may be mixed with water in all proportions.
Exp. 33.-To one-third of the nitric acid obtained in the last experiment add an equal bulk of water. The solution thus obtained will be
still intensely acid, as may be proved by its action on vegetable colors.
Litmus is a blue coloring-matter, prepared from various lichens, and
used in dyeing. Unsized paper, colored with a solution of litmus in
water, is a convenient test for many acids, which, as a rule, change
the color ok the paper from blue to red. If the acidity of this diluted
nitric acid be now destroyed or neutralized by the addition of some
other substance of opposite quality, the point at which the liquid
ceases to be acid may be determined by observing when the blue paper
remains blue on immersion in the liquid.
To the diluted nitric acid, placed in an evaporating-dish, add cautiouslv ammonia-water (the Liquor Ammoniae of the apothecaries),
which has been previously diluted with its own bulk of water, until
the liquid no longer turns the litmus-paper red. The ammonia-water
must be added, slowly at first, and at last drop by drop, and the mixture must be constantly stirred with a glass rod. The ammonia-water
has the property of turning litmus-paper, which has been reddened by
an acid, back again to blue, as direct experiment may prove; this
property is possessed by a class of bodies called alkalies, and this reaction with litmus is termed the alkaline reaction, in contradistinction to
the change from blue to red, which is the reaction characteristic of
acids.  When mixed with nitric acid, ammonia-water produces a
compound which, when fresh and pure, has no action on vegetable
colors, and being soluble in water is not visible at this stage of the
experiment. Place the evaporating-dish on the wire gauze over the
gas-lamp, and evaporate the liquid, taking care to avoid actual ebullition, until a drop of the solution taken out on a glass rod becomes
almost solid on cooling. Extinguish the lamp, let the dish become
perfectly cold, separate the semitransparent crystals which have
formed during the cooling from the fluid, if any, which remains in the
dish, allow them to drain, dry them by gentle pressure between folds
of bibulous paper, and reserve them for use in the next experiment.
Besides water, these crystals are the sole product of the reaction;
they must therefore contain both all of the nitric acid and all of the
ammonia which is not water. The chemical name of the substance is
nitrate of ammonium.




54                  MAKING NITROUS OXIDE.
65. Wle find in this experiment a striking illustration of what
is meant by chemical combination. From two fuming liquids,
of very intense but opposite properties, there has come forth a
neutral solid, as unlike its constituents in taste, smell, and all
physical attributes, as could well be imagined. The idea of neutralization, well exemplified in this experiment, has in the history
of chemistry been very fruitful both of names and hypotheses.
Bearing  in mind the fiery constituents of the cooling, harmless
salt which we have thus synthetically prepared, we proceed to
learn by experiment whatever its decomposition may teach.
E'rp. 34. —Introduce into a Florence oil-flask, or other suitable flask
of thin glass, and of about 300 c. c. capacity, the nitrate of ammonium
obtained in the last experiment. From the mouth of the flask, placed
upon the wire gauze on the iron-stand, carry a gas delivery-tube,
No. 6. beneath the saucer in the water-pan; but interrupt the tube, at
some convenient point to interpose (by means of a cork or caoutchouc
stopper with two holes) a small bottle which can be kept cool with
water, as shown in the figure.
Fig. 23.
Heat the flask moderately and cautiously, to avoid breaking it.
The nitrate of ammonium will melt, and little bubbles will soon begin
to escape from the fused mass. The heat must now be so controlled
that the evolution of the gas shall not be tumultuous. The gas is to
be collected in bottles of 300 to 400 c. c. capacity. In the cooled
bottle, through which the gas passes, a clear and colorless liquid will
condense, which will be found on examination, if the process has been
successfully conducted, to be neither acid nor alkaline, to have neither
taste nor smell, and to be wholly volatile when heated on platinum




COMPOSITION OF NITROUS OXIDE.             55
foil,-in short, to possess all the properties of watcr, and none other.
When two bottles of gas have been filled, and enough water for testing
condensed, the delive-cry-tube may be withdrawn from the water and the
lamp extinguished, although the nitrate of ammonium be not all decomposed. The nitrate of ammonium may be entirely resolved into water
and the gas which now awaits examination, but it is difficult to push
the decomposition to actual completion without breaking the flask in
which the operation is performed. That the nitrate leaves no residue
behind, when sufficiently heated, may be proved by heating a crystal
of it on platinum foil over the gas-lamp.
66. Nitrous Oxide. —It is obvious that the colorless and transparent gas, which is the most voluminous product of the decomposition of nitrate of ammonium just accomplished, must contain
all the elements, besides those of water, which enter into the
composition of nitrate of ammonium, and therefore of its constituents, nitric acid and ammonia-water; much interest, therefore,
attaches to the determination of the composition of this gas..Exp. 35.-Insert a glowing splinter of wood into a bottle of the gas.
It will reinflame with almost as much energy as in oxygen.
If oxygen be really a constituent of this gas, it may be possible to
mix the gas with hydrogen, and effect the chemical combination into
water of the oxygen in the gas and the added hydrogen, by heating the
mixture, or passing through it an electric spark (see pp. 24, 26). If
just enough hydrogen can be added to exactly convert all the oxygen
contained in a given volume of the gas into water, the constituents
other than oxygen will be left behind for separate examination, and
we shall have determined how much oxygen a given volume of the gas
contains by observing how much hydrogen has been necessary to convert it into water, the volumetric composition of water being already
known. Now it has been found that when any volume of this gas is
mixed with an equal volume of hydrogen, in a strong tube, provided
with platinum points, like those of the U-tube already used (p. 25),
and the mixture is fired by the electric spark, a violent explosion takes
place, a dew of water condenses upon the walls of the tube, and there
remains a volume of colorless gas, precise' yequal to that with which
the hydrogen was originally mixed. On studying the properties of
this residual gas, it is found to be tasteless, odorless, a little lighter
than air, and to be neither inflammable, nor yet a supporter of eombustion; it is recognized as the pure element, nitrogen.
It follows from this experiment, and the knowledge of the
composition of water previously gained, that any measure of the




56              ATOMIC WEIGHT OF NITROGEN.
gas obtained from nitrate of ammonium contains its own measure
of nitrogen and half that measure of oxygen. The constitution of
this gas is strictly analogous to that of steam; as two volumes of
hydrogen and one volume of oxygen are compacted into two
volumes of steam, so two volumes of nitrogen and one volume of
oxygen are condensed into two volumes of this transparent gas.
As the chemical formula or symbol of water is H,O, so the
formula of this new gas is N20, and its volumetric composition
may be represented by a diagram
similar to that by which we con-    N NI
veyed to the eye the composition
of water.                                — L 2     N.Jf
67. As has been already said,  
a combination of oxygen with  L —  )
another element is called an oxide; the name of the second element is given either by an adjective which precedes the word
oxide, as in the case of this gas N(), whose name is nitrous oxide,
or by connecting the name of the second element with the word
oxide by the preposition of, as in case of oxide of iron.
From the above composition by volume, and from the known
specific gravities of nitrogen and oxygen, the composition of
nitrous oxide by weight is readily deduced. The specific gravity
of nitrogen, referred to hydrogen, is 14; that of oxygen 16;
since there are two volumes of nitrogen for eaoh volume of
oxygen, the two elements must, in any given weight of the gas,
be combined in the proportion of 28 parts by weight of nitrogen
to 16 of oxygen. The molecule of nitrous oxide, N20, must be
composed, like any other quantity of the gas, of 28 parts by
weight of nitrogen and 16 of oxygen; but precisely as in the case
of water, we conceive of the molecule as made up of two atoms
of nitrogen and one atom of oxygen, and we have already learned
that if the atomic weight of hydrogen be represented by 1, that of
oxygen must be 16. It follows, from the constitution of nitrous
oxide, that, if 16 represent the smallest proportional weight of
oxygen which exists in combination, 14 must be the corresponding
smallest weight of nitrogen when thus united with oxygen. Nitrous oxide contains 16-44ths, or 36'36 per cent. of oxygen.
68. Nitrous oxide is almost without odor, but has a distinctly




PROPERTIES OF NITROUS OXIDE.             57
sweet taste; its specific gravity referred to hydrogen is 22; it is
quite soluble in water, which at 0~ dissolves more than its own
volume of the gas, and more than half its volume at 24~. Owing
to this solubility there is a trifling loss incurred by collecting it
in the usual manner over water; this loss may be partially
avoided by using warm water in the pan. Nitrous oxide may be
obtained in the liquid state by submitting it to a mechanical
pressure of about 30 atmospheres in an apparatus cooled to 0~.
The liquid is very mobile, boils at -88s, and crystallizes at
about - 1000 when allowed to evaporate spontaneously under the
exhausted receiver of an air-pump. It is noteworthy that the
permanent gases, such as nitrogen, oxygen, and hydrogen, are
but slightly soluble in water, while all the soluble gases are liquifiable, and often the more readily liquifiable in proportion to their
solubility. A drop of liquid nitrous oxide blisters the skin like a
hot iron. By mixing the solid, snow-like nitrous oxide with the
volatile liquid called bisulphide of carbon, and evaporating the
mixture in a vacuum, the lowest temperature which has hitherto
been attained is produced; it is estimated at -140~.
Small animals immersed in gaseous nitrous oxide die after
some time, but it may be respired for a few minutes with entire
impunity by the healthy human being. The physiological effects
of this gas, when respired, vary somewhat, according to the
quality of the gas and the mode of administration; sometimes it
produces a lively intoxication, attended with a disposition to
muscular exertion and violent laughter, whence its trivial name
of laughing-gas; sometimes, on the contrary, it produces a complete insensibility, during which surgical operations may be performed without pain. When intended for respiration, great attention should be paid to the purity of the gas; carefully prepared
and judiciously administered, it is advantageously used as an anaesthetic agent, especially for operations lasting but a few seconds.
As the gas contains nearly twice as much oxygen as atmospheric
air, it does not seem strange that it should make common combustibles burn with great intensity; it forms explosive mixtures
with inany inflammable gases; it causes glowing charcoal to burst
into flame, and sulphur and phosphorus burn in the gas with
great brilliancy, if well on fire when immersed in it.




58               PREPARATION OF NITRIC OXIDE.
EiP. 36.-Place a bit of sulphur in a deflagTating spoon, and ignite
it with the least possible application of heat; then thrust it into a
bottle of nitrous oxide. It will be extinguished. Yet it would continue to burn in the air. Heat the sulphur much hotter, and again
introduce it into the bottle of nitrous oxide. It will now burn far
more brilliantly than in the air.
The air is not a chemical compound, but only a mechanical mixture
of nitrogen and oxygen; so that a body burning in the air has only to
take oxygen, which is perfectly free to join it. The case is entirely
different when the substance at whose expense the oxygen is furnished
is a chemical compound; to dismember the compound will require a
force superior to that which binds its elements together. Before the
sulphur, in this experiment, can unite with oxygen, it must detach
the oxygen from the nitrogen with which it is combined. To accomplish this, the sulphur must be hotter than it need be for simple burning
in the air. We shall soon learn that there are many chemical compounds, much richer in oxygen than either the air or nitrous oxide,
which nevertheless cannot support combustion at all, in the ordinarv
sense of the term, and this simply because the common combustibles
are quite unable to detach the oxygen from the elements with which
it is already combined.
69. Nitric Oxide. -The nitrous oxide which we have thus
studied, is a derivative of nitric acid, or, more exactly, of the
compound of nitric acid with ammonia-water, but it is only one
of several such derivatives. We proceed to investigate another
substance still more directly obtained from nitric aoid.
Exp. 37.-Place some copper
turnings or filings in a bottle arranged precisely as for generating hydrogen (see Exp. 19), and
pour upon them one half of the
nitric acid still remaining from
Exp. 32, previously diluting this
portion of acid with twice its
bulk of water. Brisk action will
immediately occur. The bottle
becomes filled with red fumes,
but when the gas disengaged is
collected over water, it is found
to be colorless. Collect four bottles, of 300 to 400 c. c. capacity,




NITRIC OXIDE.                     59
of this gas. Save the blue solution (nitrate of copper) which remains
in the generator for future use.
Exp. 38.-Dip a lighted candle into a bottle of the gas. It goes
out. Into the same bottle thrust a glowing splinter of wood. It will
not inflame.
Exp. 39. Lift a bottle of the gas from the water so that the air may
enter the bottle, and the gas may escape into the air. Red fumes, of
very disagreeable smell, and very irritating when inhaled, are abundantly produced. Bring in contact with these fumes a piece of moistened litmus-paper. It will be immediately reddened.
Exp. 40.-Thoroughly ignite a bit of sulphur in a deflagratingspoon, and introduce it into a bottle of the gas. It will not burn.
Exp. 41.-Into the same bottle thrust a piece of phosphorus as big
as a pea, burning actively. The combustion will be continued with
great brilliancy.
70. The new gas is transparent and colorless, and that it is
sparingly soluble in water may be inferred from the fact that
bottles of it may stand indefinitely over water without appreciable loss. It differs from nitrous oxide, and from all the other
gases thus far studied, in its relation to combustibles. The commonest combustibles will not burn in it at all; phosphorus may
be melted in the gas without inflaming; but when its combustion
is once started, phosphorus burns with a vividness which recalls
its burning in oxygen.
When the gas touches the air, a new compound, red, acid, and
irritating, is immediately produced; the question arises, Is it
the nitrogen or the oxygen of the air which gives rise to this
new combination? Experiment would answer this question in
favor of oxygen. If into a bottle of this new gas nitrogen were
introduced, the result would be simply negative; no visible
change and no chemical combination would take place. The
introduction of oxygen into a bottle of the gas would, on the
other hand, produce the red vapors in question, only more
vividly than air, because dilution with the inert nitrogen of the
air would have been avoided. So visible and trustworthy is this
reaction, that the gas we are studying may be used to exhibit
the presence of free oxygen in gaseous mixtures.  For example,
both oxygen and nitrous oxide reinflame a glowing splinter, and
we cannot distinguish between these two gases by this test; but




s'o               ANALYSIS OF NITRIC OXIDE.
the gas we are now studying supplies us with a means of discrimination, since it produces no red fumes with nitrous oxide.
If a sufficient quantity of the metal potassium be heated in the dry
gas till it burns, and the experiment be so executed as to allow the
volume of gas to be measured both before and after the combustion, it
will be found that one half of the volume of gas used has disappeared,
and that the half which remains possesses few or none of the qualities
of the original gas; a slight examination would convince us that we
had set free the well-known element nitrogen. The other half of the
original volume of gas has united with the potassium to form a body
which we shall hereafter be familiar with under the name of oxide of
potassium. The other half of the gas is then oxygen, and we have
found the elements which enter into the composition of this gas, and
the volumetric proportions in which they are united. One volume of
nitrogen is combined with one volume of oxygen to form two volumes
of the compound gas.
We meet here the first case of chemical combination between
two gases unattended by any condensation of the ingredients.
The molecule of the gas will be represented by the formula NO,
and its elements are united bv weight in the proportion of 14
parts of nitrogen to 16 of oxygen, because equal volumes of
nitrogen and oxygen weigh respectively 14 and 16 times as
much as the same volume of hydrogen. The gas is another oxide
of nitrogen, and is distinguished by the name nitric oxide.
When there are two or more oxides of one element, the termination ous implies less oxygen than the termination ic, as in this
case; nitrous oxide contains half as much oxygen for its nitrogen
as nitric oxide.
Nitric oxide is one of the permanent gases; it has never been
liquified. It is a very stable compound, and if perfectly dry is
not decomposed by a red heat or by the action of electric sparks.
Owing to its rapid union with oxygen and the formation of acid
products, its taste, smell, and respirability cannot be ascertaihed.
Bearing in mind the fact that certain red acid fumes, the like of
which we remember to have seen in making nitric acid, are
formed by adding oxygen to nitric oxide, we proceed to a further
study of still other oxides of nitrogen.
71. Hyponitric Acid.-When a mixture of two volumes of
nitric oxide and one volume of oxygen, thoroughly stirred to



rYPONITRIC ACID.                    61
getber and perfectly dried, is submitted to the action of a freezing
mixture of salt and ice, transparent, colorless crystals are condensed from the mixed gases; but if the least trace of moisture
has been present, the product will be an almost colorless liquid.
The vapor of this new substance has a brownish red color; from
a mixture of two measures of nitric oxide and one measure of
oxygen, two measures of the new vapor are produced. Since
two measures of nitric oxide contain one measure of nitrogen and
one of oxygen, the composition of the new substance may be
represented by the accompanying diagram; the molecule will be
represented by the formula
NO2, and the composition of              7  
the substance by weight will           I           -— I _
be 14 parts of nitrogen to 32   N             -    N  2
of oxygen. The name of this           I     ]
new body is Hyponitric acid,
a name derived from  nitric acid by prefixing the Greek VITo,
"below."  The term  is used to indicate that the substance to
which it is applied contains less oxygen than the other substance
from which the name is derived. Thus hyponitric acid contains
less oxygen than nitric acid, hyposulphurous acid less than sulphurous, and so forth. It remains to justify- this assertion respecting the comparative oxygen-contents of hyponitric and
nitric acids.
Elxp. 42.-Add to the nitric acid which remains from Exp. 32, previously diluted with twice its bulk of water and warmed over the
lamp, finely powdered litharge in small portions so long as it readily
dissolves. The operation may be best performed in an evaporatingdish, which should be only very moderately heated. The substance
sold under the name of litharge is a simple combination of the two
elements, lead and oxygen; the formula of its molecule is PbO, in
which Pb represents the least proportional weight of metallic lead
(plumbum); its composition by weight is accurately known. When
the litharge no longer dissolves with promptness, no more should be
added, and the liquid in the dish should be evaporated to dryness, at
first on the wire gauze over the lamp, but towards the end ci the
operation on a water-bath.
During this evaporation there escape into the air unchanged, water
and the excess of nitric acid which was not neutralized by the oxide




62                     HYPONITRIC ACID.
of lead. There remains a white, saline substance, which has resulted
from the combination of the oxide of lead with a portion of the nitric
acid; it is called nitrate of lead. Experience has proved that it is a perfectly dry substance, containing no water whatever; it is the raw material, so to speak, of important experiments shortly to be described.
EBp. 43.-Heat one or two teaspoonfuls of the nitrate of lead, of
Exp. 42, in an ignition-tube, and observe that the compound decomposes; deep red fumes are produced, and oxide of lead (litharge) is
left behind. If these red vapors be carried through a U-tube immersed
in a mixture of ice and salt, a portion of the vapors condense into a
liquid, while an uncondensable gas passes through the cold tube, and
may be collected by a suitable arrangement beyond the U-tube. The
condensed liquid is hyponitric acid; the gas is oxygen. If the complete absence of moisture has been secured, crystals of hyponitric acid
may often be obtained by replacing the first U-tube by a second towards the end of the distillation. In fact, this is by far the most convenient method of preparing hyponitric acid. Into the U-tube containing the liquid hyponitric acid drop a bit oI ice or a little snow;
the color of the solution changes from reddish to a greenish blue, and
two layers become visible in the liquid. The explanation of this
change will be found in the next section.
The name of hyponitric acid is now justified; for it is apparent
that when the litharge was dissolved in nitric acid, it found there,
and combined with, an oxide of nitrogen containing more oxygen
than hyponitric acid does, since when this oxide of nitrogen is
decomposed by heat, as in the experiment just described, hyponitric acid and oxygen are evolved, the litharge remaining behind
unaltered. Reserving the further discussion of the composition
of nitric acid, we may here speak of the properties and products
of hyponitric acid.
72. Hyponitric acid occurs in the solid, liquid, and gaseous
states; at — 9~ it crystallizes; above that temperature and below
220 it is a mobile liquid of sp. gr. 1'451, and of various colors at
various temperatures; it boils at 22~. The liquid acid gives off
red, acid, irrespirable vapors at the ordinary temperature. If to
liquid hyponitric acid, cooled by ice and salt, a proportionally
small quantity of ice-water be added, two layers of liquid are
formed, as has been above illustrated, the upper and least-colored
of which consists principally of nitric acid, the lower and darker
of a fluid which yields, on cautious distillation at a low tempe



ANALYSIS OF NITRIC ACID.                63
rature, a very volatile blue liquid.  This blue liquid is so unstable
that its composition and qualities are not certainly known; but
it is supposed to be, or to contain, an oxide of nitrogen differing
from all those heretofore studied, —an oxide capable of direct
derivation from nitric oxide by adding, to any volume of this last,
one-fourth that volume of pure oxygen, and therefore answering
to the formula N20,, and being composed of 28 parts by weight
of nitrogen and 48 parts of oxygen. This obscure substance is
known by the name of nitrous acid; though itself but imperfectly known, some of its compounds are not unfamiliar bodies.
73. We have already learned that the white, saline substance
called nitrate of lead can be made from oxide of lead and nitric
acid, that it contains no water in its composition, and that when
heated it is decomposed into oxygen, an oxide of nitrogen, and
the original oxide of lead. On the basis of these facts, a method
has been constructed of determining the composition of nitric acid
by weight.
To 10 grammes of oxide of lead add something more than enough
nitric acid to transform it completely into nitrate of lead; evaporate
the excess of acid, dry the residual nitrate of lead completely, and
weigh it. If W  represent this weight in grammes, W-10 is the
weight of the unknown oxide of nitrogen which has combined with 10
grammes of oxide of lead to form W grammes of nitrate of lead. From
these data the percentage composition of nitrate of lead can be calculated; 10 grammes of pure nitrate of lead invariably contain
Oxide of lead.... 6'738 grammes.
Oxide of nitrogen.. 3262,,
If now 10 grammes of pure nitrate of lead be decomposed by heat
under such conditions that the vapors which it yields shall pass over
some substance capable of abstracting all the oxygen from the vapors
without affecting the nitrogen, it will be possible to collect all the
nitrogen contained in 3-262 grammes of the unknown oxide of nitrogen, and so to determine its volume and thence its weight.
This determination is actually made by heating the nitrate of lead in
a long glass tube, in such a manner that all the vapors evolved from it
pass over a large surface of red-hot copper. The copper absorbs the
oxygen, the nitrogen passes over it unaltered and is collected in a suitable vessel and accurately measured. As the specific gravity of nitrogen at any given temperature and pressure is known with precision,
the exact weight of the nitrogen can be deduced from this volume.




64               COMPOSITION OF NITRIC ACID.
Such experiments often repeated have led to the conclusion that the
oxide of nitrogen, which with oxide of lead makes up nitrate of lead,
contains
Nitrogen..,. 25'93 per cent.
Oxygen...... 7407,,
If these two numbers be divided by the specific gravities, or unitvolume-weights, of the two gases respectively, the quotients will
express the composition by measure of the oxide of nitrogen; the
quotients are 1 8t62 and 4'63, numbers which are to each other as 1 to
2X5 or, avoiding fractions, as 2 to 5.
The molecule of this compound is therefore considered to contain
two volumes of nitrogen and five volumes of oxygen, and is represented
by the formula N.2O,; the weight of this molecule will be 108, of
which 28 parts will be nitrogen and 80 oxygen,-a composition precisely corresponding, of course, to the percentage composition above
given. This new oxide of nitro)en may be called for the present nitric
acid, though we shall soon learn to distinguish between this body and
common nitric acid.
74. The combining proportions of nitrogen and oxygen in lh-lis
oxide have been experimentally determined by weight, and not
by volume, as was the case with all the preceding oxidcs of
nitrogen.  The reason of this different treatment is to be found
in the fact that nitric acid cannot be converted into vapor without suffering decomposition-not indeed into its elements, but
into oxygen and a lower oxide of nitrogen. It is therefore, in
the present state of science, im')ossible to obtain a volume of
nitric acid gas capable of experimental resolution into nitrogen
and oxygen.
Since a large majority of the elementary bodies are non-volatile
ander any treatment in our power to employ, and since the greater
number of chemical compounds either are non-volatile, or are,
like nitric acid, decomposed by a temperature high enough to
volatilize them, the proportions in which the elements unite by
weight are of much more general value than the proportions in
which they unite by volume; and the methods of determining the
atomic weights of the elements have been studied with a thoroughness, and brought to a perfection commensurate with the fundamental importance of these proportional numbers.
75. Nitric acid completes the series of oxides of nitrogen; no




OXIDES OF NITROGEN.                65
higher oxide is known. We are now prepared to exhibit this
series in a diagram which shall present to the eye at once the
volumetric composition of the oxides, the resultant volumes after
the condensation of the ingredients, the atomic weights of the
elements, and the combining weights of the compounds. Since
the atomic weights of the gaseous elements are at the same time
their specific gravities referred to hydrogen, it will be easy to
deduce the specific gravities, or equal-volume weights, of the
compound gases, N20, NO, and NO2, from their combining weights
by dividing these weights by two. As the resultant volumes
after condensation are not known for the two members of this
series N203 and N205, we abstain from  figuring hypothetical
volumes which analogy may point to as probable, but which experiment has never demonstrated as fact.
THE OXIDES OF NITROGEN.
Nitrous Oxide.                     Nitric Oxide.
j+ =13  +m=|NOj
14,
Nitrous Acid.                    Hyponitric Acid.
_                  N
Nitric Acid.
L I JKIiW                         108




66              LAW OF MULTIPLE PROPORTIONS.
76. These five bodies are all chemical compounds; they are
definite and constant in composition; and all differ essentially
from their elementary constituents and from each other, as the
experiments we have performed with them have abundantly demonstrated. It is, therefore, obvious that two of the elements
are capable of combining in several proportions to form definite
chemical compounds; and what is here proved of two of the
(lements we shall hereafter find to be true of all, though not of
every couple; so that the series of oxides of nitrogen is but one
illustration of a most comprehensive law. The difference between
a mechanical mixture and a chemical compound does not on this
account become less marked. The possible mixtures of nitrogen
with oxygen are innumerable; the known combinations of these
two elements are only five,-two volumes of nitrogen combining
chemically either with one, two, three, four, or five volurfies of
oxygen, and with no other proportions whatsoever.  As for
volumes, so for weights; the proportional weight of oxygen in
these oxides rises by definite leaps from the first member of the
series to the last.
This definite, step by step mode of forming chemical compounds is one of the most characteristic, as it is one of the most
general facts of chemistry; no other science offers a parallel to
it; but long experience and patient labor with the balance and
measuring-glass have established it as the habitual mode in which
the force called chemical ordinarily acts.  The abstract results
of observation and experiment may be expressed in the following
proposition, often called the Law of Multiple Proportions: If two
bodies combine in more than one piroportion, the ratios in which
they combine in the second, third, an(d subsequent compounds are
definite multiples of those in which they comnbine to form the first.
While the mode of action of the chemical force set forth in
this proposition is that which has long been uppermost in the
minds of chemists, most prominent in chemical treatises, and
perhaps most important to the progress of the science in the
direction in,which it has thus far been cultivated, it should be
remembered that in the phenomena of solution, in the formation
of metallic alloys by fusion,.and in the crystallization of minerals
and other substances with,cQstant forms but variable composi



ATR A MIXTVRE.                    6
tion, the chemical force has a part to play which, if more obscure
than its ordinary manifestations, is not less real.
77. Air a mixture.-All that has been said of the distinction
between a mechanical mixture and a chemical combination finds
perfect illustration in the differences between air and the definite
oxides of nitrogen which have just been studied. Air is not, like
these, a chemical combination. The evidence that it is a mechanical mixture merely may here be appropriately presented.
The statement that air contains 1 volume of oxygen to 4 volumes
of nitrogen is not absolutely true. These proportions are never
actually found; the gases are not combined in any simple ratio
either by volume or weight; they are always mixed in the proportion of 20-81 measures of oxygen to 79'19 measures of nitrogen, or 23-10 parts by weight of oxygen to 76'90 parts of nitrogen.
The experimental processes by which these numbers have been
fixed are so perfect, that it is impossible to entertain the idea
that the gases are really mixed in the ratio of 20 measures to 80,
or I volume to 4 volumes, or the proportion of 20 parts by weight
of oxygen to 70 of nitrogen, as the formula N40 would require.
When 20'81 parts of oxygen are mixed with 79'19 of nitrogen,
there is no development either of light, heat, or electricity, such
as usually attends the formation of a chemical compound; and
the specific gravity, magnetism, and refractive power of the mixture are such as calculation would directly deduce from the numbers expressing these properties for the two constituents.
We have seen that when nitric oxide is added to nitrous oxide
no red fumes are produced (~ 70), but that when the nitric oxide
is brought in contact with air these suffocating fumes are abundantly formed, though the air contains only half as much oxygen
as the nitrous oxide. These experiments go to show that while
in nitrous oxide the oxygen is held in chemical combination, in
air it is free.
Strong positive evidence that air is a mere mixture is afforded
by its behavior towards water. All gases are soluble in water
to a greater or less extent, each one dissolving in a certain fixed
and definite proportion at any given temperature and pressure.
Thus at 15~ and a pressure of 76 c.m. of mercury, 1 volume of
water dissolves 0'0193 volume of hydrogen and 0'7778 volume
F 2




68                      AIR A MIXTURE.
of nitrous oxide. As a general rule, the pressure to which the
liquid is exposed being constant, the quantity of gas dissolved by
water is less in proportion as the temperature of the water is
high; in many cases boiling water is altogether incapable of retaining gases in solution, particularly those which are only sparingly soluble at any temperature.
Hence, by prolonged boiling, many gases can be completely
expelled from the water which held them in solution. For example, if a solution of nitrous oxide be boiled, and the gas be
collected as it escapes, this gas will be found to exhibit the characteristic properties of nitrous oxide. The nitrous oxide has not
been altered by the act of solution. It still remains a definite
chemical compound as before.  But a result very different from
this is obtained upon boiling water which has become charged
with the ingredients of atmospheric air. The gas collected in this
case is composed of oxygen and nitrogen, it is true, but not in
those proportions in which the elements are united in air.
Water does not dissolve air directlv as such, as it should do
were air a chemical compound; but it dissolves out from it a
quantity of oxygen, just as if no nitrogen were present; at the
same time it dissolves nitrogen in accordance with the solubility
of this element, and to precisely the same extent that it would
absorb it if there were no oxygen in the air. Oxygen is dissolved by water in larger proportions than nitrogen; 1 volume
of water at 15~ and under a pressure of 76 c.m. of mercury
dissolves 0.02989 volume of oxygen, but only 0 0148 volume
of nitrogen.
Exp. 44.-By means of a sound perforated cork or caoutchouc
stopper, adapt to a flask of the capacity of 1 or 2 litres a gas-delivery
tube, No. 6, long enough to reach to the water-pan in the usual way.
Upon the outer end of the delivery-tube tie a short piece of caoutchouc
tubing, to which a stopper, made of a bit of glass rod or a wooden plug,
has been fitted. Fill the flask completely with ordinary well or river
water; fill also the delivery-tube with water, and close it by putting
the stopper in the caoutchouc tube. Carefully place the cork of the
delivery-tube in the neck of the flask in such manner that no air shall
be entangled by the cork; at the same moment remove the plug from
the delivery-tube, and finally press the cork firmly into the flask. Both
flask and tube will now be completely full of water. Place the dried




NITRIC A CID.                     69
flask upon a ring of the iron stand, and invert a small bottle filled with
water over the end of the delivery-tube. Then slowly bring the contents of the flask to boiling.
As the water gradually becomes warm, numerous little bubbles of
gas will be seen to separate from the liquid and to collect upon the
sides of the flask; these subsequently coalesce to larger bubbles, which
collect in the neck of the flask. As soon as the water actually boils,
the steam will force this air out of the flask, and it will collect in the
inverted bottle at the end of the delivery-tube, the steam being meanwhile condensed as fast as it comes in contact with the cold water in
the pan. By continuing to boil moderately during ten or fifteen minutes, nearly all the air can be swept out from the flask by means of
the escaping steam. The gas delivery-tube may then be lifted from
the water-pan and the lamp extinguished.
From a litre of ordinary water about 50 c. c. of gas can usually
be obtained. This contains, of course, besides oxygen and nitrogen, a certain amount of carbonic acid; but careful analyses
have shown that it is much richer in oxygen than ordinary air,
the proportion of oxygen to nitrogen in the gas from water being
as 32 to 68 in 100 volumes, instead of 20S81 to 79'19 as in air.
78. Nitric Acid.-The nitric acid above referred to (~ 73),
whose composition is represented by the formula N205, is anl unstable solid which melts at 30~; the liquid produced boils at 470
and is decomposed at 800. The crystals of this substance are
transparent and colorless; they undergo spontaneous decomposition into hyponitric acid and oxygen, even when preserved in
closed tubes. It is obvious that this is not the common nitric
acid with which we are already familiar. It remains to demonstrate that commercial nitric acid contains water and the oxide of
nitrogen N205.  A single experiment may be made use of to
demonstrate the presence of water in common nitric acid, and at
the same time to determine the proportion in which it enters into
the composition of the sample of acid examined.
By adding to a known weight of common nitric acid (10 grammes,
for example) a weighed quantity of the oxide of lead much larger than
the acid is able to dissolve (100 grammes, for instance), and then
heating the mixture, with suitable precautions against over-heating, in
a weighed flask, the vapor of water, and nothing else, is given off;
this vapor may of course be condensed, and proved to be common
water. Since the oxide of lead contains no water, and the nitrate of




70                  HYDRATED NITRIC ACID.
lead is also anhydrous, it follows that whatever water escapes from the
flask during the heating must be derived from the 10 grammes of nitric
acid, and, further, that if the heating be long enough maintained, all
the water which the acid contained will be expelled. By weighing the
flask and its contents after all water has been thus driven out, the loss
of weight will be the quantity of water contained in the 10 grammes
of acid.
It has been already proved (Exp. 42) that the nitrate of lead,
which, with a quantity of unused litharge, constitutes the residue
in the flask, yields the oxide of nitrogen N205.  If this oxide of
nitrogen be called nitric acid, then is the common acid hydrated
nitric acid; but if the shorter name, nitric acid, be applied to the
commoner substance the commercial acid, then the oxide N2O0
must be distinguished as anhydrous nitric acid.
79. Anhydrous nitric acid unites with water in at least two
definite proportions; its molecule combines with one molecule of
water, or with fozr, to form the two hydrates represented by the
formulae HIO, N20, and 4H20, N20O respectively, wherein the
symbols of one and four molecules of water are simply placed
beside the formula of the molecule of anhydrous nitric acid.
Ponderal analysis has given us this knowledge of the composition
of these two hydrates, by actually weighing the proportion of
water combined with the oxide of nitrogen in the two cases.
When pure nitric acid is spoken of, the acid containing one
combining proportion of water to one of the oxide of nitrogen is
generally referred to. This acid is often called the monohydrated
acid, an adjective which may be applied to any substance which
is coupled with a single molecule of water. The monohydrated
acid is a colorless, transparent, mobile liquid of specific gravity
1'52, which boils at 86~ and freezes at about -50~.  Light
slowly decomposes it, and a very moderate heat resolves it, not
indeed into its elements, but into less complex compounds. It
exerts a highly corrosive action on organic bodies, and stains
tissues containing nitrogen of a bright orange colour. It should
be handled with great care, as it burns the skin like a hot iron.
It absorbs water from the air. When mixed with water, heat is
developed from the mixture, and the second definite hydrate is
formed, 4H20, N205,-a colorless, strongly acid 1;,luid! having s




TrI TERM OXIDIZING AGENT.              71
specific gravity of 1'42, and containing 60 per cent. of anhydrous
nitric acid. To this last hydrate all weaker and stronger acids
are alike converted by boiling. The ordinary nitric acid of commerce has a specific gravity of either 1'40 or 1P42, and therefore
contains either 56 or 60 per cent. of anhydrous nitric acid.
Sometimes its specific gravity is as low as 1'33 or 1'28. What
is known in commerce. as aqua fortis is a still weaker nitric acid
containing some nitrous acid and other impurities.
80. Nitric acid, especially when hot, gives up a part of its
oxygen with great facility to substances capable of combining
with oxygen. When a substance habitually and readily imparts
oxygen to other bodies with which it is brought in contact, it is
called an oxidizing agent; and, on the other hand, a substance
which habitually and readily takes oxygen out of other substances
with which it is brought in contact, is called a redatcing aogent.
When concentrated, nitric acid acts with more energy than when
diluted with water; and when mixed with strong sulphuric acid
(a substance which tends to take water from other compounds),
it becomes an oxidizing agent of intense power. We have seen
liquid nitric acid yield a part of its oxygen to copper with evolution of nitric oxide, a lower member of the series (Exp. 37), and
we have also learned that the vapor of anhydrous nitric acid will
give all its oxygen to red-hot copper, nitrogen being set free
(~ 73). M3ost of the metals are dissolved by nitric acid, with
evolution of one or other of the lower oxides of nitrogen; and
sulphur, phosphorus, arsenic, carbon, and many other less familiar
elements are converted by it into oxides.  Organic substances
are oxidized by nitric acid to very various degrees and with very
various products, according to the strength and temperature of
the acid employed.
81. Combining  Weiqhts of Chemical Compounds. —The atomic
weight of oxygen, or the weight of the least proportional quantity of oxygen which enters into combination, is the same when
it unites with nitrogen as with hydrogen.  It is a general fact,
that each element has but one least combining weight with each
and all of the other elements. The atomic hypothesis is based
upon this important fact.. This hypothesis attributes to the
imagined atom of each element a constant proportional weight.




72                  COMBINING WEIGHTS.
expressed by the same number which experiment proves to be
the combining proportion by weight of the element taken in
finite, ponderable quantities. In this sense the atom of oxygen
is said to be 16 times as heavy as the atom of hydrogen, and the
atom of nitrogen 14 times as heavy as.the atom of hydrogen.
The combining weight of a chemical compound is always equal
to the sum of the atomic weights of the elementary atoms contained in its molecule. Thus the combining weight of water,
HI20, is 18=2+16;. of anhydrous nitric acid, N205, is 108
=2 x 14+5 x16; of monohydrated nitric acid, 1120, N20,  is
126=2+2 x 14+6 x 16.
It must never be forgotten that these combining weights are
not absolute weights, but simply express the projortions by
weight in which the elements and their compounds invariably
unite. The combining weight of a3 compound is directly deduced
from the composition of the molecule, or least quantity of the
compound which can exist by itself uncombined, or take part in
any chemical process. To correctly determine this least proportional quantity often requires a large accumulation of facts, and
a just collation of these facts, such as is possible only when a
wide experience is added to a keen insight into the principles of
chemical philosophy.
82. The discussion of the true molecular formulae of chemical
compounds presents difficulties which render it entirely unsuitable for our present stage of progress; nitric acid, however, will
enable us to illustrate one of the difficulties appertaining to the
subject. When nitric acid dissolves oxide of lead (Exp. 42) the
reaction which occurs may be thus symbolized:PbO   +      2,O, N205  =   PbO, N20    +   H20
Litharge.    Nitric acid.    Nitrate of lead.   Water.
The resulting molecular formulae of nitrate of lead is not divisible
by any number but unity, and therefore cannot be made simpler
and still express the same proportional combination of its elements.
If instead of the oxide of lead we employ the oxide of silver,
we shall find it possible to express the resulting molecule of
nitrate of silver, by two formula, either of which will represent




COMBINING WEIGHT OF NITRIC ACID.            73
correctly the proportions by weight in which the elements have
combined: —
Ag20  +   H20, N20    =   Ag20, N205  +   HO0; or
Oxide of silver.  Nitric acid.    Nitrate of silver.   Water.
AgO   +   H20, N20,   -       2AgN03    +   H20.
Now there is a class of metals which enter into reactions in the
manner of silver in the above equations; and there is another class,
of which lead may be taken as an example; and the difference
between these two classes is put into the form of a definition by
the statement that silver and its analogues are uni-valent or
mono-atomic, and that lead and its analogues are bi-valent or
bi-atomic,-terms which are intended to express the fact that
one atom of any metal of the lead class is capable of uniting with
twice as many atoms of oxygen, chlorine, or any other element
as one atom of any metal of the silver class.
The use of the common algebraic signs in these formule requires no explanation. The sign of equality denotes the equality
of the sum of the atomic weights on either side of it; a numeral
on the left of a group of symbols is intended to multiply the whole
group, unless a comma divides the group, in which case the numeral multiplies that part of the group on the left of the comma;
brackets are sometimes used, as in algebra, to mark the extent
of the multiplication. The formula of nitric acid itself admits of
simpler expression:
1H20, N205 =  H-2N206 =  2HN03.
The first formula reminds us that the nitric acid, of which we
speak, may, by indirect means, be made to yield anhydrous nitric
a id and water; from the second we easily learn the proportions
in which the three elements are united by weight; but the third,
HIN0,, expresses these same proportions with precision, and is
the most concise of the three. Now, although the two formulae
HI2N,06 and HNO, express precisely the same compound of the
same three elements in the same fixed proportions by weight, the
combining weight of nitric acid is 63 if the last formula be correct, and 126 if the first represents the real molecule of the acid.
In the great majority of chemical processes in which nitric acid is
involved, that proportional weight of nitric acid is necessary which




74                 NITRIC-ACID REACTIONS.
is implied by the molecular formula H2N206; but there are not a
few cases in which the proportional weight represented by the
simpler formula HNO3 completely accomplishes the actual reaction, and is capable of representation in the algebraic form.
To illustrate the first class of cases we have the formula, just
given, of the reaction which occurs when oxide of lead is dissolved in nitric acid; a less proportional weight of nitric acid
than H12N206 or 2HiN03 will not answer the conditions of the reaction. Let us draw from the preceding sections some further
illustrations of this class of nitric-acid reactions. When copper
(whose symbol is Cu, from cuprurn) is used to set free nitric
oxide from nitric acid (see Exp. 37), the reaction is symbolized
as follows:3Cu  +   4H2N206  =   2NO   +   3CuN,06  +   4H20.
Copper.   Nitric acid. NTitric oxide. Nitrate of copper. Water.
When nitrate of lead is decomposed by heat into litharge, hyponitric acid, and oxygen, the following equation represents the
chemical change:
PbN206  =   PbO   +   2N02  +   0.
In these two reactions the nitrates of copper and lead contain two
combining weights of nitrogen and six of oxygen for each one of
the metal.
To illustrate the second class of nitric-acid reactions, we have
only to explain more fully the nature of some of the substances
with which we have already experimented. As the nitrate of
lead may be made by bringing together oxide of lead and nitric
acid, water being eliminated, so the nitrates of potassium and
sodium, from which we originally prepared nitric acid (see
Exp. 32), may be formed by an analogous,. though not identical,
reaction. The composition of common caustic potash, or caustic
soda, may be expressed in two ways. Some chemists represent
these substances as consisting of oxide of potassium, or sodium,
and water, and therefore prefer the formula K20,1120; while
other chemists divide this molecular formula by two, and represent caustic potash by the briefer formula KHO, and caustic soda
by the corresponding symbol NaHO. In these formulae K stands
for Kalium, the Latin name of potassium, and Na for Natrium,




NITROGEN AND HYDROGEN.                 75
the Latin name of sodium. If we adopt for the moment these
shorter formula, which of course express precisely the same proportional composition by weight as the longer, the reaction by
which nitrate of potassium, or sodium, may be prepared will be
written as follows:KHO  +   HNO,                     KNO3   +   H20
Caustic potash.   Nitric acid.  Nitrate of potassium.    Water.
NaHO    +       HNO,    =        NaO,3    +    H20
Caustic soda.    Nitric acid.   Nitrate of sodium.   Water.
From this nitrate of potassium, or sodium, nitric acid is prepared
(see Exp. 32) by treating it with sulphuric acid, a substance whose
composition by weight, as we shall hereafter learn, may be correctly expressed by the formula HISO4.  This reaction may be
thus symbolized:KNO3   +    H2SO4    =         KHSO4   +    HNO,
Nitrate ofpotas.  Sulph. acid.  Acid su7ph. ofpotas. Nitric acid.
By substituting Na for K the reaction with nitrate of sodium
would be represented. It thus appears that the molecule of nitric
acid, which will represent in the simplest way its reactions with
caustic potash, will not represent at all its reactions with oxide
of lead, unless two molecules are assumed to enter into every reaction with this latter substance.  The formula H2N206 is the
more comprehensive, because nitrate of potassium can be represented by the formula K N206 as accurately, if not as simply, as
by the formula KNO,.
It is noteworthy that, in the reactions above formulated, one
atom or combining proportional weight of potassium, or sodium,
changes place with one atom of hydrogen, while one, atom of lead
or copper replaces two atoms of hydrogen. These different capacities of the other elements to replace hydrogen are of great importance in chemical philosophy, and will be more fully treated of
hereafter.
83. Nitrogen and Hydrogen.-Ammonia-water, such as we
made use of in Exp. 33, when gently heated, evolves a very
pungent, colorless gas, which now claims our attention.
This gas may readily be prepared as follows: —Fill a flask of 250 to
500 e. c. capacity about half full of the strongest ammonia-water to be




76                         AMMONIA.
had at the druggist's. Close the flask by a cork provided with a
funnel-tube and an exit-tube; carry the delivery-tube to the bottom
of a tall bottle, having a capacity of at least a litre and filled with
fragments of quick-lime.
n                                     Fig. 25.
When the ammonia-water
in the flask is gently boiled,
the gas which passes off
will be deprived of moisture by the quick-lime and
will issue dry from  the
bottle; it may be collected
either over mercury, or by
displacement, as shown in
the figure (Fig. 25). The
gas is so extremely soluble
in water that it cannot be 
collected over the ordinary
water-pan; as it has little,
more than half the density
of atmospheric air, it can
be readily collected by displacement.  When thus collected, the gas
should be allowed to pass into the very loosely corked bottle, until a
piece of turmeric paper, held at the mouth, is immediately turned
brown; the delivery-tube is then withdrawn, and the mouth of the
bottle is tightly closed with a caoutchouc or glass stopper.
The gas thus obtained is transparent and colorless, possesses
an extraordinarily pungent odor which provokes tears, and has an
acrid, alkaline taste.  It will be found to be uninflammable, and
is, of course, irrespirable. It turns red litmus to blue most energetically. Its specific gravity as deduced from actual experiment
is 8'62; a litre of the gas weighs 0'7625 grm. One measure of
water at 0~ dissolves 1049 measures of the gas.
The ready solubility of ammonia gas may be exhibited as follows:
-Fill a stout glass tube, an ignition-tube for example, over mercury
with the gas; grasp the tube by the top, and, holding it upright, dip
its mouth into a vessel of water. The water will rush up the tube, if
the gas be pure, with a force which might break the tube, if too thin.
84. The solution of ammonia exposed to the air, or placed in a
vacuum, or simply boiled, loses all its gas. As its ready solubility in water suggests (compare ~ 68), the liquefaction of the gas




ANALYSIS OF AMMONIA.                   77
is not only possible but easy; the gas becomes a colorless, transparent, mobile liquid at 0~, under a pressure of 41 atmospheres,
or at -40~ at the ordinary pressure. This liquid freezes at
about — 80~. An excellent freezer, applicable on both the large
and small scale,.is now constructed, in which the cold is produced
by the rapid evaporation of liqutefied ammonia-gas.  It is the low
pressure at which ammonia becomes a liquid which renders this
machine possible.
85. But of what is this gas, whose properties are so strikingly
unlike those of any gas previously studied, composed? The
question may be answered by the following experiment:With the exit-tube of                  Fig. 26.
the drying-bottle of an apparatus fitted for the generation of ammonia gas as
shown in the figure (Fig.
26), connect a tube of hard
glass, in which a bulb has
been blown (see Appendix,
~ 4). Thrust into this bulb
a piece of the metal potas-   @1
sium; cause ammonia gas
to flow through the bulb
by heating the contents of
the flask, and then warm
the glass bulb. As soon as
the potassium melts, it becomes covered with a brownish-green film, and a gas begins to escape
which we recognize as hydrogen by lighting it at the mouth of the
tube. The burning gas is certainly not ammonia, for ammonia is not
inflammable, but, as the odor proves, it is mixed with some ammonia
which has escaped decomposition by the potassium.
To separate this ammonia, and collect the pure hydrogen, fill a testtube, about 14 c.m. long, three-quarters full of mercury; pour water
upon the mercury till the tube is full, close the tube with the thumb)
and invert it into a cup of mercury; with the outer end of the bulbtube (Fig. 26) connect a suitable delivery-tube which shall dip into the
cup of mercury and deliver the gas into the test-tube, whose upper
quarter is full of water. The gas must pass through this water, which
frees the hydrogen from intermixed ammonia.




7 8                 ANALYSIS OF AMMONIA.
It is absolutely necessary to use mercury in this experiment, because
if the delivery-tube were allowed to dip into water, the extreme solubility of ammonia in water might cause the water to suck back into
the bulb containing the heated metal, whereupon an explosion would
inevitably ensue.
86. Having thus learned that potassium will set free hydrogen
from ammonia, just as the analogous metal sodium eliminates
hydrogen from water, we shall be inclined to try upon ammonia
the same powerful agent by which we resolved water into its
elements —the galvanic current.
A glass tube (No. 1), 60 to 80 c.m. long, open at one end and closed
at the other, is bent into the form of a V; the closed limb is provided
with a platinum-wire fused into the flass, through which the wire
passes, to terminate near the bend of the V, in  Fig. 27.
a slip of platinum foil. Fill the whole of the
closed limb and nearly half of the open limb of
this tube with ammonia-water, to which has
been added a teaspoonful of a strong solution
of sulphate of ammonium in order to increase
its conducting-power; support the tube as
shown in the figure (Fig. 27), and connect with
the platinum-wire in the closed limb of the
tube the negative or zinc pole of two mediumsized Bunsen cells, at the same time inserting
a platinum-wire and plate, attached to the positive or carbon pole, in the open limb. Gas
quickly collects in the closed limb. Disconnect the battery-wires, fill the open limb with water, close it with the
thumb, and by inclining the tube transfer the gas to the open limb.
On applying a match to the gas, it proves to be inflammable, and we
recognize it without difficulty as hydrogen.
The experiment is now repeated with the electrodes reversed; the
positive pole is connected with the sealed, and the negative with the
open limb. The hydrogen, which is disengaged at the negative pole,
now escapes through the open end of the tube into the air, while a
transparent and colorless gas, previously evolved at the positive pole
in the open limb and consequently lost, is now collected in the sealed
end of the apparatus. The quantity of gas evolved at the positive pole
is comparatively small, but in half an hour enough for examination
will probably have been collected. By the same manipulation as before,




SYNTHESIS OF AMMONIA.                  79
transfer this gas to the open limb, and thrust into it a lighted match.
It is neither inflammable nor will it support combustion, and has in
itself neither taste nor smell; it is the inert nitrogen.
A cheaper and perfectly effective form of the apparatus used in this
experiment may be made by closing one end of a common U-tube with
a good caoutchouc stopper, through which the conducting wire is
thrust.
87. These experiments have conclusively proved hydrogen and
nitrogen to be constituents of ammonia; and it now becomes
desirable to prove synthetically that these two gases are the only
constituents of ammonia; which would be done, if ammonia
could be experimentally produced by the direct union of the two
gases. Unfortunately, no process has been discovered whereby
ammonia can be directly reproduced from free hydrogen and
free nitrogen. The following experiments, however, will demonstrate that ammonia is actually produced from materials which
are known to generate a mixture of hydrogen and nitrogen-or,
more strictly, which are known to be capable of generating both
hydrogen and nitrogen:Exp. 45.-Place in an ignition-tube an intimate mixture of 3
grammes of fine iron filings with 0'2 gramme of caustic potash; adapt
a delivery-tube (No. 7) to the ignition-tube, heat the contents of the
tube over the gas-lamp, and collect the gas which escapes in a testtube over the water-pan. Exanmine this gas, which will prove to be
the inflammable hydrogen.  Caustic potash, as we have already
learned (p. 74), consists of potassium, hydrogen, and oxygen; at a high
temperature, metallic iron is able to seize upon a portion of the oxygen
in this compound, setting free hydrogen, which finds no place in the
new combinations.
Exp. 46.-Heat in a second ignition-tube, similarly disposed, a
mixture of 3 grammes of fine iron filings and 0-2 gramme of nitrate
of potassium, and collect the gas as before, over water. This gas has
neither taste nor smell, and when tested with a lighted splinter it is
found to be uninfiammable, and in fact to extinguish the taper. It is
nitrogen. Nitrate of potassium contains, as has been already stated
(p. 75), potassium, nitrogen, and oxygen; at the high temperature
employed the salt is partially decomposed, the metallic iron combines
with the oxygen of the nitrous vapors formed, and their nitrogen is set
free..Ep. 47. —In a third ignition-tube, heat the same quantities of the




80                   THE NASCENT STATE.
same materials which have been used in the last two experiments, at
once and together. A delivery-tube is not necessary in this case; the
tube may be held in the wooden nippers by the open end. Neither
hydrogen nor nitrogen will be evolved as before, but instead of them
we have ammonia, whose presence may be manifested by holding a
bit of reddened litmus-paper at the mouth of the tube. The intense
alkaline reaction of the gas, and its odor, sufficiently distinguish it from
both hydrogen and nitrogen.
88. It is to be observed that the hydrogen and nitrogen, which
refuse to unite when once actually in the free state, will form a
chemical compound, as in the last experiment, at the precise instant when they issue from other compounds of which they formed
part. This fact is expressed by saying that these two gases will
enter into combination when in the nascent state-that is, at the
moment of birth.
There are numerous cases in which bodies which do not unite
under ordinary conditions are capable of chemical combination at
the instant when they are disengaged from other compounds; and
the phrase " in the nascent state " is one of some convenience,
though it must not be supposed to explain, or in any way to
account for, the phenomena with reference to which it is used.
89. The exact quantitative analysis of ammonia gas will afford
proof that hydrogen and nitrogen are the sole constituents
of this gas, and will further show the proportions in which  28
they are combined. Ammonia is completely decomposed
into its elements by heat-the heat of a furnace or the
heat produced by a continuous discharge of electric sparks.
When, therefore, 50 c. c. of the gas are placed in a eudiometer (gas-measure) (Fig. 28), and sparks are passed
between the platinum points by means of a Ruhmkorff
apparatus, the gas is finally resolved into its elements;
its volume increases until it reaches 100 c. c., or double its original bulk, when it remains constant. We know that a part, at
least, of these 100 c. c. of gas is hydrogen, and that this hydrogen can be eliminated from the gaseous mixture by introducing
oxygen in sufficient quantity to convert the hydrogen into water,
and then exploding the mixture. Were the 100 c. c. of gas all
hydrogen, 50 c. c. of oxygen would convert it into water. That
we may be sure of having enough oxygen, let us introduce into




COMPOSITION OF AMMONIA.              81
the eudiometer 50 c. c. of oxygen, and then pass a spark to explode
the mixture.
After the explosion, only 37-5 c. c. of gas will remain; 112'5
c. c. have disappeared, having been converted into water. But
of these lost 112.5 c. c. we know that 75 c. c, must have been
hydrogen and 37-5 c. c. oxygen, in accordance with the known
volumetric composition of water; consequently, the 100 c. c. of
gas, into which the original 50 c. c. of ammonia were dilated,
contained 75 c. c. of hydrogen.  After the explosion, there remained 37'5 c. c. of gas, and since we used only 37'5 c. c. of
oxygen out of 50 c. c. added, we may infer that 12'5 c. c. of
oxygen still remain in the residue from the explosion. On introducing into the eudiometer a little pyrogallic acid (an acid used
in photography) dissolved in rwater, and a few drops of a solution
of caustic potash, these 1295 c. c. of oxygen will all be absorbed
by these liquids, and there will remain 25 c. c. of a colorless gas,
which may readily be recognized as pure nitrogen. The original
50 c. c. of ammonia have therefore yielded 75 c. c. of hydrogen
and 25 c. c. of nitrogen, and the composition of ammonia, both
by weight and measure, may be fully expressed by the diagram:
111f~~ + 1   -NH17
This composition of the gas is verified by its specific gravity as
determined by experiment, namely, 8'62.
Three volumes of H weigh.. 3
One volume of N weighs...   14
Two volumes of NIT3 should weigh. 17
One volume of NH3 should weigh..    85
a number sufficiently near the result of direct experiment.
90. The knowledge, thus acquired, of the composition of ammonia will enable us to recur with advantage to some of the exo




82                 THE GROUP AMMONIUM.
periments performed in the first part of this chapter. Ammonia
gas, when dissolved in water (as in the Liquor Ammonite), must be
considered to be in combination with one molecule of water, in
the form of the compound N3,H3120 or NH5O. This compound
may be supposed to be dissolved in the water present in excess of
what is necessary to form the compound. When this water of
ammonia combines with nitric acid. as in Exp. 33, to form the
compound we have called nitrate of ammonium, a reaction occurs
precisely similar to that which takes place where caustic soda
combines with nitric acid (p. 75); but in order to bringa out the
resemblance, the elements of the compound of ammonia and water
must be so arranged as to exhibit its analogy with caustic soda,
whose formula is NaHO. For that purpose its formula must be
written (NH4)HO, so that the group-of elements NH4 shall stand
in the formula of water of ammonia where the element sodium
stands in the formula of caustic soda. The combination of water
of ammonia with nitric acid may then be represented by the
equation
(NH4)HO    +   HNO,  =   (NH4)N03    +           120
Ammonia-water.  Nitric acid. Nitrate of ammonium.   Water.
just like
NaHO  +  HNO, =  NaNO3 -F 120.
The student may write both of these reactions in the typical
manner; the simple type water, H } O, is the only one needed to
represent all these substances, whether before or after the reactions which take place between them.
91. Ammonia-water combines with nearly all the acids with
which soda is capable of combining, forming a series of compounds
in which the group of atoms NH, plays the same part which the
single atom Na plays in the corresponding compounds of sodium.
For this reason it has been found convenient to give to this group
of atoms a name bearing some resemblance to the names of
metals, and it has therefore been called ammonium. Ammonium
is known only in its compounds; many attempts have been made
to obtain it in a free state, but hitherto in vain; as soon as the
group of atoms escapes from combination, it is resolved into qmmonia and hydrogen.




SALTS OF AMMONIUM.                  83
The important compounds into which ammonium enters, commonly called the salts of ammonium, will be studied hereafter in
immediate connexion with the analogous salts of sodium and
potassium. Already, however, it will be possible to verify the
statement of ~ 66, to the effect that nitrous oxide contains " all
the elements, besides those of water, which enter into the composition of nitrate of ammonium, and therefore of its constituents,
nitric acid and ammonia-water." The reaction last given shows
that a molecule of nitrate of ammonium contains all the elements
of a molecule of nitric acid and a molecule of ammonia-water,
less one molecule of water. If the formulae of ammonia-gas and
nitrate of ammonium have been correctly determined, it will be easy
to exhibit in an equation the actual result of Exp. 34, as follows:(NH4)NO3 =  N20 +   H120.
We thus link together several distinct experiments, and confirm
the determinations previously made of the composition of nitric
acid, nitrous oxide, ammonia, and water, by showing that the
representative formulae of these substances exhibit with perfect
precision reactions already well known to us, but other than those
from which the formulae in question were originally derived.
92. Ammonia exists in very minute quantity in the atmosphere, and hence in rain-water, fog, and dew. The proportion
of ammonia in rain-water has been variously given by different
observers, from 3'49 parts to 0'744 part of ammonia in 1,000,000
parts of water.  The water of fog and dew contains a larger
proportional quantity of ammonia; on accouht of the high solubility of the gas, the proportion of ammonia in water derived
from the atmosphere is greater, the smaller the fall of water.
Ammonia is given off by putrefying animal and vegetable substances containing nitrogen; and almost every process of slow
oxidation in the presence of air and moisture is attended with the
formation of ammonia or ammonia salts. Moistened iron-filings,
if exposed to the air, become rusty; and this rust is found to
contain a small quantity of ammonia. When some of the metals,
by preference tin, zinc, or iron, are dissolved in dilute nitric
acid, the oxidation of the metal is frequently accompanied, to a
greater or less extent, by the production of ammonia. But the
chief source of ammonia and its compounds is the decomposition,
G2




84                  SOURCES OF AMMONIA.
either by putrefaction or destructive distillation, of nitrogenous
organic matter. The distillation of bones and animal refuse, for
the purpose of making bone-black, yields a large amount of ammoniacal liquor, which was formerly the principal source of the
compounds of ammonia; the horns of deer used to be thus distilled, whence the name "hartshorn."  At present, the destructive distillation of coal in gas-works furnishes the great bulk of
ammonia compounds used in the arts. The ammoniacal liquor
of the gas-works is water contaminated with tarry matters, and
holding in solution a small proportion of very volatile ammonium
salts. These volatile salts are distilled off, by application of heat,
into dilute sulphuric or muriatic acid, and the sulphate, or chloride, of ammonium thus formed is obtained from the dilute liquid
by evaporation and crystallization. These salts will be studied in
detail hereafter.
93. The solution of ammonia-gas in water is a reagent continually required, as a test, in the laboratory, and much used in
the arts. The solution is colorless, intensely alkaline, has a
caustic taste, and, when concentrated, blisters the skin. The solution is lighter than water, and so much the lighter in proportion
to the amount of ammonia it contains; for this reason its strength
may be accurately determined by its specific gravity. Tables of
the relation of strength to specific gravity may be found in chemical dictionaries. The applications of ammonia-water are numerous and various; it is an ingredient in many pharmaceutical
preparations; applied to the skin, it is an irritant, and, when
very strong, even a caustic; its pungent odor is reviving and
stimulating; in veterinary practice it is a useful medicament;
on account of its alkalinity it is used in removing grease from
cloth, and in restoring colors which have been changed by acids.
The solution is ordinarily prepared from a mixture of chloride, or
sulphate of ammonium, with slaked lime.
Exp. 48. —Mix intimately 25 grins. of chloride of ammonium, a substance generally sold under the name of sal-ammoniac, with about the
same weight of cold freshly slaked lime. Introduce the mixture into
a flask of 500 c. c. capacity, and place the flask on a sand-bath over
the gas-lamp. If a smaller flask be used, the quantities of the materials must, of course, be proportionally diminished. Close the mouth




MAKIN'G AMMONIA-WATER.                   85
of the flask with a good cork, provided with a delivery-tube so bent as
to connect conveniently, by mneans of a caoutchouc connector, with the
first of the series of small three-necked bottles (Woulfe's bottles) represented in Fig. 29.
Fig. 29.
The first of this series of bottles is smaller than the rest, and is not
filled so full of water as the others; it should be kept cool by immersion in cold water; the delivery-tube coming from the flask into this
bottle must not dip into the water at all, so that it will be impossible
for any water to suck back into the flask, should the gas suddenly
cease to come off from the dry mixture. The construction of the
apparatus is easily to be understood from the figure; the open tube
which dips beneath the water in each bottle is a safety-tube, which, by
admitting, air into any bottle in which a partial vacuum may happen to
be created by rapid absorption, prevents the contents of the succeeding
bottle from flowing back into it. In order to show the action of the
safety-tubes, the open tube in the first bottle may be closed for a
moment with the finger and the bottle shaken very gently. Water
will immediately be forced back from the second bottle through the
connecting-tube to fill the vacuum cauted by the absorption of the
ammoniacal gas; but the moment the finger is removed from the safetytube air will enter through the latter to fill the vacuum, and the water
in the connecting-tube will fall back into the second bottle.
The ammonia-gas cannot avoid four separate contacts with water as
it passes through the apparatus, so that all the gas is sure to be absorbed; the contents of the first bottle will not be as pure as those of
the succeeding. This apparatus, or modifications of it, is used on the
large scale as well as the small, in operat'os which involve the absorption of a gas by a liquid capable of dissolving it. When a large quantity of gas is continuously delivered from the generating vessel, the
absorption can be made equally continuous by successively removing




86                  EMPIRICAL FORMULM.
the bottle nearest the flask as soon as the liquid in it is saturated, and
adding a fresh bottle at the other end of the series. On heatino the
flask, the ammonia-gas will be set free from the mixture, and in this
experiment will be mostly absorbed in the first and second Woulfebottles. It is evident that the Woulfe-bottles in this experiment might
be replaced by conmmon bottles with mouths wide enough to admit
corks pierced with three holes.
The reaction between the chloride of ammonium and the
slaked lime is represented by the following equations: —
2NH4C1  +   CaI,02  =   2NH3  +   CaC12  +   2H20
Chloride of   Slaked lime.  Ammonia.  Chloride of   Water.
ammonium.                              calcium.
Chloride of ammonium is a compound which may be obtained by
bringing together dry ammonia, NH3, and dry muriatic acid gas
HC1 (see Exp. 65).
NH3 +  HC1 =  NH3HC1 =  NH4C1l.
It may obviously be regarded as a compound of the group called
ammonium, NHI, with the element chlorine; from this view is
derived the name chloride of ammonium. Slaked lime is prepared from water and quicklime, a substance which is chemically
the oxide of the metal calcium, -
CaO +  H1O =  CaO,120 =  CaH,O,.
94. In each of the last two chemical equations, two expressions
are given for a single substance; in the first case, chloride of
ammonium is formulated in two different ways, and in the second,
slaked lime.  We may now inquire into the meaning of this
diversity of expression for one and the same substance. A formula which simply represents the number of atoms of each element in one molecule of any substance, as determined by its
analysis, is called an empirical formula. The truth of such a
formula depends solely upon the correct performance of the analytical process, and upon the accuracy with which the atomic
weights have been determined. Concerning such formulae thern
is little room for difference of opinion; they express all that we
actually know of the elementary composition of any compound
body. But chemists have endeavoured to contrive formulae which
should express something more than the mere elementary com



EMPIRICAL AND RATIONAL FORMULa.           87
position by weight-which should recall the materials from which
the formulated substance was made, and prophesy the products of
its decomposition-which should not only name and number the
atoms of the substance, but should also suggest such a grouping
or arrangement of those atoms as might serve to interpret its
known reactions. Such formulse are called rational formulae.
Thus NaHO is the empirical formula of caustic soda, while
Na20,HO20 is a rational formula of the same substance, which
recalls the facts that it may be made from anhydrous oxide of
sodium and water, and that it enters into many reactions in
which the Na20 goes one way and the H,20 another.
Another rational formula of the same substance is Na } 0,
which suggests many reactions in which caustic soda is involved
either as product or ingredient-for example, the decomposition
of water by metallic sodium (see Exp. 14), which may be thus
written:
H   ~ +  Na =  Ha} 0 + H.
H                H}
We shall hereafter meet with many such reactions, in which an
atom  of sodium  replaces an atom  of hydrogen; the formula
Na } 0 suggests this large class of reactions by implying that
caustic soda is itself constituted as water in which an atom of
sodium has replaced an atom of hydrogen. ~ Such formulae as
HI20,N20,, PbO,N20,, Na0,O,N20,, and Na2,OHO, are called
dualistic, because they represent these bodies as of a dual nature
-as being made up of two oxides which were distinct before
they were brought together to form the compound, and will be
distinct when separately extracted from it; in a dualistic formula
these two distinct parts are conventionally represented as having
some separate existence within the compound itself. The supposition is not unnatural; thus, for example, common plaster of
Paris is a substance containing the metal calcium and the elements sulphur and oxygen in the proportions by weight which
are correctly expressed by the formula CaS04; but this substance
may be made by methods which suggest another formula. If we
put together quicklime CaO, and anhydrous sulphuric acid SO, in




88                  DUALISTIC FORMUL2E.
due proportions, under suitable conditions, plaster of Paris, or, as
its chemical name is, sulphate of calcium, resultsCaO +  SO3 =  CaO,S03;
or if we mix slaked lime CaO,H20 with strong sulphuric acid
H20,SO3, in proper proportions, at a suitable temperature, we
shall again obtain sulphate of calcium, and water will be eliminated:
CaO,H20 + H20, SO3 =  CaO,SO, +  2H20.
Accordingly we find that the great majority of chemists have
hitherto written the formulae of sulphate of calcium, hydrated
oxide of calcium, and hydrated sulphuric acid, in conformity with
the suggestion  of these reactions, CaO,S03, CaO,H20, and
H20,SO, respectively.  Of positive knowledge concerning the
actual grouping of the imaginary atoms which are supposed to
make up the hypothetical molecule of a compound body, we have
absolutely none, and it must never be forgotten that a rational
formula is merely a suggestion of some of the chemical processes
in which the substance formulated is capable of taking part. In
some cases the rational formula may point to the majority of all
known transformations of the substance; but generally it suggests
only a few. of the possible changes. Since a rational formula
never represents a fact, but only an hypothesis or opinion, it is to
be expected that a great diversity of rational formulae should be
in use among chemists; and this is really the case.
The dualistic view, above illustrated, has long been, and still
is, the prevailing view of the proximate composition of inorganic
compounds; but in the chemistry of the very numerous compounds which the element carbon forms with oxygen, hydrogen,
and nitrogen, a different view, called the doctrine of types, widely
obtains, and has been adopted by not a few chemists as affording
the best theoretical representation of all chemical combinations,
whether in the inorganic or organic kingdom.  According to
this doctrine every possible chemical combination may be imagined to be built upon the plan, or framed upon the type or model,
of some one of the four substances, chlorhydric acid, water, ammonia, and marsh-gas. These will all shortly be to us well-known
substances; but the most important of these types is water, at




TYPICAL FORMULr.                  89
body with whose composition we are already familiar; we are
therefore competent to write upon the type of water the formulae
of several substances with which we have already dealt, and which
are classified under this type: —
Water=   } O=one molecule; i2 } O=two molecules.
Caustic Soda — Na }    one molecule.
Ca
Hydrated Oxide of Calcium, slaked lime,= Ca  0 =1 molecule.
Monohydrated Nitric acid=N2 } O= one molecule.
Nitrate of Potassium N12 } O= one molecule.
Nitrate of Lead=-2 (NO 2) 02=one molecule.
Sulphuric acid= S2 02= one molecule.
Sulphate of Calcium S02 } 02=one molecule.
In these formulae it is to be observed that K, Na, and NO, replace one atom of hydrogen in one molecule of water, while Ca,
Pb, and SO, replace two atoms of hydrogen in two molecules of
water. Facts of this class will accumulate as we advance, and
will be the subject of future discussion. The typical notation is
doubtless capable of expressing, in a logical and consistent system,
the greater part of the reactions of inorganic as well as of organic
chemistry; but at present it finds its best application in the chemistry of the compounds of carbon, and has gained but little foothold in the great departments of mineral and industrial chemistry.
The need of rational formulae is much more urgently felt in
that department of chemistry, called organic, which treats of the
chemistry of carbon, than in the wider field of mineral and inorganic chemistry. Among the very numerous compounds of carbon there are many cases in which one empirical formula represents not one compound, but several; hence it becomes of consequence to determine, or to guess, how the atoms of a compound
are arranged, as well as to know what and how many the atoms
are.  The diversity of opinion concerning this arrangement of




90                   USES OF FORAMULS.
atoms is so great, and the possible modes of grouping the numerous atoms which often enter into organic compounds are so many,
that the number of rational formulae proposed for any organic
substance is commonly large in proportion to the thoroughness
with which the substance has been studied. For acetic acid, for
example, one of the best-known of the compounds of carbon with
oxygen and hydrogen, no less than nineteen different rational
form ulae have been proposed.
Remembering that a rational formula is never to be reoarded
as the expression of an absolute truth, but only as a guide in
classification, an aid to the memory, and a help in instruction,
and holding fast to the empirical formula as containing all the
results of actual observation and experiment, we shall endeavor
to familiarize the student with both the dualistic and typical
guesses at the hidden mysteries of chemical processes and the
unknowable structure of chemical compounds, giving the preference rather to the dualistic view, as being that which at the
present moment prevails in the great bulk of chemical literature.
and has become incorporated into the language of the chemical
arts.
Lest any doubt should suggest itself to the student's mind as
to the value of symbolic formulae, let it be observed that they
express the elementary composition of a compound much more
tersely than words can, that they are written and read more
rapidly than the sentences of the same signification would be,
and that by their brevity, clearness, and precision they greatly
facilitate the comparative study and comprehensive classification
of chemical compounds. Again, the chemical equations, of whose
construction we have already had several examples, enable us to
set forth with precision the changes which accompany complicated,
as well as simple, reactions. Thus the somewhat complex decomposition of nitric acid by copper takes definite form in the
appropriate equation which has been given above (p. 74), and the
very simple reaction by which nitric oxide yields red fumes of
hyponitric acid in contact with air or oxygen is concisely stated
by the simple equation NO + O= NO2.
The chemistry of the analysis of nitric oxide by potassium
(~ 70) is all condensed into the equation NO+K2=K20+N.




CHLORHYDBIC ACID.                   91
When a little ice-cold water is added to liquid hyponitric acid
(Exp. 43), the reaction which occurs is very concisely set forth
in the equation;Empirical:    2NO2  +   Ho20   =    HNO2  +   HNO,
Hyponitric    Water.    Hydrated    Nitric acid.
acid.               nitrous acid.
Dualistic:    4NO2  +  2H110  =  1H20, N,203 +  H20, N205.
But besides having all the advantages of a short hand, chemical
symbols are susceptible of another application of hardly less importance; they often direct the chemist beforehand to the most
perfect experiment among many similar, or point out in anticipation the possibility of certain methods of research, and the inevitable fruitlessness of others. Thus the equation
NaO, + 5Cu =  N, + 5CuO
actually directs the chemist to the due proportion of copper for
the exact decomposition of anhydrous nitric acid (~ 73); neither
four, nor six, nor any other number than five, parts of copper,
would give a perfect reaction without excess of either ingredient.
Practically, an excess of copper does no harm, and is always
used to make sure of the decomposition.
The student should endeavor, from the beginning, to familiarize
himself with the use of chemical symbols and equations, and to
this end he should invariably write the formula of every reaction
described or actually witnessed in the execution of an experiment.
CHAPTER VII.
CHLORHYDRIC ACID.
95. Mluriatic (sea-salt) acid, called in modern nomenclature
chlorhydric acid, is a liquid which has been known for centuries,
and is to-day an article of commerce, largely employed in the
useful arts. The pure acid is a gas, as ammonia is; the liquid
muriatic acid of commerce is only an aqueous solution of this




92              MAKING CRLORRHYDIC ACID GAS.
gas, and gives it up when heated, precisely as ammonia-water
yields ammonia-gas.
This operation may be conveniently performed in the apparatus
shown in Fig. 30. About
250 c. c. of the commercial
acid is poured into the flask,
which is then moderately
heated; the gas disengaged
is charged with aqueous vapor, which needs to be removed before the gas is collected.  For this purpose       _
the delivery-tube is carried
to the bottom of a bottle
filled with pieces of pumice-         i
stone saturated with strong
sulphuric acid; the moisture
of the gas is greedily absorbed by the large surface
of acid with which the gas
comes in contact, as it is forced upward through the acid-soaked stone.
The dry, colorless, transparent gas must be collected over mercury, for
it is extremely soluble in water.
96. The gas is strongly acid in taste and reaction on vegetable
colors, provokes violent coughing, and is wholly irrespirable. It
is neither combustible nor will it support combustion. The gas
is somewhat heavier than air; its specific gravity referred to
hydrogen, as determined by experiment, is 18'12, its theoretical
density being 18'25; it is possible, though not convenient, to
collect it by downward displacement.  It forms opaque, white
fumes in the air, owing to its union with, and condensation of,
atmospheric moisture. Under a pressure of 40 atmospheres, at a
temperature of 10~, chlorhydric acid gas is condensed into a colorless liquid. Its great solubility in water would lead us to expect
that it could be readily reduced to the liquid state; but, on the
contrary, it can be condensed only with difficulty.  At 0~, one
volume of water dissolves about 500 volumes of chlorhydric
acid gas; at common temperatures, something more than 400.
The specific gravity of the solution is greater than that of water,
and the more concentrated the solution the higher the specific




PROPERTIES OF CHLORHYDRIC ACID.            93
gravity; so that the strength of any sample of the commercial acid
may be ascertained by taking its specific gravity. Tables for this
use will be found in chemical dictionaries.
The avidity of water for chlorhydric acid gas may be neatly
shown by thrusting a bit of ice into a small cylinder of the dry
gas standing over mercury; the ice instantly melts, and the gas
as quickly disappears. A solution of the acid containing 20'2
per cent. of the gas, and having a specificgravity of 1'104, distils
unchanged at a temperature of about 1110; stronger solutions
than this, on being heated under the ordinary atmospheric pressure, lose gas until reduced to this strength; weaker solutions
lose water until raised to this degree of concentration. This
stable solution, which distils unchanged, is supposed to be a definite compound of the dry gas and water, whose composition the
formula HC1 + 8H20 would correctly represent.
97. We propose to answer the question-of what is chlorhydric acid composed-by a partial analysis and a complete
synthesis.
One of the elements of this gas can be isolated by a method which
we have already applied to the analysis of ammonia. It is only necessary to remove the delivery-tube from the apparatus already used to
generate the dry gas (Fig. 30), and to fix in its place a bulb-tube of
hard glass, containing a
piece of potassium.  As                 Fig. 3L
soon as the acid gas reaches
the potassium, the metal
becomes covered with a
white incrustation; and if
the bulb be now verygently
heated (Fig. 31), the potassium fuses, and taking fire,
burns with a violet light.
During the reaction the         U
chlorhydric acid is decomposed, an inflammable gas,
easily recognized as hydro-  
gen, is evolved, and may
be lighted at the end of.
the tube.
The metal sodium produces similar results, but at a much higher




94              ANALYSIS OF CHLORHYDRIC ACID.
temperature. A solution of sodium in mercury, known amongst chemists as sodium-amalgam, will however bring about the decomposition
of the acid at the ordinary temperature. This solution is best prepared by very gently heating some mercury in a glass flask, and gradually adding the sodium, cut into fragments not bigger than a grain
of wheat; the fragments dissolve with evolution of light and heat.
Why the sodium-amalgam should act at a lower temperature than the
sodium itself, is not clear, unless it be that the minute subdivision of
the sodium in the mercury gives the gas a freer contact with the metal.
Instead of potassium or sodium, as above described, metallic iron
could be employed for analyzing chlorhydric acid. To this end iron
turnings should be heated to redness in a glass tube such as is shown
in Fig. 8; a wide delivery-tube should be used.
Hydrogen is, then, one ingredient of chlorhydric acid; the other, or
others, have combined with the potassium or sodium-amalgam. The
isolation of these unknown ingredients may be   Fir. 32.
accomplished by means of the V-tube already
used for the analysis of ammonia. Into this                 f
tube liquid chlorhydric acid of specific gravity
1'1, colored with indigo solution, is introduced,
so as to fill the whole length of the sealed, and
about half the length of the open, limb; the
negative pole of the battery is connected with
the wire of the sealed limb, while the positive
pole is inserted into the open limb. Gas rapidly collects at the negative pole in the closed
limb, but at the positive pole the disengagement of gas is so slight that it would hardly
attract attention but for its intensely disagreeable odor and powerful bleaching action upon the blue liquid. The
gas in the sealed limb has no such bleaching power. When enough
gas for examination has collected in the sealed limb, it is transferred
to the open limb by the manipulation previously described (~ 86); the
gas is inflammable, and is, in short, hydrogen.
The poles are now reversed, and immediately hydrogen escapes in
abundance from the open mouth of the tube, while the liquid in the
closed limb becomes decolorized. In the course of fifteen minutes the
bleached liquid in the sealed limb begins to assume a yellowish-green
color, and the evolution of gas becomes gradually more and more
copious, so that in three-quarters of an hour the greater portion of the
tube is filled with a transparent, yellowish-green gas. As the gas is
transferred to the open limb of the tube for examination, it manifests




ANALYSIS OF CHLORHIYDRIC ACID.              95
its powerful bleaching property by decolorizing, as it passes, the portion of the acid which had retained the blue color of the indi,,o. The
tube is no sooner opened to admit a burning taper than the suffocating
odor of the gas becomes offensively perceptible; the gas proves to be
uninflammable, and it supports combustion but imperfectly, as is evidenced by the sooty cloud which is produced.
This peculiar gas, so different in properties from any gas heretofore
studied, is an element; it has been named, on account of its color,
Chlorine, from the Greek word for yellowish-green. This element is
the subject of the next chapter, where it will be fully studied. As will
there be seen (Exp. 51), chlorhydric acid, when heated with a substance called black oxide of manganese, yields chlorine in abundance,
with great facility; in fact, this acid is the source of chlorine whenever large quantities of this gas are required. Chlorine is soluble in
about one-third of its volume of cold water, —a property which explains its apparently slow evolution at the outset of the foregoing
experiment, and the more rapid disengagement of the gas when the
liquid has become saturated therewith. Chlorine is heavier than air,
and consequently very much heavier than hydrogen; the best experimental determination of its specific gravity has given the number 35-66:
but there can be no doubt that the true specific gravity of the gas is
35'5, or in other words that it is 351 times as heavy as hydrogen.
98. We have thus learned that the electric current sets free
from chlorhydric acid two essentially different gases-hydrogen
and chlorine,-and that each of these gases may be separately
evolved from muriatic acid-the hydrogen by potassium, and the
chlorine by oxide of manganese. It remains to prove that chlorhydric acid contains no other than these two constituents, and to demonstrate the proportions in
which they are united. To this end the first st p
shall be to make a partial quantitative analysis
of chlorhydric acid gas.
The instrument employed is a glass U-tube, about
50 c.m. long by 1'5 in diameter, having one sealed
and one open limb; communicating with the latter is
a small outlet-tube which may be closed by a springclip on a piece of caoutchouc tubing. The apparatus,
mounted on a convenient stand, is represented in
Fig. 33. The U-tube is first filled with mercury, and
then, the spring-clip being open, the delivery-tube of
the apparatus used to generate dry chlorhydric acid gas is passed




96            COMPOSITION OF CHLORHIYDRIC ACID.
down the open limb to the bend of the tube in such a manner that the
gas bubbles up through the mercury into the sealed limb, from which
the mercury escapes as the gas enters. When the closed limb is twothirds full, the outlet-tube is closed, the gas delivery-tube withdrawn,
and mercury is poured into the apparatus until it stands at the
same level in both limbs. The space occupied by gas in the tube is
then marked by a caoutchouc ring slipped over the tube. That portion of the open limb which is not occupied by mercury is then filled
with sodium-amalgam. By closing the orifice of the tube with the
thumb, and inclining the tube, the gas may be transferred from the
sealed limb to the other, there shaken up with the amalgam, and retransferred to the sealed limb. This thorough contact with the sodium
decomposes the gas. On removing the thumb from the mouth of the
open limb, the mercury therein falls a little, and must be further
lowered by opening the spring-clip until the mercury stands at one
level in the two limbs. When this is the case, it will be observed
that the gas is reduced to half its original volume. The gas which
remains, is, of course, hydrogen. The experiment proves that any
given bulk of chlorhydric acid contains half that bulk of hydrogen.
By availing ourselves of the known specific gravities, or likevolume weights, of chlorhydric acid and chlorine, referred to
hydrogen, we may establish a strong presumption in favor of
the supposition, that that half of any bulk of chlorhydric acid
which is not hydrogen is chlorine without admixture of any
other substance.
From the relative weight of any volume of chlorhydric acid gas 18'12
Subtract the relative weight of half that volume of hydrogen..'50
And the remainder................................ 17'G2
is very nearly equal to 17'83, the relative weight of half the
same volume of chlorine, according to the best experimental determinations. If we assume, for the moment, that any volume of
chlorhydric gas is really composed of half that volume of hydrogen and half of chlorine, and if we use the theoretical specific
gravities, which are doubtless the true ones, instead of the above
approximate determinations, which are the best which experiment has hitherto furnished, the numerical statement will be as
follows:



ATOMIC WEIGHT OF CHLORINE.              97
From the relative weight of any volume of chlorhydric acid gas 18-26
Subtract the relative weight of half that volume of hydrogen..'50
And the remainder................................ 17-75
is the relative weight of half the same volume of chlorine
17'75=-35-.  The very close coincidence between the first numerica.l statement, which is based wholly upon experiment, with the
second, which is based on the theory that chlorhydric acid is, by
volume, half hydrogen. half chlorine, is evidence enough, when
taken in connexion with the preceding experimental isolation of
chlorine from chlorhydric acid, to convince us that in any two
volumes of chlorhydric acid gas, one volume of hydrogen and one
volume of chlorine are united without condensation.
The formula of the molecule of the acid will therefore be HC1,
in which C1 is the symbol of chlorine. The following diagram
represents the composition of this important compound, both by
volume and weight:II   +              HC 365
99. The atomic weight of the new element, chlorine, is hereby
determined. Hydrogen and chlorine unite by equal volumes to
form this single stable compound, chlorhydric acid; and the proportions in which the two elements unite by weight are directly
deducible from the proportions in which they unite by volume
and the known specific gravities of the two gases. Indeed it also
admits of direct proof, by appropriate experiment, that 36-5
parts by weight of chlorhydric acid gas invariably yield 35'5 parts
by weight of chlorine and 1 part by weight of hydrogen; and,
since it matters not what the absolute weight of these parts may
be, millionths or millions of grammes, the mo'ecale of chlorhydric
acid, the least proportional weight in which it is conceived to
exist uncombined, must be composed, like any other quantity of
the acid, of 35'5 parts by weight of chlorine to 1 of hydrogen.
But we conceive of this molecule as consisting of one atom of
chlorine and one atom of hydrogen; the chlorine atom, therefore,
weighs 35'5 times as much as the hydrogefi atom.
100. If it were entirely inconceivable that another substance,
H




93              SYNTHESIS OF CHLORHYDRIC ACID.
not identical with chlorine, should have precisely the same specific
gravity as chlorine, the reasoning by which we have just arrived
at the composition of chlorhydric acid would be not only convincing, it would be entirely conclusive; it would do more than
establish a very strong presumption, it would furnish a complete
demonstration. But it is not inconceivable, though in the highest
degree improbable, that there should exist a body, different from
chlorine, yet possessing the same specific gravity, and not to be
detected by the qualitative tests to which we subjected the acid;
and we therefore welcome the perfect demonstration with which
the synthesis of chlorhydric acid supplies us.
This synthesis is readily effected; but as the experiment involves the
preparation of chlorine, the actual performance of the experiment will
be best postponed until chlorine has been prepared and studied in the
next chapter. The method is as follows:-Into one of two glass
cylinders, standing full of mercury upon the mercury trough, introduce
a certain volume of dry hydrogen, not so large as to fill more than
half the cylinder, and into the other cylinder bring precisely the same
volume of dry chlorine. Cover one of the cylinders from the light
with a towel, and deliver the contents of the other cylinder into the
protected one. The mixture of equal volumes of the two gases having
been thus effected, withdraw the towel, and leave the cylinder for
several hours in diffused and not too bright daylight, but sheltered
from the direct rays of the sun, which would cause an explosive union
of the two elements. Under the influence of the light the two gases
gradually combine; when the yellowish tint of the mixture has nearly
disappeared, the cylinder may be exposed to the direct influence of
the solar rays, in order to complete the reaction. Throughout the experiment there is no change in the volume of the enclosed gas; if the
temperature were constant the mercury would neither rise nor fall in
the cylinder. The chemical union of one volume of hydrogen with
one volume of chlorine is attended neither by condensation nor expansion.
That there hasibeen chemical action, resulting in the disappearance
of the properties of the original materials, is evident from the fact that
the contents of the, cylinder will no longer take fire, or bleach vegetable colors. In contact with air the new gas forms white clouds;
blue litmus paper it turnsred; and if a little water be passed up into
the cylinder, the gas is:rapidly absorbed; the taste and smell also,
and, in short, all ithe Proper.ties of this gas are those of chlorhydric acid.




MANUFACTURE OF CHLORHYDRIC ACID.           99
By the synthetical method we therefore prove that chlorine
and hydrogen are the only constituents of chlorhydric acid. On
this fact is based the chemical name of this compound. Samming
up our previous and present results, we now possess a complete
demonstration that chlorhydric acid is composed solely of hydrogen
and chlorine, united in equal volumes without condensation.
101. The muriatic acid of commerce is made from the most
abundant and cheapest of all the natural compounds of chlorine,
common salt, whose chemical name is chloride of sodium, and
formula NaCl. This substance supplies the chlorine; the necessary hydrogen is obtained from common sulphuric acid (oil of
vitriol), whose composition, as expressed in its formula H2SO4,
we have already become familiar with.
The commercial acid is obtained by heating common salt with
sulphuric acid in iron pans or cylinders, and absorbing the evolved
gas in water contained in a series of stone-ware Woulfe-bottles,
or some similar apparatus. The reaction is somewhat various,
according to the proportion of sulphuric acid employed; it may
be either of the reactions expressed in the following equations,
or may lie between them:NaCl   +    H2SO4    =        HC1    +       NaiSO4
Chloride of    Sul-phuric    Chlorhydric    Acid sulphate
sodium.         acid.           acid.        of sodium.
2NaCl   +    H12SO4    -        2HC11   +    Na2SO4
Sulphate of sodium.
In the first reaction, only one-half of the hydrogen in each
molecule of sulphuric acid is replaced by sodium; in the second,
both atoms of hydrogen are replaced. The first reaction requires
more sulphuric acid than the second, in proportion to the amount
of the product, but is accomplished with less wear of the apparatus, because a less heat suffices for the first than for the second
reaction.
We may illustrate the practical importance of the atomic
weights, by taking an actual example of each of these reactions.
Starting with 100 kilos. of salt in each case, what quantities of
sulphuric acid should be employed, and what will be the weights
of the products in each reaction? (See ~ 81.)
2




100         MANUFACTURE OF CHLORHYDRIC ACID.
The molecular weight of NaCl   is 23 + 35'5       =  58'5,,,,,,,     2SO4  is 2+32+4X16    =  98,,,,,,    NaHSO4 is 23+ 1+32+4X 16=120,,,,,,    NaSO4 is 46+32+4X16          =142,9,,,,    HC1    is 1+35'5                =  36'5
The weight of sulphuric acid needed in the two cases is ascertained by solving the following proportions:
First      585:  98   =   100k.    x(=167'52 k.)
reaction  Mol. wt. of Mol. wt. of Quantity of   Quantity of
NaCl.  1H2SO4.   NaCl used. 112SO4 required.
Second     117       98   -   100k.    x(=83'76k.)
reaction  Mol. wt. of Mol. wt. of Quantity of   Quantity of
2NaC1.   H2SO4.  NaCl used.  B2S04 required.
The weight of chlorhydric acid gas produced in the two cases will
be precisely the same; it is deduced from the proportions:First      58'5:  36S5 -      100k.     x(=62'39 k.)
reaction r  Mo1. wt. of Mol. wt. of Quantity of   Quantity of
NaCl.    ICL.    NaCl used.  HC1 produced.
reaction J Mol. wt. of Mol. wt. of  Quantity of   Quantity of
2NaCl.   2HC1.    NaCl used.  HCI produced.
The weights of the residual sodium-salts in the two cases are deduced from the proportions:First      58'5:  120   =   100 k.: x (=205'128 k.)
reaction  Mo. wt. of Mol. wt. of  Qantity of  Quantity of
NaCi.   NaHSO4.  NaCl used. NaHSO4produced.
Second     117       142   -   100 k.: x(=121'367k.)
reaction  Mot. wt. of Mo?. wt. of  Quantity of  Quantity of
2NaC1.  Na2SO4.   NaCl used. Na2SO4 produced.
In each case the sum of the weights of the materials employed is,
of course, equal to the sum of the weights of the products. If
the questions suggest themselves-how much water will these
62'39 k. of chlorhydric acid gas saturate, and what will be the
bulk of the concentrated solution so obtained?-the answers can




MANUFACTURE OF CHLlORHYDRIC ACID.        101
be easily deduced from the following data:-The strongest chlorhydric acid has a specific gravity of about 1'20, and contains
about 40 per cent. by weight of the gas. 40 k. of chlorhydric
acid gas will then saturate 60 k. of water, and it follows that
62'39 k. of the gas will saturate 93'58 k. of water. The weight
of the solution of chlorhydric acid produced will therefore be
62'39 + 93'58= 155'97 k.; and since 1 litre of the solution weighs
1'20 k., the total bulk of concentrated aqueous acid produced will
be very nearly 130 litres.
102. If the question suggest itself-why not get the hydrogen
wanted from water, H2,O, a much simpler and cheaper substance
than sulphuric acid?-the only answer is, that experience has
taught that water has no action upon salt except to dissolve
it, while sulphuric acid has power to part the two elements of
salt, and, giving hydrogen to the chlorine of the salt, to accept the
detached sodium of the salt in the place of its own lost hydrogen.
Of the nature of the play of forces by which this new adjustment in definite proportions of the atoms of five elements is
brought about, we have no distinct conception. All that we
know has been said when it is stated that water works no chemical change on salt, while sulphuric acid (and a few other substances of analogous composition) does bring about a very essential change.
In the hope of rendering these and similar facts more intelligible, many chemists have assumed that an element like chlorine,
or a group of elements like sulphuric acid, may possess a superior chemical attraction, or a greater affinity, for some elements
or groups than for others. They would explain the reaction
between salt and sulphuric acid by saying that chlorine has a
greater affinity for hydrogen than for sodium, while a part of the
sulphuric acid has a stronger attraction for sodium than for
hydrogen; and, in like manner, they would account for the absence of action between water and salt by saying that the affinity
of oxygen for sodium is no stronger than that of chlorine for
sodium. If the second of the equations above given be written
after the dualistic theory, as follows,
2NaCl + H120, SO3 =  21C1 +  Na20, SO,,
we shall perceive the basis of a still more ample explanation,




102                  CHEMICAL AFFINITY.
often given, of such reactions. The reaction above written is
said to be determined, or caused, by three affinities:-1. The
affinity of the metal for.oxygen; 2. The affinity of the hydrogen
for chlorine; 3. The affinity of the oxide of sodium for sulphuric
acid. It will be at once perceived that the contact of water with
salt gives opportunity for the play of the first two affinities; it is,
therefore, the third affinity, superadded to the other two, which
in this view actually determines the decomposition of salt by sulphuric acid.
Such speculations as these have not been altogether fruitless
in the development of chemistry, and to some minds they seem
to render the actual phenomenon more intelligible; the term
affinity is also sometimes convenient in expressing the varying
intensity with which one element grapples and holds other elements, or groups of elements; the student must not fail to distinguish, however, between the matters of fact and the matters
of speculation, in whatever stands written in chemical literature
touching affinities and their play. The best use of the ill-chosen
term affinity is as a synonyme for chemical force. Phrases in
which the term is used in this sense may contain simple statements of fact; but very firequently, especially when the word
"elective" is coupled with it, the term is used in connexion
with unprofitable hypotheses. What we actually know of the reaction between salt and sulphuric acid is comprehended in the
statement that the hydrogen of the acid and the sodium of the
salt change places in the definite proportions by weight which are
expressed in the atomic weights of the two elements.
Commercial chlorhydric acid is not pure; its commonest impurities are sulphuric acid which gets mechanically mixed with
the acid, iron derived from the iron vessels, arsenic supplied by
the impure sulphuric acid employed, the salts contained in the
water which dissolved the gas, sulphurous acid, and, not unfrequently, free chlorine.
Exp. 49.-Melt a handful of coarse common salt in a Hessian crucible in a coal fire, and pour out the liquid salt upon a brick or stone
floor. Weigh out 30 grms. of the fused salt when it has become cold,
and place it in a flask of a litre capacity, provided with a deliverytube which can be conveniently connected by a caoutchouc connector




USES OF CHLORHYDRIC ACID.                103
with a series of small Woulfe-bottles, such as is represented in Fig. 34.
Pour 60 grammes of strong sulphuric acid upon the salt, and iammeFig. 34.
diately cork the flask, place it upon a sand-bath on the iron-stand,
and connect the deliverv-tube with the Woulfe-bottles. The tubes
by which the gas enters the bottles should barely dip beneath the
water contained in them, inasmuch as the solution of chlorhydric acid
is heavier than water; the bottles should not be more than half full,
for the water becomes hot and increases considerably in bulk. As hot
water holds less gas in solution than cold water, it is not amiss to
place each three-necked bottle in a vessel of cold water. The first
Woulfe-bottle should conta'n but a small quantity of water, and the
tube coming from the flask should not dip into this water. The contents of the flask must be very gradually and moderately heated, else
a violent frothing is liable to occur, which would spoil the experiment.
The process is like that of making ammonia-water, except that the
delivery-tube passes to the bottom of each Woulfe-bottle in making
ammonia-water, because the solution of ammonia-gas is lighter than
water, instead of heavier as is the case with the solution of chlorhydric acid gas. As with the ammonia process, the solution will be
purer in the second bottle than in the first, in the third than in the
second, and so forth. Reserve the contents of the first bottle to make
chlorine from (Exp. 51). The pure acid should be preserved for use
in experiments which cannot be performed except with an acid purer
than the commercial article.
103. The uses of chlorhydric acid are very numerous. It is
employed in making chlorine, chlorate of potassium, and chloride
of lime (bleaching-powder), in preparing chloride of ammonium
and chloride of tin, in the manufacture of gelatine, for dissolving




104                     AQUA REGIA.
metals (either by itself or mixed with nitric acid); it is one of the
most useful reagents in the chemical laboratory.
Chlorhydric acid dissolves most metallic oxides, and appears to
combine with them; but on evaporating such a solution a compound is obtained which contains neither hydrogen nor oxygen,
but only chlorine and the metal. When caustic soda, for example,
combines with chlorhydric acid, chloride of sodium and water are
the products, as exhibited by the equation
NaHO  +  ICl =  NaCl +  H{20.
When the black oxide of copper is dissolved in chlorhydric acid,
the green liquid produced is an aqueous solution of chloride of
copper;
Cu0 +  2HC1 =  CuCI, ~+   10.
But though the metal may exist in solution in the form of chloride, it is quite possible to precipitate it as oxide, if it have an
insoluble oxide, by adding to the solution of the chloride a soluble
oxide of another metal capable of displacing the first. Thus, if
to a boiling solution of chloride of copper a hot solution of caustic
soda be added, the sodium and the copper change places, and the
insoluble black oxide of copper is precipitated.
CuC1l +  2NaH0 =  2NaCl +,20 +  CuO.
Chlorhydric acid is, in fact, the chloride of hydrogen, strictly
analogous in composition to the chloride of a metal like sodium,
and it takes part in double decompositions like any other chloride.
104. Aqua Regia (Royal Water).-This name was given by
the alchemists to a mixture of chlorhydric and nitric acids, because of its power to dissolve gold, the " king of metals."
Exp. 50.-Place two square centimetres of genuine gold-leaf at the
bottom of a test-tube, and pour upon the gold six or eight drops of
strong chlorhydiic acid; put a similar piece of gold-leaf in a second
test-tube, and pour upon it two or three drops of nitric acid; neither
acid attacks the gold, which remains undissolved. If the contents of
the two test-tubes be mixed together in either tube, the gold-leaf will
almost immediately dissolve.
Platinum, which like gold resists the action of both chlorhydric
and nitric acids singly applied, yields at once to the mixture of
the two acids. Both these precious metals are converted by aqua




CHL0ORNE.                      105
regia into chlorides soluble in water.  Strong chlorhydric acid is
oxidized by strong nitric acid; chlorine, water, oxides of nitrogen,
and unstable compounds containing chlorine, oxygen, and nitrogen, are the products. The decomposition is complex, but may be
roughly represented by the equation
HC1 + HNO, =  C1 +  H20 +  NO,.
The presence of nascent (~ 88) chlorine explains the energetic
conversion of metals into chlorides by aqua regia; and the strong
oxidizing effect of the liquid is further explained by the presence
of the unstable oxygen compounds which result from the reaction.
Aqua regia has indeed a very strong oxidizing power; it can
change sulphur into sulphuric acid, arsenic into arsenic acid, and
effect many other similar oxidations.
This powerful solvent is made by simply mixing the two acids,
though in various proportions, according to the use to be made of
it; the commonest mixture is composed of one part of nitric acid
and three parts of chlorhydric acid.
CHAPTER VIIT.
CH LORI NE.
105. Chlorine can readily be prepared from chlorhydric acid
by removing the hydrogen of that acid by chemical means.
Exp. 51.-In a flask of about 500 c. c. capacity furnished with a
suitable delivery-tube, place 8 or 10 grins. of coarsely powdered black
oxide of manganese; pour upon it 20 or 30 grms. of common muriatic
acid, and gently heat the mixture. Chlorine will soon be disengaged,
and may be recognized by its peculiar color. Being very heavy the
gas may best be collected by displacement in dry bottles, placed in the
open air, or in a case or box provided with an efficient draft. It may
also be collected over warm water or brine in the water-pan. It cannot be well collected over water at the ordinary temperature, since it
is rather easily soluble therein —though the difficulty may be obviated
in part by evolving the gas rapidly, or by passing the delivery-tube
to the top of the bottle in which the gas is collected. It must not be




106                  MAKING' CHLORINE.
left standing over water, since it would soon be entirely absorbed.
In experimenting with chlorine, care must always be taken not to
inhale it.
The reaction which occurs in this experiment may be thus
formulated:
MnO2 + 41HC1 = 2H20 + MnC12 +  2C1.
Black oxide of manganese is a substance rich in oxygen, which,
under certain conditions, it readily yields up to other elements.
In the case before us, the oxygen of the oxide of manganese unites
with the hydrogen of the chlorhydric acid to form water. The
chlorine of the chlorhydric acid unites in part with the manganese, and is in part left free.
In place of the black oxide of manganese in this experiment,
several other substances which readily give up oxygen may be
employed; and instead of the free chlorhydric acid of the foregoing experiment, the mixture of common salt and sulphuric
acid, which generates chlorhydric acid (Exp. 49), is often used.
The latter method has the advantage of eliminating the whole of
the chlorine from the chlorine compound used, whereas, in the
decomposition of the oxide of manganese by chlorhydric acid alone,
half the chlorine remains combined with the manganese. Moreover, when present in excess, the sulphuric acid has the effect of
drying the chlorine. The reaction may be expressed as follows:2NaC1 + 2H2S04 + MnO2 = Na2SO4 + MnSO, + 2HO0 + 2Cl1.
Another method, which has been carried out in practice upon the
large scale, is to heat a mixture of common salt and nitrate of sodium with an excess of sulphuric acid. Chiorhydric and nitric
acids are evolved, and, reacting upon one another, generate chlorine, hyponitric acid, and water:HC1  + HNO3 = C1 + NO2 +  1H20.
The hyponitric acid is absorbed by sulphuric acid, and subsequently employed in the manufacture of sulphuric acid, while the
chlorine is collected apart and employed in such manner as may
be desired.
106. Chlorine is an abundant element, and very widely distributed in nature. It exists chiefly in combination with sodium as
a chloride of sodium, which is called rock-salt or sea-salt, accord



PROPERTIES OF CHLORINE.                107
ingly as it is found in beds in the earth, or dissolved in the water of
the ocean. Since the atomic weight of chlorine is 35'5 (~ 99),
and that of the metal sodium is 23, each molecule, each gramme, or
each kilogramme of chloride of sodium contains 35-, or 60'684 per
cent. of chlorine. Accordingly in a gramme of chloride of sodium
there exists something more than 0-6 grm. of chlorine; and a
kilogramme of common salt should yield 606'84 grms of chlorine.
Every litre of sea-water will yield about five litres of chlorine gas.
Besides chloride of sodium, sea-water contains small quantities of
the chlorides of several other metals; there are numerous minerals, also, which contain chlorine.
107. At the ordinary temperature chlorine is a gas of yellowishgreen color, 2'5 times as heavy as atmospheric air. Its specific
gravity and atomic weight are 35'5. It is excessively irritating
and suffocating, even when inhaled in exceedingly small quantities. Any attempt to breathe the undiluted gas would undoubtedly be fatal. Under a pressure of 4 atmospheres at 150 it
is condensed to a yellow mobile liquid, having a sp. gr. of 1'33;
this liquid has never yet been solidified. It is soluble to a considerable extent in water at the ordinary temperature, 1 volume
of it being dissolved by half a volume of water at 15~. This solution, which exhibits the color, odor, and general chemical properties of the gas, is called chlorine-water. At low temperatures,
water dissolves a still greater proportion of chlorine; and at 00 a
definite hydrate of chlorine, C1,5H120, crystallizes out.
Exp. 52. —Fill with water the body of a retort of the capacity of
500 c. c., and without tubulature. Invert the retort and set it upon a ring, or   Fig. 35.
upon a bed of sand, with the neck
pointed upwards in such manner that
no air shall enter the body. From a
flask in which chlorine is being generated pass a long delivery-tube down
the neck of the retort to the water, so
that the chlorine may slowly bubble
through the water. The absorption of
the gas may be promoted by gently shaking the retort from time to
time. As soon as the water becomes saturated with chlorine, so much
gas will collect in the retort that the liquid will be pressed out of the




108                   CHLORINE-WATER.
body and will flow over from the neck; when this occurs the operation may be stopped.
At the beginning of the experiment, before all the atmospheric air
has been expelled from the flask in which the chlorine is generated, it
is well not to push the gas delivery-tube completely to the bottom of
the neck of the retort, but to simply immerse it in the edge of the
water, so that none of the escaping bubbles of gas shall enter the body
of the retort until it has become evident that nothing but pure chlorine
is coming over; the tube may then be immersed more deeply.
The water saturated with chlorine should be transferred to a bottle
and preserved for future use. It niay be employed, more conveniently
than the gas, to illustrate many of the properties of the element.
In sunlight, or even in ordinary daylight, chlorine-water suffers decomposition (see ~ 113), but in the dark it undergoes no change. It
should be kept, therefore, either in a cellar or tight closet, or in a
stoneware bottle, or in a bottle of black, red, or yellow glass, or in one
covered with black paper. Through the blackened glass no light can
penetrate to the chlorine-water, and through red or yellow glass few,
if any, of the so-called chemical or actinic rays can pass. The violet
rays of the spectrum are those which exhibit actinic power, and these
are stopped by red or yellow glass, which is red or yellow because it
permits the passage of only these colored rays.
108. Chlorine is a powerful chemical agent. It combines with
hydrogen with explosive violence upon being heated, or even on
being exposed to sunlight.
Exp. 53.-In a soda-water bottle, which must be screened from
strong light by wrapping it in a towel, unless direct and reflected sunlight be excluded from the room, mix equal -volumes of chlorine and
hydrogen, then remove the cork and hold the mouth of the bottle in
the flame of a lamp. A sharp explosion will ensue. Or the mixture
may be made in a phial of white glass rolled up in a thick towel and
filled in a darkened chamber. The explosion can then be brought
about by carefully rolling the phial out of its envelope into a ray of
sunlight, in a place where the fragments of glass can do no harm. In
this last modification of the experiment the phial is, of course, left
corked. The operator should stand behind a window-shutter or other
suitable screen.
Still another method is to place the bottle in a shady place, and by
means of a looking-glass reflect upon it a ray of sunlight. The moment
the beam touches it, the bottle will explode.
A mixture of the two gases may be kept in the dark for any




CHLORINE UNITES WITH METALS.              109
length of time without change; in diffused daylight they usually
unite only slowly and gradually; but in direct sunlight the union
is so instantaneous as to be attended with explosion.
109. Chlorine combines also very readily with many of the
metals, the combination being in several instances attended with
evolution of light.
Exp. 54.-Fill a bottle of at least half a litre capacity with dry
chlorine gas, by displacement; the gas should be dried by passing it
through a tube filled with chloride of calcium, as described in the Appendix, ~ 15. Gradually sift a gramme or two of very finely powdered
metallic antimony into the bottle. The metal will instantly take fire
and fall in a glowing state to the bottom of the bottle. This fire
attends the formation of a compound of chlorine and antimony, a portion of which will be seen pervading the bottle as a white smoke.
This experiment, and indeed all experiments with chlorine, should
be performed only in places where there is a current of air sufficiently
powerful to carry away from the operator the volatile products of the
reaction, together with any chlorine which may escape from the bottle.
As in the case of the union of sulphur with copper (Exp. 1),
so here it will be seen that burning, as commonly understood, is
in no wise peculiar to the union of oxygen with the other elements. In the act of chemical combination heat is always evolved,
and, of course, light as well, if particles of solid matter be present
and become hot enough to be luminous.
Since oxygen is very abundant, we are more accustomed to
witness exhibitions of its chemical action than those of any other
element; but we must not therefore lose sight of the fact that
among the elements there are several which possess chemical
power as great when brought into play, though not as frequently
exhibited as that of oxygen.
Exp. 55.-Into a small dry bottle throw loosely several leaves of the
so-called Dutch-metal (an imitation gold-leaf made from an alloy of
the metals copper and zinc), and invert over it a bottle of dry chlorine.
As the heavy gas falls into the lower bottle, the chlorine attacks the
metal, which becomes red-hot for a moment, shrivels up, and is converted into a mixture of chloride of copper and chloride of zinc. Both
these compounds are readily soluble, the chloride of copper imparting to the water a peculiar green tinge. The term chloride is used to
denote the combination of chlorine with another element, just as the
term oxide denotes a compound of oxygen.




110             CHLORINE BURNS IN HYDROGEN.
110. A burning jet of hydrogen, on being introduced into a
jar of chlorine, will continue to burn with a peculiar green light,
the two gases uniting to form chlorhydric acid.
Exp. 56.-From a gas-holder containing hydrogen, carry a glass
tube, No. 6, outwards horizontally a few centimetres, then downwards
to reach the bottom of a wide-mouthed litre bottle filled with dry
chlorine; bend the end of the tube, previously drawn to a point,
sharply upward, so that the jet of hydrogen may stream upwards
through the chlorine. Light the hydrogen jet, and insert it into the
bottle of chlorine.
By reversing the experiment, chlorine may just as well be
burned in an atmosphere of hydrogen.
Exp. 57.-In a small flask of 75 or 100 c. c. capacity, provided with
a small chloride-of-calcium tube prolonged into an upright deliverytube which is drawn out to a fine point at the top, generate a free
supply of chlorine. Inflame a jar of hydrogen, held mouth downwards, and press it slowly down upon the chlorine flask so that the
orifice from which the chlorine is issuing may be at the centre of the
hydrogen bottle, in the midst of the gas. In passing through the
burning hydrogen at the bottom of the jar, the chlorine will be heated
to the temperature necessary for its own inflammation, and it will continue to burn in the hydrogen in the same way that oxygen burns in
hydrogen under similar circumstances.
111. The heat evolved during the combustion of hydrogen in
chlorine is less intense than that produced by its union with
oxygen. When one gramme of hydrogen is burned to chlorhydric acid, there are disengaged 23783 units of heat, while
34462 units of heat are evolved when it burns to water.
112. As has been seen, chlorine is both combustible and a
supporter of combustion so far as hydrogen is concerned, and it
exhibits a strong affinity for many of the metals; but it does
not unite directly with either oxygen or carbon.
Exp. 58.-If a burning taper, or a bit of flaming wood or paper, be
thrust into a bottle of chlorine gas, the flame will become murky, and
after struggling for a moment will go out. Much smoke is at the same
time given off.
Exp. 59.-A bit of paper, attached to a wire, dipped in hot oil of
turpentine and then quickly plunged into a bottle of chlorine, will
asually take fire spontaneously, and burn with evolution of dense black




COMBUSTION IN CHlORINE.                 ll1
fime.. On account of the volatility and ready inflammability of oil of
turpentine, it should be carefully heated upon a water-bath (Appendix,
~ 17) in a porcelain dish. If by any chance the turpentine take fire in
the dish, it can be instantly extinguished by covering the dish.
Exp7. 60.-Place an inverted tall bottle full of water upon the shelf
of the water-pan and fill it two-thirds full of chlorine; then displace
the rest of the water with ordinary illuminating gas. Cover the
mouth of the bottle with a glass plate, and, removing it from the
water-pan, place it in an upright position upon the table. Remove
the cover, and touch a lighted match to the gas; fire will be propagated from above downwards, while clouds of smoke are evolved.
Hold a piece of moistened blue litmus paper in the smoke; it will be
reddened by the chlorhydric acid which has been formed.
The wax, wood, paper, turpentine, and gas of the foregoing
experiments, and indeed most of the substances ordinarily used
as combustibles, contain hydrogen and carbon. The hydrogen of
these things will burn in chlorine, will unite chemically with the
chlorine to form chlorhydric acid; but the carbon will not thus
unite with chlorine. Hence it is that, in the experiments in
question. the combustion is at the expense of the hydrogen; the
hydrogen of the candle, turpentine, and so forth alone unites with
chlorine, while the carbon is set free as lampblack or smoke.
113. Chlorine can even decompose water under certain conditions, taking away its hydrogen, while the oxygen is left free.
This occurs, for example, when a mixture of chlorine and aqueous
vapor is passed through a red-hot glass or porcelain tube filled
with fragments of the same material. So, too, when an aqueous
solution of chlorine is exposed to light, the water is gradually
decomposed, as has been stated in ~ 107, oxygen being set free,
and chlorhydric acid formed.
2C1 +  H20 =  2HC1 +  O.                   Fig. 36.
Exp. 61.-Fill a narrow-mouthed bottle, of
the capacity of at least half a litre, with water
which has been saturated with chlorine at a
comparatively low  temperature-such as is
readily obtained by immersing the receiver in
ice-water during the absorption of the gas.
By means of a perforated cork, or better, a
caoutchouc stopper, fit tightly to the bottle a
glass tube, No. 6, bent twice at right angles,




112                    USES OF CLORINoE.
one branch of which shall be long enough to reach to the bottom of
the bottle, while the other arm, made much shorter than the first,
dips into an open beaker glass half full of water.
Place the apparatus in such a position that it shall be exposed to as
much direct sunlight as possible. After a time oxygen gas will begin
to collect at the top of the bottle, and in the course of several hours,
or days, so much will have collected that it can be tested by removi'ng
the cork from the bottle and thrusting in a glowing splinter. The
liquid displaced by the oxygen flows over, through the tube, into the
beaker glass. The chlorhydric acid of course remains dissolved in the
water of the bottle.
114. The applications of chlorine in the arts depend upon that
readiness to combine with hydrogen which has just been exemplified.  By virtue of this affinity for hydrogen, chlorine acts
indirectly as a powerful oxidizing agent. It acts as a purveyor
of nascent oxygen, and is hence a much more efficient agent than
free oxygen, such as exists in the air. Its chief uses are for
bleaching cotton goods, paper stock, and so forth, and for destroying foul and unhealthy emanations.
Exp. 62. —Pour into a test-glass a quantity of chlorine-water
(Exp. 52), drop into it a small quantity of a solution of indigo, and
stir the mixture with a glass rod. The blue color of the indigo will
be immediately destroyed.
In the same way the color of litmus, cochineal, aniline-purple,
or of flowers, calico, and the like, can be readily destroyed by
immersion in chlorine-water or in moist chlorine gas.  The presence of water is essential; perfectly dry chlorine will not bleach.
Exp. 63.-Fill a glass tube, No. 1, about 20 c.m. long, with scraps
of coloured calico and bits of paper which have been written upon with
ink. Take care that the tube and its contents are perfectly dry, and
that the tube is closed at either end with a cork, through which passes
a short piece of tubing, No. 6. Place the tube in a vertical position,
and pass into it, from below, chlorine gas which has been thoroughly
dried by means of chloride of calcium (Appendix, ~ 15). The coloring-matters will not be destroyed so long as they remain dry; but if,
after the dry chlorine has been allowed to act for a few minutes, a little
water be poured in at the top of the tube, so that its contents may be
wetted, they will be bleached at once.
115. Those coloring-matters which are of vegetable or animal
origin are for the most part complex compounds of carbon, hy



HOW CHLORINE BLEACHES AND DISINFECTS.       113
drogen, nitrogen, and oxygen. When moist chlorine is brought
into contact with them, a somewhat complicated reaction occurs;
a portion of their hydrogen is no doubt taken out by the chlorine,
but at the same time some of the water which is present is decomposed, and its oxygen assists the disorganization of the compound
which is to be destroyed.
Of the hydrogenized or carburetted compounds exposed to the
action of the nascent oxygen in the foregoing experiment, those
which are most complex, and of which the elements are held
together least firmly, will of course be acted upon, burned up,
and destroyed. As a rule, the coloring-matters are far more
easily oxidized than the cotton cloth; hence they can readily. be
removed by the action of chlorine, without injury to the cloth.
But if the action of the chlorine were to be continued after the
coloring-matter had been destroyed, the cloth itself would gradually be burned up.
In actual practice, where th9 duration of the exposure of the
cloth to the chlorine is carefully regulated, and the portions of
bleaching liquor which at first remain adhering to the cloth are
completely removed by washing and by chemical treatment, the
process is perfectly safe and trustworthy as regards cotton or
even linen; but the animal fibres, such as wool and silk, are of
more complex composition than cotton and linen; they cannot
be bleached by chlorine, since this gas would attack and disorganize them.
116. In destroying noxious effluvia, chlorine either acts upon
them as upon coloring-matters, or it simply takes away hydrogen,
as in the case of sulphuretted hydrogen hereafter to be studied.
Putrid animal matter may be rendered comparatively odorless, by
sprinkling it copiously with chlorine-water; hence a solution of
chlorine finds some application in inquests and judicial investigations.
The energy with which chlorine seizes upon hydrogen may be
further illustrated by causing chlorine to act upon ammonia-water.
Exp. 64.-Into a glass tube, No. 1, about a metre long, pour enough
chlorine-water to fill it nine-tenths full, and then ammonia-water
enough to fill the remaining space. Close the tube with the thumb,
invert it and place it in an upright position upon the-water-pan. The




114            ACTION OF CHLORINE ON AMMONIA.
ammonia-water, being specifically lighter than the solution of chlorine, will flow upwards and become mixed with the latter; a reaction
will immediately ensue; some of the chlorine wvill unite with the hydrogen of a portion of the ammonia, to form chlorhydric acid, and
nitrogen will be set free. Numberless little bubbles of this gas will
escape from the liquor and collect at the top of the tube, and may be
subsequently tested with a burning match. The chlorhydric acid
formed unites with the remainder of the ammonia to form chloride
of ammonium:4NH3 + 3C1 = N + 3NH4C1.
By modifying the apparatus employed in the foregoing experiment,
so that a currenit of chlorine can be passed into a vessel containing
ammonia-water, the evolution of nitrogen can readily be made continuous, and large quantities of the gas may be collected. It would
be an excellent and easy method of preparing nitrogen for use in the
laboratory, were it not that care must be taken that the ammonia
shall always be present in considerable excess. If this precaution
were neglected, there might be formed, by the action of the chlorine
upon the chloride of ammonium, a very dangerous compound called
chloride of nitrogen. As prepared by this method, the nitrogen is
always contaminated with a certain amount of oxygen.
In the foregoing experiment, the chloride of ammonium which
is produced remains dissolved in the water. It may be recovered
by evaporating the water, or a new portion of it may be prepared
by mixing ehlorhydric acid with ammonia.
Exp. 65.-Fill one half-litre bottle with dry ammonia-gas, and
another with dry chlorhydric acid gas. Invert the latter, and place it
over the former, so that the mouth of the upper bottle shall rest upon
that of the lower. The gases will immediately unite to form solid chloride of ammonium, a dense white cloud of which will fill the bottles:NH3 + HC1 = NH4C1.
One volume of ammonia unites with one volume of chlorhydric acid,
and the gases are completely condensed to a white solid.
117. Chloride of nitrogen, the dangerous compound of chlorine
and nitrogen which has been alluded to above, is formed when
chlorine is brought into contact with a weak solution of chloride
or nitrate of ammonium at the temperature of 150 or 200. As the
chlorine is gradually absorbed, yellow oily drops of chloride of
nitrogen form upon the surface of the liquid, and soon fall to the
bottom:NHICl +  6C 1 =  4HIC1 +  NC1,.




CHLORIDE OF NITROGEN.              115
Chloride of nitrogen is a volatile yellow oil, of peculiar, penetrating odor; it is insoluble in water, and does not congeal when
exposed to cold. Its specific gravity is 1'653. It decomposes
very easily. Upon being heated to nearly 1000, or touched with
any fat or oil, with turpentine, or with various other substances,
it explodes with extreme violence; indeed it often explodes
spontaneously, without any apparent cause. A single drop of it,
exploded upon a glass or porcelain dish, shatters the vessel to
atoms. The preparation and handling of this body require the
greatest caution; it should never be prepared by the novice in
chemistry.
118. We have heretofore adduced experimental proof of every
proposition and statement so far as was possible at such a stage
of the student's progress. The chemical properties of the four
elements, oxygen, nitrogen, hydrogen, and chlorine, have been
exhibited by experiment, the composition of many of their most
important compounds has been demonstrated by analysis or by
synthesis, or by both these methods, and the chemical properties
of these compounds have been illustrated by actual experiment.
Several objects have been thus attained: —first, experimental
methods of research have been illustrated by tangible examples;
secondly, the foundation, scope, and application of important
laws of chemical combination have been explained; thirdly, four
leading elements have been minutely studied-hydrogen (the
standard atom and the unity of specific gravity for gases),
oxygen, nitrogen, and chlorine (three widely diffused elements,
each of which is the first member and prototype of an important
group of elements, many of whose properties we shall hereafter
find we have already become acquainted with in studying the
prototypes); fourthly, three compounds of these elements have
been carefully studied-chlorhydric acid, water, and ammonia
-compounds which are not only interesting in themselves, but
of great significance as types, or models, of three large groups of
compounds whose properties we have really been studying while
we studied their types.:From this point forward the student will be asked to accept
on trust many facts, drawn from the accumulated stores of the
science and resting on satisfactory evidence, the full exposition
I2




116                  OXIDES OF CHLORnI.
of which would be both tedious and inappropriate. The subjectmatter of chemistry is so vast and various that it will be necessary to select from the great mass of material only the most
valuable portions, and to dwell only on those elements and compounds which are of practical importance in the useful arts, or
which are of interest in connexion with instructive theories or
recognized laws of the science.
119. Compounds of Chlorine and Oxygen.- Free chlorine does
ot combine directly with free oxygen. But by resorting to
indirect methods several compounds of the two elements can be
obtained. As many as five different oxides of chlorine, enumerated below, have been described, though as yet some of them
are known only in combination with water or other substances,
and not in the free condition:
NAMES.                 COMPOSITION.              FORM-ULA.
By volume.   =        By weight.
Chlorine. Oxygen. o    Chlorine.  Oxygen.
]Iypochlorous acid. 2 vols. +1 vol.  2  35:5 X 2=71  16 C120
Chlorous acid........ 2 vols. +3 vols.?  355 X2=71 16X3= 48 C1,03
Hypochloric acid... I vol. +2 vols. 2  35-5  16X2=  32 C102
Chloric acid.........2 vols. +-5 vols.?  35'5X2=71  16X5=  80 CL,0O
Perchloric acid...... 2 vols. +7 vols.?  35'5x2=-71  16X7=-11  C1207
120. Hypochlorous Acid (C120).-If a small quantity of slaked
lime (hydrate of calcium) be thrown into a bottle of chlorine gas,
and the mixture be then left to itself during several hours, the
chlorine will be completely absorbed, and there will be formed
two compounds, one of which will be found to be hypochlorite
of calcium, the other chloride of calcium. The reaction may be
thus formulated:2(CaO,H,,O0) +  4C1 =  CaO,Cl20 +  CaCl2 +  21120.
This mixture is a substance much used in the arts under the
technical names " chloride of lime," and "bleaching-powder;" it
will be again referred to hereafter.
121. Hydrated hypochlorous acid may be prepared from
" bleaching-powder;" the solution has a yellowish color, an
acrid taste, and a peculiar sweet odor. When concentrated it
decomposes rapidly, even if kept upon ice.  Dilute solutions are




OLOROUS, BYPOCHILORIC, AND CHLORIC ACIDS.    117
more stable, but they decompose slowly upon being boiled. Hypochlorous acid is a powerful oxidizing and bleaching agent.
Its solution produces at once effects which are only slowly obtained when chlorine-water is employed.
Anhydrous hypochlorous acid, which may be obtained by removing the water from the aqueous solution, is a gas of pale
yellow color and offensive odor, somewhat resembling that of
chlorine. It decomposes very easily into 2 volumes of chlorine
and 1 volume of oxygen; even the warmth of the hand is sufficient to decompose it; and it is difficult to preserve it unchanged
even for a few hours.  At low temperatures, such as are produced by a mixture of ice and salt, the gas condenses to a dark
orange-colored liquid, heavier than water and very explosive.
122. Chlorous Acid,'Cl20,, may be obtained by deoxidizing
chloric acid by means of nitrous acid. When in the anhydrous
condition it is a gas of a.yellowish-green color, liquefiable by
extreme cold. It is a dangerous compound to prepare, since at
temperatures above 57~ it decomposes, with explosion, into chlorine and oxygen. It is readily soluble in water, and the solution
possesses strong bleaching and oxidizing properties.  It is a
weaker acid than chloric acid, ~ 124, but. resembles it in many
respects. With metallic oxides it unites to form compounds
called chlorites.
123. Hypochloric Acid, C102.-This very explosive compound
may be prepared by gently heating a mixture of chlorate of
potassium  and concentrated sulphuric acid.  The gas is of a
bright yellow color and aromatic odor. Upon being exposed to
daylight or to a temperature somewhat below the boiling-point of
water, it decomposes into oxygen and chlorine, the decomposition
being usually attended with explosion. The preparation of the
gas is dangerous, and should never be attempted unless upon a
very small scale. At the temperature of a mixture of ice and
salt, the gas condenses to a yellow, highly explosive liquid.
124. Chloric Acid, C120,.-In the present state of science this
is the most important of the compounds of oxygen and chlorine.
It is not known in the free state, and in the hydrated condition
has never been obtained with less than 1 molecule of water,
H20,C1205.




118                    PERCHLORIC ACID.
When a current of chlorine is made to flow into a cold dilute aqueous
solution of caustic potash, a mixture of hypochlorite and of chloride of
potassium is produced:2(K20,H20) + 4C1 = K20,C120 + 2KC1 + 2H20, —
the reaction being analogous to that between lime and chlorine, described in ~ 120. But if the conditions as to the concentration and
temperature of the solution of potash be changed-if, instead of using
a dilute solution, chlorine be passed into a moderately strong hot solution of caustic potash, or of carbonate of potassium, hypochlorous acid
will no longer be formed, but instead of it chloric acid. The reaction
may be expressed as follows:6(K20,H20) +  12C1 =  K20,C1205 +  10KCl + 6H2O.
Chloride of potassium is formed as before, but the remainder of the
chlorine is now more highly oxidized. Chlorate of potassium is less
soluble than chloride of potassium; it separates in flat tabular crystals
after the liquid has been concentrated by evaporation and cooled. It
is the substance which was employed for making oxygen in Exp. 7.
Chloric acid could be prepared directly from chlorate of potassium by
boiling a solution of this substance with fluosilicic acid. An almost
insoluble fluosilicate of potassium would be formed, and chloric acid
set free. But an easier method is to first convert the chlorate of potassium into chlorate of barium, and to liberate the chloric acid from this
salt by means of sulphuric acid, with which barium forms a remarkably
insoluble compound:BaO,C120, + H20,SO3 = BaO,SO,3 + H20,C1205.
The solution of chloric acid is separated from the insoluble sulphate
of barium by filtration, and concentrated by evaporation over sulphuric
acid in the exhausted receiver of an air-pump. By cautious evaporation the acid may be brought to a syrupy consistence, but is then
rather easily decomposed, especially if it be heated or exposed to light.
At the temperature of boiling it is rapidly converted into perchloric
acid, water, chlorine. and oxygen. It is a strong acid, and a powerful
oxidizing and bleaching agent.
125. Perchloric Acid, C1207, is formed, as above stated, when
an aqueous solution of chloric acid is boiled; being volatile it may
be distilled off and collected. A compound of this acid and potassium, perchlorate of potassium, can be obtained by heating
chlorate of potassium to a certain temperature. Perchloric acid
is a more stable compound than either of the other oxides of
chlorine. The dilute aqueous solution may be concentrated by




BROMINE.                       119
evaporation over fire, even at high temperatures. The hydrated
acid HO,C 1207 is a colorless, oily liquid, which boils at 203~,
and has a specific gravity of 1'782. It is a powerful oxidizing
agent.
The student will do well to compare this series of oxides of
chlorine with that of the oxides of nitrogen, and to note the
points in which the two series resemble and those in which they
differ from each other.
CHAPTER IX.
BROMINE.
126. Bromine is an element closely allied to chlorine. It is
found in small quantities in sea-water, and in the water of many
saline springs. 1 litre of sea-water contains from 0'0143 to
0'1005 grm. of it. As it exists in nature it is combined with
metals, bromide of magnesium being the compound most commonly met with. Bromide of magnesium is a constituent of the
uncrystallizable residue, called bittern, which remains after the
chloride of sodium has been crystallized out from the natural
brines; at several saline springs this bittern contains so large a
proportion of the bromide that bromine can be profitably extracted
from it. Most of the bromine of commerce is thus obtained.
In order to obtain bromine from the bittern, the latter is mixed with
black oxide of manganese and chlorhydric acid, and heated in a retort.
Chlorine is of course evolved from these materials in the midst of the
liquid; it reacts upon the bromide of magnesium and sets free bromine,
which distils over into the receiver as a dark-red, very heavy liquid:
MgBr2 + 2C1 = MgC12 + 2Br.
All the metallic bromides are readily decomposed by chlorine, bromine being, as a rule, a less energetic chemical agent than chlorine.
127. At the ordinary temperature bromine is a liquid of dark
brown-red color, about three times as heavy as water, and highly
poisonous. Its odor is irritating and disagreeable, whence the
name bromine, derived from a Greek word and signifying a




120                PROPERTIES OF BROMINE.
stench. It boils at about 600, but is very volatile even at the
ordinary temperature of the air.
Exp. 66.-By means of a small pipette, throw into a flask or bottle,
of the capacity of 1 or 2 litres, 3 or 4 drops of bromine. Cover the
bottle loosely, and leave it standing. In a short time it will be filled
with a red vapor, which is bromine gas. This vapor is very heavy,
more than five times as heavy as air, and eighty times as heavy as
hydrogen.
At about 7~ bromine crystallizes in brittle plates. It dissolves
sparingly in water, but is soluble in alcohol, and in all proportions in ether.
In its chemical behavior, as well as in many of its physical
properties, bromine closely resembles chlorine. Its affinity for
hydrogen, though weaker than that of chlorine, is still powerful.  Like chlorine, it is an energetic bleaching and disinfecting
agent, and it decomposes the vapor of water when passed with it
through a tube heated to bright redness, bromhydric acid and
oxygen being the products of the reaction. A lighted taper burns
for an instant in bromine vapor and is then extinguished. Phosphorns, antimony, potassium, and the like, take fire on being
thrown into bromine, in the same way as in chlorine, a bromide
of the other element being produced.
Bxp. 67.-Fit a thin cork to a large, wide-mouthed bottle; perforate the cork, and through the hole pass a tube of thin glass (No. 2)
closed at one end. The tube should reach nearly to the bottom of the
bottle, and should project two or three inches above the cork. Within
the tube place a few drops of bromine; throw in upon this a very small
quantity of finely powdered antimony, and instantly cover the mouth
of the tube with an inverted crucible or wide-mouthed phial, in order
that nothing may be thrown out of the tube by the violent action
which attends the combination. If the tube be broken, its fragments
will be retained within the bottle.
Bromine is used to a certain extent in medicine, and largely in
photography. In the chemical laboratory it is often employed,
not only for its own sake, but as a substitute for chlorine; for,
though less energetic, it is more manageable than the latter, especially in those cases where a liquid is desirable.
128. Bromhydric Acid (HBr).-In spite of the strong affinity
of bromine for hydrogen, these elements cannot readily be made




BROMHYDRIC ACID.                 121
to unite directly. A mixture of equal volumes of hydrogen and
bromine vapor cannot be made to combine with explosion by exposure to the sun's rays, or by the introduction of a burning
lamp, though a certain amount of combination occurs in the immediate neighborhood of the flame. But by immersing in the
mixture a platinum wire, kept red-hot by a galvanic current, the
two elements may be made to unite slowly; and a similar result
is obtained by passing the mixed gases through a red-hot tube.
Bromhydric acid gas can, however, be readily prepared by decomposing bromide of potassium with sulphuric acid, or, better,
with a concentrated solution of phosphoric acid. The reaction is
analogous to that in which chlorhydric acid is obtained from
chloride of sodium:2KBr +  HO,SO, =  K20,SO, +  2HBr.
If sulpharic acid be employed, a secondary reaction occurs; a
small part of the bromhydric acid suffers decomposition, and the
product is slightly contaminated with free bromine and with
sulphurous acid:
2HBr + H20,SO3 = 21120 + 2Br + SO2.
Since phosphoric acid does not thus decompose bromhydrio
acid, the latter can be obtained in a state of purity by distilling a
mixture of bromide of potassium and phosphoric acid.
129. Bromhydric acid is a colorless, irritating gas, which, on
coming in contact with the moisture of the air, fumes even more
strongly than chlorhydric acid. By powerful pressure it can be
reduced to the liquid condition, and upon being exposed to intense
cold it may be obtained in the form of a crystalline solid.
It is readily soluble in water, forming a strongly acid solution
which resembles chlorhydric acid in many respects, and, like it,
fumes in the air. A ready method of preparing the solution is
to decompose a strong solution of bromide of barium with sulphuric acid diluted with its own weight of water. The solution
of the free acid may then be separated from the insoluble sulphate
of barium by filtration or by distillation.
When the solution of this acid is mixed with nitric acid, there
is obtained another aqua regia capable of dissolvin ggold and pla



122                     BROMIC ACID.
tinum, like the mixture of chlorhydric and nitric acids, though less
readily.
130. The gas undergoes no change when passed through a
red-hot tube; but it is readily decomposed by metals like potassium at the ordinary temperature, and by tin gently heated. A
bromide of the metal is formed in either case, and there remains
a volume of hydrogen equal to half that of the original gas. Observation has shown that the specific gravity of the gas is very
nearly 40'5, or half the sum of the specific gravities of bromine
vapor and hydrogen. From this fact and the above decomposition of the gas by metals, it follows that bromhydric acid is composed of equal volumes of bromine and hydrogen united without
condensation.
131. Bromic Acid,  20O,Br2O. —Only one oxide of bromine
has been studied, and even this has never been obtained free from
water. It is bromic acid, a substance corresponding to chloric
acid in composition and properties. Its compounds, also, known
as bromates, generally resemble very closely the corresponding
chlorates.
Bromate of potassium can be obtained by the action of bromine upon
potash-lye, in the same way that chlorate of potassium is obtained by
the action of chlorine:
6(K2O,H20O) + 12Br = K20,Br20, + 10KBr + 6H20.
The bromate, which is less soluble than the bromide, can subsequently be separated by crystallization. In order to obtain the hydrated acid, bromate of barium may be decomposed with dilute sulphuric acid:BaO,Br2O, + H2O,SO, = BaO,SO3 + H20,Br20,.
The insoluble sulphate of barium is separated by filtration. The
acid solution can be concentrated to a certain extent by evaporation
at a gentle heat, but cannot readily be brought to a syrupy consistency
without decomposition. It decomposes, also, on being heated to 100~,
and in general gives up oxygen on being brought into contact with substances which readily combine with that element.
132. Hypobromous Acid, H20,Br20. —There can be no doubt
of the existence of an oxide of bromine corresponding to hypochlorous acid. When bromine is added to an excess of a solution
of nitrate of silver, a liquid of strong bleaching-properties is




IODINE.                      123
obtained, from which a yellowish, acid fluid may be distilled.
This distillate bleaches strongly, and yields, on analysis, numbers
corresponding with the above formula. When cold dilute alkaline solutions are mixed with bromine they acquire a power of
bleaching, and in general behave like the alkaline hypochlorites,
which are formed under similar conditions. So too when bromine-water is treated with red oxide of mercury; a sparingly
soluble oxybromide of mercury is formed, together with a bleaching liquor supposed to contain hypobromous acid.
The analogies which subsist between chlorine and bromine are,
however, everywhere so clearly defined that there is good reason
to believe that other oxides of bromine, corresponding to those of
chlorine, will be sooner or later discovered.
133. Chloride of Bromine.-Liquid bromine absorbs a large
quantity of chlorine-gas.when the two elements are brought
together, and there is formed a very volatile liquid called chloride
of bromine. It exhales a pungent, irritating odor, and is soluble
in water; the solution possesses considerable bleaching-power.
134. Bromide of Nitrogen is an explosive compound analogous
to chloride of nitrogen, from which it may be prepared by means
of bromide of potassium.
CHAPTER X.
IOD INE.
135. In its chemical properties iodine bears a striking resemblance to bromine, and consequently to chlorine also. It exists
in sea-water and in the water of many saline and mineral springs.
The proportion of iodine in sea-water is exceedingly small, being
even smaller than that of bromine. But iodine is obtained more
readily than bromine; for iodine is absorbed from sea-water by
various marine plants, which, during their growth, collect and
concentrate the minute quantities of iodine which the sea-water




124                 EXTRACTION OF IODINE.
contains, to such an extent that it can be extracted from them
with profit.
Upon the coasts of Scotland, Ireland, and France, where labor is
cheap, the sea-weeds which contain iodine are collected, dried, and
burned, at a low heat, in shallow pits. The half-fused ashes thus
obtained, called kelp or varec, contains, among other things, iodine in
the form of iodide of sodium and iodide of potassium. It is lixiviated
with boiling water, and the solution is then evaporated and set aside
to crystallize. The iodides, being much more soluble than the other
constituents of the ash, remain dissolved in the mother-liquor after
most of the other salts have crystallized out. If this mother-liquor,
or iodine-lye, be now mixed with a small quantity of sulphuric acid
and left to itself for a day or two, it may be freed from a further portion
of impurity; it is then transferred to a leaden retort, mixed with a
suitable quantity of powdered black oxide of manganese, and gently
heated. Iodine is set free, just as chlorine would be from chloride of
sodium under similar circumstances, and may be collected in appropriate receivers:2NaI + 2(H20,SO3) + MnO2 = Na2O,SO,3 - MnO,SO+ 2H20 + 2I.
136. At the ordinary temperature iodine is a soft, heavy,
crystalline solid of bluish-black color and metallic lustre. Its
specific gravity is 4'948. It melts at a temperature (1070) a
little above that at which water boils, and boils at a somewhat
higher temperature (178~-180~); but in spite of this high boiling-point it evaporates rather freely at the ordinary temperature
of the air, and the more rapidly when it is in a moist condition.
It may be entirely volatilized by heating it upon writing-paper.
Its odor is peculiar, somewhat resembling that of chlorine, but
weaker, and easily distinguished from  it.  Its atomic weight
is 127.
The vapor of iodine is of a magnificent purple color, whence
the name iodine, derived from a Greek word signifying violetcolored; it is very heavy, indeed the heaviest of all known
gases;. it is nearly nine times as heavy as air. The specific gravity
of the vapor is 127.
Exp. 68.-Hold a dry test-tube in the gas-lamp by means of the
wooden nippers, and warm it along its entire length, in so far as this
is practicable. Drop into the hot tube a small fragment of iodine, and
observe the vapor as it rises in the tube. If only a small portion of




PROPERTIES OF IODINE.                12.5
the tube were heated, the vapor would be deposited as solid iodine
upon the cold part of its walls.
137. Solid iodine is never met with in the amorphous, shapeless state in which glass, resin, coal and many other substances
occur.  No matter how obtained, its particles always exhibit a
definite crystalline structure. If the iodine be melted and then
allowed to cool, or if it be converted into vapor and this vapor be
subsequently condensed, crystals will be formed in either case.
Perfect crystals can be still more readily obtained by dissolving
iodine in an aqueous solution of iodohydric acid, and exposing this
solution to the air in a narrow-necked or loosely stoppered bottle;
the iodohydric acid will be slowly decomposed by the action of
the atmospheric oxygen, and, as it decomposes, well-defined crystals of iodine will be deposited.
In whichever way prepared, the crystals of iodine are octahedrons with a rhombic base, belonging to the crystalline system
called trimetric. As commonly seen, the crystals are thin,
flattened tables, distorted by excessive elongation in one direction.
138. Iodine is scarcely at all soluble in water, though enough
dissolves to impart a brown color to the water; but it dissolves
readily in alcohol and ether. These solutions are much used in
medicine, particularly the alcoholic solution, which is called tincture of iodine. When swallowed in the solid state, iodine acts as
an energetic corrosive poison; but several of its compounds, and
the element itself when taken in small doses, are highly prized
as medicaments. It is also largely employed in photography, and
is a useful reagent in the chemical laboratory.
139. As has been already stated, iodine, in its chemical behavior, resembles chlorine and bromine, only its affinities are more
feeble. It enters into combination with less energy than either
of these elements, and is displaced by them from most of its combinations. Like them, it unites directly with the metals and with
several other elements. It gradually corrodes organic tissues,
and destroys coloring-matters, though but slowly. No oxygen is
given off from the aqueous solution when this is exposed to sunlight; but the color of the solution slowly disappears, and a mixture of iodohyvdric and iodic acids is formed in it.




126                  TESTING FOR IODINE.
A singular property of iodine is its power of forming a blue
compound with starch.
Exp. 69.-Prepare a quantity of thin starch-paste by boiling 30 c. c.
of water in a porcelain dish, and stirring into it 0'5 grm. of starch
which has previously been reduced to the consistence of cream by
rubbing it in a mortar with a few drops of water.
Pour 3 or 4 drops of the paste into 10 c. c. of water in a test-tube
and shake the mixture so that the paste may be equably diffused
through the water; then add a drop of an aqueous solution of iodine,
and observe the beautiful blue color which the solution assumes. If
the solution be heated the blue coloration will disappear, but it reappears when the liquid is allowed to cool.
Dip a strip of white paper in the starch-paste and suspend it, while
still moist, in a large bottle, into the bottom of which two or three
crystals of iodine have been thrown. As the vapor of iodine slowly
diffuses through the air of the bottle it will at last come in contact
with the starch, and after some minutes the paper will be colored
blue.
This reaction furnishes a very delicate test for iodine. By its
means it has been proved that iodine, though nowhere very abundant, is very widely distributed in nature; traces of it have been
detected in land plants, and in many well, river, sand spring
waters, also in rain-water, and even in the air; indeed it would
be difficult to say where iodine is not.
In order that it may be detected by this test, the iodine must
be free or uncombined. But, as has been stated, chlorine readily
expels iodine from most of its combinations. In case, then, we
have reason to suspect the presence of a compound of iodine
(iodide of potassium, for example) in any substance, a small quantity of chlorine-water, or of some other agent capable of expelling
iodine, must be added to this substance. Once displaced from
its combination, the iodine may be at once detected by means of
starch.
Exp. 70.-Place in a test-tube 10 c. c. of water, a drop of concentrated aqueous solution of iodide of potassium, and 3 or 4 drops of the
starch-paste of Exp. 69. If the iodide of potassium be pure, no coloration will occur. Add now 2 or 3 drops of chlorine-water, and shake
the tube. The characteristic blue coloration at once appears.
In order to illustrate the extreme delicacy of this reaction, dissolve
0'14 grin. of iodide of potassium in 1 litre of water, and to this solu



DELICACY OF THE IODINE-TEST.             127
tion, which contains 1 part of iodine in 10,000 parts of water, add some
of the starch-paste and several drops of red fuming nitric acid, a reagent on some accounts better fitted than chlorine to disengage iodine
in this experiment (see ~ 150). After a time the solution will exhibit
the blue color, though in solutions so dilute as this it sometimes
happens that the coloration appears only after the lapse of several
hours.
It follows, of course, from the foregoing experiment, that the
reaction of iodine upon starch can be used as a test for those
substances which, like chlorine or nitric acid, are capable of setting free iodine, as well as for iodine itself. In the chemical
laboratory it is customary to keep on hand for this purpose a
store of paper upon which has been spread a mixture of starchpaste and iodide of potassium, prepared as follows:Exp. 71.-Dissolve 0'5 grm. of pure iodide of potassium (free from
iodate) in 100 c. c. of water; boil this solution in a porcelain dish and
stir into it 5 grms. of finely powdered starch, taking care not to burn
the starch, and stirring until the mass gelatinizes. Remove the lamp,
allow the paste to become cold, and by means of a wooden spatula
spread it thinly upon one side of white glazed paper. The paper is
then dried, cut into strips about 8 c.m. long by 2 wide, and preserved
in stoppered bottles kept carefully closed.
Exp. 72.-Place in a test-tube a small quantity of binoxide of manganese, pour upon it 4 or 5 c. c. of chlorhydric acid, heat the mixture,
and hold at the top of the tube a moistened strip of the test-paper
which was prepared in the preceding experiment. The chlorine evolved
by the reaction of the chlorhydric acid upon the binoxide of manganese sets free iodine from the iodide of potassium upon the test-paper,
and the starch is thereby coloured blue. The presence of chlorine in
chlorhydric acid is thus made apparent. By this test we might discriminate, for example, between dilute nitric and chlorhydric acids.
140. lodohydric Acid (HI). - Hydrogen and iodine do not
readily unite together directly. There is here nothing to recall
the explosive violence with which chlorine and hydrogen combine. Sunlight has no power to bring about the union of the
two elements at the ordinary temperature; but when a mixture
of hydrogen gas and iodine vapor is passed through a red-hot
tube, iodohydric acid is formed. It has been observed, also, that
spongy platinum will cause the union of the two elements even at
ordinary temperatures. Even when indirect methods are resorted




128                    IODOHYDRIC ACID.
to, it is less easy to prepare iodohydric acid than chlorhydric or
bromhydric acids.
If iodide of sodium be distilled with sulphuric acid, there will be
obtained but little iodohydrie acid; for most of that which is produced
at first will be subsequently destroyed by the action of sulphuric acid,
in the same way as happens to a less extent with bromhydric acid,
~ 128.
As fast as iodohydric acid is formed in accordance with the reaction
2NaI +  H2SO4 = Na2SO4 + 2HI,
most of it is decomposed by another portion of sulphuric acid, in a
manner which may be thus represented:2HI + H2SO4 = 2H,0 + SO2 + 2I.
Solutions of iodohydric acid can, however, be readily obtained by the
action of iodine upon a compound of sulphur and hydrogen, called
sulphydric acid. In practice, a current of sulphydric acid gas is made
to pass through water in which finely divided iodine is kept suspended
by agitation. The sulphydric acid, the formula of which is H2S,
reacts upon 21, and there is formed 2HI and free sulphur, which is
deposited.
A solution of iodohydric acid may also be obtained by distilling a
mixture of iodine, phosphorus, and much water, in which case the
phosphorus unites with the oxygen of a portion of the water, while
the iodine takes the hydrogen. Or it may be prepared by decomposing an aqueous solution of iodide of barium with an equivalent quantity of dilute sulphuric acid, and filtering off the solution from the
insoluble sulphate of barium.
141. The dilute acid obtained by either of these methods can
be concentrated, by evaporation, to a liquor of 1'7 specific gravity,
boiling at 1270, and composed of one molecule of iodohydric acid
united with 11 molecules of water.  The aqueous solution has
a sour, suffocating odor, and pungent acid taste. When concentrated it fumes strongly in the air. It cannot be long preserved
when exposed to contact with the air, for the oxygen of the air
unites with its hydrogen, and iodine is set free. At first this iodine
dissolves in that portion of the iodohydric acid which has not yet
been decomposed; but after the acid has become saturated, crystals
of iodine are deposited, as has been stated in ~ 137. The decomposition of iodohydric acid is so rapid that the pure, colorless solution of it becomes red from separation of iodine after a few




YODOJIYDRIC ACID.                 129
hours' exposure to the air, no matter whether it be minute or
concentrated. The easy decompositon of this acid shows clearly
with how mulch less force hydrogen holds iodine in combination
than it holds either chlorine or bromine.
142. The usual method of preparing anhydrous iodohydric
acid is as follows:
In the bottom ot a test-tube place a mixture of 9 parts of iodine and
1 part of phosphorus. Cover the mixture with coarsely powdered
glass, and bring about chemical union between the iodine and the
phosphorus by gently heating them. Place now a few drops of water
in the tube, and connect with it a gas delivery-tube by means of a
caoutchouc stopper. Iodohydric acid will be immediately given off;
and may be collected by displacement.
Another method is to pack a test-tube with alternate layers of phosphorus, iodine, and moistened glass-powder, and then to gently heat
the tube. The operation depends upon the formation of an iodide of
phosphorus and the subsequent decomposition of this body by contact
with water into iodohvdric acid and a compound of phosphorus, oxygen,
and water, called hydrated phosphorous acid: —
2PI3 + 6H20 = 6HI + 3H20,P203.
143. Iodohydric acid is a colorless, acid gas, of suffocating
odor; it fumes strongly in the air, and is very soluble in water.
It can be liquefied rather easily by pressure, and solidified at — 51~
to a colorless mass like ice. The gas is more than four times as
heavy as air, its specific gravity having been found by observation to be 64-11. From this fact, taken in connexion with the
striking analogy which the compound bears to bromhydric and
chlorhydric acids, it follows that the gas is composed of equal
volumes of iodine vapor and hydrogen united without condensation; for the theoretical density of a gas thus composed would be
(127+1) — 2=64, a number with which the observed specific
gravity closely agrees. The chemical effect of the small proportion of hydrogen contained in iodohydric acid is most remarkable.  Only T-1-, or less than 1 per cent. of iodohydric acid is
hydrogen, yet this very small proportional quantity of hydrogen
is competent to impart an entirely new set of properties, both to
the iodine and the hydrogen; the acid bears no resemblance to
either of its constituents.
144. lodohydric acid is a compound which decomposes easily.
E




130                     IODIC ACID.
When a mixture of the gas and oxygen is passed through a redhot tube, water and free iodine are the products. Chlorine and
bromine abstract hydrogen from it, and leave iodine free; and
the same effect is produced by many oxygen compounds which
readily part with oxygen. With many of the metals it forms
iodides, while hydrogen is set free; and it reacts upon most of the
metallic oxides, forming water and a metallic iodide.
Though the hydrogen of iodohydric acid is readily removed by
means of oxygen in numerous instances, it appears, upon the
other hand, that iodine can abstract hydrogen from most of its
combinations with the other elements. Only oxygen, chlorine,
bromine, and an element (still to be studied) called fluorine, exhibit a stronger tendency than it to unite with hydrogen. Iodine
separates hydrogen from its compounds with nitrogen, sulphur,
and phosphorus, and from many organic compounds, such as alcohol and ether, iodohydric acid being formed in each case.
145. Compounds of Iodine and- Oxygen.-Of the compounds of
iodine and oxygen, only two have as yet been carefully studied.
These correspond respectively to chloric and perchloric acids.
Compounds analogous to hypochlorous and hypochloric acids appear to exist, but have not been described with much accuracy.
146. Iodic acid (I200) may be obtained directly by oxidizing
powdered iodine with monohydrated nitric acid at a moderate
heat. After all the iodine has disappeared, and the excess of
nitric acid employed has been evaporated, iodic acid will be left
as a white residue.
Iodic acid is readily soluble in water, and crystallizes from an
acidulated solution in colorless, six-sided tables, of the formula
HIO3 or H20,I205.  It has a peculiar odor, and acid, disagreeable
taste. At the temperature of 170~, water is given off and the
anhydrous acid remains. This melts upon being heated more
strongly, and suffers decomposition.
Iodic acid readily gives up oxygen to many other substances,
or, in other words, it is easily decomposed by reducing-agents;
for example, when mixed with iodohydric acid it reacts upon it
with formation of water and deposition of iodine:10HI +  I,0, = 5H20 +  121.
All of the metals are, oxidized by it, excepting gold and platinum.




IODIDE OF NITROGEN.                131
With metallic oxides it forms compounds called iodates, which are
analogous to the corresponding chlorates and bromates in composition and properties.
147. Periodic acid (I20,) may be prepared by passing chlorine
gas through a solution of iodate of sodium mixed with caustic
soda. Chloride of sodium and basic periodate of sodium will be
formed, and the latter, being sparingly soluble in water, will be
deposited in crystals:Na20,I205 + 3(Na20,H120) + 4C1=2Na20,I207+ 4NaCl + 31120.
If now the sodium salt be collected and dissolved in water, and
the solution be mixed with nitrate of lead, a periodate of lead
will be obtained; this may be decomposed by means of dilute
sulphuric acid into periodic acid and insoluble sulphate of lead.
The latter may then be separated by filtration, and the clear solution of the acid finally concentrated by evaporation.
From the concentrated aqueous solution periodic acid separates
in colorless hydrated crystals, which, upon being carefully heated,
give off water and yield as a residue the anhydrous acid 1207.
At a still higher temperature, the anhydrous acid decomposes
and gives off oxygen. It is decomposed also by reducing-agents
in the same way as iodic acid.
The other compounds of iodine and oxygen have but little interest for us, except that they serve to increase the number of
analogies which subsist between iodine, bromine, and chlorine.
148. Iodide of Nitrogen (?).-There appear to be a number of
compounds which have hitherto been usually classed under this
title. They are produced by the action of ammonia upon iodine,
and are mostly of a highly explosive character, though their properties and composition vary to a certain extent according to the
mode of their preparation.
Exp. 73.-Place 0'25 grm. of finely powdered iodine in a porcelain
capsule, and pour upon it so much concentrated ammonia-water that
the iodine shall be somewhat more than covered; allow the mixture
to stand during 15 or 20 minutes, when an insoluble dark-brown
powder will be found at the bottom of the liquid. This powder is the
so-called iodide of nitrogen. It should be collected upon two or three
very small filters and well washed with cold water. Remove the
filters, together with their contents, from the funnels, pin them upon
bits of board, and leave them to dry spontaneously.
2




132                  CHLORIDES OF IODINE.
As soon as the powder has become thoroughly dry it will explode
upon being rubbed, even with a feather, or jarred, as by the shutting
of a door, or by a blow upon the wall or table. Though incomparably
less dangerous than chloride of nitrogen, and therefore better suited
than the chloride to illustrate the explosive character of this obscure
class of nitrogen compounds, iodide of nitrogen must nevertheless be
handled with great care, and should never be prepared by the student
except in very small quantities.
149. Chlorides of Iodine.-Iodine combines directly with chlorine in several proportions, a protochloride, IC1, and a terchloride,
C13,, being the best-known of these compounds.
The protochloride is obtained by passing dry chlorine over dry
iodine, the current of chlorine being checked at the moment when all
the iodine has become liquid. Or it may be made by distilling iodine
with chlorate of potassium, and collecting the product in a cooled
receiver.
3KC103 4+ 2I = KC104-  KIO3 + KC1 + 20 + IC1.
Protochloride of iodine is a reddish-brown, oily liquid, volatile, irritating, and of penetrating odor. It decolorizes litmus and indigo, but
does not give a blue color with starch.
The terchloride may be produced by treating iodine with an excess
of chlorine gas, or by acting upon anhydrous iodic acid with dry
ehlorhydric acid gas:I205 + 10HC1 =  2IC1, + 5H20 + 4C1.
It is a yellow crystalline solid, melting at 20~-25~. It acts upon
other substances in the same manner as the protochloride; like the
protochloride, it decolorizes indigo and does not turn starch blue.
150. A knowledge of the properties of the chlorides of iodine is
of some practical importance, since they are liable to be formed
incidentally in several chemical processes, which their presence
perturbs. Thus, in the manufacture of iodine, as described under
~ 135, the iodine-lye almost always contains a certain proportion
of chloride of sodium. It is evident that if the chlorine in this
compound were to be evolved at the same time as the iodine by
the action of the black oxide of manganese and sulphuric acid,
there would be formed a quantity of the very volatile protochloride
of iodine, which would escape condensation.  Whatever of iodine
was thus combined with chlorine would be lost to the manufacturer.
But, as has been repeatedly stated, iodine is an element which




THE CHLORINE GROUP.                133
can be much more readily expelled from its combinations than
chlorine; and in the case in point it is found that the iodine in
the mixture of iodide of sodium and chloride of sodium, which
the iodine-lye contains, will all come off before the chlorine, if
the distillation be slowly conducted. If, through irregular heating, any portion of the contents of the retort should become hotter
than the rest, and so lose all its iodine, chlorine would be disengaged from that portion, and would unite with the vaporized
iodine which fills the retort. To ensure the necessary slow and
equable heat, the retort is set upon a stove suitable for the
maintenance of a slow fire, and is provided with an agitator, by
means of which its contents may be continually stirred.
Again, in testing for iodine, as in Exp. 70, chlorine is a far
less convenient agent for setting free the iodine from its combinations than fuming nitric acid; for if the slightest excess of
chlorine be employed, the iodine will all be converted into chloride of iodine, and the starch will not be colored blue.
151. Bromides of Iodine.-There are two compounds of bromine and iodine, and their properties are analogous to those of
the chlorides of iodine.
152. Chlorine, bromine, and iodine constitute one of the most
remarkable and best-defined natural groups of elements. Whether
we regard the uncombined elements or their compounds, it is impossible not to be struck with the close analogies which subsist
between them. With hydrogen, all of these elements unite in the
proportion of one volume to one volume, without condensation, to
form acid compounds extremely soluble in water and possessing
throughout analogous properties.
1+        =   HC1
[] +      =   HIIBr
With oxygen each of them forms a powerful acid containing five




134                  THE CHLORINE GROUP.
atoms of oxygen, besides divers other compounds of obvious likeness. The compounds furnished by their union with any one
metal are always isomorphous (like-formed); the chloride, bromide,
and iodide of potassium, for example, all crystallize in cubes.
With nitrogen they all form explosive compounds. Many similar
analogies will be made manifest as we proceed to study the other
elements, and their compounds with this chlorine group.
There is a distinct family resemblance between these three
elements as regards their physical as well as their chemical characteristics; but, in all their properties, a distinct progression
is observable from chlorine through bromine to iodine. At the
ordinary temperature chlorine is a gas, bromine a liquid, and
iodine a solid, though at temperatures not widely apart they are
all known in the gaseous and liquid states. The specific gravity
of bromine vapor is greater than that of chlorine, and that of
iodine greater than that of bromine. Chlorine gas is yellow, the
vapor of bromine is reddish brown, that of iodine violet. So with
all their other properties,-chlorine will be at one end of the
scale, iodine at the other, while bromine invariably occupies the
intermediate position.
The properties of many of the compounds of chlorine, bromine,
and iodine exhibit a similar progression as we pass from the
chlorine compounds to those of iodine. For example, the specific
gravity of
Chlorhydric acid gas is.....    182
Bromhydric,,....            40'5
lodohydric,,....64'0
Chlorhydric acid can be liquefied at about — 80, and has not yet
been solidified. Bromhydric acid liquefies at about - 600, and
solidifies at about — 92~.  Iodohydric acid liquefies at about
-40~, and solidifies at about -50~.
Chlorhydric acid is a more energetic acid than bromhydric,
and bromhllydric acid is more powerful than iodohydric. The
aqueous solution of chlorhydric acid can be kept without change
in contact with air; that of bromhydric acid becomes colored
after a while, from separation of bromine; but the solution of
iodohydric acid decomposes rapidly, and much iodine is deposited.
As regards the relative chemical power of these elements, it




FLUORINE.                       135
has already been shown that the intensity of this force becomes
less as we descend from chlorine to iodine. It is easy, for example, to displace iodine from its compounds by means of bromine,
NaI +  Br =  NaBr + I,
and equally easy to displace bromine from its combinations by
means of chlorine,
NaBr +  Cl =  NaCl +  Br.
153. It is an important principle, borne out by most of the
other groups of elements, and emphatically true of the natural
family now under consideration, that, with kindred elements, the
chemical power of each is great, in comparison with that of the
related elements, in proportion as its atomic weight is low.
Among the members of a natural chemical group, chemical
energy seems to be inversely proportional to atomic weight. Thus
the atomic weight of chlorine is 35'5, that of bromine 80, and that
of iodine 127, while the chemical energy of these elements follows the opposite order.
154. It is noteworthy that elements of like character almost
always occur associated with one another in nature. Bromine
and iodine are always found in company with chlorine. That
this should be so is in nowise surprising. Those elements which
are similar in character and properties must necessarily be similarly acted upon by the natural forces to which they are exposed,
and must therefore inevitably tend to be gathered or deposited
in like places under like conditions.
CHAPTER XI.
FLUORI NE.
155. There is another substance, called fluorine, which is closely
analogous to chlorine. This element cannot be readily obtained
in the free state, and scarcely anything is known of it in that
condition. Special interest attaches to it upon this very account,




136                   FLUORHYDRIC ACID.
and many fruitless efforts to isolate it have been made. Of all the
elements, it appears to have the strongest tendency to enter into
chemical combination; at all events it is the most difficult to obtain, and to keep, in the free and uncombined condition.
It is not only difficult to expel fluorine from the minerals in
which it is found in nature, but on being set free from one compound it immediately attacks whatever substance is nearest at
hand, and so enters into a new combination. Hence it is wellnigh impossible to collect it.  It destroys at once glass, porcelain,
and metal, the materials from which chemical apparatus is usually
constructed. Vessels made of the mineral fluor-spar (a compound
of fluorine and calcium), are the only ones which have as yet
been found capable of withstanding its action. By operating in
such vessels, a small quantity of impure fluorine gas appears to
have been really obtained; but the process is difficult, expensive,
and not uniformly successful.  Little or no doubt, however, is
entertained as to the general nature of fluorine, since its compounds are closely analogous in many respects to the corresponding compounds of chlorine, bromine, and iodine.
The symbol of fluorine is Fl.  Its atomic weight is 19.  It
occurs tolerably abundantly in nature as fluoride of calcium
(CaFl2), in the mineral known as fluor-spar. Small quantities
of fluorine are found also in several other minerals, in vegetable
and animal substances, particularly in bones; and traces of it
occur in sea-water, and in various rocks and soils. It appears
to be almost as widely disseminated as iodine, though, from the
lack of delicate tests for fluorine, it is far less readily detected.
Of late years a considerable mine of a fluorine mineral called cryolite (fluoride of sodium and aluminum) has been worked in Greenland.
156. Fluorhydric Acid (F111). -With hydrogen, fluorine forms
a powerful acid corresponding to chlorhydric acid and the other
hydrides of the chlorine group. It is a more energetic acid than
either of these, but is specially characterized by its corrosive
action upon glass.  It may be readily prepared by distilling
powdered fluor-spar with strong sulphuric acid; the reaction
being analogous to that which occurs when common salt is treated
with sulphuric acid:



FLUORHYDRIC ACID.                 137
CaFl2 +  H2SO4 = CaSO4 +  2HF1.
Since the acid rapidly corrodes glass, the process must be conducted in metallic vessels.  Ordinarily, retorts of lead or platinum
are employed, and the distillate is collected in receivers made of
the same metals, and carefully cooled by means of ice.
157. The product of the distillation is a very volatile, coiorless
liquid, a little heavier than water.  It is strongly acid, emits
copious white and highly suffocating fumes in the air, boils at
150~, and remains unfrozen at -20~. On account of its corrosive
power, this substance is highly dangerous; if any of it happens
to come in contact with the skin, wounds are produced which are
very difficult to heal; a single drop of it is sufficient to occasion
a deep and painful sore. In preparing the acid, special provision
must be made for carrying away from the operator any fumes
which may escape condensation.
The acid may be kept in bottles made of lead or silver, or of
gutta percha, substances upon which it has no action.  It unites
with water with great avidity, so much heat being evolved that
a hissing noise is produced, as if a bar of red-hot iron had been
immersed in the water.  In its concentrated form the acid has a
specific gravity of 1'061, but on the addition of a certain amount
of water the density increases to 1'15, a de-te hydrate
(HIFl+2HIO) being formed, which boils at 120~, and may be
distilled unchanged. The further addition of water to this hydrate is attended with a regular decrease in density.
According to some chemists, the liquid acid obtained as above
described is not anhydrous. It is asserted that if it be distilled
with an excess of anhydrous phosphoric acid (a substance which
has a very strong affinity for water), the anhydride will be set
free in the form of a colorless, extremely irritating gas.
158. Upon metals and metallic oxides, fluorhydric acid acts
like chlorhydric acid, only more powerfully; but its most striking
peculiarity is its action upon silica and the compounds of silica,
such as glass or porcelain. If a drop of the concentrated acid be
allowed to fall upon a piece of glass, it becomes hot, boils, and
partially distils off as a fluoride of silicon, while the glass is corroded and becomes covered with a white powder consisting of compounds of fluorine and various constituents of the glass. If this




1]38            ETCHING BY FLUORHYDRIC ACID.
powder be washed away a deep impression will be found upon the
glass at the point where the acid has acted.
This corrosive power, which is possessed by fluorhydric acid
gas as well as its aqueous solution, is made use of for etching
glass. The graduations on the glass stems of thermometers and
eudiometers may thus be made with great precision and facility;
the acid is largely employed also in ornamenting glass with
etched patterns.
Exp. 74.-Warm a slip of glass and rub it with beeswax so that it
shall be everywhere covered with a thin, uniform layer of the wax.
With a needle, or other pointed instrument, write a name, or trace
any outline through the wax, so as to expose a portion of the glass.
Lay the etching, face downward, upon a bowl or trough of sheet-lead,
in which has been placed a teaspoonful of powdered fluor-spar and
enough strong sulphuric acid to convert it into a thin paste; if the
glass be smaller than the opening of the dish, it may be supported
%lpon wires laid across the latter.
Cover the glass and the top of the dish with a sheet of paper, and
then gently heat the leaden vessel for a few moments, taking care not
to melt the wax; then set the di h aside in a warm place and leave it
at rest during an hour or two. Finally melt the wax and wipe it off
the glass with a towel or bit of paper; the glass will be found to be
etched and prroded at the places where it was laid bare by the removal of thlax.
This experiment can be performed more rapidly by covering the
outside of a watch-glass with wax, tracing characters upon this layer,
and then placing the glass upon a small platinum crucible containing a
mixture of fluor-spar and sulphuric acid, which is heated over the gaslamp. The watch-glass is meanwhile kept full of water, in order to
prevent the wax from melting. In this way the etching can be effected
in the course of a few minutes.
Instead of the gas, a dilute aqueous solution of the acid may be employed in this experiment. The concentrated acid of ~ 157, diluted
with six parts of water, answers a good purpose. In this case the
etched surface will appear smooth like the rest of the glass, while in
case the gas is employed the etched portion of the glass will be dull
and rough.
159. No compounds of fluorine with chlorine, bromine, iodine,
nitrogen, or oxygen have yet been discovered, though a sulphur
compound has been obtained, as a fuming liquid, by distilling
fluoride of lead with sulphur. Fluorine is the only element of




OZONE AND ANTOZONE.                139
which no oxygen compound is known; this fact, however, will
appear less remarkable if it be remembered that, in order to obtain oxygeno compounds of chlorine, bromine, and iodine, it is
necessary first to isolate these elements, and to have them in the
free and uncombined condition. Analogy would therefore teach
that a practicable method of preparing free fluorine must be discovered before we can hope to prepare oxides of fluorine.
160. The fact that fluorine forms a powerful acid with hydrogen, connects this element with the three elements (chlorine,
bromine, and iodine) which have last been studied. Many of its
compounds with the metals are analogous in composition to the
compounds of chlorine, bromine, and iodine, and not a few of
these compounds are isomorphous with one another. It is customary thqrefore to study fluorine in connexion with the chlorine
group; but the student should remember that in several respects
it differs widely from chlorine, and that its connexion therewith
is, in any event. less intimate than that of either bromine or
iodine.
CHAPTER XIL
OZONE AND ANTOZONE.
161. Besides ordinary oxygen, such as is found in the air and
has been prepared in Exps. 5 and 7, two other kinds or forms of
this element are known to chemists. These new modifications of
oxygen have received special names, and are called ozone and
antozone respectively.
162. Several other elements, notably sulphur, phosphorus, and
carbon, occur, as oxygen does, in very unlike states, or with very
different attributes, while the fundamental chemical identity of
the substance is preserved. The word allotropism is employed to
express this capability of some of the elements; it is derived
from Greek words signifying of a different habit, or character.
This word serves merely to bring into one category a considerable
number of conspicuous facts, of whose essential nature we have




140                       OZONE.
no knowledge; there is, of course, no virtue in the word itself to
explain or account for the phenomena to which it refers.
163. Ozone is an exceedingly energetic chemical agent, which
resembles chlorine in some respects; it can therefore be advantageously studied in connexion with the chlorine group. Moreover, since ozone and antozone were for a long time confounded
with one another, and since they are really intimately related,
they should, of course, be studied together. The most natural
connexion of these somewhat obscure bodies is with oxygen; but
we are better able to appreciate what is known of the properties
of ozone and antozone now that we have become acquainted with
a number of the elements, and have made ourselves familiar with
a considerable variety of chemical processes and reactions, than
we were at the very outset, when common oxygen was necessarily studied.
164. It had long been noticed that when an electrical macbine
was put in operation a peculiar, pungent odor was developed;
but it is only at a comparatively recent period that it has been
observed that the same odor is manifested during the electrolysis
of water (~ 35), and that this odor resembles that evolved by
moistened pjpsphorus when exposed to the air. It has gradually
been made W   that the odor in each of these cases is due to the
presence of a peculiar modification of oxygen, called ozone from a
Greek word signifying to smell. This modification of oxygen was
at one time erroneously supposed by some to be a high oxide of
hydrogen, of composition H202, or H203; but this view has lately
been completely disproved.
Of the methods of obtaining ozone above suggested, that by
phosphorus will usually be found most convenient.
Exp. 75.-In a clean bottle, of 1 or 2 litres capacity, place a piece
of phosphorus 2 or 3 c.m. long, the surface of which has been scraped
clean (under water) with a knife; pour water into the bottle until the
phosphorus is half covered; close the bottle with a loose stopper, and
set it aside in a place where the temperature is 200 or 300.
In the course of ten or fifteen minutes a column of fog will be seen
to rise from that portion of the phosphorus which projects above the
water, the original garlic odor of the phosphorus will soon be lost, and
the peculiar odor of ozone will gradually pervade the bottle. After




OZONE BY ELECTRICITY.                 141
five or six hours, the bottle will be found to contain an abundance of
ozone for use in the subsequent experiments.
The chemical changes which occur during this experiment are
complicated; it will be enough to say of them that the phosphorus unites with oxygen from the air in the bottle to form an
oxide of phosphorus, which will be studied hereafter under the
name of phosphorous acid; that during this process of oxidation
a portion of the oxygen in the bottle is changed into ozone and
antozone, and that some of the ozone remains, even after many
hours, diffused in the air of the bottle.
165. It must be distinctly understood that no very large quantity of ozone is obtained in the foregoing experiment. At the
best, only a very minute proportion of it will be found in the air
of the bottle. But ozone is a substance possessing great chemical
power, and but little of it is needed in order to exhibit its characteristic properties.
If it be desired to prepare ozone by passing electric discharges
through air or oxygen, either of these gases may be sealed up in narrow glass tubes, through the centres of which are passed platinum
wires, welded tightly into the glass, as shown in Fig. 37, and a series
of sparks from an electrical machine is thrown through the gas in the
tube, during ten or twelve hours. If the experimen;be continued
longer than this, nothing is gained; for the sparks after this time appear
to destroy the ozone previously produced.
To avoid the difficulty last named, a slow current of oxygen may be
forced through a tube open at both ends, and electrical discharges may
be passed through the gas in its transit; a constant stream of ozonized
air will be thus obtained.
Instead of the sparks, the gas within the tube may be subjected to silent discharges of electricity obtained by con- Fig. 37.
necting one of the platinum wires with the ground, the.fC~
other with the prime conductor of an electrical machine,
and slowly turning the crank of the latter. By using a tube
having wires near the top, as in Fig. 37, and closing the
lower end of the tube by immersing it in a bath filled with
an aqueous solution of iodide of potassium, so that the ozone
may be absorbed as fast as it is formed, it has been found
possible, by some experimenters, to transform and remove
all the original oxygen contained in the tube.
166. Ozone is produced not only during the slow oxidation of




142                  OZONE BY OXIDATION.
phosphorus, and by the action of electricity upon air or oxygen;
a certain quantity of it appears to be produced also during other
processes of oxidation. It is readily formed, for example, during
the slow combustion of ether and of various other volatile liquids;
it can be at once produced by plunging a heated glass rod or iron
wire into a mixture of air and ether vapor.
Into a wide-mouthed bottle, a small quantity of ether is poured; the
bottle is shaken for a moment, that the air within it may become
charged with the vapor of ether; the liquid ether, if any remain, is
then poured away, and a large glass rod, or thick iron wire, heated to
about 2500, is thrust into the bottle. The rod must not be too hot, lest
the ozone formed be reconverted into ordinary oxygen; if it be insufficiently heated no ozone is produced.
During the slow oxidation of oil of turpentine, oil of cinnamon, oil
of lemons, and others of the so-called essential oils, at the ordinary
temperature of the air a considerable quantity of ozone is produced.
This may be seen in oil of turpentine which has been kept for a long
time in half-filled bottles, exposed to sunlight, and frequently opened
and shaken. The formation of ozone under these circumstances explains the familiar fact that the corks employed to close bottles containing oil of turpentine and the analogous oils are soon bleached and
corroded. At the same time, antozone is also produced in large quantity, as will J explained hereafter.
If quicksilver, to which a little water and a few drops of a solution
of indigo have been added, be shaken up violently in a large bottle full
of air, the indigo will soon be bleached as if by the action of ozone.
167. One of the best methods of preparing ozone is by treating
a compound known as permanganate of potassium with sulphuric
acid. It should be observed, however, that in this process, as in
all the others, the ozone obtained is mixed with common oxygen;
no available method of isolating ozone in a condition of purity
has yet been made known.
A small quantity of concentrated sulphuric acid is placed in the
bottom of a bottle, and a quantity of pure, dry permanganate of potassium, in fine powder, is added; the proportion of acid to permanganate
should be three parts to two, by weight. A strong smell of ozone will
be at once perceived, and the pasty mass will continue to give off
ozone for a long time.
In this case it is conjectured that a portion of the oxygen of the permanganate of potassium, the empirical formula of which is K2Mn2O8.




PROPERTIES OF OZONE.                 143
actually exists in the compound as ozone, and is given off as such when
the compound is decomposed.
168. As has been already mentioned, the chemical behavior of
ozone is analogous to that of chlorine; it bleaches and destroys
vegetable coloring-matters, and is a powerful disinfectant. Like
chlorine it instantly decomposes the iodides of the metals; upon
this property is based a ready method of testing for its presence.
Exp. 76.-Into the bottle of ozonized air (Exp. 75), thrust a moistened slip of the test-paper, saturated with starch and iodide of potas.
sium, which was prepared in Exp. 71; the paper will instantly acquire
a deep blue tint. As in the case where the test-paper was employed
for detecting chlorine (Exp. 72), so here, the reaction depends upon
the displacement of the chemically feeble iodine by the more powerful
ozone:2K1  + O = K20 + 21.
The ozone here acts as oxygen, in one sense; at all events the oxide
of potassium formed is not to be distinguished from oxide of potassium
prepared with common oxygen; but this in nowise contradicts the fact
that ozone is an extraordinarily active and energetic variety of oxygen,
inasmuch as common oxygen will not effect this decomposition.
169. Ozone is an irritating, poisonous gas; air which is highly
charged with it is irrespirable, and produces effects on the human
subject similar to those produced by chlorine. Its odor, which
has been compared to that of weak chlorine, is so powerful that
it can be recognized in air containing only one millionth part of
the gas. Its oxidizing power is intense. When moisture is present it oxidizes all the metals excepting gold, platinum, and the
platinum metals; even silver is oxidized by it at the ordinary
temperature, and becomes covered with a brown coating of peroxide of silver.  It destroys many hydrogen compounds, such as
those of sulphur, phosphorus, and iodine, the hydrogen being
oxidized as well as the element with which the hydrogen is
associated; iodohydric acid, for instance, is converted into water
and iodic acid. In the same way, free iodine is oxidized by
ozone, and if test-paper which has become blue by exposure to
ozone, as in Exp. 76, be left long in ozonized air, it will become
white from oxidation of the iodine. Ozone will even oxidize
nitrogen, at the ordinary temperature, when in contact with
water and such alkaline oxides as caustic soda, caustic potash,




144                   TESTING FOR OZONE.
or caustic lime; thus, if lime-water (a solution of caustic lime in
water) be left exposed to ozonized air, a certain quantity of
nitrate of calcium will be formed. Ammonia is oxidized by it also,
and it converts nitrous and sulphurous into nitric and sulphuric
acids.
Many salts of the metals are oxidized by it-for example, the
sulphates of iron and of manganese. A valuable test for the
presence of ozone is furnished by its behavior towards sulphate
of manganese.
Exp. 77.-Dissolve a gramme or two of sulphate of manganese in
water; soak in this solution strips of thin white blotting-paper; dry
the paper, and preserve it in a bottle. If a slip of this paper be moistened, and then hung in ozonized air (Exp. 75), it will quickly become
brown from the formation upon it of black oxide of manganese.
In like manner, most organic substances are quickly oxidized
by ozone; when substances such as sawdust, garden-mould,
powdered charcoal, milk, or flesh are thrown into a bottle of
ozonized air, the odor of ozone instantly disappears; corks and
caoutchouc tubes are attacked by it, and must not be used in
experimenting with the gas. It destroys the color of indigo, and
bleaches litmus without first reddening it.  Some organic bodies,
on the other hand, become colored when exposed to its action;
thus, the cut surface of an apple becomes brown, and fresh surfaces of certain mushrooms become blue.  Gum guaiacum also
becomes blue. Papers soaked in a dilute alcoholic solution of
gum guaiacum, indeed, are often employed as a test for ozone.
Exp. 78.-Dissolve one part of gum guaiacum in thirty parts of
ninety per cent. alcohol; add a few drops of this solution to 2 c. c. of
ordinary eighty per cent. alcohol; dip in this dilute solution strips of
thin white blotting-paper, and dry them in the dark. By exposure
to ozonized air this test-paper acquires a bright blue color.
170. By virtue of its strong oxidizing-power, ozone is of great
importance as a disinfecting agent. It destroys instantly a multitude of offensive gases, such as arise from decaying animal and
vegetable matter, and has been frequently recommended of late
as a substance well fitted for the purification of sick-rooms and
hospital-wards. Where ozone is employed for purposes of disinfection, it must be borne in mind that the action of the gas




OZONE A DISINFECTANT.                145
depends solely upon oxidation. A. given quantity of ozone can
destroy only a certain definite amount of the offensive organic
matter; wherever these emanations are incessantly generated,
ozone must be as constantly produced in order to destroy them.
This disinfecting-power of ozone is interesting in connexion with
the observed facts, that ozone is abundant in the air of pine
forests, where turpentine abounds, and that pine forests are, as a
general rule, remarkably free from  malaria.  The well-known
disinfecting-power of tar is supposed in like lanller to be partly
due to the formation of ozone during the oxidation of some of its
ingredients.
Coal-tar, mixed with plaster-of-Paris, coal-ashes, or dry earth, in
quantity sufficient to destroy its stickiness, has been found to be a very
efficient disinfectant. The dry powder obtained as above, is simply
scattered freely about the offensive locality. The coal-tar, of course,
evolves a slight odor, peculiar to itself, which tends to mask or conceal
other odors, and also acts as an antiseptic, or arrester of putrefaction;
but its chief merit does not appear to depend upon either of these properties; it seems really to destroy the gases which are evolved from
putrescent matter, and probably does so by generating ozone.
171. It is supposed that a minute proportion of ozone exists
in normal atmospheric air: at all events, there is usually present
in air a substance which exhibits the various reactions of ozone,
and behaves as ozone would if it were there. This atmospheric
ozone, which is supposed to be formed in the processes of oxidation which are always going on in nature, varies in quantity with
the locality, the season of the year, the hour of the day, and many
other circumstances.
172. Ozone is seldom found in the air of thickly inhabited localitics; it often happens that it cannot be detected in the air of cities
at the very time when it is abundant in the neighboring country.
It is often found to be abundant on the windward side of a city-,
and altogether absent from the air upon the leeward side, the inference being that it is destroyed by the exhalations which arise
from a dense population. Ozone appears to be more abundant
in the air in winter than in sumner, in cloudy than in clear
weather, and by night than by day; it has been observed to be
specially abundant at times when dew was falling heavily.  As
might be expected, comparatively large quantities of it are found
L




146               OZONE 1N THE ATMOSPHERE.
during thunder-storms, and its odor has been recognized in the
neighborhood of objects struck by lightning. Ozone is abundant
during snow-storms; and it is probable that upon its presence depends the well-known bleaching-power of newly fallen snow.
In searching for ozone in the air, test-paper containing iodide of potassium and starch, such as was prepared in Exp. 71, is usually employed. Dry slips of the prepared paper are exposed, during from six
to twenty-four hours, to a free current of air, in a place well sheltered
from light and rain. By exposure the dry paper becomes brown, and
when wetted acquires shades of color varying from pinkish-white and
iron-gray to blue. The shade of color obtained in this way is then
compared with a standard chromatic scale, which includes all the
shades possible under the circumstances; and the proportion of ozone
present in the air is thus roughly estimated.
Although observations of this kind are far from possessing that
degree of accuracy and certainty which is desirable, they have nevertheless been considered trustworthy by numerous observers, and have
given rise to much speculation concerning the functions of atmospheric
ozone, more particularly with regard to its probable influence upon
health and disease. If there be ozone in the atmosphere, it will, on the
one hand, oxidize and destroy many volatile organic substances which
are supposed to be prejudicial to health. Hence many physicians are
of opinion that the atmospheric ozone plays an important part in controlling or preventing epidemic diseases through its power of removing infectious matter from the air; and it has been noticed that with
the advent of an ozone-bearingwind such diseases have abated or ceased.
But, on the other hand, ozone is a highly irritating gas, and in the
opinion of some physicians occasions many diseases of the respiratory
organs. Numerous statements are upon record to the effect that epidemics of catarrh, colds, sore throat, and influenza have been coincident with the beginning of a spell of ozoniferous wind.
173. Ozoneiis usually considered to be completely insoluble in
water; but it has been recently ascertained that water can take
up a small quantity of it, and so acquire some of the properties of
ozone.  When ozonized air is passed through a solution of caustic
soda or caustic potash, a certain amount of ozone is absorbed at
first, perhaps by.combination with some oxidizable impurity of
the solution, but after a little time the ozone will pass through
without apparent alteration. Acids do not absorb ozone. It is
readily absorbed,ihowever, by aqueous solutions of iodide of potas



BATOZONE.                      147
sium and of pyrogallic acid, with the constituents of which it enters
into combinations not to be distinguished from those made with
oxygen.
174. At moderately high temperatures ozone loses its peculiarities and changes into ordinary oxygen; if ozonized air, such as
was obtained in Exp. 75, is made to pass through a narrow glass
tube heated to 2500, its peculiar odor, and its power of decomposing iodide of potassium will entirely disappear. The same
change occurs gradually if the tube is heated only to 100~,-or
instantly if steam be thrown into the ozonized air, so that the
whole of it can be heated at once to 1000; hence it may be stated,
in general terms, that ozone is converted into ordinary oxygen at
temperatures greater than 100~.
175. Ozone is supposed to exist as such in several of the oxides.
Black oxide of manganese, for example, is thought to contain it
as a constituent; and a method of obtaining it from permanganic acid has been already given, ~ 167. The oxygen compounds which are supposed to contain ozone are called ozonides.
The formuloe of the following compounds, recognized as ozonides,
are here given for the sake of reference:PbO2  CrO3 MnO2  Co20O  N205
Ag20, MnO, Mn207 Ni2O, Bi,26
176. Antozone (the opponent or opposite of ozone) appears
to be produced simultaneously with ozone whenever the latter is
formed, whether by electrical action or during processes of oxidation. It may even be that, as some chemists believe, ordinary
oxygen is in a certain sense a compound substance, and that
when in contact with phosphorus, and in the other circumstances
under which ozone is produced, the neutral oxygen is split or decomposed into two opposite and dissimilar modifications-we had
almost said elements-one of which is ozone, the other antozone.
It is thought that while the greater part of the ozone thus engendered enters into combination with the phosphorns, or other
substance, undergoing oxidation, a certain portion of it, together
with some of the antozone, becomes mixed with the surrounding
air, and so escapes combining with the body which is being oxidized.
Only a comparatively short time has elapsed since antozone
L 2




148                PREPARhTION OF ANTOZONE.
has been recognized as a distinct substance; hence its properties
have been less thoroughly studied than those of ozone. Many of
its characteristics and properties are still involved in great obscuritv, very various and even conflicting statements having been
published concerning them.
177. Of the methods devised for preparing antozone, the following deserve notice:
By passing dry electrized air (~ 165) through a concentrated aqueous solution of iodide of potassium, or of pyrogallic acid, all the ozone
contained in the air will be at once absorbed, and the antozone left
behind, free from any admixture of ozone.
During the slow oxidation of oil of turpentine and other volatile or
essential oils (~ 166), a considerable quantity of antozone is produced,
as well as of ozone. While most of the ozone at once combines with
the constituents of the oil, to form resins and other products of oxidation, the antozone, which does not oxidize the oil, is dissolved by it.
In what state the antozone exists within the oil is still uncertain; but
it is, in any event, very loosely held, and is readily given up to other
substances.
In the same way that ozone may be prepared, by chemical decomposition, from permanganate of potassium, a compound supposed to contain ozone (~ 167), antozone may be obtained by decomposing certain
compounds which are believed to contain this variety of oxygen-such,
for example, as peroxide of barium, BaG2.  A little concentrated sulphuric acid is poured into a small bottle, and into this acid are thrown
a number of small fragments of peroxide of barium (free from any admixture of nitrate of barium); so soon as an evolution of gas ensues,
the air of the bottle will be found charged with antozone. This reaction is sometimes capricious. Usually it occurs at the ordinary
temperature of the air; but it is often necessary to place the bottle in
a water-bath heated to 500 or 60~, in order to start the evolution ol
gas; and, on the other hand, the violence of the reaction must some.
times be allayed by immersing the bottle in cold water.
In the preparation of ozone by means of phosphorus in moist air
(Exp. 75), or by the electrolysis of water (~ 35), the antozone which
is formed at the same time with the ozone, unites with the water
present, and must there be sought. (See ~ 181.)
Antozone has been found in nature in a dark-blue variety of fluorspar from Wilsendorf, in Bavaria. Upon being rubbed, this mineral
emits a peculiar odor, which was formerly thought to be that of chlorine or of hypochlorous acid. More recent investigations have showr




PROPERTIES OF ANTOZONE.               149
that the odor is that of antozone, and that by grinding the mineral
with water the antozone can be transferred to the water.
178. Antozone is a gas, the odor of which somewhat resembles
that of ozone; there is, however, a decided difference between
the two odors, that of antozone being disgusting, while that of
ozone is merely pungent and irritating. Antozone changes at
once to ordinary oxygen on being heated. Even at the ordinary
temperature it reverts to common oxygen very readily-much
more readily than ozone. Most of the antozone usually disappears from dry electrized air in the course of an hour, or an hour
and a half; and if the air be moist, the change is still more rapid.
Ozone, on the c6ntrary, is comparatively permanent, under the
same conditions; and although when a mixture of ozone and antozone is left in contact with water in a glass-stoppered bottle, some
ozone is destroyed during the reversion of the antozone, the
larger portion of it will remain almost, if not quite, unaltered
for months. Antozone, whether moist or dry, also reverts to the
condition of ordinary oxygen on being brought in contact with
black oxide of manganese, peroxide of lead, or finely divided
platinum.
179. A very remarkable characteristic of antozone is its power
of forming fogs and clouds with water. It may even be found,
after the matter has been more thoroughly studied, that all the
fogs and clouds which occur in nature are dependent for their
existence upon the presence of antozone.
If air, charged with antozone, be made to bubble through water, it
will emerge from the water in the form of a thick white mist, similar
to that formed by the cooling of steam. The same thing occurs when
electrized air, or electrized oxygen, issues into a moist atmosphere,
though the effect is less marked when ozone is present than when it
has been removed by means of iodide of potassium. The mist produced by slowly passing antozonized air through water is heavy; it
remains hanging over the surface of the liquid, and may be readily
poured from one vessel to another. By conducting it through a tube
to the bottom of a dry, tall bottle, it displaces the air, all the while
preserving a sharply defined boundary; by gentle agitation it is easily
broken up into cloud-like masses.
When a large, dry bottle is nearly filled with this antozone mist,
then closed and left to itself, the mist gradually becomes thinner and




150                 TU1TE ANTOZONE CLOUD.
less opaque, and in the course of half or three-quarters of an hour
vanishes altogether. As the cloud thus disappears, water is deposited
upon the sides of the bottle, at first as a mere dew, but afterwards
accumulating in droplets, which finally flow together to the bottom of
thd vessel. When the air in the bottle has become clear, no antozone
can be detected in it.
It thus appears that antozone has the property of taking up water
in such a manner that the water assumes the peculiar physical conditions of a cloud or mist. While the antozone lasts the cloud is permanent; but the antozone is soon transformed into ordinary oxygen,
and as fast as this change occurs the water of the cloud is deposited
in droplets.
By passing the antozone mist through tubes filled with desiccating
substances, such as chloride of calcium (Appendix, ~ 15), the water
may be removed, and transparent antozonized air obtained, capable of
again producing a mist on being brought in contact with water. Many
strong saline solutions likewise deprive antozone of water; hence the
non-appearance of the cloud when electrized air is passed through a
strong solution of iodide of potassium; the cloud does appear, however, when the solution is sufficiently dilute.
It has been proved by experiment that electrized air can support or carry nearly twice as much moisture as ordinary air or
oxygen at the same temperature, and that this air is much more
difficult to dry than the gases with which chemists usually have
to deal. This explains how it happened that, before the discovery of the cloud-forming property of antozone, so many observers had been led to consider ozone an oxide of hydrogen. One
experimenter would pass recently electrized air through an ordinary drying-tube, such as long experience had shown to be
capable of drying common air perfectly, and would then heat the
gas; by this treatment both the ozone and the antozone would
be changed to ordinary oxygen, and the water which had been
carried through the drying-tube by the antozone would be made
visible. The remarkable capacity of antozone for moisture being
unknown, the water thus obtained was naturally enough supposed
to have been derived from some compound of hydrogen and oxygen other than water, and capable of passing unabsorbed through
the drying-tube. Other chemists, performing, as they supposed,
the same experiment, but in reality operating upon air less recently
electrized, and so containing no antozone, were, of course, unable




ANTOZONE FORMED DURING COMBUSTION.           151
to obtain any water at the point where it had been observed by
their predecessors; hence arose a series of controversies which
have only recently been composed.
180. As has been already mentioned, antozone, like ozone, is
formed in all processes of oxidation and combustion. During
combustion most of the ozone produced enters into combination
with the substance burned, while the antozone is left free, or
enters into combination with water to form peroxide of hydrogen.
When the combustion is slow or smouldering, antozone appears
in large quantities, and in presence of moisture forms the characteristic mist or cloud. Tobacco-smoke, the gray smoke of
chimneys and of gunpowder, and all such smokes are antozone
clouds,-facts which support the idea that all clouds, fogs, and
mists are caused by the presence of antozone in the atmosphere.
The oxidation of phosphorus affords a ready method of exhibiting
the antozone cloud. During the oxidation of phosphorus in moist air,
white fumes are formed, which were long a great puzzle to chemists.
Whether the phosphorus be allowed to oxidize slowly, as in Exp. 75,
or burned rapidly, as in Exp. 13, there is always produced a white mist
of very considerable permanence, which remains long after the oxides
of phosphorus, which are also formed, have been takepup and removed
by the water. This mist is the antozone cloud; it is nothing but water
held suspended by antozone.
In the rapid combustion of phosphorus, little or no ozone is left free;
all of it seems to unite directly with the phosphorus; but much more
antozone is produced when the combustion is rapid than when it is
slow. The formation of antozone in this connexion explains the fact
already alluded to (Exp. 13), that phosphorus burning with flame, in
a confined volume of air, does not wholly exhaust the latter of oxygen.
The phosphorus cannot combine with antozone, but only with ozone;
hence, when no oxygen other than that in the form of antozone remains, the combustion must cease.
During the burning of a jet of hydrogen under a bell-glass through
which a stream of air is drawn, antozone is formed, as is proved by
passing the issuing stream through water; the antozone cloud is produced without difficulty, and peroxide of hydrogen appears as a product. The formation of the antozone mist, and of peroxide of hydrogen, may be observed with any other flame if care be taken that the
air which streams over the flame be not too strongly heated. A high
temperature destroys the antozone as fast as it is formed.




152                ANTOZONE OXIDIZES WATER.
181. Besides its power of forming clouds or mists with water,
which is interesting rather as a physical than as a chemical fact,
antozone, particularly when newly formed, also unites with water
chemically, the substance called peroxide of hydrogen (see ~ 61),
whose composition is expressed by the formula 11202, being the
result of the combination.
A simple method of exhibiting the formation of peroxide of hydro(gen by the action of antozone upon water, is to place a short, narrow
tube, containing concentrated sulphuric acid, within a bottle 2 or 3
c.m. in width, furnished with a ground-glass stopper, and filled with
water nearly to the top of the tube. Small portions of peroxide of
barium are now added, at intervals, to the sulphuric acid in the tube,
elevation of temperature being avoided as far as possible; the stopper
should be replaced in the bottle after each addition of the peroxide,'Most of the oxygen evolved in this process appears, however, to be in
the ordinary inactive state, and the solution of peroxide of hydrogen
obtained is consequently extremely dilute. A better method of procedure is to pass a current of carbonic acid gas into a mixture of water
and peroxide of barium,
BaO2 +  -H20 +  C02 = BaO,CO2 + — H202
In this way a highly concentrated solution of the peroxide can be
obtained.
Another easy method of preparing peroxide of hydrogen is by the
oxidation of amalgams of lead or zinc. In this case also, as in the
preceding, the peroxide of hydrogen is probably formed by tlhe union
of antozone with water.
One hundred grammes of lead-amalgam, containing so much mercury that it shall be fluid at the ordinary temperature, is shaken in a
bottle of the capacity of a litre, together with 200 c. c. of water, acidulated with 2 grms. of sulphuric acid; the water soon becomes milky
from separation of sulphate of lead, and in the course of ten or towelve
minutes contains enough peroxide of hydrogen to exhibit the characteristic reactions of this substance.
So, too, if pulverulent zinc-amnalgam be loosely thrown into a glass
funnel, with narrow throat, and a thin stream of water be allowed to
flow through it in such manner that the metal may be at the same
time acted upon by both air and water, the water will become charged
with peroxide of hydrogen. By repeatedly pouring back the dilute
solution of the peroxide upon the amalgam, it can be very considerably
strengthened. In order to prepare the zinc-amalgam, equal weights of
zinc-filings and of mercury are placed in a beaker glass, covered with




DIFFERENCES BETWEEN OZONE AND ANTOZONE.        153
water acidulated with sulphuric or chlorhydric acid, and thoroughly
mixed by stirring with a glass rod; the acid is then poured away, and
the last portions of it removed from the amalgam by washing with
water.
This power of antozone to oxidize water distinguishes it completely from ozone, which has little or no action upon water.
182. Peroxide of hydrogen, like peroxide of barium, is supposed
to contain one atom of oxygen in the form of antozone; the peroxides of potassium, sodium, and strontium also are placed in the
same category.  They are all called antozonides.
183. Antozone can be distinguished from ozone by the following tests:Strips of paper, charged with a solution of sulphate of manganese
(Exp. 77), do not become brown when exposed to the action of antozone; on the contrary, manganese papers which have been browned
by ozone are bleached by antozone. Guaiacum paper (Exp. 78) does
not become blue in antozonized air. The yellow compound called
terrocyanide of potassium, which is converted, into red ferricyanide of
potassium by the action of ozone, is not changed by antozone. In the
absence of acids, antozone has no action upon iodide of potassium.
The chemical behavior of antozone may be conveniently studied by
resorting to its compound with water, the antozonide peroxide of hydrogen. If peroxide of hydrogen be brought in contact with an ozonide
like peroxide of lead, for example, both of the peroxides will be reduced, and there will result water, protoxide of lead, and free ordinary
oxygen. Whenever an antozonide is mixed with an ozonide, a similar
reaction occurs; the two active varieties of oxygen disappear, and
common oxygen is evolved; hence it has been assumed that ordinary
inactive oxygen is a sort of compound, resulting from the union or
neutralization of ozone with antozone. Several important tests for
antozone are dependent upon this fact of the decomposition of antozonides by ozonides.
If a liquid suspected to contain peroxide of hydrogen be shaken in a
test-tube with a small quantity of ether, the ether will dissolve the
peroxide, and will finally collect upon the surface of the liquid; on
adding to it a small drop of a solution of the ozonide chromic acid, or,
what comes to the same thing, a drop of a solution of bichromate of
potassium acidulated with sulphuric acid, the ethereal solution will
become blue.
If a liquid containing peroxide of hydrogen be added to a dilute red
solution of permanganate of potassium, this solution will be decolorized,




154                       SULPHUR.
while common oxygen will be evolved; and in the same way the brown
peroxide of lead and the red-colored salts of peroxide of iron are
bleached by it.
Another exceedingly delicate and characteristic test for antozone, or
rather for peroxide of hydrogen, the rationale of which has not yet
been well made out, is the following: —If to a solution containing peroxide of hydrogen there are added a few drops of dilute starch-paste
charged with iodide of potassium, and subsequently a very small
quantity of a solution of copperas (protosulphate of iron), iodine will
be set free, and the starch will become blue. The solution to be tested
must be as nearly neutral as possible. The addition of an acid, instead
of the copperas solution, will also bring about the same reaction,
though less readily.
184. We have thus set forth whatever is best known concerning ozone and antozone, in spite of the details into which so full
an exposition has necessarily descended, partly because the subject will evidently be one of primary importance, both theoretical
and practical, in the near future, and partly from a desire to
show the student how vague and uncertain the prospect is when
once the narrow limits of established knowledge are past and
the inquirer ventures out into the obscurity which perpetually
separates the knowledge of to-day from that which shall be
knowledge to-morrow, but also because of the impossibility,
with so obscure a subject, of making such a just discrimination
between salient and unimportant points as with a well-studied
subject is both easy and desirable.
CHAPTER XIIL
SUL PHU  R.
185. Sulphur occurs somewhat abundantly in nature, both in
the free state and in combination with other elements. fMany
ores of metals, for example, are sulphur compounds. It is a
component of several abundant salts, such as the sulphates of
calcium, barium, and sodium, and occurs in small proportion in




THE SULPHUR OF COMMERCE.                155
many animal and vegetable substances. Free sulphur is found
chiefly in volcanic districts. Generally it occurs mixed with
earthy matters; but it often forms distinct veins, and is sometimes
found in the shape of well-defined crystals of considerable size.
At the present time about nine-tenths of the sulphur of commerce
comes from Sicily.
186. Native sulphur is usually subjected to a rough purification at the place of its occurrence. This purification is sometimes effected by distilling the volcanic earth in retorts or jars
of earthenware; the sulphur being volatile, distils over, and is
collected in receivers, from which it is drawn off, from time to
time, in the liquid state; or if the earth be very rich in sulphur,
it is simply heated in large kettles and the melted sulphur dipped
off from above, while the earthy impurities settle to the bottom
of the kettle. The product thus obtained is known as crude sulphur; it comes to us in irregular lumps of a dirty light-yellow
color, and is largely employed for manufacturing-purposes.
This crude sulphur is contaminated with more or less earthy
matter. In order to purify it, it is distilled from iron retorts into
large chambers constructed of masonry, in which it is deposited
either in the form of a light powder, known as flowers of sulphur, or in a liquid state, according to circumstances. At the
beginning of the operation, while the chamber is cold, the sulphur vapor condenses as an exceedingly fine, soft, powder (flowers
of sulphur) upon the walls of the chamber. But heat is given off
as the sulphur vapor condenses, and after a while the walls of
the chamber become so hot that sulphur will melt upon them.
After this, the incoming sulphur vapor of course condenses only
to the liquid state, and a layer of liquid sulphur collects upon
the floor of the chamber. This liquid sulphur is drawn off into
wooden moulds, and thus cast into the sticks familiarly known as
roll-brimstone. It is evident that, by a little management, the
sulphur-refiner can obtain, at will, either flowers of sulphur or
roll-brimstone, or first the one and then the other.
187. At the ordinary temperature of the air, sulphur is a
brittle solid, of a peculiar light-yellow color. It has neither
taste nor smell, excepting that when rubbed it exhales a faint
and peculiar odor. Most of the odors which in everyday life are




156              CRYSTALLIZATION OF SULPHUR.
referred to sulphur are really the odors of various compounds of
sulphur, and are not evolved by the element itself. It is a bad
conductor of heat and electricity.  On being rubbed it becomes
highly (negatively) electric, and is still employed as a sour-e of
electricity in some cases. The symbol of sulphur is S; its atomic
weight is 32, being precisely twice as great as the atomic weight
of oxygen.
188. Sulphur melts easily at about 1120, a temperature not
very far above that at which water boils. A fragment of it may
even be melted by heating it on writing-paper over the flame of
a candle. It volatilizes freely at temperatures lower than its
melting-point, and boils at 440~. Indeed, as is the case with
water, it is a substance which can be brought into either of the
three states of matter without any difficulty; we can have it as a
solid, a liquid, or a gas as we please. It can readily be obtained
also in the form of crystals.
Exp. 79.-In a small beaker glass, or porcelain capsule, heat slowly
50 to 60 grms. of sulphur until it has entirely melted. Remove the
vessel from the lamp, and allow it to cool slowly until about a quarter
part of the sulphur has solidified; then pour off, into a basin of water,
that portion of the sulphur which is still liquid, breaking through, for
this purpose, the crust at the top of the liquid, if any such have formed.
The interior of the vessel will be found to be lined with transparent,
prismatic crystals.
Exp. 80.-In a test-tube, melt enough sulphur to fill one-quarter of
the tube; place the tube in such a position that its contents may cool
slowly and quietly, and then watch the formation of crystals as they
shoot out from the comparatively cold walls of the tube towards the
centre of the liquid.
Exp. 79 represents one general method of obtaining crystals.
Crystals of many of the metals, lead and bismuth for example,
can be obtained by operating in this way; it is only necessary
to melt the metal in a crucible of some refractory material,
placed in a furnace. The melted metal having then been allowed
to cool until a tolerably firm crust has formed upon its surface,
this crust is pierced with an iron rod, and the crucible quickly
inverted, so that the portion of the metal which still remains
fluid in the interior shall flow out.  Upon afterwards breaking




CRYSTALLINE STRUCTURE.                157
the crucible, crystals will be found lining the cavity of the metallic
cup which has been formed within it.
189. Exp. 80, besides illustrating the manner in which crystals form, teaches us something of the physical structure of solid
bodies. The solid mass of sulphur which is left in the test-tube
when it has become cold, is evidently nothing more than a compact bundle of interlaced crystals. If the mass be removed from
the tube, and then broken across, it will present a glistening
appearance, owing to the reflection of light from the surfaces of
the minute crystals of which it is composed. It is said to have
a crystalline structure. This crystalline structure is apt to render
a body brittle; substances which possess it are liable to break
"with the grain," or to split in certain directions determined by
the shape of the crystals, and called lines of cleavage; a stick of
roll-brimstone, for example, may be readily broken or cut across,
but not so easily in the direction of its length. The same remark
applies to many samples of metal. In all cases where tenacity is
required, it is important to counteract, or to prevent as much as
possible, the tendency towards crystallization. Thus, in manufacturing wrought iron, it is the constant endeavor of the workman to render the metal stringy or fibrous, and not crystalline,
and he seeks to accomplish this by appropriate processes of
kneading, squeezing, and rolling.
190. Another easy way to crystallize sulphur is by the method
of solution and evaporation, such as was employed in the preparation of nitrate of ammonium (Exp. 33). Sulphur is not
soluble in water, but it dissolves readily in a liquid compound of
sulphur and carbon, known as bisulphide of carbon, which being
readily volatile, quickly escapes, on exposure to the air, and so
deposits the sulphur.
Exp. 81.-Place in a test-tube a small teaspoonful of flowers of
sulphur, pour upon the sulphur 10 or 12 c. c. of bisulphide of carbon,
close the tube with a cork, and allow the mixture to stand during half
an hour, shaking it occasionally. Decant the clear liquid from the
sulphur which still remains undissolved, and pour it into a small porcelain capsule, which place out of doors, or in a draught of air, until the
highly offensive bisulphide of carbon has all evaporated. Crystals of
sulphur will then be found at the bottom of the dish.
This experiment might be modified by preparing, in the first place, a




158          TIlE SIX SYSTEMS OF CRYSTALLIZATION.
saturated solution of sulphur in boiling bisulphide of carbon, and then
allowing the clear solution to cool slowly. Crystals of sulphur would
finally be found beneath the cold liquid. The method by evaporation,
as above described, is to be preferred.
It will be noticed that the crystals of Exp. 81 are not shaped
like those obtained by the method of fusion in Exp. 79. The
two sets of crystals belong in fact to entirely different systems of
crystallization.
191. The researches of crystallographers have proved that the
crystals of natural minerals and artificial chemical substances
may all be included in six general classes of form, called systems
of crystallization. In every crystal, certain directions may be
recognized, with reference to which the bounding planes of the
crystal exhibit a more or less symmetrical arrangement. These
directions, represented by straight lines drawn through the centre
of the crystal, are called axes. The thousands of crystal-forms
which occur in nature, or are produced by art, have been divided
into six systems, or groups, by observation of the number, relative length and mutual inclination of the axes around which they
are symmetrically formed. These six systems are defined as
follows:I. Monometric (single-measure) or Regular Systemn.-The axes are
three in number, equal in length, and intersect each other at right
angles. The cube, regular octahedron, and rhombic dodecahedron,
forms of perfect symmetry, belong to this system.
II. Dimetric (tzwo-measure) System.-The axes are three in number,
and intersect each other at right angles; but one, called the vertical,
is either longer or shorter than the two lateral, which are equal. The
right square prism and square octahedron are of this system.
III. Trimetric (three-measure) System. —The axes are three in number, unequal in length and intersect each other at right angles. The
system includes the right rectangular prism, the right rhombic prism,
and the rhombic octahedron.
IV. i1lonoclinic (single-inclination) System.-The axes are three in
number, and unequal in length; and one, called the vertical, is at right
angles with one of the other two axes, which are called lateral, but
obliquely inclined to the other; the two lateral axes intersect each
other at right angles. The right rhomboidal and oblique rhombic
prisms belong to this system.
V. Triclinic (three-inclination) System.-The axes are three in num.




SULPHUR IS DIMORPHOUS.                 159
ber, unequal in length, and all their intersections are oblique, The
oblique rhomboidal prism is of this system.
VI. Hexagonal System.-The axes are four in number; three, called
lateral, lie in one plane, are equal in length, and intersect each other
at angles of 60~; the fourth axis, called vertical, is either longer or
shorter than the other three, and crosses them at right angles. This
system includes the hexagonal prism and the rhombohedron.
Under these systems of crystallization, the variety of possible forms
and dimensions is unlimited. Thus, in systems in which the axes are
unequal, the inequality may be great or small, through all degrees of
discrepancy; in oblique systems the inclination of the axes may vary
indefinitely; rhombohedrons may occur of every angle. Thus the actual forms of crystallography become exceedingly numerous, although
they all belong to a few simple types.
If the student draws in perspective, upon paper, the axes of the
several systems above described, or, better, constructs the different sets
of axes out of bits of wood or wire, he will appreciate the fact that
forms belonging to different systems are ordinarily so unlike in general
appearance as to be readily distinguishable even by those who have no
exact knowledge of the mathematical science of crystallography.
192. As a general rule, a substance crystallizes in forms belonging to only one system, and the crystalline form of a substance is something so constant and characteristic as to be one of
the chemist's most valued means of recognition and definition.
But this general rule is not without exceptions. Sulphur, as has
just been proved, may be made to crystallize in forms belonging
to two distinct systems of crystallization; and there are other
substances, not a few, which when crystallized under different
Fix 38. conditions, assume forms of two distinct systems.
Substance.s which are thus capable of assuming crystalline forms belonging to two different systems are
i    said to be dimopkhous (two-formed).  Two such different forms of the same substance often  Fig. 39.
have quite dissimilar physical properties;
they are apt to differ from  each other in
hardness, specific gravity, color, optical properties, and in their relation to heat; the
chemical properties, also, of two such different forms are seldom entirely the same.
The crystals of sulphur obtained by fusion (Exp. 79) are




160     CHANGE OF PRISMATIC INTO OCTAHEDRAL SULPHUR.
elongated oblique rhombic prisms (Fig. 38), and belong to the
fourth (or monoclinic) system. The crystals of sulphur which
are derived from its solution in bisulphide of carbon (Exp. 81)
are rhombic octahedrons (Fig. 39), belonging to the trimnetric
system. The specific gravity of tlie octahedral crystals is greater
than that of the prismatic in the ratio of 2'07: 1'91. The specific heat of the octahedral crystals is 0'163, and that of the
prismatic somewhat greater. The melting-point of the prismatic
crystals is about 120~.
The prismatic crystals of sulphur (Exp. 79) cannot be kept
for any great length of time.  They soon lose their transparency
and characteristic amber color, becoming opaque and light yellow,
like ordinary brimstone. If they be examined under the microscope it will be seen. that the prisms are now composed of a
multitude of little octahedral crystals.  The change of color and
texture is due to a rearrangement of the particles of the original
crystals, though the aggregation of octahedrons which have been
formed within the prismatic crystal still retains the shape of the
prism. If the prismatic crystals be left at rest, this change of
form usually begins in the course of a few hours; but it may be
greatly accelerated by scratching the crystals, or shaking them
together. Under ordinary circumr.stances the passage of the sulphur from  the one molecular state to the other goes on very
slowly, several years being often required for its completion; but
the change can be accomplished immediately by moistening the
prismatic crystals with bisulphide of carbon. A considerable
amount of heat is developed as the prismatic sulphur changes
into octahedral; this can readily be appreciated when the conversion is effected by means of bisulphide of carbon.
In the same way that prismatic sulphur slowly changes into
the octahedral variety at the ordinary temperature, octahedral
sulphur is gradually converted into prismatic sulphur when kept
for a long time at a temperature near its melting-point. The
change in specific gravity enables us to follow the progress of this
conversion.
Sulphur which has been melted and allowed to solidify gradually, is always in the prismatic condition immediately after the
solidification.  Roll-brimstone, for example, when fresh from the




METuODS OF OBTAINING CRYSTALS.            161
moulds, is translucent, and of a dark amber or brownish-yellow
color, like the prismatic crystals of Exp. 79; but in a short time,
often in the course of a few hours, the sticks become light-yellow
and opaque, as we find them in commerce, and are then composed, at least externally, of a mass of octahedral crystals.
193. There is still a third way of obtaining crystals of sulphur,
namely, by sublimation.  At slightly elevated temperatures, sulphur is volatile; and if the circumstances be such that the vapor
shall condense very slowly, crystals will form.  The natural
crystals of sulphur found in volcanic countries, which are often
very large and of great beauty, have been formed in this way.
These native crystals are octahedral, like those obtained by means
of bisulphide of carbon (see Exp. 81).
194. As appears from the foregoing, there are three distinct
methods of obtaining crystals:-I. By fusion; that is to say, by
the slow cooling of molten matter.  II. By solution, followed
either by removal of the solvent by evaporation or chemical means,
or by reduction of its temperature. III. By sublimation.
A familiar instance of the first method is seen in the case of
ice, as when a part of the water in. any hollow vessel freezes
slowly upon the sides of the vessel; of the second, in the manufacture of common salt; and of the third, in the formation of
frost upon a window-pane.
There is still a fourth general method of obtaining crystals,
which consists in very slowly decomposing some chemical compound of the substance to be crystallized, either by the addition
of some other chemical agent, or by means of the galvanic current. Crystals of sulphur may be formed in this way, and are
in fact sometimes found in the pipes used to convey illuminating
gas through the streets of cities, under such circumstances that
it is evident that they have resulted from the decomposition of
some one of the sulphur compounds with which coal-gas is always
contaminated.
It must not be inferred, from the above enumeration of the
ordinary methods of obtaining crystals, that either fusion, solution, or sublimation is a necessary condition of the formation
of crystals. Both in nature and in art examples occur of the
crystalline arrangement of particles within solid masses, under
-H




162                     SOFT SULPHUR.
circumstances which preclude the idea that either fusion, solution, or sublimation, in the ordinary sense of these terms, should
have occurred.
195. Sulphur behaves in a very remarkable manner on being
heated. When melted at the lowest possible temperature, 1100
to 115~, it forms a limpid liquid of a light-yellow color; but if
this liquid be heated more strongly, it begins to become viscid
and dark-colored at about 150~, and at 1700 to 200~ it is almost
black, and at the same time so thick and tenacious that it cannot
be poured from the vessel which holds it, even if the vessel be
inverted. At 3300 to 340~ it regains its fluidity in part, though
the liquid is still dark-colored, and finally, at about 440~, it begins to boil, and is converted into an amber-colored vapor. The
specific gravity of sulphur vapor, referred to hydrogen, is 32.
196. If melted sulphur, in the viscid state, or, better, that
which has regained its mobility, be suddenly cooled, a semisolid
modification of sulphur, remarkably different from the ordinary
form, will be obtained.
E2p. 82.-Place in a test-tube, of about 30 c. c. capacity, 15 to 20
grms. of coarsely powdered sulphur; melt the sulphur slowly over the
gas-lamp, and continue to heat it until it begins to boil, noting, meanwhile, the changes which the sulphur undergoes-as described in ~ 195.
Finally pour the hot sulphur, in a fine stream, into a large dish full of
cold water. There will be obtained a soft, elastic, reddish-brown mass,
which can be kneaded and moulded like wax, and drawn out into
threads like caoutchouc.
This soft sulphur cannot be preserved for any great length of
time. When left to itself, at the ordinary temperature of the
air, it slowly hardens and changes into ordinary brittle yellow
sulphur.  This change is accelerated by kneading, and is instantaneous at the temperature of 1000. In any event, a certain
amount of heat is evolved as the soft sulphur changes into ordinary sulphur. The specific gravity of sift sulphur is somewhat
lower than that of the prismatic crystals.
From the foregoing facts it appears that sulphur, like oxygen,
is capable of assuming different allotropic states. (See ~ 162.)
197. In its behavior towards solvents, sulphur presents some
curious anomalies. Some specimens of sulphur are freely soluble




MTIL  OF SULPHUR.                   163
in bisulphide of carbon, while of other samples only a comparatively small portion dissolves. We distinguish, therefore, a soluble
and an insoluble modification of sulphur.
Octahedral sulphur, the bright-yellow, translucent, native crystals for example, is completely soluble in bisulphide of carbon.
But of the soft, elastic sulphur, such as was prepared in Exp. 82,
as much as 30 or 40 per cent. is completely insoluble in the
bisulphide, whether this liquid be hot or cold.
No method has as yet been discovered of preparing pure insoluble sulphur directly; but it can always be readily obtained
by dissolving out the soluble sulphur from a mixture of the two
varieties, such as the soft sulphur above mentioned. Flowers of
sulphur contain a considerable portion of insoluble sulphur; rollbrimstone much less, though the interior of the sticks contains
decidedly more than the outside portions. It may be observed,
in this connexion, that flowers of sulphur are prepared by suddenly cooling the vapor of sulphur, while the soft variety is
obtained by suddenly cooling melted sulphur.
Insoluble sulphur undergoes no change at the ordinary temperature; but if it be kept for a long time at 1000, or if it' be
exposed to the vapor of water or alcohol, it is slowly converted
into the soluble variety.
198. For some pharmaceutical purposes, sulphur is prepared as
a powder finer even than flowers of sulphur. This preparation is
known as Nilk of sullphur or precipitated sulphur.
Exp. 83.-Place in a small flask as much flowers of sulphur as can
be taken up on the point of a penknife; pour into the flask 10 or 15
c. c. of a solution of caustic soda, and boil the mixture for some time.
Part of the sulphur will dissolve and color the liquid yellowish brown.
Pour off the clear liquid from the undissolved sulphur, mix it with
an equal volume of water and stir in dilute chlorhydric acid, added by
small portions, until a drop of the mixture placed upon litmus paper
exhibits an acid reaction. As the acid is added, the liquid assumes a
milky appearance from the separation of sulphur in the form of an
exceedingly fine powder. This powder is so light that, for a long
while, it will not subside, but remains suspended in the liquor, imparting to it a milky appearance.
Collect the powder on a small filter, wash it with water, and dry it
at a gentle heat. It will now appear as a pale yellowish-grey impalM 2




164               METALS BURN IN SULPHUR.
pable powder. If it be heated more strongly, so that it melts, the
color will become distinctly yellow, the numberless small particles of
the powder being now compacted into a single mass.
It will be remarked that the color of flowers of sulphur is
lighter than that of roll-brimstone, while the color of the precipitated sulphur is far lighter than that of the flowers.  Such
differences as these are common; they depend upon a difference
of mechanical condition, upon differences in the state of aggregation of the particles of the substances which exhibit them.
The method of pulverization by precipitation, employed in this
experiment, is a general method, applicable to many other substances besides sulphur.
199. Sulphur unites energetically with most of the other elements, such union being, in many cases, attended with evolution
of light. Most of the metals, for example, combine with it
directly, just as they do with oxygen.
Exp. 84. —Melt in an ignition-tube, 12 to 15  Fig. 40.
c.m. long, 4 or 5 grms. of sulphur, and heat the
liquid until it boils; then throw in small portions of copper filings, or fine turnings, and
observe the violent action which ensues.
Or a strip of very thin sheet copper or a coil
of fine copper-wire may be suspended in the
hot sulphur vapor, in the upper part of the
ignition-tube; it will glow vividly as it unites
with the gaseous sulphur, much in the same
way as if it were burning in oxygen gas. The
product of the reaction, in either case, is called
sulphide of copper.
Erp. 85. —Mix intimately 4 grms. of flowers
of sulphur and 7 grinms. of the finest iron-filings. Place the mixture in
an ignition-tube 10 to 12 c.m. long, and heat the upper part of the
tube over the gas-lamp. In a short time the mass will begin to glow,
as the sulphur and iron enter into chemical comlbination, and this
ignition will, of itself, pass through the entire length of the tube,
even if the lamp be withdrawn. The final product of the reaction is
protosulphide of iron.
200. As has been already shown (~~ 2, 109), phenomena of
combustion, such as are exhibited in these experiments, are
directly referable to chemical union. They are strictly analogous




USES OF SULPHUR.                  165
to the ordinary processes of combustion in which oxygen is
involved, though the technical term, combustion, is by custom
limited to the act of combination with oxygen.
Sulphur combines with chlorine, bromine, iodine, and phosphorus at the ordinary temperature, and with carbon at a red
heat. With oxygen it unites readily at a comparatively low
temperature.  When heated in the air, it takes fire at about
2500, and burns with a peculiar blue light. This easy infiammability may be readily illustrated by blowing flowers of sulphur
into the hot air issuing from the chimney of an Argand gas-lamp;
the sulphur takes fire at a considerable height above the flame.
The irritating, suffocating gas, which is produced by the union of
sulphur and oxygen, will be shortly described under the name of
sulphurous acid.
Several important practical applications of sulphur depend upon
this property of igniting and continuing to burn at a moderate
heat. It is, in fact, largely employed as a kindling material.
By means of it, other bodies less readily combustible, can be
heated to the temperature at which they continue to burn.
Hence its use upon matches and in gunpowder and fireworks.
201. In its chemical properties, sulphur is closely allied to
oxygen; like oxygen, it forms a great variety of compounds
with a wide range of different elements; and the series of compounds thus obtained is in many respects parallel with, or comparable to, the series of oxygen compounds.
It is an important raw material in the chemical arts, being an
ingredient of numerous useful compounds, such as cinnabar, ultramarine, vulcanized caoutchouc, bisulphide of carbon, chloride of
sulphur, and the various compounds of sulphur and oxygen, one
of which, sulphuric acid, is the most important chemical agent at
present employed in manufacturing industry. Sulphur is largely
employed in medicine, in the treatment of cutaneous diseases of
both men and domesticated animals, and has been of late years
extensively used in the vineyards of Europe for destroying a
parasitic fungus which infests the vines.
202. Sulphydric Acid (H,S). —.When sulphur is sublimed in
hydrogen gas, or when hydrogen is passed over melted sulphur,
combination takes place between the two elements, though very




1] 66          PREPARATION OF SULPHYDRIC ACID.
slowly and imperfectly, so that only a comparatively small quantity of the compound is obtained, even if the process be continued for a long time.  Somewhat larger quantities of it are
formed when a mixture of hydrogen gas and sulphur-vapor is
passed through a tube filled with fragments of pumice-stone
heated to about 500~.
When the two elements meet in the nascent state, they combine readily; thus, when organic bodies containing sulphur
putrefy, and when they are subjected to destructive distillation,
sulphydric acid is evolved, just as ammonia is under the same
circumstances. In either event, the product of the union is a
colorless gas of highly offensive odor, like that of rotten eggs.
An easier method of preparing sulphydric acid, or sulphuretted
hydrogen, as it is often called, is by acting upon a compound of
mslphur and iron with dilute chlorhydric acid.
Exp. 86.-In a gas-bottle (Fig. 41) put 10 or 12 grins. of protosulphide of iron, see Exp. 85; re-            Fi  41.
place the cork in the bottle and
introduce the gas-delivery-tube
into another small bottle containing cold water, letting the tube
dip 5 or 6 c.m. beneath the surface of the water. Through the
thistle-tube, pour into the gas-  /
bottle water enough to seal the
lower extremity of this tube;
then add, through the thistletube, as before, 2 or 3 teaspoonfuls of muriatic acid. and observe........
that bubbles of gas soon begin to
pass through the water in the absorption bottle.
Sulphydric acid is soluble in water to a considerable extent, and is
consequently taken up by the water in the absorption bottle. The
solution thus obtained, known as sulphuretted hydrogen-water, is
much employed as a reagent in chemical laboratories; it will serve us
here as a convenient source of sulphydric acid.
VWhen the disengagement of gas slackens, a new portion of muriatic
acid may be added through the thistle-tube, and this process continued
until the water in the absorption bottle smells strongly of the gas.
This experiment should be performed out of doors, or in a draught




PREPARATION OF SULPHYDRIC ACID.           167
of air so arranged that those portions of the gas which escape solution
shall be carried away from the operator.
203. The reaction between the sulphide of iron and chlorhydric acid in the foregoing experiment is somewhat analogous
to that which occurs in the preparation of hydrogen, ~ 50. If
metallic iron (or zinc) be treated with chlorhydric acid, hydrogen
is evolved, according to the equation
Fe + 2HC1 -- FeClo + 2H.
But if, instead of simple iron, sulphide of iron, whose formula
is FeS, be taken, sulphur will be eliminated, as well as hydrogen,
by the action of the acid, and these elements, as they come together in the nascent state, will unite to form sulphydric acid.
FeS +  2HC1 =  FeCI2 +  H2S.
Instead of absorbing the gas evolved in the foregoing experiments in water, it might be collected as such, over a basin holding but a small quantity of water, or, better, filled with warm
water or with brine, either of which absorbs less of the gas than
cold water. Unless absolutely dry, the gas cannot be collected
over mercury, since, when moist, it acts upon this metal.
204. At the ordinary temperature and pressure, sulphydric
acid is a gas somewhat heavier than air, its specific gravity being
17, referred to hydrogen; but under a pressure of about 15
atmospheres at 110, it becomes liquid. The specific              Fig. 42.
gravity of this liquid referred to water is 0'9. At
— 85~ the liquid solidifies
to a white crystalline mass.
205. Since the sulphide
of iron, employed in the
preparation of suiphydrlc
acid, is usually mixed with
a certain quantity of metallic iron, the gas is liable
to be contaminated with
free hydrogen. For all or-  
dinary purposes, the gas   -




168              ANALYSIS OF SULPHYDRIC ACID.
thus mixed with hydrogen serves as well as if it were pure; but
in some cases a gas free from hydrogen is required. In order to
prepare it, sulphide of antimony is substituted for the sulphide
of iron.
1 part of powdered sulphide of antimony placed in a thin-bottomed
flask is treated with 3 or 4 parts of chlorhydric acid of 1'1 sp. gr. and
the mixture gently heated. The apparatus may be arranged as in
Fig. 42, in which the bottle into which the gas first enters contains a
very small quantity of water; this water serves to remove any particles
of the acid or of solid matter which may have been carried over in the
current of gas. In case a dry gas be needed, a chloride-of-calcium
tube (see Appendix, ~ 15) must be interposed between the wash-bottle
and the mercury-trough.
206. The volumetric composition of sulphydric acid gas is one
volume of sulphur-vapor and two volumes of hydrogen, condensed
to two volumes. Its molecule, therefore, contains one atom of
sulphur and two atoms of hydrogen, and is strictly analogous to
the molecule of water.
The composition of sulphydric acid may be determined experimentally by heating metallic tin in a confined volume of the gas. An
ignition-tube 20 or 30 c.m. long, bent at an obtuse angle, within 5 to
6 c.m. of the closed extremity, as shown in Fig. 43, should be completely filled with dry sulphydric acid gas over the mercury-trough,
and then closed with the thumb and inverted.
Some granulated tin should be dropped into
the tube and made to lodge in the bent part,
the thumb being instantly replaced upon the
mouth of the tube the moment the tin has
entered.
The tube full of gas is now replaced in the
mercury-trough, as shown in the figure, and
about one-third part of the gas is allowed to escape by inclining the
tube so that the gas may bubble out through the mercury. The tube
and its contents are left at rest during half an hour, in order that they
may acquire the temperature of the surrounding air, a caoutchouc ring
is slipped down the tube to mark the height of the gas, and the tin is
then heated with the flame of a spirit-lamp. The hot tin combines
with the sulphur, and hydrogen is set free. The apparatus is left at
rest during another half hour, and the height of the gas in the tube
is then noted. Tf the gas employed was pure, it will be found that its
volume has undergone no change. The hydrogen which has been set




PROPERTIES OF SULPHYDRIC ACID.           169
free occupies precisely the same space as the sulphydric acid did before
it was decomposed.
It is evident from this that 1 volume of sulphydric acid gas contains
1 volume of hydrogen, or multiplying these numbers by 2, in order to
arrive at the composition indicated by our molecular formula, that 2
volumes of sulphydric acid contain 2 volumes of hydrogen. Now the
specific gravity, or unit-volume weight of sulphydric acid has been
found, by experiment, to be 17-19, that of hydrogen being 1; and if
From the weight of 2 volumes of sulphydric acid... 34'38
VWe subtract the weight of two volumes hydrogen... 2'00
There will remain.... 3238
which is very nearly equal to the unit-volume weight of sulphurvapor, 31'8, as experimentally determined.
The composition of sulphydric acid, both by volume and weight,
may, therefore, be expressed by the diagram.
li
207. The gas is very poisonous; when respired in the pure
state it quickly proves fatal, and it is very deleterious,. even
though largely diluted with atmospheric air. Small birds soon
die in air which contains only T-r.51,  of its volume of the gas, dogs
in air which contains -1, and horses in air which contains  I
of its volume. Men can support more of it, but in experimenting
with it, it is best to do so where there is a free circulation of air.
Nausea and headache are often produced when an atmosphere
even slightly contaminated with sulphuretted hydrogen has been
breathed for any length of time. In case the air of an apartment become contaminated with the gas, the disgusting smell
can readily be neutralized by sprinkling the room with chlorinewater, or by evolving a little chlorine gas by adding some dilute
acid to a small quantity of bleaching-powder.
The gas exists as a natural constituent of some mineral waters
which are thence called sulphurous, such as the Virginia Sulphur Springs, and the mineral springs at Sharon, N. Y. It is




170            SULPHURETTED-HYDROGEN WATER
also found in the air and water of foul sewers, and wherever
animal matter is undergoing putrefaction.
208. Sulphydric acid gas is readily inflammable, and, like hydrogen, extinguishes the flame of a burning candle immersed in
it. It burns with a blue flame, producing water and sulphurous
acid gas.
HQS +  30 =  H20 +  SO2.
In case it be ignited in contact with a quantity of air insufficient
to burn the whole of it, the hydrogen will burn first, and a portion of the sulphur will escape combustion.
If a tall glass cylinder be filled with sulphydric acid gas, and the
gas be lighted at the top, the flame will pass down the cylinder as the
hydrogen is consumed, and a quantity of very finely divided solid sulphur will be deposited upon the walls of the vessel.
It has already been stated that sulphur kindles very easily, and that
it has a strong affinity for oxygen; but it appears from this experiment, that hydrogen kindles still more readily, and that its affinity for
oxygen is greater than that of sulphur.
When mixed with air, in certain proportions, it is explosive;
a fact which should be borne in mind by the experimenter.
209. Water dissolves about three times its own volume of the
gas at the ordinary temperature. This solution (see Exp. 86) is
transparent and colorless when recently prepared, but, when kept
it gradually becomes opalescent and turbid from deposition of
sulphur. Oxygen from the air unites with the hydrogen of the
sulphydric acid to form water, and sulphur is set free. After the
lapse of several weeks or months, it will be found that the solution
no longer contains any sulphydric acid; it has lost its nauseous
odor, and the bottom of the bottle is covered with sulphur, the
result of the decomposition of the dissolved gas.
The aqueous solution of the gas reddens litmus slightly, like
the very weak acids. Towards metals and metallic oxides it behaves in a manner somewhat analogous to that of chlorhydric
acid and its congeners, while, with regard to metallic sulphides,
it stands in the same relation as water to the oxides, as will be
explained hereafter.
Exp. 87.-Place a drop of sulphuretted-hydrogen water (Exp. 86)
upon a bright piece of copper, lead, or silver. The metal will quickly




SULPHYDRIC ACID AS A REAGENT.            171
become black. The sulphur of the sulphydric acid unites with the
metal, to form sulphide of copper, sulphide of lead, or sulphide of silver, as the case may be, while the hydrogen escapes.
Cu + H2S = CuS + 2H.
Exp. 88.-Place in a test-tube as much litharge (oxide of lead =
PbO) as can be held upon the point of a penknife, pour upon it a teaspoonful of sulphuretted hydrogen-water, and observe that the yellow
litharge imnmediately becomes black. Sulphide of lead is formed, as in
the preceding experiment, together with a quantity of water.
PbO + H2S = PbS + H20..Exp. 89.-In place of the litharge of the last experiment, take a
very small crystal of nitrate of lead (Exp. 42); dissolve it in as much
water as will half fill the test-tube, and to this solution add a few drops
of the sulphuretted-hydrogen water. Black sulphide of lead is thrown
down as a precipitate, and nitric acid is set free.
PbO,N206 + H2S = PbS + H20O,N205.
210. Since many of the metallic sulphides are, like the sulphides of lead, copper, and silver, insoluble in water and dilute
acids, sulphuretted hydrogen is peculiarly well adapted for precipitating the metals from their solutions. After having been
thrown down as sulphides, as in the last experiment, they can be
readily separated and collected by filtration.
Though many of trie metallic sulphides are black, like that of
lead, this is not true of all. Several of them exhibit characteristic colors, by which they may be readily recognized; thus the
color of sulphide of antimony is orange, that of sulphide of arsenic
is yellow, and that of sulphide of zinc white. Upon this fact the
application of sulphuretted hydrogen as a test or reagent (that is,
as a means of detecting and identifying many metals) is in part
based.
211. In the same way that sulphuretted hydrogen can be employed for detecting the presence of metals, so, conversely, solutions of the metals, or, in some cases, the metals themselves, may
be used as tests for sulphuretted hydrogen.
Exp. 90.-Prepare a strong aqueous solution of nitrate of lead, or
better, of acetate of lead (sugar of lead). Wet strips of white paper
3 to 4 c.m. wide with this solution, and dry them in air which is free
from sulphuretted hydrogen. This lead-paper, as it is called, should be
kept for use in tightly stoppered bottles.




1 72 9DECOMPOSITION OF SULPHYDRIC ACID.
Moisten a bit of the lead-paper with water and expose it to some
source of sulphuretted hydrogen-the mouth of the bottle of sulphydric acid prepared in Exp. 86, for example, or to the fetid air of a sewer.
The paper will immediately be blackened from formation of sulphide
of lead.
The blackening of silver-ware, of watches, and cards which
have been glazed with a preparation of lead, at many mineral
springs, and other localities where sulphuretted hydrogen  is
evolved, is, in like manner, indicative of the presence of this
gas. Tests like these, which are continuous and cumulative, are,
of course, much more delicate means of detection than the mere
odor of the gas.
212. Sulphydric acid is a compound which is very easily
decomposed.  When  simply heated, it breaks up  into  its
components; and it is readily destroyed by various chemical
agents.
Exp. 91.-To a gas-bottle such as was employed in Exp. 86, containing sulphide of iron, attach a chloride-of-calcium tube (Appendix,
~ 15) and a piece of hard glass tubing, No. 4, about 20 c.m. long. To
the end of this glass tube, attach another tube bent at right angles and
dipping into a bottle of water. Pour chlorhydric acid into the gasbottle, so that sulphydric acid shall be freely generated, as seen by the
flow of bubbles through the final bottle of water. After the lapse of
some minutes, when the apparatus has become completely filled with
the gas and the last portions of air have been expelled, heat the middle
of the tube of hard glass with the flame of the gas-lamp, and observe
the ring of sulphur which will collect upon the walls of the cold portion
of the tube a short distance in front of the flame.
It will be seen in subsequent chapters that several other of the
gaseous compounds of hydrogen are decomposed, like sulphydric acid,
upon being passed through hot tubes.
The influence of oxygen, in decomposing the aqueous solution
of sulphydric acid, has been already alluded to, ~ 209; it has
been observed, moreover, that air, contaminated with sulphydric
acid soon becomes odorless of itself, oxygen uniting with hydrogen, as before, and sulphur being set free. All the oxidizing
agents (that is, substances which readily give up oxygen) decompose sulphydric acid, water being formed and sulphur deposited..Exp. 92.-Into a test-tube containing 4 or 5 c. c. of sulphuretted




PERSULPHIDE OF HYDROGEN.                173
hydrogen-water (Exp. 86), pour half as much concentrated nitric acid.
Sulphur will be deposited and nitrous fumes evolved.
Very dilute nitric acid will not thus decompose sulphydric acid.
Chlorine, bromine, and iodine vapor instantly decompose sulphydric acid, uniting with its hydrogen to form chlorhydric,
bromhydric, or iodohydric acid, while sulphur is precipitated.
1H2S + 2C1 = 2HC1 + S.
Exp. 93.-In place of the nitric acid of the preceding experiment,
pour a few drops of chlorine-water into the solution ot sulphydric acid,
and observe that the odor of the latter is destroyed.
213. Persulphide of Hydrogen (HS2,?).-This is an exceedingly unstable liquid, the composition of which is not accurately
known,though it is supposed to be analogous to that of the peroxide
of hydrogen. It can be prepared by adding a solution of persulphide of calcium to diluted chlorhydric acid.  The reaction
may be conceived to take place in accordance with the following
equation:CaS, +  2HC1 =  CaC12 +  H1S2  +  3S.
Exp. 94. Mix 75 or 100 grins. of flowers of sulphur and an equal weight
of slaked lime with half a litre of water, place the mixture in a flask
and heat it to boiling, taking care to agitate the flask so that the solid
matter may not become impacted upon it. Continue to boil for about
an hour, then filter off the liquor from the undissolved portions of
sulphur and lime. The solution thus obtained is a mixture of several
sulphides of calcium, more highly sulphuretted than the protosulphide,
but will serve the present purpose as well as if it were the pure quinquisulphide.
Pour the solution of sulphide of calcium into 250 c. c. of a mixture
of 2 parts of concentrated chlorhydric acid and 1 part of water. Persulphide of hydrogen will separate in fine oily drops, producing a milky
turbidity in the liquid. These drops soon coalesce and settle out beneath the water. A good way of collecting the persulphide is to perform the precipitation in a large glass funnel, provided with a stopper.
By carefully opening this stopper, the precipitated oil can nearly all be
drawn off without disturbing the water which floats above it.
Persulphide of hydrogen emits a peculiar, disagreeable odor,
and is very irritating to the eyes and mucous membrane. It
tastes sweet and bitter, but disorganizes the flesh wherever it




174              NOMENCLATURE-PREFIXES.
touches it. Its properties closely resemble those of peroxide of
hydrogen; it is very unstable, and is decomposed by the same
substances which destroy the oxide. It even decomposes spontaneously when left at rest for a few days; ordinary vegetable
colors are quickly bleached by it; it decolorizes also solutions of
indigo.
214. In the last section we have used, for the first time, certain technical terms which, perhaps, need brief explanation.
As has been stated in ~ 76, many of the elements are capable of
uniting with other elements in several different proportions to
form chemical compounds. Sulphur, for example, is specially apt
to form more than one compound with a single element. When
sulphur unites with a metal, the compound formed is called a
sulphide, just as a compound of oxygen and a metal is called an
oxide, or one of chlorine and a metal a chloride, —the termination ide, which always indicates combination, being added to the
first syllable of the word sulphur, or oxygen, or chlorine, and the
new word ending in ide being then connected with the name of
the metal, as in the case of suiphide of copper, Exp. 87. But
when, as in the case of calcium, there are several distinct sulphides,
it is customary to distinguish one from the other by means of
various Latin and Greek prefixes. Thus the compound which
contains one atom of sulphur and one atom of calcium is the protosulphide, or simply the sulphide of calcium, the prefix proto being
derived from the Greek word for first; the compound which contains two atoms of sulphur to one of calcium is the bisulphide of
calcium, from the Latin for twice; and in like manner we have
a tersulphide, containing three atoms of sulphur to one of calcium,
and a quinquisulphide containing five atoms of sulphur. The
compound containing the highest proportion of sulphur is often
called the persulphide. A good custom is to designate the compounds which contain more sulphur than the protosulphide by
prefixes of Latin origin, and to distinguish those which may contain less sulphur than the protosulphide by means of Greek prefixes; thus, if there were a compound of two atoms of calcium
and one of sulphur (Ca2S) it would properly be called a di-sulphide
of calcium, the prefix being from the Greek Ocs. The same prefixes are used in an analogous manner in connexion with the




COMPOUNDS OF SULPHUR AND OXYGEN.         175
words oxide, chloride, bromide, iodide, and the similar words
ending in ide.
215. Com2pounds of Sulphur and Oxygen.-No less than seven
different compounds of sulphur and oxygen have been discovered,
all of which form acids by union with water. Thus the oxide of
sulphur SO3 forms, by union with the elements of water, common
sulphuric acid HlSO4; the name sulphuric acid being indiscriminately applied to both bodies, although only that one which
contains hydrogen possesses the properties commonly described
by the term acid.
Two of these compounds, viz. sulphurous acid (SO,), and
sulphuric acid (SO,), have long been known, and these are still,
comparatively speaking, of most importance, since they are employed upon the large scale in the arts. Subsequently there were
found the compounds S,20 (hyposulphuric acid) and S202 (hyposulphurous acid); and at a still more recent period the compounds
83O5, S4O5, and S,05.
As long as only two compounds of sulphur and oxygen were
known, they were distinguished as sulphurous and sulphuric, in
accordance with the rule laid down ill ~ 70; when the two compounds, S205 and S202, containing respectively less oxygen than
sulphuric and sulphurous acids, were discovered, the prefix hypo
was resorted to as explained in ~ 71; lastly, for the later-found
acid compounds of sulphur and oxygen, the ordinary rules of
chemical nomenclature being inadequate, it was necessary to
invent a special set of names. They were all called Thionic acids,
from the Greek word for sulphur, and were then distinguished
from one another by the prefixes tri, tetra, and penta, in accordance with the number of atoms of sulphur in each.  Strictly
speaking, the compound S20,, since it contains five atoms of oxygen like the thionic acids, should perhaps follow the new rule and
be called dithionic acid, but it is still customary to retain the old
name hyposulphuric acid.
The complete list of the names of the compounds of sulphur
and oxygen is as follows:Sulphurous acid... SO,
Sulphuric acid.....             SOS
Hyposulphurous acid......             S202




176                    sULPHUROUS ACID.
Hyposulphuric acid (or Dithionic acid).. S205
Trithionic acid.......... SO5
Tetrathionic acid......... S405
Pentathionic acid......... S305
Of all these compounds, only sulphurous acid can be readily
obtained by the direct union of sulphur and oxygen.  The others
must be prepared by circuitous methods.
216. Sudphurous Acid (SO,).-This acid is produced when
sulphur is burned in the air or in pure oxygen gas.
Exp. 95. —Lig.ht a piece of sulphur in a deflagrating spoon Fig. 44.
and suspend the latter in a half-litre bottle full of air. On
examining the contents of the bottle, after the sulphur has
ceased to burn, there will be found an irritating, suffoicating 
gas having the peculiar odor which is familial as that of
a burning match.  The bottle is now full of Sulphurous
acid gas, mixed with the nitrogen originally present in the
air.
217. By burning sulphur in oxygen gas, instead of in air, as
in the preceding experiment, a much purer product could, of
course, be obtained.  But the experiment would be chiefly interesting in enabling us to determine synthetically the composition
of sulphurous acid.
If sulphur be burned in a confined volume of dry oxygen gas, it will
be found, after the combustion has terminated, and the gas has been
allowed to regain its original temperature, that the volume of the sulphurous acid produced is sensibly the same as that of the original oxygen, though its weight is twice as great. Hence 1 volume of sulphurous acid gas contains 1 volume of oxygen. Now, if from the
weight
Of l unit-volume of sulphurous acid, as determined by experiment................   32 25(
We subtract the weiglit of 1 unit-volume of oxygen... 159(39
There will remain... 16287
or not far from one-half the number, 32, which represents the real
specific gravity or equal volume weight of sulphur-vapor. Consequently 1 volume of sulphurous acid gas contains half a volume of
sulphur-vapor, besides 1 volume of oxyen. Or, multiplyiug these
numbers by 2, in order to avoid a fractional volume, it appears that
the volumetric composition of sulphurous acid is 1 volume of sulphur




PREPARATION OF SULPHUROUS ACID.             177
vapor and 2 volumes of oxygen condensed to 2 volumes of the compound gas. Or, expressed in the form of a diagram:+  1       =   SO2  64
H2o
218. An easier method of preparing pure sulphurous acid is by
depriving common sulphuric acid of part of its oxygen. This can
be effected by a variety of reducing or deoxidizing agents. For
example, when concentrated sulphuric acid is heated with metallic
copper or mercury, there are formed a sulphate of the metal, water,
and sulphurous acid:Cu +  2H2S04 =  CuSO4 +  2H20 +  S02.
Exp. 96.-Into a thin-bottomed glass flask of half a litre capacity,
put 14 grms. of copper clippings, or turnings, and 50 grms. of concentrated sulphuric acid. Attach to the flask a delivery-tube and connect
this with a series of Woulfe's bottles, such as were employed in the
preparation of chlorhydric acid (Exp. 49); heat the flask over the
gas-lamp until the acid begins to react upon the copper, then quickly
withdraw the lamp for a moment, lest the contents of the flask boil
over, and finally regulate the flame so that a steady current of sulphurous acid shall pass through the water in the Woulfe bottles. After
the first tumultuous evolution of gas has subsided, the flask can be
slowly heated without further trouble. The current of gas should be
kept up until the water of the first bottle has become saturated, or, at
the least, highly charged with the gas.
Sulphurous acid is freely soluble in water, which, at 150, takes up
something like 44 times its bulk of the gas; hence the solution obtained
as above may be used as a convenient vehicle for sulphurous acid. On
account of this easy solubility, the gas cannot be collected over water;
but it can be collected over mercury, or by displacement. Since the
gas is more than twice as heavy as air, the method by displacement is
to be recommended, if an efficient ventilating flue is at command to
carry off that portion of the suffocating gas which must escape.
Mercury is, in some respects, better than copper for use in this experiment. It affords a much more regular evolution of gas, and the
operation requires less care; but copper is usually employed on account
of its comparatively low cost.
1'




178            PROPERTIES OF SULPHUROUS ACID
Instead of copper or mercury, as in the foregoing experiment, other
reducing agents, such as sulphur or charcoal, may be employed. If 1
part of powdered sulphur be boiled with 12 parts of strong sulphuric
acid, sulphurous acid is set free, as exhibited by the following equation:
S + 2H2S04 = 3S02 + 2H20.
The evolution of gas, in this case, though steady and uniform, is
comparatively slow, as contrasted with that of the experiment in which
copper is employed; hence the process is usually less convenient than
that with copper.
If bits of charcoal or dry sawdust are heated with sulphuric acid, a
copious evolution of sulphurous acid occurs, though the gas is not, in
this case, pure, being mixed with half a volume of carbonic acid.
C + 2H2S04 = 2S02 + C02 + 21120.
For many purposes, as in the preparation of the aqueous solution of
sulphurous acid, this method with charcoal is to be preferred, on the
ground of economy and convenience of application. In the laboratory
it is perhaps more frequently employed than either of the others.
Sulphurous acid mnay also be readily prepared by heating in an ignition-tube a mixture of 4 parts of sulphur and 5- parts of black oxide
of manganese, both in fine powder, and intimately mixed. A mixture
of 3 parts of black oxide of copper with 1 part of sulphur answers the
same purpose:
2S + AnO2 = SO2 + MnS.
3S + 2CiuO = SO2 + 2CuS.
In both these cases, metallic sulphides are left as a residuum in the
ignition-tube; but if the sulphur and black oxide of manganese be mixed
in the proportion of 1 part sulphur to 52 parts of the oxide, no sulphide,
but only protoxide, of manganese will be formed.
S + 2Mn02 = SO2 + 2MnO.
219. As has been already stated, sulphurous acid gas is transparent and colorless. It is irrespirable and suffocating, and when
mixed with air, even in small proportion, occasions violent coughing. It is not inflammable, but, on the contrary, it stops combustion.
The flame of a taper is immediately extinguished on being immersed
in sulphurous acid gas, just as it is by nitrogen. A useful application
of this property of the gas is in extinguishing burning chimneys. A
handful of fraglments of sulphur being thrown upon the hot coals in
the grate, and the openings of the fire-place being closed in such a




SULPHURoUS ACID STOPS COMBUSTION.          179
manner that no air shall enter the chimney, excepting that which passes
through the fire, the chimney will quickly become filled with an atmosphere of sulphurous acid mixed with nitrogen from the air employed
in burning the sulphur, and the burning soot upon the walls of the
chimney will be immediately extinguished.
It is, of course, essential that the chimney should then be closed at
the top, so that air may be excluded and the chimney kept full of the
fire-extinguishing atmosphere until its walls shall have cooled to below
the kindling temperature of the soot.
The oxygen contained in the sulphurous acid gas is so firmly
held that combustibles are powerless to take it away under ordinary circumstances, though at high temperatures this oxygen can
be removed by means of hydrogen, carbon, and easily oxidizable
metals like potassium.
When hydrogen and sulphurous acid gas are passed together
through a red-hot tube, water is formed and sulphur deposited,
4H + SO2 = 2H20 + S,
and when sulphurous acid is passed through a tube containing
ignited charcoal, carbonic acid is produced and sulphur deposited,
as before.
C + SO2 = CO2 + S.
In case nascent hydrogen come in contact with sulphurous acid,
it will decompose it at the ordinary temperature, though in a
manner somewhat different from the foregoing.  The sulphur, as
well as the oxygen, will, in this case, combine with hydrogen,
and there will be formed sulphydric acid as well as water.
6H  +  S02 =  2H20 +  H2S.
This reaction may be made visible by putting a few drops of a solution of sulphurous acid (Exp. 96) into a gas-bottle from which hydrogen is being evolved (Exp. 19), and testing the hydrogen with a strip
of moistened lead-paper (Exp. 90) both before and after the addition
of the sulphurous acid.
220. Sulphurous acid can readily be condensed to the liquid
state.  It is, in fact, one of the most easily liquefiable of the
gases. By mere cooling to — 10~, under the ordinary pressure
of the air, it is converted into a colorless, transparent, limpid
liquid.
In preparing small quantities of the liquid, it is sufficient to lead the
N2




180                LIQUID SULPrHUOUS ACID.
gas, prepared from copper and sulphuric acid, and dried by passing it
through sulphuric acid or over chloride of calcium, into a U-tube which
is immersed in a freezing mixture of ice and salt (2 parts of pounded
ice and 1 part salt).
Liquid sulphurous acid is a rather heavy liquid, of 1'4911
specific gravity, boiling at about -100 and solidifying at — 76~,
to a colorless crystalline solid. On being exposed to the air at
ordinary temperatures, the liquid acid evaporates with great
rapidity, and consequently occasions very intense cold.  By
means of it, mercury may be frozen, and chlorine and ammoniagas liquefied.
If a quantity of the liquid acid be poured into water, the temperature of which is a few degrees above 0~, a portion of the acid will evaporate at once, another portion will dissolve in the water, and a third
portion of the heavy oily liquid will sink to the bottom of the vessel.
If the portion, which has thus subsided, be stirred with a glass rod, it
will boil at once, and the temperature of the water will be so much
reduced that a portion, or even the whole, of the water will be frozen.
221. The specific gravity of the gas, as determined by different
observers, is 32-256, or 32'443, or 32'558, instead of 32, as would
be indicated by theory. This variation is explained by the fact
that sulphurous acid, like all the easily condensible gases, ceases
to conform exactly to the lavW of Mariotte at temperatures near
to its point of condensation.  Under any given pressure, its
volume decreases in somewhat larger proportion than is the case
with air and the other permanent gases.
An important property of sulphurous acid is its power of
bleaching vegetable colors. It is extensively employed in bleaching articles of straw, wool, silk, &c., which would be injured by
chlorine.
EExp. 97.-Into a bottle in which sulphur has been burned (Exp. 95)
pour a few teaspoonfuls of a solution of blue litmus, and shake the
bottle. The litmus solution will first become red, as it would if any
other acid than sulphurous were present, and will then be decolorized.
The same property may be illustrated by holding a red rose in the
fumes of burning sulphur, or by immersing the rose in an aqueous solution of sulphurous acid (Exp. 96), and leaving it for a few minutes
until it has become white.
In the same way the stains of fruit or wine can be removed from




SULPHUROUS ACID BLEACHES.               181
clothing. A bit of sulphur is burned beneath a small open cone of
paper, which serves as a chimney, and the stain, having first been
slightly moistened with water, is held in the fumes at the top of this
chimney. The cloth should finally be carefully washed with water at
the place where it has been exposed to the sulphurous acid.
In the arts, the process of bleaching is usually conducted in large
chambers, in which the slightly moistened articles are hung while
sulphur is burned below. The damp goods absorb the sulphurous acid
and gradually become white. The presence of water is essential;
perfectly dry sulphurous acid will not bleach.
222. The manner in which sulphurous acid acts as a bleaching agent is not clear. It is remarkable that, as a general rule,
it does not actually destroy the coloring-matter; and that upon
many coloring-matters it has little or no action. Mlost of the
yellows, and the green coloring-matter of leaves, are in this latter
category, and upon litmus, cochineal, and logwood, the acid does
not act very readily. In the few instances where it really destroys the color, as in the case of the garden amaranthus, it
appears to act as a deoxidizing agent.  But in most cases it
appears to enter into combination with the coloring-matters and
to form colorless compounds. These colorless compounds of sulphurous acid and coloring-matter can be broken up, with restoration of color, by exposing them to the action of various chemical
agents capable of expelling sulphurous acid.
Exp2. 98.-Bleach a rose, as in Exp. 97, and immerse it in dilute
sulphuric acid. Then dry and warm it, so that the volatile sulphurous
acid may be driven off. The color of the rose will again appear.
In many cases a solution of caustic soda will restore the color as well
as sulphuric acid. A practical illustration of this action of alkaline
solutions is seen in the reproduction of the original yellow color of the
wool when new flannel is washed with an alkaline soap.
Sulphurous acid is a powerful disinfecting and antiseptic agent.
It retards, to a remarkable extent, the processes of putrefaction
and fermentation, and is largely employed for this purpose in
wine-making; hops and compressed vegetables are charged with
it to the same end, and it has been successfully employed for
preserving meat. It has often been employed in medicine, in
the treatment of skin diseases, as a fumigation.
223. Although sulphur will not take up more than two atoms




1.82           OXIDATION OF SULPHUROUS ACID.
of oxygen when burned in the air, or in oxygen gas, it is nevertheless a matter of no very great difficulty to cause it to take up
a third atom.
In presence of water it gradually absorbs oxygen from the air,
and is converted into sulphuric acid. Hence the aqueous solution of sulphurous acid (Exp. 96) cannot be preserved for any
great length of time, unless it be kept in very tight vessels.
224. If a mixture of sulphurous acid gas and oxygen, or air,
be brought in contact with hot platinum sponge, the sulphurous
acid will unite with oxygen, and sulphuric acid will be formed;
the same union occurs when the mixed gases are brought in contact with various other substances, such as pumice-stone, clay,
and the oxides of chromium, iron, and copper.  Several attempts
have been made to put these methods in practice for manufacturing sulphuric acid, but they have been found to be too slow,
and in the case of platinum and clay, it has been observed that
these substances soon lose their power and cease to convert the
mixed gases into sulphuric acid.
Sulphurous acid is, indeed, a deoxidizing agent of very considerable power; and is much employed in the laboratory as a
reducing agent.  It decomposes iodic acid with separation of
iodine, and nitric acid with evolution of hyponitric acid, sulphuric
acid being formed in both cases.
120  +  51120 +  5SO2 =  5(H120,S03) +  21.
H20,N20, +  802  -1 H,,S0,    +  2N02.
Exp. 99.-Charge a dry bottle, of the capacity of a litre  Fig. 45.
or more, with sulphurous acid gas, by burning in it a bit of
sulphur, as shown in Fig. 45. Fasten a shaving, or, better,
a tuft of gun-cotton, upon a glass rod or tube bent at one
end in the form of a hook; wet the shaving in concentrated
nitric acid, and hang it in the bottle of sulphurous acid.
Red fumes of hyponitric acid will immediately form about
the nitric acid, and will gradually fill the bottle.
In presence of a mixture of water and chlorine, sulphurous
acid takes up an atom of oxygen from the water, while the
hydrogen of the water unites with chlorine.
SO2 +  21H20 +  2C1 =   1H20,SO3 +  2HC1.
A similar reaction occurs between iodine and sulphurous acid, if




SULPHITES.                    183
a very large amount of water be present; in spite of the fact,
already mentioned, ~ 140, that iodohydric acid is readily decomposed by concentrated sulphuric acid with liberation of iodine,
sulphurous acid, and water.
225. Sulphurous acid, though a weak acid, forms numerous
well-defined salts by uniting with metallic oxides. These salts,
called sulphites, are of two classes, —simple or normal sulphites,
such as the sulphite of potassium K2S03 (or, dualistic, KO,SO,),
and double or acid sulphites, such as the acid sulphite of potassium KHSO3 (or, dualistic, KHO,SO2).  All these salts are decomposed by strong acids, such as chlorhydric, nitric, or sulphuric, sulphurous acid being expelled; but they are not decomposed by carbonic acid. On the contrary, the salts of carbonic
acid are decomposed by sulphurous acid; and hence it happens
that the impure sulphurous acid gas obtained by heating a mixture of charcoal and sulphuric acid can be used for preparing the
sulphites. If, for example, this gas be conducted into an aqueous
solution of carbonate of sodium, there will be obtained a solution
of sulphite of sodium, and carbonic acid will be set free.
226. Besides the solution of sulphurous acid, such as was prepared in Exp. 96, there is a definite crystalline compound of
water and the acid, which can be obtained by passing a current
of sulphurous acid gas into ice-water. This compound is very
unstable, and is destroyed at temperatures but little above 0~;
but by collecting it upon a cooled filter and then pressing the
crystals repeatedly between folds of cold blotting-paper, it has
been found possible to remove most of the mother-liquor which
adheres to them at first, and to obtain the compound in a condition of tolerable purity. The composition of the crystals appears
to be S02+15H20.
227. Sulphuric Acid. —The term  sulphuric acid is applied
somewhat indiscriminately to three or more distinct substancesnamely, to a compound of one atom of sulphur and three atoms
of oxygen, SO,, which we shall call anhydrous sulphuric acid,
and to certain compounds of sulphur, oxygen, and hydrogen,
which have been usually regarded as compounds of the anhydrous su]phuric acid, just mentioned, and water.  Of these
hydrates the most important are those of the composition H2S04




184                    SULPEUlRIC ACID.
(dualistic, H0, SO3) [oil of vitriol], and H2S,0, (dualistic,
H120,2SO,) [Nordhausen or fuming sulphuric acid].  The body,
whose formula is 1H2SO4, is one of the most important of chemical substances, and is usually the thing meant when sulphuric
acid is spoken of.  We will therefore proceed to study its properties before touching upon those of the other substances abovementioned.
Sulphuric acid is one of the most important products of chemical manufacture, and is made in enormous quantities. In the
same way that the metal iron may be said to be the basis of all
mechanical industries, sulphuric acid lies at the foundation of the
chemical arts. By means of sulphuric acid, the chemist either
directly or indirectly prepares almost everything with which he
has commonly to deal.
Sulphuric acid might be prepared by passing sulphurous acid
gas into boiling nitric acid, until all of the latter had been reduced, and finally distilling off the last traces of the lower oxides
of nitrogen which would be formed. Even if sulphur itself were
boiled in concentrated nitric acid, it would gradually be oxidized
and converted into sulphuric acid.  But neither of these processes would be economical. It can be very cheaply prepared,
however, by the action of either of the high oxides of nitrogen,
nitrous, hyponitric, or nitric acids, upon sulphurous acid, in presence of air and moisture; and this method is the one actually
followed in the preparation of sulphuric acid on the large scale.
A mixture of the gases above mentioned is effected in enormous
chambers constructed of sheet lead, a metal upon which cold sulphuric acid has little or no action.
228. The essential points of the process are, first, that SO2,
when in presence of much moisture, can take oxygen from either
NN203, NO,2 or N205, and reduce them to nitric oxide, NO, while it
is itself converted into sulphuric acid, and, secondly, that NO can
take oxygen from the air and become NO2.
In practice, the sulphurous acid is obtained by burning crude
sulphur, or more commonly a compound of sulphur and iron, known
as iron-pyrites, FeS,; the gas, together with a large excess of atmospheric air, is then conducted into the first of a series of leaden chambers into which steam is admitted. Nitrous fumes are supplied either




MANUFACTURE OF SULPHURIC ACID.               185
by allowing nitric acid to fall in fine streams through the incoming
current of sulphurous acid and air, or from the decomposition of a
mixture of salt, nitrate of sodium, and sulphuric acid, as described in
~ 105, or by heating a vessel charged with nitrate of sodium and sulphuric acid, by means of the burning sulphur.
In conformity with the principles above stated, the sulphurous acid,
as soon as it comes in contact with the steam, reacts upon the nitrous
fumes; there is formed nitric oxide gas and hydrated sulphuric acid,
which falls to the floor. But, as there is present in the chamber an
excess of air, the nitric oxide immediately unites with a portion of the
oxygen therein contained, and is converted into hyponitric acid.
This hyponitric acid immediately reacts upon a new portion of sulphurous acid, and the process thus goes on through a whole series of leaden
chambers, the very small portion of nitric acid at first taken being sufficient to prepare a large quantity of sulphuric acid. In reality, the
oxygen employed in converting the sulphurous into sulphuric acid, all
comes from the air, excepting a very little at first; the nitrous fumes
serve only as a conveyer of oxygen. The nitric oxide takes oxygen
from the air and transfers it to the sulphurous acid, which, as has been
stated in ~ 223, is, by itself and unaided, incapable of combining with
oxygen. It will, of course, be understood that, although we trace out
these reactions as if they were consecutive, they are really, so far as we
know, simultaneous.
Theoretically, a single portion of hyponitric acid would be sufficient
to effect the conversion of an unlimited amount of sulphurous into sulphuric acid; but practically this power is qualified by a variety of circumstances. It is found to be impossible, for example, to mix new
portions of air with the mixture of sulphurous acid and nitric oxide
for an indefinite period; for at a certain point these gases become so
loaded down with nitrogen derived from the air already consumed, that
they are as good as lost in it. In general the flow of gases is so regulated that all the sulphurous acid shall be oxidized, and that nothing
but nitric oxide and waste nitrogen shall pass out of the last leaden
chamber.
229. The process of manufacturing sulphuric acid can readily
be illustrated upon the small scale.
A large glass balloon, or receiver, of the capacity of several litres,
placed in a vertical position, is closed with a cork pierced with five
holes, through four of which are passed small glass tubes. All of these
glass tubes reach nearly to the bottom of the balloon, and are bent at
a right angle above the cork; one of the tubes is connected at the top
with a flask containing copper-turnings and sulphuric acid, for the




186            MANUFACTURE OF SULPHURIC ACID.
generation of sulphurous acid (see Exp. 96), another with a flask
containing copper-turnings and furnished with a thistle-tube, through
which nitric acid can be poured, for the generation of nitric oxide (see
Exp. 37), and the third with a flask containing water for the evolution
of steam; the fourth tube and the fifth hole are both left open.
Everything being in readiness, nitric oxide is generated in the small
flask fitted for this purpose; as the gas passes over into the large balloon it unites with oxygen from the air, and red fumes of hyponitric
acid are formed. Sulphurous acid is now made to pass into the balloon; this will have no action upon the red fumes, so long as there
is no water present, but the moment steam is thrown in from the third
small flask, a reaction occurs, the hyponitric acid is reduced, and the
sulphurous acid oxidized. By means of bellows, air must, from time
to time, be blown into the balloon, through the fourth glass tube, the
waste nitrogen passing off through the fifth hole in the cork.
If but little steam be employed in this experiment, a solid compound,
formed by the union of nitrous and anhydrous sulphuric acids, is liable
to be deposited upon the walls of the balloon; the appearance of this
body always indicates that the supply of steam is insufficient; it is
never formed when the proper proportion of moisture is present.
230. The sulphuric acid which collects at the bottom of the
leaden chambers is necessarily dilute, because of the'large amount
of water which must be present, in order that the reactions above
described may freely occur; moreover it would not be advantageous to allow an acid more concentrated than that of specific
gravity 1'4 to form in the chambers, since a stronger acid would
absorb and retain a considerable quantity of nitric oxide. To
make it fit for the purposes for which sulphuric acid is usually
employed, the dilute acid of the chambers must be concentrated
by expulsion of the water; to this end, it is run off into shallow
leaden pans, and there evaporated until it is of specific gravity
1'71 to 1'75. The concentration cannot safely be carried beyond
this point in ordinary leaden vessels, since the strong, hot acid
begins to attack the metal, and the temperature at which the
liquid boils is so high as to approach the melting-point of lead.
This acid of 1-72 specific gravity is somewhat extensively employed, for a variety of purposes, at the factories where it has
been prepared, but is still too dilute for transportation. It is
therefore transferred from the leaden pans to large glass retorts
set in deep sand baths, or to piatinum stills, and there evaporated
further, until it is nearly of the composition HISO4.




PROPERTIES OF SULPHURIC ACID.             187
231. The acid thus boiled down is the concentrated sulphuric
acid, or oil of vitriol, of commerce; its specific gravity is usually
about 1'83, that of the absolutely pure acid being 1P842. Besides
this slight excess of water, it contains also, in solution, a certain
quantity of sulphate of lead, and a variety of other impurities.
For most purposes, however, it will answer as well as the pure
acid. Like the latter, it is a heavy, oily, colorless, and odorless
liquid, boiling at about 330~.
Since a comparatively small amount of heat is absorbed in the conversion of the liquid acid to the condition of gas, its vapor can be very
easily condensed; in distilling the acid, the receiver need not even be
placed in cold water.
From the same cause, combined with the great weight of the liquid,
the acid is liable to boil tumultuously, the act of ebullition being
irregular and attended with violent blows or shocks. The bubbles of
vapor formed at the bottom of the retort condense almost as soon as
they are formed, and the heavy liquid above suddenly falls back to fill
the vacuum.
In distilling the concentrated acid, it is therefore best to heat only
the upper portions of the liquid in the retort; this can be effected
either by placing the retort upon a wire-grate so perforated that about
half the body of the retort can be sunk below the level of the burning
charcoal upon the grate, or by placing a layer of ashes, or of some
other bad conductor of heat, beneath the very bottom of the retort,
then piling sand around the sides of the retort outside of the ashes,
and applying heat beneath the iron pan upon which the whole is supported.
232. The common acid usually freezes at about -34~; but it
has been found possible to lower the freezing-point to -80~, by
adding a small quantity of water to the commercial acid. When
once frozen, it remains solid until the temperature rises to about
the freezing-point of water. Crystals of the pure acid melt at
about 100. At the ordinary temperature, sulphuric acid does not
vaporize, but, on the contrary, greedily absorbs water from the
air and so increases in bulk. In moist weather, its bulk may
increase to the extent of a quarter or more, in the course of a
single day, and, by longer exposure, a still larger quantity of water
will be taken up; the acid must always be kept, therefore, in
tightly stoppered bottles.




188           SULPHURIC ACID ABSORBS WATER.
Exp. 100.-Into a shallow dish of about 200 c. c. capacity, pour
about 75 c. c. of concentrated sulphuric acid; place this dish of acid
upon one pan of a balance, and upon the other pan put enough small
shot, or clean, dry sand, to exactly balance the acid. Preserve the material of the counterpoise, and place the dish of acid uncovered in the
open air; from day to day replace it upon the balance, together with
the counterpoise, and note the number of grammes or fractions of a
gramme that it has increased in weight.
If the acid were allowed to stand for a week or two in a damp place,
it might become two or three times as heavy as it was at first. From
its power of absorbing aqueous vapor, sulphuric acid is often employed
for drying gases. (See Appendix, ~ 15.)
233. With liquid water sulphuric acid unites with great energy,
much heat being evolved at the moment of combination; during
the union a certain amount of condensation occurs, the mixture,
when cold, occupying less space than was previously occupied by
the acid and the water. The water and acid may be mixed in
all proportions, being mutually soluble one in the other.
In mixing water and sulphuric acid, the acid should always be
poured into the water, in a fine stream, not the water into the acid,the water being meanwhile stirred. In this way the heavy acid has
an opportunity to mix with the water as it sinks down through it.
If, by any accident, water were to fall upon sulphuric acid, it would
float on top of it, and great heat would be developed at the point of
contact of the two liquids; if the quantities of acid and water were
large, sudden bursts of steam would be occasioned and serious damage
might arise from the scattering about of portions of the acid.
In mixing water and commercial sulphuric acid as in the following
experiment, it will be observed that the solution becomes cloudy, and
that a white powder is gradually deposited from it. This precipitate
is sulphate of lead, originally derived from the leaden pans in which
the acid was concentrated; it is soluble in concentrated, but insoluble
in dilute sulphuric acid, and is consequently thrown down when water
is added to the commercial acid.
Exp. 101.-Place in a beaker glass of about 250 c. c. capacity, 30
c. c. of water; in accordance with the directions above given, pour
into the water 120 grins. of concentrated sulphuric acid, and stir the
mixture with a narrow test-tube containing a teaspoonful of water.
So much heat will be evolved during the union of the water and the
acid that the water in the test-tube will boil.
234. If sulphuric acid be mixed with ice or snow, the latter




MIXING SULPHURIC ACID WITH WATER.           189
will be immediately liquefied. If the proportion of ice in the
mixture be small, as compared with that of the sulphuric acid,
heat will be evolved much as is the case with liquid water, though
to a less extent.  But when a large proportion of ice is mixed
with a comparatively small quantity of the acid, no heat will be
perceived, but, on the contrary, intense cold.
Exp. 102.-Place in a beaker glass of about half a litre capacity
120 grms. of snow, or finely pounded ice; pour upon it 30 grms. of
concentrated sulphuric acid, and stir the mixture with a test-tube containing a small quantity of water. The water in the tube will be
frozen.
Exp. 103.-Repeat the foregoing experiment, using 30 grms. of
snow or ice and 120 grins. of sulphuric acid. A very considerable
evolution of heat will occur, as may be seen more clearly by immersing
a thermometer in the liquid.
The result of Exp. 102 seems, at first sight, inconsistent with
the general fact that heat is always set free during chemical combination; for though chemical union between the acid and water
has evidently occurred, no heat, but cold, is manifested. The
anomaly is only a seeming one; a certain amount of heat is required, in order that the cohesive force, by which the particles of
the ice are held together, shall be overcome; hence the heat which
is really produced by the chemical combination is all absorbed,
together with much more, taken from the materials and the vessel
which contained them, during the liquefaction of the ice.
235. Besides the indefinite mixture or solution above mentioned, several crystalline compounds of anhydrous sulphuric acid
and the elements of water, of fixed composition and characteristic
form, can be prepared.
If the commercial acid be diluted with water until its specific
gravity is reduced to 1 78, and the liquid be then cooled strongly,
a substance of composition H 4SO, (dualistic, 2H120,SO3) will crystallize out in the form of large rhombic prisms. These crystals
are of sp. gr. 1'785; they melt and solidify at about 8~.
A second hydrate, of composition 3H20,S03, can be obtained
by evaporating a dilute acid in a vacuum at the temperature of
100~, until it ceases to lose weight; and another of composition
H2,0,2S03, will be described below when we come to speak of
fuming sulphuric acid.




190              PROPERTIES OF SULPIIURTC ACMID
236. Slflphuric acid is one of the most powerful acids known.
If one drop of it be diluted with a thousand times as much water,
it is still capable of reddening blue litmus. It expels most of the
other acids from their compounds, in the same way that we have
seen it expel nitric acid from nitrate of sodium in Exp. 32. At
very low temperatures, however, as at — 80~, it loses its power of
reddening litmus, and has no action upon the carbonates, though
it acts violently upon these salts at the ordinary temperature.
It is intensely caustic and corrosive, and quickly chars and
destroys most vegetable and animal substances.
Exp. 104.-Into a test-glass pour a tablespoonful of sulphuric acid
and immerse in it a splinter of wood. The wood will blacken as if
charred by fire, and the acid will become dark-colored. Wood is composed of carbon, hydrogen, and oxygen; and since sulphuric acid
unites with compounds of hydrogen and oxygen, rather than with
carbon, a portion of the latter is left free; some carbonaceous matter
is, however, dissolved by the acid and darkens it. The acid of commerce is often dark-colored, from fragments of straw or other organic
matter having accidentally fallen into it. Sulphuric acid which has
been colored by organic matter may be rendered colorless by strongly
heating it till it becomes fully concentrated.
The action of the acid upon organic matter is more rapid when moisture is present. Thus, if a few drops of oil of vitriol be poured upon
dry paper, decomposition will take place only slowly; but if a little
water be added to the acid, heat will be developed by the chemical
union, and the paper will be at once decomposed by the hot acid.
237. When heated with charcoal or with any organic matter,
sulphuric acid gives up oxygen, as has been shown in Exp. 96,
and is itself reduced to sulphurous acid; by sulphur, also, and
by several of the metals, such as copper and mercury, it is reduced in a similar way. (See Exp. 96.) Towards some metals,
such as zinc for example, its behavior is various, according as it
is concentrated or dilute. If zinc be treated with cold, dilute
sulphuric acid, the zinc simply replaces the hydrogen of the acid,
sulphate of zinc is formed, and hydrogen is set free.
Zn +- 112S04 =  ZnSO,4 +  21H.
But if zinc be heated with concentrated sulphuric acid, a portion
of the latter is reduced, as it would be in presence of copper or




SULPHATES.                     191
mercury, sulphurous acid is evolved, as well as hydrogen, and
these gases, reacting upon each other, produce sulphydric acid and
a deposit of sulphur, in accordance with the following formulae:802 +  6H  =  H2S +  21H20.
SO2 + 4I  =  S    +  2H20.
As a general rule, concentrated sulphuric acid acts but feebly
upon the metals in the cold, though, when boiled upon them, it
often behaves as an oxidizing agent.
238. With the oxides of the metals, sulphuric acid unites
directly to form the very important salts called sulphates, water
being simultaneously eliminated. Oil of vitriol, H2SO4, may, in
fact, be itself regarded as a salt, in the composition of which, hydrogen fills the same place that sodium  does in sulphate of
sodium, Na2S04, and it might well be called sulphate of hydrogen, were it not that usage has assigned to it another name.
Besides the normal sulphates, in which all the hydrogen has been
replaced by a metal, as above (or, on the dualistic hypothesis,
in which all the water has been replaced by a metallic oxide),
there is another class of sulphates, often called bi or acid sulphates, in which only half of the hydrogen has been thus replaced; as an example of these, the student will recall the acid
sulphate of sodium, NaHSO4, mentioned in ~ 101.
Acids, which, like sulphuric acid, contain two replaceable atoms
of hydrogen, and are therefore capable of forming two series of
salts, are called bibasic, in contradistinction to the monobasic
acids, like nitric acid, which form but one series of salts. There
is but one nitrate of sodium, for example, NaNO3.  It is for this
reason that many chemists object to the doubled formula for
nitric acid H 2N206, in spite of its convenience, because this formula suggests, what is not true, that one or both of the atoms of
hydrogen might be replaced by any metal which, like sodium,
potassium, or silver, replaces hydrogen atom for atom.
239. Fuming Sulphuric Acid.-The acid H2S04, above described, has been, for nearly a century, the most important of the
several varieties of sulphuric acid; but long previous to the discovery of the process of making it in leaden chambers, there
was manufactured another variety, now known as fuming sulphuric acid.




192               FUMING SULPHURIC ACID.
This fuming acid, or Nordhausen acid, as it is often called,
from the name of a German town in which large quantities of
it were formerly prepared, was at first obtained by distilling in
earthen retorts the salt now known as sulphate of iron, formerly
called green vitriol.  Hence the origin of the name oil of vitriol,
which, in England and this country, has come to be applied
solely to the common acid H2SO4, though it is still used as a
synonyme for the fuming acid by German writers.
When dry sulphate of iron is exposed to a full red heat, it
suffers decomposition; a considerable quantity of sulphuric acid
is given off and can be collected in receivers. The distillate
thus obtained, which is a dense fuming liquid of about 1'9 specific gravity, is the acid now in question. Though of far less
importance than was formerly the case, consi lerable quantities of
the fuming acid are still prepared for the purpose of dissolving
indigo and for other special uses, where an acid stronger than
the common acid is needed. It may be regarded as a solution
of varying quantities of the anhydrous acid SO3 in the common
acid H2SO4; if it be gently heated, all of the anhydrous acid will
be expelled, and common sulphuric acid will remain. So, too, if
it be exposed to the air, the anhydrous acid will be given off, and,
coming in contact with the moisture of the air, will combine
therewith to form common sulphuric acid, which, falling as a
cloud, occasions the appearance of fumes.
When the fuming acid is cooled to about — 5, a crystalline
compound of composition 1T2S,07 (dualistic, H,20,2SO,) separates
out. After having been freed from liquid acid, these crystals
melt at 35~. When pure, the fuming acid is colorless; but the
commercial article is often brown, from having been in contact
with organic matter. It is an excessively corrosive liquid, and
destroys most organic matters, even more rapidly than the common acid. On being dropped into water, a noise is emitted as
if a red-hot bar of metal had touched the water.
240. Anhydrous Sulphuric Acid (SO,).-As has been mentioned in ~ 224, this substance can be obtained by passing a mixture of sulphurous acid gas and oxygen over hot, finely divided,
metallic platinum, or over various oxides and other porous substances.




MAKING ANHYDROUS SULPHURIC ACID.             193.Exp. 105.-Prepare a small quantity of platinized asbestos as follows: dissolve about 0'25 grm. of metallic platinum in aqua regia (~ 104),
and soak in this solution as much soft, porous asbestos as will form a
loose ball of 1'5 c.m. diameter; heat the wet asbestos gently until it
has become dry, and then ignite it in the flame of the gas-lamp. The
chloride of platinum, which was formed by the solution of the metal, is
decomposed by heat, alid metallic platinum, in a finely divided condition. is left adhering to the asbestos.
Select a tribe of hard glass, No. 3, about 30 c.m. long, and, at a distance of about 10 c.m. from one end, bend it to an obtuse angle, so
that when the tube is supported upon a ring of the iron stand above
the gas-lamp the shorter bent portion can be thrust into the neck of a
receiver; in the centre of the longer portion of the glass tube pack the
platinized asbestos loosely; then force into and through the tube a current of mixed sulphurous acid and oxygen; at the same time heat over
the gas-lamp that portion of the tube which contains the platinized
asbestos, and collect the anhydrous sulphuric acid which is formed in a
perfectly dry test-tube or U-tube immersed in ice, or, better, in a freezing-mixture of ice and salt. In the course of the experiment, lift up
the receiver for a moment, pour out from it a little of the vapor of
anhydrous sulphuric acid, with which it is filled, and observe how
rapidly the heavy gas falls through the air, and the cloud which forms
as it unites with moisture.
The mixture of sulphurous acid and oxygen may be made before the
experiment in a small gas-holder, or, as well, during the progress of
the experiment in a bottle behind the asbestos tube. This bottle, which
should be of at least half a litre capacity, is fitted with a cork carrying. three glass tubes, and is connected with the asbestos tube by one
of these tubes, which reaches no lower than the cork; by the other
tubes, which pass nearly to.the bottom of the bottle, and dip beneath the
surface of a layer of common sulphuric acid, which has- been placed
in it, the bottle is connected with a flask in which sulphurous acid is
being generated (Exp. 96), and with a gas-holder containing oxygen.
The sulphuric acid in the bottle serves to dry the gases, and. the rapidity
with which the bubbles of gas pass through, the liquid, enables the
operator to judge of the proportions in which the gases are being
mixed; the flow of oxygen havilng been fixed. at a moderate rate, once
for all, the sulphurous acid will alone. need attention.
The action of the platinum, in, this experiment is obscure'; it will be
treated of under the metal platinum.
Instead of the platinized asbestos, oxide of iron,. oxide of copper, or
oxide of chromium, or, better, a mixture of the- last two can be heated
0




194       PROPERTIES OF ANHYDROUS SULPHURIC ACID.
in the tube through which the mixed gases are passing. These processes of preparing sulphuric acid are interesting from a scientific point
of view, but, as has been already stated (~ 224), they do not admit of
commercial application.
Anhydrous sulphuric acid can readily be prepared by heating the
Nordhausen acid (see ~ 239):20 or 30 grms. of fuming sulphuric acid are poured into a perfectly
dry, small glass retort; the neck of the retort is thrust into a dry, cold
receiver, and the acid is slowlv heated until it boils moderately. The
vapor of the anhydrous acid will condense and solidify in the receiver.
The anhydrous acid may also be obtained by distilling dry bisulphate
of sodium. The bisulphate is prepared by heating a mixture of 3 parts,
by weight, of dry sulphate of sodium and 2 parts of common sulphuric
acid, until the mixture fuses. All the water of the acid is thus eliminated:Na2O,SO3 + 120,SO3 = Na20,2SO3 + HO.20
The bisulphate of sodium, on being distilled in an eathern retort, will
give up one molecule of anhydrous sulphuric acid, and a residue of
normal sulphate of sodium will remain in the retort:Na20,2SO3 = Na20,SO3 + SOa.
241. As thus prepared, anhydrous sulphuric acid is a glistening white solid mass of silky, crystalline fibres, somewhat resembling asbestos; it is tough and ductile, and can be moulded with
the fingers like wax. So long as no water is present, it can be
handled without danger; when perfectly dry, it is not corrosive,
nor does it even react upon blue litmus. It unites with water,
however, with great avidity, and is converted into common sulphuric acid. It rapidly absorbs water from the air and deliquesces; at the same time it forms dense fumes; for it is volatile,
to a very considerable extent, at the ordinary temperature, and
its vapor combines with the moisture of the air.  On being
brought in contact with a small quantity of water, it combines
with it with explosive violence, and much heat is evolved. If a
bit of it be thrown into a large quantity of water, the water
hisses as if a hot iron had been thrust into it. Owing to its
great:tendency to deliquesce, the solid acid can only be preserved
in dry tubes sealed at the lamp.
The specific gravity of anhydrous sulphuric acid is 1'97. It
melpts-readilyupo i being heated; but it has been noticed that




COMPOSITION OF ANHYDROUS SULPIHURIIC ACID.   195
some samples melt far more easily than others. There appear
to be two distinct varieties of the acid;, for in some cases a temperature of 180 is sufficient to render the mass fluid, while in
others the heat must be carried even to 1000. The easily fusible
modification appears to change gradually, by keeping, into that
which is more difficultly fusible; and the latter seems to be
changed to the former by distillation. The melted acid boils at
about 350, and evolves a colorless and transparent vapor, three
times as heavy as air, which, upon coming in contact with the
air, unites with moisture and forms dense white fumes. When
brought in contact with hot lime or baryta (oxide of calcium and
oxide of barium), it unites with them directly; intense heat is
evolved, and there is formed sulphate of calcium or sulphate of
barium:
CaO + SO3 =  CaSO4 =  CaO,SO,.
242. On being exposed to a strong red heat, the vapor of anhydrous sulphuric acid splits up into oxygen and sulphurous
acid-two volumes of it yielding two volumes of sulphurous acid
and one volume of oxygen. As has been shown in ~ 217, two
volumes of sulphurous acid gas contain one volume of sulphur
vapor and two volumes of oxygen; hence it follows that the
volumetric composition of anhydrous sulphuric acid is one volume
of sulphur vapor and three volumes of oxygen, the whole condensed to two volumes. The specific gravity of sulphur vapor is
32, that of oxygen is 16, and the proportions, by weight, in which
sulphur and oxygen are united in anhydrous sulphuric acid, are
consequently 32 and 16 x 3=48, the combining weight of sulphuric acid being 32+48 = 80. The combining weight of sulphuric acid can also be readily determined in a manner analogous
to that employed in the case of nitric acid (~ 73), by saturating
with the common acid a known quantity of oxide of lead,
evaporating off the water and excess of acid, and then determining
the weight of the dry sulphate of lead which is -formed. By
subtracting from the latter the weight of the original oxide of
lead, we obtain the weight of the sulphuric acid which has combined with it. Experiment will show that the weight of this
sulphuric acid is to that of the oxide of lead in the ratio of 80
to 223.
o2




196               HYPOSULPTHUROUS ACID.
The facility with which sulphuric acid is decomposed at a red
heat (~ 242) is the basis of a very economical method of preparing oxygen gas in large quantities for manufacturing-purposes.
Commercial sulphuric acid is allowed to drop upon fragments of
red-hot porcelain, there to be decomposed, in accordance with
the formula
H2SO4 = S02 + 0 + H20,
and the products of the decomposition are then washed with
water, so that the sulphurous acid may be absorbed, the steam
condensed, and the oxygen left free. Here again, as in our
earlier experiments (~ 10), the oxygen has really been obtained
from the air; and if it were desirable, the solution of sulphurous
acid obtained in washing this oxygen, might be placed in the
leaden chambers and again be converted into sulphuric acid by
the addition of oxygen from the air.
243. Hyposulphurous Acid (S,0,) has never been obtained
in the free state, nor is any compound of it with water known;
but there are numerous saline compounds of which it makes
part, and some of these are of considerable importance in the
arts.  These salts, called hyposulphites, can be prepared in
various ways,-for example, by digesting sulphur in a hot (but
not boiling) concentrated solution of an alkaline sulphite. If sulphite of sodium be taken, the reaction can be thus formulated,
Na2O,S02 + S =  Na,0,S202.
Another method of preparing the hyposulphites is to pass a current of sulphurous acid gas through the solution of an alkaline
sulphide, until no further precipitation of sulphur occurs
2CaS -- 3SO2 = 2(CaO,S20,) + S.
When a hyposulphite is treated with a strong acid, decomposition immediately ensues; S202 breaks up into  02,+S; hence
our inability to isolate the acid.
Some of the hyposulphites will be more fully described when
we come to treat of the metals.
244. Other Conzmpounds of Sulphur and Oxygen.-With the
exception of hyposulphuric acid, S205, none of these compounds
(see ~ 215) have been very thoroughly studied; any detailed




CHLORIDES OF SULPHUR.                197
description of the methods of preparing them would be out of
place in an elementary manual.
245. C'ompounds of Sulphur and Chlorine. —Chlorine and sulphur combine with one another directly and readily, forming
several different compounds, whose properties vary in accordance
with the varying proportions of chlorine and sulphur which they
respectively contain.
246. Dichloride of Sulphur (SCI) is the best-known of the
compounds of chlorine and sulphur, and is often called simply
chloride of sulphur.
It can be prepared by passing a current of dry chlorine through a
flask or tubulated retort containing flowers of sulphur. The chlorine
is rapidly absorbed by the sulphur, and care must be taken lest the
mass become too hot. The reddish-yellow liquid, obtained as the
result of the reaction, is a solution of sulphur in dichloride of sulphur;
by distilling it the excess of sulphur can be separated.
Dichloride of sulphur is a yellowish-brown liquid of 1'68
specific gravity, and boiling, without decomposition, at 1440.
It emits a peculiar odor, which has been likened to that of
sea-plants; its vapor excites tears, and its taste is acid, acrid,
and bitter. It fumes strongly in the air, being decomposed by
the moisture of the air with evolution of chlorhydric acid. It is
decomposed by water, but can be mixed with bisulphide of carbon and with benzine. It is remarkable as a powerful solvent
of sulphur; 100 parts of dichloride of sulphur can take up about
70 parts of sulphur at the ordinary temperature; on slowly cooling the hot saturated solution, beautiful crystals of sulphur are
deposited.  Dichloride of sulphur is used in a process of vulcanizing caoutchouc, known as the cold process.
247. Chloride of Sulphur (SC12).-This compound is formed
when sulphur is treated with an excess of dry chlorine, or when
a current of chlorine is passed into dichloride of sulphur; the
dichloride requires some 278 times its own bulk of chlorine gas,
and absorbs it very slowly. Chloride of sulphur is a red liquid,
of 1'625 specific gravity.  It exhales suffocating and irritating
fumes of chlorine and the dichloride, since it slowly decomposes
when kept. The decomposition is particularly rapid in a strong
light; and so much gas is evolved that a tightly-stoppered bottle




198                     SELENIUM.
containing chloride of sulphur will explode, after a time, if it
be) placed in sunlight. On being heated, the liquid gives off so
much chlorine at 50~ that it seems to boil; but the temperature
gradually rises to 64~, which appears to be the real boiling-point
of the liquid. It is slowly decomposed by water.
The density of its vapor has been found to be 53; admitting
that the gas is composed of one volume of sulphur vapor and two
volumes of chlorine, condensed to two volumes of vapor, the calculated specific gravity of its vapor would be 51'5. In view of
the instability of the compound, the experimental result is sufficiently near coincidence with the calculated number to make it
certain that the composition of the gas is really as above stated.
248. The other compounds of sulphur with chlorine, and with
chlorine and oxygen, need not here be discussed; and the same
remark applies to the compounds formed by the union of sulphur
with iodine, bromine, fluorine, and nitrogen. The compounds of
sulphur with carbon, phosphorus, arsenic, and the metals will be
treated of hereafter.
CHAPTER XIV.
SELENIUM AND TELLURIUM.
249. These elements are rare, and of little or no industrial
importance; but to the chemist they are exceedingly interesting,
on account of the close resemblance they bear to sulphur. The
three elements, sulphur, selenium, and tellurium, constitute a
group which is equally remarkable with that formed by chlorine,
bromine, and iodine. (See ~ 152.)
250. Selenium, Se,. is never found in any considerable quantity in any one place. Traces of it occur in many varieties of
native sulphur, and in various metallic sulphides. It is now
obtained chiefly from the sulphides of iron, copper, and zinc.
These sulphides often contain minute traces of selenium, though
the quantity is sometimes so small that it can hardly be detected




PROPERTIES OF SELENIUM.               199
by the ordinary methods of analysis. When these sulphides are
burned for the purpose of manufacturing sulphuric acid, or in
metallurgical operations, the selenium goes off with the sulphurous acid produced by, the combustion, and is deposited either in
the dust-flues of the furnaces or upon the floors of the leaden
chambers at the sulphuric-acid works. In this way the selenium
from hundreds of tons of the pyritous ores is collected and concentrated to a comparatively small bulk. The deposit taken
from the leaden chambers of some sulphuric-acid works contains
as much as from 2 to 10 per cent. of selenium. The methods of
obtaining pure selenium from these deposits are founded upon
the facts that by treatment with nitric acid or aqua regia the
selenium can all be oxidized and converted into selenious acid
(SeO2), that selenious acid is soluble in water, and that when a
solution of it is treated with sulphurous acid, the selenious acid
is reduced and pure selenium deposited.
SeO2 +  2SO2 =  2SO3 +  Se.
251. In its properties and in its chemical behavior, selenium
resembles sulphur in many respects, while in others it is like
tellurium. Like sulphur and oxygen, it occurs in distinct allotropic modifications (~~ 162, 196). The precipitate obtained by
mixing solutions of sulphurous and selenious acids is of a deep
red color, almost like that of cinnabar. But, after having been
fused and suddenly cooled, selenium appears as a brilliant black
mass, amorphous, like glass, and of 4-3 specific gravity. When
fused selenium has been slowly cooled, it appears as a dark-grey,
very brittle, crumbling mass, of crystalline or granular structure
and a metallic lustre like that of lead; the specific gravity of
this variety is 4-81. The amorphous or vitreous modification of
selenium does not conduct electricity; but the granular or crystalline variety conducts it, and the more readily in proportion as it
is hotter. The specific heat of selenium, at the ordinary temperature, is 0'0746, being the same for bbth the vitreous modification and that with metallic lustre. The vitreous variety is
soluble in bisulphide of carbon, but the granular variety is insoluble in that liquid.
Selenium melts readily upon being heated, and the liquid thus




200                    ISOMORPIMSM.
obtained boils at about 7000, being converted into a dark yellow
vapor, the specific gravity of which has been found to be 82'3.
The atomic weight of selenium has been determined to be 79'5.
This discrepancy between the vapor-density and the atomic weight
is to be ascribed to the imperfection of the experimental determinations. Of itself, selenium has neither taste nor odor. When
heated in the flame of a lamp, it burns with a beautiful blue
flame and exhales a peculiarly offensive odor, like that of putrid
horseradish,-selenious acid, SeO2, being the chief product of the
reaction.
Selenium combines with most of the elements, usually in the
same way as sulphur, though not always, since it is a weaker
chemical agent than sulphur; its compounds are as a rule somewhat less stable than the corresponding sulphur compounds.
With oxygen it forms selenious acid, SeO2, and selenic acid,
SeO,,-analogous to sulphurous and sulphuric acids respectively.
Besides these, there is a lower oxide, SeO (?); it is a colorless
gas, having the strong and disagreeable odor like horseradish
before mentioned.
252. Both selenious and selenic acids form numerous salts,
which closely resemble the corresponding sulphites and sulphates
in composition and in many of their properties. Normal seleniate
of potassium, for example, K2SeO4, cannot be distinguished, by its
external appearance, from  sulphate of potassium, K2SO04,-the
crystalline form of the two bodies, as well as their texture, color,
and lustre, being identical. If solutions of these two salts be
mixed, neither of the salts can subsequently be crystallized out
by itself when the solution is evaporated; the crystals obtained
will be composed of sulphate of potassium and seleniate of potassium mixed in the most varied proportions. Bodies which
are thus capable of crystallizing together in all proportions, without alteration of the crystalline form, are said to be isomorpihous
(like-formed). The formulae of the two isomorphous salts, just
mentioned, differ only in this-that the one contains the atom
Se, where the other contains the atom S. It is therefore possible to replace 32 parts by weight of sulphur by 79'5 parts of
selenium, or 79'5 of selenium by 32 of sulphur, without changing
the crystalline form of the salts; it follows that 32 parts by




ATOMIC VOLUME-TELLURIUM.               201
weight of solid sulphur, and 79'5 parts of solid selenium, occupy
the same space. That this is actually the case may be shown
by comparing the quotients obtained by dividing the atomic
weights of the two elements by their specific gravities; these
quotients will be found to be equal, or as nearly equal as the
limits of error of the physical determinations involved will permit. The specific gravity of prismatic sulphur is 1'91, or, in
other words, one cubic centimetre of solid sulphur weighs 1'91
gramme; the specific gravity of crystalline selenium is 4.81, or
one cubic centimetre of selenium  weighs 4'81 grammes' 32
grammes of sulphur will, therefore, occupy  i3-=16'75 cubic
centimetres; 79-5 grammes of selenium will occupy 4.5=16'53
cubic centimetres. What is true of grammes, is true of any parts
by weight, and ultimately of the atoms. This quotient, obtained
by dividing the atomic weight of an element by its specific gravity, is called the atomic volume of the element; it must be borne
in mind that the standard of specific gravity for liquids and solids
is water, for gases hydrogen, and, therefore, that the atomic
volume of a solid or liquid must not be directly compared with
that of a gas. Two elements whose atomic volume is the same can
be exchanged in their compounds without alteration of crystalline
form, precisely as a brick or stone taken out of a wall can be
replaced by another of the same size and shape without changing
the form of the wall.
253. With chlorine, selenium forms two compounds, SeCl and
SeC14, the first of which is analogous to dichloride of sulphur.
With hydrogen it forms a compound, H2Se, called selenhydric
acid, or seleniuretted hydrogen, which is perfectly analogous to
sulphuretted hydrogen, H2S, but possesses a still more disagreeable odor. In its action upon solutions of the metallic salts, upon
metals and metallic oxides, selenhydric acid behaves like sulphydric acid, a selenide of the metal being always formed.
254. le7lurliumt (Te) occurs in nature even more rarely than
selenium. Sometimes it is found in the free state, but more generally in combination with the heavy metals, such as gold, silver,
lead, copper, and bismuth. It is one of the few elements with
regard to which chemists have, at times, been in doubt whether




202                 THE SULPHUR GROUP.
or not they should be classed as metals. Many of its physical
properties are similar to those of the metals, and it particularly resembles the metal antimony; but it is so intimately related to sulphur and selenium in its chemical properties, its crystalline form,
and mode of occurrence in nature, that it is now almost always
studied as a member of the sulphur group.
255. Tellurium is of a silver-white color and glittering metallic lustre. It is hard and brittle, and crystallizes very easily in
rhombohedrons. It is a bad conductor of heat and electricity.
Its specific gravity is 6'2; its specific heat is 0'04737, and its
atomic weight 128. It melts at a temperature somewhat above
the melting-point of lead, and is volatile at a full red heat, the
vapor being of a yellow colour, like that of selenium. When
heated in the air, it takes fire, and burns with a greenish-blue
flame, copious fumes of tellurous acid, TeO2, being at the same
time evolved.
256. The compounds of tellurium and oxygen (tellurous acid,
TeO2, and telluric acid, TeO,) are analogous to sulphurous and
sulphuric acids. By uniting with metallic oxides, they form numerous salts, analogous to, and isomorphous with, the corresponding compounds of sulphur and selenium. So, too, the hydrogen
compound, H2Te, is analogous to sulphuretted and seleniuretted
hydrogen, in composition and properties. With the metals it
unites directly to form tellurides. There are chlorine compounds
also, TeCl and TeCl4.
257. The close relationship which subsists between sulphur and
oxygen has been already alluded to, as well as the many points
of resemblance between sulphur, selenium, and tellurium; the
student is therefore now prepared to recognize the fact that in
oxygen, sulphur, selenium, and tellurium we have another group
or family of elements, as intimately and naturally related to each
other as are the members of the chlorine group. (See ~ 152.) It
will be seen at a glance that in passing from oxygen, at one end
of the series, to tellurium, at the other, we meet with the same
progression of physical and chemical properties that was so
noticeable in passing from chlorine to iodine. The properties of
the various compounds formed by the union of the members of
the sulphur group with other elements exhibit the same kind




COMBINATION BY VOLUME.              203
of progression; that these compounds are of analogous composition has been shown in the preceding paragraphs.
With hydrogen the members of the sulphur group unite in the
proportion of two atoms of hydrogen to one atom of the other
element; thus, H120, I2,S, HS, 2Se, HlTe. This peculiar relation
to hydrogen is an important characteristic of the group.
In this sulphur group, precisely as in the chlorine group, the
relative chemical power of each element in the family is great
in proportion as its atomic weight is low (~ 153); oxygen is, as
a rule, stronger than sulphur, sulphur than selenium, and selenium
than tellurium, their atomic weights being respectively:0 =  16, S =  32, Se =  79-5(80?), Te =  128.
CHAPTER XV.
COMBINATION BY VOLUME.
258. A comparison of the formulae representing the volumetric composition of all the well-defined compound gases and vapors
which have been thus far studied, will bring into clear view some
of the general facts relating to combination by volume.
It has been established, by experiment, that the following compounds are formed by the chemical union, without condensation,
of equal volumes of the two elements which enter into each compound:Hydrogen   Chlorine _ Chlorhydric Acid  o 1   Cl  H1l
1 vol. +  1 vol.        2 vols.      or 1   35.5= 365
Hydrogen + Bromine   Bromhydric Acid      If   Br  HBr
i vol.    IvoL  -        2 vols.    -or 1   80 = 81
Hydrogen +  Iodine _  Iodohydric Acid     H    I    HI
1vol.     1 vol.        2 vols.,or    +127   128
Nitrogen +  Oxygen      Nitric Oxide      N    O    NO
1 vol.    1 vol.  -     2 vols.,or 14 + 16 =  30
It has further been demonstrated that the following compounds
of two elements contain two volumes of one element and one




204               COMBINATION BY VOLUME.
volume of the other, but that these three volumes are condensed,
during the act of combination, into two volumes:Hydrogen + Oxygen          Steam       or  2 + 0   H20
2vols.    1 vol.         2 vols.         2    16   18
Hydrogen   Sulphur    Sulphydric Acid     H2    S   H S
vols.    1 vol.         2 vols.         2   32     34
Hydrogen  Selenium    Selenhvdric Acid    H2   Se - H2Se
2 vols. +  1ol.          2 vols.      or 2   795- 815
Hydrogen  Tellurium   Tellurhydric Acid   H2  Te  H2Te
2 vols.    1 vol.       2 vols.     or 2 + 128   130
Chlorine    Oxygen  iHypochlorous Acid orC12+0   C120
2vols. +  Ivol.          2 vols.         71   16   87
Chlorine    Oxygen _ Hypochloric Acid     C1 + 02 _ C102
1 vol.    2 vols.        2 vols., or 355  32 - 675
Nitrogen    Oxygen     Nitrous Oxide   or N2 +       N20
2vols. +  Ivol.          2 vols.      r 28   16    44
Nitrogen    Oxygen _  Hyponitric Acid      N   0  _ NO2
1 vol.  +  2 vols. -     2 vols.,or 14   32 -  46
Sulphur + Oxygen _  Sulphurous Acid        S + 02   SO2
1 ol.     2 vols.        2 vols.        32   32    64
Selenium    Oxygen     Selenious Acid     Se +  2 _ SeO2
1 vol. +  2 vols.        2 vols.      or 795  321115
Tellurium    Oxygen _  Tellurous Acid     Te         TeO2
1 vol. +  2 vols. -      2 vols.     or 128  32    160
Lastly, still a third mode of combination by volume with condensation of four volumes to two has been thoroughly studied in
the case of ammonia, and has been further illustrated in the composition of anhydrous sulphuric acid:Nitrogen +llydrogen      Ammonia           N or      NH3
Ivol.  3vols.         2vols.,or      3     17
Sulphur + Oxygen       Sulphuric Acid     S   O 0    SO3
1 vol.    3 vols.        2 vols.         32   48 =  80
Throughout these tables the unit-volume is, of course, the same
for every element and compound.  What the absolute bulk of
this unit-volume may be, is not an essential point; for the relations remain the same, whatever the unit of measure. Some
chemists have thought that an advantage was gained by using
the bulk of one gramme of hydrogen at the ordinary pressure




CONDENSATON- RATIOS.                205
and temperature, viz. 11'2 litres, as the unit-volume, while others
prefer to use the litre itself as the unit.
Three condensation-ratios are exhibited in these tables. In the
first the condensation is 0; in the second it is 1, and in the third
it is 2. The typical character of the three compounds, chlorhydric acid, water, and ammonia, is also clearly brought out; each
of these bodies represents a group of compounds which obey the
same structural law. The tables also show very clearly the fact
that very unequal weights of the compounds tabulated occupy
equal spaces, under the same conditions of temperature and pressure. The space occupied by the compound molecule is, in each
case, exactly twice the unit-volume.
259. The symbols H, Cl, 0, and N represent the relative
weights of the same volume of four elements which are gaseous
at common temperatures and pressures; the symbols Br, I, S,
and Se, represent the relative unit-volume weights of four other
elements which are not gases under the ordinary atmospheric
conditions, but which can be converted into gases at a higher
temperature.  At this higher temperature their unit-volume
weights have been experimentally determined, and from these
observed volume-weights, the unit-volume weights which they
would possess at the ordinary pressure and temperature have been
deduced. The symbols of these eight elements, therefore, represent at once the combining weights and the relative weights of
equal volumes (specific gravities) of these substances in the gaseous
state. In the present state of the science, these eight symbols
are the only ones of which this can be affirmed; tellurium would
undoubtedly make a ninth, if the relative size of its combining
weight had been experimentally determined, but until this determination has been made, the symbol Te represents only the combining weight of the element, and not its equal-volume weight as
well.
The relative sizes of the combining weights of four other elements in the state of vapor, have been experimentally ascertained.
These four elements are arsenic, phosphorus, cadmium, and mercury. When we come to study these elements, we shall find that
the symbols of arsenic and phosphorus, namely, As and P, represent only the half-volume weights of these two bodies, while the




206        COMBINING WEIGHT AND VOLUME-WEIGHT.
symbols of cadmium and mercury represent the two-volume
weights of these volatile metals. Coincidence of the combining
weight and the volume-weight has been established for eight
elements; discrepancy between the combining weight and the
volume-weight has been proved for four elements; of the remaining elements, constituting more than four-fifths of the whole
number, the equal-volume weights are wholly unknown, inasmuch as these elements have never been converted into vapor
under conditions which permit the experimental determination of
the equal-volume weights of their vapors.  For example, the
symbols Na and K represent the combining weights of these two
metals; but they can be held to represent the weights of the
unit-volumes of these metals only by pure assumption, or, at best,
on the uncertain evidence of analogies, since the unit-volume
weights of these metals, when converted by intense heat into
gases, have never yet been determined. As the great majority
of the known elements cannot be volatilized, or made gaseous, by
the highest temperatures as yet at our command, under conditions
which permit the chemist to experiment with the gases produced,
it is plain that composition by weight is, in the present state of
chemistry, of far greater practical importance than composition by
volume. The symbols of all the elements represent their combining weights, as determined by ponderal analysis; the symbols
of eight elements represent also the equal-volume weights of the
substances they stand for. These eight elements, though few in
number, are nevertheless the leading elements in inorganic
chemistry.
260. The volume of the molecule of every compound gas in the
above tables is twice that occupied by the atom of hydrogen. Two
volumes of compound gas invariably result from the chemical
combination of one volume of hydrogen with one volume of chlorine, of two volumes of hydrogen with one of oxygen, of three
volumes of hydrogen with one of nitrogen, and these instances
are but types of large classes of chemical reactions. In organic
chemistry the same law holds good for a great multitude of complicated compounds of carbon; the molecule of every organic
compound in the state of vapor occupies a volume twice as large
as that occupied by an atom of hydrogen, or, in other words,




DOUBLE OR PRODUCT-VOLUME.             207
twice the unit-volume. This doubled volume is often called the
normal or product-volume of a compound gas. Since the combining weight of a compound gas or vapor occupies two unitvolumes, it is obvious that the weight of one volume, which is the
specific gravity of the gas or vapor, is deduced from the combining
weight by dividing the latter by two. The specific gravity of a
compound gas or vapor is, therefore, one-half its combining weight.
261. Molecular condition of elementary gases.-Bearing in mind
our definitions of atom and molecule (~~ 38, 39), let us inquire
what inferences concerning the molecular condition of simple
gases in a free state can be legitimately drawn from our knowledge of the molecular condition of compound gases. To give
definiteness to our conceptions, let us assume the unit-volume of
the elements to be one litre; the product-volume of a compound
will then be two litres. Two litres, the product-volume, of chlorhydric acid gas are made up of one litre of hydrogen and one
litre of chlorine, united without condensation, and each molecule
of chlorhydric acid must contain at least one atom of hydrogen
and one of chlorine. In these two litres of chlorhydric acid there
must be some definite number of molecules; the number is, of
course, indeterminable; but let us assign to it some numerical
value, say 1000, in order tc give clearness to our reasoning. One
litre of chlorhydric acid will then contain 500 molecules, and sinco
equal volumes of all gases, whether simple or compound, are
assumed to contain, under like conditions, the same numbers of
molecules (~ 39), one litre of hydrogen or of chlorine will also
contain 500 molecules." But the one litre of hydrogen and the one
litre of chlorine, which, by uniting, produced 2 litres =1000
molecules of chlorhydric acid, must each have contained 1000
atoms of hydrogen and of chlorine respectively, for each molecule
of chlorhydric acid demands an atom of hydrogen and an atom of
chlorine. The litre of hydrogen, or of chlorine, then, contains
500 molecules, but l00  atoms,-each molecule of the simple gas
being made up of two atoms of the single element, just as each
molecule of the compound gas under review is composed of two
atoms, one of hydrogen and one of chlorine. It is clear that this
train of reasoning is independent of the particular numerical
value assumed as the number of molecules in two litres of chlor



208          fMOLECULES OF ELEMENTARY GASES.
hydric acid. If, therefore, the molecule of chlorhydric acid is
represented by the formula HC1, and the diagram
[H+    --- HC1
there is good reason to assign to free hydrogen and free chlorine
the formulae HH and C1C1, and to represent the constitution of
all uncombined gases by such diagrams as
+    =      1 Ilj~ICl =                 c1CI0
Upon these models the molecular formulae of all the elements
with which we have become acquainted might readily be written.
It is only in a free state that the elementary gases and vapors are
thus conceived to exist as molecules; when they enter into combination, it is by atoms rather than by molecules. An atom of
hydrogen unites with an atom of chlorine; three atoms of hydrogen combine with one of nitrogen.
If this view of the molecular structure of free elementary gases
and vapors be correct, perfect consistency would require that no
equation should ever be written in such a manner as to represent
less than two atoms, or one molecule, of an clement in a free state
as either entering into or issuing from a chemical reaction. Thus
instead of H 2+0=1H20, N+3H=NH3, HC1 +Na=NaCI+:H,
it would be necessary to write
2HH  +  00 =  2HO2,        NN +  3HH =  2NH3,,
2HC1 +  NaNa =  2NaCl +  HH.
We have not heretofore conformed to this theoretical rule, and
do not propose to do so in the succeeding pages, and this for two
reasons: —first, because many equations, representing chemical reactions, must be multiplied by two in order to bring them into conformity with this hypothesis concerning molecular structure; the
equations are thus rendered unduly complex; secondly, because,
in undertaking to make chemical equations express the molecular
constitution of elements and compounds, as well as the equality
of the atomic weights on each side of the sign of equality, there
is imminent danger of taking the student away from the sure




PHUdPHIORUS.                    209
ground of fact Td experimental demonstration, into an uncertain region of hypotheses based only on definitions and analogies.
The symbol Na represents 23 proportional parts by weight of the
metal sodium; of the molecular symbol NaNa, the most that can
be said is, that some strong analogies justify us in assuming, for
the present, in default of any experimental evidence on the subject, that a molecule of free sodium gas, if we could get at it,
would be found to consist of two least combining parts by weight
of sodium.  We know as much, at least, of the molecular structure of sodium as we do of four-fifths of the recognized chemical
elements.  For the present, the biatomic structure of the molecule of a simple gas or vapor in the free state must take place, in
an elementary manual, as an ingenious and philosophical hypothesis, rather than as a general and indubitable fact.
CHAPTER XVI.
PHO SP1 O RUS.
262. Phosphorus occurs somewhat abundantly and very widely
diffused in nature. It is never found in the free state, but almost
always in combination with oxygen and some one of the metals.
The most abundant of its compounds is phosphate of calcium;
small quantities of this mineral are found in most rocks and soils,
and in several localities it occurs in large beds. Phosphate of
calcium is the chief mineral constituent of the bones of animals;
it contains one-fifth of its own weight of phosphorus. The proportion of phosphorus present in most of the ordinary rocks, and
in the soils which have resulted from their disintegration, is
usually very small, and phosphorus would be an exceedingly
costly substance if we were compelled to collect it directly from
this source; but it so happens that the phosphorus-compounds
are important articles of food for plants and animals, and it is
easy to obtain through their intervention the phosphorus which
was before widely diffused,. but has been by them. concentrated.
P




210                 ORDINARY PHOSPHORUS.
Growing plants seek out and collect the traces of phosphoruscompounds which exist in the soil; the herbivorous animals in
their turn consume the phosphorus which has been accumulated
by the plants, and from the bones of animals chemists and manufacturers derive the phosphorus of which they stand in need.
Like oxygen and sulphur, phosphorus occurs in several distinct
allotropic modifications. Of these, the best-known are called respectively ordinary phosphorus and red phosphorus.
263. Ordinary phosphorus, when perfectly pure, is a transparent, colorless, wax-like solid of 1'8 specific gravity, which, when
freshly cut, emits an odor like garlic, though under ordinary conditions this odor is overpowered by the odor of ozone, which, as
has been previously stated (~ 164), is developed when phosphorus
is exposed to the air. It unites with oxygen readily, even at the
ordinary temperature of the air, and with great energy at somewhat higher temperatures (above 60~); when in contact with air
it is all the while undergoing slow combustion.
Exp. 106. Thoroughly wash a piece of phosphorus by rinsing it in
successive large quantities of water; place it, for a moment, upon a
sheet of filter-paper, in order that a portion of the water adhering to
it may be removed, then lay it upon a clean porcelain capsule, and at
short intervals press against it a slip of blue litmus paper. In a very
few moments the color of the paper will be changed to red; for the products of the oxidation of phosphorus are acid, and they are formed
with great rapidity.
If the temperature of the slowly burning phosphorus be slightly
increased in any way, the mass will burst into flame and be rapidly
consumed. On account of this extreme inflammability, phosphorus
must always be kept under water; it is best also to cut it under
water, lest it become heated to the kindling-point by the warmth
of the hand or by friction against the knife. When wanted for
use, the phosphorus is taken from the water and dried by gently
pressing it between pieces of blotting-paper.
Phosphorus must always be handled with great caution; for
when once on fire, it is exceedingly difficult to extinguish it, and,
in case-it happens to burn upon the flesh, painful wounds are inflicted, which are exceedingly difficult to heal.  Whenever phosphorsusi cut Qr brokcn, care must likewise to taken that no small




FRICTION MATCHES.                 211
fragments of it fall unobserved into cracks of the table or floor,
where they might subsequently take fire.
Exip. 107.-Nip a piece of phosphorus, as large as a small pea, between two bits of wood, in such manner that a part of the phosphorus
shall project below the wood; rub the phosphorus strongly upon a sheet
of coarse paper; it will take fire at the temperature developed by the
friction.
264. On account of this easy inflammability by friction, phosphorus is extensively employed for making matches. The matter
upon the end of an ordinary friction match usually contains a
little phosphorus, together with some substance capable of supplying oxygen, such as red-lead, black oxide of manganese, saltpetre, or chlorate of potassium. The phosphorus and the oxidizing
agent are kneaded into a paste made of glue or gum, and the
wooden match-sticks, the ends of which have previously been
dipped in melted sulphur, are touched to the surface of the phosphorized paste, so that the sulphured points shall receive a coating
of it. The sulphur serves merely as a kindling material which,
as it were, passes along the fire from the phosphorus to the wood.
By rubbing the dried, coated point of the match against a rough
surface, heat enough is developed to bring about chemical action
between the phosphorus and the oxygen of the other ingredient,
combustion ensues, the sulphur becomes hot enough to take on
oxygen from the air, and finally the wood is involved in the play
of chemical force.
Exp. 108.-Put a piece of phosphorus, as big as a grain of wheat,
upon a piece of filter-paper, and sprinkle over it some lampblack or
powdered bone-black. The phosphorus will melt after a time, and will
finally take fire. As will be more fully explained hereafter, under
carbon, the porous, finely divided lampblack has the power of absorbing and condensing within its pores much oxygen from the air; heat
is developed by the act of condensation, and, at the same time, oxygen
Is brought into very intimate contact with the phosphorus, particularly
with the vapor of phosphorus which is condensed by the lampblack
together with the oxygen, so that chemical action soon results, and
ultimately fire. Both the lampblack and the paper are bad conductors
of heat; they prevent the phosphorus from losing the heat developed
by the condensation and by the slow action of oxygen.
It is remarkable that when dry phosphorus, in very thin slices, is
2




212                  PHOSPHORESCENCE.
laid upon fine feathers, wool, lint, flannel, dry wood, or other non-conducting substances, it quickly melts, and readily inflames upon the
slightest friction, heat enough being produced by the slow combustion
of the phosphorus to fuse it, if only this heat can be retained by some
bad conductor.
265. At the ordinary temperature of the air, and still more
at somewhat higher temperatures, phosphorus shines with a
greenish-white light, as may be seen by placing the phosphorus
in the dark; hence the name, phosphorus, from Greek words signifying light-bearing. This phosphorescence is seen when an
ordinary friction-match is rubbed against any surface in a dark
room. Although the phenomena of phosphorescence and of oxidation, or slow combustion, occur simultaneously when phosphorus
is exposed to the air, it does not appear that the phosphorescence
is a consequence of the oxidation; for phosphorus shines not only
in the air, but also when placed in an atmosphere of pure hydrogen, or nitrogen, or carbonic acid, or even in a vacuum, though
the light emitted by phosphorus in these inert gases is of different
appearance from that developed in presence of oxygen.
266. In warm weather phosphorus is soft and somewhat
flexible; it may then be bent without breaking, can be scratched
with the nail and cut with a knife like wax; but at 0~ it is brittle,
and exhibits a crystalline fracture when broken. It melts at
440, forming a viscid oily liquid, which boils at about 2900, and
is converted into colorless vapor. Phosphorus can readily be distilled in a retort filled with some inert gas, like hydrogen, nitrogen, or carbonic acid. The specific gravity of its vapor has been
found to be 62'1. Contrary to all our previous experience, however, the density of phosphorus is not identical with its atomic
weight, a point which will be discussed when the compounds of
phosphorus and hydrogen are treated of.
On being heated to about 230~, out of contact with air, it is
converted into red phosphorus. (See Exp. 111.) By exposure
to light, also, phosphorus undergoes a certain amount of change;
hence it is rarely seen in the perfectly colorless, transparent condition which it exhibits when recently prepared and perfectly
pure. The phosphorus of commerce is usually of a light ambercolor. When kept for some time under water, phosphorus be



SOLUTIONS OF PHOSPHORUS.               213
comes covered with a white opaque coating, which appears to be
a result of the oxidizing action of air held in solution by the
water; the surface of the phosphorus is irregularly corroded by
this dissolved oxygen, and is thus roughened and made opaque,
in much the same way that the transparency of glass is destroyed
by grinding one of its surfaces. It is noticed, for that matter,
that the water in which phosphorus is kept soon becomes strongly
acid; for it dissolves the oxygenated compounds which are produced by the action of the dissolved air upon the phosphorus.
The specific heat of solid phosphorus is 0'1788; of liquid phosphorus, 0'2045. It is a non-conductor of electricity, both in the
solid and in the liquid state.
Phosphorus is insoluble in water, but is somewhat soluble in
ether, petroleum, benzine, oil of turpentine, and other oils; it
dissolves abundantly in bisulphide of carbon, in chloride of sulphur, and in sulphide of phosphorus.
Evp. 109.-Pour into a phial of the capacity of 80 or 90 c. c., 10
or 12 c. c. of bisulphide of carbon, and throw into this liquid a bit of
phosphorus as large as a pea. Cork the phial, and shake its contents,
at intervals, until the phosphorus has dissolved. Preserve the solution
for use in subsequent experiments.
267. From the solution in chloride of sulphur and from that in
sulphide of phosphorus, crystals of phosphorus, usually in the
form either of regular octahedrons or of rhombic dodecahedrons,
can be obtained; but, owing to the slowness with which phosphorus passes from the liquid to the solid state, distinct crystals
cannot readily be prepared by the method of fusion, unless a
comparatively large quantity of phosphorus be operated upon.
268. When a solution of phosphorus in ether, or, better, in
bisulphide of carbon, is poured upon the surface of any porous
substance and left to evaporate in the air, the volatile solvent will
quickly escape, leaving the phosphorus behind in a very finely
divided condition. In proportion as a substance is more finely
divided, the greater will be the surface which it presents to the
oxygen of the air, and the more readily will it combine with this
oxygen. In the case before us, the comminuted phosphorus absorbs oxygen very rapidly, and this chemical action is attended
with the evolution of so much heat that the phosphorus will take




214                 PHOSPHORUS A POISON.
fire, if the material upon which it has been deposited is a bad
conductor of heat.
Erp. 110.-Pour some of the solution of phosphorus obtained ir
Exp. 109, upon a sheet of filter-paper, and hang the paper upon the
iron stand in such manner that the bisulphide of carbon may freely
evaporate. The paper will soon burst into flame. It will be noticed
that the paper is not completely consumed, but that a very considerable
residue of carbon remains unburned. This depends upon the fact that
the product of the combustion of the phosphorus, phosphoric acid,
quickly covers the paper with a varnish which is not only incombustible
in itself, but is quite impervious to air.
In lack of bisulphide of carbon, this experiment can be performed
with the ethereal solution of phosphorus, prepared in the manner described in Exp. 109, excepting that common ether is substituted for
the bisulphide, and that the mixture is left to digest for a day or two.
269. Ordinary phosphorus is a violent poison, a few decigrammes of it being sufficient to destroy human life.  It is the efficient
ingredient of many preparations used for poisoning rats, cockroaches, and other vermin. Phosphorus evaporates rather freely
at the ordinary temperature of the air; and the vapor has been
found to be exceedingly injurious to persons constantly exposed
to it. The makers of friction matches are subject to a horrible
wasting disease, one of the symptoms of which is the destruction
of the bones of the jaws.
270. When phosphorus burns with flame in free air, two atoms
of it unite with five atoms of oxygen, and there is formed the
compound of P20; this highest oxide of phosphorus is called
phosphoric acid. This compound occurs in bones, and from it
phosphorus is prepared. Bone-earth, that portion of bones which
remains after all the organic matter has been burnt off in the
fire, consists mainly of triphosphate of calcium, Ca3P208.  IIn
order to obtain phosphorus from bone-earth, the calcium and the
oxygen must both be removed; the calcium is removed by means
of sulphuric acid, the oxygen by means of hot charcoal. Bones
are burnt to a white ash (calcined, as the term is), then finely
powdered and mixed with a quantity of dilute sulphuric acid.
The sulphuric acid removes two of the atoms of calcium  and
fomns sulphate of calcium, while there remains monophosphate of




MANUFACTURE OF JROSPHORUS.           215
calcium (superphosphate of lime) in accordance with the following reaction:CaP208 + 2HS04 = CaH4P208 + 2CaS0,
It will be remembered that one atom of calcium replaces two
atoms of hydrogen. (See p. 89.)
The solution of monophosphate of calcium is then filtered off
from the insoluble sulphate of calcium, and evaporated to the
consistence of syrup; the syrup is mixed with powdered charcoal,
and the mixture dried at a dull red heat; by this means a quantity of water is expelled from the monophosphate:CaH4P208 =  CaP206 + 2H20.
The porous dry mixture is finally placed in retorts of fireclay
and intensely ignited. At high temperatures, charcoal is a
powerful deoxidizing agent; it takes away oxygen from the
phosphate of calcium, and forms carbonic oxide, which goes off
as a gas; phosphorus is thus set free, and distilling over into an
appropriate receiver, is condensed under cold water, a quantity
of triphosphate of calcium is at the same time reproduced and
remains in the retort:
3CaP206 +  10C =  10C0  +  4P + Ca3P208.
If a quantity of sand (silicic acid) be added to the mixture of
charcoal and monophosphate, the whole of the phosphorus can
be expelled, —the phosphate of calcium, which would otherwise
escape decomposition, being entirely converted into silicate of
calcium,
2CaP206 + 2SiO2 +  10C =  10CO +  4P + 2CaSiO3.
Another proposed method is to pass chlorhydric acid gas over a
mixture of bone phosphate and charcoal, maintained at a red
heat, in a cylinder of fireclay. By this means all of the phosphorus is set free and chloride of calcium remains:Ca3P08 + 6HC = 3Ca + 8C ~ 6HC10= 3CaC, + 6CO + 6H + 2P.
The crude phosphorus thus obtained is remelted, and purified by
filtration, redistillation, and by chemical treatment with a mixture
of bichromate of potassium and sulphuric acid, which oxidizes
the principal contaminations. The purified phosphorus is finally




216                    RED PHOSPHORUS.
remelted and cast into the sticks or cakes in which it is found in
commerce.
271. Red Phosphorus.-This remarkable allotropic modification of phosphorus is a body as unlike ordinary phosphorus in
most respects as could well be conceived. It is of a scarlet-red
color, has neither odor nor taste, is not poisonous so far as is
known, is not phosphorescent, does not take fire at ordinary temperatures, is insoluble in bisulphide of carbon, and in general
behaves altogether differently from the ordinary modification.
Yet it is no difficult matter to change one of these modifications
into the other. For example, if red phosphorus be heated to
about 260~, in an atmosphere of nitrogen, or other inert gas, it
will pass into the condition of ordinary phosphorus without undergoing any alteration of weight, or, in other words, without absorbing or disengaging anything.
Exp. 111.-In a narrow glass tube, No. 6, about 30 c.m. long and
closed at one end, place a quantity of red phosphorus as large as a small
pea; heat the phosphorus gently over the gas-lamp, and note that a
sublimate of a light-colored substance is quickly deposited upon the
cold walls of the tube a short distance above the heated portion. This
light-colored sublimate is ordinary phosphorus, as may be shown by
cutting off the tube just below the sublimate, after the glass has been
allowed to cool, and then scratching the coating with a piece of wire;
the coating will take fire. The air in the narrow tube employed is deprived of its oxygen by the combustion of a small portion of the phosphorus at the moment of its transformation from the red to the ordinary
condition; the remaining phosphorus is thus enveloped in nitrogen and
so protected from further loss.
272. Red phosphorus is itself neither volatile nor inflammable;
it neither rises as vapor nor inflames at temperatures lower than
2600, the point at which it changes into ordinary phosphorus; at
2500 it suffers no alteration. As compared with ordinary phosphorus, it may be said that red phosphorus can be handled
without danger, and that it may be kept in bottles without
special precautions, since it is not liable to take fire by moderate
friction; but by powerful friction heat enough may be evolved to
convert it into ordinary phosphorus, and if it be even moderately
heated, by friction, or in any other way, in contact with oxidizing
agents, it instantly bursts into flame




CRYSTALLIZED RED PHOSPHORUS.            217
E1xp. 112.-In order to observe the comparative difficulty of inflaming red phosphorus, lay an inverted cover of a porcelain crucible upon
an iron triangle upon the lamp-stand; place upon the cover, which may
be 15 c.m. wide, a small bit of ordinary phosphorus, and, at a distance
of 12 c.m., the same quantity of red phosphorus; heat the cover gently
and gradually over the gas-lamp. The ordinary phosphorus will soon
inflame and burn away; but a considerable space of time i-.1l elapse
I efore the red phosphorus takes fire.
By operating in vessels filled with nitrogen, or some other gas which
had no chemical action upon phosphorus, the precise temperature at
which the red phosphorus ceases to exist can be noted, and the ordinary phosphorus obtained from it can be distilled over and collected.
273. Red phosphorus has been obtained in crystals by dissolving
common phosphorus in melted lead, and subjecting the fluid mass
to a high temperature for several hours in closed tubes. When the
lead cools, the phosphorus separates in thin crystals, which have
a metallic lustre and a black color; the crystals, however, transmit a yellowish-red light, and the thinnest of them appear, not
black, but red. These crystals of red phosphorus are generally
enveloped in the lead; but the lead may be mostly dissolved away
by dilute nitric acid, and the phosphorus crystals may thus be
obtained in a condition of comparative purity. They are not
affected by exposure to the air. These crystals are seen under
the microscope to be rhombohedrons; so that phosphorus, like the
succeeding members of the family of elements to which it belongs,
is dimorphous, presenting forms both of the monometric and hexagonal systems.
The red variety of phosphorus has been not inaptly called
metallic phosphorus, crystallized in the crystals just described, and
amorphous in the usual form of red phosphorus. The crystallized
metallic phosphorus is less volatile, and has a higher specific
gravity than the amorphous. The power of the so-called metallic
phosphorus to conduct electricity is small, if compared with that
of the common metals, but it is very much greater than the conducting-power of colorless phosphorus, for this latter substance is
generally classed with the insulators.
The specific gravity of amorphous red phosphorus is 2'14; its
specific heat is 01698. When dry, it undergoes no change at the
ordinary temperature of the air; but in moist air it oxidizes very




218           PREPARATION OF RED PHOSPHORUS.
slowly. It is easily soluble in nitric acid, which oxidizes it; and
since it is much more readily dissolved than ordinary phosphorus,
the latter can be purified from any contamination of red phosphorus, by digesting it at a gentle heat in dilute nitric acid.
274. Amorphous red phosphorus is prepared by maintaining
ordinary phosphorus, for some time, at a temperature of 2300 to
235~, either under water in an air-tight vessel, or in an atmosphere of some gas which has no chemical action upon phosphorus.
It is manufactured upon the large scale by heating ordinary
phosphorus in a cast-iron vessel provided with a gas delivery-tube
dipping into mercury outside the vessel, in such manner that, while
the expanded air and some escaping vapors of phosphorus can
pass out, no air can enter the vessel. About 200 kilos. of phosphorus are taken for a single charge; this quantity of phosphorus
is maintained during ten days or more, as nearly as may be, at
the temperature of 240~, care being taken that the heat shall, at
no time, much exceed this limit. Under these conditions, the
ordinary phosphorus slowly changes into the red variety. After
the phosphorus has been exposed during the time which the previous experience of the manufacturer has shown to be most advantageous, the apparatus is allowed to become cold, and the transmuted phosphorus is found adhering to the sides of the vessel, in
the shape of a hard, brittle, brick-colored coating, which can be
removed by means of hammer and chisel, after covering it with
water. It is ground to powder under water, and any particles
of ordinary phosphorus which have escaped change are removed
from it by means of bisulphide of carbon, or by a solution of
caustic soda, which dissolves ordinary phosphorus without acting
upon the red modification.
275. Red phosphorus is employed, to a certain extent, as an
adjunct to the so-called safety-matches. Such matches contain
no phosphorus in themselves, and will not take fire readily by
friction upon an ordinary rough surface, though they burst into
flame at once when rubbed upon a surface specially prepared
with red phosphorus. f The matter upon the tips of safety-matches
is usually a mixture of chlorate of potassium and sulphide of
antimony, made into a paste by means of glue; the surface upon
which the match is to be rubbed is composed of red phosphorus,




PHOSPHORUS A REDUCING AGENT.           210
black oxide of manganese, and glue. In favor of the use of red
phosphorus for matches are the facts, that, unlike ordinary phosphorus, it is not deleterious to the workmen who have to deal
with it, and that it is far less liable to be set on fire by accidental
friction. For these reasons, the manufacture of safety-matches
has been encouraged by the governments of several European
countries, and such matches are now much used in France and
upon other partcs of the continent, though they are manifestly less
convenient, in several respects, than the ordinary matches, which
can be ignited by friction upon any rough surface.
276. Phosphorus combines readily with many other elements
besides oxygen. The ordinary modification of phosphorus combines violently with sulphur at temperatures near the meltingpoint of sulphur, the act of combination being attended with vivid
combustion and loud explosion. Red phosphorus, on the other
hand, does not combine with sulphur at temperatures lower than
230~, and the combustion, though rapid, is not explosive. With
chlorine, bromine, and iodine, ordinary phosphorus unites directly
at the ordinary temperature of the air, the combination being
rapid and attended with inflammation.  Red phosphorus also
unites with chlorine, bromine, and iodine at the ordinary temperature; and much heat is evolved during the act of combination,
though the amount of heat is usually insufficient to produce
ignition.
Phosphorus unites directly with most of the metals also; and
several of the compounds thus formed closely resemble the socalled alloys, or compounds of one metal with another. With
hydrogen it forms several interesting compounds, which will be
described directly. From the remarkable facility with which it
combines with oxygen (~~ 263, 264), it follows necessarily that
phosphorus is a powerful reducing agent. Many oxygen compounds can be decomposed by means of it. When immersed in
the vapor of anhydrous sulphuric acid, phosphorus takes fire after
a time, and combines with the oxygen of the acid, while sulphur
is deposited. Monohydrated sulphuric acid is reduced to sulphurous acid, while phosphorous acid is formed:3H2SO, +  2P =  2H3PO,  + 3SO2,




220              PHOSPHORUS AND HYDROGEN.
A solution of sulphurous acid, on being heated with phosphorus,
yields phosphorous and sulphydric acids, as follows:SO2 +  4H20 +  2P =  2H3PO3 +  H2S.
When gently heated with chlorate or with nitrate of potassium,
or with other highly oxygenated bodies, like the peroxides of lead
and manganese, phosphorus combines with their oxygen so rapidly
that an explosion ensues; heat enough to bring about the reaction
call be developed by gentle friction, as when the phosphorus and
the other ingredient are rubbed together upon some hard surface.
(Compare ~ 264.)
Exp. 113.-Provide a bit of ordinary phosphorus, as large as a pin's
head, also an equal quantity of red phosphorus; add to each of these
portions enough finely powdered chlorate of potassium to cover the
phosphorus; fold up each of the mixtures tightly and separately in a
small piece of writing-paper; place the parcels, one after the other,
upon an anvil and strike them sharply with a hammer. They will explode with violence.
277. Comnpounds of Phosphorus and of Hydrogen.-There are
three compounds of phosphorus and hydrogen, one gaseous, PH,,
one liquid, PH,, and one solid, PH, at ordinary temperatures.
The gaseous compound, or, rather, the gaseous compound charged
with the vapor of the liquid compound, is somewhat interesting,
from the fact that it takes fire spontaneously immediately on
coming into contact with the air.
Exp. 114.-In a thin-bottomed flask of about 140 c. c. capacity, put
1 gramme of phosphorus and 115 c. c. of potash-lye of 1'27 specific
gravity, obtained by dissolving 40 grms. of hydrate of potassium in
110 c. c. of water. Pour two or three drops of ether upon the liquid
in the neck of the flask, then close the flask with a cork carrying a long
gas-delivery-tube of glass No. 5. Place the flask over the gas-lamp,
upon the wire-gauze ring of the iron stand, and immerse the end of
the delivery-tube in the water-pan, then gently heat the flask. The
ether is added to the contents of the flask in order that the last traces
of air may be expelled from the flask by the vapor which arises from
this highly volatile liquid so soon as it is warmed.
As the potash-lye becomes hot, small bubbles of gas will be seen to
arise from the surface of the phosphorus, and in a short time large
bubbles of gas will escape from the delivery-tube; each of these bubbles, as it comes in contact with the air at the surface of the water,




PREPARATION OF PHOSPHURETTED HYDROGEN.         221
will spontaneously burst into flame, and burn with a vivid light and
the formation of beautiful rings of white smoke, if the air be not disturbed by draughts. In burning, the phosphuretted hydrogen is converted into phosphoric acid and water, or, rather, into hydrated phosphoric acid; and of this product the white smoke is, of course, composed.
2PH3 + 80 = H6P208.
Exp. 115.-Place a small inverted bottle full of water over the end
of the delivery-tube from which the phosphuretted hydrogen is escaping, as in Exp. 114, and collect 50 or 100 c. c. of this gas. By single
bubbles pass the gas thus collected into a litre bottle half-full of oxygen
standing inverted upon a shelf in the water-pan. The phosphuretted
hydrogen will burn much more vividly in oxygen than in the air. In
case, however, several successive bubbles of the gas should fail to inflame on coming into contact with the oxygen, the experiment must be
interrupted and the oxygen thrown away, for the introduction of another bubble of phosphuretted hydrogen into this explosive mixture
might set fire to it and so shatter the bottle.
278. The reaction which occurs during the preparation of
phosphuretted hydrogen is chiefly between water and phosphorus.
Phosphorus and water by themselves do not react upon each other,
but when in presence of powerful bases, like soda, potash, lime,
or baryta, water is decomposed by phosphorus with formation both
of oxygenated and hydrogenized phosphorus compounds:3(K20,H120) +  8P +  6H20 =  2PH3 +  3(K20,2H20,P20).
HIypophosphite of
Potassium.
Another method of obtaining phosphuretted hydrogen is by
decomposing phosphide of calcium with water.
Exp. 116.-Prepare a number of small balls or sticks of quick lime
by moulding moistened slaked lime into these forms and then drying
and calcining the product. Select a tube of hard glass, No. 2, close it
at one end, and place two or three pieces of phosphorus as large as
peas at the closed end; fill the tube with the pellets of quick lime,
and put it in a sheet-iron trough above a wire-gauze gas-lamp, in the
manner depicted in Fig. 8. To prevent the melted phosphorus from
flowing against the quick lime, an iron nail may be laid beneath that
part of the trough which is farthest from the phosphorus. Heat to
redness the portion of the tube which contains the lime, and then cause
the vapor of phosphorus to pass over it by cautiously heating the closed
end of the tube with an ordinary gas-lamp. To ensure the success of




222        PROPERTIES OF PHOSPHURETTED HYDROGEN.
this experiment, that portion of the tube which contains the phosphorus must be heated so slowly that none of the phosphorus can
escape uncombined through the lime. After the phosphorus has all
been driven forward from the closed end of the tube, the open end of
the tube should be stopped with a cork and the lamps should be extinguished; the tube is then left at rest until it has become cold.
When a piece of the impure phosphide of calcium thus obtained is
thrown into water, it slowly decomposes with formation of hypophosphite of calcium and disengagement of phosphuretted hydrogen; the
bubbles of gas take fire as they reach the surface of the water.
279. Besides the spontaneously inflammable gas, there is another
variety of phosphuretted hydrogen which does not take fire of
itself in the air. It can be prepared in various ways-for example,
by heating hypophosphorous or phosphorous acids (~~ 286, 287),
these acids being resolved by heat into phosphoric acid and phosphuretted hydrogen:Hypophosphorous Empirical: 2HPO2 = PH, + H3PO4.
Acid.      Dualistic: 6H20,2P20 = 2PH3 + 3H20,P20.
Phosphorous   I Empirical: 4HPO, = PH, + 3HTPO,.
Acid.     ) Dualistic: 4(3H20,P203) = 2PH3+ 3(3H20,P20,).
The non-inflammable gas is regarded as pure phosphuretted hydrogen, the property of spontaneously inflaming possessed by the
other variety being supposed to depend upon the presence of
minute portions of some foreign substance; the vapor of liquid
phosphuretted hydrogen, PH,I, produces this effect; on adding
to the non-inflammable gas so small a quantity as 1o0o0th of its
bulk of nitric oxide, it acquires the property of inflaming spontaneously.
Pure phosphuretted hydrogen gas is colorless and highly inflammable; its odor is fetid, and has been compared to that of
tainted fish; it is slightly soluble in water, and can be liquefied,
but has not yet been solidified. Neither the gas nor its solutions
have any action on red or blue litmus. It is a powerful deoxidizing agent, and is, in general, easily decomposed. Most of the
metals, when heated in the gas, combine with its phosphorus and
liberate its hydrogen, just as we have seen the metal potassium
set free hydrogen from ammonia. (See p. 77.) This ready decomposition of the gas by hot metals is the basis of the method of
determining its composition by weight.




ANALYSIS OF PHOSPHURETTED HYDROGEN.       223
280. Phosphuretted hydrogen is resolved into its two elements,
and the proportional weights of the elements which enter into
its composition are simultaneously determined, by the following
process: —The gas is passed through a hard-glass tube (A,
Fig. 45), filled with copper turnings and heated to redness; the
Fig. 45.
copper retains all the phosphorus, and the hydrogen becomes free.
This last gas is carried forward through a second tube, B, filled
with oxide of copper heated to redness; the hydrogen combines
with the oxygen of the oxide of copper, and the steam thus
formed is condensed and absorbed in a third tube, C, filled with
pumice-stone soaked in sulphuric acid. (Appendix, ~ 15.) The
tubes A and (7 are weighed both before and after the experiment, and the augmentation of weight gives the phosphorus in
A and the water in C; from the weight of the water is calculated the weight of the hydrogen required to produce it. Care
must be taken that'the tube A be heated so moderately as not
to distort it, and that nothing be added to its weight by depositions from the lamp-flames used to heat it. It is also necessary
to fill the tubes with nitrogen gas before beginning the actual
analysis, and to sweep them out with nitrogen at the end. This
operation is easily performed by the aid of a small gas-holder
full of nitrogen. It has thus been experimentally proved that
any given weight of phosphuretted hydrogen contains 8'57 per
cent. of hydrogen and 91-43 per cent. of phosphorus. Now il
has been determined, as the result of many experiments and o,
a careful collation of the formulae of all known compounds of
phosphorus, that the least proportional weight of this element
which enters into combination is 31, that of hydrogen being I
The proportion,    91'43: 8'57 = 31:x,




224       COMPOSITION OF PHOSPHURETTED IYDROGEN.
gives as the value of x, 2'905. The nearest whole number is 3;
and the discrepancy may be attributed to defects of the analytical process, always specially to be feared in cases like the present, where the quantity of one ingredient is many times as large
as that of the other. A loss of matter, or error in weighing,
which would amount to only 1 per cent. of 90 centigrammes,
would cause an error of more than 11 per cent. on 8 centigrammes.  The analysis clearly points to the formula PHI3 as
representing the composition of phosphuretted hydrogen, inasmuch as for every 31 parts by weight of phosphorus, the gas
contains three parts by weight of hydrogen. This result is partially corroborated by volumetric analysis. If the hydrogen liberated from any measured quantity of phosphuretted hydrogen by
passing the gas through a tube filled with hot metal, be accurately measured, it will be found that, for every two volumes of
the compound gas, three volumes of hydrogen are set free.
Thus far the composition of pllosphuretted hydrogen has
seemed to be completely analogous to that of ammonia gas; but
at this point the analogy fails.  In ammonia, three parts by
weight of hydrogen are combined with fourteen of nitrogen, and
three volumes of hydrogen are united with one volume of nitrogen to form two volumes of the compound gas. If the parallelism  between NHIT  and PH3 were perfect, one volume of
phosphorus-vapor ought to be united with the three volumes of
hydrogen which two volumes of phosphuretted hydrogen invariably contain. The densities of phosphorus-vapor and of phosphuretted hydrogen, as experimentally determined, prove that
this is not the case. The unit-volume being that volume of
hydrogen which weighs 1,
From the weight of 2 unit-volumes of PH3 (sp. gr. = 17'09).. 34'18
Subtract the weight of 3 unit-volumes of hydrogen.... 3'00
And there remains for the weight of the phosphorus-vapor.. 31'18
The specific gravity, or relative weight of one unit-volume of
phosphorus-vapor, is 62'1, as has been already mentioned. Two
volumes of phosphuretted hydrogen, therefore, contain, not one
volume, but only half a volume of phosphorus-vapor. The atom
of phosphorus weighing 31, combines with the same quantity of




AMMONIA AND PHOSPHURETTED HYDROGEN.       225
hydrogen by weight as the atom of nitrogen weighing 14; but
the volume of the phosphorus atom is only one-half the volume
of the nitrogen atom. The combining weights and the unitvolume weights of all the elements previously studied have been
identical; but the combining weight of phosphorus must be
doubled in order to bring it into coincidence with its unit-volume
weight. The volumetric and the ponderal composition of phosphuretted hydrogen are both exhibited in the annexed diagram:
281. This difference between ammonia and phos-    1
phuretted hydrogen is completely outweighed by the   H        31 /
essential likeness in com-        + P           PH3 34
position of these two gases.
and by the other striking
analogies which exist be-   H
tween them.  When one
or more of the hydrogen atoms in phosphuretted hydrogen are
replaced by certain groups of elements, which in organic chemistry play the part of elements, compounds are obtained which,
like ammonia, neutralize acids and are strongly alkaline. Phosphuretted hydrogen itself combines with certain of the acids in
definite proportions. With bromhydric and iodohydric acids, for
example, it forms crystalline compounds whose composition is
represented by the formula- PH4Br and PH4I, —formulse which
are evidently comparable with NH4Br and NHI4I.
282. Liquid Phosphuretted Hydrogen (PI2) may be obtained
by passing the spontaneously inflammable gaseous compound
obtained in Exp. 114, through a U-tube surrounded by a mixture of ice and salt. Under these conditions, the vapor of the
liquid compound, which was diffused in the gas, condenses and
separates. Liquid phosphuretted hydrogen is colorless, has a high
refracting-power, and is not miscible with water. It does not
solidify at — 20~; when heated to 30~0 or 400 it decomposes. It
is exceedingly inflammable, and bursts into flame when brought
in contact with the air; when a small quantity of its vapor is
mingled with combustible gases, such as carbonic oxide, hydrogen, or carburetted hydrogen, these gases acquire the property
Q




226              PHOSPHORUS AND OXYGEN.
of inflaming spontaneously. When exposed to sunlight, it is
resolved into gaseous and solid phosphuretted hydrogen:5PH, =  P2H  +  3PH,.
283. Solid Phosphuretted Hydrogen (P,H?) is formed by exposing liquid phosphuretted hydrogen to sunshine, or by acting
upon the liquid with chlorhydric acid, or by dissolving phosphide
of calcium in strong chlorhydric acid. It is a compound insoluble in water or alcohol, but soluble in warm potash-lye with
liberation of gaseous phosphuretted hydrogen. It takes fire at
about 1500, and is of a yellow color, but becomes red when
exposed to light.
284. Compounds of Phosphorus and of Oxygen.-Phosphorus
unites with oxygen in four different proportions, as follows:Oxide of Phosphorus, P40.
HIypophosphorous Acid, P20.
Phosphorous Acid, P2O,.
Phosphoric Acid, P205.
All of these compounds exhibit a more or less distinct acid character, especially when combined with water, and the one containing most oxygen, phosphoric acid, is a very important acid.
285. Oxide of Phosphorus (P40).-When ordinary phosphorus
is burned in a confined volume of air or oxygen, insufficient for
its complete combustion, there will be found mixed with the
unconsumed phosphorus, after the chemical action has ceased,
a certain quantity of a red powder, which is the oxide of phosphorus now in question.
Exp. 117.-Repeat Exp. 13, and examine the red mass which remains in the porcelain capsule after it has been sunk in the water-pan
and thoroughly cooled.
Since the red oxide of phosphorus is insoluble in bisulphide of carbon, it can readily be obtained in a state of purity by dissolving in this
liquid the free phosphorus with which it is contaminated.
Although the red oxide is not spontaneously inflammable by
itself, a mixture of it with free phosphorus, such as the residue
from the preparation of nitrogen (Exp. 13), takes fire with great
ease, being even more readily inflammable than phosphorus alone.
Such.esidues must be handled with special care.




RED OXIDE OF PHOSPHORUS.             227
Red oxide of phosphorus can be obtained in larger quantities
by bringing a stream of oxygen gas into contact with phosphorus
melted under hot water.
lExp. 118.-Place about a cubic centimetre of ordinary phosphorus
in the bottom of a conical test-glass, or wine-glass, and pour upon it
hot water enough to half fill the glass; the phosphorus will melt, but
cannot burn, since the water protects it from contact with the air, and
since phosphorus by itself is incapable of decomposing water. By
means of a narrow gas-delivery-tube of glass, conduct a slow stream of
oxygen from a gas-holder to the bottom of the test-glass, so that the
oxygen shall come into immediate contact with and bubble through
the melted phosphorus. The phosphorus will burn with a vivid light
beneath the water; red oxide of phosphorus will be formed, and will
float about in the water, from which it may be separated by filtration.
In the lack of oxygen, air may be forced down upon the phosphorus;
even the impure air blown from the mouth will answer; but with air
the reaction is less intense than with oxygen; hence, when it is employed, the experiment had better be performed in a dark room.
Oxide of phosphorus has neither taste nor smell. On being
heated to 350~ to 4000, it splits up into phosphoric acid and free
phosphorus, the latter, of course, taking fire in case oxygen be
present.
286. Hypophosphorous Acid (H1P,O4= 2H3PO,).-This compound has usually been classed among the oxides of phosphorus,
on the supposition that it might be possible to obtain from it an
anhydrous oxide, of the composition P20; the oxide in question
has, however, never yet been obtained.
When ordinary phosphorus is boiled in a solution of caustic
potash, soda, lime, or baryta, water is decomposed, a compound
of phosphorus and hydrogen (~ 278) is formed, and a hypophosphite of the alkali employed remains in solution, from which it
may be separated in crystals by cautious evaporation. If baryta
be employed, the reaction may be formulated as follows:Empirical: 3BaH20,  + 8P + 61120 = -2PH, + 3BaH4P204.
Dualistic: 3(BaO,H20) + 8P + 6H20 = 2PH3 + 3(BaO,2H20,P20).
By cautiously adding sulphuric acid to the solution of the barium
salt, sulphate of barium is precipitated and hypophosphorous acid
remains in solution:
BaO,2H20,P20 +  H20,SO, =  BaO,SO, +  320,O,P20.




228                   HYPOPIOSPHITES.
By evaporating the aqueous solution, after filtration, hypophosphorous acid is left as a viscid, uncrystallizable, acid liquid, which,
on being strongly heated, splits up into phosphoric acid and
phosphuretted hydrogen.  It unites with oxygen readily, and is
consequently a powerful reducing agent.  Sulphuric acid, for
example, is reduced by it, with evolution of sulphurous acid and
separation of sulphur.
The hypophosphites are, for the most part, crystallizable salts,
soluble in water and often in alcohol also; they can usually be
preserved in dry air. Several of them have recently been somewhat extensively employed as medicaments.
287. Phosphorous Acid (P20,).-This acid is a product of the
slow combustion of phosphorus.
When phosphorus is gently heated in a very slow current of perfectly dry air, it takes on oxygen enough to form phosphorous acid,
which, being volatile, condenses upon the cold walls of the tube beyond
the phosphorus as a bulky white sublimate. By conducting the operation in a tube drawn out to a fine point at one end and almost completely closed at the other by a perforated cork carrying a narrow
tube, and carefully regulating the supply of air which is admitted into
the tube, so that just enough oxygen to form phosphorous acid, and no
more, shall come in contact with the phosphorus, a tolerably pure product can be obtained. For purposes of illustration, however, a simpler
arrangement of the apparatus may be employed, as in the following
experiment:Exp. 119.-Place a bit of phosphorus, as big as a pea, in the middle
of a piece of glass tubing, No. 2, about 30 c.m. long, and open at both
extremities; gently heat the phosphorus ulntil it takes fire, and then
extinguish the lamp. So long as the tube is held in a horizontal
position, the combustion will be so feeble and imperfect that some red
oxide of phosphorus will be formed as well as phosphorous acid. On
the other hand, if one end of the tube be inclined upwards, so that the
products of combustion can pass off and make way for the entrance of
fresh air, the combustion will become more vivid, and there will be
produced a quantity of the highest oxide of phosphorus, phosphoric
acid. If the tube were held perpendicularly, the draught of air, passing through it as through a chimney, would be so powerful that all
the phosphorus would be burned completely to phosphoric acid.
It is evident, from the foregoing, that if it were only possible to find
out the precise angle at which the tube should be inclined, and, at the




PHOSPHOROUS ACID.                  229
same time, to provide means for continually maintaining a suitable tenlperature within the tube, the phosphorus might all be converted into
pure phosphorous acid, instead of the various and mixed products
which are actually obtained.
288. Hydrated phosphorous acid, HPO3, or 3H1O,P20,, is
readily obtained, though in an impure condition, by exposing
sticks of phosphorus to moist air.
RExp. 120.-Select a piece of glass tubing, the diameter of which is
so much greater than that of an ordinary stick of phosphorus, that the
latter can readily be slipped into it; from this tubing prepare three or
four short tubes 3 or 4 c.m. long, open above and below, but drawn in
at the bottom to such an extent that a stick of phosphorus placed in
the upper part of the tube cannot pass the narrowed portion and fall
out of the tube. In each of these short tubes put a stick of phosphorus, and place them all in a glass funnel which rests upon a bottle
standing in a soup plate full of water; over the funnel and bottle place
a tall tubulated bell-jar, from which the stopper has been removed,
and allow the apparatus to stand at rest during several days in a cool
place where no damage can be done in case the phosphorus take fire.
Under these conditions, the phosphorus will slowly oxidize and
waste away (if time enough be allowed it will completely disappear),
and the mixture of phosphorous and phosphoric acids which is formed
will flow down through the tube of the funnel into the bottle beneath.
The mixture thus obtained is often technically termed.phosphatic acid.
The object of the glass tubes employed to envelope the sticks of
phosphorus is, to keep the several pieces of phosphorus from touching
one another. If two or three pieces of phosphorus were to be left in
contact, in the air, the heat generated during the oxidation of each
would be added to that derived from the others, and after a time the
mass would become hot enough to take fire spontaneously. But when
each stick of phosphorus is placed within a g-lass tube, the heat generated by its oxidation passes off harmlessly, and a dangerous accumulation of heat is very much less likely to occur than if no such system
of isolation were resorted to.
289. The fact that a collection of fragments of phosphorus is
thus liable to take fire, so well illustrates the theory of spontaneous combustion in general, and the precautionary measures
taken in the foregoing experiment to prevent the ignition of the
phosphorus point so clearly to the methods which must often be
resorted to in order to prevent the spontaneous inflammation of




230               SPONTANEOUS COMBUSTION.
many highly combustible substances, that a few words may here
be appropriately devoted to this important practical subject.
As a rule, all easily oxidizable substances, when finely -divided
and thrown into heaps, are liable to take fire spontaneously in
the air. Many oils, for example, particularly the so-called drying oils, absorb oxygen from the air and enter into combination
with it. Wherever chemical combination occurs, heat is developed, and in case the oil be poured upon some porous substance
which is both combustible and a non-conductor of heat, like wool
or cotton, paper or cloth, the heat developed during the oxidation of the oil may very readily accumulate to the extent
necessary to produce inflammation. To prevent this catastrophe,
the heap of greasy wool or other matter should be broken up as
soon as warmth is perceived in it, and its particles should be
scattered about so that air may have free access to them; the
heat will then pass off harmlessly from each of these particles
as fast as it is generated.
This process of subdivision will prove an effectual protection
if the subdivision be carried far enough; but it is a fact not to
be lost sight of, that very small parcels of some substances (a
hank of oiled twine, for example, or a handful of greasy' rags)
may take fire when all the conditions are favourable; and it is
a matter of the first importance that all such matters should be
kept in places where no harm can be done in case they inflame.
A still more familiar instance of the accumulation of heat
during chemical action occurs in the ordinary process of hay-.making, as when a cock of half-cured hay is left unopened for
any length of time; the green hay combines with oxygen from
the air, fermentation sets in, and heat is, of course, evolved; but
when the hay is scattered about the field, this heat passes off into
the air as fast as it is generated, and we cannot perceive it. On
the other hand, if, instead of the usual small hay-cocks, the
farmer were to throw a large quantity of new-mown hay into
one great stack, this stack would undoubtedly take fire if left to
itself.
In large heaps of many kinds of bituminous coal also, strong
chemical action is induced under very various conditions as regards temperature, moisture, and mechanical subdivision, and




PROPERTIES OF PHOSPHOROUS ACID.           231
the heat evolved becomes at last intense enough to kindle the
coal. Protection from the weather, exclusion of moisture, free
ventilation, and the avoiding of too large heaps are the most
effectual preventives in this case.
290. Anhydrous phosphorous acid is a white amorphous substance, which rapidly absorbs water from the air, and when
sprinkled with water, dissolves rapidly with a hissing noise; it is
volatile, and may easily be driven from one place to another, in
a tube filled with nitrogen (see ~ 287), by applying a gentle heat.
Being a product of the incomplete combustion of phosphorus, it
is necessarily combustible; when heated in the air, it undergoes
vivid combustion.
Hydrated phosphorous acid is obtained in the form of rectangular prisms, when the aqueous solution is evaporated at
temperatures not exceeding 200~. The crystals are deliquescent,
and they gradually absorb oxygen from the air; when strongly
heated, they are decomposed into phosphoric acid, and phosphuretted hydrogen, and at the same time take fire. The aqueous
solution of phosphorous acid exhibits a strong acid reaction; by
absorbing oxygen from the air, it is converted into phosphoric
acid, quickly in case the solution is dilute, but slowly if it be
concentrated. It is a powerful reducing agent; when heated
with sulphurous acid, it yields phosphoric and sulphydric acids.
Though a very weak acid, it forms salts by combining with those
metallic oxides upon which it exerts no reducing action; the
phosphites of the alkalies proper are easily soluble in water, but
the phosphites of calcium and barium can only be dissolved with
difficulty. These salts are more stable than the hypophosphites,
but are all decomposed by heat.
291. Phosphoric Acid (P20,).-As has been already stated,
this highest oxide of phosphorus is the product of the rapid
combustion of phosphorus in an excess of air or oxygen.
Exp. 121.-Dry thoroughly a large porcelain plate, a small porcelain capsule, and a wide-mouthed bottle of two litres capacity, by
warming them at a fire; place the capsule upon the plate and put in
the capsule a bit of dry phosphorus of the weight of half a gramme
or thereabouts; light the phosphorus, and cover it at once with the




232                   PHOSPHORIC ACID.
inverted bottle. The phosphoric acid, formed by the combustion of
the phosphorus, will be deposited as a white powder, like flakes of
snow, upon the sides of the bottle, and much of it will fall down upon
the plate below.
The apparatus employed in this experiment can readily be arranged
in such manner that fresh portions of phosphorus and of air can be
introduced into the bottle as fast as occasion may require; the process
will then be continuous, and any desired quantity of phosphoric acid
may be prepared by means of it.
The flocculent, amorphous, odorless powder thus obtained
unites with water with remarkable facility; if it be left in the
air for a few minutes, it deliquesces completely; and upon being
thrown into water it dissolves, with a hissing noise and development of much heat. In order to preserve it, it must be placed
in a dry tube, and the tube closed by sealing it in the lamp.
When touched to the moist tongue, it burns as if it were red-hot
metal. On account of this strong affinity for water, it is frequently employed by chemists to withdraw the elements of water
from other substances; anhydrous sulphuric acid, for example,
can be prepared from oil of vitriol by heating the latter with
anhydrous phosphoric acid. On being heated with various organic substances, such as some of the alcohols and essential oils
composed of carbon, hydrogen, and oxygen, it decomposes them
in sich manner that the oxygen of the organic substance, and as
much of its hydrogen as is necessary to form water by uniting
with this oxygen, combine with the phosphoric acid, while a
compound of carbon and hydrogen (technically called a hydrocarbon) is set free.
After the anhydrous acid has once been dissolved in water, it
cannot again be completely deprived of water by mere evaporation or ignition. When the aqueous solution is evaporated, there
is left, not the anhydrous powder, but a transparent glassy mass,
which is a hydrate, of the composition H2,P 205. This hydrate
is often called glacial phosphoric acid. It is extremely deliquescent, and, at a bright red heat, sublimes in dense white fumes.
Besides the monohydrate, there are two other hydrates of phosphoric acid, of the compositions, respectively, 2H2,,P20,  and




HYDRATES OF PHOSPHORIC ACID.             233
3H20,P,05. Of these, the terhydrate is, perhaps, the most important; it is the substance usually meant when phosphoric acid
is spoken of.
292. Phosphoric acid can be prepared, also, by the oxidation
of phosphorus, or of hypophosphorous or phosphorous acid, or
by the decomposition of some one of its salts, such as the phosphate of calcium (bone-earth).
When phosphorus is heated in dilute nitric acid, of 1'2 specific gravity, the nitric acid gives up oxygen to the phosphorus, nitric oxide and
phosphoric acid are formed, and the latter goes into solution. When
the phosphorus has all disappeared, the solution is evaporated to dryness in order to drive off the nitric acid which was employed in excess, and there is obtained a quantity of the monohydrated glacial
acid.
A product less pure than the acid prepared by means of nitric acid,
is obtained by neutralizing a solution of monophosphate of calcium,
prepared from bones in the manner already described when treating
of the preparation of phosphorus (see ~ 270), with carbonate of ammonium, evaporating the filtered solution of phosphate of ammonium
to dryness and heating the residue to low redness. Ammonia is expelled and glacial phosphoric acid remains.
293. It is a remarkable fact that each of the three different
hydrates of phosphoric acid possesses properties peculiar to itself,
and unlike those of the other two; in fact, each of the hydrates
must be regarded as a distinct acid.  The monohydratdd or
glacial acid, H20,P20,, is usually called metaphosphoric acid;
the bihydrate, 2H2O,P205, is called pyrophosphoric acid; and
the terhydrate, 3H20,P 2O, is commonly spoken of as ordinary
phosphoric acid, or simply as phosphoric acid. The three hydrates are sometimes distinguished as a phosphoric acid (meta),
b phosphoric acid (pyro), and c phosphoric acid (ordinary).
These different hydrates of the acid retain their peculiar characteristics for a considerable time when dissolved in water,
though the mono- and bihydrates change, after a while, to the
terhydrate; and in combining with metallic oxides to form salts,
they unite with 1, 2, or 3 molecules of the oxide, accordingly as
they themselves contain the elements of 1, 2, or 3 molecules of
water. There are thus formed three distinct series of salts, each
of which corresponds to one of the hydrates, as is seen in the




234           META- AND PYROPHOSPHORIC ACID.
following formulae, where M stands for any metal which habitually replaces one atom of hydrogen.
Monohydrated Acid.    Bihydrated Acid.      Terhydrated Acid.
H20,P205            2H1120,P205           3H20,P205
M20,P205            2M20, P205            3M20,P205
Metaphosphate of M.   Pyrophosphate of M.   Phosphate of M.
This behavior is very different from that of the hydrates of nitric
or of sulphuric acid; when either of the hydrates of nitric acid,
for example, is made to combine with a base,.like soda, there hs
formed always one and the same salt, nitrate of sodium. In
each of the three series of salts formed by phosphoric acid,
the acid exhibits peculiar properties.  A salt of the formula
3M20,P 20 will behave very differently towards many reagents
from a salt containing the same metal but in the proportions
M20,P205, or 2M2O,P205. As an example of the kind of differences here alluded to, it may be mentioned that, while metaphosphoric acid, on being added to a solution of albumen, will
cause the albumen to coagulate, no such coagulation can be
brought about by either pyrophosphoric acid or the ordinary terhydrate. Metaphosphoric acid gives a white precipitate when
its solution is mixed with a solution of nitrate of silver by either
of the other hydrates, unless they are first neutralized with an
alkali, in which event a white precipitate is produced by pyrophosphoric acid, and a yellow precipitate by the ordinary acid.
These peculiarities will be examined in detail when we come to
treat of the phosphates of sodium in the chapter upon sodium
and its compounds.
From the formulae given in the above table, it is apparent that
metaphosphoric acid is a monobasic acid, that pyrophosphoric acid.is bibasic, and that ordinary phosphoric acid is terbasic. Since
each of the atoms of M, in either of the formulae, can be replaced by an equivalent atom of any other metal, or by hydrogen,
it follows that the composition of some of the salts of phosphoric acid is rather complex; thus there is a phosphate of potassium  and sodium of the composition (K20,Na20,H20)P205, or
KNaHPO4.
Although we have here studied the acids which contain phosphorus, oxygen, and hydrogen, as if they were in reality composed




PROPERTIES OF PHOSPHORIC ACID.          235
of water and the anhydrous oxide of phosphorus, as the manner of
their derivation would suggest, yet it must not be forgotten that
we have absolutely no knowledge of the actual structure of the
molecules of these compounds, and that the empirical formulae
HPO3, H4P207, and H3PO4 express all that is absolutely known
of their composition.
294. In whichever way prepared, and in all its varieties,
phosphoric acid is a very strong acid. Although a less powerful
agent, at the ordinary temperature, than sulphuric acid, yet, from
being much less volatile than sulphuric acid, it can expel the
latter, and most other acids, from their compounds on being
heated with them. The behavior of the two acids towards calcium, or its oxide, furnishes an instructive example of the influence
of extrinsic or physical circumstances upon the play of the chemical forcc. When triphosphate of calcium (bone-earth) is treated
with dilute sulphuric acid at the ordinary temperature, a quantity of phosphoric acid is set free from the calcium and goes into
solution. From this result it might, at first sight, be thought
that the calcium was removed from the phosphate of calcium
simply by force of the superior chemical power of sulphuric as
contrasted with phosphoric acid; but, in reality, the water which
is present plays an important part in the reaction. Monophosphate of calcium is readily soluble in water, sulphate of calcium,
on the other hand, being well-nigh insoluble. Hence it happens
that when triphosphate of calcium is digested in dilute sulphuric
acid, monophosphate of calcium goes into solution, while sulphate
of calcium is deposited as an insoluble powder. But if the mixture of solid sulphate of calcium and of dissolved monophosphate
of calcium, thus obtained, be evaporated to dryness and the
residue be strongly heated, all the sulphuric acid will be expelled
from the calcium; it will evaporate and pass off into the air, and
nothing will finally be left in the vessel but triphosphate of calcium, precisely similar in quality and quantity to that with which
the experiment started. In the same way, if a mixture of sulphate
of calcium and glacial phosphoric acid be strongly heated, the sulphuric acid, being readily volatile, as compared with phosphoric
acid, will all be expelled from its combination with the calcium:3CaSO4 + PO20 = Ca,3P08 +3S08.




236             TERCHIORIDE OF PHOSPHORUS.
295. Chlorides of Phosphorus.-Phosphorus and chlorine
unite readily and directly even at temperatures as low as 0~,
the act of combination being attended with evolution of light and
heat. If the chlorine be in excess, as regards the phosphorus,
there will be formed a solid quinquichloride of phosphorus, while,
if an excess of phosphorus be present, a liquid terchloride of
phosphorus wili be obtained.
296. Terchloride of Phosphorus (PCI,) is a colorless liquid
of about 1'5 specific gravity, which boils at about 750. It fumes
in the air, and is decomposed by moist air. When heated in the
flame of the gas-lamp, it takes fire and burns with a bright light.
When mixed with water it decomposes, yielding clorhydric and
phosphorous acids:2PC13+6HO=6[IC1+3H20,P20,, or PCl+3H110=3HCl+H3PO,.
This reaction is particularly interesting, in view of the fact that
by means of it we are enabled to obtain phosphorous acid in a
condition of purity. It will be remembered that by the method
of direct oxidation (~ 287) it is no easy matter to obtain pure
phosphorous acid from phosphorus. But by simply treating terchloride of phosphorus with water and evaporating the solution,
so that the chlorhydric acid which results from the reaction may
be expelled, hydrated phosphorous acid is obtained as the sole
product. The process is a good example of the indirect methods
to which chemists are frequently compelled to resort.
Terchloride of phosphorus can be prepared by passing a slow
stream of dry chlorine through melted, almost boiling, phosphorus contained in a tubulated retort which has previously been
filled with chlorine in the cold, and condensing the chloride in
an appropriate receiver as fast as it distils over. The process,
like all operations with phosphorus, requires special care.
297. It will be observed that the formula of terchloride of
phosphorus is that of phosphuretted hydrogen in which all the
hydrogen has been replaced by chlorine. The two substances
have a perfectly similar volumetric composition. In phosphuretted hydrogen, three volumes of hydrogen in combination with
half a volume of phosphorus, produce two volumes of the compound gas; if, in terchloride of phosphorus,




QUINQUICHLORIDE OF PHOSPHORUS.          237
Half a unit-volume of phosphorus-vapor, weighing.    3100
And 3 unit-volumes of chlorine, weighing (35'5 X 3).   106-50
Produce 2 volumes of PC13 vapor, weighing..   137'50
One unit-volume of PC13 vapor should weigh...    6875
The specific gravity of the vapor of terchloride of phosphorus
has been found, by experiment, to be 69'12,-a number so
nearly identical with the above result of calculation as entirely
to confirm the assumption on which the calculation rests, viz.
that in terchloride of phosphorus three volumes of chlorine are
united with half a volume of phosphorus-vapor.
298. Quinquichloride of Phosphorus (PC]5) is a white or
straw-colored crystalline solid, which volatilizes at a temperature
below 100~ without previously fusing; but when subjected to
pressure, it melts at 1480, and boils at a temperature somewhat
higher. It burns in the flame of a candle with production of
phosphoric acid and evolution of chlorine. It is very deliquescent, and is decomposed by the moisture of the air; by a large
excess of water it is immediately resolved into chlorhydric and
phosphoric acids:PC  + 411H20 = 5HC1 + H3P04;
with a smaller quantity of water it yields chlorhydric acid and
oxychloride of phosphorus:PC15 + 1H20 = 2HC1 + POC13.
Sulphydric acid decomposes it in like manner, with production of
chlorhydric acid and sulphochloride of phosphorus:PC1, +  HS = 2HC1 + PSC13.
Quinquichloride of phosphorus reacts upon many organic compounds also, with formation of very interesting products, and is
hence an important agent of research in the department of
organic chemistry.
299. In order to prepare quinquichloride of phosphorus, a
current of dry chlorine may be passed into terchloride of phosphorus until the latter has been completely solidified; the
product is then distilled in a current of chlorine. The quinquichloride may be obtained directly from phosphorus in one operation, if a rapid stream of chlorine be conducted into a retort




238                     DISSOCTATION.
containing phosphorus, kept so cool that the terchloride of phosphorus at first produced shall not distil over. Again, if powdered
red phosphorus is exposed to the action of a rapid stream of
chlorine, it will all be quickly converted into the solid quinquichloride.
300. The formula above given for quinquichloride of phosphorus represents the following composition by volume:
Half a unit-volume of phosphorus-vapor, weighing..  31 00
And 5 unit-volumes of chlorine, weighing (355 X 5)   177-50
Should produce 2 vols. of PCI, vapor, weighing.. 208'50
One unit-volume of PC1, vapor ought then to weigh.   104'25
The specific.gravity of the supposed vapor of quinquichloride of
phosphorus, as determined by experiment, does not accord with
this calculated result; it is 52-81, almost exactly one-half of the
theoretical unit-volume weight above given. Iffour volumes of
vapor, instead of two, resulted from the union of half a volume
of phosphorus-vapor with five volumes of chlorine, the calculated and the actual vapor-density would coincide. But hitherto
we have never found a single compound gas in which the product volume was four unit-volumes; two unit-volumes have invariably resulted from the union of the constituent volumes,
whatever the character and number of the constituents. It
would be necessary to admit this substance as presenting an exceptional volumetric composition, were it not for a well-founded
distrust of the experimental determination of the vapor-density
of this compound. It not infrequently happens that all attempts
to determine the vapor-density of volatile compounds of two or
more elements are baffled by their splitting up, at the temperature of vaporization, into their constituent gases or vapors, which,
in the act of separating, resume their own proper volumes, however much they may have been condensed during combination.
This splitting up of compound vapors, at high temperatures,
into less complex compounds, or into the elementary constituents, is termed dissociation.  Thus, at the elevated temperature
necessary to convert quinquichloride of phosphorus into vapor,
it is probable that the quinquichloride splits into terchloride
of phosphorus and free chlorine, and that it is the specific




DISSOCIATION.                   239
gravity of this mixture which has been determined, instead of
the specific gravity of the real unaltered vapor of the quinquichloride.
Two unit-volumes of PC13 weigh..      138-24
Two unit-volumes of C1 weigh......    71
Four unit-volumes of the mixture weigh..           209'24
One unit-volume of the mixture weighs....           65231
The specific gravity which has been assigned to the quinquichloride is 52'81,-a number very nearly coincident with the above
calculated density of the mixture of terchloride-vapor and free
chlorine. At the high temperature of vaporization it is therefore probable that quinquichloride of phosphorus undergoes
dissociation into terchloride of phosphorus and chlorine; but if
this be the case, these constituents recombine when the temperature falls, for by lowering the temperature the quinquichloride is
recovered.
As we advance, we shall meet with several other examples of
the dissociation of compound gases and vapors; for the present
it will be sufficient to give one more illustration of the meaning
of this term. When equal volumes of dry ammonia and dry
chlorhydric acid gas are mixed (Exp. 65), the two gases are
completely condensed to a white solid, which we are familiar
with as chloride of ammonium. Since this ammonium-salt is
readily volatilizable, there would be no difficulty in determining
the product-volume of the compound of ammonia and chlorhydric acid, were it not for the fact that the vapor of chloride of
ammonium undergoes dissociation at the temperature of vaporization. If the real vapor of the compound could be measured, the
facts would undoubtedly be correctly represented by the diagram,
H111     +       NH3             NH11;
but the vapor of the compound is resolved into its constituent
gases at the high temperature necessarily employed, so that the
following diagram really figures the actual state of things:



240         BROMIDES AND IODIDES OF PHOSPHORUS.
HC1      +      NHI3              C1...... NH 
When the dissociated vapor cools, the parted gases recombine to
form solid chloride of ammonium.
It is obvious that the phenomena of dissociation interfere
fatally with one of the common methods of arriving at the
weight and structure of the molecule of a volatile compound;
the indirect method of getting at the volumetric composition of
a substance from its ponderal composition and the specific gravity of its vapor becomes impracticable whenever the vapor of
the compound under examination is liable to dissociation, inasmuch as experiment cannot determine beyo:ld a doubt the real
vapor-density of such a body.
301. Bromides of Phosphorus.-When a piece of phosphorus
is dropped into bromine, the two elements combine with explosive violence, the burning phosphorus being thrown about in a
highly dangerous manner. There are two bromides of phosphorus, PBr3 and PBr,, corresponding to the two chlorides.
The terbromide is liquid at ordinary temperatures and the quinquibromide solid.
302. Iodides of Phosphorus.-Iodine and phosphorus unite
directly, when brought in contact with one another, and so much
heat is developed by their union, that a portion of the phosphorus will take fire if the mixture be in contact with the air.
There are two iodides of phosphorus, both of them solid at the
ordinary temperature; their composition is respectively PI, and
PI3. It will be noticed that, while the teriodide corresponds to
the terchloride and terbromide, the other compound is a biniodide, of which there is known neither a bromine nor a chlorine
analogue.  The fact is interesting as illustrating the general
truth that, when in any group or family of elements we compare
the behavior of its several members, analogy ceases to be a sure
guide, in proportion as the individuals compared are more widely
separated in the natural series. Chlorine and bromine stand
next to one another in the family or series of elements to which




SULPHIDES OF PHOSPHORUS.             241
they belong; and as we have just seen, their behavior, as regards
phosphorus, is well-nigh identical; but iodine, one step further
removed from chlorine than bromine is, enters into new combinations not altogether conformable to those of chlorine.
303. Sulphides of Phosphorus.-There is a definite sulphide
of phosphorus corresponding to each of the oxides, and in addition to these there are two other compounds, which may be represented by the formulae P4S3 and P2S12.  Sulphur and phosphorus
may also be melted together in any proportion. The sulphides
of phosphorus are exceedingly inflammable, taking fire even
more readily than phosphorus itself, and they are all more readily
fusible than either of the two elements of which they are composed. They may be prepared by heating sulphur under water
in contact with melted phosphorus. The union of the two elements is attended with development of much heat, and sometimes
with dangerous explosions. It is well, therefore, to operate only
upon small quantities, and to add the sulphur gradually to the
phosphorus.
CHAPTER XVII.
A RSENI C.
304. Compounds of arsenic have been known from very early
times. The element is sometimes found native, but much more
frequently associated with other metals and with oxygen and
sulphur. The metals in connexion with which it most commonly
occurs are iron, cobalt, nickel, and copper. Ferruginous ores
and deposits, in particular, are rarely free from traces of arsenic.
In small quantity, arsenic is very widely distributed.
The greater part of the arsenic of commerce is prepared from
a native arsenide and sulphide of iron (arsenical pyrites') corresponding to the formula FeAsS, and from the arsenides of nickel
and cobalt.  3Metallic arsenic is obtained directly from  the
mineral of the formula FeAsS by heating. it in earthen tubes laid
R




24-2               PROPERTIES OF ARSENIC.
horizontally in a long furnace; a tube, made by rolling up a
piece of thin sheet iron, is inserted in the mouth of each earthen
retort, and an earthen receiver is luted on to this iron tube.
The arsenic condenses principally in the iron tube, in the form
of a compact, whitish, crystalline mass, which is detached, when
cold, by unrolling the sheet iron. The metal is also indirectly
obtained by reducing the arsenious acid (s203,) which results
from roasting (heating in a current of air) arsenides, like those
of cobalt and nickel; this oxide is heated with charcoal in earthen
crucibles covered with conical iron caps, or invei-ted crucibles,
into which the reduced metal sublimes. The metal obtained by
the second process is gray and pulverulent, instead of whitish
and coherent.
FeAsS = FeS + As;        As203 + 3C = 2As + 3C0.
305. Arsenic is a brittle solid, of a steel-gray color and a metallic lustre. Its specific gravity has been variously given at
from 5'62 to 5'96. Like the metals, it is a good conductor of
electricity. It crystallizes in acute rhombohedrons, and in octahedrons also, thus taking on forms of both the monometric and
hexagonal systems, as do phosphorus, the preceding member, and
antimony, the succeeding member of this family. At a dull red
heat it volatilizes without previous fusion; the vapor is colorless,
and possesses a characteristic odor resembling that of garlic.
The specific gravity of this vapor is 150, while the atomic
weight of the element is 75; arsenic, therefore, resembles phosphorus, and differs from all the other elements heretofore studied,
in that its atomic weight is not identical with its unit-volume
weight; two combining proportions by weight of arsenic occupy
the same volume as one combining proportion of hydrogen; its
symbol, As, represents its atomic weight, but only half the
weight of the unit-volume of its vapor. At the ordinary temperature the compact metal does not tarnish by exposure to dry
air, but a moistened powder of arsenic is slowly converted by the
air into a mixture of arsenious acid and metallic arsenic. At a
red heat the metal burns with a whitish flame, producing a white
smoke of arsenious acid. When thrown, in fine powder, into
ciiioxine gas, it takes fire spontaneously and is converted into




ARSENIC AND HYDROGEN.              243
chloride of arsenic (AsCl,). Bromine, iodine, and sulphur also
combine readily with arsenic, when aided by a gentle heat.
Nitric acid and aqua regia convert the metal into arsenic acid
(As2O,); chlorhydric acid has little action upon it. Dilute sulphuric acid has no action upon the metal; but the concentrated
acid has the same effect upon arsenic as upon phosphorus (~ 276);
arsenious acid is formed and sulphurous acid escapes:3H2S04 + 2As = As203 + 3S02 + 3H20.
Some fatty oils dissolve arsenic to a slight extent, as they do
phosphorus. Metallic arsenic unites by fusion with most metals,
forming alloys which the arsenic tends to make hard or brittle.
In the manufacture of shot a little arsenic is added to the lead
to facilitate the formation of regular globules.
306. Arsenic and Hydrogen.-Arsenic forms two combinations
with hydrogen; one of these is an unstable, brown solid of uncertain composition; the other is a well-known gas whose constitution is represented by the formula AsH3, and which is
therefore analogous in composition to ammonia (NH3) and phosphuretted hydrogen (PH3). The solid hydride is so obscure a
substance that nothing need here be said of it, except that it is
supposed to contain two atoms of hydrogen and one of arsenic
(AsH,?).
307. Arseniuretted Hydrogen.-This very dangerous gas may
be prepared in an impure state by decomposing, with sulphuric
acid diluted with three parts of water, an arsenide of zinc obtained by fusing together equal weights of powdered arsenic and
granulated zinc:3H2S04 + Zn3As = 3ZnHSO4+ AsH3.
As it is not possible to prepare the precise alloy ZnAs, the
arseniuretted hydrogen thus obtained is always mixed with hydrogen. The arsenide of sodium can be decomposed by water,
with evolution of arseniuretted hydrogen:
3HO0 + Na3As = 3NaHO + AsH3.
The same remark, however, applies to this reaction as to the preceding one; the product is contaminated with an indeterminate
quantity of free hydrogen. Arseniuretted hydrogen seems also
to be formed whenever the oxides of arsenic, or compounds of
R 2




244               ARSENIURET~ED HYDROGEN.
these oxides, are brought in contact with nascent hydrogen.  A
mixture of arseniuretted hydrogen and hydrogen may be readily
obtained by acting upon zinc by dilute chlorhydric or sulphuric
acid in which arsenious acid has been dissolved.
308. Arseniuretted hydrogen is a colorless gas, having a fetid
odor; even when very much diluted with air, it is intensely
poisonous, and fatal results have repeatedly followed its accidental inhalation. In experimenting with this deadly gas, the
greatest care is required not to inhale the least portion of it.
It has been condensed at — 40~ to a transparent liquid, but it
has never been solidified. The gas is soluble in water at the
ordinary temperature only to the extent of one-fifth of its volume,
and neither the gas nor its aqueous solution has any action upon
blue or red litmus-paper. In spite, therefore, of its strong resemblance to ammonia in composition, some of its physical properties are strikingly unlike those of that very soluble and intensely
alkaline gas. Arseniuretted hydrogen burns in the air with a
whitish flame, forming water and a white smoke of arsenious
acid; but if a cold body, like a piece of porcelain, for example,
be introduced into a jet of the burning gas, the hydrogen alone
will burn, and the arsenic will be deposited in the metallic
state upon the porcelain surface, forming a lustrous black spot.
This effect is precisely similar to the deposition of soot on a cold
body held in the flame of a candle. It is also decomposed when
caused to pass through tubes heated to dull redness, metallic
arsenic being deposited as a brown or blackish mirror, while
hydrogen gas escapes. This decomposition is a good illustration
of the dissociation of gases (~ 300). Chlorine in excess reacts
violently upon it, forming terchloride of arsenic (AsC1,) and
chlorhydric acid:AsH3 + 6CI1    AsC13 + 3IHC1.
When, however, chlorine acts on an excess of arseniuretted hydrogen, there are formed chlorhydric acid and metallic arsenic;
flame accompanies this reaction. The reactions of bromine and
iodine are similar to, but less violent than, those of chlorine.
We recall, in this connexion, the decomposition of ammonia by
chlorine, with formation of chlorhydric acid and liberation of




ARSENIURETTED HYDROGEN.              245
nitrogen. Arseniuretted hydrogen decomposes the solutions of
the salts of many of the heavy metals, but the products are
somewhat various; sometimes a metallic arsenide is precipitated;
sometimes the heavy metal is precipitated, while arsenious acid
remainis in the solution. As we shall shortly see, the chemical
properties of this gas are of great importance in the processes
used for detecting arsenic in cases of poisoning.
309. Arseniuretted hydrogen may be analyzed by precisely
the same method which was used for the analysis of phosphuretted hydrogen (~ 280); but the results can only be approximate, because of the extreme difficulty, not to say impossibility,
of obtaining the gas in a state of tolerable purity. The composition of the analogous gases, ammonia and phosphuretted hydrogen, and the specific gravity of the gas, lead us to the following
statement of its composition. Two volumes of the gas contain
3 unit-volumes of hydrogen, weighing  3 X 1  = 3
munit-volume of arsenic-vapor, weighing   X 150 — 75
2 unit-volumes of arseniuretted hydrogen weigh    78
1 unit-volume of arseniuretted hydrogen weighs   39
The actual specific gravity of arseniuretted hydrogen, as determined by experiment, is as nearly as possible 39, —a fact which
makes it certain that two volumes of the gas do not contain one
volume of the heavy arsenic-vapor, which is 150 times as heavy
as hydrogen, but only half a volume. Herein this gas differs
from  ammonia, but resembles phosphuretted hydrogen;  The
weight of the quantity of arsenic which combines with three
atoms of hydrogen is 75, just as the weight of the quantity of
nitrogen which combines with three atoms of hydrogen is 14;
but 75 parts by weight of arsenic-vapor only occupy one half the
space which 14 parts of nitrogen fill.
310. Arsenic and Oxygen.-Arsenic forms two well-defined
oxides, arsenious acid, As20,, and arsenic acid, As205.  The
black film which forms on the surface of the metal when exposed to the air is by many supposed to be a suboxide, while
others think it is more probably a mixture of metallic arsenic
with arsenious acid. The first of the above-mentioned acids corresponds with nitrous and phosphorous acids, the second with




246                    ARSENI'OUS ACID.
nitric and phosphoric; but arsenious acid is very stable, in comparison with arsenic acid, while the reverse is true of the analogous acids containing nitrogen and phosphorus. The element
arsenic possesses many properties which ally it to the metals; but
in its compounds its close connexion with nitrogen and phosphorus is clearly exhibited. Its oxides, for example, are both
acids, and these acids unite with the oxides of the metals proper
to form stable, crystallizable salts, which are in many cases
isomorphous with the corresponding salts containing phosphorus.
311. Arsenious Acid (As,O,).-Arsenious acid, known in commerce as arsenic, or white arsenic, is obtained as a secondary product in the roasting of arsenical ores of nickel, cobalt, and tin,
and as a principal product in the roasting of arsenical pyrites.
The volatile matters which escape from the roasted ores consist
mainly of sulphurous and arsenious acids; the first is allowed to
pass off into the atmosphere, the second condenses to the solid
state in the chambers and long passages through which the vapors
are forced to pass in order that they may deposit their arsenious
acid.  A second sublimation purifies the raw product.  According to the temperature at which the arsenious acid is sublimed
and condensed, the product is either in powder or in transparent
masses; a low temperature with sudden condensation yields a
white powder of minute crystals; a higher temperature with
gradual solidification produces a transparent glass.
312. Arsenious acid is a white solid, which occurs not only in
two conditions, one amorphous and the other crystalline, but also
in two distinct crystalline forms. When the vapor of the acid is
cooled so quickly that it solidifies at once, without passing
through the semifluid state, each particle of the solid acid
assumes more or less perfectly the octahedral form. A hot saturated aqueous solution of the acid also deposits regular octahedral crystals on cooling. The amorphous, glassy variety of the
acid changes spontaneously, when kept in contact with the air,
into an aggregation of minute octahedral crystals, thereby becoming opaque and porcelain-like in appearance. The other
crystalline form of arsenious acid is the right rhombic prism;
this form occurs much less frequently than the first, and is con



ISOMERISM.                    247
verted into the octahedral form by sublimation and by solution
in hot water.
The two varieties of arsenious acid, the vitreous and the porcellaneous, differ decidedly in physical and chemical properties, yet
they have precisely the same chemical composition; and when
either variety changes into the other, no alteration of weight, no
addition or subtraction of matter accompanies the change. The
two varieties contain the same two elements in precisely the
same proportions. When two or more compounds, which exhibit
essential differences of physical and chemical properties, are,
nevertheless, found to be identical in respect to constituent elements and their proportions, the compounds are said to be isovr,nric (equal parts).  The teim  allotropism (~ 162) properly
applies to the elements only, the term  isomerism to com2pounds
only; both terms, however, refer to one and the same unquestionable, though perplexing, truth-namely, that the widest
diversity of properties may coexist with absolute identity of
ultimate chemical constitution. Two allotropic states of the
same  element not infrequently present more striking differences than elements recognized as distinct; and among the
numerous complex compounds of carbon with which organic
chemistry deals, there are many isomeric compounds which are
so entirely dissimilar as to lead almost irresistibly to the belief
that it is of as much consequence how the atoms of a compound
are arranged as what kind of atoms they are. Arsenious acid
does not afford a very striking example of isomerism; nevertheless the properties of its two modifications are quite diverse. If
it be true that the different arrangement of atoms is the cause of
the diversity of isomeric compounds, it is evident that the differences between two varieties of a compound of only two kinds of
atoms, united in the simple ratio of 2 to 3, cannot be expected to
be so marked as the differences between isomeric compounds
which contain four or five elements united in the very complicated proportions which frequently characterize the compounds
of carbon.  Nevertheless the differences between the two isomeric compounds of arsenic and oxygen are sufficiently distinct.
The glassy acid dissolves much more rapidly in water than
the porcelain-like variety, being three times as soluble in that




248             PROPERTIES OF ARSENIOUS ACID.
liquid. The relation of the two varieties to heat is not the same;
for when the vitreous acid changes into the opaque, heat is disengaged. As this change generally takes place slowly, from the
surface towards the centre of any fragment of the vitreous variety,
the heat evolved is not perceived; but if the change be suddenly
accomplished, not only heat, but light also will be disengaged.
EExp. 122.-Dissolve 4 or 5 grms. of the vitreous acid in a hot mixture of 24 grms. of strong chlorhydric acid and 8 c. c. of water, and
let the solution cool slowly; the arsenious acid will crystallize in
transparent octahedrons, and the formation of the crystals will be accompanied by flashes of light.
The specific gravity of the vitreous acid is 3'738; that of the
porcellaneous 3'699. The opaque variety may be changed into
the vitreous by long boiling with water.  It appears, therefore,
that the arrangement of atoms which may be supposed to furnish the vitreous acid is stable only at high temperatures, and
that the arrangement of atoms which is peculiar to the opaque
acid is stable only at low temperatures.
313. Arsenious acid volatilizes without change when heated
with free access of air; if heated in contact with carbon, it gives
up its oxygen, and metallic arsenic is liberated.  Copper and
many other metals reduce arsenious acid.
Exp. 123.-Place a few particles of arsenious acid in an open tube
of hard glass (No. 5) about 10 c.m. long, and heat the acid over the
lamp, holding the tube in a sloping position; the white solid will be
volatilized, but it will immediately be deposited again upon the cold
part of the tube. By the aid of a lens, this deposit may be seen to be
crystalline.
Fig. 46.            Exp. 124.-Drop into the point
A of a drawn-out tube of hard glass,
No. 5, a morsel of arsenious acid,
and above it place a splinter of
charcoal (Fig. 46); heat the coal
red-hot in the flame of the lamp,
and then volatilize the arsenious acid. The acid will give its oxygen
to the coal, and the arsenic will be deposited in a ring on the cold
part of the tube, presenting a brilliant metallic appearance.
Exp. 125.-Throw a particle of arsenious acid upon a piece of redhot charcoal; the acid will be partly reduced, and the peculiar garlic
odor of the vapor of metallic arsenic will be perceived.




SOLUBILITY OF ARSENIOUS ACID.            249.rxp. 126.-Dissolve a few centigrammes of arsenious acid in 5 or 6
c. c. of chlorhydric acid heated in a test-tube; in the hot solution
immerse a narrow strip of clean copper; an iron-gray film will be deposited upon the copper. This coating contains metallic arsenic derived from the arsenious acid; it consists of an alloy of arsenic and
copper.
314. It is very difficult to say what the solubility of arsenious
acid in water really is. The results of different experimenters
rP[esent very wide discrepancies, due in part to the fact already
stated, that the two modifications of arsenious acid are of unlike
solubility, and in part also to the circumstance that the acid dissolves with extreme slowness. The difficulty of the determination is increased by the readiness with which either modification
passes into the other in consequence of changes of temperature;
it is quite possible that both varieties may simultaneously exist
in the same solution. A hot aqueous solution usually contains 1
part of the acid in 10 or 12 parts of water; on cooling this solution, a portion of the acid separates, leaving a solution which
contains 1 part of acid in 20 to 30 parts of water. The aqueous
solution has a feeble acid reaction.  No definite hydrate of arsenious acid is known. The acid is much less soluble in alcohol
than in water.  Hot chlorhydric acid dissolves it with facility,
and when cold retains a large proportion in solution; other acids,
even some vegetable acids, dissolve it readily when hot, though
most of them keep but little in solution when cooled. When the
solution of arsenious acid in chlorhydric acid is evaporated, a
compound of chlorine and arsenic, the terchloride of arsenic
(~ 336), is volatilized, and the solution thus loses a portion of
its arsenic. This fact is of importance in examinations for arsenic
in cases of suspected poisoning.
315. Solutions of caustic soda and potash readily dissolve the
acid, a soluble arsenite of sodium or potassium resulting from
the reaction.  From  these arsenites of sodium  and potassium
the arsenites of other metals are generally obtained by the way
of double decomposition. The arsenites are numerous, but they
are not very stable and have been but little studied.
316. Arsenious acid is oxidized and converted into arsenic
acid by digestion with nitric acid. The same transformation is




250                USES OF ARSENIOUS ACID.
brought about, but quicker, by the action of aqua regia, and by
chlorine, bromine, and iodine in presence of water.  When
iodine is added to a solution of arsenious acid mixed with a
little starch-paste, the whole of the arsenious acid is converted
into arsenic acid before any blue coloration of the starch is produced by the iodine. These facts are turned to account in the
volumetric determination of chlorine (~ 565). Sulphuretted hydrogen colors an aqueous solution of arsenious acid yellow, and
precipitates a yellow sulphide of arsenic (~ 210) from a solution
acidulated with chlorhydric acid.
Arsenious acid is a violent poison, all the more dangerous
because it has neither taste nor odor to warn the victim of its
presence; two decigrammes of it will cause death.  All the soluble salts of arsenious acid are likewise horribly poisonous. The
best antidote to the poison is a mixture of moist, freshly precipitated sesquioxide of iron and caustic magnesia.
317. Arsenious acid is largely used for the manufacture of
two green paints, an arsenite of copper and a compound of
arsenite and acetate of copper; it is applied as an oxidizing
agent in the manufacture of glass; it is used for poisoning vermin, and is consumed in considerable quantities for producing
the arsenic acid which is used in the dyeing and printing of
cloth, and in the manufacture of aniline colors; it is used in
very small doses as a remedy for asthma, and in some skin-diseases. Although the acid is so,violent a poison, it seems to be
possible, by beginning with small doses and gradually increasing
them, to accustom the human body to sustain, without injury,
doses of 2 to 3 decigrammes, or even more; the arsenic thus
taken is said to produce a plump and healthy appearance in those
who use it, and especially to increase the power of the respiratory organs.  In veterinary practice, it has been found that
arsenious acid administered to animals in this manner improves
the appearance of the skin.
318. Arsenic Acid (As,20).-This compound is produced by
oxidizing arsenious acid with nitric acid, aqua regia, hypochlorous
acid, or other oxidizing agents.
Exp. 127.-Add 4 grms. of powdered arsenious acid, little by little,
to a mixture of 4 grms. of concentrated nitric acid, and 8 grms. of




ARSENIC ACID.                    251
concentrated chlorhydric acid, contained in a small evaporating dish,
and gently heated over a lamp in a strong current of air, or'beneath
a well-ventilated hood. The liquid, which at first gives off red fumes
in considerable quantity, must be evaporated until it assumes a syrupy
consistency, resembling that of oil of vitriol. This syrupy liquid is
arsenic acid.
319. The syrupy solution thus obtained deposits, after standing for some days, at the ordinary temperature, transparent
elongated prisms or rhomboidal laminse. These crystals, heated
to 1000, first melt and then yield the terhydrate of arsenic acid
(3HO20,As2O5 = 2H3As04) as a crystalline  precipitate.  The
same hydrated acid separates in large prismatic crystals when a
concentrated aqueous solution is cooled to a low temperature.
There are two other hydrates of the oxide As205, —a bihydrate,
2H20,As205 = H4As207, and a  monohydrate, 1H20,As20, =
2HAsO3; both these lower hydrates are obtained from the terhydrate by subjecting the latter to the prolonged action of certain
temperatures. If either of the hydrates be heated to dull redness, a white amorphous mass remains, which is the anhydrous
acid, As205; this substance has no action upon litmus, and seems
to be scarcely soluble in water. After long exposure to moist
air, it slowly deliquesces, and if covered with water and soaked
for a long time, it at last dissolves, being probably converted into
the soluble terhydrate. At a full red heat it is resolved into
arsenious acid and oxygen.
320. In spite of the recognized existence of three solid hydrates of arsenic acid, there is but one aqueous solution of this
acid, inasmuch as the monohydrate, the bihydrate, and the anhydride, are all immediately converted into the terhydrate when
dissolved in water. The solution has a very sour taste' and a
strong acid reaction on vegetable colors. The concentrated liquid
is highly corrosive and produces blisters on the skin. Arsenic
acid and its salts are poisonous, but not in so high a degree as
arsenious acid and the arsenites.
321. Arsenic acid is a strong acid, capable of expelling all
the more volatile acids from their salts at high temperatures.
Its three hydrates are strictly comparable with the three hydrates
of phosphoric acid.




252                 SALTS OF A KSENIC ACID.
Hydrates      HPO3             HAsO,3         ydrates
0y    Hi P,0,             H4As07         of
Phosphoric Acid.  HPO            H AsO4   Arsenic Acid.
Either one, two, or all three of the hydrogen atoms in common
arsenic acid, H3As04, may be replaced by a metal, so that three
arseniates of any one metal may exist, as for example,
NaHAsO4             Na,HAsO4               NaAs04
Acid Arseniate     " Neutral" Arseniate    Basic Arseniate
of Sodium.           of Sodium.            of Sodium.
If an acid arseniate be suitably heated, a meta-arseniate results,
as, for example, NaAsO3 = NaH2AsO4 —'H20; if a neutral
arseniate be sufficiently heated, a pyro-arseniate results, as, for
example, Na4As207 = 2Na2HAs04 - H20; but such meta- and
pyro-arseniates, unlike the corresponding meta- and pyro-phosphates, have very little stability, take up again the molecule of
water, which the heat expelled, the moment they are brought in
contact with water, and are so changed back again into salts of
ordinary arsenic acid. The salts of arsenic acid are isomorphous
throughout with the corresponding phosphates.
322. Arsenic acid is readily reduced to arsenious acid, and,
consequently, acts in some cases as an oxidizing agent. Thus
sulphurous acid reduces arsenic acid, and is itself converted into
sulphuric acid: —
2H1AsO4 + 2S02 = 2H2S04 + As203 + H20.
Sulphydric acid gas, passed throtgh a not too concentrated solution of arsenic acid, slowly precipitates the yellow tersulphide of
arsenic, the action being assisted by heat and by the presence of
another acid. Charcoal and the metals at a red heat reduce
arsenic acid to the metallic state, just as they do arsenious acid.
323. Arsenic acid has been extensively used in calico-printing, in place of the more expensive tartaric acid, for developing
white patterns on a colored ground in the chloride-of-lime vat.
It is also an excellent preservative of animal substances, and is
accordingly used to defend the specimens and preparations of
the anatomist and naturalist from decay and from the attack of
insects.
324. Detection of arsenic in cases of poisoning.-Nearly all
compounds of arsenic are poisonous; but arsenious acid is best




DETECTION OF ARSENIC.                253
known and most easily procured, and is therefore most likely to
be met with in cases of poisoning by arsenic, whether accidental
or intentional. In criminal trials the solubility of arsenious acid
in water has often been much discussed; but this is practically a
point of little importance, for the tasteless poison is generally
administered in the solid state mixed with soup, gruel, milk, or
even with solid food. It thus sometimes happens that small particles of the poison can be found adhering to culinary vessels,
cups, plates, or spoons, or even to the coatings of the stomach
and intestines after death. If the arsenious acid is too finely
divided to be picked out in lumps, it may sometimes be separated by stirring up the mass, under examination, with water, and
leaving the heavier particles to settle. Any solid arsenious acid
that may be present will be sure to be found in the residue; it
may be washed with cold water. It is always very satisfactory
thus to obtain the solid poison in the condition in which it was
administered, because the examination is, in such cases, very
direct and conclusive. It is only necessary to try, with the white
powder thus obtained, the experiments already given to illustrate
the properties of arsenious acid (Exps. 123-126), together with
certain other discriminating tests shortly to be described.
325. It more frequently happens, however, that the arsenic
has been dissolved by the acid secretion of the stomach, and has
become intimately mixed with the food or excretions, or incorporated into the substance of the organs themselves. The examination then becomes more difficult. The reduction of arsenious
acid by copper (Exp. 126) is an available test in such cases.
To the suspected matter, if liquid, about one-sixth of its bulk of
chlorhydric acid is added, and the mixture is gently boiled.
Solid tissues must be cut into small pieces and boiled for some
time with dilute chlorhydric acid (1 part acid to 6 parts water)
until the whole is disintegrated; this solution is finally clarified
by filtration. Strips of copper gauze or foil are then immersed
in the boiling liquid; and if any gray deposit is produced, fresh
pieces of metal are added so long as the color of the copper is
perceptibly changed.  They are then removed, washed with
water, dried, folded up, placed in a dry tube of hard glass and
gently heated. Some of the metallic arsenic in the gray alloy




25F4                       DIALYSIS.
will be converted into arsenious acid, which collects on the cold
part of the tube in the form of a crystalline sublimate. To this
sublimate all tests for the identification of arsenious acid can be
applied. This mode of operation is known as Reinsch's test.
The chlorhydric acid employed must be proved to be free from
arsenic.
326. Another method of separating arsenious acid from the
organic matters with which it is mixed is that of dialysis, a process which depends upon the very different rates at which different substances diffuse through water.
EBp. 128.-Select two straight-sided bottles of clear glass about 15
c.m. deep and 8 to 9 cm. wide. Fill them seven-eighths full of distilled water, or any pure soft water. Dissolve 10 grms. of bichromate
of potassium in 100 c. c. of water; suck as much of this solution as
will fill the remaining eighth of one of the above-mentioned bottles
into a pipette (Appendix, ~ 22), and carefully convey the colored fluid
to the bottom of the bottle by bringing the fine point of the pipette to
the bottom of the bottle and then allowing the liquid to flow very
slowly out of the pipette. If time enough (5 or 6 minutes) be taken
for this process, no sensible intermixture of the two liquids will take
place during the delivery. Dissolve 10 grms. of caramel (melted and
partially burnt sugar) in 100 c. c. of water, and convey to the bottom
of the second bottle, in the same manner as before, enough of the
dark-colored solution to fill the bottle.
The two bottles are left at rest for several days in a room where the
temperature is nearly constant. Spontaneous diffusion immediately
begins; and the very different rates at which the two colored substances diffuse upwards through the water should be from time to
time observed.
327. Substances which have a comparatively high diffusive
power have generally, though not invariably, the power of crystallizing; their solutions are generally free from viscosity, and
always have taste.  Such substances are designated by the term
crystalloids.  Among crystalloids there are wide differences of
diffusive power; thus caustic potash diffuses twice as fast as sulphate of potassium, and sulphate of potassium twice as fast as
sulphate of magnesium.
Substances of very low diffusive power have little, if any,
tendency to crystallize, and affect a vitreous structure. Such




CRYSTALLOIDS AND COLLOIDS.               255
kubstances are often very soluble in water; but their solutions
have always a certain degree of viscosity when concentrated,
and are insipid or wholly tasteless. By combining with water,
these substances are apt to form  jellies.  Gelatine has been
taken as the type of this class; and they have hence been
called co7loids, a name derived from  Greek words signifying
ylete-llile. Among the colloids rank hydrated silicic acid, alumina,
starch, gums, caramel, albumen, and animal and vegetable extractive matters.
As we can separate by means of distillation or evaporation two
bodies of different volatility, so by the aid of diffusion we can separate
one substance more or less completely from another. Jellies and colloid membranes are permeable to crystalloids, but are practically impermeable by colloids like themselves. By means, therefore, of a
colloidal diaphragm, or partition, crystalloids can be separated from
colloids by diffusion. The most suitable substance for the dialytic
diaphragm is parchment paper, prepared by soaking unglazed paper,
for a few seconds, in a mixture of 6 parts of strong sulphuric acid and
1 part of water, and immediately washing it, first in water and then in
water containing amlnonia. The paper subjected to this treatment
becomes semitransparent and tough, like parchment.
A dialyzer, as the apparatus for effecting separation by diffuision is
called, consists of two gutta-percha, or wooden, hoops, one of which
should be 5 c.m., and the other 2-5 c.m. deep. The deeper hoop is
slightly conical, and the shallower must slip over the small end of the
deeper. The hoops may be fromn 15 to 25 c.m. in diameter. The
parchment paper, which is to form the bottom, must be about 8 c.m.
wider than the small end of the 5 c.m. hoop. To prepare the dialyzer
for use, soak the parchment paper for about a minlute in distilled water;
stretch it evenly over the small end of the 5 c.m. hoop and strain it
on tightly by pushing over it the 2'5 c.m. hoop. The paper must be
pressed smoothly up round the outside of the deeper hoop, and the
bottom must be flat and even.
There must be no small holes in the paper. To detect such, put
distilled water into the dialyser to the depth of 5 m.m., and place the
dialyzer on some white blotting-paper. If any wet or dark spots appear, they indicate the existence of small holes. To close such holes,
apply to the under surface of the paper about the holes some solution
of albumen, put on a small patch of parchment paper, and iron the
patch with a hot smoothing-iron. The albumen will coagulate, fix
the patch, and close the hole..Exp. 129.-Into the dialyzer so prepared pour an aqueous solution




256               DIALYSIS OF ARSENIOUS ACID.
containing five per cent. of cane-sugar and five per cent. of gumi
arabic, to the depth of about 1-25 c.m.  Then float the dialyzer on
distilled water contained in a flat basin. The volume of water in the
basin should be from 5 to 10 times as great as the volume of the fluid
in the dialyzer. The wider the dialyzer and the greater the quantity
of distilled water in the outer basin, the more rapid and effective the
diffusion.  A dialyzer 15 c.m. in diameter, serves to operate upon 200
to 2.50 c. c. of liquid; one of 20 c.m. upon 400 to 450 c. c.; one of 25
c.m. upon 600 c. c.
After the lapse of twenty-four hours, the water in the basin should
be poured into an evaporating-dish and gently evaporated over a waterbath. Pure sugar will crystallize from the solution. The sugar, a
crystalloid, has passed through the diaphragm; the gum, a colloid, has
remained in the dialvzer. It should be remarked that gum is wholly
uncrystallizable, and that the mixed solution of gum and sugar will not
yield crystals, but only an amorphous mass, whenl evaporated.
328. By means of this dialyzing apparatus, arsenious acid,
salts of the metals, strychnine, and other crystallizable poisons,
mineral and organic, can be readily separated from organic
fluids; and the process has the very great advantage of introducing no metal, chemical reagent, or other foreign substance
into the fluids under examination. After twenty-four hours, the
crystallizable poison, or a large proportion of it, will have been
transferred to the distilled water in the outer basin, and in this
solution it may be sought for by the application of the appropriate tests.
Exp. 130. —Dissolve 0'1 grm. of arsenious acid in about 30 c. c. of
hot water, and stir the solution into about 200 c. c. of milk,-ale, soup,
gruel, or other thick organic fluid; place the poisonous fluid in a 15
c.m. dialyzer and float the dialyzer on two litres of distilled water in
a clean basin. Allow the apparatus to stand at rest in a room where
the temperature is tolerably uniform for forty-eight hours. At the
expiration of this time, transfer the clear solution in the basin to an
evaporating-dish, without losing a drop, rinse the basin carefully with
distilled water, and add the rinsings to the contents of the dish; evaporate the solution over a water-bath (see Appendix, ~ 14) to the bulk
of 50 c. c. To one-third of this concentrated solution add a few drops
of pure chlorhydric acid, and apply Reinsch's test for arsenic (~ 325),
with due regard to the small scale on which the operation must be conducted. About 0'025 grm. of arsenious acid is the quantity which may
be expected to respond to the test by copper. Three-quarters of the




EXAMINATION FOR ARSENIC.               257
original decigramme should be transferred by diffusion through the
dialyzer in the course of forty-eight hours, and of this solution of this.
0'075 grinm. of arsenious acid we have taken one-third. The rest of
the solution is to be reserved for tests hereafter to be described.
329. When the arsenious acid must be sought in large organs
of the body, like the stomach, liver, or intestines, or in considerable quantities of disgusting semifluid materials, it is sometimes
necessary to utterly destroy the organic matters by processes
which cannot cause the loss of arsenic. Several methods may be
employed for this purpose.  1. The organic matter is gently
heated in a tubulated retort with strong chlorhydric acid, and
strong nitric acid is added from time to time. The organic matter is thus completely destroyed, with the exception of the fat.
A cooled receiver is connected with the retort to condense the
distillate from the hot mass. The fat is separated from the clear
liquid in the retort by decantation, and well washed with water;
these washings, together with the distillate in the receiver, are
added to the main bulk of the fluid. 2. Chlorate of potassium
may be added instead of nitric acid. 3. The organic matter,
after being made as fine as possible, is stirred up with water,
and chlorine gas is passed through the liquid until the organic
substances are partly destroyed and partly deposited in brovwn
flakes.
All these processes, and there are others based on like principles, are processes of combustion; aqua regia, chlorate of potassium, and chlorine are oxidizing agents of great power, as we
have already seen; they burn the carbon and hydrogen of the
organic materials as literally as the oxygen of the air burns the
coal in the grate. The arsenic also is oxidized and converted
into its highest oxide, arsenic acid.  Whenever chlorhydric acid
is used, and heat is applied, there is danger that chloride of
arsenic (AsCl3) may be formed; this chloride is a volatile body,
against whose loss precautions must be taken, by never allowing
the temperature of the fluids to rise much above 1000, and by
collecting any distillate which may be formed under circumstances
which make it possible for this chloride to be evolved.
330. All these methods of destroying the organic matters in
which arsenious acid is to be sought for are liable to one objec



258                 DETECTIoN OF ARSENIC.
tion. Considerable quantities, even kilogramnmes, of acids must
be used, if the quantity of organic substance to be destroyed is
large; chlorhydric and sulphuric acids very commonly themselves
contain arsenic, and since the liquids, which result from the destruction of the organic tissues, are finally evaporated to a very
small bulk, all the arsenic in several kilogrammes of the acids
employed, as well as all the arsenic which may have been contained in the bodily organs or fluids submitted to examination, will
be concentrated into a small cupful of liquid. It is obviously
necessary to demonstrate that the arsenic reactions cannot be
obtained from the same quantities of the same acids actually
employed, subjected to the same series of operations. The best
way is to conduct a parallel examination of normal animal
organs or fluids; in this examination the identical processes and'the same weights of the same chemical materials must be employed as in the examination of the suspected substances; if
arsenic is found in the latter investigation, but not in the parallel
examination of the normal animal substances, it will be quite
certain that the arsenic was not derived from the chemicals employed in the research.
331. When, by any of the processes above described, a clear
arsenical solution, free from organic matter, has been obtained,
the identification of the arsenic may be accomplished by many
methods, of which the two following will serve as examples:1. By preeipitation as sulphide of arsenic.-If the clear solution contains arsenic acid, it is necessary to reduce this oxide to
arsenious acid before the precipitation can be effected. This reduction may be accomplished by passing a slow stream of washed
sulphydric acid gas (~ 202) through the solution for several
hours, but may be immediately effected by saturating the solution with sulphurous acid gas, the superfluous gas being finally
expelled by gentle heating.  After the reduction  has been
effected, a slow stream of washed sulphydric acid gas precipitates
the yellow sullphide of arsenic from the liquid.
Exp. 131,.-Acidulate one-half of the liquid reserved from Exp.
130 with pure chlorhydric acid, place it in a small beaker glass and
pass a slow stream of washed sulphydric acid gas through the solution.
The delivery -tube of the gas shpuld be small and the current slow; a




MARSH'S TEST.                    259
piece of unglazed paper should be used as a cover, in order to keep
the beaker full of the gas. A yellow precipitate (As2S,) will appear,
indicating the probable presence of arsenic. When no more precipitate seems to form, stop the current of gas, ana let the beaker stand in
a warm place till the odor of the gas has nearly disappeared. Collect
the precipitate on a small filter (see Appendix, ~ 14), wash it thoroughly with water, and dry it.
cvp. 132.-Mix intimately the dry precipitate obtained in the last
experiment with its bulk of dry carbonate of sodium and its bulk of
dry cyanide of potassium, and introduce this mixture into a hard glass
tube (No. 5), the end of which has been closed and expanded to a small
bulb. If the precipitate stick to the filter-paper, it must be scraped
off. Warm the bulb and its contents over the lamp to expel moisture,
then wipe the tube out with a tuft of cotton on the end of a wire, and
bring the bulb to a red heat. A ring of metallic arsenic, like that of
Exp. 124, will be deposited in the tube. Preserve this metallic mirror
for subsequent study.
2. By conversion into arseniuretted hyclrogen.-When an aqueous
or acid solution containing arsenious or arsenic acid is added
to the contents of a flask in which hydrogen is being generated,
the nascent hydrogen reduces the oxide of arsenic, and there is
formed a quantity of arseniuretted hydrogen, which mixes with
the uncombined hydrogen evolved. (Compare ~ 307.) This
arseniuretted hydrogen is decomposed, with deposition of metallic
arsenic, by being passed through a red-hot tube. The undecomposed gas burns with a whitish flame; and if a cold body be held
in the flame, a spot of metallic arsenic will be deposited upon it.
(See ~ 308.) Upon these properties and reactions is based the
process for detecting arsenic known as Marsh's test.
Fig. 47.
Lxp. 133.-To a bottle prepared for generating hydrogen from
pure zinc and dilute sulphuric acid, adapt a chloride-of-calcium tube,
s2




260                      MARSH'S TEST.
and with the outer end of this drying-tube connect a tube of hard glass
(No. 4) which has been twice drawn to a fine bore and which terminates in a fine open point. (Fig. 47.) Support this long tube at
three or four points, in such a manner that the softening of the glass,
first at the point a, and then at the point b, shall not distort the tube.
By adding acid through the funnel-tube of the flask, evolve hydrogen,
and when the whole apparatus is full of hydrogen, light the gas at the
tip of the hard glass tube. By means of an efficient gas-lamp, heat
about 2 c.m. of this tube to dull redness at the point a; just beyond
the hot part of the tube, place a small sheet-iron screen, as shown in the
figure, to cut off the heat from the adjoining narrow part of the tube.
Maintain the apparatus in this condition for ten minutes, the glass
tube red-hot at one point and the hydrogen flowing steadily through
the tube and burning with a colorless flame at the point. If no deposit, or only a scarcely perceptible deposit, appears in the fine tube
adjoining the heated portion, the zinc and sulphuric acid are pure
enough for the experiment, but if a black, shining deposit appears in
the fine tube, the materials themselves contain arsenic and are, of
course, unsuitable for use in testing for this substance.
If the zinc and sulphuric acid prove to be sufficiently free from
arsenic, add to the contents of the flask a few drops of the liquid obtained by dialysis (Exp. 130). In a moment a mirror of arsenic
will be deposited in the fine tube adjoining a; when this mirror has
become large and dense, move the lamp to b, transfer the screen. and
obtain a similar mirror in the second attenuated portion of the tube;
finally, extinguish the lamp and allow the arseniuretted hydrogen to
reach the burning jet of gas at the. extreme point of the apparatus;
the white coloration of the flame will now, for the first time, be seen;
introduce into the jet a bit of cold porcelain, and obtain the characteristic black and lustrous spot of metallic arsenic; this experiment may
be repeated indefinitely and a large number of spots obtained for subsequent use. Preserve the two mirrors and a number of arsenic spots
for future study. In order to prevent the possibility of any arseniuretted hydrogen escaping into the air of the room, the jet of gas must
be kept constantly burning, and when the experiments are ended the
flask must be washed out promptly and thoroughly.
332. This method is very well adapted for the speedy and
certain detection of arsenic in green paints, such as are applied
to wall-papers, artificial flowers, lamp-shades, and the like; for in
such cases, if any arsenic is present, there is so much as to make
any traces of arsenic which may contaminate the zinc and sulphuric acid of no consequence whatever.  It is only necessary, in




DETECTION OF ARSENIC.               261
such examinations, to scrape off some of the green coloring-matter, dissolve it in dilute chlorhydric acid, and add the solution to
the hydrogen flask of the apparatus described above. Arsenic
greens instantly give enormous mirrors and spots under these
conditions.
333. In medico-legal investigations, upon whose results life
often depends, it must always be remembered that arsenic is
very widely diffused in the mineral kingdom, and that it is a
matter of great difficulty to procure reagents absolutely free
from it. The substances employed as reagents in Marsh's test
are often contaminated with it; and the acids used in destroying
organic matter may well contain arsenic enough to become visible
after the great concentration of this impurity which inevitably
occurs in the evaporation of the liquid which results from the
burning of the organic matter. The use of zinc is avoided, and
other advantages gained, by obtaining the necessary hydrogen
by the electrolysis of acidulated water. (~ 35.) When a solution of arsenious acid, acidulated with chlorhydric or sulphuric
acid, is decomposed by the electric current, the greater part of
the arsenic eliminated at the negative pole is given off in the
form of arseniuretted hydrogen, which may be examined precisely as if it were generated in Marsh's apparatus. This method
is very delicate, and seems to possess considerable advantages
over Marsh's process, but it has not yet (1870) been actually
applied in judicial investigations.
334. To describe the methods by which the analytical chemist
purifies his reagents and proves their purity, would involve descending into technical details which are unsuitable for this
manual.  Zinc and acids, pure enough for illustrative experiments, can be bought of the dealers in pure chemicals. None
but expert analysts should ever be intrusted with the chemical
investigation in a supposed case of poisoning by arsenic.  A
difficulty attending such examinations remains to be discussed
under the metal antimony, a substance which combines with
hydrogen, as arsenic does, to form a gas which is decomposed by
heat, as arseniuretted hydrogen is, with deposition of a metallic
mirror which cannot be distinguished by mere inspection from
that of arsenic. Since preparations of antimony are much em



262                 CHLORIDE OF ARSENIC.
ployed as medicines, and particularly since tartar-emetic, a salt
containing antimony, is often administered in cases of poisoning,
it is essential to find means of distinguishing between compounds
of arsenic and the analogous compounds of antimony.
335. Chloride of Arsenic.-Only one chloride of arsenic is
known, the terchloride (AsCl3), corresponding to the terchloride
of phosphorus. No quinquichloride corresponding to the quinquichloride of phosphorus is known. The chloride of arsenic is
formed by passing dry chlorine gas over finely divided metallic
arsenic placed in a retort. The combination is usually attended
with combustion; and the heat developed is sufficient to distil the
chloride over into the receiver. It may also be made by distilling a mixture of metallic arsenic and the mercury compound
called corrosive sublimate, in accordance with the following equation, in which Hg stands for mercury (Hydrargyrum):6 HgCl,2 +   2As =  3Hg2Cl12  +   2 AsC13.
Corrosive Sublimate.      Caloimel.
Terchloride of arsenic may also be procured by distilling arsenious
acid with common salt and sulphuric acid. Small lumps of fused
salt should be added from time to time to a mixture of arsenious
acid with a large excess of sulphuric acid:As,203 + 6NaC1 + 6H2SO4 = 3H110 + 2AsCl3 + 6NaHSO4.
336. Terchloride of arsenic is a dense, colorless, oily liquid,
whose specific gravity is 2.205. It boils at 1320, producing a
vapor whose density is 90'91. It evaporates freely at the ordinary temperature, producing fumes of arsenious acid. It is
highly poisonous.  The chloride is decomposed by an excess of
water into chlorhydric acid and arsenious acid, just as the chloride of phosphorus is resolved by water into chlorhydric and
phosphorous acids; this reaction is the basis of the best determination of the atomic weight of arsenic.
2AsC1, + 31120 = 6HC1 + As20,.
All the chlorine in a known weight of chloride of arsenic is converted by this reaction into chlorhydric acid; the weight of the
chlorine contained in this chlorhydric acid can be accurately determined, and the weight of the arsenic with which this quantity
of chlorine was originally combined is obtained by simple sub



SULPHIDES OF ARSENIC.                263
traction. The proportions by weight in which arsenic and chlorine combine are thus determined. In terchloride of arsenic,
as in terchloride of phosphorus, three volumes of chlorine unite
with only half a volume of arsenic vapor to produce two volumes of the terchloride vapor. Indeed all the volatile compounds
of arsenic illustrate the fact, already mentioned (~~ 305, 309),
that the unit-volume weight, or specific gravity, of arsenic vapor
is the double of its atomic weight.
337. Bromide and Iodide of Arsenic.  It is enough to say of
these two compounds that they are crystallizable solids, obtainable by the direct action of the elements upon each other, and
answering to the formulae AsBr3 and AsI3 respectively.
338. Sull2hides of Arsenic.-There' are  three well-defined
sulphides of arsenic, corresponding to the formule As2S2, As2S3,
and As2S_. The first two occur as natural minerals, reallar and
orpiment, and may also be obtained in the free state by artificial
processes; the third is known only in combination.
339. Bisu7phide of Arsenic (As2S2).-The native mineral realgar has this composition.  The compound is obtained artificially by melting arsenic with sulphur, or arsenic with orpiment
(see the next section), or sulphur with arsenious acid, in such
proportions in either case as will bring together those parts by
weight of the two elements which the above formula requires.
The commercial product is a brownish-ed opaque substance of
variable composition, generally containing free arsenious acid.
Realgar is one of the ingredients of white Indian fire, a mixture
of 24 parts of nitre, 7 of sulphur, and 2 of realgar, sometimes
used as a signal light.
340. Tersulphide of Arsenic (As2S).-This sulphide occurs native in translucent rhombic prisms of a yellow  color.  It is
obtained artificially by passing sulphydric acid gas through a
solution of arsenious acid, or an arsenite acidulated with chlorhydric acid; the sulphide falls as a bright yellow amorphous
powder, insoluble in water and dilute acids.  It melts easily, and
burns in the air with a pale blue flame; in closed vessels it may
be sublimed without change.
Under the name of orpimnent, this sulphide is used as an orange
pigment; a mixture of the sulphide with arsenious acid, called




264                  SULPHARSENITES.
king's yellow, was formerly employed as a yellow pigment.
This impure tersulphide was made by subliming 7 parts of
arsenious acid with 1 part of sulphur, a proportion of sulphur
not sufficient to convert all the acid into tersulphide. If a pattern be printed upon cotton cloth with a preparation containing
arsenious acid, and the cloth be then passed through water containing sulphydric acid, orpiment will be deposited in the fibre
of the cloth and the pattern will be brought out in orange-yellow.
We have already seen (Exp. 132) that the tersulphide of arsenic
yields a mirror of metallic arsenic when heated in a closed tube
with a mixture of carbonate of sodium and cyanide of potassium.
The sulphide is readily dissolved by a cold solution of potash,
soda, or ammonia, the oxygen of the alkali converting part of
the arsenic in the sulphide into arsenious acid, while the alkalimetal combines with the sulphur liberated; this alkaline sulphide
then unites with the undecomposed portion of sulphide of arsenic
to form a sulphur-salt, whose composition is that of an arsenite
in which the oxygen has been replaced by sulphur.
4As2S3 +  5K20 =  3(K2S,As2S3) +  2K20,AS203.
If an acid be added to this solution, no sulphuretted hydrogen is
evolved, as is generally the case when an acid is brought in contact with an alkaline sulphide, but the sulphur and arsenic recombine and are separated as tersulphide of arsenic.
3(K2S,As2S3) + 2K20,As203O  + 10C1 = 10KC1 + 51120 + 4As2S3.
341. Sulpharsenites.-The tersulphide of arsenic unites with
basic metallic sulphides in three different proportions, forming
with potassium, for example, the three salts 3K2S,As2S3, 2K2S,
As2S., and K2S,As2S3. One mode of preparing a sulpharsenite
has been mentioned in the last section; another method is to dissolve arsenious acid in an alkaline sulphydrate, in which case
one-half of the alkali is converted into arsenite:Dualistic: 2As203 + 4KIIS = K2S,As2S3 + K20,As203 + 21120.
Empirical: As203 + 2KEHS = KAsS2 + KAs0O  + H20.
The sulpharsenites of the other metals are mostly obtained from
the sulpharsenites of sodium and potassium by the method of
double decomposition. The sulpharsenites are either yellow or
red; they are obscure bodies, of no practical importance at pre



SULPHARSENIATES.                   265
sent.  They illustrate, however, two points of theoretical interest-namely, the existence of sulphur-salts which bear to
sulphides the same relation which oxygen salts bear to oxides,
and the parallelism of composition between these two classes of
salts.  We place beside each other the empirical formule of the
sulphur-salts of potassium and arsenic, and the corresponding
oxygen-salts:Sulyphuzr-salts.           Oxygen-salts.
KAsS3                   K3As03
K4As2S5                  K4As205
KAsS2                    KAsO2
342. Quinquisulphide of Arsenic (As2S,). —A sulphide of arsenic
corresponding to anhydrous arsenic acid is not known in the free
state. The quinquisulphide is known only in combination with
sulphides of the metals in sulphur-salts called sulppharseniates.
When a solution of sulphide of sodium is digested with some tersulphide of arsenic and sulphur enough to permit the formation of
the quinquisulphide, and the solution, after long standing, is concentrated by evaporation and then cooled, large colorless crystals
of sulpharseniate of sodium are obtained, which are not changed
by exposure to the air. The crystals have the composition indicated by the formula 3Na2S,As2S, + 15H20.  The sulpharseniate, 2Na2S,As2S~, may be prepared by saturating the aqueous
solution of the corresponding oxygen-salt 2Na20,As20  with sulphydric acid gas. The sulpharseniates of the alkali-metals, and
a few others, are soluble in water; but the greater number of sulpharseniates are insoluble. These insoluble salts are prepared by
mixing a solution of an alkaline sulpharseniate with a solution of
some salt of the metal whose sulpharseniate is desired. The same
parallelism is observable between sulpharseniates and arseniates
as between sulpharsenites and arsenites.




266                      ANTIMONY.
CHAPTER XVIII.
ANTIM O NY.
343. Antimony is found native, both alone and alloyed with
other metals, especially with arsenic, nickel, and silver. There
exist also a considerable number of minerals which consist of,
or contain, large proportions of the compounds of antimony with
oxygen and sulphur.
344. All the antimony of commerce is obtained from the
mineral tersulphide, Sb2S.S3 The symbol for antimony is Sb,
from the Latin name of the substance, Stibium. This sulphide
is very fusible, melting readily in the flame of a candle; it may
therefore be separated from the earthy or rocky gangue in which
it occurs by simple fusion at a low temperature.  The metal is
obtained from the sulpbide by several different processes:-1. By
adding to the melted sulphide iron nails, filings, or scraps; the
iron and the antimony change places.
Sb2S, + 3Fe = 3FeS + 2Sb.
2. By roasting the sulphide of antimony, reduced to a coarse
powder, until the greater part of the sulphur has been burnt off
and the antimony converted into the oxide; this residue is then
mixed into a paste with water, charcoal powder, and carbonate
of sodium, or some equivalent reducing flux, and heated in covered crucibles to full redness; the metal sinks to the bottom of
the crucible. 3. By fusing together a mixture of sulphide of
antimony, the scales which fall from hot iron when it is hammered (an oxide of iron), carbonate of sodium, and charcoal;
this process is a sort of combination in a single operation of the
two preceding methods. Since the sulphides and oxides of antimony and the metal itself are somewhat volatile at moderate
temperatures, it has thus far been found impossible to avoid a
considerable loss of metal during the melting, roasting, and reducing of the ore. From one-fifth to one-half of the metal is
lost, according to the skill and care of the workmen.
345. The commonest impurities in commercial antimony are




PROPERTIES OF ANTIMONY.               267
sulphur, sodium, arsenic, lead, iron, and copper. These impurities injure the antimony for many of its applications in the arts;
and the extensive use of antimonial preparations in medicine
renders the removal'of the arsenic a point of particular importance. The purification may be effected by fusing the powdered
metal, first, with a mixture of sulphide of antimony and carbonate of sodium, and, secondly, with a mixture of carbonate of
sodium and nitre. These fusions may be several times repeated;
the imnpurities are either oxidized or converted into sulphides,
and enter the slag.  Lead, however, cannot be got rid of by
these processes; this impurity is removed by fusing the antimony with oxide of antimony; the lead changes places with the
antimony in the oxide of antimony, and is converted into litharge.
346. Antimony is a brittle metal, having a bluish-white color,
a brilliant lustre, and a highly crystalline structure. The cakes
of the commercial metal usually present upon their upper surfaces beautiful stellate or fern-like markings.  Like phosphorus
and arsenic, it is dimorphous, crystallizing both in rhombohedrons
and octahedrons. The specific gravity of the metal is from 6'60
to 6'85; its atomic weight is 122. For a metal, it is a poor conductor of heat and electricity. At 450~ it melts, gives off vapors
at a low red heat, and takes fire at full redness, burning brilliantly,
with evolution of white fumes of the teroxide (Sb203).  If the
antimony is contaminated with arsenic, as is often the case, a
garlic odor, due to the presence of this impurity, may be imparted to the vapors.
Exp. 134. —Melt about 0'5 grm. of antimony by heating it on a
piece of charcoal before the blowpipe. (See Chapter XX.) Throw
the white, glowing globule into the middle of a large tray made of
coarse paper; the globule bursts into a multitude of small beads, which
fly over the paper, leaving in their trail a white, powdery oxide.
Exp. 135. —Melt a second small fragment of antimony upon charcoal as before, but, instead of throwing it from the coal, allow it to cool
there slowly. The globule will, in this case, become covered with an
efflorescence of crystals of the oxide.
The metal is not oxidized by exposure to dry or moist air at
ordinary temperatures. Nitric acid oxidizes it easily, but does
not dissolve it; the insoluble quinquioxide, or some mixed
oxide, is formed, according to the strength of the acid employed.




268                ALLOYS OF ANTIMONY.
Powdered antimony takes fire when thrown into chlorine gas,
and combines very energetically with bromine and iodine.
When finely powdered, it is dissolved by boiling chlorhydric acid,
with evolution of hydrogen; if a little nitric acid be added to the
chlorhydric, the metal dissolves easily, to form a solution of
terchloride of antimony (SbCl3). The metal, when in fine powder, is also dissolved readily by solutions of the higher sulphides
of sodium and potassium, with formation of sulphantimonites and
sulphantimoniates.
347. In spite of the strong tendency of this metal to crystallize, it can be obtained in an amorphous form by the electrolysis
of concentrated antimonial solutions. This amorphous antimony
always contains, however, 5 or 6 per cent. of terchloride of antimony and a trace of chlorhydric acid; whether these foreign
substances are retained mechanically, or not, within the mass, is
not clear. The amorphous metal has a dark steel-color, a smooth
surface, a comparatively soft texture, a lustrous amorphous fracture and a specific gravity varying fiom 5-74 to 5-83.  When
gently heated or sharply struck, the amorphous antimony suddenly manifests a great heat, the temperature rising from 150 to
230~ and upwards, and fumes of terchloride of antimony are
evolved. After undergoing this peculiar change, the metal approximates to the crystalline variety in structure, density, and
color.
348. Antimony enters into the composition of several very
valuable alloys. Type-metal is an alloy of lead and antimony,
containing about 20 per cent. of antimony. For stereotype
plates - to I of tin is usually added to this alloy. The common white metallic alloys used for cheap teapots, spoons, forks,
and like utensils, are variously compounded of brass, tin, lead,
bismuth, and antimony; for example, a superior kind of pewter
is made of 12 parts tin, 1 part antimony, and a small proportion
of copper; Britannia metal is sometimes compounded of equal
parts of brass, antimony, tin, bismuth, and lead. The value of
antimony in these alloys depends upon the hardness which it
communicates to the compounds, without rendering them inconveniently brittle.
With zinc, antimony forms two alloys having a definite crys



ANTIMONY AND HYDROGEN.              269
talline character.  The alloy containing 43 per cent. of zinc
crystallizes in silver-white needle-like prisms; it answers to the
formula Sb2Zn3.  The alloy containing 33 per cent. of zinc
crystallizes in broad plates presenting no similarity to the form
of the other alloy; it answers to the formula SbZn.   These
alloys, especially Sb2Zn,, decompose boiling water with evolution of hydrogen. The crystals of these two alloys are obtained
by the method of fusion (~ 194).  In each of these crystallized
alloys, the crystalline form may be preserved, although the proportions of the ingredients may vary considerably from the exact
atomic proportions indicated by their formulae. Thus needles may
be,obtained in which the actual proportion of antimony present
varies from 35'77 per cent. to 57'24 per cent., the exact atomic
proportion being 55'7 per cent.; and the percentage of antimony
in the plates may fall as low as 64'57, or may rise as high as
79'42, although 65-07 per cent. is the true atomic proportion.
These interesting crystalline alloys strikingly illustrate, therefore,
a principle of wide applicability, namely, that a definite crystalline form is not necessarily a guaranty of an unvarying chemical
composition.
349. Antimony and Hydrogen.-The composition of the gaseous compound of these two elements is not certainly known,
inasmuch as it has never yet been prepared free from admixed
hydrogen. When a solution of any salt of antimony is poured
into a mixture of zinc and dilute acid which is disengaging hydrogen, the antimony compound is decomposed; one portion of
the antimony, and sometimes even the whole of it, is deposited
upon the zinc, while another portion usually combines with the
hydrogen, and assumes the gaseous state. When this compound
gas is passed through a solution of nitrate of silver, a precipitate
is produced which has been found to consist of antimonide of
silver, SbAg3. Since silver is a metal which replaces hydrogen,
atom for atom, it is a natural inference that the gas which has
produced this precipitate must have the composition represented
by the formula SbH13.  This supposition derives strength from
the analogous formulae of the well-known gases ammonia, NHi3,
phosphuretted hydrogen, PH,, and arseniuretted hyrogen, AsH3
Antimoniuretted hydrogen is a colorless gas, inodorous when




270              ANTIMONIURETTED HYDROGEN.
free from arseuiuretted hydrogen, and insoluble in water and
alkaline liquids. The gas is decomposed at a red heat into antimony and hydrogen; it burns in the air with a whitish flame,
and gives off a white smoke of teroxide of antimony; when a
bit of cold porcelain is held against a burning jet of the gas, a
sooty spot of metallic antimony is deposited on the porcelain.
These reactions resemble those of arseniuretted hydrogen (~ 308).
7Exp. 136.-Dissolve 0'5 grm. of tartar-emetic (tartrate of antimony
and potassium) in about 30 c. c. of water. Add a few centimetres of
the solution thus obtained to the bottle of the apparatus represented
in Figure 48, in which hydrogen is already being generated from zinc
Fig. 48.
and dilute sulphuric acid. Antimoniuretted hydrogen will be produced,
and should be submitted to precisely the same series of operations by
which arseniuretted hydrogen was examined. (Exp. 133.) By heating the hard glass tube at a and b successively, two mirrors of antimony
will be obtained; when the gas reaches the jet without decomposition,
the white color of the flame will be observable; when a cold piece of
porcelain is pressed against the burning jet, spots of antimony will be
deposited thereon. Preserve these mirrors and spots.
Exp. 137.-Compare together the spots obtained on porcelain from
arseniuretted hydrogen (Exp. 133) and from antimoniuretted hydrogen
(Exp. 136). 1. The arsenical spot has a metallic lustre, and a brown
color, when thin; the stain of antimony has a feeble lustre, and is
smoky-black. 2. The arsenical stain disappears readily on the application of a heat below redness; the stain of antimony is volatile only
at a red heat. On account of the comparative want of volatility which
characterizes the antimony deposit, the mirrors of antimony obtained
in the glass tube (Exp. 136) are always deposited nearer the heated
portion of the tube than the arsenic mirrors are. 3. The arsenical
stains may be distinguished, moreover, from the antimonial stains by
means of a solution of "' chloride of soda " (a mixture of hypochlorite




TESTING FOR ANTIMONY.                  271
of sodium with chloride of sodium, prepared by mixing a solution of
chloride of lime with carbonate of sodium in excess, and filtering);
this solution, which is analogous to, and indeed may be replaced by,
a solution of chloride of lime (~ 120), immediately dissolves arsenical
spots, but leaves antimonial spots unaffected for a long time. For the
application of this test it is convenient to produce some spots on the
interior of a concave bit of porcelain.  4. On warming an arsenic
spot with a drop or two of aqua regia, and evaporating to dryness, a
slight residue of arsenic acid is left, recognizable by its ready solubility
in a drop of water; if to this drop of arsenic acid solution a drop of
anlmonio-nitrate of silver be added, a brick-red turbidity, due to the
formation of arseniate of silver, will be produced. This ammonionitrate of silver is prepared by adding exactly ammonia enough to a
solution of nitrate of silver to redissolve the precipitate which forms at
first. The antimony spot treated in the same way yields no such red
precipitate. 5. An antimony stain will dissolve readily in a few drops
of a solution of sulphydrate of ammonium which has become yellow
by keeping; when such a solution is evaporated to dryness, a bright
orange stain remains. The arsenical stain, on the contrary, is not perceptibly affected by the yellow sulphydrate of ammoniuml solution,
unless heat is applied.
Exp. 138.-Connect the tube of hard glass in which two arsenic
mirrors were formed, in Exp. 133, with a sulphuretted-hydrogen-generator (Appendix, ~ 19), interposing between the tube and the generator a suitable drying-tub, or bottle filled with chloride of calcium;
then transmit through the tube a very slow stream of sulphydric acid
gas, and heat the mirrors with a small gas-flame, proceeding from the
outer to the inner border of the mirrors, in the direction opposite to that
of the gas current.
Repeat the same process with the tube containing the antimony
mirrors obtained in Exp. 136.
Yellow tersulphide of arsenic is formed in one case, and orange-red
or black tersulphide of antimony in the other. When both metals are
present in one mirror, the two sulphides appear side by side, the sulphide of arsenic as the more volatile lying invariably beyond the
sulphide of antimony.
Exp. 139.-Transmit through the tube which contains the sulphide
of arsenic a stream of dry chlorhydric acid gas (~ 95), without applying heat; no alteration will take place in the yellow sulphide.
Transmit the same gas through the tube containing the sulphide of
antimony; the sulphide of antimony will immediately disappear. If
the gaseous current be then passed through some water, the presence




272                TEROXrDE OF ANTIMONY.
of antimony in the water can be demonstrated by means of sulphydric
acid (~ 210).
When both sulphides are present at once, the chlorhydric acid attacks and removes the sulphide of antimony, while the sulphide of
arsenic remains behind. A drop or two of ammonia-water, drawn into
the tube, will then dissolve the sulphide of arsenic. This solubility in
ammonia distinguishes the yellow sulphide from sulphur itself, with
which it might otherwise be sometimes confounded.
Antimony and Oxygen. —Antimony forms two well-marked
oxides, analogous to the oxides of arsenic, the teroxide or antimonious acid, Sb203, and the quinquioxide or antimonic acid,
Sb203; a compound of these two oxides Sb203, SbOj = 2Sb204,
is sometimes recognized as a distinct oxide under the name of the
quadroxide.
350. Tero.ide of Antimony.-This oxide occurs as a natural
mineral, called White Atntimony or Antimony Bloom1.  Like
arsenious acid, it is dimorphous, crystallizing in rhombic prisms
belonging to the trimetric system, and also in regular octahedrons.
The artificial, as well as the native, teroxide is dimorphous.
Antimonious oxide is produced when antimony is burnt in. the
air, or heated to full redness in imperfectly covered crucibles.
The easiest mode of getting it is to heat the tersulphide (Sb2S,)
with strong chlorhydric acid as long as sulphydric acid continues
to escape, and pour the resulting solution of the terlchloride
(SbCl3) into a boiling solution of carbonate of sodium:
2SbCl3 + 3Na.CO = Sb:03 + 6NaCl + 3CO2.
If the solution of carbonate of sodium be cold or only warm
instead of boiling, a hydrate of the teroxide is precipitated;
Sb203,H,20 = 2SbHO2.
Antimonious oxide is white or grayish-white at ordinary temperatures, but turns yellow when heated. It melts below a red
heat, and sublimes when raised to a higher temperature in a
closed vessel. When heated in the air it is partly converted into
antimonic acid. It is readily reduced to the metallic state by
ignition with hydrogen, charcoal, or potassium.  Teroxide of
antimony dissolves sparingly in water, but freely in strong chlorhydric acid; it also dissolves in a hot solution of tartaric acid, or
of acid tartrate of potassium (cream of tartar). The solution




ANTIMONTATE OF ANTIMONY.               273
obtained in the latter case contains the tartrate of antimony and
potassium (C,I4KSbO,), commonly called tartar-emetic.  Ordinary nitric acid does not dissolve the teroxide; but fuming nitric
acid and fuming sulphuric acid both dissolve it, forming solutions
which ultimately deposit shining scales of a nitrate in the one case
and a sulphate in the other.
It is obvious, from these facts, that this oxide of antimony differs from all the oxides which we have heretofore studied, in that
it is capable of reacting upon strong acids in such wise as to form
salts wherein the antimony plays very much the same part which
lead plays in nitrate of lead PbN2O6 (Exp. 42), or calcium in
CaSO, (p. 88). This truth is expressed in technical language
when we say that the teroxide of antimony is capable of acting
as a base; the oxides heretofore studied have either been acids,
like the oxygen acids of the chlorine and sulphur groups, of
nitrogen, phosphorus, and arsenic, or they have been indifferent
bodies not inclined to form definite, stable compounds by union
with other substances.
But if, on the one hand, teroxide of antimony is thus sometimes a base, on the other it also acts as a feeble acid. The artificial teroxide dissolves readily in solutions of caustic potash and
soda, forming very unstable antimonites, which are decomposed
by boiling, or mere evaporation. These antimonites are analogous to the arsenites; but it is to be observed that arsenious acid
is not only a stronger acid than antimonious, but that, unlike
antimonious oxide, it never plays the part of a base.
351. Antimoniate of Antimony or Quadroxide of Antimony
(SbO,).-This oxide occurs as a native mineral.  It may be
prepared artificially by heating strongly the quinquioxide (Sb2O,),
or by roasting the teroxide or the tersulphide, or by treating powdered antimony with an excess of nitric acid. As thus prepared,
it is white, infusible, and unalterable by heat, slightly soluble
in water, more soluble in chlorhydric acid, and easily resolvable,
by boiling with a solution of cream of tartar, into antimonious
oxide and antimonic acid. 2Sb,,4 = Sb20,Sb20O.  The oxide
may therefore be regarded as a compound of the two other
oxides of antimony; but it is sometimes considered a distinct
oxide on the ground that it yields by fusion with caustic potash,
T




274                   ANTIMONIC ACID.
or carbonate of potassium, an amorphous saline mass whose
composition answers to the formula K20,Sb204.  This salt itself.
however, if such it be, can be regarded as a mixture of an antimonite and an antimoniate:2(K20,Sb204) = K2O,Sb203 + K20,Sb2O,.
352. Quinquioxide of Antimony or Antimonic Acid (Sb,20). —
This compound is obtained as a hydrate: —1. By treating antimony with nitric acid or aqua regia containing an excess of
nitric acid. 2. By decomposing the quinquichloride of antimony,
SbC15 (~ 354), with water:2SbCl, + 5H20 = Sb205 + 10HC1.
3. By precipitating a solution of antimoniate of potassium
(K20,Sb20, + 5H1120) with a strong acid. This antimoniate of
potassium is obtained by fusing one part of antimony with four
parts of nitre, digesting the fused mass with tepid water to remove nitrate and nitrite of potassium, and boiling the residue for
an hour or two with water; the white insoluble mass of anhydrous antimoniate is thereby transformed into a soluble hydrate,
and the solution, treated with a strong acid, yields a precipitate
of hydrated antimonic acid. The hydrated antimoniate of potassium itself is a gummy, uncrystallizable salt.
The hydrated oxide, obtained by either of these methods, gives
off its water at a heat below redness, and yields anhydrous antimonic acid as a yellowish, tasteless powder, insoluble in water
and acids. At a red heat it gives off one-fifth of its oxygen, and
is converted into the quadroxide. A boiling solution of caustic
potash dissolves the oxide.
The hydrated oxide obtained by the first and third of the
above methods is not identical with that which results from the
second process. The product of the first and third methods is
called antimonic acid; the product of the second is called metantimonic acid,; a term derived from a Greek adverb which was
used in composition to denote a change of place, condition, or
quality. Aitimonie acid forms normal salts of the composition
M20,SbO, and acid salts containing M20,2Sb20,,  while metantimonic acid forms normal salts containing 2M2O,Sb205 and
acid salts answering to  the formula 2M20,2Sb0,,; the acid




TERCHLORIDE OF ANTIMONY.             275
metantimoniates are isomeric (~ 312) with the normal antimoniates.
The metantimoniates of sodium, potassium, and ammonium are
crystalline; the antimoniates of the same bases are gelatinous
and uncrystallizable.  The antimoniates and metantimoniates
of sodium, potassium, and ammonium are the only ones which
are readily soluble in water; all other antimoniates and metantimoniates are insoluble or sparingly soluble. Normal antimoniates correspond with normal nitrates:
M2O,Sb205 = M2Sb206 = 2MSbO,.
2O,N20,   = M2N206 = 2MNO3.
Normal metantimoniates are analogous to pyrophosphates:2M2O,Sb20,  = M54Sb207.
2M20,P20, = M4P207.
Antimony and Chlorine.-Antimony forms two chlorides, a
terchloride, SbCl3, and a quinquichloride, SbCl,, both of which
have their analogues in the chlorides of phosphorus, already
studied; the terchloride is also comparable with the chloride
of arsenic. The metal unites directly with chlorine on contact
(Exp. 54), and the two chlorides are bodies of considerable stability.
353. Terchloride of Antimony  (SbClM).-This chloride is
formed when chlorine gas is passed slowly through a tube containing antimony in large excess. It may also be prepared by
distilling 3 parts of antimony with 8 parts of corrosive sublimate
(chloride of mercury), or 2 parts of the tersulphide of antimony
with 4'6 parts of corrosive sublimate:
2Sb+3RgCl12=2SbCl3 + 3Hg; Sb2S3 + 3{gC1-2=2SbCl3 + 3HgS.
The easiest method of preparing this chloride is to dissolve the
tersulphide of antimony in strong, hot chlorhydric acid, or metallic antimony in the same acid, to which a little nitric acid has
been added; the resulting liquid, in either case, after evaporation to an oily consistency, should be distilled.
At the ordinary temperature, terchloride of antimony is a
translucent yellowish substance of fatty consistency, whence its
popular name, butter of antimony. It melts at 72~ and boils at
about 2000, fumes slightly in the air, is deliquescent and highly
T2




276             QUINQUICIHLORIDE OF ANTIMONY.
corrosive.'When thrown into water, it is decomposed into chlorhydric acid and teroxide of antimony. which, however, remains
united with a portion of the chloride, forming a white powder
which contains antimony, chlorine, and oxygen, but is somewhat
variable in composition. This white precipitate is redissolved
by excess of chlorhydric acid, and the solution thus obta.ned is
the most convenient that can be used for exhibiting the reactions
of antimony. The addition of tartaric acid to this solution prevents its decomposition by water.
Exp. 140.-In a flask of about 200 c. c. capacity, heat gently 0-5
grm. of finely powdered antimony with 30 c. c. of strong -chlorhydric
acid, to which 10 drops of nitric acid have been added. When complete solution has been effected, pour a little of the chloride into water,
to demonstrate the decomposition just referred to. Evaporate the rest
of the solution to the consistency of a thick syrup; it is the butter of
antimony.
The anhydrous terchloride combines with the chlorides of
sodium, potassium, and ammonium, and certain other chlorides,
to produce crystalline saline compounds, analogous in composition to those oxygen and sulphur compounds to which the term
salt is commonly applied.
354. Quinquicldoride of Antimony (SbCl,).-This compound
is formed, wltL brilliant combustion, when finely powdered antimony is thrown into chlorine gas (Exp. 54). It may also be
prepared by passing dry chlorine over warm powdered antimony,
or over the terchloride.
Exp. 141.-Fill a hard-glass tube, No. 2, 150 c.m. long with
coarsely powdered antimony, and fit one end of the tube so charged
into a tubulature of a two-necked glass receiver, the other neck of
which is connected with a source of dry chlorine. Support the long
tube at an angle of 100 or 150 with the table, so that its open end shall
be some 20 c.m. higher than the end which enters the receiver.
Keeping the tube warm throughout its whole extent, pass chlorine
slowly and continuously into the receiver. Combination takes place
in the tube and the product flows back into the receiver, where it remains in contact with chlorine; the long layer of antimony prevents
the escape of any free chlorine.  Preserve the product in a glassstoppered bottle.
The quinquichloride is a colorless, or yellowish liquid, which
is very volatile and emits suffocating fumes. Water in small




TERSULPHIDE OF ANTIMONY.              277
proportion forms with it white deliquescent crystals, but in large
quantity water decomposes the chloride into chlorhydric and
antimonic acids.
355. We are familiar with nitric acid (N20,) as an oxidizing
agent, as a substance which readily yields some of its oxygen to
other bodies with which it is brought in contact; in a perfectly
analogous sense, the quinquichloride of antimony and its analogue
the quinquichloride of phosphorus, may be said to be chloridizing
agents of great power, for they readily impart chlorine to other
substances. These two chlorides are much used in organic chemistry for preparing chlorine compounds; thus, for example, the
compound of carbon and hydrogen called ethylene or olefiant gas,
C2I4, is converted by passing through boiling quinquichloride of
antimony into an oily bichloride, C2H4C12, known as Dutch liquid.
The quinquichloride acts as a carrier of free chlorine, being itself
reduced to the terchloride.
Terbromide and Teriodide of Antimony (SbBr, and SblT). —It
is enough to mention the existence of these compounds, formed
by the direct union of the elements.
Antimony and Sulphur.-Antimony forms two sulphides, Sb2S,
and Sb2S., corresponding to antimonious oxide and antimonic
acid, and possibly an intermediate sulphide corresponding to the
quadroxide.
356. Tersulphide of Antimony (Sb2S,).-This compound exists
in the crystalline and in the amorphous state. Crystallized tersulphide of antimony is a natural mineral called grey antimony or
antinmony-glance. It is the source of all the antimony and antimony compounds of commerce. The mineral has a lead-grey
color and a metallic lustre; it is friable and very fusible, melting
even in the flame of a candle. At a white heat it may be distilled
unchanged in closed vessels, but by roasting in the open air it is
converted into a fusible mixture of teroxide and tersulphide of
antimony. This oxysulphide, after it has been fused, constitutes
the commercial glass of antimony, which contains about 8 parts
of the teroxide to 1 part of the tersulphide; the greater the proportion of sulphide, the darker the tint of the glass.
The native tersulphide is seldom pure, being generally contaminated with lead, copper, iron, and arsenic. To obtain pure




278               TERSULPHIDE OF ANTIMONY.
crystallized tersulphide of antimony, it is best to prepare it artificially by fusing pure metallic antimony with sulphur in the
required proportions by weight. The materials must be finely
powdered and intimately mixed, and the mixture thrown by
small portions into a heated crucible. The reactions of crystallized sulphide of antimony are the same as those of the amorphous sulphide, to be presently described; but they take place less
quickly, on account of the greater cohesion of the mass.
Amorphous tersulphide of antimony can be procured by several
processes, from which we may select the two simplest:-1. The
native grey tersulphide is changed into the amorphous variety
by keeping it in the fused state for a considerable time, and then
cooling it very suddenly by throwing the vessel in which it has
been melted into a large quantity of cold water.  The product
is an amorphous mass, having a conchoidal fracture, and a less
specific gravity, but a greater hardness than that of the crystalline variety. Its color, in thin pieces, is hyacinth-red; in the state
of powder, orange-brown. 2. When sulphydric acid gas is passed
into an acidulated solution of an antimony-salt (that of tartaremetic for example), a bright orange-red precipitate of a hydrated
tersulphide of antimony is formed, which may be rendered anhydrous at a moderate heat without losing its red color.
Exp. 142.-Dissolve 2 grms. of tartar emetic in 50 c. c. of water
and add to the solution a few drops of acetic acid; pass a slow current
of sulphuretted hydrogen, from a self-regulating generator (Appendix,
~ 19), through this solution for ten minutes. The precipitate is the
hydrated tersulphide of antimony. Collect this precipitate upon a
filter and wash it.
Exp. 143.-Pour a dilute cold solution of caustic soda upon the
washed precipitate of the last experiment as it lies upon the filter, and
collect the filtrate in a test-tube; if the whole precipitate does not
shortly redissolve, pour the filtrate a second time upon the undissolved
precipitate in the filter, or use an additional quantity of soda-lye, if
necessary. There is produced a mixture of sulphantimonite of sodium
and teroxide of antimony, which is soluble in the excess of soda-lye.
2Sb2S3 + 6NaHO = 3Na2S,Sb2S3 + Sb203 + 3H20.
Exp. 144.-Pour the clear alkaline solution, obtained in the last
experiment, into two or three times its bulk of dilute chlorhydric acid.
The whole of the antimony will be thrown down again as tersulphide,




RED SULPHIDE OF ANTIMONY.               279
without any evolution of sulphuretted hydrogen, because the gas
evolved from the sulphantimonite is exactly absorbed by the dissolved
teroxide:3Na2SSb2S,  + 6HC1  =  6NaCl + Sb2S, + 3H2S.
Sb2,O  + 3H2S =  Sb2S,  + 3H20.
When hydrated amorphous tersulphide of antimony is boiled
with a solution of carbonate of sodium, it is dissolved; the filtered
liquid, on cooling, deposits a reddish-brown substance, formerly
much used in medicine, and known as icermes mineral. This substance is not a definite compound, but is a variable mixture of
tersulphide and teroxide of antimony, the latter being combined
with a small portion of the alkali.  Minute crystals of teroxide
of antimony have been recognized in this mixture by microscopic
examination. A solution of cream of tartar will dissolve out the
teroxide. leaving the tersulphide. On acidulating the cold filtered
liquid, after the deposition of the kermes, with chlorhydric acid,
a particularly bright orange precipitate of sulphide of antimony,
known as the golden sulphide, is precipitated.  Artificial sulphide
of antimony can, indeed, be precipitated of almost any color between a light orange and a blackish brown.  A vermilion-red
sulphide has found some applications as a paint.
Exp. 145. —Place in a porcelain dish 10 grms. of a solution of
chloride of antimony of about 1'35 specific gravity; add to this chloride a cold solution of hyposulphite of sodium made by dissolving 15
grms. of the salt in 30 c. c. of water; heat the dish very slowly, and
stir its contents continually so long as any precipitate separates from
the liquid. The sulphide of antimony is thrown down of a brilliant
red color. The color of the precipitate is darker in proportion as the
temperature of the mixture is higher; when, therefore, a fine red is
produced, the lamp may be withdrawn, in order to prevent the color
from growing brown. The precipitate is collected on a filter, drained
thoroughly, and then washed, first with dilute acetic acid and subsequently with water.
Sulphantimonite solutions, similar to those prepared in the wet
way, may be obtained by fusing tersulphide of antimony with dry
caustic soda or potash, or with the carbonates of sodium or potassium, and boiling the residue with water. During the exposure
to air of hot sulphantimonite solutions, a process of oxidation
takes place, whereby the sulphur set free from one portion of




280             QtUINQUIS5ULPHIDE OF ANTIMONY.
the salt converts another portion into the state of suiphantimoniate,
so that on acidulation some quinquisulphide of antimony is precipitated along with the tersulphide.
Like the tersulphide of arsenic, the tersulphide of antimony is
a sulphur-acid which unites with basic metallic sulphides to form
sulphur-salts.  The artificial sulphantimonites of the alkalies
have been alluded to above; there are many natural minerals of
analogous composition; among such may be mentioned Miargyrite,
Ag2S,Sb2S3, Bournonite, 2Pb2S,Cu2S,Sb2S3, and Berthierite, 3FeS,
2Sb2S,.
357. Quinquisulphide of Antimony (Sb2S,).-This compound,
which is not native, is made by passing sulphuretted hydrogen
through quinquichloride of antimony dissolved in tartaric acid.
It may also be prepared by acidulating the solution of the sulphantimoniate of sodium, 3Na2S, Sb25: —
3Na2S,Sb2S5 +  6HC1 =  6NaC1 +  Sb2SA  +  3H S.
The quinquisulphide is an orange-yellow, anhydrous, amorphous
powder, and is chiefly remarkable for the facility with which it
unites with the sulphides of the metals to form sulphantimoniates;
on this account this sulphide is often called sulphantimonic acid.
It is readily soluble in the sulphides, sulphydrates, and hydrated
oxides of sodium, potassium, and ammonium.
The sulphantimoniates have generally the composition represented by the general formula 3M2S,SbS, =2I-M3SbS,, analogous
to that of the tribasic phosphates 3M2O,P205 = 2MPO4.  The
sulphantimoniates of sodium, potassium, and ammonium are very
soluble in water, and crystallize with facility; those of the heavy
metals are insoluble. The latter are precipitated by adding solutions of metallic salts to a solution of the sulphantimoniate ot
sodium, keeping the latter in excess. The sodium salt may be
prepared as follows:In a wide-mouthed bottle, or other vessel which can be closed, mix
thoroughly 22 grms. of elutriated tersulphide of antimony, 26 grms.
of crystallized carbonate of sodium, 2 grms. of flowers of sulphur, 10
grms. of quicklime, slaked after weighing, and 40 c. c. of water. Let
the mixture digest at the ordinary temperature for 24 hours, with
frequent stirring; then filter the liquid, wash the residue several times
with water, and evaporate the filtrate and the wash-water in a porce



BISMUTH.                       281
lain dish or clean iron pan, until a sample yields crystals on cooling. The
formation of the salt is accelerated by boiling. The whole is then left
to cool; the deposited crystals are washed two or three times with
cold water and dried under a bell-jar over a dish of an absorbent like
quicklime or oil of vitriol. The salt is sulphantimoniate of sodium
Na3SbS4+9HO20; it forms transparent, colorless or pale yellow, regular
tetrahedrons.
CHAPTER XIX.
BISMUTH —THE NITROGEN  GROUP.
358. The metal bismuth is found chiefly in the metallic state,
but also occurs in combination with sulphur, oxygen, and tellurium. It is prepared for the arts almost exclusively from native
bismuth. The process of extracting the metal from the gneiss
and clay-slate in which it generally occurs is very simple, the
mineral being merely heated in closed iron tubes, inclined in such
a manner that the melted bismuth runs down into earthen pots,
which are heated sufficiently to keep the metal in a state of fusion.
It is then ladled out and run into moulds. The impure metal,
which often contains sulphur, arsenic, copper, nickel, and iron,
may be purified by melting it two or three times with about U
its weight of nitre.
Bismuth is a tolerably hard, brittle metal, of a grayish-white
color with a reddish tinge. When pure, it crystallizes more
readily than any other metal; by the method of fusion (~ 188)
it may be obtained in most beautiful crystals, made highly iridescent by the thin film of exide which forms on their surfaces
while they are still hot; these crystals look like cubes, but are
really rhombohedrons.  Bismuth, like phosphorus, arsenic, and
antimony, is dimorphous, presenting forms both of the monometric and hexagonal systems.  The metal melts at 264~ and
expands about - L2 in solidifying; hence its specific gravity is
greater in the liquid than in the solid state. When the metal is
subjected to strong pressure, its specific gravity, normally 9'83.




282                TEROXIDE OF BISMUTH.
has been said to become less. At a high temperature bismuth
may be distilled. Of all metals it exhibits in the highest degree
the phenomena of diamagnetism. Its atomic weight is 21.0.
Exposed to dry or moist air the metal does not alter; but when
exposed to the combined action of air and water, it is superficially oxidized. When heated in the air, it burns with a bluish
flame, forming yellow fumes.  Strong chlorhydric acid acts on it
with difficulty; sulphuric acid attacks it only when hot and concentrated; nitric acid attacks it immediately, and effects complete
solution, with formation of nitrate of bismuth and evolution of
nitric oxide.
359. Bismuth promotes the fusibility of metals with which it
is alloyed to an extraordinary extent. The most remarkable
alloy of bismuth is that known as "fusible metal."  When composed of 1 part of lead, 1 part of tin, and 2 parts of bismuth, this
alloy melts at 93~'75. Solid fusible metal, like liquid water,
undergoes an anomalous contraction by heat. It expands regularly from 0~ to 350, then contracts gradually as the temperature
rises to 550, at which point it is less bulky than at 00, again
expands rapidly to 80~, and beyond that temperature continues
expanding regularly up to its melting-point. On account of this
property of expanding as it cools while still in the soft state, the
alloy is much used for taking impressions from dies; the finest
and faintest lines are reproduced with great accuracy. It is obvious that an alloy possessing such properties must be something
more than a mere mixture of the constituent metals.
No compound of bismuth and hydrogen is as yet known.
Bismuth and Oxygen-Bismuth forms two principal oxides,
a teroxide (Bi2O,) and an acid oxide (Bi205); besides these, there
is an intermediate oxide (Bi204) which may be represented as a
compound of the other two BiO,3, Bi20,=2Bi204.
360. Teroxidce of Bismuth (Bi203).-This oxide is formed when
the metal is roasted in the air, but is best obtained by gently
igniting the nitrate or. subnitrate. It is a pale-yellow, insoluble
powder, which melts at a red heat, and is easily reduced to the
metallic state by heating it with charcoal. A white hydrate of
this oxide, Bi203,H20= 2BiHO2, may be precipitated from a salt
of bismuth by an excess of ammonia. Teroxide of bismuth com



BISMUTHIC ACID.                  283
bines with acids to form the bismuth salts; in the normal salts
one atom of bismuth replaces three atoms of hydrogen, thus: —
Bi2O, +  3(H2N206) =  3H20 +  Bi2 3(N206).
Basic salts of bismuth are also known, in which a larger proportion of bismuth is present. Some of the normal salts crystallize
well from acid solutions, but they cannot exist in solution unless
an excess of acid is present. On diluting solutions of the normal salts with water, insoluble basic salts are precipitated; this
reaction recalls the behavior of antimony solutions. The nitrate
of bismuth, Bi2 3(N20) + 9H20, is the commonest soluble salt of
bismuth; it is procured by dissolving the metal in nitric acid.
To the basic nitrate, which is -precipitated when water is added
to the solution of the normal nitrate, the formula 5Bi,20,4N205
+ 9H20 has been assigned. Bismuth salts are heavy compounds,
which are colorless unless the acid be colored; they are poisonous
in large doses.
361. Quinquioxide of Bismuth, or Bismuthic Acid (Bi205).When chlorine gas is passed through a concentrated solution of
potash holding teroxide of bismuth in suspension, a blood-red
solution is obtained, from which there soon separates a red precipitate; this substance is a mixture of hydrated bismuthic acid
and teroxide of bismuth. Cold dilute nitric acid dissolves the
oxide, but does not attack the acid. The hydrated acid gives up
its water at a temperature of 1300, and the anhydrous quinquioxide remains as a brown powder, which, in contact with acids,
parts very readily with a portion of its oxygen and falls back to
the state of teroxide. The anhydrous quinquioxide may be also
converted by a gentle heat into the intermediate oxide Bi204.
Bismuthic acid combines with caustic soda and potash, but the
compounds are decomposed by mere washing. The bismuthates
are little known, and are of interest only in so far as they go to
show the feeble acid character of the quinquioxide.
362. Terchloride of Bismuth (BiCl1).-This compound may be
obtained by heating bismuth in chlorine, or by mixing the metal
in fine powder with twice its weight of corrosive sublimate (chloride of mercury) and distilling. The same substance is produced
when bismuth is dissolved in aqua regia, and the excess of acid
evaporated. It is a very fusible, volatile, deliquescent body,




284                 CHLORIDE OF BISMUTH.
which was called butter of bismuth, from its resemblance to the
butter of antimony, long before the relationship now established
between bismuth and antimony had been recognized. The terchloride is decomposed by water into chlorhydric acid, which dissolves
a portion of the chloride, and a precipitate containing bismuth,
chlorine, and oxygen, and called oxychloride of bismuth.
3BiCl3 +  3HI20  6HC1 +  Bi3Cl303.
The same oxychloride is precipitated when a solution of nitrate
of bismuth is poured into a solution of common salt. It is used
as a pigment, and is known as " pearl-white." It may be distinguished from the analogous oxychloride of antimony by the fact
that it is insoluble in tartaric acid and in potash, both of which
dissolve the antimony compound. Terchloride of bismuth forms
crystallizable double salts with the chlorides of sodium, potassium, and ammonium. These chlorine salts are analogous in
composition to, and isomorphous with, the corresponding double
chlorides of antimony.
363. Tersulphide of Bismuth (Bi2S,). —Bismuth  glance, a
somewhat rare mineral which occurs in acicular prisms isomorphous with the native tersulphide of antimony, is a tersulphide
of bismuth. The same compound may be formed artificially by
fusing the pulverized metal with one-third its weight of sulphur.
Heated in close vessels, the sulphide evolves sulphur; heated
with access of air, it forms teroxide of bismuth and sulphurous
acid. The tersulphide is also obtained as a brown-black precipitate when sulphuretted hydrogen is passed through a solution
of a bismuth salt.
Exp. 146.-Dissolve 0'5 grm. of finely powdered bismuth in aqua
regia, in a small flask; the aqua regia should be added by small portions at a time, so as to avoid an unnecessary excess of acid, and the
mixture should be gently heated. When complete solution has been
effected, pour one or two drops of the acid solution into a tumbler full
of water, and observe the precipitation of white oxychloride of bismuth
(~ 362). To the remainder of the solution add water, drop by drop,
with constant stirring, until a slight cloudiness appears; add a drop of
chlorhydric acid to clear the solution, and through the slightly acid
liquor, thus obtained, pass a slow stream of sulphuretted hydrogen
until the solution has become so thoroughly charged with this gas
that it continues to smell of sulphuretted hydrogen even after it has




THE NITROGEN GROUP.                   285
been removed from the source of the gas. Filter the brownish-black
precipitate of tersulphide of bismuth and wash it with water upon the
filter; scrape off a portion of the precipitate from the filter, with a
smooth slip of wood, and place it in a test-tube together with a few
drops of a solution of caustic soda; it will not dissolve. Test another
portion of the precipitate in the same way with a solution of sulphydrate of ammonium; it will not dissolve. The last two reactions
establish distinct differences between the sulphides of bismuth and
antimony, in addition to their difference of color. Again filter, and
wash the undissolved sulphide and heat a little of it moderately on
platinum foil over the gas-lamp; sulphurous acid will be given ofi;
and the oxide of bismuth remains; this oxide readily melts to darkyellow globules.
364. The Nitrogen Group of Elements.-The five elements
nitrogen, phosphorus, arsenic, antimony, and bismuth, form  a
well-marked natural group of elements.  In the first place, the
elements themselves exhibit a definite gradation of properties;
and secondly, the analogy in composition and properties, manifested by the similar compounds of the five elements, is most
striking and complete.
Nitrogen is a gas, phosphorus a solid whose specific gravity
varies from 1'8 to 2'2, arsenic has the specific gravity of 5-6,
antimony of 6'7, while that of bismuth rises to 9'8. The metallic
character is most decided in bismuth, is somewhat less marked in
antimony, is doubtful in arsenic, and almost vanishes in phosphorus. All four of these elements are dimorphous, presenting
forms both of the monometric and hexagonal systems. The series
of corresponding hydrides, oxides, chlorides, and sulphides which
the elements of this group form are very perfect; they prove the
general chemical likeness of the five elements:Hyd&rides. Oxides.   Oxides.    Oxides.    Cllorides.  Sulphides.
NH3    N203        N204       N,20      NC13 (?)    P2S3
PH3    P203        Sb204      P205      PC1,         AsSP
AsH3   As203       Bi2O4     As205      AsCl3        Sb2S3
SbH 3    Sb203               Sb205      SbC1l       Bi2S3
Bi203                Bi205     BiCl3
PAS
PC1s        As S5
SbC1s       SbSA




286                 THE NITROGEN GROUP.
The first four members of the group form gaseous terhydrides,
in which three volumes of hydrogen and one atom of a nitrogengroup element are combined to form two volumes of the compound gas. We have already spoken of the similarity of chemical composition, and the gradation of properties manifested by
these four hydrides. Ammonia is a powerful base, and requires
a high temperature for its decomposition; phosphuretted hydrogen is a very feeble base, while the basic character is not perceptible in arseniuretted and antimoniuretted hydrogen. Each
of the last three hydrides is decomposed by simple exposure to
heat, phosphuretted hydrogen requiring the highest temperature,
arseniuretted hydrogen decomposing at a lower, and antimoniuretted hydrogen at a still lower degree of heat. The affinity of
bismuth for hydrogen is so feeble that it does not appear to form
a hydride.
The teroxides also show a gradation of physical and chemical
qualities. Teroxide of nitrogen (nitrous acid) is a highly volatile liquid, that of phosphorus (phosphorous acid) a very volatile
solid, that of arsenic a less volatile solid, that of antimony a
solid volatile only at a full red heat, and that of bismuth a solid
which requires an extremely high temperature for its volatilization. The teroxides of nitrogen and phosphorus form with water
strongly acid liquids; teroxide of arsenic is but a feeble acid;
teroxide of antimony is sometimes an acid and sometimes a base,
while teroxide of bismuth is a decided base. Arsenious and
antimonious acids are isodimorphous. The series of quinquioxides also shows a very marked gradation of chemical energy,
especially when the compounds which they form with the elements of water are considered. Nitric acid is a powerful acid
of intense energy; phosphoric acid is still a strong acid, but
much less incisive than nitric acid; arsenic acid has the corrosive
properties generally attributed to acids, but it is chemically a
rather less vigorous compound than phosphoric acid. The phosphates and arseniates are, however, isomorphous, and the two
acids are very much alike. In the quinquioxide of antimony
the acid character becomes comparatively indistinct; and in the
so-called bismuthic acid little remains but the name.
The terchloride of nitrogen has hardly been examined, on




CARBON.                     287
account of its extreme instability. The other four terchlorides
are all volatile substances of analogous composition, since three
volumes of chlorine and one atom of the nitrogen-group element
unite to form two volumes of the compound vapor. The boilingpoints of the terchlorides of phosphorus, arsenic, and antimony
are 740, 132~0, and 223~ respectively, while that of the terchloride of bismuth is considerably higher still. All four terchlorides
are decomposed by an excess of water.
The tersulphides of antimony and bismuth are isomorphous.
The elements of this group do not form many combinations
among themselves. They combine with hydrogen, metals and
compound radicals which replace hydrogen atom for atom, and
with the members of the chlorine group, by preference, in the
proportion of 1 atom to 3; they also combine with oxygen and
the members of the sulphur group, by preference, in the proportions of 2 atoms to 3, or 2 atoms to 5.
When the qualities of the corresponding compounds which the
members of the nitrogen group form with other elements are duly
taken into account, it will be apparent that the relative chemical
power of each element of the group may be inferred from its position in the series of elements:
N =  14,  P =  31,  As =  75,  Sb -- 122,  Bi =  210.
The chemical energy of t!.ese five elements, broadly considered,
follows the opposite order of their atomic weights.
CHAPTER XX.
CARBO N.
365. Carbon is an extremely important and a very abundant
element. All organic substances, all things which have life,
contain it. Large quantities of it occur in the mineral kingdom
as well, both in the free state and in combination with oxygen
and with other elements. The various forms of coal, graphite,
petroleum, asphaltum, and all the different varieties of limestone,




288            THE THREE VARIETIES OF CARBON.
chalk, marble, coral, and sea shells contain it in large proportion.
It is found also in the atmosphere and in the waters of the globe,
and though existing therein in comparatively small proportion, it
is an ingredient not less essential than either of their other constituents for the maintenance of the actual balance of orlgauic
nature. All vegetable life is directly dependent upon the presence of the compound of carbon (carbonic acid) which exists in
the atmosphere.
366. Three distinct allotropic modifications of carbon are distinguished, namely: —. The diamond; 2. Plumbago or graphite;
and 3. Ordinary charcoal or lampblack. There are many subvarieties of the modification last named; but their peculiarities
appear to depend chiefly upon physical conditions of aggregation,
whereas each of the three principal varieties of carbon above enumerated is really different from the other two in chemical quality
or nature.
367. The element carbon, in each of its modifications, is an infusible, non-volatile solid devoid of taste and smell. But the
several modifications differ among themselves in color, hardness,
lustre, specific gravity, behavior towards chemical agents, power
of conducting heat and electricity, and in various other respects.
All the varieties of carbon, however, agree in this, that on being
strongly heated in presence of oxygen they unite with it and
form carbonic acid (CO,2); but in the comparative readiness with
which this result is attained great differences are noticeable in
the different varieties.
Lampblack and charcoal, as is well known, readily combine
with oxygen at the temperature of an ordinary fire; they burn
easily in the air. But graphite burns so slowly in air that it is
used for making the crucibles in which the most refractory metals
are melted. (See Appendix, ~ 26.) On being heated to a very
high temperature, however, in oxygen gas, graphite slowly.undergoes combustion; and the same remark is true of the diamond.
Both graphite and diamond can be consumed by nascent oxygen,
as when heated in the condition of fine powder with a mixture
of bichromate of potassium (a substance rich in oxygen) and sulphuric acid. They can be oxidized also by heating them with
nitrate or with chlorate of potassium.




DIAMOND.                       289
368. Diamoncl.-Of this first variety of carbon, little need here
be said. The physical properties which render it so valuable, its
high refractive power as regards light, and its extreme hardness,
are familiar to all. It the hardest known substance, being capable
of scratching all other substances; the name diamond is a mere
corruption of the word adamant.
Of the chemistry of the diamond very little is known. It consists of pure or nearly pure carbon, crystallized in octahedrons
and other forms of the first or regular system; its specific gravity
is about 3'55, and its specific heat 0-1469. It conducts electricity and heat but slowly; and yet it conducts heat so much
better than glass that this property is sometimes made use of
as a test to distinguish false from real diamonds. Its refractive
power on light, as compared with that of glass or rock-crystal, is
as 2'47 to 1'6.
Chemists are as yet unable to prepare this variety of carbon
artificially; the only source of it is the natuiral mineral. It was
thought, at one time, that if there could but be devised means of
fusing carbon, crystals of the diamond modification might possibly separate out from the molten liquid as it cooled; but, at
present, all the evidence goes to show that at high temperatures,
the second modification of carbon (namely, graphite), and not
diamond, is produced. If a diamond be heated white-hot between
the charcoal points of a powerful galvanic battery, it softens, and
swells up, and, after cooling, there is found a hard black brittle
mass like the coke obtained by heating bituminous coal. So, too,
carbon is soluble in melted iron, and a portion of it crystallizes
out as the iron becomes cold, but the crystals thus obtained are
crystals of graphite and not of diamond. We can, therefore, only
surmise that diamonds crystallize at a low temperature from some
unknown solvent of carbon, or, with greater probability, that
when carbon is separated by the decomposition of some one of its
compounds it is left in the diamond condition.
The diamond is not attacked by the strongest acids or alkalies,
not even by fluorhydric acid; nor is. it acted upon by any of the
non-metallic elements, with the exception of oxygen at high
temperatures. At the ordinary temperature of the air, diamond
undergoes no appreciable change during the lapse of centuries;'U.




290                       GRAPHITE.
it appears to be well nigh indestructible, in the ordinary sense of
the term. Out of contact with the air, or in an atmosphere which
has no chemical action upon it, it suffers no alteration at the
highest furnace heat.
369. Graphite or Plumbago, sometimes called "black-lead," is
familiarly known as the material of common "lead pencils."  It
is found as a mineral in nature in various localities. It occurs
both in the form of crystals (six-sided tables belonging to the
hexagonal system) and in the amorphous, massive state. In both
forms it is always opaque, of a black or lead-gray color and metallic lustre; its specific gravity varies from 1'8 to 2'1; its specific
heat is 0'201.
It conducts electricity nearly as well as the metals, and is, on
this account, much used for coating surfaces of wood, wax, plaster,
and other non-conducting materials so as to render them capable
of conducting the galvanic current and so of receiving a metallic
film such as is deposited from solutions of the metals when subjected to the action of the galvanic current; it is an important
material in the art of electro-metallurgy.  The lustre and conducting-power of graphite go far to justify the term which has
been often applied to it, metallic carbon.
370. Graphite is very friable; when rubbed upon paper it
leaves a black shining mark, whence its use for pencils. Amorphous graphite is much morefriable than the crystalline variety;
it makes a blacker mark upon paper, and is consequently preferred for the manufacture of pencils; it is, in fact, so soft and
unctuous to the touch that it is often used as a lubricant for diminishing the friction of machinery. But in spite of all this the
particles of which the masses of graphite are composed are extremely hard,; they rapidly wear out the saws employed to cut
these masses. J3y powerful pressure the dust of plumbago can be
forced into the condition of a coherent solid similar to the native
mineral.
In the air, at.ordinary temperatures, graphite undergoes no
change; hence:its use for covering iron articles to prevent their
rusting.  Byvirtue of its greasy, adhesive quality, it is easy to
cover iron with a thin, lustrous layer or varnish of it; the common stove-polishes,~for eqa.palpe, nrce composed of powdered gra



PROPERTIES OF GRAPHITE.                 291
phite.  Even at very high temperatures it is scarcely at all
oxidized by the air; it is, moreover, altogether infusible; hence
it is usefully applied in the manufacture of a highly refractory
kind of crucible, known as black-lead crucibles or blue pots.
(See Appendix, ~ 26.) An analogous application of graphite is
seen in its use as "foundry-facings," a term applied to the infusible dust which the iron-founder sifts over his mould of sand
before pouring into it the melted metal; if the hot metal were
allowed to come directly into contact with the sand, a quantity
of the latter would melt and remain adhering to the cold metal
when the casting was taken from the mould. For this purpose
coal-dust is an inferior substitute for graphite.
371. Pure plumbago is never met with in nature; when
burned in oxygen the mineral leaves from two to five per cent.
of ashes, composed mainly of oxide of iron, together with small
quantities of silica and alumina. The presence of this impurity
is so unvarying that graphite was formerly supposed to be not
carbon, but a chemical compound of carbon and iron, a carbide
of iron; this view has now been disproved, and it is known that
the iron in the native graphite exists there merely as a mechan
ical admixture.
Soft, fine-grained plumbago, suitable for the manufacture of
the best pencils, is rare; but the coarse, foliated crystallized
variety is abundant, and this may *sily be made soft and unctuous by the action of certain oxidizing agents.
Exp. 147. - Mix 7 grms. of coarsely powdered crystallized graphite
with 0'5 grm. of chlorate of potassium in fine powder; add the mixture
to 14 grms. of strong sulphuric acid contained in a porcelain dish, and
heat the whole over a water-bath as long as yellow vapors of hypochloric acid are evolved. Wash the cooled mass with water, and subsequently dry it on the water-bath.
Ignite a fragment of the dry product upon a piece of platinum foil,
and observe the extraordinary intumescence. After the graphite has
ceased to swell up, rub a little of it upon a porcelain plate and note
the condition of exquisite softness to which it has been reduced, and
the ease with which it can be moulded by pressure into any desired
form.
372. Regarded from the chemical point of view, graphite resembles the other modifications of carbon inasmuch as it is conu2




292                   GRAPmTIC ACID.
verted into carbonic acid on being ignited in oxygen, and in that
it undergoes no alteration when heated in close vessels, but
differs materially from the other varieties of carbon in its bebehavior towards several of the oxidizing agents. When graphite
is repeatedly exposed to the action of a mixture of strong nitric
and sulphuric acids, or to that of a mixture of chlorate or bichbromate of potassium and sulphuric or nitric acid, it is converted
into a peculiar acid, called graphitic acid.
This graphitic acid occurs in thin transparent crystals, somewhat soluble in water, but insoluble in water containing acids or
salts; on being heated, it decomposes, with explosion and evolution of light. By analysis, it has been found to contain carbon,
hydrogen, and oxygen, in proportions corresponding to the complex formula C11H40,; but some chemists, who regard this body
as an analogue of an acid (Si4H405) obtained by acting upon one
of the modifications of silicon with oxidizing agents, have suggested that the atomic weight of graphite may be different from
that of ordinary carbon, and that the composition of graphitic
acid could be represented by the simpler formula Gr4H40I, in
which the symbol Gr stands for graphon-provided the atomic
weight of this graphon were assumed to be 33, instead of 12 (the
atomic weight assigned to ordinary carbon).
The graphitic modification of carbon can readily be obtained
artificially.  When charcoal is added to melted iron, the iron
takes up a considerable quality of it, and if the iron be then
left to cool slowly, a portion of the dissolved carbon will crystallize out in the form of graphite; the crystals can readily be
isolated by dissolving away their metallic envelope by means of
dilute chlorhydric or sulphuric acid.
As has been already remarked, the crystals of graphite are
six-sided plates of the hexagonal system, altogether unlike the
forms of the regular system which are seen in the diamond. In
carbon, then, as in sulphur, we have a striking example of
dimorphism. (See ~ 192.)
373. Gas-Carbon.-An interesting sub-variety of carbon somewhat similar to graphite, and standing, as it were, between it
and the ordinary modification of carbon, is obtained from the
retorts in which common illuminating gas is manufactured. It




GAS-CARnON.                       293
is known as " gas-carbon," or " carbon of the gas-retorts," and
results from the burning on of drops of tar upon the interior
walls of the retort, and the long-continued heating of the crust
thus formed.
Gas-carbon is very hard, compact, and dense; it has the metallic lustre, and conducts electricity like a metal; its specific
gravity (2-356) and specific heat (0'2036) closely resemble those
of graphite. On account of its high conducting-power, it is employed in the manufacture of galvanic batteries and of pencils
for the electric lamp; its infusibility and power of resisting
chemical agents have led to the employment, in various scientific
researches, of crucibles and tubes wrought out of it; it has also
been sometimes employed as fuel in experiments where higher
degrees of heat are needed than can be obtained from charcoal
or coke. The intense heat developed by the combustion of this
substance is referable to its high specific gravity; a very considerable weight of carbon can here be burned in a small space.
As a fuel, it has the further merit of leaving scarcely any ashes.
374. Coke and Anthracite Coal are impure sub-varieties of
carbon which, from the chemical point of view, may be classed
either with graphite or charcoal, or better between the two.
They are less like graphite, however, than gas-carbon is. Coke
is the residue resulting from the destructive distillation of soft or
bituminous coal.
Exp. 148.-Put into a tube of hard
glass, No. 1, 12 or 15 c.m. in length,       Fig. 49.
enough bituminous coal, in coarse'oVwder, to fill one-third of the tube.
Fit to this ignition-tube a large delivery-tube of glass, No. 4, and support the apparatus upon the iron
stand, as shown in the figure. Heat
the coal in the ignition-tube, and:ollect in bottles the gas which will
be evolved. This gas is a mixture of
several compounds of carbon and hydrogen; for the present, it may be regarded as carburetted hydrogen,
it is, in fact, ordinary illuminating gas.
It is inflammable, like hydrogen, but burns with a far more lumi



294                  ILLUMINATING GAS.
nous flane. It is very light withal; hence many of the experiments
described in the chapter upon hydrogen may be performed with this
gas. (See Chapter V.)
As soon as gas ceases to be given off from the coal, take the end of
the delivery-tube out of the water and when the ignition-tube has
become cold, break it and examine the coke which it contains. The
coke used for domestic purposes is obtained as an incidental product in
the manufacture of illuminating gas.
In Europe, where anthracite is lacking, immense quantities of coke
are prepared for metallurgical uses, the gas resulting from the decomposition of the coal being usually thrown away.
375. Bituminous coal is a substance of vegetable origin, which
appears to have been formed from plants by a process of slow
decay going on without access of air and under the influence of
heat, moisture, and great pressure. Like vegetable matter in
general, it is composed of carbon and hydrogen, together with
small proportions of oxygen and nitrogen, and a certain quantity
of earthy and saline substances, commonly spoken of as inorganic matter. On being heated in the air, it burns away almost
completely after a while, leaving nothing but the inorganic component as ashes. But when heated out of contact with the air
(that is to say, when subjected to destructive distillation, as in
Exp. 148), the volatile hydrogen is all driven off in combination
with some carbon, either as gas or as a tarry liquid, and there
remains, as a residue, only carbon contaminated with the inorganic matters originally present in the coal.
376. Both coke and anthracite are hard and lustrous.  As
compared with charcoal, they are rather difficult of combustion;
but in suitable furnaces they burn readily, with evolution of intense heat. Both anthracite and coke, the latter in spite of its
porosity, conduct heat readily, as compared with charcoal; hence
one reason of the difficulty of kindling them. In building a
charcoal fire, the heat evolved by the combustion of the kindling
material is almost all retained by the portions of charcoal immediately in contact with the kindling agent; but in the case of coke
or anthracite a large proportion of this heat is conducted off and
diffused throughout the heap of fuel, so that no portion of the
fuel can at once become very hot. It follows that in both lighting and feeding fires of coke or anthracite, only small portions




COKE-ANTHRACTTE —-CARCOAL.              295
of the fuel should be added at any one time, lest the kindling
material, or the existing fire, be unduly cooled. Since coke is
asually contaminated with a considerable proportion of inorganic
matter, its combustion is apt to be hindered by the accumulation of ashes and consequent exclusion of air, unless special precautions be taken.
377. Charcoal or Lampblack is commonly taken as the representative of the third or amorphous modification of carbon. This
kind of carbon can be obtained in a state of tolerable purity,
either by heating in a close vessel sugar, or starch, or some
other organic substance which contains no inorganic constituents,
or by burning oil of turpentine in a quantity of air insufficient
for its complete combustion.  A convenient way of obtaining it
is to place a vessel filled with ice-water directly in the flame of a
lamp fed with oil of turpentine, so that the combustion of the oil
shall be impeded, and that soot may be deposited upon the walls
of the vessel. In either event, however, the product is liable to
be contaminated with traces of hydrogen or of oxygen, or of both
these elements, which cannot be expelled even by the application
of long-continued and intense heat.
For such illustrations as are required in this manual, charcoal
can readily be prepared from wood in the same way that it is
made for manufacturing and domestic uses, namely by subjecting
the wood to a process of incomplete combustion.
Elp. 149.-Light one end of a splinter of dry wood and slowly push
the burning portion into the mouth of a test-tube, as shown in Fig. 50.
The portion of the splinter which remains
outside the tube and in contact with free     Fig. 50.
air will continue to burn with flame, while ~
that within the tube is either extingui shed
altogether or barely glimmers as the carbon
slowly unites with the small portion of air
which can gain access to it.
-E2P. 150.- Repeat the foregoing experiment, but, instead of the test-tube, provide
a cup of sand and slowly thrust the burning splinter into this sand.
The flame -will be'extinfugished as fast as the splinter is cut off frcMn
the air by immersion in the sand, and a residue of carbon will be thus
obtained, as before.




296               PREPARATION OF CHARCOAL.
378. Whenever wood, or any other vegetable or animal matter, is not completely consumed, there is left a residue of carbon
similar to that obtained in the foregoing experiments. Incomplete combustion in such cases is really a process of destructive
distillation, or, rather, in any combustion of wood, or of bituminous coal, there is always destructive distillation at first.  When
the splinter of wood of Exp. 149, is heated in the lamp, in order
to set it on fire, there will distil off from it, in the beginning,
certain volatile compounds of hydrogen and carbon; for wood,
like coal (~ 375), is composed of carbon, hydrogen, oxygen, nitrogen, and inorganic or earthy matters, and when exposed to
strong heat it gives off in the gaseous form the volatile elements
hydrogen, oxygen, and nitrogen, together with some carbon.
The products of the destructive distillation of the splinter will
take fire and burn, and the heat generated by their combustion
will be sufficient, not only to distil the contiguous portions of the
wood, but also to bring the residual carbon to the temperature at
which it unites with oxygen.  This kindling-temperature of
carbon, it should be remarked, is considerably higher than that
at which the volatile distillate composed of carbon and hydrogen
takes fire. Now if, as in Experiments 149, 150, the burning
splinter be removed from the air as soon as the act of distillation
has been completed, but before the combustion of the carbon has
set in, the carbon will be preserved, as has been seen. So, too,
when burning wood is extinguished by pouring water upon it;
the distillatory process has occurred and has been more or less
thoroughly completed, but the combustion of the carbon is cut
short; for the water not only excludes air but absorbs so much
heat that the temperature of the fuel is reduced below the kindlingpoint. (Compare Exp. 24.)
379. Charcoal can be obtained also by distilling wood in retorts
in the same way that we have seen that coke can be procured
from bituminous coal. (See Exp. 148.).Exp. 151.-Provide an ignition-tube and a delivery-tube similar to
those employed in Exp. 148. Fill the ignition-tube with shavings or
small fragments of wood, arrange the apparatus as in Fig. 51, and light
the gas-lamp. Collect in bottles the gas which is given off from the
wood, and test it as to its inflammability by applying a lighted match,




DISTILLATION OF WOOD.                 297
After the flow of gas has ceased, remove the end of the delivery-tube
from the water, plug it so that no air can enter the ignition-tube, and
lay the apparatus aside until it has        Fig. 61.
become cold.  Finally remove the
cork from the ignition-tube and take
out the charcoal which is contained
in it. Heat a portion of this charcoal upon platinum foil and observe
the mlanner in which it burns. It
will illustrate the fact that solid substances which are incapable of evolving volatile or gaseous matter do
not burn with flame; they merely
glow.
Exp. 152.-Pack an ignition-tube with bits of wood, as in Exp. 151,
but, instead of the ordinary delivery-tube, insert in the mouth of this
ignition-tube a cork carrying a short piece of glass tubing drawn out
to a fine open point. By means of wire, tie the ignition-tube to a ring
of the iron stand, and place it over the gas-lamp. The gases evolved
from the wood will escape through the narrow tube, and on being
kindled will burn with a luminous flame. As has been already stated
(~ 57), flame is caused by burning gas.
This experiment, as well as Exps. 148, 151, illustrates the principle
of the manufacture of illuminating gas. Upon the large scale, bituminous coal, or sometimes dry wood, is distilled in large iron or clay
tubes, called gas-retorts, and the gas evolved is freed from tar and
other offensive impurities by processes of cooling and washing with
water, and by passing it through layers of lime or oxide of iron; it is
then collected in large gas-holders, from which it is pressed through
subterranean pipes, it may be for miles.
380. For use in the arts, charcoal is prepared by both the
methods above indicated; it is manufactured both by charring or
partially burning wood with little access of air, and by methodically distilling wood in actual retorts.  The first method is
employed in countries where wood is abundant, and is carried
on in the forest itself. Logs of wood are piled up into a large
mound or stack around a central aperture, which subsequently
serves as a temporary chimney, and also for the introduction of
burning substances for firing the heap. The finished heap is
covered with chips, leaves, sods, and a mixture of moistened
earth and charcoal dust, a number of apertures being left open




298              PREPARATION OF CHARCO&L.
around the bottom of the heap for the admission of air and the
escape of the products of distillation and combustion. The heap
is kindled at the centre, and burns during several weeks. When
the process is judged to be complete, all the openings are carefully
stopped, in order to suffocate the fire, and the heap is then left to
itself until cold. Kilns constructed of brick are often used instead
of the rude heaps here described. The charcoal retains the form
of the wood-the shape of the knots and the annual rings of the
wood being still perceptible in it,-but it occupies a much smaller
volume than the wood; generally its bulk does not amount to
more than three-fourths of that of the wood, and its weight never
exceeds one-fourth the weight of the wood.
Where charcoal is prepared by distilling wood in retorts, the
liquid products of distillation, namely tar and acetic acid (" pyroligneous acid"), are saved and utilized.
381. Lamp.black. —Usually when hydrogen is removed from a.
gaseous compound of carbon and hydrogen, the carbon separates
in the form of soft flakes, called lampblack or soot. In Experiment 60, we have already seen that lampblack is formed when
hydrogen is removed from carburetted hydrogen by means of
chlorine, and we know well that oxygen is capable of producing
the same result. Hydrogen is more combustible in oxygen than
carbon; hence if carburetted hydrogen be mixed with only enough
oxygen to consume its hydrogen, and the mixture be then inflamed,
the carbon contained in it will be set free. This is the way in
which lampblack is commonly formed; a lamp " smokes " when
the supply of air is insufficient to furnish oxygen to both the carbon and the hydrogen of the oil or other combustible.
Upon the large scale, lampblack is manufactured by heating
organic matters, such as tar, rosin, or pine knots, which contain
volatile ingredients very rich in carbon, until vapors are disengaged, and then burning these vapors in a current of air insufficient
for their complete combustion. A dark-red, very smoky flame is
thus obtained; a large portion of the carbon of the combustible
does not burn, but is deposited as a very fine powder precisely
similar tc that which constitutes the black portion of common
smoke. Lampblack finds important applications in the arts as a
pigment and as the chief ingredient of printers' ink.




LAMPBLACK.                       299
Exp. 153.-Fill an ordinary spirit-lamp (Appendix, ~ 5) with oil of
turpentine, light the wick and place over it an inverted wide-mouthed
bottle of the capacity of a litre or more, one edge of the mouth of the
bottle being propped up on a small block of wood, so that some air
may enter the bottle. As the supply of air is insufficient for the perfect combustion of the oil of turpentine, a quantity of lampblack will
separate and be deposited upon the sides of the bottle.
Hydrogen kindles at a lower temperature than carbon; hence
if the flame of a burning compound of carbon and hydrogen be
cooled down below the temperature at which carbon takes fire,
lampblack will be formed, even if there be present an abundant
supply of air.
-Exp. 154.-Press down upon the flame of an oil-lamp or candle an
iron spoon or a porcelain plate in such manner that the flame shall be
almost, but not quite, extinguished. The solid body not only obstructs
the draught of air, and thereby interferes with the act of combustion,
but it also cools the flame by actually conducting away part of its
heat; the temperature is thus reduced to below the kindling-point
of carbon, and a quantity of lampblack remains unconsumed and adhering to the spoon or plate. This experiment is, of course, comparable with Exps. 133, 136, in which spots of arsenic and antimony
were obtained upon porcelain, as prcducts of the incomplete combustion of their compounds with hydrogen. As has been stated in ~ 377,
lampblack thus prepared usually contains a certain small proportion
of hydrogen compounds.
382. Charcoal is altogether insoluble in water; it is odorless
and tasteless. Unlike the diamond, it is a good conductor of electricity, but a bad conductor of heat, particularly when in the state
of powder. It is a better conductor of heat in proportion as it is
denser; when strongly heated out of contact with the air it becomes heavier, harder, and closer in texture than common charcoal, and less easily combustible, but a far better conductor of
heat and of electricity than before. The pieces of charcoal, for
example, which sometimes fall unconsumed from the bottoms of
smelting furnaces, after having passed through a long column of
intensely heated fuel, are found to be peculiarly compact and
close-grained, and to conduct heat with comparative facility.
As has been seen in ~ 376, the combustibility of a fuel is
diminished in proportion as the fuel is a good conductor of heat;




300                PROPERTIES OF CHARCOAL.
and since, as has just been stated, charcoal conducts heat the
more readily in proportion as it is denser, it follows that, other
things being equal, a given sample of charcoal will take fire
more quickly if it be light than if heavy. It will appear, moreover, from the foregoing that the combustibility of charcoal should
be less, according as the temperature at which the charcoal prepared is higher. If, for example, linen or cotton rags be carbonized at low temperatures, there will be obtained a very light
variety of charcoal, called tinder, which takes fire with especial
ease. So, too, a very light and easily inflammable charcoal is
prepared for the manufacture of gunpowder by distilling light
woods, such as willow and alder, at low temperatures. Wood
charcoal takes fire easily, as compared with coke; coke as compared with anthracite, and anthracite as compared with gas-carbon. But, inversely, when the charcoal has once been thoroughly
lighted, the intensity of the heat obtained from a fire of it will be
greater in proportion as the charcoal is more dense. As has been
already stated (~ 373), a better fire can be obtained by burning
gas-carbon than can be had from coke or charcoal; for in any
given space, if the supply of air be ample, more of the dense than
of the light fuel can, of course, be burned in a given time. The
specific gravity of charcoal is about 1'6; but an ordinary fragment
of it readily floats upon water, owing to the air within its pores;
if this air be removed, as when the fragment is powdered, the
charcoal will sink at once. (See Exp. 159.). It is infusible and
non-volatile.
383. In all its varieties, charcoal is a very important chemical
agent, chiefly because of the readiness and energy with which it
combines with oxygen at high temperatures. Most of the common processes for extracting the useful metals from their ores are
based upon the affinity of carbon for oxygen.
Exp. 155. —Mix five grammes of litharge (oxide of lead) with a quarter of a gramme of powdered charcoal; place a portion of the mixture
in an ignition-tube made of No. 3 glass, and heat it strongly in the
gas-lamp. The charcoal will unite with the oxygen of the oxide of
lead, and the compound thus formed will escape in the form of gas,
while metallic lead will remain in the tube.
This experiment is analogous to Exp. 124, where arsenious acid




CHARCOAL A BEDUCING AGENT.              301
was reduced by means of charcoal. Both experiments are typical
of the manner in which hot charcoal acts upon metallic oxides.
At a white heat it removes oxygen from its combinations with
some elements which hold it with great force, such as the oxides
of sodium and potassium and phosphoric acid. Even water is
decomposed, with liberation of hydrogen, when brought into contact with red-hot charcoal..Exp. 156. —Fill a piece of iron gas-pipe, about 35 c.m. long and 1 c.m.
or more in internal diameter, with fragments of charcoal; adapt to it
a delivery-tube, as represented in Fig. 52, and support it upon a ring
of the iron stand over one or two wire-gauze gas-lamps. Attach to
Fig. 52.
the other end of the tube a thin-bottomed glass flask half full of water
and supported upon the ring of a second iron stand. Light the lamps
beneath the tube full of charcoal, and wait until it has become red-hot,
then heat the water in the flask and cause it to boil slowly. The
steam will react upon the hot carbon in a manner which may be formulated as follows:C + H20 =  CO + 2H,
and there will be formed a mixture of a compound of oxygen and carbon called carbonic oxide, and free hydrogen. Collect the mixed gas
in bottles upon the water-pan and test it as regards its inflammability
by applying a lighted match.
A certain quantity of water may even be decomposed by thrusting
pieces of brightly glowing charcoal into water. If the experiment be
performed beneath an inverted bottle of water held near the surface
of the water-pan, a quantity of gas large enough to be inflamed can
readily be obtained.
At a red heat charcoal deoxidizes bodies which are rich in oxygen so readily that there occurs a peculiarly rapid and sparkling




302                STABILITY OF CHARCOAL.
combustion known as deflagration.  Deflagration differs from explosion only in degree; it is less violent than explosion, because
the combustion is less rapid.
Exp. 157.-Mix 10 grms. of nitrate of potassium and 5 grms. of recently ignited charcoal, both in fine powder; place the mixture upon
a brick, in a current of air, or in any place where the volatile products
of the reaction can occasion no inconvenience, and touch it with a
lighted stick or red-hot wire. The charcoal will burn violently, with
brilliant scintillations, at the expense of the oxygen contained in the
nitrate of potassium.
As a modification of this experiment, heat a couple of pieces of charcoal to redness in the fire and sprinkle upon the one a small quantity
of powdered nitrate of potassium, and upon the other a little powdered
chlorate of potassium.
384. This deoxidizing power of charcoal, above illustrated, is
exhibited only at high temperatures. At the ordinary temperature of the air, the chemical energy of charcoal is exceedingly
feeble. Charcoal is, in fact, one of the most durable of substances.
Specimens of it have been found at Pompeii and upon Egyptian
mummies, to all appearance as fresh as if just prepared; the
action of the air continued through centuries has exerted no
appreciable influence upon it. Experience has proved that many
wooden articles which are to be placed in situations peculiarly
liable to cause their decay, may be protected by charring them
superficially; the carbonization of the interior of casks destined to
hold liquids, and of those portions of wooden posts and sills which
are to be sunk in the ground, or to remain on or near the surface
of the ground, are familiar instances of this custom.
Not only is charcoal unacted upon by air or water at the ordinary temperature. but there are few chemical substances which
have any action upon it unless they be hot; neither the neutral
solvents, such as alcohol and ether, nor corrosive agents, such as
chlorine or fluorhydric and chlorhydric acids, attack it in any
way. It is slowly oxidized, however, by nitric acid, and rapidly
by perehloric acid, and it dissolves, to a slight extent, in cold
concentrated sulphuric acid. At the ordinary temperature, no
one of the elements combines with it directly; but at high temperatures it unites directly, not only with oxygen, as we know,
but with sulphur as well, forming bisulphide of carbon (CS2).




CInARCOAL ABSORBS GAS.                303
At very high temperatures, as when a powerful galvanic battery
is discharged from carbon points immersed in an atmosphere of
hydrogen, carbon combines directly with hydrogen also, and there
is formed a gas called acetylene (C2112). With nitrogen it unites
to form cyanogen (CN) when a current of nitrogen gas is passed
through a column of ignited charcoal which has previously been
charged with carbonate of potassium  by soaking it in a solution
of this salt; but it will not unite with nitrogen without the intervention of the alkaline carbonate or of some other substance.
With chlorine and the other members of the chlorine group it
does not unite except indirectly; but with several of the metals it
unites directly to form badly defined compounds called carburets
or carbides. Upon the chlorides and their analogues carbon has
no action, and the same remark is true of some refractory oxides,
such as silicic acid and oxide of aluminum, which are not decomposed by charcoal, even at a white heat; but at a strong red heat
it reduces most of the sulphates to the condition of sulphides (for
exampleCaSO4 + 2C = CaS + 2CO2),
and the nitrates, chlorates, and perchlorates to the state of carbonates2(K20,N205) +  5C =  2(K20,C02) +  3CO2 +  4N.
385. A physical property of charcoal, which is of great practical importance, is its power of absorbing and condensing within
its pores a great variety of gases and vapors; it absorbs coloringmatters also, and various other substances as well, abstracting
them from solutions in which they are contained.
-Exp. 158.-Take from the fire a live coal (charccal) as large as a
hen's egog, extinguish it by covering it up tightly in a small metallic
vessel or by covering it with sand, and wait until it has become cold:
weigh it carefully upon a delicate balance, and record the weight
Place theweighed coal in such a situation (in a damp cellar, for example)
that it shall be freely exposed to moist air during twenty-four hours
again weigh it, and note the increase. Boxwood charcoal has been
found to increase in weight 14 per cent. in the course of a single day,
and any ordinary charcoal will usually increase in weight from 10 to
12 per cent.
Exp. 159.-Take from the fire, as before, a piece of charcoal which
has been heated to full redness for some time; thrust it under water, so




304                CHARPCOAL ABSORBS GASES.
that it may be suddenly cooled, and observe that it sinks in the water
and that few or no bubbles of gas escape from its pores.
Take another piece of charcoal which has long been exposed to the
air and has not recently been heated, attach to it a quantity of sheet
lead sufficient to sink it in water, and immerse the whole in a large
beaker glass two-thirds full of hot water. The mobile water will imnmediately enter the pores of the charcoal, and a portion of the air
which had previously been absorbed by these pores will be driven out,
and can be seen escaping in bubbles through the water, chiefly from
the broken ends of the coal.
In a similar way, if a piece of old charcoal be placed in a bottle full
of water standing inverted upon the water-pan, a quantity of air will
gradually be expelled from it as the water enters its pores, and will collect at the top of the bottle.
The air can at any time be quickly removed by sinking a piece of
the charcoal, by means of lead, in water of the ordinary temperature,
placing the vessel which contains the water undler the receiver of an
air-pump and pumping out the air from this receiver; as the pressure
of the atmosphere is removed from the surface of the water a multitude of bubbles of air will be seen to issue from the pores of the coal.
To the presence of air and aqueous vapor, which has been thus
absorbed, is to be attributed the snapping and crackling of old
charcoal when it is thrown upon a hot fire. Charcoal which has
been recently made, and which has not yet absorbed air and
moisture, does not thus snap and crackle; moreover it kindles
much more easily than charcoal which has long been exposed to
the air; for from the latter a quantity of gas and vapor must be
expanded and driven out before the coal can ignite, and this expansion, being, of course, attended with absorption of heat, keeps
down the temperature of the coal below the kindling-point.
It is not necessary that charcoal be freshly made or recently
heated, in order that it may kindle readily, if only it has been
kept in closed vessels, and thus prevented from absorbing moisture. In the laboratory it is well to throw the remnants of
charcoal fires into an iron ketttle, furnished with a tightly fitting
cover, in order to have always a store of easily inflammable coal
for starting new fires in small hand-furnaces.
The absorptive power of charcoal for gases can readily be shown
directly by placing recently ignited charcoal in a small confined
volume of almost any gas.




CHARCOAL ABSORBS GASES.               305
Exp. 160.-Fill a glass cylinder of 200 or 300 c. c. capacity with
dry sulphurous acid gas (Exp. 96), over the mercury trough; take
from the fire a red-hot piece of charcoal as large as can be introduced
into the mouth of the cylinder; thrust it beneath the surface of the
mercury and hold it there for a moment, in order that it may be extinguished, then pass it up into the jar of sulphurous acid. Mercury
will rise in the cylinder in proportion as the gas is absorbed, and will
soon completely fill it.
386. As one result of the enormous condensation to which
gases are subjected when thus absorbed by coal, heat is necessarily developed; the temperature of the charcoal rises as the gas
is condensed in it, and to such an extent that heaps of recently
ignited finely divided charcoal often take fire on being exposed
to the air.
Different gases are absorbed by charcoal in very different proportions. It has been found, by experiment, that one cubic centimetre of dry, compact charcoal, such as that from boxwood, will
absorb in the course of twenty-four hours:90 c. c. of Ammonia gas.        35 c. c. of Olefiant gas.
85,,,, Chlorhydric acid gas.    9'4,,  Carbonic oxide.
65,,, Sulphurous acid gas.    93,,,,  Oxygen.
55,,,, Sulphydric acid gas.    7'5,,, Nitrogen.
40,,,, Nitrous oxide gas.     650,,, aMarsh gas.
35,,,, Carbonic acid gas.      1'8,,, Hydrogen.
As one consequence of this diversity of absorbing-power, it
follows that, from a mixture of several gases, charcoal can remove
some gases more readily than others.  Thus, when recently ignited charcoal, which has been cooled under mercury, as in the preceding experiment, is passed up into, a jar of common air, it
absorbs the oxygen more rapidly than the nitrogen. But, on the
other hand, it is found that, after the charcoal has become saturated with one kind of gas, it can still take up a certain quantity
of any other gas which may be presented to it.
387. This power possessed by charcoal of absorbing gases is
evidently a particular case of the physical force called adhesion
or capillary attraction, whose manifestations are familiar to us
in the drinking up of water by a sponge, or of oil by a lampwick.  If the charcoal be moistened with water (that is to say,
X




306           COMBINATION THROUGH CHARCOAL.
if its pores be clogged by the interposition of a liquid), its absorbing-power for gases will be largely diminished.
This subject is chiefly interesting to chemists because of its
intimate connexion with the power of causing various gases to
combine with one another, which is possessed by charcoal as well
as by finely divided platinum (~~ 224, 240) and several other substances. In the same way that spongy platinum causes hydrogen
and oxygen, or sulphurous acid and oxygen, to unite with one
another, chemical combination ensues when mixtures of various
gases are brought into contact with charcoal. Though the absorbent power of platinum, as regards gases, can hardly be compared with that of charcoal, its power of causing combination is
very much greater; platinum appears to possess, moreover, in a
marked degree, the faculty of attracting and attaching to itself
small quantities of most gases.
Charcoal can be made more efficient as an agent for causing
combination, by covering it with a film of platinum. If coarsely
powdered charcoal be thoroughly impregnated with bichloride of
platinum by boiling it for some time in an aqueous solution of
this salt, and if it be then heated to redness in a close crucible,
in order that the platinum salt may be decomposed, there will be
left a residue of platinum everywhere attached to the sides of the
pores of the coal. Such platinized coal is very effective, both as
an absorbent of gases and as an agent for producing combination.
But even by itself charcoal has a very decided influence upon the
combination of gases.
If recently ignited charcoal be allowed to become charged with
dry sulphydric acid gas and then introduced into an atmosphere
of oxygen, the elements of the sulphydric acid will combine with
the oxygen so quickly that an explosion ensues, and both water
and sulphurous acid are produced.
If the coal charged with sulphydlic acid be brought into contact
with air, instead of pure oxygen, combination will occur as before,
only more slowly; in this case, however, the hydrogen alone will
be consumed, the sulphur of the sulphydric acid being set free in
the solid state.
Charcoal is much employed as a disinfecting agent.  It is
capable of removing many offensive odors from the air-such, for




CHARCOAL A DISINFECTANT.               307
example as the fetid products given off during the putrefaction
of animal and vegetable substances. Animal matter in an advanced stage of putrefaction loses all offensive odor when covered
with a layer of charcoal; and the flesh of a dead animal buried
beneath a thin layer of charcoal will gradually waste away and
be consumed without exhaling any unpleasant smell.
Exp. 161.-Place a small quantity of powdered charcoal in a bottle
containing sulphydric acid gas and shake the bottle. The odor of the
sulphydric acid will quickly disappear. In the same way, an aqueous
solution of sulphydric acid (Exp. 86) can be deodorized by filtering it
througrh a layer of charcoal.
Exp. 162.-In a shallow open basket, or in a box through the bottom of which numerous holes have been bored, spread a layer of
coarsely powdered boneblack about 5 c.m. thick, placing a sheet of
filtering-paper below, if need be, to prevent the powder from sifting
through the holes in the box. Place the body of a rat or of some other
small animal upon the layer of boneblack, and pour on more boneblack until the rat is covered with a layer of it about 5 c.m. deep.
Hang up the basket or box in a warm room, so that air may have free
access to it, and leave it at rest. After the lapse of several weeks
it will be found, on examination, that all the putrescible portions of
the animal have disappeared, and that nothing is left but a mass
of hair and bones; but in the interim no odor will have been detected arising from the decomposing animal, excepting a faint odor
of ammonia.
In the same way, water can be preserved untainted in casks which
have been charred internally; and the quality of some kinds of wine is
improved if it be stored in such casks.
In all these cases the use of charcoal as a disinfectant depends
not merely upon its mechanical ability to absorb offensive gases,
but also and mainly upon the fact that the absorbed gases are
chemically destroyed within the pores of the coal by the oxygen
which is sucked into these spaces from the air. The purifying
action depends upon oxidation, upon the burning up of the offensive gases as fast as they are formed. The charcoal is in no sense
an antiseptic, or preservative agent proper to prevent decay; on
the contrary, it actually hastens the destruction of putrescible
organic matters. Under ordinary circumstances, while in contact with the air, the pores of charcoal are, of course, always
charged with oxygen by virtue of their absorptive power. Whenx2




308               CHARCOAL REMOVES COLORS.
ever, therefore, any new gas is dragged in, and forced into intimate contact with this oxygen, it is precisely as if the gas had
been carefully collected and then subjected to the action of some
corrosive chemical agent. A great merit of charcoal as a disinfectant is, that it constantly draws in to destruction the offensive
matters around it; pans of charcoal placed about a room (the
wards of a hospital, for example) the air of which is offensive,
soon remove the unpleasant smell. Sieves of charcoal, placed
across the air-vents of sewers in such manner that the out-going
air may be filtered through the charcoal, are found to be most
efficient instruments for destroying the noxious effluvia which
commonly escape from these openings. In this case, where a
current of air is constantly passing through the charcoal filter,
the latter will preserve its efficiency for an indefinite length of
time, if only it be kept dry; for the action of the coal consists
merely in bringing about oxidation and destruction of the offensive gases of the sewer, and as fast as one portion of these is consumed a new portion can be taken in to destruction.
388. Charcoal not only destroys odors, but it removes colors
as well; and for this purpose it has long been employed in the
purification of sugar and of many chemical and pharmaceutical
preparations. Almost any coloring-matter can be removed from
a solution by filtering the liquid through a layer of charcoal.
1Exp. 163.-Provide four small bottles of the capacity of 100 or 200
c. c., and place in each of them a table-spoonful of boneblack (~ 389);
into the first bottle pour a quantity of the blue compound of iodine
and starch obtained in Exp. 69, into the second a decoction of cochineal, into the third a diluted portion of the blue liquor obtained by
dissolving indigo in Nordhausen sulphuric acid, and into the fourth a
solution of permanganate of potassium, enough of the solution being
taken in each instance to nearly fill the bottle. Cork the bottles and
shake them violently, then pour the contents of each upon a filter (see
Appendix, ~ 14), and observe that the filtrate is in each instance colorless or nearly so. In case the first portions of the filtrate happen to
come through colored, they may be poured back upon the filter and
allowed to again pass through the coal.
In the purification of brown sugar the coloring-matters are removed
in a manner similar to the foregoing, the colored syrup being filtered
through layers of boneblack. Many crystallizable organic acids and
alkaloids are purified in the same way in the chemical laboratory. But




CHARCOAL REMOVES COLORS.               309
it must never be forgotten that charcoal can absorb manv other substances besides coloring-matters; sulphate of quinine, for example,
is removed from its solutions, to a very considerable extent, by charcoal; and the same remark applies, with perhaps still more force, to
strychnine. The bitter principle of the hop, " lupulin," may be entirely removed from ale bj filtering the latter through bone-black.
The removal of metals, like gold and silver, from their dilute solutions,
by means of charcoal, appears to be a phenomenon of the same order.
In all these cases where coloring-matters, and the like, are removed from solutions, the action of the coal appears to depend
in the main directly upon the physical property of adhesion, the
subsequent oxidizing action being here far less clearly marked
than in the instances previously studied (~ 387) where gases are
acted upon.  Much of the absorbed color or other matter will
usually be found attached to the surfaces of the coal, undecomposed and unaltered. Thus, if the coal which has been charged
with a solution of indigo in sulphuric acid, in Exp. 163, be
digested in a solution of caustic soda, the latter will dissolve the
indigo and remove it from the coal; the alkaloids above mentioned, which have been removed from their aqueous solutions by
means of charcoal, can be again recovered by boiling the coal in
alcohol; and the metals can be dissolved again by means of strong
acids.
When employed to remove coloring-matters, charcoal soon becomes saturated with the color and ceases to absorb any more of
it; if the spent coal be then collected and redistilled, it will be
found that it has regained but little of its decolorizing-power; for
its pores are filled with charcoal resulting from the carbonization
of the absorbed coloring-matters and this charcoal is not porous,
but, on the contrary, compact and glistening, like the charcoal
obtained from sugar or glue, and is almost entirely destitute of
decolorizing-power.  In order to revivify their coal, the sugarrefiners digest it in chlorhydric acid, allow it to ferment in order
that the absorbed matters may be decomposed, and, finally, after
washing and drying, subject it to distillation.
389. As obtained from different sources, charcoal exhibits very
different degrees of decolorizing-power; but of the varieties commonly met with and to be procured in commerce, boneblack is
the most efficient. Boneblack is prepared for the use of sugar



310           DECOLORTZING-POWER OF CHARCOAL.
refiners by subjecting bones to destructive distillation in iron
cylinders and carefully cooling the charcoal out of contact with
the air.
Exp. 164.-Repeat Exp. 148, p. 293, but instead of bituminous coal,
charge the ignition-tube with coarsely-powdered bone. Distil as long
as gas is evolved, then remove the deliveryv-tube from the water, plug
it by, means of a bit of caoutchouc tube and a glass rod, and wait until
the ignition-tube is cold, before opening it to examine the boneblack.
Dry bones contain only about one-third their weight of organic
matter, nearly 66 per cent. of them consisting of phosphate of
calcium, in the interstices of which the animal matter is distributed; hence it happens that the charcoal obtained by distilling
bones is in an exceedingly porous and divided condition. The
decolorizing-power of bone-charcoal may, moreover, be increased
by digesting it in dilute chlorhydric acid, which dissolves out a
portion of the phosphate of calcium and leaves the charcoal even
more porous than before.
But because boneblack is most commonly employed for decolorizing, it must not therefore be inferred that it is the most
powerful decolorizer of the several varieties of charcoal. A more
efficient coal can easily be prepared by mixing nitrogenized organic
matters, such as blood, hoofs, horns, or scraps of leather, with
carbonate of potassium, distilling the mixture, and finally leaching
the product with water. Charcoal of peculiar decolorizing-power
may be prepared by igniting a mixture of 4 parts of fresh blood
and 1 part of carbonate of potassium in an iron crucible, so long
as vapors are evolved, then washing the product with water and
boiling- it with chlorhydric acid, and again washing, drying, and
igniting in a close vessel. A solution of horn-filings in caustic
potash, evaporated to dryness and ignited, also yields a very
effective charcoal. As with boneblack, the efficiency of this coal
prepared with caustic or carbonated potash appears to depend
upon the minute subdivision of its particles and the porosity resulting from the mixture of the inorganic matter.  The decolorizing-power of the charcoal obtained from vegetable substances
can be in like manner increased by mixing these substances with
lime or clay before the carbonization.  A mixture of 100 parts
pipe-clay, stirred up with water to the consistence of a thin paste,




DECOLORIZrNG-POWER OF CHARCOAL.             311
20 parts of tar, and 500 parts of powdered bituminous coal yields a
charcoal which decolorizes almost as well as boneblack. It is,
however, no very easy matter to determine which, of a number
of varieties of charcoal, is the most powerful decolorizer, since
the decolorizing-power differs with the nature of the colnringmatter as well as according to the quality of the charcoal. It
has been found, for example, that with
THIE RELATIVE DECOLORIZINGPOWER
EQUAL WEIGHTS OF CHARCOAL 
PREPARED BY           Upon a solution a Upon a solution o
indigo in sulphu- Upon a solutin of
ric acid $was    brown sugar was
Igniting a mixture of blood and carbonate of potassium....         50            20
Igniting a mixture of blood and carbonate of calcium..        18            11
Igniting a mixture of blood and
phosphate of calcium....      12            10
Igniting a mixture of glue and carbonate of potassium...    36            16
Igniting acetate of sodium...   12             9
Digesting boneblack with chlorhy-     1'9           1-3
dric acid....       1.  19        1
Ordinary boneblack..       1             1
390. Compounds of Carbon and Hydrogen are exceedingly
numerous. There are many different series of them, each member of either of these series differing but slightly from its next
neighbors in composition and properties. But all these substances
are commonly classed as organic compounds; they constitute a
special branch of chemical science called organic chemistry, and
are not discussed in works which treat of all the elements commonly met with, without special reference to any one.  This subdivision of the general subject is resorted to merely for convenience;
there is no inherent reason why the compounds of hydrogen and
carbon should not be studied in this manual as well as the compounds of hydrogen and oxygen. But since carbon is capable of
uniting with hydrogen, oxygen, or nitrogen, or with two of these
elements, or with all three of them, in the most varied proportions, there are formed so many different compounds, that it has
been found advantageous to study them by themselves.




312                 ORGANIC CHEMISTRY.
The best definition of the so-called organic chemistry which
can be given to-day, is, that it is the Chemistry of the Compounds
of Carbon. The department of organic chemistry has grown out
of ordinary chemistry solely because of the fact that the compounds of carbon with hydrogen, oxygen, and nitrogen are more
numerous, and often of more complex composition, than the compounds formed by any of the other elements. These compounds
of carbon with hydrogen, and with the other elements, are all
definite chemical compounds conforming to the law of multipleproportions (~ 76); but they count by thousands, and the mere
enumeration of their names and properties would fill a volume.
391. As examples of the series of compounds of carbon and
hydrogen, to which allusion was just now made, the following
lists may be given:Boils at o C.      Boils at o C.       Boils at ~ C.
CH2
CH                 C2114
C2R6               C3HO                C6H6.. o80
C3H1. -30~       C4,H..  50     C-H1.. 110~
C4H10o      00     CHo.. 350        C8Hlo.. 140~
C5H12.    30~     C6H2.     65~      CgH12 ~. 170~
C6H1,4 ~.   600    CH,,.   950
C7116.    900     CH16i.. 1250
The first series of the compounds above formulated is found
in petroleum, and in the oil known as coal-oil, obtained by distilling highly bituminous coals and shales at comparatively low temperatures; the petroleum used for lighting contains all these
different compounds, together with others of the same class.
The second series may generally be obtained by the destructive
distillation of various organic compounds; and the members of
the third series are found in the most volatile portion of coal-tar
-the tar obtained by distilling bituminous coal at high temperatures, as in the manufacture of illuminating gas.
It will be observed that there is a constant difference of one
atom of carbon and two atoms of hydrogen (CH2,) between any
two contiguous compounds enumerated in the lists. The boilingpoints of the several members exhibit a constant difference of 300
for each increment of CHI2, as we go down the lists, so far as has
been determined; and so with all the other physical properties,
such as specific gravity, mobility, expansibility by heat, and the




IfOMOLOGOUS SERIES.                  313
like; the intensity of these properties increases or decreases regularly, in a constant ratio, as we pass from one member of the
series to the members next in order.  Such series are called
"homologous" series (having the same proportion); they are
clearly analogous to the groups, families, or series of elements
which we have already studied in chlorine, bromine, and iodine
(~ 152), in oxygen, sulphur, selenium, and tellurium  (~ 257), in
nitrogen, phosphorus, arsenic, antimony, and bismuth (~ 364), and
are now proceeding to study in carbon, boron, and silicon. Just
as in these groups of elements the student has seen a true serial
arrangement of the different members, and has observed that the
different terms of each series differ from one another in atomic
weight and exhibit parallel differences in the intensity of their
properties, so here in each of the homologous series of hydrocarbons there have been observed similar constant differences.  In
the series of hydrocarbons, we know that there is a constant difference of composition of CH2; and this difference of composition
we believe to be at the bottom of the constant difference of physical properties; but to the cause of the constant differences in the
homologous series of elements we have, as yet, no clew.
The power of arranging numerous allied compounds into
groups or series, like those enumerated upon page 312, has been
of great service to chemists in facilitating the study of the compounds of carbon. For every such series a general algebraic
expression has been devised which serves as the name of the
series. Like any other name, this concise expression brings
before the mind of the chemist the general properties of the
series. Thus the expression CnHn,, + 2 is general for the first
series above mentioned, CnHT2 for the second, and Cn + 3H2. for
the third.
In this manual we shall describe only one of the compounds of
carbon and hydrogen, a compound which occurs ready formed in
nature.
392. Marsh-gas, or Light Carburetteci Hydrogen (CH4), is a
permanent gas which constitutes a large proportion of the ordinary illuminating gas obtained from coal. It is disengaged in
large quantities from some sorts of bituminous coal, even at the
ordinary temperature of the air, and more rapidly at higher tem



314               PREPARATION OF MARSH-GAS.
peratures. In coal minles the gas thus given off is known as
"' fire-damp;" by mixing with air, in the galleries of badly ventilated mines, it forms explosive mixtures, which frequently occasion frightful accidents when ignited through carelessness.
The gas is evolved also, in large quantities, from the mud at the
bottom of stagnant pools; it is one of the products of the putrefaction of vegetable matter under water, where the supply of air
is insufficient to oxidize the matter completely to carbonic acid
and water; hence the name marsh-gas.
In hot weather marsh-gas can be obtained by thrusting a pole into
the mud at the bottom of a pond and collecting the bubbles of gas as
they rise, by holding over them  an inverted bottle full of water.
When the bottle has been filled with gas it should be corked tightly
under water, and carried to the laboratory for examination. The gas
thus obtained is contaminated with nitrogen and with a large quantity
of carbonic acid, these gases being set free, together with marsh-gas,
during the process of putrefaction. Before the gas thus collected will
burn, it is usually found necessary to remove from it the carbonic acid;
this can be done by pouring into the bottle a small quantity of a solution of caustic soda, or some milk of lime closing and shaking the
bottle, and finally removing the stopper under water. The bottle may
then be placed upright and a lighted match applied to the gas; it will
take fire and burn with a blue flame, the size of which may be increased by pouring water into the bottle so that the gas shall be driven
out into the air.
The gas may be prepared artificially as follows:
Exp. 165. Mix together 2 grins. of
Fig.              crystallized acetate of sodium,4 grms.
<C~-                       of caustic soda, and 8 grmins. of slaked
lime. Heat the mixture gently upon
an iron plate, until all the water of
crystallization of the acetate has
been expelled and the mass has become dry and friable.  Charge an
— ~       ignition-tube 20 c.m. long with the
dry powder, heat it above the gaslamp, and collect the gas at the waterpan, as shown in Fig. 53. Carburetted hydrogen is evolved from the mixture at a temperature below
redness, and a residue of carbonate of sodium is left in the ignitiontube. The purpose of the lime is to render the mass porous and in



PROPERTIES OF MARSH-GAS.             315
fusible, or nearly infusible, so that the tube may be heated equably.
If caustic soda is heated with the acetate without addition of lime,
the tube usually breaks, even when the mixture has been dried beforehand. The reaction may be represented as follows:C2H,3NaO2   +    NaHO    =    CH4    +    Na2 CO.
Dry Acetate of     -ydrate of    Marsh- Gas.    Carbonate of
Sodium.          Sodium.                       Sodium.
393. Marsh-gas is transparent, colorless, and little more than
half as heavy as air. Next to hydrogen it is the lightest known
substance, its specific gravity being only 8. It takes fire readily
when touched with a lighted match, but is nevertheless more
difficult of inflammation than most of the other combustible compounds of hydrogen.  While free hydrogen and sulphuretted
hydrogen can be lighted by a glass rod which has been heated to
dull redness, the rod must be raised to the temperature of bright
redness, or even to a white heat, in order that it may kindle
marsh-gas. As prepared from acetate of potassium, the gas
burns with a pale yellowish-blue flame. It is rather more soluble in water than oxygen; at 0~ one volume of water dissolves
0'055 volume of it. It has never been condensed to the liquid
condition.
394. We have, thus far, observed three different proportions
in which the other elements unite with hydrogen. The members
of the chlorine group unite with hydrogen, by preference, in the
proportion of one volume to one volume; one volume of any
member of the sulphur group combines, by preference, with two
volumes of hydrogen; one atom of any member of the nitrogen
group unites, by preference, with three atoms of hydrogen. The
condensation increases in direct ratio to the increasing proportion
of hydrogen, so that, in every case, two volumes only of the resultant compound are produced. We have thus become familiar
with the fact that the space occupied by the molecule of a compound gas is always two unit-volumes (~ 258). An examination of the molecule of marsh-gas will reveal a fourth kind of
hydrogen-compound —a compound containing in the productvolume (~ 260) four volumes of hydrogen condensed. It is not
difficult to prove experimentally that two volumes of marsh-gas
yield, on decomposition, four volumes of hydrogen; but, unfor



316                ANIALYSIS OF MARSH-GAS.
tunately, it is not within our power to demonstrate, by experiment, the volume of carbon with which these four volumes of hydrogen are combined; for carbon is a solid incapable of volatilization by the intensest heat at present at our command. We are able
to determine the combining volume of each constituent of chlorhydric acid, steam, and ammonia; but we have no positive knowledge whatever concerning the manner in which carbon enters
into combination by volume. Its combining proportion by weight
can be ascertained; but it must be carefully observed that all
views respecting the volumetric composition of the very numerous
compounds of carbon are purely speculative, so far as the carbon
is concerned, until carbon shall have been actually volatilized and
its vapor weighed.
That marsh-gas really contains hydrogen and carbon may be
readily proved by bringing into play, under appropriate conditions, the strong affinity of chlorine for hydrogen. Chlorine will
set free carbon from marsh-gas, just as it liberates nitrogen from
ammonia (Exp. 64).
Exp. 166.-Fill a tall bottle of the capacity of 250 or, better, 500
c. c. with water, invert it over the water-pan, and pass marsh-gas into
it, until a little more than one-third of the water is displaced; cover
the bottle with a thick towel to exclude the light, and then fill the
rest of the bottle with chlorine. Cork the bottle tightly, and shake it
vigorously, in order to mix the gases together, keeping the bottle
always covered with the towel. Finally, open the bottle and apply a
light to the mixture. Ignition takes place, chlorhydric acid is produced, while the sides and mouth of the bottle become coated with
solid carbon in the form of lampblack, and a cloud of smoke rises into
the air. The presence of the acid may be proved by the smell, by its
reaction with moistened blue litmus-paper, and by the white fumes
which are generated when a rod moistened with ammonia-water is
brought into contact with the escaping acid gas.
Carbon and hydrogen are therefore elementary constituents of
marsh-gas. We should be glad to add the sypthetical to the
analytical demonstration, and make marsh-gas out of carbon and
hydrogen; but no means are at present known by which these
two elements can be directly combined to form marsh-gas. As
in the case of ammonia, we are obliged to rely upon the assurance of the balance that the sum of the weights of the two con



ATOMIC WEIGHT OF CARBON.             317
stituents separated from a given quantity of marsh-gas is precisely
equal to the weight of the marsh-gas submitted to analysis. The
experimental process by which this fact is demonstrated is too
complex to be profitably studied at this stage of progress, and
the fact must therefore be accepted as the result of experience.
395. It remains to show how the investigation of the composition of the product-volume of marsh-gas leads to the knowledge
of the atomic weight, or least combining weight, of the element
carbon. Marsh-gas is the compound best suited for ascertaining
the atomic weight of carbon, because experience has proved that
it contains proportionally less carbon than any other hydride of
the element. We determined the atomic weight of oxygen fromn
steam, the hydrogen compound which contains the smallest proportion of oxygen,-of chlorine from chlorhydric acid, the only
hydride of chlorine,-of nitrogen from ammonia, the hydride which
contains the smallest proportion of nitrogen. So the atomic
weight of carbon is the weight of carbon which experiment proves
to be contained in two unit-volumes of marsh-gas. By physical
determinations, the specific gravity of marsh-gas has been shown
to be 8; in other words, one unit-volume of marsh-gas weighs 8;
then two unit-volumes, or the product-volume, must weigh 16.
How much of this weight of 16 is hydrogen? This question can
be answered experimentally by ascertaining how many unitvolumes of hydrogen are locked up in twounit-volumes of marshgas.
When a series of electric sparks begin to traverse a measured
volume of marsh-gas contained in a U-tube, arranged like the
U-tube of figure 11, but without the jacket, b c, and its accompaniments, the gas is found to expand, and in a few minutes a
light deposit of carbon appears in the vicinity of the points of
the platinum wires. The decomposition of the marsh-gas proceeds slowly; so that a considerable time is required for the
execution of the experiment. If the mercury in the U-tube be
finally brought to a level in the two limbs, it will be seen that
the original volume of gas has very nearly doubled. When this
point is once attained, the continued transmission of sparks produces no further increase of the volume of the gas. The expanded
gas may be shown by the usual tests to be hydrogen.




318             ATOMIC WEIGHT OF CARBON.
This experiment is rather difficult to perform, and does not
yield perfectly exact results, for a minute portion of marsh-gas
escapes decomposition; nevertheless it establishes beyond reasonable doubt the fact that marsh-gas contains twice its volume
of hydrogen.  Two unit-volumes of marsh-gas, weighing 16,
must therefore contain four unit-volumes of hydrogen; but four
unit-volumes of hydrogen weigh 4; therefore the quantity of
carbon contained in the product-volume of marsh-gas must weigh
12, which number we admit as the atomic weight of carbon.
But it may be said, What means have we of knowing that 12
represents the least combining weight of carbon? (for the atom is
by definition the least quantity of an element which can be conceived to exist in combination).  If it were possible to demonstratethat the proportional quantity by weight of carbon which
is represented by the number 12 is just the quantity required
to fill one unit-volume when converted into vapor, we should
have the same reason for believing 12 to be the weight of one
atom of carbon that we have for considering 16 to be the weight
of one atom of oxygen, or 35'5 the weight of one atom of chlorine.
As we cannot convert carbon into vapor at all, it is impossible to
ascertain with certainty how many atoms, or how many volumes,
12 proportional parts by weight of carbon really represent; the
quantity of carbon which is combined with four atoms of hydrogen in marsh-gas may be four atoms, each weighing 3, or two
atoms, each weighing 6, or one atom weighing 12. As it is
necessary to assume some number of atoms as represented by the
proportional weight 12, that assumption which will conveniently
formulate the simplest and most familiar compounds of carbon
will be the best. Accordingly it has, of late, been assumed that
12 proportional parts by weight of carbon constitute the least
quantity of this element which is conceived to enter into chemical
combination, or, in other words, that the atomic weight of carbon
is 12. The atom of carbon, thus understood, can then fix, or combine with, four atoms of hydrogen or of any member of the chlorine group, and two atoms of oxygen or of any other member of
the sulphur group. The following substances, familiar in common life, or subjects of discussion in this chapter, may be mentioned in illustration of this principle:



TYPICAL HYDROGEN COMPOUNDS.            319
Marsh-gas... CH4        Carbonic acid... CO2
Chloroform... CHC13   Bisulphide of carbon. CS2
Chloride of carbon CC14
If it were assumed that 12 proportional parts by weight, the
relative quantity of carbon in each of the above-mentioned compounds, represent two or four atoms or volumes of carbon, instead
of one, every one of these formulae would become less simple.
There are a great multitude of compounds which contain a larger
proportional quantity of carbon than the compounds cited above;
but the greater proportional quantity is always some multiple of
12 proportional parts, or, in other words, is always all integral
number of carbon atoms each weighing 12.
In marsh-gas we have thus found a new term in the series of
hydrogen compounds. Marsh-gas is an example and type of the
hydrides richest in hydrogen; so far as we yet know, hydrogen
does not form, with any element whatsoever, any compound
whereof the product-volume contains more than four atoms of
hydrogen united with one atom of the other element. The following brief table comprehends all the principal types of hydrogen compounds, beginning with chlorhydric acid, the poorest
hydride, and passing through water and ammonia, intermediate
compounds, to marsh-gas, the hydride richest in hydrogen:Chlorhydric acid....... HC1 = 2 volumes.
Water.......0 = 2,,
Ammonia.......... HN = 2,,
Marsh-gas........ H4C = 2
Each of these types of hydrogen compounds characterizes a group
of chemical elements. The first type is characteristic of the chlorine group; the second, of the sulphur group; the third, of the
nitrogen group; and the fourth, of the group which we shall soon
know as the carbon group.
396. The compounds of carbon and hydrogen are of great
practical interest, since the flame of all ordinary lamps and fires
results from their combustion. Any allusion to their properties
at once suggests the influence which these properties exert in
the usual methods of obtaining light and heat, and necessitates a
more complete discussion of the subjects of flame and combustion
than we have had hitherto. We shall recur to this subject in a




320                  ILLUMINATING GAS.
subsequent section, after having studied the oxides of carbon.
For the elucidation of the subject of combustion, ordinary illuminating gas, which is a mixture of many hydrides of carbon, will
serve as well as any pure hydrocarbon. A brief description of
this product may here be given.
About 94-100ths of the volume of purified coal-gas consists of
a mixture of marsh-gas, free hydrogen, and carbonic oxide,-the
marsh-gas usually amounting to about one-third part of the
whole gas.  These non-luminous, or very feebly luminous gases,
serve as carriers of the six or seven per cent. of real lightproducing ingredients which are contained in the gas.  This
mixture of light-giving ingredients is exceedingly complex. The
vapor of benzole is, no doubt, one of the most important of these
ingredients.  Some of the higher homologues of marsh-gas lend
their aid; and a hydrocarbon of composition C2HT2, called acetylene, is important and very generally present. Sometimes a little
olefiant gas (C211,) is present, as well as other compounds of the
same homologous series; but the old view, that this substance
constitutes the chief luminiferous ingredient of coal-gas, is no
longer admitted.
For all practical purposes, we can here regard this mixture
of gases as carburetted hydrogen.  That it contains hydrogen,
can readily be shown by holding a cold, dry bottle over a burni'g jet of it, and observing that water is a product of the combustion; and that it contains carbon can be seen by holding in
the flame a piece of cold porcelain, and noting the deposition of
soot.  Coal-gas is only about half as heavy as air.  In many
respects it resembles hydrogen, and most of the experiments
which were performed with hydrogen can be equally, or nearly,
as well performed with this gas. The student will do well to
repeat, as an example, Exp. 24, substituting, for the hydrogen,
common gas drawn from the street mains.  In the same way he
may repeat Exp. 29, mixing 1 volume of coal-gas with from 8
to 12 volumes of air.  If, instead of coal-gas, pure light carburetted hydrogen be taken, the explosion will be most violent
with 8 to 10 volumes of air; with only 3 or 4 volumes of air,
or more than 15 volumes, the mixture is not explosive; either
too much or too little air prevents the explosion.




CARBON AND OXYGEN.                   321
Compounds of Carbon and Oxygen.-There are two of these
compounds-carbonic acid (CO,), and carbonic oxide (CO).
397. C~ro'onic 4Acid (C02) is always formed when carbon or
any of its compounds is burned in an excess of air or of oxygen
gas, or in contact with substances, gaseous, liquid, or solid, which
are rich in oxygen and yield it readily to other bodies.
~Erp. 167.-Place a live coal (charcoal) upon a deflagrating-spoon,
and thrust it into a bottle full of air, or, better, oxygen gas; cover the
bottle closely, and set it aside for examination. Or invert
an empty bottle over a burning lamp or candle, so that the
products of the combustion of the lamp may be received in
it; the bottle will immediately become clouded upon the
inside from deposition of water resulting from the combus-    \ 1/,
tion (see ~ 56), and will also be filled with carbonaceous
and other gaseous products, simultaneously formed. Cover, I
the bottle and test its contents in the manner described in
the succeeding experiment.
ExLp. 168.-Pour some lime-water (a solution of common slaked
lime in water) into the bottles filled with gaseous products of combustion in Exp. 167, and shake the bottles. The liquid will become
milky and turbid, and, whien left at rest, will deposit a white powder
(carbonate of calcium). The presence of carbonic acid can readily be
detected by means of lime-water, since this insoluble precipitate of
carbonate of calcium is formed when the two substances are brought
together.
The bottles of gas obtained in Exp. 167, will, of course, contain, besides carbonic acid, a quantity of nitrogen derived from
the air which took part in the combustion, unless, indeed, as was
suggested, the charcoal be burned in pure oxygen.  It is to be
observed that, in all ordinary cases; of combustion, whether of
wood, coal, wax, or oil, there result these same gaseotis products
-carbonic acid and nitrogen; they ascend, as invisible aerial
currents, from every well-regulated' flame or fire, and are continually issuing from the chimneys of our houses, though, in the
absence of the particles of solid carbon, or of condensed aqueous
vapor, which constitute smoke, we can see no product of the
combustion.
Fxp. 169.-As was just now said, carbonic acid may be produced
also by heating carbon in contact with solid bodies which contain oxygen, such, for example, as the red oxide of mercury. Mix 11 grammes




322             PREPARATION OF CARBONIC ACID.
of red oxide of mercury with 0'33 grm. of charcoal; place the mixture
in an ignition-tube, arranged as in figure 53; heat the tube and collect
over water the gas which is evolved. Test the product with limewater, as in Exp. 168. The reaction between the charcoal and the
oxide of mercury may be written as follows:2HgO + C =  C02 + 2Hg.
The metallic mercury set free condenses in droplets upon the cold
upper portions of the ignition-tube. Here, again, as in Exps. 124, 155,
the metallic oxide is reduced by the charcoal.
398. Carbonic acid may readily be obtained from certain compounds called carbonates, several of which are abundant minerals.
Common chalk, marble, and limestone, for example, are composed of carbonate of calcium; and carbonic acid can readily be
obtained by strongly heating them, or by subjecting them to the
action of strong acids.
Exp. 170.-Place two or three grammes of coarsely-powdered marble
in an ignition-tube provided with a gas delivery-tube bent at a light
angle; place the ignition-tube upon the iron stand over the gas-lamp,
and dip the outer opening of the delivery-tube into a small bottle
containing lime-water; heat the marble strongly, and observe the white
precipitate which forms in the lime-water as the carbonic acid gas
comes in contact with it. The carbonate of calcium, thus precipitated
by bringing together carbonic acid and oxide of calcium, is chemically
identical with the chalk or marble from which the acid was expelled.
In actual practice enormous quantities of carbonic acid are
expelled from limestone in this way for the sake of the quicklime
which is left as a residue; but
the carbonic acid, thus expelled
by heat, is rarely collected, for
a more convenient method of
procuring it is to treat the lime.
stone with some acid capable of   -
expelling the carbonic acid.
Exp. 171.-In a gas-bottle of
500 or 600 c. c. capacity, arranged
precisely as for generating hydrogen (see Exp. 19), place 10 or 12
grins. of chalk or marble in small
lumps.i cover the chalk with




PREPARATION OF CARBONIC ACID.             323
water, and pour in through the thistle-tube concentrated chlorhydric
acid, by small portions, in such quantity as shall insure a continuous
and equable evolution of gas. Collect several bottles of the gas over
water, then replace the anterior portion of the delivery-tube with a
straight tube and collect one or two bottles of the gas by displacement; carbonic acid gas is half as heavy again as air. The reaction
between the carbonate of calcium and the chlorhydric acid may be
thus formulated:CaCO, + 2HC1'= CaC12 + H20 +  CO2.
When chalk is the material operated upon, sulphuric acid may be substituted with advantage for the chlorhydric acid; for the latter, being
rather easily volatile, is liable to be carried forward by the current of
carbonic acid, and to contaminate the product. When carbonic acid
is prepared for commercial purposes by the action of an acid upon carbonate of calcium, sulphuric acid is almost always the acid employed,
and marble-dust the substance acted upon:CaCO, + HSO, =_ CaSO, + H,20 +  C02.
WVith a porous material, like chalk, this action occurs readily; but in
attempting to operate upon compact varieties of carbonate of calcium
such as marble, difficulties are encountered, unless the carbonate be in
the state of powder. The sulphate of calcium, which is a product of
the reaction, is a rather difficultly soluble substance, and, being deposited upon the surface of the marble, soon covers it with a coating so
thick and impermeable that the action of the sulphuric acid upon the
marble is well nigh completely arrested.
Carbonate of calcium being cheaper than most other carbonates, is
more commonly employed than any other as a source of carbonic acid;
but it is sometimes convenient to substitute for it the carbonate of
sodium, or of potassium (saleratus), or even of ammonium, since carbonic acid is given off from these compounds very quickly and abundantly. A self-regulating gas-generator, such as is represented in Fig.
xxviii. of the Appendix, charged with large solid lumps of the commercial carbonate of ammonium and chlorhydric acid, is, perhaps, the
most convenient apparatus which can be employed for preparing the
gas upon the lecture-table.
399. At the ordinary atmospheric temperature and pressure,
carbonic acid is a transparent, colorless gas, of a slightly acid smell
and taste. It is incombustible, being already the product of the
complete combustion of carbon, and is, moreover, incapable of
supporting the combustion of most other bodies, since the oxygen
Y 2




324              PROPERTIES OF CARBONIC ACID.
contained in it is very firmly held; like nitrogen, it immediately
extinguishes burning bodies which are immersed in it.
Exp. 172.-Thrust into a bottle of the gas obtained in Exp. 171, a
lighted candle, or, better, a large flame of alcohol burning upon a tuft
of cotton; in either case the flame will be instantly extinguished.
The specific heat of gaseous carbonic acid, between 10~ and
200~, is 0'2169. Its specific gravity is 22; being thus 1'53 as
heavy as air, it can be poured from one vessel to another almost
as readily as if it were water.
Exp. 173.-Invert a bottle filled with carbonic acid upon another
bottle of equal size filled with air, in such manner that the mouth of
the upper inverted bottle shall rest upon the mouth of the lower bottle.
After the lapse of several minutes, thrust a burning splinter of wood
into each of the bottles; in the upper bottle the splinter will continue
to burn, for into this bottle the air from the lower bottle has ascended;
while in the lower bottle, now full of carbonic acid, the splinter will
be extinguished.
Exp. 174.-From a large bottle full of the gas, pour a quantity of
carbonic acid upon the flame of a lamp or candle; that is to say, hold
the mouth of the open bottle of carbonic acid obliquely over the candleflame so that the gas shall fall like water upon it; the flame will immediately be extinguished.
400. Owing to the great weight of carbonic acid, it often fails
to rise out of wells and other cavities in the earth, in which it is
generated by the decomposition or decay of organic substances.
Before allowing workmen to descend into any such place, where
there is reason to suspect the presence of carbonic acid, a burning
candle should first be lowered; if the candle be extinguished, or
even if it burn feebly, the noxious character of the air is indicated,
and measures should at once be taken to purify the locality by
ventilation or otherwise. One way of removing the carbonic acid
is to absorb it by means of some chemical agent, such as slaked
lime (hydrate of calcium) or potash-lye. The slaked lime is most
efficient as an absorbent when neither very wet nor very dry; it
should not be dry enough to be dusty, nor yet noticeably moist.
If a quantity of it be suspended in the well, or thrown upon the
floor of the cellar containing carbonic acid, this gas will quickly
combine with the lime to form carbonate of calcium, as in Exp. 168.
Another good method is to lower down a chafing-dish full of




VENTILATION OF WELLS.                 325
brightly glowing charcoal; if the carbonic acid be present in such
quantity that the test-candle has been extinguished by it, the
charcoal will, in like manner, be immediately extinguished, and,
in cooling, will rapidly absorb this gas (see ~ 385), together with
any nitrogen which may be present; a current of fresh air will,
in this way, be induced to flow into the well. If, by the candle
test, but little carbonic acid be found in the well, it would, of course,
be best to extinguish the charcoal before lowering it into the impure air; this can readily be done by covering it up tightly in an
iron kettle.
It should be distinctly understood that the accumulation of
carbonic acid in wells and caves is a consequence of the low
diffusive power of the gas (see ~ 53).  Carbonic acid mixes with
air very slowly; but when once mixed with air it has no further
tendency to settle down, or to separate itself in any way.
Exp. 175.-Over a bottle filled with carbonic acid gas, invert another bottle full of air, in such manner that the mouth of the airbottle shall rest upon that of the upright bottle full of carbonic acid.
After some hours, pour lime-water into each of the bottles and shake
them; a precipitate of carbonate of calcium will be formed in both
cases, for a part of the carbonic acid will, by this time, have ascended
out of the lower into the upper bottle. The two gases have, in fact,
become intimately mixed or blended; the heavy carbonic acid has
diffused upwards into the air, and the lighter atmospheric air has diffused downwards into the carbonic acid, just as, in a previous experiment, we have seen hydrogen and oxygen diffuse into each other.
(Fig. 20, p. 43.)
The great importance of this diffusion of gases, in the economy of
nature, is well illustrated by the case now under consideration. Whenever, as in the processes of respiration and combustion, oxygen is withdrawn from and carbonic acid thrown into the air, the carbonic acid,
and the nitrogen with which it is accompanied, immediately mix with
the surrounding air and distribute themselves through the atmosphere.
The composition of the atmosphere is thus maintained uniform all
over the globe, in spite of the constant removal from it of oxygen in
some localities, and the addition of carbonic acid in others. It is only
in confined places, where nearly pure carbonic acid is produced more
rapidly than it can pass off by diffusion, that the gas accumulates to
any appreciable extent.
The singular facility with which gases, and particularly hydrogen,




326                   DIFFUSION OF GASES.
traverse porous bodies is very strikingly illustrated by the following
experiment, which our acquaintance with carbonic acid now enables
us to perform:-Through a large glass tube, a smaller tube of porous,
unglazed earthenware is passed, and the ends of the glass tube are
tightly closed by the corks which hold the porous tube. By the tube
Fig. 56.
a, a rather rapid stream of carbonic acid is brought from a self-regulating generator into the annular space between the two tubes, while,
by the tube b, a slower current of hydrogen is introduced into the inner
porous tube. It would be expected that hydrogen should be found in
the cylinder d, and carbonic acid in the cylinder c; but, on the contrary, an inflammable gas is found in c, and in the cylinder d carbonic
acid pure enough to extinguish a burning splinter. The hydrogen diffuses almost instantaneously into the annular space, and the carbonic
acid enters the inner tube to replace the issuing hydrogen.
401. When pure, or nearly pure, carbonic acid is irrespirable;
it produces spasms in the respiratory passages, and is thus prevented from entering the lungs. When so far diluted with air as
to admit of being respired, it acts as a narcotic poison; it is, however, far less poi/onous than the other oxide of carbon, carbonic
oxide, directly to be described.
402. Carbonic acid gas is soluble in water to a considerable
extent. One measure of water, at the ordinary temperature and
pressure. will dissolve one measure of carbonic acid gas; but its
solubility increases if the pressure be increased.
Exp. 176. —Into a long-necked flask or phial filled with carbonic
acid, pour a quantity of water, close the bottle with the finger, and
shake it; immerse the mouth of the bottle in water, and remove the
finger, water will rush into the bottle to supply the place of the gas
which has been dissolved. Again place the finger upon the mouth of




SOLUBILITY OF CARBONIC ACID.             327
the bottle, shake the bottle as before, and subsequently open it beneath
the surface of the water; a fresh portion of water will flow into the
bottle to supply the new vacuum; in this way, by repeated agitation
with water, all of the carbonic acid in the bottle can be absorbed.
Owing to this solubility in water, some carbonic acid is always lost
when the gas is collected over water as in Exp. 171; but since considerable time is required to absorb all the gas, there is little objection
to collecting it over water.
403. When subjected to increased pressure, carbonic acid gas
dissolves in water much more abundantly than at the ordinary
pressure of the air; under a pressure of two atmospheres, one
measure of water will dissolve two measures of the gas; under a
pressure of three atmospheres, it will dissolve three measures, and
so on.  Water thus surcharged with carbonic acid has an agreeable, acid, pungent taste, and effervesces briskly when the compression is suddenly removed, as when the liquid is allowed to
flow out into the air; such carbonic acid water, or " mineral
water," as it is then called, flows from the earth in many localities, as at Seltzer and Saratoga; it is also prepared, artificially,
in large quantities, and sold as a beverage under the meaningless
name of soda-wvater.  Carbonic acid water possesses solvent
powers far greater than those of pure water; few minerals are
capable of resisting its long-continued action; hence the waters
of tile springs from which it issues are usually highly charged
with saline and mineral ingredients, and are often of medicinal
value.
404 The effervescent qualities of fermented liquors, such as
cider, champagne, and beer, are in like manner dependent upon
the presence of compressed carbonic acid gas.  In all these cases
the carbonic acid is of value, not only on account of the agreeable
pungency which it imparts to the beverage, but also because of
the fact that in escaping from solution and assuming the gaseous
condition. it absorbs a very considerable amount of heat, and so
cools the liquid which contained it.
405. Carbonic acid is widely diffused in nature. Traces of it
occur in the air and in water everywhere; and there are many
localities, besides the mineral springs before-mentioned, where it
issues from the earth in large quantities, notably in several vol



328                     FERMENTATION.
canic districts. It is produced not only in the actual combustion
of all substances which contain carbon, but also during the decay
and putrefaction of all animal and vegetable substances.  During
fermentation it is evolved in large quantities, and it is continually
given off during the respiration of animals.
EBxp. 177.-Put one or two table-spoonfuls of coarse meal into a
bottle of about 250 c. c. capacity; cover the meal with water, and connect the bottle, by means of a cork and glass tube, with a second bottle
filled about three c.m. deep with lime-water; the delivery-tube must
reach into the lime-water.  Out of the top of the second bottle carry
a second bent glass tube, whose open end dips into water. The bottle
containing the lime-water will, of course, be closed by the cork through
which pass the two tubes above described. Keep the apparatus in a
warm place; bubbles of gas will pass from the first into the second
flask;4hey contain carbonic acid, as may be seen from the precipitate
of carbonate of calcium which they produce.
Exp. 178.-Dissolve 2 grammes of honey or molasses in 200 c. c. of
water; fill a large test-tube with the mixture, and add to it afew drops
of bakeis' or brewers' yeast; close the open mouth of the test-tube
with the thumb, and invert it in a small saucer or porcelain capsule
filled with the diluted syrup. Place the saucer and tube, with their
contents, in a warm place, having a temperature of about 20~ or 30~,
and leave them there during 24 hours. In a short time, fermentation
sets in, and the sugar of the syrup is gradually converted into alcohol
and carbonic acid.
CCHi206 =  2C2,60 + 2CO2.
Sagar.    Alcohlol.
The carbonic acid thus formed rises in minute bubbles, causing a gentle
effervescence in the liquid, and collects in the upper part of the tube,
while the alcohol remains dissolved in the liquid. That the gas ob.
tained in this experiment is really carbonic acid, may be proved by
transferring some of it into a clean tube at the water-pan, and then
testing with lime-water.
Exp. 179.-Provide two test-glasses or small bottles; place in each
15 or 20 c. c. of lime-water; through a glass tube blow into the limewater of one of the bottles air coming from the lungs. By means of
bellows, to the nozzle of which a gas-delivery tube has been attached,
force through the lime-water of the second bottle a quantity of fresh
air. The clear liquid of the first bottle will quickly become turbid
through deposition of carbonate of calcium, while the lime-water of the
second bottle will remain clear for a long while.




RESPIRATION.                     329
This experiment may be modified by constructing an apparatus
with valves, in such manner that the air drawn into the lungs shall
be made to pass through one bottle of lime-water, while the air expired goes out through another bottle of lime-water. The contents of
the first bottle will remain clear, while the liquid in the other immediately becomes turbid.
Ordinary fresh air contains only about one 2000th part of carbonic
acid, while air expired from the lungs contains as much as 3 or 4 per
cent. of it.  In breathing, animals inhale oxygen from the air; this
oxygen combines with carbon within their bodies and is exhaled as
carbonic acid.  A crowd of men consume the oxygen of the air, just
as lamps or fires consume it; to carry off this product of animal combustion is one of the objects of systems of ventilation. Air which contains even as little as 1 or 2 per cent. of carbonic acid, exerts a very
depressing effect when breathed for any length of time.
406. From the foregoing statements it appears that', in the
several processes called "life," " fermentation," "decay," and
" combustion," there is involved chemical action; in all of these
processes oxygen from the air unites with carbon, while carbonic
acid is set free and thrown into the atmosphere.  The question
now arises, What becomes of all this vast amount of carbonic acid
which is constantly coming into our atmosphere from the respiration of animals, from fires, from  decaying and fermenting substances, from volcanic fissures, and from various other sources?
If this carbonic acid remained in the air, the latter would quickly
become unfit to support animal life. But it is not found that the
average proportion of carbonic acid in the air does increase; on
the contrary, all the evidence goes to show that there was, at
certain earlier geological epochs, more of it in the air than now.
Many geologists believe that, in the early history of our globe,
there was much more carbonic acid in the air than at present;
hence immense forests arose whose carbon is now stored away
in the form of coal; hence also the formation of enormous beds
of limestone covering many parts of the earth's surface, —processes
of which the faint continuations are now seen in the formation of
peat-bogs and coral-reefs. On the other hand, it is not found
that the proportion of oxygen in the atmosphere undergoes any
appreciable change, in spite of the enormous volume of it which
is absorbed in the pilocesses of breathing, combustion, and decay




330                 LIQUID CARBONIC ACID.
above enumerated.  For, unlike animals, plants, in breathing,
take in carbonic acid and give out oxygen. The leaves of plants
are so constructed that they can decompose carbonic acid, fix carbon for the building up of the plant, and set oxygen free. This
reciprocal action of plants and animals, tending to maintain unchanged the constitution of the atmosphere, is one of the most
wonderful adjustments of nature.
407. Carbonic acid can be liquefied by pressure, and the liquid
thus obtained can be solidified by exposure to cold. When the
gas is generated in a confined space in a strong vessel, it soon
exerts so powerful a pressure that a large portion of it condenses
to a transparent, colorless, mobile liquid, somewhat resembling
water, though it refracts light less powerfully; or it may be
liquefied by mere cooling to -106~, under the atmospheric pressure. A better way of preparing the liquid acid is to pump the
gas, by means of a forcing syringe, into a strong wrought-iron
vessel surrounded with pounded ice. The pressure can thus be
regularly and methodically increased and the receiver finally filled
with the liquid. At 0~ a pressure of 36 atmospheres is required,
in order that the acid shall remain in the liquid state. Liquid
carbonic acid does not mix readily with water or with the fixed
oils; but with alcohol, ether, petroleum, and similar liquids it is
miscible in all proportions. Its specific gravity is 0-83 at 0~; but
it expands to such an extent on being heated, that at 30~ its
specific gravity is only 0'6. It expands to a greater extent, on
the application of beat, than any known substance, even to a
greater extent than the gases; 20 volumes of it at 0~ become 29
volumes when the temperature is raised to 30~; 150 volumes, at
300, shrink to 100 volumes when the temperature is reduced to
-200. The liquid acid well illustrates the general fact that
liquids expand proportionally much more when heated under a
high pressure than under a low one.
408. If the stop-cock of a vessel containing liquid carbonic
acid be opened, in such manner that a stream of the liquid shall
be forced out into the air, a portion of it at once assumes the
gaseous state, and in so doing absorbs so much heat from the remainder that the latter solidifies, and is deposited in the form of
white flakes like snow. This snow-like substance is slowly con



SOLID CARBONIC ACID.               331
verted into gas when exposed at the ordinary pressure of the air,
and so disappears. Though its temperature is lower than -78~,
it may be handled lightly without exciting any special sensation
of coldness or pain; for the gas which it is constantly emitting is
a bad conductor of heat, and prevents it from coming into intimate
contact with the skin. When, however, the solid acid is forcibly
pressed between the fingers, it produces painful blisters, as if it
were red-hot iron. In order to use the solid acid for producing
low temperatures, it is best to mix it with a small quantity of
ether; in such a mixture quicksilver can readily be frozen, and
many gases can be liquefied or even solidified. if this mixture
be placed in the vacuum of an air-pump, a temperature as low as
-1(00~ can be obtained; and if a tube containing liquid carbonic
acid be then placed in the mixture, the liquid will speedily be
frozen to a clear transparent mass like ice.
409. Carbonic acid gas, on being heated from 0~ to 1000, does
not expand at the same rate as air and the other permanent
gases, but increases in volume to a greater extent than any of
them. Upon being heated one degree, it expands 0-003688 its
volume at 00. This behavior is in accordance with the general
rule, that those gases expand the most which are most readily
condensable to the liquid state, while those gases which have resisted all efforts to liquify them scarcely show any appreciable
differences in the rate of expansion.
In the same way, carbonic acid, like the other easily condensable
gases (see ~ 221), does not conform precisely to the law of Mariotte. At pressures greater than one-third of the pressure of our
atmosphere, its volume diminishes more rapidly, with increasing
pressure, than would be the case with air and the other permanent
gases under the same conditions.
410. Carbonic acid is one of the compound gases which can be
split by heat alone into its proximate constituents; in other words,
it exhibits the phenomena of dissociation (~ 300). When the gas
is passed through a strongly heated porcelain tube, the gaseous
mixture which escapes from the tube contains, besides undecomposed carbonic acid, notable quantities of carbonic oxide (CO) and
oxygen.
411. As has been already stated (~ 399), the oxygen in carbonic




332            DECOMPOSITION OF CARBONIC ACID.
acid is so strongly held that it cannot be withdrawn by combustibles under ordinary circumstances; but at high temperatures carbonic acid is decomposed by carbon and by several of the metalssuch as iron, zinc, and manganese, besides potassium, sodium, and
the other metals of the alkalies and alkaline earths. By the
alkali-metals the oxygen is completely removed from carboiic
acid, and carbon is set free; but by the other agents above-nlelltioned, only half the oxygen of the acid is taken away, while carbonic oxide gas is formed:CO2 + C = 2CO.
Phosphorus, also, at high temperatures and in presence of a fixed
alkali, decomposes carbonic acid in the same way, and abstracts
part of its oxygen.  So, too, if a mixture of equal volumes of hydrogen and carbonic acid be passed into one end of a red-hot tube,
steam and carbonic oxide gas will escape at the other:CO2 +  2H  =  CO +  H20.
The decomposing-power of the alkali-metals, above alluded to,
furnishes us one means of partially analyzing carbonic acid..Exp. 180.-To a gas-bottle in which carbonic acid is being steadily
evolved, according to Exp. 171, attach a chloride-of-calcium tube, and
beyond this drying-tube a short tube of hard glass, from which an
exit-tube leads into a small open bottle, as shown in Fig. 57. When
Fig. 57.
the extinction of a lighted match in the open bottle proves the apparatus to be full of carbonic acid, thrust into the hard-glass tube a bit
of potassium as big as a pea, previously dried between folds of blottingpaper, and then gently heat the potassium with a lamp. The potassium
will take fire and burn at the expense of the oxygen of the carbonic
acid, and black particles of carbon will be deposited upon the walls of
the tube. After the reaction has ceased, and the tube has been allowed




COMPOSITION OF CARBONIC ACID.          333
to become cold, place it in a bottle of water so that the saline mass
(carbonate of potassium) may dissolve; the particles of carbon will then
be seen more clearly, floating in the liquid; they may be collected
upon a filter. The potassium in this experiment may be replaced by
sodium, but in this case a somewhat higher temperature is required.
412. The quantitative composition of carbonic acid may be
readily ascertained by the method of synthesis. When a quantity
of carbon is burned in a confined and measured volume of oxygen,
it is found that the volume of carbonic acid gas produced has
sensibly the same bulk as the original oxygen. Hence we conclude that the normal or product-volume of the molecule of carbonic acid gas contains two volumes of oxygen.
Now two volumes of carbonic acid weigh 44'152, since the
weight of the unit-volume, or the specific gravity of carbonic
acid, has been found to be 22-076.
Subtracting from this weight of the product-volume of the gas 44-152
The weight of two unit-volumes of oxygen (15 969 x 2).. 31938
There remains as the weight of the carbon in the productvolume of the gas.............. 12214
The weight of one atom of carbon is 12, as we have seen above
(~ 395), and it follows that the formula of carbonic acid is CO2.
From these figures the following percentage composition of carbonic acid may be deduced:
Carbon............. 27-66
Oxygen.......                       72-34
100'00
But these results must be regarded merely as approximations to
the truth, since the deviation of carbonic acid from the law of
Mariotte (~ 409) renders it well-nigh certain that we have not
yet been able to precisely determine, by experiment, the true
weight of a unit-volume of this gas.
413. The composition of carbonic acid, however, has been determrrined with very great accuracy by burning a known weight
of pure carbon in a stream of oxygen gas and carefully collecting
and weighing the carbonic acid produced.
In order to do this, the product of the combustion of the carbon,
together with the excess of oxygen, is made to flow through U-tubes




334       SYNTHESIS OF CARBONIC ACID-CARBONATES.
(Appendix, ~ 15) filled with fragments of hydrate of potassium, a substance which absorbs only the carbonic acid; the weight of these tubes
is determined before the commencement of the experiment and again
at its close, the increase of weight during the experiment being, of
course, referable to the carbonic acid absorbed. Knowing, then, the
weight of the carbon taken and the weight of the carbonic acid into
which it has been converted by uniting with oxygen, a very simple
calculation, as before, gives us the percentage composition of the acid.
Experiments of this kind have yielded the following result:Carbon............. 27'27
Oxygen......       72'73
100 00
It therefore, appears, that, in uniting to form carbonic acid, the
elements carbon and oxygen combine in the proportion of 3 parts
by weight of carbon to 8 parts by weight of oxygen, or in the
proportion of 12 to 32. If the number 72'73 be divided by 16,
the number representing the weight of a unit-volume of oxygen,
and if the number 100 be divided by 22, the number which, in
accordance with the weight of all the evidence thus far accumulated must be regarded as the true unit-volume weight of carbonic
acid, there will be obtained in each case the same quotient, namely,
4'545 volumes; whence we conclude, as before, that any volume
of carbonic acid contains the same volume of oxygen.
414. Carbonic acid unites with the protoxides of most of the
metals to form well-defined salts called carbonates. The greater
number of the best-defined carbonates contain only one molecule
of base; but besides the normal carbonates of the general formula M20,CO2 or MCO,, there are sesquicarbonates of the formula 2M20,3C02, " bicarbonates " of the formula M2C),H20,2CO2
or MHICO,, and basic carbonates containing two, three, or more
molecules of the base to one of the acid. In general the term
basic salt is applied to all salts in which, as in the carbonates last
named, the alkaline constituent or base predominates. The appellation " bicarbonate," though ordinarily applied in the manner
indicated above, is a name of very doubtful correctness; strictly
speaking, the class of compounds to which it refers should, perhaps, be regarded as double salts of the metal and hydrogen. The
normal carbonate, and the so-called bicarbonate of sodium, for




CARBONIC OXIDE.                  335
example, differ only in that half the sodium in the normal salt
has been replaced by hydrogen in the bicarbonate:Normal Carbonate of Sodium.      Bicarbonate oj Sodium.
Na,2C03.                    NaHCO3.
Carbonic acid is an exceedingly weak acid; it fails to neutralize (~ 65) completely the causticity of oxides such as those of the
alkaline metals; the normal cai-bonate of sodium, for example, is
decidedly alkaline in its reaction and properties. The so-called
bicarbonate of sodium is also slightly alkaline; and even the solution of carbonate of calcium in carbonic acid water exhibits an
alkaline reaction when tested with turmeric paper. Almost all
the carbonates are readily decomposed by acids (even by very
weak acids), with an effervescence caused by the escape of carbonic acid. Most of them are decomposed also on being heated;
but from the normal salts of sodium and potassium carbonic acid
cannot be expelled by heat alone, however intense.
415. Carbonic Oxide (CO).-As has been stated in ~ 411, carbonic oxide can be prepared by acting upon carbonic acid with
hot charcoal.
Exp. 181.-In the middle of a tube of hard glass, No. 2 or 24, about
35 c.m. long, pack a column 15 c.m. in length of coarsely-powdered
charcoal. Place the tube upon a sheet-iron trough on a ring of the
iron stand above wire-gauze lamps, as shown in the figure. Connect
Fig. 58.
the tube either with a gas-holder containing carbonic acid, or with a
bottle in which the gas is being generated. Heat the charcoal intensely, and from time to time test the gas which is delivered at the
water-pan, as to its inflammability. Carbonic oxide takes fire on being




336            PREPARATION OF CARBONIC OXIDE.
touched with a lighted match, and burns with a bluish flame. In place
of the charcoal, small fragments of iron or of zinc may be employed in
this experiment.
Instead of gaseous carbonic acid, the solid compounds called
carbonates can be conveniently employed for preparing carbonic
oxide.
ETcen. 182.-M-{ix powdered chalk (carbonate of calcium) with anl
equal weigrht of iron or zinc filings, place the mixture in an ignitiontube, provided with a gas-delivery tube, and heat it to redness over
the gas-lamp. The metal will abstract an atom of oxygen from the
carbonate of calcium, oxide of iron or of zinc will be formed, while
carbonic oxide will pass off through the delivery-tube to be collected
at the water-pan:CaCO, + Fe = FeO + CaO +  CO.
In the same wavy, when a mixture of chalk and finely-powdered
charcoal is heated to full redness, carbonic oxide gas is given off:CaCO, + C = CaO + 2CO.
It should be mentioned, however, that in all these cases the carbonic
oxide obtained is more or less contaminated with carbonic acid, portions of which escape reduction by the metal and carbon; the carbonic
acid may always be readily removed by causing the gas to pass through
a strong solution of caustic soda or through a U-tube filled with bits
of pumice- stone saturated with soda-lye.
Carbonic oxide can be obtained also by heating charcoal with
other solid oxygen compounds, such as the phosphate of calcium,
already mentioned (~ 270), or the oxides of almost any of the
metals, provided the charcoal be in excess.
E3xp. 183. —Heat in an ignition-tube, as before, a mixture of 1 grm.
of finely-powdered charcoal and 8 grms. of red oxide of iron; collect
the gas over water, pour into the bottle a little soda-lye, close the
mouth of the bottle tightly and shake it, then open the mlouth of the
bottle under water, and finally test the gas with a lighted match.
416. Another easy method of preparing carbonic oxide is to
decompose oxalic acid by means of oil of vitriol; this is the
method usually employed in the laboratory.  Oxalic acid is a
solid vegetable acid, to be procured of the druggists; its composition may be represented by the formula H2C204. On being
heated with concentrated sulphuric acid, it suffers decomposition
in a manner which may be formulated as follows:H1O,C203 +  H20,SO, =  2H20,SO3 +  CO +  CO2.




PREPARATION OF CARBONIC OXIDE.            337
The elements of water are taken away from the oxalic acid and
united with the sulphuric acid, while the remainder of the oxalic
acid breaks up into carbonic acid and carbonic oxide.
BExp. 184.-Place in a flask of about 350 c. c. capacity 9 grms of
common oxalic acid and 53Fig. 59
grms. of concentrated sulphuric acid; connect the
flask with a bottle filled
with fragments of pumicestone saturated with a strong
solution of caustic soda, as
shown in the figure; heat
the contents of the glass
gently, and collect the gas   -
which is evolved over water
in the usual way. The car-                              I
bonic acid resulting from the
reaction will all be absorbed
by the soda-lye, and carbonic   ______
oxide will alone be delivered
at the water-pan.
None of the methods heretofore given yield pure carbonic oxide
directly; in each of the experiments we are compelled to wash
out carbonic acid from the gas obtained, if an absolutely pure
product is desired; but there are methods by which pure carbonic
oxide may be prepared without the need of any process of purification.  One of the best of these is as follows:
Exp. 185.-In a thin-bottomed flask of about 250 c. c. capacity and
provided with a suitable gas-delivery tube, heat a mixture of 5 grammes
of finely powdered ferrocyanide of potassium (yellow prussiate of potash) and 40 or 50 grms. of strong sulphuric acid. Collect the gas
over water, and test it as to its inflammability. Thrust also a lighted
splinter into the gas and observe that it will, be extinguished. The
reactions which occur between the chemicals employed may be expressed as follows:Ferrocyanide of Potassium.   Water.    Sulphuric Acid.
K4FeC6N6H603   +    3H2O   +    61H2SO0
6=          +CO  -  2K2SO4  +   FeSO4  +    3(NH4)2SO4.
Carbonic    Sulphate of    Sulphate of    Sulphate of
Oxide.      Potassium.      Iron.       Amnionium.
z




338             PROPERTIES OF CARBONIC OXIDB.
417. Carbonic oxide is a transparent, colorless gas, having
little, if any, odor; it has never yet been liquefied.  It is somewhat lighter than air, its specific gravity being 14, while that of
air is 14-5. It is but little soluble in water, and may be collected and preserved over water without much loss. It extinguishes combustion just as hydrogen does, and destroys animal
life. Unlike hydrogen and nitrogen, however, it is a true poison.
It destroys life, not negatively by mere suffocation or exclusion
of oxygen, but by direct noxious action. Even when largely
diluted with air, it is still poisonous, producing giddiness, insensibility, and finally death. It is the presence of this gas which
occasions the peculiar sensation of oppression and headache which
is experienced in rooms into which the products of combustion
have escaped from fires of charcoal or anthracite. Carb: nic oxide
is very much more poisonous than carbonic acid.  Much of the
ill repute which attaches to carbonic acid really belongs to carbonic oxide; for since both these gases are produced by' burning
charcoal, many persons are liable to confound them; but carbonic acid is, comparatively speaking, almost innocuous. Carbonic acid, it is true, is somewhat poisonous; it does not merely
suffocate, like water, or nitrogen, or hydrogen; but it is very
much less poisonous than carbonic oxide.  It has been found, by
experiment, that an atmosphere containing only 1-100th of carbonic oxide is as fatal to a bird as one containing 1-2.5th part of
carbonic acid.
Carbonic oxide exhibits neither an acid nor an alkaline reaction when tested with vegetable colors, and, in general, has but
little tendency to combine with other substances. With oxygen,
however, it combines readily at comparatively low temperatures;
an iron wire heated to dull redness is sufficient to inflame it in
the air. Unlike most other combustible gases, it contains no
hydrogen, and therefore produces no water when burned;
nothing but carbonic acid results from its burning.
Exp. 186.-To the apparatus employed for evolving carbonic oxide
in Exp. 184, attach a piece of small glass tubing drawn out at the end
to a fine point, and bent in such manner that a stream of gas may be
delivered upwards from this point. Light the gas as it flows out of
the tube,.an  hold over the pale-blue flame a clean, dry bottle. No




PROPERTIES OF CARBONIC OXIDE.          339
moisture will be deposited upon the sides of the bottle. That carbonic acid has been produced by the combustion, may be proved by
pouring a little lime-water into the bottle, and shaking it about in the
gas therein contained.
418. Carbonic oxide is a very powerful deoxidizing agent. At
high temperatures it is capable of taking oxygen away from
many of the compounds which contain that element. Hence it
plays a very important part in metallurgical operations. Much
of the reducing-action which is commonly attributed directly to
carbon, is really effected in practice through the mediation of
carbonic oxide gas.
Exp. 187.-In the middle of a tube of hard glass, No. 3, about 20
c.m. long, place a grainme of black oxide of copper (CuO); support
the tube upon a ring of the iron stand over the gas-lamp, and connect
it at one end with a flask in which carbonic oxide is being evolved,
as in Exp. 184, and at the other with a tube bent at a right angle and
dipping into a bottle which contains lime-water. After the tube, which
contains the oxide of copper, has become full of carbonic oxide, heat
it and observe that the oxide of copper is reduced, that metallic copper
alone remains in the tube, and that the carbonic acid formed has made
turbid the lime-water in the bottle.
419. The specific heat of carbonic oxide is considerably greater
than that of carbonic acid, being 0'245, while that of carbonic
acid is only 0'2103.
When a mixture of carbonic oxide and oxygen, in the proportion of two volumes of the former to one of the latter gas,
is lighted, it explodes with about the same degree of violence as
a mixture of hydrogen and oxygen (~ 58), a very considerable
amount of heat being evolved in the act of combination.
Though considerable heat is evolved during the union of carbonic oxide and oxygen, the amount is much less than that which
results from the complete combustion of charcoal to carbonic acid.
One gramme of carbonic oxide disengages in burning 2403 units
of heat (~ 55), while one gramme of wood charcoal, in burning
to carbonic acid, yields 8080 units. The same amount of heat
(2403 units) is reabsorbed when the carbonic acid, obtained by
burning one gramme of carbonic oxide, is again reduced to the
state of carbonic oxide. (Compare Exp. 181.)
The gramme of hydrogen yields, as it unites with oxygen,
z2




340             DISSOCIATION OF CARBONIC OXIDE,
34,462 units of heat; but since carbonic oxide is 14 times
as heavy as hydrogen, about the same quantity of heat is developed by the complete combustion of a given volume of carbonic
oxide as by that of the same volume of hydrogen.
420. Carbonic oxide, unlike carbonic acid, is not decomposed
when heated to redness in contact with hydrogen, charcoal, iron,
or zinc. Sodium and potassium, however, abstract the oxygen
from this gas as they do from carbonic acid.
It unites with chlorine directly, under the influence of sunlight, forming a gaseous compound, the composition of which
may be represented by the formula COC12. When left for some
time in contact with caustic potash, at the temperature of 100~,
it combines with it, and there is produced a compound known as
formiate of po'tassium:KHO + CO = CHKO,.
It is absorbed readily by solutions of the salts of dinoxide of
copper (Cu20O) in ammonia-water, and by a solution of dichloride
of copper (Cu2CI2) in strong chlorhydric acid, and can thus be
separated from  a mixture with other gases.  Melted metallic
potassium also absorbs a certain amount of carbonic oxide and
combines with it.
421. Carbonic oxide may be resolved into carbon and oxygen
by heat alone; but this dissociation occurs only under very peculiar
circumstances.
A porcelain tube is placed in a furnace where it can be raised to a
very high temperature; the ends of this tube project beyond the furnace and are closed by corks; through these corks passes, in the axis
of the porcelain tube, a very thin brass tube, and each cork carries also
a small glass tube; by one of these tubes carbonic oxide enters the porcelain tube, and by the other the products of the reaction escape from
the apparatus. Two little screens of porcelain divide internally that
part of the porcelain tube which lies in the furnace and is to be heated,
from the parts which project beyond the furnace and remain cool. A
rapid current of cold water is made to flow through the thin brass tube,
in such quantity that in traversing the tube while the furnace is in full
action the water shall not be sensibly warmed.
The apparatus being thus disposed, and the porcelain tube heated, a
slow and regular current of pure and dry carbonic oxide is passed into
the hot tube. The gas, as it issues from the tube, passes immediately




COMPOSITION OF CARBONIC OXIDE.           341
through a strong solution of caustic potash, which will absorb the carbonic acid formed, so that the experimenter can weigh the quantity of
acid produced, or through lime-water, which will demonstrate the presence of carbonic acid by becoming turbid. Carbonic acid is formed
whenever the porcelain tube is bright-red hot. A portion of the carbonic oxide is decomposed into oxygen, which unites with another
portion of carbonic oxide to make carbonic acid, and carbon, which is
deposited in the condition of lampblack upon the cold brass tube which
traverses the hot porcelain tube from end to end. The first action c(i
the heat is to set free particles of carbon from the carbonic oxide, and
all such particles which happen to fasten upon the brass tube are instantly chilled down below the temperature at which they will either
unite with free oxygen on the one hand, or reduce carbonic acid on
the other.
WVe have repeatedly used the electric spark as a means of decoen'
posing compound gases, such, for example, as ammonia (~ 89) and
marsh-gas (~ 395). It is supposed that it is the intense heat of the
spark which effects such decomposition, and, in the light of the experiment just described, it seems probable that the efficacy of the sparkcurrent is due to the fact that the few particles of gas which each spaIrk
heats intensely are immediately in contact with an atmosphere of gas
which is in constant motion and is relatively very cold.
422. The composition of carbonic acid being known, that of
carbonic oxide can readily be determined by burning a Fig. 60.
definite volume of this gas with an excess of oxygen in
a eudiometer.  If there be introduced into the eudiometer
100 volumes of carbonic oxide,
and 100,,,,oxygen,
and if through the 200,,,, mixed gases
an electric spark be made to pass, combination will
occur, and the gas which remains will occupy only 150 volumes.
If a small quantity of a solution of caustic soda be now introduced
into the eudiomneter, all the carbonic acid which has been formed
will be absorbed; there will remain only 50 volumes of gas,
which, upon examination, will be found to be pure oxygen.
If only 50 volumes of the original oxygen are thus left free,
50 volumes of oxygen must have been absorbed. It, appears,
then, that 100 volumes of carbonic oxide have united with 50
volumes of oxygen to form 100 volumes of carbonic acid; and




342            COMPOSITION OP CARBONIC OXIDE.
that the original bulk of the carbonic oxide taken has remained
unchanged. Since it is admitted, as we have already seen (~ 412),
that the product-volume of carbonic acid contains 2 volumes of
oxygen, it follows that the double volume of carbonic oxide can
only contain 1 volume of oxygen-or, in other words, that 2 volumes of this gas contain 1 volume of carbon-vapor and 1 volume
of oxygen united without condensation:C                = 
12  +      16           CO  28
It will be noticed that the addition of a certain quantity of oxygen to a measured quantity of carbonic oxide converts it into
carbonic acid without changing the original measured volume of
gas. We have often prepared compound gases from elementary
gases by this method, and in such cases there is generally a
change of voiume. We are here, however, converting one compound gas into another compound gas, and the product-volumes
of all compound gases are the same.
The specific gravity or unit-volume weight of carbonic oxide
has been found, by experiment, to be 13'97.
From the weight of two unit-volumes of carbonic oxide. 27'94
Deduct the weight of one unit-volume of oxygen... 16'
There remains the weight of the atom of carbon.... 11-94
This result accords very well with the previously given atomic
weight of carbon. It will be noticed that the specific gravity of
carbonic oxide is the same as that of nitrogen.
423. Combustion.-No- w that we have become acquainted with
carbon, hydrogen, and oxygen, and with some of the more important compounds formed by the union of these elements, the
subject of combustion can be more fully discussed than has been
possible hitherto. Unlike most of the chemical processes employed by man, which have for their object the preparation of
some tangible chemical compound or product, combustion is
resorted to merely for the sake of the heat or light which incidentally accompallies the chemical action.
As a general rule, only the compounds of carbon and hydrogen
are employed as combustibles-though there are some exceptions




LUrTNOSITY OF F'LAMES.               343
to this rule, as when the metal magnesium is burned for light,
or the heating of a sulphuretted ore is effected by the combustion
of its own sulphur. In the manufacture of sulphuric acid, the
sulphur-furnace is often so arranged that the heat from the
burning sulphur generates the steam necessary for the operation.
In the Bessemer process of making steel from cast iron (~ 633),
intense heat is evolved, partly by the combustion of the carbon
which cast iron contains, but partly also by the combustion of
iron.  The carbon compounds are peculiarly well adapted to the
purpose, since the products of. their combustion are gaseous, and
can therefore be readily removed; new portions of the combustible are thus continually laid bare, and a way opened for the
admission of fresh air.
424. In almost all cases artificial light results from the intense
ignition of particles of solid matter or of dense vapors. When
the heat, which is an invariable accompaniment of chemical combination, can play directly upon such solid or semisolid particles
with force enough to ignite them, an exhibition of light will
accompany the chemical change. The hydrogen-flame affords no
light, or as good as none. because in it nothing but a highly
attenuated gas is heated.  But when a solid body, such as the
platinum wire or the piece of lime employed in Exps. 26 and 27,
is placed in this non-luminous hydrogen-flame, intense light is
radiated from the heated solid.
Exp. 188.-Sprinkle fine iron filings into the flame of an alcohollamp, or into the non-luminous flame of the gas-lamp, and observe the
light given off by the particles of metal as they become incandescent
while passing through the flame. Or rub together two pieces of charcoal above a non-luminous flame, in such manner that charcoal powder
shall fall into the flame. Or rub the coat-sleeve beneath a non-luminous flame, or even beneath the luminous flame of an ordinary Argand
gas-burner, and observe that the particles of dust detached become
incandescent and luminous as they pass upward through the flame.
Other things being equal, the hotter the flame the more intense will
be the light emitted by the ignited solid.
425. In ordinary luminous flames, such as those of candles,
lamps, and illuminating gas, the ignited substance is carbon, or
rather a vapor or fog of certain carbon compounds containing
more or less hydrogen.




344                  LUMINOSITY OF FLAMES.
Exp. 189.-By means of caoutchouc tubing, attach to any small
gas-burner a piece of hard-glass tubing, No. 4, about 20 c.m. long, the
outer end of which has been drawn to a fine open point. Open the
cock of the gas-burner, so that gas may flow into and through the glass
tube, and light this gas as it escapes. When the last traces of air have
been expelled from the tube, heat the middle of the latter intensely
with the flame of a lamp. Part of the transparent and colorless carburetted hydrogen, of which the illuminating gas consists (~ 396), will
be decomposed by the heat as it passes through the tube, just as sulphuretted hydrogen (Exp. 91), phcsphuretted hydrogen, arseniuretted
hydrogen (Exp. 133), and antillloniuretted hydrogen (Exp. 136) are
decomposed under like conditions, and a ring of carbon will be deposited in the cold portion of the tube a short distance in front of the
flame.
In the open channel afforded by the tube it is not easy to heat the
whole of the carburetted hydrogen to the temperature necessary for
its decomposition; but by lighting the gas as it issues from the tube,
heat enough to decompose it can readily be obtained. Precisely as in
the combustion of wood (~ 378), after the fire is once started the combustible suffers decomposition; the easily inflammable hydrogen of the
gas burns first, and particles of carbon, or, at the first, of hydrocarbons
rich in carbon, are set free. These particles are heated to ignition by
the burning hydrogen, and as they pass up through the flame emit
light; finally they are themselves completely burned to carbonic acid
Upon the outside of the flame, provided there be present a sufficient
supply of air. That there are really particles of free carbon in the
flame has already been sufficiently demonstrated in Exp. 154.
If the supply of air furnished to the flame is insufficient to convert
all of the components of the gas into carbonic acid and water, then a
number of the carbonaceous particles will escape unburned, and a
smoky flame will be the result. If, on the other hand, the supply of
air is excessive, then all the carbon will be burned at the instant when
it is set free, and no light will be afforded. In the gas-lamps commonly employed in chemical laboratories for purposes of heating (see
Appendix, ~ 5), illuminating gas is purposely mixed with a considerable
volume of air before it is lighted; there is thus obtain: l an intensely
hot non-luminous flame. Such flames deposit no soot upon the vessels
which are heated in them; moreover the heat which would be consumed in heating the particles of carbon, and so producing light, is in
such flames utilized for heating-purposes.
-EWt. 190.-Unscrew the tube f (Fig. 61) from the gas-lamp constructed as described in ~ 5 of the Appendix, and light the gas as it




TE BUNSEN BURNER.                    345
escapes from the holes in the face of the screw d; the flame will be
luminous, and, if the holes are large enough to permit a rapid exit of
gas, even smoky. Extinguish the burning jet, screw on the tube f,
and relight the mixture of air and gas at the top; the flame will be
nearly colorless. Sometimes, when the
gas-cock is too nearly closed, the flne      Fig. 61.
of the miixed gas and air is liable to
pass clown the tube f, and ignite the
feeble jet of gas at the apertures in d.    f
The lamp then burns with a sickly
yellow flame, which is often tinged
with green coming from the copper in       o
the heated brass tube f. The lamp  
must be extinguished, and relit at the
top of the tube with a freer supply of
gas. When the tubef is in place, the    =
jets of gas, issuing vertically from
the face of the screw d, draw in currents of air through the side holes
near the bottom of the tubef; this air mixes with the gas rising
throughf, and at the top of this tube, where the mixture is inflamed,
the carburetted hydrogen is in intimate contact with air enough to
burn it at once and completely.
Between the two extremes which a Bunsen burner may be
thus made to illustrate, between a smoky flame, on the one hand,
and a non-luminous flame, on the other, there are two points
which hhve special significance-the point of most light, and the
point of most agreeable light. The point of most light may always
be hit upon by constructing such a burner as will just not allow
the gas to smoke.
Exp. 191.-Across the top of the chimney of an Argand gas-burner,
which is burning with a shorter flame than usual, place several narrow
strips of tin or of sheet iron, so as to obstruct the flow of air through
the chimney. The small, low flame with which the experiment began
will increase in size as the access of air is diminished, and, at last, the
whole interior of the chimney will be filled with a long, smoky flame.
The volume of gas burning at any one moment of the experiment is no
greater than at another, for the cock which regulates the flow of the
gas remains fixed and untouched; but the amount of light afforded by
the large smoky flame is manifestly greater than that yielded by the
small bright flame with which the experiment started. If any doubt
suggests itself as to this point, it will quickly be dissipated by perform



343           QQUANTITY AND INTENSITY FLAMES.
ing the experiment in a darkened room and noting the comparative
visibility of the more distant objects therein contained, first with the
one flame and then with the other; or the observer may determine at
what distances from the two flames fine print can be deciphered.
A murky flame, such as was just now obtained, before actual smoking begins, in which the largest possible number of the particles of carbon are heated, though none of them are heated very hot, yields the
largest amount of light which the particular sample of gas under examination is capable of affording. Such flames are called, technically,
quantity flames; they are better adapted than any others for lighting
streets and large halls. In practice, such flames are obtained by burning the gas at a low pressure, that is, under such conditions that it
shall be very gently pressed out into the air, so that air shall mix with
it and act upon it but slowly.
But besides this point of the maximum amount of light, there
is another, of the most agreeable light; and this is something
which each individual must determine for himself. Few persons
would choose, as a study-lamp, either the murky flame of Exp.
191, or the intense lime-light of Exp. 27; but between these two
extremes no one light is likely to suit many people equally well.
If a bright intensity flame is required, we have only to arrange
matters in such a way that air may come to the gas so quickly
and abundantly that a portion of the carbon in the gas, as well
as the hydrogen, shall be burned at once in the lower part of the
flame, and by the heat of its combustion ignite more intensely
the remaining particles of carbon. Among the very great number of gas-burners which have been devised, there may be found
those adapted to meet almost any requirement. Each kind of
burner brings the gas and the air into contact with one another
in some special way, producing a flame of convenient shape, of
peculiar economy, or of particular steadiness or brilliancy. It is
obvious that the conditions under which gas is most advantageously burned are different for different uses, and that no one
burner can be equally available under such varying and, in some
sense, antagonistic conditions. The Argand burner may, perhaps,
be made to fulfil as many of these conditions as any other; from
it there may be obtained, at will, either an intensity or a quantity
flame, as has been shown in Exp. 191.
426. The chemical composition of the gas to be burned is, of




ALL FLAS.'S GA.S-FLAMES.               347
course, an important point to be considered in the construction
of the burner.  A gas rich in carbon requires, for its combustion,
far more air than gas which is less carboniferous..xp. 192.-Place three small tufts of cotton upon an earthen plate;
moisten one with alcohol, another with petroleum, and the third with
benzin; touch a lighted match to the vapor which arises from each.
In the one case there will be seen hardly any luminous particles of
carbon, in the second a bright light; and in the third so much carbon
will be set free that, under the conditions of the experiment, a great
deal of it cannot find.air with which to unite, and consequently escapes
as smoke.
The composition of alcohol may be represented by the formula
C2H60; it contains a large proportion of hydrogen and some oxygen;
hence steam is necessarily produced when it burns; this steam spreads
or diffuses the flame, and promotes the prompt union of the alcohol
vapor with the oxygen of the air, so that few carbonaceous particles
have time to become incandescent before they are consumed. But in
benzin, the formula of which may be written C6H6, there is no oxygen,
and a far larger proportion of carbon than in alcohol; hence.the
necessity of supplying a large amount of air to the lamps in which its
vapor is burned. The best way of consuming benzin is to mix its
vapor with air in suitable proportions, and to press this mixture through
a gas-burner as if it were ordinary illuminating gas. When thus
treated, it burns without smoke, and affords a brilliant white light.
Petroleum (C2H6), like benzin, contains no oxygen, but it contains far
less carbon than benzin, though much more than alcohol; it does not
smoke like benzin, and yet it smokes so much that it cannot readily be
burned from a simple wick; it is commonly burned in lamps provided
with a special draught of air.
427. Ordinary lamps and candles are, strictly speaking, gaslamps.  In all cases their flames are composed of burning gas.
Exp. 193. —Construct a lamp as follows:-To a wide-mouthed bottle
of the capacity of about 50 c. c. lit a cork loosely; bore a hole in the
cork and place therein a short piece of glass tubing, No. 8, open at both
ends; through this glass tube draw a piece of lamlp-wicking, or any
loose twine, long enough to reach to the bottom of the bottle. It is
essential, either that the cork should fit the bottle loosely, or that there
should be a hole in the cork, in order that the pressure of the external
air may act upon the surface of the alcohol; to this end a very small
glass tube may be inserted in the cork at some distance from the tube
which carries the wick. Fill the bottle nearly full with alcohol, and,




348               THE COMPOSITION OF FLAMES.
after a few minutes, touch a lighted match to the top of the wick. The
fluid alcohol is drawn up out of the bottle by the capillary attraction
exercised by the pores of the vegetable fibre of which the wick is composed. When heat is applied to the alcohol at the top of the wick,
some of it is converted into vapor; this vapor then takes fire, and, in
burning, furnishes heat for the vaporization of new portions of the alcohol. From the top of the wick there is constantly arising a column
of gas or vapor, and upon the exterior of this conical column chemical
combination is all the while going on between its constituents and the
oxygen of the air. The dark central portion of the alcohol-flame is
nothing but gas or vapor.
Exp. 194.-Thrust the phosphorus end of an ordinary friction-match
directly into the middle of the flame of the alcohol-lamp of Exp. 193.
The combustible matter upon the end of the match will not take fire
in the atmosphere of carbonaceous gases of which the centre of the
flame consists. The wood of the match-stick, of course, takes fire at
the point where it is in contact with the outer edge of the flame, for
it is there heated in contact with air. In withdrawing the match from
the middle of the flame, it is not easy to prevent it from taking fire as
it passes through the outer edge of the flame; for the materials on the
tip of the match have been so strongly heated by radiation, during
their sojourn within the circle of fire, that they are now ready to burst
into flame immediately on coming in contact with the air; by a quick
jerk, however, the match may often be withdrawn from the flame
without taking fire.
Exp. 195.-Hold a thin wire (best of platinum, thoutgh iron will
answer well enough) or a splinter of wood across thI( iS1:,lle of an alcohol-lamp, as shown in Fig. 62. The wire will be heated to Figr. 62.
redness, and the wood will burn, only at the outer edges of
the flame where the gas and air meet; in the interior of the    A
flame the wire will remain dark and the wood unburned; for
there is no combustion there, and comparatively little heat.
If the wire be successively placed at different heights in the
flame, the size and shape of the internal cone of gas can
easily be made out; it will appear, moreover, that the hottest
part of the flame is just above the top of the interior cone of gas. As
a rule, when glass tubing, or the like, is to be heated in a flame, it
should never be placed below this point of the greatest heat.
428. The flame of an ordinary oil-lamp or of a petroleumlamp, in the same way as the flame of an alcohol-lamp, is composed of an inner cone of gas, or vapor of hydrocarbons, and an
envelope where chemical combination is going on; and a candle



CANDLE-FLAMES..4 9
flame is really the flame of an oil-lamp (Exp. 196). The presence of vapor in the candle-flame can be readily shown (Exps.
197, 198). In the candle-flame, as in that of the alcohol-lamp,
there is a cone of unburnt gas surrounded by a shell of burning
substances (Exps. 199, 200).
Exp. 196.-Touch a lighted match to the wick of a new candle; the
cotton of which the wick is composed takes fire and is at once consumed for the most part; but, in burning, the cotton gives off considerable heat, and some of the wax or tallow of which the candle is
composed is thereby melted and converted into oil. The liquid oil
ascends the wick by virtue of capillary attraction, and is converted into
vapor or gas by the heat of the cotton still burning at the stump of the
wick; this gas then burns precisely like the alcohol vapor in Exp. 193
or the illuminating gas in Exp. 189, and by the heat thus disengaged
new portions of wax or tallow are continually melted. There is always
a little cup of oil at the top of the rod of wax or tallow of which the
candle consists, and the apparatus is as truly an oil-lamp as if the oil
were held in a vessel of glass or metal.
Exp. 197.-Let a candle, best of tallow, burn until the snuff has
become long; blow out the flame, and observe the cloud of vapor which
ascends from the hot wick. Touch a lighted match to this column of
vapor, and notice that it takes fire at some little distance from the wick.
After the flame has been extinguished, the wick retains heat enough
for a few moments to distil off a quantity of gas, although there is not
heat enough generated to inflame this gas. To the gas or vapor thus
evolved is to be referred the disagreeable odor which is observed when
a candle is blown out.
Exp. 198.-Draw a glass tube (No. 5 or 6) 10 or 15 c.m. long to a
moderately fine open point; with a piece of wire bind this tube in an
inclined position to a ring of the iron stand, and place the lower end
of the tube in the middle of a candle-flame, just below the centre, so
that a portion of the gas of the inner cone of the flame may escape
through the tube; light the gas at the point at the top of the glass
tube, and observe that it will burn there steadily, if the experiment is
performed in a quiet place where there are no draughts of air.
Exp. 199.-Press down a piece of white         Fig. 63.
letter-paper, for an instant, upon the flame
of a candle until it almlost touches the wick,
then quickly remove the paper before it
takes fire, and observe that its upper surface is charred in the manner shown in Fig. 63. There will be obtained, in fact, burned into the paper, a diagram of the cylindrical




350              FORM OF LUM.NOUS FLAMES.
column ofunburnt gas and of the shell of burning matter which surrounds it. Within the charred ring the paper is unacted upon; for
that part of it was in contact only with the unburnt gas in the centre
of the flame.
Exp. 200. —Replace the paper of Exp. 199 with a strip of glass, so
held that the conical flame of the candle shall be cut across horizontally by the glass as it was by the paper in Exp. 199. Look down
fiom above through the glass into the hollow cylinder of unburnt gas
within the circle of combustion.
429. In any flame, which is rendered luminous by incandescent
carbonaceous particles, three portions can be distin- Fig. 64.
guished: —lst, the dark interior cone of gas, a, Fig. 64;
2nd, the zone of intense chemical action, b, where the    i:c
hydrogen is burning and the tearbonaceous particles are
heated to whiteness; and, finally, upon the very outside
a thin, scarcely perceptible film  of burning carbonic
oxide, c.                                                 \
430. From the study of luminous flames we pass to'
a consideration of flames employed only as sources of
heat.  In the experiments (25-27) with the oxyhydro-  
gen blowpipe, it has been already shown that a very
intense heat may be obtained by throwing oxygen into
the hydrogen-flame, and so localizing the chemical action and the
heat with which this action is accompanied. The subject may
be here conveniently studied by employing coal-gas and air in
place of hydrogen and oxygen.
Exp. 201. —lFill one gas-holder with air, and screw to it a metallic
jet, such as is shown in Fig. 65. Fill another gas-holder with ordinary illuminating gas and connect the opening of this gas-holder with
Fig. 65.
the lower opening of the metallic jet. Open the cock of the holder
which contains the coal-gas, and inflame the gas at the point of the
metallic jet. There will be thus obtained a long stream of gas burn



THE BLAST-LAMP.                    351
ing at the expense of the air which bathes its surface. The chemical
action between the oxygen of the air and the constituents of the coalgas, and the heat resulting from this action, are diffused over the entire
surface of this long flame. Without touching the cock of the holder
which contains the coal-gas, or in any way altering the amount of gas
which flows out of this holder, open the cock of the holder which contains air, so that air may be thrown into the middle of the coal-gas
flame. The latter will be immediately shortened down to almost
nothing. The constituents of the coal-gas will now all combine with
oxygen in a very small space, and the heat of combination, which was
diffused before, will be correspondingly concentrated. It is much the
same as if the coal-gas and the air had been mixed together beforehand
and then lighted. Indeed, in one of the first forms of the oxyhydrogen blow-pipe, a mixture of the two gases, such as we have exploded
in Exp. 30, was first prepared, and then forced out of a single gasholder of peculiar construction, provided with an exceedingly minute
orifice, at the mouth of which the mixed gases were burned. This
apparatus was inconvenient and dangerous, and has long since been
superseded; but it well illustrated the local concentration of heat now
under discussion.
431. The principle of the common mouth-blowpipe, of the
glass-blower's lamp (Appendix, ~ 6), and of all blasts and blowers,
is identical with that of the oxyhydrogen blowpipe, which, as has
been already stated (~ 55), is the simplest of all intense and concentrated combustions. Air, or more strictly speaking, oxygen,
is thrown into the combustible gas or fuel, in order that the combustion may go on in a small space.
The mouth-blowpipe may be used with a candle, or with any
hand-lamp proper for burning oil, petroleum, or any of the so-called
burning-fluids, provided that the form of the lamp below the wickholder is such as to permit, the close approach of the object to be
heated to the side of the wick. When a lamp is used, a wick about
1 2 c.m. wide and 0'5 c.m. thick is more convenient than a round or
narrow wick; a wick of this sort, though hardly so wide, is used in
some of the open burning-fluid (naphtha) lamps now in common use.
The wick-holder should be filed off on its longer dimension a little
obliquely, and the wick cut parallel to the holder, in order that
the blowpipe-flame may be directed downwards when necessary
(Fig. 67).
The cheapest and best form of mouth-blowpipe for chemical purposes is a tube of tin-plate, about 18 c.m. long, 2 c.m. broad at one end.




352                      THE BLOWPIPE.
and tapering to 0'7 c.m. at the other (Fig. 66); the broad end is
closed, and a little above this closed end
a small cylindrical tube of brass about 5
c.m. long is soldered in at right angles;
this brass tube is slightly conical at the
end, and carries a small nozzle or tip,
which may be made either of brass or
platinum. The tip should be drilled out
of a solid piece of metal, and should not
be fastened upon the brass tube with a
screw.  A trumpet-shaped mouth-piece
of horn or box-wood is a convenient
thouoh by no means essential, addition to   _'
this blowpipe.
E'xp. 202.-To use the mouthblowpipe, place the open end of the
tin tube between the lips, or, if the pipe is provided with a mouth-piece,
press the trumpet-shaped mouth-piece against the lips; fill the mouth
with air till the cheeks are widely distended, and insert the tip in the
flame of a lamp or candle; close the communication between the lungs
and the mouth, and force a current of air through the tube by squeezing the air in the mouth with the muscles of the cheeks, breathing, in
the meantime, regularly and quietly through the nostrils. The knack
of blowing a steady stream for several minutes at a time is readily acquired by a little practice. It will be at once observed that the
appearance of the flame varies considerably, according to the strength
of the blast and the position of the jet with reference to the wick.
When the jet of the blowpipe is inserted into the middle of a
candle-flame, or is placed in the lamp-flame in the position shown in
Figure 67, and a strong blast             Fig. 67.
is forced through the tube, a
blue cone of flame, a b, is pro-  /&  i 
duced, beyond and outside of
which stretches a more or less
colored outer cone towards c.                         -.
The point of greatest heat in
this flame is at the point of
the inner blue cone, because
the combustible gases are there supplied with just the quantity of
oxygen necessary to consume them; but between this point and the
extremity of the flame the combustion is concentrated and intense.
The greater part of the flame thus produced is oxidizing in its effect;
and this flame is technically called the o.zidizing flane. From the




OXIDIZING AND REDUCING FLAMES.             353
point a of the inner blue cone, the heat of the flame diminishes in both
directions-towards b, on the one hand, and towards c, on the other;
most substances require the temperature to be found between a and c.
Oxidation takes place most rapidly at, or just beyond, the point c of
the flame, provided that the temperature at this point is high enough
for the special substance to be heated.
A flame of precisely the opposite chemical effect may be produced
with the blowpipe. To obtain a good reducing flame, it is necessary
to place the tip of the blowpipe, not within, but just outside of the
flame, and to blow rather over than through the middle of the flame
(Fig. 68). In this manner, the flame is less altered in its general
character than in the former
case, the chief part consisting           Fig. 68.
of a large, luminous cone,
containing carbonaceous vapors in a state of intense ignition and just in the condition
for taking up oxygen. This
flame is therefore reduczing in
its effect, and is technically
called the reducing flame. The
substance which is to be reduced by exposure to this flame should be
completely covered up by the luminous cone, so that contact with the
air may be entirely avoided. It is to be observed that, whereas to produce an effective oxidizing flame a strong blast of air is desirable, to
get a good reducing flame the operator should blow gently, with only
enough force to divert the lamp-flame.
Substances to be heated in the blowpipe-flame, are supported, sometimes on charcoal, sometimes on platinum foil or wire or in platinum
spoons or forceps, and sometimes on little capsules made of clay or
bone-earth. Charcoal is especially suited for a support in experiments
of reduction.
Exp. 203.-The heat of the oxidizing flame may be well shown by
melting the extremity of a very fine platinum wire into a little ball.
To effect this fusion, bend the wire at right angles at 5 to 6 m.m. from
the end, and hold this bent end precisely in the axis of the flame, with
the angle outward and the extremity of the wire at the hottest part of
the flamle. If the wire is kept steadily in position, and the blast is
strong and the flame pure, a little knob will soon appear on the end of
the wire. The bend in the wire is made in order to keep a certain
length of the wire hot, and. so to diminish, the coaduction of heat from
the point.
2A




354                  ACTION OF CHIMN'EYS.
Exp. 204.-Place a kernel of metallic tin, as large as this o, in a
little hollow scooped out at one end of a bit of charcoal 8 to 12 c.m.
long. 3Melt this tin in the reducing flame of the blowpipe, and endeavour to preserve the metallic lustre of the fused metal untarnished.
A pure reducing flame is necessary for this purpose. A touch of the
oxidizing flame upon the metal covers its surface with a white, infusible, incandescent ash.
432. Another method of supplying the burning fuel with air
is by means of chimneys. Chimneys, whether of lamps or furnaces, are simply devices for bringing air in abundance, and therefore oxygen, into the fire; that in so doing they, at the same time,
carry off the waste products of combustion, is an incidental advantage..Exp. 205.-Light a piece of a candle 8 or 10 c.m. long and stand it
upon a smooth table; over the candle place a rather tall, narrow lampchimney of glass, the bottom of the chimney being made to rest upon
the table, and observe that the candle-flame will soon be extinguished.
No fresh air can enter the chimney from below to maintain the chemical action, and the small quantity of air which can creep down the
chimney from above is altogether insufficient to meet the requirements
of the case.
Exp. 206.-Relight the candle of Exp. 205, and again place over it
the lamp-chimney; but instead of allowing the chimney to rest closely
upon the surface o: the table, prop it up on two narrow strips of wood,
so that air can have free entrance into the chimney from below. The
candle will now continue to burn freely; for the heavy cold air outside
will continually press into the lower part of the chimney, and push
out the warm, light products of combustion, and the candle-flame will
all the while be supplied with fresh air.
Exp. 207.-Prepare several strips of " ni-    Fig. 69.
tre-paper" by soaking ordinary brown paper
in a strong solution of nitrate of potassium
and drying the product. On being lighted,
paper thus prepared will burn without flame
while emitting clouds of smoke. Light a
piece of the nitre-paper and plqce it at the
foot of the chimney arranged as in Fig. 69.
The smoke of the burning paper will instantly pass up through the chimney, and
so indicate the direction of the invisible a':r
which is all the while entering below and




CHIMNEYS CREATE DRAUGHTS.                355
passing out above at the top of the chimney, as fast as it is heated and
made lighter by the burning of the candle.
Exp. 208.-Repeat Exp. 200, and when the candle is burning
quietly, cover the top of the chimney tightly with a piece of tin or
sheet-iron, or with a strip of window glass; the candle will soon
cease to burn, precisely as if the chimney were closed at the bottom,
for, the escape of the hot products of combustion being prevented, no
air can pass into the chimney to reach the candle-flame.
It is by inducing the current of fresh air (Exp. 206), or draught, as
it is ordinarily termed, that chimneys are specially useful. They give
direction and precision both to the incoming cold air and the outgoing
hot gas. Where there is no chimney, the hot air from a lamp goes off
at a comparatively slow rate, and vaguely; through the chimney, on
the other hand, it flows straight forward and rapidly, and, of course,
a correspondingly direct and rapid current of fresh air presses in
to supply its place. Owing to this power of rapidly supplying air,
chimneys are employed upon lamps burning petroleum and other
highly carbonized oils which are liable to smoke; in general they
are made use of upon lamps in all cases where intensity flames are
required.
Exp. 209.-It is not absolutely necessary that the fresh air should
Aow into a chimney from below. Divide the
upper part of the chimney of Exp. 205, into     Fig. 70
two channels, by hanging in it a strip of
sheet-iron or tin, as a partition at the centre
of the chimney. (See Fig. 70.) Place the
chimney thus divided over a burning candle,
and observe that the candle will continue to
burn as if in a strong draught of air, although
no air can enter the chimney from below.
Hold a piece of burning nitre-paper (Exp.
207) at the top of the divided chimney; the
smoke will be drawn down into the chimney    -
on one side of the partition and thrown out
again upon the other, as indicated by the arrows in Fig. 70. It appears from this, as well as from the tremulous motion of the flame,
that a current of cold air passes down upon one side of the division
wall and supplies the required oxygen.
433. One exceedingly important point in chemical philosophy
which we have hitherto taken for granted, may now be readily
illustrated. As the result of long-continued experimentation and
the most rigid scrutiny of all chemical facts, it is admitted, as a
2A2




356                MATTER INDESTRUCTIBLE.
fundamental truth, that matter is indestructible. When substances undergo chemical change, none of their ingredients are
really destroyed, not an atom of them is annihilated; nor, upon
the other hand, is any new matter created; it is the form only of
the old substances which is changed; their weight remains in
every case unaltered.
By bringing about chemical combination between two or more
bodies, we can entirely change their appearance, their condition,
and their properties, but in every case it will be found that the
weight of the resulting compounds is precisely equal to the sum
of the weights of their components. Thus, when a candle burns
in the air and gradually disappears, none of the elements which
compose it are either lost or destroyed. Though by uniting with
the oxygen of the air, the components of the candle have been
converted into compounds which are invisible, it is nevertheless
easy to satisfy ourselves of the existence of these compounds.
Already, in Exps. 167, 168, it has been demonstrated that carbonic
acid and water are products of the combustion of the candle; and
we now proceed to show that these products weigh more than the
candle.
Exp. 210.-Take a glass tube 2 or 3 c.m. in diameter and 25 or 30
c.m. long, such, for example, as the neck of a broken retort; fit a cork
to each extremity of the tube, and at a distance of 6 or 8 c.m. from its
upper, wider end, fix across the tube a partition of wire gauze. Through
the upper cork insert a glass tube, No. 3 or 4, bent at a right angle,
and in the lower cork bore several open holes, besides a central orifice
into which a small wax taper is fastened. Fill the space between the
upper cork and the shelf of wire gauze with recently slaked lime, in
not too fine powder; replace the cork, hang the tube, with its contents,
at the end of one arm of a balance, and counterpoise the apparatus
precisely by placing shot or sand in the pan upon the other side of the
scales. Next make an "aspirating flask," by fitting to the upper orifice of a bottle with a stopcock, such as is depicted in the upper part of
Fig. xvii. in the Appendix, a sound cork, carrying a glass tube, No. 3
or 4; the tube should reach to the bottom of the bottle and be bent at
a right angle a short distance above the cork.
By means of flexible tubing, connect the upper tube of the aspirating
flask with the tribe issuing from the upper cork of the apparatus upon
the balance, and open the stopcock of the aspirator to such an extent
that water shall flow out from it slowly. Remove the lower cork from




KINDLING-TEMPERATURE.                  357
the apparatus upon the balance, in order to light the taper, then quickly
replace it and regulate the flow of water from the aspirator, so that
sufficient air shall be drawn in through the holes in the cork which
supports the taper to maintain the latter in lively combustion.
After the taper has burned during 4 or 5 minutes, close the aspirator,
remove the flexible tube from the candle apparatus, so that the latter
may again hang freely from the arm of the balance, and observe that
the weight of the apparatus is now greater than it was at the beginning of the experiment, for the lime within the tube has atsorbed the
carbonic acid and water which were produced by the combination of
the ingredients of the candle with the oxygen of the air.
434. In order that any combustible substance shall burn, or,
in other words, in order that brisk chemical action shall occur
between the combustible and the oxygen of the air, it must first
be heated to a certain temperature, and then kept at that heat.
The temperature at which any substance takes fire is known as
the kindling-temperature of that substance.
Exp. 211.-Place a small bit of phosphorus and another of sulphur,
not in contact with the first, upon a frag-ment of porcelain 6 or 8 c.m.
across, and heat them slowly over the gas-lamp, The phosphorus will
soon take fire at a temperature of 68o-70~; but the sulphur will not
inflame until the temperature of the porcelain support has risen to
about 250, as can be readily ascertained by the thermometer.
As was just now said, the degree of heat necessary to start
any fire must be kept up continually, or the fire will go out.
Whenever burning bodies are cooled below the kindling-temperature they are extinguished,-the chemical action which occasioned the appearance of heat and light ceases.
Exp. 212.-Pile up upon an iron grate, thick in metal, and supported in such manner that air may enter beneath it, several pieces of
red-hot charcoal. The charcoal will go on burning until nearly all of it
has been consumed; for the heat generated by the combustion of the
portions first burned keeps up the temperature necessary to kindle the
subsequent portions.
Upon a cold grate, similar to the one just employed, scatter about
several small pieces of red-hot charcoal, taking care that no two pieces
of the coal shall come in contact, or be placed so as to heat one
another. Each of the pieces of charcoal will soon cease to burn; for
the metallic grate is so good a conductor of heat that it removes heat
from the isolated pieces of charcoal more rapidly than these can pro



358           FLATMES PUT OUT BY GOOD CONDUCTORS.
duce it; consequently the temperature of the charcoal is soon reduced
to below the kindling-point.
A result similar to the foregoing is obtained when any fire is broken
up and scattered about to such an extent that its several portions cannot assist in one another's combustion-though if, instead of being
placed upon iron, the separate glowing coals be laid upon ashes or dry
earth, they will be extinguished only after a much longer time, for
ashes and dry earth are very poor conductors of heat in comparison
with iron. ~
In all ordinary fires the heat evolved by the combustion of
the fuel is more than sufficient to maintain a temperature higher
than the kindling-point of the fuel-though, generally speaking,
the fuel becomes at last so clogged with ashes, that the oxygen
of the air cannot get at the remaining combustible matter in
sufficient quantity to maintain lively chemical action.
435. Precisely as coals can be extinguished by placing them
upon cold metal, so flames may be put out..Exp. 213.-Upon a ring of the iron stand place a sheet of clean
wire gauze about 10 c.m. square; lower the ring so that the gauze
shall be pressed down upon the flame of a lamp or candle almost to
the wick, as shown in Fig. 71. No flame will be
seen above the gauze, but instead of flame a cloud  Fig. 71.
of smoke. The gauze is a mere open sieve; there
is nothing about it which can prevent the gas,
which was just now burning with flame above
the wick of the candle, from  passing through.
Indeed it may be seen from the smoke that the
particles of carbon which, in the orioinal undis-   i
turbed flame, were becoming incandescent, and so affording light, do
now actually come through the gauze.
The explanation of the phenomenon is sirply that the metallic
sieve conducts away so much heat that the temperature of the candleflame is reduced to below the kindling-point. That this is really so is
proved by the fact that, after the gauze has become sufficiently heated
by long-continued contact with the flame below-after it has attained
the kindling-point of the candle-gas, it will no longer extinguish the
flame. In like manner, a candle-flame may be cooled to such an extent that it will go out, by placing over it a small coil of cold copper
wire, while if the wire be previously heated the flame will continue
to burn.
If the smoke and unburned gas which has passed through the cold




SAFETY-LAMPS.                     359
wire gauze be touched with a lighted match, and so brought to the
kindling temperature, it will burst into flame.
The power of wire gauze to prevent the passage of flame has been
usefully applied in several ways, notably for the prevention of explosions in those coal-mines which are liable to accumulations of marshgas. For this purpose safety-lamps are constructed by enclosing an
ordinary oil-lamp completely in wire gauze, so that the flame within
the gauze cannot kindle any combustible or explosive gas into which it
may be carried. In case such a lamp be carried into a place filled with
explosive gas, the latter will, of course, pass into the lamp through the
meshes of the gauze and burn within the cage. This combustion gives
warning of the presence of the dangerous gas, and indicates to the
workman that he should withdraw firom the locality; the gas can then
be expelled by appropriate methods of ventilation.
The wire-gauze lamps employed in chemical laboratories (for
one form of which see Appendix, Fig. viii.) are simply applications of the same general idea.
Exp. 214.-Beneath the sheet of wire gauze of Exp. 213, place an
unlighted ordinary gas-lamp (Bunsen's burner), at such distance that
the gauze shall be 3 or 4 c.m. above the top of the lamp; turn on the
gas and light it above the wire gauze; it will continue to burn on the
top of the gauze for an indefinite period, for the gauze will, in this
case, always be kept cool by the cold gas which is continually passing
through it. Carefully and gradually lift the ring which carries the
gauze, and determine how far it is possible to lift the gauze above the
gas-jet without extinguishing the flame.
The student will remember that other experiments illustrating the
influence of cooling agencies in extinguishing combustion (Exps. 136
and 154), have already been performed. Compare also Exp. 196 and
~ 200, as regards kindling.
436. An effect somewhat similar to that produced by wire
gauze is often seen in ordinary fires. When a mass of red-hot
anthracite, charcoal, or coke is burning freely upon a grate in
the open air, there is always a blue flame of carbonic oxide
burning above the coal. This gas results from the reduction of
carbonic acid by means of hot carbon, precisely as in Exp. 181.
Air enters at the bottom of the grate and combines with the hot
coal which it finds there to form carbonic acid (CO,). This carbonic acid, as it rises through the hot coal in the middle of the
fire, is deprived by the heated carbon of half its oxygen,
CO2 +  C =  2CO,




360                   FIcAMING FIRES.
so that two molecules of carbonic oxide gas finally emerge at the
top of the coal, instead of the single molecule of carbonic acid
which was formed at first. The carbonic oxide being combustible, will at once take fire on coming in contact with the air,
provided the temperature at the summit of the fire be equal to
the kindling-temperature of carbonic oxide. But if the temperature of the fire is in any way reduced below this point, as, for
example, by throwing on too large a quantity of cold fuel, which
is, of course. equivalent to covering the fire with a sheet of wire
gauze, then the carbonic oxide will be extinguished, and, escaping
into the chimney, will produce no useful effect.
Were it not for the formation of carbonic oxide, as above mentioned, neither anthracite coal, nor coke, nor charcoal would burn
with flame after having once got well on fire; they would simply
glow, as a single live coal, or a bar of metal, glows when taken
from the fire.
437. In heating steam-boilers and other large vessels, it is
often a point of great importance to obtain from the fuel a large
flame, in order that the heat from the fuel may be quickly distributed and brought into contact with the matter to be heated.
With anthracite and coke this result is effected by placing beneath the grate, upon which the fuel is burned, a quantity of
water. From this water steam gradually rises, as hot ashes and
cinders fall into it and as heat radiates down upon it from the
fire above. The steam, as it enters the fire, is decomposed by
the hot coal (see Exp. 156), in accordance with the following
reaction,
C +  H1O  =  CO +  211,
and the combustible gases thus obtained are superadded to the
carbonic oxide which is formed in due course from the action of
air upon the coal. All these gases burn again to carbonic acid
and water above the fire, where air is thrown in to meet therm
through appropriate orifices. In this use of water as an adjunct
to the combustion of coal, the absolute amount of heat given off
by the fuel is in nowise increased; but in many instances much
heat may undoubtedly be saved by thus equally distributing and
applying it by means of flame.
438. On the other hand, furnaces are sometimes seen con



CARBONIC OXIDE SHOULD BE BURNED.        361
suming fuel under such conditions that all the carbonic oxide
produced within them escapes unburned into the chimney. In
such cases, more than two-thirds of the amount of heat which
the fuel is capable of yielding must necessarily be lost; for while
1 gramme of charcoal gives off in burning to carbonic acid
8080 units of heat (~ 55), 1 gramme of carbon in burning to
carbonic oxide gives off only 2473 units of heat. The number
last given is determined as follows. It has been found, by direct
experiment, that 1 gramme of carbonic oxide, on being burned
to carbonic acid, yields 2403 units of heat; carbonic oxide is
composed (~ 422) of one atom of carbon, weighing 12, and one
atom of oxygen weighing 16,-the weight of the molecule of carbonic oxide being consequently 28. In one gramme of carbonic
oxide, therefore, there can be only 1 2-=04286 of a gramme of
carbon; but 0'4286: 1=2403: x-=5607, whence it appears
that there is evolved by one gramme of carbon in carbonic oxide
5607 units of heat when this carbon unites with the additional
oxygen to form carbonic acid; and the difference between this
number (5607) and the number (8080) denoting the amount of
heat given off by one gramme of charcoal in burning to carbonic
acid will show how much heat is evolved by one gramme of
carbon burned'to carbonic oxide: 8080-5607=2473, as above
stated.
In order to thoroughly burn the carbonic oxide in any case,
the stove or furnace should be so arranged that a volume of air,
as large as that which has already passed through the fire, can
be constantly supplied to the carbonic oxide and nitrogen as they
emerge from the coal, and be intimately mixed with these gases
while they are still hot.
439. The amount of air needed for the complete combustion of
coal or other fuel can always be readily calculated. We have
only to determine how much oxygen will be needed by the combustible, and then how much air must be taken in order to supply
this oxygen. Let it be supposed, for example, that we wish to
learn how much air is needed in order to burn one kilogramme
of charcoal. Having learned the full significance of the formula
CO2, a moment's consideration of this formula informs us that,
for every 12 parts, by weight of carbon, 32 parts by weight of




362                 CHLORIDE OF CARBON.
oxygen are needed in order to its complete combustion, or, for
one part of carbon, 2-67 parts of oxygen. Air contains 231 per
cent. of oxygen by weight; hence the proportion 23-1: 100=
2'67: x= 11'558, from which it appears that to burn 1 kilo. of
charcoal, 11'558 kilos. of air are needed. Since the weight of a
litre of air, at the ordinary temperature, is only 1'2258 grm.,
these 11'558 kilos. of air will occupy about 9429 litres, or, in other
words, nearly 92 cubic metres. In round numbers, it may be
said that about 12 kilos., or 93 cubic metres, of air are required
to burn one kilo. of charcoal. For coke and anthracite, corrections must, of course, be made for the ashes which they contain,
as well as for a certain portion of hydrogen which may also be
present. If a gramme of pure carbon will disengage 8080 units
of heat, a gramme of well-burned coke, containing 15 per cent. of
ashes, will disengage only 686S units.
440. Chloride of Carbon (CCl4).-Chlorine does not unite directly with carbon; but several compounds of the two elements
can be obtained by subjecting compounds of carbon and hydrogen to the action of chlorine. Of these compounds, only the socalled bichloride (CCI4) need here be mentioned, the others being
usually treated of in works upon organic chemistry. Bichloride
of carbon may be obtained by the action of chlorine on marshgas, by subjecting chloroform or wood-spirit to the action of an
excess of chlorine in sunlight, by passing a mixture of bisulphideof-carbon vapor and chlorine through a red-hot porcelain tube,
or by the action of quinquichloride of antimony upon bisulphide
of carbon:
CS2 + 2SbCl, = CC14 + 2SbCl3 + 2S.
At the ordinary temperature of the air it is a transparent, colorless liquid, of pungent aromatic odor, boiling at 770, and having'
a specific gravity of 1'56. At -23~ it solidifies in the form ofcrystals of pearly lustre. The specific gravity of its vapor has
been determined to be 76'96, which would indicate that a molecule of the vapor is composed of I atom of carbon and 4 volumes of
chlorine, condensed to 2 unit-volumes:For, since the weight of one atom of carbon is......  12
And the weight of 4 atoms of chlorine is....        142
The weight of the two volumes of gas produced would be.. 154




CARBON AND NITROGEN.                363
And the weight of 1 volume would be 77, a number with which the
experimental determination very nearly agrees.
In composition, this chloride of carbon (CC14) is analogous to
the hydride CH4 which we have already studied under the name
light carburetted hydrogen, or marsh-gas (~ 392), and in the
same way that the hydride may be converted into the chloride by
acting upon it with chlorine, so, conversely, the chloride, on being
brought into contact with water and sodium-amalgam (~ 97),
may be deprived of chlorine and converted back into the hydride,-the hydrogen of the water being substituted for the
chlorine, which, like the oxygen of the water, unites with the
sodium.
441. Compounds of Carbon and Nitrogen.-With nitrogen,
carbon forms a number of highly interesting compounds, which
may be found treated of in works upon organic chemistry. Prominent among these compounds are cyanogen, CN (~ 384), and
cyanhydric acid, HCN. Cyanhydric acid corresponds to chlorhydric acid and the other hydrides of the chlorine group of elements; by its action upon metallic oxides there may be formed
a series of cyanides of the metals, corresponding perfectly with
the metallic chlorides:
Mo20 +  2H(CN) =  2M1(CN) + H0.
M20 +  2HC1   =  2MCl   +  H20.
The group of atoms (CN) which constitutes cyanogen, acts, in
fact, as if it were a single element. In the same way that the
group NH4, called ammonium, is capable of replacing a metal
like sodium (see ~ 91), so the group CN can replace chlorine and
the elements allied to chlorine. Groups or knots of atoms, such
as these, are often called conipound radicals.  Cyanhydric acid,
or prussic acid as it is sometimes called, is notorious as a violent
poison. Many of the cyanides are of great importance in the
arts, particularly in processes of gilding, plating, electrotyping,
and dyeing, as well as in the preparation of pigments. Some of
them will be described hereafter under the respective metals.
442. Bisullhide of Carbon (CS2), or sulpho-carbonic acid, as
it is often called, is especially interesting from its correspondence
with the binoxide of carbon, carbonic acid, in accordance with the




364                 BISrLPHIDE OF CARBON.
general rule that sulphur compounds are analogous to the compounds of oxygen. With the metallic sulphides this compound
forms a series of salts, of the general formula MI2CS3, or M2S,CS2,
analogous to those formed by the union of carbonic acid with
the metallic oxides, the general formula of which, as we know, is
M2CO, or M2,,C02.
Bisulphide of carbon may be obtained by bringing the vapor of
sulphur in contact with red-hot charcoal. It is a mobile, colorless
liquid of 1'27 specific gravity, which refracts light powerfully. It
boils at 450 and evaporates rapidly at the ordinary temperature of
the air. The density of its vapor is 38'19. It has a peculiarly
fetid odor, and is very easily inflammable, burning with a blue
flame to carbonic and sulphurous acids. It is not soluble in water,
but is easily soluble in alcohol, and is itself a powerful solvent of
fats and various other substances; it has been, of late years,
somewhat extensively employed as a solvent of phosphorus (~ 274)
and of chloride of sulphur (~ 246) in the cold process of vulcanizing
caoutchouc.  Mixtures of its vapor with oxygen, air, nitric oxide,
and other gaseous oxygen compounds burn, without violent explosion, with a sudden brilliant flash of intensely blue light. This
flame is remarkable for its great actinic power; it acts very energetically upon a prepared daguerreotype plate, and causes a mixture of hydrogen and chlorine to combine almost as readily as
sunlight would (Exp. 53).
Exp. 215.-Fill a tall bottle, of the capacity of 700 or 800 c. c., with
nitric oxide gas, at the water-pan; cover the bottle with a plate of glass
and stand it upright upon the table; draw the cover aside far enough
to admit of the introduction, from a pipette, of 8 or 10 c. c. of liquid
bisulphide of carbon; replace the cover and leave the bottle at rest
for a fiw minutes, in order that the vapor of the bisulphide may have
time to diffuse through the nitric oxide.  Finally, touch a lighted
match to the opened mouth of the bottle and observe the brilliant
flame which is produced.




BORON.                        365
CHAPTER XXI.
BOBO N.
443. Boron is an element of the same natural family as carbon.
It is found in nature, in combination with oxygen, as boracic
a-cid, and in combination with oxygen and some of the metalsnotably as a biborate of sodium, commonly called borax, and as a
double borate of sodium and of calcium.
In certain volcanic districts in Tuscany, jets of steam mixed
with other vapors escape continually from cracks in the soil, and
bring to the surface small quantities of boracic acid. Since boracic
acid is not volatile, in the ordinary sense of the term, at temperatures so low as 100~, it appears that it is transported mechanically
by the steam, much in the same way that dust is carried along
by a current of air. The jets of vapor, laden with boracic acid,
are made to bubble through water as they escape from the earth;
this water retains the acid, and so concentrates it to a very considerable extent. After the water has become as highly charged
with the acid as has been found in practice to be desirable, it is
run off into pans, lower down upon the hill-side, beneath which
hot jets of vapor from the earth are caused to circulate. The
excess of water is thus evaporated by heat which the earth supplies, and the solution becomes so concentrated that, on cooling,
crystals of boracic acid separate from it. About 120 millions of
kilogrammes of water are thus evaporated annually, and 1,300,000
kilogrammes of boracic acid produced without the intervention of
any artificial motor or the consumption of any fuel. After having
been purified, the boracic acid is sent into commerce as such, or
it is treated with a hot solution of carbonate of sodium, and so
converted into borax.
Besides this principal source of the boron compounds, a certain
quantity of native borax is obtained from the mud and waters of
certain lakes in Tartary, Ceylon, Thibet, and California.  In
Peru, also, a mineral composed of borate of sodium and borate of
calcium is found associated with nitrate of sodium.




366                 ALLOTROPISH OF BORON.
444. The element boron has been obtained in two distinct allotropic conditions (~~ 162, 366) comparable with those of carbon.
It can be had amorphous like charcoal, and crystallized like the
diamond.
To prepare the diamond-like modification, fused boracic acid is intensely heated with metallic aluminum; a portion of the aluminum
takes oxygen away from the boron, while another portion of the molten
metal dissolves this boron as fast as it is formed. As soon as the
solution has become saturated (that is to say, when the melted aluminum has dissolved all the boron that it is capable of dissolving), a
portion of this boron is deposited in diamond-like crystals. When the
crucible and its contents are broken up after cooling, these crystals are
found lining cavities within the mass.
The amorphous modification may be prepared by heating together
boracic acid and metallic sodium or potassium beneath a layer of fused
chloride of sodium. After having been allowed to cool, the products
of the reaction are treated with water, which dissolves away the chloride and the borate of sodium, and leaves the boron as an amorphous
powder which may be collected upon a filter.
A third variety of boron, similar to graphite, was at one time described; but recent experiments have shown that the so-called graphitoidal boron is really a boride of aluminum (A1B2).
445. Of the properties of these varieties of boron, little need
here be said; they are analogous to, and closely resemble, the
corresponding modifications of carbon, which have already been
fully described. The diamond modification is transparent, and
sometimes colorless, though usually of a yellow or reddish color.
It crystallizes in the same forms as the real diamond, and refracts
light very powerfully. Its specific gravity is 2S68. It is almost
as hard as the real diamond, being capable of scratching the ruby,
and even of polishing the diamond. It is only with extreme
difficulty that it can be burned in oxygen, since a coating of boracic
acid soon forms which protects it from further action of the oxygen. It is remarkable that at the moment of its combustion it
swells up as the diamond does when intensely heated. Amorphous boron is an infusible greenish powder, which readily takes
fire on being heated in air or oxygen. It is, in fact, necessary, in
preparing it, to take care that the filters upon which it has been
collected shall be dried at low temperatures, lest the finely divided




COMPOSITION OF BORACIC ACID.            367
horon take fire spontaneously; if exposed to the sun's rays in
summer, the filters containing boron will often take fire as soon as
they have become dry. Unlike the other modification of boron, it
is readily attacked by most chemical agents. Like the corresponding modification of carbon, it is an energetic reducing agent.
No compound of boron and hydrogen has yet been discovered,
but it unites readily with chlorine, bromine, and iodine; it can
be made to combine also with fluorine, nitrogen, and sulphur.
446. The best-known of the compounds of boron is the oxide
B203, called boracic acid. This oxide occurs ready formed in nature, as has been said, and is the sole product when boron is burned
in oxygen.  The composition of boracic acid may be determined
synthetically by burning amorphous boron in pure oxygen, or by
treating it with nitric acid. A certain definite weight of boron
being taken in the first instance, the weight of the dry boracic
acid obtained is carefully determined, and from these data the
percentage composition of the boracic acid is calculated; 100
grammes of boracic acid consist of
Oxygen.........   6871 grammes.
Boron.. 31'29,
From the specific gravities of the vapors of chloride of boron and
of fluoride of boron, as determined by experiment, and from some
other rather inconclusive considerations which need not here be
dwelt upon, chemists have been led to admit that the atomic composition of boracic acid may be represented by the formula B,20,.
If boracic acid be really composed of 3 atoms of oxygen and 2
atoms of' boron, the weight of the atom of boron will follow from
the proportion 68'71: 31'29=(16 x 3): x, in which 16 equals the
weight of an atom of oxygen and x the weight of two atoms of
boron. Upon this assumption, the weight of one atom of boron
will be 10'93.
447. Boracic acid is but a feeble acid at the ordinary temperature, and may be set free from its compounds by almost any of
the acids, excepting carbonic acid.
Exp. 216.-Dissolve 4 grms. of borax in 10 grms. of boiling water,
in a beaker glass or porcelain capsule of 30 or 40 c. c. capacity, and
add to the solution 2'5 grins. of concentrated chlorhl dric acid. After




368                PROPERTIES OF BORACIC ACID.
the lapse of some time, hydrated boracic acid will be deposited from
the solution in the form of glistening, colorless plates or scales. These
crystals contain as much as 43'6 per cent. of water; their formula is
HBO.:, or, dualistic, 3H1120, B,03.
On being heated in a clean iron spoon, the crystals will first dissolve
in the water which they contain, or, as th~ fact is usually stated, they
will " melt in their water of crystallization; " if the heat be continued,
the mass will become pasty, and will swell up as the water is expelled.
After all the water has been driven off by strong heat, the anhydrous
acid is left as a clear, viscous liquid, from which long threads of the
solid acid may be drawn out by touching the surface of the liquid with
the end of a stick or glass rod, and then gently pulling away the stick
with the matter which has adhered to it.
If the fused acid be allowed to cool,'t will solidify to a hard, transparent glass, which soon cracks in every direction and splits up into
fragments.
Anhydrous boracic acid is of about 1.8 specific gravity; it is
odorless and destitute of corrosive power; it has a slightly bitter,
but not sour, taste. It is much more soluble in hot than in cold
water, and more soluble in alcohol than in water. It imparts to
the flame of burning alcohol a peculiar green tint, which is quite
characteristic, and affords a valuable test by which the presence
of the acid may be detected. Upon litmus and turmeric, boracic
acid acts somewhat differently from other acids.
Exp. 217.-Dissolve a little of the crystallized boracic acid of Exp.
216, in a teaspoonful of alcohol in a small porcelain capsule. Set fire
to the alcohol and stir the burning solultion with a rod, or agitate it by
jarring the dish. Or moisten a tuft of cotton with alcohol, strew upon
it some powdered boracic acid, and light the alchohol. In either case
the flame of the alcohol will be of a fine green color..trp. 218.-Pour into a test-glass 20 or 30 c. c. of a solution of blue
litmus; in a small quantity of water, contained in another test-glass
or tube, dissolve a little of the boracic acid of E:p. 216; add the
solution of boracic acid to the litmus, and observe that the color of
the latter changes to a brownish wine-red, decidedly different from the
bright clear red which is obtained by the action of other acids upon
litmus. If a large quantity of boracic acid, however, be added to a
small portion of the litmus solution, the latter will be colored strongly,
as if by a powerful acid.
Exp. 219.-Dip into a solution of boracic acid a slip of yellow turmeric paper, and observe that the yellow color is changed to brown,




CHLORIDE OF BORON.                  369
as it would be by ammonia-water or by any other alkalifie solution.
None of the other acids produce a like effect.
448. When an aqueous solution of boracic acid is boiled, an
appreciable quantity of the acid goes off with the vapor of water;
but the dry acid, when heated by itself, is nevertheless one of
the least volatile of all the acids. It does slowly sublime, however,
at a white heat, and may be completely evaporated, if left for a
long time in the hottest part of a porcelain furnace. As a consequence of this fixity, or lack of volatility, it follows that boracic
acid is a comparatively powerful acid at temperatures high enough
to volatilize the ordinary acids.  On being heated with nitrates,
or sulphates, for example, it quickly expels nitric or sulphuric
acid, and unites with the other ingredients of the salt, though
either of these acids would at once decompose the borate thus
formed, if they were collected and added to it at the ordinary
temperature. Even phosphoric acid is expelled by it from the
phosphates.
449. Chloride oj Boron (BC1,) is a colorless, mobile liquid, of
1-35 specific gravity, which boils at 170. The specific gravity of
its vapor has been found to be 58'78, a result which points directly
to the formula BC13 as representing the true composition of the
compound.
From the weight of 2 vols. of chloride of boron (5878 X 2). 117'56
Subtract the weight of 3 vols. of chlorine (35'5X 3).  106-50
and the remainder will be..                11-06
a number almost precisely equal to the weight of one atom of boron
as paeviously determined (~ 446).
Upon being mixed with water, chloride of boron decomposes, with
formation of boracic and chlorhydric acids:2BC13 +  3H20 =  B2O; +  6HC1.
Chloride of boron may be prepared by slightly heating amorphous
boron in an atmosphere of chlorine, or more readily by passing
a current of chlorine over a mixture of anhydrous boracic- acid
and charcoal, heated to redness in, a, porcelain tube. In presence
of the hot charcoal which stands ready to take oxygen from the
boracic acid, chlorine can takelboron aw-ay from boracic acid; and,
2 R




370                  FLUORIDE OF BORON.
conversely, in presence of the chlorine, ready to combine with the
boron, carbon can take away oxygen:
B,20  -  3C +  6C1 =  2BCl, +  3CO.
The method here described, of converting an oxide into a chloride
through the intervention of carbon, is a method of very general
applicability, and is often employed for the preparation of chlorides
of the metals.
450. Fluoride of Boron (BF13) is a colorless gas of 34'19 specific gravity, as determined by experiment. Upon the assumption that an atom of boron weighs 11, its specific gravity would
be (11+ 3 x 19)-2 = 34.  It fumes strongly in damp air, and,
by pressure, may be readily condensed to a colorless and very
mobile liquid. The gas is exceedingly caustic and corrosive; it
carbonizes and destroys wood and other organic substances, in
the same way as concentrated sulphuric acid. As with sulphuric
acid (Exp. 104), so here, the fluoride of boron unites with the
elements of water which are contained in the organic matter, and
the integrity of the latter is destroyed.
451. Fluoride of boron is absorbed by water rapidly and in
large quantity, 1 volume of water being capable of dissolving 700
or 800 volumes of the gas; but in the act of solution decomposition occurs as well, and there is obtained, not a simple solution
of fluoride of boron in water, but a mixture, or rather a compound, of fluorhydric and boracic acids:2BF13 +  3H20 =  B,2O3,6HF1.
The reaction is interesting in all its stages, inasmuch as it well
illustrates the vagueness and indefiniteness of a considerable class
of chemical reactions. When water dissolves fluoride of boron,
it increases in bulk to a considerable extent, and in density also,
its specific gravity rising as high as 1'77. Upon warming the
saturated solution, some fluoride of boron is again disengaged,
perhaps as much as one-fifth of all that had been absorbed; but
on continuing to heat the solutio'n, it distils over unchanged, and
the condensed liquid presents the appearance' of oil of vitriol. In
it the elements of boracic and fluorhydric acids are undoubtedly
held together in a loose condition of chemical combination. By
many chemists the compound is called fluoboric acid, though, in




FLUOBORIC ACI).                  371
order to avoid confusion, it would, perhaps, be better if the name
fluorhydrate of boracic acid were allotted to it; for when this
compound is largely diluted with water, boracic acid is deposited,
and another acid compound is left in solution, the composition of
which may be represented by the formula HF1,BF1,.  This new
acid is called indifferently fluoboric acid or fluorborhydric acid,
and the salts formed by its union with metals are called fluoborates. It is remarkable that the first-named compound, the fluorhydrate of boracic acid, upon being neutralized with alkalies,
yields, not mixtures of a borate and a fluoride of the alkali employed, but true chemical compounds, double salts, of the general
formula M20,B20,; 6MF1, whence the name fluoboric acid has
arisen. The best way of preparing the fluorhydrate of boracic
acid is to dissolve boracic acid by small portions in fluorhydric acid.
Fluoride of boron may be itself prepared by heating in a glass
flask a mixture of 1 part of fused boracic acid, 2 parts of powdered
fluorspar, and 10 or 12 parts of concentrated sulphuric acid:3CaFl2 +  B203 +  3H2SO4 =  3CaSO4 +  3H20 +  2BF1,.
On account of its easy solubility in water, the gas must be collected over mercury. Fluoride of boron may be employed as a
test to determine whether a given sample of any gas is completely
dry; if a few bubbles of it are added to the gas to be tested, the
slightest trace of moisture in the gas will be made manifest by
the appearance of white fumes of the fluorhydrate of boracic acid
above described.
452. Sulphide of Boron (B2S,) is a white, crystalline solid, decomposable by water, in accordance with the following formula:
B2S, + 31120 =  B203 +  3SHS.
It may be prepared by passing a current of sulrhide-of-carbon
vapor over a mixture of boracic acid and charcoal strongly heated
in a porcelain tube.
453. Nitride of Boron (BN) is a soft, white, amorphous solid,
tasteless, odorless, infusible, and non-volatile. It is, in general,
but little acted upon by chemical agents.
454. It will be remarked that while boron is closely analogous
to carbon in many respects, it differs from it decidedly in others.
Thus, while in their allotropic modifications the two elements are
2 B2




372                       sILICoN.
almost precisely alike in appearance and properties, and while
many of the salts of boracic acid are strikingly similar to the
corresponding carbonates, the compounds of boron are not comparable as regards their composition with the compounds of carbon. While one atom of carbon unites by preference with four
atoms of any member of the chlorine group (as in CC14), or with
two atoms of any member of the sulphur group (as in CO,), an
atom of boron unites with only three atoms of chlorine (BC13), or
two atoms of it unite with three atoms of oxygen or sulphur
(as in BO,). The exceptional character of the composition of
boron compounds will appear still more clearly in the next chapter,
where it will be shown that silicon, the third member of the carbon group, resembles carbon as regards the atomic composition
of its compounds, as well as in other respects. The atomic weight
of boron above given (~ 446) cannot be accepted as established
beyond a doubt, and it may happen that further investigation will
show that the boron compounds are really formed upon the same
type or pattern as those of carbon and silicon; but in face of the
experimental evidence now at hand, this view cannot be maintained. In the meantime the student will better understand
the physical and chemical properties of boron and its compounds,
if this element is studied in company with carbon and silicon,
which it so closely resembles, than if it were described in connexion with arsenic, antimony, and the other elements of the
nitrogen group, which form teroxides and terchlorides indeed,
but which present not the least other analogy to boron, either in
the simple or the compounded condition.
CHAPTER XXII.
SILICON.
455. Like carbon, silicon may be obtained in three distinct
allotropic conditions, which have been designated as amorphous,
diamond-like, and graphitoidal.  After oxygen, it is the most
abundant and widely diffused of all the chemical elements; at




MODIFICATIONS OF SILICON.              373
least one quarter of the solid crust of the earth is composed of it.
In combination with oxygen, it occurs in silicic acid, which is one
of the commonest substances upon the surface of the globe.
Common quartz and flint, as well as rock-crystal, agate, and the
like, are pure silicic acid. The yellow sand of sea-beaches and of
many sterile tracts of country is silicic acid contaminated with a
trace of oxide of iron.  In like manner, sandstones and a great
variety of other rocks are mainly composed of it.
In order to obtain pure silicon, the following methods may be employed:-When metallic potassium is heated with a double compound
of fluoride of potassium and fluoride of silicon, known as fluosilicate
of potassium, a violent reaction occurs, fluoride of potassium is formed
and silicon set free in the amorphous state,
2KFlSiFl4 + 4K = 6KF1 +  Si;
by washing with water, the silicon may then be readily freed from the
fluoride of potassium, which is soluble. The graphitoidkl modification
of silicon can be prepared either by heating the amorphous variety
very strongly, in which event the powder contracts upon itself and
becomes much more dense, or by melting together in a crucible a mixture of metallic aluminum, in excess, and fluosilicate of potassium; a
fusible double fluoride of aluminum and potassium.is formed, and the
silicon thus set free dissolves in the melted aluminum. If the aluminum be dissolved away, after the mass has become cold, by means of
chlorhydric acid, the silicon will be left in the form of hexagonal scales.
Still a third method must be resorted to, in order to obtain the diamond-like modification of silicon; a mixture of dry fluosilicate of
potassium, metallic zinc, and metallic sodium is thrown into a hot
crucible, the mass is covered with a layer of fluosilicate of potassium,
and the crucible covered. A lively reaction ensues, and the mixture
within the crucible fuses; the fused mass is then stirred with an iron
rod until the zinc begins to escape as vapor. The crucible is then removed fiom the fire and allowed to cool; within it there will be found
a button of metallic zinc filled with long crystals of silicon, which can
readily be isolated by dissolving the zinc in chlorhydric acid.
456. The amorphous variety of silicon is a brown powder,
which, when touched, soils the fingers. It may be melted at a
temperature not far from that at which cast iron becomes liquid.
When mixed with common salt and exposed to a degree of heat
strong enough to volatilize the salt, amorphous silicon changes to
the graphitoidal variety, as has been already remarked. Amor



374                SILICON AND HYDROGEN.
phous silicon burns readily in the air and in oxygen.  It is not
acted upon by acids, excepting fluorhydric acid, which dissolves
it with disengagement of hydrogen.
Graphitoidal silicon is very similar to true graphite; it crystallizes in lustrous hexagonal plates of a leaden-gray color, and
is an excellent conductor of electricity. It is, however, much
harder than graphite. By exposing it to an exceedingly high
temperature, it can be transformed to the diamond-like condition.
The diamond-like modification of silicon occurs in the form of
regular octahedral crystals, exhibiting a decided metallic lustre.
The specific gravity of these crystals is 2'49. They are less
hard than the corresponding crystals of carbon and boron, and
melt at the same temperature at which cast iron melts.
Either graphitoidal or diamond-like silicon may be heated to
redness in oxygen gas without burning to any appreciable extent,
for a film of silicic acid is formed which protects the remainder
of the silicon and prevents it from being consumed; but if the
silicon be heated together with a substance capable of furnishing
oxygen in presence of a base competent to unite with silicic acid
and form a fusible silicate, it can readily be oxidized and converted into a silicate.  Thus, when heated to intense redness
with carbonate of sodium, it decomposes the latter with evolution
of light and heat, and there is produced silicate of sodium, while
carbon is set free:
Na 2O,CO2 +  Si =  Na2O,SiO2 +  C.
In a similar way, if silicon be heated in very concentrated solutions of caustic potash or soda, water will be decomposed and a
silicate of potassium or of sodium formed, while hydrogen is set
free:2NaHO0 +  H120 +  Si =  Na2SiO3 +  4H.
Diamond-like silicon is not attacked at the ordinary temperature
by any of the acids, excepting a mixture of fluorhydric and nitric
acids, by which it is converted into fluoride of silicon.  Hot
chlorhydric acid gas, as well as chlorine, attacks it readily.
457. Silicon and Hydrogen (SiH4).-A  gaseous compound of
these elements, called siliciuretted hydrogen, may be obtained,
mixed with free hydrogen, by acting upon silicide of magnesium
(M3gSi) with chlorhydric acid; other methods of preparing this




OXIDES OF SILICON.                 375
gas give a purer product. It is a colorless gas, which takes fire
spontaneously on coming in contact with the air, and burns to
silicic acid and water. On being heated in a tube, out of contact
with the air, it is decomposed into free hydrogen and free silicon,
and the latter is deposited upon the walls of the tube as a shining
mirror, similar to the mirrors of arsenic and antimony obtained
in Exps. 133, 136.
Silicon and Oxygen.-Two compounds of these elements have
been discovered, though but one of them is as yet well known.
4.58. Oxide of Silicon is described as a white, amorphous,
hydrated substance, so light that it floats upon water; it may be
obtained by treating with water at 00 the compound of silicon,
hydrogen, and chlorine described in ~ 472; its composition is
probably that represented by the formula SiH2203. It is scarcely
at all acted upon by cold water; but in presence, of alkalies it
decomposes water rapidly, with evolution of hydrogen and formation of the higher oxide of silicon, silicic acid, directly to be
described. At temperatures above 300~ it suffers decomposition,
breaking up into silicic acid and siliciuretted hydrogen. It burns
brilliantly in the air, and still better in oxygen. None of the
acids, excepting fluorhydric acid, have any action upon it.
459. Silicic Acid (Si02) constitutes at least one half of the
crust of the earth. Besides being thus abundant, it is one of the
most important acids known, regarded merely as a chemical
agent.  It is found everywhere, sometimes free, as quartz and
sand, sometimes in combination with metallic oxides, in the form
of salts known as silicates. It occurs more or less abundantly
in all soils, plants, and waters, while rocks like granite and
earths like clay are largely composed of it.  The common minerals feldspar and mica are double silicates of aluminum and
potassium, and of aluminum, iron, and potassium respectively.
In plants, the silicic acid (or silica, as it is often called) is contained particularly in the outer covering of the stalks and the
husks of grain.  The cuticle of rattan, for example, contains a
large proportion of silica, and the same remark is true of most
of the grasses and grains. The value of the plant called horsetail (Equisetum) as a polishing or scouring agent depends upon
the large quantity of silica contained in it.




376                  SOLUBLE SILICIC ACID.
460. Silicic acid occurs in at least two distinct isomeric modifications,-in one of which it is completely insoluble in water and
in acids, excepting fluorhydric acid, and is only slowly soluble in
boiling potash-lye; while, when in the other modification, it dissolves easily in a solution of potash, and may be retained in solution in considerable quantities both by water and by acids. Rock
crystal, as it is found in nature, may be regarded as the type of
the first or insoluble modification, and the soluble variety may be
prepared artificially by treating the solution of some one of the
soluble silicates with chlorhydric or sulphuric acid.
Exp. 220.-In a block of charcoal 12 or 15 c.m. long by 4 or 5 c.m.
wide and thick, scoop out a small cup-shaped cavity large enough to
hold a pea; place in this cavity a fragment of amorphous quartz or
flint, and heat it intensely by means of the blowpipe (Exp. 202) during
several minutes. When the quartz has become red-hot, suddenly
throw it from the coal into a dish of cold water; it will break up into
numerous small fragments or become filled with a multitude of cracks,
and can hence be readily pulverized. Grind the broken quartz to fine
powder in a wedgwood, or, better, in an iron or agate mortar; weigh
out 1 grm. of the powder, also 2 grms. of caustic soda, and to the mixture of these ingredients add 8 or 10 c. c. of water. Boil the mixture
in a porcelain dish for an hour or two, taking care to add, from time to
time, water enough to supply that lost by evaporation; then pour the
solution into a tall, narrow bottle, and leave it at rest until the undissolved portions of silica have settled and the liquid has become clear.
The highly alkaline solution thus obtained may be regarded either as
an aqueous solution of basic silicate of sodium, or as a solution of
normal silicate of sodium in the soda-lye.
Eap. 221.-To one-third part of the solution obtained in Exp. 220
add 10 or 12 times its bulk of water, and to this solution add dilute
chlorhydric acid, drop by drop, until the liquor manifests a decided
acid reaction (Exp. 33). The liquid will remain clear, and no silicic
acid will be deposited. All the silicic acid which has been set freo
from its combination with the alkali remains dissolved in the acidulated water.
If the clear solution be placed in a dialyzer (~ 327), all the chloride
of sodium, together with the free chlorhydric acid, will pass off through
the parchment-paper in the course of a few days, and there will be left
in the dialvzer nothing but pure silicic acid and water. Aqueous solutions of silicic acid, containing from 5 to 14 per cent. by weight of the
acid, may be thus obtained. This pure aqueous solution of silicic acid




SILICA IN NATURAL WATERS.                377
exhibits a decided acid reaction. It has little or no taste, though when
applied to the tongue it occasions an unpleasant sensation. It is by no
means so permanent as the solution, above mentioned, acidulated with
chlorhydric acid. When left to itself, the aqueous solution deposits
silicic acid after a while as an insoluble gelatinous precipitate. In like
manner, it has been observed that when chlorhydric or another acid is
added to the aqueous solution of basic silicate of sodium in no greater
quantity than is sufficient to exactly neutralize the alkali, the liquor,
though it remain clear for a considerable space of time, will gradually
become cloudy and deposit silicic acid..Ixp. 222.-To one-third part of the solution obtained in Exp. 220
add at once enough concentrated chlorhydric acid to render the solution acid. All the silicic acid will immediately be thrown down as a
thick insoluble jelly.
Exp. 223.-Place in a small porcelain dish a portion of the solution
of silicic acid in acidulated water obtained in Exp. 221, evaporate
it to dryness upon a water-bath, and heat the residue, over the gaslamp, to a temperature of 180~ or 190~. Add water to the cold dry
residue, and observe that the silicic acid does not redissolve; it remains
as a fine white powder, which may be readily collected upon a filter
and there washed clean by pouring upon it several successive portions
of water. After having been once thoroughly dried, silicic acid is
completely insoluble in water.
461. By adding chlorhydric acid to a sufficient quantity of the
water of almost any spring or river, and then evaporating the
water to dryness, a small quantity of silicic acid will be found in
the residuum. The silica, in this case, may be combined with
an alkali, or may be held in solution by the action of an alkaline
carbonate.  Silicic acid which has been finely powdered, or,
better, that which has recently been precipitated, is soluble, to a
considerable extent, in aqueous solutions of the alkaline carbonates, or even of the bicarbonates, only a small proportion of the
carbonic acid being expelled by it as it dissolves. It is not improbable, therefore, that the silicic acid found in waters may be
retained in solution by force of this solvent action of the alkaline
carbonates. It is, at all events, a matter of fact that waters
highly charged with carbonate of sodium, such as flow from the
geysers or boiling springs of Iceland, contain, in solution, very
large quantities of silica, much of which is deposited as the




378                      sLICIC ACID.
water becomes cold, upon the rocks and other objects with which
it comes in contact; siliceous petrifactions are thus formed.
462. Besides the soluble and insoluble modifications above
mentioned, a distinction is made between the variety of silicic
acid which is found crystallized in nature and that which occurs
in the amorphous state. Crystallized silicic acid is anhydrous,
has a specific gravity of from 2-6 to 2'66, and is scarcely at all
acted upon by alkaline lyes. The amorphous silicic acid found
in nature contains a certain amount of water, has a specific gravity of only 241 to 2'2, and is soluble in alkaline solutions.
As it exists in the insoluble modification, either as found in
nature or as prepared by calcining the artificial product, silicic acid
is a white, tasteless solid, incapable of forming a cohesive plastic
mass with water. It is infusible, excepting at very high temperatures, such as may be obtained by means of the oxyhydrogen blowpipe, in the flame of which it melts to a colorless glass.
When melted, it is tough and viscous, and, like glass, can be
drawn out into fine threads which are exceedingly elastic, particularly if they be dipped, while white-hot, into water. Silicic
acid is not volatile by itself; but when exposed to a current of
gas or vapor at a white heat, portions of it are carried off by the
vapor; like boracic acid, it can be transported in considerable
quantities by superheated steam.
When in the soluble modification, whether in the gelatinous
hydrated condition or in the state of an air-dried powder which
has never been exposed to heat, it dissolves readily in solutions of
the caustic and carbonated alkalies, particularly if these be heated.
Some varieties of native silica, such, for example, as the soft,
pulverulent, infusorial earth found at the bottoms of many ponds
and swamps, are as readily soluble in the alkalies as that which
has been prepared artificially. All the varieties of silica can be
dissolved by heating them with alkaline lyes in close boilers under
a pressure of 4 or 5 atmospheres; in other words, even crystallized silica is soluble in the alkalies at high temperatures.
463. At the ordinary temperature silicic acid is but a weak
acid; almost any of the other acids are capable of decomposing
the aqueous solution of a salt of silicic acid and of expelling the




SILICATES.                     379
latter from its combination with the metal. But, as is the case
with boracic acid also (~ 448), at high temperatures the reverse
of all this is true, silicic acid being then capable of expelling all
acids which are more volatile than itself. Neither variety of
silicic acid is acted upon at any temperature by either carbon,
hydrogen, phosphorus, or chlorine, if these elements be taken
separately; but when exposed at high temperatures to the combined action of carbon and chlorine, chloride of silicon and carbonic oxide are produced, much in the same way that chloride of
boron is formed under similar circumstances, as has been shown
in ~ 449. When silicic acid is heated together with carbon, or
other reducing agents, in contact with metals like iron or platinum, some of it is decomposed, and a silicide of the metal is
formed. Platinum crucibles are often injured in this way by the
nascent silicon which forms when certain minerals are fused in
them.
464. Silicic acid combines with many of the metallic oxides
to form salts, many of which are of very complex composition.
Hundreds of silicates are found in nature as crystallized minerals,
and they may, in general, be formed by fusing together silicic
acid and the appropriate metallic oxide, or carbonate, in suitable
proportions. Most of the silicates are fusible, and the greater
number of them, when in the molten state, have the power of
dissolving, either an excess of base, or an excess of silicic acid over
and above the quantities corresponding to strict atomic proportions; hence it happens that mixtures of silicic acid and of bases
may be melted together in the most varied proportions. The
study of the silicates thus becomes difficult, since it is often impossible to determine whether a given silicate be really a definite
chemical compound or a chemical compound contaminated with
an excess of silicic acid or of the base; in very many instances,
moreover, it is probably true that the excess of silica, above the
normal atomic proportion, or of the base, is really held in combination by virtue of the chemical force, though it be held feebly
and indefinitely, as is the case with the ingredients of many solutions and metallic alloys. (See ~~ 49, 76.)
As a general rule, the simple silicates containing but a single
base, and no excess of either base or acid, crystallize in passing




380                 FORMUL1E OF SILICATES.
from the liquid to the solid condition; but the double silicates
(compounds formed by the union of two or more simple silicates,
and therefore containing at least two different metals) usually
solidify to a non-crystalline, homogeneous glass. All the common
varieties of glass consist essentially of such double salts of silicic
acid. In general, those silicates of the alkali-metals and of lead
which contain an excess of base are more readily fusible than the
normal or acid salts; those silicates which contain easily fusible
oxides melt at correspondingly low temperatures; and many mixtures of two or more different silicates fuse at temperatures lower
than the melting-point of either of the simple silicates of which
the mixture is composed.
The silicates of potassium and of sodium, containing I or 2 or
more molecules of base to one molecule of the acid, are readily
soluble in cold water, and compounds containing as many as 4
molecules of the acid to one molecule of the alkaline base may be
completely dissolved in water when boiled therewith for a considerable time. These acid compounds (such, for example, as the
sodium-salt of the formula Na2O, 2-4SiO,) are much employed
in the arts under the name of Waterglass or Soluble glass. When,
however, the proportion of acid is more than 4k molecules to 1
molecule of the alkaline base, the silicate may, for all practical
purposes, be regarded as insoluble in water. The silicates of the
metals not included in the alkali group (~ 544) are. in like manner, insoluble in water, in the ordinary sense of the word, though,
in the last analysis, it would be hard to find any silicious minerals,
or glasses, which are not susceptible of being decomposed and
dissolved, to a greater or less extent, by water. Some metallic
silicates may be prepared, not only by the method of fusion, but
also by adding the solution of an alkaline silicate to the solution
of a salt of the metal whose silicate is desired.
465. The most commonly occurring silicates may be referred
to three or four general classes, similar to the classes of carbonates (~414): —1st. Normal silicates, of the general formula
M2O,SiO2, such as the crystallizable silicate of sodium Na2O,SiO2,
or the normal silicate of calcium, CaO,SiO2. 2d. Bisilicates, of the
formula M2O,2SiO2,, such as the bisilicate of calcium, CaO,2SiO2.
3d. Disilicates (basic silicates) of the general formula 2M20O,SiO2,




DECOMPOSITION OF SILICATES.            381
such, for example, as the silicates of iron, 2FeO,SiO2, of magnesium, 2MgO,SiO2, of manganese, 2MnO,SiO2, and of zinc,
2ZnO,Si 2.  Also, sesquisilicates, of the formula 2M20,3SiO2,
to which the mineral meerschaum may, perhaps, be referred, its
formula being 2TMg0,3Si0 + 2H + 20.
466. Many of the natural silicates may be decomposed by
digestion with the strong mineral acids, such as concentrated
chlorhydric acid; and this is especially true of those silicates
which contain a large proportion of base, and of those containing
water of crystallization. But many anhydrous normal or acid
silicates are not decomposed by any acid, excepting fluorhydric
acid. Fluorhydric acid attacks and dissolves, not only pure silicic
acid in all its varieties, but all silicates as well, and, other things
being equal, the rapidity of its action is in proportion to the degree
of its concentration. Crystallized silica, however, is much less
readily acted upon by fluorhydric acid than the amorphous variety;
it dissolves slowly and without development of heat, while amorphous silica not only dissolves rapidly, but the act of solution is
accompanied with the evolution of considerable heat. From those
silicates which are decomposed with difficulty by chlorhydric
acid the silicic acid separates out as a soft powder; but from
those which are more readily decomposed the silicic acid separates as a hydrated gelatinous mass when the finely powdered
mineral is digested with chlorhydric acid; in this case the
mineral is said to gelatinize with acids.  Some silicious minerals
may even be completely dissolved, silica and all, in dilute chlorhydric acid.
467. It is remarkable that, while some of the hydrated silicates
(such, for example, as the class of minerals called zeolites) can no
longer be readily decomposed by chlorhydric acid after they have
been ignited, there are other silicates (such as the minerals garnet, epidote, and idocrase) which, though scarcely at all acted upon
by chlorhydric acid when in the natural condition, may be decomposed thereby with gelatinization after they have been melted.
Facts like these would seem, at first sight, to indicate that silicic
acid exists in combination, sometimes in the insoluble, and at
other times in the soluble modification, and that it may be transformed from the one state to the other without being first set free;




382              COMPOSITION OF SILICIC ACID.
but some of them admit of being explained in another way.
When any silicate, no matter whether it be decomposable by
chlorhydric acid or unacted upon by this agent, is mixed with an
excess of one of the fixed alkalies, or alkaline earths, or with a
carbonate or nitrate of either of the alkaline or alkaline-earthy
metals, and the mixture then heated to intense redness, there will
be obtained a melted or agglutinated mass. which decomposes
readily on being treated with chlorhydric acid, and yields gelatinous silica. At the high temperature to which the mixture is
exposed, the strong alkaline or alkaline-earthy bases combine with
the silicic acid, and there are formed silicates of potassium, sodium,
calcium, or the like, easily decomposable by acids. Now, in the
case of the minerals garnet, epidote, and idocrase, above mentioned (which are attacked by chlorhydric acid after they have been
melted), it may be conceived that the strong bases (such as lime,
magnesia, and protoxide of iron) which are contained in these
minerals have acted upon the silicate of alumina (also contained
in them) in much the same way that an extraneous alkali would
act if it were fused with the powdered mineral.
468. Sulphuric acid, when concentrated or but slightly diluted,
acts much more energetically than chlorhydric acid upon the silicates, apparently because of its comparatively high boiling-point.
Common clay, for example (silicate of aluminum), which is but
little acted upon by chlorhydric acid, may be completely decomposed when digested with boiling concentrated sulphuric acid,
silicic acid being set free and sulphate of aluminum produced.
This decomposition, like that of the garnet above mentioned, is
in any event very much more readily effected if the clay be first
gently roasted or calcined. The mineral feldspar also, when in
fine powder, may be decomposed by boiling oil of vitriol.
469. Contrary to the view formerly held by many chemists,
the atomic weight of silicon is now thought to be 28, instead of
21, as formerly. This change makes the formula of silicic acid
SiO2, in strict analogy to carbonic acid, instead of SiO, which was
for a long time the accepted formula, or Si203.
The percentage composition of silicic acid has been determined
by direct experiment (by oxidizing a weighed quantity of silicon to silicic acid) to be 51.92 per cent. of oxygen and 48'08




CHLORIDE OF STLICON.                  383
per cent. of silicon.  But since it is difficult in this way to
oxidize the whole of the silicon, inasmuch as portions of it become covered with silicic acid, and are so protected from the
further action of oxygen, the above result can be regarded
only as a rough approximation to the truth. These numbers
would correspond to 29'63 as the atomic weight of silicon.  For
51-92: 48'08:: (32 = 2 atoms of O): (x = 1 atom of Si = 29'63).
More accurate experiments, however, upon another plan, have
since shoi4k that the real atomic weight of silicon is 28, or very
nearly 28, as will be fully set forth in the next section.
470. Chloride of Silicon (SiCl4) is formed when silicon is
heated in an atmosphere of dry chlorine, or more conveniently
by heating together finely divided silicic acid, charcoal, and
chlorine.
The experiment may be conducted as follows:-M- ix together intimately equal weights of lampblack and dry pulverulent silicic acid,
such as is obtained by decomposing an alkaline silicate with chlorhydric
acid and evaporating the residue to dryness, as in Exp. 223. Knead
into the mixture enough oil to render it plastic, and from the thick
paste thus formed make a number of small balls or pellets; roll these
pellets in powdered charcoal, place them in a Hessian crucible, cover
the crucible, and heat in a strong fire until all the oil has been distilled
off or decomposed. The mixture of silicic acid and charcoal is thus
left in an open porous condition, well adapted for the action of the
chlorine. Place the dry pellets in a porcelain tube about 2 c.m. in
diameter, connect one end of the porcelain tube with a U-tube surrounded with a freezing-mixture of ice and salt, and to the other end
attach a flask in which to generate chlorine, taking care to interpose
between this flask and the porcelain tube one U-tube filled with pumicestone soaked in oil of vitriol, and another filled with chloride of
calcium, in order that the gas may be thoroughly dried. Place the
porcelain tube in a suitable stove or furnace; pass into it a slow current of chlorine, in order to expel the air, and then build around it a
hot fire of charcoal or coke; finally, when the tube has become heated
to intense redness, pass in the chlorine as fast as it is absorbed. The
chloride of silicon will be condensed in the U-tube as a liquid which
has usually a yellow color, owing to the presence of dissolved chlorine.
In order to purify it, the liquid may be shaken in a dry flask with a
quantity of quicksilver in which a little potassium has been dissolved;




384          COMPOSITION OF CHLORIDE OF SILICON.
after having been decanted from the mercury and subjected to redistillation in dry vessels, it will be found to be pure.
It will be noticed that the method here described, of preparing chloride of silicon, is similar to that recommended for preparing chloride
of boron (~ 449). Chlorine alone is not capable of decomposing either
silicic or boracic acids; but in presence of carbon combination of chlorine and silicon is effected, at the same time that carbon and oxygen
unite to form carbonic oxide:SiO2 + 2C + 4C1 = SiCl4 + 2CO.
471. Chloride of silicon is a transparent, colorless, v'ry mobile,
volatile liquid, of pungent, acid, and irritating odor.  It fumes
strongly in the air, boils at 590, and is of 1'52 specific gravity.
The specific gravity of its vapor has been found to be 85'74, which
accords well with the supposition that a molecule of chloride of
silicon contains 4 volumes of chlorine and 1 volume of siliconvapor condensed to 2 volumes:For if to the weig'ht of 4 volumes of chlorine (4 X 35'5) =.    142
There be added the weight of 1 volume of silicon-vapor =..   28
The two volumes of gas produced will weigh....... 170
The weight of one volume of the gas should, consequently, be equal to
170 ^ 2 = 85.
Chloride of silicon is at once decomposed by water with formation of chlorhydric acid and deposition of gelatinous silicic
acid,
SiC14 +  2H20 = SiO2 +  4HC1,
no gas of any kind being set free, and no other products being
formed. The fact is important, since it shows at once that the
composition of chloride of silicon must correspond with that of
silicic acid-that in it four atoms of chlorine simply replace the
two atoms of oxygen contained in the silicic acid.  For knowing,
as we do, the composition of chlorhydric acid (~ 98) and of water
(~ 36), we are sure that, for every atom of chlorine eliminated
from the chloride of silicon, one atom of hydrogen is required,
in order that chlorhydric acid may be formed, and that for every
two atoms of hydrogen thus taken from the water, one atom of
oxygen will be detached. But as no oxygen escapes from the
solution, all of that derived from the water must evidently have




COMPOSITION OF CHLORIDE OF SILICON.        385
combined with the silicon to form silicic acid.  Now, it is a comparatively easy matter to determine the amount of chlorine in
any solution by adding to the solution nitrate of silver, and then
collecting and weighing the insoluble chloride of silver which is
formed.  Hence, if a weighed quantity of chloride of silicon be
decomposed with water, and the amount of chlorine contained in
the solution be determined, the difference between the weight of
chlorine found and the weight of chloride of silicon taken will
give us the weight of the silicon with which the chlorine was
combined.  By careful experiments, it has been proved in this
way that chloride of silicon contains in 100 parts by weight 16'52
parts of silicon and 83'48 parts of chlorine. But 83'48 parts of
chlorine are equivalent to 18'81 parts of oxygen, for
(35.5 x 2 =  71): 16:: 83.48: (x =  18.81).
Weight of two       Weight of
atoms of C1.     one atom of O.
Hence 18'81 parts of oxygen must have combined with every
16.52 parts of silicon contained in the solution; or, reduced to
per cent., every 46' 75 parts of silicon must have combined with
53'25 parts of oxygen; and if this proportion of oxygen lead
to the formula SiOn, the proportion of chlorine will, in like
manner, require the formula SiCl4.  Fiom  the composition of
chloride of silicon, as thus determined; the atomic weight of silicon may be accurately derived (compare ~ 469):53-25:    46'75      =     32:    (x =  28'1)
Per cent. of    Per cent. of   WTeight of tto    Weight of one
Oxygen         Silicon        atoms of        atom of..... X~           O(x ygen.        Silicon.
in Silicic Acid.
Or, more directly, by the proportion
83'48:    16'52      =     142         (x =  28'1)
Per cent. of   Per cent. of    Weight of        Weight of
Chlorine       Silicon,     four atoms        one atom
- v- — Y ----     of Chlorine.     of Siliconz.
in Chloride of Silicon.
At the ordinary temperature, liquid chloride of silicon does
not act upon potassium; but if this metal be heated in its vapor,
chloride of potassium is formed and silicon set free; the reaction
2 c




386                  FLUORrDE OF SILICON.
affords, in fact, a good method of obtaining silicon in the amorphous state. There is a bromide of silicon (SiBr4) analogous to
the chloride in composition and properties; it may be prepared in
a similar way. No compound of iodine and silicon has yet been
discovered.
472. Compound of Silicon, Hydrogen, and  Cdlorine.-By
passing a current of dry chlorhydric acid gas over crystallized
silicon, heated nearly to redness in a glass tube, and condensing
the product in a receiver immersed in a freezing-mixture, there
is obtained a colorless, fuming liquid, which boils at 340; the
density of its vapor has been determined at 4'64, the calculated
value being 4'69. At a red heat it decomposes, yielding chloride
of silicon, chlorhydric acid, and amorphous silicon. Its vapor is
very inflammable, and, when mixed with air or oxygen, detonates
violently on being lighted; chloride of silicon, together with a
smoke of chlorhydric and silicic acids, are the products of this
reaction.  Water decomposes it instantly, with production of
hydrated oxide of silicon, as has been already set forth (~ 458).
473. Fluoride of Silicon (SiFl4) is formed whenever dry fluorhydric acid comes in contact with silicic acid, either free or combined. It may be readily prepared by treating a mixture of siliceous sand and fluorspar with oil of vitriol.
In a -flask of about 250 c. c. capacity, place an intimate mixture of
9 grms. of finely-powdered quartz-sand, 9 grms. of powdered fluorspar,
and 53 grmis. of concentrated sulphuric acid. Heat the flask, and collect the gas which is evolved, in tall bottles of 150 c. c. capacity, at
the mercury trough. The sand and fluorspar should both be heated
before being used, in order that they may be perfectly dry; the flask
also and the mercury in the trough should be thoroughly dried; for
fluoride of silicon is decomposed by moisture, and the bulky precipitate
of silicic acid which would be formed might clog the delivery-tube of
the apparatus or coat the glass vessels and render them opaque.
The decomposition of the gas by water can readily be shown by
causing some of it to pass out from the delivery-tube into a bottle
inverted upon the mercury-trough, one half of which has been filled
with water and the other half with quicksilver. The water in the
bottle will soon become filled with gelatinous silica.  With the
quantities of material above given, there will be obtained first a bottle
of 150 c. c. capacity of air from the flask, which should be thrown
away; a bottle of the saie.size can then be filled with sufficiently




FLUORIDE OF SILICON.                387
pure fluoride of silicon, and in a third bottle the decomposition with
water may be shown.
Instead of sand, coarsely powdered glass may be placed in the flask
as the source of silicon. In any event the glass of the flask will be
somewhat corroded; it is better to conduct the operation in a platinum
retort, if such a vessel can be had.
The reactions which occur among the materials in the flask may be
represented by the following equations: —In the first place, fluorhydric
acid is formed by the action of sulphuric acid upon the fluoride of
calcium,
CaFl2 + H2SO4 = CaSO4 + 2HF1,
and the fluorhydric acid then attacks the silica,
SiO2 + 4HF1 = 2H20 + SiF14.
The water here formed unites with the sulphuric acid which has been
employed in excess, and is thus prevented from acting upon the fluoride of silicon. The reaction last named is essentially the same as
that which occurs when glass is etched with fluorhydric acid 0gas
(~ 158). When the fluoride of silicon is brought into contact with
water, there is produced, together with the silicici acid, fluosilicic acid
(2HF1,SiF14), a compound to be described directly.
474. Fluoride of silicon is a colorless gas which fumes strongly
in the air, especially if the air be moist. It has an acid odor
and taste, extinguishes combustion though itself uninflammable,
has no action upon glass, can support a high degree of heat
without being decomposed, and may be condensed by pressure
and cold to a colorless, very mobile liquid. Its specific gravity
has been determined to be 51'98, or 3'6 as compared with air.
It can consequently be readily collected by displacement, if
pains be taken to protect it from contact with moist air by means
of drying-tubes.
The specific gravity of this gas points very decidedly to the
formula SiFl4, and consequently to the number 2B as the atomic
weight of silicon, and to the formula SiO2 for silicic acid.  For, if
a molecule of gas be composed of
One volume of silicon-vapor =......   28
And four volumes of fluorine (19 X 4) =...   76
Condensed to two volumes =..104
One volume of the gas will weigh 52, or almost precisely as much as
has been indicated by experiment.
2c2




388                    FLUOSILICIC ACID.
The behavior of the gas with water is peculiarly interesting.
If water be sprinkled into a bottle filled with fluoride of silicon,
or if the gas be conducted into water, decomposition occurs instantly, and silicic acid is deposited in the gelatinous condition.
At first sight it might seem as if the reaction were strictly analogous to that which occurs when chloride of silicon is mixed with
water (see ~ 471), and that it could be represented by the formula
SiFl4 + 2110 = SiO2 + 4HF1;
but if it be remembered that fluorhydric acid, unlike chlorhydric
acid, is capable of dissolving silica, it will be at once evident that
the reaction between water and fluoride of silicon cannot be
directly comparable with the other reaction, in which chloride of
silicon is concerned. The reaction which really occurs between
water and fluoride of silicon may be represented by the following
equation:3SiF14 +  2H20 =  SiO2 +  2(2HF1,SiFl4),
— a double compound of fluorhydric acid and fluoride of silicon,
known as fluosilicic acid, being the other product besides silica.
Fluosilicic acid may be prepared either in the manner indicated in
5 473, or more conveniently by conducting fluoride of silicon gas into
an upright bottle full of water; in this case, however, the orifice of the
tube which delivers the gas should be immersed in a layer of mercury,
3 or 4 c.m. deep, beneath the water, so that the gas may bubble up
through the mercury, and the water never come in contact with the
mouth of the delivery-tube. In this way the tube may be kept free
from the silicic acid which would quickly clog it if it opened directly
into the water. As the bubbles of fluoride of silicon escape from the
mercury, each of them becomes covered with a thick crust of gelatinous silica, which it carries with it to the surface of the water. Sometimes one of these crusts will remain adhering to the mercury and will
be gradually pr'olonged into the form of a tube extending from the
mercury to the surface of the water. Such tubes should, however, be
broken up, by stirring the liquor, lest fluoride of silicon escape through
them into the air without coming in contact with water.
Another modification of the method of preparing fluosilicic acid is to
conduct the gas into a large flask containing but little water, and to
agitate this flask so that its sides may be continually moistened with
water, although no water can come in contact with the opening of the
gas delivery-tube. The following quantities of materials have been




FLIOSILICIC ACID.                   389
found convenient in practice. In a flask of 600 or 700 c. c. capacity
place a mixture of 35 grins. of powdered sand, 35 grins. of powdered
fluorspar, and 210 grins. of concentrated sulphuric acid. Place the
flask upon a sand-bath over the gas-lamp, and by means of a deliverytube connect it either with the bottom of a tall bottle containing
150 c. c. of water and enough mercury to form the layer above described, or with the middle of a flask of 700 or 800 c. c. capacity, and
containing 150 c. c. of water. In the latter case the receiving-flask
must be agitated, as aforesaid, so that its walls may be kept moist.
When the evolution of gas has ceased, the gelatinous mass in the bottle
should be thrown upon a piece of cotton cloth, and the liquid contained
in it separated from the solid matter by pressure. The cloudy liquor
thus obtained may be rendered clear by filtering it through paper.
The clear liquid is a strong aqueous solution of fluosilicic acid, and the
solid matter is hydrated silicic acid.
475. Fluosilicic acid is known only in aqueous solution.  The
saturated solution is a transparent, colorless, fuming, and very
acid liquid, which has no action upon glass, and may consequently
be kept in glass bottles.  It cannot be distilled without suffering
decomposition, nor can it be evaporated beyond a certain degree
of concentration without breaking up into fluoride of silicon and
fluorhydric acid. If the evaporation is conducted in a glass -vessel,
the silica of the glass will retain the fluorhydric acid, and only
fluoride of silicon and water will be set free. The reaction which
occurs in this case,
SiO2 +  2(2HF1,SiFl4) =  2H20 +  3SiF,,
is just the reverse of that which takes place when fluoride of silicon is decomposed by watcr2H20 +  3SiF14 =  SiO2 +  2(2HFl,SiF14).
With bases fluosilicic acid unites to form compounds known as
fluosilicates, such as the fluosilicate of potassium, K2SiFl6 
2KFl,SiF14, and fluosilicate of barium, BaSiFl6 =  13aF12,SiFl4,
if no excess of the base be present; but if an excess of the base
be added, silicic acid will be precipitated, and the whole of the
fluorine will unite with the metal contained in the base to form
a fluoride. In the first case, where no excess of base is employed,
the reaction may be thus represented:2KX10 +  H2SiF16 =  2H120 +  K2SiFle.




390                   THE CARBON GROUP.
In the second case, where an excess of the base is present,
6KHO +  IT2SiF1  =  4H20 +  SiO2 +  6KF1.
Most of the fluosilicates are easily soluble in water; but some of
them are so nearly insoluble that the acid is sometimes employed
as a precipitant of various oxides. It is remarkable that precisely those bases, namely the alkalies, which form soluble salts
with almost all the other acids, should here yield insoluble compounds. Fluosilicic acid is in fact a good reagent for the detection
of potassium; and it is valuable as a means of removing potassium
from many of its salts when we desire to obtain in a free state the
acids which these salts contain. (Compare ~ 124.)
476. Sulphide of Silicon (SiS2) occurs in white, needle-like
crystals, unalterable in dry air. It is volatile at the temperature
of redness, and is decomposed at once by water, with deposition
of gelatinous silica and evolution of sulphydric acid:
SiS2 +  2H20 =  SiO2 +  C2H2S.
Like sulphide of boron, it may be obtained by passing the vapor
of bisulphide of carbon over a mixture of silicic acid and carbon
heated to redness.
477. As has appeared abundantly from the foregoing, the three
elements carbon, boron, and silicon constitute a distinct natural
family. This family differs in character from each and every one
of the other natural groups of elements hitherto described (~~ 152,
257, 364).  The occurrence of car.bon and silicon in three distinct and extraordinary modifications (as diamond, graphite, and
charcoal), and of boron in the diamond and amorphous modifications, distinguishes this group from all others. The eminently
refractory nature of the several members, and their fixity as regards heat and solvents, are marked characteristics of the three
elements. The vitrifiable character of the oxides of boron and
silicon (boracic and silicic acids), and of the berates and silicates
of many of the metals, is a noteworthy resemblance between these
elements. Remarkable resemblances between the hydrated carbonate, borate, and silicate of sodium often manifest themselves
to the chemical manipulator. The corresponding compounds of
carbon, boron, and silicon, with the members of the chlorine group,
and with the members of the sulphur group, have similar proper



SODIUM.                     391
ties, result from like reactions, and suffer, for the most part, analogous decompositions; but, in the present state of the science,
only the compounds of carbon and silicon can be said to have also
the same atomic constitution.  The three members of this group
are not arranged in the order of their atomic weights, as has been
the case with all the preceding groups; but if, in the future,
boracic acid comes to be written BO2, the atomic weight of boron
will be 14, and therefore higher than that of carbon. The existing anomaly in the arrangement of the group will then disappear.
The collocation of these three elements, and the order in which
they now stand, is based upon too many natural resemblances to
be lightly set aside.
CHAPTER XXIII.
SOD IU I.
478. This abundant element is chiefly found in nature in the
state of chloride, nitrate, carbonate, borate, and silicate. The
most abundant of its compounds is common salt, which is the
combination of sodium with chlorine (NaCi). Sea-water contains
two and a half per cent. of salt, and enormous deposits of the same
substance are found in the solid crust of the earth. There are
also many natural salt springs, whose waters yield on evaporation
the chloride of sodium which they hold in solution. Sodium also
occurs, in the condition of silicate, in very many common minerals
and rocks. From the soil which has resulted from the disintegration and decomposition of these minerals and rocks, and from
its soluble compounds, like common salt, sodium enters into plants,
and thence into animals. Its chloride is one of the essential
mineral constituents of the food of man and other animals. On
account of the inexhaustible abundance of common salt, this substance constitutes the chief source from which all manufactured
compounds of sodium are more or less directly derived; one other
natural sodium-containing mineral, however, deserves mention as
a source of sodium compounds-the mineral Cryolite, a double




392                  CHLORIDE OF SODIUM.
fluoride of sodium and aluminum. Nitrate of sodium (NaNO,), a
somewhat deliquescent and very soluble salt, occurs abundantly
on the surface of the soil in certain desert districts of Peru.
When heated, this salt first fuses and then undergoes decomposition. It is employed in the manufacture of nitric and sulphuric
acids and as a manure; but as a source of sodium-compounds
it is comparatively insignificant.
479. Chloride of Sodium (NaC1).-There is but one chloride of
sodium-common salt. This natural mineral is, when pure, a
colorless, transparent, anhydrous stone, which crystallizes in cubes,
dissolves readily in about three times its weight of cold water,
and possesses a specific gravity of 2-15, and an agreeable taste,
which, because familiar, is the representative or type of that peculiar savor called saline. A saline taste means a taste suggestive of
that of common salt, just as the phrase s" saline substance" characterizes a very large class of bodies which resemble more or less
in appearance and properties the longest- and best-known of all
such substances, common salt.
There are three sources of salt-the beds of the native mineral,
saline springs; and sea-water. In all cases in which the salt is obtained
from its solution in water, evaporation by fire, or by the heat of the
sun in warm, sunny climates, is necessary. When pure enough, the
rock-salt is mined like any other ore; but when it is mixed with earth
or other impurities as it lies in its natural bed, the solubility of the
chloride of sodium in water is availed of to fiee the salt from its insoluble impurities, and to facilitate the lifting of it to the surface of the
earth. Watei is let into the bed of salt, and allowed to remain there
till it has become saturated; the brine is then pumped out and evaporated. Some natural brine-springs contain so small a proportion of
salt that some cheaper mode of evaporation than by fire is essential to
their profitable working. Such waters are concentrated by a process
termed graduation. The brine is pumped up to a sufficient height, and
then allowed to trickle slowly over large stacks of fagots, which are
sheltered by a roof from rain, but are freely exposed to the prevailing
wind. The brine, thus diffused over a very large surface, is rapidly
concentrated by the draft of air. By repeating the process a moderate
number of times, a weak brine may be brought to a degree of concentration at which evaporation by fire may be employed. Since almost
all brines contain sulphate of calcium (CaSO4) in solution, the surfaces
of the fagots employed in the graduation process become covered with




SOLUBILITY OF SALT.                   393
a stony coating of this comparatively insoluble salt. If the strong brine
is boiled down rapidly, a fine-grained table-salt is obtained; if it is
slowly evaporated, a hard, coarsely crystallized salt is the product.
During the earlier stages of the evaporation a deposit is formed, consisting principally of sulphate of sodium  and sulphate of calcium.
Finally there remains a thick mother-liquor, from which no more
chloride of sodium will crystallize, but which contains the more soluble
salts of the original brine, such as chloride of calcium, and chloride
and bromide of magnesium, besides a large proportion of common salt
which cannot be separated from the liquor. Such mother-liquors are
sometimes so rich in mnagnesium-salts as to be advantageously worked
for these substances, and they are also sometimes profitable sources of
bromine. Considerable quantities of magnesium-salts and of bromimne
have also been extracted from concentrated sea-water after all the
available chloride of sodium has been withdrawn. The salt of commerce generally contains a small proportion of chloride of magnesium,
which makes it slightly deliquescent and bitter.
Exp. 224.-Heat a few crystals of coarse salt on a piece of sheet-iron
over the gas-lamp. The crystals will decrepitate forcibly, and the
greater part of the salt will be thrown off the plate; what remains will
melt as the temperature rises, and if the heat be strong enough, it will
finally volatilize. The decrepitation is due to little particles of water,
mechanically enclosed in the crystals, which, when expanded by heat,
burst the crystals asunder.
Pip. 225.-Dissolve 9 grms. of fine salt in 25 c. c. of water at about
20~. Add to the solution another gramme of salt; it will not dissolve.
Bring the solution to boiling; the added gramme of salt will barely
dissolve. Chloride of sodium is scarcely more soluble in hot than in
cold water, wherein it differs from the great majority of soluble salts.
Evaporated brines deposit their salt with almost equal facility when
hot and when cold; but the hot liquors will hold in solution a much
greater proportion of the salts with which the chloride of sodium is
associated, than the cold brines could retain. In the process of evaporation by fire, the associated magnesium, calcium, and sodium-salts
therefore do not crystallize with the common salt, but remain in the
hot mother-liquor.
Exp. 226.-Expose a saturated solution of salt in winter weather
to a temperature of — 10~. Large, transparent, six-sided tables, which
contain a considerable proportion of water chemically combined, will
crystallize from the solution. The warmth of the hand is sufficient to
destroy this crystalline compound; the water separates, and the crystals are resolved into a mass of minute cubes.




394                  SULPHATE OF SODIUM.
480. The uses of common salt are manifold. Since it is a constituent of almost all kinds of food, and essential to the life of
animals, it is not surprising that salt exists in small quantities in
almost every spring, soil, plant, and animal. The antiseptic quality of salt is applied to the preservation of fish, meat, and wood.
Salt is extensively employed in glazing earthenware, its volatility
at furnace-heat (Exp. 224) combining with other qualities to fit
it for this use. Immense quantities of salt are consumed in preparing sulphate of sodium, from which in turn common "soda"
(carbonate of sodium) is made. From the carbonate of sodium
thus obtained the greater number of other sodium compounds are
prepared. Salt is also the source from which chlorhydric acid
and chlorine are derived (~~ 101, 105).
481. Bromide and Iodide of Sodium (NaBr and NaI).-These
salts bear a close resemblance to the chloride of sodium; they
both crystallize in anhydrous cubes, and both occur native in seawater, though in minute proportion. Many marine plants appropriate iodide of sodium from sea-water; sea-weeds are therefore
the commercial source of iodine (~ 135).
482. Sulphate of Sodium (Na2SO4).-This compound is made
in great quantities from common salt and sulphuric acid as a preliminary step in the manufacture of carbonate of sodium.
The process has two stages. The first operation is performed in
large, covered, cast-iron pans, capable of holding 250 kilos. of salt, and
an equal weight of sulphuric acid of the density of 1-7. A very gentle
heat suffices to disengagoe from such a mixture enormous volumes of
chlorhydric acid gas; this gas, which would be injurious to vegetation
if suffered to escape into the air, is all absorbed by being passed through
vertical stone towers, filled with lumps of coke, over which water is
kept trickling. The reaction in the iron pan is by no means complete,
much chloride of sodium remaining undecomposed. The reaction at
this first stage may be represented as follows:2NaC1 - 112S04 = NaCl + NaHSO4 + HC1.
The pasty mass is then pushed into an adjoining fire-brick chamber,.which is strongly heated by flues from a furnace. The acid sulphate
of sodium of the last reaction decomposes the remainder of the salt,
and a further quantity of chlorhydric acid is disengaged to be condensed by the water in the coke-towers, while sulphate of sodium
remains:-    NaCl +NaHSO4 = Na2SO4 + HC1.




GLAUBER S SALT.                    395
The sulphuric acid used in this process is not stronger than can readily
be made by evaporation in leaden pans (compare ~ 230); it is always
made on the spot. The crude, weak, chlorhydric acid, which is the
incidental product of the manufacture of sulphate of sodium, has of
course some value in the arts; a portion of the product is often immediately consumed in the same factory in the manufacture of" bleachingpowder" (~~ 105, 120). In some works the smoke and gases from the
fires pass through the coke-towers in which the chlorhydric acid is
absorbed; in others, the products of combustion are not suffered to
mix with the liberated chlorhydric acid, but are conducted by separate
flues around the pans and chambers to be heated, and thence into the
main chimney. The latter mode of construction is the best.
The sulphate of sodium resulting from  this process is a white
anhydrous salt which dissolves easily in water at 30~.' When a
strong solution of the anhydrous salt, made at this temperature,
is cooled, there separate large colorless crystals of a transparent
salt, bitter and cooling to the taste. This salt, long known as
Glauber's salt, contains, besides the elements of sulphate of sodium, ten molecules of water; it therefore answers to the formula
INNaSO4,10H20.  There is another hydrated sulphate of sodium,
which contains only seven molecules of water. These hydrated
salts effloresce in dry air, and crumble into an opaque powder of
the anhydrous salt.
Exp. 227. —Place 10 grms. of crystallized Glauber's salt in a warm,
dry place. When it is completely converted into a white powder,
weigh the residue. Since Glauber's salt is more than half water, the
dry residue will not weigh more than 4-5 grins.
Exp. 228.-In a flask holding about 250 c. c., heat 50 c. c. of water
to a temperature of 33~, and keep the water at this temperature, as
determined by a thermometer immersed in it. Add to the warm water
161 grms. of crystallized Glauber's sal-t. If this saturated solution be
made hotter, the anhydrous sulphate of sodium crystallizes out in octahedrons of rhombic base; if, on the other hand, the solution be suffered
to cool, crystals of the common hydrated Glauber's salt appear. This
example forcibly illustrates the general fact that the relations of water
to other bodies are greatly affected by temperature.
Exp. 229. —Dissolve 10 grins. of crystallized Glauber's salt in water
of which the temperature has been previously observed; during solution the temperature falls; cold is produced, in consequence of the
expenditure of some of the heat of the mixture in overcoming the cohesion of the crystallized salt. Dissolve a like quantity of effloresced




396               SUPERSATURATED SOLUTIONS.
Glauber's salt (anhydrous sulphate of sodium) in a small bulk of water;
heat will he developed. A part of the water is solidified by combining
with the anhydrous sulphate to form the hydrated sulphate, and the
heat which before kept that quantity of water fluid, being set free to
do other work, raises by a certain amount the temperature of the
mixture.
Exp. 230.-Dissolve 50 grms. of Glauber's salt in 25 c. c. of water
in a small flask, by heating the contents of the flask until the liquid
boils. Cover the mouth of the flask loosely with a card or piece of
glass, and allow the liquid to cool, at perfect rest, to the ordinary temperature. No crystals will be deposited from the fluid, although the
water holds in solution a much larger quantity of salt than it could
dissolve at the atmospheric temperature. Such a solution is said to be
supersaturated. The crystallization of such a solution may generally
be brought about, almost instantaneously, by jarring the vessel which
contains it, or by permitting some foreign body, like a glass rod, a wire,
a crystal of the salt, or a grain of dust, to come in contact with the
fluid. By touching with a stick or glass rod the clear solution prepared as above described, this sudden crystallization will be strikingly
illustrated, the whole mass becomihg solid.
Crystallized sulphate of sodium is rapidly soluble in chlorhydric acid,
with great depression of temperature. A convenient refrigerating mixture may be prepared in climates where ice is dear, by pouring 5 parts
of the commercial acid upon 8 of the crystallized sulphate. The effects
of solution on temperature, and the phenomena of supersaturated solutions, though well exhibited by sulphate of sodium, are by no means
peculiar to this substance; they are manifested to a greater or less
degree by a large number of salts.
483. The Double Sul7phate of Sodium and Hydrogen (NaHS04)
is a very acid salt, to which the name of bisulphate is commonly
applied.  When heated it first gives up a molecule of water, and
subsequently, at a higher temperature, a molecule of anhydrous
sulphuric acid; the anhydrous salt is therefore employed as a
convenient source of the teroxide of sulphur.
2(NaHSO4) =-1O 0       S3 +  Na+SO4
The formation of this double sulphate marks an intermediate
stage in the making of sulphate of sodium and chlorhydric acid
from common salt; and it is the residue of the nitric-acid manufacture whenever nitrate of sodium is employed.
484. Carbonate of Sodium (Na2CO3).-The manufacture of this




CARBONATE OF SODIUM.                  397
substance constitutes one of the most important branches of chemical industry.  Immense quantities of it are consumed in the
fabrication of glass and soap, in the preparation of the various
compounds of sodium, and in washing, both by the manufacturer
of cloth and in the household. The ashes of sea and sea-shore
plants were formerly the source of the carbonate of sodium, but
it is now chiefly made from common salt by a process called, from
the name of its French inventor, the process of Leblanc.
The first staoe of this process we have already studied; it consists
in the preparation of sulphate of sodium from  common salt. The
second stage consists in the reduction of this sulphate to the condition
of sulphide of sodium in presence of carbonate of calcium; by interchange of metals there results sulphide of calcium and carbonate of
sodium. The sulphate of sodium is ground up with an equal weight of
chalk and rather more than half its weight of coal, and the mixture is
thoroughly melted by the flame of a reverberatory furnace. The black
mass, which is called "black ball " or " black ash," is cast into blocks
in iron wheelbarrows, cooled, broken up, and systematically washed
with warm water until all the soluble portions are extracted. The,black solution is evaporated in large iron pans by the waste heat of
the reverberatory furnaces. The residue contains some caustic soda,
mixed with the carbonate; the residue is therefore mixed with about
one-seventh of its weight of sawdust, or like material, and roasted in
a reverberatory furnace. The product of this heating is the soda-ash
oi commerce; it is almost white, and generally contains about 80 per
cent. of pure anhydrous carbonate of sodium.
The so-called crystals of soda are obtained by dissolving the crude
soda-ash in hot water, and suffering the hot solution to cool in large
paans. In the course of five or six days, large transparent crystals are
formed which contain 62'93 per cent. of water, and correspond to the
formula Ta2CO3,10H1120. The impure mother-liquor, drained from the
crystals, is used in the manufacture of caustic soda. These crystals
effloresce in the air; they have a disagreeable taste, called alkaline,
are soluble in very large proportion both in hot and cold water, and
even melt at a moderate temperature in their own water of crystallization. The crystals readily part with all their water, and the dry
residue melts at a bright-red heat; this residue is anhydrous carbonate of sodium, purified by the process of crystallization which it
has undergone. In this case, as in all others, the process of crystallization consists essentially in the aggregation of like particles; the
strong tendency is to exclude heterogeneous particles, or, in other




398                BICARBONATE OF SODIUM.
words, impurities, from the crystallizing structure. There is no more
universally applicable and valuable means of purification than the
process of crystallization.
The purchaser of carbonate of sodium for the sake of the alkali
which it contains will prefer soda-ash to soda-crystals, unless the
purity of the material be an important consideration. The crystals
are purer than the ash; but more than half their weight is water,
which must be transported at the cost of the consumer. Effloresced
crystals are more advantageous to buy by weight than crystals which
have not lost their water by exposure to the air. There are several
hydrates of carbonate of sodium, of different solubilities.
485. Double Carbonate of Sodium and Hydrogen (NaHCO.).When masses of crystals of hydrated carbonate of sodium (sodacrystals) are exposed to an atmosphere of carbonic acid gas, they
absorb carbonic acid with an evolution of heat sufficient to expel
the greater part of their water of crystallization. A white powder
remains, whose dualistic formula is Na20,I20,2C02, whence its
most familiar name-bicarbonate of soda.  This substance is one
of the ingredients in most of the artificial yeasts used for raising
bread, cake, and puddings, and is known to grocers and cooks
as "soda," although the constituent which is really utilized is its
carbonic acid. This carbonate of sodium and hydrogen is much
less soluble than the carbonate of sodium; it may be washed with
cold water until it is freed from the sulphates and chlorides which
generally contaminate the carbonate. The chemist resorts to this
process in order to prepare from the purified bicarbonate pure
sodium-salts. At a low red heat the double carbonate loses its
water and half its carbonic acid, and is converted into the normal
carbonate (Na2,CO).  If the aqueous solution of the double carbonate be heated, it loses one qua.ter of its carbonic acid.
Weigh out two separate and equal portions of bicarbonate of sodium.
Heat one of these portions for a few minutes in an iron spoon or porcelain capsule, to a low red heat. Then wrap each portion in a piece
of paper and introduce each little roll into a cylindrical jar filled with
mercury and standing inverted on the mercury trough. By means of
a curved pipette (see Appendix, ~ 22) pass equal quantities of dilute
sulphuric acid into each of the jars. The moment the sulphuric acid
comes in contact with the carbonates of sodium, carbonic acid is
evolved, and upon comparing the volumes of gas yielded by the two
samples, both of which contain the same quantity of sodium, it will be




YEAST POWDERS.                   399
found that the carbonic acid evolved from the unheated bicarbonate is
twice as great as that yielded by the portion which was ignited.
Bicarbonate of sodium may be deprived of its carbonic acid by
almost any acid or acid salt. In the experiment just described
sulphuric acid displaced the gaseous carbonic acid; tartaric acid,
or an acid tartrate like cream of tartar (tartrate of potassium),
will effect the same displacement, as will also the acid sulphate of
sodium (NaHSO4), the acid phosphate of calcium (CaO,2HI20,P,O,),
or common alum.  "Soda powders" should be made of bicarbonate of sodium and tartaric acid.  " Rochelle powders" consist of
bicarbonate of sodium in one paper and cream of tartar in another;
when these two materials are mixed in water, carbonic acid is
set free, and a double tartrate of sodium and potassium, called
Rochelle salt and used as a purgative, remains in the liquid.
When bread or cake is "raised" with "soda" and cream  of
tartar, the escaping carbonic acid is the agent in puffing up the
dough, and the same Rochelle salt remains in the bread. Tartaric acid anid cream of tartar having been dear in late years, a
cheaper chemical yeast powder has been made from acid phosphate of calcium; when this substance reacts within the dough
with bicarbonate of sodium, there remains in the bread a mixture
of the phosphates of sodium and calcium. Alum is sometimes
used for the same purpose. It is necessary to employ for such
purposes, in connexion with the bicarbonate, acids or acid salts
which are solid, and not so corrosive as to be obviously dangerous
and harmful.
There exists a native sesquicarbonate of sodium called t~rona or
natron; it is a saline efflorescence, always contaminated with the
sulphate and chloride of sodium, and less soluble in water than
the carbonate, but more soluble than the bicarbonate of sodium.
All the carbonates of sodium have an alkaline reaction on vegetable colors; carbonic acid is too weak an acid to overcome the
intensely alkaline reaction of caustic soda.
486. Suit)hides of Sodittzn.-In the conversion of sulphate of
sodium into carbonate of sodium (~ 484) one step is the reduction
of the sulphate to the sulphide of sodium by coal. The oxygen
of the sulphuric acid used in making the sulphate is thus combined with carbon and lost; the sulphur of the acid goes to the




400                  SULPHIDES OF SODIUM.
waste-heap in a useless combination with calcium. Herein lies
the wastefulness of the Leblanc process, which the experience of
fifty years has failed to remedy.
Exp. 231. —Mix a little powdered anhydrous sulphate of sodium,
obtained by drying Glauber's salt, with as much powdered charcoal,
and make the mixture into a paste with a drop of water. Place a little
ball of this paste, as large as a small pea, in a depression in a piece of
charcoal, and heat it strongly with a blowpipe. The mixture effervesces and finally melts into a brownish mass. The glowing coal talkes
all the oxygen out of the sulphate, and the carbonic oxide, which
results from the union of the carbon and oxygen, in escaping causes
the effervescence which occurs; the sodium and sulphur alone remain
united. The reaction may be thus formulated:
Na2SO4 + 4C = 4CO + NaS.
If the brownish solid be removed from the coal, p':ced in a watch-glass,
and moistened with dilute chlorhydric or sulphuric acid, sulphydric
acid gas will be evolved, as may be recognized by the smell, or by the
use of lead-paper:Na2S + 2HC1 = 2NaCl + H2S.
Sulphur unites with sodium in more than one proportion; and
it may indeed be doubted whether the actual reaction in the
above experiment is so simple as the formula represents it to be.
There probably exist five distinct sulphides of sodium-Na2S,
Na2S2, Na2Ss, Na2S4, and Na2S5.  All these sulphides have an
alkaline reaction to test-paper, and evolve a more or less distinct
odor of sulphuretted hydrogen.  When they are brought into
contact with an acid they are decomposed, sulphydric acid escapes,
and a white precipitate of finely divided sulphur falls, in every
case except that of the first sulphide, Na.S.  Besides these sulphides, a compound of sodium, hydrogen, and sulphur (NailS)
is known, which is perfectly analogous in composition to the combination of sodium, hydrogen, and oxygen with which we are
already familiar under the name of caustic soda (NaHO).
487. Sodium (Na).-The element sodium is never found uncombined in nature, for the reason that in its elementary condition it cannot exist in contact with either air or water.  It is,
however, artificially prepared from the carbonate of sodium without serious difficulty, and it might be produced in considerable
quantities, if there were any large use for the element.




METALLIC SODIUM.                    401
A mixture of 20 parts of carbonate of sodium, 9 parts of coal, and
3 parts of chalk, placed in an iron bottle or cylinder, is heated to a
very high temperature in a suitable furnace; a narrow iron tube connects the bottle, or cylinder, in the furnace, with a flat sheet-iron box
outside the furnace; th's box receives the sodium which distils from
the hot mixture, and a hole in the front lower corner of the receptacle
permits the melted metal to fall into a vessel containing petroleunmnaphtha, beneath which the sodium can be preserved. The reaction is
theoretically a conversion of all the oxygen in the carbonate of sodium
into carbonic oxide, partly by a new combination with the carbon in
the carbonate, and partly by union with the carbon which the coal
supplies,
NaCO3 + 2C = 3CO + 2Na;
but the amount of sodium practically obtained from a given weight of
the materials is by no means as large as this simple formula would indicate. The chalk has no chemical effect, but is practically essential to'he success of the operation. It prevents the mass from melting, and
the gas which it gives off when heated assists in sweeping the sodium
out of the bottle into the receiver. To purify the crude sodium thus
obtained, it is melted under naphtha and cast into ingots in iron
moulds. It must be kept under naphtha in tightly closed bottles.
The properties of the element, sodium, are very curious. The
substance, when freshly cut, or when melted under naphtha or in
an atmosphere artificially deprived.of oxygen, has the brilliant,
white metallic lustre of silver. Though possessing so eminently
this characteristic property of the class of bodies called metals,
and being like them  a good conductor of heat and electricity,
sodium is far from resembling the ordinary metals in other respects. Thus it is lighter than water, having a specific gravity of
only 0'972, whereas the common metals are dense and heavy;
again, it is as soft as wax at common temperatures, and melts at
a temperature below that of boiling water, while it has none of
the comparative permanence which characterizes lead, tin, copper,
silver, gold, and other familiar metals.  If exposed to the air,
even for a few seconds only, it tarnishes, and soon becomes cove. red
with a coating of oxide. Instead of being quenched by water, it
takes fire when thrown into warm water. We have already seen
that it decomposes cold water (Exp. 14), setting free its hydrogen,
and combining with its oxygen.
Etr. 232.-Cover the bottom of a large bottle (at least a litre
2D




402                  HYDRATE OF SODIUM.
bottle) with hot water, drop in a piece of sodium as large as a small
pea, and immediately cover the mouth of the bottle with a card or
glass plate. The heat of the chemical combination between the sodium and the oxygen of the water is sufficient to inflame the hydrogen; the escaping hydrogen carries with it a small portion of the
volatilized sodium, and therefore burns with an intensely yellow flame
which is very characteristic of sodium-compounds. The metal swims
rapidly about on the surface of the water, and is completely converted
into caustic soda; at a little interval, after the flame has ceased to
burn, a globule of caustic soda, which has escaped solution, bursts, and
scatters in all directions; the mouth of the bottle should always be
covered to avoid the possible projection of particles of hot soda out of
the bottle. The water in the bottle, tested with litmus paper, will be
found to possess a strong alkaline reaction. If the bit of sodium be
previously wrapped up in muslin, it will take fire in cold water or even
on ice. The muslin prevents the sodium from moving about, and the
heat of combination is therefore concentrated upon one spot. The
same effect may be produced without the muslin on cold water made
viscid with gum.
At a high temperature sodium will remove oxygen from almost
all bodies which contain it, whether solid, liquid, or gaseous.
Hence the necessity of preserving the metal under some liquid
which, like naphtha, contains no oxygen. Sodium enters directly
into combination with all the elements of the chlorine and sulphur
groups, and is capable of withdrawing these elements from nearly
all the compounds into which they enter. In this substance,
therefore, the chemist possesses a very potent agent for effecting
chemical transformations.  Its atomic weight is 23.
488. Hydrate of Sodium (NaHO).-When sodium is burnt upon
water, a solution of hydrate of sodium, possessing an intensely
alkaline reactign, remains behind; but in practice the hydrate is
generally made from the carbonate.
Exp. 233.-Dissolve 100 grms. of crystallized carbonate of sodium
(~ 484) in 400 c. c. of water. Slake 20 grms. of quicklime with water
enough to.make the slaked lime into a cream. Boil the solution of
carbonate of sodium in,_a wron pan, and add to it, little by little, the
cream oflime, until a smalltp9rtion of the liquor, filtered off, produces
no effervescence when poured into dilute chlorhydric or sulphuric acid.
Thecaldium of the lime replans the sodium in the carbonate of so



CAUSTIC SODA.                    403
dium; a white insoluble precipitate of carbonate of calcium is formed,
and hydrate of sodium r1emains in the solution:Na2CO, + CaH202 = 2NaHO + CaCO,.
When this reaction is complete, no carbonic acid remains in the clear
solution, which, therefore, causes no effervescence when mixed with an
acid. When this test shows the reaction to be accomplished, extinguish
the lamp; cover the pan, and let the carbonate of calcium settle to the
bottom of the vessel. After several hours, draw off the clear supernatant liquid into a bottle by means of a siphon. A weaker lye may
be obtained by boiling up the residue in the iron pan once more with
water, and again decanting the clear liquid from the deposited precipitate.
The lye thus prepared contains a very little lime, inasmuch as the
hydrate of calcium is not completely insoluble in a dilute solution of
hydrate of sodium. If the lye is needed stronger, it is only necessary
to concentrate it by evaporation in a clean iron pan to the requisite
strength. Since soda-lye attacks ground glass with facility, the stoppers of bottles in which the lye is kept are apt to get stuck so fast that
they cannot be removed. Such a bottle may be best closed with a
caoutchouc-stopper, or a cork soaked in melted wax or paraffine.
In order to obtain the solid hydrate of sodium, its solution must be
evaporated in a silver dish (since iron would color the concentrated.
lye) until the liquid in the dish flows smoothly like oil at a temperature
near to a red heat. This thick, oily liquid solidifies when turned out
upon a cold plate of metal or poured into metal moulds.
Fused caustic soda is a white, somewhat translucent mass, whose
composition corresponds to the formula NaHO. Heat will separate from it no more water;'if the attempt be made, it volatilizes
in caustic vapors without change of constitution. The commercial
caustic soda always contains much more water than the formula
indicates. The solid hydrate is very soluble in water, and greedily
absorbs both water and carbonic acid from the air, until the formation of a coating of the non-deliquescent carbonate of sodium
arrests the process by protecting the enclosed hydrate. It is the
prototype of the class of bodies called bases. It colors litmus blue
and turmeric brown, and when mixed in due proportion with
oxides of the opposite quality, called acid, a saline compound is
formed which is neither acid nor alkaline, and which may bear no
2D2




1404       REPLACEMENT-DIRECT COMBINATION.
more resemblance to its proximate constituents than bread bears
to flour and water, or rust to iron and oxygen.
From such reactions between acids and hydrate of sodium, water
is always disengaged simultaneously with the saline product, and
the reaction may almost always be as well considered an interchange of place between hydrogen and some other element as an
act of combination between one oxide and another oxide, both of
which, or.one of which, contain also hydrogen. The following
formulae will illustrate the meaning of this statement:NaHO + NHO3 = NaNO, + H,20, or
NaH     + N022O =               1} + o
2NaHO  + H2S04 =  Na,2S4 + 2H20, or
2Na} 0 +          2}~ = 0a2}; +'}2
0222
NaHO  + C2H402 =  C2H3NaO2 + H20.
Acetic Acid. Acetate of Sodium.
While recognizing the frequent occurrence of such reactions as
are above represented between hydrated oxides, it must not be
forgotten that many anhydrous saline compounds can be made by
the direct combination, under appropriate conditions, of two oxides
which contain no hydrogen. By heating one molecule of hydrate
of sodium, or 40 parts by weight, with one molecule, or 23 parts
by weight, of sodium, an oxide of sodium is obtained which contains no hydrogen; but this body has none of the properties described by the adjective alkaline, any more than the anhydrous
teroxide of sulphur possesses the properties suggested to the mind
by the term "acid ":NalRO  + Na = Na20 + H.
Now the. very same sulphate of sodium which results from the
second of the above reactions, may be prepared by bringing together this anhydrous oxide of sodium and anhydrous sulphuric
acid:Na20 + S03 = Na2S04,
There exists another anhydrous oxide of sodium, corresponding
in composition to the formula NaaO0, and the same sulphate of




PHOSPHATES OP SODIUM.               405
sodium can be made by heating this oxide with sulphurous acid
gas:Na202 +  S02 =  Na2804.
These facts show that a knowledge of the constituents from
which a salt may be made is not sufficient to establish any presumption concerning the molecular constitution of the salt itself.
The preparation of solid caustic soda, for household and other
uses, has lately become a considerable industry. Soap is made by
boiling together grease or oil with caustic soda or potash; sodalye yields a hard soap, potash-lye a soft soap. If the maker of
soap starts from the carbonate of sodium or potassium, he must
first make a solution of caustic lye by the method of Exp. 233,
and then boil the lye, so obtained, with the grease which is the
other essential ingredient of soap. The soap-maker is saved the
trouble of converting the carbonate into the hydrate of sodium or
potassium by the maker of the solid caustic alkalies, which need
only to be dissolved in water to yield the requisite lye. The solid
alkali is commonly put up for transportation in sheet-iron canisters, of all sizes. The manufacturer of caustic soda directly from
the sulphate of soda has an advantage, in that he can avail himself of the caustic soda which the " black ash " always contains.
He is not obliged, first to convert this caustic alkali into carbonate,
and then to remove all the carbonic acid by lime; he can therefore dispense with the heating with sawdust which is necessary
in the manufacture of soda-ash.
489. Phosphates of Sodium.-There are three sets of phosphates
of sodium: —1. Common phosphates, which contain three atoms of
hydrogen or an equivalent metal; 2. Pyrophosphates, which occupy an intermediate place between common phosphates and metaphosphates; 3. Metaphosphates, which contain only one equivalent of hydrogen or an equivalent metal. (See ~ 293.)
3H,O,P20,  =  2HPO;  3NaO,P05 -=  2NaPO,
2H20,P2,O  =  H4P207;  2NaO,P20,  =  Na4P,07
H2O,P20 -=  2HPO,;  Na2O,P20,    -  2NaPO3.
The most familiar of the ordinary phosphates is the rhombic
phosphate of sodium, of the formula Na2HPO,V12H20, which is
the salt commonly called phosphate of sodium.




406               PYROPHOSPHATE OF SODIUM.
Exp. 234.-Digest 8 grins. of powdered, white, burnt bones with
32 c. c. of water and 6 grms. of sulphuric acid, until a uniform paste is
produced; strain the mass through a piece of muslin, stir up the residue with water, and squeeze the liquor through the cloth filter. Evaporate the filtered solution considerably, again filter off the separated
sulphate of calcium, dilute the filtrate with water until it measures 48
c. c., and gradually neutralize the acid solution with solid carbonate of
sodium. A slight excess of carbonate may be added, and the solution
ei aporated until it crystallizes; the crystals may be almost freed from
adhering sulphate of sodium by washing them. By recrystallizing the
salt it may be obtained in large, transparent rhombic prisms, which
are decidedly efflorescent. Thev have a cooling, saline taste, are soluble
in four parts of cold water, and readily melt in their water of crystallization; their solution has a faint alkaline reaction.
Exp. 235.-Add to a solution of the purified crystals of the last experiment a few drops of a solution of nitrate of silver; a yellow precipitate appears, and the liquid becomes distinctly acid in its reaction
with litmus. The precipitate is phosphate of silver, and the acidity is
due to the simultaneous liberation of nitric acid in the solution:
Na2HP04 +  3AgN03 = 2NaNO,    + HNO.3  +  Ag3P04.
From the rhombic phosphate two other terbasic.phosphates may be
prepared: —by adding caustic soda, a terbasic, or termetallic, phosphate,
Na3PO4,12H20; by adding phosphoric acid, a so-called " acid " phosphate, NaH2PO4,H20..Exp. 236. —Heat 4 or 5 grins. of rhombic phosphate of sodium in a
porcelain crucible to bright redness. The water of crystallization first
escapes, then another portion of water is driven off at a higher heat,
the residue melts, and on cooling solidifies again to an opaque, white,
substance-the pyrophosphate of sodium, Na4P2O7,:
2(Na2HP04,12H20) = 24H20 +  1120 +  Na4P207.
Extp. 237. —Dissolve the new salt of the last experiment in water
and add a few drops of a solution of nitrate of silver; a chalky-white
precipitate of pyrophosphate of silver will be produced, while the
supernatant liquid is neutral:Na4P207 + 4AgoNO, = 4NaNO, + Ag4P207.
The student will observe, in passing, that the pyrophosphate of
sodium dissolves in water with more difficulty than the ordinary
phosphate. From its aqueous solution the pyrophosphate crystallizes ill prisms which contain ten molecules of water of crjstallizuation, Na4P207, 10H20.  These crystals are not effilorescent, but




METAPHOSPHATE OF SODIUM.                407
permanent in the air. Two of the atoms of sodium in neutral
pyrophosphate of sodium can be replaced by hydrogen, and we
thus obtain a salt of the formula Na2H,2P207 which, like the neutral
pyrophosphate. throws down the white pyrophosphate of silver
from a solution of nitrate of silver, but, unlike the neutral salt,
leaves behind a liquor acidified with free nitric acid.
Exp. 238. —Mix a hot solution of 18 grms. of common phosphate of
sodium with a concentrated solution of 3 grms. of chloride of ammonium; common salt remains in solution, and transparent crystals of a
complex phosphate of sodium, ammonium, and hydrogen separate from
the liquid as it cools:NaHPO4 + NH4C1 = NaCl + NaNH4HPO4.
This complex triphosphate is known as microcosmic salt; it contains
only one atom of sodium.
Exp. 239.-Heat the microcosmic salt obtained in the last experiment to redness. Ammonia and water are driven off with a great
bubbling, and a glassy substance is left which corresponds to the formula NaPO3, and is called the metalphosphate of sodium. This salt
cannot be obtained by this process in crystals; it is deliquescent and
very soluble.
Exp. 240. —Dissolve in water some of the metaphosphate just made,
and add a few drops of the solution to a solution of nitrate of silver.
The precipitate produced is white, but gelatinous; its appearance
is quite different from that of the pyrophosphate of silver. The metaphosphate of silver is soluble in an excess of metaphosphate of sodium,
so that a sufficient quantity of the sodium-salt will redissolve the
silver-salt precipitate which the first addition of the sodium-salt produced. It may be observed, that the solution of metaphosphate of
sodium possesses a feebly acid reaction.
The metaphosphate of sodium exists in several distinct modifications; or, in other words, there are several sodiumn-salts which
correspond to the formula NaPO3.  The most noteworthy modifications are the glassy salt just prepared, a crystallizable modification, and a variety which is almost insoluble in water though
soluble in acids.  Pyrophosphates and metaphosphates are alike
converted into common ter-metallic phosphates by long boiling
with water in an acid liquid, or by fusion with an excess of caustic alkali. The metaphosphate may also be converted into the
pyrophosphate by mere heat in presence of water. The meta



408                        BORAX.
phosphate, when dried at 1000, retains a certain portion of water;
but this water does not enter into the constitution of the salt,
for, on again dissolving the salt, the solution gives the usual reactions of the metaphosphate. If, however, the water-containing salt be heated to 150~, the water enters into the constitution
of the salt, which becomes the " acid " pyrophosphate of sodium,
yielding, on solution, the reaction of the pyrophosphates:2 NaPO, +  1120 =  Na2H2P207.
We have here a striking illustration of the fact that two substances, which contain precisely the same elements in the same
proportions, may have quite different properties in consequence of
the different arrangement of their elements, or, in other words,
of their different molecular constitution.
490. Borax (Na2B,O7, 10I2,=Na20,2B20,+ 10H20). - The
dualistic formula of this useful salt accounts for its chemical
name —biborate of sodium. The greater part of the borax used
in the arts is prepared by dissolving the native boracic acid of
Tuscany (~ 443) in a hot solution of carbonate of sodium. After
all the carbonic acid has escaped, the hot liquid is allowed to
settle and cool. Either of two different crystalline forms of
borax may be obtained, at will, by suitably regulating the density
of the solution and the temperature at which crystallization takes
place. One of these forms, prismatic or common borax, contains
10 molecules of water of crystallization; the other, called octahedral] borax, only 5. Both varieties must be purified by recrystallization. The large crystals, in which borax is demanded for
the market, can only be obtained by operating on very large
masses of the salt,-a remark which applies to most of the crystalline salts which are used in the arts. The crystals of octahedral borax are harder and less fragile than those of ordinary
borax. They are unalterable in dry air, but in a moist atmosphere they absorb water and are converted iuto common borax.
When heated, they fuse to an anhydrous glass with less intumescence than common borax, and without cracking.  For these
reasons, octahedral borax is better than prismatic borax for
many purposes, and its smaller proportion of water (30 per cent.
against 47 per cent.) diminishes the cost of transport. Never



BORAX AS A TEST.                   409
theless the prismatic borax is generally preferred, probably because it is sold at a less price, weight for weight.  The prismatic
crystals are soluble in about half their weight of boiling water,
but in twelve times their weight of cold water; as they are
slightly efflorescent, they are generally covered with a white
dust. Borax has a feebly alkaline taste and reaction. When
heated it bubbles up, loses its water, and melts below redness
into a transparent glass; this glass dissolves many oxides of the
metals, acquiring thereby various colors characteristic of these
oxides. Hence borax is much used as a blowpipe test for determining the presence of certain oxides of the metals.
Etp. 241. —Make a little loop, about as large as this 0, on the end
of a bit of fine platinum wire 6 or 8 c.m. long. Make the loop whitehot in the blowpipe-flame, and thrust it while hot into some powdered
borax; a quantity of borax will adhere to the hot wire; reheat the
loop in the oxidizing flame; the borax will puff up at first, and then
fuse to a transparent glass. If enough borax to form a solid transparent bead within the loop does not adhere to the hot wire the first
time, the hot loop may be dipped a second time into the powdered
borax. When a transparent glass has been formed within the loop of
the platinum wire, touch the bead of glass, while it is hot and soft, to
a speck of black oxide of manganese no bigger than the period of this
type; reheat the bead with the adhering particle of oxide in the oxidizing flame; the black speck will gradually dissolve, and on looking
through the bead towards the light, or a white wall, when the black
oxide has disappeared, the glass will be seen to have assumed a
purplish-red color.
The same experiment may be performed with oxide of iron, which
imparts to the glass a yellow color, or with oxide of copper, which
imparts a bluish-green color. The oxidizing flame must be used in
both these cases, as with the oxide of manganese.
The power which borax possesses of dissolving metallic oxides
suggests an explanation of its use in brazing, and in soldering
the precious metals. The solder will only adhere to a bright and
clean metallic surface, and the borax which melts with the solder
removes from the pieces of metal the film of oxide which would
otherwise prevent the adhesion of the solder. Borax is also used
by the assayer and refiner as a flux. In making enamels and
glazes, it is frequently added for the purpose of rendering the




410                  SILICATES OF SODIUM.
compound more fusible, and it is largely employed in fixing colors
on porcelain.
There is a normal, or neutral, borate of sodium, NaBO2 or
Na2O,B203, which crystallizes with various quantities of water;
and other borates of sodium are known; but the " biborate" is
the only one of any practical importance.
491. Silicates of Sodium may be prepared by dissolving silicic
acid in caustic soda, as in Exp. 220, or by fusing together silicic
acid and carbonate of sodium, or a mixture of silicic acid, sulphate of sodium, and carbon.  From alkaline solutions, the single
crystallizable silicate of sodium, Na SiO,, can readily be obtained
in hydrated crystals. The silicate of sodium of commerce, called
waterglass or soluble glass, is, however, a much more siliceous
silicate, of varying composition. Some samples of it have very
nearly the composition expressed by the formula Na2O,2SiO2,
while other specimens approximate to the formula Na2O, 4SiO2.
Between these extremes, the acid and the alkali unite in all
possible proportions to form compounds, all of which are soluble
with more or less difficulty in boiling water.
The normal silicate of sodium (Na2SiO,) is readily soluble in
cold water, and, like carbonate of sodium, which it closely resembles, may even be melted in its own water of crystallization;
but the acid salt of commerce is as good as insoluble in cold
water. Waterglass may, however, be completely dissolved by
long-continued boiling in water, and the solution thus obtained is
largely employed by calico-printers and by soapmakers. Waterglass is also used for hardening porous stones, or even for binding sand into artificial stone, for painting rough woodwork to protect it from the weather, and for diminishing the combustibility
of wood, canvas, and other coarse stuffs, such as are used for the
decorations of theatres. It has been suggested that the interior
surfaces of wooden-roofed railway bridges might be protected
from the sparks of the locomotive by washing them with a solution of waterglass. When employed as paint, the coating of
waterglass may be washed over with a solution of chloride of
ammonium, which decomposes the silicate, with deposition of free
silicic acid; or the silicate may simply be left to the decomposing
action of atmospheric carbonic acid.  The chloride of sodium




WATERGLASS.                     411
resulting from the decomposition in the one case, and the carbonate of sodium in the other, are subsequently washed away by
the rain. Combustible matters, when covered with a coating of
silica, or of silicate of sodium, as above, are prevented from burning freely, in the same manner that the carbon of the paper in Exp.
110 was kept from burning by the coating of phosphoric acid.
492. The chief use of silicate of sodium, however, is as a component of common glass. The various glasses of commerce are
mixtures of a highly siliceous silicate of sodium, or of potassium,
or of both these substances, with silicates of other metals, such as
calcium, aluminum, and lead. In green bottle-glass the easily
fusible silicate of iron replaces in part the silicate of sodium or
of potassium.
It is a peculiarity of the alkaline silicates that, in changing
from the liquid to the solid state, they pass through an intermediate pasty or viscous stage, and finally solidify in transparent
amorphous masses, totally devoid of crystalline structure. While
in this pasty, ductile state, these silicates may readily be moulded
into almost any form; and transparent vessels might doubtless be
made from the acid alkaline silicates without admixture of other
materials. The alkaline silicates, however, are, by themselves,
far too easily acted upon by air and moisture to admit of being
used as substitutes for ordinary glass. But it has been found
that, by combining the alkaline silicates with the silicates of
certain other metals, such as calcium, there may be obtained
compound glasses which, while they retain the plasticity of the
alkaline silicates as well as their amorphous character and transparency, are capable of resisting the action not only of air and
water, but even of acids and alkalies, to a very great extent.
Though the ordinary glasses are so difficultly attacked by water
that they may, for most practical purposes, be regarded as altogether insoluble, it is nevertheless true, as has been stated in
~ 464, that glass may be partially dissolved by long-continued
contact with water, particularly if the water be hot and the glass
in the condition of powder. After the smooth surface of glass, as
it comes from the fire, has once been removed, the corrosion of
the glass goes on more rapidly.
It is remarkable that mixtures composed of several different




412                        GLASS.
silicates melt at temperatures considerably lower than the mean
of the melting-points of their several ingredients.
The different varieties of glass vary in composition, according
to the purposes for which they are prepared; they cannot be regarded as chemical compounds, being really indefinite mixtures of
various acid silicates. The composition of window-glass may be represented approximately by the formula NaO, 2SiO2; CaO, 2SiO2;
and that of the hard, Bohemian glass. suitable for ignitiontubes, by the formula 2(K20, 3SiO2); 3(CaO, 3SiO2). The lustrous "flint" glass, employed for the nicer kinds of household
ware, may, on the other hand, be represented by the formula
2(K20, 2SiO,); 3(PbO, 2Si02); it is prepared from the purest
materials attainable. Bottle-glass is composed of the silicates of
calcium and aluminum,together with a small proportion of the
silicates of iron, of potassium, and of sodium; in this glass, as in
the other varieties above formulated, small quantities of various
other silicates, such as the silicates of magnesium and of mainganese, almost always occur.
In the preparation of the cheaper kinds of glass the materials
are melted in large open crucibles of refractory clay; but the better sorts of glass, such as flint-glass, are made in pots so covered
that no smoke or dust from the fire can come in contact with
their contents. In both cases the thoroughly melted mixture is
kept in the liquid state until it has become perfectly homogeneous
and until all bubbles of air have escaped from it; it is then
allowed to cool to the temperature at which it possesses the
peculiar, pasty, ductile, condition in which it admits of being
blown, pressed, and moulded.
493. After glass has been moulded into the shape desired, it
must still be subjected to a process of annealing before it can be
used. Glass which has been suddenly cooled after fusion is extremely brittle; and in general all glass which has been quickly
cooled after heating is far more fragile than that which has been
allowed to cool gradually. The operation of annealing is nothing
but a process of slow baking, at a temperature which, though not
so high as the melting-point of glass, is nevertheless high enough
to allow the particles of softened glass to move among themselves, and to come into easy and natural positions as regards




IYPOSULPHITE OF SODIUM.               413
one another; after the baking follows a process of slow cooling,
during which the heated material contracts uniformly in all
directions as it assumes the dimensions proper to it when cold.
Unlike the silicates of the alkali-metals, most metallic silicates
have a tendency to assume crystalline form on cooling, and it is
not difficult to bring about crystallization of silicate of calcium,
or silicate of aluminum, from ordinary glass, particularly from
the coarser kinds, such as bottle-glass. Glass, in which some of
the constituents have thus crystallized, has the appearance of
porcelain  it is said to be devitrified.  The devitrification may
readily be shown by imbedding bottle-glass in sand, heating the
glass almost to the melting-point, and then allowing it to cool
slowly. In annealing some kinds of glass care must be taken
not to heat the ware too strongly, lest it be devitrified during the
process.
494. AMelted glass, like melted borax (Exp. 241), is capable of
dissolving small quantities of many of the metallic oxides, a transparent and often colored silicate of the metal being formed, which
imparts its hue to the entire mass of glass. In this way, glass
may be obtained of almost any desired color. The green color
of bottle-glass is due to silicate of protoxide of iron; but richer
shades of green may be obtained by using protoxide of copper or
oxide of chromium. Dioxide of copper gives a ruby-red color,
and oxide of gold various shades of red, inclining to purple.
The oxides of uranium, of antimony, and of silver yield yellow
glasses; oxide of cobalt affords a beautiful blue, and the binoxide
of manganese a violet glass; while mixtures of the oxides of cobalt
and of manganese impart to the glass a black color.
495. Hyposulphite of Sodium  (NaS20,O,, 520).-This easily
crystallized and tolerably permanent salt is of great use to the
photographer, because its aqueous solution is capable of rendering soluble the chloride, bromide, and iodide of silver, compounds
much employed by the photographer, and very insoluble in water.
The photographic paper or glass, uniformly coated with some
silver compound, is exposed to light in the camera or press, and
then immersed in a strong solution of the hyposulphite, which
forms with the silver compound, in those parts of the picture
which have not been acted upon by the light, a double salt which




414                      POTASSIUM.
is soluble in water. This double salt and the superfluous hyposulphite must be washed away by soaking the picture several
hours in water which is constantly renewed. Hyposulphite of
sodium is also used as an " antichlore," or agent for removing
the last traces of chlorine, or hypochlorous acid, from substances
which have been bleached therewith. The salt may be best prepared by digesting sulphur with a solution of sulphite of sodium.
Exp. 242.-Dissolve 20 grms. of crystallized carbonate of sodium in
30-40 c. c. of water. Place the solution in a small Woulfe-bottle, and
pass sulphurous acid gas (Exp. 96) through it until all the carbonic
acid is expelled from the carbonate and effervescence ceases. The
liquid then holds in solution sulphite of sodium (Na2SO3). Pour this
solution into a bottle which can be tightly closed, and add to it 3 or 4
grms. of finely powdered sulphur; let this mixture stand corked up for
several days in a warm place; the sulphur will gradually dissolve, and
form a colorless solution, which on evaporation will yield crystals of
hyposulphite of sodium. Time may be saved by keeping the solution
of sulphite of sodium hot, but not boiling, during the digestion of the
sulphur. The reaction has been already formulated (~ 243).
CHAPTER XXIV
P O T AS I U M.
496. The proximate sources of sodium-compounds are the sea,
and salt springs and deposits. Potassium-compounds, on the
other hand, are derived indirectly from the soil. Arable soils are
produced by the weathering and gradual decomposition of the
common granitic rocks. Into the composition of these rocks there
enter two minerals, called feldspar and mica, which are mixed
silicates of potassium or sodium and aluminum or iron. The
element potassium thus becomes a normal constituent of the
earthy food of plants. The soil itself is not directly available as
a source of potassium-salts, because no cheap and easy method
has yet been devised for separating the potassium-compounds




CARBONATE OF POTASSIUM.                415
from the other ingredients of the soil. Plants, however, are able
to pick out and assimilate the potassium-salts from these rocks
and soils; so that by burning the plants and extracting the ashes
with water a soluble potassium-salt is obtained. Plants thus concentrate the potassium from out great masses of earth, and make
it accessible to us. The salt which is obtained from the ashes of
plants by washing and evaporation, is called potash, or, if refined,
pearlash, and it is from this substance that the bulk of potassiumcompounds are obtained.
Exp. 243.-Place a handful of wood-ashes on a filter, and pour hot
water over them, collecting the filtrate in a bottle and returning it upon
the ashes two or three times, in order to obtain a stronger solution.
To exhaust the ashes of their potash, they must, of course, be treated
with successive portions of hot water. This solution has a strong alkaline reaction upon test-paper. A few drops of it, poured into a testtube containing a little dilute acid, occasion a brisk effervescence, a
reaction from which we readily surmise the truth, that the potassiumsalt contained in the solution is the carbonate of potassium. Proof that
the gas evolved is carbonic acid can readily be obtained by conducting
the gas into lime-water, as in Exp. 170. By evaporating the rest of
the solution to dryness in a porcelain dish, we obtain a small sample of
crude potash. By treating this potash with a quantity of cold water,
insufficient to dissolve any but the most soluble portions of the mass,
letting the mixture stand some time, and evaporating the partial solution to dryness, a whiter, purer carbonate is obtained, the pearlash.
497. Carbonate of Potassium  (K2CO,) is a hygroscopic and
very soluble salt.  When exposed to damp air it becomes moist,
and finally deliquesces. In this respect it does not resemble
soda-ash, which is not hygroscopic, and is, for this reason among
others, better adapted than potash for transportation, storing, and
most commercial uses. Carbonate of potassium fuses at a red
heat, but cannot be decomposed by heat alone. At a red heat it
is decomposed by silica, as is also the carbonate of sodium, carbonic acid being expelled with effervescence, whilst the silica
unites with the alkali. Advantage is taken of this property in
the analysis of minerals which contain a large quantity of silica,
and are not easily decomposed by acids. The finely powdered
mineral is fused with about three times its weight of carbonate of
sodium or of potassium, or, better, with thrice its weight of a




416                 HYDRATE OF POTASSIUM.
mixture of 5- parts of carts of carbonate of sodium with 7 parts of carbonate of potassium. The mixed carbonates produce a more
fusible mixture than either alone (~ 492). The fused mass is
then treated with dilute chlorhydric acid, which decomposes the
alkaline silicates, and dissolves all the bases of the mineral which
were before combined with the silica.
Carbonate of potassium was the most important source of alkali,
until Leblanc's process made soda cheaper than potash. It is still
largely consumed in the manufacture of soap, glass, caustic potash, and other compounds of potassium; but sodium-salts have,
to a great extent, displaced potassium-salts in commerce and
the arts.
498. Carbonate of Potassium and Hydrogen (KHCO,).-This
salt, commonly called the "bicarbonate" of potassium  (K20,
H20,2C02), is prepared by passing a current of carbonic acid
through a strong solution of carbonate of potassium; crystals of
the bicarbonate will be deposited; which are permanent in the
air, and require about 4 parts of cold water for solution.  When
the solution of this salt is long exposed to the air, or boiled, it
loses one-fourth of its carbonic acid; when the dry salt is fused,
it loses half its carbonic acid, and is converted into the carbonate.
It is a valuable salt to the chemist and the apothecary, because it
can be readily obtained in a state of purity; when itself made
from a refined carbonate of potassium, it may be advantageously
used as the material from which to prepare other pure compounds
of this important element.
499. Hydrate of Potassium (KHO).-The manufacture of hydrate of potassium, from carbonate of potassium, resembles, in
every detail, the preparation of caustic soda from carbonate of
sodium (Exp. 233). The carbonate of potassium is dissolved in
10 or 12 times its weight of water, and decomposed by a cream of
lime: carbonate of calcium is precipitated, and hydrate of potassium remains in solution. All that has been said of the concentration of the solution of hydrate of sodium (~ 488) is true, also,
of hydrate of potassium.
Hydrate of potassium is a hard, whitish substance, possessing
a peculiar odor and a very acrid taste. Like the hydrate of sodium, it rapidly absorbs moisture and carbonic acid from the air;




CAUSTIC POTASH.                    417
and since the carbonate of potassium  thus formed is a deliquescent salt, this change will go on until the entire mass of hydrate
is converted into a syrup of the carbonate; whereas, in the case
of hydrate of sodium, the absorption of water and carbonic acid
is soon arrested by the formation of a coating of non-deliquescent
carbonate of sodium upon the surface of the lump of hydrate. In
chemical industries and speculations, the question of success or
failure often turns on such points as this; the advantage of a new
material, for example, often depends upon just such differences as
this between caustic soda and caustic potash.
500. The hydrate of potassium, cast into small sticks, is employed by physicians as a cautery,-a use which illustrates forcibly
its destructive effect upon animal and vegetable matters. Like
hydrate of sodium, its solution destroys ordinary paper, and
cannot be filtered except through asbestos, or gun-cotton. A
clear solution is best obtained by decantation from the subsided
impurities. All vessels made of materials which contain silica
are attacked by this caustic solution; and even platinum is slowly
oxidized in its presence; vessels of gold and silver resist it best.
This hydrate, like that of sodium, forms soaps with oils or fats;
the sodium-soaps are hard, the potassium- soft. At a high temperature hydrate of potassium volatilizes without change; heat
alone cannot decompose the caustic alkalies.  In the chemical
laboratory, solutions of caustic potash and caustic soda are in
frequent use for absorbing acid gases, such as carbonic acid, and
especially for separating the hydrates of other metals from solutions of their salts.
Exp. 244.-Dissolve a crystal of blue vitriol (sulphate of copper) in
a few centimetres of cold water, and add to the solution several drops
of a solution of caustic soda (or potash). The hydrate of copper is
precipitated as a delicate, blue, insoluble powder, while colorless sulphate of sodium (or potassium) remains in solution.
CuSO-  +   2NaHO   -   CuH20O   +   Na2SO4.
Sulphate of Copper.       Hydr'ate of Copper.
Exp. 245.-Place in a small flask 4 or 5 grms. of chalk or marble
(carbonate of calcium), and 7 or 8 c. c. of water; then cautiously add
chlorhydric acid, little by little, until effervescence ceases and the
chalk is dissolved. When the effervescence is not violent, the flask
2 E




418                     STRONG BASES.
may be warmed to facilitate the process of solution. A rather concentrated solution of chloride of calcium will thus be obtained.
CaCO, +  2HC1   =  CaCl2  +  H20  +  C02.
Carbonate of Calcium.   C6qoride of Calcium.
Add to this solution a few drops of a solution of caustic soda, which is
free from carbonic acid. A white precipitate of hydrate of calcium
will immediately appear, since this hydrate is insoluble in the menstruum, while chloride of sodium will be found in the clear solution.
CaCl2  +  2NaHO  =  CaH202  +  2NaC1.
Hydrate of Calcium.
501. On account of this power of precipitating other hydrates
from solutions of their salts, the caustic alkalies are often called
strong bases, as he is the strongest wrestler who throws his adversary; but this term " strong" is applied to bases and acids so
confusedly as to be frequently a hindrance rather than a help in
classification. Thus, if the reaction which occurs in the preparation of caustic soda (Exp. 233), between carbonate of sodium
and hydrate of calcium, be compared with the reaction last given
between chloride of calcium and hydrate of sodium, it will be seen
that in the first calcium displaces sodium, while in the second
sodium displaces calcium; in the one case hydrate of sodium is
eliminated from the reaction, and in the other hydrate of calcium.
So of acids; we have before had occasion to remark that, of two
acids, now the one and now the other will be stronger, according to the temperature at which the contest between them takes
place, or other extrinsic conditions. Thus sulphuric acid is at
certain temperatures capable of displacing phosphoric acid or
boracic acid, but at high temperatures both these acids displace
it (~~ 294, 448). It is obvious that the definition of the term
"strength " must be very vague and unsatisfactory when applied
to relations thus capable of actual reversal. Two general principles, however, have been arrived at through the comparative
study of such reactions; these are: —1st. When from all or part
of the elements of any mixture of liquefied materials a substance
can be formed which is insoluble in the existing menstruum, that
substance will separate in the solid state; 2nd. When from  all
or part of the elements of a solid or liquid mixture a substance
can be compounded which is volatile at the existing, or induced,




ALKALIMETRY.                    419
temperature, that substance will be eliminated in the gaseous
state. In either case such interchange of atoms as may be essential to the formation of the eliminated substance, takes place, and
the remaining elements necessarily adjust themselves to new relations. Such insoluble precipitates often present peculiarities of
color or texture by which they may be recognized; and such
volatile gases may frequently be identified by their color, odor,
or specific gravity, or by the chemical effects which they are
capable of producing. If every chemical element were known to
yield, under attainable conditions, a characteristic precipitate, or
to evolve a peculiar and recognizable gas, the analytical chemist
would possess the means of detecting every element with certainty. This is to a great extent the case, and chemical analysis
is chiefly based upon a knowledge of the degrees of solubility
and volatility which belong to a great variety of chemical substances with whose appearance and prominent properties the
analyst has previously made himself acquainted; of these substances many are common, but not a few rare and useless, except
to serve the purpose of the analyst.
There exists an anhydrous oxide of potassium, K20, and also
a peroxide. The anhydrous oxide K2O is a gray, inodorous, hard,
brittle solid; it melts a little above a red heat, but volatilizes
only at a very high temperature.
502. Alkalimetry.-Since the value of the carbonates and hydrates of sodium and potassium, as they are manufactured and
consumed on the large scale in the chemical arts, is generally dependent upon the amount of alkali which they contain ready to
enter into chemical combination, it is important to have some
quick and easy method of determining how much available alkali
any sample of these substances really contains. The impurities
which most frequently contaminate the carbonate of potassium
are the chlorides and sulphates of potassium and sodium, silicie
acid, lime, alumina, and the oxides of iron; the commonest impurities of carbonate of sodium are the chloride and sulphate
of sodium, as might readily be inferred from consideration of the
process by which the carbonate is manufactured. Some sulphite
of sodium also is not infrequently present in commercial soda-ash.
Both carbonates are apt to contain small proportions of the hy2 E 2




420                  VOLUMETRIC ANALYSIS.
drates; but as the hydrates are quite as valuable for most uses
as the carbonates, weight for weight, this admixture is not inconvenient. In the common methods of testing the carbonates, the
hydrates present are estimated as available alkali, but are made
no separate account of.
It would be foreign to our purpose to enter upon the details of
alkalimetry; the process consists essentially in ascertaining how
much dilute sulphuric acid of a known strength is required to
neutralize exactly a known weight of the sample examined. The
requisites are a graduated burette (Appendix, ~ 21), a standard
acid, pure carbonate of sodium wherewith to prepare this acid,
and a colored solution, sensitive to both acid and alkali, to indicate the point of neutralization. The following experiment will
give some idea of the manipulation required in this sort of analysis,
which, on account of its rapidity, is of very general application
in technical chemistry.  The general method is called volumetric,
because, when once the standard liquids are prepared, quantitative results are obtained, not by weig'hing, but by measuring the
bulk of liquid consumed in the testing.
Exp. 246.-Weigh out accurately 5 grms. of pure anhydrous carbonate of sodium; transfer it to a flask or beaker having the capacity of
about 400 c. c.; dissolve it in about 200 c. c. of water, and color the
solution blue with about 2 c. c. of a violet tincture of litmus. To prepare this tincture, digest 1 part of litmus in 6 parts of water, on a
water-bath, for several hours; filter; divide the blue liquid into two
equal portions, and stir one half repeatedly with a glass rod dipped in
very dilute nitric acid, until the color just appears red; then add the
other blue half, together with one part of alcohol, and keep the tincture
in a small open bottle. In a stoppered bottle the tincture fades.
Mix about 60 grms. of strong sulphuric acid with 500 c. c. of water
and let the mixture cool (~ 233). Fill a 50 c. c. Mohr's burette (Appendix, ~ 21) up to the 0 mark with the cold dilute acid. Place the
flask or beaker containing the soda solution beneath the b4ette, gently
press the spring-clip, and allow the acid to flow gradually into the soda
solution, stirring the while, until the color of the liquid changes to a
Wine-red; then place the flask or beaker over a lamp, and bring the
liquid to ebullition; the dissolved carbonic acid will be driven out, and
the liquid will again become blue; more acid is then added to the
nearly boiling fluid, the vessel being occasionally placed over the lamp,
until the color of the liquid becomes red, slightly inclining to yellow.
~ —o —   -— ^'-b'  "^'"




VALUATION OF POT- AND SODA-ASH.              421
When the point of saturation is approaching, add the acid two drops at
a time, and after each fresh addition dip a fine glass rod into the fluid,
and make with it two spots on a slip of fine blue litmus-paper, reading
the volume each time, and marking the number of centimetres between
the two spots.  Continue this operation until the spots on the paper
appear distinctly red; then dry the paper and take for the correct
number that figure which stands between those two spots which just
remain red when dry.
For some eyes, tincture of cochineal possesses great advantages over
tincture of litmus as a means of recognizing the point of neutralization.
The tincture of cochineal is prepared by digesting 3 grms. of powdered
cochineal in a mixture of 50 c. c. of alcohol and 200 c. c. of water, at
the ordinary temperature, for several days. The tincture, Which may
be either decanted or filtered from the residue, has a ruby-red color.
The caustic alkalies and the alkaline carbonates change the color to a
violet-carmine; solutions o; strong acids and acid-salts make it orange;
to carbonic acid it is nearly indifferent. 10 or 15 drops of the tincture
are sufficient to color 200 c. c. of liquid.
Dilute the acid which remains of the original 500 c. c. with enough
water to give a fluid of which exactly 50 c. c. are required to saturate
5 grms. of pure carbonate of sodium.  This dilution is effected as follows:-Suppose that 40 c. c. of the acid have proved sufficient to neutralize 5 grmins. of the pure carbonate, then 10 c. c. of water must be
added to every 40 c. c. of the acid. 400 c. c. of the original dilute
acid remain; now, as 40: 460= 10: number of c. c. water to be added
= 115 c. c. Dilute, therefore, the acid with 115 c. c. of water, using
some of this water to rinse the burette which contains the undiluted
acid, the measuring-glass which may have been used to ascertain the
extent of volume of the original acid, and any other vessel into which
it may have been temporarily poured. Of this diluted acid, 50 c. c.
should exactly neutralize 5 grms. of pure anhydrous carbonate of sodium. It may be tested by again weighing out 5 grms. of the pure
carbonate, and repeating the volumetric determination precisely as
above described. If the work has been well done, the standard acid
will be found ready for use.
Weigh out 5 grms. of common soda-ash, and dissolve in about
200 c. c. of water whatever of the sample is soluble; repeat upon this
liquid, without filtering, the volumetric operation of the last experiment. The number of half c. c. of standard acid used gives directly
the percentage of pure anhydrous carbonate of sodium which the
sample contains. If,50 c. c. or 100 half c. c. are used, the sample is
pure carbonate of sodium; if 40 c. c. or 80 half c. c. are used, the




422                      POTASSiUM.
sample is eighty per cent. pure carbonate, and twenty per cent. water
and other impurities.
503. When an acid has been prepared, of which 50 c. c. exactly
neutralize 5 grms. of pure carbonate of sodium, it is a matter of
simple calculation only to determine how much hydrate of sodium,
how much carbonate of potassium, and how much hydrate of
potassium  the same 50 c. c. acid will neutralize.  The following are the proportions required, the atomic weight of potassium
being 39'1:106:         80    -          5: x - 3'773;
Combbining weight of   Corn. wt. of   Grn1s.          Grins.
Na2CO3.         NatII:O2. 1a9CO,.              Na2H.02.
106:        138'2   =         5: x -- 6519
Combining weight of   Cornm. wt. of    Grins.         Grms.
Na2CO3.          K2CO3.         Na2CO3.          K2CO3.
106:        1122   =          5: x -= 5292.
Combbinig weight of   Corn. wt. of     Grins.         Grinms.
NaCO3.          K21120          Na CO,.        K2He02i
The following table shows the quantities of these four substances
which are equivalent in neutralizing- or combining-power:5'000 grmins. of carbonate of sodium
3~'773                              1are saturated by 50 e. e.
63519,,,, carbonate of potassium    standard acid.
5'292,,,, hydrate,,,,
It is apparent from this table that a given weight of carbonate of
sodium will saturate more silicic acid or more fat, or will give off
more carbonic acid than the same weight of carbonate of potassium
and that the hydrate of sodium is also more efficient, gramme for
gramme, than the hydrate of potassium.  These facts are to be
inferred at once from the atomic weights of sodium and potassium,
which are 23 and 39'1 respectively; they have had their weight
in bringing about the rapid substitution of sodium-salts for potassium-salts in the chemical arts.
504. Potassium (K).-This element, like sodium, is made from
its carbonate by heating intensely a mixture of the carbonate and
charcoal, in accordance with the reaction,
K2CO3 +  2C =  2K  +  3CO.




POTASSIUM DECOMPOSES WATER.             423
The apparatus employed is similar to that described in treating
of sodium (~ 487); and as potassium closely resembles sodium, the
same general method is followed, and the same precautions are
observed. A second distillation of the crude potassium is absolutely essential, because, if it be neglected, a black, detonating
compound of uncertain composition is formed, which explodes
violently upon the slightest friction.
Potassium is a silver-white substance, of very brilliant lustre,
which is brittle at 00, soft as wax at ordinary temperatures, fuses
at 62~.5, and is volatile. at a red heat. In its soft state, two clean
surfaces can be welded together like iron. It is lighter than
water, having a specific gravity of only 0'865. It is almost instantaneously oxidized by air and water at the ordinary temperature, and, when heated, by nearly every gas or liquid which contains oxygen. Like sodium, it must therefore be collected and
kept under naphtha, out of contact with the air. At moderate
temperatures, potassium readily absorbs hydrogen, nitrogen, and
carbonic oxide, and enters into direct combination with chlorine,
bromine, iodine, sulphur, selenium, and tellurium. We have had
occasion to avail ourselves of its intense chemical energy (~~ 85,
97, 411).
Exp. 247.-Throw a piece of potassium, as large as a small pea, upon
some cold water in the bottom of a large bottle, and place a card or
glass plate over the mouth of the bottle. The potassium decomposes
the water, and evolves heat enough to kindle the hydrogen which is
given off; this hydrogen burns with a purplish-red color, imparted to
the flame by a minute quantity of vaporized solid. This color is characteristic of potassium' compounds, as a yellow color is characteristic
of sodium compounds.
505. It is not a matter of indifference from what kind of a
mixture of carbonate of potassium and carbon potassium is prepared. The material which is best adapted to its preparation is
the potassium-salt of a vegetable acid rich in carbon, which, when
decomposed by heat in a vessel from which air is excluded, yields
carbonate of potassium and a large quantity of free carbon.
While grape-juice is being converted into wine by fermentation,
a stony deposit, called "tartar," which is sometimes gray and
sometimes reddish, fastens to the bottom  and sides of the casks




424                CHLORIDE OF POTASSIUM.
which contain the fermented juice. When freed by recrystallization from  adhering coloring-matters, this crystalline and difficultly soluble substance is a white salt, acid and cooling to the
taste.  It is an acid tartrate of potassium, and is commonly called
"cream of tartar."  When this substance, or crude tartar, is
heated in a covered crucible until it ceases to emit combustible
vapors, the cooled residue is found to consist of a porous mass of
carbonate of potassium, intimately mixed with very finely divided
carbon. This mixture is the best material from which to prepare potassium; it is also an excellent flux, useful in assaying
the ores of the common metals.  In wine-producing countries
considerable quantities of excellent carbonate of potassium are
prepared from the deposits of the wine-vats, by dissolving the
carbonate out of the ignited tartar and purifying the salt, so extracted, by recrystallization.  The carbonate so obtained is the
purest source of hydrate of potassium for laboratory use.
506. Chloride of Potassium (KC1).-This salt is a subordinate
source of potassium compounds. It is extracted in considerable
quantity from the ashes of sea-weeds, and is largely used in the
manufacture of common alum, which is a sulphate of aluminum
and potassium. It occurs native, sometimes pure, but more frequently mixed or combined with other chlorides. The chloride
of potassium is capable of uniting with most other metallic chlorides, forming crystallizable double salts.  The native double
chloride of potassium and magnesium (KC1,MlgCl2,6H20) has become of late years a productive source of potassium-salts. This
mineral is dissolved in hot water; from the cooled solution the
greater part of the chloride of potassium crystallizes out, while
the chloride of magnesium  remains in solution.  Chloride of
potassium occurs also with the chlorides of sodium, magnesium,
calcium, and other salts in sea-water and brine-springs, and is
obtained as a secondary product in the preparation of chlorate of
potassium (~ 517), the purification of nitre (~ 514), and in several other manufacturing-operations. It may be prepared directly
from potassium and chlorine, or from the carbonate or hydrate
of potassium and chlorhydric acid.
Chloride of potassium  crystallizes in anhydrous cubes, looks
and tastes like common salt, is not acted upon by the air, de



IODIDE OF POTASSIUM.                425
crepitates when heated, melts at a low red heat, and volatilizes
unchanged at a higher temperature. It is somewhat more volatile than common salt, is more soluble in water, and produces a
greater degree of cold in dissolving. This chloride enters into
some highly crystalline compounds, of curious composition, of
which the product of the following experiment may serve as an
example:Exp. 248.-By the aid of a gentle heat, dissolve 6 grms. of powdered
red chromate of potassium in 8 grms. of strong chlorhydric acid,
avoidin-g evolution of chlorine. WVhen the powder is dissolved, allow
the solution to cool; flat, red prisms will crystallize from the liquid.
This compound answers to the formula KCl,CrO3. It is decomposed
by solution in water.
K20,2Cr03 + 21HC1 = 2(KC1,Cr03) + 1H20.
Red Claoronate
of Potassium.
507. Bromide of Potassium (KBr) is a very soluble salt, which
crystallizes in cubes, and closely resembles in all its properties
the chloride. Potassium and bromine unite directly, with inflammation and violent detonation, the bromide being the product of the
reaction. When bromine is added to a solution of caustic potash
until the liquid acquires a permanent yellowish tinge, a mixture
of bromide and bromnate of potassium is produced:6KHO +  6Br =  5KBr +  KBrO3 +  3H20.
The mixed salts are dissolved in water, and a current of sulphydric acid is passed through the solution to reduce the bromate:KBrO3, +  3HS =  KBr =  3120 +  3S.
The liquid is then gently warmed to expel the excess of gas,
filtered from the deposited sulphur, and evaporated until the bromide crystallizes. The salt has lately come into use in medicine
as a sedative.
508. Iodide of Potassium (KI).-This valuable medicine and
photographic material may be procured by adding iodine to a solution of hydrate of potassium until the liquid turns brown, and
gently igniting the residue obtained by evaporation.  The process and the reaction are the same as in the preparation of the
bromide, except that the iodate may be decomposed by heat alone




426                 CYANIDE OF POTASSIUM.
and does not require reduction by sulphuretted hydrogen. A
better mode of preparing the iodide of potassium is to be found
in the decomposition of iodide of iron by carbonate of potassium.
Digest 4 grins. of iodine and 2 grins. of iron filings in a stoppered
bottle with 20 c. c. of water; iodide of iron (FeI2) is formed under
these conditions by direct union of the elements. The liquid is then
transferred to a flask and boiled, and a solution of carbonate of potassium is added by small portions so long as a precipitate occurs. The
solution, filtered from the insoluble carbonate of iron, yields on evaporation crystals of iodide of potassium.
Iodide of potassium ordinarily crystallizes in semiopaque cubes.
It is permanent in the air, has a sharp taste, melts below a red
heat, and volatilizes unchanged at a low red heat.  It is very
soluble in water, and in dissolving produces a considerable fall of
temperature. Alcohol also freely dissolves this salt. The facility
with which the salt is decomposed and its iodine liberated by
chlorine, ozone, and nitric acid has been already amply illustrated
(Exps. 70, 72, 76).
Iodide of potassium is much used in medicine; it is not poisonous even in doses of 5 to 20 grms.  Its solution in water dissolves iodine to a large extent, acquiring thereby a dark-brown
color. This brown solution is sometimes used in medicine as a
vehicle for iodine; it is also useful to the photographer for removing from the skin the stains produced by nitrate of silver.
Iodide of potassium is, further, employed in the photographic process upon glass, to produce in the substance of the film of collodion a deposit of iodide of silver by double decomposition with
nitrate of silver.
509. Cyanide of Potassium (KCN).-The elements which make
part of the whitish, soluble, fusible, and deliquescent solid which
bears this name are potassium, carbon, and nitrogen. At a brightred heat these three elements will come together to form this salt
out of quite a variety of materials, and under quite various circumstances. When nitrogen gas is passed over a mixture of
charcoal and hydrate or carbonate of potassium, at a bright-red
heat, cyanide of potassium is formed. When nitrogenous organic
matters are fused with hydrate or carbonate of potassium, the
cyanide is formed.  In presence of iron scraps or filings, this




SULPHIDES OF POTASSIUM.             427
mixture produces a common cyanide, containing potassium, iron,
carbon, and nitrogen, and known in the arts as " yellow prussiate
of potash," and to chemists as ferrocyanide of potassium. If this
"prussiate " be ignited at a moderate red heat, it is decomposed
into cyanide of potassium, nitrogen, and a compound of iron with
carbon:4KCN,Fe(CN), =  4KCN  +  2N  +  FeC2.
In the blast-furnaces in which iron ores are smelted with coal
or coke, a considerable quantity of cyanide of potassium is often
produced, the nitrogen being probably derived from the torrents
of nitrogen which the blast of air carries into the furnace. Instead of igniting the yellow prussiate alone, a more economical
process is to ignite a mixture of the prussiate with carbonate of
potassium. This method saves all the cyanogen in the prussiate;
but the salt thus obtained is always mixed with cyanate and carbonate of potassium:K4Fe(CN)6 +  K2CO3 =  5K(CN) + K(CN)O +  CO,2 +  Fe.
The presence of these impurities does not injure the cyanide for
many of its uses.
Cyanide of potassium is of great use in galvanic gilding and
silvering, since the cyanides of gold and silver are both soluble in
a solution of cyanide of potassium. Its solution dissolves the
sulphide of silver, and has therefore been suggested for household use in cleaning silver-ware; photographers sometimes use it
for removing stains of nitrate of silver from the hands; but both
these applications of cyanide of potassium are dangerous and inexpedient. The cyanide is intensely poisonous, not only when
taken internally, but also when brought into contact with an
abrasion of the skin, a cut, or scratch. As a reducing agent,
cyanide of potassium has great power; it is especially useful in
blowpipe reactions (Exp. 132).
510. Sulphides of Potassium.-Potassium, heated in sulphurvapor, takes fire readily, and burns brilliantly.  There are supposed to be five different compounds of potassium and sulphur,
corresponding to the formulae KS, K2S2, K2S3, K2S4, and K,S,; and
there is a sulphydrate KHS, analogous to the hydrate KHO.
Erp. 249.-Heat, gently, a thorough mixture of 10 grms. of dry




428                 SULPHATE OF POTASSIUM.
powdered carbonate of potassium and 6 grms. of sulphur, in a covered
earthen or iron crucible, until effervescence ceases and the mass flows
quietly. The fused mass has a yellowish-brown color, and consists of
tersulphide, quinquisuiphide, and intermediate sulphides of potassium,
mixed with sulphate, and often with carbonate of potassium; it is
called "liver of sulphur."  This substance, dissolved in water, gives a
greenish solution, trom which dilute acids liberate sulphuretted hydrogen, and precipitate milk ot sulphur (~ 198). Tlhe carbonic acid of
the air is strong enough to effect this decomposition' hence the solid
substance and its solution, when exposed to the air, smell of sulphuretted hydrogen: —
KoS3 +  II2S04 = KS04 + II2S  - 2S.
The chief use of liver of sulphur is in the medical treatment of skindiseases.
511. When sulphydric acid gas is passed to saturation into a
solution of caustic potash, a colorless solution.is obtained, which
is supposed to contain the sulphydrate, KHS. It has no permanency, quickly absorbing oxygen and turning yellow.
All the sulphides of potassium are brown solids, having an alkaline reaction to test-paper, and smelling of sulphydric acid.
Their solutions are oxidized by exposure to the air, and sulphur
is deposited from them.
512. Sulphate of Potassium  (K2SO4).-This anhydrous salt
crystallizes in transparent hexagonal prisms, terminated by hexagonal pyramids, aid consequently bears some resemblance to
common quartz crystals. The salt has a strong tendency to
form double sulphates; a double sulphate of potassium and magnesium is of importance in the manufacture of potassium-salts
from sea-water. It also enters into the composition of many
of the double sulphates, which are called alums from the name of
the commonest member of the class, the sulphate of aluminum
and potassium.
513. Sulphate of Potassium and Hydrogen (K11SO4).-This salt,
commonly called the "bisulphate," is formed on a large scale as
a residuary product whenever nitric acid is manufactured from
nitrate of potassium.  When ignited, its crystals lose half their
acid:2(KIIS0) =  K2S04 +  H2S04;
and the salt is therefore sometimes used as a flux, in cases where




REFINING OF SALTPETRE.                 429
the action of a strong acid is required at a high temperature upon
salts or oxides with which it may be fused.  Platinum tools may
be cleaned by fusing this salt in or upon them.
514. Nitrate of Potassium (KN03).-This valuable salt, commonly called saltpetre, is very widely diffused as a natural product, being indeed seldom wholly wanting in a productive soil, or
in spring- or river-water.  At many localities it is found in
caverns or caves in calcareous formations; but the principal commercial source of the salt is the soil of certain tropical regions,
especially of districts in Arabia, Persia, and India, where the
nitrate is found disseminated through the upper portion of the
soil, or forming an efflorescence upon the surface, but never
extending lower than the depth to wh;ch the air can easily
penetrate.
This natural production of nitrates appears to result mainly from the
putrefactioi of vegetable and animal matters, in presence of the air
and of alkaline or earthy bases capable of fixing the nitric acid as soon
as it is formed. The well-waters of towns, contaminated by neighboring serers and cesspools, nearly always contain nitrates. The decay
of the luxuriant vegetation of the tropics, promoted by a hot sun and
a moist atmosphere, is a never-failing source of ammonia; but it is not
certain that the production of ammonia is a necessary stage in the process of nitrification.  In the artificial production of nitrates in temperate climates, the supposed natural conditions have been roughly
imitated. In the old saltpetre " plantations " of European countries,
nitrate of calcium was produced by mixing decomposing vegetable and
animal matters with cinders, chalk, marl, and so forth, moistening the
mass repeatedly with leachings of manure-heaps, exposing it freely to
the air for two or three years, and then lixiviating. The nitrate of
calcium, which was the main product, was then converted into saltpetre by double decomposition with carbonate of potassium.
The saltpetre is extracted from the earth which contains it by lixiviation, evaporation, and crystallization; but inasmuch as for most uses
it is required in a very pure state, the crude salt must generally be refined. The common impurities of crude saltpetre are chloride of sodium and chloride of potassium. In order to separate these chlorides,
advantage is taken of the fact that the nitrate of potassium is four
times as soluble in boiling water as the chloride of potassium, and six
times as soluble as the chloride of sodium. The crude saltpetre is
treated with a quantity of water, sufficient to dissolve at boiling all the
nitrate of potassium, but not all of the chloride of sodium beside. This




430                PROPERTIES OF SALTPETRE.
residual salt is scooped out of the vessel in which the solution is
effected, and the solution, after being somewhat diluted, is boiled with
a little glue, to coagulate the coloring-matters and other soluble dirt,
and sweep the liquid clean by means of the adhesive scum which rises
to the surface. The strong, clear solution is then transferred to shallow
crystallizing pans, and left at rest if large crystals are desired; if small
crystals are preferable, the liquid is constantly stirred from the moment
that crystallization begins; the saltpetre is then deposited in a crystalline powder, called saltpetre-flour.  The chlorides of sodium  and
potassium are nearly as soluble in cold water as in hot; but nitrate of
potassium is only one-eighth as soluble in water at the temperature of
the atmosphere as in boiling water. Hence the chlorides remain in
the mother-liquor, while the nitrate rapidly separates from the solution
as it cools. The crystals of saltpetre are drained, and then washed
with a solution of saltpetre saturated in the cold. This solution takes
up the adhering chlorides, but leaves the pure nitrate of potassium
undissolved.
Large quantities of saltpetre are now made by decomposing nitrate
of sodium with carbonate of potassium. When, through governmental
interference, the East-Indian supply of saltpetre is checked, this
method is resorted to with advantage. Tartar, and the ashes of the
beetroot-sugar manufacture, are good sources of potash to be applied
to this purpose. Crude nitrate of sodium contains so much chloride
of sodium, that it is desirable to purify it for this use by previous recrystallization; otherwise potash would be unprofitably consumed in
converting chloride of sodium into chloride of potassium. One of the
processes for converting the nitrate of sodium into nitrate of potassium
consists simply in adding the nitrate of sodium to a hot concentrated
solution of carbonate of potassium; a precipitation of carbonate of sodium takes place, and this precipitate is removed as fast as it forms,
until no more appears; from the cooled liquid saltpetre-flour is deposited. The carbonate of potassium may be replaced in this process
by chloride of potassium. NaNO3+KC1= KNO3+NaCL
515. Nitrate of potassium is white, inodorous, and anhydrous,
and has a cooling, bitter taste. When pure, it is permanent in
the air,-a fact of great importance, inasmuch as the chief use
of this salt is in the manufacture of gunpowder.  Were it hygroscopic, like nitrate of sodium, it would not be applicable to this
use. Saltpetre is one of those few potassium-salts which cannot
be wholly replaced in the arts by the corresponding sodium-salt.
It is very soluble in water, especially in hot water; it melts below




SALTPETRE NOT EXPLOSIVE.                431
a red heat to a colorless liquid without loss of substance, but at a
red heat it gives off oxygen and suffers decomposition.  Its
most marked chemical characteristic is its oxidizing-power. It
deflagrates in the fire with charcoal, sulphur, phosphorus, and
other combustible bodies; when ignited in contact with copper
or iron (Exp. 46), it converts these metals into oxides; and it
even oxidizes gold, silver, and platinum. It is on the oxidizingpower of saltpetre that its use in the manufacture of gunpowder
and fireworks, and in the preparation of matches, depends.
Arrange 10 grms. of pure nitrate of potassium and 20 or 30 grms. of
thin copper-turnings, or small bits of sheet copper, in alternate layers,
in a covered copper crucible, and expose the mixture for half an hour
to a moderate red heat. Dissolve out the cooled mass with water, and
let the liquid stand in a tall, closed bottle until the oxide of copper has
settled to the bottom. The supernatant liquid is a pure solution of
caustic potash; indeed this is an excellent method of preparing pure
hydrate of potassium for use in analysis:2KNO0  +  5Cu + 1120 = 2KHO + 5CuO + 2N.
Exp. 250.-Heat 10 or 12 grinms. of saltpetre, gently, in a small
porcelain dish, until it melts; pour the melted salt out on a cold piece
of iron or stone; break the fused mass into small fragments, and fill
an ignition-tube, 12-15 c.m. long, one-third full with these bits. Heat
the tube cautiously, taking pains to keep all the salt, when once melted,
in a state of fusion. At a red heat, oxygen, pure at first, is slowly
evolved, and may be collected at the water-pan; simultaneously nitrite
of potassium (KN02) is formed; at a second stage this nitrite is itself
decomposed, and the escaping oxygen is then contaminated with a
certain proportion of nitrogen. A portion of the gas collected may be
tested with a glowing splinter (Exp. 6); another portion may be mixed
with coal-gas, and, with the mixture, bubbles may be blown, as directed
in Exp. 30; the mixture will be found to be exceedingly explosive.
This experiment proves, in the first place, that saltpetre itself is not
explosive, and, in the second place, affords an explanation of the fact
that frightful explosions do often occur when storehouses containing
saltpetre are burned. Carburetted hydrogen, such as was obtained in
Exp. 151 (represented by the coal-gas in the last experiment), is
evolved from the wood-work of the burning building, wherever the
wood is heated out of contact with the air; meanwhile oxygen is given
off from the ignited saltpetre, and whenever these two gases mix in the
requisite proportions, and their mixture comes in contact with a flame,
a violent explosion inevitably ensues.




432                      GUNPOWDER.
u'zp. 251.-Dissolve 5 grins. of saltpetre in 20 c. c. of water; dip
strips of bibulous paper in the solution, and dry them; this paper,
once kindled, will smoulder away till consumed. It is used in connexion with fireworks, and in the manufacture of pastilles and aromatic
fumigating paper.
Exp. 252. -Mix 5 grms. of powdered saltpetre with 1 grm. of dry
powdered charcoal; place the mixture on a piece of porcelain and i>,nite it with a hot wire. When the deflagration is over, a white solid
will be found upon the porcelain. Dissolve this solid in a few drops
of water; the solution will be alkaline to test-paper; add a few drops
of a dilute acid; a brisk effervescence marks the escape of carbonic
acid. The nitrate has oxidized the carbon to carbonic acid, part of
which escaped with the nitrogen during the deflagration, while part
entered into combination with the potassium:4KNO,  + 5C    2KC0, + 3CO2 + 4N.
Exp. 253.-Place 30 grms. of saltpetre in a small beaker with 110
c. c. of water; insert a thermometer in the mixture, and observe the
very considerable fall of temperature occasioned by the solution of the
salt. In those countries where saltpetre is cheap and ice dear, this
property of the salt is availed of for the refrigeration of drinks.
516. Gucnpowder is an intimate mechanical mixture of softwood charcoal (~ 382), sulphur, and nitrate of potassium, in the
proportions of 70 or 80 per cent. of the nitrate to 10 or 12 per
cent. of each of the other ingredients. Though it is in no sense
a chemical compound, we may, for convenience' sake, express
the composition of gunpowder by the formula K,2N20O+S+3C,
and may roughly formulate the reactions which occur when it is
burned, by the following equation:K2N206 +  S +  3C =  3CO2 +  2N  +  K2S.
Speaking in general terms, the oxygen of the nitrate combines
with the carbon to form carbonic acid, or, at the least, carbonic
oxide, while the sulphui is retained by the potassium, and nitrogen left free. Gunpowder burns at the expense of the oxygen
contained in it; it has no need of air for its combustion, but can
be burned in any closed space-as well, for example, in canisters
under water, or tightly enclosed in the chamber of a gun, as in
free air.
From the formula, it will be seen, at a glance, that a very large
proportion of gas, as compared with the bulk of the solid powder,




PREPARATION OF GUNPOWDER.               433
must be evolved when powder is burned.  But gunpowder burns
rapidly and with great evolution of heat, so that the volume of
gas, large at any temperature, is enormously expanded at the
moment of its formation; hence it happens that the gas set free
in the barrel of a gun may be capable of occupying a thousand
or fifteen hundred times as much space as the powder which generated it. An enormous pressure is thus engendered at the spot
where the powder burns, and to this pressure some part of the
matter which confines the powder must yield.  In the case of
the gun-barrel, it is the bullet which represents the weakest, or
breaking side of the chamber in which the powder burns; but
when rocks are blasted, then the packing, or "tamping," which
represents the ball, is made so firm that it shall be stronger than
the rocky sides of the drill-hole, which is equivalent to the barrel of the gun. In case the walls of the gun can be disrupted
more readily than the firmly impacted bullet can be driven out,
then, of course, the gun bursts; and, conversely, the tamping of
a drill-hole is thrown out if it be less firm than the rock. In the
case of the gun-barrel, a part of the effect of the explosion is felt
in the kick or recoil of the gun.
Though the equation last given is useful in so far as it exhibits
the gaseous products evolved,during the combustion of gunpowder, it does not truly express the solid products of the reaction. The residue of the combustion really contains only a comparatively small proportion of sulphide of potassium; it consists
mainly of sulphate of potassium  and carbonate of potassium,
together with some hyposulphite of potassium, and a trace of unburned carbon. Enough sulphide of potassium is always present,
however, to impart the offensive odor which is perceived in washing a foul gun, and in powder-smoke.
Exp. 254.-Pulverize, separately, 23 grinms. of nitrate of potassium,
4 grins. of sulphur, and 4 grins. of charcoal. Place a drop or two of
water in a porcelain mortar, and grind into it, first, the charcoal, and
then the other ingredients, taking care to add enough water to form a
plastic dough. After the mass has been thoroughly kneaded, roll out
small portions of it between two pieces of board, into long threads, of
the thickness of a fine knitting-needle. With a knife, cut the threads
into small fragments or granules, and leave the granules in a warm
room to dry. The thoroughly dried product is gunpowder; and the
2 F




434                 CHLORATE OF POTA.SSIUM.
manipulation in this experiment does not differ essentially from th,
mode of manufacture employed in the powder-mills, excepting that the
granulation is there effected by passing the moist paste through cullenders.
The sulphur in gunpowder acts mainly as a kindling material
(~ 200). In powder intended for use in guns, the proportion of sulphur is kept comparatively low, since any excess of it would corrode
the metal of the gun..Exp. 255.-Knead together, as in Exp. 254, 7 grms. of powdered
nitrate of potassium and 1-5 grm. of moistened, finely powdered charcoal. Granulate and dry the product, as before, and compare its inflammability with that of the gunpowder prepared in Exp. 254, by
touching small heaps of each with a red-hot wire. Mixtures of charcoal and nitrate of potassium, such as the foregoing, are much used in
the manufacture of fireworks.
517. Chlorate of Potassium (KC103).-The basis of the large
use now made of this beautiful salt in medicine, in calico-printing, in pyrotechny, in the match-manufacture, and in the chemical
laboratory is its large oxygen-contents. It is an oxidizing agent
of the most vigorous description.
It may be prepared by saturating a solution of 1 part of hydrate of
potassium in 3 parts of water with chlorine, and heating the liquid
some time to the boiling-point. The ultimate result may be expressed
by the formula
6KHO + 6C1 =  KCIO3 +- 5KC1 +  311O;
but the process has two stages, which are sufficiently described in ~ 124.
The hot solution, left to itself, deposits the greater part of the chlorate
in anhydrous six-sided plates of a pearly lustre; the chloride of potassium remains in the mother-liquor. The chlorate is freed from adhering chloride by recrystallization. The success of the process depends
upon the very different solubilities of the chlorate and the chloride of
potassium. At the temperature of their saturated boiling solutions
both salts are about equally soluble; 100 parts of water will dissolve
between 60 and 67 parts of either salt; but at the ordinary temperature of the air 100 parts of water will dissolve 30-40 parts of chloride
of potassium and only 6 or 7 parts of chlorate of potassium. We find
here the explanation of the fact that chlorate of sodium has not replaced chlorate of potassium in the arts. The chlorate of sodium is
more soluble in water at all temperatures than the chloride of sodium
is, while both are exceedingly soluble, so that the two salts cannot be
separated by crystallization. This process of crystallization is the
chemical manufacturer's chief reliance in refining both his materials




PREPARATION OF CILORATE OF POTASSIUM.          435
and his products; and the purchaser of chemicals finds his best guaranty
of the purity of his commodities in the peculiar form, lustre, color, and
degree of transparency which characterize the crystals of every crystallizable and permanent chemical compound. Hence an easily crystallized permanient salt, of characteristic appearance, like chlorate of
potassium, will always have the preference over one which, like chlorate of sodium, can be crystallized and purified only with difficulty,
and is not permanent when once obtained. The chlorate of sodium is
deliquescent.
The waste product in the making of chlorate of potassium by the
process just described is chloride of potassium, a comparatively dear
salt. An economy is effected by substituting hydrate of calcium for
hydrate of potassium, and thus making the secondary product chloride
of calcium instead, of chloride of potassium; one equivalent only of the
chloride of potassium is then required instead of six of the hydrate of
potassium. An excess of chlorine is passed into a mixture of 300 parts
of quicklime, 154 parts of chloride of potassium, and 100 of water.
The mass is heated by steam, stirred with agitators, filtered, and then
evaporated nearly to dryness by steam-heat; the mass is then redissolved in hot water and set to crystallize:3CaO + KC1 +  6Cl = KC103 +  3CaCl2.
The mother-liquor, which contains all the chloride of calcium, may be
decomposed with sulphate of potassium, in which event a very finely
divided sulphate of calcium, available for " stuffing " in the manufacture of paper, is precipitated, and chloride of potassium is recovered,
to be again applied to the production of the chlorate; or the chlorideof-calcium solution may be decomposed with carbonate of sodium, in
order to precipitate a very finely divided carbonate of calcium, which
is largely employed by the pharmaceutist and perfumer. In the latter
case, chloride of sodium has to be thrown away. The whole manufacture is a good example of a technical chemical process.
518. Chlorate of potassium is easily decomposed by heat; at a
moderate temperature it yields perchlorate and chloride of potassium (~ 124); but at a red heat it is resolved into chloride of
potassium and oxygen (Exp. 7): —KC103 =KC+30.  Chlorate
of potassium is so prompt an oxidizing agent that mixtures of it
with combustible bodies often detonate violently when struck or
heated (Exps. 113, 157). These combustions are dangerous unless very small quantities be used.  It has been often proposed to
replace gunpowder by such mixtures; a mixture of the chlorate
with catechu, or some similar stable substance rich in tannin, is
2F2




436          OXIDATION BY CHLORATE OF POTASSIUM.
the most promising of these suggestions. Strong acids, like sutlphuric, nitric, and chlorhydric acids, decompose chlorate of potassium with evolution of oxides of chlorine, or of chlorine and oxygen. The decomposition is often attended with decrepitation, and
sometimes with a flashing light; combustibles, like sulphur, phosphorus, sugar, and resin, are inflamed by the gases evolved.  A
mixture of chlorate of potassium and chlorhydric acid is used in
toxicological investigations as an oxidizing agent for the destruction of organic matter (~ 329). The following formulae will elucidate some of these reactions:
3KC103 + 2HS80 = 2C102 + KC104  + 2KHS4   + H120
8KC103 + 6HN03 = 6KNO3 + 2K(C10  + 6C1 + 130 + 31120
4KCiO03 + 12HC1 = 4KC1  + 3C102  + 9C1 + 6H20.
Exp. 256.-Pour into a conical test-glass 25-30 c. c. of water, and
throw into the water some scraps of phosphorus, weighing together not
more than 0'3 grm., and 3-4 grins. of crystals of chlorate of potassium.
By means of a thistle-tube bring 5 or 6 c. c. of strong sulphuric acid
into immediate contact with the chlorate at the bottom of the class.
Then withdraw the thistle-tube. In a moment the phosphorus is
kindled, and burns with vivid flashes of light beneath the water. An
evolution of chlorine accompanies the reaction.
Exp. 257. —Rub 4 or 5 grins. of clean chlorate of potassium, free
from dust, to a fine powder in a porcelain mortar. In powdering chlorate of potassium, care must be taken that the mortar and pestle are
perfectly clean, and the salt free from organic matter, and that violent
percussion and heavy pressure upon the contents of the mortar be
wholly avoided. Place the powdered chlorate on a piece of paper,
add an equal bulk of dry powdered sugar to the pile, and, with the
fingers and a piece of card, mix the two materials thoroughly together.
Mixtures of chlorate of potassium  and organic matter are liable to
explode, if strongly rubbed or but lightly struck. Wrap the mixture
in a paper cylinder, and place the cylinder on a brick in a strong
draught of air; let fall upon the mixture a drop of sulphuric acid froul
the end of a glass rod; a very vivid combustion will ensue, with the
violet-colored flame characteristic of potassium.
Exp. 258.-Mix together, on paper, with the precautions above described, 1 grm. of black oxide of copper, 1 grnm. of sulphur, and 2'5
grms. of powdered chlorate of potassium. Place the mixture, inclosed
in a paper cylinder, on the top of a brick, and touch it with a hot
wire; it will burn vividly, and with a purple color which is prized in
pyrotechuy.




AMMONIUM-SALTS.                   437
CHAPTER XXV.
A X M 0 N I UM-S A L T S.
519. The hypothetical metal ammonium (NH4) is a device for
explaining the constitution and properties of one well-defined
class out of the several classes of compounds into which the gas
ammonia enters. This class of compounds is that which results
from neutralizing an aqueous solution of ammonia with acids, as
in the following reactions:NH3,H20 +  H2SO,  =  (NH,4)HS04 +  H20.
&ulphate of Anmonzum
and Hydrogen.
NH1,IH.  +  INOH    =  (NH4)NO,  +   1H20.
2Anitrate of Anmonznrum.
According to this hypothesis, the crystalline salts which result
from such neutralizations contain a group of atoms (NHI4) which
is analogous in its action to potassium and sodium, and which
forms salts analogous in composition to the potassium-salts. Thus
chloride of ammonium (NH4)Cl is analogous to chloride of potassium KCl1; sulphate of ammonium (NHl4)2S04 is analogous to
sulphate of potassium K2SO4, and so forth (~ 91).
All the actual evidence we possess of the separate existence
and metallic character of the group NH4 is contained in the following curious but inconclusive experiment:Exp. 259.-Pour 8 or 10 c. c. of mercury into a small flask, and
warm the mercury over a gas-lamp; drop upon the mercury six or
eight bits of metallic sodium, no one of them larger than a hemp-seed.
The sodium dissolves with some spattering in the warm mercury, and
a sodium amalgam is thus obtained. Transfer the amalgam to a tall
glass or bottle of at least 300 c. c. capacity, and pour over it a concentrated solution of chloride of ammonium. The amalgam immediately
begins to swell up, and ultimately increases to 8 or 10 times its original
bulk, in the cold, or to 20 or 30 times if the solution be hot, assuming
a pasty consistency like that of soft butter, but preserving its metallic
lustre. It begins to undergo spontaneous decomposition as soon as it is
formed, and if it is placed in water, this decomposition is quite rapid;




438                HYDRATE OF AMMONIUM.
hydrogen gas is given off in minute bubbles, and ammonia is found in
the solution. This curious substance has been supposed to be a combination of ammonium (NH4) with mercury; all attempts, however,
to isolate the ammonium have been unsuccessful. The proportion
of ammonium (?) present in the amalgam is extremely minute, notwithstanding the great change of bulk and properties experienced by
the mercury. The amalgam is said to contain only 1 part of nitrogen
and hydrogen to 1800 parts of mercury.
520. Ammonium-salts are generally isomorphous with potassium-salts. They have mostly a pungent, saline taste; they are
colorless, like sodium- and potassium-salts, unless the acids are
colored; the carbonates, and those salts which, like the chloride
and iodide, contain no oxygen, are volatile at a moderate heat
without decomposition; some salts lose their ammonia when
heated; if the acid which neutralized this ammonia is a nonvolatile substance, like phosphoric acid, it will remain behind undecomposed; others, like the nitrate (Exp. 34), yield simpler
gases than ammonia, as, for example, nitrogen or nitrous oxide.
An aqueous solution of an ammonium-salt, when exposed to the
air or evaporated, generally loses ammonia and acquires an acid
reaction; hence in crystallizing an ammonium-salt, ammoniawater must be occasionally added during evaporation. All ammonium-salts, whether solid or in solution, evolve ammonia when
heated with the hydrates of sodium, potassium, calcium and a
few other metals (Exp. 48).
Exp. 260.-W- arm  a few centimetres of a solutioni of chloride of
ammonium in a test-tube, add a few drops of a solution of caustic soda,
and boil the liquid. The gaseous ammonia can be detected by its odor.
If in any case the ammonia evolved be in so small a quantity that its
characteristic smell cannot be detected, it may be recognized by its
property of restoring the blue color to reddened litmus-paper (~ 83),
and of forming white fumes by contact with a rod moistened with
somewhat dilute chlorhydric acid (Exp. 65). The reaction may be
formulated as follows:NH4Cl + NaHO = Na1Cl + NH, + H20.
521. The solution of ammonia gas in water (NH,,IT,O20) may be
regarded as a solution of hydrate of ammonium (NH1)!TO, comparable with the solution of caustic soda, NlaHO, or caustic potash,
KHO.  This solution produces, indeed, many of the effects which




CHLORIDE OF AMMONIUM.                  439
the solutions of the caustic alkalies produce. it neutralizes acids,
displaces the oxides of many metals from solutions of their salts,
and combines with fats to form a soap; it is, in short, a powerful
base.
Exp. 261.-Dissolve a small crystal of alum in 6-8 c. c. of water in
a test-tube, and add ammonia-water until the solution, after being
well shaken, smells strongly of ammonia. A gelatinous precipitate of
the hydrate of aluminum will appear in the liquid.
Expl. 262. —Dissolve about 1 grm. of sulphate of zinc in 6-8 c. c. of
water in a test-tube; add 4 or 5 drops of ammonia-water, and shake
up the contents of the tube. A white, translucent precipitate of the
hydrate of zinc will appear. Pour into the turbid liquid in the tube
3 or 4 c. c. more of ammonia-water; the precipitate will redissolve
and the liquid again become clear. The zinc is at first displaced from
its position in the sulphate by the group (NiH4); but the hydrate of
zinc thus precipitated is soluble in an excess of ammonia-water. The
hydrate of ammonium behaves in this way with not a few salts of
metals.
Ammonium-salts are very numerous; but only the few which
are of present importance in the useful arts will be here described.
522. Chloride of Ammnonium (NHIC).-This salt, commonly
called sal-ammoniac, is found native in many volcanic regions.
When nitrogenized animal matter and chloride of sodium  are
distilled together, this salt sublimes from the mixture; the commercial supply of the salt was formerly obtained from the soot
resulting from the incomplete combustion of camel's dung.
The raw material whence ammonium-salts are now manufactured
is derived from gasworks and boneblack-factories.  Coal and bones
contain a portion of nitrogen, which, during the process of distillation,
is partially converted into ammonia (~ 92); this ammonia combines
with the carbonic acid and sulphydric acid which are likewise products
of the distillation, and these compounds are condensed into a somewhat watery liquor, contaminated with tarry and oily matters, from
which the ammonium-salts are subsequently extracted. The impure
carbonate is converted into chloride by the addition of chlorhydric
acid, or of the mother-liquor from saltworks (a liquor containing the
chlorides of magnesium and calcium); on evaporating the clarified
solution, crystals of sal-ammoniac are obtained, but they are generally
too dirty for use. They are partly freed from tarry matters by heating
them to a temperature a little below their subliming-point, but high




4- O0              SULPHATE OF AMMONIUM.
enough to drive off the tar, and are then redissolved in water; this
solution, decolorized by being filtered through animal charcoal, is
recrystallized; these crystals are sometimes further purified by sublimation.
Chloride of ammonium serves for the preparation of ammonia
(Exp. 48), and of carbonate of ammonium. It is somewhat em.
ployed in dyeing, and also in certain processes with metals, such
as tinning, soldering, and silvering copper and brass, and galvanizing (zincing) iron.  The sublimed salt forms semitransparent,
tough, fibrous masses; it is very soluble in water, and a great
reduction of temperature occurs during its solution; hence it is
employed as a refrigerant. Its taste is sharp and acrid. When
heated, it sublimes much below redness, without undergoing
fusion.
Exp. 263.-Heat a bit of sal-ammoniac on a piece of porcelain, and
observe the low temperature at which the solid is completely converted into vapor.
Exp. 264.-Place a teaspoonful of powdered chloride of ammonium
in the hollow of the hand, and pour upon it two or three teaspoonfuls
of water. The cold produced by the solution of the salt will be very
perceptible.
523. Suph7ate of Ammonium ((NH4)2SO4). —If the ammoniacal
liquor from gasworks or animal-charcoal factories be neutralized
with sulphuric acid, or if it be decomposed by gypsum (sulphate
of calcium), the sulphate of ammonium will be obtained.  In the
latter case, the impure carbonate of ammonium in the liquor,
on being filtered through powdered gypsum, yields carbonate of
calcium and sulphate of ammonium.  Another recent mode of
utilizing the ammoniacal liquor of gasworks yields a crude sulphate of ammonium; the liquor is suffered to flow down the coketowers which are now often connected with sulphuric-acid chambers (~ 228), and there absorbs all the acid-fumes which escape
from the chambers. A crude chloride of ammonium may be prepared in a similar way, by substituting ammoniacal liquor for
water in the coke-towers of sulphate-of-sodium furnaces (~ 482).
The absorbent power of the ammoniacal liquor is, of course, much
greater than that of water.
Sulphate of ammonium is colorless, and has a very bitter taste;
it is soluble in twice its weight of cold, and in its own weight of




NITRATE OF AMMONIUM.                 441
boiling water. Its crystalline form is the same as that of sulphate of potassium, and the commercial article looks very much
like sand, just as the crystals of sulphate of potassium have a
superficial resemblance to quartz crystals. It forms a considerable number of double salts, which are isomorphous with the corresponding salts of potassium.  Sulphate of ammonium is employed in the manufacture of ammonium-alum, as an ingredient
of artificial manures, and as a source of other ammonium-salts.
524. INitrate of Ammonium ((NH,)NO,).-The method of preparing this salt, and its complete decomposition by heat, have
been already described (Exps. 33, 34, and ~ 9 ). The salt crystallizes in long needles; it has a pungent taste, is soluble in less
than half its weight of boiling water, and in dissolving produces
sharp cold. Between 230~ and 250~ it is decomposed into water
and nitrous oxide; if it be heated hotter, or too rapidly, ammonia,
nitric oxide, and nitrite of ammonium (NH,)N02 are also formed.
Nitrate of ammonium is formed by the action of dilute nitric acid
on several metals, especially tin.
525. Carbonates of Amnmonium.-Commercial carbonate of ammonium  (sal-volatile) is a white, semitransparent, fibrous substance, with a pungent taste and a strong ammoniacal smell; it
is prepared, on a large scale, by the dry distillation of bones,
horn, and other animal matters. The product is purified from
empyreumatic substances by repeated sublimation.
E2p. 265. —Mix thoroughly together 10 grms. of chloride of amllonium and 20 grins. of powdered chalk; heat the mixture in a small
evaporating-dish placed upon a sand-bath. When white vapors begin
to rise forom the hot mass, place a wide-mouthed bottle over the fumning
mixture. The white sublimate which collects in the bottle is a carbonate of ammonium; chloride of calcium remains in the dish.
This experiment illustrates a second, and very common, method'
of preparing the commercial carbonate, which simply consists in
heating to redness a mixture of 1 part of chloride, or sulphate, of
ammonium and 2 parts of carbonate of calcium.  When this
commercial carbonate is dissolved in strong ammonia-water at
about 30~, a solution is obtained which yields large, transparent,
prismatic crystals. These crystals, however, have no stability;
they are rapidly decomposed in the air, giving off water and




442                SULPHIDES OF AMMONIUM.
ammonia. " Sesquicarbonate of ammonia" is the name generally
applied to this substance-a name deduced from the dualistic
formula 2(NH,)20,3C02.  The commercial carbonate approximates to the composition represented by this formula; but it is
an impure product, and, when exposed to the air, changes gradually into a more stable compound, the carbonate of ammonium
and hydrogen, or " bicarbonate" (NH4)HCO3.
This bicarbonate may be obtained by saturating the solution of
ammonia, or sesquicarbonate of ammonium, with carbonic acid;
it forms large, transparent, prismatic crystals. When exposed to
the air, it slowly volatilizes, giving off a slight ammoniacal odor.
At the ordinary temperature, it is soluble in 8 parts of water;
this solution, if heated above 36~, evolves carbonic acid. Even
at the ordinary temperature, the solution gradually becomes
ammoniacal on keeping. White crystalline masses of this bicarbonate have been found in guano deposits. It seems to be the
most stable of the carbonates of ammonium; for the other carboiiates change into it if left to themselves.
526. Sulphides of Ammnonium.-At a temperature of -18~,
two volumes of ammonia-gas combine with one volume of sulphydric acid gas to form a crystalline unstable substance, of
strong alkaline reaction, which corresponds in composition with
the sulphides Na2S and K2S.
2NH,3 +  H2S =  (NH4)2S.
BEX1. 266. —Pass a slow stream of washed sulphydric acid throutgh
100 c. c. of strong ammonia-water, until the solution has a predominating and persistent odor of sulphuretted hydrogen. This solution is
at first colorless, and contains a sulphide of ammonium and hydrogen
(NH4)HS; but when kept in contact with air it becomes yellow,
owing' to the formation of a higher sulphide of ammonium. The sollution has the property of dissolving many of the sulphides of the metals,
by forming with theni double sulphides, and is a very useful reagent
in the analytical laboratory.
The higher sulphides of ammonium  are obscure bodies, to
which the following formula have been assigned,-(NI4)2S2,
(N14)2S3, (NH4)284, (NH4)2S5, (N1H4)2S,.  With the exception of
the last, the septisulphide, all these sulphides are soluble in
water. With the same exception, they correspond with the sul



LITHIUM.                       443
phides of sodium and potassium.  They are in general unstable;
the most permanent is the septisulphide, which forms ruby-red
crystals, capable of resisting temperatures below 300~, and only
slowly decomposable by water and chlorhydric acid.
CHAPTER XXVI.
L I T H I U M-R U B I D I U M-C. SI U M-T H A L L I U M.
SPECTRUM ANALYSIS.
527. Lithium (Li). —This rare metal occurs as a constituent
of not a few minerals, especially micas and feldspars, but does
not form a large percentage of any of them. The minerals lepidolite, triphylline, and petalite usually contain from 3'5 to 5 per
cent. of lithium, and are the chief sources of the element.  In
much smaller proportion, it has been recognized in sea-water,
mineral waters, and almost all spring-waters, in milk, tobacco,
and human blood. It is therefore a widely diffused, but not
abundant, substance. The metal, which is obtained by decomposing the fused chloride by the galvanic current, has the color
and lustre of silver on a freshly cut surface, but quickly tarnishes
on exposure to the air. It is harder and less fusible than sodium
and potassium, but softer than lead; it may be welded, by pressure, at ordinary temperatures. It floats on naphtha, and is the
lightest of all known solids which include no air, its specific
gravity being only 0'59. The atomic weight of the element is
also low, namely 7.
In its chemical relations, lithium closely resembles sodium and
potassium, but is somewhat less energetic; it combines with the
same elements to form analogous compounds to those of sodium
and potassium; but the properties of these compounds, while presenting a striking general resemblance to those of the sodium and
potassium compounds, nevertheless offer some special points of
divergence from them. Thus, the hydrate of lithium (LiHO)




444                  SPECTRUM ANALYSIS.
has the same taste, causticity, and alkalinity as the hydrates of
sodium and potassium, but is much less soluble in water. The
fused hydrate attacks platinum  more energetically than caustic
soda and potash do. The carbonate and phosphate of lithium are
rather sparingly soluble in water, while the corresponding sodium
and potassium-salts are exceedingly soluble.  The chloride of
lithium (LiC1), produced when lithium burns in chlorine, or
when the hydrate or carbonate of lithium is dissolved in chlorhydric acid, crystallizes in cubes, and has the taste of common
salt; but it is more volatile than the chloride of potassium, and it
deliquesces faster than any other known salt, whereas the chlorides of sodium and potassium are almost permanent when pure.
All the volatile lithium-compounds color a gas-, alcohol-, or blowpipe-flame carmine-red. The most delicate reaction for the detection of lithium, the test which has revealed its existence in a
great variety of substances which were never imagined to contain it, is the presence of one bright line, of a peculiar red, in the
spectrum  seen on looking through a glass prism  at a flame
colored with a lithium compound.
528. Spectrum Analysis.-We have had occasion to observe
that certain chemical substances, like boracic acid and salts of
sodium, potassium, and lithium, impart peculiar colors to the
blowpipe-flame, or to any other hot and colorless flame.  If
these colored flames are looked at through a prism, a narrow
pencil of the colored light being directed through a slit upon
the prism, it will be seen that each different flame produces a
peculiar spectrum, consisting.of one or more distinct bright lines
of colored light, and bearing no resemblance to the continuous
band of rainbow-colors which constitutes the common spectrum
produced by a pencil from any source of white light. Thus the
spectrum of the yellow sodium-flame consists of a single, bright,
yellow line; the purple potassium-flame gives a spectrum containing two bright lines, one lying at the extreme red and the
other at the extreme violet end, and a second, fainter red line;
while the lithium spectrum consists of a very characteristic red
line and a fainter orange line.
These peculiar lines which characterize the spectrum of any
element are invariably produced by that element, and never by




SPECTRUM ANALYSIS.                 445
any other substance, and not only the color and number of lines,
but theirposition in the normal spectrum always remains unaltered.
When the spectrum of a flame colored with a mixture of sodium
and potassium  salts is examined, the yellow line of sodium is
seen in its place, and the red and purple lines of potassium are
as visible in their respective positions as if no sodium had been
present. This example illustrates one great advantage which
the use of the prism gives,-the unaided eye cannot distinguish
the potassium color in the presence of the intense sodiumyellow, the brighter color hiding the paler; but with the prism
it is easy to detect each of several ingredients of a mixture by
the appearance of its characteristic lines.  Now, every elementary substance, whether metallic or non-metallic, solid, liquid, or
gaseous, when heated to the point at which its vapor becomes
luminous, emits a peculiar light produced by it alone, and the
bright lines of the spectrum of this light are characteristic of this
element in number, color, and position.  Many metals require
a much higher temperature than that of the common gas-flame
to convert them into luminous vapors; but by the use of the
electric lamp all the metals, even gold, silver, and platinum, may
be made to yield peculiar spectra. The permanent gases also
give characteristic spectra when they are heated by the passage
of the electric spark; the spectrum of hydrogen, for example,
consists of one red, one green, and one blue line.
A new method of analysis, of extreme delicacy, is based upon
these facts.  Spectrum  analysis is competent to detect the
I'1i    of a gramme of sodium, or the   I-,-T —   of a
gramme of lithium, and many other elements in incredibly small
proportions. So extreme is the delicacy of the method that it
brings into plain sight minute quantities which altogether escape
the coarser process of analysis, and reveals, as substances common in familiar things, elements which were long supposed to be
of extreme rarity. Thus the presence of lithium, formerly considered a rare element peculiar to a few obscure minerals, has
been demonstrated by spectrum analysis in many drinking-waters,
in tea, tobacco, milk, and blood. A still more striking illustration of the value of spectrum analysis is to be found in the discovery of four new elementary bodies by its means.




446                RUBIDI7L.I AND CAESIUM.
529. Two new elements which closely resemble sodium and
potassium, and are in nature associated with these alkali-metals,
have been found in certain mineral waters, and in the mineral
lepidolite. One of these elements gives a spectrum containing,
among others of less mark, a superb, double, red line, and has
thence been called Rubidium; the other produces a spectrum
characterized by two beautiful blue lines, and has thence been
called Ccesium. These two new metals resemble potassium  so
closely in all their chemical properties, that it would have been
nearly impossible to detect them by the common analytical processes; yet their spectra are in the highest degree characteristic,
exhibiting bright bands which exist neither in the potassium
spectrum nor in any other known spectrum. The recently discovered metal Thallium was discovered and traced to its source
in certain kinds of pyrites by observing a splendid green line
which did not belong to any known substance.  The new metal
Indium was also detected, traced to its source in certain zinc
ores, and successfully isolated, by the help of a dark-blue line
which had not been previously observed.
The methods and processes of spectrum analysis are not applicable to colored artificial lights alone; they have been applied
with encouraging success to the lights of various quality which
emanate from the sun, the stars, and the nebula; but the details
of these observations belong rather to physics than to chemistry.
530. Rubidium and Ccesium (Rb and Cs).-These two elements
are always found together, and in association with potassium.
Though extensively diffused, they generally occur in very minute
quantities. Rubidium seems to be rather the most abundant.
Ten kilogrammes of the mineral water in which these metals
were first discovered yield not quite two milligrammes of chloride
of coesium, and about two and a half milligrammes of chloride of
rubidium. Since the original discovery of the elements, they
have been found in many other springs, in several kinds of mica
and in other silicates, and in the ashes of beet-root, tobacco, coffee,
and grapes. To separate the metals from potassium the analyst
relies on the greater insolubility of the double chlorides which
they form with platinum; if a mixture of the chlorides of potassium, rubidium, and csesium be completely precipitated by chlo



THALLIUM.                       447
ride of platinum, and the yellow precipitate be repeatedly treated
with boiling water, the insoluble residue will contain the new
metals.  The csesium is separated from the rubidium by converting a mixture of their carbonates into their tartrates, the rubidium  into the acid, or bitartrate, the cesium into the neutral
tartrate, and then exposing this mixture to very moist air. The
neutral tartrate of cresium deliquesces, the acid tartrate of rubidiumn remains solid, and the two salts are separated by filtration.
Most of the salts of rubidium and cesium are isomorphous with
the corresponding potassium-salts. Their hydrates, RbHO and
CslHO, are caustic alkalies, soluble in water and alcohol.  Their
carbonates are fusible, deliquescent, and strongly alkaline; the
nitrates (RbNO3 and CsNO,) and sulphates (ARb2SO, and Cs2SO4)
are anhydrous crystalline salts, soluble in water; the sulphates
form alums with sulphate of aluminum. The chloride of coesium,
CsC1, is a deliquescent salt, like chloride of lithium; the chloride
of rubidium, RbC1, is permanent, like the chlorides of sodium and
potassium. The fused chlorides are easily decomposed by the
galvanic current.
The metal rubidium  is white, and has the brilliant lustre of
silver, but it rapidly oxidizes in the air; its specific gravity is
1'52, and its atomic weight 85-7.  It may be prepared either by
the electrolysis of its chloride, or, like potassium, by the reduction of its carbonate by ignition with carbon and chalk.
The properties of caesium have only been studied in the amalgam with mercury resulting from the electrolysis of its chloride;
the metal itself has not been isolated.  Its atomic weight, deduced from the analysis of its chloride, is 133.  There can be no
question that the properties of both rubidium and caesium differ
from those of sodium and potassium not in kind, but only in
degree.  They are therefore classed with sodium and potassium
as alkali-metals.
531. Thalliuon (T1). —This metal was discovered, by means of
spectrum analysis, in Lipari sulphur and in the deposit in the flue
of a pyrites-burner-a furnace in which iron pyrites are roasted
for the sake of the sulphurous acid they yield. The element is
found to occur in not inconsiderable quantities in many specimens
of iron pyrites, and appears to take the place of arsenic, which is




448                        SILVER.
a common impurity of this mineral. Thallium presents the external characters of lead; it is heavier than lead, having a specific
gravity of 11'85, and it is so soft that the thumb-nail can indent
it; it is very malleable, and ductile enough to be drawn into
wire; it melts at 200~ and volatilizes at redness; its freshly cut
surface has a bluish-white lustre; but it quickly tarnishes and is
gradually oxidized in the air, so that it is best preserved under
water.  Water is not decomposed by it even at 100~.  When
strongly heated in oxygen, it takes fire and burns with a bright
green flame.
Thallium dissolves in dilute acids, with evolution of hydrogen.
There are several oxides of this metal, of which the most important is the oxide T1 0, corresponding in composition, and, to some
extent, in properties, with the oxide of sodium Na2O. This oxide
is somewhat soluble in water, and yields a caustic alkaline solution, which absorbs carbonic acid from the air, and forms a welldefined series of salts. The sulphate, T12SO4, is a soluble salt,
which forms an alum with sulphate of aluminum; the chloride,
TlC1, is only slightly soluble in water, resembling, in this respect,
the chloride of lead, and being quite unlike the soluble chlorides
of sodium, potassium, rubidium, and cesium.  The carbonate of
thallium is a soluble salt; but the sulphide of thallium, T12S, is
an insoluble blatrk powder, which resembles the sulphide of lead,
but is entirely unlike the sulphides of the alkali-metals. The
soluble salts of thallium are very poisonous. In general, the
properties of thallium are intermediate between those of lead and
those of sodium and potassium. Like the alkali-metals, it replaces
hydrogen atom for atom; its atomic weight is 204.
CHAPTER XXVII.
SILVER-THE A L KALI-METAL — QUANTIVAENCE.
532. Silver is a widely diffused and quite abundant element,
but in its mode of occurrence it differs widely from the alkalimetals which we have just been studying.  In the first place, it




OCCURtENCE OF SILVER.                449
frequently occurs native, both pure and alloyed with meerrury,
copper, and gold, — a mode of occurrence quite impossible for the
alkali-metals, because of their readiness to combine with the elements of air and water. Native silver is found in various forms,
sometimes crystallized in cubes or octahedrons, sometimes in
filaments, both coarse and fine, and sometimes in shapeless masses.
The metal more commonly occurs in combination with sulphur,
mixed with sulphides of lead, antimony, copper, and iron.  It is
from argentiferous sulphides that the larger part of the silver of
commerce is extracted; among ores of this kind the argentiferous sulphide of lead (galena) is the most abundant. Combinations of silver with selenium, tellurium, chlorine, bromine, and
iodine are also to be enumerated among silver-containing minerals;
of these the chloride (horn-silver) occurs in quantities large
enough to make it valuable as an ore of the metal.  It is noticeable that the only elements which are extracted in any quantity
from their chlorides as ores are sodium, potassium, and silver.
The chlorides of copper, mercury, and lead do, indeed, occur as
natural minerals; but as sources of those metals they have no
significance. A small proportion of silver exists in sea-water
(about 1 milligramme in 100 litres), and its presence has been
recognized in common salt, in chemical products in the making
of which salt is used, in various sea-weeds, in the ashes of landplants. in the ash of ox-blood, and probably also in coal.  In seawater it exists, as sodium and potassium do, in the form of chloride.
When silver is extracted from argentiferous sulphide of lead, the ore
is first treated for lead, precisely as it would be if it contained no silver.
The lead, so reduced, contains all the silver originally present in the
quantity of ore treated. The subsequent separation of the metallic
silver from the metallic lead depends upon the chemical properties of
lead rather than of silver, for the silver remains unaltered during the
whole process; this separation will therefore be described in the next
chapter.
The mixed sulphides which contain silver have been heretofore generally reduced by a complicated process which depends ultimately on
an amalgamation of the silver with mercury. The ore, after thorough
washing and grinding, is mixed with a portion of common salt, and
roasted for several hours; during this roasting, white fumes of arsenious
und antimonious acids are:expelled,. the, sulphides of copper and, iron




450                  EXTRACTION CF SILVER.
are partially converted into oxides, chlorides, and sliphates, and chloride of silver and sulphate of sodium are formed. The roasted product
is then reduced to a very fine powder, and agitated in revolving casks
with water and iron filings, or scraps, to which mercury is soon added.
This operation lasts about 20 hours; during it, the iron decomposes the
chloride of silver, and the mercury dissolves the silver to an amalgam;
from this amalgam the excess of mercury is first squeezed out through
leather or cloth filters, and the remainder is driven off by distillation.
The residual spongy mass is silver, alloyed with a variable proportion
of copper, derived from the ore and reduced to the metallic state by
the same steps which have reduced the silver.
This process is European; the process of amalgamation as practised
in Mlexico and South America is quite different, and the reactions which
it depends upon are somewhat obscure. The ore is not roasted, but,
after being ground to powder, moistened with water, and mixed with
from 1 to 5 per cent. of common salt, it is suffered to lie undisturbed
for some days. From 4 to I per cent. of roasted copper pyrites is then
added, together with a considerable proportion of metallic mercury,
and the mass is worked together and commingled by the trampling of
mules or horses. After an interval of two or three weeks, a second
dose of mercury is given, and after a still longer interval a third. By
this last addition, a fluid amalgam is obtained, which is separated by
washing and filtering, and distilled for the recovery of a portion of the
mercury employed, and the isolation of the silver. In this process
there is a great waste of mercury, because much of it is cbnverted into
a chloride of mercury (calomel) and lost. The recommendations of
the process are mainly these-that it requires no fuel, except for the
distillation of the amalgam, and that it leaves the silver in a condition
of great purity. The whole process, though far from economical from
the point of view of the theoretical chemist, was doubtless a legitimate
outgrowth of the conditions under which it took birth.
Various processes have been patented for the extraction of silver
without the use of the costly mercury, some of which have been successfully practised on a large scale. They depend, for the most part,
either on the comparative stability, in the fire, of sulphate of silver
when once formed, as compared with the sulphates of iron and copper,
and the consequent possibility of dissolving sulphate of silver out of the
roasted ore, or upon the fact that the chloride of silver, which results
from the roasting of the ore with chloride of sodium, may be dissolved
in solutions of the alkaline chlorides, and, indeed, in aqueous solutions
of a great many other soluble salts, though it is by itself insoluble in
water. Any.aqueous solution containing, among other things, a silver;




THE TERM METAL.                   451
salt (whether in the condition of chloride, sulphate, or nitrate is indifferent) may be decomposed by digestion with metallic copper; the
silver-salt will be decomposed, the corresponding copper-salt formed
and dissolved, and the metallic silver will be precipitated.
533. Silver (Ag).-The element, silver, is much more familiarly
known than any of its compounds; known from the earliest ages,
this metal has always been prized as much for its beauty as for
its rarity. White, brilliantly lustrous, susceptible of an admirable
polish, wonderfully malleable and ductile, the best known conductor of heat and electricity, fusible only at a very elevated
temperature and permanent in the air, whether hot or cold, wet
or dry, it represents and embodies in the completest sense all that
is commonly understood by the term metal.
This word metal cannot be strictly defined; it is a conventional
term, vaguely used because expressing a vague idea. Thus
metals would all be solid were not mercury, and perhaps cersium,
fluid; they are generally heavy; but lithium, sodium, and potassium float upon water; they have all a peculiar lustre, called
metallic; but this lustre does not characterize metals alone, for
coke and graphite, galena, molybdenite, and many other minerals
often exhibit a similar lustre; they may all be said to be opaque;
but gold may be beaten out so thin as to transmit a greenish
light. While it is not possible to define the term metal with
precision from chemical any more than from physical properties,
one general chemical fact deserves attention in this connexion.
We have seen that bodies which contain a large proportion of
oxygen, such as SO3, P2O,, N2O,, and CO,, have a common tendency to unite with other bodies which are alike in that they
contain a much smaller proportion of oxygen, such as K,,O, NaO,
PbO, and CaO, to form more or less stable saline substances. The
first class of bodies, which are usually rich in oxygen, have been
balled acids; the second class, which are usually poor in oxygen,
have been designated collectively as bases. Now those elements
which unite with oxygen to form acids alone are, as a rule, nonmetallic, and those elements which unite with oxygen to form
bases are, in the chemical sense of the term, the metals; but no
sharp line of division between metallic and non-metallic elements
can be established on this principle, inasmuch as some elements
2o2




452                 PROPERTIES OF SILVER.
which possess the other characteristics of a metal form no basic
oxide, while some metals, like antimony, form oxides which are
at one time bases and at another time acids. The metal arsenic,
for example, forms no basic oxide; and we shall hereafter meet
with another illustration of the same difficulty of classification, in
the element tungsten.
Melted silver possesses the curious property of absorbing a
large volume of oxygen (twenty-two times its bulk), from the
air, while it is liquid. This gas it gives out again on solidifying.
When a globule of molten silver is cooled suddenly, the film of
solid metal which forms upon its surface is burst open by the
escaping gas, and the liquid silver within is apt to be projected
outwards with the gas; this phenomenon is called spitting; it
often occasions a loss in silver assays. When heated on lime
before the oxyhydrogen blowpipe, silver gives off vapors which
become oxidized if the blast of gas contain an excess of oxygen;
a fine silver wire is dispersed in greenish vapors when a very
powerful electric discharge'is sent through it. Silver combines
slowly with chlorine, bromine, and iodine, and promptly with
sulphur. The tarnishing of silver is due to the formation of a
thin film of the black sulphide over the metallic surface, by combination between the silver and the sulphur of the sulphuretted
hydrogen which is often present in the air of towlls and houses.
The best solvent for silver is nitric acid diluted with two or
three parts of water; nitric oxide is evolved, and nitrate of silver
remains in solution:3Ag +  4HN03 -  3AgNO  +  NO  +  2H20.
Chlorhydric acid acts upon it but slowly; for the chloride of silver
is but slightly soluble in chlorhydric acid, whether strong or
dilute. Boiling sulphuric acid dissolves silver, and forms the
sulphate, sulphurous acid being evolved during the reaction:2Ag  +  2H2S04 =  Ag2SO4 +  2H20 +  S02.
Neither the alkalies nor their nitrates have much effect on silver,
whether they are in solution or are fused by heat; hence a silver
dish is used in concentrating the caustic alkalies, and a silver
crucible for fusing refractory minerals with the hydrate of sodium
or potassium. The specific gravity of silver is 10'5, and its atomic
weight 108.




SILVER COIN.                     453
534. The physical and chemical qualities of silver fit it to serve
as a medium of exchange, and as the material of jewellery and
plate. But as the pure metal would be rather too soft for ordinary use, it is hardened by combining with it a small proportion
of copper.  The proportion of copper in the " standard" silver
employed for coinage varies in different countries. In the United
States and in France it is 10 per cent.; in Great Britain it is 7'5
per cent.; in Germany it is 25 per cent.
Exp. 267.-Place one or two dimes in a small flask, and cover them
with nitric acid diluted with two parts of water. Warm the flask
gently in a place where there is a good draught of air; the coins will
gradually dissolve, with evolution of nitric oxide, which, on contact
with the air, produces the abundant red filmes which escape fiom the
flask; add more nitric acid, from time to time, if necessary to complete the solution. The blue solution contains both the silver and the
copper dissolved in nitric acid.
Place in the blue solution two or three copper cents, and leave the
flask at rest for some days in a warm place. Then collect the little
plates of pute silver, which have separated from the solution, upon a
filter, and wash them, first with water, and then with ammonia-water,
until the ammonia-water no longer shows any tinge of blue. This
silver, washed finally with water and dried, is well nigh pure; twothirds of it may be again dissolved in nitric acid; the solution will
contain pure nitrate of silver.
535. NTitrate of Silver (AgNO.).-This salt, as we have already
seen, is obtained in solution by dissolving silver in nitric acid.
When such a solution is evaporated to the point of crystallization,
the nitrate is obtained in transparent, anhydrous, tabular crystals,
which are soluble in their own weight of cold water, and in half
their weight of hot water. The salt fuses easily, and when cast
into cylindrical sticks is used in surgery as a caustic, under the
name of lunar caustic.
Nitrate of silver, when pure, is not altered by exposure to sunlight; but if it be in contact with organic matter, light readily
decomposes it. and a black, insoluble product is formed of no
ordinary stability.  Hence the solution of the nitrate stains the
skin black, and the salt forms the basis of an indelible ink used
for marking linen and other fabrics.
Exp. 268.-Dissolve 8 grins. of crystallized carbonate of sodium and




454                    NITRATE OF SILVER.
1 grm. of gum-arabic in 16 c. c. of hot water. Moisten a bit of linen
or cotton-cloth with this preparatory solution, dry it, and press it
smooth with a hot iron.
Dissolve 1 grm. of nitrate of silver and 0'1 grm. of gum-arabic in
1-75 c. c. of water, previously colored with India-ink.
Write with this silver solution upon the prepared surface of cloth,
and expose the writing to the direct rays of the sun for a few hours.
Then wash out the gum and carbonate of sodium with water; a vert
durable mark, which neither soap nor "soda" will obliterate, will
remain on the cloth.
Nitrate of silver is even more completely decomposed by a red
heat than nitrate of potassium, for nothing but metallic silver
remains behind; in decomposing, it gives up a large quantity of
oxygen; hence mixtures of combustibles, like sulphur, phosphorus, and charcoal, with nitrate of silver, detonate, explode, or
deflagrate when struck sharply with a hammer or touched with
a hot wire. (Compare ~~ 515, 516.)  Phosphorus, mercury, charcoal, grape-sugar, certain essential oils, and many other organic
substances reduce metallic silver from solutions o~f nitrate of
silver.  Nitrate of silver is the material from which most other
silver compounds are artificially prepared. It is largely consumed in photography.
The precipitation of metallic silver in a beautiful arborescent form
is accomplished as follows:-Dissolve 2 grins. of nitrate of silver in
60 c. c. of water, and place the solution in a test-glass; pour 2 grms.
of mercury into the liquid, and let the glass stand at rest for several
hours. The nitrate of silver may be recovered by dissolving the precipitated silver in nitric acid, and evaporating the solution.
To illustrate the decomposition of a silver solution by an organic
substance, dissolve 2 grms. of nitrate of silver in 60 c. c. of water, and
immerse in the solution a horn or ivory paper-knife, which has been
cleansed from grease with ammonia-water and rinsed in fresh water.
Let the knife remain in the solution about an hour; it will turn yellow; take it out, rinse it in water, and expose it to the direct rays of
the sun until it turns jet black; then burnish it with a piece of leather,
and the silver will appear in the metallic state.
Exp. 269.-Wrap a piece of phosphorus, no bigger than a pin's
head, with a small crystal of nitrate of silver, in a bit of paper; place
the packet on an anvil and strike it with a hammer. The explosion
will be sharp. The student will remember that nitrate of silver stains
the fingers.




OXIDES OF SILVER- FULMINATING SILVER.        455
Ex. 270.-Mix I grm. of powdered nitrate of silver with 0'2 grm.
of dry, powdered charcoal; place the mixture on a piece of porcelain,
and touch it with a red-hot wire. The mixture deflagrates, and there
remains behind metallic silver.
Exp. 271.-Add to a solution of nitrate of silver a solution of caustic
soda, until no further precipitate is produced. The brownish precipitate is a hydrated oxide of silver.
536. Oxides of Silver.-Silver probably forms three oxides,
Ag4O, AgO, and Ag202.  The first is a very unstable black
powder; the second forms with acids the ordinary silver salts;
the third is a crystalline body obtained by electrolysis of nitrate
of silver. The precipitate obtained in the last experiment is
the hydrate of the oxide Ag2O; this precipitate readily parts
with its water, and at a temperature much below 1000 becomes
anhydrous.  Unlike the oxides of sodium  and potassium, this
oxide of silver yields up its oxygen below a red heat, and
metallic silver remains-as may be demonstrated by heating the
product of the last experiment; light also reduces it, and hydrogen even at 100~ has the same effect. Oxide of silver bears,
however, certain striking resemblances to the oxides of the
alkali-metals; thus it is a strong base, uniting with strong acids
to form salts which are neutral to test-paper, and which are in
some cases isomorphous with the corresponding salts of sodium.
The oxide is slightly soluble in water, and the solution has a
feeble alkaline reaction.
The oxide is freely soluble in ammonia-water, and the solution
deposits, on exposure to the air, a black, micaceous powder which
has received the name offulminating silver, because of its explosive character. The same dangerous compound is formed when
freshly precipitated oxide of silver is digested for some hours in
ammonia-water, and it is also produced when an ammoniacal
solution of chloride or nitrate of silver is precipitated with a
solution of hydrate of sodium or potassium. It is necessary to
be aware of these facts in order to avoid the risk of producing by
accident this exceedingly dangerous substance. Its composition
is not accurately known. Friction or slight pressure, even under
water, may cause it to explode.  The student should never
venture to prepare this substance.




456             PIHOTOGRA.PrY-DAPG UERRFEOTY PE.
Exp. 272.-Fill three test-tubes one-third full of water, and pour
into each a few drops of a moderately strong solution of nitrate of
silver. Add to the first test-tube 2 or 3 c. c. of a solution of chloride
of sodium, and shake the tube violently; a dense, white, curdy precipitate of the chloride of silver will be produced. Add to the second
test-tube 2 or 3 c. c. of a solution of bromide of potassium, and shake
the tube; a yellowish precipitate of bromide of silver will be thrown
down. Add to the third test-tube 1 or 2 c. c. of a solution of iodide
of potassium, and shake up the liquid; a pale-yellow flocculent deposit
of iodide of silver will be formed.
Withdraw  fiomn each test-tube a portion of the precipitate it contains, and try to dissolve each precipitate in moderately strong nitric
acid; the attempt will fail, for these silver salts are insoluble in nitric
acid.
Withdraw from each test-tube another portion of the precipitate it
contains, and treat each precipitate with ammonia-water; the chloride of silver will dissolve easily, the bromide less easily, the iodide
with difficulty. Lastly, pour upon the remnants of the original precipitates in the three test-tubes a moderately strong solution of hyposulphite of sodium (~ 495); all three precipitates will immediately
dissolve.
Sap. 273.-Precipitate some curdy chloride of silver by adding
chloride-of-sodium solution, or chlorhydric acid, to a solution of nitrate
of silver, so long as any precipitate is produced. Throw the precipitate upon a filter, and wash it with water; then open the filter, spread
the chloride evenly over it, and place it in direct sunlight. The white
precipitate rapidly changes to violet on exposure to the sun's rays, the
depth of shade increasing as the action of the light continues. This
coloration arises from a partial decomposition of the chloride of silver,
the change of color being accompanied; by a loss of chlorine. Upon
the facts illustrated in this and the preceding experiments the main
processes of photography depend.
537. Photograjuphy.-The chemical changes which the salts of
silver undergo, whcn exposed to light, are the basis of the art of
photography-not because these are the only salts which are
affected by light, but because none are so advantageous on the
whole. There are three different kinds of photographic process
-that on silver, that on glass, and that on paper.
To produce a photograph on silver (a daguerreotype), a highly
polished silver plate is exposed in a dark box to the diluted
vapor of a mixture of bromine and iodine.  A bronze-yellow film




PHOTOGRAPHY ON GLASS.                457
of brom-iodide of silver is thus produced upon the plate, which,
at a certain stage, possesses a high degree of sensitiveness to
light. The plate is then transferred to a camera, and exposed
at the focus of the lens to the light radiated from the object to be
copied. After remaining a few seconds in the camera, it is withdrawn, and immediately exposed in a warm box to the vapor of
metallic mercury.  When the plate is taken from the camera, the
film looks as uniform as ever, and no image is visible upon it;
but the exposure to mercury vapor immediately brings out an
image. The mercury fixes itself strongly upon those parts which
have received the light, while it talkes no hold upon those parts
of the film which the light has not decomposed. A strong solution of hyposulphite of sodium is then poured over the plate, in
order to dissolve off the undecomposed brom-iodide. The highly
polished silver, beneath, forms the shades, and the amalgam  of
mercury with silver forms the lights. The plate is washed, and
a very dilute solution of chloride of gold in hyposulphite of sodium
is poured over its surface and gently warmed. A thin film of
gold, which, as it were. varnishes the picture, is thus deposited
upon the plate; another washing completes the operation. The
daguerreotype is the most perfect of photographs; but the polish
of the surface prevents the image from being seen in all lights,
and the plate is liable to be tarnished and ruined by sulphuretted
gases.
In order to get a photograph upon glass, a transparent film
capable of holding the necessary silver-salt must first be attached
to the glass plate. Collodion (a solution of a variety of guncotton in a mixture of alcohol and ether) is the material of this
film. To the collodion is added a solution of an iodide, either of
potassium, cadmium, or ammonium, or a mixture of these; the
bromides of ammonium and cadmium, or one of them, added
in the proportion of one part of the bromides to three or four of
the iodides, render the film more sensitive to yellow and red rays
-a point of importance in cloudy weather or smoky towns. The
collodion thus prepared is poured rapidly over a clean and dry
surface of plate-glass; the volatile solvents evaporate rapidly,
and as soon as the film is coherent the glass is plunged into a
bath of nitrate of silver very slightly acidified with acetic or dilute




458                    PHOTOGRAPHY.
nitric acid. This bath must be in a dark place; the plate remains in it for several minutes. A yellow layer of iodide oI
brom-iodide of silver is produced in the film, and nitrate of potassium, cadmium, or ammonium dissolves in the bath. The plate
is then exposed in the camera for a few seconds. When removed
no image is perceptible; but on pouring over the film a solution of
gallic or pyrogallic acid in alcohol and acetic acid, or a solution
of the green sulphate of iron, mixed with a few drops of a weak
solution of nitrate of silver, the image will be developed, slowly or
rapidly, according to the nature and strength of the developingliquid, the degree of exposure, and the intensity of the light. The
illuminated portions of the picture will appear, under the action
of the developer, more or less black, while the shaded portions
will retain the yellow colour of the iodide. As soon as the details of the shaded portions appear, the liquid is washed off and the
development arrested. A saturated solution of hyposulphite of
sodium is then poured over the film to dissolve off the yellow iodide of silver where it has not been affected by the light; only
the reduced portions of silver remain, and they appear more or less
opaque. The plate must finally be very thoroughly washed to
remove all traces of the hyposulphite, and then dried and varnished on the collodion side to protect the film from injury.
Concerning the nature of the change which a film of iodide of
silver undergoes when exposed to light, we cannot be said to have
dny exact knowledge. There is no perceptible alteration in the
film; there is no loss of iodine; the iodide retains its solubility
in hyposulphite of sodium; yet the impression is not of a temporary kind; for the invisible image produced on a plate may be
developed many hours afterwards, if the plate is kept in the dark
during the interval.
The photograph on collodion may be employed directly as a
positive picture, if not too strongly developed, by placing it on a
black background. Those portions which are opaque to light, or
in other words those in which silver is deposited, will reflect light,
and furnish the lights of the picture; while those on which the
light did not act, and which are therefore transparent, will appear
black from the nature of the background, and these will form the
shades of the picture. In the daguerreotype the finished picture




PHOTOGRAPHY ON PAPERL                 459
is inverted; in the collodion positive it is not inverted. If the
development be pushed further, the image becomes so strongly defined that the deposited silver will more or less completely intercept the light. The collodion side of the plate is then placed in
contact with the sensitive side of paper impregnated with chloride
of silver by a process immediately to be described, and exposed
to light in a pressure-frame. The light is arrested by the altered
parts of the collodion, but is freely transmitted by the other portions; upon the paper, therefore, the lights of the real object are
light and the shades are dark. Such a negative collodion picture
may of course be copied on a second sensitive collodion film.
Two developing-solutions, used one after the other, produce a
better effect than one. The green sulphate of iron may be used
first, and pyrogallic acid with nitrate of silver subsequently;
the iron solution must be completely washed off before the other
is added.  The picture may even be intensified by pyrogallic acid
after the plate has been washed in hyposulphite of sodium.
Photographs were made on paper long before the film on glass came
into use; bat the paper process is now chiefly confined to the printing
of positive impressions from collodion negatives on glass. The silversalt which is preferred for photographic paper is the chloride, with or
without albumen, but always accompanied with free nitrate of silver.
The paper is floated for five minutes on a solution of chloride of sodium
or ammonium; when dried, it is floated in a dark room, for five minutes, on its salted surface, in a solution of nitrate of silver; again
dried, it is fit for use. When such paper is used to obtain a positive
impression from a collodion negative, or from a paper negative made
transparent with wax or a mixture of wax and paraffine, it is exposed
to lighf under the negative to be copied, until the lights of the picture
are of a pale lilac hue, and the shades of a deep bronze color.  After
being thoroughly washed, the paper is transferred to a " toning "-bath,
which consists of a very dilute solution of bicarbonate of sodium, with
a minute proportion of chloride of gold. The picture is kept in motion while in this bath; it remains there until its shades have acquired
a deep purple-black color. It is only in those parts of the picture in
which the silver has been well reduced that this toning effect is proiuced. The picture is again washed in water, and soaked for fifteen
minutes in a solution of hyposulphite of sodium, in order to remove
all the chloride of silver which is contained in the substance of the
paper. Finally the picture must be soaked for twenty-four hours in




460                 CHLORIDE OF SILVR.water which is constantly renewed, in order to wash away every trace
of the compound hyposulphite of sodium and silver. No photograph
will keep long, unless the chloride of silver has been completely dissolved by the hyposulphite, and the compound hyposulphite washed
away with a thoroughness that leaves no trace behind. If the first
condition is not fulfilled, diffused daylight will alter the picture; if
the second condition is not complied with, yellow or brown stains will
ultimately destroy the picture.
As in every other art which embraces many details, and demands a
trained eye and hand, eminent skill in photography can, as a rule, be
acquired only by long practice.
538. Chloride of Silver (AgCl).-Native chloride of silver
occurs, sometimes in cubical crystals, sometimes in compact
semitransparent masses, which are sectile, and, from their general appearance, have given the mineral the name of horn-sill er.
Chloride of silver may be precipitated from any soluble silversalt by adding to the silver solution chlorhydric acid, or the solution of any soluble chloride; or it may be obtained by passing
over a dry silver-salt a stream of dry chlorine gas. This last reaction is the basis of a method of preparing anhydrous nitric acid.
When a stream  of dry chlorine is made to pass over perfectly
dry nitrate of silver heated to 500 or 600, the following reaction
takes place:Ag2N206 +  2C1 =  2AgC1 +  N2,O  +  0.
The characteristics of precipitated chloride of silver have been
already described (Exp. 272). The presence of an extraordlnarily minute proportion of chloride of silver renders a clear
liquid opalescent. It is easy to detect silver in a solution of
which it forms only the- I-  part, by adding to the solution
a drop of chlorhydric acid or of a soluble chloride. An admirable method of determining the amount of silver present in any
solution depends upon the insolubility of chloride of silver, the
density and peculiar curdy quality of the precipitate, and the
visibility of the smallest trace of it in a clear fluid. This method,
now generally employed in mints and assay-offices, is applicable
to the quantitative analysis of silver alloys; it is volumetric, and
depends upon the measurement of the amount of a standard solution of chloride of sodium which is required to effect the complete precipitation, as chloride, of the silver contained in a given




ATOMIC WEIGHT OF SILVEB.                461
weight of the alloy. In a solution which is acidulated with nitric
acid, and which contains no excess of the soluble chlorides, the
chloride of silver is easily coagulated into dense flocks by agitation; so that the exact point at which the precipitate ceases to be
formed is readily perceived.
Chloride of silver melts at about 2600. It is not decomposed
when heated with carbon, but is easily reduced by hydrogen
when heated in a current of the gas; zinc and iron reduce moist
chloride of silver to metallic silver; when heated with carbonates
or hydrates of sodium, potassium, or calcium, chloride of silver
gives its chlorine to the other metal, and pure silver is set free.
These methods of reducing chloride of silver, except that by hydrogen, are turned to account in the refining of silver on a large scale.
The coin, or bullion to be refined is dissolved in nitric acid, and to
the solution chloride of sodium is added; the precipitated chloride of
silver is washed until the washings are tasteless, and is then slightly
acidulated with sulphuric acid; bars of zinc are plaecl in the moist
mass, and the whole left at rest for two or three days. Chloride of
zinc and metallic silver are the products. As soon as the reduced
silver is entirely soluble in nitric acid, the reduction is complete.
The reduced metal is digested for two or three days in dilute sulphuric acid, to remove adhe-ring zinc-salts, and is then thoroughly
washed, dried, and finally melted and cast into ingots. If an absolutely pure metal is desired, the first reduet'on should be made with
pure zinc, and this refined silver may be aga:l dissolved in nitric acid,
thrown down as chloride, and reduced ag.ll fi'om tfhe washed chloride
by fusion with carbonate of calcium.
539. The reduction of chloride of silver by hydrogen is the
basis of one of the several determinations of the atomici weight of
silver; and since silver forms a large number of anhydrous salts
with acids, and has little or no tendency to form more than one
salt with each acid; the silver-salt is often the best one to prepare
and analyze whenever the combining-weight of an acid is to be
determined. But it is clear that the accuracy of these determinations depends upon the accuracy with which the atomic weight
of silver is known; hence extraordinary pains have been taken
to arrive at the true atomic weight of silver. It has been found,
by the most careful experiment, by heating chloride of silver in
a current of hydrogen, that in 132-856 parts of that compound,




462           VERIFICATION OF ATOMIC WEIGHTS.
100 parts of silver are united with 32'856 of chlorine.  If the
atomic weight of chlorine be accepted as 35'5, a simple proportion leads to the atomic weight of silver.
32'856        35'5    =      100:        =   108'07.
Amount of C1. At. Weight of Cl. Amt. of Ag. At. lWeight of Silver.
An entirely different experiment verifies this result; by burning finely divided silver in a current of perfectly dry chlorine, it
is proved that 108 parts of silver combine with 35'505 of chlorine. The following round of experiment and simple calculation
will illustrate the manner in which one atomic weight is derived
from another, and all are verified by mutual comparison. Chlorate of potassium, when heated, gives off all its oxygen and
chloride of potassium remains. Assuming that the formula of
chlorate of potassium is KC10, and that the atomic weight of
oxygen is 16, we derive the following proportion from the fact of
experiment that 100 parts of KC10 yield 39'209 parts of oxygen.
39'209        48  =    60'791:         =  74-4208.
Amount of 0.    30.    Amt. of KC1. Molecular Weight of KCl.
It is another experimental fact that 100 parts of chloride of
potassium produce, when precipitated with nitrate of silver,
192'75 of chloride of silver.
100:    192-75  =    74'4208: x         143'446.
Amt.ofKC1. Amt.ofAgCL. Al. WeightofKC1. Molecular Wt. AgCl.
But it has been determined, as above stated, that 132'856 parts
of chloride of silver contained 32'856 parts of chlorine, and
accordingly
132'856:   32'856   -=  143'446: x-   35'476.
Amt. of AgC1.  Amt. of C1.  M. Weight of AgC1.  At. Weight of C.
But if the molecular weight of chloride of silver is.. 143'446
we may deduct the atomic weight of chlorine...  35'476
and so obtain the atomic weight of silver;.... 107'970
and if the molecular weight of chloride of potassium is  74'4208
we may deduct the atomic weight of chlorine,... 35476
and so obtain the atomic weight of potassium... 389448
These numbers will be found to be very nearly coincident with
those previously given as the accepted atomic weights of these
three very important elements.




SULPIIIDE OF SILVER.                463
540. Bromide and Iodide of Silver (AglBr and AgI) are two
rather rare minerals, usually associated with chloride of silver or
with native silver.  Their artificial preparation and such of their
properties as have present importance have been already alluded
to (Exp. 272).  They are both easily fusible and insoluble in
water, but soluble in concentrated solutions of the bromide and
iodide of potassium.
541. Cyanide of Silver (AgCN) is a white powder, insoluble
in water but soluble in ammonia-water, obtained by precipitating
nitrate of silver with a soluble cyanide like the cyanide of potassium.  Cyanide of silver is soluble in solutions of the cyanides of
sodium, potassium, calcium, and other metals, forming double
cyanides of the formula MAgC2N2.  When such a solution is
subjected to the action of a galvanic battery, metallic silver is
deposited at the negative pole, in a compact, adherent layer,
while at the positive pole, where a strip or plate of metallic
silver is placed, a quantity of the metal equal to that which is
deposited at the negative pole continually dissolves. A solution
which contains =-l of its weight of silver is found to be of convenient strength for the ordinary operations of electro-plating.
542. Sulphide of Silver (Ag2S).-This compound is a principal
ore of silver. The native mineral is sometimes crystallized, in
cubes or octahedrons, and sometimes massive; it has a leaden
lustre and color, and it is so soft that a knife will cut or a die
impress it; it is fusible, and when roasted in the air yields
silver (which remains in the metallic state) and sulphurous acid
(the product of the combination of its sulphur with the oxygen of
the air).  The pure mineral is very easily recognized by these
marked characteristics.  Silver is readily tarnished by contact
with moist gaseous sulphydric acid, or with a solution of a soluble sulphide; this tarnish is the sulphide of silver (~ 533). The
sulphide may be artificially prepared by transmitting a stream of
sulphuretted hydrogen through a solution of a salt of silver.
Exp. 274.-Place in a test-glass 8 or 10 c. c. of water to which 20
or 30 drops of a solution of nitrate of silvei'iave been added, and pass
through the dilute solution a slow stream of sulphuretted hydrogen.
The black precipitate is the sulphide of silver.
Strong acids, especially when hot, dissolve or decompose this




464                  THE ALAKALI GROUP.
sulphide.  It is not soluble in solutions of the sulphides of the
alkali-metals; but by fusion it may be made to unite with many
other sulphides of metals.
543. Sulphate of Silver (Ag2SO4).-When silver is boiled with
strong sulphuric acid, the silver gradually dissolves, and there are
formed the sulphate of silver, water, and sulphurous acid:2Ag +  21H2SO4 =  Ag,2SO4 +  2H2,  +  SO,.
The sulphate is dissolved by the excess of acid, but it is deposited
in great part on the addition of water, for it is but slightly
soluble in water.  As gold is not soluble in, sulphuric acid, small
quantities of gold may be separated from large quantities of silver
or silver alloys by boiling the metal, finely granulated, in castiron vessels with oil of vitriol; silver and copper dissolve, and the
gold is left behind in a fine powder. The solution of silver is
subsequently diluted, and the silver precipitated from the solution
in the metallic state by means of metallic copper. (Exp. 267.)
Old silver coin, containing not more than   I  of gold, has been
profitably worked over by this process.
544. The A7lcali Group.-The metals which must plainly be
classed together under this head are sodium, potassium, (ammonium,) lithium, rubidium, and csesium. Two other metals are
better classed with this group than elsewhere; but their likeness
to the alkali-metals is but partial, and in many respects their properties are quite unlike those of the six metals just enumerated;
these two metals are silver and thallium. The common properties of the alkali-metals are mainly these:-tliey have the lustre
of silver, are soft, easily fusible, and volatile at high temperataures; they unite greedily with oxygen, and decompose water
with facility, forming basic hydrates which are very caustic and
intensely alkaline bodies, not to be decomposed by heat; their
carbonates, sulphates, sulphides, and chlorides, and, indeed, the
vast majority of their salts, are soluble in water; and each metal
forms but one chloride, one bromide, and one iodide; they all
form basic oxides, and never an acid oxide; they occur in nature
in modes analogous though not the same; their corresponding
salts are often, though not always, isomorphous; lastly, there is
a general, though not absolute, uniformity among the formulae of
the compounds into which these elements enter, so that, if a com



QUANTIVALENCE.                   465
pound of a given composition is proved to exist with one of these
elements, the strong presumption is that analogous compounds
with all the other elements of the group exist likewise, with properties similar though not identical.
Silver and thallium present, on the whole, so few points of
resemblance to the alkali-metals that they would not be comprehended in the smne group with them were it not for one consideration weighty enough to turn the balance when the discussion
of other properties leaves the matter in doubt. Sodium, potassium, (ammonium,) lithium, caesium, rubidium, silver, and thallium
all replace hydrogen, atom for atom. All these elements are exchangeable for hydrogen and with each other, atom for atom, and
in the present state of the science they must be regarded as the only
metals thus equivalent to hydrogen. The atom of the elements
of the chlorine group, including fluorine in that designation, and
of the seven elements above enumerated, is exchangeable for one
atom of hydrogen; it is worth one in exchange, and these elements are therefore said to be univalent, or, with less verbal precision, rnonatomwc.
545. Quanttvalence. —The chemical elements have not all the
same atom-Jixing power; thus, while an atom of chlorine combines
with only one atom of hydrogen, an atom of oxygen has the
power to drag two atoms of hydrogen into a molecule; an atom of
nitrogen holds three atoms of hydrogen in firm chemical combination, and an atom of carbon four hydrogen-atoms. In all double
decompositions an atom of sodium, potassium, or silver replaces
one atom of hydrogen, but an atom of calcium or lead two atoms
of hydrogen (~ 82). To indicate conveniently the atom-fixing
power of each element a sign is needed and a name. The conventional sign is a Roman numeral, or an equivalent number
of accents, placed above and at the right of the symbol of the
element, in case its atom be worth more than one of hydrogen;
and for the name to denote this atom-fixing power of the elements
the word " quantivalence" may be used, or the less descriptive
word " atomicity." The elements are called univalent, bivalent,
trivalent, and quadrivalent, or monatomic, cliatomic, triatomic, and
tetratomric, according as their respective atoms are capable of saturating, or holding in firm chemical combination, 1, 2, 3, or 4
2a




466                   QUANTIVALENCL
atoms of hydrogten. Thus, while the simple symbols Cl. Br, K,
Ag indicate that chlorine, bromine, potassium, and silver are
univalent, the symbols of nitrogen, antimony, and other trivalent
elements may be written N"', Sb"', &c. In the same way the
symbols 0" and Ca" indicate that oxygen and calcium are bivalent, and the symbol C"" shows that carbon is quadrivalent.
The quantivalence of many of the elements Is not yet determined with certainty; but the classification into groups of the
elements we have thus far studied rests upon the quantivalence
of the elements, as well as upon the other chemical resemblances,
which have been dwelt upon in connexion with each group. The
elements of the chlorine-group and the alkali-group are univalent;
the elements of the sulphur-group and the majority of the metals,
hereafter to be studied, are bivalent; the elements of the nitrogen-group are trivalent, and of the carbon-group quadrivalent.
It must not be supposed that the atom-fixing power of the elementary bodies is, under all circumstances, and in all compounds,
invariably exerted to the fullest extent. Were the combination
of the elements governed by any such law as this, it would evidently be impossible for any two elements to unite in more than
one proportion. Thus trivalent nitrogen and bivalent oxygen
could only combine in the proportions represented by the formula
N2"'0,", proportions which completely satisfy the atom-fixing
power of both elements. But we know that these two elements
actually form no less than five different compounds (~~ 75, 76),
of which only one is marked by the complete equilibrium of the
two elements; and this one is by no means the most stable member of the series; on the contrary, it is about the most unstable.
The student must not imagine that a bivalent element has twice
as strong affinities as a univalent element; the atom-fixing power
of an element is no test or index of the avidity with which it
seeks combination.  Chlorine, which holds but one atom of hydrogen, is competent to decompose sulphuretted hydrogen, ammonia, and marsh-gas, although sulphur unites by preference with
two, nitrogen with three, and carbon with four atoms of hydrogen.




CALCIUM.                        467
C H AP T E R XXVIIT.
CAL C I U —S TRO N T I U M-B A RI U M —LE A D.
CALCIUM.
546. This metal is a constituent of several of the commonest
and most important minerals; it forms a very considerable portion
(perhaps as much as one-sixteenth) of the solid crust of the
earth. Before considering the properties of the metal itself, let
us examine some of its familiar compounds.
547. Carbonate of Calcium (CaCO,) occurs in nature in many
different forms, called by a great variety of names, among which
may be mentioned limestone, chalk, marble, calc-spar, and coral.
There are whole ranges of mountains composed almost entirely
of limestone, while in many extensive tracts of country the soil is
calcareous and reposes upon limestone rocks. The shells of shellfish are almost entirely composed of it, and it is an important
constituent of dolomite, marl, and many other rocks and minerals.
It is formed artificially, as has been seen (Exp. 168), when carbonic acid is brought into contact with lime-water; but it is
noteworthy that carbonic acid will not unite with the anhydrous
oxide of calcium (quicklime).
Carbonate of calcium, though tasteless, is slightly soluble in
water, and the solution exhibits a faint alkaline reaction; it
is, however, rather freely soluble in water charged with carbonic
acid (~ 403).
Exp. 275. —Place in a test-tube 20 or 30 drops of lime-water, and
as much pure water; immerse in the mixture the delivery-tube of a
bottle from which carbonic acid gas is being evolved (Exp. 171). Carbonate of calcium will be thrown down at first; but after a while, as
the water in the test-tube becomes saturated with carbonic acid, the
precipitated carbonate will redissolve, and there will be obtained a
perfectly clear solution, which, in spite of the large proportion of carbonic acid contained in it, has a decided alkaline reaction. By boiling
the solution, so that a portion of its carbonic acid may be expelled, the
carbonate of calcium can be again precipitated. So, too, if the liquid
2    2




468                CALCAREOUS PETIRIFACTIONS.
be left exposed to the air, it will gradually give off carbonic acid, and
become turbid from deposition of carbonate of calcium.
The phenomena illustrated in this experiment often occur in nature.
In many districts where limestone is abundant, the well- and riverwaters are highly charged with carbonate of calcium held dissolved by
carbonic acid; the water is thus made " hard" (see ~ i560), and is, comparatively speaking, unfit for waslhing and for many other purposes.
When employed as a source of steam-power, such waters deposit carbonate of calcium as an incrustation upon the sides of the boilers as
fast as the excess of carbonic acid is expelled by boiling. This scale,
or incrustation, forms a more or less coherent coating upon the inner
surface of the boiler, and, being a very poor conductor of heat, it greatly
interferes with the heating of the water; the scale keeps the water
away from the iron sides of the boiler, and the metal, being thus unduly heated, is rapidly oxidized, or " burnt out," as the fireman correctly states it.
The formation of calcareous petrifactions, of stalactites and stalagmites, of the stones called tufa and travertine. and of many deposits oi
crystallized carbonate of calcium, is directly referable to the escape of
carbonic acid from calcareous waters. Whenever water, charged with
carbonate of calcium, flows out from the earth into the open air, or
trickles into hollows or caverns within the earth, carbonic acid is given
off in the gaseous state, and carbonate of calcium is deposited. Stalactites are the pendent masses, like icicles, which hang fiom the roofs of
caverns and the walls of cellars, bridges, and like covered ways; stalagmites are the opposite masses which grow up out of the drops of water
which fall from the stalactites above them, before all the dissolved
carbonate has been deposited. The waters of some mineral springs
are so highly charged with carbonate of calcium, that, on being exposed
to the air, they quickly deposit a considerable quantity of it upon any
solid substance with which they come in contact. In case such waters
flow over pieces of wood or other organic matter, the form of the wood
will be preserved in the cast or "petrifaction," long after the wood
itself has decayed and disappeared. Where such deposits are formed
upon a scale so large as to be of geological importance, as is the case
in some of the volcanic districts of Italy, the rock formed is called tufa
when porous, and travertine if compact.
548. Carbonate of calcium dissolves also in aqueous solutions
of several of the salts of ammonium, such as the chloride, nitrate,
and sulphate, especially if it has only recently been precipitated
and is still moist and incoherent.
Exp. 276. —Through 2 or 3 c. c. of lime-water, contained in a test



CARBONATE OF CALCIUM NOT INSOLUBLE.          469
tube, blow, by means of a glass tube, a quantity of air fi'om the lungs;
to the milky liquid obtained, add, drop by drop, a cold, saturated
aqueous solution or chloride of ammonium, until the cloudiness in the
lime-water has disappeared-that is, until the carbonate of calcium
has all been dissolved.
EZxp. 277.-Place a drop or two of a solution of chloride of calcium
in a test-tube, pour upon it several drops of a strong solution of chloride of ammonium; shake the mixture, and then add to it a few drops
of a solution of carbonate of ammlonium, and also a few drops of ammonia-water. If enough chloride of ammonium has been added to the
liquid, no precipitate will be formed in it, though, in the absence of
chloride of ammonium, a precipitate will at once be produced on mixing the other ingredients. A precipitate may, however, always be obtained by boiling the mixed solutions, unless a large excess of chloride
of ammonium be present, or unless the chloride-ot-calcium solution be
very dilute.
By repeating this experiment under varied conditions, talking note,
in each case, of the number of drops of the solutions of chloride of ammonium, chloride of calcium, and of water employed, and methodically
increasing or diminishing each of these, the student will quickly perceive the real significance of the solvent power of the ammoniacal salt,
and will appreciate the fact that, in testing for small quantities of
either lime or carbonic acid, it is necessary for the analyst to exclude
ammonium-salts from his solutions as far as may be practicable.
When boiled with solutions oi the salts of ammonium (with choride of ammonium for example), carbonate of calcium is rapidly decomposed and dissolved, carbonate of ammonium being given off, while
the chloride (or some other salt) of calcium remains in solution.
549. Carbonate of calcium is remarkable not only for the very
great diversity of external appearance which is presented by its
several massive and amorphous varieties, but it is likewise found
in a greater variety of regular crystalline forms than any other
substance; more than 150 native varieties of it have been observed
by mineralogists. As calc-spar, it occurs in rhombohedrons and
other derivative forms of the sixth or hexagonal system (~ 191);
but it is found also as the mineral arragonite, in forms of the
trimetric system, and is consequently dimorphous.
The two forms of carbonate of calcium, calc-spar and arragonite, present many differences in their physical properties.
Some specimens of calc-spar, called Iceland spar, are perfectly
transparent and colorless, and exhibit to a remarkable degree the




470               CALC-SPAR AND ARR.GONITO.
phenomena of double refraction. Transparent crystals of arragonite exhibit also the phenomena of double refraction; but arragonite has two axes of double refraction, calc-spar only one.
Crystals of calc-spar are cleavable parallel to the faces of the
rhombohedron which is the primary form of the mineral, and
masses of it may often be broken up into more or less perfect
rhombohedrons. Arragonite, on the contrary, presents two directions of distinct cleavage, parallel to the faces of a right rhombic
prism. The fractures of the two minerals are therefore quite unlike.  The specific gravity of calc-spar ranges from 2'7 to 2-75,
while the specific gravity of arragonite is generally between 2'9
and 3-3. Arragonite is considerably harder than cale-spar, but
its specific heat (0'1966) is less. When carbonate of calcium
crystallizes from hot solutions it takes the form of arragonite, but
from cold solutions it crystallizes as calc-spar. In like manner
the precipitate formed by mixing boiling solutions of chloride of
calcium and carbonate of ammonium is seen under the microscope
to consist of acicular crystals of arragonite, while the precipitate
obtained from cold solutions of the same salts is amorphous. In
either case, however, if the moist precipitate be left to itself for
some time in the cold, it will gradually assume the rhombohedral
form of cale-spar, no matter whether it was at first acicular or
amorphous.
In all its varieties carbonate of calcium is readily attacked by
acids, even if they be dilute; the action is attended with effervescence, owing to the expulsion of carbonic acid and the escape
of this gas through the liquid:CaO,CO, +  2HC1  =  CaCl1  +  CO2 +  IH20.
Limestone is readily distinguished by this reaction from other
rocks.
550. Oxide of Calcium (CaO).-On being heated, carbonate of
calcium begins to give off carbonic acid at a low red heat, as has
been seen in Exp. 170, and at full redness is completely resolved
into oxide of calcium, commonly called quicklime, and carbonic
acid.
Exp. 278.-Place a small fragment of marble upon a piece of charcoal and heat it strongly in the blowpipe-flame during several minutes.
Or throw a lump of limestone upon an anthracite fire, and leave it there




OXIDE OF CALCIUM.                  471
f:r half an hour or more. In either case, it will be found, upon examination, that the calcined product has lost the property of effervescing with acids; that it weighs less than the original iimestone, and
that it exhibits a distinct alkaline reaction when placed on wet testpaper.
For use in the arts, limestone is burned in special furnaces, of
peculiar construction, called lime-kilns, some of which are so
arranged that they may be kept in operation for years without
intermission.  When carbonate of calcium, instead of being
heated merely in quiescent air, is heated in a current of air, or
of any other gas, such as steam for example, it will give off all
its carbonic acid very easily. It has been found in practice that
limestone fresh from the quarry can be more readily burned than
that which has been long dug out of the ground and has so lost
its natural moisture; in damp weather, moreover, the burning is
said to go on more satisfactorily than when the atmosphere is
dry. If carbonate of calcium be ignited in a tube of iron, or other
metal, closed hermetically, so that no carbonic acid can escape
from the tube, the carbonate disengages carbonic acid until the
pressure of the confined gas becomes so great as to arrest the
further decomposition of the carbonate.  Under these conditions,
the temperature may be raised high enough to fuse the undecomposed carbonate; the cooled mass often presents the appearance
of fine-grained marble. If the tube in which the experiment has
been performed be very slowly cooled, the carbonic acid will be
reabsorbed.
Of the anhydrous oxide of calcium little need here be said. It
is infusible at the most intense heat at present at our command,
and is therefore used for making crucibles in which the most refractory metals are melted by the aid of the compound blowpipe.
It has no power to unite with dry carbonic acid at ordinary temperatures, but when exposed at very high temperatures to an
atmosphere of carbonic acid possessing a certain tension, some of
the gas is absorbed.  It unites with water very energetically, and
the product of this union combines readily with carbonic acid.
When lumps of quicklime are exposed to the air they slowly
absorb both water and carbonic acid, and after a while fall to
powder.  This powder is known as air-slaked lime; its composi



472                  HYDRATE OF CALCIUM.
tion may be represented by the formula CaH,2O,,CaCO,, or. dualistic, CaO,C,0; CaO,120.
551. IHydrate of Calcium (CaT202).-When water is brought
into contact with oxide of calcium, the latter swells up and falls to
powder; a large amount of heat is evolved, and there is obtained
a compound of calcium, hydrogen, and oxygen, commonly called
slaked lime, or in chemical language hydrate of calcium:CaO  +  H20 =  CaHO.22
RE:p. 279.-Place a lump of recently burned quicklime, weighing
about 30 grins., upon a large earthen plate; pour upon the lime some
15 or 20 c. c. of water, and observe how much the lime increases in
bulk as it is converted into hydrate of calcium. The heat of the mass
may be shown by thrusting an ordinary friction-match into the middle
of it; or, in case a considerable quantity of quicklime has been employed, by excavating a small hole in the dry powder and throwing in
a few grains of gunpowder, inflammation will ensue in both cases.
That much heat is evolved, may be shown also by covering the moistened quicklime with a not too tall inverted beaker g'lass or bottle, and
observing that, after a considerable amount of aqueous vapor has been
condensed upon the walls of the glass, the space within the latter will
at last become filled with a hot, invisible atmosphere of steam; when
the bottle is lifted, and the steam thus brought into contact with the
cold external air, a dense cloud or fog is immediately formed.
So much heat is developed during the union of water with lime, that
wood will quickly be brought to the kindling-temperature and inflamed,
if it happen to be in contact with large masses of these substances reacting upon one another. Fires are very frequently occasioned by the
access of water to ships or warehouses in which quicklime is stored.
It has been noticed, when large quantities of quicklime are slaked in
a dark place, that light as well as heat is evolved from the lime. Even
when quicklime is brought into contact with ice,so much heat is evolved
that the mixture sometimes becomes hot enough to boil water.
552. When hydrate of calcium is stirred into water, there is
formed not only a true solution, lime-water, which may be obtained clear and colorless by filtration (See Exp. 168), but also
a turbid liquor consisting of particles of solid hydrate of calcium
diffused through the lime-water; this liquor is known as milk or
cream of lime, according to its consistency.  In slaking lime, only
about half a part of water is really needed to convert one part of
quicklime into hydrate of calcium; but in all cases where a fine,




SLAKED LIME.                     473
smooth paste is desired, as in the preparation of milk of lime,
or of mortar, and in general whenever hydrate of calcium is
required in a very finely divided condition, it is best to pour two
or three parts by weight of water upon one part of quicklime,
so that the slaking may be quickly effected.  By using hot water
the process may be still further accelerated.  The proportions
of material given at the beginning of Exp. 279 are better adapted
than these last for illustrating the evolution of heat; but if
too little water be employed, the hydrate of calcium formed is
liable to be granular and crystalline rather than powdery. Both
milk of lime and dry powdery hydrate of calcium are largely
employed for purifying the illuminating gas made from  coal.
They remove from the gas sulphydric and carbonic acids.
Exp. 280.-Provide two gas-bottles, one arranged for generating
sulphydric acid (Exp. 86), the other for generating carbonic acid
(Exp. 171). Connect with one of the gas-bottles a tube filled loosely
with dry hydrate of calcium (Appendix, Fig. 15), and with the other
a small bottle containing milk of lime. Pour chlorhydric acid into
the gas-bottles, so that sulphydric and carbonic acids shall be freely
evolved, and test, from time to time, with lead-paper (Exp. 90), and
with lime-water (Exp. 170), as to whether these acids are completely
absorbed by the dry hydrate of calcium and the milk of lime. After a
while, change the places of the absorbing tube and bottle, so that the
milk of lime shall now be where the dry hydrate was before, and again
test the efficiency of the absorption, with lead-paper and lime-water.
In actual practice it is found that, while the dry hydrate is a more
efficient absorbent of carbonic acid than milk of lime, the latter is
capable of taking up far more sulphydric acid than the former.
553. Hydrate of calcium may be obtained crystallized, in hexagonal prisms, by evaporating lime-water in the dry exhausted
receiver of an air-pump. At a red heat it gives off its water,
and is reconverted into quicklime. The residue in this case is
left in an open, porous condition which well fits it for many
chemical purposes (see ~ 120).
It is noteworthy that hydrate of calcium is somewhat less
soluble in hot than in cold water.  If a cold, saturated solution
of lime-water be boiled, nearly half of its solid contents will be
deposited; and in case none of the water has been driven off, the
matter thus precipitated will slowly dissolve again after the liquid




474                       MORTAR.
above it has become cold. In studying this point, the experimenter must take care that the solution is kept out of contact
with the air, lest it absorb some of the carbonic acid which is
always present in the atmosphere, and become turbid from deposition of carbonate of calcium. A familiar instance of this absorption is seen in cases where milk of lime is employed for whitewashing: the loosely adherent white coating, left after the liquid
has become thoroughly dry, is no longer hydrate of calcium, but
carbonate of calcium in a more or less pure condition.
554. Slaked lime is very largely employed for making mortar,
as an ingredient of various cements, and for plastering. When
mixed with enough water to form a thick paste, it is decidedly
plastic, and admits of being spread and moulded like wax or clay.
This paste sets, as it dries, to a firm, solid mass, which, when in
thin layers, adheres firmly to any rough surfaces upon which it
may have hardened. When, however, any considerable mass of
the moist paste is allowed to solidify by itself, the dry product
will gradually crack and fall to pieces. Lime-paste cannot, therefore, be employed as a mortar unless it be mixed with some substance like sand, which shall present numerous surfaces upon
which the hardened product may adhere; by the addition of sand,
moreover, the moist lime is prevented from shrinking too much
as it becomes dry.
Mortar is commonly prepared by mixing I part of quicklime
with water enough to form a thin paste. then adding 3 or 4 parts
of coarse, sharp sand, and thoroughly incorporating these ingredients. The paste thus obtained is applied as a thin layer to the
moistened surfaces of the bricks or stones to be united. The
pasty mortar soon sets to the hard mass above described, and, on
continued exposure to the air, it slowly absorbs carbonic acid at
its surface, and is there converted into a compact compound of
hydrate and carbonate of calcium. The stone-like mass thus obtained binds firmly together the bricks or stones between which
it has been interposed. It has been asserted that the original
mortar-paste sets more firmly if it contain a certain admixture of
carbonate of calcium, than if it contain only the pure hydrate;
this admixture is, of course, produced when mortar is left for
some time in contact with the air before being used. In the




CAUSTIC LTME.                    475
course of time chemical combination occurs, to a limited extent,
between the silicic acid of the sand and the oxide of calcium in
the hardened mortar, though the process goes on but slowly;
each grain of sand finally becomes covered with a thin layer of
hydrated silicate of calcium, which contributes materially to the
solidity of the mortar. The mortar taken from old buildings
yields a certain proportion of gelatinous silica on being treated
with chlorhydric acid (~ 466).
The conversion of the original mortar into hydrocarbonate and
silicate of calcium is never completely accomplished; in the central portions of the mass, free hydrate of calcium will still be
found after the lapse of many centuries. Samples of mortar, recently taken from the Great Pyramid, were found on analysis
to contain a large proportion of the free hydrate.
555. The plastering used for finishing the walls and ceilings of
rooms is mortar to which a quantity of hair has been added to
increase its tenacity; in drying, it is, of course, subject to the
same chemical changes as ordinary mortar. By absorbing carbonic acid from the air, it is gradually converted, in part, into
carbonate of calcium, while water is set free:CaO,Ht20 +  CO2 =  CaO,CO, +  H 20.
Consequently the walls of recently plastered rooms cannot become
permanently dry, until enough carbonic acid has been absorbed
to expel the chemically combined water from their outside surfaces; hence the dampness so often perceived in new houses,
when carbonic acid first comes to be freely generated in them by
respiration and by burning lamps. In order to dry plastering, it
would, doubtless, be better to employ open fires of charcoal, or of
coke, and to deliver the products of the combustion directly into
the room which is to be dried, instead of relying solely upon hot
air, as is now usual.
556. Hydrate of calcium, like the hydrates of sodium and of
potassium, exhibits a strong alkaline reaction when tested with
moistened litmus-paper, and exerts a corrosive action upon most
organic substances; hence it is often called caustic lime.
Exp. 281.-Add a few drops of water to a small quantity of dry
hydrate of calcium, and rub it to a paste between the fingers. It will




476                 SULPHATE OF CALCIUM.
be felt that the alkali acts upon the skin; a little of the cuticle is
really dissolved.
EExp. 282. —Wrap a handful of dry hydrate of calcium in a papel;
or, better,. in a piece o-i linen or cotton cloth, and set the packet
aside for a week or two. After a while, the cloth or paper will become
rotten and friable: the caustic lime, as the common phrase is, has eaten
awsay their more corruptible portions, and has so destroyed the integrity of the whole. As a preliminary operation in tanning leather,
hides are soaked in milk of lime to loosen the hair, so that it may be
readily scraped off. The value of lime, as an ingredient of composts
to be used as manure, appears to depend, in great measure, upon its
power of hastening the decay and disintegration of organic matter.
Lime has been found to be specially valuable as manure when applied to soils rich in vegetable matter. The organic matters are decomposed or oxidized into carbonic and various other organic acids,
which unite with the lime; sometimes, under special conditions, more
or less nitrate of calcium is found among the products.
lime is important, also, from being not only the cheapest alkali,
but the cheapest of all the bases.  Since its compounds with carbonic and sulphuric acids are nearly insoluble in water, it is
largely employed for removing these acids from solutions in which
their presence is not desired; it may itself be removed from any
solution by means of the acids in question.  It is used in the
manufacture of the caustic alkalies (soda and potash), of ammoniawater and of bleaching-powders, as a flux in many metallurgical operations, in the refining of sugar, for preparing a lime
soap in the manufacture of stearine candles, and for numberless
other purposes. A noteworthy property of slaked lime is its
power of dissolving freely in solutions of common sugar.
557. Sulphate of Calcium  (CaSO4) is found native in large
quantities, as the minerals gypsum and alabaster.  These minerals contain one-fifth their weight of water; their composition
may be represented by the formula CaSO4 + 2HO0. The same
hydrated salt may be obtained by adding sulphuric acid, or the
solution of some sulphate, to a strong aqueous solution of almost
any of the salts of calcium.  This hydrated compound is the substance commonly meant when sulphate of calcium is spoken of.
The anhydrous compound is also important: it is sometimes
found in nature as the mineral anhydrite, and may be readily




PLASTER OF PARIS.                    477
prepared by heating the hydrated salt.  There is still a third
compound, the composition of which may be represented by the
formula 2CaSO4 + H20, of which, however, but little is known.
xrp. 283.-Place in a porcelain evaporating-dish, or, better, in an
iron pan, two tablespoonfuls of powdered gypsum; heat the gypsum moderately over the flame of the gas-lamp, and observe the movement of ebullition occasioned by the escaping water; stir the mixture
as long as the vapor of water is seen to escape, and then set the residue
aside to cool. The dry product is known as calcined gypsum or plaster
of Paris.
As much as nine-tenths of the water which the gypsum contains
may be readily expelled at temperatures between 100~ and 120~; but,
in order to drive off the last portions of the water, a temperature of
nearly 300~ is required. If the dry compound be heated to temperatures much higher than 300~, its particles appear to become agglutinated, and the chemical properties of the substance are somewhat
changed; the gypsum is then said to be over-burned. At the temperature of redness, sulphate of calcium melts without decomposition,
and, on cooling, assumes a crystalline structure similar to that of native anhydrite.
558. When powdered sulphate of calcium, which has been made
anhydrous at a comparatively low  temperature, is made into a
paste with water, and then left to itself; it soon sets or hardens
into a compact, coherent mass. This solidification is a consequence of the reassumption by the sulphate of calcium of the
two molecules of water of crystallization which were driven off
by heat when the substance was made anhydrous.
Exp. 284.-Place a small coin at the bottom  of a cylindrical
pasteboard pill-box a little wider than the coin; smear the coin and
the interior of the box with a thin film of oil. Mix intimately two
or three teaspoonfuls of the calcined gypsum of Exp. 283, with about
half their volume of water, in a small porcelain dish, and quickly
pour the mixture into the box, so that the coin shall be completely
covered by it. The mixture, which is of the consistence of cream,
should then be immediately stirred or puddled with a hair-pencil, or
with a tuft of cotton tied upon a stick, or with the end of the finger,
so that the bubbles of air which remain adhering to the surface of the
coin may be pressed out, and the moist paste be made to come everywhere into contact with the metal. In the course of a few minutes the
paste will solidify and become so hard that the pasteboard envelope
may be torn away from it, and the coin removed. A perfect cast or




4718                    rPLASTER-CkSTS.
copy of the stamp upon the coin will be found impressed upon the
hardened gypsum. The impression in this first cast is, of course, reversed, but by smearing it with oil and then pouring over it a new
portion of the gypsum-paste, precisely as was done with the coin, a
fac-simile of the original coin may be obtained.
Plaster of Paris is largely used in this way for taking accurate copies
of a great variety of objects.  Thus, in the process known as stereotyping, a thin paste of plaster is poured upon the surface of the printers'
types, after they have been set up and made ready for printing; the
mould thus formed is dried and baked to expel the water from the
gypsum, and is then plunged into a bath of a melted alloy of lead,
antimony, and tin, known as stereotype-metal, in such manner that,
on withdrawing the mould and allowing the metal within it to cool,
there is obtained a fac-simile of the original types. From this durable
metallic casting the page is finally printed.
As has been said above, the moist paste sets as soon as the water,
which has been mechanically mixed with the anhydrous sulphate of
calcium, enters into chemical combination with it. As in all other
instances of chemical action, so here, heat is evolved as the water and
plaster combine, as may readily be appreciated by operating upon considerable quantities of the materials. Since the plaster assumes crystalline form as it becomes hydrated, the paste increases in bulk as it
hardens, and is thus pressed into the finest interstices of the moulds.
Gypsum sets the more quickly in proportion as the temperature at
which it has been dehydrated was low. After it has been heated above
300~, it will no longer set on being mixed with water. Besides its
use in taking casts, plaster of Paris, on account of this power of combining with water, is largely employed in the preparation of stucco
and of various imitations of marble. The hydrated compound finds
application also as a manure, in the manufacture of ammoniacal salts,
and for various other purposes..Exp. 285.-That the plaster paste expands considerably at the moment of solidification may be shown as follows: —Procure a cracked
test-tube, or small flask, and fill it completely with a paste made of
calcined gypsum and water, in the proportions of 12 pts., by weight,
of the former, to 5 of the latter. In the course of 15 or 20 minutes
it will be seen that the original crack in the glass vessel has extended
in various directions, in consequence of the expansion of the mass
within it. It will be noticed, also, that the vessel feels warm to the
hand (compare Exp. 284). Finally, by breaking away the glass envelope, there may be obtained a cast of the glass vessel.
Exp. 286. —The power of sulphate of calcium to take up water,




BOILEr-SCATL.                    479
to solidify water as it unites with it to form the crystalline compound
CaIE4SOV, can be made manifest as follows:-Prepare two tablespoonfuls of a saturated aqueous solution of chloride of calcium, by
dissolving 1 pt., by weight, of the dry chloride in 1]5 part of water;
also prepare the same quantity of a saturated solution of sulphate of
sodium (Exp. 228), and, finally, mix the two solutions. Sulphate of
calcitum will be formed, in accordance with the reaction
CaCl2 + Na2SO4 = CaSO4 + 2NaC1,
and will unite with the water in which the ingredients from which it
has been formed were previously held in solution, so that an almost
solid mass of CaSO4,2H20 will take the place of the two liquids.
Ordinary hydrated sulphate of calcium is soluble in about 400
parts of water at the ordinary temperature of the air; but, like
hydrate of calcium. sulphate of sodium, and a few other salts, it
is less soluble in hot water than in cold.  When an aqueous solution of sulphate of calcium is heated to 100~ or more, a precipitate will soon be formed in it, even if the solution be very dilute
and at temipetatures as high as 140' or 150~ the anhydrous compound is completely insoluble in water.  In the same way as with
sulphate of sodium (Exp. 228), it appears that the bihydrated
sulphate of calcium cannot exist at temperatures much superior to
1000, and that above that temperature we have to deal with other
compounds of different solubility.  In other words, the water
which is held in chemical combination in ordinary unburned gypsum may be expelled by heat even when the gypsum is dissolverd
in water.  Whenever water containing sulphate of calcium in
solution is strongly heated, as in steam-boilers, there is precipi —
tated the half-hydrated compound, of composition 2CaSO4+ H0O,
which has been mentioned above. Hence the formation of incrustations, or scale, of sulphate of calcium upon the walls of boilers
fed with sea-water, or with other water containing the sulphate.
It should be remarked that the incrustation in this case does not
depend upon evaporation; the sulphate of calcium  will be deposited the more rapidly in proportion as the water of the boiler
is hot, and as more of the impure feed-water is pumped into the
boiler.
559. Besides occurring in sea-water, sulphate of calcium is a
very common impurity in spring-water.  Water which contains




480                       Tl:sIING WATER.
much of it is " hard," and is not well adapted either for washing
or for cooking.
Exp. 287. —Dissolve a small bit of soap in hot water, and add to the
solution an equal bulk of a solution of sulphate of calcium. The mixture immediately becomes turbid, and after a few moments there will
be formed a greasy, flocculent, adhesive scum upon the surface of the
liquor. Tlis precipitate is a lime soap, formed by the union of the
fatty ingredients of the soap and the base of the sulphate of calcium.
Common soap is a compound of one or more organic acids, known as
fatty acids, with caustic soda.  This soda soap is soluble in water, but
lime soap is insoluble; hence, when a soluble salt of calcium is added
to a solution of soap, precipitation occurs. When soap is added to
hard water, it will produce neither permanent froth nor cleansing
eftect, until the sulphate, or other lime-salt present, has all been
decomposed; with such waters, much soap is consumed in removing
the calcium compound, before the proper detergent action of the soap
can be brought into play.
560. An excellent process for determining the relative hardness of several samples of water has been founded upon the
behavior of water towards soap, as set forth in the foregoing
experiment:
Exp. 288.-Prepare a sample of water, of standard hardness, as follows:-Dissolve 0'5 grm. of white marble, or other pure carbonate ecalcium, in dilute chlorhydric acid, evaporate to dryness, in order to
expel the excess of acid, and dissolve the pure chloride of calcium obtained in 2 litres of water. Next prepare a solution of soap by dio'esting 7 grms. of Castile soap, or, better, white curd soap, in 1120
grins. of a mixture of 3 parts of alcohol, of 0'83 specific gravity, and 1
of pure water, until no more soap dissolves; filter the solution, and
preserve it in a tight bottle. Measure off 100 c. c. of the water, oi
standard hardness, place it in a bottle of 200 or 250 c. c. capacity, and
by means of a graduated burette (Appendix, ~ 21), or pipette, add to
it the solution of soap by portions of 1 c. c. each. After the addition
of each c. c. of the soap solution, replace the stopper in the bottle, and
shake the latter violently, then place the bottle upon its side, and
observe whether the bubbles, which form upon the surface of the
liquid, quickly disappear. So long as the bubbles disappear immediately, new portions of the soap-liquor must be added; but as soon as
a permanent froth is formed, the operation is finished. It is customary
to consider the operation completed when the bubbles persist during
three minutes. The number of c. c. of soap-liquor which has been
employed in preducing this result, is then carefully recorded.




PEOSPHATES OF CALCIUM.                481
Samples of well- and river-water may readily be compared with the
water of known standard hardness. We have only to measure off 100
c. c. of the well-water, place it in a small bottle, as above, and add to
it the soap-liquor, whose value has been determined, until a persistent
froth is produced. If it be assumed that the standard chloride-of-calcium water represent 1000 of hardness, the comparative hardness of
aniy other sample of water will follow from the proportion: —As the
quantity of soap-liquor required to produce persistent bubbles in the
standard water is to 100, so is the quantity of soap-liquor which
produces bubbles in any given sample of water to the relative hardness of the sample.
When the water under examination has a much higher degree of
hardness than 1000, it is necessary to dilute it with from 1 to 5 times
its volume of distilled water before adding the soap-liquor; for the
curdy precipitate, which would form, if soap were added to the undiluted liquid, would interfere with the formation of froth, and so
make it difficult to determine when a sufficient quantity of the soapliquor had been used.
On being ignited in an atmosphere of hydrogen, or in contact
with substances containing carbon, gypsum may readily be deoxidized and converted into sulphide of calcium: —
CaSO4 +  4C =  CaS +  4CO.
This reduction is readily effected, also, when aqueous solutions of
gypsum are left in contact with decaying vegetable matter. Since,
in this case, carbonic acid will necessarily come in contact with
the sulphide of calcium as soon as it is formed, sulphuretted
hydrogen gas will be set free, as may be perceived wherever the
mud of docks and marshes is wet with sea-water:CaS +  H20 +  C02 =  CaCO3 +  H2S.
561. Phosphates of Calcium.-There are several of these phosphates, comparable respectively with the various phosphates of
sodium (~ 489); the most remarkable among them is the triphosphate (3CaO, P20), commonly called bone-phosphate, from being
found in bones. It is the chief of the inorganic constituents of
which the skeletons of animals are composed. Small portions of
it are found in most rocks and soils (~ 262), it being a very widely
diffused, though nowhere a very abundant, substance. Considerable masses of it have been found, however, in Spain, New Jersey, and Canada, and it is the principal ingredient of some kinds
2I




482                CHLORIDE OF CALCIUM.
of guano. No matter whence obtained, it is a valuable manure
when reduced to a fine powder. Though as good as insoluble ill
water, it dissolves readily in acids and in solutions of various
organic substances.
562. Chloride of Calcium  (CaCl2) may be prepared by dissolving chalk or marble in chlorhydric acid (as in Exp. 171),
and evaporating the solution to dryness. It is produced in large
quantities in the arts by heating chloride of ammonium with
slaked lime in the preparation of ammonia-water (Exp. 48):2NH4C1 +  CaH202 =  CaCl2 +  2NH3 +  2H20.
When dried at about 200~, chloride of calcium is left as a porous
mass, which is largely employed in chemical laboratories for drying gases (Appendix, ~ 15). It absorbs water with great avidity,
and is one of the most deliquescent substances known. When
exposed to air at the ordinary temperature, it soon absorbs so
much water that it dissolves completely. At a low red heat the
anhydrous chloride melts to a clear liquid; if ignited for any
length of time in contact with the air, it suffers decomposition to
a slight extent, a little oxide and carbonate of calcium being
formed. From highly concentrated aqueous solutions there may
be obtained crystals of the hydrated compound CaCl2 + 6H20.
Slaked lime mav be dissolved in considerable q-uantity in a
boiling aqueous solution of chloride of calcium, and the filtered
solution deposits, on cooling, long, thin crystals of a compound
known as oxychloride of calcium (CaCl2, 3CaO+ 16H20), which
is immediately decomposed when treated with pure water.
563. Hypochlorite of Calcium (CaCOL22), as has been shown
in ~ 120, is a component of the substance commonly called
"chloride of lime."  This important bleaching agent is prepared
by passing chlorine gas into chambers filled with layers of finely
powdered slaked lime, in accordance with the reaction already
set forth. Chloride of lime, or bleaching-powder, is a dry, white
powder, smelling feebly of hypochlorous acid; it always contains
a certain excess of hydrate of calcium which has been unacted
upon by chlorine; it is therefore only partially soluble in
water. When exposed to the air, it slowly absorbs carbonic
acid, and, at the same time, evolves chlorine; hence its employment as a disinfecting agent. If, instead of being left to be




HYPOCHLORITE OF CALCIUMI                 483
slowly acted upon by the carbonic acid of the air, it be treated
with a dilute acid (such as vinegar), a copious evolution of
chlorine will immediately occur.
Exp. 289.-Place half a teaspoonful of bleaching-powder in a testglass, cover the powder with water, and stir in enough of a solution of
blue litmus to distinctly color the mixture. By means of a glass tube,
blow into the mixture air expired from the lungs, and observe that the
blue color of the litmus will soon be destroyed. The carbonic acid
from the lungs decomposes the hypochlorite of calcium, and the chlorine set free destroys the color.
Exp. 290.-At the bottom of a large, tall beaker, or other widemouthed glass vessel, of the capacity of two or three litres, place a
small.bottle containing 15 or 20 grins. of bleaching-powder. Cover
the beaker with a glass plate, or sheet of pasteboard, provided with a
small hole at the centre; through this hole in the cover pass a thistletube down into the bottle of bleaching-powder, and pour upon it several
small successive portions of sulphuric acid diluted with an equal volume
of water. Chlorine gas will immediately be set free from the bleaching-powder, in accordance with the reaction
CaCl2,CaC1202 +  2H12SO0  =  2CaSO4 + 2H20 + 4C1,
and, falling over into the bottom of the large beaker, will gradually
press out and displace the air therein contained, so that after a short
time the beaker will be seen to be completely filled with the green gas.
This is by far the easiest and most expeditious method of preparing
chlorine. If desirable, the bleaching-powder may, of course, be placed
in a flask, together with the acid, and the evolved gas collected at will,
by means of suitable delivery-tubes; but many of the experiments of
Chapter VIII. may be performed perfectly well in the jar of chlorine
obtained as above. The heavy gas may be ladled out of the jar with
a dipper made of any small bottle, and poured upon a solution of
indigo to show its bleaching-power.
It will be noticed, in the above reaction, that by the addition of an
acid all the chlorine of the bleaching-powder is expelled. The point
is important as bearing upon the practical use of this agent.
Exp. 291.-Soak a bit of printed calico in a half-litre of water, into
which 10 or 15 grms. of bleaching-powder have been stirred. Observe
that the color of the calico slowly undergoes change; then transfer the
cloth to another bottle filled with very dilute chlorhydric or sulphuric
acid, and take note of the rapidity with which the color is discharged.
If need be, again immerse the calico in the bleaching bath, and afterwards in the dilute acid. Finally, wash the whitened cloth thoroughly
in water.
212




484            OXYGEN FROM BLEACHING-POWDER.
564. When heated, bleaching-powder gives off oxygen, while
chloride of calcium is left as a residue. The reaction furnishes
a cheap and convenient method of obtaining oxygen. Another
method of procuring oxygen from  the hypochlorite is to mix a
solution of the latter with black oxide of manganese, red oxide
of mercury, oxide of iron, or oxide of copper, or, better, with
hydrated sesquioxide'of iron, hydrate of copper, of nickel, or of
cobalt, and to gently warm the mixture.
Exp. 292.-Fill an ignition-tube one-third full of bleaching-powder, and arrange the apparatus so that the gas may be collected
over water. Heat the tube, and observe that the gas is expelled at
a comparatively low temperature.  1 grin. of bleaching-powder yields
40 or 50 c. c. of oxygen gas.
Exp. 293.-Take as much bleaching-powder as was employed in
Exp. 292, dissolve it in a small quantity of water, filter the solution,
and place it in a small flask provided with a delivery-tube. Add to
the contents of the flask two or three drops of the solution of a cobalt
salt, connect the flask with an inverted bottle of water upon the waterpan, by means of the delivery-tube, then heat the flask to 700 or 80~,
and observe that oxygen is freely evolved.
The cobalt solution employed amounts to the same thing as hydrated oxide of cobalt, since the latter is immediately precipitated
from the cobalt salt by the caustic lime in the bleaching-powder.
The action of the oxide of cobalt, or other metallic oxide, in this experiment, appears to be somewhat analogous to that of the higher
oxides of nitrogen in the manufacture of sulphuric acid (~ 228). The
oxide of cobalt probably takes oxygen from the solution of bleachingpowder, and combines with it to form a high, unstable oxide, which
immediately decomposes again with evolution of oxygen. The oxidation and deoxidation of the cobalt compound thus goes on incessantly,
and a very small quantity of the latter is sufficient to decompose any
desired amount of bleaching-powder. It is important that the solu.tion of the hypochlorite should be filtered as above directed, lest a
quantity of it be lost by foaming over out of the flask.
565. The proportion of hypochlorite of calcium in bleachingpowder varies widely in different samples, according to the care
with which the sample has been prepared, and to the length of
time it has been exposed to the action of the air.  The bleachingT)ower, or in other words, the money value of each special sample,
should therefore be determined before it is sold. Of the several




CILORIMUETRY.                    485
methods of ascertaining the value of bleaching-powder, one of
the simplest is to determine how much arsenious acid (As203)
can be converted into arsenic acid (As,O5) by a given weight
of the sample:
As2O3 +  4C1 +  21120 =  As205 +  4HC1.
To this end a weighed quantity of arsenious acid is dissolved in
a certain definite quantity of a solution of carbonate of sodium in
such manner that each c. c. of the liquor shall contain a known
weight of arsenious acid. A considerable quantity of this standard
solution may be prepared once for all, and kept for use in tightly
closed bottles. A weighed sample of the bleaching-powder under
examination is then dissolved in water, and into this solution of
bleaching-powder the standard solution of arsenious acid is carefully
poured, from a burette, so long as the arsenious acid continues to be
oxidized and converted into arsenic acid.
In order to determine the precise moment when the oxidizingpower of the bleaching-powder solution ceases, a drop of this solution is taken out from time to time upon a glass rod and placed
upon paper prepared with iodide of potassium and starch, as has
been described in Exp. 71. When the paper no longer becomes
blue on being touched with the solution, the operation is known
to be completed. Towards the close of the experiment the solution
of arsenious acid should only be added drop by drop to the solution of
bleaching-powder, and some of the latter should be touched to the
test-paper after each addition of the arsenious acid.
The number of c. c. of the solution of arsenious acid which have
been employed is then carefully noted, and the amount of arsenious
acid contained in them is computed. From these data the amount
of chlorine in the sample of bleaching-powder is obtained by the
following proportion:Weight of    Weight of
198:     142        =   arsenious: chlorine in
Weight of one   Weight of four    acid used.   the sample.
molecule of ar- atoms of chlorine.
senious acid.
Of the other compounds of calcium may be mentioned the
fluoride (~ 155), bromide, and iodide (analogous to the chloride),
the peroxide CaO2, several sulphides (~. 213), and the phosphide
(~ 278).  Nitrate of calcium (CaN206) is a very easily soluble,
deliquescent salt, found in many soils, and in other localities
where organic matters putrefy in contact with hydrate or car



486                STRONTIUM AND BARI'M.
bonate of calcium. Sulphydrate of calcium CaS,H,2S, analogous
to the hydrate, may be prepared by boiling monosulphide of
calcium with water, or by passing sulphydric acid gas into a cold
solution of the sulphide. The solution of this substance possesses
a remarkable power of loosening hair from the skins of animals.
After a skin has been soaked for a few minutes in a strong solution of this substance, the hair may readily be scraped off with
any blunt instrument. A solution of sulphite of calcium (CaSO,)
in an aqueous solution of sulphurous acid, is sometimes employed,
under the name bisulphite of lime, to check fermentation.
566. The metal itself may be obtained by decomposing fused
chloride of calcium by means of the galvanic current, or by
heating iodide of calcium with metallic sodium in a closed iron
tube. It is a yellowish-white, lustrous, ductile metal, of 1'6
specific gravity, which suffers no change in dry air at the ordinary temperature. In moist air it oxidizes quickly, and it
decomposes water with evolution of hydrogen. At a red heat
it melts, and, if oxygen be present, takes fire and burns with a
bright light. It is a bivalent element; the weight of its atom
is 40.
STRONTIUDM AND BARIUM.
567. The metals strontium and barium closely resemble calcium in appearance and properties, and may be prepared by
methods similar to those used for calcium, as described in the
last section. The specific gravity of strontium is 2'6, that of
barium is 4. The atomic weight of strontium  is 87'5, and
that of barium 137. Like calcium, strontium and barium are
both bivalent elements.
Most of the compounds of strontium and barium are closely
analogous to the corresponding compounds of calcium. The
oxides, peroxides, hydrates, carbonates, sulphates, nitrates, phosphates, chlorides, sulphides, &c. resemble in the main the corresponding calcium salts. The hydrates of strontium and barium,
however, are more readily soluble in water than the hydrate of
calcium, while their sulphates, nitrates, and chlorides are less
soluble than those of calcium. Sulphate of barium is almost
absolutely insoluble in water, and sulphate of strontium is only




PEROXIDE OF BARrUM.                487
very slightly soluble. Sulphate of barium is found native, sometimes in considerable masses, as a very heavy white mineral
called barytes, which, when powdered, is largely employed for
adulterating white lead. The name barium comes from a Greek
word meaning heavy.
From the carbonates of strontium and barium the carbonic
acid cannot readily be driven off by heat alone, though when
mixed with charcoal, and then ignited, these carbonates may
be reduced to oxides. A better way of preparing the oxides is
to heat the nitrates strongly in a porcelain crucible or retort.
Unlike hydrate of calcium, hydrate of barium  does not give
off its water at the temperature of redness, but melts without
undergoing decomposition. From hydrate of strontium the water
may be expelled by heat, though with difficulty. Peroxide of
barium (BaO2) is of interest, since by means of it peroxide of
hydrogen (~ 61) and antozone (~ 177) may be prepared.
568. In order to obtain peroxide of barium, a mixture of 1
part of oxide of barium, and 4 parts of chlorate of potassium
may be thrown, little by little, into a crucible heated to low
redness, and the fused mass subsequently washed with water to
remove chloride of potassium; or a current of oxygen gas, or
of air, may be made to flow over oxide of barium heated to low
redness in a porcelain tube.  As thus prepared the peroxide is
never pure, being mixed with more or less protoxide. Peroxide
of barium decomposes, with evolution of oxygen, at the temperature of bright redness, and in view of this fact it was at one time
proposed to employ the substance as a means of obtaining pure
oxygen from the air upon the large scale.  A considerable number of tubes charged with protoxide of barium, having been
suitably arranged in furnaces, half of the tubes were heated to
dull redness, and a current of air was made to flow through
them, until the protoxide had been converted into peroxide;
the current of air was then transferred to the other tubes, while
the first series was put in connexion with a gas-holder and
heated to bright redness, until the second atom of oxygen had
been driven out. The second series of tubes were next deprived
of oxygen, while the tubes of the first series were put to their
Dld work of absorbing oxygen from the air. The process thlus




488                   STRONTIUM SALTS.
became a continuous one, and was really capable of furnishing
large quantities of oxygen; it has, however, been superseded by
cheaper methods (~ 242).
Strontium salts are commonly prepared from the native carbonate, a mineral called strontianite, while the various salts of
barium are obtained either from the native carbonate (twithlerdte),
or more commonly from the sulphate. The finely powdered sulphate, after having been mixed with powdered charcoal and oil,
is strongly heated in a covered crucible, and so reduced to the
condition of sulphide of barium:BaSO4 +  4C =  BaS +  4CO.
Sulphide of barium is readily soluble in water, and on being
treated with chlorhydric, nitric, or any other acid, it decomposes,
sulphuretted hydrogen is given off, and there is formed chloride
of barium, or some other salt, according to the acid employed:BaS +  2HC11   =  BaCl2 +  H2S.
Several of the compounds of barium are useful reagents in the
chemical laboratory.  Sulphate of barium is employed as a pigment by artists in water-colors, under the name p)ermanent white,
also in the finishing of paper, pasteboard, &c., and for adulterating
white lead. As a water-color it is valuable, since it is scarcely
at all acted upon by any chemical agent; but when ground with
oil it becomes translucent, and seriously impairs the opacity or
covering power of the better pigments with which it is mixed.
Compounds of barium and of strontium are employed in the
preparation of fireworks, for obtaining green and crimson flames
respectively:The green barium-flame may be well shown by mixing with the
fingers a gramme of powdered chlorate of barium with half a gramme
of flowers of sulphur, and strewing the mixture upon a glowing coal.
The green fire of the pyrotechnists may be prepared by mixing together.58 parts of nitrate of barium, 13 parts of sulphur, 6 parts of chlorate
of potassium, and 2 parts of charcoal.
To exhibit the red strontium-flame, a mixture may be prepared by
rubbing together in a mortar 30 parts of anhydrous nitrate of strontium, 10 parts of powdered sulphur, and 3 parts of sulphide of antimony; and to this mixture may be added, with the hand, taking care
to avoid all violent friction, 7 parts of powdered, fused chlorate ot




THE CALCIUM GRO UP.                 489
potassium. The mixture may then be shaken loosely upon a piece of
sheet iron and touched with a lighted stick or glowing coal.
Or the color may be shown upon a smaller scale by operating as
follows:
rExp. 294.-By means of iron wire, suspend three small bullets of
well-burned coke from a ring of the iron stand. Heat the fragments
in turn with the flame of the gas-lamp, and observe the slightly yellowish flame which will be produced in each case; then moisten one
of the pieces of coke with a solution of chloride of calcium, the second
with a, solution of chloride of barium, and the third with a solution of
nitrate of strontium, and again heat them in turn with the gas-flame.
The calcium salt will impart a reddish-yellow color to the flame, the
barium salt a green color, and the strontium salt a beautiful crimson.
Instead of the bits of coke, platinum wire might of course be employed,
as in Exp. 202.
569. As appears abundantly from the foregoing, the three
elements calcium, strontium, and barium are intimately related
one to the other, and are, as a family, clearly distinguished in
several important particulars from the metals of the preceding
group. Even before the metals of this family were discovered
and isolated, it had long been customary among chemists to speak
of the oxides of calcium, strontium, and barium  as the alkaline
earths, in contradistinction from the " alkalies," potash and soda,
upon the one hand, and the "earths," such as the oxides of
magnesium and aluminum, upon the other.
Each of the members of the alkaline-earthy group, now in
question, decomposes water, even at the ordinary temperature,
taking away its oxygen; and each of them forms two oxides-a
neutral, insoluble binoxide, belonging to the class of antozonides
(~ 182), and a more or less soluble protoxide, acting as a powerful
base; their carbonates and sulphates are all difficultly soluble,
and, like the other compounds of the three metals, are isomcrphous with one another. In all their compounds, there may be
seen the same progression of properties which has been met with
in the groups previously studied. The barium compound will
always be found at one end of the scale, the calcium compound
at the other, and the strontium compound interposed between the
two. The hydrate and the carbonate of calcium are both readily
destroyed by heat, while the corresponding strontium compounds




490                        IE).
are decomposed with difficulty, and the barium compounds only
at exceedingly high temperatures. The solubility of the oxides
diminishes as we pass from baryta to lime, while that of the sulphates and carbonates follows the inverse order. The specific
gravities of the metals are, Ca= —16, Sr-2-6, Ba=4; and their
atomic weights are 40, 87-5, and 137 respectively, that of strontium being nearly the mean of the other two. The specific
gravities of their carbonates and sulphates are as follows:CaCO3 (arragonite) =2'95, SrCO3=3'6, BaCO3,=-433; CaSO4=
2'33, SrSO4=3'89, and BaSO4=4'4. It should be observed that,
in all these cases, the specific gravities of the strontium compounds approximate closely to the mean of the specific gravities
of the corresponding barium and calcium compounds. The same
remark applies also to the specific gravities of the three metals.
LEAD.
570. Almost all the lead which is employed in the arts is
extracted from sulphide of lead, PbS, the mineral galena. This
substance is tolerably abundant in many localities, and is often
associated with sulphate of barium, fluor-spar, quartz, and other
common minerals; it almost always contains a small proportion
of sulphide of silver.  In order to obtain metallic lead from galena, this mineral is mixed with a small quantity of lime, and
then roasted at a dull-red heat in the flame of a reverberatory
furnace. A portion of the sulphur burns off as sulphurous acid.
Some oxide of lead, and more or less sulphate of lead is formed,
while much of the ore remains undecomposed. After a time, the
roasting-process is interrupted, air is excluded from the furnace,
the oxide, sulphate, and sulphide of lead are thoroughly mixed
together, and the heat of the furnace is suddenly raised. The
undecomposed sulphide of lead then reacts upon the oxide and
sulphate, sulphurous acid is given off, and metallic lead produced.
The reactions may be thus formulated:2PbO + PbS = 3Pb + SOQ.
PbSO4 +  PbS =  2Pb +  2SO2.
The lime is added for the purpose of forming a fusible slag
with any siliceous matter which may be present in the ore.
Lead is a remarkably soft metal, of bluish-white color; it can




METALLIC LEAD.                    491
be readily cut with a knife, and may even be indented with the
finger-nail; it soils paper upon which it is rubbed. Its specific
gravity is 11'4, and its atomic weight 207. It may be drawn
into wire, and beaten into sheets, though, as contrasted with most
of the other metals, it has but little tenacity. In comparison with
other metals, it is a rather poor conductor of heat and electricity.
It melts at about 325~, and contracts considerably in passing from
the liquid to the solid condition. Its specific heat is 0'0314.
Solid lead expands greatly when heated, though the heat be not
carried near to the melting-point, and the expanded metal -does
not return again to its original dimensions when cooled.  Melted
lead begins to emit vapors at a red heat, and at very high temperatures the metal may even be distilled. Lead may be obtained
crystallized in octahedrons, by slowly Cooling the molten metal.
The ready crystallization of lead furnishes a very simple method
of separating this metal from the silver with which crude lead is
almost always contaminated as it comes from the smelting furnaces. When melted argentiferous lead is allowed to cool slowly,
and is at the same time briskly stirred, a quantity of solid crystalline grains separate out after a while, and sink beneath the
liquid metal, whence they may be dipped out in cullenders.
These crystals are composed of lead nearly free from silver, while
all but a trace of the silver contained in the original lead is left
in that portion of the metal which has not yet solidified; in a word,
the alloy of lead and silver melts at a lower temperature than
pure lead. By methodically remelting and recrystallizing the
lead crystals on the one hand, and the silver alloy on the other,
it has been found profitable to extract the silver from lead so poor
that it contained less than one thousandth part its weight of the
precious metal.
When in thick masses, such as the common sheets and pipes of
commerce, lead is scarcely at all acted upon by cold sulphuric
acid, and is but slowly corroded by chlorhydric acid.  Both
these acids form, by uniting with lead, difficultly soluble salts;
and so soon as a layer of the salt has once been deposited upon
the surface of the metal, the latter is thereby protected from
further corrosion.  By hot, concentrated sulphuric acid, however,
lead is dissolved rather easily.  The best solvent of metallic




492                   OXIDATION OF LEAD.
lead is diluted nitric acid; strong nitric acid will not dissolve it
readily, since nitrate of lead is well nigh insoluble in concentrated nitric acid..
571. Oxides of Lead.-When a compact piece of metallic lead
is freshly cut, it exhibits considerable lustre, and this lustre maybe preserved unimpaired by keeping the lead in perfectly dry air,
or beneath the surface of pure water free from air (~ 47). But
by exposure to ordinary air the brilliant surface soon tarnishes, in
consequence of the formation of a thin coating of suboxide of lead;
this incrustation protects the metal beneath from further oxidation. Finely divided lead, on the contrary, soon changes completely to suboxide when exposed to the air. If the metal be fine
enough, it will oxidize instantaneously, with evolution of light
and heat, and formation of yellow protoxide of lead.
Exp. 295.-Prepare a small quantity of tartrate of lead, as follows.
Dissolve 0-25 grm. of common sugar of lead (acetate of lead) in 8 or
10 c. c. of water, also dissolve 0'1 grm. of tartaric acid in 4 or 5 c. c.
of water, and mix the two solutions. Collect upon a filter the white
precipitate of tartrate of lead which will be formed, wash it with
water, then unfold the filter and spread it out with its contents to dry
in the air, or, better, at a gentle heat upon a ring of the iron stand, high
above a single gas-flame.
Fill an ignition-tube one-third full of the dry tartrate of lead, and
heat it upon a sand-bath so long as any fumes escape, then cork the
tube tightly and set it aside to cool. Holding the cooled tube high
in the air, sprinkle its contents upon a plate, and observe that the black
powder takes fire spontaneously, and burns with a red flash. The
composition of tartrate of lead may be represented by the formula
C4H4Pb206; on being heated this substance gives off water and carbonic oxide, as may be seen by lighting the fumes which escape from
the tube, and there is left as a residue an intimate mixture of carbon
and of metallic lead, so finely divided that it inflames in ordinary air. A
spontaneously inflammable mixture such as this is called apyrophorus.
572. Lead is far more readily oxidized by the continued action
of air and water than by ordinary moist air.  When exposed to
the simultaneous or alternate action of these agents, a coating of
the white hydrated protoxide of lead is rapidly formed; but as
this compound is somewhat soluble in water, it is continually dissolved away and affords little or no protection to the lead beneath.
The corrosive action of water upon lead is modified very mate



ACTION OF WATERS ON LEAD.              493
rially by the presence of small quantities of various saline substances. Water containing traces of nitrates, nitrites, and chlorides corrodes lead more rapidly than pure water, while the
corrosive action of pure water appears to be diminished by the
presence of sulphates, phosphates, and carbonates, oxide of lead
being scarcely at all soluble in water which contains these salts in
solution. Water contaiiing a solution of carbonate of calcium in
carbonic acid, such as is frequkntly met with in nature, has been
found to have remarkably little action upon lead; in such water
a coating of insoluble, or nearly insoluble carbonate of lead is
formed upon the metal, which protects it from further action.
But, on the other hand, water which contains much free carbonic
acid dissolves away the protective coating and exposes fresh surfaces of lead to corrosion. As a general rule, the acids, even when
very dilute, greatly accelerate the oxidation of lead in the air;
and the same remark is true of organic substances and of metals,
the first by their decay, the second by the galvanic action which
their presence excites.
Since solutions of lead are poisonous, and since the metal is
employed to an enormous extent for cisterns and conduits, a
knowledge of the action of water upon lead is very important in
a sanitary point of view. The question has consequently engaged
the attention of many chemists, and has been much discussed.
It has been proved by numberless experiments that the action of
natural waters upon lead is so general that it is rare to find any
sample of water, which has been kept in a leaden cistern, wholly
free from traces of that metal. The opinion of most chemists is at
this time (1867) decidedly adverse to the use of leaden water-pipes
in houses, in spite of the fact that the metal is nowadays employed
for this purpose almost everywhere with apparent impunity.
When lead is melted in the air it oxidizes readily, with formation at first of gray suboxide, and afterwards of the yellow
protoxide.
573. Suboxide of Lead (PbO) may be prepared in a state of
purity by cautiously heating oxalate of lead at a temperature not
exceeding 300~ in a retort from which air is excluded, so long as
any gas is evolved:2PbC2O4 =  PbO2  +  CO +  3C02.




494                      CUPELLATION.
After the retort has become cold, suboxide of lead will be found
in it as a black velvety powder.  It is decomposed by acids, with
formation of salts of protoxide of lead, and separation of metallic
lead.
574. Protoxide of Lead (PbO), commonly called lith7arge, may
be obtained as a lemon-yellow powder by gently igniting nitrate,
carbonate, or oxalate of lead upon an iron plate, or in an open
porcelain crucible.  The oxide fuses at a red heat, and when
melted in vessels of porcelain or earthenware it rapidly destroys
them by combining with their silica, an easily fusible slag or
glass composed of double silicates of lead, aluminum, iron, &c.
being formed. Silicate of lead is an actual constituent of the easily
fusible variety of glass known as flint glass (see ~ 492, and Appendix, ~ 3).
In the arts, litharge is prepare.d upon the large scale by heating
metallic lead in a current of air; the color and texture of the product varies considerably according to the temperature and the
other conditions under which it has been prepared.
Exp. 296.-Heat a small fragment of lead upon charcoal in the
oxidizing flame of the blowpipe, and observe the gray film of suboxide which forms at first, and the yellow incrustation of litharge
which is obtained subsequently. The litharge may be melted if a
strong, hot flame be thrown upon it.
This property of lead, of rapidly oxidizing when heated in the air,
taken in connexion with the easy fusibility of the oxide, is the basis
of the common method of separating lead and silver in the large way,
known as cupellation.  The iridescent film  of litharge continually
formed upon the surface of the molten metal as incessantly flows off,
exposing new surfaces of the metal to the action of the air. The
silver, on the other hand, undergoes little or no change, and when the
lead has been completely burned the silver appears in all its brilliant
whiteness.
The cupel is a shallow cup or basin, composed either of marl or of
mixture of bone-ash and wood-ashes firmly compacted and beaten to
a smooth surface, which may be placed either in. the muffle of an assay
furnace, upon the hearth of a reverberatory, or in any position where
it can be strongly heated at the same time that a free current of air
plays over its surface. A charge of argentiferous lead having been
melted upon the cupel, new portions of the lead are added as fast as
the melted litharge flows off from the convex surface of the metal and




PEROXIDE OF LEAD.                   495,
malkes room for these additions, until an alloy very rich in silver has
been obtained. This alloy is then cupelled until the last traces of lead
have been removed, and the silver is left pure and glistening. In
cupelling upon the small scale, for purposes of assaying, the cupel is
made of bone-ash of such quality that the litharge may be absorbed
into the substance of the cupel, and not flow off through gutters upon
its edge, as is the case in large metallurgical operations.
575. Protoxide of lead unites readily with acids, and forms
many important salts. When in the state of powder, it even
absorbs a certain amount of carbonic acid from the air; hence
the powdered litharge of commerce always contains more or less
carbonate of lead, and therefore effervesces on being treated with
acids, as has been seen in Exp. 42. As a general rule, it is far
better to prepare the salts of lead by dissolving the protoxide in
acids, than to treat the metal itself with acids. Protoxide of
lead has a remarkable tendency to form basic salts; thus besides
the normal nitrate (PbO,N205) there is a dinitrate (2PbO,N205),
a trinitrate (3PbO,N205), a tetranitrate (4PbO,N,,05), and a hexanitrate (6PbO,N205).
Though a strong base as regards the acids, protoxide of lead
behaves like an acid towards the alkalies and alkaline earths.
For example, it dissolves readily in soda or potash lye, with formation of plumbite of sodium, or of potassium, as the case may
be. The term plumbite, like the symbol of lead, Pb, is derived
from plumbum, the Latin name of the metal.
576. Peroxide of Lead (PbO2) may be prepared by oxidizing
the protoxide-for example, by passing a current of chlorine gas
through water in which protoxide of lead is kept suspended by
agitation, or as follows:Exp. 297.-Place 8 or 10 grms. of very finely powdered sugar of
lead in a capacious porcelain dish, cover the salt with a filtered solution of bleaching-powder (~ 563), heat the solution to boiling, and
maintain it at this temperature, until the escaping vapors smell
strongly of acetic acid; then pour the contents of the dish upon a
filter, and wash with water the dark-brown powder of peroxide of
lead which has been formed.
Peroxide of lead may be easily obtained also by digesting red lead
with dilute nitric acid, as will appear in the following paragraph.
Peroxide of lead is a powerful'oxidizing agent; it readily gives




4960                      RFD LEAD.
up oxygen to many organic substances, even at the ordinary
temperature of the air. On being heated to redness, it loses half
its oxygen, and is converted into the protoxide.
Exp. 298.-In a small porcelain mortar, rub together a mixture of
1 grin. of oxalic acid, and 1 grm. of peroxide of lead. Decomposition
will occur, aqueous vapor and carbonic acid will be given off, and carbonate of lead, PbCO,, will be left as a residue. WVhen the peroxide
is thus mixed with one-eighth its weight of sugar, or with one-sixth
its weight of tartaric acid, so much heat is developed, that the mass
in the mortar glows.
Peroxide of lead is decomposed by chlorhydric acid, with liberation of chlorine, and formation of normal chloride of lead —
PbO, +  4HCl =  PbC12 +  2H20 +  2C1,
-and by hot sulphuric acid, with evolution of oxygen. It combines with sulphurous acid readily, and is often employed in the
laboratory as an absorbent of this gas; and it is noteworthy that
the product of this combination is sulphate of lead:PbO2 +  802 =  PbSO4.
It is indifferent, therefore, whether we put together peroxide
of lead and sulphurous acid, or protoxide of lead and sulphuric
acid; the product will, in either case, be common sulphate of
lead; for
PbO  +  SO3 =  PbSO4.
As a rule, peroxide of lead does not readily enter into combination with acids, though compounds of it with acetic, phosphoric, arsenic, and some other acids have been obtained. With
strong bases, however, it combines readily, forming salts known
as plumbates.
577. Red Lead or Minium.-When protoxide of lead is kept
at a low red heat for some time, in contact with air, it gradually
absorbs one or two per cent. of oxygen, and acquires a brilliant
red color. The product of this oxidation is extensively used as a
pigment and in the manufacture of some kinds of glass ware; it
may be regarded as a compound of PbO and PbO2, in varying
proportions. By digesting it for some time in dilute nitric acid,
the protoxide of lead may all be dissolved out and converted into
nitrate of lead, while the peroxide of lead is left as a residue; —
2PbO,PbO, +  2H,2N206 =  2PbNO06 +  21120 +  PbO,.




SULP' IDES OF LEAD.                  497
Exip. 299. —Heat in an iron spoon 4 grms. of litharge and 1 grm. of
chlorate of potassium, and observe that the color of the mixture soon
changes from yellow to red. Throw the cooled product upon a filter,
wash it with water, then dry it and compare its color with that of the
original litharge. In this experiment, the red lead could be obtained
as well by simply heating litharge without admixture in the air for
many hours, at a temperature just below its melting-point; time alone
is gained by employing chlorate of potassium as the source of oxygen.
For commercial purposes, red lead is obtained by heating metallic lead
in reverberatory furnaces, or, when a very pure article is needed, by
heating carbonate of lead.
When heated strongly, red lead is resolved into protoxide of lead
and free oxygen. Oxygen gas may be prepared from it in the same
manner as from oxide of mercury (Exp. 5), though at a higher temperature.
578. Sesquioxride of Lead (Pb,O,) is recognized as a distinct
oxide by some chemists, but is more generally regarded as a
compound of the proto- and peroxides, PNO, PbO2, —a plumbate
of lead.
579. Sulphides of Lead.-There are several of these compounds, but the protosulphide, PbS, is the only one whose composition is accurately known. This sulphide is the native mineral
galena (~ 570); it may be prepared artificially either by melting
together lead and sulphur in atomic proportions, or by treating
the solution of any lead salt with sulphydric acid (~ 209). The
native mineral, like the compound obtained artificially by way of
fusion, is of a leaden-gray color of 7'5 specific gravity, but the
precipitate which forms when sulphydric acid is added to the
solution of a lead salt, is black, or brown, or even red if the
solution be dilute. On account of the deep color, as well as the
insolubility of this precipitate, sulphydric acid is often made use
of as a means of detecting lead; the test is, in fact, so delicate
that solutions containing only a hundred thousandth of their
weight of metallic lead will assume a brown  color on being
charged with sulphuretted hydrogen.
Exp. 300.-Dissolve quarter of a gramme of sugai of lead in 4 litres
of water, add to the soluItion a few drops of nitric acid so that it shall
exhibit a faint acid reaction with litmus-paper, pass into the solution
a current of sulphydric acid gas until the solution smells strongly of it
and observe the brown color imparted to the fluid after some time.
2 K




498                   CHLORIDE OF LEAD.
In testing for the presence of lead in excessively dilute solutions,
such, for example, as water drawn from leaden pipes, it is well to
evaporate the liquid to a small bulk in a porcelain dish, to acidulate
the concentrated liquor very slightly with nitric acid, and then to
transfer it to a beaker glass. The liquid should then be saturated
with sulphuretted hydrogen gas, the beaker covered with a glass plate,
and left to stand during several hours in a moderately warm room.
If lead be present, it will be indicated after a while either by the
brown coloration of the liquid, or by the actual separation of a black
powder at the bottom or upon the sides of the glass.
580. Sulphide of lead is volatile at high temperatures, and is
often found in the cracks and upon the walls of smelting furnaces, in the form of crystals, which have been deposited by sublimation. By virtue of the volatility of its sulphide, lead may be
transported to very considerable distances from furnaces where
ores containing galena are roasted or reduced.  It has been
found, moreover, that growing plants are capable of taking up
the lead thus deposited, and of assimilating a certain portion of it
in their tissues. Comparative experiments made at very high
temperatures have shown that galena may lose as much as 3'7
per cent. of its weight by volatilization, while metallic lead, exposed to the same conditions, loses less than 0'1 per cent.
Sulphide of lead, like the sulphides of the alkali-metals and
those of the alkaline-earthy metals, acts as a sulphur base;
with sulphantimonic acid, for example, it unites to form a salt,
3PbS,SbS,, analogous to that formed by the union of antimonic
acid and oxide of lead, 3PbO,SbO,.
581. Chloride qf Lead (PbC12).-Mcetallic lead is but slowly
acted upon by chlorine, or by chlorhydric acid, though hot chlorhydric acid dissolves a little of it with formation of chloride of
lead and evolution of hydrogen, even when out of contact with
the air; the chloride may, however, be readily prepared by
digesting oxide or carbonate of lead in chlorhydric acid, or by
mixing the solution of almost any lead salt with chlorhydric acid,
or with a solution of some soluble chloride:
PbN2O0  +  2NaC1 =  PbC12 +  NaN206.
Chloride of lead is but sparingly soluble in cold water, and is
still less soluble in water acidulated with chlorhydric acid; hence




SALTS OF LEAD.                    499
it may read:ly be precipitated as above described, and collected
upon a filter. In hot water, however, it dissolves rather easily,
and it is somewhat soluble in concentrated acid also.
Exp. 301.-Boil together, in a small flask, 1 grm. of litharge, 14
grins. of strong chlorhydric acid, and 14 grms. of water, during 15 or
20 minutes. Pour the mixture upon a filter supported in a funnel
which has been gently warmed by holding it over the flame of the
gas-lamp, and collect the clear filtrate in a warm bottle. As the solution cools, lustrous needle-shaped crystals of chloride of lead will
form in it.
Exp. 302.-Pour off the cold supernatant liquor from the crystals of
chloride of lead obtained in Exp. 301, place the crystals upon a fragment of porcelain, dry them at a gentle heat, and finally heat them
more strongly. It will be found that the crystals melt very easily,
and that on cooling they solidify to a soft translucent horny mass,
whence the old name of this substance, horn-lead.
582. The compounds of lead with iodine, bromine, and fluorine
are analogous to chloride of lead. The iodide is remarkable on
account of its beautiful yellow color, which may readily be shown
by adding a drop or two of a solution of iodide of potassium to a
small quantity of a solution of nitrate of lead.
Of the numerous other salts of lead little need here be said.
The nitrate and tartrate have already been prepared (Exps. 42,
295); we have obtained the sulphate also as a white, nearly
insoluble powder by adding water to concentrated sulphuric acid,
and it may be had in any quantity by mixing the solution of a
lead salt with dilute sulphuric acid, or with the solution of a
soluble sulphate. Acetate of lead, one of the most important of
the lead salts, and the one most readily to be procured in commerce in a state of purity, is prepared by dissolving oxide of lead
directly in acetic acid such as is obtained in the distillation of
wood (~ 380), or indirectly by moistening plates of metallic lead
with vinegar in vessels open to the air. It crystallizes readily,
is easily soluble in water, and has a sweet, astringent taste,
whence the name, sugar of lead. Like the other lead salts, it is
highly poisonous. It is employed for many purposes in the arts,
and is in particular much used in medicine. Carbonate of lead
(PbCO,), or rather compounds of carbonate of lead and of hydrate
of lead in varying proportions, are used to an enormous extent as
2x2




500                     WHITE LEAD.
a white paint, under the general name of white lead. The composition of this substance may usually be expressed by a formula
lying within the limits PbCO,,PbH202, upon the one hand, and
3PbCO,,PbH2O,  on the other.  As contrasted with the other
white pigments it possesses remarkable covering-power and durability, and is consequently much esteemed in spite of its high
cost, its injurious influence upon the health of workmen who
have to do with it, and the fact that it is discolored by air
containing sulphuretted hydrogen.  White lead is often adulterated with sulphate of barium, with oxide of zinc, and with
gypsum.  It is usually prepared by bringing carbonic acid, obtained from decaying vegetable matter, or from the combustion of
fuel, into contact with basic acetate of lead,-the latter being prepared in this case either by mixing litharge and vinegar to the
consistence of a paste, or by exposing rolls of sheet lead to the
simultaneous action of vapors of vinegar and air, or by actually
dissolving an excess of litharge in vinegar. Sometimes the carbonic acid is made to act upon the subacetate at the very moment
when it is being formed, while at other times the acetate is prepared by itself, and subsequently treated with carbonic acid.
583. Silicate of lead is of interest from being an important
ingredient of flint glass; a certain proportion of it renders glass
lustrous and very beautiful. Such glass is, however, soft, easily
fusible, and incapable of bearing sudden changes of temperature;
it is, moreover, rather easily acted upon by alkalies. acids, and
other chemical agents, and is hence comparatively useless in the
chemical laboratory.
584. In many points of chemical behavior the compounds of
lead resemble more or less clearly the corresponding compounds
of barium, strontium, and calcium. Its compounds are moreover
isomorphous with those of the metals in question, and its atom,
like the atoms of these metals, is bivalent. Lead is therefore
classed as a member of the calcium group, though, as is the case
with fluorine in the chlorine group, it differs in some respects
from the other members of the family.  The specific gravity of
lead is 11'4, and its atomic weight 207. The specific gravity of
carbonate of lead is 6'5, and that of sulphate of lead is 6'2.




M., AGNEITUM.                   501
CHAPTER XXIX.
AG NE SI U  -Z IN C-C A D M IU.
MA GNESIUM.
585. This metal, or rather its oxide, was formerly classed with
the group which comprises the alkaline earths, but it is now
known to be more closely connected with zinc and cadmium than
with any other of the elements. It is found widely diffused, and
rather abundantly, in nature. The bitter taste of sea-water and
of some mineral waters is due to the presence of magnesium salts,
while silicate of magnesium and carbonate of magnesium are contained in a variety of minerals and in such common rocks as
dolomite, serpentine, soapstone, and tale.
586. 3Metallic magnesium  may be prepared by heating anhydrous chloride of magnesium with sodium in a crucible of porcelain or platinum, and subsequently dissolving out in cold water
the chloride of sodium which results from the reaction.  Miagnesium is a lustrous metal, as white as tin; its specific gravity is
1'75, and its atomic weight 24. It does not tarnish in dry air,
though in damp air it soon becomes covered with a film of hydrate of magnesium. It melts at a low red heat, and volatilizes
at higher temperatures; it may be readily distilled at a bright
red heat. When heated strongly in the air it takes fire and
burns with a bluish-white light of great brilliancy and high
actinic power. The metal is employed by photographers for
illuminating caverns and other places into which sunlight cannot
penetrate, and in cloudy weather it is even used by them as a
substitute for daylight. The metal can be pressed into wire or
into thin ribbons. and a considerable quantity of it is now used
in both these forms for purposes of illumination, as above stated.
Magnesium lanterns are much used in theatres for illuminating
scenery and tableaux. The white light has the advantage of
showing colors just as they look by daylight.:For scenic effects
the light may be modified by transmission through colored glass.
Magnesium is only slowly acted upon by cold water, but is




502                 OXIDE OF MAGNESIUM.
rapidly oxidized by hot water and by water acidulated with
almost any acid; oxide of magnesium is formed and hydrogen
set free.
587. Oxide of Magnesium (MgO).-There is but one compound of magnesium  and oxygen; it is obtained as a white
amorphous powder when magnesium is burnt in the air, or when
carbonate, chloride, or nitrate of magnesium is ignited.
Exp. 303.-Roll 10 or 12 c.m. of magnesium wire or thin ribbon
into a coil around a small pencil; withdraw the pencil and place in
its stead a piece of iron wire or a knitting-needle; holding this wire
horizontally, apply a lighted match to the end of the magnesium coil;
the magnesium will burn to the white oxide, which coheres in an imperfect coil, clinging to the iron wire. A portion of the oxide goes off
as white smoke. The magnesium wire for this experiment may be
procured at toy-shops as well as of dealers in fine chemicals.
The oxide is tasteless and odorless; it is soluble to a very
slight extent in water, and the solution has an alkaline reaction.
The specific gravity of the solid oxide, or magnesia, as it is often
called, varies from 3'07 to 3'2 as ordinarily prepared; but on
being very strongly ignited it becomes denser, and samples have
been prepared in this way of specific gravity as high as 3'61.
The light powdery oxide of magnesium known as "calcined
magnesia," which is prepared by gentle but prolonged ignition
of the hydrated carbonate, differs materially in several particulars from the more compact oxide obtained by calcining nitrate
or chloride of magnesium at high temperatures, or by intensely
heating the powdery oxide. Common calcined magnesia is, for
example, readily soluble in acids; but after the oxide has been
exposed to very high temperatures it dissolves but slowly even in
the strongest acids. Similar differences between the products obtained at high and at low temperatures are met with among the
oxides of almost all the metals hereafter to be studied.
588. A compact variety of oxide of magnesium, obtained by
heating the nitrate or chloride to bright redness, but no higher,
exhibits remarkable hydraulic properties.  On being wet it
quickly combines with a portion of water, and is converted into
a crystallized hydrate of compact texture, harder than marble,
and of great durability.




HYDRAULIC MAGNESIA.                503
A mixture of equal parts of the hydraulic magnesia and of
chalk, or powdered marble, made into paste with water, yields
a slightly plastic mass, which admits of being readily pressed
into any desired shape; if the moulded material be then placed
in water it will become, after some time, extremely hard and
compact (~ 591).
Oxide of magnesium is altogether infusible at temperatures
short of that of the oxyhydrogen flame. Very excellent crucibles for scientific purposes are prepared by compressing oxide
of magnesium into suitable forms. These crucibles undergo far
less change in the air than those made from lime; and like the
lime crucibles they do not unite with oxides of iron and the other
metallic oxides to form the fusible slags or glasses which are so
annoying in the ordinary crucibles, of which silicic acid is an
essential component.
589. Chloride of M2kagnesium (MgCl2) is found in sea-water
and in many saline springs. It is formed when magnesium
is burnt in chlorine gas, when a current of chlorine is passed
over a red-hot mixture of charcoal and oxide of magnesium,
and, in combination with water, by dissolving oxide of magnesium in chlorhydric acid. It is remarkable that the hydrated
chloride last mentioned cannot be made anhydrous by evaporation
and ignition without some decomposition of the chloride; oxide
of magnesium is formed and chlorine goes off in combination
with hydrogen as chlorhydric acid.
590. Sulphate of Magnesium  (MgSO,), or rather the hydrated
compound (MgSO4+ 7HO20), is largely employed as a medicament
under the name of Epsom salts. It is obtained not only from
the mineral spring at Epsom, in England, and from various other
springs, but is also prepared from sea-water, ard by dissolving
the minerals serpentine (silicate of magnesium), magnesite (carbonate of magnesium), and dolomite (carbonate of magnesium
and of calcium), in sulphuric acid. Hydrated sulphate of magnesium is a colorless crystalline salt, readily soluble in water,
and possessing the peculiar bitter taste common to most of the
soluble magnesium compounds. It is often employed in laboratories as i..e source from which to prepare other magnesium
salts,




504                        ZINC.
591. Carbonate of Magnesium (MgCO,) is found as a mineral
in nature, and with due care may be prepared artificially. As
met with in commerce, however,-the magnesia alba of the
shops, prepared by mixing hot solutions of sulphate of magnesium
and carbonate of sodium,-it is mixed with varying proportions
of hydrate of magnesium.  This compound is employed as a
medicament.  A  compound of carbonate of magnesium  and
carbonate of calcium occurs abundantly in nature as the mineral
dolomite, constituting extensive beds in various regions.  Dolomite is much more slowly soluble in acids than true limestone,
but when heated with a dilute acid it effervesces readily. When
burnt, at temperatures so low that the carbonic acid shall be
expelled only from the magnesium salt, while the carbonate of
calcium remains unaltered, dolomite affords an hydraulic cement,
preferable in many respects to ordinary lime, The product of
the calcination "sets" rapidly under water, and is converted
into a hard compact stone (~ 588). Citrate of magnesium, a
preparation made from carbonate of magnesium and citric acid,
is also largely employed as a medicament.
Of the other salts of magnesium, none are of sufficient importance to be described in this manual. Most of them are
easily soluble in water; hence the insoluble. double phosphate
of magnesium and of ammonium (MgN11 PO4 + 611HO), obtained
by adding ordinary diphosphate of sodium  to a mixture of
ammonia-water and any magnesium  salt, is of importance to
the analyst, since by means of it magnesium may be separated
from  its solutions.  It should be obsrlvcd that the ready
solubility of sulphate of magnesium is in marked contrast with
the insolubility of the sulphates of the alkaline-earthy group
of metals.
ZINC.
592. Ores of zinc occur in considerable abundance in several
localities. The metal is extracted from the carbonate, oxide,
silicate, and sulphide.  The carbonate andl sulphide are first
roasted in order to convert them into oxides, and the oxide is
then reduced by means of hot charcoal, in earthen retorts or in
crucibles provided with iron delivery-tubes. Since m, tallic zinc




GRANULATED ZINC.                    505
is volatile at high temperatures, it distils over from the retorts
as fast as it is formed and is condensed in receivers.
Zinc is a bluish-white metal of crystalline texture, brittle at
the ordinary temperature, and also when heated above 200~, but
a temperature of about 1300 or 1400 it may easily be rolled out
or hammered into sheets. The metal melts at 425~ and boils at
a bright-red heat; in presence of air the red-hot metal takes fire
and burns with a brilliant bluish-white light and formation of a
dense cloud of white oxide of zinc.
Exp. 304. —Melt 200 or 300 grms. of metallic zinc in a small
Hessian crucible, or in an iron ladle, placed in an anthracite fire.
Remove the crucible from the fire by means of appropriate tongs
(Appendix, ~ 27), and pour its contents in a very fine stream into a
pail full of cold water, taking care to hold the crucible at a distance
of 5 or 6 feet above the pail. Replace the empty crucible in the
fire, in order that it may be ready for Exp. 305.
The small thin pieces of zinc which will be found in the pail when
the water is poured away, are known as grantlated or feathered zinc.
This process of granulation may be conveniently applied to any of the
other easily fusible metals, such as bismuth, lead, or tin, when they
are required in a finely divided condition.
Granulated zinc is much used in chemical laboratories, for a variety
of purposes, but particularly for preparing hydrogen (~ 50). In order
that it may be fit for this purpose, it is best to heat the melted metal
nearly to redness before pouring it into the water; for it has been
noticed that when zinc is melted at the lowest possible temperature
and then immediately poured into water, the granules obtained are
but slowly acted upon by dilute sulphuric acid, while another portion
of the same metal, heated nearly to redness, and then granulated, is
readily soluble in the acid. If the hot metal be poured upon a warm
iron plate, it will be found to be still more readily soluble in acids
than that which has been suddenly cooled by the water.
Exp. 305.-Dry 20 grins. of the granulated zinc of Exp. 304, and
muix it intimately in a mortar with 40 grms. of crude saltpetre; remove
the empty crucible of Exp. 304 from the fire, and place it in such
position that any fumes which may subsequently be evolved from it
shall be drawn into the chimney. By means of a spoon or ladle,
project into the red-hot crucible the mixture of zinc and aaltpetre,
taking care to stand away as far as possible from the crucible. The
metal will burn fiercely, at the expense of the oxygen in the saltpetre,
for the most part, though a portion of it will be volatilized by the




506 0bTHE GALVANIC CURRENT.
intense heat of combustion, and converted into oxide of zinc in the
air. The residue in the crucible is a soluble compound of oxide ol
zinc and potash, known as zincate of potassium.
If a strip of thin sheet zinc be held in the flame of the gas-lamp,
it can readily be burned to oxide. The experiment succeeds best
with zinc leaf, which instantly burns with a vivid flame and formation of floating flocks of the white oxide. In oxygen gas, zinc burns
with peculiar brilliancy.
Zinc is not much acted upon either by moist or dry air at the
ordinary temperature; but a fresh, bright surface of it, when
exposed in a moist atmosphere, soon tarnishes and becomes
covered with a thin film of basic carbonate of zinc, which adheres
closely to the metal, and protects it from further change. Owing
to this durability, the metal is much used in the form of sheets.
Sheet iron and iron wire are often covered with a protecting coating
of zinc by simple immersion in melted zinc, and are then said
(most improperly) to be galvanized.  The specific gravity of zinc
varies from 6'8 to 7'3; its atomic weight is 65.
593. Zinc is readily attacked and dissolved by acids, with
evolution of hydrogen in most instances.  The chemical action of
dilute acids upon zinc is a very common source of that peculiar
mode of force called a galvanic current. There are few, if any,
chemical reactions which cannot be made to produce electricity,
and, in general, the more powerful the chemical action, the more
powerful is the electrical action which results.
Exp. 306.-Solder a piece of stout copper wire to one end of a strip
of sheet zinc, 4 c.m. wide by 10 c.m. long. The soldering will be
readily effected by rubbing the zinc and the wire, in the vicinity of the
proposed place of contact, with a strong solution of chloride of zinc,
before applying the melted solder. In the same way, solder a similar
wire to a like strip of bright sheet copper. Place the strips of zinc
and copper in a tumbler filled with water, acidulated with 1-12th to
1-10th its volume of sulphuric acid, in such a way that the two strips
shall not touch each other either within or without the liquid. So
long as the wires coming from the strips of metal do not touch each
other, the copper remains quiescent, while the zinc is attacked, and
bubbles of gas rise from its surface; but if the two copper wires are
brought into close contact, by means of a binding-screw, or by the
application of solder, the following phenomena occur: —1st. Minute
bubbles of hydrogen gas will be evolved from the surface of the copper




CORRELATION OF FORCES.                507
plate.  2nd. The zinc dissolves more rapidly than before; at the
close of the experiment sulphate of zinc may be recovered from the
liquid in the beaker. 3rd. This transfer of the hydrogen from the zinc
to the copper instantly ceases if the contact between the wires is
destroyed. 4th. If the two wires be connected with the two ends of
the coil of wire which surrounds the magnetic needle of the common
galvanometer, the deflection of the suspended needle will demonstrate
the fact that an electric current is passing through the wires from one
plate of metal to the other.
This conversion of chemical force into electrical force is a striking illustration of the doctrine that all physical forces are correlated. The preceding experiment well illustrates the principle
on which a large class of batteries employed in telegraphing and
in electro-metallurgy are constructed and worked, except that
the corrosion of the zinc is generally hindered by coating it with
mercury. Artificial products, like metals, acids, and saline solutions, are used to supply all the chemical force which is immediately converted into and utilized as electrical force in the useful
arts. We have not yet succeeded in realizing as electricity any
considerable proportion of the prodigious chemical force which is
incessantly active in the common processes of combustion.
594. Zinc dissolves in hot solutions of the caustic alkalies as
well as in acids; hydrogen is given off and a zincate of the alkali
formed:Zn +  2NaIHO  =,Na2ZnO2 +  2H.
When immersed in the solution of a lead salt, such as the
nitrate or acetate, zinc dissolves and lead is deposited in the
metallic state:PbN2O6 +  Zn =  ZnN206 +  Pb.
Exp. 307.-Dissolve 10 grmins. of acetate of lead in 250 c. c. of water,
add a few drops of acetic acid in order to dissolve the cloudy precipitate of carbonate of lead, which is formed from the carbonic acid in
the water, pour the solution into a wide-mouthed bottle and suspend
mn it from the cork a strip of sheet zinc. The zinc will soon be covered
with a brilliant coating of crystalline spangles of metallic lead, and
this crystalline vegetation, as it were, will shoot out or grow even as
far as the sides of the bottle. In the course of 24 hours all the lead
will have been deposited from the solution, and the latter will contain
nothing but acetate of zinc. Under the conditions of this experiment,




508                    OXIDE OF ZINC.
and as a general rule, zinc is, chemically speaking, a stronger or more
basic element than lead; it is capable of displacing lead from its coma
pounds. The growth of lead, witnessed in this experiment, is frequently spoken of as a lead tree; the experiment is often performed in
chemical laboratories for the sake of the chemically pure lead which
it furnishes.
Many other metals besides lead may be thus thrown down by zinc,
and the zinc mray itself be replaced by other metallic precipitants. The
whole series of experiments of which the one here indicated may be
taken as the type, is interesting as illustrating the general law of the
replacement of metals one by another in atomic proportions, and from
the fact that by means of these experiments the atomic weioht of
various metals may readily be determined. For example, if in the
foregoing experiment the piece of zinc be weighed before and after its
immersion in the acetate of lead, and if the precipitated lead be also
weighed, it will be found that the weight of lead obtained is, to the
weight of zinc dissolved, very nearly as 207 is to 65, the atomic weights
of lead and zinc respectively. The atom of zinc dissolved has replaced
in the solution the atom of lead which was precipitated. By the exercise of care in the manipulation, by employing boiled water free from
carbonic acid so that the addition of acetic acid to the lead salt shall
be unnecessary, and by finally drying the lead in an atmosphere of hydrogen, a close approximation to the numb:ers above given can be obtained.
595. Oxicde of Zinc (ZnO).-Tike magnesium, zinc forms but
a single compound with oxygen.  This compound may be readily
obtained by burning the metal, or by igniting carbonate or hydrate of zinc. As thus prepared, oxide of zinc is an insoluble,
white, amorphous powder, which, under the name of zitnc white,
has of late years been largely employed as a white paint. It
lacks the opacity or covering-power of white lead (~ 582), but, on
the other hand, has no injurious action upon the health of the
workmen and does not blacken or become discolored when exposed
to the fumes of sulphydric acid.  When heated in a crucible,
oxide of zinc exhibits a yellow color, but it becomes white again
on cooling.  The oxide dissolves easily in acids, with formation
of salts of zinc.
596. Chloride of Zinc (ZnCl2), obtained by dissolving metallic
zinc in ehlorhydric acid, is a compound readily soluble in water;
it is somewhat extensively employed for preserving timber, ml0io




CADMIUM.                       509
as a disinfecting fluid. It is used also by tinmen as a wash to
cleanse the surfaces of tin-plate before soldering.
597. Sulphate of Zinc (ZnSO4) is one of the commonest of the
zinc salts.  The hydrated compound, ZnSO,+71H20, known as
white vitriol, is used to a certain extent in medicine. and for other
purposes in the arts.  The action of carbon upon sulphate of zinc
differs somewhat from its action upon the sulphates previously
studied.  When a dry mixture of the sulphate of zinc and charcoal is heated to dull redness, carbonic and sulphurous acids are
evolved in the proportion of two volumes of the former to one of
the latter gas, and pure oxide of zinc remains:2ZnSO4 +  C =  2ZnO  +  2S02 +  CO2.
It would be quite possible to obtain metallic zinc friom the sulphate in one operation by employing an excess of carbon, heating
the mixture gently at first until the sulphuric acid had all been
decomposed, and then urging the fire in order to obtain the temperature requisite for the reduction of the oxide of zinc dnd volatilization of the metal. But if the mixture of sulphate of zinc
and charcoal be quickly raised to a high temperature, then sulphide of the metal is formed and carbonic oxide set free:ZnSO4 + 4C = ZnS + 4CO.
Zinc forms several valuable alloys; brass is an alloy of zinc and
copper, and German silver is a brass whitened by the admixture
of a small proportion of nickel.
CADMIIUM.
598. Cadmium is a comparatively rare metal, found associated
with zinc in nature; it is remarkably similar to zinc in its chemical relations. In the process of obtaining zinc from its ores,
the small proportion of cadmium which these ores contain comes
over with the first products of the distillation, since cadmium is
more readily volatile than zinc.
Cadmium may be prepared either from this early distillate, or from
the residues obtained when metallic zinc is dissolved in chblorhydric
acid in the preparation of chloride of zinc for manufacturing purposes.
These residues always contain a quantity oflead, which next to cadmium
is the commonest impurity of commercial zinc; and if care has been




510                PROPERTIES OF CADMIUM.
taken to keep an excess of metallic zinc in the dissolving-vat, they will
contain also all the cadmium with which the zinc was contaminated.
From either of these sources, cadmium salts may be prepared by
dissolving the crude materials in dilute nitric acid, separating the lead'by means of sulphuric acid, and throwing down the cadmium with
sulphuretted hydrogen.  Sulphide of cadmium  is a bright-yellow
powder, insoluble in dilute acids, while sulphide of zinc is readily
soluble in acids. Once isolated, the sulphide of cadmium may be dissolved in boiling, concentrated chlorhydric acid; from the solution of
chloride of cadmium thus obtained, carbonate of cadmium may be precipitated, and from the carbonate any of the other cadmium compounds
can readily be prepared.
Metallic cadmium is of a white color tinged with blue; it is
lustrous and takes a fine polish, but gradually tarnishes upon the
surface when exposed to the air.  Its specific gravity varies
from 8 6 to 8'7. It melts and volatilizes at temperatures below
redness. Heated in the air it takes fire and burns to a brown
oxide.
When combined with other metals, such as lead or tin, cadmium
forms alloys of remarkable fusibility; in this respect it far surpasses bismuth (~ 359). The most fusible alloy yet made contains
cadmium, bismuth, tin, and lead; it melts at 630-65~.
599. Cadmium is a volatile substance, and the specific gravity
of its vapor has been experimentally determined to be 56-85; the
weight of a unit-volume of the vapor is 56'85 times the weight
of the same volume of hydrogen. Now we have seen that the
specific gravity of the elementary gases and of the vapors of the
elements included in the chlorine and sulphur groups are the
same as the atomic weights of these elements. On the other
hand, the specific gravities of the vapors of phosphorus and arsenic were twice the atomic weights of these elements. Cadmium
presents still a new relation between the least combining weight
and the unit-volume weight; for the specific gravity of its vapor,
56'85, is about one-half of 112, its accepted atomic weight. The
significance of this fact may be illustrated from its chloride.
Cadmium is bivalent, and forms the chloride CdCl2, containing, as
experiment has proved, 112'24 parts, by weight, of cadmium to
71 parts of chlorine; if the unit-volume weight of cadmium were
the same as its atomic weight, two unit-volumes of chloride of




THE MAGrNESIUM GROUP.               511
cadmium would contain one volume of cadmium and two volumes
of chlorine; but were it possible, by experiment, to resolve the
vapor of chloride of cadmium into its component vapors, it would
be found that two volumes of cadmium were therein combined
with two volumes of chlorine.
The atom of cadmium when converted into vapor occupies twice
as much space as the atom of oxygen, or hydrogen, or chlorine
does; and accordingly the product-volumes of its compounds are
packed with one volume more than the product-volumes of the
corresponding compounds of oxygen or any member of the sulphur
group.  Whether the bivalent metals in general resemble cadmium on the one hand or oxygen on the other, in regard to the
relation between their vapor-densities and their atomic weights,
is a point on which experiment has thus far thrown but little
light.  Mercury resembles cadmium; but it is certainly possible
that these two elements constitute an exception to some general
rule hereafter to be proved-a rule, for example, like that which
many chemists are inclined to accept in advance of proof, namely
that the combining weights and the unit-volume weights of the
elements are normally identical.
Cadmium is so soft that paper may be marked with it; but it is
flexible, malleable, and ductile. In dilute chlorhydric and sulphuric acids it dissolves with evolution of hydrogen, though less
readily than zinc, Its best solvent is nitric acid. It does not
dissolve in the caustic alkalies.
600. From the foregoing it is apparent that the members of
the group of metals now under consideration resemble one another
with respect to volatility and several other of their physical
properties, besides being very closely related in most of their
chemical characters. The order of progression is from magnesium
to cadmium, zinc and its compounds occupying always an intermediate position.  The specific gravities of the three metals
are-Mg=1-75, Zn= 71, Cd=8-6; and their atomic weights
are-Mg=24, Zn = 65, Cd=112.  3Iagnesium volatilizes at a
bright-red heat, cadmium at a low red heat, and zinc at temperatures between these extremes. Cadmium is very fusible, melting
at about 360~, zinc melts at 425~, and magnesium at a moderate
red heat. All of these metals are bivalent; each forms but one
oxide, sulphide, and chloride.




ALUMINuM.
CHAPTER XXX.
AL U  I NUM —GLUC r N U M-C H R 0 M I U M —M A N  A N E  E —
I R O  N —C  B A L T-N I C K EL —-U R AN I U M.
ALUMINUM.
601. Next to oxygen and silicon, aluminum is perhaps the
most abundant element upon the earth's surface. It is the most
abundant of all the metals, as much as a twelfth of the solid crust
of the globe being composed of it. It occurs in enormous quantities in combination with oxygen and silicon, in all the so-called
primitive rocks, and indeed in most other rocks and soils. It
is contained in clay, marl, and slate, as well as in feldspar, mica,
and many other common minerals.
Oxide of aluminum, chloride of aluminum, and many salts of
the metal may readily he prepared artificially from the native
minerals; they have long been known to chemists, and made use
of in the arts; but the metal itself is less readily obtainable. It is
but a few years since metallic aluminum has been prepared upon
a manufacturing scale. The metal is nowadays prepared by heating metallic sodium either with chloride or fluoride of aluminum,
or with a double chloride or fluoride of aluminum and sodium. It
is a bluish-white metal, of remarkable lightness. Its specific gravity, 2'56, is about the same as that of porcelain, and only about
a quarter of that of silver. The metal is malleable, ductile, and
tenacious, and may be beaten into thin sheets, like gold and
silver, and drawn into fine wire. It melts at a temperature lying
between the melting-points of zinc and silver, but is not volatile.
It conducts electricity much better than iron, and heat even better
than silver; after having been heated, it cools very slowly. It is
remarkably sonorous, a bar of it suspended by a wire rings with
a clear musical note on being struck.
Aluminum changes little in the air, even at a strong red heat;
but when melted in open crucibles a film of oxide forms upon it




ALLOYS OF ALUMINUM.                513
that hinders the operation of casting. It is not acted upon by
water at temperatures short of a white heat, so long as it is in
the ordinary compact condition. Sulphydric acid has no action
upon it. Nitric acid, whether dilute or concentrated, has no
action upon aluminum at the ordinary temperature, but when
boiling dissolves the metal slowly. Cold dilute sulphuric acid
has scarcely any action upon it; but it is easily soluble in chlorhydric acid, either dilute or concentrated, at all temperatures.
It is soluble also in aqueous solutions of caustic potash, soda, or
anmmonia. The vegetable acids, such as acetic and tartaric acids,
exert no perceptible action upon it. Although soluble with evolution of hydrogen in aqueous solutions of the fixed caustic alkalies, aluminum is not acted upon by fused hydrate of sodium, or
hydrate of potassium; nor is it even attacked by fused nitrate of
potassium, except at temperatures high enough to decompose the
nitre so completely that it gives off nitrogen; when this limit
is reached the aluminum is immediately oxidized with incandescence.
602. Aluminum unites readily with many of the metals to
form alloys, among which that of copper and aluminum, calledalumhinum bronze, promises to be of especial importance. Aluminum bronze, composed of 90 parts copper and 10 parts aluminum,
is exceedingly hard, very malleable, as tenacious as steel, of a
beautiful golden color, and susceptible of being highly polished.
603. By uniting with the non-metallic elements, aluminum
forms only one class of compounds, of which the oxide A120, may
be taken as the type. The atom Al is trivalent, or, in other
words, it is equivalent to three atoms of hydrogen, and of the
same value as one and a half atom of oxygen. Since it would
be inconvenient to employ fractional expressions in writing chemical formulae, as well as illogical to speak'of half atoms, it is
customary to'write the formula of oxide of aluminum A1203 as
above, and not A10,, as might perhaps at first sight seem best.
For the sake of consistency, the formula of the chloride is in
like manner written A12C16 and not A1C1,. Since it contains one
and a half atom of oxygen for each atom of aluminum, the oxide
is often called a Sesqui (one and a half) oxide.
If no other element analogous to aluminum were known, if
2L




514                 OXIDE OF LTLUMTNUM.
this metal were not intimately related to glucinum, iron, ehromium, and the other metals to be considered in the present
chapter, chemists might possibly have taken the atomic weight
of aluminum at one-third of its present value, namely at 9'1 instead of 27'4. The formula of oxide of aluminum would then
have been written A120, and that of chloride of aluminum A1C1,
corresponding respectively with the formulae of the oxides and
chlorides of the alkali-metals. But, as will appear directly from
the study of the other members of the aluminum family of metals,
and particularly from the isomorphism of their various compounds, the atomic weight 27'4, and the formule first given,
must be regarded as the most probable. The atomic weight 9'1,
and the formulae derived from it, are inadmissible, since there is
no analogy between the chemical properties of aluminum, whether
simple or compounded, and those of the alkali-metals.
604. Oxide of Alztminum (A1204), commonly called Alumina,
is found crystallized in nature as the mineral corundum. The
sapphire and the ruby are also composed of this oxide together
with a little oxide of iron. It may be prepared artificially by
oxidizing the metal, or by igniting the hydrate, or almost any
oxygen salt of aluminum. Though unalterable in oxygen so
long as it is compact, powdered aluminum and aluminum-leaf
burn brightly when heated to redness in the air, with formation
of oxide of aluminum. It has been found by careful experiments
that 53'3 parts of the metal unite with 46'69 parts of oxygen to
form 100 parts of the oxide. Now, since oxide of aluminum is
isomorphous with certain oxides of iron and of chromium, which
are known to be sesquioxides, and is capable of replacing these
oxides in any proportion in their compounds (~ 252), it is inferred
that oxide of aluminum is likewise a sesquioxide. Upon this
view the atomic weight of aluminum is directly derived from the
foregoing experimental data by the equation:46'69:  533        48:     548
wt. of 3 atoms  wt. of 2 atoms
of oxygen.    of aluminum.
605. Hydrate of Aluminulm (A12H606) may be obtained as
a gelatinous, flocculent precipitate, by adding ammonia-water
to the solution of an aluminum salt, such as common alum.




ALUMINATES.                  515
When dried at a moderate heat it forms a soft, friable mass,
which adheres strongly to the tongue like clay; when dried
still more thoroughly, it forms a hard, yellowish, translucent,
horn-like substance, and at a red heat gives off all its water.
The volume of the original precipitate contracts to an enormous
extent during the operation of drying; the bulk of the final
anhydrous oxide is exceedingly small as compared with that of
the moist hydrate from which it has-been derived.
606. Anhydrous alumina may be melted in the flame of the
oxyhydrogen blowpipe.  It is neither decomposable by heat
alone, nor can it be reduced by carbon or any of the more
common deoxidizing agents. At a white heat, potassium decomposes it partially, and an alloy of aluminum and potassium
is formed.  Oxide of aluminum is insoluble in water, and after
having been strongly heated it is scarcely at all acted upon by
acids, excepting concentrated boiling chlorhydric and nitric
acids. The crystallized native oxide is insoluble in all acids.
The anhydrous oxide is insoluble in solutions of the caustic
alkalies, but dissolves readily in water after having been fused
at a red heat with either hydrate or carbonate of sodium or of
potassium.  Hydrate of aluminum, on the contrary, though
insoluble in water, dissolves easily in acids and in solutions of
the fixed caustic alkalies. Alumina is in fact capable of acting
not only as a strong base, forming well-defined salts by uniting
with acids, but it plays the part of an acid as well (compare
~ 350), and combines with the alkalies and with other metallic
oxides to form salts known as aluminates. Aluminate of potassium  (K20,A120) and aluminate of sodium  (Na2O,A120,) are
substances somewhat extensively used in the arts; the mineral
spinel is an aluminate of magnesium (MgO,A20,,); and a native
aluminate of zinc (ZnO,A120,) is called gahnite by mineralogists.
Exp. 308.-Heat a small fragment of alum with water in a testtube until it has completely dissolved, pour half the solution into
another tube, and add to it, drop by drop, ammonia-water, until the
odor of ammonia persists after the mixture has been thoroughly
shaker., Hydrate of aluminum will be precipitated, in accordance
with the reaction:A123S04 + 6(NH4)HO = A1203,3H20 + 3(NIT4)2S04.
2 X 2




516                   MORDANTST-LAKES.
Pour two or three drops of the moist hydrate of aluminum into
another test-tube and cover them with ammonia-water; no clear
solution will be obtained, for hydrate of aluminum is but little soluble
in alrmonia-water.
Pour two or three drops of the moist hydrate of aluminum into
still another test-tube, and cover them with a solution of hydrate of
sodium; the precipitate will dissolve immediately; aluminate of
sodium is formed, and this salt is easily soluble.
Exp. 309.-Take another portion of the clear solution of alum
prepared in Exp. 308, and add to it, drop by drop, a dilute solution
of caustic soda. A precipitate will soon fall, as in Exp. 308, and
if no excess of hydrate of sodium were added over and above that
necessary to form sulphate of sodium with the sulphuric acid of
the alum, this precipitate would remain undissolved; but on adding
more of the soda solution the precipitate dissolves at once, with
formation of aluminate of sodium.
607. Hydrate of aluminum combines readily with many organic coloring-matters, forming compounds which are insoluble
in water. The fibre of cotton, when impregnated with alumina,
can be made to retain colors which the cotton itself has no power
to hold; hence the use of aluminum salts as mordants in dyeing.
Exp. 310.-Boil a few crushed granules of cochineal in water until
a considerable portion of their coloring-matter has been extracted;
add to the filtered solution an equal bulk of a solution of alum, and
to the mixture add ammonia-water. A colored precipitate, consisting
of hydrate of aluminum and of the coloring-matter of the cochineal,
will be thrown down; it is the substance called carmine-lake. Similar
precipitates may be prepared by substituting almost any other organic
coloring-matter for the cochineal of this experiment. Precipitates
thus formed by the union of a metallic oxide and a coloring-matter
are all classed as lakes.
608. Chloride of Aluminum (A12C16) may be prepared, in
the same way that the chlorides of boron and silicon are prepared (~~ 449, 470), by passing chlorine over a heated mixture
of alumina and carbon.  It is formed also when hot finely
divided aluminum  is brought into contact with chlorine -gas.
Hydrated chloride of aluminum  (A12C16.12,O0) can be made
very easily by dissolving hydrated oxide of aluminum in chlorhydric acid; but the anhydrous chloride cannot be prepared by
hLeating this hydrate, since a great part of the chlorine is expelled




SULPHATE OF ALUMINUM.               517
from it, together with the water, at a low heat. When obtained
in the dry way, however, chloride of aluminum is'readily volatile.
The anhydrous chloride, prepared by the reaction,
A1203 + 3C + 6C1 = A12C13 + 3CO,
previously described, is found condensed in the cold portions
of the tube in which the materials have been heated, apart from
the residue of undecomposed alumina and carbon. It occurs
either as a flocculent powder, or as a transparent wax-like mass
of crystalline texture. It is colorless when pure, very deliquescent,
and soluble in water. When large masses of it are heated to
dull redness, a portion of it liquefies, but at temperatures near
the melting-point it volatilizes rapidly. Unlike oxide of aluminum, it may be readily decomposed by sodium and potassium at a
heat below redness, metallic aluminum being set free.
Chloride of aluminum  combines readily with several of the
other metallic chlorides, forming compounds analogous to the
chloride of aluminum  and sodium  (2NaC1,  12C16), from which
metallic aluminum is commonly manufactured.
609. Sulphate of Aluminum (A1,23SO4) is a salt largely employed in the arts.  It is commonly prepared nowadays by
acting upon hot roasted clay with sulphuric acid. Clay is a
silicate of aluminum not very easily attacked by acids so long
as it remains in the native plastic condition, but after having
been exposed for some time to a dull red heat it readily yields
its alumina to acids. A solid mixture of sulphate of aluminum
and free silicic acid obtained as the product of this reaction is
known in commerce as alum-cake. By lixiviating alum-cake,
sulphate of aluminum  may readily be obtained in solution:
and from  this solution the salt crystallizes as a hydrate, the
composition  of which may be represented  by the formula
A123S04+ 181H20.
Until a comparatively recent period, sulphate of aluminum was
sent into commerce neither in its free state nor mixed with silica,
but in combination with sulphate of potassium in the form of
alum. Common alum is a hydrated double sulphate of aluminum
and potassium; its composition is represented by the formula
Ma2K24S04 + 24H20, or K20,SO3; A1203,3SO3 + 24H,0.
It crystallizes very easily in large, compact, well-defined octa



518                        ALUaM.
hedrons, belonging to the first or regular system.  The crystalline
character of alum is important, since it is solely on account of
this character that the salt has come into such general use.
Neither the sulphate of potassium nor the water in alum plays
any useful part in the reactions for which this salt is commonly
employed. Since 100 parts of alum contain only about 36 parts
of anhydrous sulphate of aluminum, it follows that 64 per cent.
of the alum is, for all chemical purposes, simply inert matter,
which has to be transported and manipulated for the sake of the
36 per cent. of real sulphate of aluminum. The reason why
this waste of labor and loss of the potassium salt is tolerated
is twofold: —Until a comparatively recent period, sulphate of
aluminum could be more readily purified by crystallization in
alum than in any other way. At the present time, when it is
easy to obtain pure sulphate of aluminum from responsible
manufacturers, alum is still prepared because its clean, sharply
defined crystals afford a valuable criterion of purity. So long as
it is left in the condition of crystals, alum cannot be adulterated
with any foreign substance.
Potash-alum is still the common alum of the American market,
but in Europe ammonia-alum  (A12(NIk4)24SO4, 24H20) is at
present largely employed; it is there prepared by adding sulphate of ammonium obtained from the ammoniacal liquor of the
gas-works to sulphate of aluminum resulting from the action of
sulphuric acid upon clay. Ammonia-alum crystallizes almost as
easily as potash-alum; but it is remarkable that the corresponding double sulphate of aluminum and of sodium (soda-alum) is
easily soluble in water, and crystallizes with comparative diffic-Alty; hence it has never come into commerce.
Sulphate of aluminum is employed as the source of the various
compounds of aluminum  used in dyeing, calico-printing, and
paper-making. Acetate of aluminum, for example, is largely
employed by dyers, particularly for the red colors obtained from
madder, under the name of red liquor.
Exp. 311.-Dissolve 3 grms. of sugar of lead in 4 c. c. of hot water;
also dissolve 4 grms. of common alum in 6 c. c. of hot water; mix the
hot solutions and filter off the insoluble sulphate of lead which is
formed. The solution obtained consists of basic acetate of aluminum,




SILICATES OF ALUMINUM.                519
together with some sulphate of aluminum and all the sulphate of potassium of the original alum. Such solutions are preferred in practice to
those containing normal acetate of aluminum, to prepare which a much
larger proportion of acetate of lead would be required than has been
given above.
Exp. 312.-Soak a small piece of cotton cloth in the solution of
acetate of aluminum prepared in Exp. 311, and another piece of similar cloth of equal size in pure water. Hang up both pieces to " age,"
best in a moist and warm atmosphere, for a day or two. During the
process of ageing, a portion of the acetic acid escapes from the salt on
the cloth, and there is left within and upon the fibres of the cloth a
quantity of hydrate of aluminum, or at least a mixture of highly basic
acetate and sulphate of aluminum. This deposit is the true mordant.
When cloth impregnated with it is soaked in a solution of coloringmatter, the coloring-matter unites with the alumina precisely as in
Exp. 310, and is thereby firmly attached to the cloth. It should be
mentioned that several other oxides, besides the oxide of aluminum,
are capable of acting as mordants; the sesquioxides of iron and of
chromium for example, as well as the binoxide of tin, are largely used
as mordants.
Exp. 313.-Place a quantity of a solution of extract of logwood in
two small evaporating-dishes, heat the liquor to 400 or 500, then place
the mordanted cloth in one dish, the ununordanted cloth in the other
and boil the liquor in both dishes. Continue to boil during 10 or 15
minutes, then take out the pieces of cloth and wash them thoroughly
in water.  It will appear that the coloring-matter remains firmly
attached to the mordanted cloth, while the cloth which has received
no mordant can readily be washed clean or flearly clean.
610. Silicates of Aluminum.-Of all the aluminum compounds
the silicates are by far the most important. Clay in all its varieties is a hydrated silicate of aluminum, usually mixed with
an excess of silica, besides other impurities derived from the rocks
from whose decomposition the clay itself has been formed.  The
purer kinds of clay, such as kaolin or porcelain clay, are products
of the decomposition of feldspar, a mineral composed of silicon,
aluminum, potassium, and oxygen in the proportions Al120,,K20,
tSiO2.  When exposed to the atmosphere, many varieties of
feldspar gradually decompose, an alkaline silicate is washed away,
and silicate of aluminum remains. Clay is remarkable on account of its plasticity when moist, of the facility with which it




520            BRICKS -EARTH ENWARE- GLAZES.
is converted into stone-like masses when strongly heated, and of
its infusibility when pure.
Earthenware, bricks, and ordinary pottery are made from common clay, by mixing the clay with water enough to form a plastic
paste, which is then moulded into any desired form, dried and
intensely ignited. The porous ware resulting from this operation
may be glazed, and so made impermeable to liquids, by coating it
over with some fusible substance, such for example as a mixture
of litharge and clay, and again heating it so intensely that the
coating shall melt to a glass, which either fills up the pores of the
clay, or at the least stops their openings. Porcelain proper, and
the better kinds of stoneware, are made from the purest varieties of
clay, and are glazed with feldspar. Common stoneware, such as
is used for jugs, beer-bottles, and the like, is covered with the socalled salt-glaze: —Moist chloride of sodium is thrown into the
kiln in which the ware is baking, and being volatilized by the
intense heat, comes in contact with the hot stoneware; decomposition ensues; the water and the chloride of sodium are both
decomposed, silicate of sodium is formed, and by mixing with the
silicate of aluminum, forms a smooth hard glaze upon the surface
of the ware.
For all vessels which are to be employed for chemical or culinary purposes the hard and durable salt-glaze is very much to be
preferred to the lead-glaze prepared from litharge and clay; for the
lead-glaze is readily acted upon by many chemical agents, and is
liable to impart its poisonous properties to articles of food which
have been left in contact with it.
Fire-bricks, crucibles, and similar refractory articles fitted to
support very high temperatures without undergoing fusion, are
prepared from pure varieties of clay, free from iron, lime, or magnesia, but containing an unusually large proportion of silica.
Some varieties of fire-clay contain as ]nuch silica as is represented
by the formula Al2O3,6SiO2, while the composition of many of
the common clays may be approximately represented by the formula A120,,3SiO2, or better by the formula Al203,2SiO2. Ill the
manufacture of fire-bricks, and of many varieties of potters' ware.
it is usual to incorporate with the original clay a certain proportion of foreign matter which prevents the moulded article




HYDRAULIC CEMENT.                  521
from shrinking too much as it dries, and from cracking. In firebricks the coarse powder obtained by pulverizing old fire-bricks
is employed for this purpose; in Hessian crucibles it is very easy
to detect numerous grains of quartz sand; and, in general, finely
powdered flint or quartz, as well as previously baked clay, is
used for the same purpose in many varieties of pottery.
611. Silicate of aluminum is moreover a very important ingredient of the common hydraulic cement employed to replace lime
mortar in constructions exposed to the action of water. It has
been found that by carefully burning some varieties of impure
limestone, containing from 10 or 12 to 30 or 35 per cent. of clay,
and mixing the product with water, there is obtained, in place of
ordinary mortar, a cement capable of " setting" or hardening,
even under water, to a compact stone. Hydraulic cements may
readily be prepared artificially by mixing with quicklime a
suitable proportion of roasted clay, or by heating mixtures of clay
and limestone; in fact, some of the best cements now in use
are artificial. A porous volcanic stone called pozzolana, from
the vicinity of Naples, consisting of silicates of aluminum, calcium,
and sodium, was much used by the Romans to the same end.
When powdered and mixed with ordinary lime the pozzolana
yields an excellent hydraulic mortar. Tn many Roman ruins it
may be seen to-day far more perfectly preserved than the bricks
which it cements.
When treated with water hydraulic limes simply absorb the
water and form a slightly plastic paste without greatly increasing
in bulk; they do not slake or evolve much heat like ordinary
quicklime.  The moist paste soon begins to set, and is then
ready for application. In order that the cement may harden properly under water it should not be submerged before it has begun
to set; it should in any event be kept moist until it has become
hard; otherwise it is liable to remain loose and porous.
The solidification of hydraulic limes appears to depend upon
the formation of insoluble hydrated compounds of lime with silicic
acid and alumina. Cements which contain from 25 to 35 per
cent. of clay solidify in the course of a few hours; but those in
which the proportion of clay is no more than 10 or 12 per cent.
become hard only after the lapse of several weeks.




522                     CHROMIUM.
A mixture of hydraulic cement with coarse gravel constitutes
the material known as concrete, employed for the foundations of
buildings, and by. the ancients for walls which have proved to
be of great durability; this mixture soon concretes or hardens
into a firm mass, well nigh impermeable to water. The influence
of magnesia in the preparation of hydraulic mortars has already
been indicated, in ~ 588.
GLJCINUIM.
612. Glucinum is a rather rare metal, found, together with
aluminum, in the emerald, in beryl, and a few other minerals.
It closely resembles aluminum in its physical properties, and
forms compounds analogous in composition to those of aluminum,
and of similar chemical deportment. Like aluminum, metallic
glucinum may be reduced from its chloride by means of sodium
or potassium. There is but a single oxide of glucinum, G120,, and
a single chloride, GI2C16. The salts of glucinum have a sweet
taste, whence the name, from a Greek word meaning sweet. The
atomic weight of glucinum is 14, and its specific gravity 2'1.
CHROMIUM.
613. Chromium is nowhere found in very large quantities, nor
is it very widely disseminated in small portions like iodine and
fluorine, but it is nevertheless found in sufficient abundance to
admit of its compounds being rather extensively employed in the
arts. The chief ore of chromium is a compound of oxide of chromium and oxide of iron (FeCr204) called chrome iron-ore. Metallic chromium may be reduced from its oxide by means of
intensely heated charcoal, and from its chloride by means of
sodium, potassium, magnesium, or zinc; but it has as yet been
little studied. Its specific gravity is about 7, and its atomic
weight 52'5.
614. Oxides of Chromium. -Chromium  forms three welldefined oxides: —a protoxide, CrO; a sesquioxide, Cr203; and a
teroxide, CrO3, called chromic acid. Besides these there is a compound of the protoxide and sesquioxide (Cr,30,=CrO, Cr20,),
another of the sesquioxide and teroxide (Cr20,,CrO3=3CrO2), and
an ill-defined compound containing more oxygen than chromic




OXIDES OF CHROMIUM.               523
acid, usually spoken of as perchromic acid (Cr2O,). Both the
protoxide and the sesquioxide are bases, corresponding respectively
to oxide of magnesium and oxide of aluminum; but the protoxide
and all its compounds rapidly absorb oxygen from the air, with
formation of the sesquioxide. On account of this instability, they
are rarely prepared, and are mainly interesting from their analogy to the compounds of the protoxides of manganese, iron,
cobalt, and nickel, hereafter to be studied. The sesquioxide, on
the other hand, is a stable compound, closely resembling oxide of
aluminum. Chromic acid is a strong, well-characterized acid,
which combines with bases to form a great number of salts.
The most important of the chromium compounds is the bichromate of potassium; this salt is readily procurable in commerce,
and is the source from which all the other compounds of chromium are commonly derived. Bichromate of potassium is itself
prepared by heating finely powdered chrome iron-ore with carbonate and nitrate of potassium in a reverberatory furnace. The
sesquioxide of chromium of the ore is oxidized and converted into
the teroxide, chromic acid, which displaces the carbonic acid of
the carbonate of potassium.
615. Sesquioxide of Chromium (Cr20,). —Compounds of this
oxide are more commonly met with than any other of the chromium salts, except the salts of chromic acid. As has been stated
already, compounds of the protoxide must be regarded merely as
chemical curiosities. By adding ammonia-water to the solution
of a salt containing the sesquioxide, a bulky green precipitate
of hydrated sesquioxide of chromium is thrown down, which,
when collected and ignited, leaves the anhydrous oxide as a brightgreen powder, unchangeable at the highest furnace-heat. It is
employed in the decoration of porcelain, and is a valued pigment
much used in painting and printing, under the name chrome green.
616. Chlorides of Chromium.-There are two of these compounds, the protochloride (CrC12) and the sesquichloride (Cr2Cl6).
The latter compound is the more important, and is the substance
usually meant when chloride of chromium is spoken of. Hydrated sesquichloride of chromium, obtained by dissolving the
hydrated sesquioxide in chlorhydric acid, is the chloride most
commonly met with.




524                     CHROMIUM SALTS.
617. Sulphate of Chromium (Cr203,,3S03,) is sometimes prepared
in the pure state; but, like sulphate of aluminum, it ordinarily
occurs in combination with sulphate of potassium or sulphate of
ammonium, as a double salt, called chrome alum.  Chrome alum
is a compound of a beautiful violet color, crystallizing in welldefined octahedrons of the same form as the crystals of ordinary
alum; its composition also corresponds to that of common alum,
the formula of the chromium salt being Cr2K24S04 + 24Ht0, or
K20,S,0; Cr203,3S03 + 24H20.
Exp. 314.-Dissolve 15 grms. of powdered bichromate of potassium
in 100 c. c. of warm water; cool the solution, and then add to it 25
grms. of concentrated sulphuric acid; cool the liquor again, and pour
it into a porcelain dish, surrounded with cold water; slowly stir into
the mixture 6 grms. of alcohol, and set the whole aside. In the course
of 24 hours the bottom of the dish will become covered with well-defined octahedral crystals of chrome alum.
The alcohol in this experiment deprives the chromic acid, of the
bichromate of potassium, of half its oxygen, and is itself converted for
the most part into acetic acid and water.
618. It is remarkable that the salts of sesquioxide of chromium,
as well as the oxide itself, occur in two isomeric conditions. One
modification is known as the green, the other as the violet modification. As a rule, the violet compounds crystallize readily,
while the green compounds do not. In the preparation of chrome
alum it is important to guard against the formation of a green,
soluble sulphate of chromium, which does not crystallize. In
general, if the solution of a salt of the violet modification is heated
nearly to boiling, the salt passes into the green modification and
becomes uncrystallisable. Hydrated sesquioxide of chromium, as
obtained by adding a caustic alkali to the cold solution of a chromium salt of either modification, is readily soluble both in acids
and in cold solutions of caustic soda or potash; but on boiling
the green alkaline solution, all of the chromium is precipitated
as a hydrate of the green modification.
Exp. 315.-Place in a test-tube a few drops of a dilute solution of
chrome alum, or of some other salt of sesquioxide of chromium; add.
drop by drop, a solution of hydrate of sodium, until the precipitate
which forms at first is completely redissolved. Boil the clear solution, and observe that the precipitate which again forms in the liquor
is no longer soluble in alkalies.




CHROMIC ACID.                     525
By means of this reaction, oxide of chromium may readily be separated from oxide of aluminum; for, as has been seen in Exp. 309, alumina is readily soluble in alkaline solutions, and is not precipitated
therefrom by boiling.
619. Chromic Acid (CrO3) may be obtained by decomposing
bichromate of potassium with sulphuric acid.
Exp. 316.-Mix 40 c. c. of a cold, saturated, aqueous solution of bichromate of potassium with 50 c. c. of oil of vitriol, in a small beaker
standing in cold water, and observe that chromic acid is deposited in
crystalline needles. It is remarkable that in sulphuric acid of 1'55
specific gravity, such as is obtained in the foregoing mixture, chromic
acid is well nigh insoluble, though it is readily soluble both in water
and in strong sulphuric acid. Cover the beaker, and set it aside for
some hours; finally pour off the supernatant liquor with care, scrape
out the chromic acid with a glass rod, and place it upon a dry, porous
brick, under an inverted bottle, in order that the sulphuric acid which
adheres to it may be absorbed. Preserve the dry crystals in a glassstoppered bottle. Chromic acid deliquesces rapidly when exposed to
the air. It is easily brought to the condition of sdsquioxide of chromium, both by heat and by reducing agents, and is hence an oxidizing
agent of considerable power.
Exp. 317.-Shake up in a small bottle enough strong alcohol to
moisten its sides; then throw in half a gramme or less of chromic
acid; a portion of the alcohol will be oxidized so quickly, and with
evolution of so much heat, that the remainder will take fire and burn
in the air.
Several of the salts of chromic acid, as well as the acid itself,
are employed as oxidizing agents. A mixture of bichromate of
potassium and of sulphuric acid, for example, is employed for
bleaching certain fats. From the chromates, both oxygen and
chlorine may be conveniently prepared.
Exp. 318.-Heat a mixture of 6 grmins. of powdered bichromate of
potassium and 9 grins. of concentrated sulphuric acid in a small flask,
provided with a delivery-tube leading to the water-pan, and collect
the oxygen which is freely evolved:K20,2CrO3 + 4(H20,S03) = Cr203,3S03 + K20,SO3 + 4H2O + 30.
Exp. 319.-Place a mixture of 1 grm. of powdered bichromate of
potassium and 6 grms. of chlorhydric acid of 1.16 specific gravity, in
a flask provided with a delivery-tube, as in Exp. 318. Heat the flask
gently for a few seconds until its contents begin to react upon one




526                      CHEOMATES.
another; then quickly remove the lamp and attend to the collection
of the chlorine, which will continue to be evolved without further
heating:K20,2CrO3 +  14HC1 = Cr2C1l + 2KC1 + 71120 + 6C1.
620. Chromates.-As has been already indicated, bichromate
of potassium is the commonest and the most important salt of
chromic acid. It is the material from which most of the other
compounds of chromium are prepared, and is itself important in
dyeing and calico-printing. It has of late years been used in the
art of photolithography.
When a mixture of gelatine and bichromate of potassium is exposed
to light, the chromic acid is reduced, and an insoluble compound of
gelatine and sesquioxide of chromium is formed. In practice, albumenized paper is covered in a dark room with a mixed solution of
bichromate of potassium and gelatine, then dried, pressed smooth, and
kept always in the dark until wanted for use. If a sheet of this prepared paper be placed beneath a negative photographic picture (obtained in the usual way) and exposed to light for a short time, the
chromic acid will be reduced in such wise that a positive picture will
be obtained upon the gelatine paper. In this positive, as taken from
the press, the parts acted upon by the light will be brown, while the
other portions of the sheet retain their original yellow color. The
positive is then washed with water in such manner that'the unchanged
portions of gelatine and of bichromate are dissolved away, and an insoluble, clearly defined inipression of the original picture is left upon
the paper. By means of pressure, the design is then transferred to
the lithographic stone, and from the stone any desired number of
copies may be printed upon paper with ink in the usual way.
Besides the bichromate of potassium, there are several other
chromates important in the arts or useful to the analyst. The
normal chromate of potassium (K20,CrO,) is a yellow salt, readily
obtainable by adding a molecule of carbonate of potassium to one
of the bichromate:K20,2CrO3 + K20,CO2 = 2(K20,CrO3) + CO2.
It is isomorphous with normal sulphate of potassium (K2S04),
chromic acid, like sulphuric acid, being bibasic (~ 238). The
salt is hence easily adulterated. Chromate of barium is insoluble
in water and in acetic acid; chromate of strontium is soluble in
acetic acid, though nearly insoluble in water; while chromate of




IMANGANESI.                    527
calcium is soluble both in water and in acetic acid; hence an
easy method of separating compounds of either of the three
metals from mixtures which contain compounds of all three.
Chromate of lead is the pigment called chrome-yellow; it may
easily be prepared by mixing solutions of bichromate of potassium
and acetate of lead. An orange-colored dichromate (2PbO,CrO3)
may be obtained by boiling together yellow chromate of lead and
slaked lime in the proportion of two molecules of the former to
one of the latter. This process is used to fix a permanent orange
upon calico. A still more brilliant color may be obtained by
fusing one part of the yellow chromate of lead with five parts of
nitre; chromate of potassium and dichromate of lead are formed,
and the former may be washed away. Chromate of mercury, of
a brick-red color, may be precipitated by adding bichromate of
potassium to nitrate of protoxide of mercury, or, of an orangeyellow color, by adding the potassium salt to the nitrate of dioxide of mercury.
MANGANESE.
621. Black oxide of manganese, such as has been employed
in the preparation of oxygen and of chlorine (~~ 14, 105), is a
tolerably abundant mineral. Small quantities of manganese exist
also in a great number of other minerals and rocks; so that the
element is really very widely diffused in nature. It is often
associated with ores of iron. By heating oxide of manganese
very strongly with charcoal, it may be reduced to the metallic
state, though not readily. The metal is of a grayish-white color,
and is very hard and brittle.  It oxidizes quickly when exposed
to the atmosphere; it melts only at the strongest heat of a blast
furnace.  The specific gravity of manganese is 8, its atomic
weight is 55. It slowly decomposes water at the ordinary temperature, and dissol, cs readily in dilute sulphuric acid with evolution of hydrogen.  Like iron, it combines with carbon and
silicon. Metallic manganese is not used in the arts; and the
alloys which it forms with the other metals are of no commercial
importance, except that a small proportion of manganese is present in a peculiar kind of iron largely used for making steel.
622. Oxides of Manganese.-Six well-defined compounds of




528                 OXIDES OF MANGANESE.
oxygen and manganese are known; two of them are bases, two
are acids, and two may be regarded as salts, formed by the union
of the oxides one with the other.  Protoxide of manganese
(MnO) is a powerful base, while sesquioxide of manganese (Mn03,)
is but a weak base.  31anganic acid (3lIn03) and permanganic
acid (MIn207) are well characterized as acids, though they are
known only in combination; they have never been obtained in
the free anhydrous state.' On the other hand, binoxide of manganese (AMn,O3,1MnO3=33MnO2) and the red oxide (MnO,Mn2O,
= Mn304) are both neutral or indifferent bodies; they exhibit
neither acid nor basic properties.
623. Protoxide of Manganese (MnO) may be obtained by heating carbonate of manganese out of contact with the air, or by
heating either of the higher oxides of manganese to redness in
contact with charcoal or hydrogen. The protoxide is itself reduced to the metallic state by these agents only at a white heat.
It unites freely with acids to form salts of considerable stability.
The crystallized sulphate 3MnSO4+ 51120 and the chloride M-nCl2
+ 4H110 are commonly employed in the laboratory.  Both of
them may be prepared from the residues obtained in the preparation of chlorine and oxygen (~~ 105, 626). Hydrated protoxide of manganese may be precipitated from the chloride as
follows:-.Erp. 320.-Dissolve a small crystal of chloride of manganese in
water; add to the solution soda-lye until the liquor exhibits a distinct
alkaline reaction when tested with litmus-paper. Collect the gelatinous white precipitate upon a filter, and observe that it soon becomes
brown as it absorbs oxygen from the air; the brown product is sesquioxide of manganese.
Exp. 321.-Heat a portion of the precipitated hydrate of Exp. 320
to redness upon a fragment of porcelain; it will slowly absorb oxygen,
and change to the deep-brown-colored sesquioxide.
Exp. 322.-To a solution of chloride of manganese, such as was prepared in Exp. 320, add a few drops of sulphydrate of ammonium (~ 526).
A flesh-colored precipitate of sulphide of manganese (MnS) will fall
down. Like the hydrate above described, this precipitate soon becomes
brown by exposure to the air. It is often prepared by the analyst
when testing for manganese.
624. Sesquioxide of Manganese (Mn2O,) occurs -in nature is




SESQUIOXIDE OF MANGANESE.              529
the minerals braunite and manganite. It is prepared artificially
by roasting the protoxide obtained from chlorine-residues, and
is itself used to a considerable extent in the preparation of
chlorine:Mn20, +  6HC1 =  2MnC12 +  3H20 +  2C1.
It combines with acids to form a series of unstable salts analogous to the sesquisalts of iron, though far less permanent.
A solution of the sesquisulphate, for example, Mn20a,3S03, is
reduced to the condition of protosulphate by mere boiling. In
like manner the sesquichloride Mn2Cl6 is doubtless formed when
the protochloride is treated with cold chlorine, or the sesquioxide
is digested in cold chlorhydric acid; but the salt is decomposed
with extreme readiness, and splits up into free chlorine and the
protochloride even when but slightly heated. In the preparation
of chlorine from the sesquioxide as above formulated, there is no
doubt an intermediate reaction,
Mn2O, +  6HC1 =  Mn,,l. +  3HaO,
before the final breaking up of
Mn2Cl6 into 2MnCl, +  2C1
625. Of the salts of the sesquioxide, the double compound
of sulphate of manganese and of potassium, known as manganese alum, is one of the most interesting; it is of analogous
composition to ordinary aluminum alum, and is isomorphous
with this body, as it is with the corresponding alums of iron
and chromium.  The series of double salts known as alums,
admirably illustrates the relationship of the several members
of the group of metals now under discussion, and the law of
isomorphism as well. It is interesting to observe, moreover,
that the name alum, originally applied specifically to the compound of sulphate of aluminum and of potassium, has. with
the growth of chemical knowledge come to have a generic signification.  Several salts are now classed as alums, into the
composition of which neither aluminum nor potassium enters.
The following list enumerates some of the best-known. potas
sium alums:
Common alum  K= R2S0AI3S0~ +  24H20,.
Chrome alum- =  K S0,Cr23S04 +  24H20;,
2 )s




530                       ALTMS.
Manganese alum  =  K2S04,Mn23SO4 +  24H0,
Iron alum  =  K2S04,Fe35SO4 +  24H20.
But as has been stated in ~ 609, the potassium in these compounds may be replaced by any metal isomorphous with putassium. There are ammonium alums and sodium alums corresponding to each of the potassium alums above enumerated,
and there is evidence that potassium may be replaced in these
alums by the rarer alkali-metals. Some alums, on the other
hand, are composed of mixtures in various proportions of alkalimetals, and of the metals capable of forming sesquioxides. Besides
these true alums, there are allied bodies which contain no alkalimetal whatsoever; such, for example, are the following:Aluminum iron alum  =  FeS04,A123S04 +  241120,
Aluminum magnesium alum  =  MgSO4,Al23SO4 +  24H20,
Aluminum manganese alum  =  MnS04,A123SO4 +  24H20,
but these affiliated alums do not crystallize in the octahedral
form which is characteristic of the alums proper.
626. Binoxide of Manganese (lnO2) is a black compound
found abundantly in nature, and largely employed in the arts for
the purpose of evolving chlorine from chloride of sodium or
chlorhydric acid (~ 105), as well as for decolorizing glass. It
may readily be prepared artificially from the lower oxides by
the action of oxidizing agents. By itself, at the ordinary temperature, binoxide of manganese is an inert chemical substance,
though at higher temperatures it has considerable oxidizing-power.
At a strong red heat it gives off one-third of its oxygen:3MnO2 = Mn304 +  20.
Formerly oxygen was often prepared in chemical laboratories
by heating the black oxide of manganese in iron retorts; but the
process has long been superseded by more convenient methods.
The oxide, Mn304=MnO,Mn20,, which is left as a residue in
this experiment, corresponds in composition with the magnetic
oxide of iron, an important ore of iron. This oxide is the most
easily obtained by artificial means of all the oxides of manganese;
it is produced when the protoxide or its nitrate or carbonate is
strongly heated in the air, or when either of the higher oxides is
intensely ignited.




MANGANTC ACID.                    531
Black oxide of manganese is insoluble in nitric acid, but is
decomposed by strong hot chlorhydric acid, with formation of
protochloride of manganese and free chlorine, as has been already
explained (~ 105), and by hot concentrated sulphuric acid with
evolution of oxygen:MnO, +  H2SO4 =  MnSO4 +  H20 +  0.
Exp. 323.-In a small glass flask, provided with a suitable deliverytube, heat a mixture of 15 grms. of powdered black oxide of manganese
and 10 grms. of concentrated sulphuric acid, and collect the gas over
water in the usual way.
After all the available oxygen has been obtained in this experiment,
and the flask, together with its contents, has been allowed to cool,
pour 15 or 20 c. c. of water into the flask, boil the mixture; pour it
upon a filter, and evaporate the filtrate to dryness upon a water-bath,
taking care to stir it constantly when nearly dry. Hydrated sulphate
of manganese (MnSO4,4H20) will be obtained as a reddish-white
powder.
EExp. 324.-For the sake of comparing the old process of making
oxygen with methods now in use, charge an ignition-tube, such as
was used in Exp. 7, to one-third of its capacity, with black oxide of
manganese, connect it with the water-pan in the usual way, heat it
strongly over the gas-lamp, and observe the comparatively slow rate
at which oxygen is evolved from it.
627. Manganic Acid (MnO,) has not yet been obtained in the
free state; it is known only as it occurs in combination with potash
or some other base. Of the manganates, those of potassium,
sodium, and barium are the best-known; they are isomorphous
with the corresponding chromates, sulphates, and seleniates. The
alkaline manganates are important compounds to the analyst.
Exp. 325.-Place upon a piece of platinum-foil as much dry carbonate of sodium as could be held upon half a pea; mix with it an
equal quantity of powdered nitrate of potassium and a bit of binoxide
of manganese as large as the head of a small pin. Fuse the mixture
in the outer blowpipe-flame, and observe the bluish-green-colored
manganate of sodium which is produced.
Exp. 326.-Melt together in an iron ladle over an anthracite or
charcoal fire, 10 grms. of hydrate of potassium, and 7 grms. of chlorate of potassium; stir into the pasty liquid 8 grms. of very finely
powdered black oxide of manganese, and maintain the mixture for
a short time at a temperature just below visible redness, taking care
2 x 2




532                 CHAMELEON MINERAL.
to stir it frequently with an iron rod. When the crumbly mass has
become cold, place some of it in a test-tube with a small quantity of cold
water and shake the tube. As soon as the solid particles have settled,
there will be seen a clear green liquid, which is a solution of manganate of potassium.
Exp. 327.-Pour off half of the green solution of manganate of
potassium into another short test-tube, and leave it open to the air;
the green color of the solution will gradually change to blue, then
to violet and to purple, and finally to ruby red. The red color is
that of a solution of permanganate of potassium, into which the manganate is converted by exposure to the air. The intermediate colors
are merely mixtures of the manganate green and the permanganate
crimson. On account of these remarkable changes of color, the name
chameleon mineral has been applied to manganate of potassium, and
by this term it is still commonly known.
Manganate of potassium is a very unstable salt, especially when in
solution; it may be readily decomposed in a great variety of ways.
It breaks up into permanganate of potassium and binoxide of manganese when the aqueous solution is mixed with a large quantity of,
water, and even strong solutions are rapidly decomposed in the same
way by boiling:3K2MnO4 + 2H20 = K2Mn208 + MnO2 + 4KHO.
By means of acids, the change from manganate to permanganate
may be almost instantaneously effected; but by the presence of an
excess of alkali the decomposition is always greatly retarded.
Exp. 328.-Add a few drops of sulphuric acid to the remaining
portion of the solution of manganate obtained in Exp. 326, and
observe that a quantity of the red permanganate of potassium is
immediately produced.
628. Permanganic Acid (Mn207), or rather its hydrate H2Mn208,
may be obtained in aqueous solution by decomposing permanganate
of barium with sulphuric acid. The solution bleaches powerfully,
and the acid is rapidly destroyed by organic matter and other
reducing agents. Of the compounds of this acid, that with
potassium is by far the best-known.'xp. 329.-Place 300 c. c. of water in a porcelain dish, heat it to
boiling and add to it by portions the remainder of the powdered
green manganate of Exp. 326; from time to time add small portions
of hot water to replace that which evaporates, and continue to boil
until the green color of the solution has changed to deep violet red,
and the manganate of potassium has all been changed to permanganate.




PERMANGANIC ACID.                   533
In case the manganate contains a large excess of free alkali it cannot
readily be converted into permanganate by boiling; it will therefore
often be found necessary to neutralize with nitric acid a portion of the
alkali which is in excess. As soon as the transformation has been
completed, pour the mixture into a tall bottle, leave it at rest until
the binoxide of manganese and other insoluble matters have settled;
then decant the clear liquor into a glass-stoppered bottle, and preserve
it for use in subsequent experiments. The insoluble deposit may be
again boiled with water and allowed to settle; the clear liquor thus
obtained may be added to that previously prepared.
In order to obtain crystals of the permanganate, a clear solution
like that above described should be rapidly evaporated to a small bulk,
then decanted from the binoxide of manganese which is precipitated
during the process, and set aside to cool. Needle-shaped crystals of a
dark purple-red color will soon be formed; they are soluble in 16 parts
of water at 15~, and are permanent in the air. It is well to purify the
first crop of crystals by washing them with a little cold water, then
dissolving in the least possible quantity of boiling water, and again
crystallizing in the cold. Neither the crystals nor the solution should
ever be brought into contact with paper. Decantation will ordinarily
be sufficient in order to separate the crystals from the mother-liquor;
but if filtration be necessary in any case, an asbestos filter should be
employed (Appendix, ~ 14).
629. The permanganates are isomorphous with the perchlorates
(~ 125), and the potassium salts of the two acids are capable of
crystallizing together in all proportions. These compound crystals
are red-colored when they contain much perchlorate of potassium,
but are black if they contain as much as half their weight of the
permanganate.
In the same way that perchloric acid is a more stable acid than
chloric acid, so permanganic acid is less readily decomposed than
manganic acid. Both manganic acid and permanganic acid, however, give up oxygen to other substances with remarkable facility, and are much used as oxidizing agents. Even a piece of
wood or paper thrown into the green or red solution of a manganate or permanganate, will quickly abstract oxygen from the
solution and destroy its color. In filtering the colored solutions,
paper is consequently inadmissible, as has been stated in Exp.
329; asbestos, sand, or some other inert filtering-material must
be resorted to. Permanganate of potassium is largely employed




534                          nioN.
for disinfecting putrid water and animal or vegetable matters
in a condition of putrefaction. A solution of it, such as has been
prepared in Exp. 329, is of great use in volumetric analysis,
especially for testing the value of iron ores.
IRON.
630. Although iron is one of the most widely diffused and
most abundant of the metals, it is rarely found native in the metallic state. Meteors, however, fall upon the earth from outer
space, which consist mainly of metallic iron, contaminated with
several other elements in small proportions. Minerals containing
iron occur in great numbers; and there are indeed few natural
substances, whether organic or inorganic, in which iron is not
present. It is found in the ashes of most plants, and in the blood
of animals. The natural compounds of iron which are available
as ores of the metal, are chiefly oxides and carbonates. The most
important varieties of these ores of iron are the following:-1.
Magnetic iron-ore, the richest of the ores of iron, containing when
pure, 72'41 per cent. of iron, and not infrequently approximating
closely to this composition in large masses. 2. Red Hcenmatite,
consisting, when pure, of anhydrous sesquioxide of iron containing 70 per cent. of iron; this ore often yields from 60 to 69 per
cent. of the metal. 3. Specular iron-ore, which is a crystalline
form of the same anhydrous sesquioxide of iron. 4. Limonite, or
Brown iron-ore, which consists essentially of hydrated sesquioxide
of iron, containing 59'89 per cent. of iron; yellow ochre is a
clayey variety of this very abundant ore; the numerous ores
classed under this head yield from 25 to 55 per cent. of iron. 5.
Spathic ironl-ore, or Carbonate of iron, which contains in its
purest state 48-27 per cent. of iron, but is so generally contaminated with manganese, calcium, and magnesium as to yield very
various quantities of iron, ranging from 14 to 43 per cent. 6. Clay
iron-ore, a name applied to a mixture of clay and carbonate of
iron, which occurs very abundantly in the coal-measures; as this
ore is a mixture in uncertain proportions, it yields various percentages of iron, ranging from 25 to 40 per cent.
From  the richer iron-ores, like the magnetic and specular
oxides, a very excellent iron can be obtained by simply heating




THE BLOOMARY.                    535
the broken ore with charcoal in an open forge fire, urged by a
blast. The ore is deoxidized by the carbon of the fuel, anti the
reduced iron is agglomerated into a pasty lump called a " bloom,"
while the earthy impurities contained in the ore combine with a
portion of the oxide of iron to form a fusible glass or slag. The
spongy bloom is freed from slag and rendered homogeneous and
solid by hammering while still red-hot; by reheating and hammering, the iron is then converted into bars or shaped into any
other desired form. This process is not economical in the chemical sense, for much iron is lost in the slag, and much fuel is
burnt to waste in an open fire; but when well conducted, it yields
an admirable quality of iron; and since the original outlay for
the construction of a bloomary is small, and repairs upon it are
always easy, the method has many advantages in regions where
transportation is dear while rich ores, charcoal, and water-power
abound. The bloomary process, in its crudest form, is easily
practised by people possessing but little mechanical skill and no
chemical knowledge; it is undoubtedly the oldest method of extracting iron from its ores.
631. In the extraction of iron from its common ores, the metal
is usually obtained, not pure, but in a carburetted fusible state,
known as cast iron or pig iron. The main features of the process are, first, a previous calcination or roasting to expel water,
carbonic acid, sulphur, and other volatile ingredients of the ore;
secondly, the reduction of the oxide of iron to the metallic state
by ignition with carbon; thirdly, the separation of the earthy
impurities of the ore by fusion with other matters into a crude
glass or slag; and, lastly, the carbonizing and melting of the reduced iron. With the purer kinds of iron-ore, the preliminary
calcination is not always essential; but with the majority of ores
it is very desirable; not unfrequently all the drying necessary is
effected in the upper part of the blast furnace itself, within which
the three last steps of the process always take place.
The blast furnace for iron consists essentially of a huge cylindrical structure of masonry, 15 to 25 m. in height, and 5 to 6 m.
in diameter at tLe central portion of the cylinder, but contracted
to a less diameter both at the top or throat and at the bottom or
hearth. Air is forced in at the bottom of the furnace to support




536                 THE BLAST FURNACE.
the combustion, and it has been found advantageous in the majority of cases to heat this blast of air to about the melting-point
of lead before it enters the furnace. The reduction of the oxides
of iron being effected by the carbonic oxide resulting from the
combination of carbonic acid with hot carbon (Exp. 181), it is
not difficult to calculate the amount of carbon and the amount of
air requisite to reduce to metallic state the iron contained in a
given weight of an iron-ore of known composition.  Thus the
formula of specular iron-ore, or of red haematite, is Fe203, and,
since the atomic weight of iron is 56, these ores are 70 per cent.
iron; accordingly the following quantities are equivalent one to
the other:
1  =   1429    -    0'3214       =     0 4285   -    1'865.
Iron.    Fe0O,.    Requisite weight  Weight of oxygen    Air.
of carbon in   requisite to convert
state of CO.    so much C to CO.
For every kilogramme of iron produced, nearly two kilogrammes
of air must be supplied, and at least 1 kilogramme of fuel, merely
to accomplish the chemical reaction. The reduction of the oxide
of iron, however, is not alone sufficient to secure the metal; ironores almost always contain earthy admixtures, consisting chiefly
of silica, clay, and carbonate of calcium; and these substances are
so intimately mixed with the reduced metal, that it is essential
to melt them before the iron can separate by virtue of its greater
specific gravity. Any one of these substances taken alone is infusible at the temperature of the furnace; they must be converted
into fusible double silicates; and as it is rarely the case that the
natural impurities of an ore are present in the proportions requisite for the formation of such double silicates, it is generally
necessary tp mix with the ore a substance intended to effect this
result, and therefore called the flux. With ores in which the
earthy ad ixture is chiefly calcareous, the flux must be clay or
some siliceous material; but in the more frequent case of ores containing clay or silica, the flux will be limestone or quicklime.
In either case, a fusible double silicate of aluminum and calcium
is the essential constituent of the slag.  With siliceous ores there
is another reason for the addition of lime; the double silicate of
aluminum  and iron is very fusible, and a'considerable quantity




THE BLAST FURNACE.                 537
of iron might be lost in the slag, were not lime enough added to
prevent the formation of this iron-containing silicate. Sometimes
both calcareous and siliceous ores are within reach of an ironfurnace, and the smelter, by mixing the two varieties in due proportion, may avoid the necessity of adding a flux.
The blast furnace is charged at the top with alternate layers of
the fuel (which may either be charcoal, anthracite, or coke), the
ore, and the flux, which is generally lime; these materials are
constantly supplied at the top, and air is constantly supplied in
immense quantities at the bottom of the furnace, the actual weight
of the air forced in being greater than the sum of the weights of
the ore, the fuel, and the flux. Where the blast first touches the
ignited fuel, carbonic acid is formed; this gas, rising with the
unused nitrogen through the furnace, comes in contact with whitehot carbon, and is reduced to carbonic oxide (Exp. 181). The
layers of solid material thrown in at the top of the furnace gradually sink down, and as soon as a stratum of ore has descended
sufficiently to be heated by the hot mixture of nitrogen and carbonic oxide it becomes reduced to spongy metallic iron, which,
mixed with the flux and the earthy impurities of the ore, settles
down to hotter parts of the furnace, where it enters into a fusible
combination with carbon, while the flux and earthy impurities
melt together to a liquid slag. The liquid carburetted iron settles
to the very bottom of the furnace, whence it is drawn out, at
intervals, through a tapping-hole, which is stopped with sand
when not in use. The viscous slag flows out over a dam, so placed
as to retain the iron, but to permit the escape of the slag which
floats on the iron, as fast as it accumulates in sufficient quantity.
The fusion of the materials in the lower part of the furnace requires a great heat; and the amount of fuel consumed in getting
this high temperature is much greater than the amount requisite
for the reduction and carbonization of the metal. As charcoal
is a much purer carbon than coal or coke, iron smelted with charcoal is generally purer than that smelted with coal; but as
charcoal crumbles under great pressure, the furnaces in which
charcoal is used are usually much smaller than those intended
for anthracite or coke. The consumption of fuel in smelting
1000 k. of iron varies with the nature of the furnace, the blast,




538                      CAST IRON.
and the fuel, between 500 k. and 3000 k. The gases which issue
from the mouth of the blast furnace are charged with an enormous heating-power; for besides being themselves intensely hot
they contain, even after having effected the reduction, a large
proportion of combustible gases, such as carbonic oxide, carburetted hydrogen, and hydrogen. This gaseous mixture takes fire
whenever it comes in contact with the air; a part of its heat may
be utilized in heating the air-blast and generating steam.
Two distinct varieties of cast iron exist, which differ in color,
texture, and fusibility; these are white cast iron and gray cast
iron. White iron is hard and brittle, of crystalline texture and
shining fracture. Gray iron is slightly malleable, and has a
granular texture; its fracture may be either coarse or fine-grained,
and minute particles of black graphite are visible upon the broken
surface. White cast iron melts at a lower temperature than gray,
but does not become so liquid as the gray. Gray cast iron, when
rapidly cooled, is converted into white iron; when a casting is
made in an iron mould, the layer of metal in contact with the
mould is chilled and converted into the hard white iron, while the
interior of the casting will retain the condition of the stronger
gray iron. Excellent shot and shell for rifled cannon have lately
been cast on this plan. The chief chemical difference between
white and gray cast iron consists in the different condition of the
admixed carbon. In white iron the carbon seems to be dissolved
in or combined with the iron, while in gray cast iron, on the
other hand, the greater part of the carbon seems to be mechanically
diffused through the solid iron, in the state of graphite. The two
varieties, however, shade off into each other through a great variety of intermediate mixtures. White cast iron rusts much more
slowly than gray cast iron. When white iron is heated with
strong chlorhydric acid it entirely dissolves, but the combined
carbon enters into combination with a portion of the nascent
hydrogen, forming hydrocarbons which impart a peculiar smell to
the gas evolved. Gray iron does not wholly dissolve in hot chlurhydric acid; a residue of graphite remains; but the gas evolved
has the same smell as the gas evolved from white iron. When
bars, plates, or implements of common cast iron are exposed
to the slow action of dilute acids or of saline solutions such




IMPURITIES OF IRON.               539
as sea-water, the iron is not infrequently wholly dissolved
away, while the graphitic carbon remains untouched-a light,
soft, sectile substance which often retains the color and form
of the original article, but possesses neither hardness nor tenacity. The largest proportion of carbon found in cast iron is
5 75 per cent.; this large percentage occurs in a lustrous variety
of white iron which contains manganese and is called specular
iron. In gray iron the amount of carbon varies from 2 to nearly
5 per cent.
Silicon, sulphur, phosphorus, manganese, and copper are very
common impurities in cast iron. The silicon comes from silica
deoxidized in the furnace; its amount varies from 0'1 to 3'5 per
cent. There is more of it in gray than in white iron, and more in
hot-blast iron than in cold-blast. Sulphur is almost always present in cast iron, but only in very small quantity; its presence
is supposed to conduce to the formation of white iron. The presence of phosphorus to the extent of I or 2 per cent. is not uncommon, and does not injure iron intended for castings, inasmuch
as the phosphorus makes the iron more fusible, and more liquid
when melted. Manganese is frequently present in cast iron, as
is not unnatural considering the common association of manganese
ores with iron-ores. Cast iron containing manganese appears to
be especially suitable for the production of steel.
The production of malleable or " wrought" iron from  cast
iron, consists essentially in burning out the carbon, silicon, sulphur, and phosphorus which cast iron contains. This oxidation
of the impurities of cast iron is effected either by blowing upon
the melted metal with an air-blast in a small charcoal furnace
called a " finery," or by stirring the melted iron in a reverberatory furnace in which the fuel does not come in contact with the
metal, and into which air can be admitted at will; the latter process, now much the most important method of manufacturing
wrought iron, is called "pn ldctdling."  In puddling it is customary
to add to the charge of pig iron a quantity of iron scale or other
oxide of iron. The oxidation of the silicon, carbon, phosphorus,
and other impurities is effected partly by the air and partly by
the oxide added to the charge; the carbon burns to carbonic
oxide, which heaves the seething mass as it escapes and burns in




540                      STEEL.
jets of blue flame; the other impurities form, with the oxides of
iron and manganese, a cinder or slag which ordinarily contains
sulphur, silicic acid, and phosphoric acid. When the cast iron is
so far decarbonized as to be pasty in the fire, it is gathered into
lumps on the end of an iron bar and carried from the furnace to
a hammer or squeezer which expresses the liquid slag and welds
into a coherent mass the tenacious iron. The hammered lump
may be reheated and rolled or forged into any desired shape.
The waste of iron in converting cast into malleable iron amounts
to from 13 to 30 per cent.
Ordinary malleable iron has a gray color, and a specific gravity
of about 7'6. Though less malleable than gold and silver, its
malleability is very great, and the greater the purer the metal
and the higher the temperature to which it is raised. At a red
heat, separate pieces may be firmly united by hammering or
rolling; the operation is called welding. Sulphur is said to render wrought iron brittle while hot or "red-short;" silicon and
phosphorus render iron brittle at the ordinary temperature, or
"cold-short," in technical phraseology. These common impurities of cast iron are therefore very prejudicial to wrought iron.
Wrought iron is hammered or rolled while in a doughy condition;
and the uniform, close, fibrous texture which is valued in malleable iron depends much upon the nature of this mechanical
treatment, and the extent to which it is carried. Common malleable iron still contains from 0'25 to 0'5 per cent. of carbon; the
smaller the amount of carbon the softer the iron. Wrought iron
dissolves almost completely even in dilute acids; but the hydrogen
evolved has the peculiar smell attributed to the presence of carbonaceous vapor.
632. Steel. —This invaluable substance is in composition intermediate between cast and wrought iron, containing less carbon than
cast iron, but more than wrought. It may be made from wrought
iron by heating bars of iron to redness for a week or more in contact with powdered charcoal in close boxes from which air is carefully excluded. Though the iron is not fused, nor the carbon
vaporized, yet the carbon gradually penetrates the iron and alters
its original properties; when the bars are withdrawn from the
chests in which they were packed, the metal has become fine-grained




TEE BESSEMER PROCESS.                 541
in fracture, more brittle, and more fusible. The bars, however, are far from uniform in composition, the outside being more
highly carbonized than the interior; they are apt to show blisters
of various sizes on the surfaces, and the steel thus prepared is
called "' blistered " steel. To obtain steel of a uniform quality, it
must be cast into ingots. This process of preparing steel is called
the " cementation " process; it is a curious instance of chemical action between solid materials which are apparently in a state of rest.
Since the materials used in this process are the purest attainable
-the best iron and the best charcoal-the steel obtained is of
the best quality. Cheaper methods of preparing an inferior steel
are, however, of great industrial importance. If the " puddling"
process for preparing malleable iron should be arrested when the
cast iron had lost from one-half to two-thirds of its carbon, the
product would be an impure steel, impure because the silicon,
phosphorus, sulphur, and other impurities of the cast iron would
only be incompletely removed. Nevertheless there are uses in
the arts for a steel of this quality which may bc cheaply manufactured.
633. A new and very rapid method of preparing cast steel
directly from cast iron is that known as the Bessemer process.
From two to six tons of cast iron, previously melted in a suitable
furnace, are poured into a large covered crucible, made of the
most refractory materials, and swung on. pivots in such a manner
that it can be tipped up and emptied by means of an hydraulic
press. Through numerous apertures in the bottom of the crucible
a blast of air is forced up into the molten metal; an intense
combustion ensues involving the carbon in the iron and a portion
of the metal itself, and generating a most intense heat, which
keeps the mass fluid in spite of its rapid approach to the condition of malleable iron. Such a quantity of specular iron or
white cast iron is then added to the iron in the crucible as is
necessary to give carbon enough to convert the whole mass into
steel, and the melted steel is immediately cast into ingots. Six
tons of cast iron can thus be converted into tolerable steel in
twenty minutes. This steel is suitable for the manufacture of
axles, cranks, rails, boiler-plates, and many other articles in which




542                    OXTDES OF MIPON.
great strength should be combined with hardness. A pure steel
cannot at present be made by this process, inasmuch as the combustion in the crucible does not get rid of the sulphur and phosphorus in the cast iron nearly as perfectly as does the puddling
process; for the same reason the manufacture of wrought iron
by this method, though the original object of the invention, has
been thus far found impracticable.  It deserves mention that the
nailer who keeps his nail hot, while hammering it, by a carefully
regulated blast of cold air, applies the chemical fact involved in
Bessemer's process. This ancient practice was indeed a prophecy
of Bessemer's invention.
634. The two qualities of steel which are of greatest importance
are its hardness and its elasticity.  These qualities are developed by quickly cooling the heated metal; the delicate processes by which steel tools and springs are hardened, tempered,
and annealed are exceedingly curious, but are rather physical
than chemical phenomena.  Many implements are sufficiently
well made by converting their exterior surfaces into steel, leaving
the interior of cast or wrought iron. Thus cast-iron tools may
be heated with oxide of iron to remove a part of the carbon from
their exterior and thus coat them, as it were, with steel. Tires
for wheels are well made of wrought-iron bars which have been
superficially converted into steel by the cementation process;
such tires combine the toughness of malleable, iron with the
hardness of steel.
635. Oxides of Iron.-There are several definite compounds of
iron and oxygen. The best-known of these oxides are the protoxide (FeO), or ferrous oxide, as it is often called, and the sesquioxide (Fe203), often called ferric oxide, and sometimes spoken
of as peroxide of iron. There is another oxide, ferric acid FeO3,
which is an exceedingly unstable substance, known only as it
exists in combination with potassium, as ferrate of potassium
(K2FeO,), or with some other powerful base.  Besides these
oxides, there are several compounds of intermediate composition,
which may be supposed to result from the union of ferrous and
ferric oxides, in various proportions; they are called collectively
ferroso-ferric oxides; the most important among them is the




FERROUS AND FERRIC OXIDES.             543
magnetic oxide Fe04-= FeO,Fe 08, which is the black oxide
formed when iron is oxidized at high temperatures in oxygen
gas, in air or steam (~~ 9, 18, 34).
636. Ferrous Oxide or Protoxide of Iron (FeO). —This compound is not easily obtained pure, since it absorbs oxygen from
the air with great avidity and thus becomes contaminated with
the sesquioxide. But by dissolving a ferrous salt (that is, a salt
of protoxide of iron) in recently boiled water, and adding to the
liquid a solution of caustic alkali, which has likewise been boiled
to expel air, there will be precipitated a white ferrous hydrate,
provided the operation be conducted out of contact with the air.
If this hydrate be exposed to the air, as when the solution of a
ferrous salt is mixed with the alkali without the precautions
above enumerated, it will rapidly absorb oxygen and will exhibit various shades of light green, bluish green, and black, till,
finally, it assumes the red color of hydrated ferric oxide (Exp.
341). The anhydrous oxide obtained by igniting ferrous oxalate
in close vessels, absorbs oxygen so rapidly that it takes fire when
brought into contact with the air.
Hydrated ferrous oxide is readily soluble in acids, forming salts
known as the protosalts of iron or ferrous salts; many of these
salts are of a pale green color; like the hydrate, they rapidly
suffer decomposition by absorbing oxygen fiom moist air.
637. Ferric Oxide (Fe20O).-This oxide, called also red oxide,
sesquioxide, or peroxide of iron, occurs very abundantly and
widely distributed in nature. Several of its varieties have been
already mentioned as ores of iron (~ 630).  It may be obtained
also by igniting metallic iron or either of the lower oxides or
hydrates in contact with the air. For use in the arts, it is prepared by igniting ferrous sulphate with or without addition of a
small proportion of nitrate of potassium, or by roasting the native
hydrate (yellow ochre). The better sort, known as rouge, is
largely employed for polishing glass and jewelry, and all grades
of it are extensively used as pigments. Red ochre is impure
ferric oxide. As commonly met with, the oxide is amorphous
and has a red, brown, or nearly black color, according to the
method of its preparation. At a full white heat, it gives off a
portion of its oxygen, and magnetic oxide of iron is formed. It




544                OXIDATION BY TRON-RUST.
is easily reduced to the metallic state by hydrogen gas, even at
temperatures below redness, and by carbon and carbonic oxide at
a red heat, as has been stated in ~ 631. Ammonia gas reduces
it also at a red heat.
Bxp. 330.-In the middle of a tube of hard glass, No. 3, 10 c. m.
long, and provided at both ends with corks carrying short, straight
delivery-tubes, place a teaspoonful of red ochre or ignited iron-rust.
Attach to one end of the tube a hydrogen-generator or gas-holder
provided with a chloride-of-calcium drying-tube, and connect with
the other end a U-tube. Support the tube containing oxide of iron
upon a ring of the iron stand, cause a current of hydrogen to flow
through it, immerse the U-tube in a bottle of cold water, and finally
heat the oxide of iron. The hydrogen will combine with the oxygen
of the red oxide of iron, water will be formed and will condense in the
U-tube, while finely divided metallic iron will be left behind. After
the reduction has been completed, allow the tube to become cold, and
then scatter its contents through the air upon an earthen plate. They
will take fire and burn again to the condition of red oxide.
Erp. 331. —Repeat Exp. 330, using carbonic oxide instead of hydrogen, the products will be iron and carbonic acid instead of iron
and water.
638. The facility with which red oxide of iron gives up oxygen, taken in connexion with the readiness with which metallic
iron and the protoxide take on oxygen, is a fact of great practical
importance.  It has been found that organic substances may be
more rapidly incinerated by heating them in the air in contact
with a small quantity of ferric oxide than in air alone; the oxide
of iron appears to act as a carrier of oxygen, as it is alternately
reduced by the combustible and again oxidized by the air. Even
at ordinary temperatures, and with the hydrated oxide, the same
reactiofns are witnessed, though in a less degree.  The iron
nails employed in the construction of ships, bridges, fences, or
shoes, actually corrode, "eat up" or "burn out" the organic
matter in contact with them, by absorbing oxygen from the air
and transferring it to the carbon compound with which they are in
contact. The rotting of canvas by iron-rust, or of a fishing-line
by the rusty hook, are familiar instances of corruption by rust.
These reactions doubtless play an important part in the formation of soils by the oxidation of vegetable remains.




TYIDRhAT OF IRON,                   545
In the same way ferric oxide converts sulphide of calcium
(CaS) into sulphate of calcium (CaSO4) at the expense of the
oxygen of the air. A useful cement has been prepared by mixing
the residual oxysulphide of calcium of Leblanc's soda process with
an equal weight of the ferric oxide left as a residue in burning
iron pyrites for sulphuric acid.  Hydrated sesquioxide of iron
(FeO3,:3H,0) may readily be prepared by adding an excess of
ammonia-water to the solution of almost any ferric salt.
Exp. 332.-Cover a teaspoonfuil of fine iron filings or small tacks
with three or four times as much dilute sulphuric acid in a small
bottle; wait until the evolution of hydrogen ceases, then decant the
clear liquor into a small flask or beaker, add to it a few drops of strong
nitric acid, and heat it to boiling. The liquor will soon be colored
dark brown by the nitrous fumes resulting fiom the decomposition of
the nitric acid, which are for a short time held dissolved by the liquid;
but this deep coloration soon passes away, and there is left only the
yellowish-red color of the ferric sulphate which has been formed.
Add to the solution ammonia-water, until the odor of the latter persists after agitation, and collect upon a filter the flocculent red precipitate of ferric hydrate.
Besides this normal hydrate (Fe,20,3H120) there are several
other ferric hydrates, containing smaller proportions of water.
They are found in nature, and may be obtained by heating the
normal hydrate, or by suffering it to remain for a long while
under water, or by boiling it for some time in water.
639. Generally speaking, hydrated sesquioxide of iron is easily
soluble in acids, though some peculiar varieties of it dissolve only
with difficulty.  The anhydrous oxide also dissolves in acids,
though less easily in proportion as it has been more strongly ignited; its best solvent is concentrated boilinll chlorhydric acid.
By long-continued heating at 3000 or 320~, ferric hydrate can
be deprived of all its water, and still be readily soluble in acids;
but when heated to dull redness, this powder suddenly glows
brightly for a moment and contracts in bulk, without either
losing or gaining weight, and is then attacked' by acids only
slowly and with difficulty. It has been observed, however, that
the ignited oxide may still be dissolved' rather easily by a hot
mixture of chlorhydric acid and ferrous chloride, protochloride of
tin, zinc, or some other reducing aggnt..
2 it,




546              PROPERTrES OF FERRTC HYDRATE.
Ferric hydrate is somewhat used as a mordant in dyeing, and
is largely employed for purifying coal-gas. As has been stated
under arsenic, the recently precipitated hydrate acts as an antidote to arsenious acid, since when given in sufficient quantity it
forms a basic arsenite of iron scarcely at all acted upon by water.
Exp. 333.-Dissolve half a gramme of arsenious acid in 40 or 50 c. c.
of boiling water. Divide the solution into two portions, and stir into
one of these portions a considerable quantity of moist ferric hydrate,
such as was obtained in Exp. 332; filter the mixture, acidulate the
filtrate with chlorhydric acid, and test it for arsenic by means of sulphydric acid (~ 340).
If a sufficient quantity of ferric hydrate has been employed, no precipitate of sulphide of arsenic will be obtained in the filtrate, though
on adding a drop of chlorhydric acid, and afterwards sulphydric acid,
to the original solution of arsenic, an abundant yellow precipitate will
be at once thrown down..Ecp. 334.-Fill a tube, 30 c.m. long, with alternate tufts of cotton
and loose layers of dry ferric hydrate, pass a slow current of sulphydric acid (Exp. 86) through the tube, and observe that the oxide
gradually becomes black; it appears to be converted into ferrous sulphide, while water and sulphur are set free:Fe2O3,,3H20 + 3H2S = 2FeS +  S +  6H20.
After a good part of the ferric oxide has become black, remove the
contents of the tube to a porcelain plate, and leave them exposed to
the air in a place where no harm can be done in case they take fire.
By the action of the air, sesquioxide of iron will be reproduced and
sulphur set free within the mass; some sulphurous acid is given off at
the same time, and heat is evolved, as will readily be perceived if the
quantity of material be large. The following equation, though it does
not fully express the complete reaction which really occurs, may still
serve to give a general idea of these chemical changes:2FeS +  50 = Fe203 + S +  SO0.
In practice the impure illuminating gas is made to pass through
layers of ferric oxide,,often. made porous by an admixture of sawdust;
as soon as the oxide ceases to absorb sulphuretted hydrogen, it is
"revivified" by forcing or drawing through it a current of fresh air or
by spreading it in the~,air. The oxide is thus used over and over again,
until so much sulphur' has accumulated within it as to interfere, mechanically, with its absorbent power. The sulphur may readily be
recovered from this mixture by distillation.; or the spent oxide may be




SULPHIDES OF IRON.                  547
used instead of pyrites for making sulphuric acid, wherever enough of
it can be obtained to repay the trouble of collecting.
The ferric salts are usually of a yellowish-brown or red color
when hydrated, though some of them have a violet tinge; they
are white when dry. The normal salts are usually soluble in
water and deliquescent; and there are numerous soluble basic
salts, besides other basic salts which are insoluble.
In the ferric salts iron plays the part of a trivalent element
like aluminum, while in the ferrous salts it is bivalent like calcium or lead. The ferric salts closely resemble salts of aluminum,
and are for the most part isomorphous with them. Besides acting as a base, ferric oxide, like oxide of aluminum, combines
with several of the more powerful bases to form salts called
ferrites; magnetic oxide of iron, for example, may be regarded
as the ferrite of iron.
It is remarkable that ferric oxide may be displaced from many
of its compounds by the protoxide. of iron; thus, when hydrated
ferrous oxide is added to a solution of ferric sulphate, hydrated
sesquioxide is precipitated:Fe203,3SO, +  3(FeO,I1,O) =  Fe20O,3H2O  +  3(FeO,SO,).
In like manner, carbonate of barium precipitates anhydrous
ferric oxide from ferric salts, but upon ferrous salts it has no
action:
Fe2C16 +  3BaCO3 =  Fe,03 +  3BaC12 +  3CO,.
Sulphides of Iron.-There are several sulphides of iron, the
most important of which are the protosulphide (FeS) and the
bisulphide (FeS2), found native as iron pyrites.  There is a
sulphide, Fe3S, which is magnetic like the oxide to which it
corresponds.
640. Ferrous Su7phide (FeS) is a substance of great value to
the chemist as the cheapest source of the important reagent
sulphydric acid (~~ 202, 210).
This sulphide may be prepared by igniting pyrites in a covered
crucible, by rubbing roll brimstone against a white-hot iron bar, or by
fusing together sulphur and iron turnings. The second method is to
be recommended if the student have ready access to a blacksmith's
forge. The sulphide is a waste product in those chemical works where
2N2




548                 PROTOSULPHIIDE OF IRON.
sulphate of lead, obtained from dye-houses, is reduced to the metallio
state by fusion with iron and coal. In the laboratory it may be prepared as follows:E'xp. 335. —Heat a common Hessian crucible to redness in a fire of
coke or anthracite, and project into it from an iron spoon successive
small portions of a mixture of 7 parts of iron turnings and 4 parts of
powdered sulphur, replacing the cover of the crucible after each addition of the mixture. The sulphur and iron combine with great energy,
and the sulphide formed melts down to the liquid state. Since the
molten sulphide is capable of dissolving both iron and sulphur, according as the one or the other may be present in excess, it is impossible to
prepare a pure protosulphide by this method. But the product obtained as above described, though of variable compositiofi, answers
perfectly well for all ordinary purposes. When the crucible has become half-full of the molten sulphide, remove it from the fire, pour
out its contents upon a brick floor, and, if more of the sulphide be
desired, replace the crucible in the fire and proceed as before.
Where comparatively large quantities of the sulphide are required,
it is well to bore a hole through the bottom of a plumbago crucible
and set the latter upon the grate-bars of the furnace in such manner
that the hole may remain open; fill the crucible with iron turnings;
heat it to redness and throw in lumps of sulphur upon the hot iron.
As fast as sulphide of iron forms it will melt and flow through the
hole in the crucible into the ashpit below, which should be kept clean
to receive it. In the preparation of sulphide of iron, wrought iron
should always be employed. From the filings of cast iron, but little,
if any, of the fusible sulphide can be prepared.
The foregoing experiment illustrates the practical methods of malking
ferrous sulphide; but several other reactions which produce it are of
scientific interest (compare Exp. 85).
Eap. 336.-Arrange a bottle for generating sulphuretted hydrogen,
as in Exp. 86, but in place of the delivery-tube in Fig. 41 attach to
the bottle a jet for burning the gas. After the air has been completely expelled from the bottle, light the sulphydric acid gas at the
jet, and hold in its flame a piece of fine iron wire; the iron will burn
to ferrous sulphide, and if the wire be held in the axis of the flame, so
that a considerable portion of it shall be kept red-hot, the globule of
sulphide of iron formed will melt and flow backward upon the wire as
fast as the end of the latter is consumed.
Exp. 337.-Mix 20 grms. of fine iron filings, 14 grms. of flowers of
sulphur, and 7 grTms. of water in a small bottle, and heat the mixture
gently upon a sand bath, or set it aside in a warm place. Chemical




aRON PYRITES.                    549
action will soon set in, much heat will be evolved, and in the course
of half an hour the mixture will become black from formation of sulphide of iron. If the porous black sulphide be left exposed to the air,
it will absorb oxygen, and will be partially converted into ferrous
sulphate.
A firm packing or lute for the joints of iron vessels is prepared by
mixing together 60 parts of fine iron filings, 1 part of flowers of sulphur, and 2 parts of powdered chloride of ammonium. The mixture
is made into a stiff paste with water, and immediately applied to the
iron. It soon becomes hot and swells up, and sets to a hard compact
mass, while ammonia and sulphuretted hydrogen are disengaged.
Exp. 338. —)Dissolve a small crystal of ferrous sulphate (copperas)
in water, and add to the liquid a drop or two of sulphydrate of ammonium (~ 526). Black sulphide of iron will be thrown down (compare
Exp. 334).
The finely divided ferrous sulphide obtained in the wet way, as in
the last two experiments, dissolves much more quickly in acids than
the compact sulphide obtained by the way of fusion; in contact with
acids it evolves gas so tumultuously that it would be inconvenient as
a source of sulphydric acid.
The black earth between the stones of the pavements of cities, and
at the bottoms of drains and cesspools, owes its color to sulphide of
iron formed by the putrefaction of sulphuretted compounds in contact
with ferric oxide contained in the earth.
641. Bisulphide of Iron (FeS2) occurs abundantly in nature
as the well-known mineral iron pyrites. Two distinct forms of
it are met with:-the yellow cubical pyrites, crystallized in forms
of the monometric system; and the white pyrites or marcasite,
which crystallizes in trimetric forms. A third variety of sulphide
of iron, called magnetic pyrites, is of different composition from
the foregoing, and contains less sulphur than the bisulphide.
Iron pyrites appears to have been sometimes formed in nature by
the deoxidation of sulphates, such as the sulphate of calcium, by
means of organic matter in presence of chalybeate waters. The
formation of pyrites has often been noticed in solutions of sulphate of iron into which organic matters have fallen. But bisulphide of iron may be readily formed in the dry way also.
The compact forms of yellow pyrites, whether natural or artificial, are permanent in the air; but when finely divided the
mineral oxidizes rather easily, with evolution of considerable heat,




550                  FERROUS CHLORIDE.
White iron pyrites oxidizes rapidly in the air, no matter whether
it be compact or friable. The spontaneous combustion of many
kinds of coal is due to the oxidation of iron pyrites disseminated
through the combustible. Alum and copperas are often prepared
from pyritous shales, either by firing heaps of the shale artificially, or by allowing the heaps to talike fire spontaneously through
oxidation of the pyrites, and then regulating the combustion so
that the largest practicable yield of sulphate of iron or of sulphate
of aluminum shall be obtained. So long. as the temperature of
the burning pyrites remains comparatively low, ferrous sulphate
and sulphuric acid are the principal products, the latter uniting
with the alumina of the shale, if such be present; when the heap
has become cold, the sulphates can be separated by lixiviating
the mass with water. When pyrites is roasted at higher temperatures, as in the manufactuire of sulphuric acid, sulphurous
acid is given off, and ferric oxide left as the principal residue.
When distilled in close vessels, one atom of the sulphur in
iron pyrites is expelled, and ferrous sulphide remains. Sulphur
has sometimes been prepared in this way in a dearth of native
sulphur.
642. Ferrous Chloride (FeC12) may be obtained by passing
chlorine or dry chlorhydric acid gas over hot iron; in case chlorhydric acid be employed, hydrogen will be evolved. As commonly
met with, however, the chloride is in the form of a hydrate
(FeCl2+4H20) obtained by dissolving metallic iron in dilute
chlorhydric acid. It crystallizes easily, forms double salts by
uniting with many other chlorides, and may be deprived of its
water without decomposition when heated carefully out of contact with the air.
643. Ferric Chloride (Fe2Cl6). —As obtained by burning metallic iron in an excess of dry chlorine, this compound occurs in
anhydrous, glistening scales, which volatilize easily when heated.
It dissolves readily in water, with evolution of heat, and deliquesces rapidly in the air. It hisses when thrown into water.
Once dissolved in water, it cannot be freed from the water by
evaporation, since chlorhydric acid goes off with the water, and a
basic compound of ferric oxide and ferric chloride remains. Hydrated ferric chloride may readily be obtained by boiling a solu



FERROUS ST.LPHATh.                  551
tion of ferrous chloride with a small proportion of nitric acid, or
by passing chlorine gas through a solution of ferrous chloride.
From concentrated solutions, ferric chloride crystallizes with
several different proportions of water. Ferric chloride combines
with many of the metallic chlorides to form double compounds,
among which the ammonium salts are perhaps the mo6st stable.
644. Ferrous Su7phate (FeSO4).-A hydrate of this compound,
of composition FeSO4 +7H20, usually called copperas or green
vitriol, is the most common of all the compounds of iron. It may
readily be prepared by dissolving metallic iron or protosulphide
of iron in dilute sulphuric acid. On the large scale it is commonly
prepared by roasting iron pyrites at a gentle heat, or aluminous
shales containing pyrites in the manner already indicated. Sometimes, however, it is manufactured directly from metallic iron
and sulphuric acid; and it is obtained as a secondary product in
certain metallurgical operations where copper is precipitated, by
means of iron, from a solution of copper.  The reaction is analogous to those employed for obtaining pure silver (Exp. 267) and
pure lead (~ 594).
Exp. 339.-Dissolve 5 grmins. of common blue vitriol (sulphate of
copper) in 50 or 60 c. c. of water, acidulate the liquor with a few drops
of sulphuric acid, pour it into a bottle, and place in it a rod of thick
iron wire. Copper will immediately begin to be precipitated as a
coating upon the iron, and in the course of an hour or two will be
completely removed from the solution. The original blue color of the
solution will disappear and be replaced by the faint green color of
copperas, while a spongy mass of metallic copper will be obtained:CuSO_ + Fe = FeSO4 + Cu.
Decant the solution of ferrous sulphate from the precipitated copper,
place in it a fragment of iron, and evaporate it to a small bulk; pour
the concentrated solution into a wide-mouthed phial, cork the' phial
tightly, and set it aside in a cool place; the liquid will be converted
into a mass of copperas crystals.
It has been proposed to prepare copperas from the "finery
slag" of the puddling-furnaces (where cast iron is converted into
wrought iron) by treating this slag with dilute sulphuric acid.
The finery slag consists chiefly of basic silicate of protoxide of
iron (2FeO,SiO,).
645. When perfectly pure, the crystals of ferrous sulphate arr




552            PROPERTIES OF FERROUS SULPIATE.
compact, transparent, and of a bluish-green color; but in dry air
they effloresce and become covered with a white incrustation, the
color of which subsequently changes to rusty brown through absorption of oxygen.  The common commercial article is of a
grass-green color, and is contaminated with more or less ferric
sulphate. Besides the common hydrate containing 7 molecules
of water, there are hydrates which contain 4, 3, and 2 molecules
of water respectively. From all of these hydrates the water can
easily be expelled by heat, and if the anhydrous salt thus obtained
be still further heated it will decompose; two stages in this decomposition may be formulated as follows:I. 2FeSO4 =  SO,2 +  Fe203,SO3.
II. Fe,O,,SO3 = Fe203 + SO,.
Basic ferric sulphate is at first formed, while sulphurous acid is
given off; and finally the ferric salt is itself decomposed into anhydrous sulphuric acid and ferric oxide. Upon this reaction the
preparation of Nordhausen sulphuric acid depends (~ 239).  Like
ferrous hydrate, ferrous chloride, and all the ferrous salts, moist
copperas, or all aqueous solution of copperas, rapidly absorbs oxygen from the air.
Exp. 340.-Pour a solution of copperas into an open capsule and
leave it exposed to the air for a day or two; the solution will gradually
become yellow as the oxidation proceeds, and after a while a rusty
precipitate of ferric oxide, or of highly basic ferric sulphate, will fall.
The oxide of iron which separates under these conditions is not readily
soluble in dilute acids. It appears to be an isomeric modification of
the easily soluble hydrate which is precipitated from cold ferric solutions by alkaline lyes. At all events the sulphuric acid of the copperas is insufficient to dissolve all of the ferric oxide formed during its
oxidation. In most cases where a ferrous salt is to be converted into
a ferric salt, it is best to add a certain proportion of free acid to the
mixture, in order to prevent the separation of the oxide.
A difficultly soluble deposit, similar to the foregoing, may readily bo
obtained by boiling an exceedingly dilute solution of almost any of the
soluble ferric salts. It is possible that these sediments should be regarded rather as highly basic salts than as mere hydrates. Their
inertness may perhaps be due to the presence of small proportions of
the acids of the salts fiom which they have been derived, still held in
chemical combination; but there is at present less evidence in favor ol
this view than of the one previously stated.




INr.                          553
l'xp. 341.-To a teaspoonful of a solution of copperas add a few drops
of soda lye, and observe that the hydrate rapidly absorbs oxygen, and
changes color as has been set forth in ~ 636.
Exp. 342.-Mix a few drops of a solution of copperas with a drop or
two of a solution of tannic acid, such for example as tincture of nutgalls, or of oak- or hemlock-bark; a light violet-colored precipitate
will be formed and will remain suspended in the liquid; by exposure
to the air this color soon changes to black. The violet precipitate is
ferrous tannate, and the black precipitate ferric tannate; if these finely
divided precipitates were produced in liquids made slightly viscous by
the addition of gum or sugar, they would remain suspended in the
liquor, which could then be used as writing-ink.
Ink may be prepared as follows: —Powder separately 12 grms. of
nutgalls, 5 grins. of copperas, and 5 grms. of gum-arabic. Boil the
nutgalls two or three hours in a flask with 75 c. c. of water, taking
care to add hot water, by small portions, to supply that lost by evaporation. Allow the mixture to settle, and decant the clear liquor into a
clean bottle. Dissolve the gum-arabic in a small quantity of water,
and mix the mucilage thoroughly with the solution of nutgalls. Dissolve the copperas in another portion of water, and incorporate this solution with the mixture of nutgalls and gum. Add enough water to
make the volume of the mixture equal to 100 c. c. Preserve the ink
in a tight bottle. If the color of the product be lighter than is desired, the liquid may be left exposed to the air until it has acquired a
deeper tint.  When first applied to paper, the color of fresh ink is
comparatively pale, but the writing darkens gradually in proportion as
it absorbs oxygen.
In the course of the foregoing experiments, dip a small piece of cotton cloth in the solution of nutgalls, and allow it to become dry; then
dip it in the solution of copperas and hang it up in damp air. Black,
insoluble tannate of iron will be so firmly precipitated in and upon the
fibres of the cloth, that it cannot be washed away.
The experiment illustrates one general method of dyeing, by means
of which blacks and grays of various shades may be applied to cloth or
leather, though in practice other astringent dye-stuffs, such as catechu,
cutch, or gambier, are commonly employed in place of nutgalls.
Ferrous sulphate is largely employed in dyeing, sometimes directly,
as in the foregoing experiment, but often as the source of other compounds of iron which are employed as mordants; ferrous acetate, for
example, obtained by decomposing ferrous sulphate with acetate of
calcium, is a compound much used by dyers. It should be remarked,
however, that acetate of iron is sometimes made directly by dissolving
scraps of iron in vinegar or pyroligneous acid (~ 380).




554                  NITRATES OF IRON.
646. Ferric Su7phate (Fe23S04) is interesting chiefly from its
analogy with sulphate of aluminum. Like sulphate of aluminum,
it combines with the sulphates of the alka-li-metals to form welldefined alums. (Compare ~ 625). Ferric salphate occurs as a
waste product in the mother-liquors from which copperas and
alum have crystallized. By drying these liquors and igniting
them, red ochre of excellent quality can be obtained, in accordance with the second reaction of ~ 645. Faming sulphuric acid
is commonly manufactured nowadays by distilling pure ferric
sulphate, instead of copperas as formerly at Nordhausen. The
ferric salt is obtained by dissolving ferric oxide in weak
sulphuric acid, and evaporating the solution to dryness; the
residue of ferric oxide left after the ignition of the sulphate is
thus reconverted into ferric sulphate, and the sulphate is used
over and over again as often as it is decomposed.
647. Ferrous Nitrate (FeNO6 + 61120) is a compound of considerable scientific interest, which may readily be procured by
dissolving ferrous sulphide in cold dilute nitric acid, or by decomposing a solution of copperas with an equivalent quantity of
nitrate of barium. It may also be obtained, mixed with nitrate
of ammonium, by dissolving iron in cold dilute nitric acid. The
metal dissolves without evolution of gas, in a manner which may
be thus formulated:4Fe +  10ON03 =  4FeN206 +  (NH4)N03 +  3H20.
The aqueous solution of ferrous nitrate decomposes readily when
heated, and in warm weather changes spontaneously to a ferric
compound.
648. Ferric Nitrate (Fe23N206) may be obtained in hydrated
crystals containing 18 molecules of water, by dissolving metallic
iron in nitric acid, of 1-29 specific gravity, till the liquor has
taken up about 10 per cent. of the metal, and then adding an
equal volume of nitric acid of specific gravity 1'43. The solution
will deposit, on cooling, rhombic prisms of ferric nitrate, which
are sometimes colorless, but often of a faint lavender-blue color.
They are slightly deliquescent, and very soluble in water, but are
only slightly soluble in cold nitric acid. By adding nitric acid to
a syrupy solution of ferric nitrate, there may be obtained another




CYANIDES OF IRON.                  555
hydrate, containing only 12 molecules of water, crystallized in
cubes or square prisms.  By mixing a solution of ferric nitrate, or,
for that matter, almost any other of the normal ferric salts, with
recently precipitated ferric hydrate, or by partially abstracting
the acid of the salt by means of an alkali, deep-red solutions of
various basic compounds may readily be obtained. A basic ferric
nitrate is employed in dyeing, under the name iron mordant.
649. Silicates of Iron.-Several native silicates of iron are
known; but none of them are of special interest. The green
tinge of ordinary glass is due to the presence of a ferrous silicate,
and by increasing the proportion of the ferrous salt, a deep bottlegreen colour may be imparted to glass. Ferric silicate, on the
other hand, has comparatively little coloring-power, though when
a considerable quantity of it is present it imparts a yellow color
to glass. It is sometimes used for coloring porcelain. To destroy the green color of the ferrous silicate, binoxide of manganese, or some other oxidizing agent, is often added to glass in
the process of manufacture; the ferrous silicate is thus converted,
for the most part, into ferric silicate, and a nearly colorless glass
produced.
650. Cyanides of Iron.-There is a ferrous cyanide (Fe(CN),),
known as a yellowish-red precipitate, which takes up oxygen
and becomes blue when exposed to the air; and a ferric cyanide
(Fe2(CN)6) has been obtained in solution.  But by far the bestknown of the cyanides of iron are certain double compounds,
which constitute the peculiar pigments known collectively as
Prussian blue.  Common Prussian blue, for example (Fe7C,~N18 +
18H20), may be regarded as a double compound of ferrous and
ferric cyanides, 3Fe(CN)2, 2(Fe,(CN)6) + 18H20; it may be prepared as follows:Exp. 343. —Add to an exceedingly dilute solution of almost any
ferric salt, such, for example, as the ferric sulphate of Exp. 332, a
drop of ferrocyanide of potassium (~ 509). A beautiful blue precipitate
will form, and will remain suspended in the liquor for a long while.
Another variety of Prussian blue, known as Turnbull's blue, may be
obtained by mixing a solution of red prussiate of potash, known to
chemists as ferricyanide of potassium, with a solution of copperas or
other ferrous salt.
Since the yellow prussiate of potash will give no blue coloration




556         DISTTI'GUISHING THE TWO OXIDES OF IRON.
with ferrous salts, and since the red prussiate yields no blue with
ferric salts, it is evident that the two solutions may be used as
tests by which to detect the presence of ferrous and ferric salts,
respectively, in any solution.
P'xp. 344.-Soak a piece of cotton cloth in a solution of ferric sulphate (Exp. 332), and then immerse it in an acidulated solution of
yellow prussiate of potash. Prussian blue will be precipitated upon
the cloth and will remain firmly attached to it. Prussian blue is
largely employed in dyeing and calico-printing in a variety of ways.
Now that we have discovered a ready means of detecting
ferrous and ferric salts, it will be well to determine experimentally
how easily the members of either of these classes may be changed
to salts of the other class.
Exp. 345.-Dissolve 4 or 5 grms. of iron tacks or wire in dilute
chlorhydric acid in a test-tube, pouring off the liquid from time to
time as it becomes nearly saturated. Test a few drops of the solution
first with ferro- and then with ferricyanide of potassium, in order to
prove that it is pure ferrous chloride. Boil the rest of the liquid with
a few drops of nitric acid to convert it to ferric chloride, and determine
when the conversion has been completed by testing as before. Finally,
divide the ferric solution into three portions. Through the first portion
pass sulphydric acid gas; sulphur will be deposited and ferrous chloride
formed,
Fe2C16 + H2S = 2FeC12 + 2HC1 + S;
to the second portion add small fragments of protochloride of tin, until
a drop of the mixture, tested with the ferrocyanide, will no longer give
a blue coloration,
Fe0C1l +  SnCl2 = 2FeCl2 +  SnC14;
boil the third portion with a fragment of metallic zinc, and determine
the fact of reduction as before,
Fe2C16 + Zn = 2FeCl2 + ZnCl2.
By leaving either of these reduced solutions in the air, or by heating
them with a little chlorate or nitrate of potassium, nitric acid, or
other oxidizing agent, they may be readily converted again to the
condition of ferric salts.
COBALT AND NICKEL.
651. Cobalt and nickel are two metals remarkably similar
to one another both in physical and chemical properties. They
are found together in nature in the same ores, in combination




COBALT AND NICKEL.                  557
with sulphur and arsenic, and are both ingredients of meteoric
iron. They can be reduced from their oxides by charcoal and
by hydrogen at high temperatures, and the metals thus obtained
can be melted about as readily as pure iron. Both cobalt and
nickel resemble iron more closely than any other common metal;
they are very tenacious, hard, aind refractory; like iron they are
magnetic, and when hot they may be forged; they rust less
readily than iron, but resemble it closely in most of their
chemical properties.  The atomic weights of cobalt and of nickel
are identical; the same number (58'8) applies to both. The
specific gravities of the two metals also are equal or nearly so,
varying in different samples from 8'2 to 8'9. Cobalt is not used
in the metallic state; but several of its compounds are remarkable
for the beauty of their color, and find important applications in
the arts as pigments, especially for coloring glass and porcelain.
A blue glass containing silicate of cobalt, obtained by fusing
oxide of cobalt with ordinary glass, is largely employed, under
the name of smalt, as a vitrifiable pigment. This coloration may
readily be exhibited by adding a minute particle of any cobalt
compound to a borax bead (~ 490) upon a loop of platinum wire,
and again placing the bead either in the oxidizing or in the
reducing flame of the blowpipe.  Nickel, on the other hand, is
used in the metallic state as an ingredient of various alloys, of
which the alloy known as German silver, composed of copper,
zinc, and nickel, is one of the most important. A whitish alloy,
obtained by adding nickel to copper, is sometimes employed for
coin of low denominations.
652. Both cobalt and nickel form protoxides (CoO and NiO),
protochlorides, and protoxide salts, like those of iron, except that
the protosalts of cobalt and nickel are far more stable than the
salts of protoxide of iron; so that the protoxides of cobalt and
nickel must be regarded as the principal oxides of these metals.
Like iron, chromium, and the other metals of the family now
under discussion, cobalt and nickel also unite with oxygen to
form sesquioxides (Co23O and Ni203), and these sesquioxides, or
at least the sesquioxide of cobalt, combine with bases to form
salts; but these salts and the sesquioxidcs themselves are comparatively unstable bodies; they are far more easily decomposed




558                      URANxIUM.
than compounds of the protoxides of cobalt and nickel, or than
compounds of the sesquioxides of the other metals of the group.
Hence, in the matter of nomenclature, the salts of the protoxides
of cobalt or nickel take precedence of the salts of the sesquioxides.
When, for example, nitrate of cobalt is spoken of, nitrate of
protoxide of cobalt is the substance referred to; whereas when
nitrate of iron or of chromium is mentioned, without further
specification, we must infer that the nitrate of the sesquioxide
is the substance meant. The use of terms in the loose manner
referred to in the foregoing examples is of course always to be
deprecated; but, in order to avoid the chance of being misunderstood, some chemists have extended to all metals having two
salifiable oxides the use of the terminations ous and ic, which
has been exemplified under iron by the terms ferrous and ferric
oxides.  Thus the terms cobaltous and cobaltic oxides, and
nickelous and nickelic oxides have been applied by some writers
to the oxides of cobalt and nickel; and there is at present a
tendency to adopt and amplify this system of names; but they
are as yet too little employed in the literature of science to find
appropriate place in an elementary manual.
URANIUI.
653. Uranium is a rare metal, found in but few localities. It
can be reduced from its chloride by means of hot potassium,
but not from its oxide by means of hydrogen. Metallic uranium
is of a steel-white color, and is somewhat malleable; it does
not oxidize in air or in water at ordinary temperatures, but
burns brilliantly when strongly heated in air. It dissolves in
chlorhydric or sulphuric acid, with evolution of hydrogen, and,
in general, is closely analogous to iron and manganese in its
chemical behavior. The atomic weight of uranium is 120; its
specific gravity is 18'4.
There are two principal oxides of uranium, capable of uniting
with acids to form salts (a protoxide UrO, and a sesquioxide
Ur20,), and two other intermediate oxides, formed by the union
of the proto- and sesquioxides in different proportions. The
sesquioxide also plays the part of a weak acid towards strong
bases. Uranium is never used as a metal; but compounds of it




SESQUIOXIDE SALTS.                559
are somewhat extensively employed for coloring glass, and to a
certain extent in photography also. Sesquioxide of uranium
imparts a beautiful greenish-yellow color to glass, and the glass
thus colored is to a high degree fluorescent; the protoxide, on
the other hand, gives a fine black, highly esteemed for painting
porcelain.
654. The salts of sesquioxide of uranium are remarkable in
that they constitute an exception to the general rule, that, to
form a normal salt, as many molecules of the acid are required
as there are atoms of oxygen in the base employed. The normal
sulphate of calcium, for example (~ 241), may be formed by the
union of CaO and S03, and the normal sulphate of sesquioxide of
iron is composed of Fe203 and 3S03; but in sulphate of sesquioxide
of uranium we find only Ur.O3,SO3, and analogous formulhe express the composition of the nitrate and other salts of this oxide.
But in spite of this peculiarity, uranium has many properties in
common with the other members of the sesquioxide group of
metals. One characteristic, for example, of this aluminum-iron
group, which is shared by uranium, is that the sesquioxides are
capable of uniting with acids, not only in the fixed and definite
proportions requisite for the normal, crystallized salts already
described, but also in very numerous indefinite proportions to
form soluble basic compounds, incapable of crystallization for the
most part, and solidifying in tough shining masses like gum
when their solutions are allowed to evaporate spontaneously in
the air. Nitrate of iron, for example, may be made as basic as
the compound 8Fe2O3,N20,, and still be soluble in water; and
between this limit, on the one hand, and that of the crystallized
normal salt (Fe03,3N,20,+ 18H,0) upon the other, sesquioxide
of iron and nitric acid can combine chemically in every conceivable proportion. The compounds of sesquioxide of iron with
other acids, and the nitrates and other salts of the sesquioxides
of the other metals of the group, all behave in a similar way, the
compounds of uranium being no exception to the rule. This tendency to form soluble, gummy, polybasic sesquisalts, so strikingly
exhibited by the members of the group of elements now under
discussion, is evidently one of those obscure manifestations of
the chemical force which we have already met with When dis



560                THE SESQUIOXIDE GROUP.
cussing the phenomena of solution (~ 49), and the law of multiple
proportions (~ 76, end).
655. The most important point of difference between uranium
and the other members of the sesquioxide group of elements is
the fact, already alluded to, that one molecule of its sesquioxide
unites with but one molecule of base to form crystallized salts,
whereas the sesquioxides of the other members of the group all
unite with acids in the proportion of one molecule of base to
three molecules of acid to form their normal crystallized salts.
Since the alums are formed by the union of a normal sulphate
of some metal of the alkali group with a teracid sulphate of some
metal of the sesquioxide group, there is no such thing as a
uranium-alum, because the teracid uranium-sulphate is wanting.
Sesquioxide of uranium, in fact, behaves among bases somewhat
as metaphosphoric acid does among acids; it stands in much the
same relation to the other teracid bases of its class, as metaphosphoric acid to the ordinary terbasic phosphoric acid.
656. The Sesquioxide Group.-The bond of union between
the metals included in this class is the fact that they all form
salifiable sesquioxides. Most of them form also salifiable protoxides; and if we arrange the metals in the order of their
atomic weights,
Gl = 14, Al = 27-4, Cr - 525, Mn = 55, Fe = 56,
Ni = 58-8, Co - 58'8, Ur = 120,
it will be apparent that the sesquioxides of the metals at the
head of the list are the most stable of the sesquioxides, and that
the protoxides of nickel and cobalt are the most stable of the
protoxides, while with manganese and iron both forms of oxide
are well represented; uranium does not conform to this arrangement.  Glucinum and aluminum have no protoxides at all;
and the protoxide of chromium is very unstable. Some of the
metals of the group' are usually bivalent, others trivalent, while
others are both bi- and trivalent.
The class of salts called alums affords strong evidence of
the existence of a natural relation between the members of the
alkali group, on the one hand, and the members of the sesquioxide group, on the other.  These highly crystallized isomorphous'salts are all moulded upon one pattern, and their atomic




ATOMIC VOLUME OF ALUMS.              561
volumes (~ 252) are very nearly equal, as the following table will
illustrate for some of the alums:Alums.            Atomnic      Specific    Atomic
Weig-lht.    Gravity.    Volume.
KA1S2O,J121HO        474'5       1'722       275'6
NaAS208,12H1120       458'4       11641       279'2
(NH4)AlS208,12H20        453'4       1'621       279'6
KCrS208, 12H20       40396       1'845       270'7
(N1H4)CrS208,12H20       478'5       1-736       275'5
(NH,)FeS208,12H20        482'0       1'712       281'4
It is a fact not unworthy of notice, that the compounds of this
group of metals, with the exception of those of glucinum and
aluminum, are for the most part colored, independently of the
colors of the substances with which they are united. The metals
of the sodium, cailrim, and magnesium groups produce colorless
compounds, unless when joined with an acid possessing a color of
its own. Glucinum and aluminum produce, in like manner,
colorless compounds; but the oxides, hydrates, chlorides, bromides,
iodides, sulphides, and oxygen-salts of chromium, manganese,
iron, nickel, cobalt, and uranium are all more or less colored in
themselves, and every color of the spectrum, from the violet at
one extremity to the red at the other, can be matched from
among the innumerable tints exhibited by the various compounds
of the last six members of the sesquioxide group.
657. With the members of the group now under discussion
are commonly classed a number of rare metals, more or less
nearly related to aluminum and iron. They are all, however,
of subordinate interest, and need only be named in this manual.
The following is a list of these elements, together with their
symbols and their atomic weights, so far as the latter have
been determined:-Yttrium, Yt, 61'6; Erbium, Er,112-6; Terbium,
Tb=(?); Zirconium, Zr=67'2; lNorium, No=(?); Cerium,
Ce=92; Lanthanum, La=92'8; nidymium, Di=95; Thorium,
Th=231'5 (?).
20




5632                EXTRACTION OF COPPEM
CHAPTER XXXI
COPPER AND MERCURY.
COPPER.
658. Though by no means one of the most abundant metals,
copper is nevertheless very widely diffused in nature, and is
largely employed by man. Traces of it exist in almost every
soil, whence it is taken up by plants, in which it may almost
always be detected by refined testing. Traces of it have repeatedly been found also in the various animal organs and
secretions.  Many natural waters contain minute quantities of
copper; its presence may often be recognized in the deposit of
oxide of iron which separates from chalybeate waters. Since
the metal occurs native in many localities, several of its valuable
properties were early recognized and made use of. Long before
the discovery of methods of reducing iron from its ores, tools
and weapons made of native copper were employed by many
barbarous nations.
Besides occurring in the native state, copper is found in a great
variety of combinations; the most common of its ores, however, is
the sulphide, or rather a compound of sulphide of copper and sulphide
of iron in varying proportions, known as copper pyrites. The processes
of obtaining copper from its ores vary greatly, according to the quality
of the ore. The oxides and carbonates may be readily reduced by
heating the ore in contact with some carbonaceous material and a
flux suitable to remove the impurities of the ore. The treatment of
ores containing sulphur is far more complicated. Such ores are roasted
in the first place, in order to convert a considerable portion of the
sulphides of copper and iron into oxides; a proper flux is then added
to the roasted ore, and the whole is melted down in either a reverberatory or blast furnace. The oxide of copper formed by roasting is
reconverted into sulphide, while much of the sulphide of iron which
had escaped oxidation before is now changed to oxide and passes off
in the slag. Sulphide of copper comparatively free from iron is thus
obtained; in other words, the copper ore is very much concentrated
by the operation. If need be, the concentrated product is subjected




PROPERTIES OF COPPER.                563
to a series of roastings and meltings, until it has been almost completely freed from sulphide of iron and other impurities. The pure or
nearly pure sulphide of copper is then roasted in a current of air, until
a certain proportion of the sulphide has been converted into oxide.
Finally the mixture of sulphide and oxide is strongly heated to a
temperature at which its ingredients react upon one another in such
manner as to yield sulphurous acid and metallic copper:Cu2S + 2CuO = SO2 + 4Cu.
Sometimes copper is obtained by precipitating it with iron from
solutions of its salts, as has been shown in Exp. 339. The copper
thus thrown down by iron is known as cement-copper, and is frequently obtained from the drainage-water of certain mines, in which
a small proportion of sulphide of copper is oxidized by the air to sulphate of copper, and so carried into solution. In some localities, lowgrade copper-ores are lixiviated with chlorhydric acid, obtained as a
waste product from the manufacturers of soda-ash, and the copper
solution subsequently made to flow over fragments of scrap iron.
The common method of assaying copper-ores is another application
of the precipitation of copper by means of iron.
659. topper is a rather hard metal, of a well-known red
color; it is very tenacious, ductile, and malleable. The specific
gravity of the metal when free from air-bubbles varies between
8'92 and 8'95. Copper melts less readily than silver, but more
readily than gold; its melting-point has been estimated to be
about 1170~.  At an intense heat it volatilizes, though for all
ordinary purposes it may be regarded as non-volatile. It is one
of the best conductors of heat and electricity known.  Its specific heat is 0'09515. Copper combines with oxygen far less
readily than iron. Even at a bright-red heat it is not capable of
decomposing water, excepting to a very slight extent. Finely
divided copper, however, soon becomes oxidized on being exposed
to the air, though, as is well known, solid masses of the metal
suffer little or no change, at the ordinary temperature, in air free
from sulphydric and carbonic acids. When strongly heated in
the air, copper quickly becomes covered with a coating of black
oxide of copper (see Exp. 12).  Metallic copper is not very
readily acted upon by acids, excepting those rich in oxygen.
The weaker acids, such as acetic acid, have no action upon it,
unless air be present, in which event the metal is soon corroded;
202




564                   ALLOYS OF COPPER.
and the same remark applies to dilute chlorhydric and sulphuric
acids. Finely divided copper, however, slowly dissolves with
evolution of hydrogen in hot concentrated chlorhydric acid; and
in hot oil of vitriol the metal dissolves readily, as has beeln  en
in the preparation of sulphurous acid. (See Exp. 96.)
Copper is readily soluble in somewhat diluted nitric acid, such
as is commonly found in commerce (see Exp. 37); but the
strongest nitric acid, of specific gravity 1'52, does not act upon
it.  When immersed in such acid the metal remains bright, and
no bubbles of gas arise from  its surface.  The phenomenon is
explained by the fact that nitrate of copper is insoluble in monohydrated nitric acid, though readily soluble in water and in
dilute acid. Ammonia-water and many salts, such as chlorids
of sodium and the various salts of ammonium, corrode copper
rather rapidly when in contact with air. Finely divided copper
takes fire in chlorine gas; and at a red heat the metal unites
directly with bromine, iodine, sulphur, silicon, and the various
metals. It does not appear to unite directly with carbon or with
nitrogen at any temperature.
Several of its compounds with other metals are of great importance in the arts.  Brass and the yellow  metal used for
sheathing ships are alloys of zinc and copper; bronze, gun-metal,
and bell-metal are alloys of tin and copper; and various compositions are produced by mixing these alloys with brass.  German
silver in its various forms is an alloy of nickel, zinc, and copper;
and copper is an essential ingredient of all the common coins,
implements, and ornaments of gold and silver.
660. Dioxide of Copper (Cu20) is sometimes found in nature
as Ruby copper; it may readily be obtained by heating protoxide
of copper with finely divided metallic copper, or other reducing
agents; so, too, when masses of metallic copper are gently heated
in the air, they become covered with a thin film of the dioxide.
Exp. 346. —Dissolve in a test-tube a few drops of honey or a bit of
grape-sugar in a little water. Add to the solution two or three drops
of a rather dilute solution of sulphate of copper, and then pour in
enough soda-lye to redissolve the precipitate which is at first produced by the lye.
Slowly heat the clear blue solution, and observe that a yellow pre



PROTOXIDE OF COPPER.                  565
cipitate of hydrated dioxide of copper soon separates, first at the
uppermost part of the column of liquid, but soon in all parts of the
tube, as its contents become sufficiently hot. When the liquor is
heated to boiling, the hydrated yellow precipitate changes after a
time to anhydrous red dioxide.
Most of the dilute acids decompose dioxide of copper with
formation of salts of the protoxide and separation of metallic
copper. But it dissolves in concentrated chlorhydric acid and in
ammonia-water, forming colorless solutions. The ammoniacal
solution may be employed as a test for the presence of oxygen in
any mixture of gases; oxygen is immediately absorbed by the
solution and a compound of protoxide of copper and ammonia
(Exp. 352) of characteristic deep blue color is formed. Dioxide
of copper is employed to a certain extent for coloring glass
ruby-red.
661. Protoxide of Copper (CuO) may be prepared by heating
the metal or the dioxide in a current of air, or by igniting carbonate, hydrate, or nitrate of copper.
Exp. 347.-Bind a bright copper coin with wire, in such manner
that a strip of wire 8 or 10 c.m. long shall be left projecting from the
coin; thrust the free end of the wire into a long cork or bit of wood,
and by means of this handle hold the coin obliquely in a small flame
of the gas-lamp. A beautiful play of iridescent colors will appear
upon the surface of the copper, particularly if it be moved to and fro.
Thrust the hot coin into water, and observe that it is at this stage
covered with a coating of red suboxide of copper. Replace the coin
in the lamp and hold it in the hot oxidizing portion of the flame
(Exp. 195); it will soon become black from the formation of protoxide of copper. After a rather thick coating of oxide has been
formed, again quench the coin in water; the black coating or scale of
oxide will fall off, and beneath it will be seen a thin film of the dioxide firmly adhering to the metal. This film of dioxide is intentionally produced upon the surfaces of many copper implements by the
manufacturers. If the coin were heated long enough it would all be
converted, first into the red dioxide, and then into black protoxide of
copper. The scales which fall off when hot metallic copper is beaten
or rolled, like those obtained from the coin in this experiment, always
consist of a mixture of the two oxides.
Exp. 348. —Evaporate to dryness in a porcelain dish upon a sandbath some of a solution of nitrate of copper prepared from copper as
in Exp. 37. There will be left as a residue a green basic nitrate of




566                  PROTOXIDE OF COPPER.
copper. Place a small quantity of this residue upon a fragment of
porcelain, and ignite it until red nitrous fumes are no longer given offl'
Pure protoxide of copper will be left upon the porcelain.
Though no oxygen can be expelled from protoxide of copper
by mere exposure to heat, all its oxygen may, nevertheless, be
removed with great facility by means of reducing agents. Oxide
of copper is, in fact, one of the most convenient oxidizing agents
in the chemist's possession, and is largely employed to this end
in the analysis of organic compounds.  When heated with carbonaceous substances, it converts all their carboIl into carbonic
acid; and, in like manner, hydrogen is immediately oxidized by
it and converted into water. Since carbonic acid and water can
readily be collected and weighed, and since their composition is
accurately known, the determination of the amounts of carbon
and hydrogen in any substance, through the agency of oxide of
copper, is merely a matter of mechanical detail.
Exp. 349.-Repeat Exp. 330, with the exception that a teaspoonful
of black oxide of copper (Exp. 348) is placed in the tube instead of the
iron rust. Water may be collected in the U-tube as before, and metallic copper will be left in the reduction-tube. But, unlike the easily
oxidizable iron, the reduced copper will not take fire in the air.
662. Protoxide of copper is soluble in most acids, with formation of salts which are blue or green when hydrated, but white
when thoroughly dried. From the solutions of most of these
salts hydrated oxide of copper may be precipitated by means of
any of the strong soluble bases.
Exp. 350.-Place in a test-tube, or small bottle, 8 or 10 c. c. of a
cold dilute solution of sulphate of copper, and add to it enough of a
solution of caustic soda to render the mixture alkaline to test-paper.
A light-blue precipitate will fall; hydrate of copper is insoluble in
water and in soda-lye.
Exp. 351.-Repeat Exp. 350, with the difference that the solutions
of caustic soda and sulphate of copper are both heated to boiling and
are mixed while hot., Instead of the blue hydrate, black protoxide of
copper will now be thrown down; for hydrate of copper readily parts
with its water when heated, even if it be all the while immersed in
water; it does not again combine with water after it has become cold.
Instead of mixing boiling solutions of the alkali and copper salt, the
moist precipitated hydrate of copper of Exp. 350 might be changed




SULPHIDES OF COPPER.                567
to black oxide by simple boiling; but the transformation would be,
comparatively speaking, slow, and the experiment less striking than
the one here described.
Exp. 352.-Again repeat Exp. 350, but instead of soda-lye add to
the copper salt ammonia-water, drop by drop, and shake the tube after
each addition of the anlmonia. Hydrate of copper will be precipitated
as before, in accordance with the reaction
CuSO4 + 2(NH4)HO =  (NH4)2SO4 + CuH202;
for, as has been said, this hydrate is insoluble in water; but, since hydrate of copper is readily soluble in ammonia-water, the precipitate
will redissolve as soon as more of this agent than is needed to decompose the copper salt is added. The ammoniacal solution of copper has
a magnificent azure-blue color.
663. The Sulphides of Copper (Cu2S and CuS) are interesting
from their occurrence as ores, and from the reactions already
briefly explained, which are so important in the industry of
copper-smelting (~ 658). The protosulphide, as obtained by adding sulphydric acid to acidulated solutions of the salts of copper,
is an important substance to the analyst; it is a black powder,
insoluble in water, in dilute acids, and in alkaline lyes.
664. Dichloride of Copper (CuC12) may be obtained by treating a mixture of protoxide of copper and finely divided metallic
copper with concentrated chlorhydric acid, or by boiling a solution
of the protochloride with sugar or some other reducing agent.
It is a white compound, insoluble in water, but soluble in strong
chlorhydric acid, and in aqueous solutions of chloride of sodium,
chloride of potassium, and many other chlorides.
665. Protochloride of Copper (CuCl2) is formed when copper is
burned in an excess of chlorine. It may readily be prepared in
the hydrated condition (CuC12+2HI20)- by dissolving oxide, hydrate, or carbonate of copper in chlorhydric acid, and evaporating
the solution upon a water-bath. Anhydrous chloride of copper
is brown; but the hydrated salt forms green, needle-shaped crystals. The concentrated aqueous solution is green; when diluted
with water it becomes blue, but turns green again on being boiled.
The dry salt fuses when heated, and at a red heat gives off half
its chlorine and is changed to the dichloride. Chloride of copper
is soluble in alcohol and ether; if some of the alcoholic solution




568                  SULPRATE OF COPPER.
is poured upon a tuft of cotton and then ignited, it will burn with
a green flame, which is characteristic of copper.
666. Sulphate of Copper (CuSO4) has been already obtained
in solution as a secondary product in Exp. 96. It may also be
readily prepared by dissolving oxide of copper, copper-scale for
example, in moderately dilute sulphuric acid. Much of it was
formerly prepared in this way for use in the arts; it is an incidental product also in the process of refining gold and silver, arid
in certain metallurgical operations. As it crystallizes from aqueous
solutions, sulphate of copper holds in chemical combination 5 molecules of water, and may then be represented by the formula
CuSO4+ 5HO20. This hydrated salt, known as blue vitriol, is the
commonest salt of copper; most of the various pigments and other
preparations of copper, medicinal or chemical, are made from it;
and it is itself used to a considerable extent by dyers and calicoprinters, and largely by the electrotypers.
Exp. 353. —Tie a piece of bladder over one end of a lamp-chimney,
or over the mouth of a wide-mouthed bottle or beaker, off which the
bottom has been somewhat evenly broken. Solder a piece of thick
copper wire to a strip of stout sheet zinc, just wide enough to enter
the chimney or bottle, and a little longer than the bottle is deep.
Place the zinc in the bottle or chimney, and sink the bottle or chimney,
with the closed end down, in a beaker or large tumbler containing a
strong solution of sulphate of copper. Fill the bottle or chimney with
di ute nitric acid, attach to the copper wire a clean medal or coin, of
which one side has been varnished, and the other rubbed over with
plumbago, and bend the wire so that the medal or coin may be immersed in the sulphate-of-copper solution contained in the beaker or
tumbler. Thorough contact must be secured between the copper wire
and the unvarnished side of the coin or medal. In a few hours a coherent film of solid, malleable copper will be firmly deposited on that
face of the coin or medal which was not protected by the non-conducting varnish. The shell of copper may be readily detached from the
coin or medal, the plumbago ensuring the ready separation of the two
metallic surfaces; it is a perfect reverse of the object copied.
This experiment illustrates on a small scale the important art of
electro-metallurgy. Plating in gold and silver, as well as in copper,
is extensively performed by a process perfectly analogous t0 that of
this experiment. Woodcuts, type, medals, maps, and engravings are




ACETATES OF COPPER.                 569
accurately copied by means of this deposition of metals from their solutions under the action of the galvanic current.
It is remarkable that the blue color of sulphate of copper depends upon the presence of water.
-Exp. 354.-Heat 1 c. c. of powdered blue sulphate of copper upon
a piece of porcelain; as it loses its water, the light-blue powder will
turn white. A drop of water upon the anhydrous powder will restore
its blue color.
If concentrated sulphuric acid be poured over the blue crystals, it
will abstract water from them, and a quantity of the white anhydrous
salt will be formed.
Since the protoxides of iron and of copper are isomorphous, or
rather since the metals iron and copper are capable of replacing
one another in many of their compounds, it is not surprising that
the sulphates of iron and copper should crystallize together to
form a compound which may contain almost any proportion of
either salt.  This isomorphous mixture of the two salts was
formerly largely employed in the. arts, and is still somewhat used
to meet the requirements of old receipts for dyeing. In a similar
way, sulphate of copper crystallizes together with the sulphates
of nickel, cobalt, zinc, and magnesium.
667. Nitrate of Copper (CuN,20) may be obtained crystallized,
by allowing the blue solution, obtained in Exp. 37, to evaporate
spontaneously in dry air. It is interesting as an example of a
salt ready to give up oxygen on slight provocation. If a piece
of tin-foil about 20 c.m. square be twisted firmly around a rather
large crystal of nitrate of copper, then pierced in several places
with a needle, and moistened with water, or with a few drops of
common spirits of wine, a powerful reaction will soon ensue. The
tin will be oxidized, at the expense of the oxygen of the nitrate
of copper, much heat will be evolved, and smoke, or even flame,
produced.
668. Acetates of Copper are formed by the action of acetic acid
upon metallic copper, exposed to the air. They are commonly
called verdigris.  Purified verdigris is the normal acetate of copper; and common verdigris is a hydrated basic acetate.  Verdigris is ulsually prepared by packing plates of copper between
woollen cloths steeped in vinegar; but sometimes, in wine-growing countries, the refuse of the wine-press is suffered to ferment




570                     MERCURY.
in contact with the copper plates. From time to time the verdigris is removed from the surface of the copper, the plate of
metal being again and again subjected to the action of acetic acid
so long as any of it remains.
In ordinary language, the term verdigris is often incorrectly
applied to the green coating of carbonate of copper, which forms
upon copper long exposed to damp air, or to the rust formed
upon copper by the combined action of air and almost any acid.
A compound of acetate of copper and arsenite of copper constitutes the beautiful and vivid green color known as Schweinfurt
green.
MERCURY.
669. Small globules of metallic mercury are sometimes found
in nature; but the principal ore of this metal is the sulphide
(HgS), called cinnabar. From this sulphide the metal is readily
extracted, by distilling a mixture of it and quicklime, or ironturnings, in cast-iron retorts. The sulphur is retained by the
lime or iron, as the case may be, while metallic mercury passes
off in the state of vapor into receivers containing water, beneath
which it condenses to the liquid state.
A rougher method of manufacture is to heat the coarsely
broken sulphide on a perforated brick arch, by a quick fire of
brush-wood; the sulphur in the ore is kindled, and, by its combustion, maintains the heat necessary to continue the distillation.
The liberated mercury is condensed in wide and long earthen
pipes, which slope first down, and then up.
670. At the ordinary temperature of the air, mercury is a
brilliant, mobile liquid, of 13'6 specific gravity, which vaporizes
slowly, even at ordinary temperatures, and boils at about 360~.
The vapor-density of mercury does not coincide with its atomic
weight. As is the case with the metal cadmium (~ 599), the atomic
weight of mercury is the weight of two unit-volumes of its vapor,
and is therefore double the vapor-density, instead of identical
with it. Into the product-volume of any compound of mercury,
one more volume is condensed than would be contained in the
product-volume of a corresponding compound of oxygen or sulphur. The atomic weight of mercury is 200, and two unit



MERCURY.                      571
volumes of its vapor weigh 200 times as much as one unit-volume
of hydrogen; its vapor-density should, theoretically, be 100; and
the results of experiment approximate closely to this number.
(Compare ~ 259.)
When cooled to -390~4 mercury freezes. In solidifying,
liquid mercury contracts considerably, and there results a ductile
metal of tin-white color and granular fracture, which may be cut
with a knife. Perfectly pure mercury undergoes no change in
air or in oxygen gas at the ordinary temperature, even when
shaken about in the gas for a long while; but if mercury containinog traces of foreign metals, such, for example, as that ordinarily met with in commerce, be exposed to the air, a gray pulverulent coating will, after a while, appear upon its surface.
This coating is composed of oxides of the contaminating metals
mixed with finely divided metallic mercury.  A similar gray
powder of finely divided mercury may be obtained by triturating
mercury with sulphur, tallow, and a variety of other substances,
or simply by shaking it with water or oil of turpentine. When
heated in the air to temperatures near its boiling-point, even
pure mercury absorbs oxygen, and is converted into the red oxide
(~ 672). Metallic mercury combines directly with chlorine, bromine, iodine, and sulphur.
Chlorhydric acid has no action upon mercury, not even when
it is hot and concentrated. Dilute sulphuric acid has scarcely
any action upon it; but hot concentrated sulphuric acid converts
it into solid sulphate of mercury, while sulphurous acid is evolved
(Exp. 96). Nitric acid, even when dilute, dissolves it easily.
Large quantities of mercury are used in extracting gold and
silver from their ores, for silvering mirrors, and in the process of
fire-gilding. Preparations of mercury are employed also as medicaments, and for various purposes in the useful arts.  The
fluidity of the metal makes it valuable in the construction of
certain philosophical instruments, of which the thermometer and
barometer are familiar examples.
There are two oxides of mercury —a black dioxide, and the red
protoxide; each of these oxides unites with acids to form a peculiar class of salts.
671. Dioxide of Mercury or Mercurous Oxide (HgO0) is best




572                RED OXIDE OF MERCURY.
prepared by decomposing one of its salts, calomel for example
(~ 674), by means of caustic soda. Though a rather powerful
base when in combination, it decomposes readily when in the free
state, mere exposure to light or to gentle heat being sufficient to
decompose it into metallic mercury and the red oxide. With
acids it combines readily, with formation of salts of the dioxide
of mercury, or mercurous salts as they are often called.
672. Protoxide of Mercury or Mfercuric Oxide (HgO) may be
prepared by heating metallic mercury in the air as above described,
or, better, by heating nitrate of mercury at a temperature high
enough to drive off oxides of nitrogen, but at the same time too low
to decompose the oxide of mercury, or, again, by precipitating the
solution of a salt of protoxide of mercury by means of a caustic
alkali. As obtained by the first two methods, it is a compact,
granular, almost crystalline, glistening powder, of bright brickred color; but when prepared by the last method it is, when dry,
a soft, light-orange-colored powder. Very considerable differences in the chemical deportment of these red and yellow varieties
of the oxide have been noticed, though the differences are hardly
so great as are usually found between the isomeric modifications of
other substances. The precipitated yellow oxide is, for example,
more readily decomposed by heat and by chlorine than the red
oxide. The red oxide, however, is the substance known as oxide
of mercury in commerce and the laboratory.
Considered as a source of oxygen, red oxide of mercury is
peculiarly interesting, since by means of it oxygen may be derived
directly from the air (~ 12); but it neither affords the gas
cheaply, nor yields an abundant supply. Since red oxide of
mercury contains only a single atom of oxygen for each atom of
mercury, and since the atomic weight of mercury (200) is comparatively high, any given weight of the oxide can, of course, contain but a small proportion of oxygen: —for
216:      16    =       1:    0'074
Wreight of      Weight of      Grin.        Grm.
a mo'ecule       an atomn
of HgO.           of O.
Although it contains so small a proportion of oxygen, compared
with the nitrates and chlorates commonly employed for effecting




CALOmEL.                      573
oxidation, yet, from the facility with which it gives up its oxygen,
mercuric oxide is still an oxidizing agent of considerable power.
If a small portion of it be mixed with a little sulphur, and then
heated, the mixture will explode; so, too, if it be mixed with a
small fragment of phosphorus, and struck with a hammer upon
an anvil, a similar explosion will ensue; the violent action depends upon rapid oxidation in both cases.
Mlost acids unite readily with oxide of mercury to form salts,
often spoken of as mercuric salts. Both the oxides of mercury
are, like the protoxide of lead (~ 575), remarkable for the facility
with which they form basic salts. In general, the compounds of
mercury unite with one another readily to form a great variety
of double compounds and abnormal basic salts, such as the oxychlorides xHgO,HgC1,, and the chlorosulphide 2HgS,HgCl2. The
properties of several of these complex substances are interesting;
but none of them fairly fall within the scope of this manual.
673. Disulphide of Mercury (Hg2S) is a compound nearly as
unstable as the dioxide; but the protosulphide (HgS) is a permanent substance of considerable importance in the arts. Artificially prepared for use as a pigment, it is known under the name
of vermilion. It is the most important ore of mercury, as has been
already stated (~ 669).
674. Mercurous Chloride (HgCl), commonly called calomel, is
extensively used as a medicament.
It may be prepared either by heating together a mixture of metallic
mercury and corrosive sublimate (~ 675) until the dichloride sublimes,
Hg + HgC12 = 2HgCl,
or by subliming an intimate mixture of equal parts of sulphate of dioxide of mercury and common salt,
Hg2SO4 + 2NaCl = 2HgC1 + Na2SO4.
By the way of precipitation it may be made by mixing together solutions of common salt and nitrate of dioxide of mercury:Hg2N206 + 2NaC1 = 2HgC1 + Na2N206.
In place of the corrosive sublimate used in the first method, intimate
mixtures of sulphate of protoxide of mercury and common salt, or of
common salt and black oxide of manganese or sulphate of sesquioxide
of iron may be substituted.




574                  MERCURIC CHLOR1DE.
Calomel is a heavy white powder, which volatilizes at temperatures below redness without previous fusion.
The vapor-density of calomel is, by calculation, 117'75.
Weight of one atom, or two unit-volumes, of mercury. 200'00
Wreight of one atom, or one unit-volume, of chlorine.  355
Weight of two unit-volumes of mercurous chloride.. 235'5
Weight of one unit-volume of mercurous chloride.. 117-75
The vapor-density has been determined by experiment at
120-49. By slowly cooling its vapor, prismatic crystals of calomel may readily be obtained. Unlike metallic mercury, calomel
does not volatilize at ordinary temperatures. It is tasteless, odorless, and as good as insoluble in water. Alkaline lyes decompose
it readily, and it is slightly soluble in many saline solutions.
675. Mercuric Chloride (HgC12), better known by the name
corrosive sublimate, may be prepared by burning mercury in an
excess of chlorine gas, by dissolving protoxide of mercury in
chlorhydric acid, or by dissolving metallic mercury in aqua regia,
and evaporating to dryness.
In practice, it is commonly prepared by sublimation, by carefully
heating an intimate mixture of sulphate of mercury and common
aalt:HgSO4 + 2NaC1 = HgCl2 + Na2SO4.
Another method is, to add concentrated chlorhydric acid to a strong
Loiling solution of nitrate of dioxide of mercury, as long as a precipitate is formed, and to subsequently boil this precipitate with a quantity of chlorhydric acid equal to that used in preparing it:HgNO3 + 2HC1 = HgC12 + H20 +- NO2.
Beautiful crystals of mercuric chloride will be deposited as the hot
solution cools.
676. Mercuric chloride commonly occurs in commerce, in translucent crystalline masses; but crystals of it may readily be obtained, by careful sublimation, as well as by slowly cooling hot
solutions., It melts at about 265~, forming a colorless liquid
which boils at 2930; the fumes are acrid, and, like the salt itself,
exceedingly poisonous.
The vapor-density of corrosive sublimate is, by calculation,
135.




MERICURIC IODIDE.                575
Weight of one atom, or two unit-volumes, of mercury.. 200
WTeight of two atoms, or two unit-volumes of chlorine.. 70
Weight of two unit-volumes of mercuric chloride...  270
Weight of one unit-volume of mercuric chloride.... 135
The best experiments assign to the vapor of the salt the density of 141. Mercuric chloride is rather easily soluble in water
and alcohol; with the alkaline chlorides it unites to form salts
which are easily soluble in water. These double salts are so
numerous and well defined, that they are regarded as chlorine
salts comparable with the sulphur salts (~ 340) and the oxygen
salts. In this view, protochloride of mercury would be called
chloromercuric acid, and its compounds with the alkalies chloromercurates. The compound NaCl,HlgCl2, for example, would be
called chloromercurate of sodium, and the compound NaCl,2HgC12
the bichloromercurate. Corresponding double chlorine salts are
formed by the union of the chlorides of gold and of platinum with
the chlorides of other metals and compound radicals, as will appear in the sequel.
677. Mercuric chloride unites with many organic substances to
form compounds insoluble in water and imputrescible.  It coagulates albumen, for example, and the more perishable portions
of wood; hence the employment of raw whites of eggs as an
antidote in cases of poisoning by corrosive sublimate, and the use
of the mercury salt for preserving wood, —a purpose for which it
would, no doubt, be largely employed were it not for its high
cost. Collections of dried plants, and of other objects of natural
history, are preserved both from decay and from the attacks of
insects by brushing over them a solution of the chloride in alcohol.
It is worthy of mention that the compound of albumen and chloride of mercury, though insoluble in water, is soluble in an excess
of albumen.
Mercurous and mercuric bromides, iodides, fluorides, and cyanides are, in general, analogous to the corresponding chlorides.
Mercuric iodide undergoes remarkable changes of color when
heated or subjected to friction.
Exp. 355.-Dissolve half a gramme of iodide of potassium in a small
quantity of water; also dissolve 0'4 of a gramme of corrosive sublimate
in a little water, and mix the two solutions. Collect upon a filter the




576                 NITRATES OF MERCURY.
beautiful red precipitate which is formed, wash it carefully with water
and dry it in the air. Place a portion of the dry red powder in a porcelain capsule; invert over the capsule a small glass funnel, and heat
the capsule moderately upon a sand-bath; the iodide will melt, sublime,
and finally be deposited upon the cold walls of the funnel in yellow
crystals. On rubbing these crystals with a glass rod, their color will
change again to red. Indeed the change of color often occurs of itself
as the crystals cool, without friction. The composition of the iodide is
neither changed by the sublimation nor by the friction; the change of
color is due to a change of crystalline form-mercuric iodide being dimorphous, and exhibiting a red color in its octahedral form, and a yellow
color when crystallized in rhombic prisms.
The change of coloration may be shown in another way, by dissolving some of the precipitated iodide in alcohol. The alcoholic solution
is colorless and4 appears to contain the yellow modification of the
iodide; on pouring it into water, iodide of mercury is precipitated as a
yellow powder, which soon changes to red.
678. Su7phates of Mercury.-There is a sparingly soluble sulphate of dioxide of mercury (Hg2SO4), a normal sulphate of the
protoxide (HfgSO4), and a basic sulphate of the protoxide (of composition 3HgO,SO,). Normal mercuric sulphate may be prepared
by dissolving metallic mercury in an excess of boiling concentrated
sulphuric acid, and evaporating the solution to dryness. It is the
material from which many other compounds of mercury are derived. It is decomposed by water; an insoluble trisulphate is
thrown down, while but a small proportion of mercury remains
dissolved in the dilute sulphuric acid which is formed.
679. Nitrates of Mercury are numerous. There are at least
four nitrates of the dioxide, and as many of the protoxide namely
the normal salts and three basic salts in either case. Both of the
normal salts are soluble in water, and are commonly kept in the
laboratory as examples of the mercury salts. The nitrate of the
dioxide is prepared by digesting an excess of metallic mercury in
cold moderately strong nitric acid.  The solution should be kept
in closed bottles containing a few globules of metallic mercury.
The nitrate of the protoxide may be readily obtained by dissolving
red oxide of mercury in an excess of nitric acid.
680. Amalgams.-Mercury unites with most of the other
metals, forming alloys, many of which are pasty, or liquid when
the proportion of mercury contained in them is large. These




TITANIUM.                     577
alloys are commonly called amalgams, in contradistinction to the
ordinary solid alloys of the other metals, in which mercury has
no place. The liquid amalgams are true solutions of other metals,
or of solid amalgams, in the fluid mercury. The so-called silvering of mirrors is an amalgam of tin.
Mercury may be detected in almost any soluble salt of the
element by introducing into a solution of the salt a piece of clean
copper.
Exp. 356.-Place a drop of a solution of either of the nitrates or
chlorides of mercury upon a clean copper coin and rub the liquid over
its surface. A white coating of metallic mercury will be deposited
upon the metal.
681. Copper and mercury are classed together partly because
of certain resemblances between the two metals, but also because
neither of them falls naturally into either of the other groups of
elements. They are alike in that they are not readily acted upon
by air, excepting at high temperatures, that they do not decompose water at any temperature, and that they both form two salifiable oxides, and two chlorides of analogous composition. They
are both acted upon in the same way by nitric and by sulphuric
acids, the acid being reduced to a lower degree of oxidation,
while the metal is dissolved, as has been seen in Exp. 96. As
the formulse of their compounds have doubtless already suggested,
mercury and copper are univalent, like the alkali-metals, in the
mercurous and cuprous compounds, but bivalent in the mercuric
and cupric compounds.
CHAPTER XXXII
T I T A N 1 U M —T I.N,
TITANIUM;
682. This comparatively rare metal; is fbund in several minerals, Such as rutile and titaniferous iron, in the condition of titanic
acid, TiO2. None of its compounds are employed in the arts, and
2 P




578                         TIN.
the element itself is here mentioned mainly on account of the analogies which it bears to tin. Titanic acid is isomorphous with stannic acid (~ 685), and resembles it closely in its chemical deportment. Sesquioxide of titanium, Ti20,,, corresponds to sesquioxide
of tin, Sn203 (~ 685); and in the same way that the latter may be
regarded as a stannate of tin, SnO,SnO2, the titanium compound
may be considered a titanate of titanium, TiO, TiO2.
The bisulphide of titanium, and the bichloride, bibromide, and
bifluoride, correspond in like manner to the tin compounds.
TIN.
683. Though by no means widely diffused in nature, and though
ores of it occur in but few localities. tin is one of the metals which
have longest been known to man. The fact admits, however, of
ready explanation; for the specific gravity of the ores is high, and
the metal is easily reduced from them by simple heating with
charcoal. From the manner of the occurrence of many of these
ores, in the beds of torrents, it is evident that their great weight
would be likely to attract attention, and that their behavior towards fire would soon be noticed. The simplest possible metallurgical operation, and the one most likely to suggest itself to
savage men, is the heating of a heavy stone in a wood fire.
The principal ore of tin is the binoxide, called tin-stone. In
order to extract the metal from it, the tin-stone is mixed with
powdered coal and heated upon the hearth of a reverberatory
furnace in a reducing flame. The reduced metal melts readily,
and is then run out of the furnace into iron moulds. Tin is a
lustrous white metal, soft, malleable and ductile, though not very
tenacious. Its ductility varies greatly with the temperature; at
1000 the metal may be drawn into thin wire, but at 2000 it is very
brittle. When a bar of tin is bent, it emits a peculiar crackling
sound, and if the bending be repeated, the metal becomes (decidedly warm.  These phenomena appear to depend on the
disturbance of interlaced crystals contained in the bar, and upon
the friction of these crystals one against the other. Tin alwavs
exhibits a great tendency to assume crystalline form in passing
from the liquid to the solid condition.  Upon this peculiarity is
founded a method of ornamenting tinned iron.




TIN.                         579
Exp. 357.-T-eat a piece of common tinned iron over the gas-lamp
until the tin has melted, thrust the plate into cold water in order that
the tin may harden quickly, then remove the smooth surface of the
metal by rubbing it first with a bit of paper moistened with dilute
aqua regia, and then with paper wet with soda-lye. By this treatment
there will soon be laid bare a new surface covered with beautiful crystalline figures, like frost upon a window-pane. The plate should then
be washed thoroughly with water, dried quickly, and covered with
some transparent varnish.
The same crystalline structure can be brought out, though less conspicuously, by removing the outside polished surface of almost any
piece of tin plate by means of warm dilute aqua regia, without first
heating the plate as in this experiment.
Tin does not tarnish in the air at ordinary temperatures, no
matter whether the air be moist or dry; but when strongly heated
it oxidizes rapidly, and even burns with a brilliant white light.
The specific gravity of tin is about 7'3; its atomic weight is 118.
It melts at about 230 —at a lower temperature than any of the
other common metals. At very high temperatures it is slihhtly
volatile.
On account of its brilliant lustre, and its power of resisting
atmospheric action, tin is largely employed for coating other
metals,-copper, for example, as in ordinary pins, cooking-vessels,
and bath-tubs-and iron, as in common sheet-tin, of which the
so-called tin ware is manufactured.
Exp. 358. —Thoroughly clean the surface of a copper coin, or of a
small piece of sheet-copper, by means of dilute sulphuric acid; place
the copper over the gas-lamp, and melt upon it a bit of tin as large as
a pea. Rub the melted tin over the copper with a rag. It will not
adhere to the copper; for although the latter was once carefully cleaned,
it afterwards became coated with oxide of copper in such manner that
the tin could not come in contact with the metal.
Repeat the experiment as before; but when the tin has melted, strew
over the copper some finely powdered chloride of ammonium. On now
rubbing the tin against the copper, the two metals will adhere firmly.
The chlorine of the chloride of ammonium unites with the copper of
the oxide of copper to form fusible, volatile chloride of copper, while
ammonia and water are set free, as may be perceived by the odor. The
excess of tin should be wiped off with a rag, so that a smooth surface
may be left upon the coin.
2P2




580                  PROTOXIDE OF TIN.
When in contact with dilute acids, or with alkalies, tin slowly
absorbs oxygen from the air and goes into solution. Of the
strong acids, nitric acid acts upon it violently, with formation of
insoluble hydrated binoxide of tin; a certain amount of water is
decomposed as well as the nitric acid, in this reaction, and some
nitrate of ammonium is formed; the ammonium comes from the
union of the nascent hydrogen and nitrogen of the deoxidized water
and nitric acid (see ~ 92). Hot concentrated chlorhydric acid gradually dissolves tin, and gives off hydrogen. Boiling concentrated
sulphuric acid converts it into sulphate of tin, with evolution of sulphurous acid; but dilute sulphuric acid has no action upon it out
of contact with the air. When heated with concentrated soda
or potash lye, tin slowly dissolves, with formation of soluble stanxnate of sodium or of potassium, and evolution of hydrogen.
There are two prominent oxides of tin, a protoxide and a bioxide, besides intermediate oxides compounded of these two,
The binoxide occurs, moreover, in combination, in different isomeric modifications.
684. Protoxide of Tin (SnO) may be obtained as a black
powder by heating its hydrate in an atmosphere of carbonic acid
or other inert gas. The hydrated protoxide is prepared by adding
the solution of an alkaline carbonate to a solution of tin in chlorhydric acid; the hydrate is thrown down as an insoluble precipitate while carbonic acid escapes:SnCl2 +  Na2CO, + HF20 =  SnH202 +  2NaC1 + CO2.
The hydrate rapidly absorbs oxygen from the air when moist,
but is tolerably permanent when dry. The anhydrous oxide
undergoes no change in air at the ordinary temperature; when
touched with a glowing coal, it takes fire and burns vividly,
being converted into the binoxide. The hydrate also burns in
the same way when lighted. It is remarkable that the anhydrous oxide is more readily soluble in acids than the hydrate;
but in alkaline lyes only the hydrate is soluble. Most of the
salts of protoxide of tin greedily absorb oxygen from the air and
from many oxygenated substances. They are much employed as
reducing agents.
Exp. 359.-To 5 or 6 c. c. of a solution of corrosive sublimate add
a few drops of protochloride of tin, and heat the mixture; a gray




BINOXIDE OF TIN.                  581
powder will separate; it is metallic mercury, very finely divided,
which has been reduced from the mercuric chloride. The powder
boiled with chlorhydric acid agglomerates into visible globules:HgCI2 — + SnC12 = SnCI4 + Hg.
The protochloride of tin has abstracted the chlorine from the chloride of mercury.
685. Binoxide of Tin, or Stannic Acid (SnO2).-This oxide
occurs in nature as the principal ore of tin, as has been stated in
~ 683, and may readily be prepared artificially by roasting the
metal in a free current of air, or by igniting hydrated binoxide of
tin. It is insoluble in water, acids, and alkalies, and in general
is not readily acted upon by chemical agents. When fused with
caustic soda, however, it combines with it to form a soluble compound. Like the hydrated oxides of phosphorus and antimony,
hydrated binoxide of tin (SnH2O,) is remarkable for the different
chemical properties it exhibits when prepared in different ways.
As obtained by heating metallic tin with concentrated nitric acid,
it is almost absolutely insoluble in some acids, and dissolves only
with difficulty in others. The hydrate obtained by precipitation
from solutions of bichloride of tin and of an alkaline carbonate,
is very readily soluble in acids. By ignition, the soluble hydrate
is converted into insoluble anhydrous oxide of tin.
The difference in chemical behavior above mentioned is not
confined to the hydrates alone, it exists as well in the compounds
formed by the union of these hydrates with other substances,
both acids and bases; hence the names stannit acid, applied to
the soluble hydrate, and metastannic acid, applied to the insoluble modification.  The generic terms stannate and metastannate applied to the compounds of these two varieties of the
oxide are employed precisely as in the case of phosphoric and
antimonic acids (~~ 293, 352). Both modifications are soluble
in alkaline lyes, but the one representing metastannic acid is far,ess readily soluble than the other.
The stannates of the alkali-metals, sodium and potassium,
crystallize readily from their aqueous solutions; but the corresponding metastannates do not crystallize, they are insoluble in
saline solutions, and may be precipitated as gelatinous masses by
adding almost any neutral salt of an alkal:-metal to their aqueous




582                   SULPHIDES OF TIN.
solutions.  Among the stannates, that of tin (SnO,SnO2= Sn208)
is worthy of mention, since it is often described as the sesquioxide of tin. Stannate of sodium is somewhat extensively employed in the printing of mousselines-de-laine.
A good method of preparing it is to boil granulated tin, or
scraps of tinned iron, in a solution of litharge in an excess of
caustic soda:
Sn +  NaPb203 =  Na2SnO, +  2Pb.
Or a mixture of caustic soda, nitrate of sodium, and metallic tin
may be melted in an iron kettle, a certain portion of chloride of
sodium, or of stannate of sodium, from a previous fusion, being
added to the mixture in order to mitigate the force of the
reaction.
686. The Sulphides of Tin (SnS, Sn2S3, and SnS2) correspond
to the oxides. The protosulphide (SnS) is precipitated as a darkbrown powder when sulphydric acid is added to the solution of a
salt of protoxide of tin. The bisulphide (SnS2) when prepared in
the dry way is a beautiful yellow compound, known as mosaic
gold, or bronze powder, and is somewhat employed in decorative
painting.
Exp. 360.-Prepare a quantity of tin amalgam by heating together
in a glass flask 12 grms. of granulated tin and 6 grins. of mercury.
Rub the amalgam in a porcelain mortar together with 7 grms. of
sulphur and 6 grms. of chloride of ammonium until the different ingredients have been thoroughly incorporated one with the other.
Place the mixture in a small, long-necked, glass flask, and slowly
heat it to low redness upon a sand-bath. After an hour or two there
will be found at the bottom of the flask a quantity of bisulphide of
tin, in the condition of soft, beautiful, golden-yellow powder, of flaky
texture, while in the neck of the flask there will be found a deposit of
chloride of ammonium contaminated with sulphur, sulphide of mercury, and protochloride of tin.
Instead of the amalgam and the proportions of the other ingredients
as given above, there may be heated in the flask an intimate mixture
of 2 grms. of dry protosulphide of tin, half a grm. of sulphur, and
1 grm. of chloride of ammonium.
The part played by the chloride of ainmonium in these experiments
is not well understood; it is known only that the presence of this salt
promotes the formation of a brilliant golden-colored product, though




PROTOCHLORTDE OF TIN.                583
there is no evidence that the chloride either undergoes or produces
any chemical change. It is by no means improbable, however, that
by volatiliLing at the right moment the chloride of ammonium may
so mode-rate the heat engendered by the combination of the sulphur
and the tin that the temperature of the mixture is prevented from
reaching a point at which the bisulphide would be decomposed.
The Chlorides of Tin are perhaps more important than any
other compounds of the metal.
687. Protochloride of Tin (SnCl2) is obtained by dissolving
granulated tin in boiling concentrated chlorhydric acid.  On
evaporating the solution, and allowing it to crystallize, prismatic
hydrated needles are obtained of the composition SnC12 +2H2O.
These crystals are largely used by dyers and calico-printers,
under the name of tin-salt. Protochloride of tin, whether in the
condition of crystals or in solution, rapidly absorbs oxygen from
the air, and is converted into a mixture of bichloride of tin, and
an insoluble oxychloride. It must therefore be kept in tight
packages. The pure salt is completely soluble in a small quantity of water; but when this solution is mixed with a large
quantity of water, it decomposes; a highly acid solution of protochloride of tin in chlorhydric acid remains in solution, while a
precipitate of oxychloride of tin (SnO,SnC1,+ 21120) subsides.
Protochloride of tin is a powerful reducing agent, as has been
shown in Exp. 359; it combines readily either with oxygen or
with chlorine, and is frequently employed to remove these elements from their compounds. By means of it, the oxides or
chlorides of arsenic, antimony, gold, silver, or mercury may be
reduced to the metallic state; the salts of sesquioxide of iron,
and of protoxide of copper, may be reduced to the degree of
protoxide and of dioxide respectively, while acids such as chromic and manganic are reduced to the condition of basic oxides.
Sulphurous acid is reduced by it in such wise that a precipitate
of sulphide of tin is formed, when solutions of protochloride of
tin and of sulphurous acid are mixed; it converts blue indigo to
white indigo, and is capable of abstracting oxygen from a host of
other substances.
At the temperature of 1000, all the water may be expelled
from the crystallized salt; but some of the chlorhydric acid is




584                  BICHLORIDE OF TIN.
liable to go off at the same time, so that it is not easy in this
way to obtain the anhydrous salt in a state of purity. A better
way of obtaining the anhydrous salt is to heat together equal
weights of finely divided tin and corrosive sublimate:HgCl2 + Sn = SnCl2 + Hg.
The dry chloride of tin remains as a residue, while metallic
mercury goes off. The anhydrous salt may itself be distilled at
a full red heat.
Protochloride of tin unites with many of the metallic chlorides
to form double compounds, which may be called chlorostannites.
688. Bichloride of Tin (SnC1,).-When anhydrous, this compound is a fuming, volatile; colorless liquid, of 2'7 specific gravity.
It does not solidify at -20~, but boils at 1150. When exposed
to the air it gradually absorbs water, and, after a while, hydrated
crystals are formed. When mixed with about one-third its
weight of water, it solidifies to a mass of hydrated crystals,
much heat being at the same time evolved. These crystals are
readily soluble in a small quantity of water; but when treated
with. much water they decompose, hydrated binoxide of tin is
precipitated, and free chlorhydric acid passes into solution.
In order to prepare anhydrous bichloride of tin, chlorine gas
may be passed over hot chloride of tin, or melted metallic tin;
or an intimate mixture of 1 part of tin filings and 4 or 5 parts
of corrosive sublimate may be distilled in a retort. The hydrated
bichloride and in general all solutions of the bichloride are obtained, either by dissolving tin in dilute aqua regia, by passing
chlorine into solutions of protochloride of tin, or by heating the
protochloride with chlorhydric acid to which a little nitric acid
has been added. The anhydrous salt may be prepared from the
hydrate, by distilling the latter with concentrated sulphuric acid,
which retains the water.
Like the protochloride, bichloride of tin is largely employed in
dyeing. It combines also with the alkaline and other chlorides
to form salts, known as chlorostannates.  The substance called
pink salt, in commerce, employed in the preparation of pink colors
upon calicoes, is a chlorostannate of ammonium, 2NH4Cl,SnCl4.
There are, of course, many other salts of tin, but none of them
are of sufficient interest to be mentioned here.




MOLYBDENUM.                    58.5
689. The alloys of tin are important. The composition of
bronze, bell-metal, &c. has been already mentioned under copper
(~ 659), that of stereotype-metal under antimony (~ 348), and
that of tin amalgam under mercury (~ 680). Of the other alloys
of tin, those formed by its union with lead are most remarkable.
Plumbers' solder consists commonly of equal parts of lead and tin,
though some kinds of it contain only one-third their weight of
lead, and others only one-third their weight of tin. Pewter is
composed of tin, together with a small proportion of lead.
690. With tin and titanium may be classed the two exceedingly rare metals Columbium (niobium) and Tantalum. The
principal source of columbium is the rare American mineral
columbite. Tantalum is procured from the Scandinavian mineral
tantalite. J The four metals form binoxides and volatile chlorides
containing four atoms of chlorine, and are therefore sometimes
quadrivalent. In this respect the tin group differs from all the
groups heretofore studied, excepting the group composed of
carbon, boron, and silicon. Tin and titanium have a like mode
of occurrence in nature.
C HAPTER TXXXIII.
M 0 L YB D E N U M —  A N A D I U M-T U N G S T E N.
MOLYBDENUM.
691. This rare element is generally found in nature in combination with sulphur, as bisulphide of molybdenum. This bisulphide is a mineral closely resembling graphite and galena in
appearance. The name molybdenum is derived from a Greek
word sometimes applied to galena. Molybdenum is a white
metal almost as lustrous as silver, of specific gravity 8'6. Its
atomic weight is 96.
There are three oxides of molybdenum-a protoxide (MoO)
and a binoxide (MoO2), both acting as bases, and a teroxide
(MoO,), which is a strong acid known as molybdic acid. Molyb



586               VANADIUM-TUNGSTEN.
date of ammonium is a salt much valued by the analyst, since by
means of its solution very small quantities of phosphoric acid may
be detected; a double compound of molybdate and phosphate of
ammonium is deposited as a yellow crystalline precipitate.
VANADIUM.
692. Vanadium is a metal somewhat resembling molybdennm
on the one hand, and having certain analogies with chromium
upon the other. It appears to be related to the members of the
nitrogen group also. Though seldom found in large masses,
it appears to be rather widely diffused in nature, traces of it
often accompanying the ores of iron, for example. It has four
oxides, viz., V2 02- V2 03- V2 04 and V2 05. The last acts
as an\ acid and forms salts by combining with bases.,The
atomic weight of vanadium is 51.3.
TUNG STEN.
693. The element tungsten is far less rare than the other members of the group now under discussion. It is found in considerable quantities in combination with oxygen, iron, and manganese,
in the mineral wolfram, whence the Latin name of the element
wolframium, and the symbol W. The mineral scheelite also contains tungsten, in combination with oxygen and calcium. Metallic
tungsten may be reduced from its oxides by means of hydrogen
gas at a bright-red heat, or by charcoal at a white heat. It is a
hard iron-gray metal, of specific gravity 17'6, and very refractory.
Its atomic weight is 184. The metal has been employed to a
certain extent in the preparation of steel; a small quantity of it
added to steel has been found to greatly increase the hardness
of the steel, and to impart to it other valuable properties.
694. There are two oxides of tungsten: —a binoxide (WO2),
which does not unite with acids to form salts, but acts rather as
an acid; and a teroxide (WO3) called tungstic acid. Tungstic
acid by uniting with bases forms a large number of salts, many
of which are of very complex composition. The most important
ore of tungsten, wolfram, is a mixture in varying proportions of
the tungstates of the protoxides of iron and of manganese. The
general formula of the mineral may be written RO, WO3, in which




GOLD.                      587
R stands for either iron or manganese; but there are nevertheless
two special varieties of the mineral, one tending to correspond
with the formula 2(FeO,WO,), 3(MInO,WO,), in which the proportions of iron to manganese are as 2 to 3, and the other to the
formula 4(FeO,WO,); MnO,WO,, in which the relation of the iron
to the manganese is as 4 to 1. The first variety, richer in manganese, is of lower specific gravity than the variety rich in iron.
But, since the atomic weights of iron and of manganese are nearly
equal, the proportion of tungstic acid is almost absolutely the
same in both varieties of the mineral. Wolfram is a very heavy
mineral, its specific gravity being as high as 7'3. Indeed the
name tungsten is derived from Swedish words meaning heavy
stone.
Tungstate of sodium has been employed to a small extent for the
purpose of rendering cotton and linen uninflammable. If a weak
solution of the tungstate be added to the starch employed to stiffen
light fabrics, the cloth therewith impregnated may be exposed to
fire without inflaming; it will simply be slowly charred.
The compounds of tungsten are remarkably similar to those of
molybdenum. The metal resembles both molybdenum and vanadium in forming an acid teroxide, a binoxide, and a volatile terchloride. Like molybdenum and vanadium, it decomposes water
at high temperatures.
CHA PTE R XXXIV.
GOLD AND PLATINUM.
GOLD.
695. Though generally found only in small quantities, gold is
very widely diffused upon the surface of the globe. Traces of it
may be found beneath the sandy beds of most rivers, and it
occurs in many of the crystalline rocks and in the soils resulting
from their decomposition. Many varieties of iron pyrites, in particular, contain appreciable quantities of gold, and silver is never




588                 PROPERTIES OF GOLD.
found in nature altogether free from it. It occurs in the lead
and copper of commerce, as well as in the ores from which these
metals are derived and in many of the salts obtained from them,
and has been detected in various other metals; it is, in short,
almost everywhere. The chief source of the metal as an article
of commerce is native gold; this is sometimes found in a condition
of purity, but is usually alloyed with more or less silver. It is
collected, either directly by mechanically washing away the
lighter substances with which it is associated, or, in the case of
poorer ores, the gold is dissolved out chemically by means of
quicksilver, and is subsequently recovered from the amalgam by
filtration and distillation.
The separation of gold from the rocks and sands in which it
occurs is a process attended with much labor; hence gold is one
of the costliest of metals. The price of a gramme of gold is about
sixteen times that of a gramme of silver, and twice as great as
that of a gramme of platinum.
696. Pure gold is remarkable as being the most malleable of
the metals, and as being the only metal of a decided yellow color;
also for its softness, which is nearly as great as that of lead. It
has, however, much tenacity, and may be drawn into extremely
fine wire; 1 grm. of gold can be made to yield as much as 3 kilometres of wire. The metal can be beaten into leaves which are
not more than one ten-thousandth of a millimetre thick. Very
thin leaves of gold are transparent, transmitting a green polarized
light.
Next to platinum, gold is the heaviest of the ordinary metals;
its specific gravity varies from 19-26 to 19'37, according as it
has been more or less compressed. Its atomic weight is 196.
It melts somewhat less readily than copper or silver, at a temperature estimated to lie between 12000 and 1250~. Its power
of conducting heat and electricity is greatly inferior to that of
silver. It is not volatile to any great extent at the melting temperature; but at higher temperatures, such as it is subjected to
in the ordinary processes of melting and refining, the metal wastes
considerably; and at the temperature obtained by the oxyhydrogen blowpipe the metal goes off as a thick vapor.
697. In the air, gold undergoes no change at temperatures




REFINING OF GOLD.                  589
lower than its melting-point; and upon this fact, taken in connexion with the beautiful color and lustre of the metal, and its
comparative rarity, its principal uses depend.
On account of this indestructibility, gold was regarded by the
earlier chemists as the king of metals; together with platinum
and silver, it is still spoken of as a noble metal. Few chemical
ag ents, excepting melted metals, have any action upon gold.
None of the common acids, when taken singly, can dissolve it,
though the metal is completely soluble in a mixture of chlorlhvdric and nitric acids (~ 104), and is not completely insoluble
in nitric acid contaminated with nitrous or hyponitric acid. The
elements chlorine and bromine, however, unite with it in the cold;
and when hot it is attacked by phosphorus and arsenic.
As commonly met with in coins or jewelry, gold is far from
being pure; coin, for example, usually contains at least 10 per
cent. of copper.
In order to prepare pure gold, a piece of coin may be dissolved in
aqua regia (Exp. 50), the solution evaporated to dryness upon a waterbath, in order to expel the excess of acid, the residue taken up with
water and filtered, to remove any chloride of silver which Imay be
present, and the gold finally precipitated as a brown powder, by adding
to the solution some sulphate of protoxide of iron dissolved in water:2AuCl, + 6FeSO4 = 2Au + Fe2C18 + 2(Fe20,,3S0,).
The powder may then be collected and dried, and, if desirable, melted
and cast into solid masses.
Upon the large scale, fine gold is obtained from its alloys by removing the baser metal by means of either sulphuric or nitric acid.
When an alloy of gold and silver, or copper is boiled with concentrated sulphuric acid in iron kettles, the silver and copper dissolve with
evolution of sulphurous acid, while the gold remains undissolved, and
the iron vessel is not acted upon. In order to recover the silver from
the solution of mixed sulphates, sheets of copper are placed in this
solution, and the silver is precipitated upon them, as has been shown
in Exp. 267.
The solution of sulphate of copper is then evaporated, and the salt
obtained in the crystallized condition fit for sal'e.
The treatment of the alloy of gold and silver with. nitric acid is
based upon the fact that silver is soluble, while gold is insoluble, in
this acid. But it has been found necessary, in order to obtain a complete separation of the two metals, that the proportion of silver to




590                       GILDING.
that of the gold in the alloy should be as much as 2 (or 3 to 1; otherwise portions of the silver would, after a while, become so covered
with gold as to be protected from the action of the acid, and the two
metals could not be completely separated from one another. In practice, whenever the alloy to be treated is found to contain more than a
quarter of its weight of gold, enough silver is added to reduce it to
this proportion; hence the term quartation, by which this method of
parting gold and silver is commonly known.
Finely divided gold obtained by precipitation, as above indicated,
is employed to a considerable extent for gilding porcelain. The surface to be gilt is first painted with an adhesive varnish, then covered
with a mixture of the gold powder and a fusible enamel, and exposed
to intense heat; on being subsequently burnished, the gold takes a
high polish.
There are two series of gold salts, corresponding to the two
oxides-a protoxide AuO, and the teroxide AuO3. These oxides
are rather acids than bases; the teroxide in particular unites
with many metallic oxides to form compounds known as aurates.
The chlorides, bromides, and iodides of gold also readily combine
with other metallic chlorides to form chloraurates, chloraurites,
and the analogous bromine and iodine compounds.
698. Terchloride of Gold (AuC1,) is the compound of gold most
commonly employed in the laboratory. The manner of preparing
it has been already indicated in ~ 697. It serves as a valuable
test for tin.
699. Gilding.-There are several methods of attaching a film
of metallic gold to surfaces of the baser metals.
In the old process of fire gilding, the object to be gilt was first
heated to redness, then washed with dilute acid to cleanse its surface,
and with a solution of nitrate of mercury in order to amalgamate it
slightly; it was then rubbed with a pasty amalgam composed of two
parts of gold and one part of mercury. After a portion of the gold
amalgam had thus been attached to the surface of the article to be
gilt, the latter was heated to drive off the mercury, and the gold left
upon it was polished with a burnishing-tool.
In the more modern method of electro-gilding, the object to be gilt
is attached to the negative pole of a galvanic battery, a bar of gold is
fastened to the positive pole of the battery, and both the object to be
gilt and the bar of gold are placed in a mixed solution of cyanide of
gold and cyanide of potassium. Under the action of the current, the
solution is decomposed; gold is deposited from it upon the object at




ALLOYS OF GOLD.                   591
the negative pole of the battery, while the other ingredients of the
solution go to the positive pole, there to dissolve gold from the bar,
and thus make good to the solution the metal it has lost. Compare
Exp. 353.) Articles of silver, copper, bronze, brass, or platinum, may
thus be gilt directly; but with iron, steel, or tin, it is necessary first
to immerse the article attached to the battery in a solution of cyanide
of copper and of potassium, in order to cover it with a film of copper
to which the gold may adhere.
Even without a battery, gold can be deposited upon silver or copper
by placing either of these metals in a hot solution of the double cyanide
of gold and potassium. Copper trinkets are also sometimes gilt by
boiling them in a liquor prepared by mixing a solution of chloride of
gold with a solution of an alkaline carbonate.
In general, the compounds of gold have but few properties which
are of chemical interest. What has been said of the permanence of
the metal implies, of course, that it is a weak chemical agent, having
but little affinity for other substances.
700. Alloys of Gold.-Gold unites with most of the other
metals; but its most important alloys are those of copper, silver,
and mercury.  Pure gold is so soft that articles of jewelry made
of it would quickly wear out if used; such articles, as well as
coins and watches, are therefore always made of gold which has
been alloyed with copper, in order to increase its hardness. The
standard alloy for coin in this country and in France is nine
parts by weight of gold to one part of copper; in England it is
eleven parts of gold to one of copper. These alloys, as well as
the alloys of silver and gold, are more fusible than pure gold, but
less ductile. Native gold is an alloy of gold and silver, the proportion of the latter metal varying from 0'2 to 62 per cent.
Amalgams of gold play an important part in the metallurgy of
gold (~ 695), and in the process of fire gilding above described.
PLATIN U.
701. Platinum is a metal which, like gold, has little affinity
for the other chemical elements. It is commonly found in the
native state, alloyed with gold and with other metals. Like gold,
it is obtained by washing away the earth and sand with which it
is found mixed. It is a very heavy metal, the specific gravity of
cast platinum being 21'15. Its atomic weight is 197'4. The
color of platinum is intermediate between the white of silver and




592                      PLATINUM.
the gray of steel; its lustre is far less brilliant than that of silver.
It is as soft as copper, very malleable and very tenacious; it may
be drawn into wire so fine that its diameter is only l 1,2,  of a
millimetre. It is not fusible in ordinary furnaces, but may be
fused in the blowpipe-flame (Exps. 26, 203), and is nowadays
melted in considerable quantities in lime crucibles by means of a
blowpipe-flame obtained from common coal-gas and oxygen. At
very high temperatures it may be volatilized.  Like wrought iron,
platinum admits of being forged and welded at temperatures far
below its melting-point. When heated, it expands less than any
other metal, and is hence well adapted for the construction of
apparatus in which metal and glass must be fused together. It
conducts heat and electricity much less readily than gold, silver,
or copper, standing in this respect not far from iron.
702. Platinum does not oxidize in the air at any temperature,
nor is it attacked by any of the common acids taken separately;
in aqua regia (~ 104) it dissolves slowly-much less readily than
gold. Chlorine-water dissolves it, but neither bromine nor iodine
has any action upon it. When heated to redness in the air, in
contact with the fixed caustic alkalies or alkaline earths, it is
slowly corroded, in consequence of the formation of an oxide
which unites with the alkali.  Phosphorus and arsenic unite
readily with hot finely divided platinum, forming very fusible
compounds; sulphur also combines with it, though far less readily.
At high temperatures, platinum is easily acted upon by silicon
(compare ~ 463).  A  platinum  crucible should consequently
never be placed in direct contact with a hot mixture of a carbon
compound and silicic acid. If the crucible is to be heated in a
coal fire, it should first be placed in an earthen crucible lined
with some infusible earth, such as magnesia.
With most of the other metals platinum unites readily, forming alloys which in many instances are more fusible than platinum itself; hence, in using platinum vessels in chemical experiments, care must be taken not to touch the platinum while
hot with easily fusible metals, or to place in the hot vessels any
reducible compound of a metal. Most of the alloys of platinum
are not only fusible, but they are also soluble in acids. Platinum
which has been alloyed with 10 or 12 times its weight of silver,




CHLORIDES OF PLATINUM.              593
for example, is as completely soluble in nitric acid as the silver
itself.
From its comparative inertness as a chemical agent, taken in
connexion with its infusibility, platinum is an extremely useful
metal to the chemist. It is employed in the scientific laboratory
for crucibles, evaporating-dishes, stills, tubes, spatula, forceps,
wire, blowpipe-tips, and the like; and, in the manufacture of oil
of vitriol, large platinum stills, together with cooling-siphons of
the same metal, are employed in the process of concentrating the
acid.
703. A remarkable property of platinum, of inducing various
gases to combine chemically one with the other, has already been
repeatedly alluded to and illustrated (~~ 224, 240, 387). This
power of causing combination is possessed even by clean surfaces
of the ordinary solid metal, though to a much greater degree by
spongy platinum (Exp. 364), and still more by the very finely
divided powder known as platinum black (~ 706).
Platinum forms two series of compounds, corresponding respectively to the protoxide PtO and to the binoxide PtO2. Its chlorides are well-defined compounds; but with the oxygen acids it
forms comparatively few salts, and none of these are at present
of much importance.
704. Protochloride of Platinum (PtCI,) is a compound insoluble
in water, obtained by carefully heating the bichloride to 230~ upon
an oil-bath. It dissolves in alkaline lyes, and the solution thus
obtained may be used for making platinum black (~ 706). At a
red heat chloride of platinum is completely decomposed to metallic
platinum and chlorine. With the other metallic chlorides protoohloride of platinum unites, to form compounds known as chloroplatinites; the general formula of these compounds is 2MCl,PtCl,.
705. Bichloride of Platinum (PtCl4) is the platinum compound
most commonly employed in the laboratory. It is a deliquescent
substance, readily soluble in water, alcohol, and ether; the
aqueous solution is of a reddish-brown color.. When heated to
2300, or thereabouts, the salt loses half its chlorine, as has been
already stated. The aqueous solution of bichloride of platinum
is much used as a test for potassium and! ammonium, and for preparing certain organic compounds suitable for analysis.
2 Q




594                   M,^ATIru,sPON'G..Exp. 361.-Cut half a gramme, or more, of worn-out Ilaftinum foil
or wire into small fragments, and boil them with a teaspoonful of aqua
regia so long as the metal appears to be acted upon, then decant the
liquid into a porcelain dish, add to the fragments of platinum another
teaspoonful of aqua regia, and proceed as before, repeating the treatments until all the metal has dissolved. By the repeated action
of successive small portions of the solvent, platinum and other comparatively speaking insoluble substances can be dissolved much more
readily than if all the liquid necessary for its solution were added at
once. Evaporate the solution to dryness upon a water-bath, take up
the residue with water, and preserve the solution in a bottle provided
with a glass stopper.
Exp. 362.-Pour a teaspoonful of a solution of chloride of potassium,
or of almost any other salt of potassium, into a test-tube, acidulate
the liquid with chlorhydric acid, and add to it a drop of the solution
of bichloride of platinum obtained in the preceding experiment. A
yellow, insoluble powder will soon be precipitated. It is a double
chloride of potassium and platinum, and its formula may be written
2KCl,PtCl4. This test serves to distinguish potassium from sodium,
and, if need be, to separate potassium from solutions in which it is
mixed with sodium; for the double chloride produced with chloride of
sodium and bichloride of platinum is easily soluble in water.
Exp. 363.-Repeat Exp. 362, but substitute chloride of ammonium
for the chloride of potassium. A yellow precipitate, similar to that
obtained in Exp. 362, will separate immediately, or, if the solutions
employed are dilute, after a short time. The composition of this precipitate may be represented by the formula 2NH4C1,PtCl4. Again
repeat the experiment, and this time take enough of the platinum
solution and of the chloride of ammonium to make half a teaspoonful
of the yellow precipitate, taking care that at last there shall be a slight
excess of free chloride of ammonium rather than of chloride of platinum in the supernatant liquid. Allow the precipitate to settle, separate it from the clear liquid by decantation, and dry it partially at a
gentle heat. When the precipitate has acquired the consistence of
slightly moistened earth, transfer it to a cup-shaped piece of platinum
foil, and heat it to redness in the gas-flame, as long as fumes of chloride of ammonium continue to escape. All the chlorine, hydrogen,
and nitrogen will be driven off, and there will remain upon the foil a
gray, loosel3 coherent, sponge-like mass of metallic platinum; it is
called platinum sponge.
Exp. 364.-Hold the dry platinum sponge of Exp. 363 in a stream
of hydrogen or of common illuminating gas issuing from a fine jet.




PLATINUM BLACK.                   595
The metal will soon begin to glow, and in a moment will become hot
enough to inflame the mixture of air and gas in contact with it. Before friction-matches were employed, this property of spongy platinum,
of inflaming hydrogen, was sometimes made use of for striking a light.
The mode of action of the platinum in this experiment is obscure; it
has already been alluded to in ~ 387.
From platinum sponge, solid articles of platinum may be manufactured by compression. If the spongy platinum be first rubbed to powder under water, the particles of metal of which it is composed can be
readily compacted into solid bars by subjecting the ipowder to powerful
pressure in appropriate moulds. The pressed bar is then heated intensely in a coke-fire with strong draught, and forged by striking it
with the hammer upon its ends-the process of heating and forging
being several times repeated, until the bar has become sufficiently condensed. The metal may then be wrought into any desired shape by
heating and hammering, in the same way as any other malleable metal.
This process of working platinum was for a long time the common
method, and is still employed to a certain extent.
706. Platinum Black is a term applied to metallic platinum
even more finely divided than the sponge above described.
By dissolving protochloride of platinum in hot concentrated potash-lye, and pouring into the hot liquor alcohol, by small successive
portions, platinum will be thrown down as a black powder looking
like soot. The powder should be freed from the supernatant liquor by
decantation, and then boiled successively with alcohol, chlorhydric
acid, potash-lye, and water, in order to free it from all impurities.
A capacious vessel must be chosen for the reaction of the alcohol
upon the alkaline solution of chloride of platinum; for much carbonic
acid is generated while the components of the alcohol are reducing
the solution of platinum, so that lively effervescence occurs. Platinum
black is capable not only of absorbing and storing up many times its
own bulk of oxygen gas; it is also capable of giving away this oxygen
to many other substances. If easily oxidizable liquids, such as alcohol
or ether, are dropped upon platinum black which has previously been
exposed to the air, the liquids will be oxidized and converted into new
substances, while the powder becomes red-hot from the heat evolved
during the act of oxidation.
707. Besides forming with the chlorides of potassium and ammonium the insoluble compounds above described, bichloride of
platinum unites with many other chlorides, both of metals and of
organic radicals, to form analogous salts of the general formula
2a2




596               THE PLATINUM METALS.
2MC1,PtC14, or MC2,,PtCl4. These compounds are commonly
called chloroplatinates; by means of them the composition and
combining weights of many organic compounds have been determined. It is only necessary to ignite a weighed portion of the
chloroplatinate, and to weigh the residue of pure platinum which
is left after the organic matter has all been driven off, in order
to ascertain how much platinum is contained in the compound.
This fact having been determined, the quantity of the organic
radical, or rather of the chloride of the radical, which was combined with the chloride of platinum in the chloroplatinate, may
be readily calculated.
708. With gold and platinum are classed several rare metals
which are never found except in association with platinum, and
which closely resemble that metal. They are commonly called
platinum metals, and the group may be appropriately termed the
platinum group. The whole group consists of Rhodium (atomic
weight = 104), Ruthenium (104), Palladium (106'5), Gold (196'7),
Platinum (197'4), Iridium (198), and Osmium (199). Palladium
is used to impart to brass gas-fixtures a peculiar reddish tint,
sometimes called salmon-bronze. Iridium is used for the very
hard tips of gold pens. Osmium forms, among other oxides, a
volatile compound OsO4, whose vapors are intensely poisonous.
The metals of this group are noble metals; they withstand the
action of the atmosphere; none of them are acted upon by nitric
acid, though they dissolve in chlorine and in aqua regia. Their
oxides part with all their oxygen when simply heated, leaving
the metal behind.




SYMBOLS AND ATOMIC WEIGHTS.              597
CHAPTER XXXV.
ATOMrO WEIGHTS OF THE ELEMENTS-CLASSIFICATION.
709. An alphabetical list of the sixty-five recognized elements,
with their symbols and atomic weights, is here given for convenience
of reference. The names of those elements which are so rare as
to be at present of little importance are printed in italics:Aluminum. A1. 274        Mercury          Hg. 200
Antimony,. Sb. 122           Molybdenum. Mo. 96
Arsenic.. As. 75         Nickel.. Ni. 68'8
Barium.      Ba. 137       Nitrogen.      N.     14
Bisnmuth.. Bi. 210        Norium.. No.?
Boron.. Bo. 11          Osmnim.       Os.199
Bromine.       Br. 80        Oxygen. 0. 16
Cadmium.. Cd. 112         Palladium.. Pd. 1065
Ccesium.. Cs. 133        Phosphorus. P. 31
Calcium    o. Ca. 40          Platinum., Pt. 197'4
Carbon.. C. 12          Potassium.   K. 391
Ccerium.. Ce. 92        Rhodium..,104
Chlorine.. Cl. 35'5       Rubidium.      Rb. 85'7
Chromium.. Cr. 52'5         Ruthenium.       Ru. 104
Cobalt.       Co. 58'8      Selenium    o   Se. 79'5
Columbium (Nio-                 Silicon.. Si. 28
biunl).. Nb. 94        Silver,. Ag. 108
Copper.. Cu. 63'4        Sodium,. Na. 23
Didymnium.. D. 95          Strontium.. Sr. 875
Erbilum.   E.112'6       Sulphur.       S. 32
Fluorine.. Fl. 19        Tantalum.. Ta. 182
Glucinum.. Gl. 14         Tellurium.. Te. 128
Gold... Au. 196          Terbium. Tb.?
Hydrogen *. H.  1          Thallium.. TI. 204
indium.. In. 72         Thorium.. Th. 231'5(?)
Iodine.. I. 127        Tin.. Sn.118
Iridium.. Ir. 198       Titanium.. Ti. 50
Iron...     Fe. 56        Tzunsten.. W.184
Lanthanum. La. 92-8       Uranium.. Ur. 120
Lead.. Pb.207         Vanadium.. V. 51'3
Lithium.      Li.  7        Yttrium.       Yt. 61-6
Magnesium. Mg. 24          Zinc... Zn. 65
M1anganese.. Mn. 55          Zirconium..     Zr. 672




5i98                 NATURAL GROUPS.
710. In the following table the elements are arranged in what
are believed to be natural groups. Without accepting any one infallible criterion of classification, or insisting upon any systematic
arrangement of the elements in groups with that strenuousness
which is apt to make classification rather a hindrance than a help,
the student may provisionally use this subdivision of the elements
into groups as a help in remembering facts, as a guide to the
prompt recognition of general properties and general laws, and as
a suggestive compend of his whole chemical knowledge:Fluorine...  19    Glucinum..  14
Chlorine.,.  35*5  Aluminum..  27'5
Bromine....  80    Chromium..  525
Iodine.      127    Manganese..  55
Iron...             56
Oxygen..                16    Cobalt...  588
Sulphur....  32    Nickel...  58'8
Selenium...  79 5  Yttrium...  61'6
Tellurium..   128    Erbium... 112'6
Terbium...?
Nitrogen...  14    Zirconium...  67'2
Phosphorus...  31    Norium...?
Arsenic..            75    Cerium..    92
Antimony...     122    Lanthanum..  92'8
Bismuth.   *.   210    Didymium...  95
Uranium..      120
Carbon....  12    Thorium... 2315 (?)
Boron...  11
Silicon..  28    Copper...  63'4
Mercury... 200
Hydrogen..      1
Lithium.      7    Titanium...  50
Sodium... 23    Columbium..  94
Potassium...  39'1  Tin.... 118
Rubidium...  857  Tantalum.          182
Silver...       108
Coesium.... 133    Molybdenum..  96
Thallium... 204    Vanadium..    51'3
Tungsten... 184
Calcium...  40
Strontium... 87-5  Rhodium... 104
Barium..           137    Ruthenium.. 104
Lead...207    Palladium.. 106'5
Gold.... 196
Magnesium.,.  24    Platinum... 197'4
Zinc...  65    Iridium... 198
Cadmium,,. 112    Osmium.,. 199




ATOMIC HEATS OF THE ELEMENTS.            599
711. Atomic Heats of the Elements.-The power of heat to
cause changes of temperature is not the same for any two substances, but varies with the nature of the substance submitted
to its action. Each chemical element is peculiarly affected in
this respect by heat. The quantity of heat needed to raise the
temperature of a certain weight of water from 0~ to 10 being called
unity, the quantity of heat required to raise the temperature of
the same weight of any element by the same amount is the specific
heat of that element (~ 31).  In the second column of the following table will be found the specific heats of a number of representative elements, selected from each group of elements except the
carbon group, and arranged in the order of their atomic weights.
(Compare ~ 710.)
Name.  Slpecfic Heat. Atomic Weight. Atomic Heat.
Lithium..  0'94080           7            6'59
Aluminum'21430        27'5           5'89
Sulphur..' 20259          32             6S48
Iron......'  11380          56            6-37
Copper....'09515         63'4           6'04
Arsenic..' 08140          75            6'11
Silver....    05701        108            6'16
Cadmium'05669        112            6-35
Tin..,...'05623         118            6'63
Iodine...'05412        127            6-87
Tungsten..'03342         184            6'15
Gold......   03244         196            6'36
Lead....'03140        207             6'50
Bismuth.    03084          210             6'48
The preceding table contains only solid substances. It has been
found that the specific heat of the same body is commonly greater
in the liquid than in the solid state, and always less in the gaseous
than in the liquid state. Accordingly in instituting any comparison between different bodies, based on their specific heats, it is
essential to compare them in the same physical condition, solids
with solids, liquids with liquids, gases with gases. The second
column of the following table contains the specific heats of ihe
four elements which are gaseous at the ordinary atmospheric
temperature and pressure.




600            ATOMIC HEATS OF THE ELEMENTS.
cname.  Specfc Heat. Atomic Weight. Atomic Iieat.
Hydrogen..  3'4090          1           3'4090
Nitrogen..  0'2438        14           3'4132
Oxygen....' 2175          16           3'4800
Chlorine..' 1210         35'5         4'2955
On comparing together the numbers in the second and third
columns of the preceding tables, it will be noticed that the lower
the atomic weight of an element is, the higher is its specific heat,
and vice versa. In the fourth column of the tables, under the
name of atomic heat, will be found the product of the specific heat
by the atomic weight of each element.  The atomic heats of the
elements represent the quantities of heat which are required to
cause equal alterations of temperature in atomic proportions of the
several elements. But the tables show that these quantities of
heat are nearly the same for each and all of the solid elements
compared in the first table, and are again approximately equal
for the gaseous elements which are grouped in the second table.
This striking principle is deducible from the foregoing considerations-namely, that while the capacities for heat of the same
weights of the various elements are very different, the capacities
for heat of the atoms, or atomic proportions, of the elements are
nearly identical, provided that the elements compared be in the
same physical condition. In other words, those weights of the
elements which are assumed to represent the relative weights of
their atoms require approximately the same amount of heat to raise
them through an equal number of degrees of temperature; while
the amounts of heat required to raise equal weights of the elements
through an equal number of degrees are expressed by very different numbers (the specific heats). It is essential, however, that
the elements compared should be in the same physical condition.
Tt is true that the numbers representing the atomic heats show
considerable discrepancies; but when it is remembered that there
are unavoidable errors attaching to the determinations both of the
specific heats and of the atomic weights, that many of the elements cannot yet be obtained in a condition of purity, and that
the two factors of the product (specific heat and atomic weight)
vary in the proportion of 1 to 30, it will be seen that the accordance is distinct enough to indicate the existence of a general




ATOmC HEATS OF COMPrOUNDS.               601
law. The atomic heats of carbon, boron, and silicon, however, do
not conform to this law.
712. Atomic Heats of Compounds. —The same kind of relation
between the specific heat and the molecular weight obtains within
certain classes of compound bodies of like constitution. Only
those compounds which have an analogous atomic constitution can
possess approximately the same molecular heat; and there are
many admitted exceptions even to this rule. The following table
contains the specific and atomic heats of a number of inorganic
compounds, arranged in groups according to their chemical composition.
Specific    Molecular  Atomic
Foula.  eat.  Weight.    Heat.
1. OXIDES MO.
PbO.......... 0'05089         223        11'35
HgO..........   05179        216        11'19
CuO........   14201          79'5      11'19
ZnO............'12480        81        10'11
2.       M203.
Fe203..........    16695     160        26'71
Cr203.......,'17960      153        27'47
Bi,203.........'06053        468        28-33
Sb2O..........'09009      292        26'31.
8. CHLORIDES MC1.
NaC1.......... 0'21401         58'5      12-52
KC1..........'17295          74'5      12'88
IgC1..........    05205       235.5      12.26
CuCl..........'13827         99        13.69
4            MC12.
MgC12...........19460          95        18.49
ZnCl.........     13618       136        18-52
PbC]2.........'06641      278        18'46
MnC12..........    14255     126        17'96
SnC12..........'10161      189        19'20
5.           MC14.
SnC14.........'14769       260        38837
TiC14.'1...,.  19145          192        36'76




602                VALUE OF SPECIFIC HEATS.
Formula.          Specific    Molecular  Atomic
Heat.     Weight,    Beat.
6. SULPHIDES MS.
FeS............ 013570       88        11 94
NiS..........   12813       90'7      11'62
CoS............   12512       90'7      11'36
SnS...........'08375        150        12'56
7.          MS2.
FeS2..........'13009        120        15'61
SnS2..........'11932        182        21-72
MoS2..........'12334        160        19'73
8.          MA2S.
Sb2S3.........'08403      340       28'57
BiS..........   06002        516        30'97
9. CARBONATES MX2COa.
K2CO3.......... 0'21623      138        29'84
Na2CO........'27275       106        28'91
10.            MCO,.
BaC03........'11038         197        21'74
SrCO3..........'14483       147'6     21'38
FeCO3........'19345        116        22'44
11. SULPHATES AI2.O04.
K2S04.......... 019010       174       83'08
Na2SO4......'23115         142        32'82
12.           MSO4.
CaSO4..........  *19656      136       26-73
MgS04........'22159         120        26'59
PbSO4..........'08723        303        26'43
There are considerable departures from equality in the atomic
heats of the bodies of similar constitution grouped in the above
table; but when it is observed that the two factors of the product
are very wide apart, the close coincidence in most cases will seem
much more noteworthy than the occasional discrepancies.
713. Value of Specific Ieats.-The determination of the specific heat of a substance is a difficult experiment in Physics. It
is obvious that, if the laws concerning atomic heat which have
been above illustrated were invariable and certain, the physicist




TWO SETS OF ATOMIC WEIGHTS IN USE.       603
would determine for the chemist the atomic weights of the elements in experimentally fixing their specific heats; he would
also indicate the true classification of chemical compounds by
determining the specific heats of these compounds. To obtain
the atomic weight of any element, it would suffice to divide the
atomic heat common to all the elements in the same physical state
by the specific heat of that element. To determine the class of compounds to which a given compound belonged, it would be sufficient to know its atomic heat. A given carbonate, for example,
would belong to the class M2CO3 or to the class MCO3, according
as its atomic heat were about 29 or about 22.
Although the principle of the equality of the atomic heats of
the elements and of compounds of like constitution is not established with that certainty which would give to the experimental
determination of the specific heat the conclusive weight just indicated, yet this equality certainly affords a very strong presumption in favor of the correctness of the atomic weights of the
elements contained in the above tables, and of the molecular
formulke on which the classification of compounds in the third
table is based.
714. Two Sets of Atomic Weights in use.-The student will
find in many books on chemistry weights, other than those used
in this manual, assigned to the majority of the elements, and
variously called combining, equivalent, or atomic weights. To
all the elements, except those of the chlorine, nitrogen, and
alkali groups, and the elements boron and gold, are assigned
weights which are respectively the halves of the numbers given
in ~~ 709, 710. The weight assigned to sulphur, for example,
is 16, to lead 103'5, to cadmium 56, to iron 28, to copper 31'7,
to tin 59, to oxygen 8, and so forth. If these smaller numbers
were accepted as the true atomic weights of the elements to
which they are respectively assigned, it is obvious that the relation of equality between the atomic heats of all elements in the
same physical state, presented in the tables of ~ 711, would disappear; furthermore, many of the relations indicated in the table
of ~ 712 would be no longer visible. The atomic heats of the
solid elements would be divisible into two classes, in one of
which the atomic heat would be about double what it was in the




604            DISCUSSION OF ATOMIC WEIGHTS.
other. If the atomic weight of lead were 103'5, the formula of
the white chloride of lead (~ 581) would be PbC1, analogous to
NaCl, and there would be no clear reason why its atomic heat
should not be that of the chlorides of the formula MC1, namely,
about 12'5; but its atomic heat would be only 9'23. If the
atomic weight of iron were 28, that of potassium being 39'1,
there would be no assignable reason for the marked difference
between the atomic heats of the two carbonates; the formulae of
the carbonates of iron and of potassium would be alike-either
both M2C0,CO or both MCO3.
Another consequence of using the smaller atomic weights must
not pass unnoticed. If the atomic weight of oxygen is 8 and of
sulphur 16, the coincidence of the atomic weight and the unitvolume weight of those eight elements for which this equality
has been affirmed (~ 259) ceases to be true; and the simple rule
that the molecule of every compound gas or vapor occupies a
volume twice as large as the combining volume of hydrogen,
oxygen, chlorine, and so forth, will lose the universality which
constitutes its chief value.
The student will remember that the determination of the least
combining proportion by weight of any element which cannot be
converted into vapor is not a matter of direct experiment simply
(~~ 395, 603). The natural analogies of the element and its
compounds, the greater or less simplicity of the formulae which
result from one assumption or another, the indications of isomorphism, and of atomic heat, have all to be consulted. That the
best guides have thus far failed to lead to an unquestionable
determination of the real least combining weights (or atomic
weights) of the majority of the elements, may be inferred from
the diversities of usage on this subject in chemical literature.
In order to indicate that they mean the atomic weights which
have been given in this manual, many chemists write the symbols of those elements for which two different weights are in use
with a line drawn through them, thus 0, S, Fe, Pb. This practice is almost essential in periodical publications to which writers
of different theoretical views contribute.
715. In the midst of the doubts and discussions which to-day
envelope chemical theories, the student will do well to remember




FACT AND THEORY.                  605
that all these questions lie without the sphere of fact. They do
not affect the actual composition or properties of a single element
or compound; they are questions of interpretation, classification,
and definition. The existence of atoms is itself an hypothesis,
and not a probable one; all speculations based on this hypothesis,
all names which have grown up with it, all ideas which would
be dead without it, should be accepted by the student provisionally and cautiously, as being matter for belief but not for knowledge. All dogmatic assertion upon such points is to be regarded
with distrust. The great majority of chemists, devoted to the
applications of chemistry in mineralogy, metallurgy, dyeing,
printing, and the manufacture of chemicals, remain completely
indifferent to discussions of chemical theories. Hence the student
will find that in technical chemical literature the older notation
and the corresponding smaller atomic weights are almost invariably employed.
Theories, however, are of great importance to the progress of
the science and to the clear ordering of the ground already won.
It is, on this account, very much to be wished that the great
attention now devoted to the discussion of the best methods of
representing symbolically the constitution of chemical substances
and the changes to which they are subject, may result in the
elaboration of a system good enough to command general acceptance.








APPENDIX.
COHEMICAL MANIPULATION.
1. Glass tubing. —Two qualities of glass tubing are used in chemical
experiments-that which softens readily in the flame of a gas or spiritlamp, and that which fuses with extreme difficulty in the flame of the
blast-lamp. These two quilities are distinguished by the terms soft
and hard glass. Soft glass is to be preferred for all uses except the
intense heating, or ignition, of dry substances. Fig. I. represents the
most convenient sizes of glass tubing, both hard and soft, and shows
also the proper thickness of the glass walls for each size.
Fig. I.
8  7   6    5    4       3        2           1
2. Cutting and cracking glass.-Glaws tubing and glass rod must
generally be cut to the length required for any particular apparatus.
A sharp triangular file is used for this purpose. The stick of tubing,
or rAl, to be cut is laid upon a table, and a deep scratch is made with
the file at the place where the fracture is to be made. The stick is
then grasped with the two hands, one on each side of the mark, while
the thumbs are brought together just at the scratch. By pushing
with the thumbs and pulling in the opposite direction with the fingers,
the stick is broken squarely at the scratch, just as a stick of candy or
dry twig may be broken. The sharp edges of the fracture should invariably be made smooth, either with a -wet file, or by softening the
end of the tube or rod in the lamp. (See Appendix, ~ 3.) Tubes cr
rods of sizes four to eight inclusive may readily be cut in this manner;
the larger sizes are divided with more difficulty, and it is often neces



11               CUTTING AND CRACKING GLASS.
sary to make the file-mark both long and deep. An even fracture is
not always to be obtained with large tubes. The lower ends of glass
funnels, and those ends of gas delivery-tubes which
enter the bottle or flask in which the gas is generated,    Fig. I
should be filed off, or ground off on a grindstone, obliquely (Fig. II.), to facilitate the dropping of liquids
from such extremities.
In order to cut glass plates, the glazier's diamond
must be resorted to. For the cutting of extremely thin
glass tubes and of other glass ware, like flasks, retorts, and bottles, still
other means are resorted to, based upon the sudden and unequal application of heat. The process divides itself into two parts —the producing of a crack in the required place, and the subsequent guiding of
this crack in the desired direction. To produce a crack, a scratch
must be made with the file, and to this scratch a pointed bit of redhot charcoal, or the jet of flame produced by the mouth blowpipe, or
a very fine gas-flame, or a red-hot glass rod may be applied. If the
heat does not produce a crack, a wet stick or file may be touched upon
the hot spot. Upon any part of a glass surface except the ed:ge, it is
not possible to control perfectly the direction and extent of this first
crack; at an edge a small crack may be started with tolerable certainty
by carrying the file-mark entirely over the edge. To guide the crack
thus started, a pointed bit of charcoal or slow match may be used.
The hot point must be kept on the glass from 1 c.m. to 0'5 c.m. in
advance of the point of the crack. The crack will follow the hot
point, and may therefore be carried in any desired direction. By
turning and blowing upon the coal or slow match the point may be
kept sufficiently hot. Whenever the place of experiment is supplied
with common illuminating gas, a very small jet of burning gas may
be advantageously substituted for the hot coal or slow match. To
obtain such a sharp jet, a piece of hard glass tube, No. 5, 10 c.m. long,
and drawn to a very fine point (see Appendix, ~ 3), should be placed
in the caoutchouc tube which ordinarily delivers the gas to the gas,
lamp, and the gas should be lighted at the fine extremity. The burning jet should have a fine point, and should not exceed 1'5 c.m. in
length.  By a judicious use of these simple tools, broken tubes,
beakers, flasks, retorts, and bottles may often be made to yield very
useful articles of apparatus. No sharp edges should be allowed to
remain upon glass apparatus. The durability of the apparatus itself,
and of the corks and caoutchouc stoppers and tubing used with it, will
be much greater, if all sharp edges are removed with the file or, still
better, rounded in the lamp.




HEATING GLASS TUBES.                 tU
3. Bending and closing glass tubes. —Tubing of sizes five to eight
inclusive can generally be worked in the common gas- or spirit-lamp;
for larger tubes the blast-lamp is necessary (see Appendix, ~ 6).
Glass tubing must not be introduced suddenly into the hottest part of
the flame, lest it crack. Neither should a hot tube be taken from the
flame and laid at once upon a cold surface. Gradual heating and
gradual cooling are alike necessary, and are the more essential the
thicker the glass; very thin glass will sometimes bear the most sudden
changes of temperature, but thick glass and glass of uneven thickness
absolutely require slow heating and annealing. When the end of a
tube is to be heated, as in rounding sharp edges, more care is required,
in consequence of the great facility with which cracks start at an edge.
A tube should therefore always be brought first into the current of
hot air beyond the actual flame of the gas or spirit-lamp, and there
thoroughly warmed before it is introduced into the flame itself. If
a blast-lamp is employed, the tube may be warmed in the smoky
flame, before the blast is turned on, and may subsequently be annealed
in the same manner; the deposited soot will be burnt off in the first
instance, and in the last may be wiped off when the tube is cold.
In heating a tube, whether for bending, drawing, or closing, the tube
must be constantly turned between the fingers, and also moved a little
to the right and left, in order that it may be uniformly heated all
round, and that the temperature of the neighboring parts may be duly
raised. If a tube, or rod, is to be heated at any part but an end, it
should be held between the thumb and first two fingers of each hand
in such a manner that the hands shall be below the tube or rod, with
the palms upward, while the lamp-flame is between the hands. When
the end of a tube or rod is to be heated, it is best to begin by warming
the tube or rod about 2 c.m. friom the end, and thence to proceed slowly
to the end.
The best glass will not be blackened or discolored during heating.
The blackening occurs in glass which, like ordinary fint glass, contains oxide of lead as an ingredient. Glass containing much of this
oxide is not well adapted for chemical uses. The blackening may
sometimes be removed by putting the glass in the upper or outer part
of the flame, where the reducing gases are consumed, and the air has
the best access to the glass. The blackening may be altogether avoided
by always keeping the glass in the oxidizing part of the flamer
Glass begins to soften and bend below a visible red heat. The condition of the glass is judged of as much by the fingers as the eye; the
hands feel the yielding of the glass, either to bending, pushing, or
pulling, better than the eye can see the change of color or form. It




iv        ]ENTDING, DRAWING, AND CLOSING GLASS TUBE3.
may be bent as soon as it vields in the hands, but can be drawn out
only when much hotter than this. Glass tubing, however, should not
be bent at too low a temperature; the curves made at too low a heat
are apt to be flattened, of unequal thickness on the convex and concave
sides, and brittle.
In bending tubing to make gas-delivery-tubes and the like, attention
should be paid to the following points: 1st, the glass should be equally
hot on all sides; 2nd, it should not be twisted, pulled out, or pushed
together during the heating; 3rd, the bore of the tube at the bend
should be kept round, and not altered in size; 4th, if two or more
bends be made in the same piece of tubing (Fig. III., a), they should
all be in the same plane, so that the finished tube will lie flat upon the
level table.
When a tube or rod is to be bent or drawn
close to its extremity, a temporaryv handle may
be attached to it by softening the end of the 
tube or rod, and pressing against the soft glass A|
a fragment of glass tube, which will adhere
strongly to the softened end. The handle may
subsequently be removed by a slight blow, or by the aid of a file. If a
considerable bend is to be made, so that the angle between the arms
will be very small or nothing, as in a siphon, for example, the curvature cannot be well produced at one place in the tube, but should be
made by heating, progressively, several centimetres of the tube, and
bending continuously from one end of the heated portion to the other
(Fig. III., b). Small and thick tube may be bent more sharply than
large or thin tube.
In order to draw a glass tube down to a finer bore, it is simply
necessary to thoroughly soften on all sides one or two centimetres
length of the tube, and then, taking the glass from the flane, pull the
parts asunder by a cautious movement of the hands. The larger the
heated portion of glass, the longer will be the tube thus formed. Its
length and fineness also increase with the rapidity of motion of the
hands. If it is desirable that the finer tube should have thicker walls
in proportion to its bore than the original tube, it is only necessary to
keep the heated portion soft for two or three minutes before drawing
out the tube, pressing the parts slightly together the while. By this
process the glass will be thickened at the hot ring.
To obtain a tube closed at one end, it is best to take a piece of
tubing, open at both ends, and long enough to make two closed tubes.
In the middle of the tube a ring of glass, as narrow as possible, must be
made thoroughly soft. The hands are then separated a little, to cause




BLOWING BULBS.                      V
a contraction in diameter at the hot and soft part. The point of the
flame must now be directed, not upon the narrowest part of the tube,
but upon what is to be the bottom of the closed tube. This point is
indicated by the line a in Fig. IV. By with-    F   IV
drawing the right hand, the narrow part of
the tube is attenuated, and finally melted off, -
leaving both halves of the original tube closed   -
at one end, but not of the same form; the
right-hand half is drawn out into a long point,
the other is more roundly closed. It is not possible to close handsomely the two pieces at once. The tube is s ldom perfectly finished
by the operation; a superfluous knob of glass generally remains upon
the end. If small, it may be got rid of by heating the whole end of
the tube, and blowing moderately with the mouth into the open end.
The knob, being hotter and therefore softer than any other part, yields
to the pressure from within, spreads out and disappears. If the knob
is large, it may be cut off with scissors while red-hot, or drawn off by
sticking to it a fragment of tube, and then softening the glass above the
junction. The same process may be applied to the too pointed end of
the right-hand half of the original tube, or to any misshapen result of
an unsuccessful attempt to close a tube, or to any bit of tube which is
too short to make two closed tubes. When the closed end of a tube
is too thin, it may be strengthened by keeping the whole end at a red
heat for two or three minutes, turning the tube constantly between
the fingers. It may be said in general of all the preceding operations
before the lamp, that success depends on keeping the tube to be heated
in constant rotation, in -order to secure a uniform temperature on all
sides of the tube.
4. Blowing bulbs and piercing holes in tubing.-If the bulb desired
is large in proportion to the size of the tube on which it is to be made,
the walls of the tube must be thickened by rotation in the flame before
the bulb can be blown. If the bulb is to be blown in the middle of a
piece of tubing, this thickening is effected by gently pressing the ends
of the tube together while the glass is red-hot in the place where the
bulb is to be; if the bulb is to be placed at the end of a tube, this end
is first closed, and then suitably thickened by pressing the hot glass up
with a piece of metal until enough has been accumulated at the end.
The thickened portion of glass is then to be heated to a cherry-red,
suddenly withdrawn from the flame, and expanded while hot by
steadily blowing, or rather pressing air, into the tube with the mouth;
the tube must be constantly turned on its axis, not only while in the
flame, but also while the bulb is being blown. If too strong or too
2 Pt 2




Vi                          LAMPS.
sudden a pressure be exerted with the mouth, the bulb will be extremely thin and quite useless. By watching the expanding glass, the
proper moment for arresting the pressure may usually be determined.
If the bulb obtained be not large enough, it may be reheated and enlarged by blowing into it again, provided that a sufficient thickness of
glass remain.
It is sometimes necessary to make a hole in the side of a tube or
other thin glass apparatus. This may be done by directing a pointed
flame from the blast-lamp upon the place where the hole is to be,
until a small spot is red-hot, and then blowing forcibly into one end
of the tube while the other ena is closed by the finger; at the hot spot
the glass is blown out into a thin bubble, which bursts or may be easily
broken off, leaving an aperture in the side of the tube.
It is hoped that these few directions will enable the attentive student
to perform, sufficiently well, all the manipulations with  Fi  V.
glass tubes which ordinary chemical experiments require.
Much practice will alone give a perfect mastery of the
details of glass-blowing.
5. Lamps.-The common glass spirit-lamp will be understood without description from the figure (Fig. V.). This  -
lamp does not give heat enough for most ignitions; for
such purposes a lamp with circular wick, of some one of
the numerous forms sold under the name of Berzelius's
Argand Spirit Lamp (Fig. VI.), is necessary. These argand lamps
are usually mounted on a lamp-             Fig. VI
stand provided with three brass
rings; but the fittings of these
lamps are all made slender, in
order not to carry off too much
heat. When it is necessary to
heat heavy vessels, other supports
must be used.
Whenever gas can be obtained,
gas-lalmps are greatly to be preferred to the best spirit-lamps.
For all common chemical experiments, except a few for which ignition-tubes must be prepared or
in which considerable lengths of
tubing must be heated, the gaslamp known as Bunsen's burner is the best lamp. The cheapest and
best construction of this lamp may be learned from the following




LAMPS.                         vii
description with the accompanying figures. (Fig. VII.) The single
casting of brass a b comprises the tube b through which the gas enters,
and the block a from which the gas          Fig. VII.
escapes by two or three fine vertical
holes passing through the screw d, and
issuing from the upper face of d, as
shown at e. The length of the tube
b is 4'5 c.m. and its outside diameter
varies from 0'5 c.m. at the outer end
to 1 c.m. at the junction with the           e
block a. The outside diameter of the
block a is 1l6 c.m., and its outside   c
height without the screws is 1'8 c.m.
By the screw c, the piece a b is at-  7-7tached to the iron foot g, which may be 6 c.m. in diameter. By the
screw d, the brass tubef is attached to the casting a b. The diameter
of the face e, and therefore the internal diameter of the tube f, should
be 8 m.m. The length of the tubef is 9 c.m. Through the wall of
this tube, four holes 5 m.m. in diameter are to be cut at such a height
that the bottom of each hole will come 1 m.m. above the face e when
the tube is screwed upon a b. These holes are of course opposite each
other in pairs. The finished lamp is also shown in Fig. VII. To
the tube b a caoutchouc tube of 6 to 7 m.m. internal diameter is attached; this flexible tube should be about 1 m. long, and its other
extremity should be connected with the gas-cock through the intervention of a short piece of brass gas-pipe screwed into the cock.
In cases where a very small flame is required, as, for example, in evaporating small quantities of liquid, a piece of wire gauze somewhat
larger than the opening of the tube f should be laid across the top of
the tube, and its projecting edges pressed down tightly against the sides
of the tube before the lamp is lighted. In default of this precaution
the flame of a Bunsen's burner, when small and exposed to currents &f
air, is liable to pass down the tube and ignite the    Fg. VII.
gas at d.
A lamp to give a powerful flame 8 or 10 c.m. long,  ~' ~
suitable for heating tubes, may be very simply constructed by boring two holes, entering the side and  0       a,
issuing at the upper face, through a block of compact
hard wood, 10 c.m. by 6'5 c.m. by 5'5 c.m., and fitting short pieces of brass tubing into the holes so
formed. To                                  the tubes at the the
caoutchouc tubes which deliver the gas, and from the tubes at the top




Viii              BLAST-LAMPS AND BLOWERS.
the gas issues under a sheet-iron funnel closed at the top with wire
gauze. Above this gauze, the mixture of gas and air is to be lighted.
The iron funnel will be readily understood from the accompanying figure,
and the following dimensions: —Length of the wire gauze 10 c.m.;
width of the gauze 5 c.m.; width at a b 9 c.m.; height of the line a b
from the table 8'5 c.m.; whole height of the funnel 21 c.m. A partition parallel to a b divides the funnel into two equal parts from the
gauze to the level of a b.
6. Blast-lamps and Blowers.-For drawing, bending, and closing large
glass tubes, a blast-lamp is necessary. The best form is that sold under
the name of Bunsen's Gas Blowpipe. Its construction and the method
of using it may be learned from Fig. IX.; a b is the pipe through
Fig. X.
Fig. IX
Fig. XI.
which the gas enters, e is the tube for the blast of air; the relation of
the air-tube to the external gas-tube is shown at d; there is an outer
sliding tube by which the form and volume of the flame can be
regulated.
If gas is not to be had, a lamp burning oil or naphtha must be employed. Fig. X. represents a common tin glassblower's lamp, suitable
for burning oil. A large wick is essential, whether oil or naphtha be
the combustible.
For every blast-lamp a blowing-machine of some sort is necessary.
To supply a constant blast it is essential that the bellows be of that
onstruction called double. Figs. XI. and XII. represent two forms




BLOWERS.                           ix
of blowpipe-table; their heig'ht is that of an ordinary table, from
which dimension the other proportions may be inferred. A small
double-acting bellows is now made for the use of dentists, which is
available at any table by the help of a caoutchouc tube to conduct the
blast to the jet. These bellows are too small to give a steady flame
of large size, but will nevertheless answer for most of the glassblowing necessary in the execution of the experiments described in
this manual.
Where an abundant supply of water is at command, the following
blowing apparatus is very convenient. A tin pipe, a b (Fig. XIII.),
ig'. XIII.
Fig. XII.                                    U
about one metre long and about 13 m.m. in diameter, has two smaller
pipes, 12 to 16 c.m. long, soldered into it; these snmall pipes are 8 il.m.
in diameter; one, c d, is inserted at right angles 12 c.m. from the end,
the other, ef, 2'5 c.m. lower, at an angle of 45~. The lower end of
the tube passes through the cork of a wide-mouthed five-litre bottle,
extending rather more than halfway down. Two glass tubes also
pass through the cork of the bottle-a short small tube g, No. 4, which
should reach some 16 c.m. above the cork, but should not project into




X                        CAOUTCHOUC.
the bottle, and a larger tube A, No. 2, extending to the bottom of the
bottle. The outer end of the tube h bends over and is connected by
caoutchouc tubing with a straight tube of equal diameter. This last
arrangement forms the siphon. To the tube g a caoutchouc tube, g b,
is attached to convey the blast to any desired point. To produce a
blast, the water-cock is connected with the tube c d by means of a
caoutchouc tube. When the water is turned on, the caoutchouc tube
g i is closed for a moment with the thumb and finger. This starts the
water through the siphon, and immediately a continuous and powerful
blast of air rushes out through the tube g i, and may be carried directly
to the blowpipe. The siphon must be capable of carrying off a larger
stream of water than that which is allowed to enter, so that there shall
never be more than 3 or 4 c.m. of water in the bottle. By regulating
the water-cock, the proper supply of water may be determined.
The same apparatus may be used as an aspirator. When the instrument is to be used to draw air through any apparatus, the tube g i is
closed by inserting a glass rod; the upper end of the tube a b is closed
with a cork, and the tube ef is connected with the apparatus through
which the current of air is to be drawn. The force of the current
of air is to a certain degree affected by the size of the tube a b; to
diminish the effective calibre of this tube, in case a gentle current of
air is required, a glass rod as long as the tube may be passed down
the tube through a cork inserted at a. The same apparatus may thus
be made to produce a gentle or a powerful current of air.
7. CaoutchoutC.-Vulcanized caoutchouc is a most useful substance
in the laboratory, on account of its elasticity and because it resists so
well most of the corrosive substances with which the chemist deals.
It is used in three forms:-(1) in tubing of various diameters comparable with the sizes of glass tubing; (2) in stoppers of various sizes
to replace corks; (3) in sheets. Caoutchouc tubing may be used to
conduct all gases and liquids which do not corrode its substance, provided that the pressure under which the gas or liquid flows be not
greater or its temperature higher than the texture of the tubing can
endure. The flexibility of the tubing is a very obvious advantage in a
great variety of cases. Short pieces of such tubing, a few centimetres
in length, are much used, under the name of connectors, to make
flexible joints in apparatus of which glass tubing forms part; flexible
joints add greatly to the durability of such apparatus, because long
glass tubes bent at several angles and connected with heavy objects,
like globes, bottles, or flasks full of liquid, are almost certain to break
even with the most careful usage; gas-delivery-tubes, and all conBiderable lengths of glass tubing, should invariably be divided at one or




CORKS.                          Xi
more places, and the pieces joined again with caoutchouc connectors.
The ends of glass tubing to be thus connected should be squarely cut,
and then rounded in the lamp, in order that no sharp edges may cut
the caoutchouc; the internal diameter of the caoutchouc tube must be a
little smaller than the external diameter of the glass tubes; the slipping
on of the connector is facilitated by wetting the glass. In some cases
of delicate quantitative manipulations, in which the tightest possible
joints must be secured, the caoutchouc connector is bound on to the
glass tube with a silk or smooth linen string; the string is passed as
tightly as possible twice round the connector and tied with a square
knot; the string should be moistened, in order to prevent it from
slipping while the knot is tying.
Caoutchouc stoppers of good quality are much more durable than
corks, and are in every respect to be preferred. The German stoppers
are of excellent shape and quality; the American, being chiefly intended for wine-bottles, are apt to be too conical. Caoutchouc stoppers
can be bored, like corks (see the next section), by means of suitable
cutters, and glass tubes can be fitted into the holes thus made with a
tightness unattainable with corks. German stoppers may be bought
already provided with one, two, and three holes. It is not well to lay
in a large stock of caoutchouc stoppers; for, though they last a long
time when in constant use, they not infrequently deteriorate when
kept in store, becoming hard and somewhat brittle with age. These
stoppers must not be confounded with the very inferior caps which
were in use a few years ago.
Pieces of thin sheet caoutchouc are very conveniently used for
making tight joints between large tubes of two different sizes, or between the neck of a flask, or bottle, and a large tube which enters it,
or between the neck of a retort and the receiver into which it enters.
A sufficiently broad and long piece of sheet caoutchouc is considerably
stretched, wrapped tightly over the glass parts adjoining the aperture
to be closed, and secured in place by a string wound closely about it
and tied with a square knot.
8. Corks.-It is often very difficult to obtain sound, elastic corks of
fine grain and of size suitable for large flasks and wide-mouthed bottles.
On this account, bottles with mouths not too large to be closed with a
cork cut across the grain should be chosen for chemical uses, in preference to bottles which require large corks or bungs cut with the
grain, and therefore offering continuous channels for the passage of
gases, or even liquids. The kinds sold as champagne-corks and as
satin corks for phials are suitable for chemical use. The best corks
generally need to be softened before using; this softening may be




XU1             PUTTTNG TUBES THROUGH CORKS.
effected by rolling the cork under a board upon the table, or under the
foot upon the clean floor, or by gently squeezing it on all sides with
the well-known too] expressly adapted for this purpose, and thence
called a cork-squeezer. Steaming also softens the hardest corks.
Corks must often be cut with cleanness and precision; a sharp, thin
knife, such as shoemakers use, is desirable for this purpose. When a
cork has been pared down to reduce its diameter, a fiat file may be
employed in finishing; the file must be fine enough to leave a smooth
surface upon the cork; in filing a cork, a cylindrical, not a conical
form should be aimed at.
In boring holes through corks to receive glass tubes, a hollow
cylinder of sheet brass sharpened at one end is a very convenient tool.
Fig. XIV. represents a set of such little cylinders
of graduated sizes, slipping one within the other  Fig. XIV.
into a very compact form; a stout wire, of the
same length as the cylinders, accompanies the        []
set, and serves a double purpose. Passed transversely through two holes in the cap which ter-       I
minates each cylinder, it gives the hand a better   c
grasp of the tool while penetrating the cork; and
when the hole is made, the wire thrust through
an opening in the top of the cap expels the little
cylinder of cork which else would remain in the
cutting cylinder of brass.  That cutter whose
diameter is next below that of the glass tube to
be inserted in the cork is always to be selected;
and if the hole it makes is too small, a round file
must be used to enlarge the aperture; the round
file, also, often comes into play to smooth the rough sides of a hole made
by a dull cork-borer. Cutters which have been dulled by use may be
sharpened by filing or grinding down their outer bevelled edges and
then paring off any protuberance or roughness which may remain
upon the inside of the edge with a sharp penknife. A pair of small
callipers is a very convenient, though by no means essential tool in
determining which size of cutter to employ. A flask which presents
sharp or rough edges at the mouth can seldom be tightly corked, for
the cork cannot be introduced into the neck without being cut or
roughened; such sharp edges must be rounded in the lamp.  In
thrusting glass tubes through bored corks, the following directions are
to be observed: —(1) The end of the tube must not present a sharp
edge capable of cutting the cork. (2) The tube should be grasped
very close to the cork, in.order to escape cutting the hand which holds




SUPPORTS FOR VESSELS.                  xiii
the cork, should the tube break; by observing this precaution the
chief cause of breakage, viz. irregular lateral pressure, will be at the
same time avoided. (3) A funnel-tube must never be held by the
funnel in driving it through a cork, nor a bent tube grasped at the
bend, unless the bend comes immediately above the cork.  (4) If the
tube goes very hard through the cork, the application of a little soap
and water will facilitate its passage; but if soap is used the tube can
seldom be withdrawn from the cork after the latter has become dry.
(5) The tube must not be pushed straight into the cork, but screwed
in, as it were, with a slow rotary as well as onward motion. Joints
made with corks should always be tested before the apparatus is used,
by blowing into the apparatus and at the same time stopping up all
legitimate outlets. Any leakage is revealed by the disappearance of
the pressure created. To the same end, air may be sucked out of an
apparatus and its tightness proved by the permanence of the partial
vacuum. To attempt to use a leaky cork is generally to waste time
and labor and to insure the failure of the experiment. When, however, a leak is only discovered during the actual progress of the experiment, it is sometimes possible to save the experiment by using a
lute; for this purpose wax applied with a warm knife, or a paste made
of rye-meal and water may be used; common sealing-wax also is
sometimes a useful makeshift.
9. Irron Stand, Sand-bath, and Wire Gauze.-To support vessels over
the gas-lamp, an iron stand is used consisting of a stout vertical rod
fastened into a heavy cast-iron foot, and three iron rings of graduated
sizes secured to the vertical rod with binding-screws; all the rings may
be slipped off the rod, or any ring may be adjusted at any convenient
elevation. The general arrangement is not unlike that of the stand
which makes part of the Berzelius lamp (Fig. VI.), although simpler
and cheaper. As a general rule, it is not best to apply the direct
flame of the lamp to glass and porcelain vessels; hence a piece of
wire gauze is stretched loosely over the largest ring, and bent downwards a little for the reception of round-bottomed vessels; on this
gauze, flasks, retorts, and porcelain dishes are usually supported. In
a few cases, in which a very gradual and equable heat is required, the
wire-gauze is replaced by a small shallow pan, beaten out of sheetiron, and filled with dry sand. This arrangement is called a sand-bath.
With the aid of annealed iron wire, the iron stand may be made available for supporting tubes over the lamp. Crucibles, or dishes, too small
for the smallest ring belonging to the stand, are conveniently supported
on an equilateral triangle made of three pieces of soft iron wire twisted
together at the apices; this triangle is laid on one of the rings of the




X1T                   PNEUMATIC TROUGH.
stand. An iron tripod (that is, a stout ring supported on three
legs) may often be used instead of the stand above described; but
it is not so generally useful, because of the difficulty of adjusting
it at various heights; with a sufficiency of wooden blocks wherewith
to raise the lamp or the tripod as occasion may require, it may be
made available.
10. Pneumatic Trougqh. —The pneumatic trough is a contrivance
which enables us to collect and confine gases in suitable vessels, and
to decant them from one vessel to another. Its efficiency depends on
the pressure of the atmosphere, which, as we know, is capable of supporting a column of water 10'33 metres long or a column of mercury
76 c.m. long (see ~ 7), provided that the liquid column be so arranged
that the atmospheric pressure shall be fully felt upon the foot of the
column, but not at all upon its head. If a tube, closed at one end
and open at the other, and of any length less than 10'33 m., be completely filled with water, and then inverted so that its open end shall
dip beneath some water held in a basin or saucer, the tube will remain
full of water when the thumb or cork, which closed the open end while
the immersion was effected, is withdrawn. What is true of a tube is
equally true of a bell, or other vessel closed at one end, of any diameter
or shape, provided its height be not greater than 10'33 m.; and the
principle which applies to water is equally applicable to mercury,
except that the height of the mercury column which the average
atmospheric pressure can hold up is only 76 c.m., because mercury is
13'596 times as heavy as water. If a few bubbles of any gas insoluble
in water should be delivered beneath the open end of a tube, or bell,
thus standing full of water in apparent defiance of gravitation, the gas
would rise to the top of the tube, by virtue of being lighter than the
water, and the exact volume of water displaced by the gas, small or
large, would drop into the basin or saucer beneath. If the gas were
thus delivered continuously beneath the tube or bell, we should finally
get the tube or bell full of gas, without admixture of air, and sealed
at the bottom by the water in the basin or saucer. If mercury were
the liquid, the operation would be precisely the same, except as
regards the height of the tube or bell. Even this difference of possible
height is not noticeable in practice, because bells and bottles more
than 50 cam. high are very seldom used with either liquid. On account
of its costliness, mercury is rarely used, unless the gas to be collected,
or experimented upon, be soluble in water. A trough for mercury is
made as small as possible for the same reason. It is obvious that the
object of a pneumatic trough may be accomplished under a great
variety of forms. Any bucket or tub with a hanging shelf in it may




PNEUMATIC TROUGH.                    XV
be made to serve. It will be sufficient to describe two convenient
forms of the apparatus.
A cheap pneumatic trough is represented in Fig. XV. Its materials
are durable and its capacity sufficient.      Fig. XV.
It consists of two pieces, 1st, a stoneware pan, about 30 c.m. in diameter on
the bottom, with sides sloping slightly outwards and rising to the height of about 10         J 
c.m.; 2nd, a deep flower-pot saucer about
15 c.m. in diameter, with one hole bored
through the middle of the bottom, and a
second arched hole nipped out of its rim;
this saucer is inverted in the pan. If this
second piece be made expressly for this purpose, it should be made
about 5 c.m. high, and its interior should be rounded to the hole in the
centre, while the outside is left flat like the flower-pot saucer. For
the saucer may be conveniently substituted two slabs or blocks of any
stone, like soapstone, sandstone, or marble, made of even thickness and
laid side by side in the water-pan with an interval between them which
permits the gas-delivery-tube to come beneath the mouth of an inverted
bottle or cylinder supported on the two blocks over the intervening
crack. To use this apparatus, the pan is filled with water to a level
about 2 c.m. above the top of the inverted saucer; the bottle, cylinder,
or bell which is to receive the gds is completely filled with water from
a pitcher or water-cock, then closed with the hand of the operator or
with a flat piece of glass or wood, inverted into the pan, and placed
on the saucer over the hole in its centre; the end of the gas-deliverytube is thrust through the side hole in the saucer; and the gas, rising
through the centre hole, bubbles up into the bottle or cylinder placed
to receive it. While one bottle is filling with gas, another is made
ready to replace it; and when the first is full, it is pushed off the centre
hole of the saucer, and the second bottle is brought over the hole. A
bottle full of gas may be removed from the trough by slipping beneath
the mouth of the bottle a shallow plate or dish, and then lifting plate
and bottle out of the pan together in such a manner that water enough
to seal the mouth of the bottle shall remain in the plate. The gas in
one bottle may be decanted upwards into another, by filling the second
bottle with water, and then carefully inclining the bottle containing the
gas so as to bring its mouth under the mouth of the bottle which is
full of water, keeping the mouths of both bottles all the time beneath
the surface of the water in the pan. If the gas which has been collected is heavier than air, a bottle of it may be withdrawn from the




XVi                     COLLECTING GASES.
water-pan and got at for use, by simply slipping a flat piece of glass or
wood beneath its mouth so as to close it rather tightly, and then
standing the bottle, mouth upward, upon the table. If the cover be
then removed from the bottle, the gas will not flow out, though it will
slowly diffuse into the air. As the water with which the bottles or
cylinders are filled falls into the pan when displaced by gas, it is possible
that the pan may become inconveniently full if many large bottles are
used; this difficulty must be remedied by dipping water out of the
pan, and so restoring the true level.
Where considerable quantities of gas are frequently to be handled,
and large vessels are therefore necessary, a large apparatus, shown in
Fig. XVI., is much more conve-              Fit. XVI.
nient than the small pan, which
suffices for all common experiments. The form of this larger
pneumatic trough and the mode
of using it will readily be understood from the figure; the depth
and width of the tank or well
must be determined by the size
of the bells and cylinders which
are to be sunk in it, and the length and breadth of the shallow part or
shelf by the number of bells or jars of gas which are likely to be in
use at any one time. The deep groove in the shelf permits a glass or
caoutchouc tube to pass without. compression under a bell whose rim
projects over the groove. Such a trough is best made of zinc or lead.
It is very convenient to have it sunk in a table, and permanently provided with a water-cock and drain-pipe.  A chief merit of this instrument is that the glass vessels used can be filled with water by sinking
them in the well much more conveniently than from a pitcher or
water-cock.
A pneumatic trough for mercury may be made either of wood, iron,
or stone. For all common uses, it is very well cut out of a solid block
of compact hard wood, which will not split.  Small cylinders or bells
only can be used, and the well of the trough should be scooped out
but a little larger than the bell or cylinder selected, with its principal
dimension horizontal, and its bottom curved to fit the cylindrical bell
which is to be laid in it; the shelf, too, should have but a small area,
sufficient only for four or five bells of 3 or 4 c.m. diameter.
In using a pneumatic trough, of any construction or dimensions, the
student should be on his guard against two difficulties of possible occurrence-against the sucking back of the liquid in the trough into the




SUCKING BACK AND LEABAGEE.               XVii
ga3-generating apparatus, and against the leakage sometimes induced
by the pressure created by thrusting the gas-delivery-tube deep under
water or mercury. The first of these difficulties is the most serious.
When the flow of gas from a heated flask or tube is suddenly arrested,
in consequence of some reduction of temperature, or from any other
cause, it often happens that the volume of gas in the generating apparatus contracts, and the cold water or mercury firom the trough rises
in the delivery-tube to fill the void; if the contraction is so considerable
as to suffer the cold liquid to penetrate into the hot flask or tube, an
explosion almost inevitablyv ensues, which fractures the apparatus, if it
does no worse damage. In collecting over water a gas somewhat soluble in that liquid, this danger is especially imminent. The occurrence
of such accidents may be effectually guarded against by paying attention to the following directions:-(1) Whenever it is proposed to stop
an evolution of gas which has been going on from a hot flask or tube,
withdraw the delivery-tube from the water before extinguishing the
lamp, and shake off from the bent end of the tube the drops of water
which are apt to adhere to it; the lamp may then be safely put out,
for air can enter the apparatus through the open tube. (2) When the
flow of gas from a hot apparatus is observed to slacken, watch closely the
escape of the gas from the delivery-tube, and as soon as any tendency
to reflux of water is detected, lift the delivery-tube quickly out of the
water, or, better, slip off the caoutchouc connector, which should always
be found between the flask and the water-pan on every such piece of
apparatus; if there be no connector, the cork must be loosened in the
neck of the flask. Air will thus be admitted to the hot flask or tube.
These precautions apply more particularly to the cases where gas is
evolved from dry materials, as in making oxygen or nitrous oxide;
when a liquid is contained in the generating flask, a safety-tube is a
sure protection against the danger of sucking back. The atmospheric
pressure can force air into a flask, in which a partial vacuum has been
created, through the safety-tube, by lifting and displacing a column of
the liquid whose height is the length of that portion of the safety-tube
which dips beneath the liquid. Unless the liquid in the flask be extraordinarily dense, the force required to do this will be very much less
than that required to lift a column of water whose height is determined
by the elevation of the highest point of the delivery-tube above the
level of the water in the pan.
When the gas coming from the generating flask has to force out and
keep out of the delivery-tube a column of water measured from the
lowest point of the tube to the surface of the water in the pan, a pressure determined by the height of this column is established upon the




XV111                    OAS -1fOLDERS.
interior of the flask and upon every joint of the apparatus. Hence an
apparatus will sometimes leak, and refuse to deliver gas at the desired
point, when its delivery-tube is deeply immersed, while it does not
leak if the tube merely dip beneath the surface of the water. With
mercury the pressure of a few centimetres is very considerable, on
account of the high specific gravity of the fluid, so that this difficulty
is more likely to occur with this metal than with water. Tight joints
prevent the occurrence of this difficulty. A partial remedy is to dip
the delivery-tube as little as possible below the surface of the fluid in
the trough.
11. Gas-holders. —A small gas-holder, very convenient for many uses,
is made from a common glass bottle in the following manner: —A
(Fig. XVII.) is a bottle of 4-6 litres     Fig. XVII.
capacity; through the cork in its
neck pass two glass tubes (No. 6), of
which one reaches the bottom of the                      B
bottle, while the other merely penetrates the cork; with the outer end
of the first tube a caoutchouc tube c             b
is connected, with the outer end of
the second a common gas-cock a.
The bottle being first completely',
filled with water, the apparatus                 I'
which generates, or contains, the'I
gas to l)e introduced into the holder    I t]
is connected with the tube carrying
the cock a; this cock is open. As
the gas presses in, the water mounts
in the long tube, and flows out by
the siphon c. In order to relieve
the gas from this pressure at the
beginning, it is only necessary to
suck a little at c. The tube c should
of course be thrust into a sink or
drain-pipe.
To get gas out of the bottle, thus charged, the cock a is closed, and
the flexible tube c is lifted up and connected, as shown in the figure,
with a bottle of water B placed on a shelf, or stand, somewhat above
the bottle A. When the cock b is opened, the gas in A is pressed upon
by the weight of the superincumbent column of water, and may therefore be made to issue at will from the cock a. The higher B is placed
above A, the greater will be the force with which the gas will issue.




GAS-HOLDERS.                      XiX
if a moderate or easily regulated water-pressure is at hand, supplied
by city water-works or a reservoir' in the upper part of the building,
the bottle B is unnecessary, and the flexible tube c may be connected
with such a water-supply whenever gas is to be pressed out of the
gas-holder A.
W Aen larger quantities of gas are to be stored for use, a metallic gasholder, whose construction and proportions are shown in Fig. XVIII., is
advantageously employed. The open cistern
B is supported over the vessel A on two co-   Fig. XVIII.
lumns c c, and two tubes a and b; of these
tubes the first, a, reaches from the bottom
of B nearly to the bottom of A, while the          B
second, b, starts from the bottom of B and
just enters the arched top of A without
projecting into it; d is a short large tube,               a
sloping upwards and outwards, and capable
of being tightly closed with a cork or caoutchouc stopper; g is a glass gauge to show
the height of the water in the vessel A; e    f 
is the discharge-pipe. To fill the gas-holder 
with water, close d, open the stopcocks a, 
b, and e, and pour water into the cistern B;
the water entering A will expel the air
through b and e; when the water begins      b      y
to flow through e, close that stopcock and  d'i
expel the rest of the air through b. The:
gas-holder may now be filled with gas by
displacing the water in the following manner: —Close all the stop.,
cocks, withdraw the cork or stopper from d, and introduce the tube
which delivers the gas through that opening. A.short piece of eaout~.
chouc tubing makes the best end for the gastdelivery-tube; but glass
tubing will answer the purpose if the end be slightly bent upwardi.
The water flows out at d as fast as the gas.eegters, and the gas-holdoea
should therefore stand in a shallow metal tray provided with a drain,..
pipe. When the desired quantity of gas.has.been introduced, close A4.
To draw gas out of a gas-holder of this construction, the cisterni.;filled with water and the cock a is, opened; under the pressure thuds
established the gas may be drawn of Fthrough e, or allowed to. ris'
through b into bottles or bells. filled with water and held- over the,
mouth of the tube b in the cistern B; in this last case B answers the
purpose of a pneumatic trough4.
This gas-holder may be. choapJr.made of zinc; any gasiUteer. cam
2 -s.




XX                        GAS-]EOLDERS.
supply the necessary stopcocks; care must be taken that the glass
tube which constitutes the gauge is fitted air-tight to the gas-holder.
The stopcock e need not end in a screw; tubes may be as well connected with it by caoutchouc. The available pressure under which
the gas in the holder streams out at e is of course limited by the elevation of B above A, which must always be moderate. WNVhen a
stronger pressure is desirable, as in getting the oxyhydrogen blowpipeflamne, for example, a heavier water-column may be obtained by screwing a tall tube with a capacious funnel on the top of it into the tube a,
where it opens into the bottom of the cistern B. A piece of common
iron or copper gas-pipe about a metre long, answers this purpose very
well; the funnel at the top should hold two or three litres, and must
be kept'full of water from a cask or tub provided with a cock and
placed just above the funnel. Where a water-supply, with moderate
pressure, is obtainable, it may be used to keep the funnel full, or to
replace the funnel altogether, if directly connected with the tube a. A
gas-holder, measuring not more than 50 c.m. in total height, is not too
heavy to be portable, and during the process of filling may be placed
over a tub; but a gas-holder of much larger proportions is better made
a fixture, and provided in a permanent manner with drain-pipe and
water-supply. The,gas-holder thus described is that which is the most
generally useful; it may be charged from any
glass flask, retort, or bottle, without any pres-   g. XIX.
sure being exerted upon the glass vessel; and
unused gas contained in any sort of bell, bottle,
or flask can be very readily transferred to suchn
a gas-holder without waste and with very little
trouble. 
A cheaper gas-holder may be made on the
plan of the large gas-holders, improperly called
gasometers, used in gas-works.  Fig. XIX. represents a gas-holder of this sort.  Over a tank
of water, which may be a cylinder of zinc as
shown in the figure, or a headless pork- or oilbarrel, or any other water-tight tub, is ba,1',nc ed
by pulleys and weights a tight bell of zinc, not
too large for complete immersion in the tank.
The U-tube, shown in the figure, which may
be either of lead or brass, serves both to introduce and deliver the gas. To fill such a gasometer, open the cock, lift the counterbalancing weight, and let the bell
sink into the water; then connect the vessel from which the gas is




DEFIAGRATING-SPOON.                    XXi
delivered with the tube of the holder, counterpoise the bell, and the
gas coming from the generator will gradually lift the bell out of the
water. To force the gas out of the holder it is only necessary to remove
the counterbalancing weight; the weight of the bell forces out the gas;
and if this pressure be not sufficient, additional weights may be
placed on the top of the bell. Gas-holders of this construction, unless
very small, are too heavy, when filled with water, to be carried about;
but this difficulty may be obviated, when economy is not specially to
be regarded, by placing within the lower cylinder, or tank, a second
air-tight cylinder as a core, so as to leave only a narrow space between
the inner and outer cylinders for the water into which the upper bell
dips. Elegant, but not cheap, gas-holders are thus made, which are
convenient for some uses, but are not so generally to be recommended
as those of the construction first described. The vessel from which a
gas-holder with counterpoised bell is charged is always subjected to
some pressure, slight if the pulleys, cords, and weights are in perfect
order, but more frequently considerable on account of the difficulty of
maintaining such an apparatus in perfect condition.
12. De.fagrating-Spoon.-The little cup which holds combustible
material, to be burnt in a bottle or jar of gas, is called a deflagratingspoon; it may be cheaply made by hollowing a hemispherical cup out
of a cube of chalk about 2 c.m. on a side, and attaching a stout iron or
brass wire to the chalk, in such a manner that the cup will be right
side up when hung by the wire in a jar of gas; the upper end of this
wire should be straight, that it may be thrust through the cork or piece
of wood which covers the mouth of the bottle or jar. The piece of
chalk may be replaced by a bit of the cylindrical chalk crayon commonly used with blackboards. A piece of crayon 1-5 c.m. in length
will make a sufficient spoon. A small cupel is a convenient ready-made
substitute for the chalk cup. Brass deflagrating-spoons are also to be
had of philosophical-instrument-makers.
13. Platinum Foil and Wire. —A piece of platinum foil about 3 c.m.
square is useful in experiments involving the evaporation of a drop or
two of liquid and the ignition of the residue; it is an essential tool in
qualitative analysis. The foil should be at least so thick that it does not
crinkle when wiped; and it is more economical to get foil which is too
thick than too thin, for it requires frequent cleaning. A bit of platinum
wire, not thicker than a No. 10 needle, and 20 c.m. long, will last a long
time with careful usage. No other metal, and no mixture of substances
from which a metal can be reduced, must ever be heated on platinum
foil or wire, for platinum forms alloys with other metals which are
much more fusible than itself. If once alloyed with a baser metal, the
2s2




xXii                       FILTERING.
platinum ceases to be applicable to its peculiar uses. PlatinUtm may
be cleaned by boiling it in either nitric or chlorhydric acid, by fusing
the acid sulphate of sodium or potassium upon it, or by scouringg it
with fine sand. Aqua regia and chlorine-water dissolve platinum;
the sulphides, cyanides, and oxides of sodium and potassium, when
fused in platinum vessels, slowly attack the metal.
14. Filte;ing.-Filtration is resorted to in order to separate a finely
divided solid from a liquid. The filter may be made of paper, cloth,
tow, cotton, asbestos, and other substances. Paper is the substance
oftenest used. A good filtering-paper must be porous enough to filter
rapidly, and yet sufficiently close in texture to retain the finest powders; and it must also be strong enough to bear, when wet, the pressure of the liquid which must be poured upon it. For delicate experiments it is also necessary that filtering-paper should contain no soluble
salts, and but a very small proportion of incombustible material, which
would remain as ash were the filter to be burned. Filtering-paper
(which is generally sold in sheets) is first cut into circles of various
diameters, adapted to the various scales of operation    Fig. XX.
and quantities of liquids to be filtered. To prepare a
filter for use, one of these circles is folded over on its
own diameter, and the semicircle thus obtained is
folded once upon itself into the form of a quadrant;
the paper thus folded is opened so that three thicknesses shall come upon one side, and one thickness
upon the other, as shown in Fig. XX.; the filter is
then placed in a glass funnel, the angle of which
should be precisely that of the opened paper, viz. 60~;
after being wetted, it is ready to receive the liquid to be filtered. The
paper may be so folded as to fit a funnel whose angle is more or less than
60~; but this is the most advantageous angle, and glass funnels should
be selected with reference to their correctness in this respect.
Coarse and rapid filtering can be effected with cloth bags, Fig. XXI.
also by plucgnino the neck of a funnel loosely with tow or
cotton. If a very acid or very caustic liquid, which would
destroy paper, cotton, tow, or wool, is to be filtered, the best
substances wherewith to plug the neck of the funnel are
asbestos and gun-cotton, neither of which is attacked by
such corrosive liquids.
The glass funnel which holds the filter generally requires
an independent support; for it is seldom judicious, or possible, to support the funnel directly upon the vessel which
receives theJfiltrate, as the clear liquid which runs through the filter is




DRYING GASES.                    Xxi
called. The iron stand may be used for this purpose; and wooden
stands, of a similar construction, adapted expressly for holding funnels,
are very convenient and not expensive. In general, care should be
taken that the lower end of the funnel touch the side or edge of the
vessel into which the filtrate descends, in order that the liquid may
not fall in drops, but run quietly down without splashing. Sometimes
there is no objection to thrusting a funnel directly into the neck of a
bottle or flask; but in this case an ample exit for the air in the bottle
must be provided (Fig. XXI.).
15. Drying Gases.-It is often desirable to remove the aqueous vapor
which is mixed with gases, collected over water or prepared from
materials containing water. It very seldom happens that a gas can be
prepared at one operation in so dry a state as to contain no vapor of
water; this vapor must ordinarily be removed by a subsequent or
additional process. Experience has shown that some gases are more
easily dried than others; thus air, hydrogen, and common oxygen are
thoroughly dried with great ease, but gaseous mixtures which contain antozone only with great difficulty; chlorine is three times as
hard to dry as carbonic acid. These and similar facts must be borne
in mind in constructing drying-apparatus. The common drying-process depends upon bringing the moist gas into contact with some liquid
or solid which greedily and rapidly absorbs aqueous vapor. The three
substances most used for this purpose are concentrated sulphuric acid,
chloride of calcium, and dry quicklime. Sulphuric acid may be used
in two ways: the gas may be made to bubble through a few centimetres depth of the liquid acid, or it may be forced to pass through
the interstices of a column of broken pumice-stone which has been
previously soaked in the acid. The latter method is the most effectual,
because it secures a more  thorough contact of the gas with the hygroscopic acid than is possible durino
the rapid bubbling of the light gasFig. XXII.
through a shallow layer of the dense
liquid.  The column of fragments
of pumice-stone may be held in a
U-tube, arranged like that shown
in Fig. XXII.; but the vertical cylinder shown in the same figure is
better adapted for this use, because
the acid as it becomes dilute from 
absorption of moisture, gradually
trickles from the pumice-stone, and 4
is apt to collect in such quantity at




XXiv             CHLORIDE-OF-CALCIUM U-TUBE.
the bottom of the U-tube as to completely close the tube. In preparing
the upright cylinder for use, the portion below the contraction is not
filled with pumice-stone; it receives the drippings from the pumicestone column. The gas to be dried enters by the lower lateral opening,
and goes out at the top of the cylinder. Though especially adapted
to the column of acid-soaked pumice-stone, this cylinder may very
well be used with either of the other drying-agents, chloride of calcium or quicklime. Either of the forms of drying-tube represented in
Fig. XXII. may be employed with these latter substances; in charging
the horizontal tubes, bits of loose cotton-wool should first be placed
against the exit-tube to prevent any particles of the chloride of calcium, or quicklime, from entering that tube; pieces of the perfectly
dry solid are then introduced in such a way that the tube may be compactly filled with fragments which leave room for the gas to pass very
deviously between them, but offer no direct channels through which
the gas could find straight and quick passage. Quicklime must be
charged much more loosely than chloride of calcium, because of its
great expansion when moistened. Fused chloride of calcium is not so
well adapted for drying gases as the unfused substance. It is not at
all necessary that the fragments of chloride of calcium, or quicklime,
should be of uniform size. When the tube is nearly full, a plug of
loose cotton should be inserted before putting in the cork. A chlorideof-calcium tube, once filled, will often serve for many experiments;
whenever out of use, its outlets should be covered with paper caps; or,
better, caoutchouc connectors may be slipped upon the exit-tubes, and
bits of glass rod thrust into these connectors. The moisture of the air
is thus kept from the chloride of calcium. The dimensions of dryingtubes are of course very various; the bulb-tube shown in Fig. XXII.
is seldom used with a greater length than 25 c.m.; when this form of
tube is employed the gas should invariably enter by the end without
a cork, where the small size of the tube permits direct
connexion with a common gas-delivery-tube by means Fig. XXIII.
of a caoutchouc connector; the other horizontal tube,
shown in the figure, may be of any length; but whenever a great extent of drying-surface is necessary, U-tubes
have the advantage of compactness; for many can be         0
hung upon one short frame. The upright cylinder may
be from 25 c.m. to 40 c.m. in height. A good U-tube,
with an addition which has the merit of economizing
chloride of calcium, is shown in Fig. XXIII.; the addition consists of a short test-tube, into which the tube
by which the gas comes in barely enters; a quantity of water is often




wATER-BATH.                       XXV
caught in this test-tube which otherwise would wet and spoil a considerable amount of chloride of calcium; the little test-tube may, of
course, be taken out and emptied at will.
The choice between one or other of the three drying-substances is
determined in each special case by the chemical relations of the gas to
be dried; thus ammonia-gas, which is absorbed by sulphuric acid and
by chloride of calcium, must be dried by passing it over quicklime,
while sulphurous acid gas, which would combine with quicklime, must
be dried by contact with sulphuric acid.
Fig. XXV.
Fig. XXIV.
16. aSp)ring Clip and Screw Comvressor.-These are very convenient
substitutes for the ordinary stopcock, and as such are in constant use
in the laboratory. Their form, and the manner of their use, will be
readily understood from the figures. As glass stopcocks are expensive
and fragile, and metal stopcocks are usually out of the question, because so many gases and liquids attack the common metals, these
excellent substitutes are used whenever a caoutchouc tube is not inadmissible; they cannot be used unless a bit of elastic tubing can be
inserted into the apparatus which requires a cock.
Another effective mode of temporarily stopping or partially closing
a caoutchouc tube, is to slip over the tube a common brass ring of
about the same diameter as that of the tube, and then to thrust a
slightly conical plug of hard wood or ivory between the ring and the
flexible tube.
17. Water-bath.-It is often necessary to evaporate solutions at a
moderate temperature which can permanently be kept below a certain
known limit; thus, when an aqueous solution is to
be quietly evaporated without spirting or jumping,  Fig. XXVI.
the temperature of the solution must never be suffered to rise above the boiling-point of water, nor
even quite to reach this point. This quiet evaporation is best effected by the use of a water-bath-a 
copper cup whose top is made of concentric rings of
different diameters to adapt it to dishes of various
size (Fig. XXVI.). This cup, two-thirds full of water,




XXVi                     ION BETORT.
is supported on the iron stand over the lamp, and the dish containing the solution to be evaporated is placed on that one of the several
rings which will permit the greater part of the dish to sink into the
copper cup. The steam rising from the water impinges upon the
bottom of the dish, and brings the liquid within it to a temperature
which insures the evaporation of the water, but will not cause any
actual ebullition. The water in the copper cup must never be allowed
to boil away. Whenever a constant supply of steam is at hand, as in
buildings warmed by steam, the copper cup above described may be
converted into a steam-bath by attaching it to a steam-pipe by means
of a small tube provided with a stopcock.
A cheap but serviceable water-bath may be made from a quart
milk-can, oil-can, tea-cannister, or any similarly shaped tin vessel, by
inserting the stem of a glass funnel into the neck of the can through
a well-fitting cork. In this funnel the dish containing the liquor to
be evaporated rests. The can contains the water, which is to be kept
just boiling. On account of the shape of the funnel, dishes of various
sizes can be used with the same apparatus.
The copper vessel first described may be conveniently employed
when it is required to expose substances to a constant temperature
higher than 1000. For this purpose the cup is filled with oil, wax,
paraffine, or a solution of chloride of zinc or chloride of calcium; the
flask or dish containing the substance to be heated should, in this case,
be immersed in the fluid to about two-thirds of its depth; a thermometer must be used to indicate the temperature of such a bath.  7V hen
oil, wax, or paraffine is used, the temperature must not be carried so
high as to burn or decompose these organic matters, else a very disgusting vapor will be produced.
18. Iron Retort.-A retort, made of iron, of the form shown in
Fig. XXVII., is a convenient utensil
in making large quantities of oxygen,       Fig. XXVII.
and in preparing illuminating-gas or
marsh-gas. The iron top is fitted to the
retort with a ground joint, fastened by r
screw-clamp. When the top is removed,
the whole inner surface of the retort is
exposed-a decided advantage wherever
the residue left in the retort after use is
solid. A retort of about 300 c. c. capacity is amply large for most uses. A small iron kettle makes a serviceable
retort; the lid must be luted on, and the nose becomes the exit-tube.
19 &elf-regulating Gas-generator.-An apparatus which is always




SELF-REGULATING GAS-GENERATOR.             XXVii
ready to deliver a constant stream of hydrogen, and yet does not
generate the gas except when it is immediately wanted for use, is a
great convenience in an active laboratory or on a lecture-table. The
same remark applies to the two gases sulphydric acid and carbonic
acid, which are likewise used in considerable quantities, and which can
be conveniently generated in precisely the
same form of apparatus which is advan-       Fig. XXVIIL
tageous for hydrogen.  Such a generator
may be made of divers dimensions. The
following directions, with the accompanying figure (Fig. XXVIII.), will enable the
student to construct an apparatus of convenient size. Procure a glass cylinder 20
or 25 c.m. in diameter and 30 or 35 c.m.
high; ribbed candy-jars are sometimes to
be had of about this size; procure also a
stout tubulated bell glass 10 or 12 c.m.
wide and 5 or 7 c.m. shorter than the cylinder. Get a basket of sheet-lead 7-5 c.m.
deep and 2'5 c.m. narrower than the bellglass, and bore a number of small holes in
the sides and bottom of this basket. Cast
a circular plate of lead 7 m.m. thick and of a diameter 4 c.m. larger
than that of the glass cylinder; on what is intended for its under side
solder three equidistant leaden strips, or a continuous ring of lead, to
keep the plate in proper position as a cover for the cylinder. Fit
tightly to each end of a good brass gas-cock a piece of brass tube 8 c.m.
long, 1'5 to 2 c.m. wide, and stout in metal. Perforate the centre of
the leaden plate, so that one of these tubes will snugly pass through
the orifice, and secure it by solder, leaving 5 c.m. of the tube projecting below the plate. Attach to the lower end of this tube a stout hook
on which to hang the leaden basket. By means of a sound cork and
common sealing-wax, or a cement made of oil mixed with red and
white lead, fasten this tube into the tubulature of the bell glass airtight, and so firmly that the joint will bear a weight of 2 or 3 kilos.
Hang the basket by means of copper wire within the bell 5 c.m. above
the bottom of the latter. To the tube which extends above the stopcock attach by a good cork the neck of a tubulated receiver of 100 or
150 c. c. capacity, the interior of which has been loosely stuffed with
cotton. Into the second tubulature of the receiver fit tightly the delivery-tube carrying a caoutchouc connector; into this connector can
be fitted a tube adapted to convey the gas in any desired direction




XXViii                    GLASS WARE.
To charge the apparatus, fill the cylinder with dilute acid to within
10 or 12 c.m. of the top, place the zinc, sulphide of iron, or marble, as
the case may be, in the basket, hang the basket in the bell, and put
the bell-glass full of air into its place with the stopcock closed. On
opening the cock, the weight of the acid expels the air from the bell,
the acid comes in contact with the solid in the basket, and a steady
supply of gas is generated until either the acid is saturated or the
solid dissolved; if the cock be closed, the gas accumulates in the bell,
and pushes the acid below the basket so that all action ceases. In
cold weather the apparatus must be kept in a warm place. For generating hydrogen, sulphuric acid diluted with four or five parts of water
is used; for sulphydric acid, sulphuric acid is diluted with fourteen
parts of water; for carbonic acid, chlorhydric acid diluted with two or
three parts of water is to be preferred.
20. Glass Retorts, Flasks, Beakers, Test-tubes and Test-glasses.-All
glass vessels which are meant for use in heating liquids must have
uniformly thin bottoms. Tubulated retorts are much more generally
useful than those without a tubulature. As retorts are expensive in
comparison with flasks, they are less used than formerly.
The neck of a flask should have such a form that it can be tightly
closed by a cork, and the lip must be strengthened to resist the force
used in pressing in the cork, either by a rim of glass added on the
outside, or, better, by causing the rim itself to flare outward. The
actual edge of the rim must never be sharp or rough, but always smooth
and rounded by partial fusion. The thin-bottomed flasks in which
olive-oil is sometimes imported from Italy are excellent for chemical
uses; their edges always require, however, to be rounded in the lamp.
These Florence flasks, which are very much to be recommended both
on the ground of cheapness and of durability, may be cleaned by
soaking them 24 hours in a weak caustic lye, and then washing them
with boiling water.
Beakers are thin flat-bottomed tumblers with a slightly flaring rim.
They are to be bought in sets or nests which sometimes include a large
range of sizes. The small sizes are very useful vessels; the large are
so fragile as to be almost worthless. Up to the capacity of about a
litre, beakers are to be recommended for heating liquids whenever it is
an object to have the whole interior of the vessel readily accessible.
Test-tubes are little cylinders of thin glass, with round, thin bottoms,
and lips slightly flared. Their length may be from 12 c.m. to 18 c.m.,
and their diameter 1 c.m. to 2 c.m.; they should never have a diameter so large that the open end cannot be closed by the ball of the
thumb. To hold the tubes upright a wooden rack is necessary. Be



MEASURING-GLASSES AND BURETTES.             XXiX
sides the row of holes to receive a dozen test-tubes bottom down, the
rack should have a row of pegs on which the test-tubes may be inverted when not in use; in this position the water in which they are
rinsed drains off, and dust cannot be deposited within the tubes. Testtubes are much used for heating small quantities of liquid over the gasor spirit-lamp; they may generally be held by the upper end in the
fingers without inconvenience, but if a liquid is to be boiled for some
minutes in a test-tube, the tube must be held in wooden nippers, or in
a strip of thick folded paper, nipped round the tube and grasped between the thumb and fore finger just outside the tube. The wooden
nippers above mentioned are made of two bits of wood about 30 c.m.
long hinged together at the back, and at once connected and kept apart
by a sliding steel or brass spring, somewhat like those used on certain
pruning-shears and some kinds of steel nippers. When a liquid is
boiling actively in a test-tube, it sometimes happens that portions of
the hot liquid are projected out of the tube with some force; the
operator should therefore hold the tube in an inclined position, rolling
it to a slight extent between the thumb and fore finger in order that
all sides of the tube may be equally exposed to the flame; while thus
using the tube, he should be careful not to direct it either towards
himself or towards any other person in his neighbourhood. Test-tubes
are cleaned by the aid of cylindrical brushes made of bristles caught
between twisted wires, like those used for cleaning lamp-chimneys:
they should have a round end of bristles.
Two precautions are invariably to be observed in heating glass and
porcelain vessels of whatever form: first, the outside of the vessel to
be heated must be made perfectly- dry; secondly, the temperature must
not be raised too rapidly. When a large flask or beaker containing a
cold liquid is first warmed over a lamp, moisture almost invariably
condenses upon the bottom of the vessel: this moisture should be
wiped off with a cloth.
Stout conical glasses with strong stems and feet are convenient
for many uses not involving the application of heat. They are called
test-glasses, and may be had of various shapes and sizes. It is obvious
that cheap wine- or beer-glasses and common jelly-tumblers would
answer the purposes which these test-glasses serve.
21. Measuring-glasses and Burettes.-Measuring-glasses, divided into
cubic centimetres, are made in the cylindrical form, and also in the
flaring shape common in druggists' measuring-glasses; the cylindrical
form is to be preferred. Such a glass of 250 c. c., or better of 500
c. c. capacity, is a very useful implement; a flask holding just one
litre when filled to a mark upon its neck is also convenient. Smaller




XXX                    READING BURETTES.
quantities of liquid are measured with burettes. Mohr's burette is
the most generally useful of all forms of this instru-  Fig. XXIX.
ment (see Fig. XXIX., the right-hand instrument);
it is a graduated tube drawn to a small bore at the
bottom; a caoutchouc connector is slipped upon the       o
bottom of the tube, a short bit of tube drawn to a
fine point is thrust into the lower end of the con-      10
nector, and a spring clip nips the connector between
the two glass tubes. The spring clip closes the bottom of the burette, but it can of course be opened      30
at will to permit the liquid in the burette to flow or
40
drop out. The caoutchouc affects injuriously some
of the liquids which are used in burettes; so that this
common form of Mohr's burette is not always applicable. To avoid this difficulty the instrument may
be made with a glass cock, but is then rather costly.
Gay-Lussac's burette is available whenever the caoutchouc in Mohr's
burette is objectionable. The construction of Gay-Lussac's burette is
plainly to be seen in the figure (see Fig. XXIX., the left-hand instrument); a narrow tube runs up beside the large graduated tube to the top
of the latter, and the liquid can be poured out in drops by gently inclining the instrument. Its fraoility is a serious objection to this form
of burette; the danger of breaking off the small tube is lessened, if a
small piece of cork be inserted between the two tubes at the top, and
a string tied round them both. A wooden foot in which it may stand
upright upon a table is a convenient addition to Gay-Lussac's burette.
Mohr's burette must be held upright in a suitable screw clamp, or
be fastened on to a wooden frame in such a manner that the tube shall
be vertical, firm, and at the same time easily detached. The fineness of
measurement may be increased, without impairing the distinctness of the
scale, by reducing the bore of the burette. For delicate work, burettes
divided into tenths of a cubic centimetre are employed. The way in
which the reading-off is effected is a matter of importance in using a
burette; it is essential, 1st, that the burette should be vertical; 2nd,
that the eye should be brought to a level with the surface of the fluid;
3rd, that a fixed standard should be adopted of what is to be considered the surface. If a burette, partly filled with a liquid, be held
between the eye and a white wall, the surface of the liquid presents
a light line which is nearly level, and just below this line a second
line, which is dark and curved with the convexity downward. If a
sheet of white paper be held immediately behind the tube, these two
lines, though somewhat altered in appearance, are still distinctly




FILLING BURETTES.                   XXXI
visible. They may be made still more distinct by using instead of
white paper a card half white, half black, with a straight dividing line
between the two colors. On holding this card with the white hah
uppermost, and the border line between white and black fiom 2 to f
m.m. below the lowest visible dark line, two zones are brought out,
a light zone and a dark zone, and the lower limit of the dark zone is
made very distinct. Care must be taken to hold the card invariably
in the same relative position, since, if it be held lower down, the lower
border of the dark zone will move higher up. In practice the lower
border of the dark zone is read off as the surface of the liquid, this
being the most distinctly marked line. There is one exception to this
rule; when an opaque solution of permanganate of potassium is to be
measured in Gay-Lussac's burette, the upper border of the dark zone
must be held to be the surface of the liquid; in this case it is best to
place the burette against a white background.
The zero of the graduated scale on a burette is always near the top
of the tube. In order to fill Mohr's burette, the point of the instrument is dipped into the liquid, the spring clip opened, and a little liquid,
sufficient at least to reach into the burette tube, is sucked up by applying the mouth to the upper end; the spring clip is then closed, and
the liquid poured into the burette through the upper end until it has
risen a little above the zero-line. By opening the spring clip, the
liquid is then allowed to drop out until the exact level of the zero-line
is reached. The instrument is then ready for use. When a quantity
of liquid has been allowed to flow out of a burette thus filled, and the
operator desires to read off the amount used, he must wait a few
moments to give the particles of fluid adhering to the sides of the
emptied portion of the tube time to run down. This remark applies
to all forms of burette. Erdmann's swimmer is an excellent addition
to Mohr's burette. It is a cylindrical glass float of such a width as
nearly to fill the burette, but yet so loosely as to float Fig.XXX.
freely up and down with the liquor. To set the instrument
at zero a ring cut round the swimmer is brought to coincide with the line 0 engraved on the burette. The absolute height of the liquor in the burette is to be disregarded.
In order to read the height of the liquor in the burette at
any time, it is only necessary to note that line on the scale
with which the mark cut round the float coincides.
22. Pipettes.-Pipettes are tubes drawn to a point, and
sometimes furnished with a bulb or a cylindrical enlargement. They are chiefly used to suck small quantities of fluid out of
a vessel without disturbing the bulk of the liquid. Figure XXX. re



gXi1l                  PORCELAIN WARE.
presents three forms of pipette; the form with the lower end bent
upwards is used to introduce liquids into a bell or bottle of gas standing over mercury. Pipettes graduated into cubic centimetres, or holding
a certain number of cubic centimetres when filled to a mark on the
stem, are often convenient.
23. Washl-bottle.-A wash-bottle is a flask with a uniformly thin
bottom, closed with a sound cork or caoutchouc stopper, through
which pass two glass tubes as shown in Fig. XXXI. The outer
end of the longer tube is drawn to a moderately
fine point.  A short bend near the bottom of this
longer tube in the same plane and direction as the
upper bend is of some use, because it enables the
operator to empty the flask more completely by inclining it. By blowing into the short tube, a stream of
water will be driven out of the long tube with considerable force. This force with which the stream is
projected adapts the apparatus to removing precipitates
from the sides of vessels as well as to washing them on
filters. For use in analytical operations, it is often  (
convenient to attach a caoutchouc tube 12 or 15 c.m.
long to the tube through which air is blown; this
flexible tube should be provided with a glass mouthpiece about 3 c.m.
long. As the wash-bottle is often filled with hot or even boiling
water, it may be improved by binding about its neck a ring of cork, or
winding closely round the neck a smooth cord. It may then be handled
without inconvenience when hot.
24. Porcelain Dishes and Crucibles.-Open dishes which will bear
heat without cracking are necessary implements in the laboratory for
conducting the evaporation of liquids. The best evaporating-dishes
are those made of Berlin porcelain, glazed both inside and out, and
provided with a little lip projecting beyond the rim. The dishes made
of Meissen porcelain are not glazed on the outside, and are not so
durable as those of Berlin manufacture; but they are much cheaper,
and with proper care last a long time. The small Berlin dishes will
generally bear an evaporation to dryness on the wire gauze over the
open flame of the gas-lamp; the Meissen dishes do not so well endure
this.severe treatment. Evaporating-dishes are made of all diameters,
from 3 c.m. to 45 c.m.; they should be ordered by specifying the diameter desired. The large sizes are expensive, and not very durable;
they should never be used except on a sand-bath. Dishes of German
earthenware are as good as porcelain for many uses, and are much to
be recommended in place of the large sizes of porcelain dishes.




CRUCIBLES.                      Xxxiii
Very thin, highly glazed porcelain crucibles with glazed covers are
made both at Berlin and at Meissen near Dresden; they are indispensable implements to the chemist. In general, the Meissen crucibles are
thinner than the Berlin, but the Berlin crucibles are somewhat less
liable to crack; both kinds are glazed inside and out, except on the
outside of the bottom. Crucibles should be ordered by specifying the
diameters of the sizes desired; they are to be had of nearly a dozen
different sizes, with diameters varying from 2 c.m. to 9 c.m. The
smallest and largest sizes are little used; for most purposes the best
sizes are those between 3 c.m. and 5 c.m. in diameter. As the covers
are much less liable to be broken than the crucibles, it is advantageous
to buy more crucibles than covers, whenever it is possible so to do.
Porcelain crucibles are supported over the lamp on an iron-wire triangle; they must always be gradually heated, and never brought suddenly into contact with any cold substance while they are hot.
25. Rilngs to support round-bottomed vessels.-It is often necessary to
support globes, round-bottomed flasks, evaporating-dishes, and round
receivers in a stable manner upon the table or other flat surface. For
this purpose rings are used made of braided straw, or of straw wound
about a core of straw, or of tin wound with listing or coarse woollen
cloth. The material of which these rings are made, or with which
they are covered, ought to be a substance which does not conduct heat
well, because one of the chief uses of these rings is to receive hot
vessels just removed from the lamp or sand-bath. A hot flask or
dish would almost certainly be broken, if it were placed upon the cold
surface of a good conductor of heat. The student must never touch
a hot vessel with cold water, or bring it into sudden contact with a
surface of marble, iron, copper, or other conductor of heat.
26. Crucibles. —For use in a coal fire there are three good kinds of
crucibles, each of which has its own merits which recommend it for
certain purposes. The Hessian crucibles are sold in nests containing
from 3 to 10 crucibles; there are 10 sizes, which vary from 3 to 25
c.m. in height. They generally have a triangular form, and will withstand a very high temperature, if they are warmed before being put
in the fire. They are not sold with covers; but covers may be bought
separately, or a triangular piece of soapstone may be very conveniently
used as a cover for these crucibles. Hessian crucibles are cheaper
than any other kind, and are therefore the most used. The French
crucibles, called Beaufay crucibles, are admirable, but too dear for
common use. They have a tall, narrow form, a smooth surface, and a
small lip; covers for the crucibles are sold separately. The crucibles
ire sold, not in nests, but singly; there are 22 sizes, which vary from




Xxxiv                      MORTARS.
4 to 40 c.m. in height. They are highly refractory. Plumbago crucibles are used for the fusion of the most refractory metals-gold, silver,
copper, brass, steel, iron, and so forth; they resist better than any
other crucibles the combined action of a very high temperature and a
strong flux; and as they are not liable to crack, they may often be
used many times without risk. Their first cost is higher than that
of any other crucibles. Crucibles are mainly used for the fusion and
reduction of metals; but there are also many chemical compounds
which can only be prepared at the very high temperatures which by the
use of crucibles we are able to command. Although crucibles often
withstand the most sudden changes of temperature, it is nevertheless
expedient as a general rule to heat up a crucible gradually, and to previously warm a charge which is to be placed in a crucible already hot. If
a cold crucible is to be introduced into a fire, it should first be placed in
the least hot part of the fire and gradually brought into the hottest part.
27. Tongs and Pincers.-Hot              Fi. XXXII.
crucibles are handled by means
of tongs of various shapes and
sizes, according to the weight and
nature of the vessels to be lifted.
Fig. XXXII. represents two good
forms of stout iron tongs for lifting
large crucibles out of a coal fire. The manner of using them is readily
understood from the figure.
Small porcelain crucibles are handled, when   Fig. XXXIII.
hot, by means of small steel or iron tongs, such             a
as are represented in Fig. XXXIII.
Small steel pincers (jewellers' tweezers) are
applied in the laboratory to a great variety of uses.
28. Mortars.-Iron, porcelain, and agate mortars are used by chemists to reduce solids to powder. An iron mortar is useful for coarse
work, and for effecting the first rough breaking up of substances which
are subsequently powdered in the porcelain or agate mortar. If there
be any risk of fragments being thrown out of the mortar, it should be
covered with a cloth or piece of stiff paper, having a hole in the
middle through which the pestle may be passed. Pieces of stone,
minerals, and lumps of brittle metals may be safely broken into fragments suitable for the mortar by wrapping them in strong paper,
laying them so enclosed upon an anvil, and striking them with a heavy
hammer. The paper envelope retains the broken particles which
might otherwise fly about in a dangerous manner and be lost.
The best porcelain mortars are those known by the name of Wedge



PULVERIZING.                     XXXV
wood-ware; but there are many cheaper substitutes. Porcelain inortars will not bear sharp and heavy blows; they are intended rather for
grinding and trituration than for hammering; the pestle may either be
formed of one piece of porcelain, or a piece of porcelain cemented to a
wooden handle; the latter is the less desirable form of pestle. Unglazed porcelain mortars are to be preferred. In selecting mortars, the
following points should be attended to:-1st, the mortar should not be
porous; it ought not to absorb strong acids or any colored fluid, even
if such liquids be allowed to stand for hours in the mortar; 2nd, it.
should be very hard, and its pestle should be of the same hardness;
3rd, it should be sound; 4th, it should have a lip for the convenience
of pouring out liquids and fine powders. As a rule, porcelain mortars
will not endure sudden changes of temperature. They may be cleaned
by rubbing in them a little sand soaked in nitric or sulphuric acid, or,
if acids are not appropriate, in caustic soda.
Agate mortars are only intended for trituration; a blow would break
them. They are exceedingly hard, and impermeable. The material is
so. precious and so hard to work, that agate mortars are always small.
The pestles are generally inconveniently short,-a difficulty which may
be remedied by fitting the agate pestle into a wooden handle.
In all grinding-operations in mortars, whether of porcelain or agate,
it is expedient to put only a small quantity of the substance to be
powdered into the mortar at once. The operation of powdering will
be facilitated by sifting the matter as fast as it is powdered, returning
to the mortar the particles which are too large to pass through the
sieve.
29. Spatula.-For transferring substances in powder, or in small
grains or crystals, from  one vessel to another, spatulle and scoops
made of horn or bone are convenient tools. A coarse bone paperknife makes a good spatula for laboratory use. Cards free from glaze
and enamel are excellent substitutes for spatulce.
30. Thermometers.-Thermometers intended for chemical use must
have no metal, and no wood or other organic material upon their outer
surfaces; their external surfaces must be wholly of glass. The best
thermometers are straight glass tubes, of uniform diameter, with cylindrical instead of spherical bulbs, and having the scale engraved upon
the glass; such instruments can be passed tightly through a cork, and
are free from many liabilities to error to which thermometers with paper
or metal scales are always exposed. A cheaper kind of thermometer,
having a paper or, better, enamel scale enclosed in a glass envelope, wili
answer for most experiments.
31. Furnaces.-For all common fusions, an anthracite or coke fire
2T




XXXVi                THE METRICAL SYSTEM.
in an ordinary cylinder stove will suffice. The chafing-dish, or open
portable stove, such as is used by plumbers for example, is very convenient for operations which require less heat. The clay buckets used
as open furnaces are better than the iron ones, because they hold the
heat better.
Charcoal is the fuel used in these open fires. A very useful accompaniment to these portable furnaces is a piece of straight stove-pipe,
about 60 c.m. long and 10 c.m. wide, and flaring out below like a funnel until it is wide enough to cover the top of the furnace. This contrivance powerfully increases the draught, and is used to urge the fire
during kindling, or to intensity it while a fusion is in progress. With
a furnace of this description there is no difficulty in keeping a small
crucible white-hot for a short time.
THE METRICAL SYSTEM OF WEIGHTS AND MEASURES.
The metrical system, employed in the aflairs of every-day life by
most of the nations of continental Europe and by scientific writers
throughout the world, is based upon a fundamental unit or measure
of length, called a metre. This metre is defined as the 40-millionth
part of the circumference of the earth, or, in other words, of a "great
circle " or meridian; its length was originally determined by actual
measurement of a considerable arc of a meridian; but the various measurements heretofore made of the length of the earth's meridian differ
slightly from each other, and it is to be expected, and indeed hoped,
that the steady improvements of methods and instruments will make
each successive determination of the lenoth of the meridian better
than, and therefore different from, the preceding.  It is therefore
necessary to define the standard of length, by legislation, to be a certain
rod of metal, deposited in a certain place under specified guaranties,
and to secure the uniformity and permanence of the standard by the
multiplication of exact copies in safe places of deposit.
From this single quantity, the metre, all other measures are decimally derived. Multiplied or divided by 10, 100, 1000, and so forth,
the metre supplies all needed linear measures, and the square metre
and cubic metre, with their decimal multiples, supply all needed
measures of area or surface on the one hand, and of solidity or capacity on the other.
From the unit of measure to the unit of weight the transition is
admirably simple and convenient. The cube of the 1-hundredth of the
linear metre is, of course, the millionth of the cubic metrx; its bulk is




THE METRICAL SYSTEM.,                 XXXVii
about that of a large die of the common back-gammon board. This
little cube of pure water is the univeisal unit of weight, a grammne,
which, decimally multiplied and divided, is made to express all weights.
The numbers expressing all weights, from the least to the greatest,
find direct expression in the decimal notation; the weights used in
different trades only differ from each other in being different decimal
multiples of the same fundamental unit; and in comparing together
weights and volumes, none but easy decimal computations are ever
necessary.
The nomenclature of the metrical system is extremely simple; one
general principle applies to each of the following tables. The Greek
prefixes for 10, 100, and 1000, viz. deca, hecto, and kilo, are used to
signify multiplication, while the Latin prefixes for 10, 100, and 1000,
viz. deci, centi, and milli, are employed to express subdivision. Of the
names thus systematically derived from that of the unit in each table,
many are not often used; the names in common use are those printed
in small capitals. Thus, in the table for linear measure, only the
metre, kilometre, centimetre, and millimetre are in common use,the first for such purposes as the English yard subserves, the second
instead of the Enolish mile, the third and fourth in lieu of the fractions
of the English foot and inch.
LINEAR MEASURE.
Metre.
MTLTLTIETRE      =           0001 or 1-1000th of a metre.
Divisions... CFNTIMIETRE     -=           0 01 or 1-100th of a metre.
Decimetre                    0-1  or 1-lOth  of a metre.
Unit............METRE  1
f Decametre        =          10'
Multiples... tHectometre      =         100'
l KILO3IETRE               1,000'
SURFACE MEASURE.
(Millimetre square     =       0000,001 of a metre square.
Divisions...  Centimetre square           0'000,1  of a metre square.
Decimethe square     =        001      of a metre square.
Unit............METRE SQUARE             1'        metre squar'.
CUBIC MEASURE.
Cubic Metre.
(Cubic Millimetre        =                  0'000,000,001
Divisions... Cubic Centimetre       -                  0000,001
Cubic Decimetre        -                   0001
Unit.............CUBIC  ETRE                         i.
U~nit. CUBIC METRE                 =            -    1
Cubic Decametre        =               1,000'
Multiples... Cubic Hectometre       =           1,000o000
Cubic Kilometre        -        1,000,000,0002 T 2




Xxxviii                THE METRICAL SYSTEM.
The table for land-measure we omit, as having no connexion with
our subject.  For the measurement of wine, beer, ci., grain, and similar wet and dry substances, a smaller unit than the cubic metre is
desirable. The cubic decimetre has been selected as a special standard
of capacity for the measurement of substances such as are bought and
sold by the English wet and dry measures. The cubic decimetre thus
used is called a litre.
CAPACITY MEASURE.
Litres. Cubic }Metre.
f Millilitre   =   0'001 = 0000,001 = 1 cubic centimetre.
Divisions...  Centilitre   -   0'01  = 0000.01
Decilitre         0'1   = 0'000,1
1nit...........LITRE          1    = 0'001    = 1 cubic decimetre.
r Decalitre       10    = 0'01
Multiples..  HECTOLITRE =   100    = 0'1
Kilolitre    = 1,000'  1            = 1 cubic metro.
The table of weights bears an intimate relation to this table of
capacity. As already mentioned, the weight of that die-sized cube, a
cubic centimetre or millilitre of distilled water (taken at 4~, its point of
greatest density) constitutes the metrical unit of weight. This weight
is called a gramme.  From the very definition of the gramme, and from
the table of capacity-measure, it is clear that a litre of distilled water
at 4~ will weigh 1000 grammes.
WEIGHTS.
Grammes.
MI LLTGRAIME =     0'001
Divisions... CENTIGRAn ME =    001
DECIGRAMME         0.1
Unit............ GRAMIFE  1 = 1 cubic centimetre of water at 4~.
Decagramme  =    10'
Multiples.. Hectogramme =   100
Kilogramme  = 1,000' = 1 cubic decimetre of water at 4~.
The simplicity and directness of the relations between weights and
volumes in the metrical system can now be more fully explained.
The chemist ordinarily uses the gramme as his unit-weight, and for
his unit of volume a cubic centimetre, which is the bulk of a gramme
of water. For coarser work, the kilogramme becomes the unit of
weight, and the corresponding unit of measure is the litre, which is
the bulk of a kilogramme of water.  In commercial dealings, in manufacturing-processes, and, above all, in scientific investigations, these
simple relations between weights and measures have been found to be
an inestimable advantage. The numerical expressions for metrical
weights and measures may always be read as decimals.  Thus 5-126




THERMOMETERS COMPARED.              XXXiX
metres will be read five metres and one hundred and twenty-six thousandths, and not five metres, one decimetre, two centimetres, and six
millimetres. The expression 10'5 granmmes is read ten and five-tenths
grammes; just as we say one hundred and five dollars, not ten eagles
and five dollars; or sixty-five cents, not six dimes and five cents. All
computations under the metrical system are made with decimals alone.
The abbreviations commonly met with in. chemical literature are:m.m. for millimetre;          c.m. for centimetre;
m. for metre;                 c. c. for cubic centimetre;
grm. for gramme;              kilo. for kilogramme.
The use of the metrical system of weights and measures in the
arts and trades has been legalized both in the United States and in
Great Britain.  The United States coin, composed of copper and
nickel, of the denomination 5 cents, and date 1866, is 2 c.m. in diameter, and weighs 5 grms.
TABLE for the Conversion of Degrees on the Centigrade Thermometer
into Degrees of Falhrenheit's Scale.
Cent.   Fahr.  Cent.   Fahr.  Cent.   Fahr.  Cent.   Fahr.
-50   -58~0  -26   -28.2  - 8          i6   173       55.4
-49   -56'2  -28   -18'4  - 7          19'4  14       572 i
-48   -54'4  -27   -16'6  - 6          21'2  15       59'0
-47   -52'6  -26   -14'8  - 5          23'0  16       60'8
-46   -50'8  -25   — 13'0  - 4         24'8  17       62'6
-45   -49'0   -24   -11'2  - 3         266-  18       64'4
-44   -47'2  -23   - 9'4  - 2          28-4  19       66'2
-43   -45'4  -22   - 7-6  - 1          30'2  20       68'0
-42   -43'6  -21   - 5'8         0     32'0  21       69'8
-41   -41'8  -20   - 4'0  + 1          33'8  22       71'6
-40   -40'0  -19   - 2'2         2     35'6  23       73'4
-39   -38'2  -18   - 0'4         3     37'4  24       75'2
-38   -36'4  -17   + 1'4         4     39'2' 25       77'0
-37   -34'6  -16         3'2     5     410  26        78'8
-36   -32'8  -15         5'0     6     42'8  27       80'6
-35   -31'0  -14         6'8     7     44'6  28       82'4
-34   -29'2   -13        8'6     8     46'4  29       84'2
-33   -274  -12         104      9     48'2  30       860
-32   -25-6  -11        12 2    10     50'0  31       87'8
-31   -23'8  -10        14'0    11     51'8  32       89'6
-30   -22'0  - 9        15'8    12     53'6  33       91'4




xl                 THERMOMIETERS COMPARED.
TABLE (continued).
Cent.   Fahr.  Cent.   Fahr.  Cent.   Fahr.  Cent.   Fahr.
3a       9302   78      172'4  122      25'6  166       330'8
35       95'0   79      174'2  123      253'4  167      332'6
36       968 1 80       176.0  124      2552  168       334.4
37       986 1 81       177'8  125      257'0  169      336'2
38      100'4 i 82      179'6  126      258'8  170      33880
39      102'2  83       181'4  127      260'6  171      339'8
40      104'0   84      183'2  128      262'4  172      341'6
41      105'8   85      185'0  129      264'2  173      343'4
42      107'6   86      186'8  130      266'0  174      345'2
43      109'4   87      188'6  131      267'8  175      347'0
44      111'2   88      190'4  132      269'6  176      348'8
45      113'0   89      192'2  133      271'4  177      350'6
46      114'8   90      194'0  134      273'2  178      352'4
47      116.6   91      195.8  135      275.0  179      354.2
48      118'4   92      197'6  136      276'8  180      3.560
49      120'2   93      199'4  137      278'6  181      357'8
50      122'0   94      201'2  138      280'4  182      359'6
51      123'8   95      203'0  139      282'2  183      361'4
52      125'6   96      204'8  140      284'0  184      363'2
53      127'4   97      206'6  141      285'8  185      365'0
54      129'2.~ 98     208'4  142      287'6  186      366'8
55      131'0   99      210'2  143      289'4  187      368'6
56      132'8  100      212'0  144      291'2  188      370'4
57      134'6  101      213'8  145      293'0  189      372'2
58      136'4  102      215'6  146      294'8  190      374'0
59      1382 1 103      217'4  147      296 i6 191      375'8
60      140'0  104      219'2  148      298'4  192      377'6
61      141'8  105      221'0  149      300'2  193      379'4
62      143'6  106      2228  150       302'0  194      381'2
63      145'4  107      224'6  151      303'8  195      38380
64      147'2  108      226'4  152      305'6  196      384'8
65      149'0  109      228'2  153      307'4  197      386'6
66      150'8  110      230'0  154      309'2  198      388'4
67      152'6  111      231'8  155      311'0  199      390'2
68      154'4  112      233'6  156      312'8  200      392'0
69      156'2  113      235'4  157      314'6  201      393'8
70      158'0  114      237'2  158      316'4  202      39.56
71      159'8  115      239 0  159      318'2  203      397'4
72      161'6  116      240'8  160      320'0  204      399'2
73      163'4  117      242'6  161      321'8  205      401'0
74      165'2  118      244'4  162      323'6  206      402'8
75      167'0  119      246'2  163      325'4  207      404'6
76      168'8  120      248'0  164      327'2  208      406'4
77      170'6  121      249'8  165      329'0  209      4082




THERMOMETERS COMPARED.                   xli
TABLE (continued).
Cent.   Fahr.  Cent.   Fahr.  Cent.   Fahr.  Cent.   Fahr.
0       0      0        0      0       0      0
210     41080  238     460'4  266      510-8  294      561f2
211     411-8  239     462'2  267      512-6  295      563-0
212     413-6  240     464'0  268      514'4  296      564-8
213     415-4  241     46,5'8  269     516'2  297      566-6
214     417-2  242     467-(  270      518-0  298      56;84
215     419-0  243     469-4  271      519'8  299      570.2
216     420-8  244     471-2  272      521-6  300      572-0
217     422-6  245     473'0  273      523-4  301      573'8
218     424-4  246     474'8  274      525'2  302      575-6
219     426-2  247     476'6  275      527'0  303      577'4
220     428'0  248     478'4  276      528'8  304      579'2
221     429'8  249     480'2 1 277     530'6  305      581'0
222     431 6  250     482'0  278      532'4  306      582'8
223     433;4  251     483'8  279      534'2  307      584'6
224     435'2  252     485'6  280      536'0  308      586'4
225     437'0  253     487'4  281      537-8  309      5882
226     438'8  254     489'2  282      539'6  310     590'0
227     440'6  255     491'0  283      541'4  311     591'8
228     442'4  256     492'8  284      543-2  312     593'6
229     444'2  257     494'6  285      545'0  313     595'4
230     446'0  258     496'4  286      546-8  314     597'2
231     447'8  259     498'2  287      548'6  315      599'0
232     449'6  260     500'0  288      550'4  316      600'8
233     451'4  261      501'8  289     552'2  317      602'6
234     453'2  262      503'6  290     554'0  318      604'4
235     455'0  263    56054  291       555'8  319      6062
236     456'8  264      507'2  292     557'6  320     608'0
237     458'6  265     509'0  293      559'4




TABLE FOR CONVERSION OF GRAMPMES INTO GRAINS AND CENTIMETRES INTO INCHES. 
The equivalents in Engrlish weights and measures of those metrical weights and measures which are used in
chemistry can be readily found by the aid of the following table, which is available not only for grammes and
centimetres, but, by mere change of the position of the decimal point, for all decimal multiples or subdivisions of
these quantities:Grammes        1        2|   4                       5        6         7         8         9
into Grains.  15-4346 3008692  463038   61'7384   77-1730   92'6076  108"0422  123'4768  1389114
Centimetres     1        2        3         4                   6        7         8         9
into Inches.   3937079 *7874158 1 1811237 1 5748316 19685395' 23622474 2-7559553 3 14966~32 3 5433711 
One cubic metre            = 35'31660 cubic feet.    One pound avoirdupois           = 7000' grains.,,,,  decimetre (a litre) = 61-02709,,  inches.,,,   troy                  = 5760',,,,  centimetre        =  0-06103,,,,,, ounce avoirdupois           =  437.5,,,, litre                  -  0-22017 imp. gallon.,,,,  troy                   =  480 
=  0'88066,, quart.,, imperial gallon              =  277'274 cub. ins.,  1'76133, pint.,, old English wine-measure gall. =  231-,,,
(Still used in the United States.)




INDEX.
*** Compound names are to be found under tile first word of the namechloride of sodium, for example, under chloride, sulphuric acid under sulphuric.  Occ. stands for occurrence, prep. for preparation, prop. for properties, comp. for composition. The numbers refer to pages. The Roman
numerals refer to the Appendix.
ACETATE of aluminum, 518.            Aluminum, abundance of, 512.
lead, 499.                           alloys of, 513.
Acetates of copper, 569.                        bronze, 513.
Acetic acid, prep. from wood, 298.              prop. of; 512.
Acid, meanings of the term, 53, 191,  Alums, atomic volumes of, 560.
45!.                                      class of, 529.
Acids, monobasic, bibasic, &c., 191.   Amalgam of ammonium, 437.
Affinity, use of the term, 101.                  gold,.588, 590.
Air, comp. of, 6-10.                             sodium, 94.
not an element, 10.                          tin, 577.
displacing of, 5.               Amalgams, 576.
a mixture, 67.                  Ammonia, analysis of, 77.
physical prop. of, 6.                      comp. of, S1.
presence of, 4.                            electrolysis of; 78.
quantity necessary for complete            occ. of, 83.
combustion,;361.                       physical prop. of, 76, 77.
reaction with nitric oxide, 59, 67.        prep. of, 75.
weight of, 5.                              solubility in water, 76.
Alabaster, 476.                                sources of, 84.
Alcohol, from  the fermentation of             synthesis of; 79.
sugar, 328.                       Ammoniacal liquor of gas-works, S4.
Alcohol-lamp flame, 347.             Ammonia-water precipitates metallic
Alkalies, the caustic, 4(12, 416, 418.                oxides, 439.
definition of; 53.                          prep. of, 84.
Alkalimetry, 419.                                   uses of; S4.
Alkaline-earths, 489.                Ammonium, 82.
Alkaline reaction, 53.                           amalgam, 437.
Alkaloids, organic, absorbed by char-            the 3upposedmetal,437.
coal, 308.                         Ammonium-salts, prop. of, 438.
Allotropism, 139.                                     test for, 594.
Alum, 517.                           Analysis, definition of, 3.
in bread, 399.                         qualitative, 419.
ammonium, 518.                                    a method of seAlum-cake, 517.                                           paratilig  Ca,
Alumina, 514.                                             Sr, and Ba,
Aluminates of the alkalies, 515.                          526.




xliv                          INDEX.
Analysis of organic compounds, 566.  Arsenic acid, hydrates of, 251.
volumetric, 420.                        prep. of, 250.
Animal charcoal, decolorizing power             reduction of, 252.
of, 309.                                      salts of, 252.
Animals and plants, reciprocal ac-              uses of, 252.
tion of, 330.                    Arsenical pyrites, 241.
Annealing, 412.                     Arsenious acid, antidote for, 250.
Anthracite, 293, 294.                             a poison, 250.
conducts heat, 294.                    reduction of, 248.
getting flame from, 360.                solubility of, 249.
Antimoniate of antimony, 273.                     sources of. 246.
Antimoniates, 275.                                two varieties of, 246.
Antimonic acid, 274.                              uses of, 250.
Antimoniuretted hydrogen, 269.      Arseniuretted hydrogen, analysis of,
Antimony, alloys, 268.                                      245.
alloys with zinc, 269.                         decomposed
amorphous, 268.                                by heat, 260.
bloom, 272.                                    prep. of, 243.
distinguished from arse-                       prop. of, 244.
nic, 270.              Asbestos, platinized, 193.
extraction of, 266.      Atom, definition of, 29.
glance, 277.             Atomic heats of compounds, 601.
mirrors, 270.                           elements, 599.
occ. of, 266.            Atomic volume, 201.
prop. of, 267.           Atomic weight, analogy a guide in
the  commercial metal,    determining, 514.
267.                   Atomic weights, defi nition of, 31.
vermilion, 279.                          mutual comparison
Anrozone, 139, 147.                                  of, 462.
distinguished from ozone,                practical use of, 100.
153.                                   table of, 597.
fogs and smokes, 149, 151.               two sets in use, 603.
makes air hard to dry, 150.              value of, 64.
oxidizes water, 152.      Atomicity of the elements, 465.
prep. of, 148.
Antozonides, 153.                   BARIUM, 486.
Aqua regia, 104.                            compounds, 486-488.
Argand burner, 345, 346.                    flame, 489.
Arragonite and calc-spar compared,  Barytes, 487, 488.
469                               Beakers, xxviii.
Arseniates, isomorphous with phos-  Bell-metal, 564.
phates, 252.                     Bessemer steel, 541.
Arsenic, detection of the poison, 252.  Biatomic metals, 73.
determination of its atomic   Bibasic acids, defined, 191.
weight, 262.              Biborate of sodium, 408.
distinguished  from   anti-  Bicarbonate of ammonium, 442.
mony, 270.                             potassium, 416.
eating, 250.                             sodium, 398.
extraction of, 241.         Bichloride of platinum, 593.
greens, 260.                            tin, 584.
mirrors, 260.               Bichromateof potassium, 523.
occ. of, 241., use in
prop. of, 242.                           photolithography,526.
unit-volume weight, 242.    Binoxide of manganese, uses of, 530.




INDEX.                           XIV
Binoxide of tin, 581.              Bromic acid, 122.
Bismuth, extraction of, 281.       Bromide of ammonium, used  in
makes fusible alloys, 282.                photography, 457.
prop. of, 281.                         arsenic, 263.
salts, 283.                             cadmium, used in phoBismuthic acid, 283.                              tog:aphy, 457.
Bisulphate of potassium, 428.                   nitrogen, 123.
Bisulphide of carbon, 363.                         potassium, 425.
and mtric oxide,              silicon, 386.
364.                        silver, 463.
Bisulphide of iron, 549.                        sodium, 394.
Bisulphite of calcium, checks fermen-  Bromides, formation of, 120.
tation, 486.                               of iodine, 133.
Bivalent metals, 73.                            phosphorus, 240.
Black ball, 397.                    Bromine, mother-liquors of saltBlack-lead, 290.                                works, a source of, 393.
Blast furnace, iron, working of, 535-         occ.,of, 119.
538.                                        prop. of, 119.
Bleaching by chloride of lime, 483.           uses of, 120.
chlorine, 112.         Bronze, 564.
ozone, 143.            Bronze-powder, 582.
sulphurous acid, 180.    Bulbs, blowing, v.
Bleaching-powder, 482.              Bunsen's burner, 344, vi.
Blistered steel, 541.               Burettes, xxix.
Bloom, 535.                                 how to fill, xxxi.
Bloomary, 535.                                     read, xxx.
Blowers, viii.
Blowpipe with gas and air, 350.     CADMIUM, extraction of, 509.
Blowpipe, Bunsen's gas-, viii.               makes fusible alloys, 510.
mouth-, 351.                       prop. of, 510.
oxidizing flame of, 352.           unit-volume weight half
oxybydrogen, 46.                     its atomic weight, 510.
reducing flame of, 353.   Casium, 446.
Boiler-scale, 468, 479.             Calcareous waters, 468, 479.
Bone-ash, 214.                      Calcined magnesia, 502.
Bone-black, prep. of, 310.          Calcium-flame, 489.
Bone-phosphate of calcium, 481.             light, 47.
Bones, distillation of, 84, 310.            the metal, 486.
Boracio acid, comp. of, 367.        Calc-spar and arragonite compared,
extraction of, 365.      469.
prop. of, 368.         Calomel, 573.
strong at high        j Candle-flames, 349.
temperatures, 369.    Candles are gas-burners, 347.
Borax, as a blowpipe-test, 409.     Caoutdnouc in sheets, xi.
two forms of, 408.                      stoppers, xi.
uses of, 409.                           tubing, x.
Boron, allotropic states of, 366.   Carbon, allotropic modifications of,
occ. of, 365.                          288.
resemblance to carbon and           atomic weight of, 317.
silicon, 372.                      its atom hypothetical, 318.
Brass, 564.                                compounds with nitrogens
Bread, raising with chemicals, 399.          363.
Breaking up stones, xxxiv.                 dimorphous, 292.
Britannia metal, 268.                      for galvanic batteries, 293.
Bromhydric acid, 120.                      occ. of, 287.




xlvi                           INDEX.
Carbon, vapor-density hypothetical,  Carbonic oxide,..ep. of, 335.
316.                                    prop. of, 338.
Carbonate of ammonium,commercial,                 a  reducing  agent,
441.                                   339.
calcium, dissolves in wa-  Carburetted hydrogen, decomposed
ter containing   by heat, 344.
carbonic acid, 467.  Carmine-lake, 516.
dissolves  in   Cast iron, impurities of, 539.
ammonium-                varieties of, 538.
salts, 468.    Caustic alkalies, strong bases, 418.
occ. of, 467.    Caustic potash, 417.
lead, 499.                     soda, 403.
magnesium, 504.                      advantageous use  of,
potassium, prop. of, 415.              405.
pure, source   Cement copper, 563.
of, 424.     Cement, Roman, 521.
uses of, 416.   Cerium, 561.
and  hydro-  Chalk, 322, 467.
gen, 416.   Chameleon-mineral, 532.
sodium, sources  and   Charcoal absorbs gases, 303.
uses, 396.                     different gases in
and hydrogen, 398.                    different proporCarbonates, 334.                                     tions, 305.
alkaline, impurities of,         causes combination of gases,
419.                             306.
Carbonic acid, 321.                         a disinfectant, 307.
absorbents for, 324.           hot, decomposes steam, 301.
comp. of, 333.                 physical prop. of, 300.
decomposed by                  prep. of, 296, 297.
carbon, 335.              a reducing agent, 301.
potassium, 332.           removes colors, 308.
dissociation of, 331.          satbility of, 302.
formedin combustion,  Chemical changes, 2.
321.                  combination, 32.
fermentation,328.                       and solution
respiration, 328.                       compared,38.
generator, 323.                force, 32.
liquid, 330.                   strength, 418.
obtained from carbo-  Chemistry, subject-matter of, 1.
nates, 322.         Chili-saltpetre, 51.
occ. of, 327.         Chimneys create draughts, 355.
prop. of, 323.                  on fire, how to put out,
solid, 330.                       178.
solubility of, 326.             use of, 354.
sources of, in the air,  Chlorate of potassium, an ingredient
329.                                       of explosive
synthesis of, 333.                           mixtures,
test for, 321.                              435.
Carbonic oxide, comp. of, 341.                           prep. of, 434,
dissociation of, 340.                       435.
economy in burning   Chloraurates, 590.
it, 361.           Chlorhydric acid, 40.
heating value of,339,                analysis of, 92.
361.                               comp. of, 96.
a poison, 338.                       electrolysis of, 94.




r~Drx.                         xlvii
Chlorhydric acid gas, prip. o, o2l, -ith,cf gold, 589, 590.
92.                          iodine, 132.
Chlorhydric acid, prep. i.,, 1,.         iron, 550.
prop. of, 92.                  mercury, 573, 574.
synthesis of, %.               phosphorus, 236.
uses of, 103.                            terchloride,
Chloric acid, 117.                                          236.
Chloride, definition of term, 109.                       quivquice-lrof aluminum, 516.                                 ride, 237.
ammonium, 86.                      platinum, c93.
dissociation            selenium, 201.
of, 239.              tin, 583.
prep.  of,   Chlorimetry, 485.
114.       Chlorine, action on ammonia, 113.
sources of,           atomic weight of, 97.
439.                bleaches, 112.
uses of, 440.         burns in hydrogen, 110.
arsenic, 262.                    combustion in, 110.
boron, 369.                      decomposes water, 111.
bromine, 123.                    disinfects, 113.
calcium, 482.                    explosive mixture with hy
used for dry-              drogen, 108.
ing  gases,            occ. of, 106.
482.                   oxidizing action of, 112.
tube, xxiv.              physical prop. of, 107.
carbon, 362.                     prep. of, 105, 106.
gold, used in photogra-          prep. from bleaching-powphy, 457, 459.                   der, 483.
lead, 498.                       prep. with bichromate of
lime, 482.                          potassium, 525.
magnesium, 503.                  unites with metals, 109.
nitrogen, 114.                   -water, 107.
potassium, 424.         Chloroform, formula of, 319.
silicon, decomposed by   Chloroplatinates, 596.
water, 384.    Chlorous acid, 117.
prep. of, 383.    Chromates, 526.
silver, effect of light on,  Chrome-alum, 524.
456.                    -green, 523.
ocC. of, 460.             -iron ore, 522.
reduction  of,            -yellow, 527.
461.            Chromic acid, 525.
solvents for, 456.   Chromium, occ. of, 522.
sodium, 392.          - Cinnabar, 570.
solubility of,    Citrate of magnesium, 504.
393.          Classification, 598.
sulphur, 197.           Clay, 519.
zinc, 508.              Cleavage, 157.
a disinfectant, 509.  Coal, bituminous, 294,
Chlorides of antimony, 275.             distillation of, 293.
terchloride,       varieties of, 293, 294.
275.        Coal-gas, comp. of, 320.
quinquichlo-          purification of, 473, 546.
ride, 276.   Coal-tar, 312.
chromium, 523.          Cobalt, 556.




xlviii                         INDEX.
Cochineal, tincture, preparation of,   Crystals: four methods of obtaining,
421.                                 161.
Coke, 293, 294.                      Cupellation, 494.
conducts heat, 294.            Cyanhydric acid, 363.
Cold, greatest artificial, 57.       Cyanide of potassium, formation of,
Collodion. 457.                                              426.
Colloids, 255.                                            in blast fur.
Columbium. 585.                                              naces, 427.
Combination by volume, 203.                                uses of, 427.
Combining weights of compounds,                 iron, 555.
71.                      silver, 463.
of organic com-   Cyanogen, 363.
pounds, determination    DACUFRREOTYPE, 456.
Of, 596.       Decolorizing power of charcoal, 308and   unit-vo-      311.
lume weights   Deflagrating-spoon, xxi.
compared,      Deflagration, 302.
205, 206.      Deodorizing by charcoal, 306.
Combustion, definition of, 15.                      chlorine, 113.
ordinary, 342-362.       Destruction of organic matter in cases
spontaneous, 229.          of poisoning, 257.
Compound radicals, 363.              Detection of arsenic, 252-262.
Condensation-ratios, 205.            Developers, a term of photography,
Condenser, 34, 35.                     458.
Cooling flames by good conductors,    Devitrification, 413.
358.                               Dialysis, 254.
Copper, alloys of, 564.                       applied to detection of poit.
extraction of, 562.                     sons, 256.
occ. of, 562.                Dialyzer, 255.
prop. of, 563.               Diamond, 289.
pyrites, 562.                Dichloride of copper, 567.
Copperas, 551.                                    sulphur, 197.
Cork-cutters or borers, xii.         Dichromate of lead, 527.
Corks, xi.                           Didymium, 561.
Correlation of force, 507.           Diffusion of carbonic acid, 325.
Corrosive sublimate, 574.                        gases, 43, 326.
antidote for, 575.                   rapidity of, 325.
Cream oi tartar, 424.                Dimorphous substance, defined, 159.
Crucibles, Beaufay, Hessian, plum-   Dioxide of copper, 564.
bago, xxxiii.             mercury, 571.
porcelain, xxxii.          Direct combination, 404.
how to use, xxxiv.         Disinfectant, charcoal, 306.
material of, 520.                      chloride of zinc, 508.
Cryolite, 391.                                    chlorine, 112, 113.
Crystalline form, a  guaranty  of                 hypochlorite  of  calpurity, 435.                   cium, 482.
structure, 157..               ozone, 144.
Crystallization, the chiel means of              permanganate of potaspurification, 397,                sium, 533.
434.                            sulphurous acid, 181.
by fusion, 156.       Displacement, collection of gases by,
six systems of, 158.     12.
Crystalloids, 254.                   Dissociation, 238.




INDEX.                            xlix
Distillation, the process of, 34.   Ferrous sulphate, 551.
Dolomite. 503.                                    dyeing blackwith, 553.
Drying-apparatus, xxiii.                    sulphide, prep. of, 547 —549.
gases, xxii.                 Filtering, xxii.
substances used for,   Filters, how to fold, xxii.
xxiii.                Filtrate, definition of, xxii.
plastering, 475.             Fire-brick, 520.
Dualistic formulae, 87.              Firedamp, 314.
Fireworks, a green color, 488.
EARTIENWARE, 520.                                purple, 436.
Effervescing liquids, 327.                       red, 488.
Electrolysis of ammonia, 78.                     white, 263.
chlorhydric acid, 94.   Flame, dissection of, 348-350.
marsh-gas, 317.               intensity, 346.
water, 24.                    luminosity of, 343.
Electro-metallurgy, 568.                   oxidizing, 352.
batteries, 507.          put out by good conductors,
silver-plating,            357 -359.
463.                   quantity, 346.
use of graphite,         reducing, 353.
290.                   structure of, 48, 348-350.
Electro-plating, 463, 568, 590.     Flaming fires, 360.
Element, definition of, 3.          Flasks, xxviii.
Elements are bodies incapable of de-   Florence fla ks, xxviii.
composition, 33.                  Fluoboric acid, 371.
Empirical formulae, 86.             Fluorhydric acid, 136.
Epsom salts, 503.                                    action  on silica,
Equations, chemical, 90.                               137.
Erbium, 561.                        Fluoride of boron, 370.
Erdmann's swimmer, xxxi.                       calcium, 136.
Etching glass, 138.                            silicon, its reaction with
Evaporating-dishes, xxxii.                       water, 388.
Explosion of oxygen and hydrogen,              prep. of, 386.
48.                                          volumetriccomp. of, 387.
Explosions in coal-mines, 314.      Fluorine, 135.
hard to get and keep, 136.
FELDSPAR, 375, 519.                           occ. of, 136.
Fermentation, 328.                  Fluosilicates, 389.
Ferric chloride, 550.               Fluosilicic acid, prep. of, 388.
hydrate, used in purifying                  a reagent for potacgas, 546.                                   sium salts, 390.
nitrate, 554.                               a use of, 118.
oxide, 543.                   Flux, used in smelting iron-ores,536.
corrodes organic mat-   Fluxes-alkaline carbonates, 415.
ter, 544.                    borax, 409.
reduction of, 544.        cyanide of potassium, 427.
salts, 547.                          limestone, 536.
sulphate, 554.                       sulphate of potassium  and
Ferricyanide of potassium, 555.              hydrogen, 428.
Ferrocyanide of potassium, 427.            tartar, ignited, 424.
Ferrous chloride, 550.              Formulve, empirical and rational, 86.
nitrate, 554.                         molecular, 208.
oxide, 543.                 Foundry-facings, 291.
salts,changedintoferric,556.   Free gases exist as molecules, 208.




1                               INDEX.
Freezing-apparatus, 77.               Graphite, use in electro-metallurgy,
mixtures, see Refrigerating.             290.
Friction-matches, 211.               Graphitic acid, 292.
Falminating-silver, 455.              Graphon, 292.
Fusible alloys, 510.                  Gray iron, 538.
Furnaces, xxxv.                       Green vitriol, 551.
Grey antimony, 277.
GALENA, 490, 497.                     Group, the alkali, 464.
Galvanic current, 506.                           calcium, 489.
Gas, illuminating, 293, 297, 320.                carbon, 390.
Gas-burners, 346.                                chlorine, 133.
-carbon, 292.                                magnesium, 511,
-generator, self-regulating, xxvi.           nitrogen, 2S5.
-holders, xviii.                              platinum, 596.
-lamps for heating, 344.                     sesquioxide, 560.
Gay-Lussac's burette, xxx.                       sulphur, 202.
German silver, 557, 564.                         tin. 5S5.
Gilding, 590.                         Groups, principles concerning, 135.
Glass of antimony, 277.                       table of, 598.
Glass, colored, 413.                  Gun-metal, 564.
comp. of, 412.                  Gunpowder, comp. of, 432.
etching of, 138.                            manufacture of, 433.
not insoluble, 411.                         products of combustion
varieties of, 412.                            of, 432.
Glass beakers, xxviii.                Gypsum, 476.
cutting and cracking of, i.
flasks, xxviii.                 HARD water, 468, 479.
retorts, xxviii.                Hartshorn, 84.
tubing, bending, drawing and    Heat, relation to  chemical force,
closing, iii.-v.                 418.
heating of, iii.               unit of, 45.
sizes and qualities of, i.   Homologous series, 313.
Glauber's salt, 395.                  Hydrate of aluminum, 514.
solubility of, 395.                calcium, 472.
supersaturated  solu-              copper, 566.
tion of, 396.                   iron, 545.
Glazes-lead, feldspar, salt, 520.                potassium, 416.
Glucinum, 522.                                   sodium, prep. of, 402.
Gold, alloys of, 591.                                  type of bases, 403.
coin, 591.                      Hydraulic cement, 504, 521.
occ. of, 587.                   Hydrocarbons, 312.
prop. of, 588.                  Hydrogen, chemical effect of, 129.
salts, 590.                               derived from water, 22.
separation from silver alloys,            diffusive power of, 43, 44.
464.                                    explosive  mixture with
silver   and                 oxygen, 4S, 49.
copper, 589.             heating power of, 45.
solution of, 589.                         inflammable, 44.
Golden sulphide of antimony, 279.               lightness of, 42.
Graduation, 392.                                peroxide of, 50, 51.
Gramme, definition of, 19, xxxviii.              physical prop. of, 41.
Graphite, 290.                                  precautions in  making,
artificial, 292.                         40.
in cast iron, 538.                     prep. of, 39.




INDEX.                                i
Hydrogen, standard of specific gra-   Iron, cast, impurities of, 539.
vitv for gases, 41.                varieties of, 538.
Hypobromous acid, 122.                      extraction of, 535.
Hypochloric acid, 117.                      galvanized, 506.
Hypochlorite of calcium, 482.               mordant, 555.
Hypochlorous acid, 116.                    occ. of, 534.
iyponitric acid, comp. of; 61.             ores —brown, clay, red hoeprep. of, 62.                 matite, magnetic, spathic,
prop. of, 62.                 specular, 534.
Hypophosplhites, 228.                      puddling of, 539.
Hypophosphorous acid, 227.                 pyrites. 549.
Hyposulphite of sodium, prep. of,          replaces copper, 551.
414.              wrought, 539.
use in pl:o-   Iron retort, xxvi.
tography, 413, 456.   Iron stand for supporting vessels,
Hyposulphites, 196.                     xiii.
Hyposulphurous acid, 196.             Isomerism, 247.
Isomorphism, 200.
ICE, crystals of, 21.
lighter than water, 21.          KAOLIN, 519.
Ignition-tube, definition of, 8.      Kelp, 124.
manner of using, 8, 11.   Kermes mineral, 279.
Illuminating gas, 293.                Kindling-temperat. re, 357.
description of, 320.
manufacture of,      LArES, 516.
297.               Lampblack, 295.
purification of,473,             manufacture of, 298.
546.               Lamp-flames are gas-flames, 347.
Incrustations in boilers, 468, 479.   Lamps for laboratory use, vi-viii.
Indelible ink, 453.                   Lanthanum, 561.
Indestructibility of matter, 356.     Laughing-gas, 57.
Indian fire, 263.                     Lavoisier, his analysis of air, 9.
Ink, 553.                             Lead, action of acids on, 491.
indelible, 453.                                 water on, 492.
Intensity flames, 346.                      classed with the calcium group,
Iodic acid. 130                               500.
Iodide of arsenic. 263.                     crystallization of, 491.
lead, 499.                         extraction of, 490.
nitrogen, 131.                    forms basic salts, 495.
potassium, 425.                    metallic, prop. of, 490.
silver, 463.                      testing for, 497.
sodium, 394.                      tree, 507.
Iodides of phosphorus, 240.                 use of, for water-pipes and cisIodine, extraction of; 124.                   terns, 493.
manufacture of, 132.           Leblanc's process, 397.
occ. of, 123.                                  wastefulness of,400.
prop. of, 124.                 Light, action of, on silver salts, 458.
reaction with starch, 126.                       bichromate of po-,
testing for, 126.                                 tassium, 526.
uses of, 125.                  Light carburetted hydrogen, 313,
lodohydric acid, prep. of, 128.       Lime, air-slaked, 471.
prop. oi, 128, 129.         the cheapest base, 476.
Iodo-starch paper, 127.                     heat evolved in slaking, 472.
Iridium, 596.                               milk or cream of, 472.
2U




11                             INDEX.
Lime, slaked, absorbs carbonic and   Mercury, tendency to form  double
sulphydric acids, 473.                 compounds, 573.
uses of, 476.                          unit-volume weight half its
Limestone, 467.                                atomic weight, 570.
Litharge, 494.                               uses of, 571.
Lithium-flame, 444.                Metal, meaning of the term, 451.
occ. of, 443.              Metantimoniates, 275.
resembles sodium and po-  Metantimonic acid, 274.
tassium, 443.            Metaphosphate of silver, 407.
Litmus-paper. 53.                                   sodium, 407.
tincture, prep. of, 420.    Metaphosphoric acid, 233.
Liver of sulphur, 428.             Metastannic acid, 581.
Luminosity of flames, 343.          Meteoric iron, 534.
Luminous flames, form of, 350.     Metre, definition of, xxxvi.
Lunar caustic, 453.                Metrical system of weights and measures, xxxvi.
MAGNESIA alba, 504.                Mica, 375.
crucibles, 503.           Microcosmic salt, 407.
hydraulic, 502.           Minium, 496.
Magnesite, 503.                    Mohr's burette, xxx.
Magnesium light, 501.              Molecular condition of elementary
the metal, prop. of, 501.            gases, 207.
occ. of, 501.                      hypothesis. 30.
salts,from mother-liquor   Molecule, definition of, 29.
of salt-works, 393.   Molybdate of ammonium, 586.
Manganate of potassium, 531.        Molybdenum, 585.
Manganates, 531.                    Monoatomic metals, 73.
Manganese, occ. of, 527.            Mordant, use of, 519.
Manganese-alum, 529.               AMlordants. 516.
Manganic acid, 531.                 Mortar, 474.
Marsh-gas, 313.                     Mortars, xxxiv.
analysis of, 316.        Mosaic gold, 582.
prep. of. 314.           Multiple proportions, law of, 66.
two vols. yield four vols.  Muriatic acid, 40, 91.
of hydrogen, 315.                   manufacture of, 99.
Marsh's test for arsenic, 259.
Matches, 211, 219.                 NASCENT state, 80.
Matter, indestructibility of, 355.  Neutralization, 54.
Measuring-glasses, xxix.           Nickel, 556.
Mercuric chloride, 574.            Nitrate of ammonium, 441.
an antiseptic, 575.                     prep. of, 53.
forms doublechlo-           calcium, 485.
rine salts, 575.          copper, 569.
iodide, changes of color of,        lead, decomposition of, 62.
575.                                   made, 61, 62.
oxide, 572.                          potassium, occ. of, 429.
as asourceof oxygen,                     prop. of, 430.
572.                        silver, 453.
Mercurous chloride, 573.                           stains of, removed,
oxide, 571.                                426, 427.
Mercury, alloys of, 576.                     sodium, 392.
detection of, 577.        Nitrates of iron, 554.
extraction of, 570.                  mercury, 576.
pneumatic trough, xvi.    Nitre, 429.




INDEX.                              liii
Nitre-paper. 354, 432.               Oxide of zinc, 508.
Nitric acid, analysis of, 63.        Oxides, definition of, 14.
anhydrous, 69, 70.                of antimony, 272.
prep. of, 460.                     teroxide, 272.
combining weight of, 73.                     quadroxide,273.
monohydrated, 70.                            quinquioxide,
prep. of. 52.                                           274.
solrces of, 51.                     arsenic, 245.
uses of, 71.                        bismuth, 282.
Nitric oxide, analysis of, 60.                         teroxlde, 282.
prep. of, 58.                              quinquioxide,
a test for free oxygen,                                283.
59, 60.                          chlorine, series of, 116.
Nitride of boron. 371.                         chromium, 522.
Nitrogen, a constituent of air, 10.            cobalt, 557.
dilutes the oxygen in air,           iron, 542.
18.                                lead, 492.
obtained from air, 16.               manganese, 527.
physical prop. of, 17.               mercury, 571.
prep. by copper, 16.                 molybdenum, 585.
phosphorus, 17.              nickel, 557.
widely diffused, 18.                 nitrogen, series of, 65.
Nitrous acid, 63.                              platinum, 593.
oxide, comp. of, 55, 56.               silver, 455.
physical prop. of, 57.           sulphur, series of, 175.
prep. of, 54.                    tin, 580.
supports combustion,             tungsten, 586.
58.                            uranium, 558.
Nomenclature, prefixes derived from    Oxidizing agents, defined, 71.
Greek and Latin nu-   Oxychloride of bismuth, 284.
merals, 174.         Oxygen, abundance and importance
the prefix hgypo-, 61.            of, 1,5.
the termination ide,           burning charcoal in, 13.
174.                                 iron in, 13.
the terminations ous                   phosphorus in, 13.
and ic, 60.                          sulphur in, 13.
Nordhausen acid, 192.                        burns in hydrogen, 50.
Noriurm, 561.                                a constituent of air, 9.
explosive mixture with hyOCIRE, red, 543.                                drogen, 48, 49.
yellow, 543.                          physical properties of, 12.
Oil-bath, xxvi.                              precautions in making, 11.
Oil of vitriol, manufacture of, 184-         prepared from bichromate of
186.                                                       potassium,
Organic analysis, 566.                                       525.
chemistry, defined, 312.                         bleaching-powOrpiment, 263.                                             der, 484.
Osmium, 596.                                             chlorate of poOxidation by platinum-black, 595.                          tassium. 11.
Oxide of aluminum, 514.                                  peroxide of bacalcium, 470.                                      rium, 487.
magnesium, 502.                                 sulphuric acid,
silicon, 375.                                      196.
silver, a strong bise, 455.         supports combustion, 13.




liV                            INDEX.
Oxyhydrogen blowpipe, 46.           Phosphoric acid, anhydrous, affinity
Ozone, 139.                                          of, for water, 232.
atmospheric, 145.                           glacial, 232.
a disinfecting agent, 144.                  hydrates of, 232.
distinguished fiom antozone,                 strength of, 235.
153.                                      terbasic, 234.
influence on health, 146.    Phosphorous acid, 228.
prep. by electricity, 141.                    hydrated,229,231.
ether, 142.          Phosphorus, allotropism of, 210.
permanganate of po-             burnt under water, 436.
tassium, 142.                 combinations of, 219.
phosphorus, 140.                common, 212.
prop. of, 143.                          comparison of red with
resembles chlorine, 143.                  common. 216.
tests for, 143, 144.                    compounds with hydroOzonides, 147.                                   gen, 220.
inflammability of, 210.
PALLADIUM, 596.                                manufacture of, 214.
Paraffine-bath, xxvi.                          occ. of, 209.
Parchment paper, 255.                          oxides of, 226.
Pearlash, 415.                                 poisonous, 214.
Pearl-white, 284.                              red, 216.
Perchloric acid, 118.                              converted into comPeriodic acid, 131.                                  mon, 216.
Permgnganate of potassium, 532.                    crystallized, 217.
a disinfectant,                oxide of, 226.
533.                        prep. of, 218.
Permanganates, 533.                                used   on  safetyPermanganic acid, 532.                               matches, 218.
Peroxide of barium, prep. of, 487.             shines in the dark, 212.
hydrogen, prep. of, 152.            solutions of, 213.
lead absorbs sulphurous             unit-volume weight of,
acid, 496.                      224.
an oxidizing agent,  Phosphuretted hydrogen, analysis of,
495.                                   223.
prep. of, 495.                          comp. of, 224.
Persulphide of hydrogen, 173.                          liquid, 225.
Petrifactions, calcareous, 468.                        prep. of, 220.
siliceous, 378.                                    from
Pewter, 585.                                               phosphide
Phosphate of magnesium and ammo-                           of calcium,
nium, 504.                                   221.
silver, 406.                               solid, 226.
sodium,ammonium,and   Photography, 456-460.
hydrogen, 407.                    on glass, 457.
Phosphates of calcium, 481.                        paper, 459.
sodium, 405.           Photolithography, 526.
rhombic, 405.   Physical changes, 1.
pyro-, 406.    Pincers, xxxiv.
meta-, 407.     Pink salt, 584.
varieties of, 234.       Pipettes, xxxi.
Phosphide of calcium, 221.          Plants and animals, reciprocal acPhosphorescence, 212.                 tion of, 330.
Phosphoric acid, anhydrous, 231.    Plaster of Paris, 477.




INDEX.                               lV
Plaster-casts, 477.                  QUANTITY flame, 345.
Plastering, comp. of, 475.           Quantivalence, 465.
dampness of, 475.         Quartation, 590.
Platinized asbestos, 193.            Quartz, 375.
Platinum, alloys of, 592.            Quicklime, 470.
black, 595.                           manufacture of, 471.
cleaning, xxii.            Quicksilver, 570.
foil and wire, xxi.        Quinquichloride of antimony, 276.
induces combination, 182,                     phosphorus, 237.
193, 306, 593, 595.                                dissociamelting of, 353.                                 tion of, 238.
metals, 596.               Quinquichlorides of antimony and
occ. of, 591.                phosphorus, chloridizing agents,
prop. of, 592.               277.
sponge, prep. of, 594.     Quinquioxide of antimony, 274.
uses of, 593.                              bismuth, 283.
vessels,  precautions  in    Quinquisulphide of antimony, 280.
using, 592.
working, one method of,   RADICALS, compound, 363.
595.                     Rational formulae, 87, 89.
Plumbate of lead, 497.               Realgar, 263.
Pneumatic troughs, construction and    Red lead, manufacture and uses of,
use of, xiv-xviii.                   496.
Porcelain, 520.                      Reducing agent, defined, 71.
dishes, xxxii.             Reduction of metals by carbonic oxide,
Potash, obtained from ashes of plants,                      339.
415.                                                    charcoal, 300,
Potassium-flame, 423.                                       321.
metallic, 422.             Refrigerating mixture-ice and salt,
occ. of, 414.                                       180.
salts, test for, 594.                             nitric oxide
Product-volume, defined, 207.               and bisulphide of carbon, 57.
Protochloride of copper, 567.                             saltpetre, 432.
tin, 583.                            sulphate of sodium
a reducing a-              and chlorhydric acid, 3t)6.
gent, 583.     Reinsch's test for arsenic, 253.
Protoxide of chromium, salts of, 523.   Replacement, 41, 404, 508.
copper, 505.             Respiration, 329.
lead, 494.               Retort, iron, xxvi.
manganese, 528.                  use of, 34.
tin, 580.                Retorts, xxviii.
Prussian blue, 555, 556.             RIevivification of animal charcoal, 309.
Prussic acid, 363.                   Rhodium, 596.
Puddled steel, 541.                  Rochelle-powders, 399.
Puddling iron, process of, 539.      Rock-crystal, 376.
Pulverizing, xxxv.                   Rouge, jewellers', 543.
Pyrites, 549, 562.                   Rubidium, 446.
Pyritous shales, use of, 550.        Ruby, 514.
Pyroligneous acid, 298.              Rust, of tin, iron, mercury, &c., conPyrophorus, of carbon and lead, 492.   tains something derived from the
Pyrophosphate of silver, 406.          air, 7-9.
sodium, 406.        Rusts are oxides, 14.
Pyrophosphoric acid, 233.            Ruthenium, 596.
Rutile, 577.




Ivi                            INDEX.
SAFETY-Ramlp, 359.                  Siliciuretted hydrogen, 374.
-matches, 218.                Silicon, abundance of, 372.
Saline taste, and substance, 392.           allotropic conditions of. 372.
Salt, decrepitation of, 393.                atomic weight of, 382, 385.
glaze, 394.                            compound  with  hydrogen
manufacture of, 392.                      and chlorine,'86.
solubility of, 393.                    prep. of the three modificasources of, 391.                         tions of, 373.
uses of, 394.                          prop. of, 373, 374.
Saltpetre, 429.                     Silver, atomic weight of, 461.
manufactured from nitrate         coin, 453.
of sodium, 4'30.               extraction of, 449.
not explosive, 431.               melted, absorbs oxygen, 452.
plantations, 42*9.                occ. of, 448.
purification of, 429.             precipitation of, 454.
Sand-bath, xiii.                           quantitative determination of,
Sapphire, 514.                               460.
Saturated solutions, 37.                   separation from lead by crysSchweinfurt green, 570.                      tallization, 491.
Screw compressor, xxv.                     separation from lead by cupelSelenhydric acid, 201.                       lation, 494.
Selenites and seleniates isomorphous       solvents for, 452.
with sulphites and sulphates, 200.   Silvering of mirrors, 577.
Selenium, occ. of, 198.              Slag, finery, 551.
prop. of, 199.                  of iron furnaces, 536.
sources of, 199.           Smalt, 557.
Serpentine, 503.                     Soap, hard, 405.
Sesquicarbonate of ammonium, 442.         soft, 417.
Sesquioxide, the term, 513.          Soda-ash, manufacture of, 394, 397.
of chromium, 523.            -crystals, 397.
salts  of,       grocers', 398.
524.          -powders, 399.
iron, 543.                -water, )27.
manganese, 528.       Sodium-amalgam, 94.
Shot, arsenic added to, 243.                -flale. 402.
Silicates, 379.                             occ. of, 391.
alkaline, soluble, 380.             prep. of, 400.
classes of, 380.                    prop. of, 401.
decomposition of, 381.      Solder, 585.
fused with alkaline car-   Soldering, use of chloride of zinc in,
bonates, 415.                               506, r)509.
in glass, 412.                               sal-ammoniac in,
of aluminum, 519.                             579.
iron, 555.               Soluble glass, 380.
sodium, making of, 410.   Solution and chemical combination
Siliceous petrifactions, 377.                   compared, 38.
Silicic acid, composition of, 382.           defined. 36, 37.
double salts of, 380.             heat either absorbed or set
in natural waters, 377.             free during, 395.
isomeric  modifications           how to apply the solvent,
of, 376.                          37, 594.
occ. of, 375.                     of gases in water, 67.
prop. of, 378.           Spatultb, xxxv.
soluble, 376.            Specific gravity, definition of, 20.




INDEX.                             lvii
Specific heat, definition of, 20.    Sulphate of silver, 464.
heats, value of, 602.                   sodium, manufacture of,
Spectrum-analysis, 444.                                   394.
Spongy platinum, 594.                                  and  hydrogen,
Spontaneous combustion, 229.                              396.
ofcoal,550.             zinc, 509.
Spring clip, xxv.                                    reduction of, 509.
Stalactites, 468.                   Sllphates, 191.
Stalagmites, 468.                              of iron, 551-554.
Stannate of sodium, prep. of, 582.               mercury, 576.
Stannates, 581.                     Sulphide of arsenic, precipitation of,
Stannic acid, 581.                                       258.
Starch-paste, prep. of, 126.                           reduction   of,
Steam, physical prop. of, 21.                             259.
superheated, 21.                         boron, 371.
voulluetric comp. of, 29.                lead, volatility of, 498.
Steel, Bessemer, 541.                           silicon, 390.
blistered, 540.                           silver, 463.
cast, 541.                     Sulphides of ammonium, 442.
puddled, 541.                              antimony, 277.
tempering, 542.                                      tersulphide,
Stereotype-metal, 268.                                       277.
Stone, artificial, 410.                                    quinquisulStove-polish, 290.                                          phide, 280.
Straw rings, xxxiii.                             arsenic, bisulphide, 263.
Strontianite, 488.                                     tersulphide,
Strontium, 486..                                    263.
compounds, 486-488.                          quinquisulflame, 489.                                    phide, 265.
Stucco, 478.                                     calcium, 173.
Suboxide of lead, 493.                           copper, 567.'Substitution, defined, 41.                      iron, 547.,*
Sugar, purification of, 308.                    lead, 497.
Sugar of lead, 499.                              mercurv.y, 573.
Sulphantimoniates, 280.                          phosphorus, 241.
Sulphantimonic acid. 280.                        potassium, 427.
Sulphantimonites, 278, 279.                     sodium, 399.
Sulpharseniates, 265.                            tin, 582.
Sulpharsenites, 264.                 Sulphites, 183.
Sulphate of aluminum, 517.          Sulphur, change of prismatic into
ammonium, 440.                       octahedral, 160.
barium, 487.                       crystallization of, 156,157.
calcium, 476.                      dimorphous, 159.
reduction   of,           extraction of, 155.
481:                    a kindling material, 165.
stufIing for pa-          melting of, 156.
per, 435.               metals burn in, 164.
chromium, 524.                     milk of, 163.
copper, 568.                       occ. of, 154.
lead, 499.                         salts, 264.
magnesium, 503.                         compared with oxy.
potassium, 428.                           gen salts, 265.
and  hydro-             soft, 162.
gen, 423.             solutions of, 163.




Iviii                           INDEX.
Sulphur, uses of, 165.               Symbols, chemical, 41.
Sulphuretted hydrogen, see Sulphy-   Synthesis, definition of, 3.
dric acid.                         Systems of crystallization, 158.
Sulphuric acid, absorbs water, 188.
action on metals, 190.   TABLE of atomic heats of compounds,
organicmat-                             601.
ter, 190.                           gaseous eleanhydrous, 1102.                              ments, 600.
decompo-                            solid  elesitioll                          ments, 599.
of, 195.           atomic weights and symprep. of,               bols of the elements, 597.
193.            for conversion of centigrade
prop. of             into  Fahrenheit degrees,
194.              xxxix-xli.
bibasic, 191.                for conversion of French into
freezing by, 189.              English   measures   and
fuming, 191.                   weigllts, xlii.
how to distil, 187.          of groups of the elements, 598.
how to mix with wa-   Tables of metrical weights and meater, 188.              sures, xxxvii, xxxviii.
hydrates of, 189.      Tantalum, 585.
importance of, 184.    Tar, product of the distillation of
manufacture of, 184-     wood, 298.
186.                 Tartar, 423.
Sulphurous acid, antiseptic, 181.    Tartar-emetic, 273.
bleaches, 180.       Tartrate of lead, 492.
comp. of, 176.       Tellurium, compounds of, 202.
liquid, 180.                    occ. of; 2C01.
oxidation of, 182.              prop. of. 202.
prep. of, 176, 177,   Terbium, 561.
178.               Terbromide of antimony, 277.
prop. of, 178, 179.    Terchloride of antimony, 275.
stops combustion,                  bismuth, 283.
179.                             phosphorus, 236.
Sulphydrate of calcium, 486.         Teriodide of antimony, 277.
Sulphydric acid, 165.                Terminations ous and ic, 558.
analysis of, 168.    Teroxide of antimony, 272.
comp. of; 169.                             both an acid
decomposition  of,                           and a base,
172.                                       273.
inflammable, 170.                bismuth, 282.
in mineral waters,.                    a base, 283.
169.               Tersulphide of antimony, 277-280.
poisonous, 169.                     bismuth, 284.
prep. of, 166.       Test-glasses, xxix.
prop. of, 167.        Test-tubes, xxviii.
pure, 168.           Thallium, 447.
as a reagent, 171.   Thermometers, xxxv.
solution in water,   Thermometer-scales, centigrade and
170.                 Fahrenheit, compared. xxxix-xli.
Supersaturated solution, 396.        Thionic acids, 175.
Sw-immer, Erdmann's, xxxi.           Thorium, 561.
Symbolic formulae, value of, 90.     Tin, alloys of. 585.




INDEX.                            liX
Tin, crystallization of, 578.       Water, distillation of, 33.
extraction of, 578.                    electrolysis of, 23.
occ. of, 578.                          hardness ofl 468, 480.
prop. and uses of, 578, 579.           method of determining hardTin-salt, 583.                               ness of, 480.
Tin-stone, 578.                            occ. of; 33.
Titanic acid, 577.                         removal of gases from, 35.
Titanium, occ. of, 577.                    the common solvent, 37.
resembles tin, 578.              standard of specific gravity,
Tongs, xxxiv.                                                   20.
Tubing, sizes and qualities of, i.                            heat, 20.
Tungstate of sodium, 587.                  symbol of, 32.
Tungstates, 586.                           synthesis of, 25-29.
Tungsten, 586.                             volumetric comp. of, 29.
steel, 586.                Water-bath, xxv.
Tungstic acid, 586.                 Waterglass, 380.
Type-metal, 268.                               uses of, 410.
Types-chlorhydric acid, water, and   Wedgewood mortars, xxxiv.
ammonia, 205.              Weights and measures, metrical,
doctrine of; 88.               xxxvi.
Typical hydrogen compounds, 319.    Weights, table of, xxxviii.
White iron, 538.
UNIT of heat, 45.                         lead, 500.
Unit-volume not absolute, 204.            vitriol, 509.
the bulk of 1 grm. of   White-wash, 474.
hydrogen, 204.'Wire gauze, use of, xiii.
a litre, 205.           Witherite, 488.
weights, 205.           Wolfram, 586.
Univalent metals, 73.               Wood, distillation of, 296.
LTranium, 558.                      Woulfe-bottles, 85.
salts, peculiarity of, 559.    Wrought iron, 539.
impurities of, 540.
VANADIUM, 586.                                    making of, 539.
Vegetable parchment, 255.
Ventilation of wells, 324.          YEAsT-powders, 399.
Verdigris, 569.                     Yellow-metal, 564.
Vermilion, 573.                     Yttrium, 561.
Volume, combination by, 203.
ZINC, alloys with antimony, 269.
WASH-BOTTLE, xxxii.                      extraction of, 504.
Water, analyzed by iron, 23.             granulated, 505.
sodium, 22.            prop. of, 505.
blue color of, 19.                 replaces lead, 507.
densest at 4~, 19.                 white, 508.
dissolves air, 35.           Zirconium, 561.
what it dissolves from air 68.




ORDER-LIST OF CHEMICALS.
The quantities here given are the quantities which one person will
use in performing the numbered experiments of the manual according
to the directions. In ordering chemicals for a class of several students,
a small reduction should be made upon the multiplied quantities. Less
than twenty times the assigned quantities will suffice for twenty
students. Teachers may get some idea of the cost of these chemicals
by referring to the price-lists of the " Druggists' Circular," published
monthly at 36 Beekman St., New York, price 13 cents.
Alcohol......             1 oz.  Chlorhydric (muriatic) acid 14 lbs.
Alum.......    i oz. Cobalt, nitrate of.   8 grains
Ammonia-water (Aqua Am-          Cochineal....   15 grains
monia)......   6 oz. Copper filings....   2 oz.
Ammonium, chloride of. 2 oz.  Copper, oxide of... 30 grains
Antimony, chloride of, a          Copper, sulphate of..     oz.
strong solution of...    oz.  Dutch metal.  several leaves
Antimony, metallic..     oz. Ether.....   a few drops
Antimony,  tartrate  of          Fluor-spar.......          oz.
and potassium (Tartar          Gold-leaf....   isq. inch
emetic).... 40 grains  Graphite......   oz.
Arsenious acid.....        oz.  Gum arabic....         4 oz.
Asbestos..    a few fragments  Gum guaiacum... 15 grains
Barium, chloride of.. 8 grains  Iodine.....  10 grains
Bismuth.10 grains  Iron filings.2 oz.
Bleaching powder....    lb. Iron, red oxide of....      oz.
Bone-black....         3 oz.  Iron, sulphate of...    oz.
Borax.......   oz.  Iron, sulphide of....    oz.
Bromine........           oz. Iron wire (fine piano wire) 1 foot
Calcium, chloride of.. 14 oz.  Lead, acetate of.... 1 oz.
Calcium, sulphate of (Plas-      Lead, peroxide of.  15 grains
ter of Paris)...      lb.  Litharge......   1 oz.
Carbon, bisulphide of.. 1i oz.  Litmus.....    15 grains
Caustic potash (solid).  4 oz.  Logwood, extract of. 15 grains
Caustic soda (" concentrated     Magnesium wire..   4 inches
lye").... oz.  Manganese, black oxide of 1- oz.




ORDER-LIST OF CHEMICALS.                     lxi
Manganese, sulphate of 15 grains  Sodium, acetate of.   80 grains
Mercury, chloride of (cor-         Sodium, carbonate of (anrosive sublimate).. 8 grains              hydrous)..  i oz.
Mercury, metallic... 34 oz.  Sodium, carbonate of (Sal
Mercury, red oxide of.    i oz.              Soda)....        lb.
Nitric acid.....   1 oz.  Sodium, hyposulphite of. i oz.
Nutgalls......   oz.  Sodium, phosphate of.   1 oz.
Oxalic acid......          oz.  Sodium, sulphate of...        lb.
Phosphorus, 1 stick, 34 inches long  Starch.......  4 oz.
Platinum, scrap.   10 grains  Strontium, nitrate of.. 8 grains
Potassium, 1 piece as large as a  Sulphur, flowers of..  6 oz.
pea.                   Sulphur, roll brimstone.    i lb.
Potassium, bichromate of   1i oz.  Sulphuric acid....  2 lbs.
Potassium, chlorate of.. 2 oz.  Tartaric acid...   15 grains
Potassium, iodide of..  Y  oz.  Tin foil......              1 oz.
Potassium, yellow prussiate of i oz.  Tin, granulated...    i oz.
Salt...                 2 oz.  Turmeric paper, 2 or 3 square in.
Saltpetre......    5 oz.  Turpentine, oil of...    oz.
Silver, nitrate of... 30 grains  Zinc, granulated...    3 oz.
Sodium, 4 pieces of the size of a  Zinc, solid metal or scraps   J lb.
pea.                     Zinc, sulphate of.  15 grains.
ORDER-LIST OF UTENSILS.
The following list includes the utensils which one person will need
in performing all the numbered experiments in this manual. The
principal articles of steady consumption are glass tubing, retorts, flasks,
corks, and filter-paper. Many of the other articles, once obtained, last
a long time. It is evidently not necessary to provide all this apparatus
for every member of a large class. Six retorts, as many gallon bottles, six Woulfe-bottles, four soda-water bottles, two measuring-glasses,
two mortars, two pipettes, four wire-gauze lamps, one blast-lamp and
bellows, one blowpipe jet, three or four pieces of platinum foil, two
thermometers, one burette, one pair of scales, and one set of weights,
will suffice, if used with method, for a class of twenty or thirty. The
wholesale druggists keep suitable cheap bottles and porcelain ware.




lxii                ORDER-LIST OF UTENSILS.
Many of the articles may be made by a tinman or gas-fitter.  For the
rest, teachers may consult the priced catalogues of the dealers in philosophical apparatus and chemical ware.
Glass tubing (App. Fig. 1) -          [All these bottles should be of a
2 sticks about 3 feet long of No. 1  very common quality.  The glass
3  4"    "    "        "    "  2  should be tolerably white.]
2  "    "    "         "    "  3  Funnels1  "    "    "         "    i"  4       1 three inches in diameter.
1  "    "    "         i"    "  5       1 two inches in diameter.
3  "    "     "        " 4  "  6  3 Woulfe-bottles of about 8 oz.
2  "          "       "  " "  7    capacity.
1  "    "    "         "    "  8  Glass flasks1 tube about 1 foot long and I1 1 flat-bottomed of 14  pint cainches in inside diameter.            pacity.
1 tube about 1 foot long and 1 inch   1 flat-bottomed of 4 oz. capacin inside diameter.                   ity.
1 tube about 3 feet long and 1 inch   1 flat-bottOmed of 2 oz. capacin inside diameter.                   ity.
[If ignition-tubes can be bought    3 cheap round-bottomed flasks
ready-made, a  dozen  may  be           such as olive oil used to be
bought instead of 1 stick of No. 1      imported in.
and 1 stick of No. 2.]              1 thistle-tube.
Retorts -                           1 Mohr's burette.
1 of 12 oz. capacity without  2 soda-water bottles, stout.
tubulature.                    2 wine-glasses.
1 of 6 oz. capacity, with glass  3 test-tubes.
stopper.                       1 chloride of calcium tube.
1 receiver of 12 oz. capacity, with   1 graduated, measuring-glass  of
tubulature.                         250 c. c. capacity.
Phials and bottles -                1 pipette.
1 glass-stoppered, 4 oz.         1 nest of beakers, of which the
2 narrow-mouthed, gallon.          largest is of about 500 c. c. ca1 wide-mouthed, i gallon.          pacity.
1  "       "       quart.        3 bits of window-glass, 3 inches
1  "       "       pint.           square.
1 narrow-mouthed, pint.          Porcelain dishes4 wide-mouthed,   8 oz.             1 about 21 inches in diameter.
3  "       "       4 oz.            1 about 31 inches in diameter.
2  "       "       2 oz.         1 small iron mortar.




ORDER-LIST OF UTENSILS.                    lxiii
1 Wedgewood  mortar  about 4  1 piece of platinum wire 8 inches
inches in diameter.                long and not thicker than the
1 Bunsen gas-lamp (or spirit-lamp    wire of a very small pin.
where gas is not to be had).     1 stone-ware milk-pan.
1 iron ring-stand.                 1 flower-pot saucer.
1 piece of iron wire-gauze about  1 Hessian crucible of about 8 oz.
4 inches square.                   capacity.
1 piece of' fine copper gauze about  2 common plates, of which one a
2J inches square.                  soup-plate.
2 wire-gauze lamps  (App., Fig.  2 spring-clipsorspring clothes-pins.
VIII.).                          1 thermometer (glass).
1 piece of thin iron gas-pipe about  I pair of small scales  (" Tea"15 inches long and i inch in di-   scales for grocers' use).
ameter.                         1 set of gramme weights, 1 grm
1 iron sand-bath 42 inches in di-   to 250 grms.
ameter.                         Corks - an assortment of three oi
1 sheet-iron trough 8 inches long    four sizes, small for ignition
and 1; wide.                       tubes, medium  for flasks, and
1 glass-blower's lamp or Bunsen's   large for the 8 oz. and 4 oz.
gas blast-lamp.                    wide-mouthed bottles.
1 small double-acting bellows.    Caoutchouc tubing -
1 mouth-blowpipe.                      2 feet of i inch.
1 triangular file.                     3 feet of:a inch.
1 round file.                          2 feet to fit glass tubing No. 6.
1 pair of jewellers' tweezers.         1 foot to fit No. 7.
1 jet for oxyhydrogen  blowpipe,  1 pair of wooden nippers.
p. 46.                          J quire of common filter-paper.
1 piece of platinum-foil 1i inches i sheet of parchment paper.
square.
i piece of finest platinum wire 4
inches long.
THE END.