£)tate CoIIeB^ ot Agriculture at Cornell iHnibcrSitp 3ti)aca. M. $. ILtfirarp Cornell University Library QD 151.H75 A text-book of Inorganic chemistry. 3 1924 003 042 862 Cornell University Library The original of tiiis book is in tine Cornell University Library. There are no known copyright restrictions in the United States on the use of the text. http://www.archive.org/details/cu31924003042862 WORKS OF PROF. A. F. HOLLEMAN, Professor Ordinarius in the University ol Amsterdam, Netherlands, PI HUSHED BY JOHN WILEY & SONS. A Text -book of Inorgranic Chemistry. Issued in English in co-operation with Hermon Charleb Cooper. Fourth English edition, com- pletely revised. 8vo, viii + 505 pp., 79 figures. Cloth, S2 50. A Text°book of Orgranic Chemistry. Edited by A. Jamieson Walker, Ph.D. (Heidel- berg) B.A., Head of the Department of Chemistry, Technical Collntj;o, Derby, England, assisted by Owen E. Mott, Ph.D. (Heidelberg), with the co- operation of the author. Third English Edition, partly rewritten. 8vo, xx + 599 pp. 80 figures. Cloth, S2.50. A companion volume to the preceding, and form- ing with it a comprehensive treatise on pure Chem- istry. A Laboratory Manual of Organic Chemistry for Be- ginners. An Appendix to the Author's Text-book of Organic Chemistry. Edited by A. Jamieson Walker, Ph.D. (Heidelberg), B.A , Head of the Department of Chemistry, Technical College, Derby, England. With the co-operation of the author. First English Edi- tion, revised. 12mo, xiv-f-78 pp. Cloth, SI. 00 net. A TEXT-BOOK OF INORGANIC CHEMISTRY. BY DE. A. F. HOLLEMAN, Professur Orduiarlus in the University of Amsterdam; Emeritus Professai Ordinariits in the University of Groningcn, Netherlands, and Fellow of the Royal Academy of Sciences, Amsterdam. IssDEB IN English in Cooperation with HEEMON CHAELES COOPEE. FUURTH KXtJLISH EDITIOX, (JUMP LET ELY REVISED. TOTAL ISSUE, THIRTEEN THOUSAISTD. NEW YORK: JOHN WILEY & SOT^S. London: CHAPMAN & HALL, Limited. 1912 Copyright, 1901, 1902, 1905, 1908, 1911. BY HERMON C. COOPER. First, Second, and Third Editions entered at Stationers* Hall. THE SCIENTtFIC PRESS ftOOERT DRUMMOND AND COMPJ enOOKLVN. N. V. PREFACE TO THE FOURTH EDITION. The present edition represents a thorough revision of the work by the Dutch author and the American collaborator. It profits by the author's experience with the frequent editions in other languages but is independent in composition. \'ery many of the descriptive portions have been rewritten, notably those on the sulphur oxides and acids, rare gases, nitrogen oxides and acids, sodium hydroxide and carbonate, radio-active elements and platinum, as well as the sections on thermo- chemistry, colloids and the iron-carbon system, while the sub- jects of stability and the reality of molecules and atoms furnish new material. The chapter on metal-ammonia compounds is reprinted as approved by Professor Werner for the third edition. Notwithstanding the appearance of differential formulae in the book, it is believed that a student who is unfamiliar with the calculus should have little difficulty in understanding the meaning and use of such formulae, provided he is willing to take the author's word for the solutions of the equations. Independent students may well be cautioned against regard- ing any text-book as infallible. Even in a book with a world market, such as this one enjoys, undergoing many revisions by the author and by collaborators in other nations, and being frequently reviewed critically by the journals, there will, probably, always be some textual errors and some passages whose lucidity could be impro\-e(l. Readers can therefore render great service bj- reporting all unsatisfactory passages to the publishers. Thanks are herewith expressed to my colleague. Professor H. ]\IoNMOUTH Smith, for constant aid in detecting errors. References in the text to " Org. Chem." refer to the companion volume of this work, Holleman's " Text -book of Organic Chem- istry," translated by Walker and ]Mott. H. C. Cooper. !Syracusb ITxiversity, October, 1911. CONTENTS. Light-face figures refer to pages; heavy-face figures to paragraphs. p.\nK Introduction (1-5) 1 Physic-\l and Chemical Phenomena (6) 3 Chemical Operations (7) 5 The Elejients (8) 7 Oxygen (9, 10) ;» Law of Henry, 11; Oxidation, 12; Analytic and synthetic methods, 13. Hydrogen (11-13) 13 Oxyhydrogen blowpipe, 15; Detonating-gas, 1.5: Reduction, 16. The Coxser\'ation op Matter (14) 16 Water (15-19) 17 Physical properties, 20; Natural water, 20; Composition of water, 22. Compounds and Mixtures (20) 25 Phenomena Accompanting the Formatio.v or Decomposition- op a Compound, 27. Explanation of the Constant Composition op Compounds; Atomr Theory (21-23) 27 Law of constant composition, 27: Atoms, 28; Molecules, 28: Law of multiple proportions, 29; The Atomic Weights of the Elements, 29; Chemical Symbols and Formul.e, 30. Stoichiometric.al Calculations (24) 31 Chlorine (25-35) 33 Catalytic action, 34; Hydrogen chloride, 37; .Xcids, bases and salts, 39; Composition of hydrochloric acid, 41; Law of Gay-Lussac, 43; Avogadro's hypothesis, 44; Molecular weights, 45; Rules FOR Determining Molecular and Atomic Weights, 47; Kinetic theory, 47; General gas equation, 48; The Reality of Molecules and Atoms and Their Absolute Weight, 49. Ozone (36, 37) 50 Formula of, 52; Allotropism, 53. Hydrogen Peroxide (38, 39) 53 Status nascendi, 54. Molecular Weight prom the Measurement of the Depression of the Freezing-point and Elevation op the Boiling-point (40-43) . .57 Semi-permeable membranes, 57; Osmotic iircssure, 58; Ppepfer's experiments, 50; Isotonic solutions, 62; Formula of hydrogen peroxide, 65. vi CONTENTS. PAGE Bromine (44, 45) 65 Hydrogen bromide, 67. Iodine (46-48) 69 Hydrogen iodide, 71. Dissociation (49-51) 72 Reversible reactions, 73; Equilibrium, 73; Reaction velocity, 75; Law of chemical mass action, 75; Unimoleoular and bimolecular reactions, 76. Fluonne (52, 53) 79 Hydrogen fluoride, 81. Compounds of the halogens (54-62): with each other, 83; with oxygen, 83. Nomenclature (63), 92. Summary of the halogen group (64), 93. Electrolytic Dissociation (65, 66) 94 Ionic equUibrium, 98; Strength of acids and bases, 99; Hydrolysis, 101. Sulphur (67-93) 102 AUotropic modifications, 104; The Transition Point, 106; "Sta- ble," "Metastable," and "Labile," 108; The Phase Rule OF GiBBS, 109; Hydrogen sulphide, 115; Solubility product, 119; Hydi'ogen persulphide, 120; Compounds of sulphur with the hal- ogens, 121; Valence, 122; Compounds of sulphur with oxygen, 124; Oxygen acids of sulphur, 131; Volumetric analysis, 146. Selenium and Tellurium (94, 95) 148 Selenium, 148; Tellurium, 150. Summary of the oxygen group, (96) 151. Thermochemistry (97-104) 152 Law of Hess, 153: Chemical Affinity, 156; The Displacement OF Equilibrium, 160; Passive Resistances, 161. Nitrogen (105-130) 162 The atmosphere, 165; Argon, helium and companion elements, 170; Compounds of nitrogen and hydrogen, 174; Compounds with the halogens, 179; Hyth-oxylamine, 180; Compounds with oxj'gcn, 181; Oxygen acids, 187; Derivatives of the nitrogen acids, 195; Other nitrogen compounds, 198. Phosphorus (131-154) 199 Hydrogen compounds, 204; Halogen compounds, 209; Oxygen com- pounds, 211; Acids, 212. Arsenic (155-164) 221 Hydrogen arsenide, 223; Halogen compounds, 226; Oxygen com- pounds, 226; Oxy-acids, 228; Sulphur compounds, 229; Sulpho- salts, 230. Antimonij (165-169) 231 Hydrogen antimonide, 232 ; Halogen compounds, 233; Oxygen com- pounds, 234; Sulphur compounds, 236. COXTKXTS. ^i; PACK Bismuth (17(K174) 23f; Suniinarv of the nitrogen group (175), 1230. Carbon ll7()-18!n 241 Allotropic forms, 241; Molecular and atomic weight, 24t); Com- pounds with hydrofioii, 24.S: Compounds with oxygen, 249; ( Ither carbon compounds, '2!\ry. The flame, 2r;(i. Silicon (l'JO-r,l()) ... 2(11 Hydrogen silicide, 2ti2: Halogen compounds, 263; Oxygen com- pounds, 2(i5; Silicic acids, 266; Colloids, 26S. Gennnniuni (197) . . 273 Tin (198-202) 274 Stannous compounds, 276: Stannic compounds, 279. Lia-I (203-206) . 2sl Oxides, 2S3; Halogen compounds, 2S5; Other lead salts, 286; Sum- mary of the carbon group (207), 2S7. Methods of Deteemixing Atomic \Veight.3 (208-212) .• 2ss Law of DuLDXO and Petit, 289; Law of Neimann, 291; Law ol MiTSCHERLICH, 292: ExPERlME.NT.VL DETERMINATION OF EqU1V.4- LENT Weights, 293. The Periodic Sv.'-tlm ' of the group, 458. Manganese (300, 301) 458 Manganic acid and permanganic acid, 460. iron (302-308) 463 Iron-carbon system, 466; Ferrous compounds, 473; Ferric com- pounds, 474. Cobalt and Nickel (309-312) 477 Cobalt, 477; Xiclcel, 479. Platinum Metals (313-316) 481 Ruthenium, 482; Osmium, 483; Rhodium, 483; Iridium, 484; Palladium, 484; Platinum, 485. Metal-ammonia Compounds. W'erxer's Extensions op the Notion or Valence (317-318) 486 INORGANIC CHEMISTRY. INTRODUCTION. I. Chemistry is a branch of the natural sciences, — ^the sciences which deal with the things on the earth and in the outside universe. The knowledge of these things is obtained by observation with our senses, this being the onlj' means we possess. It is well to understand, therefore, that we know not the things themselves, but simply the impressions which they make upon our sense-organs. ^^^len we see an object, we perceive, in reality, only the effect on our retina; if we feel the object, it is not the body itself but the excitement of the sensory nerves of touch in our fingers that we are made aware of. Hence it may be fairly asked whether the objects of which we are cognizant are reallj' just as we perceive them, or whether they even exist at all outside of our person. The natural sciences leave this problem out of consideration — its solution is the task of speculative philosophy. In reahty they are not concerned with the objects, which in themselves we cannot know, but with the study of the sensations that we receive. The sensations take the place of the objects, and we regard them as such. 2. The Scientific Investigation of Things. — AMiat is to be understood by the tenn? In the first place, a most accurate description of the objects. From a study of this it is found that manj' objects resemble each other to a greater or less degree, and it is therefore possible to make a classification, i.e. an arrangement of like objects into groups and a separation of the various groups from each other. By the descriptive method we are finally able to divide the natural sciences into Zoology, Botany, Mineraloni/ and Astronomy. 3. In the second place, scientific investigation includes the 2 INORGANIC CHEMISTRY, [§§ 3- study of the relations which the objects bear to each other; in other words, the study of phenomena. The heavenly bodies move towards each other; water turns to ice on cooUng; wood burns when heated. It is the task of the natural sciences to accurately observe and describe such phenomena, i.e. to ascertain in what way the heavenly bodies change their relative positions, what conditions affect the freezing of water, what becomes of the burn- ing wood, under what conditions it can burn, etc. The description of the phenomena leads to a different division of the natural sciences than the description of objects, viz., a divi- sion into Physics, Chemistry and Biology, the latter being the study of vital processes, and including Physiology, Pathology and Thera- peutics. 4. The human mind, in piirsuing the scientific study of nature, does not feel contented with the accurate description of objects and phenomena; it seeks also for an explanation of the latter. The various attempts at explanations constitute the most im- portant part of science. When, for instance, we see that a ray of light in passing through a piece of Iceland spar is split up into two other rays of different properties, we strive to account for the phenomenon. When copper is heated in the air, it turns into a black powder; the question again arises, why this thing is so. In searching for an explanation of the phenomena we thus endeavor to penetrate deeper into the essence of things than is possible by direct observation. Although the phenomena themselves are found to be unchangeable, our explanation of them may be modi- fied as our knowledge increases. The transformation of copper into a black powder on heating in the air was formerly explained by the supposition that something left the metal; subsequently, when the phenomenon was better understood, by assuming that the copper takes up something from the air. Scientific investigation pursues in general, then, the following course: A phenomenon is observed and studied as carefully as possible. Thereupon an explanation of it is sought. A hypoth- esis is set up. From this conclusions can be formed, some of which can be tested by experiment. If the latter really leads to the expected result, the hypothesis gains in probability. If it is subsequently found to explain and Imk together a whole series of phenomena, it becomes a theory. 6.] PHYSICAL AND CHEMICAL PHENOMENA. 3 The nineteenth century was an era of great prosperity for scientific inquiry. For numerous phenomena explanations have been found which possess a great degree of probabihty. Still it cannot be denied that the present theories penetrate only a little into the real essence of things, and the investigator very soon stumbles upon questions whose explanation does not at present e\'en seem to be a possibility. The chemical process that goes on when copper — to retain our former example — is heated in the air is weU known. However, the deeper question, why the action talces place just so and not otherwise, or why the resulting powder is black, still awaits a satisfactory answer. 5. We observed in the preceding paragraph that the natural phenomena are found to be unchangeable. The movement of the planets, for example, stiU takes place in the same manner as in the times of the Ptolemies; whenever water turns to ice the same increase of voliune is to be observed; the crystal form of common salt, whenever and wherever examined, is invariably the same; from the burning of wood the same products are always obtained; the microscopic structure of the leaves of one and the same plant is never found to vary. This general principle Inds its expression in the phrase, constancy of natural phenomena. Every one is con- vinced of its truth, and it is tacitly accepted as the basis of every natural scientific investigation. If, for example, one has measured the angles which the faces of a soda crystal form with each other, he considers it certain that all soda crystals must show the same angles, at whatever time or place they may be measured. If it has once been determined that pure alcohol boils at 78° under normal pressure, it is forthwith assumed that this must be the case with all alcohol, no matter how it may be obtained or when and where it n)ay be tested. PHYSICAL AND CHEATICAL PHENOMENA. 6. It was stated above (§3) that the description of phenomena leads to a division of the natural sciences into Physics, Chemistry and the study of vital processes (Biology). In defining the province of Chemistry Biology may be left out of consideration; however, it is desirable to compare the field of Chemistry with that of Physics. In general it may be said that Physics deals with the 4 INORGANIC CHEMISTRY. [§§ 6- temporary, Chemistry with the lasting, changes of matter. By matter or substance we understand the objects without reference to their form. Iron, marble, sand and glass are kinds of matter, or substances, independent of their external shape. A couple of illustrations may make this conception of temporary and lasting changes clear. (o) A platinum wire glows when held in a colorless gas-flame. On removal it cools off and no change is ^-isible. This is a physical phenomenon; the change, the glowing, is of a temporary sort. 80 soon as the cause of the change is removed, the wire resumes its original condition. When some magnesium wire is held in the flame, it biuns with the emission of a brilliant light and turns into a white powder, which is wholly different from the substance magnesium. Here a lasting change has occurred; we have to do with a chemical phenomenon. (6) Again, wc may take two white crystallized substances, naph- thalene and cane-sugar, and heat each separately in a retort with receiver. The naphthalene at first melts; on continued heating it begins to boil, then distils o^-er and condenses in the receiver. The distilled naphthalene resembles the undistilled in e^-ery respect. The substance has, as a result of heating, undergone physical changes — melting, change to vapor and, finally, return to the solid state. The cane-sugar behaves differently. Here also a melting is observed at first, but soon the sugar turns darker; a brownish liquid distils o\'er; a peculiar odor is noticeable and at last there remains in the retort a charred, porous mass. The cane-sugar suffers a lasting change on being heated. In this case we have a chemical change, (c) As a third and last example we may consider the behavior of a metallic wire on the one hand and that of acidulated water on the other, when an electric current passes through them. The wire displays other properties so long as the current is on. If the latter ceases, the wire returns to its original condition. This is a physical action. In the acidulated water, however, the current induces an evolution of gas, and this gas arising from the water has properties entirely different from those of the water. A lasting change in the substance has occurred; a chemical action has taken place. A sharp distinction between physical and chemical phenomena is often — as will be seen later — very difficult to make. 7.J CHEMICAL OPERATIONS. 5 CHEMICAL OPERATIONS. 7. In order to avoid repetitions it seems advisable at this point to describe briefly some of the commonest chemical operations. Solution. — When sugar, salt or saltpetre, for example, is put into water, the solid substance disappears and its taste is taken on by the water. The substance has dissolved in the water. There is a definite limit to the solubility of each of these, for, if the tem- perature is kept constant and more of the substance is gradually added, a point is finally reached when the water will take up no more. The solution is then saturated. The solubility of most solids increases with the temperature. Moreover it is very differ- ent with different substances, varying all the way from solubility in all proportions to imperceptible solubility. Thus cane-sugar is dis- solved in large quantit}- by water, while sand is practically insolu- ble in it. Liquids can be either miscible in aU proportions (water and alcohol) or only partially soluble in each other. When, for instance, water is shaken A^ith a sufficient quantity of ether and allowed to stand, two liquid layers are formed ; the water has dissoh-ed some ether and the ether some water. In most cases the solubility of liquids in each other also increases with the tem- perature. In the case of gases solubility decreases with rising temperature. Separation of a Solid and a Liquid. — This may be accomplished by filtration. A funnel is lined on the interior with "filter-paper " and the mixture poured upon it. The solid is retained on the paper while the liquid passes through. Decantation is a less com- plete method of separation, since more liquid remains with the solid b}' this method than by filtration. However, it is evident that neither method affords a really complete separation. This can only be accomplished by washing, i.e. b\' replacing the por- tion of the liquid wliich remains between the solid particles by another liquid. If the liquid of the mixture be a salt solution, pure water is ^•ery effecti%-e. It is obvious that by repeating the washing several times the salt solution can be wholly remo\-ed. Suppose that 1 c.c. salt solution remains between the particles of the solid and that 9 c.c. water is then added. The solution is thus rediiced to one-tenth of its original concentration. If 1 c.c. of this dilute solution again remains with the solid and another INORGANIC CHEMISTRY. §§7- 9 c.c. water is added, the concentration is then 10"^, or one- hundredth of the original concentration; after six such operations it would be only 10~^, or one millionth of the original, so that the separation is practically complete. Crystallization. — If a solution is saturated in the warm and is then allowed to cool, the dissolved substance frequently separates out in the crystallized state. Advantage is often taken of this for purifjdng crystaUizable substances. Distillation (Fig. 1). — ^This operation is frequently made use of in working with liquids. The liquid is placed in a flask or a retort and heated to boiling. The escaping vapor is cooled to Fig. 1. — Distillation. a liquid in a condenser. The latter consists of a sufficiently wide tube encased in a jacket, through which water flows to keep the inner tube cold. The condensed liquid is collected in the receiver. It is readily seen how volatile substances can be separated from non-volatile ones by distillation, e.g. water from salt, since the former distil over and the latter remain in the distilling-flask. However, Uquids of different volatility can also be separated in this manner. Take, for example, a mixture of alcohol and water. The more volatile constituent, alcohol, passes over for the most part in the early stage of the operation; towards the end the less volatile, water. If the two distillates are collected separately, an approximate separation results. A few repetitions of this so-called fractional distillation bring about a practically complete separation in many cases. «.] THE ELEMENTS. 7 Sublimation. — Certain solids, e.g. camplior, when heated (at ordinary pressure), turn to vapor without melting. If this vapor comes in contact with a cold surface, the substance is deposited in the solid, crystallized state. It is evident that we have here another method of separating some substances. THE ELEMENTS. 8. WTien a substance (§ 6) is subjected to various influences, such as heat, electricity, or light, or is brought in contact with other substances, it is verj- often split up into two or more dis- similar components. As an example let us take gunpowder. Water is added and the whole is stirred well and gently warmed; after a while it is filtered, and that which remains on the filter is found to be no longer gunpowder, for it is unexplosive. On e\'apo- rating the filtrate a white crystalline substance, saltpetre, remains. The undissolved part is dried and then shaken with another sol- vent, carbon disulphide. After a time the mixture is filtered, as before, and there is left on the filter a black mass, consisting of charcoal powder. The carbon disulphide of the filtrate e^■aporates and leaves yellow crystals of sulphur. Thus we see that, Ijy suc- cessive treatment with water and carbon disulphide, gunpowder can be separated into tlircc substances, ^-iz. carbon, sulphur and saltpetre. The two former are incapable, even wlaen subjected to all the agencies at our command, of division into different com- ponents. Not so with saltpetre, for when the latter is heated strongly a gas is given off in which a glowing wooden splinter is at once ignited. When the evolution of gas ceases, a substance remains which gives off red fumes on treatment with sulphuric acid, something that saltpetre does not do. Saltpetre can evidently be broken up still farther l)y heating. If we subject all sorts of substances to a successive treatment with reagents of the most different kinds, we finally discover cer- tain ones that cannot be resolved into simpler substances by our present means. Such substances are called elements. Although the nimiber of substances, according to § 6, may be considered as infinitely great, experience has taught that the number of elements is small. There are aljout eighty. As our methods of examination improve, it may quite possibly 8 INORGANIC CHEMISTRY. [§ 8 be found that the substances which the chemist of to-clay regaivls as elements have no right to the name. Therefore, when we use the word "element," it is to be regarded as a relative term, dependent on the extent of our knowledge and the means at our command. In the history of chemistry some cases are to be found where sub- stances, once belie^•ed to be elements, were subsequently decom- posed. The exact number of elements cannot be definitely stated, be- cause, on the one hand, not all the substances that possibly exist may be \\-ithin our reach,* and, on the other hand, it is doubtful whether certain substances now regarded as elements cannot be divided by means already known. On the inside of the back cover will be found a list of the elements now known. As may be seen from tliis list, the metals are included in the elements. Together with them we find a number of other sub- stances, as oxygen, sulphur, phosphorus, etc., that are classed under the term non-metals, or metalloids. To the latter class belong many very important substances, e.g., oxygen, an element that combines with almost all others, causing what is called combustion. Oxygen is present in a large amount in the air. Another non- metal is carbon, which is present in all organized substances, and is therefore a constituent of every animal and plant. Sulphur, which burns with a blue flame, giving off a pungent odor, and chlorine, a greenish-yellow gas of very disagreeable odor, which combines readily with most metals, are also non-metals. The elements occur in very unequal proportions in the part of the earth accessible to us. Oxygen, which occurs in air, in water, and in the solid part of the earth's crust, is very preponderant, composing approximately 50% of these portions of the earth which have been investigated. The elements silicon, aluminium, iron, calcium, carbon, magnesimii, sodium, potassium, and hydrogen, together with oxygen, make up 9990 oi the earth's crust. There remains, therefore, only 1% for all the other elements. Some of these are quite common, e.g., lithium, but they almost always * Of the interior of the earth only a very small part is known. If we think of the earth as about the size of an orange, the deepest mine-shafts would not even penetrate the thin yellow exterior layer of the orange skin . §9.] oxy(;i:X. 9 occur in \ery small quantities. {~)lheis, like niobium and tantalum, are found in relatively very small amounts and in isolated places, With the aid of spectroscopy (§§ 263-265), it has been ascer-- tained that the heavenly bodies contain most of the elements found in our earth, and also some others. OXYGEN. 9. Under ordinarj' conditions of temperature and pressure, oxygen is a colorless and odorless gas, whose most noticeable property is its ability to set glowing substances on fire with the evolution of much light and heat. A glowing splinter of wood, for example, when introduced into an atmosphere of oxygen, begins at once to burn brightly. This action is ordinarily used as a characteristic test for the identification of oxygen. This gas can be obtained in various ways. There are many substances which are known to evolve oxygen on heating. (1) Mercuric oxide, when heated strongly in a retort (Fig. 2\ yields oxygen, which can be collected by means of a delivery-tube Fio. 2. — Preparation of Oxygen from Potassium Chlor.^te. opening under the mouth of a cylindrical receiver filled with water. The inside of the retort becomes covered with drops of mercury. (2) The same apparatus can be used in making oxygen from potassium chlorate (chlorate of potash), as well as from pntuxsiuin nitrate (saltpetre), potasswim permanganate, and many other sub- stances. The preparation of oxygen by heating potassium chlorate is a metho 1 frequently used in the laboratory. 10 INORGAXIC CHEMISTRY [§9- Some substances give off oxygen when heated together with others, as in the following cases : (3) Potassium dichromate or manganese dioxide, when heated with sulphuric acid ; (4) Zinc oxide, when heated in a current of chlorine. The atmospheric air consists principally of oxygen and nitrogen. The following method for separating these gases was employed by Lavoisier in 1774. He introduced some mercury into a retort A (Fig. 3) with a long, doubly-bent neck that opened under a bell- jar P filled with air and resting in a dish R of mercury. He then Fig. 3. — Absorption of Oxygen by Mercury. heated the retort steadily for several days, keeping the mercury almost boiling. As a result, a part of the air in P disappeared, and the gas remaining was found to possess other properties than air — it was nitrogen. At the same time the mercury had been partially transformed into a red powder, mercuric oxide. On heating the latter more strongly oxygen was obtained. Oxygen in now prepared from liquid air (c/. § 109) . The physical properties of oxygen, besides those already men- tioned, are as follows: Its specific gravity, assuming the density of air to be 1, is 1.10535. A liter of oxygen at 0° and 760 mm. Hg pressure weighs 1.4290 g. Oxygen can be liquefied; the difficulties in obtaining it on a large scale in the liquid state have now been completely overcome. Apparatuses for lique- fying oxygen have been constructed by Hampson and by LiNDE, a description of which is to be found in text-books of physics. The critical temperatm-e of oxygen is —118°, and its critical pressure .10 atmospheres. Liquid oxygen 10] OXYCFX. 11 luis a spocilic iriavity of 1.124 (based on water) and a boiling-point of - 182.95° at 745.0 mm. pressure. Its color is ligVit blue. It can be preserved for some time at ordinary pres- sure, with the aid of a so-called vacuum-flask (Fig. 4.) The latter is a vessel enclosed in an air-tight jacket, the space be- tween the walls being evacuated. 100 1. water at 0° dissolves 4.89 1. oxygen. The gas is also somewhat soluble in alcohol and in molten silver. When the silver solidifies, the oxygen — a volume about ten times that of the metal — suddenly escapes from solution, caus- ing peculiar elevations on the surface of the silver ("spitting" of silver). We remarked above (§7) that the solubility of gases in liquids diminishes \vith increasing tem- perature. A very remarkable law expresses the relation that exists between the solubility of a i;as and its pressure, namely, the solubility is propor- tioiwl to the pressure. This is the law of Henry. PiQ 4 Vacuum- Thus, when the pressure becomes a-fold, the solu- FL.iSK. bility also becomes «-fold. As the mass of a gas which is present in a certain volume is hkewise proportional to tl\e pressure, the law of Hexry can also be expressed thus: The volume of a gas dissolving in a certain quantity of a liquid is independent of the pressure. This law is rigid when the solubility of the gas is small; when the solubility is large, for instance 100 volumes in 1 of the liquid, its de-s-iations are considerable. Still another formulation of this law is of value in understanding certain of its applications: The concentrations of the dissolved and undissolved portions of a gas bear a constant ratio to each other. By "concentration" is meant the quantity of the gas in grams per unit volume (cvibic centimeter). 10. Among the chemical properties of oxygen the most promi- nent is its vigorous support of combustion. The following are interesting examples : Charcoal glows in air only moderately and without much evolu- tion of light. In oxygen, however, it burns with a bright glow. Sulphur, which burns in air with only a small flame, burns in oxva;en with an intense blue light. Phosphorus burns in oxysrcn 12 IXORGAXIC CHEMISTRY [§§ 10- with a blinding white light. A steel watch-spring that has been heated to redness at one end and put into oxygen, burns with scintillation. Zinc also burns in it with a dazzling light. In all these and analogous cases the oxygen, as well as the burning mate- rial, disappears during the combustion, while new substances are formed. A lasting change therefore takes place and we have to do with a chemical process. The product of burning charcoal is found to be a gas that makes lime-water cloudy and is unable to support combustion; it is called carbonic acid gas. Sulphur also yields a gas; it has a pungent odor and is called sulphur dioxide. Phosphorus produces a white flocculent powder, phosphorus pent- oxide. When iron burns, a black cindery powder is formed, called "hammer-scale,'' because it composes the sparks that fly from the anvil. The question now arises as to what really occurs in the above cases. In the first place, it has been found that the weight of the product of combustion is greater than that of the substance burned. The increase in weight of the substance during burning can in many cases be easily demonstrated. For instance^ a horseshoe magnet that has been dipped in iron filings may be hung on the lower side of a scale-pan and balanced by weights put in the other pan. The iron filings may be burned by passing a non-luminous flame under them a few times. On cooling, the scale-pan attached to the magnet sinks. In a similar way one may demonstrate the increase of weight in the burning of copper. In order to prove the increase of weight in a case where only gaseous products are formed, a candle may be burned and the combustion products, carbon dioxide and water vapor, collected by letting them pass over unslaked lime, with which both unite. Closer investigation has revealed the fact that the increased weight is due to the presence of oxj'gen, as well as the burned substance, in all combustion products. The latter are compounds of these substances with oxygen. The participation of oxygen in the burning of zinc, for example, may be proved by heating the combustion product, zinc white, in a tube and leading over it chlorine gas, whereby oxygen is driven off. The compounds of oxygen are called oxides, and the act of this combination is known as oxidation. When substances burn in the air, it is only the oxygen which aombines with them. Nevertheless, the nitrogen of the air is heated 11-] HYDROGKX. 13 and thus takes a inirt of the heat evolved in the combustion. Therefore the temperature of a burning object cannot rise so high as in pure oxygen, and, since the emission of light increases very rapidly as the temperature rises, combustions in oxygen are for this reason much brighter than in air. There are two general methods of ascertaining what elements exist in a compound. According to the one method the compound is decomposed and the elements composing it thereupon deter- mined. This is the analytic method. According to the other, the synthetic method, the composition is found by combining different elements to form new substances. In the above-described experiment (§ 8) of Lavoisier the composition of the red powder is learned by decomposing it at high temperature, whereupon it separates into only mercury and oxygen. Inversely it was possible to obtain the red powder by heating pure oxygen and pure mercury together at a lower temperature. The former is an example of analysis, the latter of synthesis. HYDROGEN. II. Hydrogen is a colorless and odorless gas that is rarely found on the earth in the free state. The gases of some volcanoes contain it and it can also result from processes of decay. In combination with other elements, however, hydrogen is very widely distributed and occurs in very large amounts (§ 8). Hydrogen can be prepared in various ways. In the first place, hydrogen compounds can be broken up. (1) Water is decomposed by the electric current, hydrogen being evolved at the negative pole (cathode). The ordinary methods of preparing hydrogen depend on the indirect decomposition of hydrogen compounds, i.e. their reaction with other substances. The following are examples of this sort: (2) The action of zinc on dilute sulphuric acid (§ 89). This is the ordinary method. For the preparation of hydrogen in the laboratory the apparatus shown in Fig. 5 is often used. A contains granulated zinc (or iron nails) and B dilute hydrochloric acid or sulphuric acid. AMien the cock C is opened the acid flows through D to the metal and the evolution of hy- drogen commences at once. The cock being closed again, the gas still 14 IXORGAXir ( HEMISTRY. TIus is facilitated by continues to come off and forces back the acid, changing the relative levels of A and B. (3) The action of zinc or aluminium fil- ings on caustic potash or slaked lime. (4) The action of sodium or potassium on water or alcohol. (5) Magnesium powder, when boiled with water, also evolves hydrogen, especially when some chloride of magne- '*'• sium is dissolved in the water, because such a solution dissolves the magnesium oxide which forms on the surface of the metal. Likewise, red-hot iron decomposes water with the liberation of hydrogen (compare § 305). 12. The -physical 'properties of hydrogen are these: It is the- lightest of all known substances, its specific gravity (air= 1) amounting to only 0.06949. One hter of hydrogen at 0° and 760 mm. Hg. pressure weighs 0.0899 g. Its Ughtness renders it useful for inflating balloons. It is very hard to liquefy, because its critical temperature lies only 30-32° above the absolute zero- (—273°). On the other hand, the critical pressure is only 15- atmospheres. Liquid hydrogen is colorless. It boils at - 252.5°. Its specific gravity, with reference to water, is only 0.07 at it& boiling-point and 0.086 at its freezing-point, being therefore- considerably less than that of all other known hquids. Dewar. further succeeded in bringing hydrogen to the solid state by allowing the liquid to evaporate quickly at 30-40 mm. pressure- The melting-point of solid hydrogen is about 16° (absolute tem- perature). The heat of evaporation of liquid hydrogen is very high, being 200 cal.; for this reason a flask containing liquid hydrogen soon becomes covered with a layer of liquid air, •which drops down and soon partially solidifies. Hydrogen is slightly soluble in water, 100 I. water dissolving; 2.15 1. of the gas at 0°. Alcohol takes up somewhat more. 13. Chemical Properties. — Hydrogen does not unite with as large a number of elements as oxygen. At a higher temperature 13.] HrDROGKX. I5, it displays a strong tendency to unite with oxygen, burning with an almost colorless and a very hot flame to form water. This liroperty serves for the identification of hydrogen gas. When a current of hydrogen is directed upon very finely divided platinum (spongy platinum or platinum black, § 316), the hydrogen is ignited (§25). The Ingh temperature of the hj-drogen flame is made use of in fusing platinum, quaitz, etc. Such a flame is kno\\Ti as an oxi,hydrogen flame. An apparatus (oxijhydrogen blovpipe) like that represented in Fig. 6 is required for producing it. The hydrogen enters at IF and passes out at a, where it is lit. O.xygen is blown into the fiaine at S. Thus the gases do not mix till they reach the flame, and the possibilitj- of an explosion is avoided. Fig. G. — O.XYHYDROGEx Blowpipe. A mixture of hydrogen and oxygen, especially in the proportion of 2 vols. H and 1 vol. ( ) (detonating-gas), when ignited, turns instantaneously to steam; in otlier words, it explodes. This ex- periment can, however, be perfonned harmlessly by using a wiile- mouthed cylinder of not too great dimensions. A loud report is heard in this case, because the steam at the moment of its fonna- tion occupies a much larger ^'olume at the high temperature of the combustion than the mixture of the original gases, and as a result the air is suddenly ejected with violence. "When the explosion occurs in a closed vessel, no sound is lieanl (<■/. e.g. Fig. 13, p. 25). The temperature to which detonating-gas must be heated to explode is found to be about 700° At a lower temperature com- bination l)etween hydrogen and oxg}-en also takes place, but not instantaneously, as in explosions; the lower the temperature, the slower the process. When, therefore, no change in cold detonating- gas is observed even in the cnurse of several years, we must attribute the fact to the extraordinar}- slowness of the process at -ordinarv temperatures. A simple calculation will make this plain. BoDENSTEiN observed that, when detonating-^as is heated at 509° for 50 minutes, 0.15 of the whole is changed to water. Now it is a general rule that, when the temperature sinks 10°, a chemi- cal reaction becomes about twice as slow; at li)!)° it would thus 16 INOROANIC CHEMISTRY. [§§ 13- take 100 minutes till the 0.15 part of the gas had formed water. At the ordinary temperature, say at 9°, it would be 50X2'''' minutes, that is about 1.06X10" years. The same can be said of all chemi- cal reactions. When we see that wood, sulphur, etc., burn quickly at higher temperatures, we must admit that oxidation takes place also at ordinary temperatures, though so slowly that we cannot perceive it. Moissan, however, succeeded in proving that charcoal at 100° and sulphur at ordinary temperatures are oxidized very slowly in a current of oxygen. Hydrogen is not only able to unite with free oxygen, but it also has the power to withdraw oxygen from many of its compounds. The action of hydrogen on a compound is called, in general, reduction. This action is often a very useful means of determining whether a compound contains oxygen, since the latter, if present, iviU usually unite with the hydrogen to form water. Copper oxide may serve as an example of the application of this method. A little is placed in a tube, hydrogen is led over it, and heat is then applied; one soon sees the black oxide change to red copper, and water depositing in drops on the colder parts of the tube. Many other oxides can be similarly reduced, e.g. iron oxide, lead oxide, etc. THE CONSERVATION OF MATTER. 14. The quantitative relationships in oxidizing and reducing processes, such as have been discussed in § 13, i.e. the relations of the masses of the substances participating in the changes, may be used to elucidate a very important law. A definite amount of copper powder, for example, may be placed in a tube and the weight of the tube with the powder ascertained. Oxygen is then led over the copper at a high temperature. The apparatus should be so arranged that the volume of the oxygen which combines with the copper can be measured. When the oxidation process has proceeded for some time, the tube containing the oxidized copper is allowed to cool and then weighed. The weight is found to have increased, and the increase is just equal to the weight of the volume of oxygen used up. Thereupon hydrogen is passed through the tube with the copper oxide and heat applied. Here also arrange- ments should be made for measuring the volume of hydrogen con- sumed in reduction. The reduction is allowed to go on until all the copper oxide is transformed back to copper. When the tube and powder are subsequently weighed, they will be found to have re- 15.] WATh'H. 17 assumed their original weight. The water that forms can be absorbed by a substance hke quickHme or concentrated sulphuric acid and weighed. It will be found equal in weight to the loss of weight of the copper oxide on changing to copper plus the weight of the consumed hydrogen. In these cases, therefore, the combined weight of the reacting substances before and after the reaction is the same. Coppers- consumed oxygen weighs just as much as copper oxide; copper oxide + consumed hydrogen weighs just as much as copper + water; and, finally, the regained copper weighs just as much as that origi- nally taken. The substances can be changed into different states, but their weight remains unaltered. This phenomenon is observed without exception in chemical actions, and we therefore accept as a law the statement that matter is indestructible, or that no matter can be lost or gained. This principle was introduced into chemistry by Lavoisier (174:3-1794;. The old Greek philosophers were already firmly convinced of the impossibility of producing or destroj'ing matter In aU ages this belief has been the basis of philosophic thought. To Lavoisier is due the credit of having demonstrated the practical application of the principle of the indestructibility of matter. He assumed that gravity is an inseparable attribute of all matter — concerning which a great deal of doubt still existed — and that the combined weight of the substances concerned must therefore be the same before and after a chemical reaction. The theory of knowledge teaches that the principle of the indestructi- bility of matter Ues originally at the basis of our thinking. It is entirely incorrect to suppose that it was established by experimentation; on the contrary, we test the correctness of our experimental results by ascertain- ing in how far they conform to this principle. This can be easily under- stood in the above case of the oxidation and reduction of copper. In per- forming this experiment one finds that the weight of copper + oxygen is not exactly equal to that of the copper oxide formed. Even after several repetitions slight differences are still found. Because we feel that there must be absolute equality, we attribute these differences to imperfections in our instruments, and we consider our instruments improved if they enable us to approach nearer the complete equality of the weights before and after the experiment. Nevertheless, we are unable to really observe an absolute equality. WATER. 15. Water was regarded as an element for many centuries. Not until 1781 did Cavendish discover that, when a mixture of hydro- 18 IXORGAXIC CHEMISTRY. [in gen and air or oxygen explodes, water is formed. Being, how- ever, a supporter of an erroneous theory (§ 106), he failed to realize the importance of his discovery. Lavoisier in 1783 repeated this experiment and comprehended it as a synthesis of water, as we still do to-day. With the aid of the apparatus pictured in Fig. 7, this synthesis can be easily demonstrated. The hydrogen is generated in the Fig, 7. — Combustion of Hydrogen". two-necked (AVouup) bottle from zinc and sulphuric acid. In order to free the gas from water vapor, it is passed through the horizontal tube, which contains chloride of calcium or bits of pumice-stone soaked in sulphuric acid. The dry gas is ignited and, as it bui-ns, water is gradually deposited on the walls of the bell-jar. A mixture of hydrogen and oxygen unites to form water when illu- minated with ultra-\'iolet light. In addition tn this direct synthesis from its elements there are other ways of obtaining water. For example, manj^ compounds, such as the blue crystals of copper vitriol, give off water when heated. The formation of water by the action of hydrogen on oxygen compounds was illustrated (§ 13) in the reduction of copper oxide. On the other hand, it is also produced by the action of oxygen on certain hydrogen compounds. This is seen, for example, in the burning of alcohol. Finally, water can result from the reaction of a hydrogen com- pound with one of oxygen. This is the case when ammonia gas (§ 111) is led over hot copper oxide. 16. WATER. 19 The sj-nthetic methods of im-paiiug water, such as the aliow- named and many others, possess, however, merely theoretical importance. Even when water is wanted in a perfectly pure state, natural water is resorted to. This contains solids and gases in solution, which must be eliminated. Its purification is accom- plished by distillation. An apparatus well suited to this purpose IS shown in Fig. 8. High pressure steam and electricity are often used for heating instead of the flame. Water is placed in the retort A, which rests over the fireplace, and boil- ed. The dissolved gases are first driven off; the hot steam follows, passing through the dome B into the condensing coil ("worm") C, which is cooled by water in the \-essel D. The condensed water, now pure, flows down into the bottle; the solid 5^^ g _pjj^jj,p^„j,jj ^^ -^^^^j^ ^^ jj^g^^^^^^^^ substances that were dissolved in the water remain in the retort. The cooler D b supplied with cold water through a tube, entering near the bottom, while the heated, and therefore specifically lighter, water flows out near the top. The steam thus meets with cooling- water of a lower temperature as it passes down the worm, and is in this v,-ay very completely condensed (principle of the counter-current). A single distillation is usually insufficient for the complete elimination of all gaseous and solid coiistitueiits. For this purpose the operation must be repeated in an apparatus of platinuiii (tin is Irss satisfartoiy) with a condensing coil of the same metal, and only the middle fraction collected. An excellent criterion for tlic purity of water is to be found in the measurement of its elei-tricul resistance. Aery pure water conducts the electric current scarcely at all. Koulrauscii found the conductivity at 20 IXORGANIC CHEMISTRY. [§§ 15_ 1S° of the purest water obtainable to be /c=0.038xlO"'° expressed in reciprocal ohms; by this is meant the conductivity of a body a column nf which 1 cm. long and 1 cm. square in cross-section has a resistance of 1 ohm. The magnitude of the resistance of such water is better un- derstood by comparing it with resistance of copper. 1 cu. mm. of this water has at 0° the same resistance as a copper wire of the same cross- section and 25 million miles long; it could be strung around the earth's equator one thousand times. The slightest traces of salts or even con- tact with the atmosphere cause a market increase in its conductivity. PHYSICAL PROPERTIES. i6. Water at ordinary temperatures is an odorless, tasteless liquid, showing no color in thin layers. On looking through a layer 26 meters thick. Spring observed a pure dark-blue color. The- thermometer-scale of Celsius is fixed according to the physical constants of water, its freezing-point being called 0° and its boiling- point at 760 mm. pressure 100°. These two points are dependent on the pressure. An increase of pressure lowers the freezing-point (0.0075° per atmosphere). This is the reason why ice melts under high pressure. Water possesses the very uncommon property of having a maximum of density (minimum of volume) at a definite temperature. The volume of almost all other substances increases- with rising temperature, but here it diminishes up to 3.945°, above which temperature water expands as heating continues. During the transformation of water to ice the volume increases considerably. One vol. water at 0° yields 1.09082 vol. ice of the same temperature. The specific heat of water is greater than that of a vast majority of other substances. Its latent heat of fusion is 79 Cal., its latent heat of vaporization 536 Cal. Water is extensively used as a sol- vent. Numerous substances dissolve in it to a greater or less degree. There are many liquid substances that mix with water in all pro- portions, and many, also, which do not. (See § 7.) The remarkable physical properties of water play a very important lole in nature; this subject is extensively discussed in physics, meteor- ology, and geology. NATURAL WATER. 17. Water, as it occurs in nature, is by no means chemically pure. It may contain solid matter in suspension as well as sub- stances, either solid or gaseous, in solution. The purest natural water is rain-water. This has really passed through a natural process of distillation, the water on the earth's surface being vapor- 17.] NATURAL WATER. •21 ized by the sun's heat and condensed again by contact with colder portions of air, whereupon it falls in the form of rain. Neverthe- less it contains dust particles (in large cities more, of course, than in the country) and gases from the air, as well as traces of ammo- nium salts. Spring- and well-waters contain in 10,000 parts about 1-20 parts of solid matter, consisting largely of lime salts. Well-water that con- tains much lime is called hard (§ 259). Well-water also contains some carbonic acid and air in solution, both of which give it its refreshing taste; distilled water tastes flat. Natural water is used extensively for drinking purposes. When it comes out of a soil that is contaminated by decaying organic matter, as is the case in many large cities, it is injurious to health, principally on account of the presence of bacteria. It can be freed from these by filtration through a Pasteur-Chamberland porcelain filter (Fig. 9). This consists essentially of a hollow cylinder of porous porcelain (called a "candle") A, through whose walls the water is forced by its own pres- sure. The lower end of the candle opens into the nozzle. In large cities it has been found much more practicable to purify the well- or river-water at the central station and to pipe it thence to the various houses. Epidemic diseases have really decreased remarkably since the introduction of the methods of modern sinitary science. A water which contains so many substances in solution that it has a definite taste or a therapeu- tic effect is called a mineral water. There are very many kinds of mineral waters, differing accord- ing to the amount and kind of dissohed matter they contain. We distinguish between saline waters containing common salt, bitter waters with magnesium salts, sulphu- rous waters with sulphuretted hydrogen, carbonated waters with carbonic acid, chalybeate waters with iron, and many others. Detailed analyses of the mineral waters of numerous watering-places are accessible in jyorks on balneology. Sea-Hunter contains about 3% of salts^ of which 2.7% is common Fig. 9.— Pasteuh- Ch AMBERLAND Filter. INORGAXIC CHEMISTRY. m 17- salt. A large number of elements, viz., about thirty, have been found in sea-water, although the most of them exist there only in extremely small quantities. It was stated above (§ 16) that pure water is blue. The color of the rivers, lakes and seas varies, however, through many nuances from pure blue to brown. This variation is due principally to the presence of more or less brownish-yellow humous (marshy) substances or an extremely fine floating slime. Both conditions can produce a brownish-yellow color. It is easily seen how the combination of blue and yellow or brown may bring about the various blue, green or brown tints in natural waters. COMPOSITION OF WATER. i8. Decomposition. — It was stated above that water can be obtained by direct combination of hydrogen and oxygen; inversely, it can be decomposed into these same elements. In the flask ^4 ( Fig. 10) some water is heated till it boils vigorously. A strong electric current is then sent through the wire a c 6, so that the fine platuium wire c glows intensely. This heat partially decomposes the Fig. 10. — Decoiiposition' of Water by Glowing Platik water vapor into hydrogen and oxygen, which pass out through the tube d and are collected in the c.vlinder C. This gas mixture is nothing but the explosive mixture (§ 13) of hydrogen and oxygen, as can be easily proved b\' applying a flame. 19.1 COMPOSITION OF WATER. 23 Many metals decompose water on contact, the hydrogen being set free and the metal uniting with the oxygen. Potassium and sodium effect tliis decomposition at ordinary temperatures (§ 11); iron, zinc and otlier metals require a higlier temperature, iron, e.g., acting at red heat. 19. Let us now study the quantitati^■e composition of water, i.e. determine the relative amounts of hydrogen and oxygen present. For this purpose both the analytic and synthetic methods can be used. (a) The Analytic Method. — When an electric current is passed through water to which has been added a little sulphuric acid, the water is decomposed. If the gases evolved at the electrodes are collected separately, it is found that for every 1 vol. oxygen 2 vols, hydrogen are given off. A suitable apparatus for this experiment is shown in Fig. 11. Since 1 liter of hydrogen weighs 0.0899 g. and 1 liter of oxygen -weighs 1.4296 g., both at 0° and 760 mm. pressure, the weights of 2 vols, hydrogen and 1 vol. oxygen must bear to each other the ratio of 2X0.0899 : 1.4296, or 1 : 7.94.3. (6) The Synthetic Method. — As early as 1820 the reduction of copper oxide by hydrogen was emploj-ed for this purpose by Berzelius; in 1834, also, by DuiiAs and St.\s. A weighed amount of care- fully dried copper oxide is heated in a current of hydrogen and water is formed, which is collected and weighed. The weight of the oxygen given up by the fopper oxide is found from the difference between the weight of the copper oxide used and that of the resulting copper. The weight of the hydrogen contained in the water coUecterl is therefore equal to the difference in weight of water and oxygen. The apparatus used for this experiment is represented in Fig. 12. In A the hydrogen is generated from zinc and dilute sulphuric Fig. 11, — Electrolysis OF Water. 24 INORGANIC CHEMISTRY. [§§19- acid. It is then passed through the permanganate solution in the wash-bottle B to free it from impurities, and also through the U- tubes C, D and E, containing calcium chloride, sulphuric acid and phosphorus pentoxide, respectively, for drying it. In F is placed the copper oxide, which is carefully weighed together with the tube. The water that forms is condensed in G, the U-tube H being attached to absorb any escaping water vapor. At the completion of the experiment, F, with its contents, is again weighed, likewise G and H; the differences in weight indicate the amount of water Fig. 12. — Synthesis of Water after Dumas and Stas. formed. Dumas and Stas found in this way that 100 parts (by weight) of water consist of 11.136 parts of hydrogen and 88.864 parts of oxygen, or, in other words, that the mass-ratio of these elements is 1:7.980, a relation which agrees with that ob- tained in (a) within the range of the unavoidable experimental error. Another synthetic method, which is especially adapted to the lecture-table, consists in mixing hydrogen and oxygen and deter- mining in what volume-ratio these gases unite. For this pur- pose an apparatus (Fig. 13) described by Hofmann is best em- ployed. Hydrogen and oxygen in different proportions by volume are introduced into the arm of the U-tube, which can be closed by 20] COMPOUNDS AND MIXTURES. 25 a stop-cock at the top: the cock is thereupon closed and the open arm tightly stoppered with a cork. The mixture is then exploded by an in- duction spark, the volume of air en- closed on the other side acting as a cushion to moderate the severe shock on the mercury, which might otherwise break the apparatus. It is found that only when the volumes of hydrogen and oxygen bear to each other the ratio 2:1 does the entire gas mixture dis- appear, a slight coating of tiny drops of water appearing in its place on the inside of the glass. In case more hy- drogen or more oxygen than the ratio calls for is let into the tube, the excess is found to remain after the explo- sion. From these experiments, analytic and synthetic, it follows that icatcr has a constant composition; it consists of 2 vols, of hydrogen and 1 vol. of oxygen, or of 1 part, by weight, of hydrogen to 7.943 parts of oxygen. Fig. 13. — Hofjiann's Ap- paratus FOR THE Syn- thesis OP Water. COMPOUNDS AND MIXTURES. 20. In water we have become acquainted with a substance which is different in many and important respects from the elements of which it is composed. We have further seen that the elements in it bear to each other a fixed relation by weight. Such substances are known in verj' large number. Copper oxide, mercury oxide, sulphuric acid, potassimn chlorate, common salt, soda and many others already mentioned belong to this class. In each ef these, no matter how obtained, we discover by analysis or syc thesis a definite proportion between the elements composing it. Such sub- stances are called compounds. In addition to the characteristics mentioned — difference of prop- erties from those of the elements and constant composiOon — we find that the compounds also have constant physical properties. 2G INORGANIC CHEMISTRY. [§§ 20- Uiidor the same pressure water always has the same melting-point and the same boiling-point, in whatsoever way it may have been obtained; salt always crystalhzes in the same crystal system; soda, at a definite temperatm-e, always requires the same amount of water for solution, etc. When elements or compounds are brought together without any chemical action on each other taking place, we have a mixture of these elements or compounds. The number of possible mixtures is, of course, unlimited. They are distinguished from compounds by the following characteristics : In a mixture the properties of the components reappear in many and maportant respects. Gunpowder, for example, is a mixture of sulphur, charcoal and saltpetre. The latter is soluble in water; sulphur dissolves in carbon disulphide; charcoal is insoluble in both. These properties are still evident in the constituents of gunpowder. In a mixture of sulphur and iron filings one can detect with a micro- scope the yellow grains of sulphur and the black particles of iron. The iron can be drawn out with a magnet; the sulphur dissolved ■out by carbon disulphide. If, however, a mixture of 7 parts iron and 4 parts sulphur is heated, a glow passes through the powder and a compound of both — iron sulphide — is formed, whose prop- erties are entirely different from those of its elements. It is non- magnetic and insoluble in carbon disulphide and under the micro- scope only a homogeneous scoriaceous mass is seen. The constituents of a mixture, since they stUl preserve their properties, can often be separated from each other by mechanical means, e.g. by the use of microscope and tweezer.s, liy sifting, by treatment with solvents, by washing, etc. In a mixture the ratio of the constituents can vary in all pro- portions. There are, for example, many sorts of gunpowder, dis- tinguished from each other by the proportions in which their con- stituents are mixed. When 1 part sulphur and 100 parts iron, or, on the other hand, 1 part iron and 100 parts sulphur, are mixed, we have in either case a mixture of both elements, possessing hardly the same, but at least analogous, properties. IVIoreover, a mixture often has no constant physical properties. Water has a constant boiling-point; the boiling-point of a mixture of benzene and turpentine, however, rises gradually as the more volatile component, benzene, distils off. The melting-point of 21 J COMPOSITION OF COMPOUNDS.— A iOMIC THEORY. 27 sulphur is constant and can be accurately determined; that of a mixture of tin and leail differs according to the proportion of the elements and is in man>' proportions not at all sharp, there being only a softening instead of real fusion. In the examples cited here the distinction between a compound and a mixture is well marlvcd. There arc, however, other instances where this is r.ot the case and where it is therefore very difficult to know whether one is dealing with a compound or a mixture. We shall meet with many examples of this later. There is, however, one way whereby a compound can be distinguished from a mixture, viz., by ascertaining whether or not the substance, prepared in different ways, has a constant composition. PHENOMENA ACCOMPANYING THE FORMATION OR DECOMPOSITION OF A COMPOUND. The most common phenomenon of this sort is an elevation or depression of temperature, i.e. an evolution or absoiption of heat (caloric effect). Sometimes the rise of temperature is so great that light Ls produced. A decomposition or a combination can be so violent that it causes an explosion. In other instances electricity may be produced b}' chemical action. All these facts may be com- prised in this statement: Chemical action results in a change in the energy-supply of the reacting substances. EXPLANATION OF THE CONSTANT COMPOSITION OF COMPOUNDS.— ATOMIC THEORY. 21. It •was stated that constant composition is the distinctive characteristic of a chemical compound. The proportions in which elements unite to form a certain compound are always the same. This Law of Constant Composition (definite proi)i)rtions) was finally established by Proust in the beginning of the nineteenth century, and at aVjout the same time DALTo>f offered an expla- nation of it which is still accepted and may be considered as the foundation of theoretical chemistry. This explanation involves a hypothesis as to the constitution of matter. It is possible to regard matter as infinitely divisible; according to human concej^tion the smallest particle that can really be obtained is still capable of division into an infinite number of others. However, even tlie ancients were of the opini(m that there 28 INORGANIC CHEMISTRY. [§§21- must be somewhere a limit to the divisibility and that we must finally arrive at particles incapable of fuither division, the atoms. In the fifth century B.C. there existed a school of philosophy, that of the Eleatics (so called from the city of Elea), whose most prominent representative was Parmbnides. He taught that everything that exists cannot be otherwise conceived than as unchangeable; trans- formation of the existent, which was thought to have never originated and to be at the same time unalterable, was held by them to be incon- ceivable. These theses they regarded in a certain sense as axioms, i.e. statements of truths which are accepted without proof. Daily experience teaches one nevertheless that transformation does occur in that which exists, a fact that led them to suppose that everything observed by men is merely appearance. Three theories were proposed in the same century which aim to form a bridge between the doctrine of the unalterable existent and the experience that points toward continuous change. These theories originated with Empedocles, Anaxagoras, and the Atomists, Ledcippus and Democrites. The immutability of the existent is disposed of by ascribing it to extremely small unchangeable and indestructible particles; every change is thought to depend on the movement of these smallest integral particles toward or away from each other. EMPED0CL"f,5 aud Anaxagoras assume in this connection an infinite divisibility; the Atomists, on the contrary, regard the world as built up of indivisible particles, atoms, all of which consist of the same primordial substance but differ in form and size. Now Dalton has used this conception of the ancients regarding the atom to explain the fact that the combining weights are con- stant. The atoms of the various elements, he assumes, have dif- ferent weights; the atoms of the same element are alike in weight. A compound of two elements is therefore produced by the associa- tion of atoms of these elements. Such a combination of two or more atoms is called a molscule. It is obvious that these supposi- tions lead directly to the law of constant pro-portions; for, if copper oxide is formed by an atom of copper uniting with an atom of oxygen to make a molecule of copper oxide, its composition must, according to the above hypothesis, be constant. Dalton deduced another conclusion from his hypothesis, and confirmed the same experimentally. He observed that oxygen unites not only with one very definite amount of nitrogen oxide, but also with twice as much, not, however, with any intermediate 22.] THE ATOMIC WEIGHTS OF THE ELEMENTS. 29 amount. He also showed by the investigation of marsh-gas and olefiant gas, both of which are made up of only carbon and hydro- gen, that the former contains twice as much hydrogen to a certain weight of carbon as the latter. It is readily seen how such observa- tions can be explained on the basis of the atomic theory; in one case 1 atom of carbon is in combination with n atoms of hydrogen; in the other with 2« atoms. Ihe cbscrvations of Daltox were subsequently confirmed and extended, especially by Kehzelius. The following statement is therefore now accepted as a law: When two elements combine to form more than one compound, the different weights of the one element which unite with one and the same weight of the other element bear a simple ratio to each other. This is the Law of Multiple Proportions. THE ATOMIC WEIGHTS OF THE ELEMENTS. 22. The absolute weight of the atoms is only approximately known (see § 35). Nevertheless, their relative weights, i.e., the weights of the atoms of the various elements, when that of a certain element is arbitrarily fixed, have been determined in a variety of ways (§§ 208-210). These relative weights are known as atomic weights. It is now customary to take the atomic weight of oxj'gen as 16.00. The atomic weights of the remaining elements then have the values that are given in the table on the inside of the back cover of this volume. The acceptance of 16 as the atomic weight of oxygen has a historic reason. For a long time hydrogen was taken to be 1 ; it was believed that the ratio of the atomic weights of hydrogen and oxygen was 1 : 16. Inasmuch as the atomic weights of most elements are determined from the composition of their oxy- gen compounds, the basis is really 0=16 and not H = 1. This made no difference, so long as the proportion H:0=l:16 was con- sidered accm-ate. Even when the ratio was later found to be a different one (according to investigations of Morley and of W. A. Notes the ratio 1:15.88 may now be regarded as very accurately determined), it was stUl the simplest plan to preserve = 16 as the basis, since a change would necessitate a complete recalculation of all the atomic weights, and this necessity would 3U IXORGAMC CHEMISTRY. [§§22- moreover recur as often as a new refinement of methods of inves- tigation brought about a change in the ratio H:0. A few years ago there was estabhshed a permanent inter- national commission whose duty it should be to revise the table of atomic weights critically every year. Those values are accepted as the " international atomic weights " which appear to be the most probable among the determinations that have been pub- lished. The atomic weights in the table are carried out to as man}' decimal places as may be accepted with certainty. For many purposes, however, it is sufficient to use round numbers, such as K=14, Br = 80, etc. Besides the atomic weights, we quite frequently use equivalent weights. These are the weights of the elements which combine with a unit amount of a certain standard element. One part of hydrogen combines, for instance, with 35.5 parts of chlorine and mth 8 parts of oxygen. These amounts of hydrogen, chlorine and oxygen are equiv- alent to each other. The atomic weight is either equal to the equivalent weight or a multiple of it. CHEMICAL SYMBOLS AND FORMULA. 23. The relative, or atomic, weights are expressed by symbols, that were introduced by Berzelius and are of great convenience in the representation of compounds and the formidation of chemical reactions. The s)-mbols whose derivation is not at once apparent are taken from the Latin names of the elements; e.g., Sb from stibium, Au from aurum, Cu from cuprum, Hg from hydrargyrum, Pb from plumbum, Sn from stannum, Fe from ferrum, and Ag from argentum. A symbol stands not onlj' for the element concerned, but also for the relati\-c weight of an atom of that element. If the atomic weight of copper is 63.57 and that of oxygen 16.00, the symbol Cu uidicates 63.57 parts by weight of copper, the symbol 16.00 parts by weight of oxygen. It has been deter- mined that in copper oxide one atom of copper is combined with one atom of oxygen; copper oxide is therefore represented by the formula CuO, which expresses, first, that we are dealing with a compound of copper and oxygen, and, second, that 1 atom (63.57 parts by weight) of copper is united in it to 1 atom (16.00 24.] STOICHIOMETRIC AL CALCULATIONS. 31 parts by weight) of oxygen. ^lany compounds contain several atoms of the same element. This is indicated by placing the proper figm-e to the right of and below the symbol. Stdphuric acid, for example, contains 2 atoms of hydrogen (H), 1 atom of sulphui' (S) and 4 atoms of oxygen (0) in the molecule. Its formula is, therefore, H2SO4. Chemical actions can be verj'' simply represented by the use of these formula; thus, the decomposition of mercuric oxide into oxygen and mercury by HgO = Hg+0; that of potassium chlorate into oxygen and potassium chloride by KCIO3 = KCl + 30; Potass, chlorate. Potass, chloride. the generation of hydrogen from zinc and sulphiiric acid by Zn + H2SO4 = 2H + ZnSOi. In such equations the same atoms and the same number of each must appear on both sides, in accordance with the principle of the Indestructibility of Matter. STOICHIOMETRICAL CALCULATIONS. 24. If the formulae of the compounds are known — the means of ascertaining these will be discussed in detail later — and the atomic weights of the elements composing them also known, it is ^ery eas}-^ to calculate the weights that enter into reaction in all chemical changes. A couple of examples may serve to make this clear. 1. It is required to know how many liters of oxygen at 0° and 760 w,m. 'pressure can he obtained by heating 1 kilogram of mcrniric oxide. The atomic weight of mercury is 200, that of oxygen is 16; mercuric oxide, HgO, is, therefore, 200 + 16. Out of these 216 parts by weight of mercuric oxide 16 parts of oxygen can be ob- 32 INORGANIC CHEMISlRi. [§§ 24r- tained by heating, i.e. from 1 kilo (=1000 g.) can be obtained 1000X16^^^ Q^ g. Since 1 1. oxygen at 0° and 760 mm. pressure 74 07 weighs 1.4296 g., 74.07 g. occupy a volume of ][4^g=51.8 L 2. How much water can he formed from the hydrogen obtained by the interaction of 1 kg. zinc and the corresponding amount of sulphuric acid? The reaction of zinc and sulphuric acid is expressed by the equation Zn + H2SO4 = ZnSOi + 2H and the combustion of hydrogen to form water by the equation 2H + 0=H20. From these equations it follows that the hydrogen formed by the action of 1 atom of zinc yields 1 molecule of water. For every atom of zinc we obtain, therefore, 1 molecule of water. The atomic weight of zinc is 65, the molecular weight of water 18; therefore 65 parts of zinc correspond to IS parts of water. 1 kg. zinc must . ,, 1000X18 _„„ yield gg =276.9 g. 3. Hoir many grams of potassium, chlorate are necessary to pro- duce enough oxygen to oxidize 500 g. copper to copper oxide f The reactions concerned are KC103=KCl+30 and Cu+0=CuO. Hence 3 atoms of copper can be oxidized with the oxygen derived from 1 molecule of potassium chlorate. For every 3 atoms of copper 1 molecule of potassium chlorate must be consumed. The molecular weight of the latter substance is 39.10+35.6 + 3X16=122.56; the atomic weight of copper is 63.57; for every 122 56 63.57 parts of copper — ^ — = 40.85 g. potassium chlorate are o .If • 1 XT .nn ■ 500X40.85 therefore required. Hence 500 g. copper require „„ _ — -= 00 .01 321.5 g. potassium chlorate. In most chemical computations gram molecules are employed, these being the molecular weights of the substances in grams. The abbreviation mole has been suggested by Ostwald for this 25.] CHLORINE. 33 term. Thus " 1 mole " copper oxide means 63.57+16.00 = 79.57 grams of it. The molecular weight In milligrams is called a millimole. In the same way we may speak of a kilo mole, etc. CHLORINE. 25. Chlorine does not occm- free in nature, since it acts upon the most diverse substances at ordinary temperatiires. In com- pounds, however, it occiors extensively. Common table salt is a compound of sodium and chlorine. A'arious other metallic chlor- ides are also met with in nature. Chlorine gas can be obtained by the direct decomposition of certain chlorine compounds; thus: 1. By the electrolysis of Iiydrochloric, or muriatic, acid (i.e. a solution of hydrogen chloride, Hf!l, in water). Chlorine is given o£f at the positive pole (anode), hydrogen at the negative pole (cathode). The indirect decomposition of its compounds offer?, as in the case of hydrogen (§ 11), the most practicable methods of obtaining the element. They are all based on the oxidation of the hydrogen of hydrocliloric acid, whereby water is formed and chlorine liber- ated. 2. Commercially, as well as in the laboratory, manganese dioxide, jMn02, is frequently used as the oxidizing agent: :\In02 + 4HC1 = :\InCl2 + 2H2O + 2C1. It is very often con^'enient to generate the hydrochloric acid from salt and sulphuric acid in the same vessel with the manganese dioxide. The two reactions thus proceed simidtaneously: I. XaCl -1- H2SO4 = XaHSOi + HCl. II. 4HCl-|-Mn02 = MnCl2 + 2H20 + 2a. 3. Other commonly used oxiilizing agents are chloride of lime and potassmm dichromate; e.g. K2Cr207 + UHCl = 2KC1 + CrjCle + 7H2O + 6C1. 34 INORGANIC CHEMISTRY. [§§25- 4. The oxygen of the air can also serve as the oxidizing agent: 2Ha + 0=H20+Cl2. For this purpose a mixture of G0'>^ of air and 40% of hydrogen chloride at about 430° is passed over porous bricks which are soaked with copper sulphate solution. About 707c of the hydro- gen chloride is converted into chlorine. This method, which is known as the Deacon process, is used commerciall}'. The copper sulphate serves as a catal3'zer. The progress of chemical changes is often modified by the mere presence of a substance which has the same chemical com- position after the reaction as at the beginning. Such a substance is termed a catalyzer and the action which it exerts is called catalysis, or catalytic action. The quantity of the catalyzer necessary to exert a perceptible influence is often very small. Tills is the case, for example, in the combination of hydrogen and oxygen in the presence of platinum as a catalyzer (§ 13, p. 16). A minute trace of platinum sponge brought into contact with detonating-gas accelerates the combination to such a rate that the reaction takes place very quickly and can even become explosive. In the Deacon process a small quantity of copper sulphate suffices to bring into reaction unlimited quantities of hydrogen chloride and oxygen. At the temperatui-e of <430° there is practically no reaction between oxygen and hydrogen chloride without the catal}-zer. That there must, nevertheless, be a reaction, although a very slow one, can be demonstrated by the same reasoning as in § 13. The catalyzer therefore does not cause a reaction, but only accelerates it. Ostwald compares its action to that of oil on the axles of a machine which move with very great friction. When oiled, the ma- chine will go much faster, notwithstanding that the force of the spring (here the energy of the chemical reaction) has not changed. A f mother point in the analogy is that the oil is not consumed. In most cases of catalysis it can be proved that the catalyzer takes part in the reaction but at the end of it reappears in its original condition. In the platinum catalysis of detonating-gas, 27.] CHLORIXE. 35 for example, the metal unites witli the oxygen, whereupon the resulting compound reacts with the hydrogen, giving water and metallic platinum. The phcuomonon of catalysis is universal. OsTWALD thinks it probable that there is no kind of chemical reaction that cannot be influenced catalytically and that there is no substance, element, or compound, which cannot act as a catalyzer. Catalyzers may accelerate or retard reactions; at present, li<)we\-or, much more is known of the first than of the second k'ud. 26. Physical Properties. — Chlorine is yellowish-green (hence its name, which is derived from ;ifA oj/jos, greenish-yellow) and has a disagreeable odor. Its specific gra^-ity is 2.45, taking air as unity, or 35.46, based on 0=16. 1 1. chlorine weighs, therefore, 3.208 g. at 0° and 760 mm. pressure. At —34° it becomes licjuid under ordinary pressure; at —102° it solidifies and crystallizes. Its critical temperature is 146°. Liquid and solid chlorine are yellow. Chlorine gas dissoh'es in about one-half its volume of water. The aqueous solution bears the name "chlorine-water." It can, therefore, not be collected over water, but a saturated salt- solution may be used, in which it is only slightly soluble. The most convenient way to fill a vessel with it is by displacement of air, the gas being conducted to the bottom, where it remains and dri^-cs out the air above, because the chlorine is denser. 27. Chemical Properties. — E\-en at ordinary temperatures, chlo- rine combines with many elements and acts on many compounds. If perfectly pure chlorine is mixed with an equal volume of hydro- gen, the two unite in direct sunlight, causing an explosion. If the chlorine is impure or the sunlight fliffused, combination occurs slowly. When a hydrogen flame is introduced into chlorine gas, it continues to burn, with the formation of hydrogen chloride. Many metals combine with chlorine with the evolution of light, e.g. cop- per (in the form of imitation gold-leaf), finely powdered antimony, molten sodium, etc. The precious metals are in general quite resistive to chemical action. They are, liowevcr, attacked by chlorine and changed to chlorides, i.e., chlorine compound.s. Gold, for instance, dissolves in chlorine-water, forming gold chloride. 36 INORGANIC CHEMISTRY. [§§27- Chlorine also unites readily with many non-metals, e.g. phos- phorus, which burns in it with a pale flame to phosphorus chloride. The tendency of chlorine to unite with hydrogen — its so-called chemical attraction, or affinity, for the latter — is so strong that chlorine abstracts the hydrogen from many hydrogen compounds in order to combine with it. A strip of paper dipped in turpentine burns with a sooty flame when introduced into an atmosphere of chlorine; the chlorine unites with the hydrdgen of the tiu"pentine and sets the carbon free. A burning candle continues to burn in chlorine, depositing soot (carbon) and forming hydrogen chloride. If sulphuretted hydrogen gas, HjS, is passed into chlorine-water, hydrochloric acid and sulphur are formed. Water is also decomposed by chlorine, oxygen being liberated : 2H20 + 2Cl2=4HCl + 02. This reaction takes place under the influence of sunlight, but proceeds very slowly. It can be conveniently demonstrated as in Fig. 14. A retort is filled with dilute chlorine-water, inverted and exposed to the sunlight. After a few days a bubble of gas collects at the top of the retort, and. on investigation with a glowing splinter, it is found to be oxygen. Fm. 14. — Slow Decomposition of Watek by Chlorine. Upon this decomposition of water depends the bleaching and disinfect- ing action of chlorine and those substances which generate chlorine. In bleaching, the coloring matters — usually of an organic nature— are oxi- dized by ox^'gen to colorless substances. Bacteria are killed by oxida- tion. Ordinary atmospheric oxygen does not produce these effects. Lit- mus, for instance, which is rapidly decolorized in moist chlorine gas, is 28.] HYDROGEN CHLORIDE. 37 totally unaffected by the air. The particularly energetic action of the oxygen that is produced from water liy chlorine is explained by assuming that it exists in an atomic condition, fhi' status nascens, regarding which more will be said hitcr (§ 38). Perfectly dry chlorine has no bleaching power. If water is saturated with chlorine at 0°, crystals are deposited, of the composition CL + SH;,0, chlorine hydrate. At a higher temperature these are wholly decomposed into chlorine and water. HYDROGEN CHLORIDE, HCl, and HYDROCHLORIC ACID. 28. Hydrochloric acid, of the formula HCl (§ 31), is a gas, occurring in nature in the free state, e.g. in the gases of some volcanoes. It forms an important, although small, part of the gastric juice of man and other animals. Some of its methods of formation have been already given ( §27), e.g., by direct sjrathesis from its elements under the influence of light. It is quite remarkable, however, that ultraviolet rays decompose hydrogen chloride even at ordinary temperatures. We saw also (Z. c.) that hydrogen chloride is formed by the action of chlorine on hydrogen compounds. Moreover, it can also result from the action of hydrogen on some chlorine com- pounds, e.g., silver chloride, AgCl, and lead chloride, PbCl2, when heated in a current of hydrogen, yield metal and hydro- chloric acid: AgCl+H=Ag-^HCl. The ordinary method of preparation is by the action of a chlorine compoimd on a hydrogen compound, viz., that of salt (sodium chloride) on concentrated sulphiudc acid: NaCl + BSOi = NaHS04 -|- HQ. Sodium Sulphuric chloride. acid. This method is employed technically as well as in the labora- tory. 38 INORGAXIC CHEMISTRY. [§§28- The above reaction takes place at ordinary temperatures. If the sulphuric acid is to be completely used up, i.e. if all the hydrogen of the siilphuric acid is to go off with the chlorine of the salt as hydrochloric acid, the temperature of the reaction must be raised (c/. also § 226) : 2XaCl+ HjSO, =Na2S0, +2HC1. 29. Physical Properties. — Hydrogen chloride is a colorless gas with a pungent odor. Its critical temperature is +52.3°; the critical pressure 86 atmospheres. Liquid hydrogen chloride boils at —83.7°; the solid melts at -111.1°. Specific gravity of the gas = 1.2696 (air = l); 1 1. HCl at 0° and 760 mm. pressure weish* l.Oo.'lo gr. For obtaining; hydrogen chloride in a pure state Moissan has elabor- ated a method which is generally applicable to gases, since low tempera- tures are easily attainable by means of liquid air. The freshly generated gases contain in most cases moisture and other impurities. The gases arc first dried by being passed through one or two wash-bottles placed in a bath of a lower tempei'ature than —50°. At that temperature the tension of water \'apor is i)raetically zero. The gases dried in this way are now condensed by strong cooling to the solid state. Air can then be pumped out of the vessel If the temperature is now allowed to rise, the solid mass melts first; the resulting liquid, when vaporized, gives the l)crfeetly pure gas. The gas fumes strongly in the air, forming a cloud with the moisture of the air. It is very soluble in water, 1 vol. water at 0° being able to absorb 503 vols. HCl gas. The aqueous solution of the gas is called "hydrochloric acid," * also muriatic acid. It is manufactured commercially on a large scale (§ 226). Hydrochloric acid is employed almost exclusively in the form of this aqueous solution. A solution saturated at 15° contains 42.9% HCl and has a specific gravity 1.212; it fumes vigorously, in the air. The ordinary pure "concentrated" or "fuming" muriatic acid of commerce usually has a specific gravity of 1.19 and contains about 38% HCl. ■ The gas itself is often called "hydrochloric acid gas." 30.] HYDROGEN CHLORIDE. 39 Hydrogen chloride does not obey the law of Henry (§ 9) in its behavior towards water, for its solubility in this liquid is not at all proportional to the pressure. The larger part of it is absorbed in water without reference to the pressure, and an increase of pressure causes only a small increase in the solubility. Such con- duct indicates that a change in the compound has occurred; just what this change consists in we shaU soon have occasion to con- sider (§§ 65, 66). The saturated solution of hydrogen chloride in water gives off HCI on warming. On distilling it a fraction is obtained that boUs con- stant at 110° and contains 8 mols. H^O to 1 mol. HCI, corresponding to about a 20% solution of HCI. A solution of the same concentration and boiling-point results from distilling a more dilute hydrochloric acid, enough water boiling off to raise the concentration to the above value. 30. The chemical properties of hydrogen chloride are found to be quite different when it is in a perfectly dry condition, e.g. con- densed to a hquid, than when it is dissolved in water. In the former case it does not act on metals nor change the color of blue litmus. In the latter case just the contrary is true. Zinc, iron, and other metals, when dipped in the aqueous solution of hydrogen chloride, are vigorously attacked, hydrogen being given off. Blue Utmus is turned red by the solution. Moreover, even dilute solu- tions taste som-. Now, there are a lot of substances that undergo a similar change of properties when they are brought in contact with water, and whose aqueous solutions possess about the same properties as those that are described here for hydrochloric acid. The nature of this change will be discussed later on (§ 65). It should be stated here, however, that these substances have a common name. They are called acids. Acul.t hare one or more hydrogen atoms that can be replaced hi/ metals. The compounds 0} metals that are formed by such substitutwn arc c(dlrd salts. Salts can result not only from the direct action of metals on acids, but also from the interaction of acids and bases. The term "bases" includes compounds of the general type M()H, wlicrc M represents a metal. Most of them have an alkaline taste and turn red litmus 40 INORGANIC CHEMISTRY. [§§30- blue. When sodium is dropped into water, hydrogen is generated, and a base, sodium hydroxide, is formed: Na + H20 = XaOH + H. If this liydroxide is now treated with hydrochloric acid, sodium chloride and water are produced: NaOH + HCl = NaCl + H2O. If we indicate an acid by the general formula AH and a base by MOH, the formation of salts from the interaction of the two may be represented thus: M0H + HA = MA+H20. A third way of forming salts is by the action of an acid upon a metalUc oxide, e.g. ZnO + H2SO4 = ZnS04 + H2O. zinc Sulphuric Zinc oxide. acid. sulphate. In general, the bases are built up from metals, the acids froir metalloids. When hydrochloric acid is added to a solution of a silver salt, fov instance to silver nitrate, a decomposition of this salt takes place according to the equation AgNOs + HCl = HNO3 + AgCl. Silver nitrate. Nitric acid. Silver chloride. The silver chloride is insoluble, and is precipitated as a white, curdy mass. In this reaction the hydrochloric acid has liberated the nitric acid from its salt. It is also possible to liberate a base from a salt by the addition of another base: AgNOa + NaOH = AgOH + NaNOs. Silver Sodium hydroxide. nitrate. Such reactions are called single, or simple, decompositions. 31.] COMPOSITION OF HYDROCHLORIC ACID 41 Now it can also happen that two salts exchange their metals when brought together: NaCl + AgNOg = AgQ + NaNOs, Sodium chloride, SO that two other salts are obtained. Such a reaction between salts is called a double decomposition. We shall later have occasion to study the laws governing both of these decompositions. COMPOSITION OF HYDROCHLORIC ACID. LAWS OF GAY-LTJSSAC AND AVOGADRO. 31. The composition of hydrochloric acid is determined by the following experiments : (a) When strong hydrochloric acid (a more than 23% solution) is subjected to electrolysis in a suitable apparatus (see below) it is observed that equal volumes of hydrogen and chlorine are evolved. (6) Equal volumes of chlorine and hydrogen unite to form hydrochloric acid without leaving a remainder of either element. 2 vols. HCl are formed. Since the weight of 1 vol. CI is 3.5.46 (0=16), hydrochloric acid must consist of 1 part by weight of hydrogen combined with 3.5.46 parts of chlorine. In the electrolysis of hydrochloric acid sticks of charcoal are ordinarily used, because platinum, the substance employed in most other electrolyses, is attacked by chlorine. The apparatus of Fig. 11 is also impracticable, since the solubility of chlorine in water increases with rising pressure more rapidly than that of hydrogen, and equal volumes of both gases are therefore not obtained. In its place we use an apparatus suggested by LoTHAR Meyer (Fig. 15), by which the compression of the chlorine by a steadily rising column of liquid is avoided. In .1 hydrochloric acid is electrolyzed and the hydrogen and chlorine are collected in the cylinders BB, which are filled with a saturated sodium chloride solution. The collected gases arc thus under diminished ])iessure. The combination of equal volumes of chlorine and hydrogen can be carried out in a thick-walled tube, that is filled with the gases and then exposed for a day to diffused sunlight. Since the success of the experi- ment requires the use of the exact proportions of chlorine and hydrogen 42 INORGANIC CHEMISTRY. f§,31- and their absolute purity, the gas mixture is prepared by electrolysis in the dark and exposed to the action of light immediately after the tube is filled. Fig. 15. — Electrolysis of Hydrochloric Acid. The fact that hydrochloric acid gas j'ields a volume of hydrogen equal to half its own volume can also be shown in another way. When perfectly dry hydrogen chloride is treated with sodium amalgam — a solution of sodium in mercury — the sodium combines with the chlorine, setting iiydrogen free. The volume of the latter is then found to be half as large as that of the hydrochloric acid taken. Hydrogen and chlorine thus unite in a very simple ratio by volume (1:1), and the volume of their product also bears a very simjjle ratio to that of the components (2:1:1). In discussing the composition of water (§ 19) we already remarked that oxygen and hydrogen combine in a very simple ratio by volume, viz., 1:2. By carrying out this synthesis at a temperature above 100°, so that the steam is not condensed to water, it is found, further, that the volume of resulting steam bears a simple ratio to the volumes of its components, viz., that 1 vol. + 2 vols. H gives 2 vols. H2O. The following arrangement serves this purpose (Fig. 16). The explosive mixture is introduced into the closed arm B of the U-tube over mercury. B is surrounded by a glass jacket, through which the vapor of boiUng amyl alcohol (generated in A), whose temperature is about 130°, is passing. This vapor is condensed in C. As soon as the gas mixture 31.] COMPOSITIOX OF HYDROCHLORIC ACID. A?, has reached this temperature an induction spark is flashed through, and it is found that the volume of steam formed is two-thirds that of the mixture. Fig. 10. — Determixatiox of the Volume Relations betweex Steam axd its co.mponents. What was found abo^-c to be true for hydrochloric acid and for water is a general principle. Gaseous elements combine in simple proportions by volume, and the volume of the products forincd — in the gaseous state — also bears a simple ratio to the volumes of the com- ponents. This law was discovered by G.\y-Lus.s.vc in 1808. This law, together with the atomic theory of D alton, leads to important conclusions. In order to investigate the matter, let us assume that the formula of hydrochloric acid is HC'l; in other words, that an atom of hydrogen is in combination with an atom of clilorine. Since one volume of hydrogen unites with one volume of chlorine to form the compound, it follows from the above as- sumption that eijual ^'olumes of chlorine and hydrogen contain the same number of atoms. If the formula were otliorwise, e.g. H„Clm, the numbers oi atoms in equal A'olumes of hydrogen and chlorine would be in the ratio of n:m. In the synthesis of water 2 vols, of hydrogen and 1 vol. of oxy? gen yield 2 vols, of steam. If the formula of water be H2„0p, the numbers of atoms in equal volumes of hydrogen and oxygen must bear to each other the ratio n:p. Given, therefore, the rclati\c numbers of atoms in equal gg^ 44 INORGANIC CHEMISTRY. [§ 31- volumes and the volume ratio in which the gases unite, we can determine the formula of the resulting compound. As to the number of atoms in equal gas volmnes, there was at first much uncertainty. Since aU gases behave exactly alilce to- wards changes of pressure or temperature, it was reasonable to suppose that the number should be alilce for all gases; but this was soon shown to be incorrect. In the synthesis of water 3 vols. (2 voLs. H+ 1 vol. 0) give 2 vols, of steam; hence the niunber of atoms per unit volume must be different for steam than for the uncom- bmed elements. However, all difSculties were overcome by a hypothesis, which Avogadro enimciated in 1811, to the effect that equal volumes of all gases at the same icmpcraliire and pressure con- tain the same number of molecules. Avogadro further supposes that the molecules of oxygen, hy- drogen, chlorine, and other elements toiisis: of two atoms. The union of hydrogen and chlorine is then exp'ained thus; Out of a molecule of each, two molecules of hydrochloric acid are formed: H2 + Cl2=2Ha. 1 vol. 1 vol. 2 vols. The total number of molecules thus remains the same after the combination and, since the entire volume has suffered no change either, there must be just as many molecules present in each of the two volumes of hydrochloric acid as in each of the volumes of hydrogen and chlorine. The combination of hydrogen and oxygen takes place thus: 2H2 + 02=2H20. 2 vols, 1 vol. 2 vols. Every molecule of oxygen has spht up into its two atoms, and each of these unites with two hydrogen atoms. The number of water molecules becomes therefore twice as great as that of the oxygen molecules and equal to that of the hydrogen molecules; but, since the volume of steam is also double that of oxygen, there must be in equal volumes just as many water molecules as oxygen molecules and hydrogen molecules. 32. It follows from the above that Avogadro's hypothesis is of importance in two respects: (1) in furnishing us a means of ascer- 32.1 COMPOSITIOX OF HYDROCHLORIC ACID. I') taining the relative weights of the molecules of gaseous substances; (2) in putting us in a position to form an idea of how many atoms there are in the molecules. Let us examine both points more olosel}-. As to (1): Since equal volumes of gases under the same conditions contain the same number of molecules, the ratio of the weights of these volumes gives us at once the ratio of the m.olecular weights. If the specific gravity of steam is 9, based on 0=16, and that of hydrochloric acid is 18.25, the ratio of the molecular weights of water and hydro- chloric acid is 9: 18.25. The determination of the specific gravity of gases and vapors, the vapor density, becomes therefore of the greatest importance to chemistry. The practical method of pro- cedure is described in Org. Chem., § 12. For the determination of the specific gravity of gases, see §212. As to (2) : In order to understand how Avogadro's hypothesis can furnish an idea of the number of atoms which the molecules oC elements and of compounds contain, let us return to the example of the synthesis of hydrochloric acid. 1 vol. hj-drogen unites with 1 vol. chlorine to form 2 vols, hydrochloric acid. According to the above law there must be just as many molecules present in the two volumes of hydrochloric acid as there were molecules of hydrogen and chlorine together. It is evident that this is only possible in case the molecules of hydrogen and of chlorine •divide into two parts. For, if the chlorine and the hjdrogen molecules consisted of onl)' one atom each, the vohmie of hydro- chloric acid could not, in accordance with Avog.vdro's law, be double that of each of its elements, but would have to be equal to it. It therefore follows that an even munber of atoms must be present in the chlorine and in the hydrogen molecules; whether or not this number is two, as Avogadro assumed, can evidently not yet be determined; we shah therefore represent the molecules of hydrogen and of chlorine by Hox and C^y. From the synthesis of water the same conclusion is reached in regard to the oxygen molecule: 2 vols, hydrogen unite with 1 vol. oxygen to form 2 vols, steam. In each of these two volumes of steam there must be, according to Avogadro's law, just as many molecules present as in the one volume of oxygen. This is likewise impossible unless the oxygen molecule splits into two parts, each of which combines 46 INORGANIC CHEMISTRY. [§§33- with a molecule of hydrogen, so that we obtain Ti2xOz as the for- mula of water and O2, as that of the oxygen molecule. 33. The formula for hydrochloric acid, for water and for the molecules of hydrogen, chlorine and oxygen can be fully established, if the values x, y and z are known. These can be ascertained gen- erally in the following way : x must be at least equal to 1 ; if this is the case, the molecule of hydrogen becomes H2. That a smaller number of atoms is impossible is shown by the synthesis of hydro- chloric acid. The vapor densities of a series of hydrogen com- pounds, as compared with that of hydrogen, are then determined, from which we can find their molecular weights, based on the hydrogen molecule as unity. Thereupon these compounds are analyzed and the amount of hydrogen calculated that is repre- sented in the different molecular weights. It will then be found that in no case is the amount less than half of that in a molecule of hydrogen. The following table gives some examples: Substance. Sp. G. (H==l). Quantity of H. Hydrogen chloride 18.25 0.5 Hydrogen bromide 40.5 0.5 Hydrogen sulphide 17 1 Ammonia gas 8.5 1.5 Methane 8 2 Ethylene 14 2 Water 9 0.5 Since, therefore, no compound contains less than half a. molecule of hydrogen, the atomic weight of hydrogen must be half its molecular weight, i.e. the formula of the hydrogen mole- cule is H2. Similarly it is found that the oxygen molecule is O2, that of chlorine CI2; in other words, that .r, y, and z are all equal to 1. The following table illustrates the case of oxygen: Substance. Sp. G. (H = l). Quantity of 0. Oxygen 16 16 Water 9 8 Sulphur dioxide 32 16 Nitric oxide 15 8 Carbon monoxide 14 8 Carbon dioxide 22 16 3-1.] DETERMINIXG MOLECULMi A\D ATOMIC WEIGHTS 47 RULES FOR DETERMINING xMOLEC'ULAR AND ATOMIC WEICIITS. 34. When the atomic weight of oxygen is taken =16, its molecular weight is 2 X 16 = 32. If the specific gravity of another gas based on oxygen = 16 is a, the molecular weight of this gas becomes 2a. The following rule ha.s therefore been prescribed for the determination of the molecular weight. Deteriuinc the vapor density of the compound, based on oxygen= 16, and multiplij the result by J; the product is the molecular weight. For determining the atomic weight the following holds good, according to § 3o; Determine the cumjxisilion of molecular amounts of as many eompounds of the element us possible; the smallest amount of the element that is found in any instance is the atomic weight. AvoGADRo's hypothesis has been confirmed from a physical stand- point. It is at present one of the principal laws of chemistry and physics. Let us briefly examine, among others, the physical arguments in its favor. The molecules of bodies, solids as well as liquids and gases, are in con- stant motion, the intensity of which increases and decreases with the temperature. In diiferent substances at the same temperature there must be a definite relation between the intensities of the molecular mo^•cnlents. This relation has been sui/c'cssfully worked out from the theory in the i-ase of gaseous substances. It has been shown that in all gases at the same tem- perature the mean kinetic energj' of translation of a molecule is the same. The pressure which a gas exerts against the walls of the \'essel is caused by the impact of the molecules. We will call the number of mole- cules in a volume of the gas n, the mass of each molecule m and their mean velocity u. It is then clear that the gas pressure — the nbove explanation of its cause being accepted — must be proportional to n and m. Moreover the pressure must also be proportional to tr, for if the velocity were increased the enclosing walls would receive more impacts from the molecules moving to and fro, and every impact would also become stronger. The gas pressure p is therefore proportional to the product nmu^; the theory says that p = lnmu^, or w=^^^- In this expression mu'' is t\vice the kinetic energy of translation of molecules, which is the same for all gases at the same temperature. If then p is made the same for the different gascs,^^^, or 71, the number of molecules per unit volume, must be the same for all gases. The laws of Boylk, Gay-Lussac and Avogadro (we refer to the expansion law of Gay-Luss.u) can be expressed in a single 48 INORGANIC CHEMISTRY. [§§ si- comprehensive formula, which is worthy of note because of its frequent use in physical chemistry. The laws . of Boyle and Gay-Lussac are represented by the equation PV PV=RT, or ^=R, in which P is the pressure, V the volume and T the absolute tem- perature, of the gas and /2 is a constant which depends on the quantity and tlie nature of the gas under consideration. The value of R, however, becomes the same for all gases, if molecular amounts of them (one mole each) are taken. For, according to AvoGADRo's law, the volume of one mole of every gas is the same under the same pressure and temperature. In the above equa- tion, then, V is constant for all gases and, since we have already made P and T the same in each case, it is evident that R must liave a constant value. In other words, if we deal with molecular amounts, the equation PV = RT becomes a general expression of the la^ivs of Boyle, Gay-Lussac and Avogadko. The value of R may be calculated as follows: Let us consider 1 mole oxygen at 0° and 760 mm. pressure. Since 1 1. oxygen under these conditions weighs 1.4290 g., the volume "T of 1 molis V = ^-^2^ = 22.893 1. = 22393 c.c. If a correction is applied because oxygen does not exactly follow the gas laws of Boyle and Gay-IjUssai', we obtain 22412 c.c. The pressure of 760 mm. mercury corresponds to a pressure of 1013.25 g. per sq. cm., i.e. P= 1013.25. At 0° the absolute temperature is 273° (more strictly 273.09°). Substituting these values in the above expression for R, we obtain PF_ 1013.25X22 412 T 273.09 ' -^^^^^ in c.-g. units. If the pressure is expressed in millimeters of mer- cury, R becomes „ 760x22412 The product PV also represents the external work which is ' lone when a gas under constant pressure P increases its volume \>y V (on being heated, for instance) or when a gas being generated :i:i.] THE REALITY OF MOLECULES AND ATOMS. 49 under the pressure P comes to orcupy a volume 1 '. For, if we sup- pose that the gas is enclosed in a cylinder of 1 sq. cm. transverse section having atone of its ends a piston, the increaseof the volume must cause a veight P to move through V cm. One calorie (gram- calorie) =4 18',)0 gram centimeters. If this is substituted in the equation n' = 8.'U55r, the latter becomes PV = ' 'A\ or, very approximately, PV = '2T. This latter form also is a common one of expressing the combined gas laws. It gives the external work in calories that is done when 1 gram molecule of any given sub- stance is converted into the gaseous state at the absolute tem- perature T. Since 1 gram molecule of a gas has a volume of 22.4 1. at 0° and 760 mm. pressure, 1 c.c. under these same conditions contains -— , or 0.0446, millimoles. THE REALITY OF MOLECULES AND ATOMS AND TIIEIU ABSOLUTE WEIGHT. 3S. The law of Avogadro teaches that equal volumes of all gases contain the same number of molectiles. If we take a gram molecule of every gas, which, as we just saw, has a volume of 22.41 1., it follows at once that there must be the same number of molecules in every case. The number of molecules in 1 gram molecule of any given gas is thus a universal constant; it is often represented by N. The determination of this constant N, i.e., the absolute number of molecules contained in the gram molecule, has been worked out in recent years by several widely different methods, all of which have yielded approximately the same restilt, viz., 70X1022. Back in 1875 van der Waals in his famous treatise on the continuity of the gaseous and the liquid states, calculated the value of N to be between 40 and 90 X IO22, which is of the same order of magnitude as the present more accurate value. All the methods are of a physical character, so that a full description of them is inappropriate here; nevertheless, to show the diversity of these methods, we may mention that the above value of A^ has been obtained from (1) the law of van der Waals; (2) the Browniajb oaovement (the irregular movement which solid parti- 50 JXOJi-GAXIC CHEMISTRY. [§§35- cles, having the dimensions of the order of a micron (;u= 0.001 mm.) or smaller, exhibit when they are suspended in a liquid. It has now been proved that these movements are due to the impacts of the molecules) ; (3) the diffusion velocity of dissolved substances ; (i) the refraction of light in the atmosphere, causing the blue color of the sky; (5) the electric charge of the ions (§266); (6) the life period of radium (§267); (7) the energy of the infra-red spectrum. The significance of these investigations for all departments of natural science is extraordinarily great. When we see that so many wholly independent methods lead to the same absolute numbei- of molecules, 70x10^2, in a gram molecule there is no room left for doubt of the actual existence of molecules. So long as we had only rough and discordant approximations of this number, we could accept the assumption that matter is built out of mole- cules and atoms as an exceedingly useful hypothesis, while yet doubting the real existence of atoms and molecules. The satis- factory proof of their actual existence has put the knowledge of matter in general on a secure foundation. OZONE. 36. As earh' as 17.S5 van iLviiUM observed that when an elec- tric spark passes through oxygen a peculiar " garlic-like " odor is given off, and a. bright mercury surface is at once made dull. ScHoNBEiN investigated this phenomenon more carefully, and found that it is due to the formation of a peculiar substance, which he called ozone. This proved to be oxygen existing in a special condition. The fact that it really consists of nothing but oxygen is shown bj'' its formation from perfectly dry oxj-gen under the influence of electric discharges, e.g., induction sparks. The amount of ozone thus formed is nevertheless small. It is greater when silent discharges are used. As ozone is formed from oxygen by ultraviolet light, the formation of ozone by silent discharges may be caused by the ultraviolet light accompanying them. This is one of the best ways of obtaining ozone, although the maximum j'ield is only ■5.6';7,. However, if the oxygen is cooled by liquid air and then submitted to the silent discharge at a pressure of 100 mm. Hg, it is wholly converted into ozone. Fig. 17 represents an apparatus constructed by Behthelot for the preparation of ozone at the ordinary pressure and tem- perature. The wide tube /, together with the supply-tube d -36.] OZONE. 51 Fig. 17. -Preparatio.n or Ozone. and the exit-tube e, are sunk in a vessel of sulphuric acid, into which the pole of the inductor h is dipped. The other wire a of the latter ends in a tube c, which is slipped down inside / and is almost entirely filled. The silent discharge between the two bodies of sulphuric acid thus passes through a thin layer of oxygen and has a powerful ozoniz- ing effect. Ozone is formed in manj' reactions, such as the slow oxidation of moist phosphorus; also in a small quantitj^ «"hen hydrogen bums in an atmosphere of oxygen. The oxygen that is ob- tained by the electrolysis of dilute sul- phuric acid always contains it. Ozone is also given off by the decomposition of permanganic acid that is set free in the reaction of potassium permanga- nate and concentrated sulphuric acid iff. also § 52). \\ hen oxygen is subjected to a very high temperature (e. g. flame temperature) it is partially converled into ozone, and the more so the higher the temperature (§103). It is necessary, however, to cool down the ozonized gas very rapidly, because the velocity of defomposifon of ozone is verj' great, especially at high temperatures. An instantaneous cooling can be accomplished by directing the flame (of hydrogen, carbon monoxide, acet3'Iene or other gas) upon the surface of liquid air, which has a temperature of — 180°. That the generation of ozone has no con- nection with the combustion, but that it is caused only by the high tem- perature to which the oxygen is raised by the flame, may be proved by the fact that an incandescent platinum wire or Nernst glower, dipped in Uquid air also generates ozone. The formation of ozone is also obser\-ed when a rapid current of dry air or oxygen is allowed to impinge against a hot Nernst glower. When the air contain.s moisture almost no ozone is formed, the product being hydrogen peroxide (§.')4). Physical Properties. — At ordinary temperatui'cs ozone is a gas ; it has a peculiar odor, which is one of the most delicate tests for its presence. One part of ozone can still be detected by its odor in 500,000 parts of air. In tlic liquid state it is indigo-bluo. Ozone boils under normal pressure at —119°. 52 INORGANIC CHEMISTRY. [§§3fr- Chemical Properties. — Ozone is characterized above all by its ability to oxidize vigorously at ordinary temperatures, especially in the presence of moisture. Phosphorus, sulphur, and arsenic are oxidized to phosphoric acid, sulphuric acid, and arsenic acid, respectively, ammonia to nitric acid, and silver and lead to per- oxides; e.g., the metallic surface of silver, especially when heated to above 240°, become blue when ozonized air is directed against it. Iodine is deposited by ozone from a solution of potassium iodide: 2KI + H20 + = 2KOH + 2I. Organic substances are strongly oxidized by ozone, hence no apparatus containing it should have connections of rubber. Dye- stuff solutions, like indigo and litmus, are decolorized (by oxida- tion). Ozone effectively destroys micro-organisms, and is there- fore used successfully in the sterilization of drinking-water. The detection of ozone, especially in quantities too small to be recog- nized by the odor, is a difficult matter because several other oxidizing sub- stances, such as chlorine or bromine in the presence of water, the oxides of nitrogen, hydrogen peroxide and stiU others, give closely analogous reac- tions and furthermore, their smell at high dilutions somewhat resembles that of ozone ; hence it becomes necessary to first prove their absence. The tests for ozone are usually executed by moistening strips of filter-paper with the reagent and dipping them in the gas containing ozone. The reagents used for this purpose are lead sulphide and thallous hydroxide. The strips of paper are first moistened with dilute solutions of the nitrates of these metals and then exposed to hydrogen sulphide and ammonia fumes, re- spectively. Lead sulphide is oxidized by ozone to lead sulphate, thus turn- ing from black to white; thallous hydroxide, which is white, is converted to brown thallic hydroxide. However, these changes of color also occur with the other oxidizing agents mentioned. A strictly characteristic test for ozone is the violet color produced with an acetic acid solution of tetramethyl- ■p-p'- diamido-diphenyl-methane (an organic compound.) Nitrogen dioxide gives a straw-yellow color, chlorine and bromine a dark blue, while hydrogen peroxide produces no coloration at all. Ozone is stable at ordinary temperatures, but is easily changed to oxygen on heating. It is slightly soluble in water. 37. Formula of Ozone. — The foi'mula of ozone has been determined by Ladenburg in the following way. A glass globe with two cocks was first weighed when filled with pure oxygen and then when containing ozonized oxygen. After reducing both weights to the normal temperature and pres- sure the globe in the latter case was found to be a mg. heavier. This increase of weight is due to the replacement 38.] HYDROGEN PEROXIDE. 53 of a certain number of oxygen molecules by the same number of ozone molecules. The volume that the ozone occupies in the gas mixture can be determined by absorbing it in turpentine. Suppose this to be V c.c, when reduced to normal pressure and temperature. The weight of this V. c.c. ozone can be represented by the weight of an equal volume of oxygen + a mg. and must be, therefore, {v X 1.43 + a) mg., 1.43 mg. being the weight of 1 c.c. oxygen at normal pressure and temperature. Hence the weight g oil c.c. ozone is vX1.43 + a ? = • In one of his experiments Ladenbxjrg found a = 16.3 mg. and T = 26.0 c.c, hence ^=2.06 mg. 1 c.c. ozone thus weighs ^J^^ 1.45 times as much as an equal volume of oxygen, or very nearly 1^ times as much. The molecule of ox3-gen being O2, th^t of ozone must be represented by O3. In an oxidation by ozone the volume of the ozoniferous gas remains unchanged. Only the third atom in O3 has oxidizing power, not all three atoms of the molecule. In ozone we have become acquainted with oxj^gen that is dif- ferent from the ordinary' kind. This phenomenon is also seen in other elements; it is called allotropism. HYDROGEN PEROXIDE, H.Oo. 38. This compound is usually prepared by treating barium per- oxide with dilute sulphuric acid: BaOj + H2SO4 = BaS04 + H2O2. Barium Insoluble, peroxide. In a very concentrated slate it can be obtained by direct distillation in vacMoi a mixture of sodium piTOxide and sulphuric acid: N;i ( ), + l\;^< )., =Nu.,S( ),+ IIjOj. Hydrogen peroxide is also formed in many other ways; e.g. together with ozone (§ 36) in the slow oxidation of phosphorus; by the combustion of hydrogen, when the llame is cooled by a piece of ice. The formation of ozone has been often detected when hydrogen in the nascent state comes in contact with oxygen mole-^ cules. We suppose that in the moment just after hydrogen is set 54 INORGANIC CHEMISTRY. {535. free, its atoms have not yet united to form molecules, so that the individual atoms possess unusual chemical activity. This is the general conception of the status nascendi. Thus Traube has observed the following instances of the production of hydrogen peroxide: Zinc fiUngs, when shaken with water and oxygen or air, give hydrogen peroxide since the zinc and the water generate a small quantity of hydrogen, which unites with the oxygen. Palla- dium-hydrogen behaves likewise when brought in contact with water and air. In this case it is the hydrogen released from the palladiimi that unites with the oxygen. Many metals, such as copper, lead and iron, yield hydrogen peroxide on being shaken with air and dilute sulphuric acid, for the same reason as in the case of zinc and water. Finally, the peroxide is formed in the electrolysis of water, when a current of air or, better, oxygen passes over the negative electrode (at which hydrogen is evolved). Hydrogen peroxide is also formed at very high temperatures from steam and oxygen (§ 103); just as in the formation of ozone under the same conditions (§ 36), a rapid cooling is necessary in this case also, else the compound decomposes. The formation of hydrogen peroxide in the combustion of hydrogen has been shown in the following way: A hydrogen flame was allowed to burn at the mouth of a bulb tube containing a little water. By means of a ver)- rapid current of air the flame was blown into the bull), causing a very sudden cooling of the mixture of steam and air. After a time the water in the bulb gave the tests for hydrogen peroxide. As a further analogy to the case of ozone it has been shown that the formation of hj^drogen peroxide has no connection with the combustion, for on directing a fine stream of water upon an incandescent Xerxst glower some hydrogen peroxide is generated in the water. Physical Properties. — In the pure anhydrous condition hydrogen peroxide is a colorless, slightly viscid liquid, having a specific gravity of 1.4:5S4 at 0°, based on water at 4°. (A density calculated on this basis is indicated by ^4°.) It becomes sohd at a low tem- perature and melts at —2° Chemical Properties. — Hydrogen peroxide, when whoUy free from impurities, especially from suspended particles of solid mat- ter, Ls rather stable and can be distilled in vacuo; when impure, it decomposes, however, into water and oxygen, as it also does in dilute solution. In the latter state it is more stable in. the §38.] HYDROGEN PEROXIDE. 55 presence of traces of acid than in the incsciicc of liases. It is an interesting fact that it decomposes rapidly in contact with powdered substances, apparently without acting upon them. Finely divided silver, gold, ]ilatinum (platinum black), and esjiccially manganese dioxide decompose it with effervescence (due to escaping oxygen"). Even rough surfaces have a disturbing effect; P.ruhl oljscrvcd, for instance, that a concentrated solution of hydrogen peroxide evolves oxygen when poured upon ground glass. All these actions must be reKardcd as catalytic accelerations of the ordinarily very slow decomposition of hydrogen peroxide. The effect of hsat is here, as elsewhere, to accelerate the reaction; concentrated preparations, when warmed, often decompose so rapidly as to cause an explosion. The oxidizing action of hydrogen peroxide is an important chemical property. This is always due to the surrender of an oxygen atom, which effects the oxidation, while water remains. Lead sulphide, PbS, is oxidized by a weak solution of hydrogen peroxide to lead sulphate, PbSU4; sulphuretted hydrogen, H2S, is converted into water and free sulphur. Barium, strontium and calcium hydroxides, Ba(0H)2, yr(0H)2 and Ca(0H)2, are pre- cipitated by dilute hydrogen peroxide from their solutions as peroxides of the general formula ^102 -"aq.^ The colorless solu- tion of titanium diordde in dilute sulphuric acid is turned orange-red by hydrogen peroxide — lemon-yellow by traces of it — on account of the formation of yellow trioxide, TiOs. This is a delicate test for hydrogen -peroxide. Other tests are found in the following oxidation reactions: Potassium iodide starch-paste is at once turned blue by hydrogen peroxide in the presence of a little ferrous sulphate, FeS04. The ferrous sulphate carries the active oxygen of the hydrogen per- oxide to the potassium iodide. As a result two atoms of iodine are set free, the ferrou.s .sulphate being oxidized at the same time. According to Manchot a higher oxide of iron is formed in this reaction. A very characteristic reaction is this: Chromic acid solution (H2Cr04), when treated with hydrogen peroxide, is changed to a higher oxide (see § 295) which is blue in aqueous solution and may be taken up by ether if shaken with the latter. This test is, however, less delicate than the two preceding ones. ' Aq. (aqua), a frequently used abbreviation for water of crystallization or hydration. 56 INORGANIC CHEMISTRY. [§§ 38- A third group of chemical effects of hydrogen peroxide depends on its reducing power. When silver oxide is introduced into a solution of hydrogen peroxide, a vigorous evolution of oxygen occurs, water and metallic silver being formed at the same time. Potassium permanganate solution loses its color when mixed with a hydrogen peroxide solution acidulated by sulphuric acid, oxygen being given off rapidly: 2KMn04 + 3H2SO4 + 5H2O2 = K2SO4 + 2MnS04 + 8H2O + 5O2. The brown peroxide of lead, Pb02, is reduced to reddish-yeUow lead oxide, PbO. Ozone and hydrogen peroxide jaeld water and oxygen; when dilute, they are, however, able to exist side by side. There is a test for hydrogen peroxide, depending on its reducing power, which is even more dehcate than those described above. A mixed solution of ferric ctJoride and red prussiate of potash has a red color. On the addition of hydrogen peroxide Prussian blue is precipitated. Traces of the peroxide turn the solution green. The reaction fails in the presence of free acid. The abihty of so powerfully oxidizing a substance as hydrogen peroxide to act also as a reducing-agent can be explained as follows: One of its two ox3'gen atoms must be loosely joined to the mole- cule, since it is easily given up. All the substances which are reduced by hydrogen peroxide, also haxe one loosely held oxygen atom; sUver oxide, potassiiun permanganate, ozone and others give up their oxygen at rather low temperatures. It is therefore possible that the mutual attraction of the oxygen atoms, which tends to make them form oxygen molecules, is stronger than the force by which they are held in hydrogen peroxide on the one band, and the respective oxygen compound on the other. Uses of Hydrogen Peroxide. — The colors of old paintings are often restored by means of it. The darkening of them is due in many cases to the transformation of white lead sulphate, PbS04, to black lead sulphide. The latter is readily oxidized by hydrogen peroxide bacl<; to white lead sulphate. Hydrogen peroxide is also of value in bleaching ivory, silk, feathers, hair, bristles and sponges. It is also important in analysis^ For therapeutic purposes at 30% solution of hydrogen peroxide is pre- pared by Merck which is perfectly pure and is obtained by vacuum dis. il- lation from a more dilute solution. Before use it is strongly diluted. It ■10.] DETERMINATION OF MOLECULAR WEI TIT. 57 has the advantage of not being subject to decompos'tion. The idiifentra- tion of a solution of hydrogen peroxide isgeiiorally cxpi-rssed in the volumes of oxygen that it can evolve; thus, for a 3% solution it is ten volumes. 39. The coDipoyiiion of hydrogen peroxide T.\as established by Thexard as earl}- as ISlS. He first concentrated it in a vacuum and then introduced a weighed amount of it, enclosed in a \vd\, into a graduated barometer-tube over mercury. The vial was then broken and its contents decomposed by heating the tube from without or allowing finely powdered manganese dioxide to rise in the tube. It was thus found that very nearly 17 jiarts of hydrogen peroxide by weight yield s parts of oxygen, water being also formed. t)ne atom of oxygen (16 ])arts by weight) is therefore obtained from .34 parts of hydrogen peroxide, the remaining 18 parts forming water; in other words, hydrogen peroxide is 1 molecule H2O-I-I atom 0. Ihe peroxide therefore contains one atom of oxygen to every li}-dr(igen atom. Its simplest formula (the so-called empirical jormuhi) ia then Hy the electrolysis of pure anhydrous hydrofluoric acid in which potassium fluoride has been dissoh'ed to make the liquid a conductor. The manner in which Moissan accomplished this is interesting. A mixture of about 200 g. anhydrous hydrofluoric acid and 60 g. hydrogen potassium fluoride is introduced into a copper U-tube (Fig. 22) of a capacity of about 300 c.c, which has two lateral exit-tubes. The open ends of the U-tube are closed -with stoppers FF, made of fluor spar and wrapped in very thin sheet platinum. Ihe cyhndrical electrodes tt of platinum-iridium pass through the stoppers and are held in place by the copper screws EE, which fit tightly to the ends of the U-tube, wth the helji of a band of lead P. During the electrolysis the apparatus is kept at the constant tem- perature of -2.3° (by boiling methyl chloride). The free fluorine, which is given off as a gas at the positive electrode, is first passed through a platinum vessel that is cooled by a mixture of solid carbon dioxide and alcohol, in order to condense the acid fumes which were carried 80 IXORGAXIC CHEMISTRY [§§ 52- over with it. The last traces of the acid are removed by conducting the gas through two platinum tubes containing sodium fluoride, which absorbs the hydrofluoric acid. The free fluorine gas was collected by JI DISS AN in a platinum tube, whose two ends were closed with plates of fluor spar so that one could look through. Fig. 22. — Preparation op Fluoeinb by Electrolysis. (After Moissan.) Later Moissan found that perfectly pure fluorine attacks glass but very slowly, so that the gas may be collected in glass vessels. Physical Properties. — Fluorine is a gas with a very pungent odor and a greenish-yellow color, which is somewhat paler than that of chlorine. As a liquid it boils at — 187° and is bright yellow. ■It can be condensed in a glass vessel. When cooled by liquid hydrogen it freezes to a white mass, that melts at —223°. The specific gravity of the gas is 19 (0=16), that of the liquid 1.14 (water =1). Chemical Properties. — Of all the elements now known fluorine has the strongest tendency to form compounds. It combines with hydrogen in the dark at ordinary temperatures in an explo- sive manner. JIoissan demonstrated this with the help of the above apparatus by reversing the electric current while fluorine 53.] FLUORINE. gl was being generated; thus a mixture of hydrogen and fluorine was formed, which at once exploded. As low as -252.5°, solid fluorine unites with liquid hydrogen immediately, producing a flame. Finely divided carbon ignites instantaneously in fluorine gas, form- ing CF4. ^A'ith sulphur, red phosphorus, Hme and other sub- stances huorine reacts vigorously even at — 187°. Fluorine com- bines « ith most metals instantly and violently; it does not unite with oxygen, even when it is heated with the latter to 500° or mixed in the liquid state with liquid oxygen at — 190°. The alkali metals (potassium and sodium) and the alkaline-earth metals (calcium, strontium, barium) take fire in fluorine gas at ordinary temperatures with the formation of fluorides. Finely divided iron glows faintly in it. Copper be- comes covered with a layer of copper fluoride, CuF2, which pro- tects it against farther corrosion; hence the possibihty of employ- ing this metal for fluorine generators. Gold and platinum are not attacked by fluorine, — a rather striking fact, since these metals are acted on by chlorine, which otherwise displays a weaker chemical affinity. Fluorine reacts readily with hydrogen compounds; e.g. water is decomposed by it at ordinary temperatures into hydrofluoric acid and strongly ozonized (as high as 1J:% by volume) oxygen. It sets chlorine free from potassium chloride, forming potassium fluoride. The molecule of gaseous fluorine is expressed bj- the formula F2. Its vapor density being 19, the molecular weight is 38. Inasmuch as no fluorine compound contains less than 19 g. fluorine per gram-molecule, but frecjuently a multiple of this amount, the atomic weight of fluorine becomes 19 and its molecular formula F2. HYDROGEN FLUORIDE, or HYDROFLUORIC ACID, HF. 53. This compound was discovered by Scheele in 1771 upon heating together fluor spar and sulphuric acid : CaFz + H2SO4 = CaS04 + 2HF. This is still the usual method of preparing the substance. A mixture of powdered fluor spar and dilute sulphuric acid is distilled in an apparatus of platinum or lead, since glass is instantly attacked by hydrofluoric acid. The distillate is an aqueous solu- S2 IXOROAXIC CHEMISTRY. [§§53- tion of the acid, which for the above reason must be preserved ia bottles of lead or caoutchouc. ]^y direct s}-nthcsis from its elements (§ 52) hydrofluoric acid may also be obtained. Another method is by the action of hj-dro- gen on a fluorine compound; e.g. silver fluoride, when heated in a current of hydrogen, gives hj'drogen fluoride. Still other methods are l^y the action of fluorine on hydrogen compounds f§ .52) and by the direct decomposition of certain com- pounds, such as hydrogen potassium fluoride, KF-HF, which sphts up on heating into the two fluorides. This last reaction is made use of when anhydrous acid is sought. Physical Properties. — Anhydrous hydrofluoric acid is a color- less liciuid at ordinary temperatures. It boils at 19.5° and solidifies at - 102.5°. Sp. g. (H= 1) = 0.9879 at 15° It has an extremely pungent odor and is very poisonous when inhaled. It is very soluble in water. Chemical Properties. — The aqueous solution of hydrogen fluoride, the "hydrofluoric acid" of commerce, possesses entirely the character of an acid; it evolves hydrogen with most metals, the precious metals, howeA'er, and also lead, being unaffected by it. Tlu' fluorides of the metals are, in general, soluble in water; some, ho\\■p^'er, such as those of copper and lead, dissolve with difficulty, while those of the alkaline earths (Ca, ,Sr and Ba) are insoluble. It is a peculiar characteristic of the alkali fluorides that they are able to cnml^ine with a molecule of the acid, forming double fluorides like that described above, KF-HF. This characteristic is probably due to the fact that in aqueous solu- tion the molecule of hydrofluoric acid is ITjF,. The formation of such dduhlo molecules is often observed for acids (especially organic acids) . It is called association. Thus liquid water consists in all probability of H,02 molecules. The most unportant property of the gas for practical purposes is that it attacks glass (c/. § 193). As a result it finds extensive use in etching glass. Cllass may be etched in two ways — with a solution of the gas or with the gas itself. In the first case the etching is shiny and transparent; in the second dull. The glass object is covered with a coat of wax in v.'hich the figures or letters which one desires etched on the glass may be drawn with a stylus. Then the object is either dipped in dilute 54.] HYDROFLUORIC ACID. ,S3 hydrofluoric acid for a while or set over a leaden dish which contains a mixture of sulphuric acid and calcium fluoride kept slightly warm by a low flame. Only the places where the coating was removed are attacked, so that, when the latter is subsequently dissolved off (by turpentine or alcohol), the etch-figure is visible. A[oiss.\N has proved that glass is also attacked by perfectly uiy hydrofluoric acid gas. The formula oj hydrofluoric acid gas is HI'', which can be deter- mined in exactly the same way as was done for the analogous chlorine and bromine compounds. Compounds of the Halogens with each other. 54 The halogens, or salt-formers, i.e. the elements fluorine, chlorine, bromine and iodine (so-culled because they form salts with metals by direct combination), can unite with each other to foiin rather unstable compounds. In general the most stable of these compounds arc those whose component halogens show the greatest dissimilarity. Iodine unites with fluorine to form a comjjound IF5, which can exi.st even in the gaseous state. There is also a BrF^; chlorine and fluorine, however, do not combine with each other. Chlorine and bromine at low temperatures give an unbroken series of mixed crystals (§ 212, 2), but form no compound. With iodine chlorine gives two compounds, ICl and ICI3. It depends on the quantity of chlorine present as to whether the former or the latter is obtained. ICl is a reddish brown oil that eventually yields crystals meltinp; at 24.7°; it boils at 101.3° Water decomposes it into iodic acid, iodine and hydrogen chloride. It exists m two modifications. ICI3 crystaflizes in long yellow needles and on fusing dis- .sociates almost completely into ICl and Cb. In a small quantity of water it dissolves almost unchanged; but a larger quantity of water decomposes it partiallv into hydrogen chloride and iodic acid. Bromine and iodine give only one compound, Brl, which is considerably dissociated in the liquid as well as .in the gaseous state. Oxygen Compounds of the Halogens. With the exception of fluorine, the halogens are known to form various oxygen compounds, ha\'ing the common property of instability, i.e. of being easily decomposed. IMost of them can combine with water, forming acids. Oxides which show this latter property are called acid anhydrides. The acids which aic thus formed from the halogen oxides contain each but one hydro- gen atom, and this can >)(■ replaced by a metal. Acids contain- ing one hydrogen atom which can lie thus substituted are called monobasic. 84 mORQANIC CHEMISTRY- [§§ .55- HYPOCHLOROUS OXIDE. CHLORINE MONOXIDE, CI2O. 55. This compound can be prepared by passing chlorine over dr}' mercuric oxide at a low temperature: 2HgO + 2CI2 = CI2O + HgO ■ HgCla. Hypochlorous oxide is a brownish-yellow gas at ordinary tem- peratures. It can be condensed by strong cooling to a dark-brown liquid, which boUs at -1-5°. It is an extremely dangerous sub- stance, especially in the liquid state, since the slightest mechanical disturbances make it explode ^-igorously, breaking up into its elements. It is possible to distO. it without decomposition, only when everything with which it comes in contact is entirely free from dust (organic matter). It acts upon sulphur, phosphorus and compounds of carbon with explosive violence. The composition of this compound was determined by Balard in the following way: He introduced 50 vols, of the gas into a tube over mercury and decomposed it by gently warming. He thus obtained a mixture of chlorine and oxygen which occupied somewhat less than 75 vols. After the chlorine was removed by caustic potash, 25 vols, remained, i.e. 50 vols, chlorine were present, the slight difference which was observed being ascribable to the fact that a Httle chlorine had united with the mercury in the tube. 1 vol. hypochlorous oxide yielded therefore 1 vol. chlorine and i vol, oxygen. This indicates the formula CI2O: 2Cl20=2Cl2-l-02. 2 vols. 2 vols. 1 vol. The vapor density of the compound was found to be 3.03 (air = 1) , or 43.63 (0 = 16) . Its molecular weight therefore becomes 87.26, corresponding to the formula CI2O (2C1 = 71; = 16; suin = 87). HYPOCHLOROUS ACID, HCIO. 56. When chlorine monoxide, CI2O, is passed into water, it is ftbsorbed; the solution contains hypochlorous acid: Cl20 + H20=2HC10. 56.] HYPOCHLOROUS ACID. 85 This compound is known only in luiucous solution. Its com- position is studied in its salts. The same aqueous solution can also be obtained by adding finely powdered mei'curir oxide to chlorine water. IlK() + --Hn2 + Hp = HgCl2 + 2G10H. Upon distillation a pure aqueous solution of the acid is obtained. Still another method of preparing the acid solution is to lead chlorine into the solution of a Ijasc, e.g. potassium hydroxide, at' the ordinary tempt^rature, whereupon a salt of hypochlorous acid (hypochlorite) is formed: 2K( )H +CI2 = Kfl +KC10 +1120. By carefully treating the hj'pochlorite with the equivalent amount of nitric acid the hyj^ochlorous acid is set free and can be separated from the salts Ijy distillation. ^^'hen concentrated, the aqueous solution of hypochlorous acid has a golden color. It is unstable; only dilute solutions can be distilled without decomposition. It oxidizes vigorously, breaking up into oxygen and hydr(jchloric acid: 2C'10H = 21-101 + 02. On the addition of hydrochloric acid all the chlorine of both compounds is set free: HC10 + HCl=ri2 + H20. The hypochlorites act just like tlie free acid, since the presence of very weak acids, e.g. the carbonic acid of the air, serves to liberate hypochlorous acid. They are therefore extensively employed as Meaching agents (§ 27). A solution of potassium hypochlorite (cau dv Jarcllc) is used for this purpose, but chloride of lime (" bleaching p(jwdei-," § 2.")S) deserves particular notice. The latter is obtained by treating lime with chlorine at ordinary tem- S6 IXORGAXIC CHEMISTRY. [§§ 56- poratures. The bleaching action of hypochlorous acid is twice as great as that of the chlorine which it contains would be, if the latter were to act in the free state : 2Cl+H20 = 2HCl+0 and 2C10H = 2HCl+20. However, it shoidd be remembered that two atoms of chlorine were necessary to form the one HCIO molecule: 2K0H + Clo = KCl + KCIO + HgO. On shaking an aqueous solution of hypochlorous acid with mercury a brownish-yellow precipitate of mercuric oxychloride, nHgO-HgClj, is formed, which is insoluble in hydrochloric acid. Chlorine water, on the other hand, when shaken with mercury, gives white mercuric chloride, HgClj (subUmate). These reactions enable us to distinguish between the two substances. In a dilute aqueous solution of chlorine we have the following equilib- rium: Cl^ + HjOj^^HCl + HClO, as is shown by the facts that the solution reacts distinctly acid toward litmus and that the hypochlorous acid can be separated from the hydro- chloric acid by distillation. The difference in the action of chlorine water and a solution of hypo- clilorous acid on mercury is due to the fact that in the above equiUbrium the system Cl^ + H^O is by far the predominant one. . CHLORINE DIOXIDE, ClOj. 57. This gas is formed when potassium chlorate, KCIO3, is treated with concentrated sulphuric acid. Chloric acid is at first set free and this decomposes as follows: 3HCIO3 = HCIO4 + 2CIO2 + H2O. Chloric Perchloric acid. acid. Chlorine dioxide is a dark-yellow gas. It can be condensed to a liquid, which boils at 9.9° and solidifies at —79° to a yellow crystalline mass. It has a peculiar odor resembling chlorine and burned sugar. Chlorine dioxide is extremely explosive; warming, jarring or contact with organic substances causes it to explode with vio- lence. Light slowl)' decomposes it. 57.] CHLORIXE DIOXIDE. yj The following experiments give one an idea of the vigor with which it causes oxidation. (1) When finely powdered susar is mixed carefully with potassium chlorate and a drop of concentrated sulphuric acid is added, the whulc mass bursts into flame. The chlorine dioxide set free makes the sus^ar burn at ordinary temperat.uic. (2) Place a few pieces of j-ellow phosphorLis and some crystals of potassium chlorate under water and allow a few drops of concentrat(>d sulphuric acid to flow down on the two substances. The phosphorus at once burns under water \\ith a brilliant light. Cliliirinc dioxide is soluble in water. Such a solution can he easily prepared by floating a little porcelain cup in a lar;;e crystal- lizing-dish with a flat brim ami containing 22U c.c. water, putting into the cup 12 g. potassium chlorate and adding a cooled mixture 01 44 c.c. concentrated sulphuric acid and 11 c.c. water. The cry.stallizing-dish is then covered with a glass plate. The chlorine dioxide evolved dissob-es in the water, forming a yellow solution. When a base is added to a chlorine dioxide solution, a chlorite (§ .5S) and a chlorate are formed: 2K0H + 2C1(J2 = KCIO2 + KCKl,, + H2O. Pot. chlorite Pot. chlorate. This reaction proceeds very slowly in dilute solution. The composition of chlorine dioxide was determined by Gay- Lu.ssAC as follows: He allowed the gas to flow through a capillary tube with three bulbs. I3y heating the part of the tube in front of the bulbs he decomposed the gas, the action lacing non-explo- sive in so narrow a space. Thus there was obtained in the bulbs a mixture of oxygen and chlorine in the same proportions as they are contained in the compound. The chlorine v,-as absorbed by potash and the residual gas (ox^-gen) was passed over into a measuring-tube. The capacity of the bullis being known, it was possible from these data to calculate the volume ratio of oxygen and chlorine. It was found that 100 vols, of the oxide yields 67.1 vols, oxygen and M2.9 vols, chlorine. This ratio is very close to that of 1:2, represented by the formula CIO2: 2f']()2 = ri2 + 2(32. 2 vols. 1 vol 2 vols. This formula is also confirmed by the vapor density, which was found to be .34..) at 10. .5°, while the formula CIO2 demands 3.5..5 + 2X16 __gg-. 88 IXORGAXIC CHEMISTRY. [§§58- CHLOROUS ACID, HCIO.. 58. This acid is unknown in the pure state. Its sodium salt is formed by the action of sodium peroxide solution on a chlorine dioxide solution: 2a02 + \a,Oj = 2Xaa02 + O^. The sUver salt, AgClOj, is a yellow crystalUne powder, as is also the lead salt, Pb(C102)2; they are both difficultly soluble in water, and break up even on warming to 100° in an explosive manner. The anhydride of chlorous acid, corresponding to the formula CI2O3, is not known. CHLORIC ACID, HCIO3. 59. The chlorates of potassium or barium are the usual start- ing-points for the preparation of chloric acid. When dilute sulphuric acid is added to the solution of the barium chlorate, barium sulphate is precipitated and a dilute solution of chloric acid is obtained, which may be filtered off from the sulphate and dried in a vacuum desiccator over concentrated sulphuric acid. In this way a 40% solution of the acid may be obtained. On concentrating it any farther, decomposition takes place, oxygen being evolved and perchloric acid formed. The concentrated acid is a powerful oxidizing agent; w^ood or paper ignites when brought in contact with it. It oxidizes hydrochloric acid, chlo- rine being given off; further sulphuretted hydrogen, sulphurous acid and others, even in dilute solution. The follo'wing reaction is very characteristic of chloric acid. Allien indigo solution is added to a dilute solution of the acid, the former is not decolorized ; however, on the addition of a little sulphurous acid the color dis- appears, since the chloric acid is thereby reduced to lower oxides. The salts are all soluble in water, that of potassium being somewhat difficultly so, however. The com-position of chloric acid was ascertained by Stas from an analysis of silver chlorate. An accurately Aveighed amount of the latter was reduced by a solution of sulphurous acid to silver chloride and this was filtered off and weighed. Since he knew from previous investigations the exact composition of silver chloride, the analysis of the silver chlorate was complete. Stas found thus that silver chlorate consists of 60.] PERCHLORIC ACID. 89 Silver 56. 3948% Chlorine 18.5257% Oxygen 25.0795% Total 100.0000% The atomic weight of silver is 107. NS; that of chlorine 35.46; that of oxygen 16.00. We then find that the ratio of the atoms in this salt is 107.SS -^'••^-' 35.46""-^"'' 16.00 ~^-^^' i.e. very close to 1:1:3, from which it follows that the empirical formula of the salt is AgCUJs, that of the acid itself HCIO3. PERCHLORIC ACID, HCIO4. 60. This compound is obtained by distilling potassium per- chlorate with an excess of sulphuric acid of 96-97.5' ^ in vacuo: KCIO4 + H2SO4 = KHSO4 + HCIO4. Pure perchloric acid boils at 39° under a pressure of 56 mm. Hg, and has a specific gra\-it3' of 2)4-- = 1.764 at 122°. It is a colorless hquid which does not solidify on being cooled with solid carbon dioxide and alcohol (about —80°). It decomposes slowly, taking on a dark color. With water it forms difterent hydrates; the best known of them is the monohydrate, HCIO4H2O, which melts at 50°; with more water a thick oily liquid is formed, similar to concentrated sulphuric acid. Such an acid contains 71.6% HCIO4; it distils without change in composition at 203° and has a specific gravity of 1.82. The dilute solution of the perchloric acid is stable. In the concentrated state perchloric acid is a very strong oxidizing agent. "\Mien a little is dropped on wood or paper, these ignite with explosion. Xery painful flesh-wounds are pro- duced by it. When dilute, it does not, however, I'clcase its oxy- gen nearly so readily as chloric acid. It can, for example, be gently warmed with hydrochloric acid without giving off chlorine, and it is not reduced by sulphurous acid. By these facts and by its \ielding no chlorine dioxide with sulphuric acid it may be distinguished from chloric acid. 90 INORGANIC CHEMISTRY. [§§ (30- The salts of perchloric acid, perchlorates, are all solu- ble in water; that of potassium and especially that of rubidium are, however, very difficultly soluble in cold water. The composition of perchloric acid has been determined, as in the case of chloric acid, by the analysis of a salt, in this instance the potassium salt. A weighed amount of the latter is heated to drive off all the oxygen. The loss in weight indicates the amount of the latter. The analj^sis of the remaining potassiuni chloride, KCl, shows the amounts of potassium and chlorine. From these data it is found, in the same manner as with chloric acid, that the empirical formula of the salt is KCIO4, that of the acid, therefore, HCIO4. Chlorine heptoxide, CljO,, is the oxide corresponding to perchloric acid: 2HC104-I-l20=Cl20,. It may be obtained by slowly adding perchloric acid to phosphorus pent- oxide cooled below — 10° By distillation on a water bath the oxide is obtained as a colorless liquid, which boils at S2°. It is more stable than the other oxides of chlorine ; it neither attacl'cs paper nor acts on sulphur or phosphorus in the cold. OXYGEN COMPOUNDS OF BROMINE. 61. Although no compounds with oxygen alone are known, there are two oxygen acids, viz., hypobromous and bromic. Hypobromous acid, HBrO, can be obtained in the same way as HCIO, namely, by shaking up bromine water and mercuric oxide together. The dilute solution can be distilled in vacuo, and has properties entirely anal- ogous to those of hypochlorous acid. Bromic acid, HBrOg, can be obtained from the barium salt with sul- phuric acid or from the silver salt vnih bromine-water : 5AgBr03 + 3Br2 + mfi = 5AgBr + GHBrOg. Insol. It is also formed when chlorine is passed into bromine-water: Br^ + 501,, + 6H,0 = 2HBr03 + lOHCl. It corresponds in its behavior with chloric acid. !Many reducing-agents, such as hydrogen sulphide and sulphurous acid, are able to extract aU its oxygen. Most of its salts are difficultly soluble in water. When heated, they give up all their oxygen. 02.] (KWGEN COMPOUXDS OF lODlKE. 91 OXYGEN COMPOUNDS OF IODINE. 62. When iodine is iiitrcKlucrcl into a cold dilute solution of caustic potash 01' soda, a colorli^ss liquid is obtaineil, which has other properties when fresh than it has latrr. When freshly i):i'par,il it decolorizes indigo solution and iodine is liberated on the atldition of very weak aeids. Later on these two properties disaiipcar. It is therefore to be supposed that a hypo-iodate KIO is first formed, and that this is changed slowly to KI and KIO3. A( the boiling-point the change takes place almost instantly. Iodine pentoxide, I2OS, is the anhydride of iodic acid, since it can be obtained by heating tins acid to 170°, 2HI03=H2O + l2<»,-„ and yields the same acid when dissoKetl in water. It is a white crystalline substance, which 1 creaks up into its elements at 300°. Iodic acid, HIO3, is prepared by the oxidation of iodine with nitric acid, or, Ijettcr, witli nitrogen pentoxide. 3I2 + IOHNO3 = 6HI(»3 + 10x0 + 2H2(:). Nitric acid. Nitric oxide. Iodic acid is crystalline and easih- soluble in water. It is a power- ful oxidizing-agent . setting free chlorine from hydrochloric acid, for example. 2HI( >,H + lOHCl =I2 + ''''C12 + 6H20. It reacts instantaneously with hydriodic acid, all the iodine of both compounds being precipitated: .■iHI+HI03 = -ll20+6I. The salts of this acid, the i d a t e s , are in general not very- soluble in water; howe\-cr, thuso of the alkah metals dissolve rather easily. On heatmg iodic acid with concentrated sulphuric acid oxygen is evolved and the compound L< >,. iodine dioxide, is formed. This is a lemon-yellow, crystalline powder that breaks up into its elements above 130°. With hot water it reacts quickly to form iodine and iodic acid ; .-.!,(), 4-4H/)=SHI03-M2. Periodic acid, HK )4 +2H2O, is formed by the action of iodine on perchloric acid: HCIO4 + 1 + 21120 = HI()4--'I1,0+C1. It is a colorless .:iystalliiii' solid that is entirely decomposed at 140° into iodine pentoxide, oxygon and water (J 11.")). 92 INORGANIC CHEMISTRY. [§§63, NOMENCLATURE. 63. The system of naming the various halogen oxygen-acids is a general one, which is also used for the acids of other elements. The best-known acid usually has the suffix -ic, e.g. chloric acid, phosphoric acid, sulphuric acid, etc. Acids that contain more oxygen have in addition the prefix per-, thus perchloric acid and persulphuric acid. Acids containing less oxygen have the suffix -ous, e.g. chlorous acid, sulphurous acid, phosphorous acid, etc. Those which contain stiU less oxygen have the suffix -ous and also the prefix hypo-, e.g. hypochlorous acid, hyposul- phurous acid and hypophosphorous acid. The names in use in pharmaceutical chemistry (see the National Phar- macopoeia) follow the Latin. Thus we have Acidum sitlphuriciim (sul- phuric acid) and Acidum sulphurosum (sulphurous acid). The names of the salts of the best-known (-ic) acids end in -ate, e.g. potassium chlorate, -sulphate, -phosphate. The salts of the -ous acids have the ending -ite, as potassium chlorite, -sulphite, -phosphite. The salts of hypo- -ous acids are called hypo- -ites; thus sodium hypochlorite, -hyposulphite, -hypophosphite. The names of the anhydrides correspond to those of their acids. In naming oxides the name of the element with or without the ending -ic is used, unless there is more than one oxide. \Miere there are two oxides, the name of the one with the more oxygen ends in -ic, that of the other in -ous, e.g. mercuric oxide, arsenic oxide, mercurous oxide, arsenious oxide. An oxide with less oxygon than the -ous compound is given the prefix hypo-, and one with more than the -ic oxide the prefix per-, as in the case of acids, thus hypochlorous oxide, lead peroxide. In some cases, for the sake of euphony, the suffix is added to the Latin instead of the English stem, as cuprous, ferric, etc. For historifiil reasons many names now in use do not conform to this system. In some instances the oxide first discovered took the suffix -ic, and those subsequently discovered were named accordingly, as in the case of the nitrogen oxides (§ 119). It is not uncommon to speak of oxides of the general formula M2O3 as sesquioxid.es. 64.] SUMMARY OF THE HALOGEN GROUP. 93 A much more rational system is to indicate the number of atoms of oxygen by the Latin or Greek numeral, e.g. chlorine protoxide, or monoxide, iodine pentoxide, etc. SUMMARY OF THE HALOGEN GROUP. 64. It is evident from the foregoing descriptions that the properties of the halogens and their compounds possess great similarity among themselves. A closer study reveals the fact that the increase of atomic weight is accompanied by a gradual change of physical and chemical properties. For example, let us notice the physical properties: ci. Br. I. Atomic weight Melting-point Boiling-point Sp. g. (liquid or solid). Color 19 -223° -187° . 14aiquid) pale green- ish-yellow. 35.46 -102 - 33 1.33 greenish- yellow 79.92 - 7 -t-63 3.18 brown 126. 'J2 -1-113 -1-200 4.97 violet- black It is seen that the values of the physical constants increase on the whole with the atomic weight. The purely metaUoid char- acter of the first three is also found in iodine, although in the case of the latter an external characteristic of metals, viz., metalUc lustre, is at once noticeable. The affinSty for hydrogen decreases as the atomic weight increases. 'We saw that fluorine combines with this element e^•en in the dark and at very low temperatures in an explosive manner; iodine unites with hydrogen directly only at a high temperature and the compound is easily decomposed by heat. Inversely, the ox}'gen compounds are the more stable the higher the atomic weight of the halogen. W'liile a halogen of low atomic weight displaces a halogen of higher atomic weight from its hydroger compound, e.g. HI-FC1=HC1 + I, the halogen with higher atomic weight can on the other hand replace one with lower in its ox}'gen compounds, setting that other one free: KC103 + I = KI03 + C1. 94 IXOKGAXIC CHEMHri'RY [§ G5^ ELECTROLYTIC DISSOCIATION. 65. In § 30 it was stated that the properties of an aqueous solution of hydrogen chloride differ widely from those of the dry gas. It was also stated there that many other substances undergo a similar change of properties when they are dissoh'ed in water. We may now consider the nature of this change. If we investigate the freezing-point depression of the aciueous solution of an acid, base or salt of known concentration, we find that the depression does not correspond with that calculated from the accepted molecular weight (§ 43). The freezing-point depres- sion and the boiling-point elevation are both greater than they should be. A 1% sodium chloride solution would, for example, 19 be expected to show a depression of p—y = 0.325°, the molecular depression for water (§ 43) being 19, i.e. .Ll/=19, and the molec- ular weight of sodiiun chloride 58.5. In reality, however, the depression is found to be L9 times as great, namely, 0.617°. As the osmotic pressure is proportional to the freezing-point depres- sion (§ 42), it must also be greater than the calculated amount. The fact at once occurs to us that gases, to which dissolved substances have been found to show close analog}^, also exhibit a similar phenomenon. In numerous instances the pressure exerted by a definite weight of gas occupying a definite volume at a definite temperature is greater than the calculation indi- cates. — This is but another phase of the observation that the vapor density of some gases is abnormally low at certain tem- peratures (§ 47). — This is explained by assuming a breaking up of the gas molecules; the number of particles free to move about is thus increased and accordingly the pressure becomes greater. This phenomenon is known as dissociation (§ 49). In the case of abnormal osmotic pressure we are led to a similar explanation by assuming that the molecules are split up into several independent particles. A difficulty^ arises, however, when we try to conceive the nature of this division. In solu- tions of salts in water it would be possible to assume a hydrolytic separation, i.e. into free base and free acid (p. 104), which would necessarily be complete in dilute solutions of the salts of strong acids and bases, inasmuch as the osmotic pressure of such solutions in concentrations of tV normal (§ 93), for instance. 65.] ELECTROLYTIC D W.sYX 'M T/O.V. 9,1 amounts io double the calculated piossui-o. There are, however, serious objections to such a hypothesis. In the first place, it has never yet been possible to separate such a solution Ijy diffusion into the free base and free acid which it is sujjposc^d to contain. A second and still more serious objection is that an acid or base in an aqueous solution by itself exerts an osmotic pressure greater than that calculated. liere, however, hydrolytic dissociation is impossible. The eiuestion as to the real nature of the di\'ision has found its answer in a consideration of the relation which exists between the abnormal osmotic pressure and the transmission of the electric current. Akrhexius observed that only those substances which conduct the electric current in aqueous solution, namely, acids, bases and salts, show the above-mentioned abnormalities in osmotic pressure. When these substances are dissolved in another liquid than water, the resulting solution is a non-conductor, but at the same time its osmotic pressure again assumes the normal. These facts enable us to perceive the connection between the apparently disconnected phenomena of abnormal osmotic ])rcssure and elec- trolytic conduction. In order to understand this relation it is necessary to know the usual explanatiim of electrolytic conduction. Let us take hydrochloric acid as an example. Perfectly dry hydrochloric acid gas is a non-conductor, as is also perfectl}- pure water. However, when the gas is dissolved in water, a solution is obtained which transmits electricity verj' well. Evidently a certain reaction must have resulted from the mixing of the water and the In-drogen chloride. We were led to surmise this above (§ 29), when it was founrl that this gas solution does not oljcy Henry's law. Since during the transmission of the current the hydrogen chloride is broken up into hydrogen and chlorine while the water remains unchanged, it must be assumed that the hydrogen chloride molecules are the ones which have undergone a change. The phenomena of elect rol^-tic conduction now find their com- plete explanation in the assumption that the change which the hvdrochloric acid unflcrwcnt consisted in a separation of its mole- cules into electrically chai'^ed atoms (ion-^) (Jj '2Cu). This separa- tion may have been cdmplcte or partial, tlie extent depending iipiMi the concentration among other things. \Mien a current passes through the solution, the negatively charged chlorine ions (the- 96 INORGANIC CHEMISTRY. [§65- anions) are drawn toward the positive electrode (anode); they become electrically neutral on contact with the latter and escape from the liquid. Similarly the positively charged hydrogen ions (cations) wander toward the negative electrode (cathode). In this way conduction goes on, the undivided molecules having no part in it. This division of the molecules is known as electro- lytic dissociation, or ionization. The existence of free ions in the solution of an electrolyte is demonstrated by Ostwald in the following manner. The tube ahcd, Fig. 23, is nearly filled with dilute sulphuric acid. The r •am m m ^-. Fig. 23. narrowed portion be is about 40 cm. long. A rod of amalgamated zinc is lowered into a to serve as the positive electrode, while a platinum wire is fused into d a,t p for a negative electrode. If connection is made with a batterj' of ten accumulators, there is an immediate evolution of hydrogen at p. The passage of the current through the liquid results in the formation of zinc sul- phate around the bar in a: Zn + H2SO4 = ZnS04 + Ha. Now if this hydrogen has to pass through he to p, it must cover the 40 cm. in a very brief space of time. However, it has been shown both by investigations which cannot be described here and by calculus that this migration would take many hours. The hydrogen appearing at p as soon as the circuit is closed cannot, therefore, come from a; the most natural explanation is to sup- pose that there are already free ions in the neighborhood of p and that they are discharged by the current and given off from the liquid as free hydrogen. ToLMAN has shown that, when a long tube containing a solu- tion of an alkali iodide is rotated as the spoke of a wheel at 3000- 65.] ELECTROLYTIC DISSOCIATION. 97 5000 revolutions per minute, the outer end becomes negative with respect to the inner end. The solution must therefore contain positive and negative components which can move independently of each other. The iodine ions, being much heavier than the alkali ions, would naturally accumulate at the outer end. In order that this hypothesis of dissociation into ions may also account for abnormal osmotic pressure, it must be assumed that the ions are independent particles, jUst as free to move as the molecules are supposetl to be. The number of freely moving particles in the same volume is thus increased. Hence, whether the amount of ionization is calculated from the electrical con- ductivity or from the osmotic pressure, the result should be the same according to the above hypothesis. This is found to be the case. Supposing that every molecule yield.? n ions by the dissociation and that the dissociated portion of the whole number of molecules is y, the number of freely moving particles is The osmotic pressure must therefore be [l + r(n—l)] times as great as in the case of imdivided molecules. If this pressure p is p,, in the latter case, then P=pll + rin -1)1 wherefore ^ {n-i)p, ^^> From the electrical conductivity we are able to find the value of t in the following way: As the dilution becomes greater, the molecular con- ducliidty increases. By this term we mean the specific conductivity of the solution multiplied by the number of litcis in A\hich a gram-molecule of the substance is dissolved As the dilution is gradually increased, the molecular conductivity approaches a definite limit. Since the conductivity is only due to the dissociated molecules, it may be assumed that, when this limit is reached, all the molecules are broken up into ions. If the molecular conductivity for infinite dilution is represented by X^ and that for the dilution v (1 gram-molecule in v Mters) by A„, it is evident that r=^ (2) The following table shows the agreement of the values calculated 98 INORGANIC CHEMISTRY. [§§65- by the two methods. The values opposite ;-„ were calculated from the observed freezing-point depressions and those opposite ye from the con- ductivities of the salt solutions. The concentration throughout is 1 g. per liter. KCl. NHjCl. KI. NaNOa. 7 , r 0.R2 0.86 0.88 0.84 0.90 0.92 0.82 0.82 66. Ionic Eqidlihriiim. — In a case of electrolytic dissociation we have an equilibrium to deal with, namely, that between the un- dissociated molecules on the one hand and the ions on the other. In the case of a monobasic acid this equilibrium may be repre- sented by AHiz±A' + H-, ffhere A' is the acid radical (anion) and H" the hydrogen ion (cation). For a base we have M0H5=tM'+0H'. We may apply here the equilibrium equation deduced in § 49.. Given a gram-molecules of AH per unit-volume, of which x are divided into two ions each, then the equilibrium is represented by a—x=Kx^. From this equation it necessarily follows that the dissociation is diminished by the introduction into the solution of a substance with like ions (just as the addition of hydrogen or iodine reduces the dissociation of hydrogen iodide gas, § 50). This effect may be produced on a salt in solution by the addition to the solution of a salt of the same base or a salt of the same acid. The equation then becomes a — x=K-x(x + p), p being the concentration of the added ion. K can only remain constant provided x diminishes. It also follows that the degree of dissociation depends on the con- centration. If the latter be increased n-fold, we have from the above equation n(a—x) =Kn~x^, or (a — a;) =i?-n-a;^. 156.] ELECTROLYTIC DISSOCIATION. 99 If n is>l, X must diminisli, i.e. the ionization decreases with increasing conecntration. If n isms (Fig. 26). Of the various methods Inr I he determination of the transition 108 LYORGAXIC CHEMISTRY. [§§ 68- point a convenient one is the dilatometric method. It is based on the change of volume (specific gravity) which a body usually undergoes on passing through the transition point. In measuring this a dilatometer is used, an instrument which may be compared to a thermometer of very large dimensions. After rhombic sulphur, for example, has been placed in the dilatometer the latter is filled with a chemically indifferent liquid (kerosene, linseed oU) and put in a large water bath; the temperature is then slowly raised. Below the transition point the volume is seen to slowly and steadily increase "nith the temperature on account of expansion; as soon as the temperature gets a trifle above 95.4°, however, a marked increase of volume is observed, even if the temperature be maintained constant; thereupon expansion again proceeds gradually, as before, if the tem- perature is allowed to rise. The marked change of volume indicates the transition of the rhombic sulphur into the mono- clinic modification. "STABLE," "METASTABLE," AND "LABILE." These terms are coming to be so frequently used in chemistrj' that they need to be distinctively defined. They are borrowed from mechanics, for which reason it is desirable that they be employed in chemistry in the same sense as in mechanics. In the latter an equilibrium is called labile (apt to slip) when the slightest displacement suffices to transpose the body into a new position of equilibrium. An example is afforded by a cone standing on its apex. It cannot recover from even the slightest disturbance, but gets further and further from the vertical position and finally tumbles over. A labile condition .is thus really a limiting case which cannot actually be realized; not even for the cone, though its apex were a mathematical point resting on an absolutely hard surface. All actually occurring equilibria are stable; but there can be different degrees of stability. When a material cone is stood on its apex its equilibrium has very little stability. On the contrary 71.] THE PHASE RULE OF GIBBS. 109 a beam resting on the ground with its largest surface down represents a very stable equilibrium. However, if the beam is stood up on end, its equilibrium becomes less stable. Like the cone resting on its apex, the beam -wlU. have a tendency to go over into a more stable condition. In mechanics there is no need of giving such conditions a special name, but in thermodynamics and chemistry they call for special designa- tion, and the term applied to them is metastahle. We have an example in undercooled water, something that can exist, but has the tendency to go over into a more stable condition, namely, into ice. Therefore undercooled water is said to be metastable. It follows from the above that expressions such as, " labile compounds " (e.g., for CIO2), or " the substance exists in a labile condition," are to be avoided. The word " labile " should be replaced by " metastable." Strictly labUe condi- tions are impossible; nevertheless they may possess great theoretical interest, such, for instance, as the case of the continuous transformation of liquid into gas below the critical temperature, which, though it cannot be carried out, has led to very valuable theoretical considerations in the hands of \'an der Waals and others, as may be seen in the larger text-books of physics. THE PHASE RULE OF GIBBS. 71. The -phase rule treats of the equilibrium in heterogeneous systems, i.e. systems that can 1)0 separated mechanically into unlike parts. A saturated salt solution in contact with solid salt is a heterogeneous system, for it consists of sohd salt, the solution and vapor; that is, of three parts, mechanically sepa- rable. Each of the.se parts in itself is homogeneous, i.e. each part has the same composition throughout. A gas mixture is always homogeneous, as is also a solution. These homogeneous parts, separated by hmiting surfaces and of which a heterogeneous system is made up, are called by Gibbs 110 IXURGAXIC CIIEVISTHY. [§§ 71- phases. Water and its vapor constitute two phases; ice, water and steam three phases. A heterogeneous system can never have more than one gaseous phase, because all gases are miscible in all proportions; it may, however, consist of different liquid phases, in case it contains immiscible liquids. The number of these liquid phases is seldom more than two; that of the solid phases is un- limited. A further conception, introduced by Gibbs, is that of the components of a system. If the system is composed of only one element, then this element is the only component. Systems made up of one compound have in most cases this compound as the only component. A system consisting of molten and gaseous sulphur, or of water and steam, has but one component. In this case all phases have the same composition. There are systems, however, in which this is not the case; viz. systems that are made up of more than one component. We select as the components those 'compounds of which the smallest number is necessary to form the different phases. The choice of such compounds may be somewhat arbitrary but their number is always fully defined. Let us consider, for example, the system Glauber's salt- water. This salt has the composition, Na2S04- IOH2O. In order to determine the composition of the phases that are possible here (solid salt, solution, vapor) it is best to choose Na2S04 and H2O as components. We might indeed take Glauber's salt itself as one of the components; but then, in case the solid phase was the anhydrous salt, it would be necessary to regard water as a negative part of it, which is undesirable. Sulphuric acid and sodium hydroxide are not components, because they do not occur independently in any phase, neither are they found in any other relation in the phases than as a part of the salt itself. It is a property of the components that they can occur in some of the phases in varying proportions (e.g. in saturated and unsaturated solutions). Let us now take, for example, a saturated solution of salt and water in a vessel that is closed with a movable piston. Under this solution let there be a httle solid salt, above it the vapor of 71.] THE PHASE RULE OF GUiBS. Ill the solution. The system consists manifestly of two substances and three phases. So long as the temperature remains constant, the vapor of the salt solution possesses a definite tension. If we increase the volume by raising the piston, a definite amount of water will evap- orate; since the solution was saturated, the result will be that a little salt will be deposited; in the end the quantities of vapor, solution and salt will therefore have altered, but the composition of each phase will remain the same. The tension, and hence also the concentration, of the vapor remain unchanged, since the temperature is constant; there is likewise no change in the con- centration of the salt solution. The same is true in case the volume be diminished. It therefore follows that the equilibrium of such a system is independent of the quantities of the various phases. It is dependent only on the temperature chosen; if this is constant, the whole system is defined. Or, if we should select an arbitrary value for the composition, the temperature and pressure would be fully defined. It is therefore evident that the system is completelj' defined as soon as one of these magnitudes is arbitrarily chosen. The system has onlj' one degree of freedom. Such an equilibrium possesses the following characteristic: .\t a given constant temperature the vapor pres- sure is definite. Under an e\-en slightly greater or smaller pressure one of the phases will gradualh' and completely dis- appear, provided the temperature remains constant. On in- creasing the pressure the gaseous phase wholly condenses, so that only solution and salt remain. A decrease of pressure results in the complete evaporation of the solution, vapor and salt only being left. The same is true when the pressure remains constant and the temperature varies. An entirely different behavior is shown by a system made up of an unsaturated salt solution and its vapor. At a constant temperature and a definite position of the piston the vapor tension has a definite value, as in the former case. If, however, the volume of vapor be changed, the tension will correspondinglj vary, for, if the volume be increased, for example, more water will evaporate, the solution will become more concentrated and the vapor tension of course lessen. Tlierefore for every definite 112 INORGANIC CHEMISTRY. [§ 71- temperature there are not simply one but infinitely many pres- sures under which this system can be in equilibrium. The result is that the shghtest change of volume or pressure does not necessitate the disappearance of one of the phases. Two magnitudes may be chosen arbitrarily before the system is fully defined; it has two degress of freedom. It is evident in this example that the number of degrees of freedom increases by one when the number of phases decreases by one. The phase rule expresses a relation between the numbers of the components S, the phases P, and the degrees of freedom F. It is of the following form: F + P = S + 2, or, in words, The sum of the number of the degrees of freedom and the number of the phases of a system exceeds the number of com- ponents by two. Let us apply the phase rule in the first place to water, a system of one component; the sum of the degrees of freedom and the phases must therefore be three. In the following graphic representation. Fig. 27, the tempera- Solid Liquid I Gaseous 0' Fig. 27. tures, t, are plotted as abscissae, the pressures, P, as ordinates. Let us first consider liquid water above 0°. The number of the phases is two (liquid and vapor); the system has therefore only one degree of freedom, or, as we say, it is univariant. To every temperature there corresponds a definite vapor tension. The ordinates of every point in the line OB indicate these vapor ten- 71.] THE PHASE RULE OF GIBBS. US sions. If the pressure at a certain temperature were greater than that indicated by the ordinate, the gaseous phase would com- pletely disappear. The line OB therefore represents the boundary between the liquid and gaseous phases for the various temperatures and pressures. Every point in the area COB represents the liquid, every point in AOB the gaseous, phase. Only the points of the line OB iiulieate the temperatures and corresponding pressures, at which both phases are coexistent. The hne OB therefore ends on one side at 0°; its other end is at the critical temperature, since at this point vapor and liquid become identical. — Let us now allow the temperature to fall below 0°. The liquid phase dis- appears and ice takes its place. The system remains unvariant, however, for the number of phases is unchanged. The ordinates of the points on the line OA again give the vapor tensions of ice for different temperatures. For the same reason as above 0.4 is the boundary line between the solid and gaseous phases. Only along this line are the two coexistent. The line OA extends to the absolute zero, since the gaseous phase then disappears. The melting-point of ice depends somewhat on the pressure, being lowered 1d}' increasing pressure 0.0075° per atmosphere. Both phases, ice and water, will therefore be coexistent along the line OC, which shows a A-er}- considerable rise of pressure for a very slight fall of temperature. In this case also a change of pressm'e at a constant temperature, or the re^-erse, involves the complete disappearance of one of the phases. The line OC will end at a point w-here the liquid and solid phases become identical, i.e. where the whole system turns to a homogeneous amorphous mass. The location of this point has not yet been ascertained. The point (about +0°.01), the melting-point of ice at the pressure of the vapor, is, according to the above mode of repre- sentation, the point of intersection of the three lines which sepa- rate the phases and along which two phases are coexistent. It is called a triple point. Only in this point is it jiossible for the three phases to exist side by side; it is the common point of the areas which represent regions of the three phases. When three phases are coexistent the .system has no degree of freedom; it has become non-variant. In the case of sulphur we have one substance and four possible phases: rhombic, raonoclinic, liquid, gaseous. Fig. 27 makes 114 INORGANIC CHEMISTRY. [§§71- plain the relation between these phases. Below 95.4° sulphur is. rhombic; the two phases are rhombic sulphur and vapor. The line 0.1 forms the boundary between the two regions. At 1)5.4° the rhombic phase passes into the monoclinic phase. The ordi- nates of the Une OB represent the ^-ajjor pressure of monoclinic sulphur at the temperatures 95.4°-120° The two crystalhzable phases can exist side by side at the point (the transition point). According to researches l)v Reicheb this transition point depends on the pressure; an increase of pressure of one atmosphere raises it about 0.05°. The boundary between the crystallized phases is therefore furnished by a line OC, which sIkjws that a \-ery slight rise of temperature is followed Ijy a ^•ery considerable increase of pressure. At we have therefore a triple point, i.e. a point com- mon to both crystaUized phases and the gaseous phase. At B, the melting-point of monoclinic sulphur, there is a sec ond triple point, wliich is wholl}' analogous to the melting-point of ice. linally, it should also l^c noted that the line BC, ^A'hich separates the liquid and the solid phases, must indicate a rise of melting-point for an increase of pressure, since sulphur melts higher the greater the pressure. The lines OC and BC are not parallel but intersect, according to Tajimann's experiments, at 151° and 12C1 atmos- pheres. .'\s the sum of the phases and degrees of freedom is also three with sulphur, the phase rule indicates that all four phases cannot exist in the presence of each c)ther at the same time, not even when the system has become non-variant. At the triple point neither the temperature nor the pressure can be changed without altering the kind of equilibrium. Here the system is 7ion-variaid. Along the lines 0.1, OB and OC it is univariant. When the state of the system is rejjresented by a point within one of the areas it is divaiiant, consisting then of only one phase. In the succeeding chapters we shall have occasion to concern ourselves with systems of more than one component. 72.] HYDROGEN SULPHIDE. 115 HYDROGEN SULPHIDE, SULPHURETTED HYDROGEN, H^S. 72. This gas occurs in nature chiefly in \iilcanic regions. Cer- tain mineral waters, especially the so-called "sulphur springs," contain it. It is also found as a putrefactive product of organic bodies. Hydrogen sulphide can be obtained from its elements by synthesis. They unite almost completely when heated together for a long time (about 16S hours) at 310°. It can also be obtained Ijy the action of hydrogen on sulphur compounds, as well as b}^ the action of sulphur on compounds of hydrogen; the reduction of sil\'er sulphide, Ag2S, with hydrogen at high temperatures illustrates the former case, while the boiling of turpentine oil with sulphur is an example of the latter. None of the abo\-e methods is adapted to the preparation of the gas in the laboratory. For this purpose the interaction of a sulphide with a hydrogen compound is employed, iron sulphide and dilute acids being generally used: FoS + 2HC1 = FeClo + HoS . In order to have sulphuretted hydrogen always at hand, it being in con- stant demand in analytical work (c/. § 73), a very convenient apparatus was devised by Kipp, which can be used for the generation (at ordinary temperatures) of other gases as well. Its construction is shown in the figure (see next page) . The lower globe is joined to the basal portion by a narrow neck, while the upper globe tapers into a long tube, which fits tightlj' into the lower globe and extends nearly to the bottom of the generator without com- pletelv filling the neck. The iron sulphide is put into the middle portion and the dilute acid is poured into the upper portion, the stopcock remain- ing open. As soon as the basal part is filled with the acid the cock is closed and the top part is half filled with more acid. When the cock is opened the liquid sinks in the top part and rises into the middle i)ortion, where it reacts with the iron sulphide to produce hydrogen sulphide, which escapes through the cock. On closinp; the latter the gas continues to be evolved till it forces the liquid back out of the part containing the iron sulphide. The reaction thus ceases automatically and the generator is ready at any time to supply new quantities of gas on opcninp; the cock, till either acid or sulphide is exhausted. The spent acid can he let out through a stoppered opening near the bottom. On account of the free iron usually present in iron sulphide, the gas 116 INORGANIC CHEMISTRY. [§§72- prepared in this manner contains some hydrogen. Perfectly pure hydrogen sulphide is obtained by warming antimony sulphide, SbjSg, with concentrated hydrochloric acid. Fig. 2!). — Kipp Generator. Physical Properties. — Hydrogen sulphide is a colorless gas of disagreeable odor, when diluted reminding one of rotten eggs. Under a pressure of about 17 atmospheres it becomes liqiiid at ordinary temperatures; liquid hydrogen sulphide boils at —61.8'* and freezes at -85°. 1 1. H^S gas weighs 1.5392 g. at 0° and 760 mm, pressure. The gas is rather soluble in water, 1 vol. water dissolving 4.37 vols. H,S at 0° ("hydrogen sulphide water "). Chemical Properties. — Hydrogen sulphide is combustible and yields on combustion either sulphur dioxide and water or water and sulphur, according to the air supply: H,S + 30 = H20- ■so. H,S + = H,0 + S. In aqueous solution it is slowly oxidized by the oxygen of the air, sulphur being set free; this decomposition is aided by light. In order to preserve hydrogen sulphide water, it must be pre- pared from boiled (air-free) water and put into a dark bottle, filled entirely and closed air-tight. The latter condition is best met by placing the bottle, stopper downwards, in a glass of water. It is poisonous; as an antidote verA' dilute chlorine may be inhaled. 73.] HYDROGEN SULPHIDE. 117 Hydrogen sulphide is a powerful reducing-agent. Eromine water and iodine solution are decolorized by it with separation of sulphur (§§ -15 and 48). X'arious oxygen compounds are transformed by hydrogen sul- phide into compounds with less oxygen, e.g. chromic acid is reduced in acid solution to a chromic salt (§ 2',)5). Fuming nitric acid acts so ^•igorously that a sliglit explosion occurs. "When hydrogen sidpliidc is passed over lead dioxide, the gas ignites, while the oxide is reduced. Concentrated sulphuric acid, H2SO4, is also reduced; hence it cannot be used for drying the gas. Hydrogen sulphide possesses the character of a weak acid; when it is passed over zinc, copper, tin or lead, the corresponding sulphides are formed and hydrogen is set free. Composition of Hydrogen Sulphide. — A'^'hcn a bit of tin is heated in dry hydrogen sulphide — in a tube over mercury — tin sulphide and h^'drogen are formed. After coohng it ls seen that the volume of hydrogen is just as great as that of the hydrogen sulphide. The same result is obtained when a platinum wire is heated to redness (by an electric current) in the dry gas, causing the latter to break up into its elements. Since the hydrogen molecule is H2, there must be two atoms of hydrogen in the hydrogen sulphide molecule. Now the specific gravity of hydro- gen sulphide has been found to be 1.1912 for air = l, or 17.15 for ^^^^^' ^ , „ _,, 1 1 ii Fig. 30. — Decomposition or H.S. = 16. The gram-molecule there- fore weighs 34.30 g., and, as it contains 2 g. hydrogen, there remains for sulphur 32.3 g. This figure is A'cr}' close to the atomic w(-i;;lit of sulphur, hence there can only be one atom of sulphur present in the molecule of hydrogen sulphide. "We thus conclude that the formula is H^S. 73. Use of Ili/drofjcn Sulphide in Analifsis. — H}'drogen sulphide finds extensive use in cjualitative analj-.sis. A large number of metals are precipitatcfl by it from ackl solutions as sulphides, viz., gold, platinum, arsenic, antimony, tin, silver, mercury, lead, bis- muth, copper and cadmium, and also certain rare olcinciits. Some of these sulphides have a characteristic color, e.g. the oran.iic- red antimony sulphide, Sb2.'-^3, the yellow cadmium sulphide. lis INOROAXIC CHEMISTRY. [§73 CdS, the brown stannous sulphide, SnS, the yellow stannic sulphide, Sn82, and the yellow sulphides of arsenic, AS2S3 and AS2S5. The rest of the sulphides named are black. Other metals, such as nickel, cobalt, iron, manganese, zinc, chromium, aluminium, etc., are not precipitated by hydrogen sulphide from acid solution but are precipitated by ammonium sulphide. Still other metals, such as barium, strontium, calcium, magnesium, and the alkalies, are not precipitated from their solutions even by ammonium sul- phide, so that we therefore possess in sulphuretted hydrogen and its ammonium compound a means of separating these elements. An answer to the question, why some elements are precipitated from acid solution by hydrogen sulphide and others are not, is furnished by the ionic theory. Let us take, for example, a dilute solution of copper sulphate, into which hydrogen sulphide is being passed. Copper sulphate is almost entirely ionized, hydrogen sulphide only to a very small degree (3). We therefore have in the solution : Cu-+ S04"+2m-+ (JS" + (l-o)Ii2S, the cations being represented by a point and the anions by a line above and to the right, and the number of these points or lines indicating the ionic valence (§ 76). Some of the copper ions and sulphur ions will then unite to form undissociated molecules, CuS, which are only slightly soluble in water and are therefore precipitated. As S-ions thus disappear, the equilibrium between hydrogen sulphide and its ions is dis- turbed; new H2S molecules are then split up into ions, so that there are again S-ions present, which can unite with copper, and so on. The action proceeds according to the equation: CUSO4 + H2S = CuS + H2SO4, Insol. or, if only the ions which take part in it are represented: Cu-+ S" = CuS. This takes place quantitatively if the copper solution is dilute and no considerable amount of anj- strong free acid was added. However, if these conditions are not fulfilled and, as a result, the 73.] HYDh'OCKX SVLPUIDE. 119 concentration of the hydri)!ien ions is rather high, the presence of these ions reduces the ionization of IloS so much (§66) that no precipitate can ho formed. The apphcation of the mass- action law to the case is \ery simple. Copper sulphide, when in contact with water, dissolves to an cxti-cmely small extent; in this solution we have the equilibrium: Cu-+ S"^CuS. If the concentrations of the two ions are a and h, and that of the undissociated copper sulphide is c, we have the equation ab = k- c, k being a constant for a fixed temperature (§ 66). The product ab has a definite value for every saturated solu- tion (since c is definite). This ^-alue is known as the solubility product of the sul^stance in question. If in any case the product ab is less than this value, none of the substance can separate out, because the solution will then be unsaturated; if, howe\'er, the product is i^reatcr than the solubility product, the substance will be precipitated. As soon then as IJie concentration of the S-ions becomes so ■small (because of the reduction of the ionization of h}'drogen sul- phide by the Il-ions of the acid) that it makes the value of ab smaller than that of the solubility product for copper sulphide, no precipitate will be formed. If, however, the liquid is diluted, the concentration of the H-ions decreases; then, if hydrogen sulphide is passed in, the concentration of the S-ions increases. The ^•alue of the solubility product can in this way be exceeded, in which event copper sulphide will be precipitated. If a .small cjuantity of stroni^ acid be added to a precipitate of copper sulphide suspended in water, only a ^-ery small amount of the sulphide will dissoh-e; to be sure, the H-ions of the strong acid will remove a part of the S-ions, yielding some undissociated hydrogen sulphide, so that in order to establish equihbrium a trace •of copper sulphide must go into solution; but soon the point will be reached when so many Cm- and S-ions are again in the solution that the value of the solubility product is readied. After this moment no more copper sulphide goes into solution. Since the value of the solubility product is ^'ery low, tlie solubility of the 120 INORGANIC CHEMISTRY. [§§ 73- sulphide in dilute strong acids is very slight; this accounts for the practically complete precipitation of the copper sulphide. On the other hand, if the solubihty product of a sulphide is greater, as in the case of iron sulphide, the addition of sulphuretted hydrogen to the solution of an iron salt, e.g. ferrous sulphate, FeS04, will cause no precipitate of iron sulphide, and iron sulphide will, unhke the previous case, be dissolved by dilute strong acids. When hydrogen sulphide is led into a solution of ferrous sulphate to the point of saturation, the concentration of the S-ions is, on account of the shght ionization of hydrogen sulphide, not great enough together with that of the Fe-ions to reach the solubility product of iron sulphide, hence no precipitate forms. ^Moreover, when dilute hydrochloric acid is added to iron sulphide, the H-ions and the S-ions form undissociated H2S molecules and the concen- tration of the S-ions therefore becomes too small in comparison with the value of the solubihty product to pre^'ent solution ; hence in the presence of enough acid all the iron sulphide goes into solution. It now becomes clear, too, why iron solutions are precipitated by ammonium sulphide. This takes place according to the equa- tion FeS04-f- (NH4)2S = FeS4- (XH4)2S04. In this case there are no H-ions in the solution to act on the iron sulphide. The reason for the non-precipitation of metals like barium, etc., either Ijy sulphuretted hydrogen or anunonium sulphide lies in the easy solubility of their sulphides. Hydrogen Persulphide. 74. If afiolution of sodium sulphide, NasS, is digested with sulphar, the sulphur dissolves and the liquid contains compounds called poly- sulphides and having formulae from Na^Sj up to Na,S5, according to the amoimt of sulphur employed. On pouring such a solution into cold dilute hydrochloric acid an oil separates out which, by distillation under low pressure, yields two compounds, having the formulae HjSj and H2S3. The hydrogen disulphide is at ordinary temperatures a yellowish, water- clear Uquid with a consistency somewhat like that of water. With alkalies it decomposes violently. Under ordinary pressure the liquid boils with partial stabihty at 74-75°. Its specific gravity is 1.376. Its fumes attack the eyes and mucous membranes vigorously. 75.] COMPOUNDS OF SULPHUR WITH THE HALOGENS. 121 Hydrogen trisulphide at ordinary tcmjjeratures is a bright yellow liquid somewhal more mobile than olive oil. Its spccifir gravity at 15" is 1.49(i. The odor reminds one of sulphur chloride and lamphor. The liquid solidifies at —52° On warming, it turns darker, becomes more viscid and, at about 90°, begins an active evolution of hydrogen sulphide. .\lkalies produce vigorous decomposition. The sensitiveness of these compounds toward alkalies is so great that they can only be prepared and kept in glass vessels whose inside surfaces have been previously freed from traces of alkali by treating with an acid. Compounds of Sulphur with the Halogens. 75. If chlorine is conducted over molten siilphur, sulphur monochloride, S2GI2, is formed. Its formula is based on its vapor density and analysis. It is a yellow liquid of a \evy disagreeable, pungent odor, which excites one to tears. It boils at 139° and possesses in a high degree the ability to dissolve sulphur — as much as 0(5' , at (irdinary temperatun^s. This solution is a thick syrupy liquid. It is used in the vulcanizing of rubber. Sulphur monochloride is slowly decomposed b}' water: 2So( 'I. + 2H2() = SOa + 3S + 4IiCl. Two other compounds of sulphur and chlorine are known, SCI, and SO4. Sulphur dichloride, SCU, is formed, slo«i}', when sulphur monochloride is mixed with liquid chlorine. The mixture has a yellow color at first but after a few days it turns red. A determina- tion of the vapor density of the red substance and of its lowering of the freezing-point of acetic acid or benzene leads to the formula SCI2. It should, however, be borne in mind that a mixture of sulphur monochloride and chlorine, S2('l2 + Cl2, must give the same molecular weight as the compound ><{%. The existence of the SClo compound is ] 'roved not only by the above-mentioned change of color of mixtures of sidphur monochloride and chlorine but also by the following observations: (1) The composition of ihc va])or given off from fresh mixtures of sulphur monochloride and chlorine is entirely di.Terent from that of the \apor given off after the mix- ture has turned red. (2) .Mixtures of sulphur monochloride and 122 INORGANIC CHEMISTRY. [§§75- chlorine decrease in volume, and this diminution is greatest when the composition corresponds to S2CI2 + CI2. (3) On distilling under 4 mm. pressure 80-90% of the liquid could be distilled over almost constant at —24° At ordinary pressure the compound boils at about 59-60° with decomposition. Sulphur tetrachloride, SCI4, can be obtained as a fine white powder, apparently not crystalline, when a chlorine-sulphur mixture of the composition S + CI4 or S + Cle is cooled to —75° (where it solidifies) and then centrifuged. The tetrachloride melts at about —33°, giving off chlorine abundantly. The exact tem- perature could not be determined. Here, too, the form of the freezing-point curve indicates unquestionably the formula of the substance, viz., SCI4. Cf. § 237. With bromine an! iodine sulphur gives analogous com- pounds. Fluorine unites with sulphur to form a gas of the formula SFe and of rather surprising properties. It is colorless, odorless and incombustible. At — 55° it solidifies with the formation of crystals. Notwithstanding its high percentage of fluorine it is chemically so indifferent that it almost resembles nitrogen in this respect (see p. 164). For instance, it is not decomposed by fused alkalies nor by copper oxide at dull-red heat. It can be heated with hydrogen without yielding hydrogen fluoride. ^.loreover, sodium can be fused in sulphur hexafluoride without losing its metallic surface, the gas not being attacked by the metal tUl the boiling- point of the latter is reached. VALENCE. 76. Certain elements have the property whereby their atoms can combine with only one atom of another element. The halogens on the one hand and hydrogen on the other are able to form only compounds of the type HX(X = halogen). This property of the atoms is called univalence. In the case of other elements hke oxygen and sulphur each atom can enter into compounds with two univalent atoms (exam- ples: H2S, H2O). These are therefore called bivalent. The number of univalent atoms that can combine with one atom of a given element serves in an analogous way as a measure of valence in general. An atom of nitro,'T;en, for instance, unites 70.J VALEXCK. 11?, with three atoms of hydrogen; nitrogen is therefore trivalent; carbon is quadrivalent, ete. ,11 The valence is ordinarily- indicated by lines, as in 0^ and H ^H N — H, each line representing a valence unit (unit bond). The valence of one and the same clement may be different according to the nature of the univalent elements with which it combines. Sidphur, for instance, can only unite A\ith two hydrogen atoms, but ^\ith univalent chlorine it forms the compound SC'l4, ■\\'ith fluorine even Sl'e- The valence of sulphur in these cases is therefore four and six. The preparation of sulphur compounds with more than six univalent atoms has not yet been accomplished; hence its maximuni valsnce is six. The halogens are univalent towards hydrogen, but in relation to each other they display more than one valence, as may be seen from the compounds IClj and ICI5 ; in the compound CI2O7 (§ 60) the maximum valence of chlorine can c\-en be assumed to be seven. It has been very generally observed that when the maximum valence of an element is an even or an uneven number, its lower valences are of the same sort; the halogens and sulphur illustrate this. However, these are exceptions to this rule. The valence ulso depends upon the temperature. 'We .sliall soon see that SO3 dissociates at a high temperature into 8(_)., and oxygen; while sulphur is sexivalent towards oxygen at lower temperatures I 8:^:0 I, if is only quadri- valent towards oxygen above 700° ( S . ) . The valence must also depend \ 0' on the prfss'ire, fur the latter exerts a great influence on the dissociation. The basis for the above sort of formulip is the idea, borrowed from organi? chemistry, that the atoms of a molecule may not assume any conceivable arrangement whatsoever, but that there is a definite order in every molecule. For some extensions of the idea of \'alence see § 317. Valence of Ions. — In the solution of an(dcctrolytc the sums of all the p<>siti\'e and all the negative awiounts of electricity must be equal, for the solution acts as electrically neutral. In a .solu- tion of hj'drochloric acid tlic posit)\e charge of the H-ions must 124 INORGANIC CHEMISTRY. [§§76- be nmnerically equal to the negative charge of the Cl-ions and, since the same number of both ions are present, each Cl-ion must carry a charge equal, but opposite in sign, to that of an H-ion. In a sulphuric acid solution, however, the two li-ions together must possess just as much positive electricity as the S04-ion nega- tive electricity. The SO4" ion is therefore called bivalent in re- spect to the hydrogen ion. It is readily seen how the valence of other ions can be determined in an analogous manner, for it is equal to the nmnerical A'alue of their electrical charge, that of the hydrogen ion being taken as unity. Compounds of Sulphur with Oxygen. 77. Of those containing only the two elements three are known, viz., S2O3, SO2, and SO3. Especial importance attaches itself, however, only to SO 2 and SO3;' the two others have been little .studied. Sulphur Sesquioxide, S,Os. This is obtained when sulphur is treated with its trioxide. It is a blue liquid, which congeals to a malachite-green mass and is soluble in fuming sulphuric acid, giving a blue solution. On being warmed it breaks up into sulphur and the dioxide : l'S203=3S02+S. Water decomposes it vnth. the formation of sulphur, sulphurous acid and polj'tliionic acids. SULPHUR DIOXIDE, SULPHUROUS ANHYDRIDE, SO2. 78. This gas occurs in nature in volcanic gases. It is formed when sulphur burns in the air or in oxygen; the well-known odor of burning sulphur is due to it. A little trioxide is also formed by this combustion. The laboratory method of preparation con- sists in decomposing sulphmic acid with copper. 2H2SO4 + Cu = CUSO4 + SO2 + 2H2O. For this purpose concentrated sulphuric acid is heated with cop- per tiu"nings, no action taking place at ordinary temperatures. The process can be explained by supposing that at the high tern- 78.] SULPHUR DIOXIDE. 125 perature of the reaction copper is oxidized by sulphuric acid to copper oxide with the evolution of sulphur dioxide: Cu + H,y O4 = SO2 + H2O + CuO. The copper oxide reacts of course with a second molecule of sulphuric acid, producing copper sulphate. The reduction of concentrated sulphuric acid by heating with charcoal is also a convenient method of preparation: 2H2SO4 + C = 2II2O + 2SO2 + CO2. However, as this equation shows, the gas is obtained mixed with one third of its ^•ollmle of carbon dioxide, from A\hich it cannot be separated directly. ]\Icreover, sulphur dioxide can be obtained bj- the action of oxygen on sulphur compounds, thus, e.g. by the roasting of pyrite in a current of air: FeSo + 3O2 == 8O3 + FeS04. Pyrite. This reaction is employed on a large scale in the commercial manufacture of sulphuric acid. The action of sulphm- on oxygen compounds also yields sul- phurous oxide, e.g. heating copper oxide or manganese dioxide with sulphur: 2CuO + 23 = CuaS + SO2 ; MnOa + 2S = MnS + SO2. Finally, the dioxide is also formed by heating an oxygen com- pound (CuO) with a sulphur compound (CuS) : CuS + 2CuO = 3Cu+S02. Physical Properties. — At ordinary temperatures and pressures sulphur dioxide is a gas. It has a peculiar taste and odor. It is easily liquefied, the boiling-point being —8°. Its evaporation produces a marked depression of temperature, sometimes extend- ing to —50°; at —76° it becomes solid. Liquid sulphur dioxide dissolves many salts, in some cases with a cliaracteristic color. It 126 INORGANIC CHEMISTRY. [§ 7S- is very soluble in water; at 0° 1 vol. H2O dissolves 79.79 vols. SO2, at 20° 39.37 vols. SO2. Boiling the solution expels all the gas (§ S3). Chemical Properties. — Sulphur dioxide is an acid anhydride; its aqueous solution has an acid reaction and behaves in general like that of an acid (§ S3). It is easily oxidized by oxidizing- agents to the trioxide. This occurs, for instance, when a mixture of sulphur dioxide and air or oxygen is passed over hot r.pongy platinum or platinum asbestos. In aqueous solution this oxidation takes place readily at ordinary temperatures. The oxidation of the dioxide can also be brought about by chlorine-water, bromine and iodine : CI2 + 2H2O + SO2 = H2SO4 + 2HC1 ; also by chromic acid, which is reduced to chromium sulphate, or by potassium permanganate, which is reduced to a mixture of manganese and potassium sulphates, and therefore loses its color: 2KMn04 + 5SO2 + 2H2O = K2SO4 + 2]\InS04 + 2H2SO4. Lead peroxide glows faintly in a current of sulphur dioxide and is reduced to lead sulphate — from brown to white: Pb02 + S02=PbS04. It is to its reducing action that the bleaching effect of sul- phurous oxide on vegetable coloring-matters is due. A red rose, for example, loses color in it. The gas probably reacts with water, setting hydrogen free, which latter effects the reduction and hence the bleaching : S02 + 2Ii20 =H2S04 + H2. In this case, therefore, bleaching depends on a reduction; as a matter of fact the color returns in many instances, when the bleached article is exposed to the oxidizing action of the air. Silk, wool and straw, i.e. substances that cannot stand the chlorine bleaching, are whitened commercially with sulphurous oxide. It also finds use as an antiseptic. 7S.) SULPHUR DIO.\IJ)li. Vll The reduction of iodic acid liv sulphur dioxide is sometimes employed as a test for the latter. For this piu-pose strips of paper are dipped in a solution of potassium iodale and starch, which tui-ns blue in the presence of sulphur dioxide — iodine being set free (§ '17). If the reaction is carried out in dilute solution, a peculiar ph(!nomenon is observed; the blue color of starch imlidc docs not appear directly when the solutions of sulphur dioxide and iodic acid are mixed, but is with- held for a certain number of seconds (definite for every concentration at constant temperature), when it suddenly appears. The following reactions come into play: I. 3SO,ar| + HIO3 = SHjSO^aq + HI. The hydriodic acid thus formed is at once oxidized by the iodic acid still present : II. oHi + HIO3 = aiip + 61. So long as sulphur dioxide is present, it reduces the iodine in this dilute solution to hydriodic acid: III. 21 + SO.aq + 2H2O = H^SO.aq + 2HI. Xot until all the dioxide is used up by the reactions I and III does the free iodine suddenly appear according to II. There are some substances which are able to extract oxygen from sulphur dioxide, i.e. the latter can also act as an oxidizing- agcnt. Ignited magnesium ribbon continues to burn in sulphur dioxide, forming magnesium oxide and sulphur. Hydrogen sul- phide and sulphur dioxide have respecti\-ely an oxidizing and a reducing effect on each other, which follows mainly the ec^uation: 2H2S + S()2 =2H20 + 3S. Sulphur dioxide is decomposed by electric sparks into sulphur and the trioxide. The action of the electric sparks is to be ascribed solely to the sudden and enormous rise of temperature which thc\' produce and the rapid cool- ing that immediately follows, for the gas particles which ha\'c become heated by the sparks arc immediately cooled again by surrounding objects. As a result the products formed do not have time to react in the opposite direction. The correctness of this view Avas demonstrated by LJT. Claire Deville with the help of an ajjparaf us -which made it pos- 128 INORGANIC CHEMISTRY. [§§78- sible to cool objects very rapidly from a very high temperature. This apparatus, the cold-hot tube, consists of a rather ■nide porcelain tube, which is heated to a bright glow in a furnace and which contains a concentric thinner metallic tube, through which cold water is forced so rapidly that the tube maintains a low temperature. When Deville introduced sulphur dioxide into the space between the two tubes, it was seen after some time that the inner tube, wh'ch was made of silver-plated copper, had turned black because of the formation of silver sulphide, while at the same time the formation of sulphur trioxide could be detected (by its producing sulphuric acid with water, a precipitate being given by barium chloride). The composition of sulphurous oxide can be determined in the following manner: AMien sulphur burns in oxygen no change of volume is observed after cooling. Therefore just as many mole- cules of sulphurous oxide have been formed as oxygen molecules consumed. The sulphurous oxide molecule must therefore con- tain two atoms of oxygen. The specific gravity of the gas has been found to be 2.2639 (air = l), or 32.6 fO=16), so that its m.olecular weight is 65.2. If jve subtract 2X16 from this for two atoms of oxygen, there remains 33.2 for sulphur, the atomic weight of which is 32. We thus see that only one atom of sul- phur is present in the molecule of sulphurous oxide and that the formula of the latter is 802- SULPHUR TRIOXIDE, SULPHURIC ANHYDRIDE, SO3. 79. This compound is found in a small amoimt in the fumes of burning sulphur (§ 78). As was stated above, oxygen and sulphur dioxide unite to form the trioxide in the presence of platinized asbestos. On the other hand, the trioxide breaks up into the dioxide and oxygen at an elevated temperature, so that the formation of the trioxide from the dioxide and oxygen is evidently a reversible process, which is expressed thus: 2802 + 02^2803. If we call the pressure of SO 2 in the equilibrium condition pi, that 79. SULPHUR TRIOXIDE. 129 of O2 P2 and that of SO3 ps, it follows from §51 that the equilibrium relation is expressed by Pi^P2 = Kps', where K is the equilibrium constant. The combination of sulphur dioxide and oxygen is easily accomplished (in the presence of platinum) at about 500°; that is to say, the above equilibrium is shifted almost wholly to the right at this temperature. If the temperature is raised, the dissociation of trioxide begins and at about 1000° it is com- plete. The union of SO2 and O2 also occurs under the influence of ultra- violet rays. These rays are best produced by a quartz-mercury arc lamp. The gases that are to be exposed to the rays must also be contained in quartz vessels, since glass is opaque to ultraviolet rays. Furthermore, an equOibrium 2S02-(-02?=»2S03 also establishes itself under the action of these rays ; for not only is the union of SO2 and O2 incomplete, but SO3, on the other hand, breaks up under the same experimental conditions, jaelduig the same equilibrium mixture. This Ught equilibrium differs, however, in many respects from the " tem- perature equilibrium." In the first place, SO3, in the presence of platinum, does not begin to dissociate perceptibly until 300°. The light influence, however, is evident even at room temperature. Like the effect of catalyzers, the action of light is retarded by sufRcieutly careful drying of the gases. There is an optimum moisture content for the con- tact process, but the action of light is effective even when the gases are passed through the illumination vessel in a very moist condition. The light equUibrium is not perceptibly affected by a marked change of tem- perattire, but it is very sensiti\'e to varying intensity of illumination. Just exactly as the dissociation increases with rising temperature, so it increases as the illumination grows stronger. The reader can form an idea of the extent of the decomposition from the observation of Coehn and Becker that, with a mercury lamp consuming 9 amp., the equi- librium established itself when about 35% of the SO3 was decom- posed. Sulphur trioxide can also be obtained by heating certain sul- phates; in the arts ferric sulphate is thus used: Fe,(S04)3 = Fe203+3S03. 130 INORGANIC CHEMISTRY. [§§80- " Fuming sulphuric acid" (oleum) is a solution of sulphur trioxide in sulphuric acid; the anhydride can be obtained from it b}' distillation. 80. Physical Properties. — Perfectly dry sulphur trioxide melts at 17.7° and boils at 46°. It looks much like ice but usually appears in another modification, viz., long as1jostos-like needles with a silky lustre. These crystals ha-\-e no sharp melting-point but sublime on heating. This modification is the stable one, for the other goes over into it spontaneously. This transformation is greatly accelerated by traces of water. The asbestine modifica- tion consists of double molecules (^03)2, the glacial form of simple molecules (SO3). This is shown by the depression of the freezing- point of phosphorus oxychloride. The first is therefore called a polymer of the second. It is also worth noting that the SO3 modification is very readily soluble in concentrated strlphuric acid, while the other, (803)2, dissolves with diffieulty. Chemical Properties. — Suljjhur trioxide unites very easily -with water to form sulphuric acid: S03 + H20=H2S04. It therefore fumes vigorously when exposed to moist air. On introducing it into water, combination and great evolution of heat, accompanied by sizzhng, results. It reacts energetically with many metallic oxides also, forming sulphates. Baryta, for exam- ple, glows in contact with it. Yvlien its vapor is passed through a red-hot tube, it is decomposed into the dioxide and oxygen. Composition. — ^The decomposition just mentioned permits us to establish the composition of sulphuric oxide. The dissociation products, S()2 and O2, are formed in the volume ratio 2:1. Now the specific gravity of sulphuric oxide is 2.75 (air = l), from which the molecular weight is calculated to l^e 79.1. This figure corre- sponds to the formula 803(32 + 3x16) and it also harmonizes with the above dissociation; for it is clear that 2 vols. SO3 must then yield 2 vols. 8O2 and 1 vol. O2 : 2803=2802 + 02. 2 vols. 2 vols. 1 vol. 82.] OXYGhJX ACIDS OF SULl'UUli. 131 Oxygen Acids of Sulphur. 8i. Sulphur forms ;m unusually large number of acids with cxygon and hydrogen, namely nine. They are as follows: 1. Thiosulphuric arid HjSjOg. 2. ITvposulphurous acid H.,S,,0,. 3. Suli)hurous acid HjSO,. 4. .'■Sulphuric acid H,S(\. Ti. Persulphuric acid HjSjOj. 6. Dit'iionic acid H2S2O5. 7. Trithionic acid H^SgO^. 8. Tetrathionic acid HjSPu. 9. Pentathionic acid H2S5O1J. It is an important fact, however, that of these nine acids only sulphuric acid has really been isolated; all the others are known only in aqueous solution or in the form of salts. The two hydro- gen atoms which each of these acids possesses are both replaceable by metals; they are therefore dibasic acids. AMth such acids it is pos.sible that just one of the hydrogen atoms be replaced by a metal. The salts thus formed are called acid salts. By different methods, e.g. the cryoscopic method, it is found that the aqueous solution of dibasic acids AH2 contains chiefly the ions H' and HA'; it is only when these solutions are very dUute that the anion HA' sphts up further into H' and A" In the case of the 1X2 A salts, however, there is an ionization into 211' + A"; but in that of the acid salts MHA the ions are chiefly M" and HA'. How far the anion HA' is split up does not depend merely on the concentration, but also to a considerable degree on the strength of the acid, HA' being more ionized in strong than in weak acids of the same concentration. THIOSULPHURIC ACID, HjSjOj. 82. This acid can only exist in dilute aqueous solution and is even then very unstable, decomposing completely in a short time. The salts are, however, st:ible and can be prepared in the following wavs: 132 IXORGANIC CHEMISTRY. [§§82- 1. By boiling the solution of a sulphite with sulphur j Na2S(l3 + S = Na2S203; Sodium sulphite. or S(V' + ^'=S203", only the anion being changed. 2. By the oxidation of sulphides in the air* 2Ca^^o + 302 = 2CaS203. Calcium disulphide. 3. By the action of sulphur dioxide on the solution of a sul- phide: 4Na2S + 68O2 = 4Na2S203 +S2. The most important salt is the sodium thiosulphate, formerly and even yet often called sodium hyposulphite, or, abbreviated, "hypo." It is very soluble in water; the solution, when used in excess, has the property of dissolving readily the halogen com- pounds of silver, hence its extensive use in photography (§ 247). It is easily oxidized by oxidizing-agents, usually to the sulphate. This takes place ^^ith potassium permanganate, nitric acid and chlorine, for example. Practical use is also made of this latter property by employing sodium thiosulphate as an antichlor in bleaching, i.e. to remove the last traces of chlorine which cling to the bleached material very obstinately and have an injurious effect. When a dilute acid is added to a dilute solution of sodium thiosulphate, the following decomposition takes place: Na2S203 + 2HC1 = 2NaCI + H2O + SO2 + S ; or S2O3" + 2H- = HSO3' + H- + S. ilph id. Anion of sulphur- ous acid. It may be, however, that the ions first unite partially to form H2S2O3, which spUts up into H2O, S and SO2. It is an interesting fact that in this decomposition in a dilute solution the sulphur precipitate is not at once visible, being first noticeable after some seconds, or even minutes, according to the dilution. It was formerly supposed that the thiosulphuric acid remained entirely unchanged until the appearance of the sulphur and the decomposition first began at this 84.] HYPOSULPHUTtOUS AXD .SULPHUROUS ACIDS. 133 moment. This is, ho\\c\(>r, incorrect; for when a dilute solution of thiosulphate is treated with an equivalent amount of dilute acid and the solution again neutralized before the appoaraufc of the sulphur deposit, it is found that the latter appears nevertheless after some time. A certain part of the free thiosulphuric acid must therefore have already decomposed, but the sulphur was in a so very finely divided dtate in the liquid that it could not at once be detected, — not until it 'lad gathered together to form larger particles. Hyposulphurous Acid, HjSjOj. 83. As early as the ISth centur}^ it was observed that zinc is dis- solved by a solution of sulphur dioxide in water without the evolution of hydrogen. ScHtJTZENBBRGER was, however, the first to show that a particular acid is formed thereby. A salt of tliis acid is produced by the action of zinc on a solution of acid sodium sulphite, XaHSOa, or by the electrolysis of such a solution, (he nascent hydrogen acting as a reducingr- agent.. H3rposulphurous acid, as well as its salts, is characterized by a vigorous reducing power. It precipitates the metals from solutions of sublimate (HgClj), silver nitrate and copper sulphate. Iodine solution is bleached by it ■with, the formation of hydrogen iodide; indigo is reduced to indigo- white. The solution is also very easily oxidized by free oxygen. It is therefore used to determine the amount of oxygen dissolved in water. For this reason it must be kept in well-stoppered vessels. Bernthsen succeeded in preparing the solid sodium salt, which proved'to have the composition Xa2S204 + 2H20, so that the acid itself has the formula H2S2O4. This salt was isolated by preparing a concen- trated solution of it and precipitating it by the addition of a suitable amount of solid common salt. The above formula is also confirmed bv a direct synthesis of the sodium salt by Moissan, who obtained it by the action of dry sulphur dioxide on sodium: 2Na + 2SO, = Na,^204. SULPHUROUS ACID, H2SO3. 84. It is taken for granted that the aqueous solution of sulphur dioxide contains sulphurous acid, 112^(^3, for this solution reacts acid, conducts the electric current, gives salts with bases and evolves hydrogen with some metals, e.g. magnesium. The solution of sulphur dio.xide in water does not conform to the law of IIionry (§ 9) at ordinary temperatures, which proves that a combina- tion with the solvent has taken place. At higher temperatures, 134 LXORGAXIC CHEMISTRY. [§§S4- however, the solution obeys this law pretty well. A fact in con- firmation of this is that all the sulphur dioxide can be expelled from the solution by boihng it, the combination being then wholly destroyed. The compound H2y03 itself has, however, not yet been isolated. The salts ha^-e the composition ^IqSOs and MHSO3 (M being an atom of a univalent metal). The acid salts are almost all soluble in Abater, while of the neutral salts only those of the alkahes are soluble. The acid sodium sulphite, NaHSOs (sodium bisulphite), is frequently employed in organic chemistry. Sul- phites in solution gradually absorb oxygen from the air, form- ing sulphates. It is a very strange fact that minute quantities of organic substances, e.g. only 0.1% of alcohol and as little as 10~5 gram molecule of stannous chloride, greatly hinder this oxida- tion. We have here one of the few examples of a retarding catalytical action. On the other hand, traces of copper sul- phate considerably accelerate the oxidation. SULPHTJRIC ACID, HjSO,. 85. Sulphuric acid is the most imp()it:in1 arid of sulphur. It can be obtained in various wavs; in the first place by direct synthesis from its elements. According to § 79 sulphur trioxide can be formed directly from sulphur and oxygen, and this yields sulphuric acid on the addition of water. The acid can be obtained from its salts by distilling them mth phosphoric acid. Its formation from the action of ox3'gen on sulphur compounds is illustrated by the oxidation of an aqueous S02-solution by t,he air. (Jn the other hand the action of sulphur on oxygen compounds may also give sulphuric acid; thus it is formed when concentrated nitric acid, HXCJ.-j, is boiled with sul- phur; and again, potassium sulphate is formed by heating sulphur with saltpetre (KN(33). 86. For the commercial manufacture of sulphuric acid two processes are now in use, the lead-chamber process and the con- tact process. Enormous amounts of the acid are produced by these two methods. The 1 e a (1 - c h a m b e r process is based on the follow ing reactions: 1. the oxidation of sulphur dioxide by nitric acid in the presence of water; 2. the oxidation Ijy the oxygen in the 86.] SULPHURIC ACID. 135 air of lower oxides of nitrogen formed from the nitric acid in tlie previous reaction. Tliese are parti}- rec:on\crted to nitric acid and partly changed to certain stages of o.xidation of nitrogen which oxidize sulphur dioxide anew to sulphuric acid. By this last process the lower nitrogen oxides are again formed, but are soon reoxidized by atmospheric oxygen ami so on. One might suppose that a certain amount of nitric acid would suffice to con- vert unlimited amounts of sulphur dioxide into sulphuric acid ■with the aid of the air. In practice this is not true, however; for the nitrogen oxides are to a small extent still farther reduced by sulphurous oxide, so that nitrous oxide or nitrogen are formed, and these are no longer able, under the conditions of the indus- trial process, to combine with oxygen. The chemical processes which lie at the basis of the manu- facture of sulphuric acid will be taken up a little later (§ I'-JS). From a technical standpoint the lead-chamber process falls into three separate parts: 1. The preparation of sulphur dioxide; 2. The oxidation of sulphur dioxide; 3. The concentration of the resulting acid. (1) The material for the production of the dioxide is sulphur or pyrite (iron pyrites, FCS2). Sulphur }-ields a purer acid than pyrite; that prepared from the latter almost always contains arsenic. The roasting of the pyrite is carried on in furnaces, the construction of which varies considerably. In all of them, how- ever, the sulphur dioxide lea^•es the furnace mixed with a good deal of air. The furnace gases pass through a canal in w^hich the dust particles carried along by the draught are deposited. (2) The oxidation of the sulphurotis acid is carried out in a structure consisting chiefly of three parts, the Glover tower, the lead chambers, and the Gay-Lussac tower. The gases enter the bottom of the Glover Tower, which is made of sheet lead lined with acid-proof brick. It is filled with lump stone, over which is laid a layer of smaller pieces of coke. On top of the tower is a reservoir for collecting the nitroso sulphuric acid (see below) that comes from the Gay-Lussac tower and the lead chambers and is to be concentrated in the Glover towci-. It flows down over the stone in the tower from a horizontally revolv- ing tube. From the Glover tower the gases enter the lead cham- 136 INORGANIC CHEMISTRY. [§86. bers. These are three or four in number and have a total capacity of 4000-5000 cubic meters. Their form is that of a parallelopiped, whose cross-section is nearty a square. Lead has been chosen as the material for the walls of the chambers, because it is the only one of the common metals which is only slightly attacked by sulphuric acid and the substances used in its manufacture. The lead chambers are connected with each other, with the Glover tower and with the Gay-Lussac tower by means of lead pipes. The first two chambers are also furnished with openings for introducing steam. The oxidation of dioxide to trioxide having been accomplished in the lead chambers, the residual gas, principally nitrogen, passes to the Gay-Lussac tower. Usually this is entirely filled with coke. On top of the tower is a reservoir containing 60°-62° sulphuric acid (Baume, see § 88), which comes from the Glover tower. The Gay-Lussac tower serves to collect the nitrous vapors that are still present in the gas as it leaves the lead chambers. These vapors dissolve in the stilphuric acid, forming the nitroso sulphuric acid which is used in the Glover tower. In this way the loss of nitric acid is much reduced. Let us now examine the task that befalls each of these three — the Glover tower, the lead chambers and the Gay-Lussac tower. The gases that come from the pyrite furnace consist of a mix- ture of sulphur dioxide and air, a larger proportion of the latter than is required for the oxidation. They have a temperature of about 300° when thej' enter the Glover tower. A, at the opening, w. The gas current rising in the tower meets an acid mixture flow- ing down from above. The latter consists of the nitroso acid from the Gay-Lussac tower, diluted with the acid (chamber acid) (3) The acid produced in the chambers contains about 67% H2S()4 (.53° BaitiMe). In this condition it is employed directly in the manufactiu-e of fertilizers ("superphosphate"). For almost all other purposes it must first be concentrated. Ordinary sulphuric acid of commerce is of about 66° B. (B.=Baume), i.e. 96-98% H2SO4. It is prepared from the chamber acid by evap- orating it first in lead pans to about 78% (60° B.) and finally in a platinum vessel. Mi.] SVLPHVHIC ACID. y,il Tliis crude sulphuric aciil of commerce ("oil of vitriol") still contains various impurities and is usually more or less brown in color because of bits of straw (from the packing of the carboys) falling in and charring. It can be purified by (Uluting it, where- upon the thst^olvcd lead sulphate is precipitated, and then stirring in a little barium sulphide. The latter produces insoluble barium sulphate, and also hyih'ogcn sulphide, ■\\hich precipitates any arsenic or lead (§ 206) still present. The arid is then decanted from the deposit, concentrated, and finally distilled. The contact process . — It has already been stated that sulphur dioxide unites with oxygen directly to form the trioxide and that the combination is considerably accelerated by the cata- lytic influence of platinized asbestos. This simple reaction is the basis of the "contact process." In practice, however, air is used instead of pure oxygen. The process falls into four separate parts: 1. The preparation of a mixture of sulphur dioxide and air; 2. The purification of this mixture; 3. The formation of the trioxide; 4. The combination of sulphur trioxide with water to form sulphuric acid. (1) Tlie purification of the gas mixture is much the same as in the lead-chamber process. For reasons which will soon be made clear it is found necessary to conduct the roasting in the presence CI a large excess of oxygen. While the equation 2SO2 + 02 = 2803 demands only 1 vol. O2 for each 2 vols. SO2, the gases are usually mixed in the ratio of 3 vols. O2 to 2 vols. SO2. {2} The platinized asbestos acfrs efficiently only wlien the furnace gases are absolutely pure, i.e., when the mixture consists simply of sulphur dioxide and air. The complete purification of these gases has been a problem of exceptional difficulty, but has been accomplished through the perseverance of KNiETacii of the "Badische Anilin- und Sodafabrik," the great chemical factory at Mannheim,, German}'. In the first place the furnace gases must be wholly freed from dust, else the catalyzer would soon become sc coated as to lose its activity. In order to determine when the gas is really dust-free it is subjected to the "optical test," i.e., it is passed through a tube closc^d at both eiuls with glass, and is 13S L\ ORGANIC CHEMISTRY. [§S6- examined with the eye to see whether it is perfectly transparent and free from nebulous masses. Even when this optical test is quite satisfactory the catalyzer suffers a loss in activity if the gas is not entirely free from arsenic compounds; the least traces of the latter have an injurious effect. The presence of arsenic compounds in the furnace gas is due to the occurrence of arsenic in the pyrites (§ 86, 1) used for roasting. Knietsch has finally succeeded in completely eliminating the arsenic compounds by blowing steam into the gas mixture. (3) As already set forth in § 79, the equilibrium 2SOj + 0,?^2S()3 is expressed h\- the equation Pl2/), = A>32. According to this equation the formation of sulphur trioxide is more complete in the presence of an excess of either sulphur dioxide or oxygen, for as pi or p^ increases ps must also increase. Since the object in view is to convert the dioxide as completely as possible into the trioxide, it is advantageous to provide a large excess of oxj-'gen. This explains why more than the theoretical amount of oxygen is taken. Compare (1). The equilibrium must also depend on the pressure, for, if this is increased n times, the equation becomes: {n.pi)~np2=Kn"p3^, or npi^p2 = Kp3^, from which it is evident that at a higher pressure (n>l) the for- mation of the trioxide is more nearly complete (§ 102, 5). The manufacturer does not find it necessary, however, to employ high pressure, which would involve, moreover, a great complica- tion of the apparatus. If it is desired that the combination of sulphur dioxide and 0x3'- gen should be as complete as possible, the temperature must be kept at about 400°. Since, however, the heat of formation of the trioxide is great, viz., SO2 + O = SO3 + 22,600 Cal, 86.] SULI-IIUUIC ACID. 139 the apparatus must be cooled. This is done most practicably by the aid of a fresh portion of the gas mixture, as the next paragraph sets forth. The construction of the apparatus is as follows: The tubes ab (Fig. 31) contain the platinized asbestos h, supported on little sieves (shown in the middle tul)es). The purified furnace ^ases fir.st pass around the outside of the tubes and are thus warmed to Fk;. .31. — CONT.VCT-PKOCE.SS APPARATUS. the desireil temperature at the heat expense of the rjis system within. AMien the proper temperatui-e is reached the f!;ases are allowed to enter the tubes, where sulphur trioxide is formed with the evolution of more heat. By increasing or diminishing the rate of f^ow of the gas current, the temperature can be regulated very satisfactoi-il}-. When the opo-ation is started the apjjaratus must first be warmed t(j 400°. (4) The reaction Ijetween sulphur ti'ioxide and wati'i- is an enertretic one. Xevertlieless, llie iiiannfacture of sulphuric acid 140 lA'OHOANIC CHEMISTRY. [§§86- from these two compounds involved some difficulty, inasmuch as sulphur trioxide fumes invariably escaped when this substance was introduced into water or dUute sulphuric acid. Onl}' when sul- phuric acid of 97-98% is used as the absorbent and care is taken to keep the acid at this concentration b}' the simultaneous addition of water does a complete and immediate absorption occur. This is due to two circumstances: first, that traces of water change sulphur trioxide into the asbestine modification (§ SO), which is only slowly absorbed by sulphuric acid; second, that at the concentration of 97-98% H2SO4 the system a:S03 + i/H20 has a minimum of vapor tension, which is very low. 87. Physical Properties. — The pure compound, hydrogen sul- phate, is an oily liquid at ordinary temperatures, solidifying at a low temperature and melting again at + 10.0°. Its specific gravity in the Uquid state (15°) is 1.8500. Chemical Properties. — ^The concentrated acid obtained by dis- tillation is not the simple compound H2SO4, for it still contains about 1.5% of water. In order to prepare the absolute!}- pure acid the distilled product must be mixed with the theoretical amount of sulphxn: trioxide. When pure sulphuric acid is heated, it begins at 30° to give off fumes of sulphur trioxide; this continues until the boiling-point, 317° at 750 mm. Hg. pressure, is reached, when an acid with 1.5% water distils over. On heating the vapor of sulphuric acid above the boihng-point, it begins to break up into water and the anhydride; this dissociation is complete at 450°, for the vapor density at that temperature is found to be 25.1, while that of SO3 + H2O is theoretically 24.5. When sulphuric acid is mixed with water, a strong evolution of heat occurs. The mixing must therefore be done with great care, particularly in glass vessels, the acid being poured in a fine stream into the water and the liquid being steadily stirred. On mixing them in the reverse way, by pouring the water into the sulphuric acid, the intense heat that is produced may cause the glass to crack. However, when the acid is mixed with ice in a certain proportion, a strong cooling follows. The mixing of sulphuric acid and water is attended by a con- traction, i.e. the volume of the dilute acid is smaller than the sum of the volumes of water and acid. It is known that sulphuric acid is able to form hydrates with water (§ 237). ST.] SULPHURIC ACID. 141 Sulphuric acid is a strong dibasic acid, but not as strong as hydrochloric acid, for, while the latter is ionized to 95% at a dilution of 0.1 gr. mol. per 1., sulphuric acid at the same dilution is only ionized to 55%, into 2H'+S()4" At higher concentra- tions HSO4' ions also exist. It acts cm many metals, giving off hydrogen. This action is made use of, as stated above, in the preparation of hydrogen; the acid must, however, be dilute, for when it is too strong or warmed, the hydrogen gen- erated partially reduces the sulphuric acid so that the gas given off contains hydrogen sulphide. Sulphur dioxide also is formed when hydrogen is led into hot sulphuric acid. It is upon this action that the reaction of copper with hot concentrated sulphuric acid depends (§ 7S). Mercury, silver and certain other metals are similar to copper in their behavior. Platinum and gold are not attacked by the acid. Sulphuric acid makes holes in paper, hnen, dress goods and the hke, when dropped on them. It has a destructi\c, charring effect on organic substances in general. This is due in many cases to the great tendency of the acid to unite with water, which makes it not only deprive other substances of the water they con- tain, but even withdraw the hydrogen and oxj^gen from organic compounds to form water. On the other hand, sulphuric acid gives up oxygen to many organic substances, being itself reduced. In order to detect free sulphuric acid in vinegar, for example, the liquid is evaporated on a water-bath with a little sugar. Free sulphuric acid, if present, chars the sugar during the concentration. The most of the salts of sulphuric acid (sulphates) are soluble in water. Barium, strontimu, and lead sulphates are insoluble, while calcium sulphate (gj'psum) is slightly soluble, but only to a very small degree. The formation of barium sulphate, BaS04, serves as a characteristic test for sulpliurlc acid, or, as we may better say, for the ion SOi". The sulphates are in general ver}- stable. They can, for instance, be heated to v^fjf.high temperatures without decomposi- tion. The acid salts lose water on heating, and pass over into pyrosulphate S'ijji ^■^^ru\l 'irfi IIIO'II 'I7 1 /' Mill ill Ki-.'HM my I J. -I 2NaHS0i = H2O + NaoSoOy. Sodium nvroaulnhate. 142 IXORGAXIC CHEMISTRY. [§§ 87- If these pyrosulphates are heated still higher, they give off sulphur trioxide and form neutral salts: XaaSaOy = N"a2S04 + SO3. 88. Uses. — Sulphuric acid is of enormous practical value, its uses being most varied. It is employed in the preparation of almost all other mineral acids from their salts. In the manufacture of soda after Le Blanc it is used in astonishingly large amounts and in nearly all other branches of chemical industry it is of some service or other. In the laboratory it is often employed as a drying-agent. A moist substance is dried verj- thoroughly when placed in a closed apparatus near a dish of the concentrated acid. For this purpose special pieces of apparatus are constructed, called desiccators. The determination of the concevtration of sulphuric acid is an operation that is freciuently necessary. Ordinarily the specific- gravity is made use of, for this can be determined rapidly ^vith an areometer. There are tables so prepared that the i)roportion of II2SO4 or SO3 in a dilute acid whose specific gravity and tempera- ture are kno^^^l can be quickh' read. Baume, a chemist of the latter part of the eighteenth century, constructed an areometer with an arbitrarv^ scale, the zero point of which indicates pure water and the point 10 being reached in a 10% salt-^-olution. .All the diAasions are equal. 100% II2SO4 would then be represented b}' the line 66.6. In the arts the strength of sulphuric acid is still given as so mam' " degrees Baume." Fuming sulphuric acid is the name of a sulphuric acid that contains sulphur trioxide in solution. It is obtained by dissolving the oxide in concentrated sulphuric acid. Fuming sulphuric acid is a thick oily liquid, which fumes vigorously in the air, throwing off the trioxide. Sp. g. = 1.85-1. 00. CHLORIDES OF SULPHURIC ACID. 89. When phosphorus pentaehloride acts on sulphuric acid a com- pound SO3HCI, chlorosulphonic acid, is formed : H2SO, + PC'U = SO3HCI + POCI3 + HCl. The same compound results from the direct union of sulphur trioxide and hydrochloric acid. It is a colorless liquid, which fumes vigorously on 91.] CHUlKlDF.S l)F SVl.PIIUIilC ACID. 143 exposure to the air. Sp. g. =1.7(1(1 at 18°. Boiling-point, 15.S°. On the addition of water a ^'iol('at reaction occurs, producing hydroclJoric acid and sulphuric acid. .^t^HCl + H,0 = ILSO, + IIC'I. 90. ,\ fompouiui, SO, ('I2, sulphuryl chloride, is obtained by the direct union of sulphur dioxide and chlorine, most easily by first saturating camphor with sulphur cUoxide (which readily dis.s(>lvcs in it) and then passing chlorine over it. The camphor remains unchanged. Sulphuryl chloride is a colorless liquid, which boils at G9.1°, has a penetrating odor, fumes strongly in tlie air and has a specific gra\'ity of 1.6674 at 20°. The addition of a Httle water con-^-erts it into chlurosulphonic acid and hydrochloric acid, much water to sulphuric and hydrochloric acids: SO2CI2+ H,0=Sr)^HCl+ HCl. SO,Cl2+2H20= H2SO, +2HC1. These decompositions of suljiliurj-l chloride can be represented in the following way : iCI^TlftOH .OH SO, -^ =80,/ +2HC1. - - jCl + n OH -\0H In the place of the U\o chlorine atoms we have, therefore, two OH (hydroxyl) groups entering. For this reason it is assumed, in close- analogy with the methods of organic chemistr}-, that sulphuric acid con- tains two hydroxyl groups. Sulphuryl fluoride can be obtained by the direct miion of sulphm- dio-xide and fluorine. It has the same remarkable stability as the com- pound SF, (I 7.5). It is a colorless and odorless gas, liciuid at —52° and solid at —120°. It can be heated with water in a sealed tube to 1.50° without undergoing decomposition. Alkalies absorb it, though veiy slowly. Sodium can be fused in it without being attacked. Persulphuric Acid, H2S2OJ. 91. The potassium salt, KjS^Oj, or, still better, the ammonium salt, (NH,)2S20j, of this acid can be obtained by the electrolysis of a cold saturated solution of the corresponding sulphate in sulphuric acid of I..3 sp. g. In such a solution we may assume wc have the io;is K' and HSO/; the latter are discharged at the anode and can then unite to form H2S2O8, which forms with the K " ions present the difficultly soluble potassium salt KzSzO,. This sejjarates out as a white crystalline mass. lU INOBOAXIC CHEMISTRY. [§§91- However, the combination of two HSO, groups only tabes place when their concentration at the anode is quite high; for if this is not the case there is more opportunity for secondary reactions, such as a union with water to form 2H2SO, and 20H, the latter of which is decomposed into H2O and 0. Such a high concentration at the anode is reached by using a very small electrode. The electric current therefore has a high density at the anode; that is, a large quantity of electricity must pass through a small surface. The effect thereof is that this large quantity discharges a great many HSO,' ions into a small space, or in other words, produces enough HSO^ groups to make the concentration ver}' high there. As low as 100° it decomposes in the following way: 2Is:2SA=2K2SA+02. K-pyrosulphate. The bai-ium salt of persulphuric acid is soluble in water, as are also most of the other known salts. The action of 100% hydrogen peroxide on sulphur trioxide or on chlorsulphonic acid yields Cako's acid; S03+HA=H2S05, yOH .OH SO2 = SO2 +HC1. ^Cl+HA \o-OH It crj'stallizes very prettily and melts at about 4.5° with slight decom- position. Caeo's acid reacts with another molecule of chlorsulphonic acid according to the equation X»H X)H HO SO2 +Cl-S020H=S0j I , \o-OH \o.O-SO2 forming persulphuric acid, which can be obtained in this way pure and crystallized, v,\th a melting-point of 60° (attended by slight decom- position). A solution of C-iRo's acid in sulphuric acid can be prepared in a simple way by mixing H2O2 with an excess of strong sulphuric acid. On the basis of this method of formation Baeyer gave the compound the name sulpho-mono-peracid. It has very strong oxidizmg powers. It sets iodine free from potassium iodide, oxidizes sulphur dioxide to trioxide, and ferrous to 92.] POLYTHIONIC ACIDS. 145 ferric salts and also precipitates the higher oxides of silver, copper, manganese, cobalt, and nickel from solutions of salts of these metals. On the other hand, it neither bleaches permanganate solution nor oxidizes solutions of chromic and titanic acids; in these respects it is distinguished from hydrogen peroxide, to which it otherwise shows much similarit^•. POLYTHIONIC ACIDS. 92. I'nder this name are grouped four acids of the general formula H2^)i06, in which the number of sulphur atoms, n, can be 2, 3, 4 and 5, and this determines the names of the individual acids. Dithionic acid, HjSjOg. The manganese salt of this acid is obtained when finely powdered manganese dioxide is suspended in water and sul- phurous oxide passed in: 2S02 + Mn02 = MnS208. From this barium salt the dithionic acid can be liberated by sulphuric acid. The solution can b=) concentrated in vacuo over sulphuric acid till its specific gra^nty reaches 1..347: farther concentration or warming results in a decomposition: H2SA = H2SO,+S02. Trithionic acid, H2S3O8. Potassium trithionate is formed when a solution of potassiimi thiosulphate is saturated with sulphur dioxide: 3SO2 + 2K2S2OJ - 2K2S3O8+ S. The free acid is unstable ; even at ordinary temperatures it decomposes in a dilute solution into sulphur, sulphurous oxide and sulphuric acid : Tetrathionic acid. Its salts result from the action of iodine on the solu- tion of a thiosulphate. K2S,0,+ 21 - 2KI + K,8,0„- The acid itself can be obtained (also only in dilute solution) by adding suiphuric acid to the barium salt, which is prepared in an analogous manner. In dilute solution it is quite stable; in the concentrated state it breaks up into sulphur, sulphurous oxide and sulphuric acid. Pentathionic acid. On mixing solutions of sulphur dioxide and hydrogen sulphide the principal reaction is a mutual oxiflation and reduction of these compounds with thr' .separation cf sulphur (§ 78). The action is, however, much more complicated, inasmuch as polythinnic acids, among them ])entathionic acid, are formed in addition at the same time. The mixture of IL.S.aq and SOj.n(| is known as " Wackenrodicu's liquid." ' WeU-crystaUized salts of pentathionic acid liavu been prepared. UG INORGANIC CHEMISTRY. [§93. Use of Sodium Thiosulphate in Volumetric Analysis. I o d m e t r y. 93. On adding sodium thiosulphate to an iodine solution, the intensely brown liquid loses its color, sodium iodide and sodium tetrathionate, two colorless compounds, being formed: 2Xa2S203 + 21 = Na2S406 + 2NaI ; or, writing only the ions that take part in the reaction: 2S203"+2I = S406" + 2r. The disappearance of the color is thus due to the fact that the molecules of iodine are transformed into ions by taking up two negative charges from 2^2* >■/' t^pon this fact a method is based for determining the amount of free iodine in a solution. This is done b}' allowing a solution of sodium thiosulphate, whose concentration (Hire) is known, to flow drop by drop into a definite volume of iodine solution. (For letting out a certain amount of liquid a pipette (Fig. 32) is commonly employed.) The color gradually brightens and finally a point is reached when the liquid is only slightly tinged and the addition of another drop causes the color to entirely dis- appear. This transition can be ^'ery accurately detected. The iodine molecules have now entirelj' disappeared. Since according to the above equation a molecule of thiosulphate is consumed for each atom of iodine, the percentage of iodine in the solution can be calculated from the amount of thiosulphate used. To make the calculation of the result of such a determination (titration) as easy as possible the thiosulphate solution is so stand- ardii'ed that it bears a certam relation to an equivalent of iodine ( = 127 g.), i.e. a certain amount bleaches exactly this much iodine. "Normal solution " is a name aj^plied to a solution containing the equi\'alent weight (§ 23) in grams (gram equivalent) in one liter. Frequently use is also made of a J, J, j\ or a twice, thrice, etc., normal solution. Normal hydrochloric acid contains 36.5 g. HC'l, normal sulphuric acid 49 g. H2SO4 (= i gram molecule), a normal iodine solution 127 g. iodine, per liter. Detailed direc- J 9 i.] von; METRIC ANALYSIS. 147 tions for preparing such solutions can be found in the text-books of anal3'tical chemistry. In order to determine readily the volume of thiosulphate solu- tion that is required in the analysis, use is made of a burette (Fig. 33), a glass tube that is divided into ^\- c.c. and closed at the lower end with a. glass sto])-ci)ck or with a rubber tube and pinch-clamp. In titrating the iodine solution the thiosulphate Fig. 32. — Pipette. Fig. .3:'>. — BinjETTEs axb Suppokt. solution is allowed to flow out slowly and, finally, drop l)y drop, while the liquid is being stirred. Exaniplr. For 50 c.c. of an iodine .solution whoso strcngtli is to be determined 27.30 c.c. y^ normal thiosulphate solution was necessary before the color completely disappeared. Rccjuired the number of grams of iodine contained in 1 liter of this solution. 1000 c.c. -j\ normal XajSoO-j solution (see above) decolorizes jij equivalent of iodine ( = 12.7 g.); 27.3 c.c. therefoi-e decolorizes 12 7 . . g. iodine. This amount is contained in 50 c.c. of 27.3 X 1000 148 IXORGAXIC CHEMISTRY. [§§93- the iodine solution in question. Hence 1 liter of the latter con- tains- •20X27.3X12.7x10-3 =6.SS42 g. iodine. ^'arious other sul^sta nf cs which liberate iodine from potassium iodide can be determined by titrating the amount of iodine dis- placed; for example, chlorine and bromine may be thus determined, since they sot free the equivalent amount of iodine from potassium iodide solution. SELENIUM. 94. Selenium was discovered by Beezelius in 1817. It took its name from ae'Krjvr) (the moon) , because it possesses great sicnilaritjr to the element tellurium (named from tellus=i]\e earth) discovered a short time previously. It is rather widely distributed in nature, but it occui's only in small quantities. It is found native, is frequently found in pjn-ite and also appears in some rare minerals. When this sort of pyrite is employed in sulphuric acid manufacture, the selenium C(jllects in the "chamber-mud" of the lead chambers; from this it is usually obtained. The process is as follows : The selenium deposit is heated with nitric acid, which oxidizes the selenium to selcnic acid, Ii2Se04. The solution thus obtained is first boiled with hydrochloric acid, whereby selenious acid, H2Se03, is formed with the evolution of chlorine. This latter acid is then reduced by means of sulphurous oxide to selenium, which separates in amorphous red flakes. Selenium displays analo2;y with sulphur in many respects; for instance, in occurring in various allotropic conditions. According to Saunders, there is an amorphous red modification, that is soluble in carbon disulphide. From this solution the selenium separates as a second modification, which is the red crystalline selenium, fusing at 170^-180°. Then there is a metallic form fusing at 217°. This modification appears when amorphous seleniimi is heated to 97°, at which point a sudden and marked rise of tempera^ ture occurs; or when molten selenium is suddenly cooled to 210" and kept for a time at that temperature. In this metallic state selenium has a metallic lustre, is insoluble in carbon disulphide and conducts electricity. Its conductivity strangely depends very much on the intensity of its illumination, however. The melting-point of selenium is 217°, its boiling-point 680°, As in the case of sulphur the ^'apor density decreases with rising 94.] SELENIUM. 149 temperature till about 1400° is reached, when it remains constant. At this temperature it is found to be SI. 5 (11=1), corresponding to a molecular weight of 163.0. Now since the atomic weight of selenium, as deduced from the vajjor density of its compounds, is 7S.0, the abo^'e molecular weight agrees xevy closely with the formula Sco. Hydrogen selenide, H^Se, can be obtained directly from its elements, as these unite at 400°. Analogously to hydrogen sul- phide, it can also be got by the decomposition of iron selenide, FeSc, with hydrochloric acid. At a high temperature hydrogen selenide dissociates into its elements. Its properties are only slightly acidic and it is more poisonous than sulphuretted hydrogen. The heav}' metals are precipitated from their solutions as selenides by it. An aqueous hydrogen selenide solution becomes turbid on standing because of the selenium that separates out. Two chlorine compounds, Se2Cl2 and SeCU, are known. The latter is much more stable than the corresponding sulphur com- pound, SCI4 (§ 75). Selenium tetrachloride is solid and sublimes without decomposition; dissociation does not begin until 200° is reached. Selenium dioxide, ,Se02, is the only oxide of selenium known. It results from the burning of selenium in the air. The extremely disagreeable odor which arises is not a property of the dioxide, however, but is probably due to the formation of another oxygen compound of selenium which has not as j-et been isolated. Sele- nium dioxide forms long white needles that sublime at 310°. Selenium dioxide is an- acid anhydride; on dissolving it in water an acid, selenious acid, IloSeOa, is formed, which can be isolated (unlike sulphurous acid). This acid crystallizes in large colorless prisms. On being heated it breaks up into water and anhydride. Sulphur dioxide or stannous chloride reduce it to free selenium, which is deposited in red flakes: HaSeOs + 2SO2 + H2O = 2H2SO4 + Se. Sulphuretted hydrogen precipitates from the solution selenium siilphide, SeS, insoluble in ammonium sulphide. When chlorine is passed into the solution of selenious acid or 150 IXORGAAriC CHEMISTRY. [§§94- when bromine is added to it, selenic acid, H2Se04, is formed. In the pure state this is a crystalline solid, melting at 58°. The 95% solution of it is an oily liquid, which has the appearance of sul- phuric acid. The barium salt of the acid, like that of sulphuric acid, is extremely difficultly soluble. On boiling with hydrochloric acid, selenic acid is reduced to selenioiis acid with the evolution of chlorine. Tellurium. 95. Tellurium is of rare occurrence; it is known in the rative condi- tion and also in combination -wiih bismuth, and T\'ith gold or silver (in sylvanite, or graphic tellvrvim). It is found chiefly in Transylvania and in the Altai mountains, and also in Boulder Co., Colorado. In the amor- phous condition tellurium is a black powder, but after fusion it is silvery white, of a metallic lustre and a conductor of heat and electricity. The vapor densit)', as in the cases of selenium and sulphur, decreases vnih. increasing temperature and does not remain constant till about 1400°; it then corresponds to a I'ej molecule. Hydrogen telluride, HjTe, results from the action of hydrochloric acid on zinc telluride, ZnTe. The product thus obtained contains more or less hydrogen. It is very poisonous, and dissociates readily. From solu- tions of the heavy metals it precipitates their tellurium compounds (t 6 1 1 u r i d e s) . Tellurium dioxide, TeOj, is formed on burning tellurium in the air. It is very difficultly soluble in water. Tellurous acid, H2Te03, is obtained by dissolving tellurium in nitric acid. It dissolves in water with great difficulty and breaks up on warming into TeOj and H^O. Telluric acid, H2Te04, is prepared by fusing the metal or the dioxide with soda and saltpetre and separating the acid from the tellurate formed. The compound H2Te04 + 2H20 crystallizes out from the aqueous solution; it loses its water of crystallization at 100°. The free telluric acid, H2Te04, prepared in this wa)- is a white powder, difficultly soluble in cold water. Telluric acid has only feebly acid properties. Selenium and tellurium both combine with potassium cyanide, when they are fused with it, forming compounds corresponding to KCKS, viz., KCXSe and KCXTe. Nevertheless, while potassium t e 1 1 u r o-cyanide is at once decomposed by the oxygen of the air with the separation of teUm-ium, potassium s e 1 e n i o-eyanide is more stable and does not ^compose ■srith the separation of selenium until it is boiled with hydro- 96.1 SUMMAKY OF THE OXY(JE.\ GUOUF. 151 chloric acid. We have here a means of detecting aelenium in the pres- ence of tellurium and of separating the two. SUMM.VRY OF THE OXYGEN GROUP. 96. The elements oxygen, sulphur, selenium and tellurium, lilte the halogens, form a natural group, particularly in two respects ; their compounds correspond to a general type and their physical and chemical properties vary gradually with increasing atomic weight. Their hydrogen compounds have the formula RH2, their oxygen compounds and their acids the formulae ROo and H2R()3, and also RO3 and H2RO4. Ozone may be considered with reference to these tj^jes as analogotis to sulphur dioxide; O-Og ozone; S-02 sulphur dioxide. The following table shows the gradual change, or pi'ogression, of the physical properties : 0. S. Se. Te. Atomic weight Specific gravity. . . . 16.00 1.124 (at -181°) 32 . 07 1.95-2.07 119.5° 450° yellow 79.2 4.2-4.8 217° 680° red 127..-. 6.2 452° Boiling-point Color -181.4° light blue white heat black As the atomic weight increases, the values of the physical con- stants also increase, as the table shows. At the same time the external appearance approaches that of the metals; in tellurium the metaUic appearance is quite marked. The instability of the hydrogen compounds increases from oxygen to telliu-ium; the strength 0/ the oxygen acids diminishes rapidly, sulphuric acid belonging to the strongest, and telluric acid to the very weak, acids. It should also be noted that all of these elements appear in allotropic modifications. 152 }.\ORGAN-IC CHEMISTRY. [§§9T- THERMOCHEMISTRY. 97. It was stated above (§ 20) that a chemical combination or decomposition is accompanied by an evolution or absorption of heat, in other words by a heat change, or caloric effect. In many cases this caloric effect has been carefully measured. The work of Berthelot and of Thomsen along this line has been especially fruitful. That part of chemistry which deals particularly with these caloric effects is called thermochemistry. The caloric effect is always given for molecular amounts of the reacting substances, since in this way only is it possible to compare substances from a chemical standpoint. Hence, when the heat of formation of water is said to be 69.0 calories (kilogram calories), it is implied that this number of calories is evolved by the union of 2 g. hydrogen with 16 g. oxygen : 2H + O = H2O + 69.0Cal. In this equation H and stand for gram atoms. In expressing a caloric effect it is necessary to indicate the state- of matter of the reacting and the resulting substances, in so- far as this is not self-evident, because the latent heat of fusion. or vaporization must be taken into consideration. The above amount, 2H + O-H2Oiiquid = 69.0 Cal., refers to the formation of water and its conversion to a liquid. It therefore includes the heat of condensation. Since this amounts to 0.536 Cal. per gram, it would in this case (for 18 g.) be 9.6 Cal.;. hence the caloric effect of the combustion of hydrogen to steam at- 100° is 2H + 0=H20gaa + 58.4 Cal. The caloric effect is also influenced by the state of matter in which the substances react, i.e., whether solid, liquid, or gas, inasmuch as solution is almost always accompanied by a heat change. In the formation of sodium chloride by the mixture of dilute solutions of sodium hj^droxide and hydrochloric acid (this being indicated by aq after the formulae of the substances) the caloric effect is: NaOHaq + HClaq = NaClaq + B./) + 13.7 Cal. 99.] THERMOCHEMISTRY. ] ,53 However, when the salt is prepared by passing hydrochloric acid gas into a dilute solution of the base, the equation is as follows: NaOHaq + HClgas = NaClaq + H2O + 31 . 1 Cal. We thus obtain 13.7 Cal. as before, but increased by the heat of solution of gaseous hydrochloric acid in a large amount of water, viz., 17.4 Cal. The heat of formation of chemical compounds must be equal to their heat of decomposition, but have the opposite sign. Were this not the case, heat would be lost or gained when a compound is formed and then decomposed so as to return to the original con- dition, and such a result would be at variance with the Law of the Conservation of Energy. Experience has shown that in the formation of most compounds heat is generated, but that in many cases heat is absorl)ptl. Chem- ical actions of the first sort arc called exothermic, those of the second endothermic, reactions. An example of the second sort is the synthesis of chlorine monoxide: 2Cl + = Cl20ga^-15.1 Cal. 98. For the determination of the caloric effect various methods are in use. Only those actions are suitable for thermochemical measurements •which complete themselves quickly. In measuring the caloric effect in the case of liquids or solutions, as, for example, the heat of neutralization of acids and bases, the heat of solution or of dilution, etc., an ordinary calorimeter is generally used, such as is employed in physics for the method of mi.xtures, the same precautions being taken in order to secure accurate results. The heat of combustion of a substance is usually measured with the calorimetric bomb of Berthelot-Mahlee. This is the usual method with organic compounds. 99. The Law of HESS. The entire ailoric effect (the n-liole amount of energy) produced by the transformation of one chimical system into another is independent of all intermediutc stages. This law is a direct conse(iuence of the principle of the con- servation of energy. If Hi;ss's law did not hold, energy would have to be gained or lost in the transition from one system to another and the subsequent return to Ihe initial condition, wiiii-h 154 INORGANIC CHEMISTRY. [§99- is contradictory to the above principle. A few examples will serve to make this law more clearly understood. (a) A dilute solution of sodium sulphate can be prepared from sodium hydroxide, sulphuric acid and water in various ways. For instance, two gram-molecules of the base can be treated at once with dilute sulphuric acid; or one gram-molecule of the base can be mixed with the acid at first and the second added afterward. Accordingly we get the following caloric effects: (1) 2NaOH -l-HaSOiaq -NaaSO^aq -2H2O = 31.4 Cal. jNaOHaq-hHaSOiaq -NaH,S04aq-H20 = 14.75 ^' { NaOHaq -I- NaHSOiaq - NaaSOiaq - H2O = 1 6 . 65 Total. 31.4 Cal. (6) From ammonia, hydrogen chloride and water a dilute solu- tion of ammonium chloride, NH4CI, can be prepared, either by letting dry ammonia gas combine with dry hj^drogen chloride gas and dissolving the resulting ammonium chloride in water or by first dissolving ammonia and hydrogen chloride in separate por- tions of water and then mixing the solutions. In the first case we have the equations: NH3gas + HClgas-NH4Clsoiid = 42.6 NH4Clsoiid-haq-NH4Claq = - 4.0 38.6 Cal. in the second case: NHs + aq-NHgaq = 8.82 HCl-f-aq-HClaq =17.13 NHgaq -1- HClaq - NH4Claq = 12 . 45 38.40 Cal. The final effects in the two cases are found to be alike within the limits of experimental error. With the help of Hess's law the determination of the caloric effect Ls rendered possible in many reactions which cannot be dealt with directly or are unsuitable for calorimetric measure- ments. In general this is done by making thermochemical meas- 100.] THER.UO('HI-:\nsrRY. 155 urements for a series of processes in which the reaction plays a part and finally calculating the caloric effect of the reaction as the single unknown, as will be more fully explained in the examples below. Suppose it wvi-Q required to find the heat of formation of hydro- gen sulphide. Tliis compound can be formed directly from its elements (§ 72), but the reaction Is unsuitable for thermochemical study. We will therefore start with the system, H, S, and 0, and consider the two wa}'s b\- which it can form water and sulphur dioxide: (1) hydrogen and sulphur are burned directly to water and sulphur dioxide: (2) (a) hydrogen and sulphur combine and (6) the resulting hydrogen sulphide is burned to water and sulphur dioxide. Since we started with the same system and in the end reached the same result in each case, the caloric effect must be the same according to Hess's law, so that, if we measure (1) and (26), we can equate (1) and (2) and solve for (2c0, thus: Heat of combustion of 2H + heat of combustion of S = heat of formation of HjS + heat of combustion of HjS. (2H + -H,0) + (S + 20 -SO,) = (2H+S -H^S) + (H,S + 30 -SO, -H^O). (JS.O + 69.26 = a-+ 13.3.4(5; •. a-=(S+2H-H.,S)=3.8. 100. In using these values of the heat of formation and heat of decomposition it should be noted that they do not represent the amounts of heat liberated by the combination of atoms to form molecules, but that the heat of decomposition of the molecules of the elements (i.e. the amount of heat required to break these molecules up into atoms) is always included. When, for example, chlorine unites with hydrogen to form hydrochloric acid, 22.0 Cal. are given off. That which is measured is the total caloric differ- ence between the initial system H2 + CI2 and the 2HC1 formed from it. In the indirect determination of a licat of formation with the help of Hjoss's law the calculated caloric effect also includes the heat of decomposition of the molecules of the ele- ments. In the determination of tlie heat of formation of hydrosivn l.jG IXORGAMC CHEMISTRY. [§§100- sulphide, for instance, in the above way the caloric effect of the combustion of this gas is composed of the following parts: 2(2H + S - H2S) + 3(20 - O2) = 2SO2 + 2H2O + p Cal. ; that of the combustion of hydrogen of the following: 2(2H - H2) + (20 - O2) = 2H2O + q Cal. ; that of the combustion of sulphur of (2S-S2)+2(20-02) = 2S02+rCal.; (20 — O2), etc., indicating the heat of decomposition of molecules of the elements. The heat of formation of hydrogen sulphide is r + q—p. Deduc- ing the value of r + q—p from the above equations, we have r+g-p=(2S-S2)+2(2H-H2)-2f2H + R-H2S), from which it follows that the heats of formation of the sulphur and hydrogen molecules are included in the heat of formation found. Chemical Affinity. 101. When a compound is formed, we attribute the phe- nomenon to the affinitj' which exists between the combining sub- stances. The term "affinity" comes down from an age when it was thought that only those substances could combine with one another which were in a certain agreement with each other (were " in love with each other," as Empbdocles and later also Glauber expressed it). This affinity was originally considered as a force. Thomsen, for example, defined it as the force which holds the parts of a compound together. Concerning the magnitude of this force our knowledge was for a long time only qualitative. If the substances AB and C interacted to form AC and B, it was said that the affinity of A for C was greater than that of A for J5. Comparative study of such reactions led to the arrangement of a series of the elements in decreasing order of affinity; but the absolute, or even relative, magnitude of these affinities was as it were a closed !;o;)l:. ITi-i: - it "-as a great step forward when §101.] THERMOCHEMISTRY. I.j7 Berthelot developed a method of measuring affinity. He con- sidered that the quantity' of heat liberated in the formation of a chemical compound was a measure of the affinity satisfied by the action. Thus affinity came to be regarded no longer as a force, but as an amount of work. We Imow that when water is decom- posed by the current from a dynamo, work must be done in order to drive the dynamo and also to split up the water molecules; and, conversely-, when hj'drogen and oxygen unite, heat, or in other words energ}', is produced. A mixture of hydrogen and oxygen can be compared with a lifted stone ; both possess potential energy. When the stone falls, its potential energy is transformed into kinetic energy. When hydrogen combines with oxygen the potential energy of the system is converted into heat. Since he regarded this heat effect as a measure of the driving force of any chemical reaction, Berthelot was led to propose his principe du travail maximum, viz., that of all the chemical processes which can proceed without the application of energj' from an outside source that one always occurs which involves the greatest evolu- tion of heat. However, this principle did not prove to be universal])' applic- able. The very existence of endothermic compounds is at variance with it, for the heat effect of a reaction involving an endothermic compound would be greater if that compound were not formed. Further, the rapidly increasing number of knowTi equilibrium reactions throws doubt on the principle, for, if in an equilibrium A+Bt^AB the direct reaction (-^) is exothermic, the opposing reaction (^) must be endothermic. Yet, even though the principle could not be accepted as a general truth, chemists had to admit that in very many cases it represented the facts, that is, it contained a considerable amount of tiTith. Y.vn't Hoff succeodetl in putting things in their proper light. The amount of heat liberated in a chemical reaction represents the total change of the energ\- of the system, and this is what Berthelot rosarded as a measure of the affinity. Van't Hoff rejected this notion and showed that it is the "free energy" gained in a reaction which must be regarded as a measure of the affinitv. By "free energy" we understand the greatest amount of work which the reaction is capable of doing. Now, in order 1,"),S INORGAXIC CHEMISTRY. [§101. to measure the force with which an action tends to proceed, we often make use of an opposing force of knoTvii magnitude, which is just great enough to stop the action. If this opposing force is too small, the internal driving force of the system will overpower it and thereby do a certain amount of work, and this amount of work will be the greater, the greater the counter force that is overcome, or in other words, the smaller the difference between this counter force and the driving force of the system. For measuring affinity we can thus make use of the simple mechanical notions which serve for the measurement of forces in general, as, for instance, in an ordinary weighing. We oppose the force to be measured with another of known but variable magnitude and allow the latter to change until equilibrium is established. There is then equality between the known force and the force to be measured. The free energy is in general not equal to the total energy that comes into play in a reaction; but frequently the difference is not great, as, for instance, in reactions between solid com- pounds or in solution. Herein lies the explanation of the excep- tions to Bbethelot's principle as well as the reason for its agreement with experiment. The total energy-content of a body consists, according to Helm- HOLTZ, of free and bound energy. The free energy alone is capable of transformation into other forms of work. The bound energy is involved in such changes as those of state. Wlien ice melts a con- siderable amount of heat is absorbed which cannot be transformed into work, but only seems to increase the molecular movements of the water molecules. The bound energy of Avater is therefore greater than that of ice at the same temperature. Similarly, there are various other processes where the bound energy is changed. It can be proved theoretically that in every action proceeding of its own accord the free energy must decrease. In the case of an exothermic reaction the e\-olution of heat is due in part to the decrease of the free energy of the system. Further, the bound energy can at the same time either be partly converted into heat, remain unchanged, or increase less than the decrease of free energy calls for; however, if the decrease of free energy in a reac- tion is less than the increase of bound energy, the whole caloric effect must be iiegati\e, which is to say, that the reaction is endo- thermic. §101.] THERMOCHEMISTRY. 150 . For the measurement of affinit\- it is tlu'i'(>loi-c iiocossary to determine this maximum work or fi-ec energy which iw involved in chemical i-oaclioiis. Two moans aiv available, one the deter- mination of the elect i()m()ti\'e force that can be create(l l)y it, and, secondly, the determination of the oi^uililirium constant of the reaction in question. We shall learn in the chapter on electrochemistry that reactions can in many cases be conducted so as to produce an electric current. If the reaction is reversible, it can be brought to a stop Ijy sending a current of the same energy through the system in the opposite direction. The energ}- of an electric current is represented 1j\' the product of two factors, the amount of electricity (expressed in coulombs) and the electromotive force (expressed in volts). Now the decomposition of an equiv- alent amount of each compound requires, according to Faraday's law, the same amount of electricity, namely, 96,540 coulombs per equivalent weight; -whence it follows that the electromotive force must be proportional to the afi&nity; in other words, that the electromotive force is a measure of the affinity. Accordingly, the affinit)' which seeks to bring about a chemical transformation must be opposed b\' an electromotive force just great enough to prevent the reaction. This electromotive force is then the exact measure of the affinity whose action it prevents. The free energy or maximum amount of woi-k which the reaction produces is accordingh' equal to the energy of the electric current produced. The second general method for measuring affinity is applicable in all cases involving a chemical e(iuilibrium. "\A'e learn from thermodynamics that the equilibrium constant A' and the maxi- mum amoimt of work A done by the reaction, bear the following relation to each other; .l=.RTlog, K, when unit concentrations of the reacting substances are in- volved. R is the gas constant (J; 35) and T the ab.solute temperature. A' is also dependent on (i.e., a function of) the temperature. The student will find it inteiesting to leai'u fi'om the ap- propriate text-books of physical cheniisti;\- iiow thi-se two 160 IXOBGAMC CHEMISTRY. [§§ 102- methods are utilized for the j calciilation of affinity in a variety of special cases. The Displacement of Equilibrium. 102. When two systems are in equilibrium with each other (e.g., 2H2 + 02^2H20),the position of this equilibrium is depend- ent on various circumstances. The relationship is expressed by the rule of Le Chatelier: When any system is in physical or chemical equilibrium, a change in one of its equilibrium factors "produces a change in the system, whose effect is opposite to that of the former change. This rule, or theorem, which can be called the principle of the resistance of the reaction to the action, furnishes us with a con- venient means of foretelling in many instances the direction which a reaction will follow. Some examples may be given to illustrate the rule. (1) AVhen a system of water and ice is subjected to increased pressure, the ice melts; that is, that process goes on which involves a contraction, for by this contraction the- system diminishes the pre.-isure exerted on it. (2) ilonoclinic sulphur, when compressed near the transition point (the temperature of equilibrium for ordinary pressure) , passes o^•er into rhombic sulphur, since this process involves a lessening of volume, and in the end also a diminution of pressure, as in the previous case. (3) ^^'hen a solution is diluted the osmotic pressure decreases according to Boyle's law; in the case of a solution of an electro- h'te dilution will be followed by further dissociation, since this increa.ses the osmotic pressure. (4) ^Yhen a liquid i.s heated, more vapor is formed; since the vaporization absorbs heat its effect is opposite to that of the heating. (5) In partially dissociated N2O4 an increase of pressure drives back the dissociation, while diminution of pressure increases the dissociation. The former change carries with it a pressure decrease, the latter a pressure increase. 103. Van't Hoff's principle of mobile equilibrium is a special case of Le Chateliek's rule, but was derived from thermodvnamics 104.] THERMOCHEMISTRY. 101 independently. It says: An equilibrium between two different Mates of flatter (si/stems) diaphires itself under eonstnid pressure bi/ f 11 a —. — of temperature to that one of the lieo si/steins irhose formation rise •' ' J .1 J —, j- heui. A few oxanuilos will servo to make this clear. n allowing an electric spark to pass through, the hydrogen and oxygen unite to form water, which is deposited on the sides of the ^■essel. Inasmuch as 2 vols, hj'drogen combine with 1 vol. oxygen, one-third of the volume that disappeared must have been oxygen. 109. These and other methods of investigation have shown that tJte composition of the air is luvrli/ conslniit. In all parts of the earth, as well as at the highest altitudes which balloons have reached, it c(jnsists of 20. Sl^^, oxygen and 79.10'^^ nitrogen by volume; and 23.01', "" " 76.99^; " " weight. The obser\ed variations from this ratio amount to hardly iCf^p. jMoreo\'er, tin' ciinipositicin docs not appear to change with time; our present analyses agree with those of Du.mas and JVjussinuault made in 1841. This result seems surprising at first thought, because oxygen and nitrogen are constantly being removed from the air and again returned to it and it does not necessarily follow, indeed it is rather an improbabihty, that the losses and gains will exactly balance. 168 lyOiiGAXIC CHEMISTRY. [§109 The oxygen passes through the following cycle: Free oxygen is consumed in all sorts of oxidations of which the mineralization of organic matter is the most important. By the term "minerali- zation " is meant the oxidation of the residues of plants and animals by the oxygen of the air with the aid of bacilh. The carbon of these residues is oxidized to carbon dioxide; the nitro- gen, phosphorus, sulphur and other elements return to the "mineral" state, as nitrates, sulphates, etc. Along with this process there are the other oxygen-consuming processes of the respiration of animals and plants and the burning of fuels, carbon dioxide being formed in all cases. This carbon dioxide is em- ployed by the plants in their process of assimilation, the oxygen in it being again given back to the air. It will therefore depend on the relative magnitude of this process as to whether just as much oxygen gets back into the air as was previously taken up in the formation of carbon dioxide. The oxygen which serves for other oxidations does not necessarily return to the air. Different investigators have attempted to estimate the amount of carbon which annually enters into the cycle of organic life. Dubois calculated that every year the plants assimilate 11S..5 million million kilograms COj, which is ahnost to of the total carbon dioxide in the atmosphere. The amount of CO2 given off by the entire animal world is estimated at 2..5 million million kilograms, which brings us to the startling result that only about 2% of the existing plant material is engaged in the cycle with the animal life. All the rest of the carbon dioxide required by the plants comes from the process of mineralization. The amount of carbon dioxide produced by the burning of coal, etc., is estimated by Crbdner at 1.3 million million kilograms. Nitrogen passes through a cycle too. Most of the nitrogen that occurs in the form of organic compounds in animal and vege- table tissues remains in the combined state after the death of the organism, either as ammonia or as nitric acid or in other nitro- genous products. During the process of decay the combined nitrogen is partially liberated; in the burning of plant and animal remains all of it is set free. On the other hand, certain plants, the Leguminosce, are able by sjonbiosis with bacteria to absorb free nitrogen from the air directly. There are also bacteria which, acting alone, can assimilate nitrogen. Moreover, in storms some §109.] THE ATMOSPHERE. 169 nitrogen combines with oxygen, and again, silent electric discharges, sucli as must frequently pass between earth and clouds, cause the nitrogen to enter into combination. Here the question again arises whether as nuich conies back to the air as goes out. From what has been said it is sufficiently clear that it would be a mere coincidence if exacth' as much oxygen should happen to be withdrawn as is given back. Approximate compensation probably takes place, but, c\-en if it should not, the atmosphere is so vast that its composition would be only slightly affected in the course of centuries. The following calculation will convince one of the soundness of this argument; Tlie normal atmospheric pressure is 760 mm. mercury; this is due to the weight of the air and the moisture in it. Granted that the pressure of the latter a^'erages 10 mm., we have 750 mm. left for the pressure of the air its;>lf ; i.e. the weight of the air is equal to tliat of a layer of mercury 7-")0 mm. thick extending over the entire surface of the earth. This weight can be calculated thus: The volume of the space between two concentric spheres is 4-R'r, if R is the radius of the inner sphere, and r the thickness of that space. The radius of the earth (7?) is, on the average, 6,370,284 m. ; »■ is 0.7-') m.; therefore, taking into consideration the specific gravity of mercury (13.50), Ave have for the desired weight of mercury or air 5.2X10" kilograms. Since 1 m.' air at 0° and 760 mm. pressure weighs 1.2032 kg., the above weight corresponds to a volume of air of 4X10" m.' (at 0° and 760 mm.) or |X10" =8X10" m.^ of o.xygen. lii comparison with this the amount of oxygen which is Avithdrawn from the air in breathing^ burning, etc., is very small, as maj' be seen from the figures on the preceding page for the quantities used by animals and plants. Since, on the other hand, the assimilati-\-e process of the plants yields a considerable amount in addition, th3 variations in the proportion of oxygen in the air must obviously be impercenti')!':' with our present analytical methods. The air is a mixture. It cannot be a compound of nitro- gen and oxygen for the following reasons: (1) the ratio of nitrogen to oxygen is different than it would be for a compound of the two elemeuts, for in the latter case it would have to correspond to the ratio of the ato:iiic wei;;hts or a nuiltiple of the same; (12) by mixing nitro;j;eii and oxygen in the ratio in which thcN exist hi air a synthetical air is obtained which is in c\-ery respect like that around us. (This exclutles the possibility of air containing a per- ceptible amount of a compound of the two elements in addition to 170 IXORGAXIC CHEMISTRY. [§§109- ficc nitrogen and free oxygen.) (3) The ratio of the solubilities of the oxygen and the nitrogen of the air in liquids is the same as that calculated from the solubilities of the pure gases oxj^gen and nitrogen, after taking into account their partial pressures. This could not l>e the case if the air contained a compound of iixygen and nitrogen; (4) when liquid air lioils the first part of the distillate is chiefly nitrogen. The liqttcfuclion of air is now carried on in commerce. The methdds used by Lixde and by Hampsox are based on the same principle, namely, cooling the air by expansion. Further details may be found in text-books on physics. Liquid air is \'er3' mobile and has a bluish tint. It is usually somewhat cloudy because of suspended particles of ice (congealed atmospheric moisture) and solid carbon dioxide. These may be remo\-ed b}- filtration through filter-paper. It boils at about — 190°. It is now extensively used in producing, and demonstrat- ing the effects of, -\'ery low temperatures. When carbon dioxide, for example, is led into a flask containing liquid air, it falls in the solid form like snow-flakes. In spite of its low temperature liquid air can be poured upon the hand ^'sdthout danger; it does not c^-en feel cold (on account of the Leydenfrost phenomenon). Liquid air is much richer in oxygen than the gaseous air of the atmosphere, containing about 50'~"c,. If a glowing splinter is dipped into the liciuid, the wood begins to burn ver\- vigorously producing a ^dolent reaction. It can be preserved for a rather long time in vacuum flasks. B}- fractional distillation of liquid air practically' pure oxygen and nitrogen can be obtained. According to Erdmaxx pure nitrogen is obtained in the cooling down cif liquid air, whereupon nitrogen crystallizes out. ARGON, HELIUM AND COMPANION ELEMENTS. 110. Argon. Despite the fact that air had been already analyzed times without nimiber, it was first discovered in the course of investigations by Rayleigh and Raiis.vy in 1894 that there are other elements in the air than nitrogen and oxygen. One of these, named argon by its discoverers, is even found to the e.xtent of 0.9% by volume, or 1.2ff, by weight. It was on account of its extraordinary resemblance to nitrogen that it no.] ARGON, HELIUM AAD COMPANION ELEMENTS. 171 WHS so long oxcrlooked. The first imlication of its presence was the observation that the specific gravity of the nitroKcn isolaled from the air is somewhat higher than that of the nitrogen jnv- pared from ammonimn nitrite and other compounds. 1 hter of nitrogen from air weighed 1.2.')72 g., while the same amoimt from chemical compounds weighed 1.2r)21 g., in both cases at 0° and 760 mm. There must therefore be anothcM' gas heavier than nitrogen, mixed in with the nitrogen of the aii-. One of the simplest methods for obtaining argon from the air is to heat air with a mixture of 1 g. magnesium, 0.2.3 g. sodium and o gi-. freshly ignited lime. On account of the high temperature free cah'ium is formed: Mg + CaO = MgO + Ca, and it is in such a finely divided condition that it absorbs oxygen greedily and also nitrogen, so that only argon is left. .-Vrgon can also be isolated with the help of calcium carbide. When calcium carbide (better, mixed with 10% calcium chloride) is brought in contact with air at al)out S()l)°, it absorbs both oxvgen and nitrogen: 2CaC., + O, = 2t'a(_) + 4C ; CaC^ + N, = CaCN, + C. This is a suitalile method for preparing argon in large quan- tities. After argon had been once discovered it was found elsewhere than in the atmosphere; some mineral waters contain it in solution, pertain rare minerals yield it when heated, etc. Argon is a colorless, odoiless gas, having a vapor density of 19.957. It has been (Minclensed to a coloiless liquid, that boils at — ls(i.0°. b\- cooling with boiling (ixya,en and compressing to about oO.e atnKjsphercs; it sohdifies at — IS!). 6° It is somewhat more soluble in \\ater than is nitrogen (0.()57SO parts in 1 vol. at 0° and 760 mm. pi-essiii-e). .\s to its chemical nature, it is interest- ing that no one has }et sncceed('(l in prei)aring a c(nnpound of argon. It is certain that what is now called argon is neither a mixture nor a compound, but an element. The boilin,g-point and the melting-point are constant, and the va)>or ])ressure of, argon li!«>- 172 INORGAXIC CHEMISTRY. [§§110- wise remains constant during the liquefaction, so long as any gas is present. Moreover, when a certain volume of argon is three- fourths dissolved in water, the undissolved gas shows exactly the same spectrum as the dissolved. All of the above are charac- teristics of a homogeneous substance. The extraordinary stability of the gas in the presence of all sorts of reagents is a strong argument against its being a compound. III. After the di.scovery of argon R.vmsay and Tkavers detected four other rare gases in the atmosphere, though their quantity is very small. These are helium, neon, krypton, and xenon. In a spectroscopic inves- tigation (§ 26.5) Norman Lockyer had detected in the atmospheres of the sun and many fixed stars considerable quantities of a gas unknown on the earth; he named it helium. In 189.5 Kamsay and Travbrs succeeded, however, in obtaining it in small amounts on heating the rare mineral cleveite. Afterward it was also met with as a companion of argon in certain other, chiefly uraniferous, minerals as well as in mineral springs, for instance, those of Bath; and at last it was also discovered in the air. At ordinary temperatures helium is a colorless gas. It is of all gases the most difficult to condense; yet Kamerlingh On'nes recently achieved the task. HeUum boils at 54° absolute temperature (— 269°C.). By quickly evaporating it a temperature of ^ 1 .6° absolute was reached, which is the lowc:!t thus far attained. The critical pressure is 5.3 atmos- pheres. In water helium is less soluble than argon . Its relation to radium is discussed below in connection with the latter element. Hehum and neon (0.00086 wt. % of air) are found in the most volatile part of liquid air. Dewar proved that hehum and neon can be isolated directly from the air by bringing the air in contact with ignited charcoal at —185° The charcoal has the curious property of condensing in its pores all the other gases of the air, and a gaseous residue is here obtained which shows clearly the spectral lines of He and Ne. While hehum and neon were found in the most volatile part of the air, krypton and xenon were obtained, on the contrary, from the residue, after a large quantity of liquid air had been allowed to evaporate slowly. Their separation was rendered possible by the fact that krypton still has a rather large vapor tension at the temperature of Hquid air, while the vapor tension of xenon is then imperceptible. Both these elements occur only in extremely small amounts in the atmosphere. Krypton makes up 0.028%, xenon 0.005% (by weight) of the air. In the following table some of the data of these elements are given. The elements form a natural group. §111.] ARGOX. HELIUM, A.\D COMJ'AMOX ELEMENTS. 17:5 Density (0=l(i) Atomic weight Boiling-point at 760 mm.. llcliu l.!)S 4°:il)s 10.1 20 •_> I 19. iU 39,S,S ,Mi.ll°ubs, I\r\\'|)ton. 41 4r> S2 9 121.9° abs. XctKin. or. 1 130.2 163.9°abs. These gases ha\p three properties in common whii-li are worthy of- mention here. In the first place they display characteristic spectral lines in Plucker tubes (S 2(i3), whereby it has been possible to recognize them and to judge of their purity. In the second place, no one of these elements has been found to enter into combination with other elements; they ma}' therefore be considered nuUixalciit. In the third place, their molecule consists of only one atom. This fac't could not be discovered In the ordmary wa.\-, described in §§33 and 34, because of the entire absence of compounds for investigation. It has, however, been possible to ascer- tain it from the molecular heat of the gases. This is the amount of heat that must be imparted to a gram molecule of a gas in order to raise its temperature one degree. This quantity of heat differs, accorduig as the gas is under constant pressure t.ir under constant volume. It is greater in the first case because under constant pressure the gas expands on heating and so does work, which evidently in^-oh'es an expenditure of heat. We saw in § 34 that for one gram molecule of a gas the equation PV = 2T is applicable, the 2T expressing in calories the external work done when a gas under constant pressure P increases its volume by V, or when a gas being generated under the pressure P comes to occupy a volume T^ If the temperature is raised one degree we have PF = 2(7' 4-1); for each gram molecule of gas extra work is therefore done equivalent to 2 calories. The molecular heat at constant pressure is thus 2 cal. more than that at constant volume. From the kinetic theory of gases it can be deduced that the molecular heat of a monatomic gas at constant pressure is 5 cal. At constant \-olume it must be 2 cal. less, or 3 cal. The ratio of these quantities of heat is therefore 5:3 = 1.66. When the molecules of the gas consist of more than one atom, more heat is absorbed for the same rise of tem- perature, because heat is then used not only for the movement of the molecules, but also for that of the atoms in the molecule. The ratio then becomes 5+m:3+m, if m is the additional heat. The resulting ratio is thus less than 1.66. By determining this ratio (which can be found from the velocity of propagation of sound in the gas by a well-known i)hysical formula) we can ascertain whether the gases are monatomic or poly- atomic. For the gases of this group the ratio v/as found to be 1.66, proving that their molecules contain only one atom. 174 INORGANIC CHEMISTRY. [§112. Compounds of Nitrogen and Hydrogen. 112. Until recent years only one compound of hydrogen and nitrogen has been known, viz., ammonia, NH3. At present, how- ever, we know five: the others being hydrazine, N'jH^, hydrazoic acid, N3H, and the compounds of the latter with ammonia and with hydrazine (NH3 • N3H and NjH^ ■ N3H) . Of these five compounds, however, ammonia is by far the most important. AMMONIA. The material now used for obtaining ammonia is the " ammonia liquor" of the gas-factories and coke ovens. The gases that are given off in the dry distillation of coal are passed through water, which dissolves the ammonia. In order to obtain a pure ammonia, the ammonia liquor is heated v^ith mUk of lime and the expelled ammonia is led into concentrated sulphuric acid. In this wa}' crystallized ammonium sulphate is obtained. It is purified bj^ recrystallization and again distilled with lime to recover the free ammonia. Ammonia can be prepared synthetically by the following methods. The direct synthesis from the elements was given above (§ 107). There are also examples of its formation by the direct decomposition of its compoimds. Thus we obtain it by heating the ammonia compounds of certain salts, as xCaCl.^ • ^NH3 and a;AgCl • //NH3. A number of organic compounds yield nitrogen in the form of ammonia on heating. Moreover, ammonia results from the action of hydrogen on certain nitrogen compounds, as, for example, when nitric acid, HISr03, comes in contact with nascent h}'drogen (generated from zinc or iron filings and dilute sulphuric acid), or when a mixture of nitric oxide, NO, M-ith hydrogen is passed over platinum black: 2N0 + 5H2 = 2NH3 + 2H2O. The formation of ammonia by the action of free nitrogen on hydrogen compounds has not been brought about, but the gas can be produced by the interaction of a hydrogen compound with a nitrogen compound. An illustration of this is the decomposition of magnesium nitride by water: MggNa -1- 3H2O = 2NB3 -I- 3Mg0. § 112..] AMMOMA. 175 The putrefaction of organic matter (frcccs, urine, etc.) evolves ammonia. By the action of oli^ctric sparks on moist air am- monium nitrate is produced. Tliese hist two methods of for- mation are responsible for the slight traces of ammonia in the air. For the formation of ammonia fi'om calcium cyanamide, see Org. Chem., § 2GC). Physical Projiciiics.— Ammonia, at ordinary temperatures is a gas with a characteristic odor, that excites one to tears. Its specific gravity is s..") (0=16) or 0..")S!) (air = 1) ; 1 1. XH, at 0° and 7G0 mm. pressure weighs 0.7(]193 s. It can be easily liquefied; it boils at —33.7° and becomes solid at —7-5°; it then forms white translucent cr\'stals. It is extremel}- s(.iluble in water; at 0° and normal pressure 1 vol. H^O dissolves Ills ^'()ls., or ().S7.5 parts Ijy weight, of XH3. The specific gravity of the solution of ammonia in water grows smaller as the concentration increases. The evaporation of liquid ammonia involves a considerable dei)rc3- sion of temperature. This is the principle of most of the ice- machines now in use. Chcmicnl Propi iiics.—'n\e characteristic property of this com- compoimd is that it combines with acids directly to fonn salts: NH3 + HC1 = XH4C1. XH3 + H.\()3=NH4N()3. Ammonium Ammonium chloride. nitrate. 2XH3 + H2SO4 = (XH4),„S( )4. Amnionium sulphate. In these salts (which are almost all readily soluble in water) the atomic group NH4 plays the part of a metal; they correspond in every respect to the compounds KCl, KXO3, K2S(J4, etc. The group, or radical, XH4 has been gi^■en a particular name; it is called ammonium. Mine than one attempt has been made to isolate this ammonium, but ah\ays in ^-ain. However, when sodium amalgam comes in contact witli a concentrated ammo- nium chloride sohUion, the mercury swells to a soft spongy mass that rapidly decomposes at ordinary temperatures into ammonia and hydrogen and is in all probability, therefore, ammonium amal- gam. If sodium amalgam is allowed to leact with ammonium iodide dissolved in liquid ammonia at -30°, a hard metallic mass 176 INORGANIC CHEMISTRY. [§§ 112- is obtained, which swells with rising temperature because of decom- position into mercury, hydrogen (1 vol.) and ammonia (2 vols.) : 2NH4=2NH3 + H2. The aqueous solution of ammonia reacts strongly basic; so do the moist fumes of ammonia. AVe must therefore assume that this solution contains a compound NH4OH, ammonium hydroxide, and hence also the ions NH4 and OH in analogy with other soluble bases, e.g. potassium hydroxide, KOH. As a matter of fact, however, the solution of ammonia conducts the electric current much more poorly than a solution of sodium hydroxide of equiva- lent concentration (§ 2?A). Ammonium hydroxide has not yet been isolated. When the solution of it is evaporated, NH4OH splits up into NH3 and H2O. Concordant herewith is the well- known fact that ammonia can be entirely expelled from its aqueous solution by boiUng. Ammonia does not burn in the air but does in oxygen; in addi- tion to water and nitrogen traces of ammonium nitrite, NH4NO2, and nitrogen dioxide, NO2, are also formed. A mixture of ammonia and oxygen explodes violently when it is ignited. The oxygen conveyed by soU bacteria may also cause the oxidation of ammonia, producing nitric acid. Chlorine takes fire when passed into ammonia, forming nitrogen, No, and hj^drochloric acid; the latter then unites with the remaining ammonia to form sal am- moniac, NH4CI. The hydrogen of ammonia is replaceable by metals. Magnesium, e.g. bums in ammonia, forming magnesium nitride, Mg3\2- When ammonia is conducted over hot potassium or sodium, potassium amide, NH2K, or sodium amide, NH2Na, is formed. These and analogous metal compounds are decomposed by water, yielding ammonia again and also metal oxide or hydroxide. At high temperatures (produced by induction sparks) ammonia s]Hits up almost completely into its elements, the volume being doubled: 2NH3 = N2 + 3H2. 2 vols. 1 vol. 3 vols. On the other hand, nitrogen and hydrogen can unite to form ammonia under the influence of induction sparks (§ lOG). Equilibrium 114.] HYDRAZhXE, OR DIAMIDE. 177 is reached when 3^ of ammonia is formed in the gas mixture, Nj + SHj, i.e., 6'" of the theoretical yii'ld. Tliis is the reason why ammonia cannot be split up by electric sparks to more than 07%: 2NH3F yiTh'OGEX WITH THE HALOGENS. 17'.t soluble ill strong nuneral acids. They are also very explosive, hence e.xtremely dangerous, the .sodium salt being the least so. An aqueous ITi solution of the acid is only U.OOS ionized; it is thus a rather weak acid; it gives off iiydrogen in contact with many metals, e.g. Zn, Fe, Cd, and Mg. It is characteristic of the metal hydrazoates (or "azides") that the)' crjstallize anhj-drous and yield the pure metal when heated. Compounds of Nitrogen with the Halogens. ii6. When chlorine gas is allowed to act on a concentrated solution of ammonium chloride, most conveniently by inverting a flask full of chlorine over the warm (30°-40°) solution, oily drops are formed, which are li(^st collected in a leaden saucer placed under the miuitli of the flask. These drops contain some hydrogen as well as nitrogen and clilorine. T^y treating with chlorine once more pure nitrogen trichloride, NCl.j, is obtained as a yellowish oil with a disagrt'cable pungent odor and a specific gravity of 1.65. This is one of the most dangerous of substances, because it explodes in a most \-iolent manner, not only on contact with certain organic substances (e.g. turpentme) , but verjr often spontaneously. It dissoh'es in carbon disulphide, benzene and othei- solvents, form- ing yellow solutions. These solutions are relati\'ely harmless; they decompose in the sunlight. Concentrated h3^drochloric acid decomposes nitrogen trichloride according to the equation: Xria + 4HC1 = XH4CI + 3CI2; aqueous ammonia also breaks it up in a simUar way: X( 'I3 + 4XH3 = 3NH4CI + N2. Nitrogen trichloride is strongly endothermic: X + 3Cl-NCl3=-41.9Cal. When a .solution of ^0(hum azide, XaXj, is nii.\cd with .a solution of sodium hypochlorite in the iclatiou of molecule for molecule and the mixture is acidified, tlic liquid assumes a yellow color and gives off a colorless gas with an odor like that of hypochlorous acid and ha\-ing the composition NjC'l, showing it to be chlorazide. On being passed 180 INORGANIC CHEMISTRY. [§§ 116- into caustic soda it forms sodium azide and sodium hypochlorite in equivalent amoimts : N,C1 +2NaOH = NaN3 +NaOCl +H2O. The chlorazide is likewise extremely explosive. 117. Nitrogen Iodide. — If a solution of iodine in potassium iodide is mixed with ammonia solution, a precipitate is usually obtained of the composition NI2H; if the conditions are slightly altered another compound, X2I3H3 (i.e. XH3 + NI3), is deposited which breaks vip on continued treatment with water into ammonia and nitrogen tri-iodide. These compounds are likewise very explosive. Another method is to digest pulverized iodine with ammonia water. The product so obtained is still more explosive, often exploding even when damp or when it is being washed with water or by the action of hydrochloric acid. In the presence of ammonia solution it is stable. Nitrogen iodide is decomposed by dilute hydrochloric acid, forming ammonia and chlorine iodide: NH2l + HCl=NH3 + ICl. Nitrogen iodide Ls also decidedly endothermic. Hydroxylamine, NH2OH. 118. Hydroxylamine is a reduction product of many oxj'gen compounds of nitrogen intermediate to the formation of ammonia; e.g. it is formed when tin acts on dilute nitric acid. Here the nascent hydrogen effects the reduction: HNOs + 3H2 = .\H3( J + 2H2O. It is manufactured by ihe aiectrolytic reduction of nitric acid dis- solved in sulphuric acid. The free hydroxylamine is best prepared by heating the phosphate. It is a crystallized solid, melting at 30° and boiling under 60 mm. pressure at 70° When heated in the air it explodes with a yellow flame. 119.1 NITROUS OXIDE. 181 Hydroxylamine is easily soluble in water; its solution reacts strongly alkaline. It forms salts in the same way as ammonia, i.e. by direct addition of the acid: NliaOH-HCl, NHaOH-HNOa, etc. These salts are rather stable; the hydrochloride, however, must be i)reser\-ed o\-cr lime, else it slowly decomp()S(>s, for the following reason. The salt is split up to a ^•cry small degree into hydrochloric acid and hydroxylamine. Now free hydrochloric acid accelerates catalyticall}' the decomposition of the salt. A\'Tien, however, the hydrochloric acid is absorbed by the lime, the decomposition becomes so slow that it is imperceptible. The free hydroxylamine and its aqueous solution are somewhat unstable, especially in the presence of alkalies; it decomposes easily into ammonia, water and nitrogen. A further characteristic of hydroxylamine is its great reducing power; it precipitates reddish-yellow cuprous oxide from an alkaline copper solution at ordinary temperatures, even when strongly diluted; mercuric chloride, HgCl2, is reduced to calomel, Hg2Cl2; silver nitrate to silver, etc. The following reaction is also peculiar: A solution of ferrous sulphate is precipitated with an excess of sodium hydroxide and warmed; if hydroxylamine (or one of its salts) is now added to the green fen'ous h\-droxide, red ferric hydroxide is formed verj' quickly, the hydrox- ylamine being reduced in this alkaline solution to ammonia. On acidify- ing, an acid solution of a ferric salt is obtained; if this is treated with a hydroxylamine salt, it is suddenly decolorized because of reduction to ferrous salt, the hydroxylamine being now in the oxidized condition in the acid solution. Compounds of Nitrogen with Oxygen. Those included under this title are: nitrous oxide, N2O; nitric oxide, N(0; nitrogen trioxide, (ir nitrous anhydride, N2O3; nitrogen dioxide, NO2, or tetroxide, N2O4, and nitrogen pcntoxide, or nitric anhydride, N2O5. NITROUS OXIDE, N2O. 119. This compound cannot be obtained directly from its ele- ments; the ordinary method of preparation consists in heating ammonium nitrate to about 2.50°: NH4N03=N20 + 2H20. 182 INORGANIC CHEMISTRY. [§§ 119- This method is analogous to that of preparing nitrogen from ammonium nitrite (§ 105). If the nitrate is heated above 250°, the gaseous product partially decomposes. Physical Properties. — Nitrous oxide is a colorless and odorless gas, which when liquefied boils at —87° and solidifies at —102°. The evaporation of the liquid produces a great depression in the temperature, which may even reach —140° under reduced pressure. Its specific gra-\dty is 1.52 (based on air), or 21.89 for 0=16. 1 1. No*.) at 0° and 760 mm. pressure weighs 1.9657 g. It is rather soluble in water (1 vol. H2O dissolves 1.305 vol. N2O at 0°); hence it must be collected over hot water. In alcohol it is still more soluble. Chemical Properties. — Nitrous oxide supports combustion. Phosphorus, carbon and a glowing splinter burn in it as in oxygen. A mixture of nitrous oxide and hj^drogen explodes like detonating- gas when it is ignited, only not quite so loud. These properties might lead one to confuse it ^ith oxygen on a superficial examina- tion. However, it is verj^ easily distinguished from the latter by the fact that it gives no red fumes when mixed with nitric oxide (§ 120) and always leaves residual gas (nitrogen) after a combustion. A faintly burning piece of sulphur is moreover extingmshed by nitrous oxide. Nitrous oxide is endothermic: 2N + 0— N20= -17.7 Cal. Berthelot has made the general observation that endothermic substances can suffer an explosive decomposition; in this case this may be brought about by toucliing off the gas with fulminating mercur}'. It is easy to explain Berthelot's observation. When an endothermic substance decomposes, heat is evolved. Now, we saw in §§ 13 and 104 that chemical reactions are accelerated in a very high degree bj^ rise of temperature. Suppose that a sudden decomposition is caused at a certain point in a mass of an endo- thermic compound. The heat given off raises the temperature of the surrounding molecules and they too split up suddenly, evolving still more heat, and so on. The whole mass will thus reach a condition of sudden decomposition, that is, it will explode. To bring this about it is only necessary that the first impulse be vigorous enough for the sudden decomposition of so many mole- cules that the heat evolved is sufficient to raise the surrounding ones to the temperature of decomposition. 120.] NITlilC OXIDE. 1S3 Composition. — Under the protracted action of induction sparks the gas splits up into a mixture of nitrogen and oxygen, the volume of which is half again as groat as that of the nitrous oxide. When potassium and sodium arc burned in the gas, potassium and sodium oxides respecti^•cly are formed, together with nitrogen; the gas volume after cooling is unchanged. Both of these observations point lO the same formula, N^O, and this is confirmed bj' the fact that the relati^'e density of the gas, \Ahich should theoretically be was found to be 21.89. NITRIC OXIDE, NO. 120. This gas is only obtained by the reduction of nitric or nitrous acid. The ordinary method of preparation is by allowing copper to act on nitric acid or else by covering copper (in the form of thin sheets) with a saturated solution of saltpetre and adding concentrated sulphuric acid drop by drop (§ 127) : 3Cu + SHX03 = 3Cu(X03)2 + 4H20 + 2XO. In this reaction the hydrogen, which would be expected to be given off, reduces another portion of the acid. In order to prepare nitric oxide by the reduction of nitric acid or a nitrate a boUing-hot solution of ferrous chloride, FeCU, in hydrochloric acid is found very satisfactory; the ferrous chloride is converted into the ferric chloride, FeCls, by the reaction: HNO3 + SFeCU + 3HC1 = SFeClg + 2H2O + NO. Perfectly pure nitric oxide is obtained by treating a mixture of yellow prussiate of potash and potassium nitrite with acetic acid: 2K,Fe (CX)„ + 2KNO2 + 4C,H,0j = K^Fe^CCX )i2 + 4KC,H302 + 2H3O + 2N0. Yellow prus- Pot. ni- Acet. acid. Ked prua- Pot. acetate siate. trite siate. Physical Properties. — Nitric oxide is a colorless gas, whose specific gravity has been found to be 1.039 (air = l). It can be condensed to a blue liquid, which boils under ordinary pressure at 184 INORGANIC CHEMISTRY. [§§ 120- — 153.6°. The critical temperature is —93.5°, the critical pressure 71.2 atm. It is not very soluble in water, but dissolves easily in a solution of ferrous sulphate, FeS04; strange to sa}', this solution is quite dark brown in color, although the ferrous salt solution is pale green and nitric oxide colorless. The compound which is formed here has not been isolated, but it has been shown to consist of FeSO^ and NO in equimolecular proportions. Chemical Properties. — It is characteristic of this gas, above all other properties, that it combines with oxygen immediately, forming nitrogen dioxide, a reddish-brown gas. On heating it with hydrogen no explosion occurs; the mixture burns with a white flame, forming water and nitrogen. If burning phosphorus is introduced into the gas, it continues to burn; a lighted candle is, however, extinguished; sulphur and charcoal do not burn in it either. A mixture of nitric oxide and carbon disulphide burns with an intensely luminous blue flame, that is very rich in chemically effective rays. Nitric oxide is a strongly endothermic compound; it can be made to explode by fulminating mercury (§ 1 19). According to § 103 XO must be formed at a high temperature. Nernst proved that the reaction N2 + 02<=i2NO accords strictly with the law of mass-action (§ 49); from whichever side one starts, the results are in agreement with those calculated, assuming both reactions to be bimolecular. See further § 127. The formation and the decomposition of NO are much slower than in the case of ozone. Accordingly a short heating of air, followed b)' a rapid cooling, produces ozone, while a slower heating and cooling yield NO, ozone being broken up during the extended period of cooling. The following experiment illustrates this: When moist air is directed with a velocity of less than 7 m. per sec. against an incandescent Nernst filament (§ 291) XO is formed; a more rapid current gives ozone. Composition. — When sodium is heated in contact with a measured amount of nitric oxide, sodium oxide and nitrogen are formed; the latter takes up exactly half the volume of the original gas. The specific gravity of nitric oxide is 15 (H = l), hence its molecular weight is 30. According to the above decomposition the gas con- tains one atom of nitrogen (14 parts by weight). There remain for the oxygen, therefore, 16 parts by weight, i.e. just one atom. Hence the formula is XO. 122.] XITROGEX DIOXIDE AXD TETROXIDE. 185 Since nitrogen is trivalent or quinciuivMlcnt (tho latter in ammonium salts, e.g. XHjCl) and oxygen is bivalent, it must be assumed that there is a free valence bond in X(^, i.e. -X =0. The same applies to NO,. Free bonds like these ai-e very I'are. Nitrous Anhydride, NjO,. 121. Upon allowing nitric acid of 1.3 specific gravity to react with arsenic trioxide and drying very carefully the gas that comes off and finally condensing this gas, a liquid is obtained of the composition N2O3. At ordinarj' temperatures the liquid is green but below -2° it is deep indigo-blue. When cooled by liquid air it solidifies in dark blue crystals. 2 X 14 +3 X16 The compound N2O3 should have the vapor density ^ = 3S. In a perfectly dry state its vapor density was found to vary in a series of experiments between 38.1 and 62.2, so that the compound appears in that condition to be partly polymerized. The least trace of moisture, however, causes a dissociation into NO2 and NO. For instance, it was found to be sufficient merely to leave a small bulb full of it with a capillary tube open a few seconds in the air; upon resealing the tube the vapor density was found to have fallen to 28. 2. This very striking property of water, whereby even the slightest trace of it brings about dissociations which are not observed in the perfectly dry state, wiU be met mth in several examples in later chapters. The phenomenon was discovered by Brereton Baker. NITROGEN DIOXIDE AND TETROXIDE, NO, AND NjO^. 122. Nitrogen dioxide Ls formed from nitric oxiilc plus oxygen, or more conveniently by heating well-dried lead nitrate: PbCXOs) 2 = PbO + + 2XO2. When so prepared it is a very deep-brown gas. On leading it into a strongly cooled vessel it condenses to a bright-yellow liquid, which solidifies at -20° to colorless crystals, that melt at —12°. The color becomes darker on warming and at +2lj° the licjuid begins to boil, changing back again into the brown yas. The vapor density of this gas at 26° is found to be oS.O, wliilc that cal- culated for N:.()4 is 4o.',) and that for NOa 22.<) (H = l). Since the value found is betwo(>n the two, it may be assumed that at this temperature the vapor consists partly of N-^t )j molecules and partly of NOo molecules. A simple calculation indicalos the ])ercenla.n(! 186 INORGANIC CHEMISTRY. m 122- of the former to be 34.4'^,. As the temperature rises, the vapor density steadily decreases till about 150° is reached, when it becomes constant at 22.9. There is evidently complete dissociation of N2O4 molecules in this case, Xo04^2X02, 1 vol. 2 vols. and, inasmuch as the color of the gas grows darker, we must sup- pose that NO2 is dark brown, while N2O4 is colorless, which is true of the latter in the solid state. This supposition is supported by the fact that not only can the degree of dissociation be estimated from the intensit}' of the color, but that it can even be measured quantitatively in this way. According to § 51 the equilibrium between the two gases is expressed by the equation P — x = kx^, where P is the total pressure of the gas mixture and x that of the dioxide, k being a constant. From this equation it follows that the dissociation (at a constant temperature) depends on the pres- sure (§ 51), which has been shown to be the case. This a!.i0 fol- lows from the theorem of Le Chatk.lier (§ 102). On bringing nitrogen tetroxide in contact with water or, better, with alkalies, nitrous and nitric acids are formed; we may therefore consider it as a mixed anhydride of these two acids : JJfj3>o+H20 = N02-OH + NO-OH. -'■^ ^ Nitric acid. Nitrous acid. Both NO2 and N2O4 possess strong oxidizing power; many substances burn in their vapor; they precipitate iodine from solu- ble iodides. The composition of nitrogen dioxide follows from its synthesis- equation, 2NO + O2, and from the vapor density. Nitrogen Pentoxlde, N2O5. 123. This compound can be obtained by the action of chlorine on silver nitrate or by distilling fuming nitric acid with phosphorus pent- oxide. It is a colorless crystalline solid. It inelts at 30°, and at 45-50° breaks up, sivinj; off brown fumes. If the heating takes place rather 125.] HYJ'ONITROifS ACID. 1,S7 rapidly the decomposition is explosive in nature; sometimes a spon- taneous ex|ilosioa (akes i)la('e, lience it can not be kept long. As nitrogen pentoxide is strongly endothermic, its sjiontaneous explosion must be explained in the same way as is indicated in § 119. Only wo must conclude in this case that the decomposition at ordinary temperatures is vigorous enough to sufficiently heat the neighboring molecules. It unites with water, forniina; nitric acid with the evolution of much heat. As might be expected, it has strongly oxidizing properties. Phos- phorus and pota'^sium, for instance, burn with great brilliance in the slightly warmed anhvd.ride. The composition of nitrogen pentoxide is ascertained by hAating with powdered copper; the amount of nitrogen evolved corresponds to the formula X^Oj Oxygen Acids of Nitrogen. 124. Four acids of nitrogen are known: hijponitrous acid, H2N2O2; nUrohijdroxiilavdmc acid, H^NoOs; nitrous acid, HNO2; nitric acid, HXO3. The nitrous acid is known only in dilute aqueous solution; nitrohydroxylaminic acid is known only in its salts; but the others are known in the pure state. Only certain ones of the above nitrogen oxides can be regarded as acid anhydrides. The pentoxide is undoubtedly one and the tetroxide may be considered as a mixed anhydride of nitric and nitrous acids (§ 122). Xitrogen trioxide gives a solution of nitrous acid when mLxed with ^\■ate^ at a low temperature; however, this solution undergoes a decomposition slowly at ordinary, more rapidly at higher, temperatures, nitric acid and nitric oxide being formed: 3TIN( )2 = HX( )3 + 2N(J + HoO. The acid corresponding to nitric oxide, NO, is nitrohydroxyl- aminic acid. However no one has yet been able to obtain this acid fi(jm nitric oxide and w atcM-. The same is true for nitrous oxide, to which hyponitrous acid corresponds. Hyponitrous Acid, HjNjOj. 125. This acid is formed wIiimi nitidgen trioxide is introduced into a methyl-alcoholic solution of hydroxylamhie. The I'lee acid does not liberate iodine from potassimn iodide at cmce; the reaction is dcl.iyed 1S8 INORGANIC CHEMISTRY. [§§ 125- for a time, probably on account of a decomposition, by wMch nitrous acid is formed. Hyponitrous acid belongs to the class of weak acids; its aqueous solution is a poor conductor. Both neutral and acid salts of this acid are Icnown. Nitrohydroxylaminic Acid, H^N^Oj. This acid does not exist in the free state, being known only in salts. Its sodium salt is obtained by mixing an alcoholic solution containing sodium alcoholate and hydroxylamine with ethyl nitrate: C^HjONOs + N H2OH = C2H5OH + H2N2O3. Ethyl Nitrate The alcoholate is added in order to convert the free acid directly into its sodium salt. If the attempt is made to liberate it by adding a stronger acid, it is immediately decomposed according to the equation: Na.XA +2HCl = 2XaCI +2X0 +H2O. The sodium salt, heated in aqueous solution, gives sodium nitrate and nitrous oxide. When the sodium salt is heated dry until it begins to melt, it is decomposed into nitrite and hyponitrite: 2Xa2N,03 = 2XaN02 +Na2NA- NITROUS ACID, HNOj. 126. It was remarked above that this acid is only known in dilute solution at ordinary or low temperatures; its salts are, however, stable. In order to prepare them we usually employ potassium or sodium nitrate, which gives off oxj^gen when heated and is converted into nitrite. This decomposition takes place more readily if lead is added during the heating as a reducing agent: 2KN03 = 2KN02 + 02. i-''J NITRIC ACID. ]89 Its salts are all easily soluble in water, with the exception of silver nitrite, AgN02, which is rather difficultly soluble at ordinary temperatures; it is obtained as a yellow crystalline precipitate, when not too dilute solutions of silver nitrate are mixed with a nitrite. The addition of strong sulphuric acid to a nitrite at once pro- duces red fumes; in this way a nitrite can be distinguished from a nitrate, for the latter does not produce them. It may be assumed that in this reaction free nitrous acid is primarily formed; this is, however, broken up directly into water and nitrogen trioxide, the latter of which at once sphts up again into iS^02 + N0; thereupon the nitric oxide unites immediately with the surrounding oxygen to form dioxide. The red fumes thus consist soleh' of nitrogen dioxide, NO2. On treating a very dilute nitrite solution with the equivalent amount of sulphuric acid a dilute solution of free nitrous acid is obtained. This solution can act either oxidizing or reducing. As examples of the former action we have the Uberation of iodine from a solution of potassium iodide, the oxidation of sulphurous acid in dilute solution to sulphuric acid, the oxidation of ferrous sul- phate, I'"eS( )4, to ferric sulphate, Fe2(S()4)3, and the conversion of the j-eUow to the red prussiate of potash. In all of these cases lower oxides of nitrogen, chieflj^ nitric oxide, are formed. An example of its reducing action (in which nitrous acid is oxidized to nitric acid) is the bleaching of potassium permanganate, K^In04, in sul- phuric acid solution: 2KMn( )4 + 5HX()2 + 3H2SO4 = K2SO4 -I- 2MnS04 + 5HNO3 + 3H2O. This last reaction offers a means of determining quantitatively (volumetrically, see § 93) the strength of a dilute solution of nitrous acid. miRIC ACID, HNO3. 127. This is the best known acid of nitrogen. It is manufac- tured on a large scale, since its uses are many and varied; in the organic dyestuff industry, for example, large quantities are employed. The commercial process of manufacture depends on the decom- position of Chili saltpetre, NaNOs, by strong sulphuric acid: NaNOs + H2SO4 = NaHSOi + HNO3 One of the simplest methods of carrying it out is as follows: 190 IXORGAXIC CHEMISTRY [§ 127- In the cast-iron retort {C, Fig. 35), saltpetre and sulphuric acid (chamber-acid) are mixed in proportions corresponding to the above equation, a shght excess of sulphuric acid, however, being added, because this makes the residue easier to remove from the retort. The retort is connected ^\ith a row of earthenware bottles (EE') containing a little water. These receive the distilled acid. The last bottle connects with a coke tower through which water is jSiil n * ■' ' --.- "' ■■}ii^^JJ^jiS&!i&^^%), Fig. 35.' — Manufacture op Nithic Acid. trickling down to dissolve the uncondensed acid vapor. By this process a liquid of a specific gravity of 1.35 and containing 60% acid is obtained. If the saltpetre is previously dried and concen- trated sulphuric acid is used, a nitric acid of sp. g. 1.52 and almost 100% pure can be obtained. In some cases two molecules of saltpetre are used to one of sulphuric acid. If heat is moderately applied, the reaction pro- ceeds according to the above equation, but on heating to a higher temperature the acid sodium sulphate that is formed acts on the second molecule of nitrate, also forming nitric acid: NaNOg + NaHSOi = NagSO* + HNO3. A large part of the nitric acid, howe-\-er, dissociates at the same time as follows: 2HX03=2N02 + H20-t-0. The N02-fumes dissolve in the distillate. The liquid thus obtained is red and its specific gravity is 1.52-1.54; it fumes strongly in the air and is known as "red fummg nitric acid." 127] NITRIC ACID. 191 For some years the distillation of saltpetre with sulphuric acid has been carried on "n a vacuum. The yield of acid in such a case approaches closely to the theoretical and the product obtained is entirely free from nitrous fumes. An entirely distinct method for the industrial preparation of nitric acid was inv nted a few years ago by Birkeland and Eydi;. They make use of the nitrogen and oxj-gen of the atmos- phere. The problem of making nitric acid from this rather in- exhaustible source has been studied for many years, but these men are the first to handle it with success on a commercial scale. The solution of the problem became a really pressing matter, because the principal material Tor the preparation of nitrogen compounds, Chili saltpetre, bids fair to be exhausted in thirty- five j'ears, and saltpetre has not only a large significance in the industrial world but a still larger one in agriculture as a nitro- genous fertilizer. The method of Birkeland and Eyde is based on the long established fact that oxides of nitrogen are formed in an electric arc burning in the air. The reason why previous investiga- tions did not succeed lies in the fact that an ordinary electric arc has too small a volume, and therefore, cannot let a sufficient quantitj' of air pass. This dif- culty is now obviated bj- mount- ing the arc between the poles PP of a ^•ery powerful electric magnet EE (Fig. 36). The arc is produced between two hol- low bars of coppei-, which are kept cool by circulating Avater in them. "When an alternate current is used for producing the arc, the latter spreads out Fiq. 36.— Diagram op Birkeland and in the shape of a flat disc that E^'db Nitbic Acid Apparatus. reaches a diameter of 2 m. in the industrial form of the apparatus; the tension employed is 5000 volts. This flame disc is inclosed in a box through /''' -■^\'* 192 INORGANIC CHEMISTRY. [§ 127. which a rapid current of air is forced, and the contact with the flame is sufficient to form somewhat more than 1% of NO. Instead of broadening out the electric arc to a sun-shaped disc by the action of powerful magnets, Schonherr (Badische Anilin- und Sodafabrik) forms an arc in the inside of an iron pipe through which air is passed. Under these circumstances the arc is developed in a peculiar manner. When the current is turned on, the arc forms at the first instant in the lower part of the metal pipe, between the pipe itself, which serves as an electrode, and a second electrode, which is separated by only a few millimeters from the lower end of the pipe. Forthwith, however, the arc is carried along upward in the pipe by the current of air, which is given a tangential motion as it is passed into the pipe, so that the arc comes to occupy the portion of air along the axis of the pipe and does not touch the wall of the pipe (or the efHux end of the pipe or a specially devised separate electrode) until a considerable distance from the lower electrode is reached. Thus there is established in the axis of the pipe a continuous and quietly burning column of light of very powerful actinic effect. In this long-drawn-out arc the passing air is partially transformed into nitric oxide. This is quickly chilled by contact with the wall of the pipe, which is- extemall}' exposed to the atmosphere, and so prevented from redecomposition. The gaseous product is half again, if not twice, as rich in nitric oxide (yield about 2%) as by the Birke- land-Etde process. The NO must be looked upon as the primary product, which subsequently unites with oxygen to form NOj, the latter being carried to water-absorption towers much like the Gat-Lussac towers in sulphuric acid plants. The NO2 cannot be the primary product, for it dissociates at about 600° into XO and Oj. With the water the NO, forms nitric acid and nitrous acid : ■'2 N2O4 + HoO = HNO3 + HNO2. The latter yields NO2 and NO, however, when the liquid becomes more concentrated: 2HNO2 = H2O 4- NO2 + NO. I 127.1 NITRIC ACID. 193 NO is once more converted into NO2 and the NO2 again gives nitric acid; eventually all is converted into that acid. Instead of marketing sodium nitrate, to duplicate the Chili saltpetre, calcium nitrate is produced hy saturating the nitric acid with lime and the resulting calcium nitrate is used for fertilizing and other purposes. Nitrites are also manufactured directly by leading N2O4 into caustic : N2O4 + 2K0H - KNO3 + KNO2 + H2O. The nitrate and nitrite are separable by fractional crystallization. Physical Properties. — Absolute nitric acid, i.e. the compound HXi^a in the pure state, is prepared by distilling the nearly pure acid of commerce (sp. g. 1.5) with concentrated sulphuric acid in vacuo. The liquid distillate has a specific g^a^•ity of 1.559 at 0° and becomes solid at —40°; it boils under ordinary pressure at S6°, but with partial decomposition. Chemical Properties. — Xitric acid, especially when pure, is a rather unstable compound; at ordinary temperatures it is decom- posed by sunlight to a slight extent, turning yellow on account of the small amount of nitrogen dioxide formed. At an elevated temperature the acid also breaks up, decomposition into nitrogen dioxide, water, and oxygen being complete at 260° When strong nitric acid is subjected to repeated distillation under atmospheric pressure, its boiling-point gradually rises, while the acid becomes proportionately weaker, until finally a GSVt acid is obtained, which boils constant at 120.5°. The same mixture is obtained when one starts with dilute acid and distils it. In both cases the boiling- point of the original liquid is lower than that of the product; it rises during the boiling to a maximum at 120.5°- AVe have here, therefore the case of a liquid mixtui-e witli a maximum boiling-point, which is discussed in Org. Chem., §22. The mixture of hydrogen chloride and water also has a maximum boiling-point (110°). Nitric acid is xery extensi\'ely ionized in aqueous solution; it is one of the strongest acids known. When it comes in contact with metals, the salts of nitric acid (nitrates) are formed, but without any evolution of hydrogen, since part of the acid present is reduced by the nascent hydrogen. The nitrates are all easily soluble in water. The action of nitric 194 INORGANIC CHEMISTRY. [§§ 127- acid on the metals is not the same in all cases. It does not attack gold or platinum. Silver, mercury, and copper are only imper- ceptibly dissolved at ordinary temperatures, but on warming they dissolve with the evolution of nitric oxide. This and the other NO-compounds are powerful catalyzers in the dissolving of the above-named metals, for nitric acid which is perfecty free from them does not dissolve these metals, while the reaction immediately begins as soon as a little of these substances is added. It may be supposed that on warming nitric acid traces of NO-compounds are formed, which together with the elevation of the temperature accelerate the reaction. Iron, zinc, and magnesiiun reduce nitric acid to nitrous oxide and even to ammonia. Under the action of iron filings and dilute sulphuric acid the reduction of nitric acid to ammonia in dilute solution is quantitati\'e. There are also ^'arious denitrifying bacteria known, Bacillus pyocyaneus being the best studied of them. Nitric acid frecjuently acts as a powerful oxidizing agc;':t, especially at an ele\'ated temperature If sulphur Ls boiled with it, the sulphur is converted to sulphuric acid, similarly phosphorus to phosphoric acid. A glowing piece of charcoal dropped upon the concentrated acid continues to burn with a bright glow. In all these cases the highest oxidation stages are formed. Nitric acid is used particularly in the organic branches of chemical industry. The corn-position of nitric acid can be deduced from that of its anhydride. A weighed amount of the latter is introduced into water; nitric acid is formed, which is neutralized with baryta water. By evaporation it is possible to determine how many parts Ijy weight of barium oxide, BaO, combine with the anhydride. It is found that 1.53.37 parts (=lBaO) combine with l(jS.(J2 parts (= IXX)-,) of the anhydride; the formula of barium nitrate thus becomes Ba(X03)2, hence that of nitric acid itself must be HNO,. Pemitric Acid, HNO,. Pernitric acid is formed in very dilute aqueous solution by the oxidation of nitrous acid with hydrogen peroxide: 2H,0, +HNO2 = HNO, + 2H2O. Nitric acid does not yield it when treated in the same way; on the contraiy, the pernitric acid breaks up e^•en in a cold dilute aqueous 128.] DEinVATniiS OF TIIK NlTltdGKX ACIDS. VX) solution inside of alumt an Iiour roniplclely into nitric acid and liydrogen peroxido : HN(J,+I1,()-II.\C)3+H,02. Pernitric acid has the \'cry cliaractcristic propertj' of liberating bromine from potassium bromide solutions, something that neither hydrogen peroxide nor nitrous acid nor nitric acid does. Derivatives of the Nitrogen Acids. 128. In discussing the manufacture of sulphuric acid (§ 86) we alread\' referred to the chamber crystals, HSO5N. They are formed in the lead chambers in case not enough steam is supplied. The following eciuation expresses the action that takes place: iSOa + Xo* )4 + + H2(.) =2S( )^^'H.. The ordinary method of preparing this substance is by conduct- ing carefully dried sulphurous oxide into cooled fuming nitric acid: S()2 + HN().3=.S(J5NH. The crystalline mass obtained is spread out on porous earthenware to allow the adhering liquid to be alisorbed. The chamber crystals have the appearance of a coarse crystal- line, colorless mass; they melt at 73°. They are at once decom- posed by water into sulphuric and nitrous acids: SO5XH + H2O = HoSOi -h HXO2. For this reason the compound is considered as the mixed anhy- dride of sulphuric and nitrous acids. According to § 90 the struc- ture Sr)o- heating phosphorus with lead in a sealed tube; the phosphorus dissolves in molten lead and crystallizes out in the metallic form on cooling. The specific gra^'ity of the red phosphorus varies bet«'een 2.1 and 2. 28, according to the duration and temperature of the heating. It is probable that it represents a solid solution of white phosphorus in metallic phosphorus. The ratio ill which these two modifications occur in red phosphorus changes with the temperature; an equilibrium establishes itself between them for CA'ery temperature. At about .500° the proportion of white phosphorus reaches a minimum. .Accordingly only the white and the metallic phosphorus are to be regarded as well characterized allotropic modifications. 134. Chemical Properties. — Phosphorus has a great affinitv' for manjr elements : it combines directly with all elements except nitrogen and carbon, the combination occurring \^-ith great vigor in many cases, e.g. when phosphorus is brought in contact with sulphur or bromine. Certain compounds of the metals (phos- phides) are known, which are called phosphor bronzes (§ 199). Especialljr characteristic of phosphorus is its very strong affinitj^ for oxygen; yellow phosphorus takes fire in the air at 40°, so that contact with a hot glass rod Ls sufficient to ignite it. The burning is accompanied by a vigorous evolution of hght and heat, phosphorus pentoxide, P2O5, being formed. On account of this strong affinity for oxygen phosphorus is a powerful reducing-agent. Sulphuric acid, when warmed with it, is reduced to sulphur dioxide; concentrated nitric acid oxidizes it with explosive violence; dilute acid evolves nitrous fumes, oxidizing the phosphorus to phos- phoric acid. ^lany metals are precipitated by phosphorus from their salts, phosphides being formed to some extent. Silver nitrate, for instance, gives silver and sUver phosphide, AgsP, with phosphorus; on warming phosphorus with a solution of copper stilphate copper phosphide, CU3P2, is deposited. 135. The slow oxidation of phosphorus by oxygen at ordinary 135.] PHOSPHORUS. 203 temperatures is accompaiiiccl by the emission of a bluish light. This luminosity of phosphorus is Acr}' plain in the dark. This phenomenon is due to various circumstances, some of which are very mysterious. The oxidation, and hence the luminosity, is prevented by the presence of traces of certain substances, such as hydro- cai'bons, ammonia, etc. Further, the luminosity depends on the tem- perature; below 10° it is extremely weak. The gas pressure has a peculiar influence; at ordinary temperatures phosphorus does not emit light in pure oxygen of atmospheric pressure, but if the pressure is reduced, a point is reached at which luminosity commences; this is at 666 mm. for 1.5°, and at 760 mm. for 19.2°- The oxidation is there- fore more vigorous in dilute oxygen (i.e. oxygen mixed with another gas, sucli as nitrogen) than in concentrated. Vx'n ]\[\rum observed as early as 170S that a piece of phosphorus laid on wadding (which serves as a poor conductor of heat) in a closed vessel shines the more brightly as the oxygen is pumped out, and may even take fire in very dilute gas. The fact that oxidations are more energetic under reduced oxyi^en- pressure has been observed in many other cases. Sec § 137. Detection of Phosphorus. — Poisonings liy j-ellow phospliorus - occur now and then. In order to detect it in such cases, u-e is made of its luminosity. For this purpose the contents of the stomach, which are to be tested for phosphorus, are diluted ■\\ith •water in a distilling-flask, connected with a condenser by a tube doubly bent at right angles. On heating the flask w-ater (Mstils o-^-er with a little phospliorus vapor; if the whole apparatus is placed in a dark room, a luminous ring is noticed during this distillation at the place where tlic steam is condensed, i.e. where the phosphorus vapor einncs in contact with air in the condenser. The distillate contains phosphoric aci 1 (Mitscherlich's test). Use. — Phcjsphorus is used cliieHy for the manufacture of matches. The matches in use to-day may be'classed as nafety matches iind " strike rinijirhcre" matches, which latter may l>e of the (a) parlor single-dip tyjje, or (b) double-dip with combustible bulb or (c\ double- dip with safety bulb. In the SwecHsh s:\fety matches the' hciul consists chiefly of a mi.xture of potassium chlorate and antimony sulphide. They are lighted by striking them on a surface coated with rci\ phos- •Mi INORGANIC CHEMISTRY. [§§135- phori;.s. On account of their requiring a special ignition surface and their f[uick burning and dropping they are not as popular in America as the strike-anywhere matches. Of the latter the double-dip safety bulb matches are best because they can be ignited by friction of the tip only, not by side friction. Phosphorus, in some form, is used in the production of aU matches. Yellow phosphorus is generally utilized in the production of American matches, but since it is very injurious to the health of the workmen, most European countries forbid its use and employ non-poisonous substitutes. The red phosphorus and phosphorus " sesqui-sulphide, " PiSs, are the chief non-poisonous substitutes for yellow phosphorus and are to be utilized in American strike-anywhere matches as fast as practicable. Compare § 133. Compounds of Phosphorus and Hydrogen. There are three compounds of phosphorus and hydrogen known : (1) gaseous hydrogen phosphide, PHg (also called phosphine); (2) liquid hydrogen phosphide, P2H4; and (3) solid hydrogen phosphide, (P2H)6. HYDROGEN PHOSPHIDE. PHOSPHINE, PH3. 136. This compound can be prepared from the elements by bringing phosphorus together with zinc and dilute sulphuric acid, i.e., with nascent hydrogen; when thus prepared, it is mixed with a large quantity of hydrogen. The generation of hydrogen phosphide by heating phosphorous and hypophosphorous acids is another example of its fonriation by the direct decomposition of phosphorus compounds: 4H3P03=PH3 + 3H3P04. Phosphorous Phosphoric acid. acid. The ordinary method of preparation is by the action of phos- phorus on caustic potash: P4+3KOH+3H2O =PH3+3H2KP02. Pot. hypophos- phite. The reaction is really more complicated than this equation indicates, for in addition hydrogen, P2H4 and other substances are formed. (See also § 144.) i:56.] if)'i)i:ni;i:\ I'liosrii iDh:. 20." By reason of the presence of irusi'ous TJI,, wliich is sclf-inllaminablc, each bubble of gas igr^ites as it brcal^cs into tlie air, forriiing usually a smoky riug of jihospliorus penloxide (Fig. 37). On account uf thii innammability the vessel in which the gas is generated from phosphoruF and caustic potash must be as full df lii|ui(l as jiossibli'. Moreover, tlit delivery-tube (preferably with a wide niouth) must o]irn in (/v//;;t water, in onler that it may not become cloggcil with particles of phfisphorus carried over. Ry jiassing the gas through hydrochloric acid or alcohol,' the hydrogen iihosjihitle is freed from V-^^^i and is then no longer spon- taneously combustible. Fig. :-!7. — Piif.caratiov of IIyduoiif.x Phosphide. According to Senderens phosphiuc lan be \er)- advantageciusly ]jrepared by heating red phosi)horus at 240°-250° in a current of steam. Perhaps phosphorus or inctaphosphorous acid is loinieil prima rib'. No method of producing hydrogen phosphide by the action of hydrogen on phosphorus compounds is known; however, we have one by the interaction of h}'drogen compoimds and phos- phorus compounds. Calcium phosphide, when decomposed by water or dilute hydrocliloric acid, forms hydrogen phosphide; Ga3P2 + 6HC1 = 3CaCl„, + 2PH3. The phosphides of zinc, iron, tin, and magnesium arc decom- 206 INORGANIC CHEMISTRY. [§§136- posed by dilute acids with the formation of hydrogen plios- pliide. The perfectly pure gas is now best obtained by condensa- tion and subsequent fractionation at a very low temperature, as was described in § 29. Physical Properties. — Hydrogen phosphide, or phosphuretted hydrogen, PH3, is a gas at ordinary temperatures; it becomes liquid at — S3° and solid at —133.5°. It has a peculiar disagree- able odor, that reminds one of spoUed fish. It is slightly soluble in water, more so in alcohol. Sp. g. = 17 (0= 16). 1 liter weighs 1.5293 g. at 0° and 760 mm. pressure. 137. Chemical Properties. — Hydrogen phosphide is very poison- ous; it burns very easily, yielding phosphoric acid. In the presence of oxygen of ordinary pressure it remains unchanged; if, however, the pressiue is diminished, an explosion results. This conduct reminds one of phosphorus, which is luminous (because of oxidation) only below a certain limit of pressure (§ 135). The combustion of hydrogen phosphide may be expressed by the equation: 2PH3 + 4O2 = P2O5 + 3H2O. Accordingly the reaction would be hexamolecular (§ 50). Van DER Stadt demonstrated by a method, similar to that referred to in § 51, that the first stage of the reaction is bimolecular and corresponds ^'ery closely to the following equation: PH3 + 02=H2 + P02li, Met a phos- phorous acid. if the gases slowly diffuse into each other in a dUuted condition. In general, experience has taught that the mechanism of a reaction is decidedly simple and that chemical processes are almost always mono- or bimolecular. Accordingly, when the quantita- ti\'e course of a reaction is represented by an equation indicating the participation of several molecules, it is probable that several intermediate reactions are involved. Hydrogen phosphide can unite with halogen-hydrogen acids directly to form compounds of the type PH4X (X=haolgen), in analogy with ammonia. The best known of these compounds 137.] HYDROGEX PHOSPHIDE. 207 is PH4I, phosphonium iodide, a colorless, well-ciyst:illizcil com- pound, which is formed when diy hydrogen phosphide and hj-ilro- gen iodide are mixed. In contact with water it l)rcaks up into PH3 and HI; the former escapes as a gas, while the latter remains dissolved in water. Phosphonium iotlide is ver}- unstal)lc. This is even more the case with phosphonium bromide, which is also a solid, but is completely dissociated into the two hydrogen compounds, PH3 and HBr, as low as 30°. Phosphonium chloride is dissociated even at ordinary temperatures and pressures and can only exist below 14° or under more than 20 atni. pressure. Considering these properties, it is not surprising that phosphonium, PH4, — like ammonium, — should be impossible to isolate. Xo other acids except those mentioned unite \\'\\\\ hydrogen phosphide. The general behavior of the latter thus shows that it is very much less basic than ammonia. It is for this reason that PH,I is decomposed by water. The weak basic character of phosphonium hydroxide allows the iodide to be hydrolyzed into PHjOH and HI, whereupon the PHe ascertained in the following way: on being decomposed by water they }-ield phos- phoric or phosphorous acid and a halogen acid, so that the cjuan- tities of phosphorus and halogen present can be found Ijy deter- mining the amounts of these acids. ;\Ioreo^'cr, the molecular weight can be obtained 1)}- measuring tlie vapor donsil}-, though it must be borne in mind, howe-ver, that comj^ounds of tlie type PX5 are usually dissociated in the gaseous state. Oxygen Compounds of Phosphorus. 144. Three compounds of this chiss arc known: phosphorus trioxirh, V^^h', phosphorus tvirnxidc, Pi.< >_, : and phosphorus pent- ox Id c, or phosphoric anhydride, V-^' )^. Only the last is of an)- great importance. Phosphorus Trioxide, P2O3. This compound is produced when phosphorus burns in a slow cur- rent of dry air in a tube. The princijial product is phosphorus pent- oxide, which can be collected by a wad of glass fibers. The phosphorus trioxide passes through as a vapor and is condensed in a well-cooled 212 INORGANIC CHEMISTRY. [§§ 144- tube. It is a white waxy substance when thus formed, but it can also be obtained in crystals; the latter melt at 22.5° and boil at 173.1° (in a nitrogen atmosphere). The vapor density has been found to be 109.7, while that calculated for P^Og is 110. On being heated to 440° it is decomposed into red phosphorus and phosphorus tetroxide. It turns yellow in the light, which explains the fact that phosphorus pentoxide sometimes takes on a yellow color. It dissolves sloAvly in cold water forming phosphorous acid; with hot water it produces red phosphorus, self-inflammable hydrogen phosphide and phosphoric acid in a vigor- ous reaction. AYhen heated to 50°-60° in the air it takes fire and burns to the pentoxide. Phosphorus Tetroxide, P2O4, is obtained from the P2O3 compound, as was stated above. It forms colorless glistening crystals, that break up in water into phosphorous and phosphoric acids. In this respect its conduct is analogous to that of nitrogen tetroxide, wliich yields nitrous and nitric acids with water. PHOSPHORUS PENTOXIDE, P2O5. This compound is the product of the combustion of phosphorus in oxygen or an excess of dry air. It forms a white, voluminous, snow-lilve mass, that takes up water rapidly to produce phosphoric acid. It is the most powerful desiccating-agent known. Morley ascertained that it dries the air down to 1 mg. water vapor in 40,000 1. air. It exists in two modifications, both of which are formed simultaneously in the above process. The one is crystal- line, subliming at 250°; the other amorphous and not volatile below red heat; the vapor condenses crystalline. When heated above 250° the crystalhne modification passes over into the amorphous form. Heating with charcoal reduces it to phosphorus. The vapor density of phosphoric anhydride at bright redness was found to correspond to the formula (P205)2. Acids of Phosphorus. 145. Only two of the above described oxides of phosphorus, viz. P2O3 and P2O5, form corresponding acids; these oxides ca.n unite with different amounts of water to form acids. From P2O5 we have: 145.] ACIDS OF PHOSPHOROUS. 213 P2O5 + H2O = 2HPO3, mctaphosphoric acid, P205 + 2H20 = H4P:.C)7, pyrop)iosphoric acid, and P2*^5 + 3H20 = 2H3P04, orthophosphoric acid. From the other oxide two acids can be derived: metaphos- phorous acid, HPO2, and phosphorous acid, H3PO3. Besides these there are two acids of phosphorus, whose anhydrides are unknown, viz. hypophosphorous acid, H3PO2, and hypophosphoric acid, H4P206. The relation between ortho-, meta-, and pyrophosphoric acids can be shown in another way, which leads us to make some general obser- vations. It was remarked in § 141 that phosphorus pentachloride is transformed by water into phosphoric and hydrochloric acids. The action of water on the pentachloride may be regarded as consisting first of a substitution of all five chlorine atoms by hydroxyl: P| Cl5+5H| OH =5HC1 + P(0H)5. This compound, which would strictly be regarded as orthophos- phoric acid, is unknown; a molecule of water is at once split off, form- ing the ordinary phosphoric acid, H3PO4, which we are accustomed to call orthophosphoric acid. In a similar way the metaphosphoric acid can be derived from the acid P(0H)5 by the splitting off of two molecules of water: OjH OjH |0 H /OH |0H POH-»OP^OH; P0|H->02P-0H; OH OH |<)H Metaph9sphorio H Orthophos- OH phoric acid. while the pyrophosphoric acid can be regarded as 2P(OH)5-3H20: OlH HJO |0H H_0| /OH /OH POH HOP-^OP^-OH PP— OH OH HO \U ti tl\ U Pyrophosphoric acid. Orthophosphoric acid can also be deri\-cd from phosphorus oxy- chloride : OP|Cl3 + 3H|OH^ 0P(0H)3. 214 INORGANIC CHEMISTRY. [§§ 145- This way of looking at them makes plain not only the connection between the different acids, but also their structural formula. The same method can be applied to many other cases. As an example we may select the pcr-iodic acids. In § 62 onh- one was mentioned. There are salts, however, of various per-iodic acids, e.g. MIO4, M3IO5, WjIOu, etc. These can be derived from a hypothetical acid ICOH), in which iodine is joined to as many hydroxyls as correspond to its maximum valence. MjTOe would come from I(OH)j — IHjO; M3IO5 from I(OH)j-2H20; and MIO, from I(OH),-3H20. ORTHOPHOSPHORIC ACID, H3PO4. 146. This acid can be obtained by direct synthesis from its elements; phosphorus burns to the pentoxide and the latter yields the acid on dissoh'ing in water. Its formation b\' the action of nitric acid on phosphorus was mentioned in § 134. It can also be obtained by the oxidation of compounds containing phosphorus and hydrogen; phosphine and the lower acids of phosphorus are oxidized to phosphoric acid. Ordinarily this acid is prepared by the oxidation of phosphorus with nitric acid or by liberating it from its salts, particularly the calcium salt, Ca3(P04)2. The latter is stirred into the theoretical amount of dilute sulphuric acid, forming calcium sulphate, which is onljr slightly soluble in water, and phosphoric acid, which goes into solution. On evaporating this solution the acid remains. At ordinary temperatures orthophosphoric acid is a crystalline solid. It melts at 38.6°, is odorless and extremely soluble in water, forming a strongly acid solution. It has the character of a strong acid; however, it is consider- ably less ionized than hydrochloric acid; a solution of 1 mole phosphoric acid in 10 1. water contains about one-fourth as many hydrogen ions as hydrochloric acid of the same molecular con- centration. It is ionized chiefly into H" and H2PO4'. It generates hydrogen with metals, all three hydrogen atoms being replaceable by metallic atoms; it is therefore tribasic. Three classes of salts are possible and known to exist; these are the primary, secondary and tertiary salts. Of the alkali salts all three kinds are soluble; of the alkaline earth salts only the primary, the tertiary and secondary being insoluble. The other phosphates are insoluble in water but are dissolved by mineral acids. 146.] ORTHOPIIOSPIIORIC ACID. 215 This latter property is due to the fact that phosphoric acid is a weaker acid than the strong mineral acids, hydrochloric, nitric and sulphuric. On treating an insoluble phosphate with one of these acids, e.g. hydrochloric, undissociated mole(mles of phos- phoric acid are formed in the liquid; the more hydrochloric acid, the more the association, since the hydrochloric acid reduces the ionization of phosphoric acid. H2PO4' and H' ions thus disappear and, in case enough hydrochloric acid is aiklcd, the concentration of the H2PO4' ions remaining will not be great enough together with that of the metal ions present to reach the value of the solu- bihty product; hence all the phosphate must dissolve (§ 73). For the same reason, as a general rule, salts that are insoluble •in water will only dissolve in acids that are stronger than the acid of the scdt. The only exception to this is the case when the value of the solubihty product of the insoluble salt is ^'ery small, examples of which we have seen in certain sulphides (§ 73). "^ATien heated to 213° orthophosphoric acid gives off water, forming mainly the pyro-acid but als(3 a littl(> meta-acid through- out the reaction. The pyro-acid on the other hand is converted by further heating into the meta-acid. With silver nitrate orthophosphates give a yellow precipi- tate of sUver phosphate, Ag3P04, soluble in nitric acid and ammo- nia. In the case of a primary or secondary phosphate, the pre- cipitation is not complete, since nitric acid is liberated in the reaction : Na2HP04 -F3AgN03= Ag3P04 +2NaN03 +HNO3, or, expressed in ions : HPO4" -f 3Ag- ^ Ag3P04-h H-. If, however, an excess of sodium acetate is added, the precipi- tation is practically complete. The reason for this is obvious. By the addition of acetate the acetic anions C2H3O2' are forced to combine with the H' ions, for acetic acid is only very slightly ionized and its ionization is, more- over, considerably lessened by the excess of sodium acetate. The result is that in the equilibrium HP04"-|-3Ag f^ Ap3P04-|-H- the H' ions are removed. The inverse reaction ♦- is then no longer 216 INORGANIC CHEMISTRY. [§§ 146- possible, and the direct reaction -^ must therefore become com- plete, or in other words, all the phosphoric acid is precipitated as silver phosphate. It was stated above that the alkali salts of phosphoric acid are soluble in water. These aqueous solutions differ markedly in reaction. The solution of a primary salt, KH2PO4, is acid, that of a secondary salt feebly alkaline, and that of a tertiary salt strongly alkaline. The cause of this variation must be more fully explained. The acid reaction of a salt such as KH2PO4 must be attributed to the fact that its anion, H2PO4' (analogous to the anion HSO4'), is capable of splitting up into the ions H' and HPO4", the former producing the acid reaction. The feebly alkaline reaction of a salt like K2HPO4 is accounted for by hydrolysis (§ 66). Such a salt is extensively ionized in dilute solution into 2K' and HPO4". However, while H3PO4 is rather highly dissociated (into H' and H2PO4'), H2PO4' is but slightly ionized into H' and HPO4". In this case H2PO4' behaves as a weak acid. Hence, if there is a large proportion of HPO4" ions in a solution, they will tend to unite with H' ions, because the system H' + HP04"<=>H2P04' is only in equilibrium when the right-hand side preponderates. The necessary H' ions are supplied by the water, which is split up to a ver}' slight extent into H' and OH'. But when the H" ions unite with HPO4" ions we have a surplus of OH' ions in the solution and the latter takes on an alkaline reaction. Entirely analogous is the explanation of the strongly alkaline reaction of the tertiary phosphates, such as K3PO4. Their aqueous solutions contain the ions PO4'", which have a still stronger tendency to unite with H' ions than the HPO4" ions. The PO4'" ion, there- fore, causes the presence of an even larger proportion of OH' ions, not compensated by H' ions, so that the result is a strongly alkaline reaction. Phosphoric acid is precipitated from an ammoniacal solution by a magnesium salt as white crystaUine ammonium magnesium phosphate, NH4MgP04+6H20. Another very characteristic test for phosphoric acid is that in nitric acid solution a finely crystal- line, yellow precipitate is produced by ammonium molybdate, especially on warming. This precipitate has approximately the composition 14Mo03-|-(NH4)3P04+4H20, i.e. it is an ammo- US] PYROPHOSPIiORIC ACID. 217 niiim phospho-molybdate. Precipitation in acid solution is of great advantage here, since most of tlic phosphates are soluble only in acids. 147. One method of producing this acid was given in the pre- ceding paragraph. In preparing it, it is more practicable, howe\cr, to heat the secondary sodium phosphate (the ordinary sodium phosphate of commerce), because in this case only one molecule of water can be driven off from two molecules of the salt: L>Na2HP()4==H2(.) + Na4r2t>7- After being heated the sodium pyrophosphate is dissolved in water and lead acetate is added to precipitate lead p^Tophosphate, which is then decomposed with hydrogen sulphide. P^-rophosphoric acid can be obtained from its solution as a colorless vitreous mass by evaj)oration in a vacuum at a low tem- perature. When dissolvetl in water of ordinary temperature, the acid remains unchanged for cjuite awhile; on warming tliis solu- tion, especially after the addition of a little mineral acid, it is converted in a few hours into orthn-acid (§ 145). All four hydrogen atoms are replaceable l)y metals; we should therefore expect to find four classes of salts. In reality only two are known, M4P2O7 and ;\l2ll2l'2<'7- The neutral, as well as the acid, salts of the alkalies are soluble in water; the neutral salts of other bases are insoluble, the acid salts chieflj' soluble. Pyrophosphoric acid is distinguished from the ortho-acid by the fact that solutions of its .salts give a wliilc precipitate, AgjPo* )t, with silver nitrate, and from the meta-aeid liy not coagulating albmnen and giving no precipitate with barium chloride. METAPHOSPHORIC ACID, HPO3. 148. This acid is obtained by heating the ortho- or the pyro- acid till no more water passes off, or l)y heating ammonium jjho.s- phate (NH4)2HPr)4. .Moreover, on dissolving phosphorus pent- oxide in cold water, the product is at first chiefly meta-acid. At ordinary temperatm-es metaphosphoric acid is a ^■itre(lus solid (hence the name glacial phos])hayic acid), which can be melted and easily drawn out into threads. On being heated strongly 21S INORGANIC CHEMISTRY. [§§ 148- it volatilizes without breaking up into water and pentoxide. When boiled in aqueous solution it goes over into orthophosphoric acid. It is very deliquescent; use is made of this property occasionally. iMetaphosphoric acid is monobasic, corresponding to the formula HPO3. Its alkali salts only are soluble in «'ater. In solution the meta-aeid can be distinguished from the ortho- and the pyro- acids by its ability to coagulate albumen and give white precipi- tates with chlorides of barium or calcium. The vapor of this substance at bright-red heat consists chiefly of H2P2O6 molecules (di-metaphosphoric acid) , ■\\hich are apparently liable to undergo partial dissociation and even to lose a small quantity of water. There are salts of various acids known, which must be regarded as pol^TTiers of metaphosphoric acid, e.jT. KjPjOu, potassium di-metaphos- phate; there exist also tri-, tetra-, and hexa-metaphosphates, i.e. salts of the acids H3P30a, 'iij?f0,2, and HgPgOjg. Hypophosphoric Acid, H4P2O1,. I4Q- When sticks of phosphorus are suspended in a solution of sodium acetate in such a way that only 0.5 cm. is exposed above the level of the liquid and the temperature is kept between 6° and 8°, the phosphorus oxidizes slowly and the difficultly soluble acid .sodium salt of hypophosphoric acid, NajH^PzOe +6H2») soon begins to crystallize out. It can be purified by crystallization from a dilute solution of acetic acid. If this salt is dissolved in water and barium chloride added, a precipitate of barium hypophos- phate is formed, from which an aqueous solution of the free acid can be obtained by means of dilute sulphuric acid. This can be evaporated at 30° to a sirupy consistency without decomposition and, when left in a vacuum, yields crystals of the acid. At an elevated temperature and in the presence of a mineral acid phosphorous and phosphoric acids are formed. This behavior justifies the consideration of h3rpophos- phoric acid as a mixed anhydride of the two last-named acids: OH HO OH OH O POH HO P -* OPOH POH 10H_J0 \^/ Phosphoric Phosphorous Hypophosphoric acid. acid. acid. 151.] PHOSPHOROUS ACID. 'JM) However, it has not yet been ]>ossiblo to prepare the hypopho?phoric acid by melting together the other two acids. I'roni the determination of the molceular weight of the methyl ester it seem.s probable that the formula of the acid i.s HjPO, and not H^PjOj Metaphosphorous Acid, HPOo. 150. This compound was -stals of phos- phorous acid and are converted into the latter by the action of water vapor. 151. In § 149 it was mentioned that this acid is formed by the slow oxidation of phosphorus in moist air. It is more easily pre- pared by decomposing phosphorus trichloride with water; PCI3 + 3H2O = H3PO3 + 3HC1. The hydrochloric acid can be expelled by evaporating at 180° and the phosphorous acid crystallizes out on cooling. The melting-point of phosphorous acid is 70.1°. It is a very hygroscopic sub.stance Heating decomposes it into phosphoric acid and phosphine. It has a strong reducing action, being itself oxidized to phosphoric acid. The oxygen of the air acts on it very slowl}'. It precipitates the metals from solutions of gold chloride, mercuric chloride, silver nitrate, etc. A characteristic reaction is the reduction of sulphur dioxide to sulphur, which takes place at ordinary temperatures, -when solutions of the two substances are mixed. In spite of its thrcv hydronon atoms, phosphorous acid act.s as a dibasic acid. As we have already observed, the itmization of polybasic acids sometimes affects only one W ion at first, llie others being split off with increasing difficulty. According to ()stwald it may be supposed that ionization beyond 211' and HPO;/' is in this case so difficult that the acid seems to be only dibasic. The phosphites are not oxidized b}- the air, but they yield to the action of 220 IXOTiGAXIC CHEMISTRY. [§§152- oxidizing-agents; e.g. they liberate the precious metals from their salts, as does also the acid itself. Heating breaks them up into hydrogen, pyrophosphates and phosphide. The double phosphites give precipitates with baryta- or lime-water. Hypophosphorous Acid, HjPOj. 152. Salts of this acid are produced by heating phosphorus with caustic soda, lime-water or baryta-water (§ 1.3(i): 3Ba(OH)2-l-8P-l-6H20 = 3Ba(HjP02)2-}-2PH3. It can be set free from these salts by sulphuric acid; the aqueous solution is concentrated at 80°-90° and then cooled strongly, whereupon the acid crystallizes out. Melting-point, 26..5° On being heated at 130°-140° the acid splits up into phosphorous acid and phosphine; at a somewhat higher temperature the latter acid yields phosphine and phosphoric acid. The equations are: 3H3P02 = 2H3P03-HPH3; 3H3P03 = 2H3PO,-|-PH3. Hj-pophosphorous acid is a very strong reducing-agent. Gold, silver and mercury are precipitated from solutions of their salts by the free acid as well as its salts. Sulphur dioxide is reduced to sulphur at ordinary tem- peratures. In these reactions the acid itself is converted into phosphoric acid. It is distinguished from phosphorous acid by its behavior towards copper sulphate solution; when it is warmed with the latter, a red precipi- tate of copper hydride, CuoH^, is formed. Hypophosphorous acid is mono- basic. Compounds of Phosphorus and Sulphur. 153. Various compounds of this sort are known; all of them are obtained by warming the two elements together. As the reaction is very vigorous with yellow phosphorus, the red form is usually employed. The compound P2S5. which is of service in organic chemistry, is a yellow crystalline substance, melting at 274°-276° and boiling at 518°. On being warmed with water it yields phosphoric acid and sulphuretted hydrogen. P2S5 unites with 3 molecules of K2S to form a sulphopfiosphate, KsPS^, i.e. a phosphate whose oxygen is replaced by sulphur. Several compounds containing a halogen in addition to phcsphorus and sulphur are known, e.g. PSClj. This phosphorus snlphochloriile can be pre- pared by treating phosphorus pentachloride with hydrogen sulphide, a method 155.] COMPOUNDS OF PHOSPHORUS AXD KITROOES. 221 analogous to that of forming the oxychloride finiri the peutachloride and water. A more convenient method is by the action of the pentachloride on the pentasulphide, which carries out the analogy, to oxy-compounds still farther (§ 142) : 3PCl5+P,S,=5PSC]3. It is a colorless liquid, boiling at 12.')°. A\'atcr decomposes it into phosphoric acid, hydrochloric acid and hydrogen sulphide. Compounds containing Phosphorus and Nitrogen. 154. The compounds of this class are also numerous. Among them are amidophosphoric acid, OP^tji , and diamiduphosphoric acid, Ol-T OP j^Tjj X . As their names indicate, these compounds behave like acids. If dry ammonia is conducted over phosphorus pentachloride, a white mass is obtained which consists supposedl\' of ammonium chloride, XHjCl, and a compound PCl3(NH2J3. ^^'ith water it forms phosphamide, PO(XH)(XH,), a white insoluble powder. On being boiled with water secondary ammonium phosphate is formed: PO(NH)(NH,) +3H,0 =0p5J^.jj^) The name phospham is given to a compound P3H3Xg, which is formed from the product of the action of ammonia on phosphorus pentachloride, when it is heated in the absence of air till no more ammonium chloride fumes appear. It is insoluble in water. A\'hen fused with potassium hydroxide, it breaks up as follows: P3H3N8 + 9K0H + 3H2O -SKjPOi + 6NH3. By the interaction of P2S5 and NH3 it is easy to obtain a com- pound PjNj, phosphorus nitride. ARSENIC. 155. Arsenic occurs in nature in the free state — native. More frequently it is found in combination with sulphur (reuhjar, Xs-iS-^, and orpiment, Ai^iS^) and with metals {arftcnopi/ritc, or inispickcl, FeAsS,and cobaltite, Co XnH) ; also with oxygen as A.S2O3 ((irscnolite). The extraction of the element from these minerals is simple 222 INORGANIC CHEMISTRY. [§ 155- Arsenopyrite yields arsenic on mere heating, the latter subliming. Arsenolite is reduced with carbon: 2As203 + 6C = As + 6CO. Physical Properties and Allotropic Conditions. — The condition in which arsenic usually occurs is the crystalline. It then has a steel-gray color and a specific gravity of 5.727 at 14° and is a good conductor of electricity. It sublimes under ordinary pressure without melting; under increased pressure, howe\-er, it melts at 500° By sublimation in a current of hydrogen a second crystal- lized form can be obtained together with a black modification, which according to Retgees is also crystallized. An amorphous modification results from the decomposition of hydrogen arsenide by heat, the arsenic appearing as a dark brown deposit on the sides of the glass. Finally there is a yellow modification which is fomied when arsenic vapor is condensed in a dark room by liquid air. This yellow arsenic is very sensitive to light; even at the tem- perature of liquid air (— 180°) — at which it is stable in the dark — it is converted into the black modification by the light of a Welsbach burner. It is a remarkable fact that a solution of the yellow modification in chlorine is much more stable toward light and heat than is the pure substance. Such solutions are obtained in concentrations up to 7%; when they are cooled yellow arsenic crystallizes out. The relation between yellow and black arsenic is very analogous to that between yellow and red phosphorus, except that in the case of arsenic the yellow form is much less stable. At an elevated temperature (360°) all the modifications pass over into the ordinary crystalline form. Vapor Density. — The lemon-yellow vapor of arsenic has a density of 10.2 (air=l) at about 860°, which makes the molecular weight 293.8. At 1600°-1700° the vapor density is less by half, being 5.40. Since the atomic weight of arsenic is 75, its molecule therefore contains four atoms at about 860° and two at 1600°-1700°. Chemical Properties. — Arsenic is not affected by dry air at ordinary temperatures; in moist air it becomes covered with a coating of oxide. At 180° it burns with a bluish flame to the oxide AS4O6, giving off a peculiar garlic-like odor. At an elevated 156.] HYDROGES ARSENIDE. -Jjlj temperature it combines witli many elements directly; it unites with chlorine without the aid of heat, producing scintiUations. HYDROGEN ARSENIDE. ARSINE, AsHg. 156. I)irect synthesis from the clcTnents is not possible with this compound. It is formed when almost any arsenic ccjnipound comes in contact with nascent h}-(lrogen (zinc + sulphuric acid). "When thus prepared it contains considerable hydrogen, however. Pure arsine is obtained l)y treating zinc arsenide or sodium arsenide with dilute sulphuric acid : As2Zn3+ 3H2S04= 2ASH3 + 3ZnS04. Physical I'rnpcrtics. — Hydrogen arsenide is a gas; it liquefies at —40°, but does not solidify as low as —110° Sp. g. = oS.9 (H=l). It must be handled with great care, as it is very poison- ous. Fortunately its presence can be easily detected Ijy its pecuhar, disagreeable odor. Citcmicnl Properties. — Arsine can be decomposed into its ele- ments by heat. If the gas is ])asscd through a hot glass tube, arsenic is deposited on the sides in the form of a metallic mirror. Induction sparks also decompose it. By the latter means it can be shown that the resulting volume of hydrogen is U times as large as that of the gas itself, in accord with the formula AsHs. It is an endothermic compound, As-h3H-AsH3= -36.7 Cal., and has been made to explode by fulminating mercury (§ 119). Hydrngen arsenide burns with a pale flame, yielding water and arsenious , from a ^■ery concentrated solution of silver nitrat IS e: AsH3 + CiAgX( );( =^.\s.\g:j ■ 3AgX03. This is decomposed by the addition of ^\aleI■ into arsenious acid, nitri^' acid and metallic silver, the latter being depositeil. 224 INORGANIC CHEMISTRY. [§§ 15&- This reaction is called Gutzeit's lest. It is usually carried out in the following way: A drop of 50% AgNOj solution is placed on a piece of filter paper and the moist spot is held over a test-tube containing some zinc, dilute sulphuric acid and the substance to be tested for arsenic. A plug of cotton is inserted near the top to protect the paper from being spattered by the effervescing solution. If arsenic is present, the spot becomes yellow, and turns black when moistened with water. Composition of Arsine. — If arsine is passed over hot copper oxide, water and copper arsenide are formed. The ratio of hydrogen to arsenic in arsine is determined from this reaction. For 1 part (by weight) of hydrogen 24.97 parts of arsenic are obtained. The molecular weight of the compound, as found from the specific gravity (see above), is 77.9; since the atomic weight of arsenic is 75, the formula of arsine must be AsHs. Detection of Arsenic. 157. The majority of arsenic compounds are very poisonous. Several of them are of practical use and hence are on the market, e.g. white arsenic, AsjOg (rat-poison); orpiment, AsjSj; Schweinfurt green, or copper arsenite. Poisonings with these substances happen occasion- alh'. Some arsenic compounds, because of their pretty green color, are still used, though much less than formerly, in d3'eing tapestries, portieres, and the like. Rooms in which these are hung usually con- tain particles of arsenical matter, which are injurious to the health. Further, a certain species of mould, penicillium brevicaule, which is sometimes found in such tapestries, has the power of generating volatile and very poisonous arsenic compounds. The chemist is therefore quite frequently called upon to analyze a given sample (of dyed materials or the like, or the contents of a stomach) for arsenic. For this pur- pose a method has been devised which enables him to detect wdth cer- tainty extremely small amounts of arsenic. It involves the following operations: The organic substance in question is at first disintegrated as well as possible, usually bj^ digestion with hj'drochloric acid on the water bath, a little potassium chlorate being added from time to time. Thus the arsenic compound is oxidized to arsenic acid. When the chlorine has been expelled by warming and the liquid has been filtered, hydrogen sulphide is passed in for some time at a temperature of about S0° to precipitate the arsenic as sulphide. The sulphide is then dis- solved in nitric acid (in rase the presence of antimony is suspected it must first be removed) ; this solution is evaporated to drT,'ness to get rid of the excess of acid, the dry residue is dissolved in water, and this 157.] DETECTION OF ARSENIC. liquid is then tested in the Marsh apparatus, a simple form of which is shown in Fig. 3S. This consists of a small flasl<, in which hydrogen is generated from zinc and sulphuric acid; the liquid to be investigated is poured down the thistle-tube; if arsenic is present, arsine is formed. The mixture of hydrogen and arsine is dried by calcium chloride in the wide tube and then enters a tube of hard glass, which is narrowed at several places and drawn to a point at the further end. As the gas leaves the ta- pering end, which is bent upward, it is lighted. Thereupon the tube is heated with a flame on the near side of a narrowed place. The arsine is broken up and arsenic is deposited as a bright metallic mirror in the narrowed part. From the extent and thickness of the deposit one can estimate the number of milligrams of arsenic present. If the hydro- gen arsenide is not heated, it passes on to the flame and is burned. A Fig. oS. — Marsh Apparattts cold porcelain dish held in the flame is soon coated with a deposit of -arsenic, which is readily soluble in sodium hypochlorite solution (sodium arsenate being formed). This solubility enables us to dist'nguish arsenic from antimony. Arsenic is very A\idely distributed, although in small amounts; hence we always have to reckon with the jiossiblity of traces of it being present in the reagents and glass utensils of the laboratory. In order to test this a "blank experiment" is performed, i.e. all the operations are carried out with duplicate amounts of the required chemicals but without the addition of the substance to be analyzed. Xot until the materials used are proved to be free from arsenic is it permissible to use them in an actual test. Whether or not textile fabrics and the like haw been dyed with Schweinfurt green (copper arsenitc) can be determined casilj^ Ijy the GuTZEiT test. Another method is to use the aliovc-mentioned pnii- cillium brevicaule. This is cultivated on bread which is soaked with 226 INORGAXIC CHEMISTRY. [§§ 157- the liquid to be tested for arsenic. The least trace of the latter reveals itself by a characteristic garlic-Uke odor, caused by the evolution of arsenical gases. Compounds of Arsenic with the Halogens. 158. Three arsenic-halogen compounds of the type AsXs arc known; viz., the pentachloride, AsCls, the penta-iodide, Aslj, and the pentafluoride, AsF.,. Aside from these only compounds of the type AsXij are known. Arsenic trichloride, AsL'ls, can be obtained by direct synthesis or by the action of hydrochloric acid on white arsenic. The latter way is analogous to the formation of metal chlorides from the oxide and hydrochloric acid. This compound is a colorless oily liquid having a specific gravity of 2.205 (d^). It freezes at —18° and boils at 130.2° It is extremely poisonous. "\Mien exposed to the air it throws off dense white fumes. With a little water it forms an oxychloride, As(0H)2Cl; with much water hydrochloric acid and arsenious oxide. In this latter system a rise of temperature results in partial re-formation of the trichloride, which is volatile with the water vapor. The following equilibrium seems to exist: AS2O3 + 6HCI rf 2ASCI3 + 3H2O. Oxygen Compounds of Arsenic. Two such compounds are known: AS2O3, arsenious oxide, and AS2O.5, arsenic oxide. ARSENIOUS OXIDE, AS2O3. 159. .\rsenious anhydride (commonly called "arsenic" or "white arsenic") is found in nature. It is formed by the com- bustion of arsenic in air or oxygen and bjr the oxidation of arsenic with dilute nitric acid. It is manufactured commercially by roasting arsenical ores; the oxide volatilizes and is condensed in brick-walled chambers, where it collects as a white powder ("arsenic meal ") -.It is refined by sublimation from iron c^'linders. Physical Properties. — Arsenious oxide is an odorless solid, that does not melt under ordinary pressure, Ijiit subhmes. Under higher pressure it is possible to melt it. At 800° its ^^apor density 159] AliSENIOUS OXIDE. 227 is 198 (0 = 16), which makes the molecular formula AS4O6. Abo\e this temperature dissociation begins and at 1800° the vapor density corresponds to the formula As2( )■.}. By the ebullioscopic method (ekn-ation of the boiling-point) the molecular formula has been found to be AsjOq at 205° also (in boiling nitrobenzene). ]'arious Modi}icutionf:. — Arscnious oxide is known in a ^■itreous form as \\ell as in crj'stals of the regular and monoclinic systems. The vitreous modification is produced when the compound is sub- limed or heated to the sublimation-point. Sp. g. =3.738. After stand- inj for some time at urdinary temperatures, this form becomes white hl-;e porcelain because of conversion into isometric cry.stals. The latter form is better obtained by dissolving the vitreous modification in ^\ater or hydrochloric acid and letting it cr^'stallize out. During the crystal- lization the strange phenomenon of bright luminescence is observed, which is caused by the breaking of the crystals. This phenomenon, which is also noticed in other crystallizations, is called tribolumines- cence. The transformation of the amorphous into the regular variety is accompanied by the evolution of heat (5.3.30 Cal.). The monoclinic form is obtained by conducting the crystallization above 200° instead of at ordinary temperatures. If the lower half of a sealed glass tube containing arsenious oxide be heated above 400°, it will be found after cooling that the lower heat"d part cont:Hiis vitreous, the middle mono- clinic, and the upper octahedral, arsenious oxide. Since the transformation of amorphous into crystallised arsenious oxide takes place even at ordinary temperatures (rapidly at 100°) and with the evolution of heat, the octahedral form is to be regarded as the stable one at ordinar}- temperatures; the glassy form is only able to exist at those temperatures, because the ^'clocitj- of transformation is then very small. According to the above, if octahedral arsenious oxide is gradually warmed, we have first a transformation into monoclinic and then another into amorphous arsenious oxide. The transition tem- peratures have not yet been determined. Chemical Properties. — Arsenious oxide is easily reduced to arsenic ; for example, by heating with charcoal or nascent hydrogen. It is also easily oxidized to arsenic oxide and is therefore useful as a reducing-agent. This oxidation can be brought about by chlorine, bromine (bromine- water) , iodine solution, potassium perman- ganate, strong nitric acid, etc. It is slightly soluble in water; the solution has a salty metallic taste and a weak acid reaction. In acids it dissolves much more easily, because it acts towai-ds them as a basic oxide. It was stated above (§ 158) that a solution of 22S INORGANIC CHEMISTRY. [§§ 159- the oxide in hydrochloric acid gives off arsenious chloride. White arsenic is a rank poison; freshly precipitated ferric hydroxide serves as an antidote. ARSENIC OXIDE, {f^^O^:^. i6o. This compound cannot be prepared hke the correspond- ing phosphorus compound by burning arsenic in the air, for the oxidation goes no farther than to arsenious oxide. The higher oxide can only be prepared by heating arsenic acid in the air: 2H3As()4-.3H20=As205. This arsenic anhydride is a white glassy substance, that dis- s(jlves in water slowly, going o^-er into arsenic acid. By heating with carbon it is easily reduced to arsenic. At an elevated tem- perature it lireaks up into (jxygen and arsenious oxide. Its molec- ular weight is not known; the formula AsoCJg is simply empirical. Oxyacids of Arsenic. Two of these are known: ursenious acid, H3A^(33 (only in aqueous solution and salts) and arsenic acid, H3ASO4. 161. This acid exists in the aqueous solution of the anhydride. It stUl remains to be discoA^ered, however, which hydrate, H3ASO3, HA-^(_)2 or some other, is present. On evaporation the anhydride and not the acid separates out. This acid forms three classes of salts, according as one, two, or three of its hydrogen atoms are replaced by metals; it is therefore tribasic. Certain salts are known which are derived from a meta-arsenious acid, HASO2. The salts of the alkalies are soluble in water; those of the other metaLs are not, but dissolve easily in acids, however. A neutral arsenite- solution gives a yellow precipitate of silver arsenite, AgaAsOs, with silver nitrate. The solution of the free acid is easily oxidized to arsenic acid by iodine solution: HgAsOj + I^-hHjO = H3AsO,+2HI. Such a solution can therefore also be employed for the titration of iodine (§ 9.3). 163.] ARSENIC ACID. 229 ARSENIC ACID, H3ASO4. 162. This acid is most easily obtained by the oxidation of a solution of arsenious acid by warming it with nitric acid. On concentrating the solution the compound 2H3ASO4 + H2O separates out (below 15°); this substance gives off its water of crystallization at 100° and yields orthoarsenic acid, H3As(34, which crystallizes in fine needles. When heated further it gives off water (at 180°) and goes over into pyroarsenic acid, H4AS2O7, which separates in the form of hard glistening crystals. On being heated still higher the latter compound gives up another molecule of water, the final product being white crystalline metu-urKcnic acid, HA.s( )3. This conduct is completely analogous to that of phosphoric acid; however, metaphosphoric acid cannot be con^•e^ted into the anhy- dride by heat as can arsenic acid (§ 160). The pyro- and meta- arsenic acids are stable only in the solid state; when treated with water they are converted into the ortho acid, the transformation being much quicker than with the corresponding phosphorous acids. Orthoarsenic acid is easily soluble m water. Its salts, the arsenates, exist in three classes; of the tertiary only those of the alkalies are soluble in water. The reactions of arsenic acid are very similar to those of phosphoric acid (§ 146); in this case also a mixture of ammonia, ammonium chloride and magnesium sul- phate {magnesia mixture) precipitates a white crystalline am- monium magnesium salt, Mg(NH4)As04-l-6H20. Ammonium molybdate produces a yellow finely crystalline precipitate, whose composition and appearance correspond to those of th« phos- phorus compound. The precipitates formed with sUver nitrate are, however, unlike in color : Ag3P04 is yellow, Ag3As04 reddish brown. Sulphur Compounds of Arsenic. 163. Three are known: arsenic disulphide (realgar), AS2S2; arsenic trisulphide (prpiment), AS2S3; arsenic pentasulphide, AsoSs. Arsenic disulphide, AS2S2, occurs in nature as realgar (§ 155). It forms beautiful ruby-red cri'stals of a specific gravity of 3.5. It is used as a pigment. It is manufactured artificially by fusing sulphur and arsenic together; the resulting products vary in composition, howevo'-. 230 INORGANIC CHEMISTRY. [§§ 163- ARSENIC TRISULPHIDE, AS2S3. Arsenic is precipitated from tiie acid solution of arsenious oxide by sulphuretted In-drogen as sulpiride; in this respect too it beha\'es as a heav}- metal. In the above reaction arsenic tri- sulphide is deposited as an amorphous }-ellow powder. A pure solution of arsenious acid gives no precipitate with sulphuretted hydrogen, but simply a yellow liquid (§ 196). Arsemc trisulphide occurs in nature as orpiment (§ 155), having a laminated crystal- line structure; it owes its name to its beautiful golden lustre. By fusing artificial arsenic trisulphide a product is obtained which is \'ery similar to the natural orpiment, but has a lower specific gravity (2.7 instead of 3.4). Commercially the trisulphide is prepared by fusing white arsenic with sulphur; the product still contains the oxide, however, and is therefore poisonous. Arsenic trisulphide is insoluble in water and in acids. ARSENIC PENTASULPHIDE, AS2S5. After sulphuretted hydrogen has been led into a \varm acidu- lated solution of arsenic acid for some time, arsenic is precipi- tated as an amorphous 3-cllow powder of the composition AS2S5. The latter is also obtained by fusing arsenic trisulphide with the required amoimt of sulphur. In the absence of air it can be sublimed without decomposition. It is insoluble in water and in acids. SULPHO-SALTS OF ARSENIC. 164. The trisulphide and the pentasulphide of arsenic dissolve easily in alkali sulphides, forming salts of sulpho-acids: As2S3 + 3K2S=2K3AsS3; Pot. Rulph- arsenite. As2S5+3K2S=2K3AsS4. Pot. sulph- arsenate. The formation of these sulpho-salts can be regarded as analo- gous to that of an oxy-salt from a basic oxide and an acid anhy- dride, e.g. : BaO + S03 = BaS04. 165.] ANTIMONY. 231 The trisulphide and the pentasulphide are therefore to be con- sidered as sulpho-anhydrides of those sulpho-acids. The sulpharsenates can also be obtained from arsenic tri- sulphide witli the aid of an alkali poh'sulphide: AS2S3+ 1^283= 2KASS3. Pot. sulpho- meta-arsenate. This reaction can be explained bj- supposing that the arsenic trisulphide is converted into the pentasulphide by the excess of sulphur, just as the trioxide is oxidized to the pentoxide. They are also produced by treating an arsenate with hydro- gen sulphide: K3ASO4+ 4H2S = IV3ASS4+ -IHzO. The sulpharsenates and sulpharsenites of the alkalies dis- solve easily in water and can be obtained in the crj'staUine form from the solution; those of the other metals are insoluble The free sulpho-acids are unknown. On the addition of an acid to the solution of a sulpho-salt, the liberated sulpho-acid brealvs up into hj'drogen sulphide and arsenic tri- or jxntasulphide. ANTIMONY. 165. Antimonj^ occurs in nature in slibnite, Sb2S3, as well as in man}- less common minerals. Stibnite was known to the ancients. In Japan it is found in magnificent large crystals. Antimony was frequently employed bj^ the alchemists. Basilius Valentinus in the latter part of the fifteenth century described its extractidn from stibnite in a monograph entitled "The tri- umphal cur of .Viitimoniuni." The element is at present obtained from stibnite by two processes. In one tlie mineral is roastoil, being thus transformed into antimonious oxide. This oxide is tlien reduced with charcoal to metallic antimony: I. 2Sb2S3 + 902 = 2Sb203 + 6S02; II. 2Sb203+3C =4,Sb + 3C02. The other method is to fuse the mineral with iron: Sb2S3 + 3Fe = 2Sb + 3FcS. 232 IXOBGANIC CHEMISTRY. [§§ 165- The crude antimony thus obtained usually still contains arsenic, lead, sulphur, etc. It can be refined by fusing with a Uttle salt- petre, the impurities being oxidized. Physical Properties. — Antimony is silvery-white and has a high metalhc lustre and a laminate-crystalhne structure (rhombo- hedral) ; as a result of the latter it is very brittle and can be easily pulverized. Sp. g. = 6.71-6.86. Melting-point, 629°; boiling- point, 1440° Mexschixg and V. Meter succeeded in determin- ing the vapor density at 14.37°, i.e. shghtly below the boiling-point, and found that the molecule, unlike that of phosphorus or arsenic, consists of less than four atoms. Like arsenic antimony has a black and a yellow modification; the latter is obtained by passing air into liquid stibine, cooled to —90°. It is even less stable than the yellow arsenic. Chemical Properties. — At ordinary temperatures the element is not affected by the air; when heated, it burns with a bluish- white flame to the trioxide. It combines with the halogens directh', producing scintiUations (§ 27). It is dissolved by hydrochloric acid, although very slowly, with the evolution of hydrogen, thus asserting its metallic character. Aqua regia dissolves it readily. Uses. — Antimony is a constituent of various alloys. The most important of these is type-metal, from which printer's type is made. Its approximate compostion is lead (50%), antimony (25%) and tin (25%). HYDROGEN ANTIMONIDE, STIBINE, SbHj. i66. Stibine is formed when nascent hydrogen acts on a solu- ble antimony compound. It is best prepared by treating an alloy of one part of antimony and two parts of magnesium with dilute hydrochloric acid. The product consists principally of hydrogen, but contains 10-14% SI3H3. If this gas mixture is passed through a U-tube and the whole is plunged in hquid air, stibine con- denses to a white sohd mass, that soon melts after the tube is removed from the hquid air. It vaporizes to a relatively stable gas. The least trace of oxygen, however, causes some antimony to be deposited. If an electric spark is passed through stibine gas it explodes, antimony is set free and the volume of hydrogen liberated is found to be In times that of the stibine, which is in accord with the formula SbH:j. It is also decomposed rapidly by heating the containing \-essel abo^•e 150°. 167.] HALOGEN COMPOUNDS OF ANTIMONY. 233 Stibine has a characteristic musty odor, quite unhke that of phosphine or arsine. AVlicii the mixture of hydrogen and stibine evolved from the alloy of antimony is heated, as in the Mahsu exi)criment (§ 157), it produces a metallic mirror and, when ignited, the flame gives a spot on cokl porcelain similar to that of arsenic, but differing from the latter in its darker color, insolubility in hypochlorite solution and less volatility when heated in a current of hydrogen. Stibine precipitates a black powder from silver solution, consisting of a mixture of silver and silver antimonide, AgsSb. The decomposition of stibine has been carefully investigated by Stock. He arrived at the conclusion that the decomposition velocity in clean glass vessels at room temperature proceeds at first with extreme slo^yness but increases more and more as more antimony separates out. Furthermore, mirrors of antimony produced by heating stibine and mirrors of black antimony made by subliming antimony in a vacuum and condensing the Vapor at the temperature of liquid air, and mirrors of sublimed metallic antimony, all had different effects. The effective- ness of the mirrors varied not only with the size, but in large measure also with the form of the antimony surface. It was fornid that the stibine dissociation in the layer adsorbed by antimony was proportional to the mass of the la\'er ; under this assump- tion the progress of the dissociation could be calculated theoretically, and was found to agree well with the experimental results. Since the amounts of adsorbed gas depend on the surface tension and, therefore, on the form of the adsorbing surfaces, the explanation of the influence of the different antimony mirrors is obvious. The fact that the walls of the vessel influence the velocity of a reaction has also been estabhshed in many other instances. Halogen Compounds of Antimony. 167. Two compounds of this clement with chlorine are known: SbC'ls and SbClj. Antimony trichloride, SljCls, is obtained by treating antimony sulphide or oxide with concentrated hydrochloric acid. It forms a colorless laminar-crystalline mass, which is so soft that it was formerly known as "antimony butter" {btihjnnn. aritiimniii). Its melting-point is 73.5° and its boiling-point '2'l?,Jf: its va]ior density 7.S (air = l) makes the formula SbCl.j. It dissolves in water containing h\'(.lrochloric acid. Water 234 INORGANIC CHEMISTRY. [§§167- decomposes it, forming difficultly soluble oxychlorides. The composition of the precipitate depends on the amount and the temperature of the water used in the decomposition. There is evidence of the existence of the compounds S1)()C1 and Sl)40sCl2 (=2SbOCI, SbzOa), both of which crystallize. The pre- cipitated oxychlorides on being repeatedly boiled with water eventu- ally lose all their chlorine and go over into the trioxide, SboOs. Powder of Algaroth, once used in medicine, is obtained by the decomjwsition of antimony trichloride with water and has nearly the same formula as the second of the ab(>\-e-mentioned ox)-- chlorides. Antimony pentachloride, SbC'ls, is prepared by heating anti- mony in a current of chlorine or treating fused trichloride with chlorine. It is a yellow, fuming, ill-smelling liquid, which crys- tallizes at —6°. Waen heated it dissociates into the trichloride and chlorine. It unites with water, forming SbCl5-H20 and SbCl5-4H2(). Hot water decomposes it into hydrochloric and pyroantimonic acids. Oxygen Compounds of Antimony. i68. Three are known: antimony trioxide, Sb203, antimony tetroxide, Sij2(->4, and antimony pentoxide, SbaOs. Antimony trioxide occurs as a mineral, senarmontite. It can be obtained by burning antimom' in the air, as well as by the oxidation of antimony with dilute nitric acid. It is dimorphic, occurring in both regular and rhombic crystals. It is a light yellow crystalline powder, ahriost insoluble in water. It volatilizes at 1560°; the vapor density at this tem- perature corresponds to the fonnula Sb406. It is insoluble in sulphuric and nitric acids but easily soluble in hydrochloric and tartaric acids and in alkalies. On being heated in the air it turns to the tetroxide. The corresponding hydroxide is Sb(0II)3. This hydrate sepa- rates out when tartar emetic (see below) is decomposed with dilute sulphuric acid. It gives up one molecule of water readily and passes o\'er into the hydroxide SbO -OH, meta-antimonious acid. The latter is more easily obtained by treating a solution of the trichloride with soda solution: 169.] -iXTIMOA'Y PJiNTOXlDE AM) ASTIMONIC ACID. 2:j.j 2SI.1CI3 + 3Xa.( 'O3 + HaO - 2S1)0 • OH + 6NaCl + SCOg. It appears as a white precipitate, which is converted into an- timonic oxide by boiling with water. This meta-antimonious acid is dissolved by alkalies, forming salts of the acid. C)iio of them which has been obtained crystallized is tlic sodium nicta- aatimonitc, XaSbOo + SH^tJ. The latter is difficultly soluble in water, and decomposes on concentration of its solution. (3n the other hantl, antimony hydroxide displaj'S basic proper- ties \jy unitiuij; with acids to form salts. There are salts known of Sb(()H)3, as well as of SbO-OH. Examples of the former kind are the crystallized antimony sulphate, Sb2(S< )4)3, and the nitrate, Sb(X03)3. In analogy with other triA-alent metals double salts are known, e.g. KSb(S(J4)2. As to tlie salts derived from SbO-OH, we may look upon the group St)( > as taking the place of a uni- valent metal. Thus Sl)r)OH may be compared with KOH. For this reason the group (SbO) has been gi^•pn the name aniimonyl; one of its salts is antunonyl sulphate, (SbO)2S04. The most familiar antimonyl compound is tartar emetic ^ potassium antimony! tartrate, /g|jQ)t'4H406 + iH2(), which is employed in medicine. See Org. Chem. § 192. ANTIMONY PENTOXIDE AND ANTIMONIC ACID. Antimonic acid, H3Sb04, is obtained by warming antimony with concentrated nitric acid and also by decomposing the penta- chloride with water. It is a white powder, almost insoluble in water and nitric acid; ne\-crtheless, when moist, it turns litmus paper red. On heating saltpetre with powdered antimony the potassium salt of meta-antiinonic acid, KSbOa, is formed in an explosive reaction, ^^dlen this is boiled with water it dissolves, protlucing mono'potassium orthoantimoniate, KH2Sb04; on fusing with potash -potassium pyroantirnoniate, K4Sb207, is formed, which dissolves in water, giving 2K0H and K2ll2Sb207+6Il20. In the case of antimony, as in that of phosphorus, we meet with three kinds of acids belonging to the highest stage of oxidation : their f ormuliE correspond to those of the analogous phosphorus compounds. 236 LXORGAXIC CHEMISTRY. [§ 168- Antimony pentoxide, Sb205 (molecular weight unknown), can be obtained by heating antimonic acid at 300°. It is a yellow amorphous powder, soluble in hydrochloric acid. If heated strongly it gives up part of its oxygen and goes over into anti- mony tetroxide, Sb204, a white powder that turns yellow on heating but resumes its original color on cooling. This tetroxide can be regarded as antimonyl meta-antimoniate, SbOs-SbO. Sulphur Compounds of Antimony. 169. Antimony trisulphide, Sb2S3, is found in nature (§ 165). It can be made by leading hydrogen sulphide into a hydrochloric acid solution of the trichloride, from which it is deposited as an amorphous red powder. It can be melted; on cooling it crystal- lizes and takes on the appearance of stibnite. Antimony pentasulphide, SbaSs, is precipitated when hydrogen sulphide is passed into the acidified solution of antimonic acid. It is more easilj- obtained by the decomposition of sodium sulpli- antimoniate with dilute sulphuric acid. It forms an amorphous orange-red powder, which splits up into sulphur and the trisulphide on being strongly heated. It is insoluble in dilute acids; boiling- hot concentrated hydrochloric acid dissolves it, forming antimony trichloride, hydrogen sulphide and sulphur. In aqueous solutions ^f alkalies and their sulphides it dissolves easily with the formation of sulphantimoniates, M3SbS4 The best known of these is sodium sulphantimoniate, Na3SbS4+9H20 (" Schlippe's salt"). It can be obtained by boiling antimony trisulphide with sulphur and caustic soda solution. It crystallizes in large colorless tetrahedrons, is easily soluble in water (1 part by weight in 2.9 parts water at 15°) and reacts alkaline. It is decomposed by acids, depositing pentasulphide; even carbonic acid causes this, hence the crystals become covered with a j-ellowish-red coating of pentasulphide after having stood some time in the air. The free sulphantimonic acid is not known. BISMUTH. 170. This element belongs undoubtedly among the metals, so far as its physical character is concerned; its chemical properties also class it with them in almost every respect, inasmuch as its oxides are mainly basic in their behavior. 171.] COMPOUNDS OF BISMUTH. 'I'M It is found chiefly in the native state; but a sulphide, Bi2S3, bismuth glance and a telluride, Mradymitc, also occur in nature. Bismuth is obtained from the latter by roasting to the oxide Bi203 and reducing -with charcoal. The nati^•e metal is usually very pure. If refining is necessary, the fused metal is allowed to flow over a hot, 'Somewhat inclined iron plate, so that the impurities are oxidized. The amount of bismuth found in nature is not very great. Physical Properties. — Bismuth is externally very similar to antimony; it is crystallized and very brittle and has a metallic lustre, but differs from antimoii}' in' having a redchsh-white color. Sp. g. = 9.S'23. It melts at 2N(i.3° and boils at 1420°. It can be distilled in a current of hydrogen. Chemical Properties. — At ordinary temperatures bismuth is unaffected by the air. On being heated it turns to the trioxide. It combines with the halogens directly. It is not attacked by h}-drochloric or sulphiu-ic acid at ordinary temperatures, but nitric acid dissolves it readily to form the nitrate. On being heated with sulphuric acid, it gives off sulphur dioxide and forms the sul- phate. No hydrogen compound of bismuth is known. Bismuth is employed in the manufactiue of easily fusible alloys such as are used in making casts of woodcuts, stereotypes, etc. The most common of these alloys are Newton's metal (8 bismuth, 5 lead, 3 tin; melting-point 94.5°), Rose's metal (2 bismuth, 1 lead, 1 tin; melting-point 93 75°) and Wood's metal (4 bismuth, 2 lead, 1 tin, 1 cadmium; melting-point 60.5°). Halogen Compounds. 171. Compounds of the type BiXs only are known. Bismuth chloride, BiCla, is formed by direct synthesis from the elements, but more easily by dissolving bismuth in aqua regia. It is white and crystallized. Its melting-point is between 225° and 230° and its boihng-point at 435°. Its vapor density, 11.35 (air=l), gives it the formula Bids. On being dissolved in a little water it forms a sirupy liquid ; an excess of water gives bismuth oxychloride, BiOCl, and hydrochloric acid. This oxychloride is a white powder, in- soluble in water but soluble in acids. 238 INOIiGANIC CHEMISTRY. [§§ 172- Oxygen Compounds. 172, Three oxides are known: BiO, Bi203, and Bi02. Bismuthous oxide, BiO, is obtained by adding an allvaline stannous chloride solution to a solution of bismuth chloride. It is deposited as a dark-brown precipitate ot BiO. ^^'hen heated in the air it smolders like tinder. It is doubtful whether this precipitate is a homogeneous substance or a mixture of 61303 with finely divided bismuth. Bismuth trioxide, 61203, is the most familiar oxide of this ele- ment. It has strictly basic properties. In order to prepare it we can heat the nitrate or carbonate or we can precipitate the hydroxide from the solution of a bismuth salt by means of a base and heat the precipitate. If a boiling solution of a bismuth salt is treated with caustic potash, the trioxide separates out in glistening needles of microscopic dimensions. Like the corresponding oxides of arsenic and antimony, it is dimorphic. Bismuth dioxide, BiOj, has been little studied; it is a reddish-yellow powder. Bismuth pentoxide, BizOs, appears as a reddish-brown powder, which is very unstable and evolves oxygen on heating, as it also does when warmed with sulphuric acid. Hydrochloric acid does not convert it into the corresponding pentachloride, BiClj, but produces the trichloride BiClj and free chlorine. Hydroxides and Salts. 173. Bismuth hydroxide, Bi(0H)3, is obtained by precipitating a bismuth salt with an alkali. It is an amorphous white powder, insoluble in potassium hydroxide or ammonia. At 100° it goes over into the compound BiO -OH with the loss of a molecule of water. Both of these hydroxides are wholly basic in character. The salts derived from Bi(0H)3 are called neutral, those from BiO OH basic. The neutral nitrate, Bi(N03)3, is obtained by dissolving bis- muth in nitric acid. It crystallizes with five molecules of water in large translucent triclinic prisms. It is deliquescent. The addition of much water converts it into the basic nitrates, several of which are known. By treating the neutral nitrate with about 20 parts of boiling water a product is obtained whose composition is not perfectly constant for different preparations, but corresponds nearly to the formula (Bi203)6- (N205)5-(H20)9, or 2Bi0XO3 17o.] SUMMAh'V (IF THE MTJ^XiKX (SUOUP 23! I + Bi(N03):j + 3Bi(OH)3. This is the bismuth subnitrate, which is used in medicine. Bismuth sulphate, Hi2(S().|)3, is <)l)taiiir(l as an amorphous white sulistance when the metal is lieatcd with concentrated siil- phuric acid. With water it forms a l)asi<' sulphate, Bi2(OH)4S()4. Sulphur Compounds. 174. Bismuth trisulphide is found in nature (§ 170); artificially it can he pre[)ared 1)\- heating bismuth with sulphur or by leading hydrogen sulphide into tlie afpieous solution of a bismuth salt. In the latter case it comes down as an amorphous black powder that is easily soluble in warm dilute nitric acid. It is insoluble in alkalies and their sulphides, hence forms no sulpho-salts. When heated with an alkali sulphide solution to 200° it takes on a crys- talline form similar to tliat of the natural mineral. .-^r.M.M.\RY OF THK XITItoOKX ClIOIT 175. Like the lialnnens and the elements of the oxygen gi'nup, the elements just .5- In other words, the elements of this group are trivalent or pentivalent. We find here, just as in the gi-oups prex'iously studied, that, as the atomic weight increases, a gi'adual change occurs in the phvsical properties. This is shown by the following small table: .\. P. \-. Sl,. Bi. Atoiuif flciglil. 14.01 ,31 74 !i(i 120.2 L'OS.O Specific (rr:i\iU-. .s,s.-, 1.8-2 1 4,7-.-) 7 (i 7 '.1 s 1 Water- 1 1 liquid Melting-point. . + 44 4° ca. Ndll 112:1=' L'.Mi-" Boiling-point. , . . -194 i'' + 27S° 1440° 1420° Color colorless yclliAN or red gr.'iy white pink In the chemical properties, also, regular A-ariations are (n be observed, all of which can be siuiimed uj) in the general statomi rt 1240 IXORGAXIC CHEMISTRY. [§§ 175- that the metalloid character gives way to the metallic character as the atomic weight increases. Nitrogen forms either indifferent or acid-forming oxides only; so does phosphorus; arsenic, on the contrary, displays a very feebly basic character in arsenious oxide, since this oxide forms the trichloride with hj'drochloric acid, the trichloride reacting inversely with water, howe^-er, and breaking up into hydrochloric acid and arsenious oxide. In antimony trioxide this basic character is a little stronger; some salts and double salts of it with acids are known. The corresponding chlo- ride does not suffer an immediate hydrolytic dissociation with water, but oxychlorides are formed, which require a great deal of water to convert them entirely into the trioxide. While the highest oxides of arsenic and antimony have strictly acid prop- erties, with bismuth the acidic nature has practically disappeared; the oxide Bi2(J3 has exclusively basic properties and the higher oxide Bi205 acts like a peroxide, giving off oxygen readily (it generates chlorine with hydrochloric acid) and going over into the lower oxide 61203. Bismuth trichloride, BiCls, gives the oxy- chloride, BiOCl, with water and this is not decomposed by an excess of water. In the hydrogen compounds, too, the gradual change of the properties is very apparent. Consider the stability for example: ammonia requires a very high temperature for decomposition; phosphine and arsine a much lower temperature; stibine is unstable at ordinary temperatures when it comes in contact with oxygen, and the hydrogen compound of bismuth is so unstable that the conditions for its formation and existence have not yet been ascertainable. A similar change is noticeable in their ability to form XH4 ions in aqueous solutions; it is strong in ammonia, much weaker in phosphine and wholly absent in arsine and stibine. In the sulphur compounds a progressive change of color is observed. P2S5 is bright yellow, AS2S5 deep yellow, Sb2S5 red and 61285 black. The first three are sulpho-anhydrides of sulpho- acids (§ 164); bismuth sulphide is not, however, thus displaying .again the more basic nature of bismuth. 176.] ALLOTROPIC FORMS OF CARBON. 211 CARBON. 176. Carbon occurs in nature both free and combined. In combinatidn it is found in large quantities in the salts of carbonic acid, above all in calcium carbonate, limestone, which is of the widest occurrence and is even known to form great mountains. Farther, carbon is one of the constituent elements of animals and plants. It is found in these in numerous compounds. Still larger is the number of artificially prepared carbon compounds. The com- pounds of carbon exceed in number all other compounds together. For this reason and because of the peculiarities of the carbon compoimds it is customary ti 1 treat them by themselves, as " organic chemistry." Ho\vc\'er, that we may be able to obtain a general survey of the elements, it is deemed ad\isable to discuss certain compounds of carbon in inorganic chemistry as well. AUotropic Forms of Carbon. We know of three: diamond, graphite and amorphous carbon. (a) Diamond. — Lavoisier found, in 1773, that this mineral can be burned to carbon dioxide. In 1814 Davy proved that, when diamond burns, nothing else than this gas is formed, so that diamond must be pure carbon. Furthermore when the carbon dioxide given off by the combustion of diamond is absorbed by sodium hydroxide, a soda is produced which is in every respect identical with ordinary soda. . Indeed, it has been found possible to manufacture diamonds from amorphous carbon (see below). The diamond crystallizes in the isometric system. Usually it is colorless, but yellow and black diamonds are also known; the black ones are called carbonado. The specific gravity of diamond is 3.50-3.55. It is a poor conductor of heat and electricity. The refractive index is very high: n=2.42. The diamond is so hard that it scratches all other substances. If it ig subjected to a very high temperature in the absence of air, it gradually turns to graphite. It resists the action of the strongest oxidizing-agents, e.g. a mixture of nitric acid and potassium chlorate. In 1893 MoissAN succeeded in making diamonds artificially, although they were very small, the largest being about 0.5 mm. in 242 IXORGAyiC CHEMISTRY. [§176- dianieter. His method consists essentially in dissolving carbon in molten iron at a high temperature and then cooling it rapidly. Fig. 39. — Artificial Diamoxds (JIag.mfied). This is accomplished as follows: Iron is brought in contact with pure carbon (sugar charcoal) in the electric furnace at a high tem- perature. After the iron has become saturated with carbon at about 3000°, the fused mass is suddenly cooled by pouring it into a hole drilled in a copper block, which is kept cold by water, and at once covering the cavity with an iron stopper, ^'^^len the iron is all cold it is dissoh'ed away by acids, lea^'ing the carljon which did not combine with the iron. This residual carbon consists partly of small diamonds, which are identical with the natural diamond in hardness, crystal form, etc. Fig. 39 presents an enlarged view of some artificial specimens; they display the same properties as the rough natural diamonds, particularly the rounded edges and angles and the striations. The formation of the diamond by this method has been explained by Bakhuis Roozeboom as follows : In all probability the transition of diamond into graphite is endothermic. For this reason diamond is the more stable form at lower temperatures, graphite at higher ones, in analogy to the rhombic and monoclinic modifications of sulphur. But, while the velocity of transformation of monoclinic sulphur is fairly great at low temperatures and the monoclinic sulphur can thus exist only for a short time below its transition point, the transition velocity of graphite into diamond is practically zero for temperatures below 1000°. Carbon that has crystallized from molten iron in the form of graphite cannot, therefore, pass over into diamond. The rapid cooling of the molten iron, however, 17G.] ALLOTROPIC fOUMS OF CAHIiOX. 243 has the effect of briuij;in,<;- the rarlidii into the region of tcni])er;iture ill which clianioiul is the stahie modification; it can therefore separate in this form from its sohition. The electric furnace that Mmssw uscil for tlicsc and numerous other experiments is \-cry simple in cdiisti ui'tioii. It slightly hollowed nut on its lower side se as tu rcllccf the heat ray^ on to the crucible. Fig. -tO shows a cros.s-scctinn of an electric furnace, Fig. 41 a picture of the same apparatus in operation. The temperatures obtained in the electric furnace are as follows. Current of 30am]x"'res and .■),') \-(jUs with a stcain-enginc of 4 H.P., 22.")0^ " 100 " 4.30 4.1 " S " 2.")00° "50 " 3000° Fig. 11 — ^[oi,s~an's Fccn \ck ix < Ii'er \ticix. {.\fter Mdi.ssan.) 244 INORGANIC CHEMISTRY. [§§ 176- The last-named temperature can however only be maintained for a brief period, as the unslaked lime soon melts and flows like water. At 2500° the Hme becomes crystalline in structure after a few minutes. (6) Graphite is also crystallized carbon. Unlike diamond, it is very soft and opaque and a good conductor of heat and electricitj'. Sp. g. = 2.09-2.23. As was stated above, graphite can be pre- pared artificially by the crystallization of carbon from molten iron and by heating diamond strongly. There are \'arious kinds of graphite. If graphite is treated with a mixture of perfectly dry potassium chlorate and very concentrated nitric acid, it turns to a yellow crystallized -substance containing hydrogen and oxygen, in addition to carbon, and called graphitic acid. This substance is peculiar in that it decomposes explosively on heating and yields a large volume of extremely fine amorphous carbon. Graphite is used in the manufacture of lead pencils, crucibles, electrodes, polishes, etc. (c) Amorphous Carbon. — This is obtained Sn the purest state by charring sugar. The resulting mass is boiled with acids to remove the mineral matter and finally heated red-hot in a current of chlorine for quite a while to remove all the hydrogen. It can also be prepared from soot. Amorphous carbon is opaque, black and infusible. At the highest temperature that Moissan could reach with his furnace by employing a current of 2000 amperes and 80 volts (obtained with a 300 horse-power engine) it was barely possible to make carbon sublime. The sublimate was graphite. Amorphous carbon has a specific gravity of 1.5-2.3. ^^arious sorts of amorphous carbon are known. They are probably different allotropic modifications, or mixtures of such. Gas carbon and coke are obtained as residues in the drj^ distilla- tion of coal. They conduct heat and electricity. Wood charcoal is very porous and can condense large quantities of gases in its pores, e.g. 90 times its own volume of ammonia (see also § 111). When warmed or when the pressure is reduced, these gases all escape again. Bone-black is obtained b}- heating bones away from air; tlie resulting black mass is treated with hydrochloric acid to remove the phosphates and carbonates present. It has the power of absorbing coloring-matter and certain salts, e.g. lead salts, from liquids. The charcoal obtained from the dry distilla- tion of sugar is noted for its peculiar lustre. These different 177.] ALLOTIiOPIC FORMS OF CARBON. 245 sorts of charcoal do not consist of pure carbon but contain other substances in small proportions. It is a general rule that carbon conducts heat and electricity belter the longer it has been exposed to a high temperature. 177. The various kinds of carbon all find their respective uses. Soot, or lamj)black, serves for the preparation of India ink and black paint. Gas carbon (coke), being a good conductor of electricity, is used in the electrical industrj-. Wood charcoal is used in the manufacture of gun- powder; animal charcoal, or hone-black, as a water-filter to remove color- ing-matter, ill-smelling gases or injurious salts (lead salts) from drink- ing-water ; it is also employed in enormous quantities in sugar refineries to decolorize sugar liquids. By far the most important use of carbon is as a f u e 1 . The heat generated by the burning of coal warms our houses, drives our steam- engines, etc. The principal kinds used as fuels are charcoal, coke, anthracite coal, hitummous coal, brown coal (lignite) and peat. Charcoal (wood charcoal) is made on a large scale by the colliers. Long sticks of wood are piled in a large heap, covered with sod and ignited at the bottom. The wood smolders away slowly and becomes completely charred. This " charcoal-pit " process is not at all economical, inasmuch as all the volatile products are lost; it is carried on exten-. sively (Fig. -i'2). but it is being more and more replaced by the dry dis- tillation of wood from iron retorts, in which process the gaseous and tarry products are recovered. Fig. 42. — Charcoal Pit. Coke is the residue in the refoits of the gas factories after the coal has been deprived of its volatile products Ijy heating. It is also manu- factured on a large seale for metallurKieal and other purposes. Coke is thought by many to liave a great future as a fuel, since it is a hard- burning smokeless fuel, manufactured from the cheap soft eoal. 246 INORGAXIC CHEMISTRY. t§§ 177- Peat and the various coals owe their origin to the same geological process, the slow decay of plant-remains in the absence of air. Peat is the youngest formation and anthracite coal the oldest. During this transition carbon dioxide and methane, CH^, are given off and the residue becomes richer in carbon and poorer in hydrogen and oxj'gen than the corresponding chief constituent of plant tissues, cellulose. The follow- ing table shows this: Carbon. Hydrogen. Oxygen. Cellulose Peat . . . 50.0% 60.0 67.0 85.8 94.0 6.0%, 5.9 5.8 5.8 3.4 44.0% 34 1 Brown coal 27.2 Caniiel coal 8 3 Anthracite coal 2.6 The plants of which these formations originally consisted are different. Peat appears from its structure to have come chiefly from swampy growth.s, mosses and the like; mineral coal from extinct plants, gigantic horsetails (equiseta), lepidodendra and sigillaria;. Molecular and Atomic Weight of Carbon. — Chemical Properties. 178. The carbon molecule probably contains a large number of atoms. It has not yet been possible to determine how large this number is. It is supposed that graphite has a larger number of atoms to the molecule than amorphous carbon, and diamond more than graphite, since graphite and diamond are less easily attacked by chemical reagents and because they are denser. A determination of the vapor density of carbon is of course out of the question. The measurement of the melting-point depression that carbon produces in iron is also impracticable; however, it is known that even a small percentage of carbon causes a considerable lowering of the melting-point of iron (see § 304). It can be shown in the following way, however, that the number of atoms in the carbon molecule must be very great. B}- the oxidation of amorphous carbon with potassium permanganate mellitk acid is formed, which contains 12 carbon atoms to the molecule. This makes it quite probable that the carbon mole- cule contains at least 12 atoms, for in the oxidation of organic substances the products almost always contain either a smaller or the same number of carbon atoms to the molecule. For the 178] MOLECULAR AXD ATOMIC WEIGHT OF CARBOS. 247 following reason it is, however, to be supposed that the number of atoms in the carbon molecule is much greater than 12. When 7narsh-gas, CH4, is passed through a red-hot tube, vthi/lcne, C2H4, is formed among other things. If this is then treated in the same wa}', acetylene, C'jHo, L-^ obtained, and from this again benzene, OeHg. On conducting benzene vapor through a glowing tube, yiaphth-alene, CioHg, pyrene. CieHio, etc., are formed. If either of the latter is heated still higher (in the absence of air) carbon is deposited. We thus see that as the temperature rises the num- ber of carbon atoms in the molecule steadily increases. The final product of these operations, carbon, will therefore probably contain a considerably larger number of atoms in its molecule than naphthalene or ]jyrene. Carbon can unite directly \vith many elements. At ordinary temperatures it coml):iics with fluorine only. ^I(1Is,s.\n intro- duced lampblack into fluorine gas, and the carbon commenced to glow; when fluorine was present in excess carbon tetra fluoride, CF4, was formed. Hydrogen combines with carbon directly to form acety- lene and a small quantity of marsh-gas, wlien an electric arc is ]iassed between two carbons in an atmosphere of hydrogen. Of all the many compounds consisting of only carbon and hydrogen these are the only ones which can be obtained by direct synthesis. I'nder analogous conditions carbon unites with chlorine to inrm perchloroethane, C2CI6, and hexachlorobenzene, CeCle- Oxygen unites with carbon at an elc\'ated temperature to fiirni carlion monoxide, CO, or carbon dioxide, CO^, according as carbon or oxygen is in excess. If sulphur vapor is passed over red-hot coals, carbon disulphide, CSo, is produced. The elements of the nitrogen group, X, P, As, Sb and Bi, do not combine with carbon directly. Silicon and car- bon unite at the temperature of the electric furnace to form CSi, carborundum, which is so hard that it can be used as a powder for polishing glass and [jrecious stones. ^loissAN also found that many metals are able to coml)ine with carbon at a very high tempeiature, forming earhides. This was previously known to be true of iron and certain other metals. 248 INORGANIC CHEMISTRY. [§§ 179- The difference in the behavior of these carbides towards water is interesting. Iron carbide is unaffected by it; calcium carbide gives acetylene, C2H2; aluminium carbide yields methane; other carbides give mixtures of the two hydrocarbons; uranium carbide produces methane and also liquid and solid hydrocarbons. 179. The atomic weight of carbon has been determined with great accuracy by Dumas and Stas. The averages for the different series of experiments, each of which showed httle variation, were as follows: Ratio by weight of carbon to oxygen in carbon di- oxide from the combustion of: Natural graphite 2.9994:8.0000 .\T-tificial " 2.999.5:8.0000 Diamond 3.0002:8.0000 The ratio of carbon to oxygen in carbon dioxide is thus very close to 3 : 8. As the specific gravity of carbon dioxide points to a molecular weight of 44 for this gas, it must contain, according to this ratio, 12.00 parts by weight of carbon and 32 parts of oxygen. The formula is therefore Cj:02. Inasmuch as no carbon compound is known whose molecular weight includes less than 12 parts of carbon, we have CO2 as the formula; hence the atomic weight of carbon must be 12.00 for 0=16. Compounds with Hydrogen. 180. Carbon and hydrogen form a very large number of com- pounds {hydrocarbons), which are more fully discussed in organic chemistry. Two of them will be treated here briefly. Methane, also called marsh-gas and fire-darnp, is the only hydro- carbon with just one atom of carbon. It occurs in nature in volcanic gases; moreover, it gushes out of the ground in the neigh- borhood of the oil-wells at Baku and various places in America. It is an important constituent of " natural gas." It owes the name "marsh-gas" to the fact that it arises from swamps, especially when the decaying vegetation at the bottom is stirred up. It is called "fire-damp " because it occurs in coal beds (§ 177), from which it escapes when they are broken up. It forms a violently explosive mixture with air, which is frequently the cause of mine explosions. For its modes of formation and its physical and chem- ical properties reference should be had to Org. Chem., § 29. 181.] COMPOUNDS WITH OXYGIiX. 24'.) i8i. Acetylene, C2H2, is a colorbss gas of a disagreeable odor. It is soluble in an equal volume of water at 18° ;iiid bepoiucs liquid at I.S° under 83 atmospheres. Its hydrogen atoms are replaceable by metals. It is manufactured by decomposing calcium carbide with water: CaC2 + 2H20 = Ca(OH)j + C,H2. Calcium carbide is prepared by heating coke with unslaked lime (CaO) in the electric furnace. The calcium formed by the action of carbon on Ume unites with carbon at the high temperature of the fur- nace to form CaCj. Acetylene burns with a vivid flame on coming out of a smaD orifice under pressure. Since it can be prepared from calcium carbide pretty cheaply, it is used rather extensively in small systems for illuminating purposes. When mixed with air and ignited it explodes vehemently; the compounds with metals are also explosive. It is endothermic and can be exploded by fulminating mercury. The combustion of acetylene is another illustration of the rule of § 137, that reactions are in most cases of a simpler nature than the chemical equations indicate. The equation here is: 2C2H3+ 502= 40)2 +2HA According to this equation the combustion should be septimolecular. Bone and Cain proved, however, that the reaction has more than one stage, the first stage being represented by the bimolecular equation: C,,H,+02 = 2CO+H2; CO and H2 then burn further to CO, and 11,0. From a kinetic standpoint, it is quite conceivable that polymolecular reactions should be rare, for the probability of a large number of mole- cules coming together in just such a way that a reaction can take place is indeed very slight. The reaction is more likely to proceed in a way which involves the interaction of only verj' few molecules. Compounds with Oxygen. Three oxygen compounds of carbon are known: carbon monoxide, CO, carbon dioxide, CO2, and carbon suboxide, C3C2. For the latter compound, see Org. Chem. § 166. CARBON MONOXIDE, CO. 182. This gaseous compound is always formed when carbon biu-ns in a limited supply of air or oxygen. A number of carbon compounds also yield carbon monoxide when burned under this same condition. It can also be obtain(Hl by the action of carbon on oxygen compounds, e.g., by heating zinc oxide, ZnO with 250 INORGAXIC CHEMISTRY. [§§ 182- carbon. On passing steam over red-hot coals a mixture of hydrogen and carbon monoxide is produced: This mixture goes by the name of water-gas. It is used on a large scale for heating and lighting, especially in America. For the latter purpose it is charged with the vapor of hydrocarbons rich in carbon, since its own flame is not luminous. The use of the incandescent gas- light (§2'il) makes this " carburetting " unnecessary. Water-gas con- taining 50': i, of carbon monoxide is very poisonous (Org. Chem. § 241). Carbon monoxide is also formed by the reduction of carbon dioxide with red-hot carbon: C + C02=2CO. This reaction is Umited b}- the reverse one and we have here a case of balanced action expressed by C-hC02?^2CO. In view of the caloric effect of the reaction, 2CO=C + CU2-h3900 Cal., an elevation of temperature must, according to Le Chatelieh's rule (§ 51), increase the amount of carbon monoxide; a depression of temperature, the opposite. Experience has shown this to be actuall}' the case. As the temperature rises the quantity of carbon mondxide increases rapidly and at 1000° there is still a very small amount of dioxide. At 445°, on the other hand, practically all the carbon monoxide is changed into carbon dioxide and carbon. This result is surprising, because the same change should also occur at lower temperatures; nevertheless, carbon monoxide seems perfectly stable at ordinary temperatures, e^•en as high as 200°. The cause of this phenomenon must, as in analogous cases, be sought in the very great retardation of the velocity of the reaction 2C()^C02 4-C when the temperature sinks. On using certain catah'zers, e.g. finely divided nickel, the velocity of the reaction 2C0— >C02-f-C becomes measurable as low as 256°. These measurements ha^•e shown that the decomposition of carbon monoxide into carbon dio.xide and carbon is not a bimolecu- lar reaction, as would be expected from the above equation, but 183.] CARBOy MOXOXIDE. 2ol a unimolecular one. To explain this it may bo suggested that the decomposition takes place in two stages: I. CO=C + 0; II. C0 + 0=C02. If we assume that the second stage has an infinite velocity, it is only the first that is really measured, i.e. a unimolecular reaction. The reduction of salts of carbonic acid also furnishes a method of preparing carbon monoxide. If chalk (CaCOs) or magnesite (MgCOs) is heated with zinc dust, pure carbon monoxide is formed: CaCOa + Zn = CaO + ZnO + CO. Physical Properties. — Carbon monoxide is a colorless, odorless gas of a specific gravity of 0.967 (air = l). It is hard to condense, its critical temperature being —139.5° and its critical pressure 35.5 atmospheres. It boils at —190° and solidifies at —211°. It is only slightly soluble in water. 183. Chemical Properties. — Carbon monoxide burns with a characteristic blue flame to carbon dioxide. It can unite with chlorine dii'ectly to form phosgene, COCI2, and also with sulphur (at an elevated temperature) to form carbon oxysulphidc, COS both compounds are gaseous. Again, it unites directly with nickel and iron, giving the compounds Ni(C0)4 and Fe(CO)j (§§ 21-1 and 311). On account of its tendency to combine with oxygen, it displays strong reducing power, especially at high temperatures. Thus metallic oxides, like FeoOa, CuO, etc., arc easily converted into the metals when hot. ."^ome compounds are reduced l>y carbon mon- oxide even at ordinary temperatures. Palladium is precipitated from an aqueous solution of palladious chloride and an ammoniacal silver solution (prepared by dissolving sih-er oxide in ammonium hydroxide to the point of saturation) is turned black by the gas on account of formation of the metal. Both of these reactions serve for the detection of carbon monoxide. An ammoniacal cuprous chloride solution absorbs the gas because of the formation of a compound, Cu2Cl2-C()-l-2H2(>, which can be isolated in the crystalline state but decomposes again very readily. The composition of carbon monoxide can be determined by exploding a mixture of the gas with oxygen. It is then found that 2 vols. CO unite with 1 vol. O? to fonn 2 vols. COg. This together with the vapor density establishes the fonnula ns CO. 2,V2 JXORGAXIC CHEMISTRY. [§§ 18!- CARBON DIOXIDE, CARBONIC ACID ANHYDRIDE, COj. 184. This compound occurs not only by itself but also in com- bination. It is a regular constituent of the air (§ 106); many mineral waters contain the free gas; in some places of the earth (in the Dog's Grotto at Naples and the famous Poison Valley in Java) it comes up out of the ground and it is also found in volcanic exhalations. The most minerals and rocks contain numerous extremely small cavities, partly filled with liquid carbon dioxide. Combined, it occurs in large quantity in the carbonates of lime and magnesia (§ 176). Carbon dioxide results from the combustion of carbon in an excess of oxygen and also from the direct decomposition of many salts of carbonic acid {carbonates) by heat: 2XaHC(.)3 = Na2C03 + H20 + C02; CaC03 = CaO + C02. Sodium bi- carbonate. Moreover, it is formed when a carbonate is decomposed by an acid: Na2C03 + 2HCl = 2NaCl + H20 + C02. By the action of oxygen at high temperatures all carbon com- pounds are burned with the formation of carbon dioxide. It is also produced by the action of carbon on oxygen compounds, e.g. by heating powdered charcoal with an excess of copper oxide; finally also by the interaction of carbon compounds and oxygen compounds. This latter action is the basis of the general method for determining the proportion of carbon in organic substances; thc}- are heated together with copper oxide and the carbon dioxide formed is absorbed in a weighed amount of caustic potash. PJn/ftlral Properties. — Carbon dioxide at ordinary temperatures and pressures is a gas with a somewhat pungent odor and taste. Sp. g. = l..")2fl (air=l). It is thus about half again as heavj^ as air, so that in those places where it comes out of the earth, as in the Dog's Grotto at Naples, it stays in a layer close to the ground and a dog, for instance, is suffocated while a man can breathe with comfort. Carbon dioxide is easily condensed, becoming liquid at 0° under 35 atmospheres pressure. Its critical temperature is 31.35° and its critical pressure 72.9 atm. Liquid carbon dioxide (" liquid carbonic acid ") is manufactured in great quantities and 184.J (ARBOX DIOXIDE. 253 brought on to the market in steel bottles (bombs). It is a very mobile liquid, which is not miscible with water in all proportions. If the liquid is allowed to escape from the bomb into a coarse linen bag (by inverting the bomb and opening the va.\ve), part of it vaporizes, absorbing hereby so much heat that the remainder solidifies in white flakes. A mixture of this solid carbon dioxide with ether, alcohol or acetone is often used as a freezing-mixture; it enables us to obtain a temperature of —80°, and even —140° in vacuo, ^\\-^en liquid carbon dioxide is cooled down in a sealed tube, it congeals to an icy mass, which melts at —65° At 15° carbonic acid gas dissolves in its own volimie of water (more accurately 1.0020 vol.); at 0° in 1.7967 vol. In alcohol it is still more soluble. Chemical Properties. — Carbon dioxide is a -\'ery stable com- pound; it is only decomposed by intense heat (see § 182) or by the continued action of induction sparks, breaking up into oxygen and carbon monoxide. This decomposition never completes itself, for just so soon as a certain amount of these gases have been formed, they reunite with explosion. At the moment before the explosion the amount of carbon dioxide still present becomes no longer suffi- cient to dilute the mixture of oxygen and monoxide enough to hinder an explosion; the explosive limit is reached. Carbon dioxide cannot be farther oxidized; it is therefore not combustible. In general it cannot support combustion either. There are, however, certain substances that take up oxygen from it when hot; if carbon dioxide is mixed with hydrogen and passed through a red-hot tube, carbon monoxide and water are formed; when led over glowing carbon or when heated with phosphorus it is reduced to carbon monoxide. If a burning magnesium ribbon is lowered into carbon dioxide, the oxide of the metal is formed and free carbon is deposited; the same thing happens when sodium or potassium is heated in dry carbon dioxide. The aqueous solution of carbon dioxide reacts shghtly acid; it is supposed that this solution contains a compound H2CO3, of which many salts are known. This acid, carbonic acid, has not yet been isolated in the free state, however, since it gives off gaseous carbon dioxide ("carbonic acid gas") when its solution is boiled or frozen. If its salts (carbonates) are treated with an acitl, no H2CO3 is obtained either, for it breaks up forthwith into water and carbon dioxide. Carbonic acid is a very weak acid; it is 2.51 IXORGANIC CHEMISTRY. [§§ 1S4- liberated from its salts by almost every other acid. By adding hydrochloric acid to a carbonate H" ions are introduced into the liquid and they unite with the CO3" ions to form integral H2CO3 molecules. These, however, break up largely into water and carbon dioxide, the latter of which can only remain in solution up to a certain amount at a constant pressure, so that all in excess of this passes out. As a result the concentration of the H2C,;, are strong electrolytes. A solution of such a salt, therefore, con- tains a large number of COs" ions, part nf which must unite with the H" ions of the water in order to establish the equilibrium be- tween carbonic acid and its ions. The result of this is that other molecules of water must be split up into ions in order to com- pensate the loss of H' ions. This leaves in the liquid a certain number of ( )H' ions, which are not balanced by an equal number of H' ions. The liquid therefore acquires an alkaline reaction. The carbonates of the other metals are insoluble in water; however, the acid carbonates are mostly soluble. Calcium carbo- nate, e.g., dissolves in water containing carbonic acid. The solu- tions of such acid carbonates give off carbon dioxide on merely boiling, howcA-er, and the neutral carbonates are precipitated. In the solid state also the acid carbonates give off carbonic acid gas very readily on warming. Composition of Carbon Dioxide. — In connection with wliat was stated in § 179 it is an important fact that no change of volume occurs when carbon burns in an excess of oxygen: C + 02=C02. 1 vol. 1 vol. When a \'cry concentrated solution of potassium carbonate is elec- trol>zed with high current density at 30°-40°, potassium percarbonate, KiC,* )6, is formed at the anode. In aqueous solution it sets free iodine from KI solution at once, which serves to distinguish it from HjO,, since a dilute solution of the latter liberates iodine only very slowly. 185.] OTHER CARBON COMPOUNDS. 2.')5 Other Carbon Compounds. 185. Cyanogen, (CN)2, can be prepared liy heating mercuric cyanide, Hg(CN)2, or by treating a solution of potassium cj-anide with copper sulpliate solution. It is possible that first cupiic cyanide is formed and that this at once breaks up into cuprous cyanide and cyanogen: 4KCX +2CuS04 = 2K2S04 + Cu2(CN)2+ (CN)2. Cyanogen has a penetrating odor. When hquefied it boils at — 20.7°. It is unaffected l)y high temperatures. It dissolves in Water, but the solution deposits amorphous brown flakes after a while. It burns with a purple-tinged flame according to the equation C2X2 + 202 = 2C02 + X2. The reaction, however, is not trimolecular, the first stage being C2N2 + 02 = 2CO + N2. i.e. a bimolecular process. This was proved by Dixon by determining the velocity of propagation of the explosion of mixtures of cyanogen and oxygen. When explosive gas mixtures are introduced into a long tube and their explosion started at one end (by an electric spark, for example) a flame results, which is propagated through the tube with a definite and measurable velocity. Berthelot called this self-propagating flame the explosion irarc. Dixon ignited ii mixture of 1 vol. cyanogen and 1 vol. oxygen, obtaining after the explosion carbon monoxide and nitrogen; the velocity of the explosion wave was found to be 2728 m. per sec. Thereupon he mixed 1 \-ol. cyanogen with 2 vols, oxygen in one instance and with 1 vol. oxygen and 1 vol. of an indifferent gas in another instance; in both cases tlie velocity of the explosion wave was nearly the same, viz. 2321 m. and 2.{9.S m. per sec. It is plain, therefore, that tlie second volume of oxygen influenced the explosion wave in the same way as the indifferent gas, viz. as a diluent _ The conclusion may be drawn that in the explosion wa\e itself only carbon monoxide and» nitrogen are formed, c\cn in the presence of an excess of oxygen. However, since the tube contains only carbon dioxide and nitrogen after the combustion, it must \>c assumed that the cimibustion of carbon monoxide to carbon dioxide is a secniidary process. ■.")6 INORGANIC CHEMISTRY. [§§ 185- Hydrogen cyanide, HON (prussic acid), is important in inor- ganic chemistry because of the numerous complex salts which it forms. Those of the alkalies are soluble in water and crystallize beautifully; see § 308. The salts of the alkaline earths and mercuric cyanide are also soluble in water, the other salts in- soluble. The Flame. i86. A flame is produced by the burning of a gas; solids, like iron, carbon, etc., burn without a flame. If a flame is observed during the burning of mineral coal, a candle or the like, it is due Fig. 4;i. — Reverse Flame. Fig. 44. — Potassium Chlorate FlaME. to the fact that at that high temperature gaseous decomposition- products are formed, which biu-n. If a gas burns in the air, it is •called a combustible gas and the oxygen of the air is called the sup- porter of the combustion. These expressions in common use are only relative terms; it is possible to light the oxygen and have it biu-n with a flame in a gas which is ordinarily called combustible. This phenomenon is illustrated in a way by the reverse flame. This can be easily obtained with the aid of the apparatus of Fig. 43. A lamp-chimney is fitted with a two-hole cork at its lower end. Through 1S7-] THE FLAME. 257 the narrower hole of the cork a small tube a is inserted for conducting in the gas; through the wider hole a tube h for the admission of air. The chimne}' 5s first removed and the gas coming out of tube a lighted and so regulated as to produce a small flame. Then the chimne)' is replaced; the flame continues to burn quietly, inasmuch as plenty of air is supplied by the wider tube. Thereupon the gas supply is gradu- ally increased and at a certain moment the small flame at the end of a is extinguished and a large pale flame appears at the end of 6; it is air burning in the gas which fills the ehimne}-. This is the reverse flame of air in illuminating-gas. At the same time the excess of gas escaping at the top ignites in the outside air, so that the apparatus presents both a direct and a reverse flame at the same time. That it is reallj' air that burns at the mouth of h is proved by introducing a tiny gas-flame by means of the tube c into the flame of the wide tube 6; the small flame continues to burn. Substances that give up oxygen are capable of burning when sur- round°d by a combustible gas. The experiment can be carried out with potassium chlorate as follows: Illuminating-gas is conducted into a glass cylinder (Fig. 44) and lighted at the top, where the cyUnder is covered by a thin piece of metal with a hole in it. A httle potassium chlorate is then lowered into the flame by means of a deflagrating spoon and heated till oxygen comes off freely. If the bowl is then dipped down in the cylinder, the o.xygen burns with a very luminous flame, which is colored violet-blue by the vaporization of some potassium salt. We saw above (§ 27) that a hydrogen flame continues to burn in chlo- rine with the formation of hydrochloric acid; on the other hand chlorine can also be made to burn in hj'drogen. For this purpose a cyUnder closed at the top is filled with hydrogen and lit at the lower edge. A tube through which chlorine is supplied is then brought in contact with this flame and inserted in the cylinder. The chlorine burns on. 187. A flame ma}- be luminous or non-luminous. It gives light when solid particles are suspended in it. An ordinary gas- flame is luminous because particles of carbon, set free by the com- bustion, are made to glow. On introducing a cold object into the flame they are deposited as soot. The light of the "Welsbach incandescent gas-light is produced bj- the glowing incombustible mantle (§ 291). Such flames gi-\'e a continous spectnun (§ 263). Jlany gases, which yield onlj' gaseous products on burning, gi\'e either a very faint light or none at all, e.g. hydrogen, carbon monoxide, etc. However, when Ii^-ilroiipn burns in oxygen of 20 atmospheres 258 INORGAXIC CHEMISTRY. 18S- pressure, its flame is strongly luminoiis. Other incandescent gases, such as the vapors of certain metals, can render a flame Imuinous even at ordinary pressure, imparting to it a definite color. Colored flames of this sort give a line spectrum (§263). A gas-flame, whose light is due to incandescent particles of car- bon, is made non-luminous by mixing the gas with ah before the combustion. This is the principle of the Bdnsen burner (Fig. 4-i), which is used in all laboratories and Cjuite extensively also, with some variation or other, in heating and cooking apparatuses (gas stovesl. The Bunsen burner consists of a base in which is a tube for supply ing the gas, which escapes from a narrow orifice at a. Here it mixes with air that enters through the lateral holes in c, the proportion of air being regulated by the collar b. This mixture burns Avith a colorless flame when ignited at the top of c. The opinion was originally held that the loss of luminosit}' of the flame is due to the oxjigcn of the air, the latter causing the complete combustion of the carbon particles. As has since been shown, however, the dilution of the burning gas with nitrogen also has a part in it: if illuminating-gas is mixed with two or three times as much nitrogen, the former burns with a colorless fiame. Fig. 45. — Bunsen Burner. Witla the aid of a wire gauze a burning gas mixture can be cooled so low that the combustii:)n cannot propagate itself through the gauze; in other words, the flame does not get through the gauze {Vi!^. 46). If gas is allowed to flow out of a Ik'XSEX burner and a wire gauze is held across the current a short distance from the oriflce, the gas can be lit above the gauze without the flame springing back to the burner. It was by experiments such as these that Davy was led to discover h's miner's xafetij-lamp. As Fig. -17 shows, this consists of an oil-lamp, the flame of which is surrounded by a wire cage. A combustible gas mixture may catch fire ins'de of the lantern, but the fire cannot pass through the gauze to the outside. 188] THK FLAME. 251) i88. The temperature of the flame is much lower than we might Buppose. Since, when hydrogen burns in oxygen, 57.2 l^g.-calories are produced by every IS g. of the mixture, and the specific heat of Fig. 16. — Effect of a AVire G.\uze on a Flame. Bteam is 0.48, this amount of heat ought to raise tlie IS g. steam to a temperature of 57 2 ""-■^ - =6600°. In reaHty the temperatm-e 0.01SX0 4S of the flame does not exceed 2500°. This difference between calculation and observation is due to the fact that on account of dissociation only a partial combination of hy- drogen and oxygen takes place in any part of the flame. The temperature of 6600° could indeed be obtained at anj' point, if the gases united there completely and instantaneouslj''; but this is im- possible, for above 1300° the formation of the com- pound is checked b}'' the opposite process, the dis- sociation of steam. Therefore what occurs must be this: oxygen and hydrogen, when brought to- gether at the aperture, combine and effect a certam rise of temperature; in the same measure as the system in equilibrium (liydrogen, oxygen, steam) cools off, fresh portions of the gases unite. Thek combustion cannot therefore take place at any particular point but must be gi-adual throughout the whole extent of the flame and at any one point the temperature cannot ex- ceed a certain limit, which is determined by the degree of dissocia- tion of the combustion product. Fi'.. 47.— Miner's Safety-lamp. 260 INORGANIC CHEMISTRY. 189- 189. Zones of a luminous flame. Let us take a candle-flame, for example. In the central zone (1 in the diagram Fig. 48) Fig. 48. — Zones of a Lujiinohs Flame. there is no combustion. The stearin of the candle is here con- verted by the heat of the flame into volatile combustible products. In a large candle this can be proved in the manner shown in Fig. 48. The narrow tube conducts off the inflammable gases and they can be lit at the outer end. The hollo wness of a flame can be demonstrated in various ways; in a Bunsen burner, for instance, by placing a match-head in the center, where it does not ignite, or by holding a thin platinum wire across a flame; the wire only g'.ows at the edges of the flame. The dark central zone of the flame is next surrounded by the luminous zone (2). Here the volatilized hydrocarbon is decom- posed with the separation of carbon, because the air supply is insufficient for complete combustion. These carbon particles become incandescent and so make the flame luminous. Finally there is the blue outer zone (3), in which the glowing particles of carbon are burned by direct contact with the air. It radiates very little hght. The amount of solid carbon in a flame which is raised to incandes- cence and hence gives light is very small, as the following calculation shows. The substances in burn'ng illuminating-gas which break up with the liberation of carbon are chiefly benzene and ethylene. The former makes up about 1, the latter about 4, per cent by volume of the gas. If we assume that the benzene is completely broken up and 190.] SILICON. 2()1 the ethylene only half, then the total amount of carbon deposited by 1 liter of burning illuminating-gas is about 54 mg. The volume of the luminous part of a gas flame with a consumption of 150 liters per hour amounts to about 2 c.c. (reduced to 0°), so that the mass of solid 2x54 iiicandescent carbon present in it is only .„„„ =0.1 mg. 1000 SILICON. 190. This element in combination with oxygen is one of the principal constituents of the earth's crust (§ 8). In the free state, however, it does not occur in nature, being found almost exclusively as silica, SiOo, or in the silicates. Sand and the many varieties of quartz are different forms of natural silicon dioxide; the number of silicates is very large. Free silicon is obtained by heating sodium fluosilicate, Na^SiFe, with sodium: NaoSiFg -!- 4Na = 6XaF + Si, or by heating sodium in an atmosphere of silicon tetrafluoride : 4Na+Sir4=4NaF + Si. The sodium fluoride can be removed by water. Another method, which is far easier, is to mix 400 g. almninium filings with 500 g. sulphur and 360 g. sand. This mixture is ignited, whereupon it burns with a large flame. The mass fuses and becomes white-hot. When cooled it consists principally of aluminium sulphide and free silicon. It is then treated ^^■ith dilute hydro- chloric ecid, which decomposes and dissolves the sulphide, leaving the silicon behind. (Kuhne method.) Alloiropic Forms. — The silicon obtained Ijy the two first -named methods is a brown amorphous powder; it can be fused under a layer of molten sodium chloride and obtained crystal- line on cooling. The latter form is best prepared by KtniMo's method. The crystals are regular, black, and of a high lustre. If silicon is heated in the electric furnace, it vaporizes and con- denses again in small globules, mixed with a little gray powder and some siHca. Chemical Properties. — Silicon takes fire only wlien heated in 1'62 INORGAXIC CHEMISTRr. [§§19 1- the air to a very high temperature, burning to silica. It unites with fluorine at ordinary temperatures, the combustion being marked by a glow; combination with chlorine takes place on gently warming. At an elevated temperature silicon combines with nitro- gen and some metals; these silicides have been prepared mainly by MoissAN in his electric furnace. It is indifferent towards sulphuric, nitric and hydrochloric acids. Hydrofluoric acid dissolves it, however, with the evolution of hydrogen. H}'drogen chloride gas reacts with it at a high tem- perature, forming silicon tetrachloride and sUico-chloroform. It dissolves in a hot solution of sodium or potassium hydroxide, pro- ducing hydrogen and a silicate: Si + 2K0H + H2O = KaSiOs + 2H2. Hydrogen Silicide, SiH4. 191. This gas is obtained by adding freshly prepared magnesium silicide to hj^drochloric acid. The magnesium silicide is prepared by heating sand with an excess of magnesium powder, or better by fusing 40 parts of anhydrous magnesium chloride witli a mixture of 35 parts of sodium fluosilicate, 10 of sodium chloride and 20 of sodium. The hydrogen silicide so obtained is mixed with hydro- gen. A purer product results from heating an organic derivative of silicon, tri-ethjd silicoformate : 4SiH (OC2H5) 3 = 3Si (OC2H5) 4 + SiHi. Hydrogen silicide, or silicon tetrahydride, is a gas, which becomes liquid at —1° under a pressure of 100 atmospheres. It has a disagreeable odor. It takes fire in the air; each bubble that escapes from the generator forms a cloudy ring of hydrated silica. If, howeA'er the hydrogen silicide is perfectly pure, it does not ignite spontaneously in the air at ordinary temperatures except under reduced pressure. The spontaneous ignition in the air is caused by the presence of small quantities of other compoimds, probably also composed of silicon and hydrogen. We have therefore in this case phenomena similar to those in the case of hydrogen phosphide (§ 136). Heat decomposes the hydrogen silicide readily into Si and 2H2. It bums in a chlorine atmosphere and is decomposed by an alkali solution according to the equation: SiH4 + 2K0H + H2O = iK-, + KsSiOs. 192] HALOGEN COMPOUNDS OF SILICON. 263 Silico-ethane, SijH,, is formed by the decomposition of magnesium silicide by hydrochloric acid. It is a gas, which can be liquefied below Halogen Compounds of Silicon. 192. Silicon tetrachloride, SiCU, is prepared by heating silicon in a current of chlorine at 300°-310°. It is a colorless liquid with the specific gravity 1.52-11 at 0° and the boiling-point 59.6°. It is instantly decomposed by water, forming hydrochloric acid and hydrated silica. Silico-chloroform, SiCljH, is obtained, together with a large quantity of silicon tetracliloride, on heating silicon in a current of hj'drochloric acid gas (§ 190). From this mixture it is separated by fractional distilla- tion. It is a colorless, strongly smelling compound which fumes in the air, boils at 34°, and is decomposed by water. By the action of dark electrical discharges on a mixture of dry hydrogen and silico-chloroform vapors chlorine-silicon compounds are formed of the order 8in(-'l;,i + 2, e.g., perchloro-siUco-ethnne, SijCls, etc. Silicon tetrafluoride, SiF4, can be obtaineil by warming a mixture of sand and calcium fluoride with concentrated sulphuric acid: 2CaF2 + SiOo -h 2Yl2i?Oi = SiFi + 2CaS04 -l- 2H,0. It is a colorless gas with a very pungent and suffocating odor; it condenses under 9 ntm. pressure or by cooling to —160°. ^Mien perfectly dry, it does not attack glass. .Silicon fluoride is also formed by the action of hydrogen fluoride on silicates; the silica is first set free from them and then attacked in the way just described. Glass-etching (§ 53) depends on tliis action. Bv the repeated treatment of silicates with hydrous hydrofluoric acid all the silicic acid is driven off as silicon fluoride. The bases which were in combination with the silicic acid are left behind in the form of fluorides. They c:m be transformed into sulphates by warming with sulphuric acid and then converted into a form suitable for analysis. We ha\-e here a vei-y useful means of determining the metals present in the silicates. Water decomposes silicon fluoride as follows: 3SiF4 -t- 3H,0 = H.Sil )3 + 2HoSiF6. 2^4 INORGANIC CHEMISTRY. [§§ 192- The compound H2SiF6 is called hydrofluosilicic acid; it is known only in aqueous solution. If the latter is concentrated by evaporation, silicon tetrafluoride escapes but hydrogen fluoride stays in solution. When the concentration corresponds to 13.3% H2SiF6 the vapor contains 2HF to iyil'4; but dilute solutions yield a vapor which contains much more hydrogen fluoride. If, therefore, a concentrated solution of hydrofluosilicic acid is par- tially evaporated, the residual liquid is able to dissolve silica because of the presence of free hydrofluoric acid. On the other hand, a dilute solution, after partial evaporation, leaves a residue, from which silicic acid is deposited, because the excess of silicon tetra- fluoride which it contains is decomposed by water according to the above equation. The decomposition of silicon fluoride by water is usually demon- strated in the following way: The compound is generated in the pre- scribed manner in a flask (Fig. 49), whereupon it is conducted through Fig. 49. — Preparation or Hydrofluosilicic Acid. a doubly-bent glass tube into a cylindrical jar containing a little mer- cury (into which the tube opens) and on top of this some water. Every bubble of gas that rises from the mercury into the water generates in the latter a cloud of silicic acid. If the glass tube opened directly in water, it would soon become stopped up because of this decomposition. 193.] OXYGJiX COMPOUXDS OF SILICON. 26 J The solution of hydrofluosilicic acid reacts acid; it dissohcs metals with the evolution of hydrogen and behaves in all respects like an acid. A hydrate, H2SiF6 + 2H2O, is known in the solid state. It melts at 19°, and is obtained by leading silicon fluoride into con- centrated lu'ilrofluoric acid. Most of the salts of hydrofluosilicic acid arc soluble in water; the potassium salt is difficultly so, how- ever, and the barium salt is insoluble. Hydrofluosilicic acid is used in hardening objects made of gyp- sum (this is due probably to the formation of calcium fluoride) and also in analytical chemistry. Oxygen Compounds of Silicon. 193. Only one such compound is known: silicon dioxide, or silica. SILICA, SiOj. This compound occurs in astonishingly large quantities and in a great number of varieties in the solid crust of the earth. It is found crystallized as rock crystal, quartz (when colored brown, caUed smoky quartz), amethyst (the more beautiful sorts being used for ornament), tridymite, onyx, cat's-eyc, etc. Sand is largely silica; sandstone also belongs here and so does jasper (usually colored red with ferric oxide and having a conchoidal fracture). Opal is an amorphous variety, containing varying amounts of water. Silica can be prepared artificially as an amorphous white powder by heating sihcic acid. Physical Properties. — In the crystallized state silica is very hard and insoluble in water and has a specific gravity of 2.6. It is A-ery difficultly fusible; in the oxy hydrogen flame it softens and passes over into a vitreous modification. When heated strongly this can be drawn out into extremely fine threads that are so tenacious and display so regular a torsion that they are frequently used in sus- pending magnets, etc., in physical instruments. It can be made to boil vigorously in the electric furnace; the vapor condenses in woolly flakes. Quartz that has been fused has a very small coefficient of expansion (about Vi7 of that of platinum); this explains why objects made of it can endure very sudden changes of temperature. They can be heated very hot and then thrust 266 INORGANIC CHEMISTRY. [§§ 193- into cold water at once without cracking. They are attacked only by metallic oxides and at a high temperature. Recently it has become possible to utilize fused (vitreous) quartz for the manu- facture of chemical apparatus. It is interesting that quartz vessels are transparent to ultraviolet rays, which is not the case with glass vessels. Chemical Properties. — ^Especially in the crystallized condition silica is very little acted upon by acids except hydrofluoric acid (§ 193). Fused alkalies dissolve it, forming alkali silicates. It can be reduced by carbon in the electric furnace, carborundum (§ 178) being formed. It is also reduced by heating with magnesium (§ 190). Silicic Acids. 194. When a solution of potassium or sodiiun silicate {water- glass) is treated with hydrochloric acid, a very voluminous, gelat- inous mass separates out; this consists of hydrous silicic acid cor- responding to the general formula Si02aq. After being washed with water and dried in the air it is a fine white amorphous powder of the approximate composition HaSiOs. Freshly precipitated silicic acid is slightly soluble in water, but more so in dilute hydro- chloric acid If, therefore, water-glass is introduced into an excess of hydrochloric acid, the silicic acid stays in solution; it can be sepa- rated from the sodium chloride simultaneously formed, by the fol- lowing process : The solution is put into a piece of parchment tubing, which is tied at both ends, and the whole submerged in pure water, the latter being frequently renewed. It is found that the salt goes through the parchment, but that the silicic acid does not. This process is called dialysis and any arrangement for carrying it out is known as a dialyzcr. Graham found that crystallizable substances in solution (crystalloids) are able to pass through such a membrane, while other substances, which he called colloids, are not. In the latter class are glue, gums, gelatine, albumen — in short, many amorphous substances occurring in the animal and vegetable kingdoms. The silicic acid which separates from the colloidal solution dries in the air to a white amorphous powder, still containing a 195.] SILICIC ACIDS. 267 good deal of water, however. The water can be slowly extracted in a sulphuric acid desiccator. Since silicon tetrachloride is changed to silicic acid by water, just like phosphorus pentachloride to phosphoric acid (§ 145), we can consider the compound as the basis from which the remaining silicic acids are deri-^-ed. The latter can in general be represented by the formula ?reSi(0H)4— nH2U. These polysilicic acids themselves have not been isolated, but many of their salts and double-salts are known, which occur as minerals in nature. The silicates of potassium and sodium are soluble in water, those of the other metals insoluble, as are also most of the double silicates of the alkalies. In the soil hydrous silicates are found whose b."ises are usually lime and alumina. In contact with alkali salts these undergo a double decom- position, an insoluble potassium aluminium silicate, for example, being formed together with chloride of calcium, which is taken off by the under- ground water. This phenomenon is said to be caused by the absorptive power of the soU; it plays an important role in the determination of soil- values. It is this that holds back the potash, an invaluable nutrient, which is furnished to the soil in the form of jiotassium salts and would otherwise be quickly washed off by the rain because of its solubility. The soluble phosphates are "absorbed" by the soil in the same way. This is mainly to be ascribed to the lime they contain, with which iasoluble tri- or dicalcium phosphate is formed; to some extent this ab- sorption may be caused also by basic lime silicates. Silicon Compounds of Other Elements. 195. Silicon sulphide, SiSj, is produced when carbon disulphide vapor is led over a mixture of charcoal and silica at red heat. It forms long, silken needles, which are broken up by water into SiOj aq and hydrogen sulphide. Silicon nitride, SizNj, a white amorphous substance, results from the heating of silicon in an atmosphere of nitrogen. (For metal silicides cf. § 190.) 26S INORGAXIC CHEMISTRY. [§ 196. COLLOIDS. 196. In silicic acid we have become acquainted with a sub- stance that occurs in a special form, viz., as a colloid. A con- siderable number of such substances is now known, and the study of them has latterly been so active and prolific that a brief recapitulation of the principal results is fitting at this point. Graham discovered that in aqueous solution a number of substances, principally amorphous materials, such as the glues, albumen and dextrin, have a very small power of diffusion, quite contrary to most salts. Accordingly he distinguished between colloids and crystalloids. Subsequent investigations served to increase greatly the number of colloids, i.e., substances of small diffusibility. Gradually the view developed that the colloidal condition is not something peculiar to certain com- pounds, but that all sorts of substances, even the crystalloids, can be obtained colloidal by suitable treatment. Hence the colloidal state is now regarded as a general property of matter. Just as we have substances in the solid, liquid and gaseous states, so we can also transform them into the colloidal state. The question that' arises first is: How may this conditioH be brought about? The following methods serve the purpose: In the first place, colloids may be prepared by simply dis- solving certain substances, such as glue, in water. Secondly, they are formed in many cases instead of pre- cipitates, wlicn no ions are present. For example, if hydrogen sulphide is passed into a solution of arsenic trioxide, there results, instead of a precipitate of AS2S3, a yellow liquid containing the arsenic sulphide in colloidal solution. However, if the arsenic trioxide solution is first acidified with a little hydrochloric acid (a highly ionized substance), the AS2S3 separates out as a yellow precipitate. Again, we may take mercuric cyanide, a compound that is hardly ionized at all in aqueous solution. If a solution of it is treated with hydrogen sulphide, which is also a very feebly ionized substance, the mercuric sulphide that is formed is retained in colloidal solution; yet, the usual precipitate can be obtained b}' adding previously a small amount of a strong mineral acid. § 196.] COLLOIDS. '2m A third way of preparing colloids is by dialysis, a process described in connection with silicic acid. In this way hydrosols (see below) of ferric oxide, aluminium oxide and many other substances can be obtained. Ferric oxide hydrosol, for instance, is formed when ferric chloride, FeCls, is dissolved in water, and just a little less ammonia added than would produce a precipitate, and the whole then dialyzed. The ammonium chloride, NH4CI, and hydrochloric acid (resulting from a partial hydrolysis of FeCla in aqueous solution) pass through the membrane, while Fe^.O.-j aq. remains inside in colloidal solution. A fourth method is the comminution, or dusting, of metals imder water. This is accomplished by connecting wires or rods of platinum, gold and other metals with the poles of a 110-volt circuit; if the wires are moved toward each other under water, a small arc is formed when they are a short distance apart, and dark clouds of the metal proceetl out into the liquid from the cathode. The liquid is then filtei'ed; the coarser liits of metal remain on the filter and the filtrate is a clear, dark-colored solu- tion containing the metal as hj'drosol. Metals can often be converted into the colloidal state 1j}' treat- ing a very dilute solution of one of their salts with certain reduc- ing-agents at ordinary temperature. Thus from a very dilute gold chloride solution the colloidal gold can be prepared by the addition of phem-lhydrazine hydrochloride or acetylene. Finally, it is worthy of note that by means of protective colloids many substances can be obtained colloidal when other means fail. For instance, if a silver nitrate solution and a potassium bromide solution, each containing about 1% of gelatine, are mixed together, the silver bromide is not precipitated, but comes out colloidal. " Collargol," a therapeutic preparation,' is a silver colloid, made stable by a protective colloid. Colloids can be divided into two groups, reversible and irreversible. The reversible colloids comprise among other sub- stances the agglutinants, as they are called, — gelatine, agar-agar, albumins, starch, etc. When they are mixed with water the_\- swell up and on being gently warmed form a solution. When cooled they geaiinize, i.e., they congeal to a soft, viscous mass which retains all the solvent water. The solution itself is called 270 ISORGANIC CHEMISTRY. [§ iflg. a hydrosol (or, in case alcohol is the solvent, an aZcosoQ. and the gelatinized mass a hydrogel. When the solvent water is extracted from a reversible colloid by evaporation at a low temperature a hydrogel is at first formed, which still contains a great deal of water. This water is partially lost on exposure to the air. More rapidly in a desiccator, — and its vapor tension does not differ perceptibly from that of pure water. When, however, a certain stage of dehydration is reached the vapor tension begins to diminish. If water is added to the hydrogel before this stage is reached, a hydrogel is again obtained with the same properties as originally. The process of s o 1- and g e 1-formation is thus a reversible one. Other interesting properties are attached to hydrosols. Crystalloid salts, for example, diffuse in them, — even in the con- gealed mass, — almost as easily as in water. If a piece of jellied agar-agar is immersed for some time in a dark blue ammoniacal solution of a copper salt, the agar-agar becomes stained through- out its entire mass. Colloids, on the contrary, do not diffuse. This can be shown by a colloidal solution of Prussian blue, which does not penetrate at all into the agar-agar, as above. The electrical conductance, too, is practically the same for a gel containing crystalloid salts in solution as for an aqueous solu- tion of the same salts at like concentration. Oftentimes large amounts of crystalloid salts can be added to a reversible hydrosol without the formation of gel. This is very different with the second class of colloids, the irreversible colloids, for they are in many cases very sensitive to additions of salts. When the salt is added the irreversible hydrosol begins to appear cloudy, and a precipitate is formed which cannot be reconverted offhand into hydrosol. Irreversible hydrosols can be prepared in the various ways already mentioned. They comprise the colloidal metals, sul- phides, hydrated oxides, etc. Most of them are mobile liquids, in contrast to many reversible colloids, such as glue. The quantity of an electrolyte that is just sufficient to pre- cipitate an irreversible hydrosol is connected with the valence of the electrolyte, the quantity decreasing rapidly with n- creasing valence. The AsoSs hydrosol is just coagulated by § 196.] COLLOIDS. 271 71 mill-equivalents of NaCl per liter, 2.0 of MgCla and 0.39 of AICI3. Certain irreversible colloids are capable of mutually pre- cipitating each other; othei's are not. The hydrosols of ferric oxide and arsenious sulphide give a precipitate when mixed, but the hydrosols of gold and arsenious sulphide mix without precipitation. These phenomena have been shown to be connected with the behavior of the substances toward the electric current. If a solution is introduced into a U-tube supplied with electrodes at the xipper ends and a strong current (say 110 volts) is passed through it, the colloid is seen to separate out and wander either to the anode or to the cathode. At one of the two electrodes an aqueous layer appears, which is entirely free from colloid and is separated sharply from the hydrosol. This convective trans- ference, or electrical endosmose, is by no means to be confused with the ionic migration in electrolytes. For, while in the elec- trolytes there is an electrical opposition between the dissocia- tion products of the dissolved substance, the electrical opposi- tion exists in this case between the colloid and the solvent. In general, mutual precipitation is only possible with colloids whose electrical charges are opposite with respect to that of a common .solvent. The colloidal state must be regardetl as a very fine divi- sion, or distribution, of one substance in another. This follows from the great analogy which exists between suspensions (e.g., clay and water) and colloids. Both can be separated out by centrifuging and both display the Tyndall efTect. This effect may be described as follows: When a beam, of light passes through a body of air that is free from dust it is invisible transversely; the gas is "optically a vacuum." But, so soon as dust particles enter the air, the path of the beam can be followed through the dispersion of the light by the particles. Optically vacuous liquids and optically vacuous solutions of crystalloid salts can also be prepared. But if a beam of light is passed through a hydrosol the path of the beam can be seen. The hydrosol is therefore not an optical vacuum; it must con- tain floating particles, but these are so small that they can not be seen even with the best microscopes. 272 IXORGAXIC CHEMISTRY. [§§ 196- However, Siedextopf and Zsigmoxdy have succeeded in rendering these sub microscopic particles visible with an apparatus that thej' call the ultmmicroscope. In it the hydrosol is illuminated transversely so that the luminous rays do not blind the eye of the observer. The submicroscopic particles bend (dif- fract) the light rays in ah directions, so that with sufficiently intense illumination the light effect produced by each individual particle comes within the range of microscopic visibility and can be separately observed without however revealing its form. Suspensions resemble colloids further in that they ■ exhibit electrical endosmose and can be precipitated b}' the addition of electrolytes. When a liquid is distributed through another in exceedingly small drops we have an emulsion. The most familiar example is milkj an emulsion of butter fat. There is reason for assuming that many reversible colloids are extremeh' fine emulsions; for instance, an emulsion, like a gelatine solution, cannot be coag- ulated by the addition of an electrolyte. The knowledge that in the colloidal state we have to do with a vei'j- fine distribution of one substance in another has led to the introduction of a new set of terms. The substance distributed as a colloid is now generally spoken of as the disperse phase, distributed in the dispersion medium. Further, the words dis- persoids and emidsoids are replacing the word "colloids." It was formerly thought that a sharp distinction must be drawn between colloids and real solutions. Graham spoke of two different worlds of matter. In contrast to the true solu- tions colloids exhibit practically no diffusion, no vapor pressure lowering, no boiling-point elevation or freezing-point depres- sion, — in short, no osmotic phenomena. The researches of recent yeai's have, however, shown that essential differences do not really exist. To begin with, we have come upon many cases of transition between colloidal and real solutions. Furthermore, it was previously observed by Lobry de Bruyn that salt solu- tions can be separated bj- centrifugal force into portions of unlike concentration. More important still, the investigations of EixsTEix, Perrix, Svedberg and others have shown that colloids, just like true solutions, are subject to the osmotic laws. 197.] GERMANIUM. 273 From the molecular-kinetic point of view there is no difference, according to these investigations, between a " dissolved molecule " and a "suspended particle"; consequently a mechanical sus- pension must exert exactly the same osmotic pressure as a " true solution " of the same number of particles per unit volume. The fact that the colloidal solutions display no properties cor- responding to osmotic pressure is simply due to the fact that at the same concentration the number of freely moving particles in solutions is enormousl.y greater than with the colloids, i.e., an individual colloid particle has gigantic dimensions as compared with those of a molecule. This makes the freezing-point lower- ing, etc., so slight that it cannot be measured b}' present exper- imental means. In order to test the applicability of the osmotic laws to colloidal solutions we are therefore forced to employ indirect methods. The methods employed are associated with four phenomena: (1) the translatory and rotatory movements of the particles. (2) diffusion ; (3) the change of concentration under the influence of gravit}-; and (4) the local temporary changes of concentration. The possibilit}- of testing the osmotic laws by such measurements is a result of developing formula for these phenomena that are deduced upon the assumption that the osmotic laws are applic- able. These investigations also serve to corroborate the reality of atoms and molecules, supporting the information gained in many other ways, as has been set forth in § 35. GERMAOTUM. 197. This element is of extremely rare occurrence. It was discovered by \\'iNKLER in an argentiferous mineral, argyrodite, GeSj.4Ag,S, found in Freiberg, in Sa.xony. Germanium forms grayish-white octahedrons with a metallic lustre and a specific gravity of 5.409 at 20° It melts at 900°. At ordinary temperatures it is tmaffected by the air; at rod heat it burns, forming white fumes of germanium oxide, (IcOj. Two series of compounds of this element are known, which arc derived from the oxides Ge() and GeOj; the ons compounds are easily oxidized to the higher form, germanic acid. The hydrogen compounds, GcH, and GeHClj, are known. Germanic chloride, GeCl,, can be prepared directly from the elements. It is broken up by water forming Oe((>H\. 274 IXORGAXIC CHEMISTRY. [§§ 197- Germanium dioxide, GeOz, is produced by heating tlie corresponding hydroxide, or by roasting the element or its sulphide or by treating it with nitric acid. It is a white powder of a specific gravity of 4.703 at 18° and is unaffected by heat. Germanium disulphide, GeS2, separates as a white precipitate when hydrogen sulphide is passed into the solution of germanium dioxide in strong hydrochloric acid. In moist air it decomposes, giving off hydrogen sulphide. It dissolves in alkalies and alkali sulphides to form sulpho-salts. For germanium cf. also § 218. TIN. 198. This metal is not very widely distributed on the earth; in some places, however, it is found in quite large quantities. The principal tin mines of Europe are those in Cornwall; even the Phoenicians obtained tin there. The most important present locali- ties are on the group of islands lying east of Sumatra (Banca, Bil- liton, Sinkop, etc.). There the metal occurs in the form of tin- stone {cassiteriie, Sn02); it is found in quadratic crystals, which are usually colored brown or black by a small amount of iron. In order to extract the metal, the ore is at first roasted, to eliminate any sulphur or arsenic it may contain, and then reduced with car- bon. The tin thus obtained is refined by liquation, i.e. by fusing again at a low temperature and pouring it off from the less fusible alloy of tin with iron and arsenic. It is then melted once more and stirred with a wooden pole (branch of a tree), whereby the oxide still remaining is reduced. The Banca tin is nearly chemic- ally pure. Physical Properties. — ^Tin is a sUvery-white metal, melting at 232.7° and volatilizing between 1450° and 1600°. Sp. g. =7.293 at 13°. It has a crystalline structure which can be made visible by moistening with hydrochloric acid, whereupon peculiar frost- hke etch-figures are produced on the surface (tin-moiree) . When tin is bent, a characteristic crackling sound (cry of tin) is heard, which is probably caused by the grating of the crystal faces on each other. Tin is very malleable and ductile; it can be beaten into very thin leaves {tin-foil) at the ordinary temperature, and at 100° it can be drawn out into wire. At a very low temperature and in contact with an alcoholic pink-salt solution (§ 201), tin passes spontaneously into another modification, gray tin, which has a lower specific gravity, 198.] TIX. , ?75 5.8. Above 20° this form changes back to white tin. If the latter is brought in contact with graj' tin at ordinary temperatures (below + 20°), it turns very slowly into gray tin, falling to powder, probably because of the increase in volume (this phenomenon is called the "tin-disease"). If it is not in contact with the gray modification, the transformation does not take place at all at ordinary temperatures, or at least not for centuries. Evidently there is a transition point of the two forms at 20°, and we are forced to the odd conclusion that, except on warm summer days, tin is in the metastable condition. The reason why tin, even in contact ^\ ith gray modification, passes so slowly into that form at ordinary temperatures is that the velocity of transformation is small in the neighborhood of the transition point; it is accelerated on moving away from that point. When the temperature sinks this acceleration is counteracted, however, by the retardation that all reactions undergo by a lower- ing of temperature. In manjr cases, therefore, there must be a maximum of the \'elocity of transformation, such as we have here at — 4S°; below that temperature the transformation again becomes slower. Ordinary tin crystallizes in the tetragonal system. In addition to t'.ie gray modification there is also a third one, the rhombic modification. The transition point tetragonal<=±rhombic is about 170° This point was determined in a unique way, namely, by measuring the velocity- of flow of the metal under high pressure. For this purpose the solid metal was placed in a cylinder ha\ing a hole in the bottom, and the quantity of metal was mea.sured that was forced out under constant jire^- sure in the unit of time. In general, this quantity increases rapidly with rising temjierature, but with tin it was found to diminish considerably when the temperature reached alniut 200° This may be taken as a proof that the metal has another (third) modification. At 200° tin is so brittle that it can be easily pulverized. Chemical Properlirs. — Tin is unaffected by the air at ordinary temperatures; if heated stroiiKlx', it burns with an intense white light tn tin oxide, SnO-,. Hydrochloric acid dissolves it, forming stannous chloride and h\-drogen. It is also attacked by nitric acid (§ 201). A boiling solution of caustic soda or potash converts it into a stannic acid salt (s t a n n a t e) with the eA'olution of hydro- gen: Sn + 2K0H + H2O = KoSnOa + 2H2. 276 INORGANIC CHEMISTRY. [§§ 199- In the presence of weak acids (acetic acid) and alkalies it is very stable. 199. Uses. — On account of its permanence tin is used as a pro- tecti^-e co-\'ering for metals which are attacked by the air and the above-named agencies. Many kitchen utensils are " tinned." Sheet iron is covered with a layer of tin, to protect it from rusting (§ 279), and is then known as tin-plate, or shcct-tin. This is done by simply dipping the sheet iron, which has been cleaned by hydrochloric or sulphuric acid, in molten tin. Many alloys of tin are in use. Solder consists of tin and lead (in the ratio 2:1 or 1:1 or 1:2), and is harder than either of its components but more easily fusible. The alloys of c o p p e r and tin are called bronzes; their composition varies according to the purpose they serve. At present the bronzes usually contain a little lead and zinc as well. Bronze is hard and tough, can be easily worked and fuses to a mobUe liquid, hence it is particularly suitable for casting. Gun metal contains 90% copper and 10% tin; bell metal 20-25% tin, the rest being copper. Phosphor bronze is prepared by fusing copper with tin phosphide (§ 202). The result- ing mass is remarkably homogeneous and contains 0.25-2.5% phosphorus and 5-15% tin. Its great hardness and firmness render it especially valuable for certain parts of machines (axle-bearings). Silicon bronze contains silicon in place of phosphorus, is very hard and conducts electricity well, hence it is used for making telephone wire. Tin amalgam forms the metallic coating of mirrors. Compounds of Tin. Tin forms two sets of compounds; thej^ correspond to the oxygen compounds, stannous oxide, SnO, and stannic oxide, Sn02. STANNOUS COMPOUKDS. 200. Stannous chloride, SnCl2, is prepared by dissolving tin in hydrochloric acid: Sn-|-2HCl = 2SnCl2-)-H2. It cr\'Stallizes with two molecules of water, which are given off at 100° It is very readily soluble in water (1 part in 0.37 at ordinary temperatures). Anhydrous stannous chloride is white and trans- parent; it melts at 250° and boils at 606°. A little above the 200] STANXOUS COMPOVXDS. 277 boiling-point the vapor density corresponds to the formula Sn2Cl4; above 900°, however, to SnClo. The aqueous solution acts strongly reducing. It absorbs oxy- gen from the air witli the partial formation of basic chloride (a, white powder), if the liquid is not too acidic: SSnCla + H2O + = SnCl4 + 2Sn(0H)Cl. Basic chloride. But if the liquid is strongh- acid the tetrachloride SnCU is also formed in this oxidation. This same basic chloride also results from hydrolytic dissociation, ■when a neutral stannous chloride solution is strongly diluted. SnClj +aq =Sn(OH)Cl + HC1 +aq. The reducing power of stannous chloride is further seen in its action on potassium permanganate, potassium dichromate, cupric chloride, mercuric chloride, etc., all of which are converted into lower stages of oxidation in acid solution. It may be remarked here that, from the ionic point of view, oxidation amounts in many cases to raising an ion to a higher positive potential, and reduction to the reverse. Let us consider, for instance, the reaction between stannous chloride and mercuric chloride. This can be expressed bj- the equation SnCl2 + HgCl2 = SnCl4 + Hg. Stannous chloride is oxidized to stannic chloride; at the same time mercuric chloride is "reduced" to the metal. Written in ions, this equation becomes Sn" + Hg" = Sn-- + Hg; that is, the electrical charge of the mercury ion is taken by the bivalent tin ion, the former losing its electrification. Another example is the action of chlorine on stannous chloride, by which the latter is "oxidized" to stannic chloride: SnCl2 + a2 = SnCl4. The ionic reaction is Sn" + 2CI' + Cl2 = Sn-- + 4Cl'. 278 INORGANIC CHEMISTRY. [§§ 20C- Tin takes up two more positive charges, but this necessitates that the two CI atoms become ions; they thus require two negative charges; but when these are formed two positive charges are ob- tained at the same time. However, the Sn"" and CI' ions unite to form stannic chloride, SnCU, which is a very weak electrolyte (c/. § 201). In the preparation of chlorine, hydrochloric acid is "oxidized " by manganese dioxide : MnOa + 4HC1 = MnCla + 2H2O + CI2, or Mn02 + 4H- + 4C1' = Mn • • + 2C1' + 2H2O + CI2 ; the positive charge of the four H' ions is thus transferred, half to the manganese and the rest serving to discharge two chlorine ions, i.e. to equalize their negative charges. Various double salts of stannous chloride are known, e.g. SnCl2-2KCI; SnCl2-2NH4Cl. Stannous hydroxide, Sn(0H)2 is precipitated when a solution of stannous chloride is treated with soda : SnCl 2 + NazCOa + H2O = Sn(OH) 2 + 2NaCH- CO2. This hydroxide is insoluble in ammonia, but soluble in alkalies; when the latter solution is boiled, tin is deposited and alkali stan- nate, e.g. KoSnOs, formed. The hydroxide is also soluble in acids, thus displaying a basic as well as an acidic nature. Such compounds are able to give hydroxyl ions (Sn" + 20H') on the one hand and hydrogen ions (Sn02" + 2H') on the other. They are termed amphoteric compounds. Stannous oxide is obtained by heating the hydroxide in a cur- rent of carbon dioxide; it is a dark-brown powder, which takes fire in the air, burning to stannic oxide, Sn02. Other salts of stannous oxide than the above-mentioned stannous chloride are also known. The sulphate, for instance, is obtained by dis- solving the hydroxide or the metal in dilute sulphuric acid. It forms a basic salt readily. Stannous sulphide, SnS, is precipitated as an amorphous brown powder when hydrogen sulphide is passed into the solution of stan- nous salts. It is insoluble in potassium sulphide, K2S, but it 201.] STAXXIC COMPOUXDS. 271) dissolves to form a sulpho-stannate wlicn brought in ccmtact with the poly sulphide of ammonium or potassium, K2Sx(x=2-5). SnS + K2J^o = K2SnS3. Stannous sulphide can also be prepared by fusing tin with sul- phur. It then forms a bluish-gray crystalline mass. STANNIC COMPOUNDS. 201. Stannic chloride, SnCU, was prepared as early as 1605. It was named spiriius jumans Lihavii, after its discoverer. It is obtained by the action of chlorine on tin or stannous chloride. Stannic chloride is a liquid which fimies strongly in the air; it boils at 113.9°, and has a specific gravity of 2.234 at 15°. When brought in contact with a little water or on taking up moisture from the air, it goes over into a semi-solid, cr3-stallized mass, SnCU -31120, the so-called iin-buttcr. A fresh solution of stannic chloride is a very poor conduc-tor of electricitj-. However, the con- ductivity increases slowl}' at ordinarj-, faster at higher, tempera- tures; after several days it reaches a maximtmi. In the case of more dilute solutions this maximum is much higher. These facts can be explained by assuming that stannic chloride is but feebly ionized and that it reacts with water in the following way: SnCU + fflzO ?=± Sn(0H)4 + 4IiCl; in other words, that it undergoes hydrol3tic dissociation. It is the liberated hydrochloric acid that causes the conductivit}-. The solution contains tin hydroxide in the colloidal state. The water has thus split up the stannic chloride into a basic hydroxide and an acid. Stannic chloride forms well-cr^-stallized double salts with the alkah chlorides, e.g. SnCl4-2KCl and SnCl4-2NH4Cl. The latter is known as pink salt (because of its color) and is used as a mor- dant in dyeing. Tin tetrachloride also unites with the chlorides of the metalloids to form crystallized substances, e.g. SnCl4-PCl5; SnCU • POCI3 ; SnCU • SCI4, etc. It combines with hydrochloric acid, forming a leafy-crystalline mass, H2SnCl6-6H20, which melts at 9°. Tin fluoride, SnF^, itself is not known, but there is a compound, KjSnFg, which corresponds to potassium fluo-silirate; the salts of hvdro- fluostannir acid are isomoiphous with thn annloirous siliron compounds. 280 INORGANIC CHEMISTRY. [§§ 201- Stannic oxide, Sn02, can be prepared synthetically by heating tin in air. It is an amorphous white powder, insoluble in acids and al- kalies ; the latter, however, dissolve it when fused, forming stannates. Stannic Acid and Metastannic Acid. — The hydroxides corre- sponding to Sn02 have onlj- very weakly basic properties; here the acidic properties are prominent. The normal hydroxide, Sn(OH) 4, is unknown, but there is a hydroxide of the empirical composi- tion H2Sn03(=Sn(OH)4— H2O), corresponding to carbonic acid, H2CO3. Strangely enough this exists in two modifications, which differ from each other both chemically and physically; they are called stannic and metastannic acids. The stannic acid is precipitated when ammonia is added to an aqueous solution of stannic chloride or hydrochloric acid to a potas- sium stannate solution. This precipitate reacts acid when moist and is soluble in concentrated hydrochloric and nitric acids, as well as in alkalies. It gradually changes into metastannic acid. Metastannic acid is generally prepared by treating tin with strong nitric acid; it is then formed in a vigorous reaction as a. dense white powder. Metastannic acid is insoluble in sodium hydroxide, but nevertheless unites with it to form sodium metastan- nate ; this is dissolved by water, although with difficulty, but is insolu- ble in the caustic soda solution. When boiled with hydrochloric acid, metastannic acid goes over into a chloride, which is insoluble. in the concentrated acid but soluble in water. This solution does not contain the ordinary tin chloride, but another one, m e t a-t i n chloride, having, however, the same composition, SnCl4. It is distinguished from the ordinary stannic chloride by giving a jollow coloration with stannous chloride solution; the solution of the ordi- nary chloride does not do this till after some time, during which the metachloride is formed in it. Stannic acid and the corresponding chloride thus pass over into the meta-compounds spontaneously; on the other hand, metastan- nic acid can be converted into the ordinary tin compounds by boil- ing it for some time or fusing it with a caustic alkali. The difference between stannic and metastannic acids was pointed out by Berzelius as early as^tlie beginning of ttie nineteenth century. They are both colloids. The salts of metastannic acid have in general a very com- plicated composition, similar to the polysilicates (§ 195), for which reason metastannic acid is regarded as a polymer of the ordinary stannic acid, i.e., that its molecule is represented by (HjSnOj)^, stannic acid itself being H.SnOj. 203.] LEAD. 281 Of the salts of stannic acid, the sodium stannate, Na2Sn03 + SHoO, is especially well known. It comes on the market under the name of " preparing-salt " and is used as a mordant in dyeing. It is made by fusing tin-stone with caustic soda and crystallizes in hexagonal crystals, which are more soluble in cold than in warm water. Purple of Cassius is obtained when a mixture of the hydrosols of tin dioxide and gold is precipitated by adding some such electrolyte as am- monium chloride. This mode of formation proves that the substance is not a compound of the two components, as wn.s formerly believed, but only a mixed gel. 202. Stannic sulphide, SnS2, falls out as a yellow amorphous powder, when hydrogen sulphide is passed into the acid solution of a stannic compound. It can be synthesized by heating tin amalgam with sulphur and ammonium chloride, being thus obtained in the form of transparent golden lea^•C's and known as nurum musivuin, or mosaic gold; it is used for gilding. Stannic sulphide is a sulpho- anhydride; the corresponding sulpho-acid, HoSnSu, is not known in the free state, but exists in the form of salts. Sodium sulphostannate, Na2SnS3 + 2H20, crystallizes in color- less octahedrons. "When its solution is treated with an acid, stan- nic sulphide is precipitated. Tin phosphide serves, as was stated above, for the manufacture of phosphor bronze. Of the various compounds of tin and phosphorus, the best known is the compound SugP. It forms a coarsely crystalline mass, which melts at 170°- LEAD. 203. Among the lead ores the most important is galrnilc (PbS); it occurs in isometric crystals (cubes) of a graphitic color. ()ther ores are nriissitr (PbCO.3), crocoite (PbCr()4), iridfenite (Ph^loOi), etc. For the extraction of the metal galenite is used almost exclu- sively. This is roastpfl to convert the sulphide partially into oxide, and partially into sulphate: PbS + 30=PbO + S()2; PbS + 2()2 = PbS()4. In roasting care is taken that a considerable portion of the ore remains as sulphide. On farther heating, the latter reacts with the oxygen compounds in the following way: 2PbO-|-PbS=3Pb + S()o; and l'l)S()4 + PbS=2Pb + 2SO:;. 282 IXZRGAA'IC CHEMISTRY. [§§203- Physical Properties. — Lead is a soft ductile metal of a bluish color. On exposure to the air it loses its lustre rapidly, becoming coated with a very thin layer of oxide. It has a specific gravity of 11.254, melts at 327°, and boils at 1525°- Chemical Properties. — The thin coating formed by the oxide on the brilhant surface of the metal protects the lead from further attack by the air. If, however, it is prepared in a A'ery finely divided state, e.g., by heating lead tartrate or citrate in the absence of air, it takes fire in the air even at ordinary temperatures. (Other metals can be reduced in a similar way to a fine state of division, whereupon they ignite spontaneously in the air. A substance which exhibits this phenomenon is called a pyrophorus.) When lead is melted, it becomes coated with red oxide of lead; by constantly removing the latter, the lead can be entirely oxidized. A compact mass is unaffected by sulphuric or hydrochloric acid, but, when finely divided, it reacts to form the corresponding salts. Nitric acid easily dissolves it to form the nitrate. Acetic acid and various vegetable acids attack it; since all lead salts are very piosonous and very serious effects result from chronic poisoning with insignificant but successive amounts, it is not admissible to use tin containing lead in tin-plating vessels for use in the kitchen. Zinc and iron precipitate the metal from solutions. A piece of zinc becomes covered with a dendritic crystalline mass (" lead- tree ")• This reaction can be expressed by: Zn + Pb--=Zn-+Pb, i.e. zinc is changed into the ionic condition, and the lead ions are discharged. How it comes about that one metal thus assumes the electrical charge of another may be explained by a hypothesis of Nernst. His supposition is that every metal on coming in con- tact with water or a solution tends to send positive ions into it. This emission of ions continues until the positive charge acquired by the solution and the negative charge created on the metal balance by their mutual attraction the tension (called the electrolytic solution-tension) with which the ions are driven into the solution. This tension differs considerably for different metals; for zinc it is much greater than for lead. When, therefore, a strip of zinc is dipped in a lead solution it forces zinc ions into the solution and the zinc thus becomes much more negatively charged than would 204.] OXlDhS OF LEAD. 283 a piece of lead by the omissidu of lead idns. The lead ions are therefore attracted by the zinc and discharged, i.e. lead is precipi- tated from the solution. This jmxess stops only when all the lead of the solution has been replaced by zinc. Distilled water, from whi('h the air has been entirely remo^•cd by boiling, has no effect on lead, but the simultaneous action of air and water produce lead hydroxide, which is somewhat soluble in water. This hydroxide is conxerted into insoluble basic carbonate by carbonic acid. From a hygienic standpoint these properties of lead are of vast impor- tance, because drinking-water is almost universally conducted through pipes made of lead or material containing lead ("compo-pipes")- The absorption of lead from such pipes !>>' water and the continuation of the process depends in a large measure on the proportion of salt in the water. As a rule, the less of salts it contains, the more lead it takes up. Rain-water, which is aliimst entirely free from solid matter, but contains ox}gen, carbon dioxide and traces of ammonia, is, therefore, most likely to dissolve lead. The lead eave-troughs, etc., which were once extensively used, should, therefore, be rejected, in case the rain- water is used for drinking Well-water usually contains acid calcium carbonate and gypsum; as a result, the lead pipes soon become coated with an insoluble layer of lead sulphate and basic carbonate (as well as calcium carbonate), so that after a while the lead can no longer be absorbed by the water. Lead is used for many purpnsos. not only in the elemental con- dition, but also in the form - weight of calcium. The next question is, whether 40 is really tlic 21)2 INORGANIC CHEMISTRY. {%% 209- atomic weight of calcium or a multiple or submultiple of the atomic weight. Now the molecular weight of calcium sulphate must be 40 + 32 + 4 X 16 = 136, no matter whether 40 is the relative weight of one or of more than one calcium atom. The molecular heat is therefore 136 XO.2 =27.2. The atomic heat of sulphur in compounds is about 5.4 and that of oxygen about 4.0; consequently the molecular heat of the SO, group in its solid compounds is 5.4+4X4.0 = 21.4. For the atomic heat of calcium we- have the remainder, 27.2 — 21.4=5.8. It therefore follows that the formula of calcium sulphate must be CaSO,, which means that 40 is the atomic weight of calcium; for, if the atomic weight were a multiple or submultiple of this number, we should have found for the atomic heat of the metal a number much farther from the average atomic heat of the elements, 6.4, than 5.8. The value of the atomic weight calculated from Neumann's. law therefore serves merely to decide what multiple of the equiva- lent weight must be taken ; for this purpose the number so obtained is sufficiently accurate. 210. 3. The law of Mitscherlich. The crystal form of com- poimds hairing analogous chemical composition is the same: or, in other words, compounds of analogous chemical composition are isomorphous. The compounds KCl, KI, KBr, e.g. are analogous in composition; they all crystallize in cubes. H2KPO4, H2KASO4, H2(NH4)P04 also have an analogoiis composition and 2,11 crystallize in the tetragonal system. The analogous compour.ds KCIO4 and KMn04 both crystallize rhombic. If two compounds have been proved to be isomorphous, it is very probable that their composition is analogous, whereupon the atomic weight is readily found. Let us, for example, take the case of manganese, supposing its atomic weight to be unltnown; now potassium permanganate is isomorphous 'wdth potassium per- chlorate, which latter is known to have the formula KCIO4. Analysis has shown the formula of potassium permanganate to be KMnx04, X being unknown, for 39 parts (by weight) of potassium (1 atom) are combined with 64 parts of oxygen (4 atoms) and 55 parts of manganese (x atoms). From its isomorphism with KCIO4 it follows that its formula must be KMn04 (i.e. x = l), hence 55 is the atomic weight of manganese. 211.] DETERMIXATIO.X OF EiJl'IVALENT WEIGHTS. 2!t3 In determining the atomic weight of zinc we could use the iso- morphism of the crystallized sulphates of magnesium and zinc. The formula of the former is MgS04 + 7H20. On the basis of the analysis of zinc sulphate and the isomorphism mentioned we ha^•e the formula ZnS04 + 7H20, from which the atomic weight is ob- tained in the same way as above. The law of isomorphism was discovered as early as 1819. Since at that time the law of Avogadro received little attention and the determination of the specific lieat was in many cases impossible, the phenomena of iso- morphism were the most important means of getting information regarding the value of the atomic weight. Subsequently its importance for this pur- pose lessened, mainly because simpler means were found, but also because it proved to be very difficult in many cases to decide whether two sub- stances are isomorphous. Moreover it was found that certain substances of entirely different composition are isomorphous. A very delicate test for isomorphism is the fact that a supersaturated solution can be made to crystallize, not only by an extremely small amount of the dissolved substance itself ("sowing," or "inoculation"), but by bodies that are isomorphous with it. Experimental Determination of Equivalent Weights. 211, In the methods a!)0\'<-' descrihrd the (juestion is one of determining •which multiple of the equivalent weight is the atomic weight. In order to establish the atomic weights with accuracy the equivalent weights must be determined with the greatest possible precision. The solution of this problem, which is one of fundamental importance, since all the numerical relationships of chemical reactions are based on the atomic weights, has been the object of numerous investigations in the preceding century and to-day it is still only partially accomplished. The first atomic weight table dates from Dalton in 1805. The figures given in it were scarcely more than rough approximations. Berzelius (1779-1848) in the first and second decades of the century determined a long series of equivalent numbers, after having been first obliged in most cases to work out reliable analytical methods. The atomic weights at which he arrived were in general use for many years and really differ from the more accurate ones now employed by hardly more than r fraction of a per cent. Exceedingly accurate "atomic weight determinations" were undertaken by Stas (1813-1891). The ten atomic weights determined by him, viz. those of .'Vg, CI, Br, I, K, Na, Li, S, Pb, and N, are in most cases accurate to within a few units in the second decimal place. The researches called for most exhausting and persistent labors during a long period of years. 2!:I4 IXORGAXIC CHEMISTRY. [§§ Lll- In the last decade atomic weight determinations have been carried out on a scale of much greater refinement by Morlby, Richards, Guye and others and the accuracy of the values has been extended another decimal place, so that now not a few of the atomic weights are established with certainty to within a few units in the third decimal place. In determining atomic weights either purely chemical or physico- chemical methods may be employed. Both have been greatly perfected in these latter investigations and they will now be described in a few parar graphs. As for the purely chemical methods, there are four conditions which are essential to an accurate determination of an atomic weight: (o) A suitable substance must be found which can be prepared perfectly pure. (6) This compound must contain in addition to the element under study only elements of accurately known atomic weight, (c) The valence of the elements in this compound must be well defined. It is not permissible, for example, that the substance be a mixture of two stages of oxidation, (d) The compound selected must be adapted to an exact analysis, or else its exact synthesis from the weighed elements must be possible. Notwithstanding the simplicity and legitimacy of these demands it is often difficult to satisfy them. The preparation of a compound in the pure state is among the most difficult of operations, if by purity we mean the reduction of the impurities to a 10 ~* part of the whole. It was formerly beUe\'ed that this could be readity accomplished by recrystallization, but now we know that every substance that separates out in a solid phase has u. tendency to retain upon its surface or in its interior a part of the other substance contained in the phase out of which the solid separated. All precipitates or crystals from aqueous solutions contain water that is not in chemical combination. Even the splendid glistening silver crystals that are obtained in the electrolysis of a silver nitrate solution and are apparently perfectly dry and pure contain not only water but silver nitrate as well. Silver chloride, precipitated from a solution of sodium chloride by silver nitrate, may have included traces of NaCl, AgXOa or NaXOj, even after a thorough washing. Potassium chlorate, though much less soluble than potassium chloride, contains nevertheless 0.027% of the latter after repeated recrystallizations. One of the most troiiblesome sources of error in all quantitative researches is the unsuspected presence of hygroscooically held water, since it is not at all easy to detect by chemical tests and causes no essential change in the external appearance of the substance containing it. The analysis of a substance resolves itself in most cases into a separation of its components in the form of other compounds and weighings of the latter. For example, in order to determine the silver content of silver nitrate the metal is thrown down as silver chloride and the latter is weighed, whereupon the quantity of silver can be calculated from the known silver content of the chloride. The analyst generally finds it also necessary to convert one com; ound into anothe" tiuanlitativelv. '"he modern in esti 'a- 212.] DETERMIXATIOX OF EQUIVALEXT WEIGHTS. _".),') tions of atomic weights have also taught us that tliis is often a very difricult problem. Among other sources of error in this connection are the solu- bility of the so-called "insoluble" substances and the solubility of glass. It has long been known that substances like silver chloride, barium sul- phate, etc., are not strictly insoluble; but their solubility has first received proper attention in connection with the recent atomic weight determinations. In working with glass \'es.sels it is impossible to a\oid silicic acid as an impurity. Recognition of this fact has led to the u.se of vessels of (luartz or, better still, of platinum, which has proved to be an important refinement of method. 212. Physico-chemical methods have found application in the determination of the volume weight of gases. One of the most fruitful of modern physical concepts is that of the ideal gas, whose expansion at constant pressure or pressure increase at constant volume both have a coefficient for a tempera- ture change of one degree of exactly 1/273.08. Moreover, the ideal gas is in strict accord with Boyle's law. A gram molecule of such a gas at ti^ and 760 mm. Hg pressure would occupy -i volume of 22.412 I. However, the actual gases are more compressible and expansible than the ideal gas; hydrogen and helium are the only ones that are less compressible. For this reason 22.412 1. of an actual gas at 0° and 700 mm. Hg contains a little more than one gram molecule. If we let \ + X represent the number of gram molecules of an actual gas which are contained in 22.412 1., the molecular weight of the gas becomes 22.412 G il/ = - 1 -HA where G is the weight of a liter of the gas under normal conditions. The establishment of the atomic weight of a gas thus resolves itself into the accurate determination of the magnitudes G and k. The methods for ascertaining the exact weight of a giAcn ^-olume of a gas have undergone important improvements in recent years. The agreement of the values found by the different investigators is within ±0.0001. While formerly the gases were weighed in huge globes, some containing as much as 21 I., later investigators have been able to reduce this volume to l)etvveen one liter and half a liter, or even less. Nevertheless the concordance between the ^•arious series of determinations was improved, because the corrections for the small globes were much less. An additional correction was a|)plied for the contraction of a globe on evacuation, due to the external pressure of the atmosphere reducing the volume slightly; the buoyancy effect of the air i.s somewhat less for an evacuated globe than for one filled with gas. In order to remove completely the layer of air that has been condensed on the inner surface of the globe, it is necessary to evacuate the latter repeatedly to as low a v-acuum as possible and to fill it witli the ga.s whose density is to be determined, great care being taken meanwhile to exclude the air. Furthermore, the purification of the gases to be weighed is much better 296 INORGAXIC CHEMISTRY. [§§ 212- accompUshed by first liquefying them and then removing the impurities by fractional distillation at a low temperature. The determination ot the quantity X can be accomplished in four different ways, which are found described in the larger physics manuals. It should be noted, however, that they have not yet attained the exactness that characterizes the methods of determining G. In the cases of the less easily condensed gases, like H, N, O, and CI, however, very accurate determinations have already been made. From an experimental standpoint these physico-chemical methods have a decided superiority over the purely chemical methods in that physica' measurements only are carried out after the gas has been obtaii ;^d pure. Vll the uncertainties that are involved in chemical transformations are thus avoided ; and upon such transformations every purely chemical determi- nation of an atomic weight is based. THE PERIODIC SYSTEM OF THE ELEMENTS. 213. In studying the elements v?hich we have considered so far, ■we have found that they can be arranged into groups of elements according to their valence, the elements of each group sh.'.ving great similarity in the types of their compounds. The physical and chemical properties of the elements of such a group are found to change progressively as the atomic weight increases. The question now arises whether all elements can be thus arranged into groups; the reply is afBrmative. In the course of the last century there was no lack of attempts to arrange the elements into groups of similar elements. Doebereiner called atten- tion to a simple relation between the atomic weights of kindred elements as early as 1817, and in 1829 he presented the doctrine of trimls. i.e. he showed that there are different groups of three elements each, which have a great similarity among themselves and a constant difference in the atomic weights, e.g. CI, Br, I; Ca, Sr, Ba, etc. In the year 1865 ths law of octaves was proposed by Newlands, he having discovered that, if the elements are arranged according to increasing atomic weight, after an interval of seven elements an element follows which has properties analogous to those of the first, i.e. the first, eighth, fifteenth, etc., are similar. In 1869 I\1exdei.eeff and Lothar Meyer almost sinaultaneously reached con- clusions which are comprehended by the term "periodic system." If we arrange the elements according to increasing atomic weight, thus: H 1 Li 7 Be 9.1 B 11 C 12 N 14 16 F 19 Na23 Mg24 Al 27.1 Si 28.4 P 31.0 S 32 CI 35.4 213,] THE PERIODIC SYSTEM OF THE ELEMENTS. 207 we see that there is a (jnuluul variation in the properties of elements in a horizontal line; after fluorine, h(iAve\er, a small increase in the atomic weight in-\-oh-es a siaklcn change of properties. Moreover those elements %\hich are in the same \crti(al column show great similarity, as we saw abo^c in the cases of carbon and sUicon, nitrogen and phosphorus, etc. This regular change makes itbclf evident in the valence toward oxygen, which rises from one (with Li and Xa) to two (Be, Mg)^ hrcc (B, Al), four (C, Si), five (N, P), six (S) and seven (CI in CI2O7). fhe valence toward hydrogen or a halogen increases, however, from one (Li) to four (C) and then falls again to one (F). A similar regular change is to be observed with reference to the physical properties, e.g. specific gravity and atomic volume. Na Sp. gravity 0.97 At. volume 24 By atomic volume we understand the atomic weight divided by the wwlglit of the unit volume (based on water of 4° as 1); it is therefore the number of cubic centimeters occupied by a gram-atom. Here we observe an increa.se of the specific gravity up to alu- minium, then again a decrea.se to chlorine, whUe the atomic vol- mne, on the other hand, decreases from the beginning of the series to aluminium and then increases. This steady change of the same physical properties is also observed in the compounds of the above elements. For the oxides, e.g. we have: 1.75 Al 2.67 Si P (red) 2.49 2.14 S a (liq.> 2.06 1.33 14 10 11 14 16 27 Sp. gravity. At. volume. 2.S MeO 3.7 A1..03 4.0 SiOj 2.6 P2O5 2.7 SO3 1.9 0,0, ? 22 00 25 45 55 S2 ? Aloreover, if we viTite down a series of elements according to. increasing atomic weight, beginning with another univalent metal, we discover irregularities of exacth' the same sort as the above. The following series may ser^'e as an example of this: Ac rd In Sn Sb Te I Atomic wt 107. S 112.4 115 119.0 120.2 127.6 127.0 Sp.gravity 10.5 S.O 7.4 7.2 0.7 0.2 4.9 Heio also we find the same gradual rise of valence from silver, which is univalent, to sejitivalent iodine, the pr()!j;ressi\(' tran^i- 298 INORGANIC CHEMISTRY. [§ 213. tion from metal to metalloid and a continuous decrease in specific gravity. But more; if we put this last row under the first two: H 1 Li 7 Be 9.1 B 11 C 12 N 14 16 F 19 Na 23.1 Mg 24.4 Al 27.1 Si 28.4 P 31.0 S 32.1 CI 35.5 Ag 107.9 Cd 112.4 In 115 Sn 119.0 Sb 120.2 Te 127.6 I 127.0, it is apparent that the elements in the same vertical columns belong to a group. This has been demonstrated for the last four columns in preceding chapters; it will be proved for the others later on. In the light of these facts, we are led to the conclusion that the physical and chemical properties of the elements are functions of their atomic weights; and when we consider the series beginning with lithium and sodium, and note that in each instance, after a difference of about 16, there follows another element with corre- sponding properties, we are led to the supposition that these properties are periodic functions of the atomic weights. By a function we understand in general a dependent relation between two or more magnitudes, of such a sort that, when one changes, the other does likewise. In the equations y=a±x; y=ax; y=x^, etc., 2/ is a function of .r. A periodic function requires that the same value appear for one magnitude in regular intervals as the other magnitude steadily increases. An example of this kind is presented by the gonio- metric functions, as y =sin x, etc., for every time x increases by 2:r, y comes to have the same value again. If we desire to substantiate the conclusion just stated, we shall have to in^-estigate first the length of each period, in other words, determine how many elements inter\-ene in the table, according to increasing atomic weight between two T\'ith analogous properties. It has already been shown that for the elements as far as chlo- rine, a period always includes seven elements. After chlorine comes potassium (39), which thus falls into the coluimi under sodium. The following elements, K39.2 Ca40.1 Sc44.1 Ti48.1 V51.2 Cr52.1 MnSS.Os correspond very well with the preceding series, Na23.1 Mg24.4 AI27.1 Si28.4 P31.0 S32.1 C135.5, §213.] THE PERI<>D/C SYSTEM OF THE ELEMENTS. 299 at least so far as the ^'al('lK■c and the form of the compound are concerned (AI3O3 and S(-203, Ti02 and Si02, Iv2Cr()4 and K2S(J4, KMnOi and KCIO4), although the similarity of these elements in other respects is not vor}- marked. The elements following manganese, viz., Fe 55.9, Co 59.0, Ni 5S.7, however, do not fit in at all under K, Ca, Sc; but if we pass these by there follows another series of se^"en elements, which corresponds to the one beginning with potassium: Cu63.6 Zn65.4 Ga 70 Ge72.5 As 75.0 Se 79.2 Br 80.0. "We therefore reach the conclusion that, after the first two periods of seven elements ending with chlorine must come one of seventeen elements (two of seven each, and three elements placed at the side), if the elements in the same vertical column are to correspond in their properties. This large period of se\entoen elements can, therefore, be arranged under the preceding .^mall period of seven elements in the following way: SMALL PERIOD. Xa23.1 :\rg24.4 .U27.1 Si2S.4 P31.0 S32.1 C13.5.5 L.\RGE PERIOD. K39.2 Ca40.1 Sc44.1 Ti48.1 V51.2 Cr.52.1 Mn.j.lO Fer).).9 Co59.0 Xi58.7 Cu63.6 Zn65.4 Ga70 Ga72..5 As7.").0 Sc79 2 BrSO.O. In order to arrange in periods the elements whose atomic weights exceed eighty, it is again ncccssar}- to assume large periods, and, moreover, to leave several places vacant. In this manner we arrive at the scheme known as Mendkleeff's table (see p. 301). As to the jjosition of hydrogen in this table opiiiions are divided. JIendeleeff placed this element in the first group, above lithium; its chemical properties indicate A\ithout doubt that it belongs with these metals. On the other hand, Orue Massdn has ])resciited arguments for placing it at the head of group MI, as is done in this table. These arguments are as follows: (1) The molecule of hydrogen contains two atoms, as docs a halogen molecule, while the molecule of an alkali metal consists of one atom. (2) The very low boiling-point of liydrDgcn indicates a similarity to the halogens; moreover the boihng-points of the alkah metals fall with 300 IXORGAXIC CHEMISTRY. [§213. increasing atomic weight. (3) The difference between the atomic weights of the elements of a horizontal series is, on the average, 3. By placing hydrogen in group ^'II it differs by 3 from the next element, helium; but it is then also in good agreement with fluorine, for the mean difference in atomic weight between the successive elements of a column is 16. The difference F — H =18, while Li -H =6, i.e., there is no analogy in the latter case. (4) Liquid and solid hydrogen have no metallic properties. (5) The most important argument for placing hydrogen m the first group is based on its relation to the acids, which may be regarded as salts of hydrogen. But in organic compounds chlorine can replace hydrogen without essentially altering the nature of the substance. Tliis "organic" argument thus offsets the ''inorganic'' one, based on the analogy of acids and salts. .\j3 to the elements discovered in the atmosphere since 1894, viz., helium (at. ^^t. 4), weo;/ ('20,\ argon (40 , krypton (81.6) and xenon (128), it is clear that they form a natural group, for their properties display great analogy (see § 110). Since they are not able to form compounds vAth other elements, they can be regarded as nullivalent. In that case their group could find a place after the eighth, or before the first, group (compare the table, page 294), thus forming a bridge, or transition, from the strongest electro-negative, to the strongest electro-positive, elements. However, it must be noted that argon with an atomic weight of 40 precedes potassium with an atomic weight of 39. As may be seen from Plate I, their atomic volumes fit into Lothar Meyee's curve very well. Group "S'lII, as has been said, owes its origin to the setting at the side of the elements included in it, for by this means the corresponding elements of groups I-VII could be brought under each other. It will thus be of importance to the system, if the nine elements of this group display so much analogy to each other that the grouping of them together appears actually justified. Now this is really the case, as is seen from the following study of their properties : 1. All these elements are of a gray color and difficultly fusi- ble; indeed, osmium is one of the hardest of all metals to fuse (2500°) ; iridium melts at 1950°, wrought iron at 1500°, etc. The melting-point of iron is higher than that of cobalt, and the latter higher than that of nickel. A similar fall of this constant is found with ruthenium, rhodium and palladium, and also with osmium, iridium and platinum. 2. Their atomic volumes are small in comparison with those § -'13.] THE PERIODIC SYSTEM OF THE ELEMENTS. 301 o 05 (>1 Oi ■ CO (N zj at -5 'i IS 1— t M > o 00 CO CD O 2 IN lO 1 s fa i o o IN o CO 1 2 Eh 1— ( i 00 o fa s l-H 3 1 1 s (—1 1> go ^ -^ ■-ri fa ^ t— f CO CO CO CJ CM ^ 1 1 o l-H eg 1 1 > O O CD O CO ■71 o - ci ^ o 1 o 1 i 1 ^ CO D >■ •* l-H 2; O .— 1 CO fa ■ -r TH [^ § 5' — — X 1 o o 1 ^' S i o o CO CI CO CC TH ■ CD O l-H t~ S 1 ^ Tti c^ CO IN 1— i a s" o l-H l-H m 1—1 C"! -t^ ci fa >-' H 1 1— i I— 1 CO fen CQ 1 K CD IN IN ^ CO o to 1 o lo lO 00 ■* 00 CO 2 erf <: l-H 1 IN 00 1 IN 05 CO l-H m 1 g 1 1 o 1 c •- c 5 * S » ^. "^"^ >Cd; Snle in water than K„SiFn, the solu- bility of KjEsF,, must also be greater than that of K^SiF,. Properties of germanium discovered by Winkler. 1. At. wt. =72.5. 2. Sp. g. =5.409 at 20°. 3. At. vol. = 13.3. Allvyl compounds were obtained. Ge(C2H5), boils at 160° and its specific gra\itj' is a little less than 1. GeO, lacks entirely the basic prop- erties which are found to a lim- ited extent in SnC)... The specific gra\ity of GeO„ is 4.703 at 18°. GeOj is easily reduced to the metai by heating with carbon or in hydrogen. G&S^ dissoI\-es readily in XH|.SH. GeCl,, is liquid, boils at 86°, and has a specific gra\'ity of 1.887. 10. K,SiF„ is almost insolulile: K/icF, (lis.solves in 34 parts of boiling Mater. 308 i:f ORGANIC CHEMISTRY. \B 218- Use of the Periodic System in Correcting Atomic Weights. 218. In the group of the platinum metals the atomic weights were previously determined by Berzelius and Fremy to be as follows: Os = 199, Pt 198, Ir 197. In this order the metals named do not, however, fit into the system, for osmium should be the first of the three on account of its analogy with iron and ruthenium, i.e. it should have the smallest atomic weight of the three. On the other hand, platinum, which is more akin to palladium, ought to have the highest atomic weight. A painstaking investigation by Seubeet showed that in reality osmium has the atomic weight 191, iridium 193.0, and platinum 194.8, which order is in harmony vith the system. 219. Graphic representation. — The fact that an arrangement of the elements according to increasing atomic weight also makes the gradual change of the physical properties apparent can be seen most clearly from a graphic representation (as proposed by LoTHAR Meyer), in which the atomic weights are the abscissas and the atomic \'olumes the ordinates (see Plate I at end of book). The first thing we notice in the curve is the regular rise and fall of atomic volumes. At the beginning of each period the atomic volume is at a maximum; it reaches a minimum half way through the period (in the large periods at group VIII) and then increases again. In the descending portions are the ductile, on the ascending the brittle, elements. On the ascending portions and at the maxima are, further, the gaseous ani the easily fusible elements; on the descending and at the minima the difficultly and •ser}- difficultly fusible. On the descending portions are the electro-positive, on the ascending the electro-negative, elements. The peiiodicity of the elements thus becomes very evident. 220. The discovery of the periodic system once more thrust into the foreground of interest one of tlie oldest of questions, viz., that concerning the uititij of matter. The striking connection between all the properties of the elements and their atomic weights leads unavoidably to the assumption of a primordial or ground substance. As to the nature of this substance no evidence is at present obtainable. Peout, in 1815, regarded hydrogen as such a sub- stance. He observed that the (then accepted) atomic weights of many elements, based on hydrogen as unity, are whole numbers. Later, very accurate atomic weight determinations, particularly those of St-a.s, which 221.] THE PERIODIC Si\STEM OF THE ELEMENTS. 30» ■were undertaken with the purpose of testing this hypothesis, demonstratfl with certainty, however, that it was untenable. Cf. S 267. 221. The periodic system ot the elements is one of the most important discoveries in the field of inorganii- chemistry; it can never lose its impor- tance, though it is gradually becoming more c\idciit that the sj"-tem in its present form represents the relations of the elements to each other merely in an approximate way and is only a crude first attein[)t at a real system. There are indeed serious objections to the periodic sys- t e m. These objections concern, in the first place, the positions which certain elements occupy in the system, and which agree very poorly with their properties. This is the case, e.g., with gold and copper, which, indeed, show some analogy with hthium and sodium in their OMS-compounds, but otherwise differ decidedly from the latter elements. The same is true of the metals of the cerium group (Ce, La, Nd, Pr, etc.). The analogy with the other elements of the same group is feeble. If the reply be made that these elements are still too Uttle known because of their rarity and the great expenditure of energy and pains required for their investigation, it must be protested that the same is also true of those better known, such as cerium, and that if these elements should disclose themselves as complexes of several (which is not improbable), it is very doubtful whether there would be a place in the system for elements with approximately those properties which we are at present acquainted with in lanthanum, neodymium, and praseodymium, imless more than one element were put in a place, as has been proposed. A second objection concerns the inability to fit all elements into the system. This is particularly the case with tellurium. Its physical and chemical properties put it without doubt in the sulphur group, and here there is a space for it, if its atomic weight were only about 125, or at least smaller than that of iodine (126.92). Nevertheless repeated and careful investigations have fixed its atomic weight at 127. .5. The same difficulty presents itself with cobalt and nickel. According to their atomic weights the four elements Fe, Co, Ni, Cu must be arranged as follows: Fe, 55.8.5, Ni 58.68, Co 5.S.97, Cu 63.57. But the order which best corresponds with their properties is the one given first, cobalt belonging more strictly with iron, nickel with copper. When we recall that the newly discovered elements of the argon group fit into the system very satisfactorily and that radium, too, finds a place in it, we have Uttle reason to beUeve that the usefulness of the Periodic System is in any way exhausted. 310 IXORGAXIC CHEMISTRY [§§ 222- LITHIUM AND SODIUM. Lithium. 222. This metal is not found free in nature; in combination, however, it is very widely distributed, although always in small amounts. Man)' mineral waters contain it. It occurs chiefly in the silicate lepidolile, or lithia-mica, also as the phosphate in triphylite and in company with alu- minium, sodium, and fluorine in amblygoniic. Finally, lithium is met with in the ashes of certain plants, such as tobacco, indicating that it is also contained in the soil. With the aid of the spectroscope it can be detected in very many minerals. Lithium can be obtained from lepidolite by the following very simple process : The mineral is fused and then poured into cold water, whereupon it becomes very brittle and its silicates are brought into such a condition that they can be decomposed by hydrochloric acid. The finely powdered mass is boiled with hydrochloric acid and the metals Ca, Al, Mg, etc., precipitated by soda from the resulting solution (after filtering off the silica), lithium and the other alkali metals remaining in solution. By evaporation a salt mixture is obtained from which the lithium chloride can be isolated by extraction with alcohol, the insoluble chlorides of sodium and potassium remaining behind. Sletallic Uthium is prepared by electrolysis of the fused chloride or a concentrated solution of this salt in pyridine. Xext to solid hydrogen, it is the lightest of all solid substances, its specific gravity being only 0.59, so that it floats on coal-oil. It is silvery-white, but tarnishes very rapidly in moist sir. Jlelting-point 180°. When heated in the air it burns with an intense white light to the oxide; at ordinary temperatures it is not so readily oxidized as sodium and potassium. It decomposes water with the evolution of hydrogen-, the heat generated is sufficient, however, to melt the metal. Lithium oxide, Li20, and hydroxide, LiOH. The former is obtained by heating the nitrate strongly. It dissolves in water slowly, forming the hydroxide. The latter is a \vhite, crystalline substance, similar to caustic soda; it dissolves in water, producing a strongly alkaline solution. Lithium chloride, LiCl, crystallizes anhydrous in regular octahedra; below 0° it takes up two molecules of water of crystallization however. It dissolves very easily in water and deliquesces in moist air. Lithium carbonate, LijCOj, unlike the carbonates of the other alkalies, is difficultly soluble in water (100 parts of water at 13° take up 0.769 part) ; hence it can be precipitated from the concentrated solution of the chloride bv ammonium carbonate. 223.] LITHIUM AND SODIUM. 311 Lithium phosphate, LijPOi, is likewise very sparingly soluble in wnter (1 part in 2539 parts of water), although the phosphates of the other alkalies are freely soluble. The formation of this salt serves as ;i. test for lithium. The lithium spectrum consistp of two red bands, one of wmch in par- ticular is easy to recognize. Sodium. 223. Sodium occurs in nature in enormous quantities and is very widely diffused. It is a constituent of countless silicates and, as a result of rock decay, gets into the soil, wlience it enters the plants and finally reaches the animal organism. Tlie nitrate is known as Chili saltpetre, the chloride as roek-sult or halite, the carbonate as soda : the cryolite (ice-stone) of Greenland is a sodium aluminium fluoride. Common salt, XaCl, constitutes tlie main part of the saline matter in sea-water. Certain bodies of water such as the Dead Sea of Palestine, and the Great Salt Lake in N'orth America, are almost saturated solutions of common salt. The metal was first obtained by Daa v in 1.S07 by the elec- trolysis of molten sodium hydroxide. CiAY-Li'Ss.vc and Th£nard got it b)' Iieating sodium h}-droxide with powdered iron to white heat. The first named method is the one now generally employed for its commercial manufacture, inasmucli as electric power can be obtained quite cheaply. For this purpose sodium hydroxide is heated a little above its melting- point. The sodium formed at the cathode is kept away from the anode by an iron net. At the anode hydroxyl groups are liberated, which yield water and oxygen. The latter escapes but the water dissolves in "the molten mass and comes in contact with the sodium at the cathode, causing half of it to be changed back to sodium hydroxide, while hydrogen is -evolved. As a result the maximum yield of metal for a given <|\umtity of electricity is only 50%. If the temperature gets too high during the electrolysis, sodium dissolves in the molten mass and is oxidized at the anode. Sodium is silvery-white, melts at 95.6° and boils at 900°, turning at the latter temperature to a colorless vapor. At ordinary temperatures it is very soft, so that it can be readily cut with a knife. It can also be easily pressed through a small hole, coming out in the form of wire. Sp. g. 13.5° at =0.9735. Sodium, like the other alkaU metals (Li, K, Rb, Cs), dissolves in liquid ammonia. If one of these metals is introduced into liciuid ammonia, the bright surface of th:; metal becomes tarnished with an indigo-blue color, which soon turns to a pretty metallic red. The metal then liquefies and forms a bronze-colored solution, which is deep blue at greater dilution. If 312 IXORGANIC CHEMISTRY. [§ 223- the pressure of the ammonia over the solution is diminished or the tem- perature raised, the solution gives off ammonia and deposits copper-red, crystaUine masses. When there is no longer a Uquid solution these masses also lose ammonia and the metal is left behind in the crystalline form. Ruff has demonstrated that we do not have here a case of compounds being formed between the metal and the ammonia, but that the copper-red masses are mixtures of metal and saturated hquid solution; for the solution can be pressed out, leaving a compact piece of metal. The molecule of sodium contains only one atom, as is proved by the depression of the freezing-point of its solution in tin. A great many metals have this same property. In moist air the bright surface of a freshly cut piece tarnishes rapidly, but in air that has been dried with phosphorus pentoxide it keeps its metallic lustre for days. Sodium can be heated in the air to melting and even still higher without catching fire. It ignites only when heated strongly, whereupon it burns with a very bright yellow light (especially in an atmosphere of oxygen). With, water it generates hydrogen, sodium hydroxide being also formed. If it is held firmly in one place during this process (e.g. by laying it on a piece of filter-paper floating on water, or upon ice), the hydrogen takes fire because of the locahzation of the heat produc- tion. Sodium finds extensive use in the laboratory and in the arts. Because of its strong reducing-power it is often used to obtain the elements from their oxides; magnesium and aluminium were formerly obtained thus. In organic chemistry it is also frequently employed for \-arious purposes. OXIDES AND HYDROXIDES OF SODIUM. 224. On burning sodium in dry oxygen a mixture of two oxides^ Na20 and Na202, results. Sodium oxide, Na20, is obtained pure by the partial and slow oxidation of sodium with ox}'gen under reduced pressure and removal of the excess of metal by distillation in a vacuum. The oxide dissolves slightly in the metal and after distilling off the latter the oxide is left in the crystalline form. It is white; it dissolves in water, forming sodium hydroxide, NaOH, and giving off much heat. The peroxide, Na202, is obtained by heating sodium in a current of oxygen till no more oxygen is absorbed. With S mols. •224] OXIDES A\D HYDROXIDES OF SODIUM. 313 water it forms a liydrate, Nii2<>2 + 8Ho(). Since it yields hydrogen peroxide with dikite acids and is a vigorous oxidizing-agent it is manufactured commercially. Sodium hydroxide, NaOH, caustic soda, is formed, together with metallic sodium, \\-hen sodium monoxide is reduced in a cur- rent of hydrogen. The ordinary method of preparing caustic soda consists in boiling soda with slaked lime: ^"a2C03 + Ca(OH)2 =2XaOH + CaC03; or 2Na- + CO3" + Ca- • + 20H' = 2(\a- + OH') + CaCOg. As the solubilit}^ product (§ 73) of CaCOa molecules is very small, the ions Ca" and Ci)/' must unite and the carbonate of lime sinks to the bottom. In order to make the decomposition of sodium carbonate complete, a slight excess of slaked lime is added. NoA-er- theless, the solution does not contain an appreciable quantity of calcium h^'droxide. The reason of this is clear: In the solution there are a large number of OH-ions: as a result the number of Ca-ions can only be very small, for the value of the solubility product of calcium hydroxide is reached with even a veiy low concentration of the latter ions. Sodium hydro.xide is now manufactured on a large scale by the electrolysis of concentrated brine. This method yields an almost chemically pure hydroxide and it dominates the market with users of high-grade caustic. Three types of electrolytic processes are in operation: the diaphragm type, the amalgaviation type, and the bell type. In the first type the cathode and anode compartments are separated by a diaphragm. In the Griesheim process, a suc- cessful lepresentative of this type, the cathode is of iron and the anode ferrous-ferric oxide, Fe304. that has been fused at 2000°- 3000° and cast into plates. (This makes an anode unaffected by chlorine.) In a freshly charged bath containing only chloride solu- tion the current is carried mainly by sodium and chlorine ions; but as fast as the sodium is liberated at the cathode and reacts with water to form hydroxide and free hydrogen, the ions of the hydroxide participate in the transport of the current. The sodium atoms are reliberated and again react with water to form hydrogen, while the discharged hydroxyl ions yield watc^r and oxygen. As a net result of the electrolysis we have, so to speak, an intentional decomposition of alkali chloride accom- 314 INORGAXIC CHEMISTRY. [§§224- panied in an increasing measure by an unintended decomposition of water. On this account the chloride electrolysis cannot be carried to completion; in practice the process is interrupted as soon as the alkali hydroxide concentration gets up to 8%. The caustic cathode liquid is then replaced by fresh brine and the former is evaporated in vacuum pans to a concentration of 50%, whereupon the undecomposed chloride separates out and is returned to the electrolytic cell. The diaphragm process most favorabl}' known in America is the TowNSEND process. The compartment containing the (graphite) anodes occupies the center of the cell and on each side i_; a diaphragm of thin asbestos. The cathodes, of woven wire, ;est closel_v against the outside of the diaphragm and the cathode compartment is filled with warm oil, that is kept in lively cir- culation by the escaping hydrogen. The freshly formed caustic liquor trickles down the side of the cathode and is drawn off from beneath the oil. The constant removal of the hj^droxide enables the electrolysis to be carried jjractically to completion and the yield approaches the theoretical. A high current densitj^ is employed. The Castner process is the most extensively used of the amalgamation type. Its cell is divided into three compartments, the two outside ones containing brine and the carbon anodes, while the middle one contains the caustic liquor. A layer of mercury covers the bottom of the whole cell. In the brine com- partments the mercury acts as cathode, taking up the sodium to form amalgam. The amalgamated mercury is transferred to the middle compartment where "it is decomposed by water to form sodium hydroxide. This caustic solution is drawn off and fresh water introduced at a regulated rate. On evaporation a very pure sodium hydroxide results. The efficiency of the proc- ess is enhanced by conserving the electrical energy liberated in the decomposition of the amalgam. In one successful bell process the anode consists of some carbon supported in a bell which is suspended in the brine and has an exit tube at the top for piping off the chlorine. The cathode is a cylindrical piece of sheet iron encircling the bell. For critical discussions of the relative merits of these and 225.] SALTS OF SODIUM. 315 competitive proeesses .the reader must refer to the technical journals or the most recent works on industrial chemistry. The hydroxide is obtainetl by evaporation, whereupon it is generally cast into sticks for the market. It is radiate-crys- talline and \'ery hygroscopic. It dissolves in water with the evolution of considerable heat. Sodium hydroxide is a very strong base; it is used in thi' arts for numerous purposes, among others the manufacture of soap. SALTS OF SODIUM. 225. The sodium salts are df great industrial importance; many of them are prepared in enormous quantities. The starting- point for their manufacture is usually cDunnon salt. Sodium chloride, NaCl, common khU, is found in large masses as rock-salt, e.g. at Stassfurt and Reichenhall and in Galicia, where it }s dug out by miners. Farther, large, amounts are obtained from sea water and the water of salt wells. Three methods are employed to remove salt. In sufficiently warm countries (e.g. ^lediterranean coast, central New York State) the brine is led into flat basins of very large surface area ("salterns," or "salt covers "). In these the water is removed by solar evaporation and the salt crystallizes out. Any gypsum that may be present separates out first, whereupon the brine passes to further basins and yields pure salt. Later on, the other salts separate out; these are sometimes worked up commercially. In countries with a cold climate (e.g. on the shores of the White Sea) the water is allowed to freeze in flat basins. The ice that forms is free from salt so that the remaining liquid is more con- centrated. In countries of the temperate zone the sea water is con- centrated by letting it evaporate spontaneously from a greatly enlarged surface. This is done by the "graduation" process (Fig. TiO). Bundles of fagots are piled up together in a "rick," above which a trough with small outlet-holes runs from end to end. The brine is pumped up into the trough and trickles down from along the entire length of the latter upon the brush; in this way the surface of the salt solution is greatly enlarged and tlic evapdrafion is mailc 31G INORGANIC CHEMISTRY. [§225 much more rapid. A very concentrated .brine flows out at the bottom. The salt is obtained from this concentrated solution by boiling ("salt-boiling"). Common salt is almost equally soluble in hot and cold water, hence it does not crystallize out on cooling but falls out at the same rate as the saturated brine evaporates, even while Fig. 50. — Gradu.\tion Process. boiling hot. The salt obtained from the first crj-stallization is of course impure, containing small amounts of magnesium salts, which render it hygroscopic. In order to purify it, it Ls redissolved in water and again precipitated by evaporation. Chemicallj' pure sodium chloride is obtained by passing hydro- chloric acid gas into a saturated solution of the salt or treating the solution with the concentrated acid. The sodium chloride is deposited because it is less soluble in hydrochloric acid than in water (§ 205). Common salt crystallizes in cubes; when the solution evaporates I -'-"••] SALTS OF SODIUM. 317 slowly the well-knowTi hollow four-sided pyramids, or hopper- crystals (Fig. 51), are formed. Sp. g.=2.16. :M.-pt.=776°. 100 parts of water dissolve 36 parts NaCl at 0°. 39 parts at 100° ; a saturated solution contains about 26*^0 of salt. The crystals frequently enclose some of the mother-liquor; for this reason they decrepitate on heating. On cool- ing below —10° a saturated solution depos- its crystals of the composition NaCl-l-H20, which lose their water at 0°. Chemically pure sodium chloride is not hygroscopic. It is insoluble in absolute alcohol. Fig. 51.— Hori'Eii- CRYSTAL OF S.VLT Sodium bromide, NaBr, and sodium iodide, Xal, are more soluble in water than the chloride. From hot solutions they crystallize in anhydrous cubes, below 30° in monoclinic crystals with 2 mols. 11,0. Sodium bromide is difficultly soluble, sodium iodide easih' soluble, in alcohol. Sodium thiosulphate, Na2S203oH20, spoken of sometimes as sodium hyposulphite or " hypo>" is employe'stallization. After the black ash has been leached out as far as practicable, it is cast aside as " tank waste." The most valuable con.stituent of the latter is calcium sulphide; it is conserved by treating the 320 INORGANIC CHEMISTRY. [§226. wayte with water and carbon dioxide to form hydrogen sulphide, which is then oxidized to sulphur and the latter to sulphuric acid. The process in its entirety is thus represented by the following equations: 2XaCl + H2S04= Na2S04H-2HCl; Na2S04 + 2C = 2CO2 + NaaS; Na2S + CaCOs = CaS + NazCOg ; CO2 + H2O + CaS = CaCOs + H2S ; H2S + 2O2 = H2S04; or, summed up: 2XaCl + C02 + H20 = 2HCl + Na2C03. The process is noted for its high efficiency, since all the by-products are worked up. Nevertheless, this process, which for a long period of >ears practically controlled the industrial market, is now almost wholly superseded by the other one and a few years more will probably see its entire abandonment. (2) The ammonia-soda process ofSoLVAY. This process, which originally presented numerous technical difficulties, is now so perfected that about ninety-five per cent of the total soda production is by the Solvay process. The chemistry of the Solvay process is very simple. Ammonia and carbon dioxide are led alternatel}- into a cold concentrated salt solution under pressure. The following reaction then takes place: NaCl + (NH4) HCO3 = NaHCOs + NH4CI. The acid sodium carbonate (" bicarbonate ") so formed sepa- rates out, inasmuch as it is very difficultly soluble in the cold concentrated ammonium chloride solution. It is broken up, on heating, into soda and carbon dioxide, the latter of which is carried back to be used over. The ammonium chloride solution is di.stilled with lime, whereby ammonia is recovered. The process as a whole may be represented by the following equations : § 226.] SODIUM CARBONATE. 321 2i\aCH-2NH3 + 2CO. + 2H,.0 = 2NH4Cl + 2NaHC03; 2NaHC03=H20 + C02 + Na2C03; 2NH4Cl + CaO=2NH3 + H20 + CaCl2; CaC03 = CaO + C02; or, summed up: 2NaCl + CaC03 = NaaCOa + CaGla. Ill the S(iL\Ai process there is formed together with the soda an equivalent amount of calcium chloride, for which there is only a limited market (chiefly for refrigeration; cf. § 2.58), so that one \'aluable con- stituent of salt, the chlorine, is largely lost. Attempts to substitute magnesia for the lime so as to be able to utilize the resulting magnesium chloride for the reco\ery of hydrochloric acid or chlorine have not been commercially successful. Some soda is also manufactured by carbonating the electrolyt- ically prepared sodium hydnixide. Sodium carbonate crystallizes at ordinarj- temperatures with ten molecules of water of crystallization in large transparent monoclinic <'r}-stals, which soon turn white and dull from loss of water (efflorcsi cncc). They melt at 60° in their own water; on continued warming the hydrate Xa2C"(>3 + 2H2-s occur when a salt of a weak liase and a strong acid or a ' salt of a weak acid and a strong liase is formed. In case both acid and base are weak the hydrolysis will be all the greater. WTiich reaction the solution of such a salt will give depends on the relative strengths of the acid and base. Soda is used commercially on a large scale, particularly in the soap and glass industries. It is the " washing-soda " of the household. 322 INORGANIC CHEMISTRY. [§§226- Acid sodium carbonate, XaHCOs (bicarbonate of soda), is obtained as a primary product in the Solvay process. It dis- solves in 10-11 parts of water at room temperature and reacts alkaline. On being genth- warmed it breaks up into carbon dioxide, water and soda; this decomposition occurs even on warm- ing the aqueous solution, and when a current of air is passed through the concentrated solution at ordinary temperatures car- bon dioxide escajx's. It is used extensively in baking-powders and is the salcratus of commerce. Sodium, silicate (sodium water-glass) is prepared, among other ways, by fusing sand with Glauber's salt and charcoal. This 3'ields a \'itreous mass, which is dissolved by boiling water. The concentrated solution has the consistency of glue. It finds use as a fixative in calico printing, as well as for impregnating inflam- mable textiles like theater decorations, etc.; it is also used for " filling " soaps. The sulphides of sodium correspond to those of potassium and are prepared in the same way (see § 2ol). Sodium borate : cf. Borax (§ 283) . POTASSIUM. 227. Compounds of potassium occur in nature very extensively but not in such large quantities as those of sodium. Potassium exists principally in the silicates, especially feldspar and mica. Upon the deca}- of these minerals it Ls carried into the soil and thence into the plants, to which potassium compounds are indis- pensable. Potassium salts are also found in sea-water. The largest source of them, howc^-er, is the Stassfurt " Abraum , salts " (§ 44), mainly double-salts of potassium and magnesium, such as camallite, JigCk'-KCl-eKaO, Jccinite, :\IgS04 ■ KCl • 3H2O, etc. The large amiounts of potassium in the feldspars makes its recovery from them a very enticing problem. The metal was first obtained by D.wy by the electrolysis of molten caustic potash. One of the commercial methods is to ignite a mixture of carbonate of potash and powdered charcoal (preferably charred acid potassium tartrate). The extraction of the metal is thus analogous to that of sodium; in the preparation of potassium, however, potassium carbom/I, Co(OK)6, may be 228. ] OXYGEX COMPOUNDS OF POTASSIUM. 323 formed under certain circumstances, a substance which acquires €xplosi^'e properties on exposure to the air. Potassium has a silvery-white metahic lustre and is almost as soft as wax at ordinary temperatures. Sp. g. =0.S75 at 13°. It melts at G2.5° and boils at about 720°, forming a green vapor. The mirror-like surface of the metal immediately becomes dull in the air; when heated in the air it burns with an intense \'iolet light. Water i^ decomposed by it with great vigor, the heat evolved being sufficient to ignite the escaping hydrogen and drive the piece of potassium around on the water. Oxygen Compounds of Potassium. 228. Potassium oxide, K2O, is formed by oxidizing potas- sium by the method described in § 224. It is a white substance, which unites with water to form the hydroxide with the evolution of much heat. Potassium peroxide, KO2, is produced together with the mon- oxide on burning potassium in the air. It is dark yellow. In con- tact with water it yields potassium hydroxide, hydrogen peroxide and free oxygen. Potassium hydroxide, KOH, results from the action of potas- sium on water and is generally prepared in the same manner as sodium hydroxide, viz., by treating potassium carbonate solution with milk of lime, Ca(0H)2. It can also be obtained by heating saltpetre with powdered copper (forming copper oxide and potas- sium oxide), and adding water; the copper oxide can be removed by filtration. The hydroxide usually comes on the market in sticks. The commercial product (" caustic potash ") is obtained chiefly by the first method and usually contains a little sulphate, chloride, etc., besides the carbonate which is gradually formed by the action of atmospheric carbon dioxide. It can be purified by treating with strong alcohol, which dissolves only the hydroxide; after filtering, the alcoholic solution is evaporated in a silver dish. Caustic soda is also purified in this way. Potassium hydroxide is one of the strongest bases. In the solid state it greedily absorbs water and carbon dioxide from the air and finally deliquesces to a concentrated solution of potassium carbonate, whUc sodium hydroxide under these conditions turns to a solid 32-1 INORGANIC CHEMISTRY. [§§ 22S- white mass of soda. For this reason caustic potash is a much more suitable absorptive agent for carbon dioxide in analyses than caustic soda, for the use of the latter might easily cause a stopping up of the apparatus. Caustic potash is used especially in the manufacture of soft soaps. Potassium Salts. 229. Potassium chloride, KCl, occurs at Stassfurt in the min- eral sylvitc. It crystallizes in cubes and melts at 730°. It is easily volatilized at elevated temperatures. 100 parts H2O dissolve 25.5 parts KCl at 0°, 57 parts at 100° Like its sodium analogue, potassium chloride is precipitated from its saturated solution by hydrochloric acid. It unites with many salts to form double salts. Potassium bromide, KBr, is important therapeuticall}-. It is prepared lay mixing bromine with a potassium hydroxide solution, the bromide and bromate being formed; the bromate is reduced by heating the salty product with powdered charcoal. Potassium bromide crystallizes in cubes and dissolves readily in water. Potassium iodide, KI, also of medicinal value, can be prepared like the bromide and also in the following manner: Iodine and iron filings are mixed together under water, whereupon a solu- tion of the compound Feslij is formed; on treating this with a potash solution the oxide Fe304 is precipitated, carbon dioxide escapes and potassium iodide is left in solution; the salt is then obtained by filtration and e-\-aporation. It crystallizes in cubes and is ^-ery soluble in water: 1 part H2O dissolves 1.27S parts KI at 0°. On exposure to light or the air the cr3-stals gradually turn yellow because of the separation of iodine. It was remarked in § 4G that iodine, though only slightly soluble in water, dissolves to a much greater extent when the water contains potassium iodide. This is due to the formation of I3' ions in the latter case. That the iodine has entered into combination may be concluded in the first place from the fact that the addition of iodine to an aqueous solution of potassium iodide does not cause a further depression of the freezing-point; the number of molecules is thu3 unchanged, or, in other words, iodine has combined with potassium iodide: in the second place, from the fact that carbon disulphide takes up nearly all the iodine from an aqueous solution of the latter when it is shaken with the solution, but only a small proportion when the same operation is 229. POTASSIUM SALT$. ;jj,-j performed with a solution of iodine in a dilute aqueous solution of potassium iodide. The distribution ratio for iodine between water and carbon disulphide is 1:410. If, therefore, we divide the con- centration of the iodine in carbon disulphide by 410 we obtain the concentration of the free iodine in the potassium iodide solution. Subtracting this from the total concentration of the iodine in this solution, wc have the amount of combined iodine. It is found that IKI, or, rather, II', has taken up II2. I3' ions are thus formed in the solution. Xevertheless, a solution of iodine and potassium iodide in water behaves in many cases as if all the iodine were present in the free state, e.g. when it is titrated with sodium thiosulphate. This must be explained by the supposition that in the liquid we have the equilibrium: I3'^l' + l2. If the free iodine is removed, the equilibrium is disturbed; a new portion of I3' must therefore split up, and so on till it is entirely consumed. Potassium fluoride, KF, possesses a peculiar property, which is lacking with the other halogen compounds of potassium: it com- bines with hj'drofluoric acid eagerly, forming the double halide KFHF. Potassium cyanide, KCX (often also written KCy), is manufac- tured on a large scale by fusing yellow prussiate of potash with potash : K4re(CN)6 + K2CO3 = 5CNK + KCXO + CO2 + Fe. The cyanate of potassium KCXO is reduced by the iron to potassium cyanide also. It is -\'ery soluble in water, forming a strongljr alkalme solution. On account of its great tendency to form double- salts, it is employed in electro-metallurgy. It is also used in extracting gold from its ores (§ 248). Potassium chlorate, KCIO3, can be obtained by passing chlo- rine into a h o t solution of caustic potash (§ 56). It is now pre- pared almost exclusively by the electrolysis of a hot solution of sodium chloride. If in the Oriesheim process (§ 224) for the manufacture of caustic alkali the clrctrolysis is continued after a certain amount of hydroxide has formed, the oxygen liberated 326 INORGANIC CHEMISTRY. [§§ 229- at the anode oxidizes the sodium chloride to chlorate, NaClOs. The latter is converted into the potassium chlorate by treatment with potassium chloride. The advantage of this method is that sodium chlorate is much more soluble and does not retard the electrolytic process b}- separating from the solution, as does potassium chlorate. Potassium chlorate is a well-crystallized salt, which is used for the preparation of oxygen (§ 9) ; furthermore, it is used in the manufacture of matches and fireworks, and also medicinally as a remedy for sore throat. On being heated it gives up oxj^gen, part of the salt being at the same time converted into potassium perchlorate, KCIO4. The last-named salt is difficultly soluble in water. It is sometimes found in crude Chili saltpetre, rendering the latter unfit for use in fertilizing various cultivated plants. Potassium sulphate, K2SO4, is obtained by the action of sul- phuric acid on potassium chloride. It crystallizes in beautiful, lustrous rhombic prisms and dissolves with some difficulty in cold water (1 part in 10 parts H2O at room temperature). It is used principally for the preparation of potash according to the Le Blanc method. Acid potassium sulphate, KHSO4, is very soluble in water; it melts at 200°, losing water and going over into potas- sium pyrosulphate, IV2S2O7. The latter breaks up into potassium sulphate and sulphur trioxide on heating. Potassium nitrate, KWOs, is widely distributed in nature, — although usually found only in small amounts, — for it is formed wherever nitrogeneous organic bodies decay in contact with potas- sium compounds. This is the basis of an artificial method of preparing saltpetre, which method was formerly much used. Another process of manufacture depends on the double decom- position of Chili saltpetre with potassium chloride, which is obtained in large quantities at Stassfurt: KCl + NaNOs = KNO3 + NaCl. For this purpose hot-saturated solutions of the two salts are brought together. As sodium chloride is much less soluble than saltpetre 231.] POTASSIUM SALTS. 327 at the temperature of boiling water, it is the first to crystallize out on evaporation, but when the solution is cooled the saltpetre comes out first, for it is much less soluble than sodium chloride in cold water. Potash saltpetre crystallizes in anhydrous prisms, either rhom- bohedral or rhombic according to the temperature. In the neigh- borhood of the melting-point the former is the stable variety at ordinary temperatures the latter. The location of the transi- tion point of the two forms has not yet been determined. 100 parts H2O dissolve 13.3 parts KNO3 at 0°, 247 at 100°. It melts at 338°; farther heating breaks it up into potassium nitrite and oxygen. It has a cooling taste. 230. Potash saltpetre is consumed in large quantities in the manufacture of gunpowder. This is a mixture of sulphur, charcoal and potash saltpetre, the proportions varying in different countries, but being in most cases 75% KNO3, 10% !^ and 15% charcoal. 231. Potassium phosphates. — The three potassium salts of phosphoric acid are known. The}- are very soluble in water. Potassium carbonate, K2CO3, potash. — This salt ^^•as formerly obtained solely from wood-ashes, these being soaked in water and the strained liquor evaporated. At present it is manufactured extensively from potassium chloride after the Le ISlanc process. Another source of potash is the molasses of the beet-sugar fac- tories, that contains the potassium salts in which the sugar-beet is rich. At the Xeustassfurt salt mine it is made from potassium chloride by a patented process as follows: Magnesium carbonate, MgCOj -31120, is suspended in a solution of potassium chloride, and carbon dioxide is led in, whereupon the following reaction takes place: 3 JIgCOj . 3H2O + 2KC1 + CO2 = MgCU + L'MgCOg ■ KHCO3 ■ RJO. The potassium magnesium carbonate separates out and is broken up by heating to a temperature not exceeding 80° into magnesium car- bonate and potash. The former salt is again obtained with three mole- cules of water of crystallization, which form is the only one suited for the above reaction. Potassium carbonate is a white powder, which ticliquesccs in the air and is very soluble in water (1.12 parts K^^'Oa in 1 part 328 INORGAXIC CHEMISTRY. [§§ 231- H2O at 20°)"; the solution has a strong alkaUne reaction. The salt melts at 838°. It is used in the preparation of soft soaps and hard glass (potash-glass). Potassium silicate, potassium water-glass, is formed when sand is fused with potash. Different salts of this sort are described. They dissoh'c in water, forming a thick, mucUaginous mass which on drying turns to a \-itreous, and finally opaque, product. Potassium water-glass is used for the same purposes as sodium water-glass. Sulphides of Potassium. Potassium monosulphide, K2S, is prepared by reducing^ potassium sulphate with charcoal. It dissolves in water very readily and crystallizes out with five molecules of water. It absorbs oxygen from the air, going over into the thio- sulphate and hydroxide: 2K2S + H2O + 2O2 = K2S2O3 + 2K0H. Acids react with it, liberating hydrogen sulphide. Potassium hydrosulphide, KSH, is obtained by saturating a caustic potash solution with hydrogen sulphide: KOH + H2S=KSH-|-H20. It Ls very soluble in water, the solution reacting alkaline; on t'\'aporation in vacuo the solution deposits crystals of the com- jjosition 2KSH-I-H2O. With potassium hydroxide it forms the monosulphide : KSHH-KOH=K2S + H20. Potassium polysulphides. — "\'\Tien a solution of potassium mono- sulphide is boiled with sulphm-, we obtain the compounds K2S3, K2S4, K2S5. A mixture of these substances is also obtained by fusing potash with sulphur; besides these it contains the sulphate- and the thiosulphatc and is called hcpar sulphuris {" liver of sul- phur ") because of its liver-brown color. These polysulphides are 2^-'.] RUBIDIUM A.\'D CESIUM. 329 decomposed by aiitls with the evolution of hydrogen sulphide and the separation of sulphur: KoS, + 2HC1 = 2KC1 + HoS + (x - 1)S. Rubidium and Caesium. 232. These elements are widely distributed, but always occur in extremely small amounts. The silicate lepidolite, or lithia mica, fre- quenth' contains a little rubidium. The exceedingly rare mineral pullux from the isle of Elba is a silicate of aluminium and caesium, and contains about 30T( caesium oxide. In general these elements are found where- ever potassium salts are met with: in mineral springs, in the Stassfurt salts (cariiallite contains rubidium), etc. They were discovered by BuNSEN and Kirchhoff in 1860 with the aid of spectrum analysis (§ 264) and obtained their names from the most important hnes in their spectra (rubidus = dsiTk red; ccestMs = sky-blue.) The spectrum lines were used as a test in the separation of these elements from the others; after tr\-ing a possible method of separation the two sa\-unts would see which portion showed the lines of these elements the brightest; this portion was then examined further. In order to separate them from the large amount of potassium salts with which they generally occur, they are converted into chlorides p.nd the solution is evaporated, whereupon the dry residue is extracted ■nnth strong alcohol. Almost all the sodium chloride and potassium chloride remains behind, while the chlorides of rubidium and caesium dissolve. Platinum chloride is then added to precipitate KjPtClp, RbzPtCle, and CsjPtClj; the solubiUty of these double salts in water is quite different (at 10° 100 parts H,() dissolve 0.9 parts K,-salt, 0.154 libo-salt, and 0.05 Csj-salt), so that they can be very well separated by fractional extraction with boiling water. The rubidium iron "alum" is particularly well suited for the purification of rubidium salts, and especially for their seii;;i'ati(jn from potassium salts, since it is readily soluble in hot, and only slightly soluble in cold, water, and moreover crystallizes beautifully: potassium iron alum, on the other hand, is very soluble even in i-old water. The metals rubidium and ciesinni arc best obtahied by heating their hydroxides with lalcium tilings in a vacuum. The metals then distil 330 hWHGAXIC CHEMISTRY. 1232- off. Rubidium has a silvery lustre, melts at 38.5, and has a specific gravity of 1.522 at 15°. The metal oxidizes very rapidly in the air or in oxygen, forming dark brown crystals of the peroxide, Rb02. On being heated in current of hydrogen it yields the hydroxide and free oxygen : 2Rb02 + 2H2 = 2RbOH + H^O + 0. Rubidium oxide, when prepared in the same way as NajO, is obtained as transparent, pale yellow crystals which turn golden yellow on heating but lose this color again on cooling. The hydroxide is a ^-ery strong base ; its salts show much similarity to the analogous potassium compounds; they are in several instances less soluble, e.g., Rb-alum, Rb-perchlorate (§ 60), etc. C a; s i u m is a silvery-white metal; sp. g. 1.85; m.-pt. 26.5°; b.-pt. 670° It soon takes fire on exposure to the air. The oxide, CS2O, obtained in the same way as the other alkali oxides, is crystallized and is orange-colored at room temperature but almost black at 250°- The salts of ctesium are very simOar to those of rubidium ; some of them are even less soluble, and are therefore used for the preparation of pure cjesium compounds. This is particularly true of the platinum double-salt already mentioned and the caesium alum and the acid tartrate. Rubidium bromide and iodide, and even more so the corresponding compounds of csesium, have the property of combining with two atoms of bromhie and iodine, forming yellow or brown crystalhne compounds, e.g. CSI3; these metals can thus be trivalent. SUMMARY OF THE GROUP OF ALKALI METALS. 233. The gradual change of the physical properties of these metals with increasing atomic weight is made plain by the follow- ing table: Atomic weight. Specific gravity Melting-poioit. . Boiling-point. . . Atomic volume. Li Na K Rb 7.00 23.00 39 10 85,45 0.59 0.97 0.865 1.52 180.0° 97.6° 62.5° 38.5° <1400° 877° 757° 696° 11.8 23.7 45.3 56.7 132.81 1.85 26°-27° 670° 71.9 234.] AMMONIUM SALTS. :',?,] The specific gravity increases with the atomic weight, as does also the atomic vohime; on the other hand, there is a fall in the melting- and boiling-points. From a chemical standpoint wc notice, in the first place, the same general type in the compounds, showing that all these ele- ments are univalent. The hydroxides all have the formula R( )II, the halogen compounds RX, etc. The salts of them all, even the carbonates and phosphates, arc soluble in water (although in different degrees), the carbonates with basic reaction. The metals all oxidize very readily in the air. On the other hand, we cannot o\'erlook the fact that the metals potassium, rubidium and csesium, which are ^'ery similar to each other, differ from sodium and lithium in many respects. The last- named metal, as we shall see in the sorjuel, displays annlo.iry with magnesium in several important points, thus differing from the metals of its own group. A slight di\-ergence in the behavior of the fu-st members of a group from that of the rest is found to characterize almost all of the groups, ^^"e may recall carbon, for instance, the first member of the fourth group, which differs dis- tinctly from silicon and the rest in the ability of its atoms to unite with each other; also fluorine with its soluble silver compound. Still other examples of this sort will be met with later. Sodium differs from the sub-group, K, Rb, Cs, in the solubility of its salts. The sodium salts are almost all readily soluble in water; this is true even of the platinum double-salt, XaaPtCle. the acid sodium tartrate and others. Soda crystals effloresce, while potash deliquesces in the air. The spectra of sodium, on the one hand, and the other alkali metals, on the other, are entirely dis- similar. Ammonium Salts. 234. In the description of ammonia (§ 112) it was already observed that it combines with acids directly, forming salts which are \-ery similar to those of ]iotassium, and in which the group, NH4, the ammonium group, is assumed to exist. In connection with the alkali group a description of a fe-w ammonium salts may find a place. The aqueous solution of ammonia must, because of its electrical conductivity and its alkaline reaction, contain NH4 and OH' ions ;J32 IXORGANIC CHEMISTRY. [§234- and hence also undissociated molecules of ammonium hydroxide, NH4OH. While solutions of the alkalies, KOH, NaOH, etc., conduct the electric current very well, this is not the case with an ammonia solution; it is a poor conductor. A 0.1 normal solution contains only 5% of ionized NH4OH molecules, while a solution of potassium hydroxide of the same concentration is 91% ionized. An aqueous solution of ammonia may be presumed to contain: (1) free ammonia, XH3; (2) hydrates of ammonia, XHs-n aq.; (3) ammonium hydroxide, XH4OH; (4) the ions XH4 and OH'. The existence of these hydrates in addition to free ammonia reveals itself in the behavior of ammonia solutions on being shaken with chloroform. According to Berthelot's law (Org. Chem., § 25) the distribution ratio of the ammonia between the two solvents should be constant. Eut this is not the case. Therefore, just as the deviations from Henry's law lead us to conclude that a dissolved gas exists in a special condition, so wc can apply a simi- lar explanation to the exceptions to the Berthelot law; for Henry's law Ls really the expression of the distribution ratio of a gas between a liquid and a vacuum, while tlic other law has to deal with the distribution ratio between two liquids. Since in the case of ammonia this deviation is observed only when one of the two liquids is water, we are obliged to conclude that there Ls a combination of the ammonia with the water. The reason for assuming the existence of hydrates instead of ammo- nium hydroxide, XH4OH, is a double one. We have, on the one hand, the analogy between the behavior of ammonia and amines and, on the other, the entirely abnormal beha^•ior of the organic quaternary ammonium bases. For, while the aqueous solutions of primary, secondary and tertiary amines are weak electrolytes, as is the case with ammonia, the solution of a quater- nary base, on the contrary, conducts electricity as well as a solu- tion of potassium or sodium hydroxide. We may thus conclude that if ammonium hydroxide, NH4OH, could reach the same concentration in solution as a quaternary base it would display just as great a conductivity as the latter. Unlike the quaternary base, however, it breaks up principally into ammonia, XH3, and water, for the quaternary base cannot be thus decomposed. That an aqueous ammonia solution really contains at least §234.] AMMONIUM SALTS. 33:5 an appreciable quantity of the hydroxide is evidenced by the existence of its ions. These necessitate the establishment of the equilibrium NHiOH^NHi+OH', however far the point of equilibrium may be displaced toward the right. The great tendency of ammonium hydroxide to break up into ammonia and water is the reason for the yevj feeble basic reac- tion of an aqueous solution of ammonia, for undoubtedly ammo- nium hydroxide, so far as it is formed, is extensively ionized, like the strong bases. This view is supported by various observa- tions, among them the neutral reaction of the ammonium salts of strong acids, such as the chloride and the nitrate, and also the alkaline reaction of the carbonate and the cyanide, in har- mony with the simUar alkaline reaction of the corresponding salts of the alkali metals. Amm onium chloride, NH4CI, sal ammoniac, is obtained from the ammonia hquor of the gas factories (§ 112). The ammonia is expelled by warming and absorbed in hydrochloric acid; this solution is evaporated and the sohd residue sublimed, wherebv the salt is obtained in compact fibrous masses. It dissolves in 2.7 parts of cold, and in 1 part of boihng, water and crystallizes out of the solution in small, usually feather-like groups of octa- hedrons or cubes. It has a sharp saline taste. Ammonium chloride vaporizes easily, dissociating into ammonia and hydrochloric acid, as is shown bj' the vapor density', which at 350° is only half as great as calculated. This dissociation can be easily demonstrated in the following manner : Introduce into a tube sealed at one end a little ammonium chloride and, not far from this, a piece of blue litmus paper. In front of the latter is pushed a plug of asbestos wool and finally a piece of red litmus paper. The chloride is then heated. Since hydrochloric acid has a smaller diffusion velocity than ammonia the latter passes through the wad first and colors the red paper blue; as a result an excess of hydrochloric acid is left at the other end and it reddens the blue paper placed there. It is a remarkable fact, discovered by Baker, that perfectly dry ammonium chloride (having stood for a long time in a desiccator over resublimed phosphorus pentoxide) has the normal vapor densit.y. On the other hand, the same investigator found that 334 INORGANIC CHEMISTRY. [5§234_ similarly dried ammonia gas and hydrochloric acid gas do noi unite to form ammonium chloride (§ 38). Traces of water thus produce a marked catalytical acceleration, both of the formation and of the decomposition of ammonium chloride. We liave here an illustration of the general rule that when one part of the system in a reversible reaction is accelerated by a catalyzer the other must be lil clings to it. At present, however, it is made by dry distilling a mixture of calcium carbonate and ammonium chloride or sulphate. The product is a mixture (moleculo for molecule) of acid salt, NH4HCO3, and ammonium carbamate, NH2-C02NH4 (this latter being the neutral salt minus IHoO). From its composition, (NH3)3 (002)2 'HaO, it takes tlie name (imnwniuni sc.iqiiirnrlxjmilc, On passing ammonia gas into a concentrated a( peony sohition of it the neutral salt, (NH4)2C03, separates out as a crystalline powder; it smells strongly of ammonia and passes slowly o\'er into the acid salt, NH4HCO3, a white odorless powder, which is scarcely soluble in water. This acid salt is also formed directly from the sesqui- carbonate, as the latter gi\'cs off carbon dioxide and ammonia in the air (hence the odor of ammonia) and goes o^-e^ into the first- named salt. Ammonium sulphide is extensively used in analysis (§ 73). A solution of ammonium hydrosulphide (or sulphydrate), NH4SH, is obtained by saturating aqueous ammonia with hydrogen sulphide; it is a colorless liquid, which soon turns yellow because of the for- mation of ammonium poh'sulphidcs. The oxygen of the air oxidizes part of the hydrogen sulphide and thus sets free sulphur, which combines with ammonium hydrosulphide to form polysulphidcs. These polysulphidcs are also obtained by dissoh'ing sulphur in a solution of ammonium hydrosulphide. On mixing 2 vols. XH3 gas and 1 vol. H2S gas at — 18° a white crystal- line mass is obtained, which decomposes at ordinary temperatures into NH^SH and XH3. The compound XH,,SH separates out crystalline when hydrogen sulphide is ]iasscd into alcoholic ammonia. As low as 45° it is completely dissociated into equal volumes of XH3 and HoS. SALT SOLUTIONS. 235. Every solid substance is soluble in every liquid; howe\'er, the proportion which dissolves can ^■a^y all the way from zito to infinity. If only an infinitesimal amount of the solid goes into solution, we say ordinarily that the substance is " insoluljle " in the liquid; there can be no doubt, however, that, if our means of inves- tigation were sufficiently impro\ed and large enough cpiantitics of liquid were taken, the solubility would be perceptible. This has already been demonstrated in many cases of so-called insoluble substances (§ 210). Even when we ctmfine our attenticm to aqueous solutions of salts (including acids and bases) we find the 336 INORGANIC CHEMISTRY [§§235- same infinite difference in solubilit}' that is observed between sub- ^tances in general. Substances such as sand, barium sulphate (§ 262), silver iodide, etc., are "insoluble "; others, like sulphuric acid, are able to dissolve in any given amount of water. With regard to the solubility of salts the foUo-vidng practical rules are worth remembering : Potassium, sodium and ammonium salts are soluble. — Normal nitrates, chlorates and ace- tates are soluble. — Normal chlorides are soluble (except AgCl, HgzCb, and PbClj). — Normal sulphates are soluble (except those of Ba, Sr, Ca, and Pb) . — H y d r o x i d e s are insoluble (except those of the alkaUes and alkaline earths) . — Normal carbonates, phos- 11 h a t e s, and sulphides are insoluble (except those of the alka- lies). — Basic salts are insoluble. — Acid salts are soluble if the acid ilself is soluble. The solubility, i.e. the maximum relative amount of salt that can go into solution, is a function of the temperature and the pressure. In the great majority of cases the solubility increases with the temperature. If the temperature is plotted on the axis of abscissas and the amount of salt which dissolves in one hundred parts of water is plotted on the ordinate axis, a solubility curve is obtained (Fig. 52) which shows at a glance the variation of the solubility with the temperature. For some salts, e.g. potassium nitrate, the solubility increases very rapidly with the temperature; for sodium chloride it remains practically constant. In certain cases, such as those of calcium hydroxide and calcium sulphate (within certain limits of tempera- ture) the solubility decreases with rising temperatm'e. These phenomena are connected, as has already been explained, with the heat of solution, i.e. with the caloric effect which accompanies the process of solution, and in the manner expressed by van't Hoff's principle of mobile equilibrium (§ 103). In fact saltpetre, for instance, whose solubility increases very rapidly with the temperature (see Fig. 52) dissolves in water with a considerable absorption of heat. 236. The term heat of solution has various meanings. We are obliged to distinguish between (1) the caloric effect of dissolving a salt in a very large amount of water; (2) the caloric effect of dissolving a salt in an almost saturated solution; and (3) the total heat of solution, i.e. the whole caloric effect of dissolving a salt in water until the solution is saturated. As a rule these three magnitudes will have dissimilar values, indeed their algebraic signs may be opposite. This is the case, 236.] SALT SOLUTIONS. 337 for instance, with the compound CuClj -21120; 1 g.-mol. dissolved in 19S g.-mols. lip at 11° gives a caloric effect of +3.71 Cal.; 19.56 g.-mols. in the same amount of water, —3.129 Cal. The heat of solution to which v.vn't Hoff's principle applies is that of the salt in its saturated solution. A\"e have here the system: r 80° fto" 00" 00° 70° Fig. 52. — Solubility Cttrves. 80- 100° C. salt '+ saturated solution; when the temperature changes, the equi- librium is displaced, i.e. salt either goes into solution or crystallizes out, the latter action producing just as large a thermal effect numerically as dissolving in the saturated solution, but vnth. the opposite sign. Since this was not taken into consideration when the matter was first dis- 338 INORGAXIC CHEMISTRY. [§§ 236- cussed, it was believed that there were exceptions to the principle, but closer investigation has proved the contrary. In some cases the solubility of a salt at first increases gradually with rising temperature and then steadily decreases, so that the solubility curve has a maximum (c/. Fig. 53). In full agreement with van't Hoff's principle the heat of solution is negative in thp ascending portion of the curve, zero at the maximum and positive in the descending portion. In the case of gj^jsum, CaS04 -21120, for instance, the maximum was found to lie at about 38° and at that point the heat of solution was actually proved to be 0.00; at 14° it is -0.36; above 35°, +0.24. The effect of pressure on the solubility is at the most -^-ery slight, but it is in entire accord with the principle of Le Chatelier. ^ Anmionium chloride, for instance, dissolves „ With expansion; therefore its solubility lessens with increasing pressure (1 % for an increase of 160 atm.). Copper sulphate, which dissohes with contraction,, has its solubility increased 3.2% by an increase of 60 atm. pressure. 237. It was formerly thought that the terms " soh-ent " and " dissolved substance " (" solute ") should be kept distinct. How- ever, it has since developed that there is no essential difference between the components of a solution, and that aqueous solutions arc therefore better defined as " licpid complexes, one of whose components is water," than as " water in which substances are dissolved." The interchangeability of the terms " solvent " and " solute " is evidenced first of all by the phenomena attending the cooling of salt solutions. Let us consider, for instance, a nearly saturated solution of potassium chloride at a definite temperature. We have in it two substances (KCl and II2O) and two phases (§ 71), hence two degrees of freedom. We will suppose that the solution is then cooled; potassium chloride crystallizes out forthwith and, as three phases are then present, the system becomes univariant We recall that changes in the quantity of any phase have no effect on such a system; therefore, if more salt is introduced into the system, the concentrations of saturated solution and vapor are unaffected. 237.] SALT SOLUTIONS. 339 This is none the less true when water is added or the vapor ■\-ohime increased, so long as the three phases remain. On cooling still farther, more potassium chloride is gradually- deposited until a point is reached below which the entire liquid congeals to a mixture of salt and ice. This point is known as the cryohydric, or eutectic, point. There are now four phases, — salt, ice, solution and A-apor; — hence the system has become nonvariant. The opinion was formerly held that at this point a chemical com- pound between the salt and water (a " iTvohydrate ") came into exist- ence. That it is only a matter of mixtures can be seen in the case of colored salts (K^CrOJ, for instance, with a microscope; moreover the composition of these so-called hydrates may differ in case the solidifi- cation takes place under a different pressure. If we start with a dilute potassium chloride solution as another example, and cool it, we ha\-e ice formed at a definite temperature and a univariant system established, ice being the third phase required. Below this point the solution can he regarded as saturated in respect to iec, just as it could be considered saturated in respect to the salt in the previous case; for an increase of the soUd phase (ice) does not now cause a displacement of the equilibrium (§ 71) any more than the addition of the solid (salt) did in the pre\ious instance. The addition of potassium chloride causes part of the ice to go into solution (i.e. melt) ; for the dissolving of more salt increases the concentration of the solution. Therefore, if the temperature is kept constant, ice must melt in order to restore the solution to its previous concentration. It is therefore perfectl}' analogous to the addition of water to a satu- rated potassium chloride solution in contact with the solid salt, in which case also the solid phase goes into solution. If the tem- perature rises, more ice dissohes; if it falls, more crystallizes out — just as with rising temperature more potassium chloride goes into solution and with sinking temperatiu'c more crj-stallizes out. On farther cooling more and more ice will Ijc deposited until, in this case also, the cryohydric point is reached, below which the whole system solidifies to a mixture of salt and ice. The analogy is therefore complete. Tlje cryohydric point is, according to this view, the point of intersection of two cur\es, \'iz.: the solul^ilitj' curves of salt and of ice in the salt solution. 340 INORGANIC CHEMISTRY. [§237. Another argument against the assumption of any essential difference between solvent and solute is found in the behavior of the solutions of certain hydrous salts, e.g. CaCl2 -61120. A satu- rated solution of CaCl2 in water at 30.2° has exactly the com- position CaCU-GHoO. At this temperature, therefore, the h3'drate melts to a homogeneous liquid. // e ith c r H2O or CaCU is added, there is a deposition of CaCl2 ■ 6H2O on cooling, for the addi- tion of either causes a depression of the point of solidification (freez- ing-point) of CaCl2-6HoO. In the first case this hydrate is in equilibrium with a liquid which contains more water than the hydrate does and which is therefore called an aqueous solution in the ordinary sense. In the second it is in equilibrium with a liquid which contains more CaCl2 than CaClo -61120 and must therefore be regarded as a solution in CaCl2. On examining the solubUity curves of various salts (c/. Fig. 52) it is found that they are in general regular; however, in one of the curves (sodium sulphate) a sudden change of direction is noticed. This is often observed with salts that contain water of crystalliza- tion. Taking sodimn sulphate as an example, the phenomenon may be explained thus : It has already been remarked (§ 225) that this salt has a transition point at the temperature of 33°, Na2S04 • IOH2O being transformed into Na2S04 and IOH2O. Up to 33°, therefore, we have the hydrous salt as the solid phase; above this temperature the anhydrous salt. This change must necessariLy involve a sudden bend of the solubility curve. Below 33° the curve represents the solubility of Na2SO4-10H2O, above 33° that of Na2S04. We can therefore also regard the point of inflection of the curve {at S3°) as the point of intersection of the curves for Na2S0i- IOH2O and Na2S0i. In s-odiuTii sulphate the special case appears where the sol- ubilitj' of the anhj'drous salt decreases with rising temperature and hence the solubility curve falls as the temperature rises above 33°. In the light of the above the solubility of a substance which has a transition point is the same for both modifications at this point. This must always be the case; it can be demonstrated in the same w&y as in § 70, where it was shown that the vapor pressures become equal at the transition point. Indeed the same figure can be employed, if it is borne in mind that the solubility of a metastable modification is always greater than that of the stable modification at one and the same temperature. Inversely, more- over, we have here a means of determining the transition point. § 2370 SALT SOLUTIONS. 341 100 iHYOROLISjiii In general, as Ostwald has pointed out, the solubility of any substance whatevei- is dependent on the condition in which it exists. The solid phase determines the equilibrium, not only in \'irtue of its chemical composition but also Ity the particular modification in which the solid substance is present. Thus, e.g., each of the various forms of the same polymorphous substance or different hydrates of the same salt has its own soKibilit}', other things being equal. In a hydrous salt we maj^ have the case where there are various hydrates, which are connected with each other by ti-ansition points. A salt -with m+n molecules of water of crystallization passes over at a definite temperature into another with m molecules, for ex- ample. The latter may, at a higher temperature, have a second transition point (to anhydrous salt). At each of these points the solubility cur-\-e will show a bend, because the solid phase changes; the curve will therefore assume some such form as that of Fig. 54. Let us examine such a solubility curve a little more closely. At 0° (A i\\ Fjg. o-i) we will suppose that we have pure water and ice, to start with, and that small portions of salt are then gradually dissolved. If the ice ti phase is to be preserved, ^o the temperature must be allowed to sink, for a salt solution has a lower freezing-point than pure Mater. We therefore move along the curve AK. Soon a point K is reached when no more salt dissolves, since all the water has now turned to ice. Here, therefore, we have a mixture of ice and solid salt, or, in other words, the cryohydric point. If we wish to bring more salt into solution after 7v is reached the temperature must l^c raisetl. Tlic ice phase then, of course, disappears and in its place wc> ha^e the salt with m + n molecules of water of crystallization as solid phase. If the temperature is steadily raised and the solution is kept constantly saturated liy the addition of this salt, wo mo\-e along the curve KB. At B, however, we meet the transition point from the salt witli m + n mols. H2O to the one with m mols. H2O; hence the soluliility om the standpoint of the ionic theorj^ the following theory of indicators has been ad^-anced: If a couple of drops of the indicator are introduced into an acid solution, the ionization of the indicator, ^^•hich is only very slight, is reduced by the great excess of acid to practically zero. If a base is then added, the H-ions of the acid to be titrated are removed bj- the OH-ions. Howe^'er, if the acid is very strong, enough H-ions remain in the liquid up to the last to prevent an^-thing like an extensive ioniza- tion of the coloring-substance; not until the first excessi^'e drop of alkali is added do the anions of the coloring-substance come into existence, the alkali compound of the latter being strongly disso- ciated. The change of color is therefore sharply defined, for it is due to this tUfference in color of the non-ionized molecule and the anion. On the other hand, if the acid is a weak one, there will not be enough H-ions present when the end of the titration is nearly reached to prevent a slight ionization of the coloring-substance. As a result we shall have in the solution not only the undissociated coloring-substance but its anions as well, even before the titration is completed, — in other words, tJie change of color becomes more gradual and hence the end reaction more indefinite. The effect will be the same if the alkah employed contains carbonate. In that case near the end of the titration the solution will only contain carbonic acid, which is very weak; consequently the color change will not be sudden. It is for this reason that in titrating soda solu- tions (see § 240) the carbonic acid must be expelled by boiling. If a weak acid is to be titrated, it is necessary, according to the above, to select an indicator which is much less ionized even than the acid itself and whose allcali salts are sufficiently ionized to produce a distinct change of color. A very suitable one for this purpose is phenolphthalein. Acetic acid, for example, can be sat- isfactorily titrated with it, if a strong base is employed, for the reasons set forth above. On the other hand, in case a weak base is to be titrated, phenolphthalein is not so satisfactory. Ammonia 242.] COPPER. 3.')3 does not color a phenolphtlialein solution till a considerable excess is added, because at the great dilution in which the ammonium- phenolphthalein compound exists in a titration it is almost com- pletely spHt up by hydrolysis (§ 66). If a weak base is to be titrated, an indicator must be selected which is a relatively strons;- acid, for then the salt of the coloring- substance will be hydrolyzed only to a limited extent, e^■en near the termination of the titration (i.e. when the concentration of the base has become weak), and tlie color of its ions will therefore still predominate. For such a titration a strong acitl (e.g. hydro- chloric or sulphuric acid) must be used, in order that the first drop after tlie point of neutralization is reached may diminish the elec- trolytic dissociation of the coloring-substance and so gi^'C the solu- tion the color of the non-ionized molecules. Methyl orange is an indicator that ans\\'ers these requirements; it serves \ery well in the titration of ammonia. All other indicators are intermediate to these two extremes (phenolphthalein and methyl orange) as re- gards their ionization, and their applicability is determined accord- ingly. COPPER. 242. This metal occurs native in America, China and Japan, forming regular crystals. Other copper minerals are cuprite (Cu2- splits up into its elements. On being wanned with water it yields 2Au and AuCls. Aurous iodide is formed (like cuprous iodide) on treating a solution of the chloride with potassium iodide. The gold double cyanide, KCN-AuCN, is prepared by dissolving auric oxide in potassium cyanide; it is used in gold-plating. Of the oxy-salts of aurous oxide only a few double salts are known. Auric Compounds. Auric chloride, AuCls, can be obtained by dissolving gold in aqua regia or by the action of chlorine on the metal. It forms a dark red crystalhne mass, which deliquesces rapidly. On the evaporation of its solution it partially decomposes into chlorine and aurous chloride. By evaporating with hydrochloric acid long yellow needles are obtained, consisting of a compound AuCls -HCl, which can be regarded as chlor-avric acid, ilauy salts of this acid are known to exist, e.g. KCl • AuCls + 2^H20 and NH4CI.AUCI3 + H2O, as well as many chlor-aiu^ates of organic bases. These double salts give the ordinary tests for gold, hence this acid either forms no complex ion AuCU' or is very unstable. Auric chloride is also soluble in alcohol and in ether. Auric oxide, AU2O3, can be prepared bj^ precipitating auric chloride with magnesia. The latter can be removed from the pre- cipitate with concentrated nitric acid, the auric oxide remaining as a brown powder, which breaks up at 250° into its elements. If the precipitate produced by magnesia is treated with dilute nitric acid, a reddish-yellow powder of the formula AUO3H3 is obtained, which cUsplays acid, instead of basic, properties. Salts of this auric acid are known, which are derived from the com- pound Au(0H)3 — H20=AuO-OII. Potassium aurate, foi exam- ple, has the formula I^u02-l-3Il20 and crystaUizes in yellow needles. Many other salts are also known; the above-mentioned precipitate with magnesia, for example, can be looked upon as the magnesium salt of auric acid, Mg(Au02)2- Auric sulphide, AU2S3, is precipitated from gold solutions by hydrogen sulphide. It is very dark brown and soluble in am- monium sulphide. Gold is precipitated from its solutions in the metallic form by 251.] AURIC COMPOUNDS. 371 various reducing agents. Ferrous sulphate (§ 248), oxalic acid and acetylene water serve very well for this purpose. Hydrogen peroxide precipitates gold quickly in alkaline solution. 251. For many centuries the alchemists endeavored to produce gold from the baser metals. It is needless to say that their efforts were never rewarded. The chances of this hope being realized must at present be regarded as very slight, since gold is an element. Inasmuch, however, as our conception of an element is relative (§8), i.e. it depends on the extent of our mastery over natural forces, the impossibility of decom- posing gold or synthesizing it from other elements is by no means abso- lutely established. (See § § 266-7) . Although we now ascribe to every metal fixed, unalterable properties, it might well have seemed possible to the alchemists, with their more limited knowledge, that the properties of the metals could vary. None of the metals except gold occur pure in nature; they have to be extracted from oxides or sulphides, which frequently contain various impurities. The metals thus obtained had no definite properties; distinction ^A-as made between various sorts of lead, copper, etc. The mutabilit}' of the metals may be said to have been the first principle which observation taught; indeed, when a piece of metal is fused with small amounts of various other substances, its properties (color, etc.) really do change. Moreover, at the time of the alchemists the present concept "element " was not jTt established; this was first introduced by Boyle (1627-1691). Before then, the doctrine of Aristotle was very generally accepted, according to which all substances are made up of air, fire, earth and water. In order to produce gold it therefore seemed only necessary to deprive the baser metals of certain properties and substitute others. As to the metals themselves the idea was prevalent in alchemistic circles that mercury was the primordial substance and that it had undergone various changes. Before gold could be made from it it must be made refractory and of a yellow color. Xot a few alchemists were convinced, moreover, that the success of the "great work " depended on the cooper- ation of a higher power. SOniAKY OF THE GROUP. 252. The metah copper, silver and gold form a bridge from the difTicultly fusible metals, Xi, Pd, Pt (Group VIII), to the easily fusible, Zn, Cd, Hg (Clroup II); their melting-points are between those of the two groups. The following brief table sum- marizes the phj'sical constants of these metals as well as those of the related elements, lithium and sodium: 372 INORGANIC CHEMISTRY. ■ 252- y Li Na Cu Ag Au Atomic weight Specific gra\'ity .Melting-point Color 7.00 0.59 180 white 23.00 0.97 97.6 white 63.57 8.94 1083 red 107.88 10 5 961 white 197.2 19.33 1063 red The analogy in the chemical properties is chiefly apparent in. the -ous compounds. These have the type R2O for the oxygen compounds and RX for the halides. The -ous halides of Cu, Ag and Au are all white and insoluble in water; they are isomorphous with sodium chloride. iloreover, there are certain analogies in solubility. Lithium carbonate and hydroxide are less soluble in water than the corre- sponding sodium compounds; copper carbonate and hydroxide are insoluble, while the corresponding silver compounds dissolve to some extent. The sulphate of sodium (third horizontal series) crystallizes preferably with IOH2O, that of copper (fifth series) with 5H2O, while silver sulphate (seventh series) is anhydrous. The oxygen compounds exhibit a gradual decrease in stability. Li20 and Na20 are unaffected by high temperatures, but CuO is transformed into CU2O, and the oxides of silver and gold break up even at comparatively low temperatures into their elements. However, it must be admitted that the analogy between these elements is not so great as in other groups. Their difference in. ^'alence is especially striking and, moreover, there is little simi- larity in the properties of the higher stages of oxidation. This is one of the weak parts of the periodic system. BERYLLIUM AND MAGNESIUM. I. Beryllium (Glucinum). 253. This is one of the rarer elements. It occurs in the mineral beryl, Al203-3Si02 + 3(BeOSi02); that variety of beryl which is colored green by traces of a chromium compound is the gem called emerald, or smaragd. Chrysoberyl has the composition BeO-AljOj. Almost all the beryllium compounds are made from beryl. This is disintegrated by fusing with potassium carbonate. The fused mass, after cooling, is treated with sulphuric acid to precipitate the silica. Most of the aluminium is then removed by crystallization in the form of alum. 253.] BERYLLIUM AND MAGNESIUM. 373 as this is sparingly soluble in oold water, while beryllium sulphate remains in the mother liquor. The latter is then mixed with a hot solution of ammonium carbonate to precipitate aluminium and iron, beryllium still remaining in solution. After acidifying with hydrochloric acid, the beryllium is precipitated as the hydroxide by ammonia The metal was obtained by heating the double' lluoride BeF2-2KF with sodium. It is a malleable solid with the specKic gravity 1.04. It does not decompose water, (■\-en at 100°. At ordinary temperatures it is permanent in the air. Hydrochloric and sulphuric acids diaoolve it readily with the evolution of hydrogen; dilute nitric acid docs not attack it so readily. Beryllium is also dissolved easily by caustic potash and soda vrith the evolution of hydrogen and the formation of salts having the formula BeCOR),. The hydroxide thus behaves as a weak acid towards strong bases. These properties correspond to those of aluminium; in § 215 attention was already called to the analogy between these two elements. This analogy also characterizes their compounds, e.g., beryllium carbide yields pui'e methane with water, just like alu- minium carbide (§ 17S). Only one oxide nf beryllium is known, BeO (§ 215). It Is a white powder, which after ignition is difficultly soluble in acids (like -MjOg). It is obtained by heating the hydroxide, Be(0H)2, which is precipi- tated from solutions of the salts as a white gelatinous mass. When freshly precipitated, it is easily soluble in alkalies, anunoniuni carbonate and dilute ai'ids. On being heated with water, dilute ammonia solution or dilute alkali solution, or on being ignited, or even on standing for some time, it " grows old " and loses these properties. Heating with ten-fold normal solution of an alkali hydroxide " rejuvenates" even the " oldest" beryllium hydroxides, which are dissolved only .slowly by warm concen- trated hydrochloric acid. Beryllium hydroxide is distinguished from aluminium hydroxide in two respects: it dissolves in ammonium carbonate (see above) and is precipitated from the solution in caustic soda or caustic potash by prolonged boiling. Beryllium sulphate, BeSO, ci ystallizes with four or seven molecules of water, in the latter case being isomor,^hous with M^iSOj-THjO. The dovible salt BeSOj-KjSOj-SHjO is (like alum) sparingly soluble in cold water. Beryllium chloride, BeClj, must be prepared from the oxide by heating with charcoal in a current of chlorine. Its vapor density corresponds to the formula Bed, It crystallizes with 4H2C). Beryllium carbonate is soluble in water. It loses carbon dioxide very easily. The beryllium salts taste swcci, licncc llir name glucinuni (or jilycin- iuni), which is common in France anil Anicrica. 374 INOROANIC CHEMISTRY. [j§ 254- II. Magnesium. 254. This element occurs as carbonate, sib'cate, and chlo- ride in considerable quantities. Magnesite is MgCOa, dolomite MgCa(C03)2. Among the silicates containing magnesium we have talc and soapstone, H2Mg3Si40i2; serpentine {asbestos), H4Mg3Si209; meerschaum, H4Mg2Si30io. It is found in smaller amounts in many other silicates, e.g. hornblende (asbestos), augite, tourmaline. Other salts found in nature are camallite, MgCl2 ■ KCl ■ 6H2O, kie- serite, MgS04-H20, and kainite, MgS04- KCl -31120 (Stassfurt Abraum salts). Upon the weathering of the silicates the mag- nesium goes into the soil, whence it is absorbed by the plants (to which this element is invaluable) and finally taken into the animal body. The metal is manufactured on a large scale, since it is employed for illumination in photography, pyrotechnics, etc., on account of the intense light (flash-light) produced by its combustion. At present it is prepared mainly by the electrolysis of fused mag- nesium chloride or carnallite in a cast-steel crucible, which serves as cathode; gas carbon is used for the anode. It is also obtained by heating the double chloride MgCl2-NaCl with sodium. It is silvery- white and has a high lustre. Sp. g. =1.75. It is malleable and ductile and comes on the market in the form of wire or ribbon as well as powder, but the ribbon frequently contains zinc. It melts at 651° and boils at 1120°. It is quite permanent in the air, since it soons becomes coated with a thin cohesive film of the oxide; at an elevated temperature it burns to magnesia, MgO. When it is heated red-hot in a limited supply of air, a large part is converted into the nitrite, Mg3N2, a yellowish-green substance. Boiling water decomposes it slowly with the evolution of hydrogen. It dissolves readily in acids but is unaffected by alkaUes. It is a powerful reducing-agent, reducing sihca (§ 190), for example; moreover, when ignited, it burns in water vapor. Magnesium oxide, MgO, magnesia, is the only oxide of mag- nesium known. It results from the combustion of the metal or from heating the hydroxide or carbonate. It is a white, very hght powder, which is employed in medicine under the name mag- nesia usta. With water it forms the hydroxide Mg(0H)2. Magnesium hydroxide, Mg(0H)2, is precipitated from solutions of magnesium salts by alkalies. It is slightlj' soluble in water and 255.] MAGNESIUM SALTS. 375 turns red litijius blue; howe\'cr in an excess of alkali its ionization is so diminished that it becomes practically insoluble. It is a weak base, but is strong enough to absorb carbon dioxide from the air. It dissoh-cs readily in an aqueous solution containing ammonium salts. According to Ostwald, this is to be explained as follows: The solution of an anmionium salt contains a large quantity of XIl4-ions. When a substance is introduced into the solution, which gives off OH-ions, as does magnesium hydroxide, these XIl4-ions unite with OH-ions to form NH4OH, or rather NH3+H2O (c/. § 234). As a result of this reaction OH-ions dis- appear. In order to restore the equilibrium between the undis- solved magnesium hydroxide and the solution, more of this hy- droxide must go in solution, but again the freshly formed OH-ions are taken up b}- the NH4-ions. If sufficient t )f the latter are present, this process will go on till all the magnesium liydroxide has entered into solution. It now becomes clear why, on the other hand, the solution of a magnesium salt is not precipitated by ammonia in the presence of a sufficient quantity of ammonium salt. MAGNESIUM SALTS. 255. Magnesium chloride, ^[gCU, crystalhzes with six mole- cules of water and is ■N'cry hygroscopic. The deliquescence of common salt is due to tlie magnesium .salt it u.sually contains. ( )H On evaporating the aqueous solution the basic chloride, Mg _,, , and hydrochloric acid are formed; sea-water cannot be used in boilers because of the magnesium salt it contains, for the hydro- chloric acid set free attacks the iron. Many double salts of mag- nesium chloride are known. It can be obtained anhydrous by heating the double chloride MgC'l2-XH4Cl-6H20, when it forms a laminar-crystalline mass, which melts at 708° and distils without decomposition at bright red heat. Careful study of the decomposition of magnesium chloride by oxygen and by steam has shown that a reversible reaction is involved in each case: 2MgCl2+(),rf2.MgO+2ri2; .^^g('U + H,(),-=i^[g(■) + 2HCl. In the former reaction a rise of temperatuie displaces the equilibrium toward the right, although l)elu\v .500° the velocity is still very small. 376 IXORGAXIC CHEMISTRY. [§§ 255- Iii the second process the composition of the gaseous equilibrium mix- ture at 700° has been found to be 90% HCl + 10% HjO. Magnesium sulphate, MgS04 • 7H2O, Epsom salt, finds use in medicine. It is very soluble in water. It loses 6 mols. H2O at 150°, and the seventh above 200°. In this respect it behaves hke other sulphates, e.g. ZnS04 -71120, FeS04-7H20, and those (i nickel and cobalt, which are, moreover, isomorphous with it. A further analogy between these sulphates appears in the fact that with sulphate of potassium or ammonium they form double salts of the same type, K2S04-MgS04 -61120, which are also isomorphous. Magnesium ammonium phosphates, MgNH4P04 • 6H2O, serves for the precipitation of magnesium as well as of phosphoric acid. It is not wholly insoluble in water, but does not dissolve in ammonia, the reason for which is again to be found in the reduc- tion of the ionization. Completely analogous to this compound is the corresponding arsenate, MgNH4As04 - 6H2O. Magnesium carbonate. — From solutions of magnesium salts soda precipitates a basic carbonate, Mg(OII)2-4MgC03 -41120. The carbon dioxide liberated holds part of the magnesium in solu- tion as acid carbonate. This precipitate is known as magnesia alba. The neutral carbonate can be prepared from it iDy sus- pending magnesia alba in water, passing in carbon dioxide and allowing to stand; in_time the salt MgCOs-SHaO crystalhzes out, which is, hri^-^c^-cr, readily split up hydrolytically by water, form- ing basic carbonate again. CALCIUM, STRONTIUM AND BARIUM. I. Calcium. 256. This element is one of the ten principal constituents of the earth's crust (§8). Particularly the carbonate is found in large quantities in nature, limestone, calcite, aragonite, marble and chalk, all being forms of it. An earthy deposit containing a certain amount of calcitun carbonate is termed marl. Calcium silicates and especially calcium double salts constitute the major portion of the siliceous rocks. There are also extensive beds of calcium phosphate, phosphorite, apatite, etc., particularly in Spain and Florida. Calcium occurs as sulphate in the form of gypsum and alabaster. Moreover, in the animal kingdom large quantities 257.] OXIDES A.\D HYDROXIDIIH OF CALCIUM. Ml of this element are found. The skeletons of veitebrates are rliielly phosphate and carhonate of calcium; the shells of mollusks con- sist of calcium carlionate, as do also ejrfishells. As for the plants, lime is one of their indispensable inorganic constituents. Metallic calcium can lie obtained by electi-olj-sis of a fused mixture of calcium chloride and calcium fluoride. Such a mixture melts much lower than the single salts (§ 2.37). The lower temperature makes the separation of the metal easier and prevents its combustion. Calcium is a silvery-white metal, which melts at S00°; it is soft enough to cut and is malleable, but less so than potassium and sodium; it has a ci-^'stallinc frac- ture. Sp. g. = ]..")2. It is relati\ely little affected by oxy- gen, chlorine, bromine, and iodine, all of which leact with the metal only at a higher temj^erature than the ordinary one. In a current of air calcium unites with both oxyucn and nitrogen (§ 110). AA'ith hydrogen it forms a compound ('aH2, which is also prepared commercially by passing hydrogen into molten calcium. The calcium hydride reacts with water most vigorously raH2-|-H20 = CaO + 2H2. Since 1 kilo of the hydride evolves about 1 cubic meter of h\xh'ogen, it constitutes a very suitable material for generating hydrogen for aeronautic purposes, especially in out-of-the-way places. OXIDES AND HYDROXIDES OF CALCIUM. 257. Calcium oxide, Ca(J, (quick-lime, unslaked lime) is pre- paxed commercially by " burning " limestone or moUusk shells. The limestone is mixed with coal and the latter is set on fire; the heat of the burning coal decomposes the cai'ljonate of lime into calcium oxide and carbon dioxide. The kilns are usually con- structed in such a way that the burned lime can be drawn out at the bottom while the mixture of fuel and limestone is fed in at the top, so that the process fecontinuous. In the United States " long-flame " periodic kilns are generally used because they are simpler and fuel is inexpensive. Calcium oxide is a white amorphous powder, which requires the temperature of the electric arc furnace for fusion (§ 170). On 37 S INORGANIC CHEMISTRY. [§§257- being heated strongly with an oxy-hydrogen flame it emits an intense white hght (§ 13). It absorbs water and carbon dioxide from the air; as a result the chunks of lime, which are hard and solid when they come from the kUn, gradually cnumble to fine powder. Calcium hydroxide, Ca(0H)2, (slaked lime) is obtained by " slaking " quick-lime with water. Its formation is attended by the evolution of much heat. It is only sparingly soluble in water (forming lime-ioater) , but more soluble in cold water than in warm. The solubilitVj is howeA-er, sufficient to make the precipitation of this hydroxide by ammonium hydroxide impossible, for the con- centration of the hydroxyl ions of the latter is too small together with that of the calcium ions present to reach the value of the solubility product of calcium hydroxide. At red-heat it is recon- A'erted into the oxide. Mortar. — Calcium hydroxide is used in masonry. For this purpose quicklime is mixed with water and sand so as to form a thick paste, called mortar, which is thrown in between the stones. After some time the mass becomes as hard as stone; this is due to the com-ersion of the hydroxide into the carbonate by the action of the carbon dioxide of the air. The sand makes the mass porous, so that the process of hardening extends inward; the older the wall the harder the mortar. The formation of cal- cium silicate appears to plaj' only a minor role in this process. If the lime contains more or less magnesia it is difficult to slake; it is therefore less adapted to masonry purposes and is called " poor," or " lean," in contrast with the pure, easily slaked " fat " lime. Cement contains, besides lime (50-60%), principally silica (ca. 24%) and alumina (ca. 8%). It is made by burning a mixture of limestone, clay and sand. In some places, e.g. Brohlthal in the Rhine region, such a mixture occurs as " tuffstone," which yields cement directly on burning. Cement after being mixed with water sets very firmly in a short time; this is due, in all probabilit}', to the fact that on treating it with water calcium aluminate is dissolved and the solution slowly deposits a hydrous aluminate, which is much less soluble and causes the setting of the cement. At the same time insoluble calcium aluminium silicates are formed. 258.] SALTS OF CALCIUM. 379 Calcium peroxide, CaO:.-SH2(>, is deposited when lime-water is treated with hydrogen peroxide solution. It gives up oxygen on heating. SALTS OF CALCIUM. 258. Calcimn chloride, CaC'lo, is obtained liy dissolving the hydroxide or carbonate in hydrochloric acid. It can crystallize with various amounts of water. The hydrate CaCl2-6H20 forms large crj-stals. Calcium chloride is ^■cry hy<;roscoinc and is there- fore frequently used to dry gases or to absorb water dissolved in organic liquids (ether, carbon disulphide, etc.). It melts at 719° It unites with ammonia to form CaClo-SXHs; hence it cannot be used to dry this gas. TVTien crystallized calcium chloride is mixed 'uith ice the temperature falls considerably, even reaching — 4.'^..5° Such a mixture i.s called a cooling- or jreezing-mixturc and is often employed for producing low temperatures. Besides calcium chloride and ico, many other such mixtures are known; the one most frequently used is that of common salt and ice, with which a temperature of —21° can be obtained. Ice is not absolutely necessary; for instance, if solid ammonium nitrate is added to its own weight of water, a temperature of —15.5° can be produced. In order to understand why such mixtures become so cold we must recall § 237. Suppose that ice is introduced into a saturated salt solu- tion of 0°, solid salt being present at the bottom so that the liquid remains saturated. The system Kolulion + ice is not in a state of equi- librium at 0°, for the salt solution has a freezing-point much lower than 0°. It cannot therefore continue in this state, but, if it is to be in equilibrium with ice as solid phase, the temperature must sink, and this is only possible as the ice melts, by which process heat is changed into the latent condition. If enough ice is pre.'^ciit, it lan, by melting, continue to withdraw free heat from the system till the cryohydric jHjint is reached; for only at or below that point can ice and salt exist ]icrma- nently side by side. It follows, therefore, that the cryohydric tempera- ture is the lowest that can he reached by the mixture. In § 237 it was .shown, further, that there is no essential diifercnce between the two components of a solution; this is also seen on considering cooling- mixtures containing no ice. For instance, when ammonium nitrate is added to water, the solution has a freezing-point much lower tlian (1^^- Here it is the great absorption of heat in dissolving the salt, that causes the fall of temperature necessary to establish the c(iuilibrium. If this 3S0 IXORGAXIC CHEMISTRY. [§ 258. fall is to be considerable, the solubility of the salt must of course be great. la this case also the cryohydric point is the lowest tempera- ture that can be reached by the mixture. Chloride of lime is a name given to a product obtained by saturating slaked lime with chlorine at ordinary temperatures. Just what compound is formed here is not yet definitely known although the matter has been frequently investigated. There is, OCl however, much e^'idence in favor of the formula C'a, whereupon the mass becomes hard. This is the basis of the application of gypsum in the manufacture of casts, etc. The "setting" depends upon the relatively high solubility (about 1%) of this half-hydrate, on account of which it forms a solution supersaturated as to gypsum (C'aS()4-2H2<.); solubihty about 0.2' () and gypsum is deposited. Another very essential fac- tor in the setting is the filamentary character of the precipitated gypsum, a property which is entirely lacking in the case of calcium hydroxide, for which reason slaked hme does not hold together. The credit of having explained the conditions governing the exist- ence of the above-mentioned modifications as well as of having deter- mined the positions of their transition points is due to Van't Hoff. The investigation was rendered the more difficult because of the retard- ation phenomena which obscure the true situation. It was found that the half-hvdratc is to be regarded as a metastable modification, because, for one reason, the temperature at which it goes over into soluble anhy- drite is lower than that at which the dihydrate liises all its wat(n-, while in general the loss of water by hydrates proceeds stop by steji with rising 382 IXORGAMC CHEMISTRY. [§§ 258- temperature. Moreover the greater solubility of the half-hydrate, as compared -nith that of the dihydrate, is an additional reason. There is thus the same relationship here as between the metastable crystals NajSO^-yHjO and the salts Na^SO^-lOH^O and NajSO^, except that in this latter case the transformation from metastable to stable modifi- cation takes place very easily on touching the heptahydrate wdth a crystal of the decahydrate, while the half-hj^drate of calcium sulphate, even in contact with the dihydrate, retains its identity indefinitely. Calcium nitrate, Ca(N03)2, results from the decay of nitrog- enous organic substances in the presence of lime. It crystallizes with four molecules of water. The anhydrous salt deliquesces in the air and dissolves readily in alcohol. It is converted into salt- petre by potash or potassium chloride (§ 229). Calcium phosphates. — The tertiary salt, Ca3(P04)2, is insoluble in water, as is also the secondary salt, Ca2H2(P04)2, The primary salt, CaH4(P04)2, however, is readily soluble; it is employed in large c^uantities as an artificial fertilizer, under the name of "super- phosphate." This superphosphate is manufactured by thoroughly mixing ground phosphorite (or bone meal) in a cast-iron mixer with chamber acid according to the proportions of the equation Ca^(P0,)2 + 2H,S0, =CaH,(P0,)2 +2CaS0,. The mass, which is at first semi-solid, soon becomes solid, since the calcium sulphate that is formed takes up the water contained in the chamber acid to form crystals. When superphosphate is mixed with soil the primary calcium sulphate goes into solution and, since every soil contains lime, it is forthwith reconverted into insoluble secondary or tertiary phosphate. Appar- ently nothing has been gained toward "making the phosphoric acid soluble." However the phosphate is now diffused widely in the soil and is therefore much more accessible to the roots of the plants than if the soil had been mixed with tertiary phosphate only. 259. Calcium carbonate, CaCOa, is dimorphous, occurring rhombohedral as calcite and rhombic as aragonite. When the solution of a calcium salt is treated with soda, calcium carbonate is at first precipitated in an amorphous, very voluminous and more soluble form; after a short time, however, it turns to a f.ncly crj'stalline pov.-dcr. It is very slightly soluble in water, but more •259.] SAI/rs UF CALCIUM. :iSi extensively so in water containing carbonic acid, since the acid calcium carbonate is then formed. The latter decomposes when the solution is boiled, carbon dioxide escaping and crj-stalline neutral carbonate being deposited. Hardness of TT'a^er. — Almost e-\-ery ri^•er- or spring-water holds more or less lime in solution. The lime is present as sulphate or as acid carbonate. Such a water forms but little, if any, lather with soap; the fatty acids of the soap form white insoluble salts with the lime, so that water containing much lime is not good for washing. Such a water is termed hard in contrast with a water that is free or nearly free from lime, which is called soft. If the hardness is due to acid carbonate (also called "bicarbonate" of lime), it disappears on protracted boiling, calcium carbonate being precipitated. In such a case we speak of temporary hardness. In metallic boilers and similar vessels the carbonate, of lime that is deposited adheres firmly to the sides ("boiler- scale"). If the hardness of a water is due to gypsum, which is only partially removed by boiling (§ 236), it is spoken of as permanent hardness. When heated, calcium carbonate breaks up into lime and carbon dioxide. We have here a case of complete heterogeneous equilibrium (§ 71), for the susbtances are CaO and COo and the phases CaO, CaCOa and C()2. This is confirmed by experiments, which show that the concentration of the gaseous phase (the dis- sociation tension) at a definite temperature is constant and there- fore independent of the amount of each phase. Complete decom- position into lime and carbon dioxide can only occur, therefore, ■uhen the gaseous phase is removed (as in lime-burning, § 257) or when its tension is kept below the dissociation tension. On the other hand, if tlie tension of the carbon dioxide is greater than the dissociation tension, calcium carbonate cannot decom- pose. Under these circumstances it is possible to fuse calcium carbonate; on solidification it assumes a Caco-Yco v^ crystalline structure and becomes viarhlc. In the adjoining Fig. 60, let AB repre- sent the dissociation curve of calcium car- bonate in a coordinate system Pt. Only . along this cur\'e arc the three phases in „ „„ e(]uilibrium with each other; under any other conditions one of the phases disappears ■384 IXORGAXIC CHEMISTRY. [§§'259- and we enter either the region of the phases CaO+C02 or that of CaCOs + COo. GLASS. 260. Calcium silicate is chiefiy important because it is a con- stituent of almost all sorts of glass. Glass is a mixture of silicates of the alkalies with calcium silicate or lead silicate. The alkali sihcates are soluble in water, amorphous and easily fusible. The calcium silicates, however, are insoluble, very hard to fuse and frequently crystallized. By fusing both together an insoluble amorphous transparent mass of moderate fusibility is obtained, which is glass. It is prepared by fusing a mixture of clean sand, lime and soda in refractory crucibles. The properties of glass depend primarily on the quality of the materials and secondarily on the proportions used. By varying these two conditions it is easy to obtain grades of glass varying widely in fusibility, hardness, lustre, refractive power, etc. There are very many different sorts in use. Some of the most important .are the following: Soda glass {window-glass) is a soda-lime sUicate. It is readily fusible and is used for most purposes of the household. Potash glass (crown glass, Bohemian glass) consists of a silicate of potassium and calcium. It is very difficult to fuse and is therefore extensively used for chemical purposes (combustion tubes, etc.). Lead glass (flint glass) is a silicate of potassium and lead. It is softer, more easily fusible and highly refractive aiid takes on a beautiful lustre when polished. It is therefore used for optical instruments and fancy glassware ("cut glass"). Besides the substances mentioned many others are used in glass factories to impart particular properties to the glass. The addi- tion of boric acid or the partial replacement of lead with thallium gives lead glass a still higher refractive index. An admixture of alumina, AI2O3, prevents or hinders chemical utensils of glass from becoming brittle and allows the replacement of part of the alkali by lime. Certain metallic oxides form colored silicates and are therefore mixed in with the furnace charge to color the glass (cobalt, blue; chromium or copper, green; uranium, yellow- green fluorescent, etc.). The addition of bone-ash, Ca3(P04)2, iliO ] GLASS. or tin oxide gi\'es a milk}--whitc opaque glass. Tiic following- table shows the percentage composition of various kinds of glass, as determined bv analysis: Window-glas Bottle-glass Crown glass Flint glass Plate glass 7(1 (i4 74 Kit ) Na20 CaO PbO AI2O3 and re203 15 r.', ') 4 21 9 I'.i . . . i 14 31 17 H 5 "Water has in general very little effect on glass; nevertheless it attacks it somewhat. Old window-panes ha^'e a peculiar irides- cence, due to surface weathering. As it is ^-ery important in exact analyses to know how much glass can be dissolved from the utensils, careful in^•estigations have been carried out, the results indicating the following: "\Mien the glass is new a rela- tively large amount goes into solution; this amount gradually decreases in the course of a few weeks to a minimum. At the first the alkali in particular is dissolved from the surface and the resulting solution then acts as a solvent for the silicic acid. To prepare glass vessels so that they are almost wholly unaffected by water they are subjected to a jet of steam for a quarter of an hour or left for se\'eral weelcs full of water, the water being renewed occasionally. Thus there is formed on the surface a thin layer, rich in silica and lime, which protects the inner portion from the action of the water. The dissolving action of water on the alkali of glass can be readily shown by agitating finel\- powdered glass in water. The liquid at once turns phenolphthalein bright red. Glass is a tyjiical amorphous substance. Such substances are often defined as liquids with a ^-ery high internal friction and the behavior of molten glass on cooling is an excellent illustration of this definition. At high temperatures molten glass is a thin liquid; if the temperature is allowed to sink, the consistency of the glass becomes tougher, so that between the wholly liquid 3S6 INORGANIC CHEMISTRY. [§§ 260- and the wholly " solid " states, there is a continuous series of half-liquid states. As it is thus impossible to find a temperature limit to the applicability to glass of the laws of solutions, e.g. the law of diffusion, it seems rational to consider the " solid " amorphous state as liquid, in contradistinction to the crystalline state, which latter is truly solid, having very different properties from liquids. Solid solution. — This term was introduced by Van't Hoff to apply to a solid homogeneous mixture. The best example is to be found in mixed crystals, including isomorphous mixtures (§ 210). Thus, for instance, when a molten mixture of silver and gold solidifies, the components do • not separate, but solidify together in homogeneous crystals of the same composition as the melt. (See Fig. 68, III.) The term " solid solution " is applied to this and somewhat similar solid mixtures, because they exhibit some of the properties of liquid solutions, e.g., in mis- cibility relationships. Glass represents an amorphous type of solid solutions, of which the constituent silicates are the integral components, but, as intimated in the preceding paragraph, many are inclined to regard the amorphous solid solutions as pseudo- solid solutions, i.e., really undercooled liquid solutions. II. Strontium. 261. This is one of the very widely diffused elements. Clarke showed that in most of the rocks containing calcium this latter metal is accompanied by small quantities of strontium and barium. The principal strontium minerals are strontianite, SrCOs, and celestite, SrS04. Its compounds are very analogous to those of calcium. The metal has been obtained by the electrolysis of fused strontium chloride. Its specific gravity is 2.5. In its properties it corresponds to calcium throughout. Strontium oxide, SrO, is formed on igniting the hydroxide or carbonate. The temperature required for the complete dissocia- tion of the latter is higher than that for the corresponding calcium compound. The hydroxide, Sr (OH) 2 -81120, is more soluble in water than calcium hydroxide. The chloride, SrCl2 -61120, is hygroscopic, like that of calcium. It is soluble in alcohol and 20-2.] BARIUM. ;iS7 can, with the aid of the latter, be easily separated from barium chloride, which is insoluble in alcohol. Strontium sulphate is much less soluble than calcium sulphate; at 16.1° 1 part SrS04 dissolves in 10070 parts H2O (CaS04, 1 part in 543 at 15.2°). In a mixture of alcohol and water it dissolves to an extremely small extent. Strontium nitrate, Sr(N03)2, is insoluble in alcohol; this forms the basis of separating it from calcium nitrate, which dissoh-es in alcohol. Strontium salts are used in pyrotechnics because of the beauti- ful crimson color thej' impart to a flame. III. Barium. 262. This element occurs combined as barite, or heavy spar, BaS04, and as ivitluritc, BaCOs, in considerable quantities. In preparing the other barium salts it is merely necessary to dissoh'e the latter mineral in the proper acid. Barite, however, must first be reduced b}- ignition with charcoal. This can be accompUshed in the electric furnace: (1) 4BaS04+4C =BaS+3BaS04+4CO; (2) 3BaS04+BaS = 4BaO+4S02. The metal is, in this case also, obtained by the electrolysis of the fused chloride. Another method is to heat the oxide with magnesium. Barium decomposes water vigorously even at ordinary temperatures. Sp. g. = 3.75. Barium oxide, BaO, is obtained by igm'ting the nitrate or hydroxide at a high temperature. It unites very readily with water to form the hydroxide, Ba(0H)2, which is rather soluble in water (yielding hanita-mitcr) , and crystallizes from the hot solu- tion on cooling in pretty laminae, which contain eight molecules of water. Barium peroxide, Ba02, forms on heating the oxide in a cur- rent of oxygen or air. When it is introduced into dilute sulphuric acid, barium sulphate is precipitated and hydrogen peroxide left in solution. If baryta- water is again added, the hydrate Ba02 • 8H2O crystallizes out. Barium chloride, BaCl2 -21120, is not hygroscopic like the chlorides of strontium and calcium. The nitrate crystallizes anhydrous. 388 INORGANIC CHEMISTRY. [§§■262- Barium sulphate, BaS04, is characterized b}- an exceedingly small solubility in water and acids; at 18.4° 1 part dissolves in 429,700 parts H2O. It is used as a filler and as a pigment under the name of " permanent white," or blanc fixe. Barium carbonate yields carbon dioxide only at very high temperatures, prolonged heating at 1450° being required for complete decomposition. SCMMARY OF THE GROUP OF THE ALK.\LINE EARTHS. The following small table summarizes the phj-sical properties of the elements of this group: Be Mg Ca Sr Ba 1 Atomic weight Specific gravity .-Vtomic volume 9.1 1,64 5.6 white 24.32 1.75 13 S white 40 09 1 5S 25 . 2 white 87.62 2.5 34.9 white 137.37 3.75 36.5 white As to the specific gravity we observe that only in the cases of Ca, Sr and Ba is a steady increase noticeable. In respect to the chemical properties, it has already been remarked that these elements act only as bivalent; all compounds of the groujD therefore have the same formula type. In the solu- bility of the sulphates a gradual decrease is to be observed with rising atomic weight. Just as in the first group three elements K, Rb, Cs, exhibit a particular kinship, so here calcium, strontium and barium are closely related in their properties, while the two other members of the group are unlike them in many respects. Beryllium displays analogy with aluminium in certain points just as lithium does with magnesium. 2(i3] SPECrh'OSCOPY. 3S!I SPECTROSCOPY. 263. If the light from an ordinary gas flame or the \\'elsi3ach incantlescent Hght is broken up by a prism, there is projected a continuous series of perfectly blended colors from red throufiii yellow, green, and blue to violet. This phenomenon is called ;i spectrum, and since it is unbroken, a continuous spectrum. W r have previously remarked that the luminosity of a gas-flame is due to incandescent solid particles of carlwn. It has been found to be a general rule that incandescent solids give a continuous spectrum. With incandescent gases it is different. If, for instance, we split up the light from a Bunsen flame, in which salts of sodium calcium or other metals are volatilized, we see only a few narr^v bands of light in certain j)laces, the rest of the spectrun^ being dark. This is termed a line spectrum. Everj- element lias its own peculiar spectrum lines. If the spectrum of the incandescent ■vapors of a mixture of elements is carefuUj' examined, it is found Fig. 61. to contain all the characteristic hues of each element. Since it is only necessary to volatilize extremely small amounts of sub- stances in order to show their lines, it is readily seen how important the spectrum-analytical methods introduced by Bunsen and KiRCHHOFP must be. For the examination of spectra a number of instriunents have- been constructed, varying according to the particular object in view. For chemical analysis the apparatus of Vogel or that of John Browning is now very generally used. It is a small direct- vision spectroscope which gives a very bright spectrum and has a sufficient dispersion. At the end B (see Fig. 61) is the slit whicli can be made narrower or wider by turning the rim D. The small mirror m serves to throw light through the hole P on to an auxil- iary prism, in order to compare the spectrum of the light which is to be analyzed with that of a known source. At the left end 390 INORGANIC CHEMISTRY [§§263- is the ocular through which the spectrum is seen. For further information text-books on physics should be consulted. In order to examine the spectra of metals it is necessary to convert the latter into the form of vapor at a high temperature. There are different ways of doing this. One is to introduce salts of the metals into a colorless flame by means of a platinum wire. The heat dissociates halogen salts and in the case of oxysalts con\-erts them into oxides, which are reduced to the metallic condition by the hot gases of the flame. This method is very satisfactory for some elements, e.g. those of the alkali and alkaline earth groups, when there is plenty of material. In other cases a flame spectrum of this sort is not so good as a spark or an arc spec- trum, for with the latter it is possible to detect with accuracy extremely small amounts of a substance. Other advantages of the latter spectra are their greater light intensity, the greater con- venience in execution, and the like. Moreover, at the high tem- perature here prevailing most elements exhibit spectra which cannot be obtained with the gas-flame. A spark spectrum can be obtained in a very simple manner, thus: Into the bottom of a little glass cup (n, Fig. 62), about 15 mm. wide is fused a platinum jC^^^^'^ "wire, which ends in a tube g containing mercury t and is thus connected with the negative pole of an r — — t induction coil; it is incased in a conical capillary \ 1"'^ tube X, beyond which the wire projects about 0.5 mm. d\) At the opposite end is the positive electrode in the form of a platinum wire, which, with the excep- tion of the short end d, is fused into a glass tube; the latter is fltted into the cork a. If some of the salt solution is poured into the cup about half way up the negative electrode, the liquid is drawn up to the end of x bj^ capillarity and every spark volatilizes a tiny portion. In this way there is no loss of material and the sparks are very uni- form, so that the observation of the spectrum can Fig. 62. be continued at length. For the study of the spectra of substances which are gaseous at ordinary temperatures the P l u c k e r-H i t t o r F (Geissler) tubes are used (Fig. 63). The gases are sealed up in them in a very dilute condition. On connecting one of these -'l>4.] SPECTROSCOPY. 31)1 with the poles of an induction coil, the whole tube is illuminated most intensely in the narrow portion. This part is placed ver- tically in front of the slit of the spectruscopc. Some substances have the pn)pert>' of al)S(irbing certain colors and transmitting others. If the solution of such a substance is placed before tlie slit of a spectroscope and the light of a contin- uous spectrum allowed to pass through it, dark bands or lines are observed in the spectrum. A number of substances have very characteristic absorption spectra. 264, The spectroscope is one of the most delicate means we have of detecting many substances. This is readily seen on con- sidering how small an amount of the substance under examina- tion is \'olatilized by the sparks. We arrive at numbers like Fig. 63 0.3X10"^ mg. sodium, for instance, as the least amount that can be detected. It has thus been possible to disco^•er elements which occur only in company with large amounts of others and would therefore have been ^■ery difficult to find in the ordinary way. BuNSEN and Kirchhoff themselves found ca'sium and rubidium in this way in Diirkheim mineral water. In order to obtain these elements from it in the form of chlorides, it was necessary to evaporate 44,000 kg. water, whicli yielded 16.5 g. of a mixture of the chlorides. Other very rare elements which were discovered by spectrum anah'sis are thallium, indium, gallium, ytterbium and scandium. The spectra of the elements differ greatly in appearance, as may be seen at once from Table II (Frontispiece). The numbers indicate wave lengths of light expressed in hundredth microns (10~5 mm.). Certain metals, such as sodium, thallium and indium, exhibit only one distinct line when their flame spectra are examined with a spectroscope hke the one described above. If a sparking current or an electric arc is employed for the vola- tilization of the substance and the spectroscope is one giving strong dispersion, many more lines become visible. It is further found on photographing spectra that there are still more lines in the infra-red and ultra-violet portions, which are invisible to the eye. Present-day spectroscopic studies deal, therefore, almost exclusively with carefully prepared photographs. 392 IXOROAXIC CHEMISTRY. [§§--^04-- The number of .spectral lines increases rapidly as we proceed to elements of the higher groups of the periodic system. While lithium, sodium and potassium give 20, 35, and 41 lines, re- spectively, the spectrum of barium contains 163 lines and that of iron more than 5000 lines. Among these lines there are certain ones which, in virtue of their position (color) and intensity, are specially characteristic of an element, like that of the yellow line in the case of sodium, the green line of thallium and the blue lines of indium. For purposes of identification of such elements these prominent lines are generally observed directly in the spectral apparatus. Nitrogen is an example of a substance that gives a band spec- trum when it is examined in the manner described in § 263. 265. The position of the spectrum lines was formerly indi- cated numerically according to an arbitrary scale. Now it is expressed with the aid of the wave length;^, 10"'' being taken as a unit and the unit being called the Angstbom unit after the physicist who introduced it. The wave length of the sodium Di line was found to be expressed by 5896.16 such units. The visible part of the spectrum comprises the wave lengths of about 7500-4000 A.U. Thanks to the researches of Rowland, Michelson, Kayser and RuNGE, and others, the wave lengths of a very large number of spectrum lines have been determined with great exactness, so that one is encourag'ed to attack the question whether in the apparently very promiscuous distribution of lines in the spectra there is such a thing as order. Balxier was the first to show that this is the case in the hydrogen spectrum. The formula X = A-2 — -, where A is a constant (3646.13) expresses the wave lengths }. of the lines of the spectrum of the element with very close approximation, provided m is substituted by consecutive whole numbers begin- ning with 3. The spectra of other elements have been examined for similar regularities, chiefly by Rydberg and by Kayser and Runge, and it has been found that the regularities are in all of the cases more complex than for the hydrogen spectrum. It would lead us too far to enter upon a discussion of these questions, which 265.] SPECTROSCOPY. 393 properly belong to Physics, but a few of the interesting results are worth mentioning here. Rydberg, who has devoted particular attention to the spectra of the alkalies, introduced into his formulae the reciprocal of the wave length, the oscillation frequency u, which represents the number of wave lengths per centimeter. In the spectra of the alkalies he found three series of lines, whose oscillation fre- quencies can be expressed by the formula -Vo In this formula Nq is a constant having the same value for all these metals and all the series; no and ,u, however, are two constants that have different values for each series of each metal. For m we substitute again consecutive integers, as in Balmek's formula. This last formula is, moreover, a special case of that of Rydberg, since Balmer's formula can be trans- formed into 1 , -Vo' (where no' = A and No' = 4:A), into which Rydberg 's formula is also transformed when fi = 0. The values of the constants of these different series were found by Rydberg to have still further definite relationships. The formula of Kayser and Runge isy = A + Bm,-^ + Cm~-^, A in which A, B and C are constants and m consecutive whole num- bers. It represents the wave length of the lines in many cases more faithfully than does the formula of Rydbkrg; however, there is no relation between the constants A, B and C of the different series. • The spectral lines of the alkalies also exhibit the pecuharity of consisting of double lines (doublets) or triple lines (triplets), the wave-length differences being constant for each series. Similar series of lines, whose oscillation frequency can be expressed by one of the above formulae, are found in the spectra of some of the elements. In the spectra of many others, how- 394: IXORGANIC CHEMISTRY. [§§ 265- ever, they are lacking. In their place we find in the spectra of lead, tin, arsenic, bismuth and others a constant difference between the oscillation frequencies of a considerable number of their lines. Such investigations as these are prompted by the notion that a knowledge of the laws which govern the distribution of the spectral lines of one and the same substance on the one hand, and the variation in the distribution from substance to sub- stance on the other hand, would throw some light on the nature and kinetic condition of the atoms. With the aid of spectroscopy it has been possible to deter- mine what elements are present in the heavenly bodies. When light from the latter is passed through a prism, line spectra are obtained and these lines correspond in position to those of terres- trial elements. The composition of sunlight has been especially the object of a most extensive study. The spectrum of that body contains numerous black lines, known as Fhaunhofee lines. The theory of this phenomenon is explained in Optics. By comparing the Fraunhofeb lines with the spectra of ter- restrial substances it has been found that the sun's atmosphere contains chiefly Fe, Na, Mg, Ca, Cr, Ni, Ba, Cu, Zn and H (the latter in enormous quantitj-). Moreover, for 450 lines of the iron spectrum there are found to be corresponding dark lines in the sun's spectrum. On the other hand, the solar spectrum displays countless lines which are not j'et identified in terrestrial spectra. We are led to presume that many of the elements to which these lines are due will also be revealed on the earth by more careful research, especially when we consider what a small part of the earth is known (see footnote, p. S}. This presumption has been strongly confirmed by the discovery of liclium (§ HI). The principal line of the latter, D, — so termed because of its proximity to the double D-line (D^D^) of sodium — was observed in the spectra of many fixed stars as well as in that of the sun before the element itself was identified on the earth. Helium was thus discovered in the stars before it was found on the earth. It is a striking fact that it occurs in exceedingly large quantities in the fixed stars (according to speotrometric ol )servations) while there is apparently only a very small amount of it on the earth. 266. THE UXITY OF MATTER.] 395 THE UNITY OF MATTER. 266. The notion that all substances are derived from a single original substance and that the variety that we observe in the material world is merely the result of a difference in arrangement and form of the smallest particles has long been prevalent. Even the old Greek philosophers had a fondness for it. However, the rise of experimental investigation was not very conducive to the idea. Boyle (1626-1691) introduced the concept element in its present form. According to this concept all substances are to be regarded as elements, which, with the means at our command, cannot be further resolved into dissimilar components. The sub- sequent de\-elopment of chemistry has shown that the number of these elements is rather large. Notwithstanding that the idea of a primordial substance lacked substantiation and was more or less discredited, it was by no means rejected, for we have been expressly reminded again and again that the substances which chemistry regards as the simple substances are only classed as elements conditionally; the pos- sibihty always remains that a so-called element may be found capable of division into dissimilar components, as has often ac- tually been the case. Impossible as it was to deny the existence of a primordial substance, the researches of the 18th, and a large part of those of the 19th, century brought to light nothing to support the idea. Not until the discovery of the Periodic System did the question again demand serious attention. This discovery was the first to supply an experimental basis for the assumption of a primordial substance (§ 220). The striking dependence of the properties of the elements on the periodic functions of their atomic weights, which finds its expression in this system, leads of itself to the thought of a fundamental substance, of which the simple sub- stances called elements may be said to be polymers, incapable of resolution by the means at our command. Another argument for the divisibility of the elemental atoms is contributed by spectroscopy. In order to explain the line spectra exhibited by many elements, we assume that the move- ments of the atoms give rise to light vibrations of definite wave 396 IXORGAXIC CHEMISTRY. [§ 266- length, which are perceived bj- us in the spectral lines. However, since the spectrum of a single element is extremely complicated, we should have to assume that the atoms engender very complex movements. The simphfied hypothesis was -then offered that it is not the entire atom but smaller particles of which the atoms are composed, that give rise by their vibrations to the different spectral lines. The physical investigations of the last decade have furnished substantial reasons for believing that the chemical atoms are not in realit}- the ultimate particles of matter but that they are divi- sible into particles approximately 2000 times smaller than the hydrogen atom. These particles, carrying with them, as they do, verj- strong electrical charges, are called "electrons." We cannot in this book do more than indicate the main observations and inferences. The hypothesis of electrons is the result of the study of cathode rays in connection with the below-mentioned investiga- tions of radio-active elements. Cathode rays are generated when the discharges of an induction coil are sent through a rarefied gas. It is assumed that from the cathode there are projected particles wth strong negative charges, electrons, which are propagated with a velocity of several thousand kilometers per second. Ac- cordingly, the cathode rays consist of a stream of these electrons. Measurements of the mass of an electron have shown that it is about -j-oViF of that of a hydrogen atom. The electrons are the same, whatever gas is contained in the apparatus and whatever electrodes are used, so that we are evidently deahng with a decom- position of the atoms into their ultimate components. The anions, according to the electron theory, consist of atoms and one or more electrons; the chlorine ion, for example, con- sists of chlorine and an electron, which latter is represented by 6. The cations are formed from the atoms by the release of one or more electrons. We may thus write : C\+d = CY and K-^ = K-. The ionization of potassium chloride in water can be represented l)y the equation: KClaq= (K-^)^+ (CH-^)aci. Certain physicists presume to be able to go a step further, Some very remarkable investigations of Rowland have shown 267.] RADIO-ACTIVE ELEMENTS, 397 that a moving electrically charged conductor exerts the same effect as an electric current, so that the latter may be regarded as consisting of very swiftly propagated, discrete electrical par- ticles. Electricity, would thus have an "atomistic" structure. Faraday's law points in the same direction. A''cording to this law the charges on the ions are either equal to or a multiple of the charge on a hydrogen ion. 'Fractions of the charge are not found. This is a very significant fact. Just as for the explana- tion of the analogous laws of Dalton we have assumed that the elements consists of atoms, so we cannot avoid inferring from Faraday's law that electricity is divided into discrete elemental particles, which are to he regarded as "atoms of electricity." Furthermore, it lias been shown that induction and other pheno- mena proceed just as if electricity had mass, for electricity has the same properties of inertia as ponderable matter. The conse- quence is that matter is identified with electricity and the elec- trons are no longer to be regarded as electrically charged mass par- ticles but as electrical charges themselves, without a material body. These hypotheses would not only unify matter but would also dispel the time-honored notion that energy and matter are distinct, RADIO-ACTIVE ELEMENTS. 267. Becquerel (liscovci-ed that uraninium emits a peculiar sort of ra}'s which are propagated in a straight line and act on a photographic plate, but are not reflected, refracted, or polar- ized. When gases ai-e traversed by them the gases become electrical conductors. Xow when uraninite (or pitchblende, a uranium-l)eai-ing mineral of very complicated composition) was investigated as to its ladiation the strange fact was brought out that the radiation of the mineral is 4. .5 times as powerful as that of its constituent metal, uranium, although only .50''^ of the mineial is uranium. Uraninite must therefore contain one or more substances ha\'ing a stronger radiating power than uranium. We ai'c indebted principally to the gifted couple, M. and AIme. CrtRiE, for the discovei-y that the emission of these special rays, or the radio-activity, is due to the pi-es(nice of very small amounts of elements, hitherto unknown and of very sur- prising properties. 398 IXORGANIC CHEMISTRY. [§§ 267- The onl}' means of control in the separation of these elements from the other compounds in uraninite after the removal of uranium was to measure the radio-activity of the product ob- tained in each operation. This was accomplished by measuring t"he conductivity of a layer of air that was exposed to the rays. Thus after numerous chemical operations the active substance was concentrated more and more. This method is comparable to that employed by Bunsen and Kirchhoff in isolating rubid- ium and caesium from the Diirkheimer mineral water, where the spectroscope (§ 232) indicated the progress of the concentration of these elements. However, the measurement of radio-activity is many thousand times more sensitive than a spectroscopic examination. Were it not for this fact, the discovery of the radio-active elements would have been impossible, because they occur in such extremely small quantities. For example, 2000 kg. uraninite residues from Joachimsthal yield only about 0.2 g. radium chloride. Radium is the best known of these elements. It is the only one that has boon isolated and whose compounds have been prepared in the pur(> state. Its spark spectrum has three very bright lines in the blue and violet and accordingly the Bunsen flame color is carmine. In its chemical behavior it shows close analogy to barium; it is separated from the latter element by fractional crystallization of the bromides, radium bromide being more difficultlj' soluble than the corresponding barium salt (this is true for all the respective salts of the two elements). With the aid of the spectroscope it can be determined whether the salt is entirely free from barium bromide. The atomic weight of the radium thus purified was found to be 226.4, which could not be raised by further fractional crystal- lization. With this atomic weight radium fits exactly into the second group of the periodic system. All radium salts are lumi- nous and excite a large number of substances, such as barium platinocyanide, BaPt(CX)4, uranyl sulphate, precious stones, and the like, to powerful fluorescence. It similarly affects the dia- mond. Genuine diamonds can thus be distinguished from imi- tations. The radio-activity of the pure bromide is about a million times that of uraninite. Mme. Curie and Debierne succeeded in 1910 in isolating the § 297.] RADIO-ACTIVE ELEMENTS. 399' element itself. Thoy electrolyzcd a solution, using a mercury cathode, and obtained a radium unialgani, from which the mer- cury was distilled oiT in a current of liydrogen. Radium is a white metal, melting at 700° Even as low as this temperature it volatilizes appreciably. It is attacked by the air and decom- poses water vigorously. Besides uranium and thorium the most important radio- active elements are polonium, actinium, ionium, and radio- thorium. Polonium is precipitated in a numlDcr of reactions with bismuth; liy hydrogen sulphide, as well as when the basic salts of bisnmth ai'c precipitated b)- water; stannous chloride precipitates it in the same way as mercury and tellurium. It is also deposited on a rod of silver or bismuth when one of these is immersed in a solution containing polonium. The radio- activity of polonium is about a thousand-fold as great as that of radium. From 1.5 tons of pitchblende Marckwald could only obtain 3 mg. polonium salt, still somewhat impure; so that polo- nium also sui'passes radium considerablj' in scarcity. Actinium occurs with the rare earth metals, particularly lanthanum, and can l)e partiall}-, though unsatisfactorily, separated from them In- fractional crystallization of the manganese double nitrate. For ionium see belo^\-; for radiothorium see under thorium. The rays emitted by radium preparations are of three sorts and are distinguished as a- ,3- and ;--rays. Quantitatively the first are predominant. All of them have the above-mentioned properties in common; they differ, however, in their penetrating power and in their behavior in the magnetic field. The a-rays are not very penetrating and are only slightly deflected in a strong magnetic field. A sheet of aluminium foil 0.1 mm. thick almost entirely stops their passage. Moreover they are com- pletely absorbed by a layer of air a few centimeters in thickness. The /?-rays are strongly deflected in a magnetic field and consist of rays of various but greater penetrating power; some kinds of /?-rays can even pass through an aluminium plate 1 cm. thick. The ^--rays are scarcely deflected at all and go through obstruc- tions with ease, several centimeters of lead being insufficient to .stop them; they form only a small part of the total radiation. The interesting thing is that these rays are analogous to those generated by electric dischaiges in highh- rarefied gases. The 400 IN ORG AX IC CHEMISTRY [j 200- ;9-rays are to be regarded as cathode rays of great velocity. They consist of negative electrons which are propagated with very great velocitj', some almost with the velocity of light (300,000 km. per sec). From the deflection which they undergo in an electrical field and a magnetic field of known intensity their mass is calculated to be (as in the case of the cathode rays) about 2^,i„„ of that of a hydrogen atom. The velocity of these electrons can also be calculated from the same data. Their (>normous velocity explains the great penetrating power of /?-rays. The a-rays resemble a sort of radiation which is also obtained by discharging electricity in a rarefied gas, viz., the canal rays of Fig. 64. — Effect of a Magnetic Field on the a-, /?-, and )--rat8. Goldstein. They behave as positively charged projectiles hurled at a great velocity (about -Jg- that of light). Their mass is aliout equal to that of a hydrogen atom, or much greater than the mass of the projectiles formed by the /J-rays and the cathode rays. Their greater size and relatively small velocity explains their slight penetrative power. In the light of the more recent investigations they appear to consist of helium atoms bearing two positive charges each, or, more specifically, having lost two electrons. The ^--rays are analogous to the X- or RoiSNTGEN-rays. These proceed from a metal plate which is placed in the path of cathode rays; they do not consist of a stream of electrically charged particles, but are regarded as a form of wave motion of the ether, which originates when electrons are projected with great ^'elocity against a solid body. The manner of detecting -the various sorts of rays follows § 267.] RADIO-ACTIVE ELEMENTS. -101 readily from the above description of their properties. Use can be made, for example, of their dissimilar penetrative power. Their .separation in a maf^netic field is diagrammed in Fig. 64. While the ;--rays suffer no deflection, the a-rays arc deflected to one side, the i^-rays very much to the opposite side. According to what has been recited in the preceding sections, we are to look upon tliese forms of radiation as evidencing a spon- taneous decomposition of the atoms of the radio-active elements. The decomposition is accompanied by a very considerable evolu- tion of heat. One gram of radium gives off about US g.-cal. per hour; for this reason radium salts have a highoi- temperature than their surroundings. Even cooling with liquid hydrogen (— 2.io°) does not stop this evolution of heat. The magnitude of the heat effect is more apparent upon comparison ^with other caloric effects attending chemical I'cactions. We now assume that the heat evolution in the decomposition of 1 g. radium is about 10^ g.-cal. On the other hand, the formation of 1 g. water from its elements evolves IXlO-^ cal., so that the first-mentioned process gi\-('s ofT 2.")0,000 times more heat than the second. The spontaneous decomposition of radio-active substances is accompanied by other phenomena. Every substance that is brought into proximity with a radium salt acquires a temporary, or induced, radio-activity, i.e., it emits the same rays as radium itself. This induced radio-activity is best observed on putting a radium salt in an enclosed space. The enclosing walls, as well as all bodies within the space, become active. It is not the radium rays that cause this effect, for a radium salt in a sealed tube emits rays without exciting any radio-activity. Rutherford discovered the cause of this phenomenon by the observation that there is a constant outflow from radio-active substances, which outflow he calls emanation. Since bodies with induced radio-activity give out rays that are identical with those of radium itself, these rays can be regarded as transformation products of the emanation of radium. Emanation behaves in many respects as a gas; it diffuses from one vessel into another, follows the law of Boyle in its compres- sion, can be condensed by cooling with liquid air and volatilized again if the temperature is allowed to rise. Neither physical nor 402 INORGANIC CHEMISTRY. [§ 267. chemical agencies are able to alter emanation. It is indifferent to temperature variation between —180° and 500°, is not absorbed by concentrated acids or alkalies, and can be conducted without change over hot copper oxide. It has the properties of a gas of the argon group. Ramsay has lately succeeded in preparing radium emanation in somewhat larger quantities. He calls this element niton. It is found to be a water-clear liquid with a specific gravity of about 5 and a boiling-point of —62° under atmospheric pressure. In glass vessels it is very highly fluores- cent. Emanation and induced radio-activity must be considered as intermediate stages in the complete disintegration of the radium atom into the above-mentioned radiations. How- ever, other substances, part of which have not been further studied, are formed simultaneously. One of them is pretty well known, viz., helium. Ramsay and Soddy have demon- strated that helium is formed in the spontaneous decom- position of radium emanation. The maximum quantity of emanation that could be obtained from 50 mg. radium bromide was conducted by them with the help of an oxygen current into a U-tube cooled l)y liquid air and the U-tube was then evacuated with a pump. A \':u'uum tube which was fused on to the U-tube showed no traces of helium after removal of the liquid air. The spectrum appeared to be that of an unknown element — presumably emanation. After the apparatus stood four days the helium spectrum appeared. This phenomenon explains the mysterious persistent occurrence of helium in radium-bearing minerals. The law governing the rise and decay of radio-active sub- stances is the same as that for a unimolecular reaction. As we have seen in § 50, the velocity S of such a reaction can be repre- sented by the equation : if we let C stand for the concentration and K for a constant. With the aid of higher mathematics this equation can be trans- formed into : C=Coe-^', J 207.] RADIO-ACTIVE ELEMENTS. 103 where Co is the initial concentration and e the base of natural logarithms. The logarithmic form of the equation is : ,C •-0 This same equation holds, as above stated, for the velocity of decomposition of a radio-active substance; in that case, however, we understand by C the intensity of radiation. This magnitude can be determined electrometrically. If a radio-active substance changed into only one new sub- stance, the phenomenon would be very easy to represent graph- ically; for upon plotting the time on the abscissa axis and the logarithm of the activity on the ordinate axis the phenomenon would be represented by a straight line. But when the substance .■1 is converted into another active substance B, and this again into a new acti\'e su))stance and so on, the situation becomes much more complicated. A graphic representation with the same co-ordinates as before would no longer 3'ield a straight line, but a rather complicated curve. Xc\'ertheless, it has been found possible to resolve these experimental curves and to calculate with certainty the number of active substances which participate in the transformations, as well as their constant K. This is not the only method of ascertaining the number and kind of the intermediate products. We can often distinguish the individual substances involved, by a study of the kind of radiation given off, certain of the substances emitting only «-rays, others only ;3-rays, and still others a mixture of all three ra}'s; indeed there are some of the substances which emit no rays at all. In some instances these active substances have been actually separated by physical or chemical means. Certain of the sub- stances are found to be gaseous; others form a deposit on solid bodies. The gaseous substances can be condensed by cooling. The best way to characterize the various radio-active sub- stances is by the exponent K of the above equation; this constant is to a large degree independent of temperature and pressure; which is not true of ordinary reactions. Frequently, however, 404 IXORGAXIC CHEMISTRY. [§ JeZ. use is made of another magnitude, related to K, viz., the period of half decay. If in the equation it K ' we take Co" 2' we have the case where the intensity of the radiation has decreased to just half. Solving for t, we have K" K ' as the length of time necessary to reduce the intensity to half. It is this magnitude t that often serves as a characterizing constant instead of K. As a result of such observations and determinations a series of transformations of radium has been worked out, as given in the table on page 405. Polonium, as we see, is brought into relationship with radium. The transformation of polonium into lead, however, is far from being established. In the other (parental) direction radium is related to uranium. Since the half-decay period of radium is about 1760 years and the age of the solid earth-crust is counted in millions of years, the radium in the earth would long since have disappeared, if it had not steadily been re-formed. It is now definitely estab- lished that uranium, a substance much slower in its transforma- tions, is the parent substance of radium. One of the most signifi- cant e\-idences of this is that in the various uraniferous and radium-bearing ores the ratio of uranium to radium is very nearly the same. In order to prove this, Boltwood and Strutt determined the uranium content of the ores by ordinary analytical methods and the radium content by collecting the emanation evolved on dissolving the ores and measuring its activity with an electrometer. Since the decomposition of uranium proceeds much more slowly than that of radium, the constant relationship shows that the radium is formed from the uranium. Indeed, § L>b7.] RADIO-ACTIVE ELEMENTS. 405 Transformation Products. Physical and Chemical Prniiortics Time of Half ' Decay IlMy.. Effective Limit of Radiation in the Air. Uranium i Uranium A' Ionium i Radium - 4.6X10' yrs. a — — 21.5 days i^.r — Produces Ra 5 10' yrs. «■.;? 2 X7 cm. At. wt., 22(i..'). rhaiactpris- 1 i (Xiton) i Radium A i Radium B tic. spectrum; produces helium continuously At. wt., < 2.22. Inert, con- densable gas Deposit of induced activity 1760 yrs. 3 St) days 3 min. a a a 3 .5 cm. 4.23 cm. 4 83 cm. Same 26.7 min. Slow _ I Radium <\ i Radium Cj i Radium D i Radium £, i Radium E^ i Radium F or Polonium i Lead(?) Same; separation l^y elec- trolysis In. .5 min. 1-2 5 min. ;'3-rays a= 7.06cm. One of the constituents of radiolead. Resembles lead 12 yrs. none — closely in properties Separable from Ra^ by elec- trolysis 6.2 days 4.8 days 3 — Sulphide insol. in acids; pre- cipitated with basic Bi- salts; pptd. by SnClj 140 days a 3 86 cm. SoDDY finally succeeded in showing that solutions of most care- fully purified uranyl nitrate came to contain radium in the course of three yeans. Furthermore, it is e(]ually well established that there are intermediate products between uranium and radium; in other words, that uranium is not the direct parent of radium. One of these intermediate products is ionium. It is closely allied chemically to thorium. It emits a- and ;9-rays, of which the 406 INORGANIC CHEMISTRY. [§§ 267- former are characterized bj' an especial!}' feeble penetrative power, their effective Hmit in the air being less than 3 cm. The percentage relationship between ionium and radium in the different ores is practically constant. The most interesting property of ionium is that it can produce radium. As for radium and uranium, so for thorium and actinium, series of successive decomposition products have been worked out. Chemical Effects of Radioactive Substcmces. — A'arious chemical reactions are brought about by the influence of radioactive substances. Among the many which ha\'c been observed we maj' mention the conversion of oxygen into ozone and of yellow phosphorus into red phosphorus, the decomposition of iodic acid and the dissociation of water into it.s elements. An aqueous solution of a radium salt is constantly giving off slight amounts of detonating-gas, amounting to 0.6 c.mm. per day per gram of radium. Furthermore, these rays have the property of developing a strong color in different substances, such as glass, porcelain, and the alkali salts; the same color can be produced bj'' cathode rays. The skin is also attacked by these rays. Occurrence of Radioactive Substances. — This is by no means limited to thorium- and uranium-bearing minerals. Indeed, it has been shown that small quantities of radioactive substances occur very extensively in nature. A charged electroscope gradually loses its electrification in the air, — a phenomenon which is traceable chiefly to the; ions in the air. Sea-water also contains slight traces of radioactive substances; so does the earth proper. Many springs that come from considerable depths are rich in radium emanation; in this respect the waters of Gastein (Austria) and Yellowstone Park are particularly noted. Radioactivity of Other Elements. — The radioactive decay of actinium emanation proceeds SXIO^^® times faster than that of uranium. The question at once arises whether there may not be other substances with a rate of decay very much slower than that of uranium. We could then conclude that there is, after all, no essential difference between the radioactive and the other elements, but that all of them suffer decay, even though in most cases the decay is so slow as to escape observation. To be sure, this is by no means proven, but there are two reasons for such 268.] Z/,VC. 40 1 a hypothesis. In the first plac(^ it would explain why so many pairs and groups of elements are found occurring together in nature: niobium and tantalum, for instance; selenium and tellurium, the platinum metals; the rare earths. The two last mentioneil are groups of elements whose properties are quite as closely allied as those of the decomposition products of radio- active elements. Secondly, a slight radioactivity has been detected in potassium and rubidium. So far as we know, the radioactive transformations are irre- versible. We can only stand by and look on; we can neither produce nor stop them. If these changes should prove to be a general property of matter, it would mean that all matter is engaged in slow decay. Inasmuch as the transformation can- not be made to go backwards, the inference could be drawn that the univei'se was constructed out of a primordial substance by some sort of a creative act. ZINC. 268. The most important zinc minerals are calamine {'H.2Zn-2tiiC> 5) , smithsonite (ZnCOs), sphalerite, or blende (ZnS), and various oxides. The principal locaUties are Silesia, England, Belgiu:n Poland, and more recently, certain districts in the United States notably southwestern ilissouri. To obtain the metal the ores a;-e roasted — the gas (SO2) from the sulphide ores is con- veiteu into sulphuric acid — yielding zinc oxide. In the older processes this is mixed with coal and heated, forming carbon monoxide and zinc. The latter distils over and collects in the receiver together with a fine gray powder, zinc dust. This "dust ■ is a mixture of zinc oxide and zinc powder and is frequently used in the laboratory as a vigorous reducing- agent. The metal is bluish-white and has a specific gravity of 6.9-7.2. At ordinary temperatures it is brittle, but at 100-150° it becomes softer; it can then be beaten into plates. At the same time the specific gravity rises to 7.2 and the metal becomes firmer. At 200° it again becomes brittle and can be easily pulverized. It molts at 418° and boils at 920°. The metallic vapor has a specific gravity of 33. S (H = l); hence its molecular weight is 67.6. Since the atomic weight, as deduced from Dulong and Petit's 408 INORGANIC CHEMISTRY m 268- law, is 65.4, the molecule in the vaporous state can contain onlj^ one atom. The same is true of the related metals cadmium and mercury. Zinc is ■ permanent in the air, since it becomes firmly coated with a protecti\e layer of oxide. Zinc dust decomposes water. When heated to boiling in the air the metal burns to zinc oxide, producing an intensely bright light. It is dissolved very easily by hydrochloric or sulphuric acid with the evolution of hydrogen; it is an interesting fact, however, that when a piece of absolutely pure zinc is placed in either of these acids no hydrogen is generated. Tf this piece of zinc is brought in contact with a platinum wire, effervescence begins at once, not from the surface Zn Zri Pb Fig. 65. Fig. 66. of the zinc, however, but from that of the wire, and zinc goes into solution. Written in ions the process is Zn + 2H" =Zn" + 2H, and its explanation is just the same as that given in § 203, for the formation of a "lead tree." In this case also the zinc drives cations into the solution with great force, itself thus assuming a negative charge, with which hydrogen ions can be discharged. The only difference seems to be that these hydrogen ions discharge themselves at the platinum instead of at the zinc. However, this difference 'is not real, since in the case of the lead tree the fresh portions of lead are deposited on the outermost parts of it. The perfect analogy is made still clearer by a somewhat modified form of the experiment: When, on the one hand, a plate of amalgamated zinc and one of platinum are connected by a metallic wire and dipped in dilute sulphuric acid (Fig. 65) hydrogen is evolved from the platinum plate and when, on the other hand, a plate of amalgamated zinc and one of lead are similarly connected and dipped in a dilute solu- tion of lead nitrate (Fig. 66) lead crystals are deposited not on the 269.] ZINC. 409 zinc but on the lead. In both cases the negative charge of the zinc goes through the wire to the other plate, on which the ions of hydrogen, or lead, as the case may be, can discharge themselves. Metallic zinc — often called speller in commerce — has numerous uses. For instance, zinc plates .are very extensively used for roof- ing. Iron is frequently coated with zinc to prevent rusting; it is then known as galvanized iron. Further, zinc is a constituent of many alloys, e.g. brass (§ 242). 269. Zinc oxide, ZnO, is usually prepared by igniting the basic carbonate. On being heated it turns yellow; on cooling, the original wliite color returns. It is employed as a pigment under the name zinc ivhilc, or Chinese while. Zinc hydroxide, Zn(0H)2, is precipitated by alkalies from the solution of a zinc salt as a white gelatinous mass, soluble in the alkalies as well as ammonia; however, the reason is different in the two cases. In the presence of alkalies zinc hydroxide behaves as a weak acid; it forms ZnO^" anions and the cations 2H\ which yield a salt Zn(0K)2 with the alkali in the ordinary way (§ 66). When treated with ammonia, however, a complex zinc-ammonia ion is formed, which is soluble. Zinc chloride, ZnCl2, can be obtained by heating zinc in a cur- rent of chlorine or by dissolving zinc in hydrochloric acid and evaporating the solution. In the latter ease some oxychloride is formed, however. Zinc chloride melts on heating and distils at 680°. It is very hygroscopic and is often used for splitting off water from organic compounds. On adding zinc oxide to a con- centrated zinc chloride solution a soft mass is obtained, which soon becomes hard because of the formation of the basic chloride Zn900° 24.32 1 7.5 >651° >Zn 65.37 6.9 418° 920° 1 112.40 8.6 322° 778° 1 200.0 13.6 -39.4° 330° 1 In respect to chemical properties it should be noted that all of these elements are bivalent, except that mercury can be con- sidered as univalent in its ous-compounds. Their sulphates unite with those of the alkalies to form double salts of the same type, RS04-R2'S04-6H20 (R' = K, \a, XH4); the beryllium double salt ahjne crystallizes with .SHoO. The hydroxides of this group are ■soluble in ammonia with the formation of complex ions, or else they yield insoluble metal-ammonia compounds (Hg). The neutral salts have a tendency to go over into basic salts. This is especially marked in mercury; in the case of cadmium it is, strange to say, very weak. With the halogen comjiounds of the three related metals Zn, ■Cd and Hg the electrolytic dissociation is small; it decreases as the atomic weight of the metal rises and is verj- slight in the case of mercury. 41S IXORGAXIC CHEMISTRY. [§ 276- ELECTROCHEMISTRY. 276. As early as the beginning of the nineteenth century, when Da^'y isolated the allcali metals by means of the electric current (|§ 223 and 227), there was known to be an intimate relation between electrical and chemical phenomena. Berzelids even went so far as to suppose that affinity could be perfectly explained by assuming that the atoms are electrically charged and that these charges are the attractive or repellent forces. The galvanic element has been for a long time a familiar means of converting chemical energy into electrical energy. However, it was not until 1889 that a theoretical explanation of the connection between chemical and electrical phenomena \\'as offered; this explanation by Xernst is not only a i,'ery satisfactory one, but it also affords an insight into numerous chemical phenomena. The key to the explanation is the concept of "electrolytic solution tension," which has already been referred to in a few instances (§§ 203 and 268). T\Tien a metal comes in contact with the aqueous solution of one of its salts a difference in potential develops between the two. This phenomenon is explained by Xehnst as follows: Just as a liquid continues to evaporate at its surface until the pressure of the vapor becomes equal to the \'apor tension of the liquid, so a salt must continue to dissoh-e in water (evaporation and solution being analogous processes) until the osmotic pressure of its solu- tion balances the solution tension of the salt. Now, according to Xerxst, every metal also has a certain tendency, dependent only on its chemical nature, to force its atoms into solution as ions. This force, called the electrolytic solution tension, comes into action when the metal is immersed in an electrolyte and its strength is the less, the more cations of the metal are already in the solution. The amount of cations sent into the solution is very small, as experi- ment .shows, — so much so that it cannot be determined by the usual chemical means. The cause of this is not that the solution tension is low,- — on the contrai-}', the latter is often very large — but that an ecjuilibrium is very soon reached, because, notwith- standing the low concentration of the ions, they carry a very high electrical charge and the negati\ely charged metal soon attracts its positive ions in the solution with such force that just as many ions are precipitated on the metal as are sent out into the ■:7r, ] ELECTROCHEMISTRY . 419 solution. If P represents the solution tension of a metal and p the osmotic pressure of the cations in the solution, there are three possibilities to be distinguished: (1) P>p. The metal then behaves like a salt in contact with its own unsaturated solution. It forces cations into the solu- tion of the electrolyte, so that the solution becomes positively charged and the metal has to take on a negative charge. An equilibrium is soon established. However, if the free positive and negati\e electricities acquired by the electrolyte and the metal are conducted away by a connecting wire the metal wiU again send cations into the solution, and this action will continue tiU p reaches the \'alue of P. (2) P=p. There can be no potential difference. (3) P when both metals have the valence n. In such a ceU with closed circuit there are differences of potential not only between metal and solution but also between the two liquids and between the two metals. Experience has shown, however, that both of the latter are very small in comparison to the former, so that they may be disregarded. Leaving the solution tensions Pi and P2 out of consideration, E therefore depends on the values of the osmotic pressure pi and P2 of the metal-ions. If p2 can be made extremely small, so that P P loge — < lege — ^, E becomes negative, i.e. the current must alter Pi P2 its direction. This can be demonstrated as follows: In a Daniell cell, in which the osmotic pressure of the zinc ions (pi) is seldom ^•ery different from that of the copper ions (P2)) the current goes from the copper through the connecting wire to the zinc, for the solution tension (Pi) of the zinc is much larger than that (P2) of the copper (see below). Now the con- centration of the copper ions can be made se\-eral powers of ten smaller by adding potassium cyanide to the copper sulphate solu- tion, for l:)y this means tlie A'ery slightly ionized complex (Cu2Cy4)" is formed (§ 243). This addition actually reverses the direction of the current. Xeither the precipitation of the copper by potas- sium hydroxide nor the precipitation by ammonium sulphide reduces the concentration of the copper ions enough to produce this effect. 276.] ELECTROCHEMISTRY. 421 Since equation (2) pan also be written ^ the last expression becomes zero, it is apparent that the electromotive force of a Daniell cell is mainly determined by the ratio of the solution tensions of the metals. A galvanic cell can be regarded as a machine driven by the electrolytic solution tensions of the metals. The introduction of this conception of solution tension and the ideas connected with it has led to an altogether clearer understand- ing of the chemical processes of galvanic cells, as well as of the way in which the current is generated in them. Galvanic cells may be divided into two classes, reversible and noti-rcrcrKiblf. The Daniell belongs to the first class. It pro- duces a eurrent because the solution tension of the zinc exceeds that of the copper. The zinc sends its positively charged ions into the sulphate solution and itself becomes negative. On the other hand, the copper ions, on passing over into atoms and pre- cipitating themsehcs on the copper plate, transfer their positive charges to the latter, which thus becomes the positive pole. Chemically the process amounts to the simultaneous solution of zinc and precipitation of copper: CuS04 + Zn=Cu + ZnS04, or in ions: Cu" + Zn = Cu + Zn". If a current is sent through the Daniell cell in the opposite direction, ions will enter into solution at the copper plate because the latter acquires a positive charge, and the zinc ions will be forced to deposit themselves on the zinc, for the reverse current charges the zinc negatively so that it attracts the zinc ions. It is therefore possible by passing a reverse current through the cell to re- store it to its original condition, — hence the term "reversible." One of the most important styles iimnnfirw'imtJB °^ reversible batteries is the accumu- lli^AnJJJi^. \ ^,3k later, or storage-battery (Fig. 69). ThLs consists of a glass jar in which lead plates are susjx'ndod so that they dip into dilute sulphuric acicl. These plates are Fig. 69. — Accdmi'LATor. 422 lA'ORGANIC CHEMISTRY. [§276- coated alternately with lead peroxide (positive) and lead sulphate (negative). The positive plates are all connected with each other, as are also the negative ones. (From a large number of such cells a battery is constructed by connecting the positive pole of each cell with the negative pole of the adjoining one.) If a current is passed through the system so that it enters at the lead peroxide plate and goes through the sulphuric acid to the other plate, lead peroxide collects on the positive plate, while on the other, the cathode plate, the lead sulphate is converted into spongy lead. By this process the accumulator is charged. There- upon, if the poles are connected (by a wire), the opposite process goes on; the lead peroxide is reduced at the one plate and the spongy lead is converted into lead sulphate at the other. During the discharge the peroxide plate is again positive, the lead plate negative. The chemical process in the accumulator cell is there- fore expressed by PbOa + Pb + 2H2S04?^2PbS04 + 2H2O. The generation of the current has been explained in variouT ways. One is as follows : The lead peroxide on the anode plate has a certain sohition tension, and hence goes into solution as nega- tively charged Pb02" ions. Thereby it of course imparts to the plate istelf a numerically equivalent positive charge. These bivalent Pb02" ions encounter positively charged Pb" ions at the cathode plate, which are being sent by it into the solution; the cathode plate charges itself negatively at the same time. The two sorts of ions now combine to form electrically neutral PbO mole- cules, which yield lead sulphate with the sulphuric acid present: Pb02"+Pb-- =2PbO; 2PbO + 2H2S04 = 2PbS04-l-2H20. Among the non-reversible cells are the Bunsen and the Leclanche. A reverse current does not restore these to their original condition and their electromotive forces E cannot be calculated by the above formula; nevertheless the general prin- ciples of the pressure theory can be applied to explain the produc- tion of the galvanic current in these cells. The arrangement of the Bunsen cell — an amalgamated zinc plate dipped in sulphuric acid and a carbon cylinder in nitric or chromic acid — is well known. From an electrochemical stand- 277.] ELECTROCIIEMISTR Y. 423 point the generation of hydrogen from zinc and sulphuric acid amounts to a transfer of the charges of the hydrogen ions of the dUute acid to the zinc atoms and an escape of hydrogen in the form of discharged molecules. In the Bunsen cell, however, most of the hydrogen ions find an opportunity to give up their positive charges to the carbon cylinder and exercise a reducing action on the nitric or chromic acid. On the other hand the zuu plate sends positively charged zinc ions into the solution to the same extent as hydrogen ions disappear, the zinc plate itself acquiring a negative charge. The Leclanche cell consists of a zinc bar in concentrated ammonium chloride solution and a porous earthenware cylinder immersed in the same solution and containing some manganese peroxide and a stick of carbon for conducting off the current. Here again the zinc goes into solution: Zn + 2NH4CI = ZnCla • 2NH3 + Ho. The hydrogen ions discharge themselves at the carbon and reduce the peroxide. In this case also the carbon is the positive, the zinc the negative, pole. 277. Just as in galvanic cells chemical energy is transformed into electrical energy, so reactions between ions in general can produce an electric current if the conditions are suitable. A few examples of this may be cited. For these experiments a cell devised bj'LfPKE is very satisfac- tory (Fig. 70). It consists of two glass vessels Zj and Z2, to the bottoms of which the platinum electrodes ki and kz are attached. The vessels are connected by means of the wide siphon E. The wires A and K lead to a gal- vanoscope. To show that elec- trical energy can be obtained by the oxidation of the stannous to the stannic chlonde an acidu- LtJPKE Cell. lated stannous chloride solution (11.2 : 100) is introduced into Zi 424 INORGANIC CHEMISTRY. [§ 277- and an acidulated normal sodium chloride solution into Z2; the siphon also is filled with the latter solution. As soon as a few drops of chlorine-water or a solution of auric or mercuric chloride are allowed to fall from a pipette upon the electrode (^2) in the salt solution, the galvanoscope indicates a current in the wire circuit from K to A. Now, in order that the bivalent ion Sn" may become- quadrivalent (Sn"") it must acquire two more positive charges and this requires the addition of two chlorine ions. These are at once supplied by the mercuric or auric chloride. The metallic ions (Hg" or Au'") are deposited on k2 and impart to the latter a posi- tive charge, which, if conducted by means of the wire circuit K, is at the disposal of the Sn" ions. If free chlorine is added, it splits up into ions, as a result of which positive electricity is imparted to A-2 and this flows through the wire circuit back to ki and raises the potential of the Sn'" ions. The fact that electrical energj' can be obtained by the neutral- ization of sulphuric acid is capable of demonstration with the same apparatus. To this end a ^-normal sulphuric acid is intro- duced into Z2 and a ^-normal potassium sulphate solution into Zi and the siphon. If a large piece of palladium foil (about 4 sq. cm.) that has been saturated electrolytically with hydrogen is placed on the platinum disk of the electrode ki and touched for a few moments with a stick of caustic potash, bubbles of hydrogen will rise from the platinum plate of the other electrode kz and the needle of the galvanoscope will indicate the passage of a power- ful current outward from fcg. The hydrogen of the palladimn foil sends positi\'e ions into the solution, which, however, forth- with unite with the OH-ions of the potassium hydroxide to form neutral water. By the emission of these positive ions ki acquires a negative potential, ^^-llich flows out through the external circuit to k2. The hydrogen ions of the sulphuric acid surrounding this electrode are thus afforded an opportunity of discharging them- selves against this negative charge so that hydrogen is given off in the free state. In the combination of chlorine (or oxygen) and hydrogen chemical energy can also be transformed into electricity. To accomplish this, two tubes sealed at the top and fitted there with platinum electrodes, reaching almost to the open end of the tubes, are filled, one \\-ith hydrogen and the other with chlorine (oxj-geu) 277.1 ELECTROCHEMISTRY. 42.") and inverted in dilute sulphuric acid. On connecting the electrodes by a wire a strong current traverses the circuit. The gases ab- sorbed in the platinum electrodes drive their ions into the sur- rounding liquid, making the H-electrode negati\'e and the C1(0)- electrode positive. The ions of hydrogen and chlorine dissolve in the dilute sulphuric acid, however. This apparatus is called Groa'e's gas battery and was known long before a satisfac- tory explanation of it could be given. It is characteristic of all these various cells that ffie reacting substances are apart from each other. In the oxidation of stan- nous chloride by mercuric chloride the latter was not put in the vessel with the stannous chloride but in the other vessel; in the precipitation of silver chloride the silver nitrate was not put with the sodium chloride solution but with the sodium nitrate solution, and so on. The reaction took place only because one sort of ions transferred their electrification wholly or in part through the wire circuit to the other electrode, where it either converted atoms into ions or raised existing ions to a higher potential or, possibly, changed ions of opposite potential sign to neutral atoms. Since we know that chemical reactions can under suitable con- ditions produce an electric ciurent, we can, conversely, regard the existence of such a ciu-rent as an indication of the occurrence of a chemical reaction. Cohen has made use of this fact in deter- mining electrically the transition points of hydrous salts and other systems. Let us take, for example, a salt which loses its water of crystallization at a definite temperature, e.g. Glauber's salt, Na2S04 -101120; this has a transition point at about 33°, where the anhydrous salt becomes capable of permanent existence. Now it is ])0ssible for the anhydrous salt to remain in contact with its saturated solution in an unstable condition after the system has been cooled a few degrees below 33°; the re\erse is also true of the hj'drous salt. Shice these solutions are in contact with different solid phases (one with Na2S( )4 ■ IOH2O and the other with Na2y(34) they do not have the same concentration; at the transition point, however, these concentrations become equal, for since both solid phases are in contact witli the solution in each case, the solubility becomes the same. In his electrical method Cohen uses the differc;ic( in concentration of the solutions wliicli are saturated in respect to the two solid phases to form a galvanic cell. This 426 INORGANIC CHEMISTRY. >277 SNa- NajSO, Na,SO, can be done as follows: In the bottom of each of the cylinders A and B (Fig. 71) there is a little mercury. A platinum vdre is fused into each cylinder and the two are connected by means of a metallic wire. On top of the mercury Ls some insoluble mercurou? sulphate; above this in A is a paste of Na2S04 -101120 and water, in B is a similar mixture of water and Na2S04. Below the transi- tion point the solution in B is in the \mstable condition and more concentrated than that in A, which is stable. The result is that sodium ions dif- fuse through the siphon from the concentrated to the dilute solution, while at the same time an equivalent amoimt of S04-ions in B combines with part of the mercury to form mercurous sulphate, the nega- tive charge of the sulphate ions being transferred to the remaining mercury. Thus an electric current is produced \\hich passes through the wire circuit from the dilute to the con- centrated solution. Its direction and intensity can be determined by inserting a galvanometer in the circuit. Now, suppose that the whole apparatus is gradually warmed; the concentrations in A and B will approach each other as the temperature nears the transition point and at this point they will become equal. The intensity of the current will therefore decrease steadily till the transition point is reached, when it is zero. If the temperature is raised stUl higher the solution in A will become unstable and more concentrated than that in B, which latter will then be the stable solution; as a result the directio*. of the current will be reversed. In this way it is possible to determine the transition point very accurately. 278. As was remarked in § 276, the electromotive force which can be obtained from chemical reactions depends in large measure on the solution tensions of the metals. A knowledge of the latter Fig. 71 278." EL&TROCHEMISTRY. 427 is therefore of very great importance. It can be acquired with the aid of the equation previously given: E = 10-*-^^ log-, n p E, the difference of potential between a metal and the aqueous solution of one of its salts can be measured. All the other quan- tities of this equation arc known with the exception of P, which can therefore be calculated, as is illustrated in the following ex- ample : The potential difference between magnesium and the normal solution (^ mole per liter) of its sulphate was determined to be 1.22 volts. The equation thus becomes: 1.22= lO-^r log — . since for magnesium n = 2. p is the osmotic pressure of the ;\Ig- ions. On the assumption that the salt is entirely split up into ions, p is 22.4 Atm., for, osmotic pressure being equal to gas pressure, 1 mole gas at 0° and 760 mm. occupies a volume of 22.4 liters (§ 34); hence, if the volume is a liter, the pressure becomes 22.4 Atni. Therefore, at 0° we have: 1 .22 = 10-* X 27:3(log P- log 22.4), or log P= 43.23, whence we have, approximately, P= 10«. The following brief table indicates some of the results for different metals : Metal. N'aleiice. Solution Ten- sion in Atmospheres. Magnesium Zinc 2 2 3 2 2 2 2 T 2 1 1 10" 10'8 10" 10' 10' 10° 10-' 10-' 10-" 10-16 10-16 Aluminium Cadmium Iron Xickel Hydrogen Copper Mercury 428 INORGANIC CHEMISTRY. [§§278- The above figures show how enormously the solution tension differs in different substances. For magnesium and zinc it is many milUons of atmospheres, for copper, mercury (ous) and silver extremely small fractions of an atmosphere. Despite the comparatively large errors in the above data, due to the difficulty of determining the potential difference between the metal and its salt solution, the order of the decimal expressing the value of P can be accepted as reliable in each instance. Some of these differences of potential between metals and their normal salt solidions are as follows: Metal. Volt. Mg Mn Zn Al Cd Fe Tl Co Ni Pb H Cu + 1.24 + 0.798 + 0.439 + 0.22 + 0.143 + 0.063 + 0.045 -0.043 -0.049 -0.129 -0.277 -0.606 -1.027 -1.048 Hg Ag The algebraic sign of these differences of potential can be directly determined from the solution tensions. The electrolyte in which zinc is immersed must assume a positive potential and the metal itself a negative potential, because no zinc solution can be concentrated enough to hinder the emission of (positive) zinc ions by the metal. On the other hand, copper must become posi- tive in respect to a copper solution, for even in the most dilute solutions the osmotic pressure of the copper ions is greater than the solution tension of the metal. 279. A knowledge of the electrochemical series of the metals in electrolytes is of great practical value. Wherever combinations of various metals, alloys, metallic crustations, etc., are exposed to atmospheric action there is an opportunity to form cells of short circuit. In general, the metal with the greatest solution tension goes into solution and the other remains intact. A piece of gal- vanized (zinc-plated) iron wire does not rust, even in those places where the plate has been worn off, as much as if it were not zinc- 280.] ELIiCTROCHEMlSTRY. 420 plated. The reverse phenomenon, that tinned iron rusts faster than iron alone, is also due to galvanic causes. If our hypoth- esis is correct the atmospheric moisture adhering to the metal must act as an electrolyte with the combination tin-iron in such a v/ay that iron becomes the dissolving (negati\'e) electrode. Iron salts must therefore be fornietl and then transformed into rust by the loss of carbonic acid. The following experiment confirms this view. Rods of iron and tin are brought in contact by a wire which connects with a galvanometer. If the metals are dipped in water, into which air and carbonic acid are passed and to which is added a trace of sodium chloride (which always floats in the air and is washed down by the rain), the needle is deflected. The iron is found to be the anode, and in the course of an hour a thin yeUow coating of rust is to be observed on it. Sheet iron is tinned, as is well known, to prevent it from rusting (§ 199). If the tin plating is scratched off at any place so as to expose the iron, the latter begins to rust very rapidly, more so even than if it were not tinned. Galvanized iron, however, does not show a trace of rust where the plating has been damaged. 280. An ion can only go out of solution when a force greater than the solution tension acts on it, just as electrically neutral molecules cannot crystallize out of ^ solution until, its osmotic pressure exceeds that of the saturated solution. The removal of an ion can be brought about by the action of an electrical force. This is the real principle of electrolysis. The separation of an ion from a solution thus requires a definite electromotive force, 27" P which must be equivalent to 10~* — ■ log — (see above) and must therefore be stronger as the solution tension is greater and the ■osmotic pressure of the ions smaller. But since electrolj'sis takes place simultaneously at both the anode and the cathode, the total force E which is necessary for an electrolysis can be found by taking the sum of the forces necessary for the separation of the cation and the separation of the anion, thus: \rJi Pi «2 " P2/ Since it is always the case that various sorts of anions and cations are present together in a solution, electrolysis can thus take place 430 INORGANIC CHEMISTRY. [§§280- when E has become large enough to separate out one of the varie- ties of cations and one of the varieties of anions present. This is the basis of a method of utihzing various electromotive forces to effect an electrolytic separation of metals. It is not the current strength which is of primary importance to the electrolytic process (as was formerly supposed), but the difference of potential between the electrodes. A very successful example of this method is the separation of copper from zinc. With a current of low voltage it is possible to precipitate only the copper from a solution containing ions of both metals; if the electromotive force is in- creased, zinc also is separated. In many cases the ions of the water are more easily separated out than those of the dissolved electrolyte. In the electrolysis of potassium hydroxide, for example, OH-ions are liberated at the anode (they are .at once decomposed, however, into water and oxygen); at the cathode it is not potassium ions but hydrogen ions (in spite of their extreme^ small concentration) which are discharged, since the solution tension of hydrogen is much less than that of potassium. 281. The dissociation tensions E for various ions are given below. The figures are based on equivalent normal solutions. The dissociation tension 'of hydrogen is fixed at zero in the table. Inasmuch as there is always an anode and a cathode, it is possible to subtract from all the values of Ei an arbitrary but. constant amount and add it to the values of E2, without affecting E(=Ei + E2). The symbol 0" represents a secondary ionization product of the hydroxyl ion: OH' = 0"+H", the existence of which, according to Nernst, we are obliged toi assume, although only to an extremely small degree. Dissociation Tensions. J?i (Cations). E2 (Anions). Ag- -0.771 I' -0.520 Cu"- 0.329 Br' -0.993 H- 0.0 0" -1.23 Pb" + 0.148 CI' -1.417 Cd" + 0.420 OH' -1.67 Zn- 0.770 SO," -1.9 HSO/-2.6 281.] ELECTRO II EMISTRY. 431 These figures lead us to important results. In the first place they enable us to know at once the dissociation tension of any combination of ions. Zinc bromide, for instance, will require 0.993 + 0.770=1.763 volts for its electrolysis; when the concen- tration of the ions is normal, the electrolysis of hydrochloric acid will require 1.417 + = 1.417 volts, and so on. It is also obvious that it must be easy to separate silver from copper electrolytically, since the difference of their dissociation tensions is almost 0.5 volt. It also appears theoretically possible to separate electrolytically iodine from bromine and bromine from chlorine. The order of the metals in the above electrochemical series is the same as that in which one metal is precipitated from its solu- tion by the succeeding ones. As soon as a trace of the dissolved metal is deposited on the other one, the two metals form with the liquid an element, which electrolyzes the surrounding solution. The formula Of p ^=10-4— log- n ° p tells us, however, that the values of E depend not only on the solution tension but also on the osmotic pressure of the cations. Very decided changes in the concentration of the salt solution would make the order of the metals a different one. For instance, it would be possible to conceive a case in which lead would not be precipitated by cadmium. The electrochemical series of the anions also brings out im- portant relations. Bromine must quicklj- liberate iodine from iodide solutions and chlorine quickly liberate bromine from bromide solutions because of the marked difference in their dissociation tensions. We see, further, that chlorine must be able to generate oxygen in acid solutions, but not so with bromine or iodine. It is also known, however, that the generation of oxygen by chlorine proceeds with extreme slowness, in sharp contrast to the rapidity with which chlorine deprives bromine of its negative charge: Cl2+2Br' = Br2+2Cl'. This is not surprising in the light of the above considerations, for the chlorine, in order to enter the ionic condition, must make use of the ion 0", of which there is only an extremely small amount 4;j-J INORGAXIC CHEMISTRY. (§§281- present. The hydroxyl ion OH', which is present in relatively much larger amount and which after the loss of its negative cliarge would also yield a quantity of oxygen equal to that of the chlorine, holds its charge more than 0.3 volt firmer than the chlorine ion in acid solution. The application of electrolysis to commercial processes is referred to in connection with the substances concerned (c/. 5,§ 223, 226, 242, 245, 248, and elsewhere). BORON. 282. This element occurs in nature as sassolite, H3BO3, fcoraaie, ^klgrCUBieOao, colemanite, Ca2B60ii -51120, and borax, Na2B407' IOH2O. It can be obtained in the elemental state by the re- duction of boric anhydride, B2O3, or borax by means of magne- sium powder. It is prepared pure by subjecting a mixture of boron chloride and hydrogen to an arc discharge between two lioron or water-cooled copper electrodes, or bringing the mixture in contact with an electrically heated graphite tube. The element melts between 2000° and 2500°, but has such a high vapor tension that it sublimes rather easily as low as 1600°- Its hardness (exceeded only by diamond), combined with its amor- phous structure, constitutes a valuable mechanical characteristic. The electrical resistance decreases with rising temperature at remarkable rapidity. It dissolves in molten aluminium and, on cooling the melt, the compound AIB12 crystallizes out. Boron takes fire in fluorine and chlorine, uniting with them directly. When ignited in the air it burns to the oxide B2O3. At a very high temperature it com- bines with nitrogen to form boron nitride, BN. It reduces many compounds, such as CuO and PbO, and decomposes water at red-heat. Heating with nitric and sulphuric acid converts it into boric acid. It is also attacked by boiling caustic alkalies (like aluminium) : 2B + 2K0H -I- 2H2O = 2KBO2 + 3H2. Boron hydride. — ^When boric anhydride is reduced with an excess of magnesium powder, magnesium boride, Mg3B2, is formed. 283.] OXYGEN (VMPOCWDS OF BORON. 133 The latter on being aiUlecl to hydrochloric acid generates an ill- smelling gas consisting of hydrogen and a little boron hydride. This gas mixture burns with a green flame. Halogen Compounds. Boron chloride, BCI3, can be prepared by direct synthesis, but, better, by passing chlorine over boron carbide. It boils at 17°. Its vapor density indicates the above formula. Water breaks it up into hydrochloric and boric acids; it was with the aid of this reaction that the composition of the compound was determined. Boron fluoride, BF3, is formed, like silicon fluoride (§ 193), when the oxide is warmed with a mixture of calcium fluoride and sulphuric acid: B2O3 + SCaFo + 3H2SO4 = 2BF3 + 3CaS04+ 3H2O. It is a gas, of which water dissolves 700-800 volumes. A solution of this concentration fumes in the air. On dilution boric acid separates out after some time; hydrofluoboric acid, HF-BFs, is left in the solution. This acid cannot be isolated in the free state but various salts of it are known. It thus displays a very close analogv to silicon fluoride. Oxygen Compounds of Boron. Boron oxide, B2O3, boric anhydride, is obtained as a vitreoua mass by igniting boric acid. It is very hygroscopic and is recon- verted by the absorbed water into boric acid. With hydrofluoric acid it forms boron fluoride. The oxide is volatile only at elevated temperatures. 283. Boric acid, H3BO3, is found in the volcanic districts of Tuscany, where jets of steam (the springs are called " fumaroles " and the jets proper " soffioni") containing a little boric acid issue from the earth. The steam is conducted into water, in which the boric acid is retained. When this liquid reaches a certain concen- tration, it is allowed to settle, whereupon it is piped off into a very 434 INORGANIC CHEMISTRY. [§283. long, flat leaden pan, which is warmed to about 50-60° by other soffioni. At this temperature the boric acid volatilizes but very little with steam and when the concentration has become great enough it crystallizes out. It is purified by converting it into borax, which is recrystallized and then decomposed by hydrochloric acid, setting free boric acid. Considerable boric acid is also made by decomposing native borates with a strong mineral acid. The volatility of boric acid with steam has for a long time been regarded as an especially interesting phenomenon, because the anhy- dride B2O3, into which it is readily converted at an elevated tempera- ture, is only volatilized with extreme difficulty. The question there- fore arises as to the particular comijound in which boric acid exists in solution and the one in which it escapes from solution. The first point can be settled by a determination of the boiling- point elevation or vapor-tension lowering of boric acid solutions. Meas- urements of this sort have shown that H3BO3 molecules exist in dilute solution. As the solution becomes more concentrated the vapor-tension lowering no longer corresponds to this formula; the decrease in the lowering indicates that the number of molecules of dissolved substance has grown less, i.e. some such change as 4H3B03 = H2B407 + 5H2O has occurred. If H3BO3 molecules volatilize with the water, the concentrations of boric acid in the solution and in the vapor must, according to Henry's law, maintain a constant ratio, independent of the amount of boric acid present. This was found to be true for dilute solutions but not for concentrated ones, which is in agreement with the experiments on vapor-tension lowering, because there also the concentration of the acid in the vapor remained proportional to its concentration in the solution. It is therefore demonstrated that the compound which escapes with the steam is boric acid, H3BO3. Boric acid crystallizes in lustrous laminse, which feel greasy and are difficultly soluble in cold water (about 3% at ordinary tem- peratures). This solution acts as a weak antiseptic, for which purpose it is frequently used. At 100° boric acid loses 1 molecule H2O, passing over into metaboric acid, HBO2. At 140° tetraboric acid, H2B4O7 (=4B(OH)3~5H20), is formed, the sodium salt oi which is borax. No salts of the normal boric acid, B(0H)3, are known, but § 283.] BUROX. 435 metaboric acid forms several. They are unstable and are converted by carbon dioxide into salts of tetraboric acid: 4NaBU2 + CO2 = Na2B407 + NagCOs. The best-known salt of boric acid is borax, Na2B407-12H2(), ■often called tinkal. At present most of the borax on the market is made by boiling colemanite, Ca2B60ii-5H20, found in California, or a similar borate, occurring in Chile, with soda. Borax swells greatly on heating; this is due to the escape of water of crystalliza- tion from the semi-molten salt. On continued heating it forms a vitreous mass. This glass has the property of dissolving metallic ^)xides, some of which give double borates of a characteristic color; hence its use in qualitati\'e analysis. The same property makes it valuable in soldering; solder adheres only to the untar- nished metal, so a little borax is placed on the surface of the metal -and heated with the soldering-iron in order to remove the rust. The dissolving of metallic oxides is easily understood, when we write Na2B407 as 2NaB02+B203; it is the boric oxide, B2O3, which can be regarded as combining with the metallic oxides to form salts. Boric acid is a weak acid; its salts are therefore hydrolyzed •quite perceptibly — ^more so, of course, as the dilution increases. This can be illustrated by a .simple experiment demised many years ago by Rose. To a concentrated solution of borax some litmus is added and then acetic acid until the litmus is just red; if the liquid is then diluted, it turns blue because the alkali is set free and boric acid has scarcely any effect on litmus. Rather interesting, also, is the Ix^havior of silver borate, which is deposited as a white salt on mixing concentrated solutions of borax and silver nitrate. When dilute solutions are mixed, how- ever, a precipitate of grayish-brown silver oxide is formed, the silver borate being almost completely hydrolyzed in the dilut-* solution. On treating a mixture of boric acid and sodium peroxide with water a perborate is formed and crystals of the composi- tion NaBOs -41120 scjjarate out. They are stable when solid but liberate oxygen from a warm solution. The solution con- tains hydrogen peroxide also. 436 LXOROAXIC CHEMISTRY. [§§284- ALUMINIUM. 284. This metal does not occur native, but in combination it is found in large quantities and very widely diffused. Corundum, including the precious stones sapphire and oriental ruby and the natural abrasive emery (all noted for their hardness), consists of alumina AI2O3, colored by traces of other oxides. Bauxite is a hydrate of aluminium and iron. Clay and kaolin (China clay) are principally aluminium silicate. Many other minerals, such as feldspar, mica, etc., contain it as a base. A peculiar aluminium mineral, cryolite or ice stone, SNaF-AlFs, is found in Greenland. The metal can be obtained from the chloride by reduction with sodium but at present it is produced exclusively by decom- posing aluminium oxide with the electric current. The most important commercial process is that of Hall (invented independently in Europe by Heroult). Alumina is dissolved in a fused bath consisting of cryolite or an equivalent mixture. The process is carried out in a large carbon-lined pot, the inner surface of which constitutes the cathode. Carbon rods immersed in the bath serve as anodes. Fresh alumina is added from time to time and the metal is drawn off at the bottom periodically. The temperature is a little above the melting-point of cryohte. A current of several thousand amperes and less than 8 volts maintains the liquidity of the bath as well as effects the electrolysis. The increased output due to improved methods has brought the price of the metal down from over $90 per pound in 1856 to- about $0.20 at the present time, and the production is steadily increasing (.34,000,000 lbs. in 1909). Aluminium is a silvery-white metal of low specific gravity (2.583). It is rigid but very ductile and malleable. It softens at about 600°, melts at 658°, and boils at about 1800°. It is permanent in the air, since it soon becomes coated with a firm thin layer of oxide. Small fragments burn with a bright light when heated in an oxygen atmosphere. It is not attacked by dilute nitric acid at ordinary temperatures and only slightly so by dilute sulphuric acid, Hydrochloric acid dissolves it readily. 285.] ALUMINIUM. 437 as does also caustic potash, hydrogen Ijcing evolved and a 1 u m i - nates formed in the latter case. Various alloys of aluininiiiw have found a place in the arts. Among them mention may be made of aluminium bronze, which consists of lopper and 5-12^f, aluminium. It can be easily cast and has a golden color and lustre. Its gi-eat firmness and elasticity render it valuable for physical instruments (balance beams) and watch springs. New allocs of aluminium are being constantly brought on the market: there is one with magnesium called inagnalium and another with tung- sten, for example. Aluminium reduces many oxides (Goldschmidt) with a vigor- ous evolution of heat (§ 293). The reduction proceeds of itself after it has been started at a certain place in the mixture. For this purpose a primer is used consisting of a mixture of oxygen-producing substances, such as KCIO3, etc., and a piece of magnesium ribbon, which is ignited with a match. The heat that is thus evolved Ls used to heat iron bolts to white-heat and also for welding railroad rails, etc. The welding is accomplished by packing the rails in a mixture of iron oxide, sand, and aluminium powder together with a special sort of cement for making it compact. When this mass is ignited it continues to burn and heats the rails to glowing. An amalgam of aluminiuia is easily prepared by introducing aluminium filings into a i'^c solution of corrosive sublimate. This amalgam decomposes water energetically at ordinary tempera- tures, liberating hydrogen and forming aluminium h3'droxide. As neither basic nor acid substances go into solution, it is a neutral reducing-ageni. The cause of this energetic reaction is due to the mrcumstance that the mercury hinders the formation of a thin firm coating of oxide over the surface of the metal, which would otherwise protect it from further oxidation. Compounds of Aluminium. 283. The only known oxkle of aluminium is alumina, AI2O3, which is formed on heating aluminium salts or the hydnixide. It is a white amorphous powder, readily soluble in acids; however, after it has been strongly ignited it is no longer soluble and nnist 438 INORGANIC CHEMISTRY. [§ 285. then be disintegrated bj- fusion with potassium hydroxide or acid potassium sulphate. It is found crystallized in nature (§284). The artificial manufacture of rubies and sapphires is accomplished by fusing amorphous AI2O3 with lead oxide at bright red heat in a Hessian crucible. Lead alumipate is first formed, whereupon the silica of the crucible causes the alumina to separate out in beautiful crystals, exactly like the natural gems. By adding a little potassium dichro- mate we get crystals having the color of the natural rubies; similarly, the addition of cobalt oxide gives sapphires. Aluminium hydroxide, Al203-nH20, is deposited as a hy- drogel (§ 195) when a solution of an aluminium salt is treated with ammonia. I;i the decomposition of the aluminates it is obtained as a white jiowder. A hydrate with a low percentage cf water, AI2O3 -21120, bauxite, occurs in France and different parts of the United Statesi in large deposits. Aluminium hydroxide is both weakly acidic and weakh^ basic in character. Its salts with acids suffer partial hj'drolysis in aqueous solution and hence react acid (§ 2.39). It dissoh-es in alkalies to form aluminates, such as AlOoK, A102Na, and AlOsXas, which are deposited in the amor- phous state when alcohol is added to their aqueous solutions. They are decomposed by atmospheric carbonic acid. Aluminium hydroxide is insoluble in water but dissolves in a solution of aluminium chloride. By subjecting this solution to dialysis, it is possible t(i get rid of the hydrochloric acid (which is present because of hydrolytic dissociation) entirely and thus obtain a colloidal solution of the hydroxide. Alumin- iimi hj'droxide does not form salts with weak acids. Aluminium chloride, AICI3, is most conveniently prepared by passing drj^ hydrochloric acid gas over aluminium fiUngs in a tube of porcelain or glass and collecting the sublimed product in a wide-mouthed bottle (see Fig. 72) . After the tube has been heated to a sufficiently high temperature to start the reaction, no further heating is rec[uired; however, it is more practicable to continue heating in order to collect the chloride in the receiver. Aluminium chloride is very hygroscopic. The aqueous solu- tion hydrolyzes so readily, depositing alumina, that it can only be preserved by the addition of an excess of hydrochloric acid. Such § 285.] COMPOUA'DS OF ALUMINIUM. 4:!!) a solution does not yield aluminium chloride on evaporation, since it decomposes completely into the hydroxide and hydrochloric acid on account of the continued removal of the latter dissociation product. The %'apor density of the chloride up to 400° corre- sponds to the formula AlzClg, above 760° to AICI3. With the chlorides of potassium and sodium, aluminium chloride forms compounds such as AlCls-KCl, whose solutions can be evap- orated without decomposition. Compounds such as AlCls-PCls, Fio. -Preparation op Aluminium Chloride. AlCla-POCls, etc., have also been prepared. In organic chemistry anhydrous aluminium chloride is of great value in synthetical work. Aluminium sulphate, Al2(S04)3 • I6H2O, is obtained by treating clay with concentrated sulphuric acid; the product is dissolved in water and allowed to crystallize. Aluminium sulphate unites with the alkali salts to form double salts of the general type: R2S04-R2'(S04)3-24H20, which are known as alums. R may be either K, Na, NH4, Cs, Rb, Tl, or an organic base ; R' may be Fe (ic) or Cr, instead of Al. The alums all crystallize in octahedrons and cubes, which often grow to large dimensions; they form mixed crystals readily. Ordinary alum (potassium alum) is used as a mordant in dyeing i'Org. Chem., § 362), but it is being gradually superseded as such by aluminium sulphate and sodium aluminate. In the vicinity of Rome the mineral ahuiile, or ahan slonc, is found, whose com- position is K(A102H2)3(S()4)2; from it a much sought variety of alum is made. Alum is also made from cryolite, etc. When two salts combine we may have one of two results: either the new salt which is formed gi'ves ions in dilute aqueous 440 INORGANIC CHEMISTRY. [§§285- solution that differ from those of the two salts, or it gives the same ions. A good example of the former ease is yellow prussiate of potash; it gives neither ferrous ions nor cyanide ions, so that it must be regarded as K4[Fe(CN)6]. Such salts are termed com- plex. The second case is illustrated by the alums. A dilute- alum solution exhibits all the reactions which characterize its com- ponents and its conductivity is the mean of the two separate salts for the same concentration. When the union is of this sort we have what is called a double salt. Between the two kinds there are salts of an intermediate nature which form not only complex ions but also the original ions to a greater or less extent. The copper- ammonia compounds (§ 244) behave in this way. 286. Aluminium silicate, kaolin, is formed in nature by the weathering of the numerous alkali-alumina double silicates, the alkali sUicate being dissolved out, lea^-ing the insoluble aluminium silicate. Clay is aluminium silicate; it is usually colored brown by iron oxide. It is the essential raw material of the ceramic industries, being used both for rough bricks and the finest china- ware; of course the better grades require better sorts of clay. Bricks are molded out of ferruginous and calcareous clays (loam) and then baked ("burned," or "fired") till they become firm. Under the head of earthenware, or porous ware (faience, majolica, etc., and common crockery) we include all articles which consist of burned clay (frequently mixed with quartz), are porous and display an earthy fracture and which are covered with a glaze of easily fusible silicates. The glaze is produced by introducing salt into the kiln. The hot steam causes the formation of hydrochloric acid and SO; luin hydroxide, which unite with the clay to form sodium alumhiium silicate. In -porcelain the pores of the earthen mass are completely filled with fused silicate, as a result of the addition of feldspar and quartz before the burning. The less of such admix- tures is present the more difficult the porcelain is to bum and the less sensitive it is to changes of temperature. Clay is the most widely diffused refractory material; it resists not only high temperatures and sudden changes of tem- peraturcj but chemical action as well. Ultramarine is a very beautiful blue pigment, which is prepared arti- ficially by heating a mixture of clay, soda, sulphur and wood charcoal in the absence of air. It occurs in nature as lapis lazuli. It is usually 288.] GALLIUM, L\DIUiM, THALLIUM. 441 regarded as a compound of sodium aluminium silicate with polysul- phides of sodium. This is indicated by the fact that it is attacked by acids with the evolution of hydrogen sulphide and the disappearance of the color, while it is unaffected by alkalies. It is still uncertain what substance gives the pigment its blue color. GALLIUM, INDIUM, THALLIUM. 287, The existence of gallium was predicted by Mendele'bff (§ 217) in the same manner as that of germanium. The hypothetical eka-alu- minium was discovered in 1875 by Lecoq de Boisbaudran in a zinc blende by means of spectrum analysis. Its spectrum consists of two violet lines. It is a very rare element. The metal is ■white, melts as low as 30° and has a specific gravity of 5.9. It is only superficially oxidized by the air and is not attacked by water. Like aluminium, it is only slightly affected by nitric acid but dissolves readily in hydro- chloric acid as well as ammonia and potassium hydroxide. It forms alloys ■n'ith aluminium, which, when the proportion of aluminiimi i-? small, are liquid at ordinary temperatures because of the depressioi. of the melting-point of gallium, and it decomposes water almost as readily as sodium. In its compounds, also, gallium displays much analogy with aluminium. The hydroxide dissolves in alkalies. The chloride, GaClj, fumes in the air like AICI3 and its aqueous solution yields hydrochloric acid on evap- oration. The sulphate gives an alum, Ga,(S0j)3- (NH,")2SO,-24R,0, with ammonium sulphate. Hydrogen sulphide precipitates gallium only from an acetic acid solution, in which respect gallium resembles zinc (§ 269). Indium has already been referred to in the discussion of the periodic bvstem (§216), so that it will be passed over here with a brief descrip- tion. It was discovered through its spectrum, a blue line. This ele< ment, too, occurs very rarely, being found in certain blendes. The m e t a 1 is white; m.-pt., 170°; sp. g., 7.42. It is permanent in the air; heated to a high temperature it burns with a blue flame to the oxide In203. The chloride InClg is hygroscopic; its aqueous solution does not decompose on evaporation. The sulphate forms an alum with ammonium sulphate. The hydroxide dissolves in alkalies. 288. Thallium is the most common of these three elements, notwith- standing it always occurs in limited amounts. It is occasionally found in the " Abraum salts" camallite and sylvUc and frequently also in different native sulphides. When the zinc blendes are roasted in sulphuric acid factories the thallium goes off with the fumes and settles in the flue dust and chamber mud. From these deposits it is obtained by boiling with 442 IXORGAMC CHEMISTRY. [§§288- dilute sulphuric acid and precipitating with hydrochloric (or better hydri- odic) acid, whereupon the sparingly soluble chloride (or iodide) is deposited. This element was also discovered with the spectroscope (Crookbs); its spectrum is a bright green line. Thallium is a soft metal, about like sodium, and has a bluish color like lead. Sp. g., 11.8; m.-pt., 290°. In moist air it oxidizes very rapidly at the surface; but it does not decompose water at ordinary tempera- tures. When heated it burns with a beautiful green flame. Sulphuric and nitric acids dissolve it readily, but hydrochloric acid acts very slowly because of the slight solubility of the chloride. There are two sets of compounds: the Ihallous compounds, derived from the oxide TljO, and the thallic compounds, from the oxide TljOj. The former resemble those of the alkalies and silver very much. This similarity shows itself, for instance, in the solubility of the hydroxide and the carbonate, whose solutions react alkaline. Moreover, many thallium salts are isomorphous with potassium salts and, like the latter, give double salts with platinum chloride, e.g. TljPtCle. Further t-he.re is an alum Tl2S04-Al2(SOi)3-24H20, as well as other double sul- phates, e.g. TUSOi-ilgSOi-GHjO, which are analogous to the corre- sponding potassium double salts. On the other hand thallium resembles silver and lead in the small solubility of its halides (the iodide is the least, and the chloride the most, soluble) and also in respect to the order of solubility of these compounds. In the thallic compounds the element is trivalent, like the other elements of the group; furthermore, hke the compounds of the latter, the thallic compounds readily form complex salts, and undergo con- siderable hydrolysis when dissolved in water. SUMMARY OF THE GROUP. 289. The five elements last considered, B, Al, Ga, In, Tl, form a natural group, in which the last three display particular similarity to each other in their physical properties. Something analogous was observed with copper, silver and gold in the first group and with zinc, cadmium and mercury in the second group. The fol- lowing table affords a brief comparison of certain physical data: B AI Ga In Tl Atomic weight Specific gravity Melting-point 11.0 2.45 2000° 27.1 2.58 658° 69.9 5.9 30° 114.8 7.4 176° 204.0 11.8 290° 290.] THE RARE EARTHS. 443 In the spectra of Ga, In and Tl it is again noticeable that the lines move towards the red end as the atomic weight increases (§ 265). As to their chemical nature it may l)e remarked that all the elements of this group arc trivalent and that the basicity of their oxides increases with rising atomic weight; boron hydroxide (boric acid) has exclusively acid properties, but the hydroxides of the other elements, even T1(0H)3, are also soluble in alkalies. As most of the lower oxides of the metals are more strongly basic than the higher oxides, it is not strange that thallous hydroxide is a strong base. THE RARE EARTHS. 290. In the middle of the periodic table (p. 301) are located a number of elements, which are classed under the term '"rare earths." There is still much uncertainty in regard to some of them, particularly as to their elemental nature. This is due in large measure to the great simi- larity between the elements and the conse(iuent difficulty in .-separating them. They may be arranged in two groups : the cerium group con- taining the elements lanthaniim, cerium, praseodymium, neodymium and samarium ; and the yttrium group containing europium, terbium, dysprosium, holmium, yttrium, gadoluiium, erbium, thulium, ytter- bium, scandium, and lutecium. These elements occur in \arious rare minerals which ha\c been found principally in Sweden and Greenland, viz., cerite, gadolinile, euxenite, orthite, etc. Since the use of the oxides of cerium and thorium in the incandescent gas-light of AuER vox Welsbacii has created a demand for them, minerals in which the rare earths occur are being ardently sought. The interesting tc'ct has developed that they are b}- no means so "rare " as was supposed. An especially rich source of these earths has been found in mouazite sand, which occurs in rather large quantities in the United States (production 900,000 lbs. annually), Canada and Brazil. It coii.sists chieliy of a phos- phate of Ca, La, Di, Y and Er, with varjing amounts of thorium silicate and thorium phosphate. In order to isolate the rare earths from these minerals the latter are powdered very finely and heated to faint-reil heat with concentrated sul- phuric acid. Thus the rare earths are changed into sulphates and the sUicic acid is converted into the iji.soluble condition. The sulphates r.re then taken up in ice-water, in which they di.'^solve much more readily than in warm water (since a difficultly soluble hydrate is formcilal a 444 IXORGAXIC CHEMISTRY. [§290. higher temperature). From this cold solution they can be precipitated with oxalic acid, their oxalates being almost insoluble even in dilute acids. Thus they are freed from Ca, Fe, etc. The oxalates are the a converted into oxides by heating. The separation of these oxides is a more difficult task. Various methods are in use, by which the separation of the eerie earths is fairly well accomplished ; but for the numerous j-tteric earths no suc- cessful method has yet been devised. Some of the methods employed are as follows: The insolubilit}' of the sulphates of cerium, lanthanum and didymium in a saturated sodium sulphate solution (by reason of the formation of double salts) is made use of to separate them from erbium, ytterbium and yttrium. The nitrates of the various metals of this group differ markedh' in their stability on heatmg; hence another method of separation has been devised, by which the nitrates are decom- posed one after another by heating and those that remain undecom- posed after each successive heating are extracted with water. A third method is the fractional precipitation of the solutions with ammonia. Further, by fractional precipitation with potassium chromate (the insoluble neutral chromates being deposited), separations can be accom- plished which are otherwise very difficult. The ytteric earths are separated after Urbain by mixing the aqueous solution of their nitrates with bismuth nitrate, which is iso- morphous with them. From this solution they separate out as mixed crystals with the bismuth nitrate, the different kinds of mixed crystals having different solubilities. If a fractional crystallization is con- ducted, certain of the rare earths accumulate in the first crystalliza- tions, others in the last crystallizations, the middle ones consisting almost wholly of bismuth nitrate. Most of the rare earth metals form only one oxide, having the formula M^Oji cerium, praseodymium and neodymium have higher oxides as well ; of these Ce02 is able to form salts. The best method of detecting these metals is by spectroscopy. The spectra of the eerie metals are satisfactorily known ; thoae of the ytteric metals are not so well known. Many of the latter are characterized by absorption bands, thus dysprosium, holmium and thulium. Others like yttrium gadolinium and ytterbium, whose oxides and salts are colorless, do not give an absorption spectrum, but their spark spectrum is characteristic. The spectra of the ytteric earths display a great many lines. Furthermore, investigations of the ultra\-iolet spectra (photographic) have furnished important information. In addition to these kinds of spectra the phosphorescence spectrum should be mentioned as an important means of investigation, especially § ■2SII1.] THE RARE EARTHS. 41."! to determine (he purity of these earths. When the earths are placed in an evacuated tube and exposed to the action of cathode rays, the earths become luminous, a phenomenon that is known as cathodic phosphorescence. The spectrum of the phosphorescence has character- istic lines. It has been proved that the perfectlj- pure earths do not show the phospherescence, but that it is caused by extremely slight admixtures of other earths. The maximum influence is caused in most cases by an admixture of 1-0.1'^ The disappearance of the phosphor- escence is therefore a means of tcllins when the earth is pure. On the other hand, however, the characteristic phosphorescence spectrum can be used to recognize some of the earths. Cerium occurs principally in cerite (as high as 60%). Its salts are colorless when pure and give no absorption spectrum (§ 263). The metal looks like iron but is soft, hke lead. It oxidizes slowly in the air, becoming coated with a black layer. At an elevated tem- perature it takes fire. An alloy of 70'^'t Cc and 30% Fe gives off sparks when it is scratched, and it is sometimes used as a substitute for matches. It forms two sets of salts, the cerous salts, which can be derived from the oxide CejOs and are colorless, and the eerie salts, derivable from CeOj, wliich are yellow or brown. Cerium is thus quadrivalent (as the existence of the fluoride CeFj-HjO also indicates) and so belongs to the fourth group of the periodic system. When chlorine is passed into an alkaline solution of a cerous salt a yellow precipitate of CVOj is obtained. Lanthanum is onh' trivalent. Its oxide LajOj and its salts are color- less when pure. Didymiutit was formerly regarded as an element, but Ateh vox Wels- BACH succeeded in splitting it up into two components, called praseody- mium and neodymium. This can be accomplished by making use of the difTerence in solubility of their potassium double sulphates in a concentrated solution of potassium sul]}hate. The praseodymium salts are green and give green solutions; the neodymium salts have an ame- thyst color and give pink solutions. The alworption spectra of the two elements differ considerably. Scandium occurs in the mineral wolframite. It is a trivalent element, like La. Its existence was predicted by Mendel£eff, who callfed it ekahiirun. Its trivalence places it in the aluminium group. The hydrox- ide Sc(()H)3 is gelatnious, but insoluble in an excess of alkah. Pure scandium can be inepared by way of the sodium double carbonate, Sc,(('03)3-4Na2C(),-6HjO. Ytterbium. — The oxide ^b,*):, is the main constituent of erbia, formerly regarded as an elementary oxide (ulitaiiied from euxenite and gadolinite), but now known to contain also the oxides of scandium, 446 INORGANIC CHEMISTRY. [§§ 29e- yttrium, erbium, etc. Ytterbia (oxide) is obtained by fractional heating of the nitrate mixture (see above). The salts of ytterbium are colorless" and give no absorption spectrum. Auer ^von Welsbach recently suc- ceeded in resohing ytterbium into two elements, which he called alde- haranium (Ad) and cassiopeium (Cp). He found their atomic weights to be Ad = 172.90 and Cp = 174.23. The salts of samarium are yellow and ha^•e a characteristic absorp- tion spectrum. TITANIUM, ZIRCONIUM, AND THORIUM. 291. These uncommon elements are related to carbon and silicon in the same way as K, Rb, and Cs are to Li and Na, and as Ca, Sr, and Ba are to Be and Mg. Titanium and zirconium still give acid-forming oxides, while thorium forms only basic oxides. Titanium displays very close analogy to sihcon ; it frequently occurs with the latter, but always in a small amount. It is best prepared pure b}' reducing TiC'l, with sodium in a steel bomb at low red heat. Sp. g. = 4.50; m.-pt. = l.s00°-18.50°. The metal is hard and tough in the cold, but can be worked on heating; it is a good electrical conductor. When heated in a current of nitrogen It burns quantitatively to the nitride TiN. Titanium dioxide, TiOj, occurs as mineral in three modifications : rutile, anatase and brookile. Titarium chloride, TiCl,, is prepared by passing chlorine over the carbide, which is prepared in the electric furnace. TiCl, is liquid and fumes in the air because of decomposition by atmospheric water into HCl and Ti(OH)(. Titanic acid, Ti(OH)„ separates out as a white amorphous powder when the hydrochloric acid solution of a titanate is treated with ammonia. This action is due to the weak basic character of ammonia and the weak acid nature of titanic acid- as a result the ammonium titanate is completely hydrolyzed (§ 239). Like silicic and stannic acids, titanic acid readily forms poly-acids (§194). It dissolves in alkalies to form titanates, which are also obtained by fusing TiOj with alkalies. On the other hand titanic acid dissolves in concentrated sulphuric acid ; it then remains in solution eveu when poured into water, because the excess of sulphuric acid hinders hydrolytic dissociation. Higher as well as lower oxides of titanium are known. The lemon-yellow oxide TiOj is formed on treat- ing the sulphuric acid solution of Ti(OH)j with hydrogen peroxide (§ 38). Zirconium occurs in nature chiefly as zircon, ZrSiO^. It is not reduced from the oxide by aluminium. It is prepared pure by heating KzZrFe withmet alHc sodium. Sp. g. =6.3. Small pieces bum bril- liantly when heated in the air. Moissan obtained zirconium carbide, CZr, from zircon directly by heating it with sugar charcoal in an 291.] TITANIUM, ZIRCONIUM, AND THORIUM. U7 electric furnace (1000 amp. and 40 volts) for ten minutes. The silicon for the most part volatilizes. If the carbide is treated with chlorine at dull-red heat, it is converted into the chloride. Zirconium chloride behaves with water in the same way as TiC'l^ anil SuC'l,. The hydroxide, Zr(OH)< is precipitated by anmionia from acid .solutions as ~a volu- minous mass. It is insoluble in alkalies, but on being fused with the latter it forms salts such as Na^ZrO, and Na,ZrO,, which arc decom- posable b>- water. The basic character of the hydroxide is apparent from the fact that it gives a sulphate, Zyii^O,),, with sulphuric acid, which can be recrystallized out of water. Zircon ia, ZrOz, emits a \-ery bright light when heated strongly. Thorium is at present obtained mainly from monazite sand; it is also found in the thn gives up its electrical charge and an atom of oxygen to the hydrogen ions of the free acid, thus yielding water and forming, together with a second Cr04 ion, the red ion Cr207: 4K- + 2Crr)4" + 2H- -F SO4" = 4K- + CtzOj" + H2O + SO4". Acid salts of chromic acid do not exist on account of this reducing effect of the hydrogen ions on the Cr04-ions, for which reason also free chromic acid, H2Cr04, is incapable of independent existence. Chromic acid is a weak acid, since its insoluble (in water) salts, e.g. those of barium, lead and silver, are readily dissolved by strong acids (§ 146). Alkali chromates are invariably obtained by fusing a chromium compound with an alkali carbonate and an oxidi zing-agent. The latter is unnecessary wnen the fused mass can be brought suffi- ciently in contact with the oxygen of the air by stirring. Chromite is worked up commercially into chromates in this way; it is cal- cined with soda and lime above 1000° in a reverberatory furnace: 2Cr203 ■ FeO + 4Na2C03 + 4CaO + 70 = 4Na2Cr04 +4CaC03 + FegOa- 454 INORGANIC CHEMISTRY. [§§ 295- The resulting sodium chromate is lixiviated and sulphuric acid is added to its solution; on evaporation sodium dichromate, Na2Cr207, crystallizes out, and this can be converted into potassium dichromate, a well-known salt, by double decomposition with potassium chloride. The fusion is much more readily accomplished when caustic potash (KOH) is used instead of soda (Xa2C(33), probably because fused potassium hydroxide absorbs oxygen from the air and forms the peroxide, thus becoming a much more acti^'e oxygen-carrier than soda. Under these circumstances the oxidation proceeds rapidly and completely as low as 500°. Potassium dichromate finds frequent use as an oxidizing-agent in sulphuric acid solution, being itself reduced to chromic sulphate : K2Cr207 + 4H2S04 = KoS04 + Cr2(S04)3-h4H20-|-30. An important commercial task is the regeneration of the chrotnio acid from such a chromium sulphate solution. The method followed in the factories at Hochst, German}', is an electrical one. The solution is elec- trolyzed between lead electrodes in a vessel containing a diaphragm (porous partition). By the action of the current chromic acid is formed -at the anode, while at the cathode h;,'ell(iw powder, prepared by heating the nitrate. The corresponding hydroxide, U(OH)e, is not known, but salts of the compound U(OH)|,— 2H;0 =U02(0H)o with acids have been prepared. Since the UOj group acts here as a bivalent radical it is called uranyl and its salts uranyl salts, e.g. UOjCNOj),, uranyl nitrate, crystallizing with 6H2O in beautiful greenish-yellow prisms. Uranium trioxide also has somewhat the character of an acid anhydride; if caustic potash and soda are added to uranyl salt solutions yellow uranales 43S INORGANIC CHEMISTRY. [§§ 298- (KjUoO, and Na2U207) are precipitated, which are soluble in acids. I'l-aninite can be regarded as the uranate of uranous oxide, Vfi^ = 2U03-U02. Both oxides are converted into this UjOg oxide by heat- ing in the air. Uranium salts are used to impart to glass a beautiful greenish-yellow fluorescence. The detection of uranyl salts is accomplished with the aid of the pre- cipitate, soluble in excess, which they give with ammonium carbonate and by tiie reddish-brown precipitate with potassium ferrocyanide. For the radioacti\-e properties of uranium see § 2B7. SUMMARY OF THE GROUP. 299. The elements chromium, molybdenum, tungsten, and ura- nium, in connection with sulphur, constitute a natural group in the periodic system. Particularly in the higher oxides there is con- sideral^le analogj- with the behavior of this metalloid. Their acids, for example, all have the formula H2RO4. Moreo^'e^ sulphur also has the ability to form polyacids (pyrosulphuric acid) although it is not so prominent as in the first-named four elements. Se^•eral of their salt.s are isomorphous. The strength of the acids decreases, as in other groups, with rising atomic weight. Another character- istic of all the elements of this group is the great abundance of formula types; it is also \'ery noticeable in the case of sulphur, whose acids are remarkably numerous. The physical properties of these elements have not yet been fully determined, but a few of them are given in the following table: Cr Mo w U Atomic weight 52.0 6.7 white 96.0 8.6 white 184.0 16.6 white > 2800° 238 5 18 7 Color Melting-point MANGANESE. 300. This element is widely diffused in nature. Its most im- portant minerals are pyrolusite, Mn02, hausmannite, Mn304, and rhodochrosite, MnCOa. The metal is of minor importance. It is best prepared by the Goldschmidt method,, i.e. by reducing pyrolusite with aluminium powder, when it is obtained as a regulus of brilliant 300.] MANGANESE. 45!) Iu8tre. Sp.g. = 7.2-8.0 ; m.-pt. = 1245°; b.-pt.= 1900°. It under- goes surface oxidation readily in moist air, which gives the regulus an iridescence, and whea finely divided dcconiposes boiling water. It dissolves in acids to form manganous salts. Manganese forms several series of compounds: the mangniinvs compounds of the type MnX2; the manganic compounds, ^InXs; manganic acid, ll-^^lnOi, which can be derived from an anhydride MnOs; pcnnanganic acid, H^In04 derivable from the oxide JI110O7 Jlost of the familiar salts of this element are derived from man- ganous oxide, MnC). This oxide, which is prepared by heating the carbonate in the absence of air, is an amorphous green powder, that oxidizes readily in the air to the higher oxide MnsO^. Manganous hydroxide, ^In(0II)2, is white when freshly precipitated from solu- tions by an alkali but soon turns brown in the air because of the formation of manganic hydroxide, j\Ino(0H)6. The solutions of manganous salts are pink (color of the Mn"- ion). The chloride, MnCU, crystallizes with four molecules of water. It can be obtained anhydrous by heating the double salt MnCl2 -2X11401 + 1120, since the hydrochloric acid set free hinders the hydrolytic dissociation of the chloride. The sulphate, ]*InS04, crystallizes below 6° with 7H2O, above this temperature with 5H2O. It forms double salts, such as K2S04-iInSO4+6H20, simi- lar to those of magnesium and iron; they are moreover isomor- phous with the latter. Manganous sulphide, ]\InS, has a pinkish-white color, which distinguishes it from all other sulphides. If ammonium chloride is added to the solution of a manganese salt, no hydroxide is precipitated by ammonia; this is analogous to what is observed with magnesium (§ 254). The solution is, howe\er, readily oxidized by the oxygen of the air and brown manganic hydroxide is deposited. The manganic ion Mn"' is only weakly basic. Its salts are almost completely hydrolyzed in aqueous solution. Tlie sulphate gives alums with ciesium and rubidium sulphates, which are also very unstable. Manganic oxide, Mn203, is obtained from any of the other oxides by heating in an oxygen current. Since dilute sulphuric acid reacts with it, giving manganous sulphate and manganese dioxide, the oxide Mn203 is often considered as MnO -11002. The 460 INORGANIC CHEMISTRY. [§§ 300- corresponding hydroxide is soluble in cold hydrochloric acid to a dark-brown solution. It is not certain whether this solution con- tains Mn2Cl6 or MnCl2 and MnCU; on being warmed it gives off chlorine and is then known to contain the manganous chloride. Mangano-manganic oxide, Mn304 or MnO • Mn203, is obtained on strongly igniting the other oxides in the air. It is a brownish- red powder. When heated with hydrochloric acid it yields chlorine. Manganese di- (or per-) oxide, Mn02, the best-known man- ganese mineral (pi/rolusite) , is commercially of great importance in the production of chlorine. In the cold it dissolves in hydro- chloric acid to a very dark liquid, probably containing the tetra- chloride, and gives off no chlorine; when warmed it decomposes into chlorine and manganous chloride (§ 25). Since pyrolusite is comparatively expensive; various methods have been devised for reconverting the manganous chloride into the peroxide. One which is of practical importance is the Weldon process. An excess of milk of lime is added to the chloride solution, whereupon air is forced through the warmed liquid. The manganous hydroxide which is precipitated undergoes oxidation and is con- verted into calcium manganite, CaMnOs (=CaO-Mn02), which settles down as a black slimy mass: MnCla + 2CaO H- = CaMnOg 4- CaClg. The calcium chloride solution is run off and the manganite is used for generating chlorine, since it acts towards hydrochloric acid like a mixture of lime and manganese dioxide. The value of the peroxide depends on the amount of chlorine it can produce with hydrochloric acid. In order to determine this, the mineral, finely pulverized, is warmed with hydrochloric acid and the evolved chlorine passed into potassium iodide solution, whereupon an equivalent amount of iodine is hberated. This iodine can be titrated with thiosulphate (§ 93). Manganic acid and Permanganic acid. 301. When manganese compounds are fused with potassium hydroxide in the air or, better, in the presence of an oxidizing-agent (potassium nitrate or chlorate) a green mass results, which is dis- solved by cold water, forming a dark-green solution. On evaporat- ing this solution in a vacuum dark-green rhombic prisms of potas- 301.] MANOANIC AND PERMAXGAXIC ACIDS. 461 sium manganate, K2M11O4, crystallize out, which have a metallic lustre and are isomorphous with potassium chromate. They dissolve in potassium or sodium hydroxide solutions without change, but are decomposotl by water with the separation of man- ganese dioxide and the formation of potassium permanganate Kiln04, the latter giving the solution a deep violet color: SKa^f n04 + SHgO = 2KMn04 + :\In02 • HgO + 4K0H. On account of these changes of color the manganate solution recei\'ed the name chamcvleon minerale, from the early chemists. Both in the solution of a manganate and in that of a perman- ganate we have the anion I\[n04; in the fomier, however, it is bivalent, in the latter univalent. This causes the difference in the properties of the two ions ; the univalent ion ]\In04' is deep red and resembles the perchloric acid ion in behavior, while the bivalent Mn04" is deep green and displays analogy to the SO4" ion of sulphuric acid. The bivalent ion ^[1104" is only stable in alkaline liquids; it is converted by water (more easily by acids) into the univalent ion: SKa^nOi + 4HNO3 = 2KMn04 +Mn02 + 4KNO3 + 2H2O, or, written in ions: 6K- +3Mn04" +4H" + 4NO3' = 2ir + 2iln04' + ilnOa + 4K' + 4N03'+ 2H2O. The reaction obviously amounts to a formation of water by the four hydrogen ions and two oxygen atoms which they extract from a bivalent anion Mn04", the latter being reduced to MnOa. Of the four negati\e charges which are required to neutralize the four positive charges of the hydrogen ions two are taken from this jMn04 anion, which is reduced to Mn02, and the remaining two from two other bivalent anions ]\In04", which thus become univalent. The transformation of potassium manganate into the permanganate is effected commercially by passing ozone into its concentrated solution: 2K2Mn04 + 03 =2KMn04+K20 + 02. The permanganate crystallizes out of the solution and the resulting mother-liquor can at once be used with a fresh ciuantity of p^Tolusite to prepare more manganate. Potassium permanganate, KMn04, cry.^tallizes in beautiful glistening greenish-black prisms of the rhombic system, which 462 INORGANIC CHEMISTRY. [§§ 301- dissolve readily in water, forming a deep-violet liquid. This salt is isomorphous with potassium perchlorate. All solutions of per- manganates display the same absorption spectrum, viz., five dark bands in the yellow and green, no matter what the base is. It is thus e\'ident that the ion ;Mn04' is really the coloring-agent. The solution of potassium permanganate acts as a powerful oxidizing-agent; in acid solutions two KMn04 molecules yield five oxygen atoms : 2KMn04 + 3H2SO4 = K2SO4 + 2MnS04 + SHgO -f- 50. The process may be regarded as a transformation of the anhydride of permanganic acid, Mn207( = 2HMn04— H2O), into two mole- cules of basic oxide, MnO, and five atoms of oxygen; thus: Mn207 = 2MnO-|-50. In neutral or alkaline solutions, however, two KMn04 molecules yield only three atoms of oxygen, manganese peroxide being deposited at the same time (transformation of Mn207 into 2Mn02-|-30): 2KMn04 + H2O = 2Mn02 + 2K0H + 30. Since in oxidations with potassium permanganate in acid solu- tion the deep color of the permanganate is replaced by the very faint color of manganous sulphate, many substances can be titrated with potassium permanganate in acid solution without an indi- cator. Ferrous sulphate is oxidized to ferric sulphate; oxalic acid goes over into carbon dioxide and water; nitrous acid in very dilute solutions is converted into nitric acid (§ 126) ; from hydrogen peroxide water and oxygen gas are formed. AU these reactions proceed quickly and quantitatively at ordinary temperatures so that they are suitable for titration. Permanganic acid is known only in aqueous solution; however, its anhydride, Mn207, can be obtained. It is prepared by carefully treat- ing dry permanganate with concentrated sulphuric acid. It is a vola- tile, bro\vnish-green, oily liquid, whose vapor explodes easily, yielding oxygen and manganese dioxide. Manganese occupies an isolated position in the periodic system. No elements are known which are related to it as the elements Mo, W and U are to chromium. Moreover, only in its highest stage of oxidation, permanganic acid, does it display analogy with 302.] IRON. ICIi the corresponding chlorine compound, HCIO4. The salts of both acids are isomorphous and both are powerful oxidizing-agents. IRON. 302. Iron is the most useful metal, and is therefore prepared commercially on an enormous scale (approximately 50,000,000 metric tons a year). It occurs only T&Tely native , e.g. in meteoric rocks. In the form of oxides, sulphides and silicates it is widely diffused in nature and is found in very large quantities. The most important minerals for the iron industry are nvignctite, Fe304, hcmatile, FC..O3, and sidcritc, FeCOs. The pyrites (FeS2, etc.) are worked up into iron after they have been roasted in the sulphuric acid factories. The metallurgy of iron is theoretically very simple; it is liased on the ability of carbon to reduce the oxides of iron to the metal at an elevated temperature. This process (smelting) is car- ried out in blast furnaces. The iron ore is first roasted (calcined) to remo^^e volatile substances (H2O, COo, S, As, etc.) and loosen up the mineral. Then it is crushed and mixed with a slag- forming suljstance (flux, see § 242), according to the grade of the ore. If the gangue, or earthy matrix, contains much silica or alumina, limestone or dolomite is employed as the fluxing-agent, but ores rich in lime or magnesia are mbj:ed with quartz or aluminous ore to effect the necessary fusion and formation of slag (silicates of Al, Jig and Ca). The blast furnace, previously warmed to the proper tempera- ture or already in operation, is charged from above with alternate layers of coke and the mixture of ore and flux, both being intro- duced in "rounds," or "charges," of definite weight. (Sometimes charcoal or anthracite is used as fuel.) The modern furnaces (Fig. 73) are built of fire-brick encased in iron and are of much lighter construction than those formerly used. They vary greatly in size but consist mainly of a long shaft tapering towards licith ends. In order to utilize the escaping hot gases iCO, etc.) an apparatus ("cup and cone") is fitted on the top to conduct them off and also allow the introduction of the charge. The air necessary for the process is forced in, hot, through pipes {Iwaers) at the bottuni 464 INORGANIC CHEMISTRY. [§§ 302- of the furnace. The burning coke produces carbon monoxide, which is the principal factor in the reduction of the ore: FeaOs + SCO = 2Fe + 3CO2. The reduced iron sinks downward and comes in contact with carbon at a high temperature; as a result some of the carbon is Fig. 73. — Blast Furnace. dissolved by it and its melting-point considerably depressed. When a definite stage is reached the fused iron is drawn off below. It is protected from atmospheric oxidation by the slag floating on it. 303.] mux. ■ 46.') The attempt tt) extract iron from its ores by electrical heating has met with success. Stassano calculates on the basis of the analj-sis of the ore the additions which will be necessary to yield a slag as nearlj^ as possible of the composition SiC)2 + 4 Base, com- presses the finely powdered material to briquetts with the aid of tar in hydraulic presses, and smelts it in a specially constructed arc furnace. 303. It was stated above that the waste furnace-gases corttain a con- siderable quantity of carbon monoxide ; therefore a large amount of heat is lost, which could be utilized by burning the monoxide to dioxide. Supposing that this incomplete reaction was due to an incomplete con- tact of carbon monoxide and ferric oxide, manufacturers increased the dimensions of the blast furnaces, particularly in England and America, a height of thirty meters being not uncommon. The ratio of carbon mon- oxide to the dioxide in the escaping gases was not affected however; it was thus demonstrated by very expensive experience that the reduction of ferric oxide b)' carbon monoxide has a limit. A study of the laws of chemical equilibrium would have led to this conclusion much more* quickly and above all much less expensively. These laws teach us that; ] . In the reduction of ferric oxide bj' carbon monoxide an equilibriun* is established between this action and the oxidation of iron by carbon dioxide. FeA + .3C0^2Fe + SCOj. 2. The ratio COiCOj must be independent of the pressure, since no change in the volume of the gas takes plai-o (§ 51). 3. This ratio ^'a^ics only slightly with the temperature, since very httle heat is generated in the reaction. An experimental investigation conducted at a few different tempera- tures and pressures would have sufRced to determine the ratio C0:C02. The result, when compared with the ratio COit'Oj of the waste gases, would thus have shown that little could be gained by an increase of the furnace dimensions. This illustrates in a very striking way the value of physical chemistry for industrial processes. Efforts are now being made to utilize the waste gases in other ways, such as by burning them under the boilers of steam engines or in wind heaters (for heating the blast air, or "wind"). In recent years it has been found that greater efficiency is attained by using the hot waste- gases directly in gas engines for motive-power. 463 • IXOROANIC CHEMISTRY. [§ 304- 304. The properties of iron are influenced in great measure by the slight admixtures which it contains, particularly by the carbon. The percentage of carbon forms the ordinary basis of classification of the different grades of iron under the heads, pig iron and malleable iron; however, in the industrial world this classification is not always adhered to. Pig iron, or cast iron, contains 2.3-5.1% carbon. It fuses very easily but there is no previous softening; hence it is not malleable. It is brittle. Pig iron is the direct product of the blast furnaces- and the iron is therefore mixed with small amounts of silicon, phos- phorus, sulphur, etc. The presence of manganese makes it coarsely crystalline and it is then known as spiegel-eisen. This is utilized mainly for steel. Refined iron, containing less than 2.3% carbon, is harder to fuse, but is extensible and malleable, and the more so the less the impurities. If the carbon amounts to 2.3-0.5%, the iron can be hardened; in this manner stcd is obtained. If there is less than 0.5%, carbon, it can no longer be hardened; this is wrought iron. It is obvious that between these main varieties there are numerous intermediate sorts, which are prepared in such a way as to suit the purpose for which they are intended. The immense commercial importance of the iron-carbon system has led to extensive investigations regarding it, notwithstanding that such investiga- tions are attended by great experimental difficulties, partly because of the very high temperatures involved. Because of these difficulties it is not yet possible to give an entirely satisfactory representation of the equilibrium conditions ooncemed. Bakhxjis Roozeboom, Chaepy, Robert-Austen and others have succeeded in workmg out the accompanying graphic representa- tion which indicates the behavior of the system in the main at least. To appreciate this diagram it is necessary in the first place to know a few general facts regarding the components that are now regarded as existing in the iron-carbon system. Distinction is made between: 1. ferrite, or chem- ically pure iron (pure wrought iron); 2. martensite (steel), a solid solution of carbon in iron. It is so regarded because microscopic studies have shown that martensite is always homogeneous in spite of its changing carbon content, which may be as high as 2%. 3. cementite (the commercial white cast iron), an iron-carbon compound of the formula FcjC; and 4. perlite, carboniferous iron (0.85%), that is seen under a high-power microscope to be heterogeneous and is regarded as a eutectic mixture of ferrite and cementite. The solidification curve of a binary system (§ 237) does not take a nor- mal course in the iron-carbon system. Three circumstances complicate the 304.] IRON. 467 situation. The first is that pure iron does not separate out of the molten mass, but that we obtain the soUd solution martensite. The second is that changes continue to occur in the cooling mass after complete soUdification ; the third that other substances separate out with very slow cooling than with quick cooling. We may consider first the case of slow cooling, where the equilibria that establish themselves between solid and liquid phases are presumably stable: Let us assume that we have liquid iron with a carbon content below 4.3%. On cooling the liquid the iron begins to solidify at a definite temperature (the point 1 1 in Fig. 74) ; however, it is not pure iron, but a solid solution of 1500° A 1 1 / 1 1 1 1 1 [ 1 I 1 1 1 1 1 1 -1 1 1 1 1 1 1 1 J 1 -1 1 1 1 B ,1000° G / Cj "■ ~ \ 4 1 t 700" 2 a - El 1 1 t 1 1 1 . 1 1 . ,1.1 . 1 r 1 1 1 1 r . r , 1 1 1 1 1 ^c- Fig. 74. — Irox-Carbon System. carbon in iron that separates out; its composition is shown in the diagram by the point dp If the carbon content of the fused iron is a different one, we have separating out at i\, for example, the solid substance, whose com- position is again given by the point d^. Thus for e\'ery solidification point of the curve AC we can find a point d„ d^, etc., that gives the composition of the sohd substance which begins to separate out. The curve AD is the geometrical locus uf these points. If, therefore, a horizontal line is drawn through the triangle ADC, the point c gives the composition of the liquid solution which solidifie's at the corresponding definite temperature (indicated by the f)rdinate) and the point d the composition of the sohd solution which begins to separate out at that temperature. At C the eutectic point is reached. Along CB graphite separates out; at C itself a mixture of graphite and martensite, the composition being given by D. The point C is at 1130° and 4.3% of carbon. The martcntite formed at this tem- perature contains 2"; of carbon. 46S IXORGANIC CHEMISTRY. [§§304- Below DC all is solid; but, as we have already explained, changes con- tinue to occur in the solid mass. For example, if martensite is heated, it breaks up with the formation of graphite. The curve DE represents the change of composition of the solid solution with falling temperature or, in other words, it represents the equilibrium between graphite and the mixed crystals (solid solution) at different temperatures. Around E, where the temperature has reached about 700°, the martensite contains only about 0.85% carbon. At the point E the formation of ferrite begins. Finally, the curve EG indicates the composition of the solid solutions from which ferrite separates out. Hence, if the martensite contains less than 0.85% of carbon, ferrite is deposited along EG, exactly as ice separates out of a dilute salt solution with falling temperature. It the cooling is sudden, other phases are formed, the limits of which are represented in the figure by lines, which are readily under- stood. Instead of the eutectic point C, at which graphite and martensite separate out, we have a eutectic point at Cj, very close to C, where cementite, FcjC, separates out with the martensite. Further, the line C^E^ represents the equilibrium between cementite and the mixed crystals (martensite). At .E, ferrite is formed together with cementite. Martensite changes over at this temperature into a eutectic mixture of these last two substances, which has the fine conglomerate structure so characteristic of eutcctics and is known as " perlite." Although this whole system shown by lines is meta- stable, it can exist for an indefinite period at ordinary temperatures because of the reduction of the velocities of reactions which might restore the stable forms. It is evident from the above that with slow cooling martensite entirely disappears. If the cooling is rapid, however, as in the hardening of steel, martensite can be brought to exist at ordinary temperatures even though it is in a metastable condition; its transformation velocity is then extremely small. If the hardened steel is reheated, it changes over partially into the soft conglomerate of ferrite and cementite; this is what takes place in the "tempering" of steel. Small admixtures of other elements have an effect on the jjroperties of iron equally as great as that of carbon. The presence of silicon has about the same effect as that of carbon, but it is less intense. Sulphur even in a small amount renders the iron brittle when hot and, therefore, useless for forging. On this account sulphurous ores as such are unsuitable for the manu- facture of iron. Phosphorus makes the iron brittle at ordinary temperatures. It should also be mentioned that as a general rule the effect of these admixtures is strongly modified by the pres- ence of others. 305. From the crude pig iron, the direct product of the blast furnace, the other varieties of iron are prepared. For this purpose 305.] IRON. 46» it must be freed from silicon, sulphur, phosphorus, etc., as well as from a large portion of its carbon. The most important process for accomplishing this commercially is the Bessemer process. The pig iron is fused and run into a pear-shaped apparatus, or con- verter (Fig. 75), in the bottom of which are holes through which air is blown in. Thus by the oxidation of silicon, manganese and a little iron and without the use of fuel the temperature is raised high enough to effect the burning of the carbon. The Bes- semer process is easier controlled if the elimination of carbon is continued past the steel stage and until molten wrought iron is formed, whereupon enough carboniferous iron is added to furnish Fig. -Converter. steel with the desired percentage of carbon. .Vt the completion of the process the converter is emptied 1)}' tipping. In some European mills n basic converter lining containing an excess of lime and magnesia is used. The phosphorus in the ore combines with the bases to form phosiiliatcs, whicli enter the slag, and this so-called "Thomas-slag" is used in large quantities as a fertilizer. The only successful rival of the Bkssemer process is the Siemens, or open-hearth, pnicess. By employing a special furnace and gaseous fuel a mixture of cast iron and wi'ought iron (together with some iron ore) in the proper proportions can be fused together 470 IXORGAXIC CHEMISTRY. )305. A basic lining can also be so as to produce a vcrj- good steel, used with this process. The increased demand for special steels, where physical and chemical conditions have to be regulated carefully, has given greater significance to the old crucible process, the steel being made in graphite crucibles in a laboratory manner but on about ten times the laboratory scale. Recently electric furnaces of the arc and induction type have been found very successful in pro- ducing '' crucible " steel and with much less labor than the Fig. 76. — Heroult Furnace. crucible process requires. The cradle-shaped Heroult furnace is shown in the accompanying combined end-view and vertical section (Fig. 76). M is the molten metal, S the slag, and E one of the carbon electrodes; B is brick lining and L a layer of magnesium silicate. As the resistance of the metal is small- compared with that of the slag and the air, most of the heat is generated at the surface, where the chemical action goes on between the slag and the metal. The furnace is eventually emptied Ijy rocking forward. Steel, however made, is a very complex alloy, containing 305] IRON. All carbon and manganese, O.lO-l.SO';^,; silicon 0.02-0.25%, sulphur and phosphorus 0.01-0.10','; and possibly copper, arsenic, alu- minium, oxygen, nitrogen, and cyanides, and is capable, as has been explained, of containing the iron and carbon in various combinations. Steel of the above description is " ordinary " steel. Recently a large market has developed for " special " steels, having new qualities, especially with respect to hardness and brittleness, and ser\'ing new purposes, notably in tools, military materials and materials of construction. They may be produced by (1) changing the physico-chemical character with respect to the iron-carbon system, (2) removing harmful occluded gases, (3) combining other elements chemically with iron or carbon or both, and (4) adding other elements to form isomorphous solutions with iron. Steel brccimes very hard and brittle, for instance, when it is suddenly cooled from a high temperature. If, however, it is then luxated for a definite ]jeii(xl and allowed to cool slowly, it becomes more or less tempered according to the temperature, i.i'.. it can be made to have any desired hardness and elasticity (within certain limits). Of the specidl (lUoy stech the nickel steels, chrome-nickel steels and chrome-vanadium steels seem to be most important. The maximum hardness of steel is reached when it contains 1-2' , carbon; if, however, some manganese (up to S%) or chro- mium (up to l*-'}) is added, a much harder modification of steel is produced. The addition of nickel gives a tougher nivvl, which is especially valuable for armor plate. Tungsten (cf. § 2'.)7) and molybdenum are also added for different purposes. In an>- case, however, a careful heat treatment is essential to develop the desired properties. The production of wrought iron from pig iron is usually accomplished by the puddling pnici'Kn. Pig iron is melted in a r('\-erberator>' furnace lined with iron ore (oxide) ; the carbon and also the silicon are oxidized (and so removed) partly by the action of the aii-, but mainly by that of the ore, which is stirred in with the metal. The violent reaction due to escaping carbon monoxide gi\('s the process the name of "pig- boiling." The iron is tiien allowed to become pa.sty, when it is worked up into large masses (Idaiinix), wliicii are rcnioNi'd and ham- mered and rolled. The cinder is thus sciucezcil out and the iron is formed into bars. 472 IXORGANIC CHEMISTRY. [§§ 305- Chemically pure iron is obtained electrolytically and by reduc- ing the oxide or chloride in a current of hydrogen. If the reduc- tion takes place at a low temperature, the resulting iron powder is pyrophoric (§ 203). It is a silverj'-white, lustrous metal with a specific gravity of 7.8-1 and a melting-point as high as 1520°- It boils at 2450°. It is the most magnetic of the metals; pure iron and wrought iron can be magnetized only temporarih' ; steel, however, permanently. Iron is permanent in dry air or in water free from air (COo). In moist air it rusts rapidly (§ 279), forming ferric hydroxide; as the rust does not form a compact film, it keeps on forming. The rusting of iron is greatly retarded by contact with water contain- ing a little alkali or salts of alkaline reaction. In a soda solution, for instance, iron remains bright. The rusting of iron in contact with water can be explained by assuming that the oxygen dissolved in water endeavors to form hydroxyl ions \vith the hydrogen ions. In order to compensate their negative potential the iron sends its positive ions into the solution; in a short time the solubility product of ferric hydroxide is reached and the latter is deposited; in other words, the iron rusts. Kow, if hydroxyl ions are previously introduced into the liquid by the addition of a base or a salt of alkaline reaction, the ionization of the water is diminished so much that the oxygen can find almost no hydro- gen ions with which to form hydroxyl ions; therefore the iron does not send any more ions (§§ 276 and 277) into the solution and rusting is greatly retarded. Iron dissolves readily in hydrochloric and sulphuric acids with the evolution of hydrogen. At red-heat it decomposes water, but the oxide is also reduced by hydrogen, so that an equOibrium results : 3Fe + 4H20?^re304 + 4H2. In nitric acid (not too concentrated) iron dissolves readily with the evolution of nitric oxide, NO, but if the iron is first dipped in concentrated nitric acid and then rinsed off it becomes indif- ferent to the action of nitric acid. Th^s so-called "passivity" of iron is probably caused by a very thin coating of oxide. Iron forms two sets of salts, the ferrous and the ferric. 306.] FERROUt< (XniPOUXDS. 473 Ferrous Compounds. 306. In the ferrous condition iron has only basic properties. Ferrous oxide, FeO, is obtained by reducing ferric oxide with carbon monoxide. It is a black powder, which oxidizes easily on warming. Ferrous hydroxide, Fe(0H)2, is precipitated from ferrous salt solutions as a pale green gelatinous substance by the addition of an alkali; it oxidizes very rapidly in the air to ferric hydroxide. Ferrous chloride, FeCl2, is formed on dissolving iron in hydro- chloric acid; it crystallizes from this solution in green monoclinic prisms containing four molecules of water. The anhydrous salt is obtained as a white sublimate when iron is heated in dr}- hydro- chloric acid gas. With potassium chloride and ammonium chlo- ride ferrous chloride forms well crystallized double, salts, e.g. FeCl2-2KCl + 2H20. Ferrous sulphate, reS04-l-7H20 (green vitriol, copperas), is the most familiar ferrous salt. It is prepared commercially, princi- pally by dissolving up the waste metal of steel-wire factories in sulphmic acid, but also by partially roasting pyrite, whereby fer- rous sulphide, FeS, is formed; the latter is left exposed to the ah-, when it oxidizes gradually to ferrous sulphate, which can be dis- solved out. It crystallizes in large, bright green, monoclinic prisms, which effloresce slightly and at the same time become coated with a brown layer of basic ferric sulphate. The double salts such as FeS04-(NH4)2S04-l-6H20, Mohr's salt, are not so liable to oxi- dation; for this reason use is frequently made of Mohr's salt to standardize permanganate solutions (§ 301). Iron vitriol has numerous uses, e.g. for making inlt, in dyeing, as a disinfectant (it absorbs both ammonia and sulphuretted hydrogen and is there- fore used to dispel bad odors), etc., etc. Ferrous carbonate is somewhat soluble in water containing carbonic acid and Ls therefore often present in natural waters (§ 17). The basic carbonate which is precipitated from a ferrous solution by soda oxidizes rapidly in the air to ferric hydroxide. TJie latter is also deposited from chalybeate waters on standing in the air for a time. Ferrous carbonate is only known as a mineral {siderite, § 302). 474 INORGAXIC CHEMISTRY. [§§307^ Ferric Compounds. 307. The ferric ion has only very slightly basic properties. Ferric salts of weak acids, such as carbonic acid, do not exist. In aqueous solution most of the ferric salts, even those of strong acids, are partially hydrolyzed. For that reason they are brown- ish-red, since this is the color of ferric hydroxide in colloidal solu- tion. On the addition of an excess of sulphuric or nitric acid this color disappears, because there is no longer any hydrolysis. From this it appears that the ferric ion itself in aqueous solution is only slightl}' colored. The ferric salts are readily converted into ferrous salts by reducing-agents. Ferric oxide, Fe203, iron sesquioxide, is formed on heating various iron compounds in the air and is manufactured by igniting green vitriol (§ 79). It is a dark-red powder and finds use as a pigment (colcothar) and in polishing glass, etc. Ferric hydroxide separates out as a reddish-brown hydrogel, Fe203+MH20, when a ferric salt solution is treated with an alkaU. The freshly precipitated hydrogel dissolves in a solution of ferric chloride or acetate. If this solution is dialyzed, a pure colloidal solution of the hj'droxide is finally obtained; from this the h}'drogel is reprecipitated by a smaU amoimt of alkali or acid. Ferrous ferric oxide, Fe304, also called ferroso-ferric oxide or magnetic iron oxide, occurs in nature as magnetite. It is produced bj' heating iron in steam (§ 305). Ferric chloride is obtained by passing chlorine into a solu- tion of ferrous chloride. It crystallizes at different temperatures with different amounts of water, being an example of the case described on p. 341 . On heating the salt hydrochloric acid escapes with the water of crystallization. Anhydrous ferric chloride can be prepared by heating iron in a current of dry chlorine. Between 320° and 440° the vapor density is approximately that calculated for FesCle; between 750° and 1050° it falls to half, indicating a splitting off of chlorine or a dissociation into 2FeCl3. The reddish-brown color of the aqueous solution of ferric chloride must be ascribed chiefly to un-ionized FeCls molecules, for the salt has this same color when dissolved in ether, in which no ioniza- tion occurs. In part, also, this color comes from ferric hydroxide, 308.] FEHliie COMPOUXDS. J7.') which is formed by hydrolytic dissociation. This dissociation increases on warming the dilute af|iicous solution, for a very dilute, almost colorless solution of ferric chloride turns reddish-brown on boiling. Wlien cooled the liquid gradually resumes its original color. Ferric sulphate, obtained by dissolving ferric oxide in sulphuric acid, forms alums, e.g. potassium iron alum, K2S04-Fe2(S04)3 + 24HoO. "V\1ien a ferrous salt is converted into a ferric salt in aqueous solution the bi-\'alent ferrous ion is transformed into a trivalent ferric ion. The oxygen required for the conversion serves to oxidize the hydrogen ions of the acid (which must be added) to water, whereupon these hydrogen ions surrender their charge to the iron ions : 2(Fe" + 2a')-f-2(H-+Cl')-l-0=2(Fe--l-3Cl')+H20. Ferrous chloride HydrochJoiic acid Ferric ciiloride Inversely, the reduction of ferric salts to ferrous salts can be explained by supposing that everj' ferric ion gives up a third of its charge to another atom and thus makes the latter an ion or neu- tralizes its charge. Salts of iron are also known which are derived from the hypothetipal oxide FeOj. They are obtained by heating iron filings with saltpetre or ■passing ciilorine into an alkaline suspension of the ferric oxide hydrogeJ. From such solutions potassium ferrate, KjFeOi, crystallizes out in dark- red prisms, isomorphous with the chromatc and sulphate of potassium. These crystals are readily soluble in water, but their dark-red solution soon decomposes with the separation of ferric hydroxide and oxygen gas. The free ferric acid is unknown. 308. Iron unites with cyanogen to form complex and unusually stable anions, viz., the fcirocyanic ion [Fe(CN)6]"" and the ferricyanic ion [Fe(CN)6]"'. Their best-known salts are potas- sium ferrocyanide, K4Fe(CN)6-3H20, and potassium ferricyanide, K3Fe(CN)6, the yellow and red prussiates of potash, respectively. The ionization of the complex it)ns themselves is so slight that they give none of the ordinary reactions for iron. For the commercial manufacture of yellow prussiatc of potash two processes are used: In the first, animal refuse (e.g. blood) is charred, yielding a black, highly nitrogenous mass. This is 470 INORGAXIC CHEMISTRY. [§§308- ignited with potash and iron fiUngs. After cooling, hot water is added and the mixture filtered; from this filtrate the yellow prussiate crystallizes out on standing. This salt is not formed until the ignited mass is treated with water, for yellow prussiate is decomposed by heat and cannot therefore be present in the ignited mass. The latter probably contains potassium cyanide, iron and iron sulphide (animal refuse always contains sulphur compounds"). These substances can interact according to the equations: 6KCN+ FeS = K4Fe(CN)6 + KaS; 2KCN+Fe + H20=Fe(CN)2 + 2KOH+H2^ re(CN)2 + 4KCN = K4Fe(CN)p. The second process is employed in illuminating-gas factories, for the unpurified gas contains a little cyanogen and prussic acid. After being freed from tar and ammonia it is passed through a washer (scrubber) containing a solution of potash in which ferrous carbonate (ferrous sulphate + potassium carbonate) is suspended. The following reactions, among others, are known to go on here: FeCOs + 2HCN^Fe (CN) 2 + HgO + CO2 ; K2CO3 + 2HCN?^2KCN + H2O + CO2. Notwithstanding that these reactions are reversible, the hydro- cyanic acid can be quantitatively fixed in this way, because the ferrous cyanide and potassium cyanide interact to form potassium ferrocyanide, which is but very slightly affected by carbon dioxide. Potassium ferrocyanide, K4Fe(CN)6-3H20, forms large sulphur- colored crystals. Its three molecules of water can be expelled by gently warming, whereupon the salt is left as a white powder. It is not poisonous. With dilute sulphuric acid it produces prussic acid on warming; with concentrated sulphuric acid it yields carbon monoxide. The free ferrocyanic acid, H4Fe(CN)6, separates out as a white crystalline precipitate when concentrated hydrochloric acid is added to a strong solution of potassium ferrocyanide. The pre- cipitate soon turns blue in the air on account of the formation of Prussian blue (and partial decomposition as well), ^'arious salts of this acid have characteristic colors and are insoluble; hence potassium ferrocyanide finds use in analysis. It is an interesting 509.] COBALT AXD MCKEL. 477 fact that this compound of iron can serve as a distinguishing reagent for ferrous and ferric compounds. The ferrous salt of ferrocyanic acid is white, but in the presence of air it passes rapidly over into the blue ferric salt {Prumnn blue — a valuable pigment). The copper salt (§ 40) is brownish-red, the zinc salt white, etc. Sodiiun nitroprusside, Na2Fe(CN)5(NO) •2H2O, is formed by the action of nitric acid on sodium ferrocyanide. It crystallizes in ruby-red prisms and is a delicate reagent for alkali sulphides, whose solutions it colors violet. Potassium ferricyanide, KaFeCCNlg, red prussiate of potash, is formed from the j'ellow prussiate by treating a solution of the latter with chlorine or bromine: K4re(CN)6 +C1 = KCl + K3re(CN)6. It appears in dark-red crystals, which are readily soluble in water. The aqueous solution is unstable. The salt is often employed as an oxidizing-agent in alkaline solution, being itself converted into the ferrocyanide: 2K3Fe(C'X)6 -I-2KUH = 2K4Fe(CN)6 -HHaO -f 0. Iron forms some very peculiar compounds with carbon monoxide: Pe(C'0)4 and Fe(C0)5. They are produced when carbon monoxide is passed over finely divided iron at 80°, or at ordinary temperatures if the gas is under pressure. Iron vessels which have held compressed illuminating-gas for some time are more or less attacked by the carbon monoxide of the gas, for if gas which has been kept in such a vessel is allowed to escape through a hot glass tube an iron mirror is formed on the inside of the tube. COBALT AND NICKEL. Cobalt. 309. The two best-known minerals of this metal are smaltite, C0AS2, and cobaltitc, or cobalt glance, CoAsS. The metal is ob- tained by calcining these minerals and reducing the resulting cobalto-cobaltic oxide, Co:jC)4, with carbon (or hydrogen). It has a pink color and a high lustre. Sp. g. s.fl; m.-pt., 1490°. It is magnetic but much less so than iron. It is indifferent to the air. Hydrochloric and sulphuric acids dis.soh-e it very slowly but it readily forms a nitrate with nitric acid. 47S IXORGAXIC CHEMISTRY. [§§309- Besides the oxide, C03O4, just referred to there are two others, cobaltous oxide, CoO, and cobaltic oxide, C02O3. The salts are all cobaltous, corresponding to the bivalent ion Co". COBALTOUS COMPOUNDS. The solutions of the salts are red; hence this is the color of the cobalt ion. The non-ionized cobalt salts are blue, e.g. the anhydrous C0CI2, the silicate, etc. This difference in color enables us to tell readily ^Yhether a cobalt salt in solution is ionized or not. Thus in concentrated solutions, for instance, all those cir- cumstances which reduce the ionization cause a change of color from red to blue, e.g. when a concentrated cobalt chloride solu- tion is warmed or treated with hydrochloric acid. That the ioni- zation is diminished by warming was mentioned in connection with cupric chloride (§ 244). Cobaltous chloride, CoCU ■ 6H2O, forms red monoclinic crj's- tals, which turn blue on heating because of dehydration. Cobalt sulphate, C0SO4 ■ 7H2O, is obtained in dark-red monoclinic prisms and is isomorphous with reS04 -71120. It forms double salts with alkali sulphates, e.g. K2SO4 -00804 + 61120. Cobalt nitrate, Co(N03)2-6H20, appears in red hygroscopic prisms. Cobalt sili- cate is A'ery deep blue; hence its use for coloring glass. Pulverized cobalt silicate serves as a pigment (smalt) in painting, etc. The- nard's blue is a pigment, obtained by igniting cobalt salts with alumina. COBALTIC COMPOUNDS. 310. Cobaltic oxide, C02O3, is obtained by igniting cobalt nitrate. It is a black powder, which passes over into cobalto- cobaltic oxide, C03O4, at red heat and at white heat yields cobaltous oxide. It has the character of a peroxide; for by the addition of sulphuric acid it is converted into a cobaltous salt with the evolu- tion of oxygen and it yields chlorine with hydrochloric acid. How- ever, in cold dilute hydrochloric acid it dissolves without generat- ing scarcely any chlorine. Like iron, cobalt also forms complex ions, of which those with cyanogen are ^•ery stable. There are cobalt salts corresponding in composition to the yellow and the red prussiates of potash; the salt K3Co(CX)6, potassium cobalticyanide, crystallizes in colorless 311.[ COBALT AND MCKEL. 47!> rhombic prisms. A peculiar complex ion occurs in the potassium cobaltic nitrite, 6KN()a('02(N02)6 + ''H,,(), or K3Co(N02)6 + /(H2O. It is formed on treating a solution of a cobalt salt with potassium nitrite and acetic acid. It is a yellow crystalline [ire- cipitate, which is very slightly soluble when potassium ions are present in excess in the liquid. Cobalt also forms numerous comple.x ions with ammonia (§ 317). Nickel. 311. Xickel occui-s in nkcolite, Xi.\s, and nickel glance, or gersdorffite , XiAsS, Especially important is the nickel sihcatc,, garnierile. H2(Xi,:\I<;)Si()4 +aqC.>), which was discovered l)y (Iak- NiER in Xew Caledonia, where it occurs in enormous quantities. Canada, too, has some rich nickel deposits. From this ore the nickel is obtained by a l:last -furnace process similar to that for iron. The discovery of garnierite marked the beginning of a new era in the nickel industry. Mucii nickel is refined electro- lytically. Xickel is almost as white as silver, is very tough and has a high metallic lustre. Sp. g. = 8.S-9.1; m.-pt.= 14.52° It is feebly magnetic. It dissolves sparingly in hydrochloric and sulphuric acids but freely in nitric acid. It is permanent in the air. It is employed in nickel-plating metallic objects and as a con- stituent of several alloys. German silver contains about 50% copper, 2.5' f nickel, and 25% zinc. The nickel coins of Germany and the United States consist of 75% copper and 25"^'^ nickel. The use of nickel to vary the properties of iron has already been mentioned (§ 305). The o.xides of nickel, XiO and Xi203, are very similar to those of cobalt. The nickelous oxide, XiO, is the only one which forms salts. Nickel chloride, XiClo -61120, yields green monoclinic prisms. When heated it turns j-ellow on account of loss of water. Nickel sulphate, XiS().i-7H2<>, crystallizing in green rhombic prisms, is isomorphous \\ith the corresponding ferrous and other salts and also forms analogous double salts. Nickelic oxide, Xi203, also bciiav(\sas a peroxide; when warmed with hydrochloric acid it yields chlorine gas and nickel chloride. 480 INORGAXIC CHEMISTRY. [§§311- Nickel carbonyl, Ni(C0)4, is formed when carbon monoxide is led over finely divided nickel at ordinary temperatures. A state of equilibrium results here, viz.: Ni+4CO^Ni(CO)4, ■nhich is displaced to the left •nith rising temperature, since the decomposition of nickel carbonyl takes place with a considerable absorption of heat (§ 103). Even as low as 60° the decomposition is of an explosive nature. It follows from the above equation that an increase of pressure (§ 122) must greatly increase the propor- tion of nickel carbonyl formed. Experiments confirming this showed at the same time that both the formation and the decom- ]iosition of this compound are very sensitive to traces of foreign substances. Nickel carbonyl is a colorless, highly refractive liquid, which boils at 43° and congeals (crystalline) at —23°. When heated in the air it burns with a very sooty flame. This compound is of aid in extracting nickel from low-grade ores. Nickel also forms a complex ion with cyanogen. On dissolving nickel cyanide in an excess of potassium cyanide the compound K2Ni(CN)4is produced; it is, however, unstable, being decomposed by hydrochloric acid with the deposition of nickel cyanide, Ni(CN)2. 312. A pecuUar property is exhibited by the sulphides of cobalt and nickel, CoS and NiS. Hydrogen sulphide does not precipitate these sulphides from acid solutions, but, once precipitated (by ammonium sulphide), they are apparently not redissolved by dilute acids. This is contrary to the general rule of § 146 (see also § 73), for the sulphide should either be precipitated by hydrogen sulphide from a feebly acid solution (e.g. CuS), which is the case when the solubility product is very small, or else, when the solubility product is larger, it should dissolve in dilute acids, as is the case with ferrous sulphide. As a matter of fact, however, no real anomaly exists here, for the rate of solubility of these sulphides is only very slow binder the usual conditions of the reaction, viz., dilute acid and room temperature. It increases ■vsith the concentration of the acid, temperature of reaction and fineness of grain of the precipitate. Nickel sulphide is soluble in alkali sulphides immediately upon its formation, but when once deposited in the solid state it is insol- uble, or nearly so. This is seen when a nickel solution is treated 313.] PLATINUM METALS. 4s I with tartaric acid and then with an excess of sodium hydroxide, no nicl<;elous hydroxide being precipitated. If hydrogen sulphide is passed into this solution a very dark-colored liquid results, from which nickel sulphide is deposited only very slowly. The same is true of cobalt in very dilute solution; in concentrated solution, however, the cobaU sulphide soon passes over into the insoluble modification and separates out. PLATINUM METALS. 313. Under this head are included the metals ruthenium, rho- dium, ■palladium, osmium, iridium, and platinum. They occur only in metallic form and are associated together in mixtures or combinations. The principal deposits are in the Ural and Caucasu.s, but smaller quantities are also found in Colombia, Brazil and Borneo. The Ural yields il.V ^, of the total production. The most important of these metals is platinum. The platinum ores usually contain admixtures of iron, gold, etc. This group falls into two subdivisions : the light metals, ruthenium, rhodium and palladium, and the heavy metals, osmium, iridium and platinum. The two sub-groups differ con- siderably in atomic weight and specific gravity: Light. Heavy. Ru Rh Pd Os Ir Ft .\tomic weight.. . . Specific gravity. . . 101.7 12.2(i 103.0 12.1 106.5 11.9 191 22.4 193.0 22.38 194.8 21.45 A complete separation of the platinum metals from each other is extremely difficult, in the first place because their properties are very similar, and in the second place because their behavior is con- siderably modified by their mutual presence — a fact which indi- cates the existence of compounds with each other. Thus, for instance, platinum dissolves readily in aqua regia while piu'e iridium is insoluble in it ; nevertheless when an alloy of the two metals is treated with aqua regia, some of the iridium is carried into solu- 482 INORGANIC CHEMISTRY. [§§ 313- tion. Further, the presence of iron (which occurs in all platinum ores) is often verj^ disturbing; for example, pure platinum solu- tions are not precipitated by soda or barium carbonate, but if iron is present more or less platinum comes down with the hydroxide of iron. In spite of these difficulties platinum, palladium, rhodium and iridium can now be purchased in a remarkably pure state. For the manufacture of platinum and the other metals of the group in the pure state the various factories employ their OAvn secret methods. In general the procedure is about as follows: The ore is first treated with aqua regia to dissolve out the major part of the "noble" metals, lea\-ing in the residue the alloy iridosmine, besides more or less sand. The ore thus consists not of one alloy, but of two: the crude platinum and the iridosmine. Both contain all six platinum metals, although in different relative amounts. The crude platinum contains, besides platinum, principally palladium, rhodium, and iridium, while the iridos- mine, as its name indicates, consists mainly of iridium and osmium. It is comparative!}' easy to separate out the osmium and ruthenium; the>' form volatile oxygen compounds and can therefore be removed by distillation. Platinum and iridium give difficultly soluble com- pounds with ammonium chloride, which in turn are reduced to the metal form by ignition. However, if the solution of the crude platinum is precipitated with ammonium chloride, the resulting precipitate is found to contain considerable amounts of rhodium and palladium, and an extended procedure is necessary for the isolation of the pure metals. On the other hand, it is not possible to precipitate the platinum and iridium completely in this way; the filtrate from the ammonium chloride contains palladium and rhodium, with smaller amounts of iridium and platinum, and, in order to work it up, a further complicated procedure is necessary. Ruthenium. 314. This steel-gray metal occurs only in very small quantities; it is hard, very brittle, and very difficult to fuse, a temperature of at least 1800° being necessary. Even when finely divided it is but very sparingly soluble in aqua regia, forming RuzCl,,, but when alloyed with platinum, it dissolves readDy. The compound RuCl, is known only in double salts. As a powder the metal oxidizes in the air to RuO and RU2O3. Ruthe- nium also forms characteristic salts, in which it plays the part of an acid. Potassium ruthenate, KjRuO^, results from fusing ruthenium with caustic potash and saltpetre. It crystallizes with IH2O in black prisms of 314.] PLATIXUM METALS. 4-;] a greenish lustre. With water it forms a darlc orange-colored solution. Its conduct reminds one of potassium manganate, for under the influeu<'e of dilute acids it is converted into potassium perruthenate, KRuOj, with the simultaneous precipitation of a black oxide, RhjOj (or RuOjV). It crystallizes in black octahedrons of metallic lustre, which dissolve in water to a dark-green solution. A peculiar compound is the tetroxide, RuO^, which volatilizes when chlorine is passed into the concentrated solution of potassium ruthenate. It can be solidified by cooling, when it forms a golden crystalline mass, fusible at 2."). 5°. There is no acid corresponding to this o.xide. RUO4 is used for the preparation of pure ruthenium. Osmimn. This metal is very analogous to ruthenium; it melts as high as 2500°. The chlorides OsCU and Ost'l^ and the oxides ()s<), Os,03 and OsO^ are known. The great similarity to ruthenium is especially noticeable in the highest oxides. Thus fusion with caustic potash and saltpetre pro- duces potassium osmiate, K^OsO,, which crystallizes from aqueous solu- tion in dark-violet octahedrons containing two molecules of water. The characteristic osmium compound is the tetroxide, OsOj, formed by igniting finely powdered osmium in the air or by the action of chlo- rine on the metal in the presence of water. The aqueous solution of OSO4 reacts neutral, but is often (^wongly) called osmic acid. It is employed in microscopy since organic substances (i.e. reducing-agents') reduce it to black osmium. Xo salts derived from OSO4 are kno^ATi. This compound is used in preparing pure osmium. Rhodium. The metal in the fused state has the appearance of aluminium and is just as extensiljle (malleable and ductile) as silver. It is prepared pure in the arts by way of the chloro-purpureo rhodium chloride, Rh(XH3)-('l3 (f/. § 317). Neither acids nor aqua regia affect it. AMieii heated in the air it is oxidized to the rhodious oxide, RhO. It is able to absorb a considerable aimiunt of hydrogen. The rhodic oxide, RhjO,,, yields salts with acids. Of the chlorides only Rh/'iy is known; this is obtained by direct synthesis as a reddish-brown substance; it forms sol- uble double salts with the alkali chlorides. The most satisfactory thermocouple for measuring high temperatures is miide of pure platinum and an alloy of platinum and rho(iium. 4S-1 IXORGAyiC CHEMISTRY. [§§314- Iriditun. This very refractory metal is obtained from iridosmine by heating in a cmTent of oxygen, when the osmium volatilizes as tetroxide. In the form of a platinmn alloy it is employed in the manufacture of "platinum" crucibles, dishes, distUKng vessels for the concentration of sulphmic acid (§ 86), etc. The prototype of the meter at Paris is made of an alloy of 90% platinum and 10% iridium. The admixture of iridium makes the platinum more indifferent to chemical agents, although at high tem- peratures the A'olatihty of the iridium is often troublesome. When pure, iridium is not attacked by aqua regia. Iridium forms two chlorides, Ir2Cl|, and IrCl^. Both of them give double salts with the alkali chlorides; e.g. IrjClg-OKCl+GHjO and IrCl4-2KCl. The former dissolves in water readily, the latter with difficulty. The tetrachloride appears as a black substance forming with water an intensely red solution. For tliis reason a platinum chlo- ride solution which contains iridium has a much deeper color than a pure solution. Palladium. 315. The silvery-white metal fuses at 1549°, i.e. more easily than platinum. When finely divided it dissolves in boiling concentrated hydrochloric, sulphuric and nitric acids. On ignition in the air it is at first oxidized, thus losing its lustre, but at a higher temperature the metalUc lustre reappears. The most peculiar characteristic of the metal is its abihty to absorb hydrogen in large quantities (occlusion). Freshly ignited palladium foil absorbs 370 times its own volume of hydrogen at room temperature. By making palladium foil the cathode in a water electrolysis apparatus the metal can be made to take up even 960 times its own volume. This absorption does not alter its metallic appearance. The absorbed hydrogen can all be expelled by heating in a vacuum. Palladium charged with hydrogen is a strong reducing-agent; chlo- rine and iodine are reduced by it (see § 200) to hydrogen chloride and hydrogen iodide, respectively, and ferric salts are reduced to ferrous salts. Palladium forms two series of compounds, the -ous PdXz, and the -ic, PdXi. A characteristic compound of the first series is paUadious iodide, Pdl2, which is precipitated by potassium iodide from solutions of -ous salts as a black insoluble substance. This reaction is occasionally used to separate iodine from the other halogens, since their palladium compounds are readily soluble. — Palladic chloride, PdCl,, is produced by dissolving the metal in aqua regia. With KCl or NHjCl it forms a difficultly soluble double chloride, K^PdCl, or (NHJ^PdCle. On the e\-aporation of its solution PdCU dissociates into PdClj. and CI2. 316.] PLATLXUM METALS. JS.") Platinum. 316. This metal is the principal constituent of the platinum ores. It fuses at aliout 1760° and is extremely malleable antl ductile, hence it can be made into xovy fine wire and very thin foil. Since it becomes soft at red heat it can be easily worked. Platinum is used in the greatest variety of A\ays, among othei's the manufacture of utensils for the chemical laboratory', in tlic distillation of sulphuric acid, in electrical apparatus, in jewelry for settings of gems, in electric furnaces, in incandeseccnt lights and dentistry. The last named use consumes about one-third of the total production, ^^'hen finely divided it absorbs oxygen (partially combining with it), a property to which is attributetl the phenomenon that numerous oxidations proceed with unusual ease in the presence of platinum. Its use in gas-lighters depends on this property. When the metal is precipitated from its solutions by reducing-agents, it is frequently obtained as an extremely fine velvet-black powiler, platinum black, ^^■hen tlie duuble chloride ;XH4)2PtCl6 is ignited, the metal is left as a porous mass — platinum sponge. At red heat a platinum partition allows hydrogen to pass through, while other gases are held back. This is due to the forma- tion of a compound or to the solubility cf hydrogen in platinum. Various substances attack platinum at elevated temperatures, eg the hydroxides, cyanides and sulphides of the alkalies; hence these substances should not be fused in platinum vessels. This also applies to lead and other hea\-y metals, for the}' form low-melting alloys with platinum. There are two sets of platinum compounds according to the gen- eral formula PtX2 and PtX4. The best-known platinum compound is chlorplatinic acid, H2PtCl6, obtained by dissolving platinum in aqua regia. When the solution is e\aporated the chlorplatinic acid is left in the form of large, reddish-brown, very hygroscopic prisms. Its aqueous solution contains the anion PtC'lo", for such an anion goes to the anode in an electrolysis; silver nitrate precipitates from the solution not silver chloride, which it would certainly do if free chlorine ions \verc present, but the compound Ag2PtCl6- Two characteristic salts of this acid are those of potas- sium and ammonium; they are very diihcultly soluble in water and insoluble in alcohol; when the aqueous solution is evajxirated 486 INORGANIC CHEMISTRY. [§§ 316- the salt remains in the form of small, but well-formed octa- hedrons of a golden hue. The potassium salt is often made use of in determining potassium when sodium is also present, the sodium platinic chloride being very soluble, even in alcohol. Of the remaining platinum compounds a few may be referred to. If a solution of the above acid, H2PtCl6, is treated with sodium hydroxide and then with acetic acid, platinum hydroxide, Pt(0H)4, is precipitated. It is soluble in strong acids and also in alkalies, so that basic as well as acidic properties must be ascribed to it (platinic acid). Salts of this acid are moreover formed when platinum is fused with alkalies. Platinous chloride, PtCl2, is produced by heating chlorplatinic acid to 200° and in small amount, also, when the solution of this acid is strongly concentrated. It is a green powder, insoluble in water. With the alkali chlorides it gives soluble double salts, such as PtCl2-2NaCl. Double cyanides of platinum with many metals are also known, e.g. K2pt(CN)4- ■3H2O, BaPt(CN)4 • 4H2O, etc. The latter has come into prominence because of its ability to make Rontgen rays visible. All these ■double salts are noted for their beautiful colors and strong dichroism. METAL-AMMONIA COMPOUNDS. WERNER'S EXTENSIONS OP THE NOTION OF VALENCE. 317. Several metals, notably those of the eighth group of the periodio Bystem, are capable of forming complex compounds -with ammonia and acid radicals. Such compounds have long been known, some having been pre- pared by the old master, Berzelius. The study of these substances occupied various investigators of the nineteenth century, especially Jorgensen. In recent years, however, this field has been explored and greatly extended by the investigations of Werner and his pupils, so that at the present time over 1700 compounds of the general type MXp(Am)q, are aheady known, M being a metal atom, X an acid radical and Am ammonia or an organic base (or even water). Chief credit is also due Werner for having taken up the theoretical study of the relationships between these compounds and, as a result, generalizations of considerable importance for the structure of inorganic compounds, especially the complex salts, have been established. The whole subject deserves a little attention at this point. Concerning the methods of preparing these metal-ammonia compounds very little of a general nature can yet be stated. It is readily appreciated that the preparation of such a large number of complex compounds calls for the most diversified synthetical methods. 317.] METAL-AMMilXlA (X'MPOl'XDS. 4S7 In order to acquire an insight into the nature of the compounds concerned we may first examine the trinitrito triammine cobalt, Co(NH3)j(N02)3, which can be obtained by mixing cold solutions of cobalt chloride, ammo- nium chloride and sodium nitrite and treating the mixture with a current of air. The compound then separates out as a difficultly soluble crystalline powder. It can be reorystallized from hot water containing a little acetic acid, without liberating nitrous acid, and is not attacked by dilute mineral acids in the cold. The electrical conductance of its aqueous solution is approximately zero. Evidently, therefore, the substance lacks the ordinary properties of a nitrite; the NOj-groups must be joined to the molecule differ- ently than in the nitrites. The ammonia, too, is otherwise combined than in the ammonium salts. This follows at once from the fact that the com- pound is a non-electrolyte: furthermore, from the fact that the action of even concentrated acids is insufficient to split off ammonia. The NHj-groups are not held by the acid radicals present in the molecule, for by energetic reactions it is possible to substitute other acid radicals for these without liberating the ammonia molecules. The supposition of an active participa- tion of the acid radicals in the linkage of the NHj-molecules is thus excluded. In deciding how the ammonias are connected to the molecule it is significant that the ammonia molecules can be replaced successively by other molecules. This shows that each ammonia molecule must be linked independently of the others in the complex molecule. About the only satisfactory explana- tion is that the NHj-molecules are attached directly to the metal atom. Werner, however, makes the same assumption for the acid radicals in order to explain their abnormal behavior. This theory, whereby the special properties of the groups in compounds of the type M.\:p(NH3)q are explained on the assumption that these groups are in direct combination with the metal, has proved to be of great importance for the classification of these compounds. The acceptance of this principle, however, necessitates an extension of the present notion of valence. Cobalt, for example, is at most trivalent in its salts and oxides; but, if we assume a direct linkage to the metal of the three NH3- and the three NOz-groups, the cobalt must ha\'e a valence of six. These new-appearing affinities of the metal atom cannot offhand be classed with the ordinary valence bonds. For this reason it Beems appropriate to assign a special name to them. In order to distinguish them from the ordinary valences, which are termed "principal" or "primary" valences, they are called "subordinate," or "secondary," valences. Compounds of the type of the trinitrito triammine cobalt, that is, ot the general formula MX3(NH3)3 have the property of combining with more ammonia molecules still. Thus there are compounds of the types MXjCNHj),, MX3(NH3)5, and M.V,(NH3)6. The addition of ammonia gives rise to a peculiar change in the function of the acid radicals, because for e\'cry addi- tional NHj-molecule that is taken on, one of the acid radicals enters the ionizable condition. If we compare the molecular conductivities in y|^ or -j-J-g- normal solution at 250°, to wit: 488 LXORGANIC CHEMISTRY. [§§ 317- Co(NH3),X, Co(NH3)jX, Co(NHp,X, Co(^R^\X^ a 402 245 117 7 with those of the salts: NajPO, MgCl, Naa a 370 249 125 we find that the first of the above complex salts must be a quaternary, the second a ternary, the third a binary, electrolyte, and the fourth a non-elec- trolyte. The chemical properties of these compounds accord perfectly with this theory. In the compound, Co(NH3)jCl3, formerly called "praseo-cobalt chloride," only one chlorine can be directly precipitated by silver nitrate; in Co(NH3)5Cl3 two can be so precipitated, while in Co(NH3)8Cl3, hexammine cobaltic chloride, all three chlorine atoms appear to be ionized in aqueous solution, since all three react at once with a silver salt solution. This is explained on the assumption that the added ammonia molecules displace the acid radicals from their immediate connection with the metal atom and themselves enter into this direct union with the metal. In order to distin- guish the ammonia molecules which are linked up in this way Werner applies to them the term ammine and devises the following formula* and names to express the constitution of the above-named compounds: M(NH3)3X3 Tri-acido triammine compounds [M(NH3)3X3], M(NH3)4X3 Di-acido tetrammine salts [M(NH3)iX2]X, M(NH3)iX3 Acido pentammine saltt [M(NH3)53qX2, M(NH3)eX3 Hexammine salts [M(NH3)JX3, the atoms or groups within brackets being regarded as in direct union with the metal atom. The ionizable groups X outside of the brackets are in indirect union with the complex and, according to Werxer, are not joined to a definite elemental atom. If we compare the composition of the various metal-ammonias, we find that the formation of complexes is not without its limitations, but that after the addition of a definite number of NHj-molecules it comes to an end. Particularly sharp is the limit in respect to the number of groups which can unite directly with an atom serving as the center of a complex radical. It is a striking fact that this limit is the same for a good many elements and is commonly six (6). It seems most likely that this limiting number is characteristic of the elemental atom and of considerable importance. Wer- ner calls it the coordination number and defines it as the maximum number of individual groups that can be directly united to an elemental atom. As aheady stated, the coordination number for most elements is 6; for a few elements, boron, carbon, and nitrogen, it is 4. 318.] METAL-A MMOMA COMPOUNDS. 4,S'J 3i8. The constitution deduced for the metal-ammonia compounds en- ables us to explain in a very simple way a series of isomerism phenomena that are characteristic of these compounds. Two isomeric compounds are known of the formula Pt(NH3)<-SO,(OH)2. One behaves as a strong base, absorbing carbon dioxide from the air like caustic alkalies and precipitating metallic oxides from their salts, but gives no reaction of the sulphate ion in aqueous solution, e.g., no precipitate of barium sulphate witli barium salts. The second compound is a perfectly neutral salt and acts as a normal sulphate, giving a barium sulphate precipitate at once with barium salts. We conclude from these reactions that the first compound gives only OH- ions; the second, on the contrary, only SO,-ions. The cause of the isomerism is explained in the following coordination formulse; b S<(°^Pt(NH3), I (OH)j >,]. and Pt(NH3). Sulphate tetrammine plato hydroxide rHo\ [hox Dihydroxylo tetrammine plat^ sulphate J so,. The researches on metal-ammonia compounds have Ijeen of great assist- ance in clearing up the structure of many complex salts, since the latter can be prepared by the gradual replacement of NHj-groups by other mole- cules. An illuminating example of this is potassium cobaltic nitrite (§ 310), whose relation to hexammine cobaltic chloride is shown by the following series of ammonia compounds: 1. [Co(NH3)JCl3 Hexammine cobaltic chloride I (NH3)3 Trinitrito triammine cobalt 2. Co XO2 (NH3)5 CL Nitrite pentammine cobaltic chloride Co (NOj), (NH3), K Potassium tetranitrito diammine cobaltite 7. [Co(N02)e]K3. PotEissum cobaltic nitrite 3. [co'^J^^^^^lci Dinitrito tetrammine cobaltic chloride (unknown) In all these compounds there is no ionizable NOj-radical. The existence of such transition series clearly brings out the constitutional relationships which must exist between the metal-ammonias and the complex salts. The formulae for these are so constructed that the groups which do not respond to analytical tests, as well as the NHj-molecules displaced by them, are in direct union with the central metal atom. The progressive substitution of NH, by a negative group or element results in the gradual decrease of the valence of the cation. While this il 490 INORGANIC CHEMISTRY [§ 318. three in Xo. 1, it has become zero in Xo. 4, i.e., the molecule has become a non-conductor. If the substitution is carried further, as in Nos. 5, 6, and 7, the complex in brackets assumes the anion rdle and its negative charge Fig. 77. increases from one to three. If the molecular conductivity for a given concentration is plotted on the ordinate axis in a diagram, we obtain for the series of compounds a curve, as is shown in Fig. 77. NumeTous analo- gous transitions of metal-ammonia compounds to complex salts have been discovered, e.g.: [Pt(NH3)JCl, [PtClJK3 [Pt(NH3)JCl3 [PtClJKjT etc. Furthermore, relationships have been found to exist between metal- ammonias and hydrates. It is a significant fact that for a large number of salts of metals which give hexammine salts with ammonia, hydrates with six molecules of water predominate. This is the case with the salts of cobalt, chromium, nickel, etc. Moreover, Jorgensen observed that com- pounds are derived from the hexammine salts by the exchange of one or two XHj-molecules for water, which compounds correspond in behavior to the hexammine salts. This follows from the fact that the molecular conductivity in aqueous solution is but little affected by this exchange, e.g.: a 402 390 380 Br,. It is also confirmed by the fact that with the removal of each water mole- cule an acid radical loses its ionizing property, just as was the case for the removal of each ammonia molecule from the hexammine salts; J 318] METAL-AMMONIA COMPOUNDS. 4i»l Aquo pentammino Chloro pentammine chromic chloride chromic chloride The addition of water has the opposite effect. The trichloro triammine cobalt, for example, is known to form hydrates with 1, 2, and 3 mols. water. Since all of the three chlorine atoms in the trihydrate have an ionizing charac- ter, while in the diyhdrate only two have this property, in the monohydrate one and in the anhydrous compound none, we are justified in ascribing to them in the following constitution: [CljCoCNHj),], rCo(Il'o) CI, rCo(H,0)Jci„ and ICo^j^^^'lci, L (NH,')J L (NH,-)J L ^J^^^sJsJ CI (NH^),. A nearly complete transition series from a metal-ammonia to a hydrate (hydrous salt) is to be found in the case of chromium: [CJc,, [c,™^»>;],,. [cSg]c,., Lacking Blue hexahj'drate All these compounds contain three chlorine atoms in ionizable linkage; accordingly silver nitrate at once precipitates all the chlorine as silver chloride, even from freshly prepared solutions of the salts. The function of water in these compounds corresponds perfectly to that of the water in the aquo-salts of the cobalt ammonias, i.e., with the release of each water mole- cule a chlorine atom sacrifices its ionizing ability. Two hexahydrates of chromic chloride, CrClj- 611,0, are known, one blue and the other green. The blue one contains three ionizable Cl-atoms, and according to its general behavior should be regarded as [Cr(OH2)e]Cl3, hexaquo chromic chloride. Upon the loss of two water molecules it goes over into dichloro tetraquo chromic chloride [Cl2Cr(OH2)JCl, which contains but one ionizable Cl-atom. If this dichloro tetraquo chloride is crystallized out of water, it adds on two molecules of the water and goes over into the green hydrate, [Cl2Cr(OH2),]CH-2H20, which is thus isomeric with the blue hydrate. The green hydrate, moreover, contains only one chlorine atom that can be pre- cipitated directly with silver nitrate. It should also be mentioned that the green hydrate can be transformed into the blue one and the water mole- cules, that condition the isomerism, shifted. Since the isomerism must 492 INORGANIC CHEMISTRY. [§318 depend on the difference in the manner of attachment of the water, Werner has called it hydrate isomerism. Various cases of isomerism that are met with in the metal-ammonia compounds are best explained as cases of stereo isomerism. Two series of compounds are known of the general formula, Pt^ ^ ' ' ; they are called platin diammine compounds and have been known since 1870. Later in- vestigations have shown that this kina of isomerism occurs only with com- pounds of the type 1It>'- The assumption can therefore be made that the six groups are arranged at the apexes of an octahedron around the central atom. Only with the type above described are two isomers possible, as the accompanying figures (Figs. 78 and 79) show. Compounds of the M ' type cannot, therefore, have isomers and, as a matter of fact, no one has yet been able to prepare such isomers. Werneb has also discovered metal-ammonias of more than one nucleus, i.e., where the molecules contain several metal atoms joined to one another in such a way as to undergo no separation either by ionization or by spon- taneous hydrolysis; but we cannot dwell longer upon the subject here. Missing Page Missing Page INDEX. The principal references are in italics Abraum salts of Stassfurt, 65, 322, 374, 441 Absolute weight of atoms, 49 Absorption, 70 Absorption spectra, 391 Accumulator, 421 Acetylene, 249 Acidimetry, 350 Acids, 39, 99 strength of, 99 Acid chlorides, UIS Acids, dibasic, 131 Acids, monobasic, S3 Acid salts, 131 Acid sodium carbonate, 322 Actinium, 399 Additive properties, 348 Affinity, 156 Agglutinants, 269 Air, 1G5 Alabaster, 376 Alchemists, 371 Aldebaranium, 446 Algaroth powder, 234 Alkalimetry, 350 AUotropisin, .5.3, 104, 148, 201, 222, 241, 261, 274 Alloys, 232, 237, 270, 355, 361, 409, 437, 471, 479, 482, 4.S3, 484 AludeLs, 410 Alumina, 437 Aluminatcs, 43.S Aluminium, 436 amalgam, 437 bronze, 437 chloride, 43.S hydroxide, 43.S oxide, 437 silicate, 440 sulphate, 439 Alums, 430 Alum stone, 430 Alunite, 439 Amalgams, 412 Amalgam, aluminium, 437 Amalgam, ammonium, 175 sodium, 313 tin, 276 Amblygonite, 310 Amethyst, 265 Amides, 19,S Amidophosphoric acid, 221 Amido-sulphonic acid, 198 Ammine, 488 Ammonia, i 74, 311 compounds of the metal.s of tne eighth group, 4S(i -soda process, 320 Ammonium, 176, 331 amalgam, 175 carbonate, 334 chloride, 333 hydroxide, 176 magnesium phosphate, 216 metavanadate, 448 nitrate, 33.',, 379 phosphate, 334 salts in anah'sis, 375 sulphate, 334 sulphide, 335 Amorphous substance, 3S5 Amphoteric compounds, 27S, 284 Amphoteric reaction, 319 Analysis, hydrogen sulphide in, 117 Analysis, volumetric. See ^'olumot- ric analysis. Anatase, 446 An.\xagoras, 28 Anglesite, 286 Anhydride, arid, S3 mixed, 1S6, 195 Anhydrite, 381 Anions, 96 Anode, 96 Antichlor, 132, 317 Antimonic acid, 236 Antimony, 231 butter, 233 pentachloride, 2.34 pentasulphidc, 236 40.3 494 INDEX. Antimony, pentoxide, 236 tetroxide, 236 trichloride, 233 trioxide, 234 trisulphide, 236 Antimonyl, 235 Antiseptic, 416 Anglesite, 286 Angstrom, 392 Apatite, 199, 376 Aqua regia, 198 Aragonite, 376, 383 Argentite, 359 Argon, 170, 300, 309 Argyrodite, 286 Aristotle, 371 Abkhenius, 95, 101 Arsenic, 221 acid, 229 meal, 226 oxide, 228 sulphides, 229, 268 sulpho-salts, 230 trichloride, 226 white, 224, ii26 Arseni-tungstates, 457 Arsenious acid, 228 oxide, 226 Arsenolite, 221 Arsenopyrite, 221 Arsine, 223 Asbestos, 374 Association, 82, 185 Atmosphere, 165 Atomic heat, 289 theory, 27 volume, 297 weights, 29 Determination of, 44, 47, 49, 288, 305 Table, of, Inside hack cover Atoms, :i8, 49 Augite, 374 Auric compounds, 370 Aurous compounds, 369 Aurum musivum, 281 AvooADEo's hypothesis, 44, 60, 293 Azides, 179 Azurite, 353 Bacteria, 194 Baeyer, 144 Baker, 185, 333, 414 Balard, 65, 84 Balmer, 392, 393 Barite, 387 Barium, 387 Barium peroxide, 53, SS7 Baryta-water, 387 Bases, 39, 102 Basilius Valentinus, 231 Baijme degrees, 142 Bauxite, 436, 438 Becker, 129 Becquerel, 397 BeU metal, 276 Bergmann, 99 Bernthsen, 133 Beethelot, 50, 152, 157, 182, 255 Behthelot's principle, 157 Berthelot-Mahlbr calorimetric bomb, 153 Beryl, 372 Beryllium, 303, 372 Berzelius, 23, 29, 30, 148, 2M), 293, 308, 418, 486 Bessemer process, 469 Bimolecular reactions, 76, 206, 255 Birkeland, 191 Bismuth, 236 chloride, 237 glance, 237 hydroxides, 238 oxides, 238 oxy chloride, 237 subnitrate, 239 sulphate, 239 sulphides, 239 Black ash, 319 Blanc fixe, 388 Blast furnace, 463 Bleaching, 36, 56, 126 Bleaching-powder, 86, 380 Blende, 407 Blooms, 471 BODENSTEIN, 15 BoiUng-point, elevation of , 61, 272 boltwood, 404 Bone, 249 Boneblack, 244 Boracite, 432 Borax, 432, 435 Boric acid, 433 Boron, 432 boussingadlt, 166 Boyle, 371, 395 Boyle's law, 47, 58, 160, 295 Brand, 199 Brass, 355 Bricks, 440 Brimstone, 103 Bromine, 65 oxygen compounds of, 90 Bronze, phosphor, 202, 276 siUcon, 276 IXDKX. 4'J5 Bronzes, 276 Brookite, 446 Brownian movement, 49 Browning, J., spectroscope, 389 BuNSBN, 329, 389, 391, 398 burner, 258 ceU, 422 Burette, 147 Cadmium, 410 Cxsium, 329 Cain, 249 Calamine, 407 Calcining, 353, 463 Calcite, 376, Jtii Calcium, 376 cai-bide, 171, 249 carbonate, 3S2 elJoride, 321, 340, 37!J cyanimide, 175 fluoride, 380 hydroxide, 37S hypochlorite, 380 manganitp, 460 nitrate, 382 oxide, 377 peroxide, 379 phosphates, 382 plumbate, 285 silicate, 384 sulphate, 338, 381 f1 perruthenate, 483 phosphates, 327 ru then ate, 482 sihcate, 328 silver cyanide, 363 sulphates, 326 sulphides, 328 water-glass, 328 Praseo-cobalt chloride, 488 Praseodymium, 445 Precious (metals), 361, 369 Preparing-salt, 281 Priestley, 165. Primary salt, 214 Primordial substance, 28, 308, 395 Proust, 27 Prout, 308 Prussian blue, 367, 477 Prussiates, 475 Prussic acid, 256 PuddUng, 471 Purple of Cassius, 281 Pyrarg3Tite, 359 Pyrite, 103, 135 Pyrites, 463 Pyrolusite, 458 Pyrophorus, 282 Pyrophosphoric acid, 217 Quartz, 265 Quick-lime, 377 Quicksilver, 410 Radio-activity, 397 Radio-thorium, 448 Radium, 50, 398 Rain-water, 20 Ramsay, 170, 172, 402 Rare earths, 443 Raschig, 177, 196, 198 Rayleigh, 170 Rays, 399 and foil. Roentgen, 400 Reaction constant, 75 Reaction velocity, 15, 75, 182, 233. See also Catalysis. Realgar, 103, 221, 229 Reduction, 16, 277 Reefs, 367 Refining, electrolytic, 354, 361, 479 Refractory material, 440 Rbgnault, 291 Regulus, 354 Reichee, 114 Retgers, 222 Reversible cells, 421 Reversible reactions, 73, 422. See also EquiUbrium. Rhodium, 483 Rhodochrosite, 458 Richards, 294 Roberts-Austen, 466 Rock crystal, 265 Rook salt, 311, 315 RoozEBooM, 242, 466 RoNTGEN rays, 486 Rose, 435 Rose's metal, 237 Rowland, 392, 396 Rubidium, 329 Ruby, 436, 438 Ruff, 312 RuNGE, 392, 393 Rusting, 428, 472 Ruthenium, 482 Rutile, 446 Rydberg, 392 Sal ammoniac, 333 Saleratus, 322 Sal mirabile Glauberi, 317 Sal soda, 319 Salt cake, 319 Salt, common, 315 Salt covers, 315 Salterns, 315 i.\D/:x. 503 Saltpetre, 32() Cliili, 31S Salts, 39 acid, 131 complex, 4-10 double, sec Double salts. primary, secondarv. and tcrliarv, ■2U Salt solutions, oo.'i, 379 Samarium, ,446 Sand, 2(55 Sandstone, 2(5,") Sapphire, 430. 43S Sassolite,, 432 SvUNDEnS, 14.S .■Scandium, 44.") Si'UEELE. SI, IGli. 199 Scheele's sreen, 3r)9 Scheelite, 4.")(i Schlippe's salt. 23(1 ScH(iNBEl.\, 49 ScHiixHERR, 1!I2 SCHUTZEN'BEGKR. 133 Sehweinfurth green, 224, 359 Scientific investigation, 1 Sea-water, 21 Secondary salt, 214 ."-^elenium, 14S ."^cmi-permeable meiubranes, .57 Senarmontite, 234 Sf.ndcren,';. 20') Sen-it ive film, 364 salts, 364 SiTiiciitine 374 .Sesquioxide, 92 .~rijcess, 367 process (steel). -169 Silica, 261, .'I!.-, Silicic acids, 266 Silicides, 262 Silico-chloroform, '2it-i -ethane, 2(J3 Silicon, 261 bronze, 276 chlorides, 263 in iron, 468 nitride, 267 sulphide, 2()7 tetrafluoride, 263 Silver, 359 bromide, 363 chloride, 363 fluoride, 363 iodide, 363 nitrate, 366 Silver, nitrite, 367 oxide, 3()2 i peroxide, 362 Iilating, 362, ,363 suboxide. 362 suljihate, 366 SlaK. 3.-j3 Small. 47S Smalt ite, 477 Smaragd, 372 Smelting, 463 j Smithsonite, 407 Soapstone, 374 Soda, 311, 319 caustic, 313 ijlass, 3S4 SoDDV, 402, 405 Sodium, 311 amide, 176 ammonium phosi)hate, 334 bicarbonate, 320, 32 J borate. ,Si'e borax. bromide, 317 carbonate, 319 chloride, 315 dichromate, 454 hydroxide, 313 iodide, 317 nitrate, 31S nitrite, 319 nitroprusside, 477 oxides, 312 phosphates, 319 silicate, 266, 3J.2 sulphantimoniate, 236 sulphate, 110, 317, 340, 345, 425 sulphides, 322 sulphostannate, 2S1 thiosulphate, 132, 317, 365 Soffioni, 433 Soil, absor])tive power, 2(57 Solder, 276 Solubility. 333 and foil. curves, 336 and foil. product, 119, 2S5, 315, 378, 472, 475, 480 Solute, 338 Solution, 5, 335 saturated, 5 solid, 202, 3S6\ 468 supersaturated, 345 tension, 418 electrolytic, 2S2, 418 SijLVAV soda process, 320 Soot, 245 Specillc heat, 290 Speclri)scf)|)e, 3S!) Spectro,sco|)y, .329, ,;,S',7, 39S, 4J1, -144 504 IXDEX. Spectrum, 389 continuous, 25/, 389 line, 25S spark, 390, 39S Spelter, 409 Sphalerite, 103, 407 Spiegeleisen, 466 Spiritus fumans Libavii, 279 Spitting of silver, 11 Spring, 20 Stable, 108 Stahl, 165 Stannic acid, 2S0 chloride, 279 oxide, 280 sulphide, 281 Stannous chloride, 276 hydroxide, 278 oxide, 278 sulphide, 278 Stas, 23, 24, 88, 248, 293, 308, 361 Stas.sano, 465 Status nascens, 37, 54 Steel, 456, 466 Stephanite, 359 Stereo-isomerism, 492 Sterling silver, 362 Stibine, 231 Stibnite, 103, 231 Stock, 233 Stoichiometry, 31 Storage battery, 421 Strength of an acid or a base, 99 Stromeyerite, 359 Strontianite, 386 Strontium, 386 Strutt, 404 Sublimation, 7 Sulpho-anhydrides, 231 Sulphomonoper acid, 144 Sulphur, 102 halogen compounds, 121 in iron, 468 oxides, 124 oxygen acids, 131 springs, 115 Sulphuretted hydrogen, 115 Sulphuric acid, 134 chlorides of, 142 fuming, 130, US anhydride, 128 Sulphurous acid, 133 anhydride, 124 oxide, 124 Sulphuryl chloride, 143 Summary of the alkali group, 330 carbon group, 287 chromium group, 458 Summarj- of the copper group, 372 gallium group, 442 group of the alkahne earths, 388 halogen group, 93 nitrogen group, 239 oxygen group, 151 platinum group, 481 zinc group, 4l7 Supercooled liquids, 107 Superphosphate, 136, 382 Suspension, 271 SVEDBERG, 272 Svlvanite, 150, 367 Sylvite, 324, 441 Symbols, 30 Talc, 374 Tammann, 114 Tank waste, 319 Tantalum, 448 Tartar emetic, 235 Tellurium, 150 Tempering, 471 Terbium, 443 Tertiary salt, 214 Testing of gold and silver, 369 Tetraboric acid, 434 Tetradymite, 237 Thallium, 304, Ul Thenaed, 57, 311 Thenard's blue, 478 Theory, 2 of indicators, 350 Thermochemistry, 152. See also Endothermic compounds. Thermocouple, 483 Thermodynamics, 48, 158, 160 Thermoneutrahty, law of, 350 Thionic acids, 145 Thiosulphuric acid, 133 Thomas slag, 469 Thomsen, 152, 156 Thorite, 447 Thorium, 447 Tin, 274 amalgam, 276 butter, 279 disease, 275 fluoride, 279 ' foil, 274 -moiree, 274 phosphide, 281 plate, 276 Tinkal, 435 Titanium, 446 Titre, 146 TOLMAN, 96 iA'vi':x. 505 Touchneedles, 369 Touchstone, 369 Tourmaline, 374 TowNSEND process, 314 Transition point, 106, 160, 227, 276, 317, 334, 340, 415, 425 Traube, 54 Travehs. 172 Triads, 296 Triboluminescence, 227 Triilymitc, 265 Trinitrito tiiamraine cobalt, 487 Triple point, 113 Triplnlite, 310 Tuffstune, 37S Tungsten, 456 ^ steel, 456, 471 Twyers, 463 Tyxdall, effect, 271 Type metal, 232 Ultramarine, 440 Ultramicroscope, 272 Ultraviolet light, 18, 37, 129, 266 Unimolecular reactions, 76, 251, 402 Unit>- of matter, 395 Univarient system, 112 Uraninite, 397, 457 Uranium, 177, 397 and foil., 457 Urbai.v, 444 Vacuum flask, 11, 170 Valence, 122, 185, 486 Valence, maximum, 123 of ions, 123 \'anadinite, 448 \'anadium, 448 ^'AN DER Stadt, 206, 219 \'an DER Walls, 49, 109 \'an- Marum, 49, 203 Van T Hope, 58, 60, 157, 381 Van't Hoff's principle of mobile equihbrium, 160, 336, 337 Vapor pressure curve, 61, 106, 112 Varec, 69 Vein mining, 367 Velocity constant, 75 of reaction, 15, 75, 182, 233. See also Catalysis. Vermilion, 417 Vitriol, blue, 358 green, 473 oil of, 137 Vivianite, 199 Vogel's spectroscope, 389 Volumetric analysis, 146, 189, 350, 460, 402 Vulcanizing rubber, 121 Wackenroder's Uquid, 145 Washing (precipitates), 5 Washing-soda, 321 Water, 17 composition of, 22 gas, 250 glass, 266, 322, 328 natural, 20, 383, 473 physical properties of, 20 purification of, 21, 283 Wavellite, 199 Welding, 437 Weldon process, 460 Welsbach, Ader von, 445, 440 incandescent gas light, 250, 257, 388, 443, 447 Werner, 486 Whitney, 347 Winkler, 273, 307 Witherite, 387 Wolframite, 456 Wood's metal, 237 WouLFF bottle, 18 Wrought iron, 466 Wulfcnitc, 281, 455 Xenon, 172, 300 Ytterbium, 445 Yttrium, 443 Zinc, 407 compounds, 409 dust, 41)7 white, 409 Zircon, 440 Zirconia, 447 Zirconium, 446 ZSIOMONDY, 272 Short-title Catalogue OF THE PUBLICATIONS OP JOHN WILEY & SONS New York London: CHAP:^IAN & HALL, Limited ARRANGED UNDER SUBJECTS Descriptive circulars sent on application. 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