I - I~~~~~~~~~~~~~~~,~ti.:'- ~::'- j ~"i; _~~~~;~ G~E ~ A.L N OTIONS OF BY J. PELOUZE & E. FREMY. TRANSLATED FROM THE FRENCH. J. B. LIPPINCOTT & CO. GENERAL NOTIONS OF CHE MI S TR Y, J. PELOUZE AND E. FREMY. TRANSLATED FROM THE FRENCH. BY EDMUND C. EVANS, M.D. P HIILADELPHIA: LIPPINCOTT, GRAMBO & CO. 1854. Entered according to Act of Congress, in the year 1854, by LIPPINCOTT, GRAMBO & CO., in the Clerk's Office of the District Court of the United States for the Eastern District of Pennsylvania. (2) PREFACE. THIS work is intended for persons who, unaccustomed to scientific studies, wish to acquire a general knowledge of Chemistry and its principal applications. Among the numerous facts which compose this science, we have chosen those which recommend themselves by their importance in the arts; these we have attempted to make clear by freeing them from formulas and details purely scientific, which we have given in other works. To render our explanations more intelligible, we have (when desirable) accompanied them with plates which faithfully represent the form and arrangement of apparatus used in laboratories and manufactories. Our object will have been gained, if the "GENERAL NOTIONS OF CHEMISTRY" contribute to develope a taste for a science which renders such great services to the industrial arts, that ignorance of its first elements is, at this day, inadmissible. (iii) CONTENTS. Page NOTE BY THE TRANSLATOR.............................. xvii INTRODUCTION......................................... 13 NOMENCLATURE.~ Nomenclature of Simple Bodies............................. 14 List of Metalloids...................... 16 " " Metals......................... 16 Nomenclature of Compound Bodies.......................... 16 Acids................................................................ 17 Hydracids........................................ 18 Oxides................................................... 18 Salts.................................................... 19 Binary Compounds of which Oxygen is not one of the Elements. 20 Alloys................................................... 21 Crystallization of Bodies............................. 21 " " by Fusion.................................. 22 " " " Volatilization....................... 22 " " " Solution................................. 23 METALLOIDSS. Oxygen.................................................. 25 Preparation of Oxygen............................... 28 Hydrogen..................................... 30 Action of Hydrogen on Oxygen...................... 31 Preparation of Hydrogen....................... 34 Uses of Hydrogen................................... 35 COMBINATION OF HYDROGEN WITH OXYGEN. Water................................................... 35 Composition of Water......................... 36 Solid Water....................................... 37 Liquid Water..................................... 39 Water in a state of Vapour........................ 39 (v) Vi CONTENTS. Page Chemical Properties of Water................ 40 The state of Water in Nature................. 41 Water called Selenitic......................... 42 Stalactites, Incrustations, and Deposites in Boilers....... 43 Air dissolved in Water.............................. 43 Distillation of Water............................... 44 Distilling Apparatus................................ 45 Nitrogen................................................. 46 Preparation of Nitrogen............................ 47 Atmospheric Air.................................. 47 Composition of the Air............................... 50 Properties of the Air -Phenomena of Combustion in the Air............................................. 53 Nitric or Azotic Acid................................... 56 Preparation......................................... 58 Uses............................................. 59 Ammonia ( Volatile alkali).................................. 59 Properties............................................. 61 Preparation........................................ 63 Uses.................................. 64 Chlorine................................................. 66 Preparation........................................ 69 Water of Chlorine................................. 69 Uses................................................ 69 Chlorhydric Acid...................................... 70 Preparation...................................... 70 Uses............................................... 72 Bromine................................................. 72 Iodine.............................. 73 Properties.................................... 73 Preparation...................................... 74 Uses.............................................. 74 Sulphur............................ 75 Natural State...................................... 75 Properties......................................... 75 Extraction of Sulphur............................. 76 Uses.............................................. 77 COMBINATION OF SULPHUR WITH OXYGEN. Sulphurous Acid......................................... 78 Uses.............................................. 78 CONTENTS. vii Page Sulphuric Acid........................................... 79 Uses........................................... 81 Acid Sulphydric........................................ 82 The Natural State of Sulphydric Acid................ 83 Uses.............................................. 83 Phosphorus............................................... 84 Properties......................................... 84 Preparation of Phosphorus............................ 88 Uses.............................................. 90 Arsenic.................................................. 90 Arsenious Acid.......................................... 91 Uses............................................... 92 Arseniuretted Hydrogen................................. 92 Preparation...................................... 93 The Detection of Arsenious Acid in cases of Poisoning... 93 Carbon................................................... 95 Properties........................................ 95 The Diamond........................................... 96 Graphite or Plumbagine -Black Lead....................... 99 Anthracite....................................... 100 Lamp-Black............................................. 101 Metallic Carbon........................................... 101 Coke.................................... 102 Charcoal................................................. 103 Carbonisation in Mounds......................... 104 General Properties of Charcoal................. 105 Animal Charcoal...................................... 108 Preparation of Animal-Black..................... 109 COMBINATIONS OF CARBON WITH OXYGEN. Carbonic Oxide............................................ 109 Preparation...................................... 110 Carbonic Acid..................................... 111 Gaseous Carbonic Acid......................... 111 Preparation...................................... 114 Liquid Carbonic Acid.................................... 114 Properties......................................... 114 Solid Carbonic Acid.......................... 115 COMBINATIONS OF CARBON WITH HYDROGEN. General Remarks on Illuminating by Gas..................... 117 Cyanogen, Prussic Acid, and Cyanurets..................... 122 Viii CONTENTS. Page COMBINATIONS OF BORON WITH OXYGEN. Boracic Acid................ -................................ 123 Uses................................................ 124 COMBINATIONS OF SILICIUM WITH OXYGEN. Silicic Acid, or Silex........................................ 124 Anhydrous Silicic Acid.................................... 124 METALS. Generalities on the Metals................................... 126 The action of Oxygen, of Atmospheric Air, and of Water, on the Metals................................................... 129 Classification of Metals...................................... 129 Metallic Oxides......................................... 131 Action of Heat on the Oxides........................ 132 " " the'Pile "..132... Oxygen ".............. 133 " " Hydrogen "........................ 133 ". " Carbon ".133 Salts.. 134 Phenomena of Saturation.............................. 135 General Properties of Salts................. 136 Action of Heat on the Salts. 136 " " Electricity on Salts.......................... 137 " " Metals on Saline Solutions................... 137 Hygrometric Action of the Air on Salts................. 138 Action of Water on Salts.................... 139 Phenomena which determine the Composition of Salts.... 141 Potassium................. 141 Properties....................................... 142 Preparation............... 143 Potash................................................ 144 Preparation...................................... 145 Uses................................................ 147 Natural State of Potash........................ 147 Nitrate of Potash........................................... 147 Natural State of Nitre................................. 149 Washing the Saltpetre Materials......................150 Boiling............................................... 151 (Boues), or Impurities............................... 151 Refining Saltpetre.............................. 151 CONTENTS. ix Page Gunpowder................................................. 152 Properties of Powder.153 Carbonization in close vessels.......................... 159 Manufacture of Powder................................ 159 The Process by Pestles............................... 160 Process of the Mill s.161 Carbonate of Potash..162 Preparation............................. 162 Uses.........................163 Sodium.................................................... 164 Chloride of Sodium..........................................164 Uses.........................,166 Rock Salt........................................ 166 Extraction of Chloride of Sodium............... 167 Borate of Soda.................................. 168 Uses................................................. 169 Carbonate of Soda................................ 169 Preparation........................................ 169 Utses................................................ 171 Calcium.....................-................................ 171 Lime....................................................... 171 Natural State of Lime................172 Extraction of Lime.................................... 173 Uses.. 174 Chloride of Lime.. 174 Preparation...................................... 175 Uses................................................. 175 Sulphate of Lime, Gypsum, Plaster. 175 Uses............................................... 177 Carbonate of Lime.......................................... 178 Properties........................................... 178 Marbles................................................... 180 Magnesia................................................... 181 Uses................................................. 182 Alumine................................................... 182 Alum...................................................... 184 Preparation.......................................... 184 Manufacture of Alum from Clay........................ 185 Uses................................................. 185 Feldspar................................................... 185 Kaolin..................................................... 186 1 A x CONTENTS..Page Clay (Argil)............................................... 186 Marl....................................................... 187 Ochre...................................................... 187 Fullers' Earth.............................................. 187 Glass...................................................... 187 General Properties of Glass............................ 188 Manufacture of Glass.............. 191 Bohemian Glass...................................... 193 Crown Glass....................................... 193 Window Glass........................................ 194 Looking-Glass.....................194 Glass for Bottles...................................... 195 Crystal.............................................. 196 Colored Glasses....................................... 19S Enamel.............................................. 199 Venetian Glass-Filagree Glass........................ 200 Millefiori (Mosaic Glass).............................. 201 Flint Glass........................................... 201 Colorless Paste (Strass)............................... 201 Avanturine........................................... 202 Painting on Glass..................................... 202 Potteries................................................... 203 Washing............................................ 204 Grinding............................................. 204 Intimate Mixture of Materials.......................... 204 Glazing.............................................. 206 Baking of Potters' Ware............................... 207 Decoration of Potteries...................................... 209 Earthen Wares....................................... 212 Bricks.. --- 212 Tiles- Paving Bricks................................ 214 Crucibles............................................. 214 Water-Coolers........................................ 215 Common Pottery.......... 215 Common Delft, or Italian Ware, with opaque Coating.... 215 Fine or English Ware, with transparent Glazing.. 216 Stone Ware......................................... 217 Hard, or China Porcelain.... 218 Tender French Porcelain.............................. 219 Tender English Porcelain.............................. 220 CONTENTS. xi Page Porcelain of Tournay................................. 220 Building Stones....................................... 220 Lime Stone.......................................... 221 Mortars.................................................... 222 Common Mortar not Hydraulic........................ 222 Hydraulic Lime and Mortars.......................... 223 Theory of the Hardening of Hydraulic Lime............ 224 Roman Cement....................................... 225 Artificial Hydraulic Lime.............................. 225 Hydraulic Mortars................................... 226 Concrete................................................... 227 Mastics....................... 22 Mastics applied cold.................................. 227 Mastics applied by Fusion.......................... 229 Iron....................................................... 229 Pure Iron........................................... 229 Iron of Commerce....................... 230 Principal Compounds of Iron....................... 233 Oxides of Iron........................................ 233 Magnetic Oxide of Iron................................ 234 Sulphuret of Iron...................................... 235 Sulphate of Iron...................................... 236 The Extraction of Iron.......................236 Preparation of the Ores................................ 238 The Catalan Method................................... 238 Manufacture of Cast or Pig Iron.............................. 239 Refining of Cast or Pig Iron........................242 Refining with Pit-Coal, by the English Process.......... 244 Different kinds of Cast-Iron.................................. 245 Black Cast-Iron................................ 246 Grey Cast-Iron.......... 246 White Cast-Iron...................................... 247 Steel............................ 248 Natural Steel.................. 250 Steel of Cementation.................................. 250 Cast-Steel........................ 251 Damask Steel......................................... 251 Zinc...................................................... 252 Properties............................................ 253 Oxide of Zinc........................................ 254 X IL...I CONTENTS. Page Extraction of Zinc. 255 Metallurgic Treatment................................ 256 Lamination of Zinc................................... 256 Uses................................................ 257 Tin........................... 257 Salt of Tin. 258 Tin Plte, or Fer Blanc.... 259 Moir6 Metallique, or Crystallized Zinc Plate.. 260 Uses of Tin.......................................... 261 Lead..................................................... 261 The Protoxide of Lead............................. 262 Properties...................................... 262 Minium............................;,263 Uses................................................ 264 Sulphuret of Lead.......................................... 264 Uses............................................... 265 Carbonate of Lead... 265 Manufacture...................................... 266 Alloys of Lead 267 Copper................................................... 268 Alloys of Copper.................................... 270." " "and Zinc....... 270 Bronze............................................. 271 Mercury.................................................. 272 Amalgams, or Alloys of Mercury.. 274 Amalgams of Tin............................. 274 " "Bismuth................................ 275 Silver.................................................... 275 Nitrate of Silver..................................... 277 Alloys of Silver.................................. 278 SC " and Copper.................. 278 Plate............................................... 279 Gold. 280 Purple Powder of Cassius.......... 282 Alloys of Gold and Copper....... 282 Amalgams of Gold 283 Alloys of Gold and Silver.......... 284 The Extraction of G old 284 Gold from Alluvial Formations... 285 Gold in Veins........ 286 CONTENTS. Xlll Page Platinum................................................. 287 Gilding.................................................. 289 Galvanoplastie, or Electrotype............................... 290 Photography...................................... 291 ORGANIC CHEMISTRY. Generalities on Organic Bodies....................... 292 Acetic Acid............................................... 294 Acetification.............................. 294 Properties of Acetic Acid....................... 298 Acetates............................................ 298 Oxalic Acid............................................... 299 Natural state....................... 299 Preparation...................................... 300 Uses.................................................301 Tartaric Acid.............................................. 301 Citric Acid................................................. 302 Lactic Acid............... 302 THE ASTRINGENT PRINCIPLES OF VEGETABLES. Tannins....................................................305 Tannic Acid, or Tannin................ 305 Ink................ 307 Tanning.................................................. 308 Acids or Fruits................ 308 Organic Alkalies........ 309 The Natural State and Extraction of the Organic Alkalies....... 310 NEUTRAL BODIES. Lignine.................................................. 313 Cellulose.................................................. 313 Incrusting Matter......................................... 15 General Properties of Wood and the Combustibles............. 316 Peat..................................................... 318 Fossil Combustibles........................................ 318 Lignite, Pit-Coal, Anthracite.......................... 318 Properties of Coal................................. 319 Processes for the Preservation of Wood........................ 320 c)* Xiv CONTENTS. Page Pyroxylin, Pyroxyle, Gun-Cotton............................. 323 Properties of Pyroxylin.............................. 324 Amidon.................................................. 327 Physiological Notions of Amidon....................... 328 Properties of Amidon................................. 329 Diastase.................................................. 332 Dextrine.................................. 332 Extraction of Amidon.................................. 334 ""Fecula.................................. 337 Gum..................................................... 338 Sugars.................................................... 338 Sugar of Milk............................................. 339 Glucose, Grape Sugar, Sugar of Amidon...................... 340 Manufacture of Glucose, and Syrup of Fecula................. 341 Sugar of Diabetes.................................. 343 Cane Sugar.................................. 343 Beet-Root Sugar..................................345 Refining of Sugar.................................. 349 Sucre Royal........................................ 352 Stamped Sugar....................................... 352 Sugar-Candy.................................. 353 Barley Sugar, Apple Sugar........................... 353 Alcoholic Fermentation - Panification.......................354 Alcohol................................................... 356 Ether...................................... 359 Manufacture of Wine.........................361 Manufacture of White Wine......................... 362 " " Red Wine............................ 364 Sparkling Wines...................... 365 Vins de Liqueur-Sweet Wines...................... 367 Disease of Wines...................... 367 Table of the quantity of Alcohol contained in certain Wines and Spirituous Drinks....................................... 369 Essential Oils.............................................. 370 Essence of Turpentine............................... 372 Resins.................................................... 373 Varnish................................................... 373 Caoutchouc............................................... 375 Gutta Percha.............................................. 377 Fat Bodies................................................ 377 CONTENTS. XV rage Stearine.................................................. 378 Margarine................................................ 379 Oleine.................................................. 379 Stearic, Margaric, and Oleic Acids........................... 380 Glycerine................................................ 380 General Properties of the Neutral Fat Bodies............. 381 Olive Oil................................................. 382 Linseed Oil......................................... 383 Tallow, or Suet......................... 384 Butter.................................................. 385 Beeswax................................................. 385 Soaps................................................... 386 Manufacture of Stearic Candles...................... 389 Coloring Matters....................................... 390 General Principles of Dyeing............................... 394 Animal Chemistry......................................... 397 Fibrine.................................................. 397 Albumen..................................399...... 399 Caseine................................................. 400 Gelatine.................................................. 401 Blood................................................... 401 Microscopic examination of the Blood................. 403 Coagulation of the Blood............................. 404 Circumstances which influence the Coagulation of the Blood........................................... 406 Milk........................................... 409 Muscular Flesh........................................... 409 Bones.................................................. 409 Bile..................................................... 410 Elementary Notions of Respiration and Nutrition....... 410 Digestion................................................ 410 Respiration.............................................. 412 Chemical Phenomena of Vegetation......................... 414 Agricultural Chemistry................................... 417 Farm Land......................................... 417 Sandy Soil O,...................... 419 Clayey Soil............... 419 Calcareous Soil...................... 420 Vegetable Soil....................................... 421 Mineral Manures.....................................422 Xvi CONTENTS. Page Marling or Manuring with Marl.....423 Plaster..........................424 Use of Salt (Chloride of Sodium,) in Agriculture.425 Drainage. 425 Manure..................426 Ploughing in Crops..........I...........429 Culture of Trefoil................ 430 Employment of Azotised. Salts......431 NOTE BY THE TRANSLATOR. IT was thought best to give the French system of weights and measures, and not to attempt to convert them into the English system in the text. This renders it necessary here to place before the reader a table of the comparative value of the French with English weights and measures. Also, a comparative scale of Fahrenheit's Thermometer with Centigrade's, which last is in universal use in France. The metre is the standard of linear measure. It is the ten millionth part of the distance from the Equator to the North Pole. This measure is marked on a bar of platinum, which is preserved in the archives of the Academy of Sciences. Thus in linear measure - The Metre is equal to 39,37 English inches, or 3 —~- feet. Decimetre, or IA of a metre, 3 i8 inches. Centimetre, or ~l of a metre, T of an inch. Millimetre,,- of a metre, T.~9 of an inch. Decametre is 10 times 1 metre, or 393[-7-k inches, or 10 yards 2 feet 7 inches. Hectometre, 100 times I metre, or 3,937 inches, or 100 yards 1 foot 1 inch. Chiliometre, 1000 times I metre, or 39,371 inches. SUPERFICIAL OR SQUARE MEASURE. Are is a square decametre, 3 1.- English perches. Decare is 10 ares, or 39 5 English perches. Hectare 100 ares, or 395 perches, or 2 acres 1 rood 35 perches. SOLID MEASURE. Stere is a cubic metre, or 35 o cubic feet. Decistere is — th of a stere. Decastre is 10 cubic steres. (xvii) XVi1i NOTE BY THE TRANSLATOR. MEASURES OF CAPACITY. Litre is the cubic decimetre. It is the standard measure of capacity in the decimal system. It contains 61.-~ 02 English cubic inches. The English imperial gallon is equal to 4.T. litres. The Decalitre is IQ litres. Hectolitre is 100 " Decilitre is Lth of a litre. Centilitre is ~-A " Millilitre is T'0 " By keeping in mind that a metre is the basis of all other measures both of length, of superficies and capacity, and that it is rather more than 3 feet and 3 inches, and that it is a decimal system of measures, the reader can readily follow the authors. WEIGHTS. The unit of weight is the gramme. It is the weight of the 100th part of a cubic metre of distilled water at the temperature of melting ice. A gramme is equal to 15r4 3 troy grains. A decigramme is equal to 1 5 43 troy grains. A centigramme is equal to 4,,ub A milligramme is equal to o1~, A decagramme is 10 grammes or 154 —4 grains. A hectogramme is 100 grammes or 1543 4 " T i A kilogramme is 1000 grammes or 15,434 " A myriagramme is 10,000 grammes or 154,340 grains. A kilogramme is equal to 2 lbs. 3 oz. 4'428 drachms avoirdupois weight. THERMOMETERS. THREE scales or methods of division of the thermometer have been adopted. They are those of Fahrenheit, mostly used here and in England; Reaumur, now seldom used; and the Centigrade scale used in this work, and most generally adopted in France and on the Continent of Europe. Properly, the point marked Zero 0~ should in all thermometers mark the freezing point of water. This is the case in the Centigrade; but on Fahrenheit's scale 0~ is 32~ below freezing. In order to have a clear idea of the division of these scales, it must be borne in mind that the starting point in both is the freezing of water, which is a constant temperature (that is under ordinary natural circumstances); the next point to be fixed in the scale is the boiling of water; this, too, at the level of the sea in an open vessel, the water being pure, is a constant temperature. These two facts being remembered, and also that the space passed over by the mercury on the scale between these two points is, on the Centigrade scale, divided into 100 equal parts called degrees, and in Fahrenheit's scale into 180 equal parts or degrees, it will be easy to compare the degrees on Centigrade with those of Fahrenheit, by always comparing those of Centigrade above or below zero with those above or below the freezing point of Fahrenheit, 32~ F., without reference to his zero. 32~ F. and 0~ C. mark the same temperature. Then by the single rule of three, it is very easy to convert any degree marked on Centigrade's scale to the same temperature marked on Fahrenheit. What degree, for example, on Fahrenheit's scale corresponds with 5~ C.? Say as follows: as 100 (the number of degrees between freezing and boiling of water on Centigrade's scale,) is to 5~ C., so is 180~ (the number of degrees between freezing and boiling of water in Fahrenheit's scale,) to the answer required, which is 9~, that is, 9~ above freezing, (or 9~ added to 32~,) or 41~ F.; if we wish to compare 5~ C. below 0~ C., or freezing, then it would be by the same rule, 32~ F. less 9~, or 23~ F. Take again 40~ C. above 0~ or freezing, this will be found equal to 72~ above freezing of Fahrenheit; that is, 72~ added to 32~ F., which is 104~ F.; if it is 40~ below 0~ C. (the freezing point), it will correspond to 72~ F. below freezing (32~ F.), which is 40~ below zero, Fahrenheit. As water when kept perfectly still may have its temperature reduced to 12~ below 00 C. without freezing, and as ice or snow when melting preserves a constant temperature, viz., 0~ C. or 32~ F., melting ice or snow is used for fixing the zero, or freezing point of water, in making thermometers. This point is, therefore, properly speaking, the zero, or starting point. (xix) NOMENCLATURE. IN relation to the nomenclature of the so-called hydracids, the following is from Regnault: - " Certain combinations of the metalloids with each other are energetic acids, which are scarcely below the most powerful oxacids; as, for example, the chloride of hydrogen, fluoride of hydrogen, etc., etc. It has unfortunately been thought proper to establish for these compounds a particular rule of nomenclature. It was thought that, in these new acids, hydrogen acted a part analogous to that of oxygen in te oxacids, and the name hydracids was given to them. This was, however, a grave error; in the oxacids, oxygen is the electro-negative element, while in the hydracids, hydrogen is constantly the electro-positive element." (That is, when subjected to the action of a feeble voltaic pile, the oxacid is decomposed, and the oxygen goes to the POSITIVE POLE; but when subjected to the same influence, the hydrogen of the hydracid goes to the negative pole, and is therefore the electro-positive element.) "The nomenclature of hydracids is, however, so generally employed that we are obliged to adopt it ourselves. The chloride of hydrogen takes the name of chlorhydric acid; the sulphuret of hydrogen takes the name of sulphydric acid. The names hydrochloric acid, hydrosulphuric acid, are often given to these same acids; but these names are more defective than the first, for they are an infraction of that general principle, according to which the name of the compound body ought always to commence with the name of the electro-negative element." REGNAULT'S COURS ELEMENTAIRE DE CHIMIE. (xx) GENERAL NOTIONS OF CHEMISTRY. INTRODUCTION. TiHE special object of Chemistry is the study of phenomena, which, taken together, enable us to characterise bodies. Bodies may be ranged in two classes: the first class comprises simple bodies; the second, compound bodies. A simple body is one from which but a single substance can be extracted; for example, from sulphur, in whatever way it be treated, we get nothing but sulphur. A compound body i's one from which two or more substances may be separated, having different properties. If oxide of mercury be heated, oxygen and mercury are obtained from it; the oxide of mercury is, then, a compound body. It is known that bodies present themselves in three states; they may be either solid, liquid, or gaseous. A large number of bodies, such as sulphur and water, are known in these three states; others, as platinum among the metals, and wax among the organic bodies, are only known in the solid and liquid states. Some, as carbon, lime, and lignine, are only solid; while others, called permanent gases, as oxygen, hydrogen, and nitrogen, are only known at this time in the gaseous state. 2 (13) 14 NOMENCLATURE. Heat, cold, compression, the solvents, are frequently used to modify the state of aggregation of bodies. Many gases are liquefied by compression or cold, or more frequently by both of these united. A gas which has been liquefied by great pressure, may even be solidified. The force which unites the molecules of solid and liquid bodies, is called cohesion. Cohesion is very great in solid bodies, feeble in liquids, and does not exist in gases; the molecules of a gas, on the contrary, have a tendency to separate from each other. Heat tends to destroy the force of cohesion; thus it often causes the fusion, and even volatilization, of solid bodies. The force which unites the molecules of simple bodies to form the molecule of a compound body, is called affinity. It is this which, in the oxide of mercury, unites the oxygen with the.mercury. Affinity plays a great part in chemical phenomena; it determines the combinations of bodies, and a great number of decompositions. All the causes which tend to destroy cohesion, such as heat, and solution in a liquid, tend equally to increase affinity; thus there is a great number of bodies which only unite together under the influence of heat, or solvents. Sometimes heat will, according to its intensity, produce different results, and destroy combinations which it at first produced.* NOMENCLATURE. NOMENCLATURE OF SIMPLE BODIES. THE number of simple bodies known at present, is sixtyone. * It may be well to explain by means of an illustration, or otherwise, what is meant by certain apparatus, or parts of apparatus, constantly nade use of in experiments and processes in this work, viz.: NOMENCLATURE. 15 Many simple bodies derive their names from some of their essential properties. The word chlorine, for example, has reference to the green color of this gas; the word bromine, to the foetid odor of this body. Retort is a globular vessel with a long neck, employed in a variety of distillations. It may be either without a stopper, as in Fig. 6, Plate 2, or provided with a stopper called a tubulum, forming, in this last case, a tubulated retort, as in Fig. 9, Plate 3. Receiver is a vessel intended for receiving the product of distillation, or for receiving and containing gases or other products over a pneumatic trough, or for placing over the plate of an air-pump, as a recipient. A receiver may be cylindrical, globular, or of any shape, of any suitable material. For some experiments and purposes, they must be tubulated, or have two mouths or necks; a bottle will sometimes answer. We have examples in Fig. 4, and in Figs.- 7 and 9. Matrass is an oval or globular vessel, with a long open neck, used for digestion, evaporation, &c. (Fig. 9 and Fig. 1.) Alembic, Fig. 10, used in laboratories (as well as in the large way), for distilling water when required in large quantities. It is composed of a bottom part, or boiler, in which the liquid to be distilled is placed, (properly called eucurbit), and a head or dome (which is in fact the alembic) which is provided with a neck or beak communicating with a serpentine or worm, which descends through a vessel of cold water to condense the steam; this condensed water, which is the distilled water, is conducted to a vessel or receiver, from which it is taken for use. This entire apparatus is called a distilling apparatus. Cupel is a small cup or vessel used in the process of refining metals, called cupellation. The cupel is porous, so as to absorb the baser vitrified metals when subjected to heat, and the precious metals remain. It is used, also, in certain experiments and chemical processes. Lute is a composition of tenacious substances, used for stopping the junctures of vessels so closely as to prevent the escape of, or entrance of gas or air. This term is also applied to an external covering of clay, sand, or other refractory substance, to protect glass retorts, or other vessels, from the effects of heat. It will be impossible for those ignorant of chemistry to obtain any, even the slightest, knowledge of the science, without going through the preliminary, and to some, no doubt, dry explanations of nomenclature and terms. These matters have, however, been so generalized in this work, that they are less dull than in most treatises. 16 NOMENCLATURE. Simple bodies are divided into two classes, metalloids and metals. Metals are distinguished from metalloids by certain physical characteristics, but more especially by their essential property of forming bases in their union with oxygen; while the metalloids, in combining with oxygen, only produce neutral compounds or acids. We will explain, further on, what is meant by the words bases, acids, neutral odies. No salifiable base is known which results from the combination of a metalloid with oxygen. The following is a list of metalloids and metalsMETALLOIDS. Arsenic,* Carbon, Iodine, Silicon, Nitrogen, Chlorine, Oxygen, Sulphur, Boron, Fluorine, Phosphorus, Tellurium. Bromine, Hydrogen, Selenium, METALS. Aluminum, Erbium, Niobium, Tantalum, Antimony, Tin, Gold, Terbium, Silver, Iron, Osmium, Thorium, Barium, Glu-cinium, Palladium, Titanium, Bismuth, Iridium, Pelopium, Tungsten, Cadmium, Lantanium, Platinum, Uranium, Calcium, Lithium, Lead, Vanadium, Cerium, Magnesium, Potassium, Yttrium, Chromium, Manganese, Rhodium, Zinc, Cobalt, Mercury, Rhuthenium, Zirconium, Copper, Molybdenum, Sodium, Didymium, Nickel, Strontium, NOMENCLATURE OF COMPOUND BODIES. The principle of chemical nomenclature which we owe to Guyton de Morveau, ILavoisier, Bertholet, and Fourcroy, * Arsenic resembles the metals fully in its phy-sical properties, but its combinations are so analogous with the corresponding combinations of phosphorus, that it is best to study these two bodies together. (Regnault.) NOMENCLATURE. 17 consists in designating the compound body by names indicating their composition, and sometimes even their properties. The principal compound bodies are, acids, oxides, salts, and those binary bodies of which oxygen is not one of the elements. ACIDS. The name acid is given to such bodies as have the property of reddening the tincture of litmus, and of combining with bases to form salts. Acids are divided into two principal groups: oxacids and hydracids. Oxacids are produced by the combination of a simple body with oxygen; they are named after the following rules:When a simple body combines with oxygen only in a single proportion to form an oxacid, the name of this acid is composed of the name which designates the simple body with the termination ic. Example: The oxacid formed by the combination of silicium with oxygen is called silicic acid. When a simple body combines with oxygen in two proportions to form two acids, that which contains the least oxygen takes the termination ous, and the most oxygen takes the termination ic. Example: The two acids formed by the combination of arsenic with oxygen, are termed arsenious acid, and arsenic acid; arsenious acid contains less oxygen than arsenic acid. When in fine a simple body combines in four proportions with oxygen, the prefix hypo, is placed before the name of the acid terminating in ous, or in ic. This prefix expresses a smaller quantity of oxygen than is contained in the acid terminating in ous, or ic, of which the name is not preceded by this prefix hypo. Example: The acids formed by chlorine and oxygen have received the following names: 2* B 18 NOMENCLATURE. Hypochlorous acid, chlorous acid, hypochoric acid, chloric acid. In these compounds the proportion of oxygen goes on increasing from the hypochlorous acid to the choric acid. There is an acid more oxygenated than choric acid, which is distinguished from it, and at the same time it is indicated that it contains more oxygen than choric acid, for the same quantity of chlorine, by adding to the word chloric the prefix per or hyper. Thus it is called perchoric acid, or hyperchloric acid. This rule has been extended to other acids, such as hyperiodic, and hypermanganic. fydracids. The name hydracid, is given to binary compound acids which are formed by the combination of hydrogen with a metalloid. The names of the hydracids, are composed of the name of the simple body, which is sometimes called a radical, followed by the termination IAydric. Thus the hydracids, produced by the union of hydrogen with chlorine, bromine, iodine, are named chlorhydric, bromhydric, iodhydric acids. It may be remarked that hydrogen never forms but a single hydracid with the same radical. OXIDES. The name oxide is given to the binary compounds of oxygen which exercise no action on the colour of litmus. Oxides are divided into two series. The first comprises those oxides which have not the property of combining with acids to form salts; these are termed indifferent oxides. In the second series are found those oxides which can unite with acids to form salts; these are termed 8aliftable bases. When a simple body forms but one oxide in combining NOMENCLATURE. 19 with oxygen, the compound is designated by the collective word oxide, which is followed by the name of the simple body; thus the combination of zinc with oxygen is called oxide of zinc. If the simple body is capable of uniting in several proportions with oxygen, the compounds which result are designated by adding to the term oxide the prefixed words, prot, sesqui, bi, and per, which express progressively increasing quantities of oxygen. Examples: Protoxide of manganese, of iron, of copper, of tin. Sesqui oxides of manganese, of iron, of chromium. Binoxide of manganese, of copper, of chromium. The name of peroxide is often given to such oxides as contain the most oxygen, and which still preserve the generic characters of oxides; thus we say peroxide of iron, peroxide of manganese. SALTS. When we make an acid act upon a base, we ordinarily see the properties of the acid and those of the base neutralize each other; thus the acid which at first reddened litmus, loses this property in proportion as we mix it with the base: in this case the acid and the base combine to form a salt. A salt is then the combination of an acid with a base. To give a name to the salt, we should have regard, 1st, to the nature of the acid; 2nd, to the salifiable base; 3rd, to the proportions in which the acid and the base have combined. Every acid, the termination of which is in ic, will form a salt whose termination will be in ate. Every acid whose termination is in ous, will form a salt whose termination will be in ite. These new names terminating in ate, or in ite, will be followed by the name of the oxide which enters into the salt. 20 NOMENCLATURE. Examples: Sulphuric acid and the protoxide of iron will give the sulphate of the protoxide of iron. Sulphurous acid and the protoxide of iron, will give the sulphite of the protoxide of iron. Hyposulphuric acid and the protoxide of iron will give the hyposulphate of the protoxide of iron. It often happens that to abridge the names of salts, the word oxide is suppressed; thus the salt formed by the combination of sulphuric acid and the oxide of lead is called sulphate of lead. Acid and basic bodies have the property of neutralizing each other more or less exactly, and of losing more or less completely their action on coloured re-agents. When the salt is brought as close as possible to the neutral state, its name is fixed according to the preceding rules, but if the proportion of the acid is greater than in neutral salts, the name acid salt is given to it. It is tus that the combination of sulphuric acid and potassa, which reddens litmus, would be called the acid su'lphate of potassa. If the base is in excess, the generic name is preceded by the preposition sub. Thus subacetate of lead, subnitrate of bismuth; the subsalts, are also called basic salts. The combinations which water forms with simple or com-i pound bodies are called hydrates. BINARY COMPOUNDS OF WHICHL OXYGEN IS NOT ONE OF THE ELEMENTS. When a metalloid combines with a metal to form a compound which is neither acid nor basic, the combination is designated by giving to the metalloid the termination uret, which is followed by the name of the metal; thus the combinations of sulphur, of chlorine with iron will be sulphuret of iron, chloruret of iron (or chloride of iron.) This nomenclature is applied also to the binary compounds CRYSTALLIZATION. 21 which result from the action of a hydracid on an oxide: in this case the radical of the hydracid takes the termination uret; thus chlorhydric acid in acting on the oxide of iron, produces the chloruret of iron; sulphhydric acid with the oxide of mercury, forms the sulphuret of mercury. If the metalloid combines with the metal in many proportions, the generic name is preceded by the term proto, sesqui, bi, trito or tri, quadri, yenta etc. per; thus, to name the different combinations of potassium with sulphur, which for the same quantities of metal contain quantities of sulphur represented by 1, 1~, 2, 3, 4, 5, would be said, protosulphuret, sesquisulphuret, bisulphuret, trisulphuret, quadrisulphuret, pentasulphuret of potassium. Certain bases, as ammonia, combine integrally with hydracids to form true salts. In this case, the salt takes the termination ate; thus, the combination of chlorhydric acid with ammonia is called chlorhydrate of ammonia. ALLOYS. The combinations of metals with each other are termed alloys. The alloys into which mercury enters are called amalgams. Thus the alloy of mercury and silver is called amalgam of silver. CRYSTALLIZATION OF BODIES. When a solid body has lost its cohesion by the action of extraneous causes, and these causes discontinue their action, the body gradually resumes its solid state: if this change of state takes place with sufficient slowness, the body presents itself in small masses, sometimes isolated, sometimes grouped together assuming geometric forms, and terminated on all sides by plane and brilliant surfaces. These small masses are designated by the name of crystals. The crystalline forms of a body are not always apparent 22 CRYSTALLIZATION. to the naked eye; often these can only be. distinguished by the aid of a lens or microscope. Substances which do not assume geometric forms are called amorphous. Artificial crystallization is brought about by different processes, which vary according to the properties of bodies; the chief of these methods we shall explain. CRYSTALLIZATION BY FSION. A fusible body can be made to crystallize by bringing it to such a temperature as will cause it to melt, and then allowing it to cool very slowly. The parts of the liquid in contact with the air, and those which touch the sides of the vessel in which it has been melted, cooling more rapidly, there is produced by the cooling, a crystalline layer which adheres to the sides of the vessel, and a solid crust which forms on the surface of the liquid, while the central part maintains its state of fluidity. If we pierce this upper crust and pour out the fluid, crystals will be formed in the interior, which are larger or smaller in proportion as it has cooled slowly, and as we have been operating on a more or less considerable mass. While the melted body is cooling it should be placed where it will be free from all vibration. It is thus that sulphur, bismuth, and a great number of metals and alloys are made to crystallize. CRYSTALLIZATION BY VOLATILIZATION. Solid and volatile bodies can be crystallized by volatilization; they are put into a retort of glass, or porcelain, or earthen-ware, according to the degree of their volatility. The beak of the retort is made to communicate with a receiver, properly cooled, and the retort is then brought to a temperature which will volatilize the body we wish to crystallize. CRY S TALL I ZAT I O N. 23 The vapours in cooling take the solid form, and give crystals which are deposited in the neck of the retort, or in the receiver. Arsenic, certain metallic chlorides, sundry salts of ammonia, or of mercury, crystallize by volatilization. CRYSTALLIZATION BY SOLUTION. There are two different methods of causing bodies to crystallize by solution. The first consists in dissolving the substance in a liquid, and in evaporating this liquid by means of heat, or else spontaneously, until the solid body is deposited. The form of the crystals is more beautiful in proportion to the slowness of the evaportion. The second method is founded on the unequal solubility of bodies in hot and cold liquids. Suppose a body much more soluble in hot water than in cold, nitre for example: if nitre is dissolved in boiling water, until the water becomes saturated, and the liquor is then allowed to cool, it will necessarily deposit a certain quantity of nitre. If it is cooled slowly, we will obtain beautiful crystals of this salt. It is thus that most salts are made to crystallize in laboratories, as the carbonate of soda, phosphate of soda, sulphate of copper, &c. Leblanc has made known a method which enables us to increase the size of the crystals at will, and gives them in a state of perfect regularity. We first choose small regular crystals, obtained in a crystallization by evaporation, or by cooling a solution; these are introduced into a glass crystallizer and covered with the same liquor from which they were deposited, which is called the mother-water; this liquor is then left to a spontaneous evaporation. As the liquor evaporates, it deposits on the small crystals a certain quantity of the salt which the mother-water contains; this deposit takes place in a manner so symmetrical that the crystal augments equally in all its dimensions. 24 CRYSTALLIZATION. We should be careful to turn over the crystals from time to time, so that they may increase equally on all their faces, and that the irregularities may be repaired. Many causes contribute to promote the crystallization of bodies. It may be said generally that a solution crystallizes more rapidly when it is agitated with a solid body, than when it is allowed to repose quietly without agitation; however, a solution which has been agitated, always gives small crystals. Thus a syrup evaporated as is usual, gives sugar in little crystals when it is agitated, and sugar candy, that is to say, sugar in large crystals, when it is allowed to evaporate slowly in a stove. When several bodies are dissolved in the same liquid, that which is first deposited is more pure and regular than the crystallization in a less dense medium. Thus the first crystals of chloride of sodium (marine salt) which are formed during the evaporation of sea-water, are more pure and more regular than those which are last produced. It happens often that a solution remains many days without giving any signs of crystallization, and will become a crystalline mass as soon as it is slightly agitated. Solid bodies will sometimes favour crystallization, and become in some sort a nucleus for crystals, which form on their surface; sometimes when a liquor refuses to crystallize, small crystals of the same nature with those which ought to be deposited being thrown in, bring about, by their presence, the crystallization of the whole mass. The nature of vessels in certain cases facilitates crystallization. It is remarked that a liquor crystallizes more rapidly in rough vessels, as earthen-ware, than in glass vessels. Vibrations exercise such an influence on the crystallization of bodies, that they not only facilitate the deposit of crystals in a liquor, but they can bring about the transformation of a solid amorphous body, into a crystalline body. It is thus that tough iron, of a good quality, which META L LO IDS. 25 does not present to the eye any appearance of crystallization, becomes in a short time crystallized, and very brittle, when it is exposed to long continued vibrations; this is a consideration of great importance in the arts. The form of crystals is not accidental, as might be supposed. An attentive examination has shown, that in general, and with certain exceptions, the same body always crystallizes in the same forms, and that the identity of form carries with it, if not identity of nature, at least an extreme analogy in its chemical properties. The exterior configuration is then an important characteristic, for the distinction and classification of bodies. Though crystalline forms may be, so to speak, innumerable, we are enabled to group them from certain characters of symmetry which determine the optical properties, and physical qualities, suitable to characterize them. These groups take the name of crystalline systems, and are six in number. These preliminary notions will enable us to understand the details which will be given on the metalloids and the metals. The arrangement we observe in this work is easy to comprehend. We shall take up each simple body successively and study it in a thorough manner when it presents any thing of interest, and we will then speak of the compounds it produces with the bodies which have been previously examined, whenever these compounds are of importance in their applications. METAL L OI D S. OXYGEN. OXYGEN was discovered by Priestley, in 1774; and shortly after by Scheele, who isolated it without having known of the labours of Priestley. 3 26 METALLOIDS. Lavoisier was the first to study the principal properties of oxygen; he made known the part which it played in a great number of chemical phenomena, and chiefly in combustion. This gas was first named dephogisticated air, pure air, and vital air; and afterwards, at the time chemical nomenclature was made, oxygen, from two Greek words, acid, and Evvuw, to produce, because it was then thought that all acids necessarily contained oxygen. Oxygen is a permanent gas, without color, taste, or smell its density is 1'10563. It refracts light the least of all gases. It is slightly soluble in water, which dissolves, at the ordinary temperature, one twenty-seventh of its volume of oxygen. Oxygen, suddenly compressed in a pneumatic briquet, or piston, developes a temperature which exceeds 2 Cent., and produces a bright light. M. Thenard has shown that in this case, oxygen causes the combustion of a certain quantity of fat matter, which is used to grease thme piston of the pneumatic briquet. Oxygen i's essentially proper to support combustion, which has given it the name of supporter of combustion. This property is characteristic of oxygen, and is demonstrated by the following experiment: a taper, or lighted allumet, which has been just extinguished, and on which there is the least fire, is immediately relighted, when it is plunged into a% jar filled with oxygen. The protoxide of nitrogen, which is a gas resulting from the combination of nitrogen with oxygen, also infiamies allumets almost extinguished, but wvithl less ra~pi-dity than oxygen; and the combustion is less brisk than in this latter gas. All combustible bodies, such a~s sulphur, charcoal, &c., burn in oxygen with greater brilliancy and rapidity, than in atmospheric air. Certain metals, even, will burn in oxygen, when their temperature is first elevated; thus, when an iron wire, having OXYGEN. 27 a piece of lighted tinder on its extremity, is placed in a jar filled with oxygen, the iron, Fig. 2, soon takes fire, and throws oat thousands of colored sparks; in this ease the iron, in uniting with the oxygen, forms the oxide of iron, which melts and penetrates to some depth into the glass of the jar. The temperature produced by the combustion of the iron in oxygen, is sufficiently elevated to cause the fusion of some globules of the iron which are found covered by the oxide. Phosphorus inflamed, when it is placed in a jar full of oxygen, burns there with so brilliant a light, that the eyes can scarcely endure it. The combustion of sulphur, of carbon, and of phosphorus, is shown by introducing into a glass jar, of the size of about two quarts, filled with oxygen, a small earthen or plaster cupel, supported by an iron wire fastened to a cork which is too large to enter into the mouth of the jar; the wire should be of such a length that the cupel will be suspended to within about four inches of the bottom of the jar. The combustible is then placed in the cupel, and lighted and introduced into the jar. One of the essential characters of oxygen, is that of supporting respiration. Animals placed in this gas will live a much longer time than in the same volume of atmospheric air. Hence the name of vital air, which was at first given to this gas. Electricity causes oxygen to undergo a peculiar modification, as Van Marum pointed out, in 1785. Under its influence, the affinities of this gas are more energetic than those of oxygen in its ordinary state. Thus oxygen electrified, attacks mercury and silver in presence of water and at the ordinary temperature, displaces the iodine contained in the iodides, combines directly with nitrogen to form nitric acid, causes the super-oxidation of the protoxide of lead, &c. 28 METALLOIDS. Electrified oxygen is odorous, its odor resembling that of phosphorus. This curious modification of oxygen has been studied by M. Schoenbein with great care; he gave it the name of ozone. M. Schoenbein recognised that the oxygen which disengages itself at the positive pole of a pile, of which the poles are plunged in water, is strongly ozoned, and that ozone is also obtained by passing moist air over sticks of phosphorus. More lately, Messrs. Marignac and de la Rive have shown that ozone is oxygen modified by electricity. This fact is rigorously demonstrated by experiments which prove that a limited volume of very pure oxygen, exposed for several days to the influence of a series of electric sparks, becomes entirely absorbable in the cold by silver or the iodide of potassium moistened. (lEd. Becquerel and Fremy.) The name of ozone ougIt henceforth to be replaced by that of electrified oxygen. Preparation of oxygen. The most simple mode of preparing oxygen is to decompose, by heat, certain metallic oxides. When the oxide of silver, or the oxide of mercury is heated, oxygen is disengaged, and silver or metallic mercury remains. These oxides are not however employed in laboratories for the preparation of oxygen, on account of their high price. The preference is given to the peroxide of manganese, which is found in nature, and is of low price. The apparatus in which this decomposition is made is composed of an earthen retort, into which about a pound of peroxide of manganese is introduced. The retort is placed in a reverberatory furnace; a tube provided with a safety tube to prevent absorption, connected through a cork or stopper with the neck of the retort, passes under a receiver filled with water, the retort is then heated, and gradually brought to a red heat. (Fig. 3.) OXYGEN. 29 At first there is disengaged a mixture of atmospheric air and carbonic acid. The atmospheric air which comes over, was in the retort, being displaced by the disengagement of the gas; the carbonic acid comes from the carbonates which the peroxide of manganese of commerce almost always contains; these carbonates are decomposed by the heat, and produce the carbonic acid which mixes with the oxygen. The property which carbonic acid possesses of rendering lime-water turbid, serves to show the presence of this acid in the oxygen. The receivers of the gas which first come over are allowed to escape, and the oxygen is not collected till it relights a taper with a slight report, and when lime-water is no longer precipitated by it. To free the oxygen from the carbonic acid which it sometimes contains, it is sufficient to agitate the gas with a concentrated solution of potassa, which absorbs the carbonic acid and leaves the oxygen pure. One-half of the oxygen which the peroxide of manganese contains, may be withdrawn from it by heating this oxide with concentrated sulphuric acid. The peroxide of manganese is an indifferent body, but there is another oxide of manganese, the protoxide, which is an energetic base; sulphuric acid causes the formation of this base, and combines with it. Oxygen may be obtained very pure, by extricating it from the chlorate of potash, which is transformed by heat into chloride of potassium and oxygen. We shall not press the importance of the part which oxygen plays in most chemical reactions. Oxygen forms one of the constituent parts of atmospheric air; without it, the phenomena of vegetation and combustion could not be accomplished. It unites itself besides to most of the bodies which we shall examine successively, and forms combinations which will enable us to complete its study. g* 30 METALLOIDS. HYDROGEN. Hydrogen was discovered in the beginning of the seventeenth century; but it has only been well known since the year 1777, when Cavendish described its principal properties. This gas was first called inflammable air, and then hydrogen (generator of water), from two Greek words, U[wp, water, and 79vvUw, to generate, because it is one of the elements of water. Hydrogen is a permanent gas, without color, taste, or smell, when it is pure; often it exhales a slight odor of garlic, due to the presence of a carburet of hydrogen, or to traces of sulphuretted hydrogen, and of arseniurctted hydrogen. It is rendered inodorous by passing it through solutions of salts of lead, of silver, or of mercury. Hydrogen is the lightest of all bodies. The density of the air being taken for unity at the temperature of 0~C., and under the natural pressure, Om076, that of hydrogen is equal to 0'06926; a litre of hydrogen weighs 0'08957 grammes. This gas is about fourteen and a half times lighter than the air (a litre of air weighs 1'2937 g.). The lightness of this gas may be shown by means of a receiver filled with hydrogen, which is taken out of the water vertically, and then turned over. The hydrogen immediately escapes, and is replaced by atmospheric air. If, on the contrary, the receiver is raised, keeping its orifice down towards the water, the hydrogen remains in it for some time. Finally, if the receiver containing hydrogen is placed in communication, by its orifice, with another receiver filled with air, inverting the two receivers in such a way that the one containing the air shall be above, and the one containing the hydrogen below, it will be seen that the hydrogen has HYDROGEN. 31 taken the place of the air, and the air that of the hydrogen. Hydrogen easily passes through bodies which would be almnost impermeable to other gases. If a sheet of paper is placed at a little distance from an orifice from which hydrogen is escaping, the gaseous current passes through the paper, without changing its direction, and it can be lighted on the opposite side of the sheet. But hydrogen does not pass through the thin pellicles of glass which are blown at the lamp. (M. Louyet.) Hydrogen, placed in a cracked glass receiver, over water or mercury, completely escapes. Hydrogen is not a supporter of combustion, but is very combustible. Thus, a taper, plunged into hydrogen, infiames the first layers of this gas, because they are in contact with the air, but is extinguished on penetrating further into the gas. Hydrogen is unfitted for respiration, without being poisonous. An animal dies in hydrogen only for the want of oxygen. This gas, introduced into the lungs, does not produce disorganization. Hydrogen is the most refractive of all gases. It refracts light about six and a half times more than atmospheric air. This gas, in burning in the air, combines with oxygen, and forms water. Its flame is yellow, and has but little brilliancy, because it does not contain any solid particles. Hydrogen is but slightly soluble in water, which only dissolves one and a half hundredths of its volume. Thus, it can be collected over water; but to obtain it pure, it must be collected over mercury, because water holds oxygen in solution, as well as nitrogen and carbonic acid, which are partly disengaged when it is traversed by a current of gas. Action of hydrogen on oxygen. Oxygen and hydrogen do not exercise any action on each other at the ordinary temperature, but at 400~ or 500~ C., these two gases combine and produce water. 32 METALLOIDS. This combination takes place in the proportion of 2 volumes of hydrogen to 1 volume of oxygen. The combination of hydrogen with oxygen is produced also under the influence of the electric spark, or of spongy platinum, and of several bodies in a state of minute division, which only act by their presence. Hydrogen in combining with oxygen produces a very intense heat. The high temperature which is produced by the combustion of hydrogen is made use of to melt certain refractory bodies. The mixture of hydrogen and oxygen, called detonating mixture, causes the fusion of platina which resists the heat of the forge. To fuse platina and other refractory bodies, place two volumes of hydrogen and one volume of oxygen in separate recipients; let these two gases come out and meet on the flame of a lamp; thus will be obtained a yellow flame of but little brilliancy, but which possesses calorific properties highly developed.* A mixture of hydrogen and oxygen enclosed in a flask, if presented to the flame of a candle, produces a violent detonation, and forms water. The electric spark also inflames the detonating mixture. The detonation is caused by the instantaneous condensation of vapour of water in contact with cold air. The water occupying a volume nearly 1700 times less than its vapour, a momentary vacuum is formed in the interior of the flask into which the air suddenly passes, and causes a detonation sufficiently strong to break the flask, which should be covered with a cloth when the experiment is made. A loud explosion may be produced also by inflaming a mixture of oxygen and hydrogen contained in soap-bubbles. To make this experiment, the detonating mixture is placed * This constitutes the oxyhydrogen blow-pipe invented by Dr. IIare. See Ogilvie, Imperial Dictionar?/, art. Blow-piipe, 7Y'rer's Chemistry. HIYDROGEN. 83 in a bladder furnished with a stop-cork carrying a fine tube, which is plunged into soapsuds, the bladder is compressed, the gas which comes out forms bubbles, which on being lighted, cause a detonation. A rapid succession of detonations in a glass tube gives rise to musical sounds. This phenomenon is shown, by enclosing in a large tube open at both ends a flame of hydrogen produced by dry gas, which escapes from the extremity of a finely drawn tube. From the series of detonations which cause the column of air to vibrate, there results a sound which varies in intensity with the diameter and length of the tube. The apparatus used for this experiment bears the name of chemical harmonica. The flame of a detonating mixture, hardly visible by itself to the naked eye, acquires a brilliancy which the eye can scarcely endure, by the contact of certain solid bodies, such as platina, and particularly lime. This light has been applied to the purposes of the microscope. The property which spongy platina possesses of inflaming hydrogen, has been applied to the construction of a briquet which bears the name of briquet for hydrogen, the invention of which is due to Gay Lussac. In this apparatus, the hydrogen is produced by the reaction of zinc on sulphuric acid and water. The gas is let out by a cock, and passes through a little grating of copper containing the spongy platinum, which causes the inflammuation of the hydrogen. The gas is disengaged in a bell-glass, which contains a cylinder of zinc suspended by a wire a little above the opening; this bell plunges into a jar half full of acidulated water, the gas presses back little by little the liquid from the bell, and in a little while completely expels it, and prevents the acid reacting on the zinc when the bell is filled with hydrogen. By this ingenious arrangement, the piece c 34 METALLOIDS. of zinc is preserved from the action of the acid, when the bell is filled with hydrogen, and will answer the purpose for a long time. (Fig. 8.) Preparation of hydrogen.-Hydrogen is extracted from water, which is composed of oxygen and hydrogen. This liquid, placed in contact with a body having a strong attraction for oxygen, is decomposed and disengages hydrogen. The metals in general having a strong affinity for oxygen, are used in the preparation of hydrogen. Certain metals, such as potassium and sodium, decompose water in the cold. A piece of potassium, introduced into a receiver filled with mercury, which contains a small quantity of water in its upper part, disengages hydrogen, and there remains in solution, in the water, oxide of potassium (potash). HIydrogen may be prepared by passing steam over iron heated to redness; the water is decomposed, its oxygen combines with the iron to form the magnetic oxide of iron, and hydrogen is disengaged. The apparatus employed consists of a porcelain tube containing pieces of iron wire, and a long furnace for heating it. The porcelain tube communicates on one side with a small glass retort, into which is introduced a small quantity of water, on the other, with a tube for the disengagement of the gas, which passes under a receiver filled with water, Fig. 6. We begin by heating the porcelain tube to a red heat, and then placing a chafing-dish of coals uLnder the retort, cause the vapour of water to pass over the iron; the hydrogen gas soon disengages itself with great rapidity. In the laboratory, hydrogen is always prepared by decomposing water by zinc, in presence of sulphuric acid. Zinc alone not having sufficient affinity for oxygen to decompose water at the ordinary temperature, sulphuric acid is added to it; the water is then decomposed by the zinc, under the influence of the sulphuric acid; its oxygen -7 XN B / |. gq*Xftt@,> z r *';ew,>e'*'!,s,. * 6 e WATER. 35 combines with the zinc to form oxide of zinc, which, uniting with the sulphuric acid, produces sulphate of zinc, which remains in solution in the water, and the hydrogen is disengaged. The apparatus consists in a flask with two tubulures, into which is introduced the zinc, in grains or strips. To one of these necks is fitted a tube for collecting the gas; to the other, a tube or a funnel for pouring in water, so placed that its extremity will be in the liquid, Fig. 5. The tube for the disengagement of the gas passes under a receiver filled with water. On pouring through the funnel a few grammes of sulphuric acid, a rapid evolution of the gas takes place, so that a large quantity may be collected in a few minutes. Uses of hydrogen. —Hydrogen is used in chemical laboratories for reducing oxides, and restoring them to the metallic state; metals reduced by hydrogen are in general very pure. Hydrogen also serves for isolating some metals from their combinations with chlorine and sulphur. Hydrogen is used for filling balloons. Sometimes, in this last application, it is often replaced by the gas produced in the distillation of coal. COMBINATION OF IHYDROGEN WITH OXYGEN. WATER. WATER, as well as air, was considered as an element until the end of the last century. Towards the year 1781, Priestley, Watt, and Cavendish demonstrated that hydrogen, in burning in the air, produced water. In 1789, Lavoisier demonstrated that water is composed of hydrogen and oxygen, and that these two gases form, in combining together, a quantity of water represented by the sum of their weights. 36 MiETALLOIDS. Composition of Water. Water is proved to be formed of oxygen and hydrogen: 1. By turning a jet of dry hydrogen under a bell-glass, its interior is seen to be covered with drops of water, the number of which increases as long as the combustion goes on. (Fig. 7.) 2. By placing water in contact with the metals which decompose it, whether with potassium, or at a high temperature, as with iron, tin, &c. In this decomposition, the oxygen of the water combines with the metals, and hydrogen is disengaged. 3. By decomposing water with galvanism, the oxygen appears at the positive pole, and the hydrogen at the negative. The volume of the first of these gases is sensibly one-half of that of the second. These experiments show that water is formed of oxygen and hydrogen, but it remains to fix the relation of these two elements exactly. To show in what proportion in volumes hydrogen combines with oxygen to form water, a mixture of pure oxygen and hydrogen is made in a graduated glass tube placed over a mercurial trough; this mixture is then passed into another tube, very thick and strong, called a eudiometer. These two gases are inflamed by the action of the electric spark. In this experiment it is shown that two volumes of hydrogen combine always with one volume of oxygen to form water. If, for example, the mixture is made of exactly two volumes of hydrogen and one volume of oxygen, it disappears entirely; if it is composed of two volumes of hydrogen and two volumes of oxygen, one volume of oxygen remains after the passage of the electric spark; and if the mixture is composed of three volumes of hydrogen and one volume of oxygen, one volume of hydrogen will remain after the explosion. Thus water is formed by the combination of two volumes WATER. 37 of hydrogen and one volume of oxygen. These proportions expressed in weight correspond to 88.888 of oxygen, and to 11.112 of hydrogen. Water may be solid, liquid, or gaseous. We shall examine it under these different conditions. SOLID WATER. Water in solidifying may either be amorphous or regularly crystallized. The crystalline form of solid water is that of a hexahedral prism of 1200, or that of an isosceles dodecahedron. These crystals have a double refraction. Snow often assumes the forms of stars with six points, each point or ray being a regular prism with six faces; sometimes, too, the centre of the star is occupied by a small hexagonal brilliant scale, and the rays of the star diverge from each of its angles. In passing from the liquid to the solid state, water increases in volume. Its density becomes 0.918, that of water at +40 centigrade being 1000. This increase of volume which water undergoes in solidifying explains, 1st. Why ice always remains on the surface of quiet water.* 2nd. Why the water contained in the cellular tissue of plants or fruits, which is solidified by a heavy frost, by its increase in volume bursts the capillary vessels, kills the *In tumbling streams of water, solid ice is often seen on the bottom or bed of the stream, because the water being agitated at a low temperature cannot follow the law above explained; its lighter and colder particles cannot remain at the surface, but the whole mass being agitated becomes of an uniform temperature, and crystals of ice form on the rough bed of the stream, and in time a great mass of ice forms at the bottom of the stream. We can by constant agitation on a cold day, solidify the entire mass of water in a bucket, which if left to itself would merely freeze in a cake over the surface; hence, also, the necessity of agitation in making ice-cream. —Trans. 4 38 METALLOIDS. vegetables in a short time, and brings on the rapid decomposition of frozen fruits. 3rd. Why fountains and vessels filled with water become broken during a hard frost, when the water which they contain solidifies; water-pipes which are not deeply buried also burst when the water which passes through them freezes. 4th. Why stones called frost-stones, which condense within them a considerable quantity of water, burst in winter in consequence of the dilatation which they undergo from the congelation of the water contained in their pores. 5th. Why water in soldifying sometimes causes the breaking of the most resisting metals and alloys. It is thus also we can burst cannons, by filling them with water, and exposing them, after having hermetically sealed them, to a temperature which will solidify the water they contain. Melting ice preserves a consistent temperature, which is taken for the point of departure in centigrade thermometers, and serves to fix the zero on their scale. The point at which water freezes, presenting often great variations, is not adapted for fixing the zero of thermometers. When water is not disturbed, we can, according to Gay-Lussac, lower its temperature to -12~ Cent. without solidifying it; if it is agitated it congeals instantly, and there is observed to be a disengagement of caloric which causes its temperature to rise rapidly to zero. The congelation of water then presents two remarkable phenomena, a disengagement of heat, and an increase in volume. Water which holds salts in solution freezes more slowly than pure water. When a saline solution suffers a partial congelation, it is the pure water which congeals in the first place, while the salts remain in the mother-water. This property is applied in cold countries to the concentration of the waters of the sea, from which ordinary salt or sea salt is extracted. WATER. 89 LIQUID WATER. Water is without smell, taste, or colour, though when seen in a considerable mass, it has a green tint. When water at the temperature of zero is submitted to the action of heat, its volume diminishes till at +4~0 Cent.; it then increases progressively until the temperature of its ebullition which is constant. At 8~ Cent., the volume of water is about the same as at zero. Its maximum of density is at 40 Cent. according to MI. Despretz. Water, considered as a solvent, is of importance in the arts, industry, and chemical analysis. Thus among the different properties of a body, its degree of solubility or insolubility in water ranks among the foremost. WATER IN A STATE OF VAPOUR. The boiling point of a liquid is always the same under the same pressure. Water under a pressure of 0.760m. boils at an invariable temperature, which serves to fix the 1000 Cent. of the thermometric scale.* Water passing from 00 Cent. to the state of vapour, increases to about 1700 times its volume. The temperature of water in ebullition varies with the pressure. Water enclosed in a sufficiently resisting case, may be raised to a very high temperature, and be prevented from boiling. This experiment is made in an apparatus called Papin's digester. M. Cagniard Latour enclosed water in very thick glass tubes, freed from air and sealed with a lamp. He has proved by bringing these tubes first to a red heat, that water can be compressed in steam into a space which is but four times greater than its natural volume.'* The boiling point of a body is the temperature at which the tension of its vapour is in equilibrium with the pressure exercised on it; by augmenting this pressure, its boiling point necessarily rises. 40 M ETA LLOIDS. The vapour of water Is without smell or colour, and transparent; its density, 0.622. Water, as is the case with all volatile bodies, gives off vapours at the lowest temperatures. This evaporation increases with the temperature. The vapour of water, when chilled, condenses and passes to the state of liquid water. This condensation is produced in the atmosphere, when it contains a quantity of vapour greater than it can keep in a state of saturation. It is thus are produced dews, white frost, fogs, rain, and snow. The vapour of water condensed in the atmosphere takes the name of fog when it is at the surface of the earth, and cloud when it is suspended at a certain height in the atmosphere. Vapour in condensing in the air forms little spheroids, which constitute vapour in the vesicular fornm. Water, to transform itself completely into an elastic fluid, requires about five and a half times more heat than to heat it from 0~ to 1000 Cent.* Thus, a kilogramme of vapour of water, at 1000 Cent., received into 51 kilogrammes of water at zero, produces 61 kilogrammes of water at 1000 C. This principle is made use of in workshops, to boil large masses 6f water contained in vessels of wood, which would be destroyed by the direct action of the fire. Lest the steam, in condensing, be hurtful, it is made to circulate in a double bottom, or in pipes which plunge into the liquid which it is intended to heat. Chemical properties of Water.-Water does not exercise any action on the coloured reagents. It is undecomposable by heat. Many simple bodies decompose it; some, as chlorine, combine with its hydrogen, and disengage its oxygen; others, as potassium, iron, &c., take its oxygen, and set free its hydrogen. Water combines in definite proportions with a great many * That is, from freezing to boiling. WATER. 41 bodies, and forms compounds which have received the name of hydrates. In uniting itself with acids, bases, and salts, water, in general, does not modify their distinctive characters. Thus, ordinarily we can study the properties of these bodies in their hydrates. However, in some cases, the water which unites itself to acids, bases, and salts, causes important modifications in their properties. The state of Water in nature.-The water which we meet with on the surface of the earth, and in its interior, is never pure. Rain-water holds in solution all the substances which are found in the air, such as oxygen, nitrogen, carbonic acid, and sometimes traces of nitric acid, of carbonate of ammonia, or of nitrate of ammonia. These last salts are especially found in the rain-water of storms. The rain-water that first falls contains, besides these foreign bodies, the dust which is in suspension in the atmosphere. Occasionally rain-water, collected with care, is very pure, and will do to replace distilled water in many chemical operations. The water of floods, rivers, springs, and wells, is less pure than rain-water. It contains chlorides, sulphates and carbonates of lime, magnesia, and sometimes of soda, potassa, and alumina. The composition of these waters varies with the nature of the ground they have passed over. They are mostly fit to drink, and to cook vegetables, and are without perceptible taste. In this case they have the name of soft waters, or drinkable waters. Sometimes waters are not fit for cooking vegetables, or for washing with soap; they are then said to be hard. Soft waters leave but a small residue on evaporation, preserve their transparency when boiled, are limpid, and without taste. They dissolve soap, or at least in its solutions give rise to but trifling precipitates. The bad quality of hard waters is owing to the presence of calcareous salts. They 4* 42 METALLOIDS. curdle solutions of soap, and cannot be'applied to all domestic uses. Hlrd waters are divided into two principal kinds. Waters called Selenitic.-The greater part of their lime is in the state of sulphate. Such are the waters of the wells of Paris, which are sometimes saturated with sulphate of lime (plaster). They are not rendered turbid by ebullition, and form abundant precipitates with the oxalate of ammonia, and chloride of barium. The hard waters of the second kind contain carbonate of lime, dissolved by en excess of carbonic acid. They render blue the decoction of Campeachy wood, are rendered turbid by boiling, and by exposure to the air, or under the influence of lime-water. They may be rendered fit for drinking, and for domestic use1st, by boiling them for some instants, and then letting them stand (the excess of carbonic acid which dissolved the calcareous salt, is disengaged, and the carbonate of lime is precipitated). 2d, by agitating them in contact with the air, which also causes the disengagement of the carbonic acid which is in excess, and the deposit of the carbonate of lime. 3rd, by treating them with lime-water until they are no longer precipitated by this reagent. In this case, the bicarbonate is transformed into a neutral insoluble carbonate of lime. The selenitic waters are rendered, if not drinkable, at least- fit for cooking vegetables, or for washing, when a solution of carbonate of soda is added. Thus is produced an insoluble carbonate of lime, and sulphate of soda. This last salt, though soluble in water, is of no inconvenience in most industrial operations. We may, by the aid of soap, render selenitic water fit for washing. A small quantity of soap is sufficient to precipitate all the lime in the state of insoluble margarate, stearate, and oleate of lime. These WATER. 43 precipitates being once formed, the soap will dissolve without further decomposition. The waters which are considered the most pure, are those of torrents which descend from granitic mountains. We ought, however, to prefer for drinking, the less pure waters, which contain a small quantity of calcareous salts. The experiments of M. Boussingault have clearly established that the lime of drinking water concurs with that which the aliments contain, in the development of the osseous system. The drinkable waters leave, on evaporation, a residue of which the weight does not in general go above from 1 to 3 decigrammes to the litre. This residue consists chiefly of carbonate and sulphate of lime, and of chloride of calcium. A litre of sea-water contains about 42 grammes of salts, of which marine salt comprises 26 to 27 grammes. Stalactites, incrustations, and deposits in boilers.When water, charged with the carbonate or phosphate of lime, is left in contact with the air, or submitted to the action of heat, these two salts are deposited, because the excess of carbonic acid, which held them in solution, is disengaged. Most stalactites, and many deposits of calcareous carbonates and phosphates, are formed by this sort of slow precipitation. The deposits which certain waters leave in water-conduits, have the same origin. The great quantities of water evaporated in steam-boilers, deposit on their sides calcareous salts, the hardness of which is one cause of their deterioration, because the hammer must frequently be employed to detach them. This inconvenience may be obviated by introducing into the water raspings of potatoes or clay, chlorhydrate of ammonia, or carbonate of soda. Air dissolved in Water.-Water which has been in contact with atmospheric air, contains carbonic acid, besides a mixture of oxygen and nitrogen. The presence of these gases is shown by filling with water a glass matrass of about two 44 METALLOIDS. litres capacity. This matrass is attached, by means of a cork fitted to its neck, to a tube also filled with water, which passes under a bell-glass filled with water or mercury; the water in the matrass is gradually made to boil, and in a short time a considerable quantity of gas is seen to disengage itself, and pass into the bell-glass. 100 volumes of water give about 3'2 volumes of gas. In analyzing air extracted from water, it is seen to be much richer in oxygen than atmospheric air, containing 32 or 33 volumes of oxygen in 100, instead of 21 volumes which are found in the atmosphere. The presence of this excess of oxygen is easily explained, for oxygen is much more soluble in water than nitrogen, and the small quantity of foreign matters held in solution in ordinary water, does not sensibly modify the solubility of these two gases. Air dissolved in water serves for the respiration of fish. If water is boiled to deprive it of air, and it is then allowed to cool in vessels hermetically sealed, a fish plunged in it dies in a short time. It is known, besides, that certain species of fish scarcely ever come to the surface of the water, and that all are furnished with gills fit for absorbing the oxygen in solution. When the proportion of oxygen contained in the water of a pond diminishes, the fish which people it soon perish. The air which water holds in solution gives to spring waters their fresh and agreeable taste. Those waters deprived of air become heavy, and of a slow and difficult digestion. Distilled water is insipid and disagreeable, but if it is agitated in contact with the air and saturated with it, it becomes fit to drink. It is thus that in ships they can use distilled sea-water, having first exposed it to the air. Distillation of Water.-The distillation of water has for its end the purification of it, by freeing it from foreign bodies which it holds in solution. I-~~~~~~~~~~~~~~~~~~~~~~~~E~ P / a,~~~~~~~ t~L itI L_ __ W; _~ 8 I'I_ =__ F~~~~~~ WATER. 45 These bodies are of two kinds: one kind gaseous or volatile, as oxygen, nitrogen, ammonia, nitrate and carbonate of ammonia, the other fixed, as salts with potash, soda, lime, magnesia, and alumine for bases. The first portions of distilled waters carry over with them the gaseous or volatile bodies, and are rejected as impure; the fixed compounds remain behind in the distilling apparatus. The distillation should be stopped, as soon as the salts held in solution in the water begin to be deposited. If the operation should be prolonged, the distilled water would contain small quantities of these same salts carried over mechanically, or even decomposed. Distilling apparatus.-There are many kinds of distilling apparatus. The most simple is composed of a glass retort and a receiver. (Fig. 9.) The water introduced into the retort, which it fills about three-fourths, is made to boil over an ordinary furnace. The steam condenses in the receiver, which is surrounded with water, carefully kept cool. The first portions of the distilled water are rejected; those afterwards collected are pure. The distillation ought not to be stopped till about 4 of the water has passed into the recipient. Water distilled in glass apparatus is sometimes slightly alkaline, because boiling water attacks glass of bad quality and dissolves traces of soda. Sometimes, also, distilled water contains a little muriatic acid, which proceeds from the chloride of magnesium, decomposed by the concentration, into magnesia and muriatic acid. This alteration of the water is prevented by adding a certain quantity of lime to the water to be distilled; this forms, with the chloride of magnesium, magnesia and chloride of calcium, which is not decomposed by ebullition. 46 METALLOIDS. The lime having the ability to absorb the carbonic acid contained in the water, ought to be used in excess: in most cases water is distilled without adding lime. Generally, distilled water is prepared in an apparatus called an alembic, Fig. 10. A copper boiler contains the liquid to be distilled; this is covered with a hood, or movable piece, which together form a sort of retort. The neck is fitted to a curved tube called a worm, which is plunged into a cooler, where is kept up a current of cold water admitted below, while the warm water passes out above, and may be used to feed the boiler. We owe to Gay-Lussac a small apparatus, which may be applied not only to the distillation of water, but to that of all kinds of liquids. This apparatus is composed of a matrass of glass (into which the liquid to be distilled is placed,) communicating with a condensing tube, which connects with a receiver. This tube passes through a cooler slightly inclined, which receives cold water from a cock, and loses its warm water by a lateral tube, Fig. 11. This excellent system of condensation may be employed usefully in the arts. NITROGEN. Nitrogen was discovered in 1772 by Dr. Rutherford. In 1773, Lavoisier recognised its existence in a free state in the atmosphere, of which it forms about four-fifths. Nitrogen is a permanent gas, without color or smell, and is unfit for respiration. It is this property which has given it the name of azote (o, privative, and god, life); but it is not poisonous, and animals die in it for want of oxygen. This gas is not fit for combustion. A lighted taper, plunged into a receiver filled with nitrogen, is instantly extinguished. The density of this gas is 0'97137. Nitrogen combines directly with but a small number of All ATMOSPHERIC AIR. 47 bodies; however, when a great number of electric sparks are made to pass through a mixture of oxygen and moist nitrogen, we get a combination known as azotic or nitric acid. It is to this reaction that we are to attribute the presence of nitric acid in the rains of thunder-storms. Nitrogen is less soluble in water than oxygen. Nitrogen exists in a great number of organic substances. The experiments of M. Boussingault have proved that the nitrogen contained in vegetables often comes from the air. This chemist has observed that certain leguminous plants, growing in a soil free from nitrogenous bodies, contain, after their development, a considerable quantity of nitrogen, which has been obtained from the air. The nitrogen which enters into the composition of animal substances, proceeds from that which their aliments contain; it results, in fact, from the experiments made on warm-blooded animals, that, in the act of respiration, the atmospheric nitrogen does not appear to be sensibly absorbed. Preparation of nitrogen.-In general, nitrogen is obtained from the atmospheric air. Place on a pneumatic trough a cork to float a small cupel of plaster containing a piece of ignited phosphorus, which must be covered with a bell glass, Fig. 4. The phosphorus, in burning, absorbs the oxygen which is in the bell glass, and we thus obtain the nitrogen. ATMOSPHERIC AIR. The ancients considered atmospheric air as an element. Its composition has been known only since the works of Lavoisier and Scheele. Brun and Jean Rey had shown, a century and a half before Lavoisier, that tin, heated in contact with air, increased in weight. Bayen, one of the contemporaries of Lavoisier, replaced 48 METALLOIDS. tin by mercury, and came to the same conclusion with Brun and Jean Rey. The increase in weight observed during the calcination of metals, did not indicate whether the air was absorbed entire or in part. Lavoisier first showed that atmospheric air was composed of two gases, oxygen and nitrogen, of which one only, oxygen, is absorbed by the metals. We will now describe the memorable experiment which led Lavoisier to the discovery of the composition of atmospheric air. He introduced mercury into a matrass, the neck of which, very long and curved, came out under a graduated bell-glass placed over a mercurial trough or basin. (Fig. 1.) The arrangement of this apparatus permitted him to determine with precision, ist. The volume of air on which he operated. 2d. The volume of the gas absorbed during the operation. 3d. That of the gas remaining. He kept the mercury heated just to ebullition during five consecutive days; and, although it was evident that, after the five days, the volume of air contained in the bell-glass did not suffer any further diminution, he still continued the experiment some days longer, after which he allowed the apparatus to cool, and showed that 100 volumes of air had been reduced to 73 volumes. A red crystalline substance was formed on the surface of the mercury; this substance was the peroxide of mercury. Lavoisier ascertained that the gas which remained in the bell, had properties entirely different from those of atmospheric air-that it was unfit for combustion, and respiration. This gas was nitrogen. He then introduced into a small retort the peroxide of mercury, which had formed on the surface of the mercury, ATMOSPHERIC AIR. 49 heated it to nascent redness, and saw that it was decomposed into metallic mercury and a gas, which was, as he has expressed it, much better fitted than atmospheric air to support combustion and the respiration of animals. This gas was oxygen. Lavoisier had then drawn from the atmosphere two different gases; one oxygen, supporting with energy combustion and respiration, the other nitrogen, unfit for combustion and respiration. After having decomposed atmospheric air, he desired to recompose it, by mixing the two gases which he had extracted from it. Ile found that the nitrogen, which remained in the graduated bell-glass, mixed with the oxygen resulting from the calcination of the oxide of mercury formed during the operation, produced a gas absolutely the same as atmospheric air. While Lavoisier was making these experiments on the composition of the atmosphere, Scheele arrived at the same results. The Swedish chemist showed that the alkaline sulphurets ababsorbed one of the elements of the air (oxygen), and left behind a gas unfit for respiration and combustion (nitrogen). Scheele's experiments have attracted less attention than those of Lavoisier, because they do not present the same evidence; inasmuch as the sulphurets will not, like the oxide of mercury, restore the oxygen which they have absorbed. We will observe, that the processes of Lavoisier and of Scheele, so remarkable in other respects, are not sufficiently precise. Thus, in their analysis of the air, these two chemists found more than 27 per cent. of oxygen, while the air really does not contain more than 21 per cent. Their processes have been perfected in this century, and brought almost to a rigorous exactitude. 5 D 50 METALLOIDS. Composition of the Air. When the air is analyzed by the most exact methods, it is found to contain in volume 20.80 volumes of oxygen, 79.20 volumes of nitrogen. And in weight, 23.10 parts of oxygen, 76.90 parts of nitrogen. These numbers are the results of experiments made by Messrs. Gay-Lussac, Brunner, Dumas, and Boussingault, which agree entirely with each other. Under ordinary circumstances, the atmosphere contains from 3 to 6 ten thousandths of carbonic acid, and from 6 to 9 thousandths of vapour of water. The analysis of air, taken at great heights by Gay-Lussac, and those which have been recently made at Paris by Messrs. Dumas and Boussingault, and repeated at Berne, Geneva, Brussels and Copenhagen, appear to establish a uniformity in the chemical constitution of the air, as to the proportion of oxygen and nitrogen which it contains. M. Lewy has however shown recently, that the air collected over the North Sea contains in weight 22.6 per cent. of oxygen, while the air over the land contains 23 per cent. of it. M. Lewy ascribes this difference to the fact that oxygen is more soluble in water than nitrogen, and that the animals which people these seas need the oxygen for their respiration. In proportion as these animals use the oxygen dissolved, the surface of the sea, which is in contact with the atmosphere, takes away a new quantity of oxygen from it. The oxygen and nitrogen in the atmosphere are found in the state of a simple mixture, and not of a combination. The proportion of vapour of water which the air contains is subject to great variations. It depends in general on the temperature of the air, and of the masses of water ATMOSPHERIC AIR. 51 which evaporate in certain localities. The proportion of carbonic acid is also very variable. According to M. Th. de Saussure, a rain diminishes the quantity of carbonic acid contained in the air. In passing through the atmosphere, the water becomes charged with carbonic acid, and carries it into the soil; this gas becomes again disengaged in proportion as the earth dries. A cold winter, accompanied with frosts which dry the earth, increases the quantity of carbonic acid in the air; a thaw diminishes it. Over great lakes the proportion of carbonic acid is less than over the surface of the land. The difference is 0.5 for 10,000 parts of air. The quantity of carbonic acid increases in inhabited places. On elevated mountains the proportion of carbonic acid is more considerable than on the plains, and does not seem to vary day or night. On the plains the variations are more marked. The proportion of carbonic acid is greater at night than in the day by 0.34 in 10,000 parts of air. These changes, which generally take place during the first hours after sunrise, proceed from the decomposition of the carbonic acid under the influence of the solar rays by the green parts of the plants. Messrs. Boussingault and Lewy have confirmed these results, and have assured themselves that the air of a city contains a little more carbonic acid than the air of the country. 10,000 volumes of air collected at Paris contained 3.190 of carbonic acid, and the air taken at Andilly, near Montmorency, only 2.989. The respiration of men and animals may be particularly mentioned among different causes which produce variations in the composition of a confined atmosphere. According to M. Dumas, a man consumes, in respiration, as well of carbon as of hydrogen, a quantity equivalent to 10 grammes of 52 METALLOIDS. carbon an hour. Air from the lungs contains, on an average, 4 per cent. of carbonic acid. Combustion is also a cause of variation. One kilogramme of stearic acid gives off, in burning in a vessel of 50 cubic metres, nearly 4 per cent of carbonic acid. Thus, many lamps will also cause the composition of confined air to vary. Organic substances, exposed to the air, decompose and transform the oxygen of the air into carbonic acid. We see, then, that many causes constantly tend to vary the composition of the air, and diminish the quantity of oxygen which it contains, by transforming it into carbonic acid; these are combustion, the respiration of animals, the spontaneous decomposition of organic matters, &c. But as the mass of the air is very considerable, the phenomena which take place at the surface of the earth, modify but feebly the composition of the air. The causes of alteration, however, being permanent, we could foresee an epoch when the atmosphere would be found sensibly changed, if the vegetation did not each year decompose the carbonic acid which is produced at the expense of the oxygen of the air. The beautiful experiments of Priestley, of Aime, and of Th. de Saussure, have proved, in fact, that the green parts of vegetables have the property of decomposing the carbonic acid, under the influence of the solar light, by appropriating to themselves the carbon, and restoring to the air the oxygen engaged in combination with it. Thus the relation which exists in the atmosphere between the oxygen and nitrogen, is kept up. In comparing the analyses of atmospheric air made by Gay-Lussac some years since, with those which have been lately made, it is established that the proportions of oxygen and nitrogen, contained in the air, have not varied. Analytical methods, however in other respects very complete, not being absolutely exact, it may be that the composition of ATMOSPHERIC AIR. 53 the air has undergone slight variations, which could not be appreciable unless in a great number of years. Properties of the Air-Phenomena of combustion in the Air.-It is known that the properties of the atmospheric air are composed of those of the two gases which constitute it. The action the air exercises on a simple or a compound body, is but the ensemble of the actions of oxygen and nitrogen on this body. As to its general properties, the atmospheric air should be considered an elastic fluid, permanent, inodorous, tasteless, and colorless, the density of which, represented by unity, serves as a term of comparison for the density of other gases. A litre of dry air, under the pressure of 0'760m., and at the temperature of 0~C., weighs 1'2937 grs. Combustion in the air results from the combination of a combustible body, or of its' elements, with the atmospheric oxygen. In every combustion, the oxygen is absorbed, and nitrogen undergoes no alteration. The products of combustion are, by themselves, unfit for combustion, and would soon put a stop to it if they were not replaced by a new quantity of air, the oxygen of which goes constantly to support the combustion which had commenced. Hence the necessity of establishing in fire-places what is called a draught, to carry on the combustion. It is known that wood burns imperfectly when the products of combustion are imperfectly carried off. On the contrary, combustion is energetic in a rapid current of air. In blowing on a body which is burning, we can increase the rapidity of the combustion so that it will burn as in pure oxygen; thus, a bar of iron heated to redness, and presented to the tuyere of a forge-bellows, burns, throwing out brilliant scintillations. The construction of ordinary bellows is based on this principle, as well as that of blowers used in workshops. 5* 54 METALLOIDS. Combustion in the air being the result of the combination of different bodies with its oxygen, it may be conceived that it ought to be arrested if the access of the air is suppressed. Thus we put out a coal by covering it with a bell-glass or enclosing it in an extinguisher. The state of division of bodies exercises a great influence on their combustibility; iron, carbon, the sulphurets, &c., which do not in general burn but at a high temperature, inflame at the ordinary temperature when they are exposed to the air in a state of very minute division. The bodies which exhibit this phenomenon are called pyrophoric bodies, their inflammation is caused by the disengagement of heat, which results from the condensation of the air absorbed by their pores. The combustion of an ignited body continues only when the heat developed by the combustion of one part of its mass brings the parts adjacent to those which are burning to the temperature necessary to make them burn themselves. On the contrary, the combustion ceases whenever the ignited body suffers such a reduction of temperature that it can no longer combine with oxygen. Thus a piece of iron, brought to a red heat, burns in pure oxygen, and is extinguished in atmospheric air, because the nitrogen of the air, by cooling it, stops its combustion. In the same way also, too rapid a current of air, directed on a candle, extinguishes it by reducing its temperature. A lighted coal is soon put out, when it is placed on a sheet of iron, which cools it. The gases, like solids, stop burning when they are in contact with bodies which cool them. Thus a fine wire gauze placed in a flame cools it, so that the flame cannot pass through it. This principle gave Sir Humphrey Davy the ingenious idea of his safety-lamp. This instrument is composed of an oillamp surrounded by a fine wire gauze. When this lamp is placed in an explosive mixture, a detonation is produced in ATMOSPHERIC AIR. 55 its interior, but the flame does not pass out, being cooled by the metallic gauze. The miner who works in the coal-pit, and often finds himself in explosive mixtures, is protected from danger when using Davy's lamp: and further, a fine platina wire placed within the lamp, becomes luminous in the explosive mixture which enters the lamp after the explosion, and guides the workman in the darkness. A flame is always produced by the combustion of a gas, or of a body which becomes volatilized by heat. The illuminating power of a flame varies with the prodqcts which are formed during the combustion. When these products remain in the flame in a gaseous form there is but little brilliancy, as is the case with hydrogen, carbonic oxide, and alcohol. But if there should separate during the combustion, a solid body which can become incandescent, the flame is brilliant. Thus the flames produced by the combustion of phosphorus and zinc are very brilliant, because they contain solid bodies, as phosphoric acid and oxide of zinc. The flames of illuminating gas, and those of candles and lamps, are brilliant, because they are chiefly formed of carburetted hydrogen, which undergoes an incomplete combustion, and gives out the carbon in a very divided state, which becomes incandescent. The presence of carbon in the flame of a lamp or candle may be shown by placing in it a strip of metal, which, by cooling it, is immediately covered with lamp-black. The presence of hydrogen renders the flame more brilliant. This gas, in burning, produces in fact a great heat, and brings to a white heat the molecules of carbon which give brightness to the flame. The light produced by a flame may be considerably augmented by placing in it solid bodies, such as platina wire or asbestos. Pieces of quicklime give to the flame of an ex 56 METALLOIDS. plosive mixture a brilliancy which the eye can scarcely endure. The quantity of air which supplies the flame influences its illuminating power. If the air comes in excess, it dims the flame by cooling it; if it is in small quantity, the combustion is incomplete, and the flame becomes smoky. It may be said that the flame attains its maximum of brilliancy at the moment when it is about giving off smoke. The current of air which nourishes a lamp is ordinarily produced by a chimney, the length and position of which is made to vary according to the appearance of the flame. The temperature of a flame is independent of its illuminating power; thus the flame of hydrogen, which gives a great deal of heat, is hardly visible. NITRIC OR AZOTIC ACID. Nitrogen combines with oxygen in several proportions. We shall only speak here of the most important combination, which has received the name of Azotic acid or Nitric acid. The name nitric acid or aquafortis is given to azotic acid. Azotic acid is liquid, without color, fuming in the air, and very corrosive; it is considered a violent poison. It acts on all organic bodies, and destroys them rapidly. A small quantity applied to the skin disorganizes it, and colours it yellow. Azotic acid produces a similar colour when it acts on most organic matters. This property is often used in the arts, for coloring yellow, feathers, silk, &c.: it often serves in analysis to recognise small quantities of azotic acid. Azotic acid acts on the tincture of litmus like an energetic acid, and reddens it promptly; it destroys all colouring matters, even indigo. The solution of indigo in sulphuric acid is generally used to recognise the presence of azotic acid in a liquid. Indigo, which resists the action NITRIC ACID. 57 of all acids, even that of concentrated sulphuric acid, is immediately destroyed and coloured yellow under the influence of a small quantity of azotic acid. Azotic acid* is decomposed, in a great number of cases, into water, into azote and oxygen, or even into oxygen, and into a compound less oxygenated than azotic acid; it is considered one of the most energetic oxidizers. Light, as well as heat, decomposes azotic acid. The action which the metals exercise on azotic acid is of the highest importance to the arts; it is that, in fact, which enables us to obtain most of the metallic salts. It may be said, in a general way, that azotic acid is decomposed by nearly all the metals; it oxidizes them by giving up to them part of its oxygen; these oxides once formed, unite to a part of the azotic acid not decomposed, so as to form the azotates. All the azotates being soluble in water, it is seen why azotic acid is generally used for attacking the metals; it is to transform them into soluble azotates. This action of azotic acid on the metals is always accompanied by the production of red vapours, called nitrous fumes, which are caused by the disengagement of a body less oxygenated than azotic acid, and which must necessarily arise, when azotic acid has given up to the metals a part of its oxygen. The principal metals attacked by azotic acid are iron, zinc, tin, lead, copper, mercury and silver. Among the metals which azotic acid does not attack, we will cite gold and platinum. To dissolve metals, which, like gold and platinum, are not attacked by azotic acid, a mixture of azotic and chlorhydric acids is employed in the arts, which bears the name of aqua regia. Aqua regia is obtained by mixing 1 part of azotic acid * As concentrated as possible, azotic acid contains 14.5 per cent. of water. 58 METALLOIDS. with 3 or 4 parts of chlorhydric acid. These two acids act upon each other,evolving chlorine: therefore, as chlorine attacks all the metals, even gold and platinum, to form chlorides of gold and platinum, which are soluble in water, we can understand how aqua regia may be employed to dissolve the metals which azotic acid by itself cannot attack. Preparation. —In the laboratory, azotic acid is obtained (Fig. 14.) by heating in a glass retort a mixture of azotate of potassa (nitre or saltpetre) and sulphuric acid. The azotate of potassa is a salt formed by the combination of azotic acid with potassa. Sulphuric acid has the property of displacing from its combination with potash, the azotic acid, which passes over in the distillation, and becomes condensed in the receiver. The process used in the large way for preparing azotic acid, is the same as that of the laboratories; only the azotate of potassa is often replaced by the azotate of soda, which is cheaper, and yields by its decomposition a greater quantity of azotic acid. This operation is always performed in cast-iron cylinders, having a capacity which permits their receiving a charge of from 100 to 150 kilogrammes of azotate of soda. (Fig. 15.) The cylinder communicates by tubes of earthen or glass with twelve or fifteen tubulated receivers with three necks, containing a small portion of water; the first receivers are often placed in basins and cooled by water. The application of heat ought to be gradual; and, towards the close of the operation, the cylinder is heated to redness. The acid which condenses in the receivers is of an orangeyellow color; to render it colorless and fit for commerce, it is boiledfor some time in glass or earthen vessels. The acid of commerce ordinarily marks 36 or 400; when it is wanted for the manufacture of sulphuric acid, it is employed at 32~. 100 kilogrammes of azotate of soda produce about 130 ::" ~~~~~~~~~~; r~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~z~~~~~~~~~~~~~~~~~i t Z;~~~~~~~~~~~~~~~~~~~~~~~~~~~~u<; In 1- i;~~~~~~~ s S~~~ x jfl~~w;"~A 1~~~~ i:iT -~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~i AMMONIA. 59 kilogrammes of azotic acid at 360, and 84 of sulphate of soda. This last salt is used in the manufacture of common glass or in the preparation of artificial soda. It contains marked quantities of iron, taken from the material of the cylinders, so that these apparatus undergo a gradual change, especially in that part exposed to the vapours; their wear is made to be more uniform by turning them occasionally. To preserve the cylinders, sulphuric acid is used which is never below 600; a weaker acid would destroy them very rapidly, because the water, in decomposing, would oxidize the iron. Uses.-Azotic acid being an energetic oxidizer, is one of the acids most frequently employed in the arts or in the laboratory. It serves, in the manufacture of sulphuric acid, to transform sulphurous into sulphuric acid; mixed with chlorhydric acid, it forms aqua regia, which will dissolve gold, platinum, &c. It is employed to convert starch and sugar into oxalic acid. It is used in dyeing; in engraving on copper and steel; in the assay of money; the polishing, and cleaning off the rust of metals and alloys. It unites with ligneous matters, which it transforms into pyroxyline (gun cotton). Azotic acid is a valuable reagent; all the salts which it forms with bases being soluble, it is used, in analyses, to dissolve nearly all the metals, and most oxides, carbonates, &c. It is employed in chemical researches to produce the phenomena of oxidation. AMMONIA (Volatile alkali). We place here a combination of azote and hydrogen named ammonia, which presents the properties of a base; that is to say, which unites with acids to form true salts. We will speak first of the circumstances under which ammonia can be formed. 60 METAL I, OIDS. Organic substances often contain azote in the number of their elements. They may give rise to ammonia1st, When they decompose spontaneously. 2d, When they are subjected to the influence of heat. 3d, When they are heated with a hydrated alkali, potash, for example, all the azote of the organic matter is then disengaged in the form of ammonia. Azote and hydrogen may unite in their nascent state, to produce ammonia. Thus, when an excess of hydrogen, and an oxygenated compound of nitrogen, are simultaneously passed over spongy platinum, moderately heated, ammonia is produced. In this case, platinum acts only by its presence. It has been shown that in this experiment, this metal may be replaced by the sesqui-oxide of iron. Many metals, and particularly tin, zinc, and iron, treated by azotic acid, produce ammonia. To explain the production of ammonia in the reaction of azotic acid on the metals, it is supposed that, under the influence of this acid, water is decomposed, and hydrogen becomes disengaged, which in the nascent state reacts on the azotic acid, and transforms it into ammonia. When azotic acid is introduced into a liquid which contains sulphuric acid and zinc, and which produces in consequence hydrogen, the disengagement of gas ceases sometimes completely, and the hydrogen combines with the azote of the azotic acid to form ammonia, which remains in the liquor in the state of an ammoniacal salt. Ammonia is also developed when iron is exposed to the action of moist air. Water is decomposed, its oxygen unites with the iron to form a sesqui-oxide, and its hydrogen produces, with the azote of the air, ammonia, which, combining with the carbonic acid of the air, gives carbonate of ammonia. Also, rust always disengages ammonia when it is mixed with potassa. Certain natural oxides of iron and of manganese, some clays, and many earthy matters, contain ammonia. It is AMMONIA. 61 found in small quantities in the waters of storms, in the state of carbonate or azotate of ammonia. Traces of ammonia constantly exist in the air: it can thus be seen that porous bodies, in which ammonia is found, have only condensed this gas. The secretions of animals, their urine, their excrements, contain ammonia, or bodies which easily transform themselves into salts of ammonia. It may be said, in fine, that hydrogen and azote have a great tendency to unite together to form ammonia, and that this combination is produced with facility when these two bodies, on leaving a combination, meet in their nascent state. Properties.-Ammonia is a colourless gas, of a sharp and penetrating odour, entirely characteristic. This gas is unfit for combustion; a candle placed in it is immediately extinguished. It is not combustible in atmospheric air, but when it is introduced by a small opening into a bell-glass filled with oxygen, it will inflame; it then burns with a yellow flame. Ammonia is unfit for respiration: it produces opthalmia, which is frequently seen among workmen who are exposed to ammoniacal emanations. The aqueous solution of ammonia, placed on the skin, produces a redness which is soon followed by a blister and a true cauterization; this property leads to the employment of ammonia to cauterize wounds made by the bites of mad animals. It is sometimes used to restore to consciousness persons who have fainted. Ammoniacal gas is not permanent. M. Bussy has shown that, in exposing it to the cold produced by the evaporation of liquid sulphurous acid, it can be liquified. M. Faraday has liquified ammonia, by exposing it to the cold produced by the evaporation, in a vacuum, of a mixture of solid carbonic acid and ether. Solid ammonia is white, 6 62 METALLOIDS. crystalline, transparent, heavier than liquid ammonia; its odour is weak, because, at this low temperature, its tension is not great. Ammoniacal gas reacts as an alkali on the red paper of turnsol and the syrup of violets. This property, which does not belong to any other elastic fluid, serves to characterise it; hence its name of volatile alkali. Ammonia is in general known by three characteristics: 1st, by its smell; 2d, by its alkalinity; 3d, by the white fumes of chlorhydrate of ammonia which it produces when it is approached with a glass tube which has been dipped in chlorohydric acid. If ammonia is made to pass through a tube filled with fragments of porcelain and heated to redness, it is partly decomposed, and gives rise to azote and hydrogen, which are found to be in the proportion of one volume of azote and three volumes of hydrogen; this decomposition takes place more readily if a small quantity of platinum is introduced into the porcelain tube. Electricity will also decompose ammonia. Ammonia is one of the most soluble gases in water. Water dissolves 670 times its own volume of it. If a receiver filled with pure ammonia is placed in contact with water, the gas is instantly absorbed, and water strikes the top of the receiver with such force as to break it; so that, in making this experiment, it is necessary to envelope the receiver in a cloth, for the force of the glass might wound the operator. (Fig. 21.) A piece of ice introduced into a receiver filled with ammoniacal gas, absorbs the gas rapidly, and is soon melted. Notwithstanding its great solubility, ammonia does not give white fumes in the air, because it does not combine with water in definite proportions. Water saturated with ammoniacal gas is used in chemical reactions, in place of gaseous ammonia, which it would be difficult to manage. This solution, which is often called liquid ammonia, AMMONIA. 63 gives up all the gas which it contains when it is heated to 600 Cent., or when it is exposed for a long time to the air: it loses it also in vacuum. Oxygen acts on ammonia under the influence of electricity, and forms water and azote. M. Kuhlman has shown that, under the influence of spongy platinum, slightly heated, a jmixture of ammoniacal gas and oxygen is transformed into azotic acid. Chlorine decomposes ammonia. If a few bubbles only of chlorine are made to pass through this gas, white fumes are soon formed, which are accompanied with the disengagement of heat and light. Chlorhydrate of ammonia, and nitrogen are produced. Chlorine also decomposes liquid ammonia, but then the action is less energetic. It is not accompanied with light. Preparation.-The preparation of ammonia is founded upon the property which the fixed alkalies possess of displacing it from its saline combinations. All the ammoniacal salts could be used without distinction in this preparation; but the chlorohydrate of ammonia, which is abundant in commerce, is preferred. The chlorhydrate of ammonia is a combination of ammonia and chlorhydric acid; when this salt is treated with a base, lime, for example, it is deprived of its chlorhydric acid, the ammonia is set free, and disengaged. Equal weights of quicklime and sal ammoniac,* mixed together, are introduced into a matrass; the mixture ought only to occupy a third, or a half part of the capacity of the matrass; the empty part is filled with small fragments of quicklime, for the purpose of drying the gas. (Fig. 22.) The matrass is fitted with a curved tube, which passes under a receiver. The chloride of calcium cannot be used for drying the ammonia, as it has the property of absorbing large quantities of this gas. The action of lime on the ammoni* Chlorhydrate or muriate of ammonia. 64 METALLOIDS. acal salt, commences at once; but it soon ceases, unless the matrass is heated by a chaffing-dish of coals. The gas ought to be collected over mercury, and it is known to be pure when it is completely absorbable by water. Instead of producing ammoniacal gas with lime and an ammoniacal salt, it may also be obtained by heating the liquid ammonia of commerce. A slightly elevated temperture is sufficient for the disengagement of all the gas it contains. The solution of ammonia (ordinarily called liquid ammonia) is prepared by receiving the ammoniacal gas, first in a bottle for washing it, in which is placed a small quantity of milk of lime to absorb the carbonic acid and foreign bodies which may come over with it, and afterwards in a series of bottles containing distilled water. As the ammoniacal water is lighter than pure water, the tube which brings over the gas ought to plunge to the very bottom of the bottle: to facilitate the solution of the ammonia, the condensing bottles are surrounded with cold water. The saturation may be considered complete, when numerous bubbles of the gas are seen to disengage themselves from the cool ammroniacal solution. Uses of Ammonia.-Ammonia is used in laboratories as a reagent. It is employed in medicine, and enters into the preparation of many irritating ointments. Ammonia is used for dissolving carmine, and many other coloring matters, and for modifying the tints of some colors, such as crimson, and Prussian blue. Scourers consume large quantities of ammonia for taking out grease-spots, and restoring colors changed by acids. The manufacturers of artificial pearls use ammonia for preparing the Essence d' Orient. This liquor is obtained by holding in suspension, in liquid ammonia, the small scales of a river fish called blayfish. This ammoniacal liquor is blown into the globules of glass; the scales fix themselves around the parietes of the glass, and produce the effect of natural oriental pearls. AMMONIA. 65 M. Thenard has proved that ammonia is serviceable in the treatment of animals affected with a disease known among farmers as hoven, especially sheep and cows which have been eating wet grass. The hoven is owing to the production of a considerable quantity of carbonic acid in the stomach and intestines of animals, which causes death in a short time. As the gas which produces hoven is carbonic acid, and as this gas is absorbable by ammonia, it is sufficient to administer to the animal some liquid ammonia, mixed with water, to cause the disappearance of hoven. Ammonia combines with all the acids to form ammoniaeal salts. The most important of these salts is that which results from the combination of ammonia with chlorhydric acid. This is called chlorhydrate of ammonia, or sal ammoniac. Sal ammoniac is used to prepare ammonia for the requirements of the arts, and for chemical laboratories. It is employed in the manufacture of medicinal sesqui-carbonate of ammonia; and is used for taking off the rust from metals, particularly copper. In this case, the ammonia of the chlorhydrate reduces by its hydrogen a part of the oxide of copper to the metallic state, and the chlorine transforms the other part into protochloride of copper, which volatilizes. Sal ammoniac is also employed in some processes of dyeing. The chloride of silver being soluble in an aqueous solution of sal ammoniac, a mixture of these two salts is sometimes used for silvering, without heat, copper and brass. Sal ammoniac is used in the extraction of platinum to precipitate this metal from its solution in aqua regia. Finally, sal ammoniac enters into the composition of a lute used for cementing iron into stone. This lute is prepared by sprinkling iron filings, first mixed with one or two hundreths of sulphur, with a solution of sal ammoniac. 6* E 66 METALLOIDS. CIILORINE. Scheele discovered chlorine in 1774. This chemist considered chlorine to be muriatic or marine acid deprived of phlogiston, and named it dephlogisticated marine acid. Then came Lavoisier, who regarded chlorine as being formed of muriatic acid and oxygen, and called it oxygenated muriatic acid. Finally, in 1809, Gay-Lussac and M. Thenard, in France, and Davy, in England, discovered that all the reactions of chlorine could be explained by considering oxygenated muriatic acid as a simple body; and, in fact, this body has never been decomposed. M. Ampere gave it the name of chlorine, which all chemists have adopted. Chlorine is a gas of a yellowish green color, as is indicated by its name, derived from the Greek XXwpos;; of a strong and suffocating smell, of a caustic taste, and of a density of 2'44. It is unfit for combustion. A lighted bougie, plunged into a bell-glass filled with this gas, burns for awhile, and goes out after having changed color. Chlorine is unfit for respiration, and besides is deleterious. A few bubbles of chlorine, introduced into the lungs, produce a violent suffocation, and may even cause lesions, followed by bloody expectoration.* Chlorine is not a permanent gas. Faraday liquefied it by heating to 33~C., in a tube closed at both ends, crystals formed by the combination of chlorine with water. Under the influence of a slight rise of temperature, this hydrate of chlorine is decomposed, and in the bottom of the tube two liquid strata are found; the inferior one is liquefied chlorine, the superior is water saturated with chlorile. Chlorine is soluble in water. One volume of this liquid * When much diluted may be used with benefit for inhalation in phthisis. (Gregory). CHILORINE. 67 dissolves 3'04 of it, at 8~C. It is at this point that the solubility of this gas is at its maximum. This solubility diminishes rapidly when the temperature increases; at 50~C., it is not more than 1'09. When the solution of chlorine is made to boil, it loses all the chlorine which it contains. This solution is of a greenish yellow, deeper than chlorine, and presents all the properties of gaseous chlorine. It is employed in the laboratories in preference to chlorine, because it can be more easily managed. It must not be prepared at a temperature too low, because the dissolving power of water diminishes rapidly by cooling below + 80C.; and at 0~, water will not dissolve more than one and a half times its volume of chlorine. The solution of chlorine ought to be protected from the light, because it is decomposed under the influence of solar radiation. The chlorine reacting on the elements of water, combines with hydrogen to form chlorhydric acid, and oxygen is disengaged. In cooling to the temperature of 2~ or 30 above zero, a solution of chlorine saturated at +80, it is soon seen to deposit crystals of hydrate of chlorine of a yellow white, the form of which appears to be that of an elongated octahedron with a rhomboidal base. Chlorine has a great affinity for hydrogen. The action of these two gases on each other, does not show itself much, when they are protected from the effects of ordinary light and heat. Subjected to the influence of electricity, heat, or solar light, these two gases immediately combine with a violent detonation. According to Mr. Draper, an artificial light, that of a candle for example, will also cause the combination of chlorine and hydrogen. This combination produces chlorhydric acid, which results from the union of equal volumes of chlorine and hydrogen, without condensation: thus, 2 volumes of hydrogen, and 2 volumes of chlorine, give 4 volumes of chlorhydric acid. Under a diffuse light, chlorine and hydrogen unite slowly 68 METALLOIDS. and without noise; but the mixture keeps indefinitely in a dark place. All the luminous rays do not cause the union of chlorine with hydrogen, the violet rays only having this property; thus we may expose to the direct light a mixture of chlorine and hydrogen, contained in vessels colored red, yellow, or green, without having the combination. When dry chlorine has been exposed for some time to solar radiation, it possesses more energetic chemical affinities; it combines with hydrogen at the ordinary temperature, and, sheltered from the light (M. Draper). Chlorine acts not only on free hydrogen, but also on a great number of compounds of hydrogen. Thus, when moist chlorine is made to pass into a porcelain tube heated to redness, water is decomposed, and chlorhydric acid and oxygen are formed. Under the influence, then, of solar radiation and heat, chlorine may take hydrogen from water to form chlorhydric acid. It acts in the same manner on certain organic matters, and deprives them of their hydrogen. This reaction of chlorine becomes particularly evident when it exercises itself on colouring matter. No coloring matter of an organic nature can resist the action of chlorine. The tincture of turnsol, indigo, ink, are destroyed by chlorine; this property has been applied by Berthollet to the bleaching of cotton and linen cloth. When chlorine acts on coloring matter, it is easily understood that hydrogen is taken away; but it also sometimes happens that chlorine causes the oxidation of matters subjected to its action. Chlorine then decomposes the water to form chlorhydric acid, and the oxygen being in the nascent state, is brought to bear on the colouring matter to destroy or modify it. Chlorine can then sometimes be employed as an agent of oxidation, sometimes as an agent of dishydrogenation. A hydrogenized body, after having been submitted to the CHLORINE. 69 action of chlorine, often keeps some chlorine, which is substituted for the hydrogen, and chlorhydric acid is formed. The affinity of chlorine for hydrogen, explains its action on miasms, and organic matters in a state of decomposition. The odor which certain substances give off in putrefying, is due to the presence of a combination of hydrogen and sulphur, called sulpitydric acid, which the chlorine decoumposes. Chlorine is then used to disinfect substances which disengage sulphydric acid: further, this acid being highly deleterious, chlorine is used to prevent the asphyxias produced by sulphydric acid. But chlorine, being itself deleterious, ought to be employed with caution. Chlorine does not act on hydrogenated substances only; it combines directly with many simple bodies, such as arsenic, antimony, and potassium, etc., which inflame when they are reduced to a fine powder, and thrown into a jar filled with chlorine. A copper wire, heated at the end, which is plunged into a bottle containing chlorine gas, burns completely, transforming itself into chloruret of copper. Preparation.-Chlorine is prepared by heating a mixture of oxide of manganese, and an acid formed 6f chlorine and hydrogen, called chlorhydric acid. The oxygen of the oxide of manganese unites with the hydrogen of the chlorhydric acid to form water, and chlorine is disengaged. It is received in bottles filled with air, as it attacks mercury. Water of Chlorine is obtained by directing a current of chlorine into a bottle filled with water. This solution is of a greenish yellow, and possesses all the properties of chlorine. Uses.-The uses of chlorine are numerous. Since the beautiful experiments of Berthollet, chlorine is used for bleaching cloths, taking out the color of paste for making paper, whitening old engravings, restoring old books, and taking out ink-spots. To whiten an engraving, or take out 70 METALLOIDS. an ink-spot, it suffices to plunge the paper for some instants into water which holds chlorine in solution; the paper is then passed through ordinary water to take out the smell of chlorine: in these experiments, chlorine destroys ordinary ink, but does not in any way attack printers' ink, which has for a base a fat body and lamp-black. CHLORHYDRIC ACID. Chlorhydric acid, which is a combination of chlorine and hydrogen, was for a long time called marine acid, muriatic acid, hydrochloric acid. This acid is gaseous and without colour; it throws out into the air white fumes; its small is irritating. It excites cough when it is introduced into the air passages; its density is 1.2474. It is not permanent. At +100 C., under a pressure of 40 atmospheres, it transforms itself into a liquid without colour, of a density of 1.27. This gas is unfit for combustion, very soluble in water, which, at the temperature of 0~, dissolves about 480 times its volume of it. The solution of chlorhydric gas takes place with such rapidity, that when a bell-glass filled with this gas is placed in contact with water, the column of liquid introduces itself with such rapidity into the bell-glass, as sometimes to cause its rupture. The presence of a small quantity of air very much retards the rapidity of this absorption. Chlorhydric acid blackens organic matters, and rapidly destroys them. It does not in general act on the metalloids; many metals, such as potassium, iron, tin, &c., decompose it, combining with chlorine, and disengaging hydrogen. The great ease with which chlorhydric acid is decomposed in the cold by iron and zinc, causes it sometimes to be employed in the preparation of hydrogen. Preparation of Chlorhydric Acid. - Chlorhydric acid is prepared by decomposing marine salt (chloride of sodium) CHLO'tROHYDRIC ACID. 71 by hydrated sulphuric acid. The water contained in the sulphuric acid is decomposed, and sulphate of soda and chlorohydric acid are formed. A few grammes of marine salt are introduced into a glass matrass; a tube for collecting the gas is fitted to the matrass, and concentrated sulphuric acid is poured in. The reaction first commences in the cold, but it is afterwards increased by gentle heat. In this preparation, marine salt, first fused and moulded, and then broken into fragments of some size, is used. If concentrated sulphuric acid is made to react on marine salt crystallized, and very fine, it will produce, at the moment when the acid is poured on, a very brisk effervescence, which will cause the mixture to rise up into the tube. Chlorohydric acid being very soluble in water, should be collected over mercury, which does not exercise any action upon it. The solution of chlorohydric acid in water, which is often called liquid chlorohydric acid, is obtained in laboratories, by means of Woulf's bottles, which serve for the preparation of nearly all the solutions of gas in water. This apparatus consists of a matrass communicating with a series of condensing bottles. As the solution of chlorohydric acid is heavier than water, the gas ought to be brought into each bottle by a tube which plunges hut little into the water; in this way the different strata of the liquid constantly mix. The liquid of the first bottle is never pure; it always contains volatile chlorides, and sulphuric acid, which have been carried over in the reaction; but the liquid chlorohydric acid contained in the other bottles is generally pure. Six parts of marine salt require about five parts of sulphuric acid to decompose them. The water which absorbs the chlorohydric acid increases 72 METALLOIDS'. in volume; so that, in commencing the operation, the bottles should not be entirely filled with water. In the arts, chlorhydric acid is prepared by decomposing marine salt by sulphuric acid, in kilns or cylinders of iron. The chlorhydric acid which is disengaged is condensed in a series of earthen jars communicating with a chimney, having a strong draught. (Fig. 15.) The chlorhydric acid of commerce (muriatic acid) is not pure; it contains generally all the salts which are found in common water, used for the solution of the chlorhydric acid gas, and besides, sulphurous and sulphuric acids, perchloride of iron, and sometimes arsenious and arsenic acids. The uses of Chlorhydric Acid.-The uses of chlorhydric acid are numerous and important. This acid, employed as a reagent, serves as a test for the salts of silver, to decompose the carbonates, the sulphurets, as a test for ammonia, &c. It serves in the arts for the preparation of chlorine and the decolorising chlorides, for the extraction of gelatine from bones, &c. It is used alone, or with azotic acid, to dissolve a great number of metals or alloys, and to prepare the metallic chlorides. BROMINE. Bromine was discovered in 1826 by M. Balard, who extracted it from the mother-waters of the salt-pits. This body exists in the sodas of sea-weed, in the water of the sea, in a great number of salt-springs, &c. Bromine is a liquid of a brownish red, very poisonous, of a penetrating and strong odor. Its name is derived from the Greek FpSios, fetid. In a chemical point of view, bromine resembles chlorine strongly; it has, like it, affinity for hydrogen, and bleaches colouring matters. It forms in uniting with hydrogen a hydracid, known as bromhydric acid. Bromine has not been yet applied to important uses. IODINE. 73 IODINE. This body was discovered in 1811 by Courtois. GayLussac has traced a complete history of it in one of the most important of his memoirs. Iodine does not exist in nature in a free state: like chlorine and bromine, to which it has strong analogy, it is always found united to sodium in marine plants, such as seaweed, the fuci, &c., in sponges, in sea-water, in some saline springs, and in the mineral kingdom in the state of iodide of silver. M. Bussy has shown its presence in the coal of Commentry, (Allier); and according to M. Duflos, iodine is also met with, mixed with bromine, in the coal of Silesia. Properties.-Iodine is solid at the ordinary temperature: its odor recalls that of chlorine and bromine, its color is of a metallic gray, it resembles plumbagine. Iodine crystallizes in rhomboidal scales, large and brilliant, and often in elongated octahedrons. Iodine fuses at 1070 Cent., and boils at 1800 C. The violet vapors which it produces in volatilizing have given it the name of iodine, from the Greek word i;"w6s, violet. When a well-dried matrass is heated, and a small quantity of iodine is thrown in, the matrass soon fills with violet vapors remarkable for their richness and intensity. Iodine is but slightly soluble in water, which dissolves about 0.007 of its weight at the ordinary temperature; but it is very soluble in alcohol, and gives it a very deep brown color. It is also very soluble in ether. These two solutions, by evaporation, deposit crystals of iodine. They are precipitated by water, which immediately separates from them the iodine, in the form of a brown powder. Iodine dissolves in the sulphuret of carbon, and gives to this liquid a violet colour. It exercises on organic substances a destructive action, 7 74 METALLOIDS. and colors yellow the skin, paper, &c. This color disappears under the influence of an elevated temperature, if the contact has not been sufficiently prolonged; if so, the organic matter is completely destroyed: iodine combines in this case with the hydrogen of the organic substance, to form iodohydric acid. Iodine, in reacting on other bodies, behaves in general like chlorine and bromine; but its affinities are more feeble, and these two metalloids displace it from most of its combinations. It slowly destroys coloring-matters, and does not decompose water under the influence of solar radiation. Among the properties of iodine, there is one which enables us to recognise the smallest quantity of this body, and which serves to characterise it. Placed in contact with starch, in the presence of water, it produces a blue combination, which bears the name of iodide of starch. This iodide loses its color at a temperature of 700 to 800, and recovers its blue color when the liquor is allowed to cool. This curious experiment is due to M. Lassaign. Preparation of Iodine.-Iodine may be obtained by decomposing an iodide by chlorine, which substitutes itself for the iodine, causing it to precipitate; but it is necessary to be careful to stop the disengagement of chlorine as soon as all the iodine is displaced, or the chlorine will combine with the iodine. The process which is ordinarily used to prepare iodine, consists in decomposing an alkaline iodide - of potassium, for example —by sulphuric acid, and peroxide of manganese. Uses of Iodine.-Iodine, free or combined with potassium, is applied in medicine in the treatment of goitres, and scrofulous diseases. It is used, also, in the preparation of daguerrean plates. SULPHUR. 75 SULPHUR. Natural State.-Sulphur is widely spread in nature, particularly in combination with the metals. It exists, in the native state, in volcanic districts. It enters into the composition of plaster, sulphates of barytes, strontium, &c. United with hydrogen, it forms part of a great number of mineral waters. It is met with, also, in certain animal substances, s'ome essential oils, &c. The most beautiful specimens of native sulphur come from Sicily, where they are found in crystals derived from the octahedron, and along with sulphate of strontium. Properties.-Sulphur, at the ordinary temperature, is a solid body, of a peculiar clear yellow, without taste or smell, acquiring by rubbing a characteristic odor: a bad conductor of heat and electricity. It is insoluble in water, and but slightly soluble in alcohol and ether. Sulphur is very brittle; a stick of sulphur, held in the hands when broken, produces a peculiar cracking, which results from the unequal dilatation of its molecules. This body, by rubbing, excites resinous or negative electricity. Its density is represented by the number 2'087. Sulphur fuses at the temperature of 1100 C., and boils at 460~ C. Its volatility enables us to free it easily by distillation, from impurities which it may contain. Melted sulphur, at 1100C., presents the appearance of a yellow liquid; and by cooling returns to the solid state and yellow color, as it was before fusion. In gradually raising the temperature of sulphur, it is seen, between 1400 and 150~C., to take a deep yellow color; at 1900~C., an orange color, and that its consistence has become viscous; at 2600, it becomes brown; at this temperature, its viscidity is such, that the matrass in which it is fused may be inverted without the sulphur pouring out. 76 METALLOIDS. In continuing to raise the temperature, the sulphur is seen to become again somewhat fluid. If, at this moment, it is suddenly cooled by pouring it into cold water, it remains ropy, transparent, preserves its brown color, and becomes elastic, somewhat like caoutchouc, so that it can be drawn out in long threads. It requires some time before the soft sulphur recovers its yellow color, and its original hardness. Sulphur crystallizes easily, and presents the singular property of taking two incompatible forms. One of these forms is the right octahedron, and elongated with rhomboidal base, the other is the oblique prism, with rhomboidal base. We sometimes see simple or compound bodies take two or more different and incompatible forms: it is to this property has been given the name of dimorphism, or polymorphism. Sulphur has a great affinity for oxygen. It burns in this gas, or in the air, at the temperature of about 1500, producing a beautiful blue flame, with a sharp, characteristic smell, which is that of matches when they are lighted. The product of this combustion is sulphurous acid, which is always found mixed with a small quantity of sulphuric acid. Sulphur burns brightly, when, after having lighted it, it is placed in a large bottle filled with oxygen. It combines directly with hydrogen to form an acid, known under the name of sulphydric acid; but this acid is never thus obtained, for sulphur and hydrogen do not easily unite except in the nascent state, that is to say, at the moment they each leave a combination. Sulphur also unites with chlorine, bromine, iodine, and most of the metals; some metals, as iron, copper, silver, will even inflame in the vapor of sulphur, and burn there with as much energy as in oxygen or chlorine. Extraction of Sulphur.-Sulphur employed in the arts is ordinarily extracted from the earthy soils called solfataras, SULPHUR. 77 which contain it in the native state. Sulphur is obtained for the most part in Sicily, which produces annually about 50 millions of kilogrammes. The ores of Sicily are very rich, and contain from 30 to 35 per 100 of sulphur. The sulphur is extracted by distillation, or simply by melting, when the mine is very rich. Two successive distillations are required to purify sulphur completely. In general, the first is made at the solfatara, in rude apparatus of earthenware. The second is made with more care, in the places where it is consumed, by means of a cylinder of metal which communicates with a chamber of masonry, where the sulphur condenses. When the distillation is conducted slowly, and when the sides of the chamber do not become heated above 1100 C., it condenses in the chamber in a very fine powder, which takes the name of flowers of sulphur: if the distillation is carried on rapidly, the sulphur melts, and may be run into moulds of wood, and we thus obtain roll brimstone. The flowers of sulphur usually contain sulphurous and often sulphuric acid; it reddens the tincture of litmus: to purify it, it must be washed in warm water, and then dried at a low heat. When we wish to prepare pure sulphur in the laboratories, the sulphur of commerce, or roll brimstone, is exposed to the action of heat in a glass retort connected with a glass receiver. This distillation presents no difficulty, and gives very pure sulphur. Uses of Sulphur..-Sulphur has numerous uses in the arts. It is often used for making moulds and medals, or for taking impressions. Mixed with carbon and nitre, it constitutes gunpowder. Transformed by combustion into sulphurous acid, it is used for bleaching wool and silk, for the preparation of sulphuric acid, &c. 7* 78 METALLOIDS. Sulphur is used in the manufacture of matches. It is classed among therapeutic agents. It is applied in medicine to the treatment of diseases of the skin. COMBINATIONS OF SULPHUR WITH OXYGEN. Sulphur combines with oxygen in several proportions. But we shall only here speak of sulphurous and sulphuric acids, which are the most important, and the only ones which are employed in the arts. SULPHUROUS ACID. Sulphurous acid is composed of sulphur and oxygen, and is the result of the combustion of sulphur in the air. This acid is gaseous, without color, and unfit for respiration and combustion; its odor, irritating and characteristic, is like that of burning sulphur. It acts upon the lungs, and excites cough. Sulphurous acid is soluble in water; this solution presents all the chemical properties of sulphurous gas. Uses.-Sulphurous acid has the property of decolorizing most coloring matters; thus, violets which are dipped in a solution of sulphurous acid, become in a short time completely white. The property which this acid presents, of acting on certain coloring-matters, is turned to account in the bleaching of silks and wool. These substances cannot be bleached by chlorine, which gives them a yellow color. The sulphurising of silk gives it a peculiar feeling, which is recognised when the silk is taken in the hand. At Lyons, to sulphurize silk, the sulphur is burnt in an apartment, of which the door and windows are closed with care. The sulphur, in burning in the air, changes into sulphurous acid, which reacts on the silk, and bleaches it; the silk is moistened, and sus SULPHUR. 79 pended on rails placed at the height of 3 metres. For 100 kilogrammes of silk, about two kilogrammes of sulphur are used. Sulphurous acid is employed in medicine in the treatment of diseases of the skin. It is also used for bleaching the isinglass, and the straw which is used in making hats. It is also used for taking out fruit-stains, and purifying infected places and lazarettos. Sulphurous acid is employed to sulphurize casks intended for preserving wine; its presence prevents the wine from forming vinegar. To sulphurize a cask, it is sufficient to burn in it a small quantity of sulphur, which changes into sulphurous acid. Sulphurous acid will extinguish a fire in a chimney. In this case, throw a sufficient quantity of sulphur on the coals on the hearth; this absorbs the oxygen of the air, and is transformed into sulphurous acid, which is entirely unfit for combustion. It is necessary to be careful, in this case, to close up as tight as possible all the openings of the fire-place, to prevent the access of air. With this view, damp cloths are generally used. SULPHURIC ACID. Sulphuric acid is a compound of sulphur and oxygen, more oxygenated than sulphurous acid; thus the termination ic has been given to it. Sulphuric acid is made in quantity by oxygenating sulphurous acid by a body which has the property of readily giving up its oxygen. The agent which is used for this purpose is azotic acid. Formerly, sulphuric acid was prepared in matrasses of glass; but the applications of sulphuric acid have become so important, that this acid is now made in large chambers of lead, the capacity of which has been carried even above 3,000 cubic metres. 80 METALLOIDS. The sulphuric acid produced in the leaden chambers is then concentrated in a platina apparatus. Sulphuric acid is liquid, without color or smell; it is almost twice as heavy as water. When it is poured into a vessel containing water, it falls at once to the bottom, and is then dissolved. Its oleaginous consistence has given it the name of oil of vitriol. This acid has a great affinity for water. This affinity is shown either by directly mixing the acid and water, when the combination causes a disengagement of heat which often exceeds 1000C., or in making sulphuric acid react on organic matters. When, for example, wood is dipped into sulphuric acid, a part of the oxygen and hydrogen of the organic matters combines to form water, and there is produced on the surface of the wood a black matter which contains less water than the wood. When sulphuric acid is exposed to the air, it absorbs moisture, increases in volume, and may thus take up as much as fifteen times its weight of water. When it is desirable to preserve sulphuric acid in its concentrated state, it is indispensable to keep it in a bottle well stopped. This affinity of sulphuric acid for water, is advantageously used to dry gases; it is sufficient, in fact, to pass them over bottles containing sulphuric acid, to have them quite dry. Sulphuric acid destroys organic matters by depriving them of their water; it rapidly decomposes animal membranes, and acts as a violent poison; a few drops placed on the skin produce deep burns. In cases of poisoning' by sulphuric acid, there should be immediately administered to the patient soap-suds, wood-ashes, and above all, magnesia suspended in water or oil. These principles are applicable to all cases of poisoning by the acids. A great number of metals, such as iron and zinc, are attacked by sulphuric acid. Under the influence of this acid, water is decomposed, and its hydrogen is disengaged, while its oxygen unites with the metal to form an oxide, SULPHUR. 81 which then combining with sulphuric acid, produces a salt. When sulphuric acid, diluted with water, is poured on iron, it forms sulphate of the protoxide of iron, while hydrogen is disengaged. Other metals, as copper, mercury, and silver, heated with sulphuric acid, form sulphates of copper, mercury, and silver; but in this case, it is not the water which furnishes the oxygen, but the sulphuric acid itself, a part of which is decomposed into oxygen, which goes to the metal, and into sulphurous acid, which is disengaged. This reaction is applied in the laboratories for preparing sulphurous acid. A mixture of concentrated sulphuric acid and copper is introduced into a matrass, to the neck of which a glass tube is fitted, which passes under a receiver full of mercury. Sulphurous acid is collected over mercury, because this gas is soluble in water. By heating the matrass, the reaction of the sulphuric acid commences, and in a short time several litres of sulphurous acid are obtained. Uses. —Sulphuric acid has many uses, and is employed in nearly all the chemical arts. Its energy and its fixedness render it useful in isolating most of the acids: it is used in the preparation of chlorhydric and azotic acids. The greater part of the sulphuric acid produced, is applied to the manufacture of sulphate of soda, which serves to prepare artificial soda, of which the arts consume a large quantity. It is also used in the preparation of alum, sulphate of iron, chlorine, phosphorus, sugar of starch, ether, stearic candles, and in the purification of oils. Finally, sulphuric acid is the acid which is most used in chemical laboratories. We find in commerce a fuming sulphuric acid, which is called sulphuric acid of Nordhausen, less hydrated than the above, which is principally used to dissolve indigo. This sulphuric solution of indigo is used to dye wool a Saxon blue. F 82 METALLOIDS. ACID SULPIIYDRIC. Sulphydric acid is a combination of sulphur and hydrogen, and was discovered by Scheele; it is often called hydrosulphuric acid or sulphuretted hydrogen. This acid is gaseous, without color; its fetid odor, which recalls that of rotten eggs, constitutes one of the characteristic properties of sulphydric acid. This acid liquefies under a pressure of about 1T atmospheres, and forms a very fluid, colorless liquid. Sulphydric acid is very deleterious; according to the experiments of Messrs. Th6nard and Dupuytren, a small bird dies immediately in an atmosphere which contains Ty oth of its volume of sulphydric acid; iTolTth will kill a dog of medium size; - o5th will cause the death of a horse. It is to the presence of sulphydric acid that we must attribute the accidents which unfortunately are of too frequent occurrence to workmen engaged in cleaning privysinks. Sulphydric acid gas is partly decomposed by heat, so that it cannot be directly obtained pure, by passing a mixture of vapor of sulphur and hydrogen through a red hot tube. It is set fire to by a lighted candle, and is transformed into water and sulphurous acid. When sulphydric acid is burnt in a narrow receiver, it deposits sulphur on the sides of the glass; but a jet of sulphuretted hydrogen, lighted in the open air, burns up completely, with a blue flame and sharp characteristic smell. Sulphydric acid is slightly soluble in water. This liquid dissolves only about three times its volume at the temperature of 10~ C. The aqueous solution keeps a long time without change, when protected from the air; under the influence of oxygen it becomes rapidly turbid, and the sulphuretted hydrogen is transformed into water and sulphur, which precipitates. ACID SULPHYDRIC. 83 Alcohol dissolves about six times its volume of sulphydric acid. Water, saturated with sea-salt, dissolves but a small quantity; thus we could use a solution of salt over which to collect sulphydric acid gas, instead of mercury. Chlorine acts at the ordinary temperature on sulphuretted hydrogen gas, and decomposes it, forming chlorhydric acid and a deposit of sulphur. This property might serve to combat the asphyxias caused by the absorption of sulphydric acid; but in this case, chlorine, which is itself deleterious, ought to be used with caution. As an antidote to poisoning by sulphydric acid, sulphurous acid might be used; that is, the gas which is produced by the combustion of sulphur in the air. This acid immediately decomposes sulphydric acid, producing water and sulphur. A great number of metals decompose sulphydric acid; some in the cold, some under the influence of heat, form metallic sulphurets, and disengage the hydrogen of this acid. Sulphydric acid blackens utensils of silver, copper, tin, and oil paintings, when privy-wells are emptied, because it forms in these different cases sulphurets of silver, copper, tin, &c., which are black; eggs cooked in silver vessels disengage during the cooking a small quantity of sulphydric acid, which produces on the surface of the metal black sulphuret of silver. The natural state of Sulp2hydric Acid.-Sulphydric acid is found free, or in part combined with earthy or alkaline bases, in the mineral waters called sulphur-waters. Organic substances which contain sulphur, produce on spontaneous decomposition sulphuretted hydrogen; it is found in marsh mud, and in pools where sea-water remains. Intestinal gases always contain some of it. Uses.-Sulphydric acid is one of the reagents most frequently used in analytical researches, to characterise and separate the different metals from each other: when it is free or combined, it precipitates, in the state of different 84 METALLOIDS. colored sulphurets, the metals which are held in saline combinations. Sulphur-waters are used in the form of baths, in the treatment of diseases of the skin, and many internal affections. PHOSPHORUS. The discovery of phosphorus goes back to the year 1669: it is due to Brandt and Kunckel, who extracted this body from the phosphates contained in the urine. In 1769, Gahn and Scheele showed the existence of a considerable portion of phosphate of lime in the bones, and made known an easy process for extracting the phosphorus from it. Properties. — Phosphorus is without color, transparent, insipid, of a feeble garlicky odor, and of a horny appearance. It is flexible, and sufficiently soft to be indented by the nail. A small quantity of sulphur will render it brittle. Phosphorus melts at 440 Cent., thus it can easily be melted in hot water. After having been melted, it remains liquid at the ordinary temperature, and even some degrees below zero (Cent). The experiment is made in a glass, where the phosphorus is melted in warm water; this water is taken away and replaced several times by cold water. Phosphorus thus is seen to remain liquid several minutes, in water of the temperature of 12~ to 15~ Cent.; it solidifies as soon as it is touched with a foreign body, and the thermometer rapidly rises. Phosphorus presents when perfectly pure, another phenomenon not less curious, which was observed by Thdnard. When it is exposed to a temperature of 700 Cent. and suddenly cooled, it becomes black. This modification is known to be due to a change of molecules; the color disappears by fusion. Of all the simple bodies, phosphorus presents, in regard to color, the most numerous modifications. When it is submitted to solar rays, either in vacuo, or in gases which PHOSPHORUS. 85 do not alter it chemically, as hydrogen, azote, etc.; it rapidly becomes red. This color is due to an isomeric transformation of the phosphorus; the phosphorus thus modified, bears the name of red phosphorus, or amorphous phosphorus. It is obtained by keeping it at a temperature of 2300 to 2500 Cent. in an atmosphere which cannot alter it chemically. Red phosphorus does not become luminous in the air, below a temperature of 2000 Cent.; it keeps in the air without change. It does not combine with sulphur in fusion; ordinary phosphorus would produce an explosion in contact with melted sulphur. Red phosphorus melts at 2500 Cent. At 2600 Cent. it passes back to the state of ordinary phosphorus. When it is kept for several days at a temperature a little below 260~ Cent., it collects together in a mass, very hard, of a violet brown, which appears even less alterable than red phosphorus. Red phosphorus is insoluble in the sulphuret of carbon, while ordinary phosphorus is largely soluble in this liquid. Phosphorus loses its transparency in water-it becomes rapidly yellow and opaque. When it is kept a long time in this liquid, it becomes covered with a thick coating of a yellowish white, which appears to be phosphorus in a peculiar molecular state, and sometimes, also, a combination of phosphorus and water, analogous to the hydrate of chlorine. (Pelouze). Phosphorus does not crystallize by fusion, but its solution in essential oil, in the sulphuret of phosphorus, or in the sulphuret of carbon, deposits rhomboidal dodecahedrons. The best dissolvent of phosphorus, is the sulphuret of carbon, which takes up considerable quantities of it. This solution ought to be handled with precaution, for in evaporating from a great surface, it leaves the phosphorus very much divided, which takes fire spontaneously. Thus a sheet 8 86 METALLOIDS. of paper, impregnated with this solution, takes fire as soon as the sulphuret of carbon is evaporated. Phosphorus may be reduced to powder, by introducing it into a bottle containing warm water, and agitating it rapidly till the water becomes cold; the phosphorus becomes divided into small drops, which become reduced to powder in solidifying. The property which phosphorus possesses, of becoming luminous in the dark, serves to distinguish it. It takes its name from two Greek words; lis, light, and qEw, I carry. Figures or letters traced on a tablet placed in the dark are luminous, and are said to be phosphorescent. The phosphorescence of phosphorus is generally considered as the result of a slow combination of this body with oxygen. The water in which phosphorus has been kept also glimmers in the dark. When it is agitated, it throws out a glimmering for a short time. Many bodies, such as chlorine, carburetted hydrogen, alchohol, ether, essence of turpentine, prevent phosphorus from shining in the dark. Phosphorus glimmers more in rarefied air than under the ordinary pressure of the atmosphere. Phosphorus is one of those bodies which has the greatest affinity for oxygen; a slight elevation of temperature is sufficient to make it burn, so that the distillation of it demands particular care. In distilling phosphorus we cannot use an ordinary distilling apparatus, which might break from the inflaming of the phosphorus, and expose the operator to the particles of the phosphorus thrown off in the combustion. This operation may be performed without danger in an atmosphere of hydrogen. A current of this gas is kept up in a small tubulated retort, into which is placed some twenty grains of phosphorus; this communicates with a receiver also tubulated, and carrying a recurved tube, the extremity of which P IIOSPHORUS. 87 plunges some millimters into a vessel filled with water. When the air of the apparatus has been driven out, and replaced by the hydrogen, the retort is heated with some live coals, and the distillation of the phosphorus goes on rapidly. Phosphorus combines at the ordinary temperature with oxygen. A stick of phosphorus, exposed for some time to the air, gives off white fumes, due to the formation of a peculiar acid, which has received the name of phosphatic acid. The heat which is developed during the formation of this acid, is sufficiently great to cause the inflammation of the phosphorus at the end of a few minutes. To preserve phosphorus, it must then be kept from the contact of the air by covering it with water. The slightest friction will cause the combustion of phosphorus, so that it must always be handled under water. The burns produced by phosphorus are dangerous, and a long time in healing, because it leaves in the wound a very corrosive acid (phosphoric), which is the result of its combustion. Phosphorus inflames, under ordinary pressure, at the temperature of 75~C., and burns with a very brilliant flame. Phosphorus produces, when burning in the air, and especially in pure oxygen, a very high temperature, aiA light so bright that the eyes can scarcely endure it. It then changes into phosphoric acid. This combustion ordinarily takes place, on plunging into a bottle of 4 or 5 litres capacity, and filled with oxygen, a cupel containing phosphorus, placed on a cork suspended by an iron wire, and then touching it with a hot wire or match. (See Oxygen, Fig. 12.) It rarely happens that the phosphorus becomes completely converted into phosphoric acid, even when it burns in pure oxygen, and in excess. It nearly always produces a small quantity of red phosphorus, which is preserved from 88 METALLOIDS. the action of the oxygen by a coating of phosphoric acid, which covers it. Phosphorus, notwithstanding its affinity for oxygen, does not act at the ordinary temperature on this gas, when it is pure and dry. If a stick of phosphorus is placed in a jar filled with pure and perfectly dry oxygen, the phosphorus will remain without acting on the oxygen, provided the temperature is not raised above 270~C.; but if the pressure is reduced, or another gas is introduced into the oxygen, the combination is soon brought about, and the oxygen is rapidly absorbed by the phosphorus. Preparation of Phosphorus.-Phosphorus is generally obtained from the phosphate of lime contained in the bones of animals. The bones are composed of carbonate of lime, basic phosphate of lime, and an animal matter used in the preparation of gelatine. At first, the bones are subjected to a calcination in contact with the air, to destroy the organic matter which they contain. After the calcination, the bones are white, and very friable; they contain about 77 parts of phosphate of lime, 20 parts of carbonate of lime, and a small quantity of other salts. They are reduced to a fine powder, of which are taken about six parts, which are mixed with water so as to form a very liquid paste, to which are added, by degrees, four or five parts of sulphuric acid. The sulphuric acid, by the aid of boiling, changes the carbonate of lime into sulphate of lime, disengaging carbonic acid; at the same time it takes a part of the lime from the subphosphate, and transforms it into an acid phosphate of lime. This last is very soluble in water, while the sulphate of lime (plaster) is scarcely soluble. In treating the mass, then, with water, the acid phosphate of lime is PHOSPHORUS. 89 dissolved, and the - sulphate of lime almost completely precipitates. The waters which hold the acid phosphate of lime in solution, are evaporated in a copper basin, or a porcelain capsule; during this evaporation, the greater part of the sulphate of lime held in solution deposits. This salt is taken away with care, and a syrupy liquid is obtained, containing the acid phosphate of lime nearly pure. This liquid is intimately mixed with the fourth part of its weight of powdered wood charcoal, and dried at a nascent red heat in an iron basin. This drying process is not arrested till the mass begins to give off vapors of phosphorus. It is then introduced into an earthen retort covered with a coating of fire-clay, which is filled to about three-fourths its volume with the mixture; this communicates by a copper tube with a large, wide-mouthed bottle, half full of water, which has a tube provided to give issue to the gas. The retort is brought to a bright red heat, and the phosphorus condenses in the bottle. The phosphorus obtained by the method we have just described, is not yet in a state of purity; it contains carbon, and other bodies which have been carried over during the volatilization. It is purified by melting it in warm water, and mixing it with bone-black in powder, which decolorizes it. It is then taken out with a spoon, and plunged rapidly into cold water, in order to get it in mass. To get rid of the black which it contains, it is tied up in a chamois'-skin,- which is plunged into an earthen vessel nearly full of boiling water. By compressing it with pincers, the melted phosphorus passes through the pores of the skin. Phosphorus is not found in commerce in mass, but in small sticks, some millimeters in diameter. These sticks are obtained by melting the phosphorus in 8* 90 METALLOIDS. water, and plunging into the melted phosphorus a glass tube slightly conical, into which the phosphorus is drawn by suction. Some water should be left in the tube to cover the melted phosphorus, and prevent it from getting into the mouth of the operator. The tube is then shut with the finger, and placed in cold water; the phosphorus is then taken out of the tube by a slight jar. These sticks, intended for commerce, are kept in water. Uses of Phosphorus.-The manufacture of phosphorus has become much extended of late years; and its price, formerly so high, is now not more than seven or eight francs the kilogramme. Phosphorus is chiefly used in the manufacture of friction matches, and in chemical laboratories, to analyse the air, and to prepare phosphurets, phosphoric acid, &c. ARSENIC. Arsenic is solid at the ordinary temperature, of a steel gray color, very brilliant when first sublimed, but altering rapidly in contact with the air. This body is easily reduced into powder; it is without taste, insoluble in water, and its texture is generally crystalline. Arsenic has no sensible odor at the ordinary temperature; heated to redness, or thrown on a lighted coal, it gives off a garlicky odor, very strong and characteristic. The density of arsenic is 5'75. This body, subjected to the action of heat, volatilizes without becoming liquid. However, arsenic may be fused by heating it in a metallic tube closed at both ends. Arsenic volatilizes at about 300~C.; its vapors, in condensing, give rise to tetrahedral crystals. It combines with oxygen, under the influence of a tempe ARSENIOUS ACID. 91 rature slightly elevated, burns in this gas with a light blue flame, and produces arsenious acid, improperly called in commerce, arsenic. A great number of simple bodies combine directly with arsenic; arsenic in powder, thrown into a bottle filled with chlorine, takes fire, and produces white vapors of chloride of arsenic. Arsenic is sometimes found in nature in a state of purity. It is prepared by subjecting metallic arseniurets to roasting, which form volatile arsenious acid; this acid, heated with an excess of carbon, is reduced, and gives arsenic which condenses in the receivers. Arsenic is used for the destruction of insects; it is reduced to a fine powder, and covered with water. Arsenic introduced into the stomach of an animal will not induce the symptoms of poisoning at once; it is supposed that in this case it becomes poisonous by being transformed into arsenious acid. ARSENIOUS ACID Arsenious acid is solid, white, its taste is acrid, nauseous, and excites the saliva; introduced into the stomach in small doses, it produces gangrenous spots, and causes death with extreme suffering. The antidotes of arsenious acid are the hydrate of the peroxide of iron, and magnesia. These two oxides saturate the arsenious acid, and form with it insoluble compounds, which have no further action on the animal economy.! Arsenious acid is volatile below redness, its vapours are without smell, this can be readily ascertained by volatilizing it on a brick heated to redness. Arsenious acid thrown on burning coals, gives off a garlicky odor, which is that of metallic arsenic. In this case the arsenious acid is reduced by the carbon. 92 METALLOIDS. If, in the distillation of arsenious acid, the sides of the condensing vessel become raised to a high temperature, the vapors of arsenious acid form, in condensing, a vitreous and transparent covering. But if the distillation is made into a receiver where the air circulates, the acid condenses in isolated octahedric crystals. Bodies having an affinity for oxygen, such as hydrogen and carbon, readily reduce arsenious acid. Arsenious acid immediately after volatilization is found in colorless plates, which have often the transparency of the crystal; if vitreous arsenious acid is kept for some time, even protected from the air and moisture, it is seen by degrees to lose its transparency, and to become transformed into a completely opaque body. Arsenious acid is obtained for commerce as an accessory product in roasting the ores of tin and cobalt, and as principal product by the roasting of arsenical iron. These operations are performed in reverberatory furnaces, which communicate with chambers'where the arsenious acid condenses. To purify it, it is sublimed a second time in iron vessels. In laboratories, this sublimation is made in glass or earthen retorts. Uses.-Arsenious acid is used chiefly in the manufacture of printed calicoes and in glass-works; it transforms the protoxide of iron into the sesquioxide of iron, which gives the glass less color than the protoxide. It is also employed in the seeding of corn, to preserve the seed from the attacks of insects. ARSENIURETTED HYDROGEN. Arseniuretted hydrogen is gaseous: it liquefies at - 300 C., but it has never yet been solidified. Its odor is disagreeable and very garlicky. It exercises no action on the ARSENIURETTED HYDROGEN. 93 tincture of turnsol. Water dissolves about one-fifth of its volume -of it. Exposed to the influence of moist air, it gives rise to water, and a deposit of black arsenic. Heat decomposes it into hydrogen and metallic arsenic. It is on this property is based the use of the apparatus of Marsh. Electricity causes a similar decomposition. It is combustible, and burns with a whitish flame; there are formed in the combustion of this gas, water, arsenious acid, and at the same time, a deposit of arsenic. Chlorine, bromine, iodine, decompose it by depriving it of its hydrogen. The action of chlorine takes place with the disengagement of heat and a bright light. The experiment ought to be made with small quantities of gas, and with a good deal of care, to avoid an explosion. Arseniuretted hydrogen is very poisonous: a German chemist, Gehlen, died, from having inhaled a small quantity of it. Preparation.-Arseniuretted hydrogen is prepared by the following processes:-. 1st. By treating an alloy of arsenic and tin with chlorhydric acid. 2d. By attacking an alloy of arsenic and zinc with hydrated sulphuric acid. 3d. By placing an arsenious solution in the presence of hydrogen in the nascent state. The gas prepared by these different methods is not pure, it contains a small quantity of hydrogen. Arseniuretted hydrogen being one of the most deleterious gases known, too many precautions cannot be taken in preparing it; the slightest leak in the apparatus would be dangerous for the operator. The Detection of Arsenious Acid in cases of Poisoning.The detection of arsenious acid in cases of poisoning is one of the greatest questions in legal medicine; the number of 94 METALLOIDS. cases of poisoning by arsenic alone is greater than that of all others put together. In the chemical researches relating to these kinds of poisoning, arsenious acid may be found, either in a free state in the alimentary substances which have produced the poisoning, or in the matters vomited, in the stools, in the folds of the stomach or alimentary canal. Often, also, it must be looked for in the different organs of the animal economy, where it has been carried by absorption; this last case presents itself when death has followed and the body has been for some days buried. When the arsenious acid is mixed with solid or liquid matters, it may in general be separated by mechanical means, by washing, or the employment of simple reagents. The body thus extracted is considered as arsenious acid when it is white, and when mixed with carbon, in a small narrow tube, and heated over an alcoholic lamp, it soon produces arsenic, which sublimes in a shining ring, with a metallic appearance. To characterise with still more certainty the sublimed arsenic, it must be taken from the tube and thrown on burning charcoal; it then gives off a garlicky odor, so foetid and characteristic, that this experiment enables us to recognise the smallest quantity of arsenic. When the compound of arsenic is mixed with organic matters, or when it has been absorbed by the organs, it is indispensable, in order to isolate the poison, to destroy completely the organic matters, because they would mask the reactions proper for recognising the arsenic; they would even, in some cases, present characters which might be confounded with those of arsenic. To destroy organic matters, concentrated sulphuric acid is usually employed; we thus obtain a black char, which is treated with nitric acid, in order to dissolve the arsenic, which thus is transformed into arsenic acid, soluble in water. After having carbonised the organic substances, and CARBON. 95 treated the carbon with distilled water, the solution which contains the arsenic acid is subjected to the reactions which best characterise the arsenic, it is then that the apparatus invented by Marsh, the English chemist, is used; it is called the apparatus of JMfarsh. The principle of the apparatus of Marsh is very simple, and rests on the following observation: If an arsenical compound is introduced into a flask containing water, sulphuric acid, and zinc-a mixture which disengages hydrogen-the arsenic combines with the hydrogen to form a gas called arseniuretted hydrogen. If this gas is then lighted, and a cold body, such as a porcelain plate, is placed in the flame, the plate will soon be covered with a deposit of arsenic, easy to recognise by its metallic aspect. Again, the presence of arsenic can be proved by passing the arseniuretted hydrogen gas through a glass tube slightly heated; in a short time, the tube is lined with arsenic, proceeding from the decomposition of the arseniuretted hydrogen. This apparatus, in the hands of experienced chemists who know how to appreciate the results obtained, enables them to prove with certainty the slightest traces of arsenic, which might have escaped other methods of investigation. CARBON. Chemists give the name of carbon to a simple body which, in the state of purity, constitutes the diamond, but which may be black and opaque, and then form coal, which is used as a combustible. Before particularly describing the properties of the principal kinds of carbon, we will first give the characters which are common to all the varieties of this body. Properties of Carbon. —Carbon is solid, inodorous, infusible, and fixed. 96 METALLOIDS. Many of its physical properties, such as color, brilliancy, hardness, density, sonorousness, the faculty of conducting heat and electricity, are eminently variable; so that one might be brought to consider the diamond, graphite, lampblack, anthracite, coke, charcoal, as bodies belonging to different species, which are, however, but varieties of carbon, since, like it, they combine directly with oxygen under the influence of heat, and produce carbonic acid, which is the characteristic property of carbon. Carbon burns in oxygen, and in the air, better in proportion as it is lighter; however, in a current of pure oxygen, and under the influence of a high temperature, the hardest and most dense carbon, which is the diamond, burns freely. Hydrogen, though it forms numerous combinations with carbon, is without direct action on this body. Sulphur, heated with carbon, distils without combining with it; but when the vapor of sulphur is passed over incandescent coal, these two bodies unite to produce a liquid known as sulphuret of carbon. The properties which we shall treat of may be considered as the characteristic properties of carbon, and are exhibited in the different species. We shall now examine each of these species in particular, commencing with the diamond, which is carbon, pure and crystallized. The Diamond. The true nature of the diamond remained for a long time unknown. In 1694, the experimental academicians of Florence found that the diamond burns in the focus of a burningglass. This fact was confirmed by Francis Etienne, of Lorraine, who substituted for the action of the lens that of a powerful fire of the forge. From 1766 to 1776, many French chemists, and especially Macquer, pointed out that the diamond, kept from the contact of the air, resists the CARBON. 97 most intense heat. At the same epoch, Lavoisier and Guyton de Morveau remarked that the diamond, in burning in oxygen, produced constantly carbonic acid; from whence they concluded that the diamond ought to contain carbon. The nature of the diamond was established by Humphrey Davy, who showed that this body gives, in burning, the same quantity of carbonic acid as pure carbon; that in this combustion it produces nothing but carbonic acid; and that, finally, the diamond, in burning in oxygen, does not make the volume of this gas vary. Davy concluded, from his experiments, that the diamond is pure carbon. The diamond is the hardest body known; it cannot be cut, except by its own dust; it scratches, on the contrary, all other bodies, even tempered steel. The diamond is used, in consequence of its great hardness, to form the pivots of fine watches, to polish fine stones, and to cut glass. The diamond is fixed and infusible, and a bad conductor of electricity. When insulated, it becomes very phosphorescent. It becomes electric by rubbing. Diamonds are generally without color, transparent and vitreous, but sometimes they present blue, yellow, rosy, or black tints. The diamond is ordinarily found crystallized; its principal crystalline forms are octahedron, tetrahedron, and dodecahedron rhomboid; the faces of the crystals are (Fig. 13) often curvilinear. The fracture of the diamond is generally lamellated, from the facility and neatness of its cleavage. Lapidaries make use of this property to advantage in working it. The diamond possesses, to a high degree, simple refraction. Newton, relying upon the property which combustible bodies possess of refracting light, was the first led to suspect the combustibility of the diamond. Its refractive and dispersive power gives to the diamond, when cut, its beautiful effects of light. The numerous attempts made to obtain the diamond by artificial processes 9 G 98 METALLOIDS. have failed. This body being infusible and fixed, it is conceived that the ordinary processes of crystallization, by fusion and volatilization, cannot be applied to it. Liquid cast iron is the only body which dissolves carbon and deposits it on cooling; but the carbon which separates from it is graphite, a black and opaque body. Geological inductions teach nothing on the mode of formation of the diamond; this body is always found in alluvials; it is disseminated through the ferruginous sands which constitute the ancient alluvials. The cutting of the diamond, unknown among the ancients, was discovered, in 1476, by Louis de Berquem. It added to its natural brilliancy by multiplying the number of its facets. The cutting of the diamond is executed by giving a rotatory motion to a horizontal plate of steel, covered with the powder of diamond (diamond dust) mixed with oil. The diamond is strongly pressed against this plate, and is thus freed from the rough particles which cover it. The diamond may be cut as a rose, or as a brilliant. The rose diamond has the under surface flat, and the upper elevated, enu dome, without table, and presenting twenty-four facets. The brilliant differs from the rose, in that the under surface is cut like the upper, and is composed of symmetrical facets, which correspond with the superior part of the diamond. The brilliant presents, above, a facet called table, which is surrounded with many oblique facets, Fig. 13. Diamonds which have a greenish crust are those which possess the most beautiful water after cutting. M. Jacquelain has recently arrived at interesting results, by submitting the diamond to the action of strong heat, produced by the pile of Bunsen. Under this influence, the diamond softens, separates into many fragments, loses its transparency, increases in volume, becomes black, and changes into a carbon entirely comparable to coke, and lighter than the diamond. Thus modi C-A It BON. 99 fled, the diamond still scratches glass, but becomes sufficiently friable to be broken up between the fingers. M. Jacquelain thinks that certain black diamonds, called Savoyard diamonds, may have been produced under circumstances comparable to those which he has described. Diamonds are principally found in the Indies, in the Isle of Borneo, and in Brazil. They are extracted by subrmitting the earth containing diamonds to the action of a current of water, on an inclined plane, composed of a table divided into compartments, in which are detained only the gravel and the diamonds, which are then separated by hand. There are diamonds called diamnonds of nature. These bodies are found in a coarse state, in a spheroidal form, and have no cleavage. It is impossible to cut them; they serve to make diamond powder. The largest diamond known, is that of the Rajah of Matan, at Borneo. It weighs 300 carats, or more than 65 grammes (the carat is equivalent to 0'202 grammes.) The regent diamond of the crown of France weighs 136 carats. It was bought for two and a half millions francs by the Duke of OrleansJ then Regent, from an Englishman named Pitt. On account of the beauty of its form, and its perfect clearness, it is estimated to be worth double its cost. Diamonds which can be cut sell for forty-eight francs the carat (230 francs the gramme), when they do not exceed this weight; when they are larger, their value increases considerably. Graphite or Plumbagine —Black Lead. Graphite is sometimes called plumbagine, or black lead. It contains 95 to'96 per cent. of pure carbon; it is crystalline, soft and unctuous to the touch, and soils the fingers and paper. It burns with as much difficulty as the diamond. Graphite is commonly crystallized in small tables, or in hexagonal spangles, marked with sufficient clearness. This 100 METALLOID S. body is found in the oldest transition formations, and chiefly in Bavaria, Piedmont, in the Pyrenees, and in England. Graphite was for a long time considered as a carburet of iron; but analysis has shown that pure graphite contains but a small quantity of iron, which often does not exceed -both; and the proportion of it varies with different species of graphite. It is now agreed to consider it as a variety of crystallized carbon. Artificial graphite may be obtained by allowing some kinds of melted iron, which are supersaturated with carbon, to cool slowly, and dissolving them in a mixture of chlorohydric and azotic acid. There remains in suspension in the liquor a crystalline body, of a metallic grey, identical with natural graphite. Plumbagine is used in the arts, reduced to a fine powder, and mixed with oil; it is applied to the surface of castings of iron, which it colors grey, and preserves from rust. It is often used to relieve the friction of carriage axles, of wheels of machinery, and even of clocks. Plumbagine is also used for polishing gun-bullets. In fine, plumbagine is used, either alone or with clay, to make black-lead pencils. Anthracite. Anthracite is a variety of carbon nearly pure, more brilliant than ordinary mineral coal, and blacker than graphite. It is found in Savoy, Saxony, Bohemia, England, and the United States. By its properties and composition, anthracite is intermediate between graphite and coal. Anthracite burns with difficulty on account of its compactness, and does not kindle unless in large masses, and subjected to an elevated temperature. Isolated pieces go out almost immediately, and do not run together like the fragments of pit-coal. Anthracite decrepitates when heated; this circumstance has prevented the use of it alone in blast CARBON. 101 furnaces, because the small firagments it produces in expanding, clog the furnace. There are two varieties of anthracite: vitreous anthracite, and comnmon anthracite. The first is purer than the second. Although the combustion of anthracite presents some difficulties, this body ought to be considered as a valuable combustible, which is of great service in the arts. Lamp-Black. Lamp-black is produced by the incomplete combustion of certain organic substances rich in carbon. When a piece of porcelain, or a metallic plate, is placed in the flame of a candle, a deposit of black from the smoke soon takes place on the body which cools the flame. Lamp-black is far from being pure carbon. It contains but about 80 per cent. of carbon, and is mixed with resinous matters, and different salts. It is much less impure when it has been strongly heated. Lamp-black is obtained by condensing, in brick chambers, or in large bags, the fumes which come from the incomplete combustion of resinous, bituminous, or fatty matters. The black thus procured is used in painting. Intimately mixed with dry linseed oil, it constitutes printers' ink. Lamp-black, mixed with two-thirds its weight of clay, forms the black crayons which are used in drawing. Metallic Carbon. The name metallic carbon is given to a carbonaceous residue which certain volatile substances deposit in passing through tubes of porcelain or iron heated to redness. This carbon is also produced in blast-furnaces, and in the manufacture of gas for illumination, In this latter case, the gaseous combinations of carbon and hydrogen, which result from the distillation of coal, 9 * 102 M ETALLOIDS. suffer a partial decomposition in passing through cylinders strongly heated, and produce metallic carbon. This carbon has often the brilliancy and the ringing sound of a metal. It is very hard, a good conductor of heat, and burns with difficulty. Coke. Coke is carbon produced in the distillation of coal. It is often called refined coal. Coal, submitted to the action of heat, gives rise to volatile products, formed principally of water, tar, and gas, and leaves for a residue coke, which has the porous aspect of pumice-stone. Its color is of an iron gray, with a half-metallic sound. It can be touched without leaving the marks of black on the hands. Coke draws moisture from the air. In dry times it gives up a part of this moisture. Coke does not easily burn, except in large masses; and, under the influence of a strong current of air, the incandescent fragments which are taken from the fire soon go out. Of combustibles, this produces in burning, the greatest heat. In blast-furnaces, it gives results which cannot be obtained with charcoal. Coke weighs less than coal, but more than charcoal. The hectolitre of coke, in pieces, weighs from 40 to 50 kilogrammes. Coke is used for domestic fires, but it is chiefly used for locomotive fires, and in the melting of metals. In the manufacture of iron, it is used in place of pit-coal, which cannot be employed in the working of blast-furnaces, on account of its easy fusion, and the large quantity of sulphur which it contains. Coke being much less combustible than charcoal, the carbonization of coal is effected much more easily than that of CARBON. 103 wood. Coal may be distilled in cylinders, as in the manufacture of illuminating gas; but in this case, the manufacture of coke is only accessory. Coal is often carbonized by a method analogous to that followed in the forests for the preparation of charcoal. In some localities, and principally in the environs of St. Etienne, coal is carbonized in heaps, of a prismatic form, which are 15 to 20 metres long, 1 metre high, 2'50 m. at their base, and 1'75 m. at their upper part. Finally, coke may be made in kilns of brick, of. various forms, which permit the black which forms in the incomplete combustion of the coal to be collected. The smoke is made to pass into a series of arched chambers of brick, were the black is deposited. As a mean, 100 parts of coal furnish 50 to 60 parts of coke. Charcoal.* Charcoal is the fixed residue left from the distillation of wood, or from its incomplete combustion. Wood, dried in the air, shows about the following composition:Carbon,...38-5 Water in combination,. 35.5 Ashes,. 1'0 Free water,. 25'0 100'0 It is thus seen that if, by distillation, the wood could be decomposed into water and carbon, there would be 3885 per cent. of carbon. But, during the distillation, we cannot avoid producing carburetted hydrogen gas, carbonic oxide, tar, and acetic acid, all which bodies contain carbon; so that * The object of coking, or making charcoal, is because the natural moisture, and the water formed by the combustion of the wood, absorb much of the beat, and thus prevent the attainment of the high temperature required in furnaces. 104 METALLOIDS. the most perfect methods give but 27 to 28 per cent. of carbon. The processes ordinarily used in forests give but 17 to 18 per cent. Charcoal is made by two different processes. The first, which is the most usual, is done in the open air, and is called the process of the forests. In the second process, distillingvessels are used,by which not only the charcoal is collected, but also the condensed volatile products, rich in acetic acid, and the spirit of wood, which are formed during the distillation of wood. We shall describe here the process of the forests. Carbonisation in mounds. —Charcoal is ordinarily made in the forests, by a process which is called earbonisation fit mounds, or carbonisation of the forests. In the process of carbonisation in mounds, three or four uprights, which form a chimney of about 0'30 m. in diameter, are placed in the centre of a plane, of a circular area. Around this, the billets of wood are arranged upright, in a circle. The large pieces are in the centre, the small ones on the exterior. The mound is covered with leaves and dirt. It is lighted by uncovering the chimney, and throwing into the centre of the mound some ignited coals, which are covered with small wood. At the base and exterior of the mound holes are made, which remain open during the carbonisation, for the introduction of air necessary for the operation. The chimney is left open for some hours, to ensure the combustion in the centre, and it is supplied from time to time with light wood, so as to form in the centre of the mound a mass of charcoal. When the combustion is sufficiently active, which varies according to the size of the mound, the chimney is stopped, and then it is left to itself for some hours. Whitish fumes then escape from the surface over the upper part, which begins to settle down. Vent-holes are made in the covering, towards its upper part. A volume of white smoke escapes for some hours; this then CARBON. 105 becomes bluish and almost transparent, which indicates that the carbonization is done at this place. New vent-holes are made 0'30m. to 0'40m. below the first, and so on, till we get near the holes at the base of the mound, which always remain open. During this operation, the mound settles down considerably, and the wood is converted into charcoal. General Properties of Charcoal.-Charcoal, made from hard wood, is dense; and very light when made with white wood. Charcoal always keeps the form of the wood from which it is made. Though friable, charcoal is very hard. It is often employed for polishing copper and bronze. The combustibility of charcoal varies with its density. Oak coal, which is very dense, kindles with more difficulty than that of black alder, which is very light. The latter is preferred for the preparation of gunpowder. The method which has been used for charring the wood, also influences the combustibility of the coal; and it may be said that coal prepared by distillation, is always lighter and more combustible than that which has been made in mounds. Charcoal does not begin to burn below the temperature of 240~ C.; but at the moment it is taken out of the mound, it is often pyrophoric. Introduced into store while it is warm, it sometimes kindles spontaneously. This inflammability is owing to the power which charcoal possesses of absorbing atmospheric air on disengaging heat. Charcoal decomposes water at a red heat. The hydrogen of the water is liberated, and its oxygen unites with carbon to form carbonic oxide and carbonic acid. This decomposition is effected in a porcelain tube filled with live coals, which have been first freed, by heat, of gases which they may contain. This tube communicates on one side with a small retort filled with water, and on the other with a glass tube, suitable for collecting the gas, Fig. 6. While the coal is incandescent, the water in the retort is made to boil, and 106 M E TALL OIDS. the vapor, in passing slowly over the charcoal, gives rise to gases, which are collected in receivers. Thus water, in its contact with the red hot coal, will give rise to inflammable gases, which are principally hydrogen and carbonic oxide. It is from this cause that the combustion of charcoal is rendered more active by sprinkling it with a small quantity of water; this property is well known to smiths. The proportion of ashes which charcoal leaves, in burning, varies with the quality of the wood from which it is made. Charcoal is very porous. It absorbs, in cooling, a great quantity of gas and vapor of water. Ordinary charcoal, exposed to the air, contains from 10 to 12 per cent. of water; it is a bad conductor of heat and electricity. By calcination, it loses the gases which it has absorbed, and is transformed into braise, or fine coals, which conducts heat and electricity. In this state, it is used to cover the lower extremity of lightning-rods, to facilitate the escape of the electric fluid into the soil. When two pieces of coal, calcined and cut to a point, are made to communicate with the two poles of a strong pile, if the extremities of the coals are made to approach gradually, there is produced a light whose brilliancy may be compared to that of the sun: this light is no stronger in the air than in vacuo. The experiment is ordinarily made in a glass vessel in which a vacuum has been first made. The point of the carbon, at the negative pole, is seen to be hollowed out; while that at the positive pole is covered with a carbonaceous deposit, which seems to indicate that there is a transfer of the carbon from one pole to the other. Charcoal has the property of absorbing coloring-matters; it can even cause the separation of a great number of inorganic bodies. It is probable that the bodies absorbed by the carbon adhere to it, and fix themselves to its surface, like the mordants and the coloring-matters to the surface of tissues. The organic matters which have been absorbed by CARBON. 107 the charcoal may be withdrawn, without having undergone any modification. The decolorizing properties of charcoal are developed in proportion to its porosity; these are particularly seen in animal charcoal, of which we shall speak further on. Charcoal has also the property of absorbing gases, without, however, combining with them. This is easily shown by extinguishing, in mercury, pieces of incandescent charcoal, and then letting them ascend into receivers where there are different gases. The gases are absorbed with great rapidity, particularly ammonia and chlorhydric acid. Theodore de Saussure has shown that the absorption varies with the nature of the gas. It may be said, in a general way, that the gases which are absorbed in the greatest quantity by charcoal, are the most soluble in water. The absorbed gases are disengaged when the coal is subjected to the action of a vacuum. The absorbing property of charcoal has been made use of in the arts. On account of this property, charcoal is used as a disinfectant, and to preserve animal matters from putrefaction. Meats may be preserved for a long time from putrefaction, by covering them with charcoal-dust, which is an excellent antiseptic. Some physicians advise the use of charcoal in the treatment of ulcers. Charcoal powder has been used with success to disinfect foecal matters, and convert them into poudrette, which may be used at once in agriculture. The decolorizing and disinfecting properties of charcoal may be used with advantage to purify the most putrid and dirty waters. With this view, it is used in fountains, or charcoal filters. But, as the water, in passing through the charcoal, loses the air which it holds in solution, it is im 108 METALLOIDS. portant, before drinking it, to aerate it, by agitating it for some time in contact with the air. Carbon is entirely unalterable. Thus, Chinese inks, and black paintings, which have charcoal for their basis, are considered indelible. At Herculaneum and Pompeii, manuscripts have been discovered, the black writing of which was perfectly visible. This ink was composed of lamp-black mixed with gum-water. Charcoal does not alter in moist ground. Acting on this property, wood which is to be placed in moist ground, is charred, for posts and piles. Animal Charcoal. The name animal black, or coal, is given to a mixture of finely-divided carbon and earthy salts, produced by the calcination of bones in close vessels. This body possesses, in a high degree, the property of decolorizing, which it owes, without doubt, to the great division of the carbon which it contains, for charcoal, compact, and not porous, exercises no action on coloring-matters. Animal black is chiefly used in the refining of sugar. It is met with, in commerce, in powder and in grains. When it has been used for a certain time for decolorizing syrup, it loses its decolorizing properties. It is revivified by boiling with water acidulated with chlorhydric acid, then with water, and afterwards calcining it, either alone or with bones, in kilns such as are used for making new black. The chlorhydric acid used in the revivification of animalblack, dissolves the lime which has been absorbed by the black in the decolorization of the syrups. Some refiners, in revivifying the black, first leave it to itself for some weeks, when a fermentation is developed in the mass, attributed to the presence of sugar and organic matters. The coal is then calcined. Animal-black may be revivified repeatedly; but, as there is always a quantity of dust CARBONIC OXIDE. 109 formed at each operation, the revivified black is passed over a sieve, which retains the coarse part, and allows the fine to pass through. Often the black is revivified by simply subjecting it to the action of steam, which is passed over it in red-hot cylinders of iron. It may also be revivified by boiling in water slightly alkaline, which dissolves the coloring viscous matters which it has absorbed; it is then washed with acidulated water. Preparation of Animal-black. —It is ordinarily prepared in large kilns, in which are placed iron vessels, filled with the bones to be burnt. A small quantity of fuel is sufficient to commence the operation; in fact, as soon as the disengaged gases take fire, the heat is sufficient for the calcination to go on of itself. When the organic matter is decomposed, the fire is withdrawn from the vessels; and, on cooling, the black is obtained in the shape of the bones used. This is then broken up in a mill, and passed over a sieve. Thus the coarse and fine black is produced. Combinations of Carbon with Oxygen. Carbon combines with oxygen in several proportions. We shall here speak of carbonic oxide and carbonic acid. Carbonic Oxide. This gas was discovered by Priestly; its true nature was established by Cldment D6sormes. Carbonic oxide is a gas without color, taste, or smell, of a density of 0'967; completely neutral, and slightly soluble in water. It is combustible, and burns with a characteristic blue flame, producing carbonic acid. It was for a long time supposed that this gas exercised but little action on the animal economy. But the researches of M. Leblanc show, 10 110 METALLOIDS. on the contrary, that this gas is very deleterious, and that an atmosphere containing TO- of it will kill a small bird. The knowledge of this fact is highly important in a hygienic point of view. In fact, carbonic oxide is generated in fireplaces whenever there is too much coal; and, if the products of combustion should escape into the room, either from the flues being in an improper condition, or if the damper of a stove should be kept closed, the carbonic oxide produces pains in the head, vertigo, and asphyxia, which were formerly ascribed, improperly, to carbonic acid. Oxygen, under the influence of heat, transforms carbonic oxide into carbonic acid. Some oxides are reduced by carbonic oxide - oxide of iron, for example. It is chiefly on this property is based the metallurgy of iron. Preparation.-Oxides difficult to reduce give rise to carbonic oxide when heated with charcoal, while those easy to reduce give rise to carbonic acid. Acting on this, carbonic oxide may be produced by heating to redness, in an earthen retort, charcoal with oxide of zinc. It is thus that Priestly discovered this gas. Carbonic oxide forms in fire-places when there is a deficiency of air. The blue flame which is seen in the upper part of a covered furnace proceeds mostly from the combustion of carbonic oxide. Carbonic oxide may be easily prepared by passing carbonic acid gas over coals heated to redness in a porcelain tube. The process usually employed in laboratories consists in decomposing, in a small matrass, oxalic acid, or the binoxalate of potassa (salt of sorrel), by an excess of monohydrated sulphuric acid. Take one part of oxalic acid, and five parts of concentrated sulphuric acid. Oxalic acid cannot exist in the conditions of the experiment in the anhydrous state; when it is heated with concentrated sulphuric acid, which dehydrates it, it is decomposed into equal volumes of carbonic oxide and carbonic acid. The carbonic acid is absorbed with potassa, and the carbonic oxide remains perfectly pure. A CARBONIC ACID. 111 like washing with potassa is necessary, in most cases, when carbonic oxide is prepared, because it is seldom it is formed unmixed with carbonic acid. Carbonic Acid. The name carbonic acid is given to a gas composed of carbon and oxygen. It exists in the air, and in all waters in contact with the air, in wells, in the galleries of coalmines, and in a great number of grottos and caves. Fermentation, combustion, the spontaneous decomposition of organic matters or that which results from the action of heat, and the respiration of all animals, throw into the atmosphere considerable quantities of carbonic acid, which vegetables incessantly decompose, under the influence of light, appropriating to themselves the carbon, and restoring the oxygen to the air. This decomposition of carbonic acid by the green parts of plants, explains the invariableness of the composition of the atmosphere, and at the same time the small quantity of carbonic acid which exists in it. Carbonic acid is found in combination with most metallic oxides, forming marbles, chalk, marls, carbonates of barytes, strontian, iron, copper, &c. We have to examine the acid in three states, gaseous, liquid, and solid.* G- aseous Carbonic Acid. Carbonic acid gas is colorless, of a slightly acid taste, and a pungent smell. It is about one and a half times heavier than the atmosphere. it gives to the blue tincture * In the soluble carbonates, the alkali is not neutralized; and many carbonates of the bases, especially of ammonia, may be obtained, in all of which the properties of the alkali predominate. It would appear from this that carbonic acid is not a true acid, although it combines with bases.-Gregory's Hand-Book/, p. 155. 112 AIETALL O IDS. of litmus a slight tinge of red, which disappears on exposure to the air, or by boiling the solution, because, under these influences, the carbonic acid is disengaged. The strongest heat does not change this gas, which, however, is decomposed, by a series of electric sparks, into oxygen and carbonic oxide; a curious phenomenon, inasmuch as under the influence of electricity, oxygen and carbonic oxide unite, and form carbonic acid. The density of this gas being much greater than that of the air, it can be poured from one receiver into another, as readily as a liquid. This fact explains many curious phenomena. Thus, at Pouzzoles, near Naples, in the Grotto du Chien, animals of low stature perish in a few minutes, while men can go into it without danger; the strata of carbonic acid contained in the interior of the grotto not rising higher than about a metre and a half, these animals suffocate, while man escapes. The phenomenon which takes place in the Grotto of Pouzzoles, may be produced artificially by plunging into a receiver of carbonic acid a close cylinder, which will drive out a certain quantity of the carbonic acid, which, when the cylinder is withdrawn, is replaced by an equal volume of air. In this way two different atmospheres are obtained, which do not mix for a considerable time; one is formed of air, the other of carbonic acid gas. A candle which burns in the first is extinguished in the second. Carbonic acid may cause asphyxias in cases which, unfortunately, are not sufficiently known. Thus, a vat filled with grape-juice in fermentation, placed at the entrance of a cellar, may disengage sufficient gas to asphyxiate people who might be inside the cellar. In such a case, if we had to rescue a person from such a place, we should first throw in some ammoniacal water, which, uniting with the carbonic acid, would neutralise its action on the economy. The cellars in the neighbourhood of Paris, certain wells, CARBONIC ACID. 113 or other excavations, often become filled with carbonic acid, the result of the decomposition of organic matters. Water dissolves about its own volume of carbonic acid at the ordinary pressure; but this solubility increases considerably with pressure. By compressing a mixture of carbonic acid and water, a liquid is obtained which contains five or six times its volume of carbonic acid. This has been applied to the preparation of waters called gaseous, and particularly to artificial Seltzer or mineral waters. These gaseous waters are made by two different processes. The first is a continuous process, in which a suction and force-pump draw from separate reservoirs water and carbonic acid, to force them again into a tight apparatus. The second process is intermittent, the carbonic acid being produced in the apparatus where the saturation is made, and is dissolved in the water by the pressure which it exercises on this liquid. In these two cases, the carbonic acid is produced by the action of sulphuric acid on chalk. Water, charged with carbonic acid, loses all the gas which it contains, when it is heated, or exposed to the ordinary temperature in vacuo. It loses it, also, but slowly, when it is left to itself in contact with the air. The property which carbonic acid possesses, of precipitating lime water, is often used to show the presence of this acid in water, or the gases. Carbonic acid, placed in contact with an excess of lime-water, forms with it a white flocculent precipitate, insoluble in water; very soluble, on the contrary, in chlorhydric, nitric, and acetic acids. This precipitate, which is carbonate of lime, being soluble in carbonic acid itself, requires for its formation an excess of limewater, which is shown with a red paper of litmus or turnsol, which should become blue. Without this precaution, carbonic acid might escape detection, just in proportion as its quantity is considerable. Carbonate of lime, dissolved in water by means of an 10* H 114 METALLOIDS. excess of carbonic acid, separates when this gas is disengaged under the influence of heat, by the contact of bodies in minute division, or by the action of the air alone. It is thus are produced calcareous deposits in steam-boilers, and in water-pipes. Preparation of Carbonic Acid Gas. —Carbonic acid is prepared, 1st, by burning carbon in an excess of air or of oxygen; 2d, by calcining carbonate of lime, which loses its carbonic acid, and becomes caustic lime; 3d, by decomposing carbonate of lime by an acid. In this case, the acid, acting on the carbonate, displaces the carbonic acid gas. This last process is generally used in laboratories. Carbonic acid gas may be collected over water or mercury. Liquid Carbonic Acid. Carbonic acid was first liquefied by Mr. Faraday, by decomposing, in a tube of glass closed at both ends, a carbonate by concentrated sulphuric acid. This chemist found that, at the temperature of 0~ C., carbonic acid liquefies under a pressure of 36 atmospheres. This mode of liquefaction has the double inconvenience of being dangerous for the operator, and of giving but small quantities of liquid carbonic acid. M. Thilorier, a few years ago, made an apparatus with which several kilogrammes of liquid carbonic acid can be made at a time. The principle of the apparatus of M. Thilorier is the same as that of Mr. Faraday, except that the carbonate is decomposed in a cylinder of iron, which will bear an enormous pressure. M. Thilorier produced liquid carbonic acid by decomposing the bicarbonate of soda with sulphuric acid. Properties of Liquid Carbonic Acid. -Liquid carbonic acid is colorless, very soluble in alcohol, ether, and the essential oils; it does not mix with water. Liquid carbonic acid is very dilatable: in passing sud CARBONIC ACID. 115 denly from ntce liquid to the gaseous state, it produces the extraordinary cold of about 1000 C. below zero. By throwing a jet of liquid carbonic acid into a metallic box pierced with holes, the vessel is seen to fill, almost entirely, with a white flaky matter like snow, which is solid carbonic acid, produced under the influence of the great cold which a part of this acid makes the other part undergo, in passing from the liquid to the gaseous state. Solid Carbonic Acid. Carbonic acid solidified as we have explained, keeps for some time in the open air, without being subjected to any pressure. Solid carbonic acid exists at a temperature of 900 C. below zero, and yet does not produce, on the organized tissues, a frigorific effect as great as one would suppose, no doubt owing to its porosity, and above all to the gaseous atmosphere which surrounds it. The intensity of the cold produced by solid carbonic acid is increased by mixing it with ether. This mixture will congeal, in a few seconds, four times its weight of mercury. Solidified mercury has the appearance of lead. M. Thilorier was able to make of it pieces of money, medals, &c., and to keep them for some time, in a mixture of ether and solid carbonic acid. The effect produced on the animal tissues by a mixture of solid carbonic acid and ether is like that of a burn. The fluids become solidified, the blood coagulates and becomes completely hardened; an active inflammation soon shows itself in the organ subjected to the influence of this excessive cold. Mr. Faraday has still further increased the cold produced by this mixture, by putting it under the receiver of an air-pump. By placing in this mixture tubes of glass or copper, in which gases may be compressed to 40 atmospheres with a 116 METALLOIDS. force-pump, Mr. Faraday has produced liquefactions and solidifications of gas which could not be produced by other methods. Combinations of Carbon with Hydrogen. The combinations of carbon with hydrogen are very numerous. Many essential oils, such as the essence of rose, of citron, of turpentine, &c., and caoutchouc, are formed entirely of carbon and hydrogen. When organic substances decompose slowly, or are subjected to the action of fire, they disengage gaseous compounds, formed also of carbon and hydrogen; these bodies are known under the names of proto-carburei and bicarburet of hydrogen. The proto-carburet of hydrogen, or the gas of marshes, arises during the spontaneous decomposition of a great number of organic matters. Muddy or stagnant waters, when they are stirred up, disengage gaseous compounds, in great part, of proto-carburet of hydrogen, mixed with nitrogen, oxygen, and carbonic acid. Proto-carburet of hydrogen is found in the galleries of coal-mines, where it is mixed with air and bicarburet of hydrogen, forming an explosive mixture. It is to its presence, especially, that are to be attributed the fire-damps, which occasion such serious accidents in mines. Proto-carburet of hydrogen is found in a state of compression, more or less great, in certain specimens of mineral salt, from which it is separated by the action of water, giving rise to a decrepitation. This gas escapes spontaneously from the earth in some places, where it is sometimes employed as a combustible; it also exhales in considerable quantities from the craters of some volcanoes. The organic matters which, by burning, produce the ILLUMINATING GAS. 117 greatest quantities of carburetted hydrogen, are coal, fat, and resinous bodies. All these substances are rich in hydrogen and carbon. The bicarburet of hydrogen, like the proto-carburet, forms in the distillation of most organic bodies; it is this which, mixed with the proto-carburet, constitutes, in great part, gas for illumination. The bicarburet of hydrogen is obtained pure by heating a mixture of alcohol and concentrated sulphuric acid. The bicarburet is colorless; it burns with a bright white flame. With oxygen or atmospheric air, it forms a mixture which explodes with greats force under the influence of ignited bodies. In this case, the hydrogen contained in this gas unites with the oxygen to form water; while the carbon, also combining with the oxygen, forms carbonic acid. These detonating mixtures are very dangerous, and are produced whenever illuminating gas becomes mixed with the air of a room. This gas is also used for inflating balloons. It unites with chlorine to form an oily substance. This reaction has given it the name of olejfiant gas. General Remarks on illuzninating by Gas. An organic matter, subjected to the action of a high temperature, is decomposed into carbon, and volatile or gaseous substances more or less combustible, whose flame is sometimes very brilliant. It is to Lebon, a French engineer, that are due the first experiments on lighting by gas. He published a memoir having for title-" Thermo-Lamps, or stoves wihic heat, light with economy, and offer, with several valuable products, a motive power applicable to every kind of rmachine." Lebon, in 1785, thought to distil wood in close vessels, to extract from it, on one side, carbon and acetic acid, and on 118 METALLOIDS. the other, gas fit for illumination. He demonstrated that coal is better than wood for making illuminating gas; but notwithstanding this observation, twenty-five years passed before coal-gas was used in the arts. The first manufactories of coal-gas were established in England by Murdoch. Attempts were made to replace coal by resins, fat oils, oils of shists,* and by mixtures of tar and steam exposed to a red heat on surfaces of coke, &c. None of these processes could sustain a competition with the distillation of coal. Coal is, in fact, abundant at a low price, and gives two products, of which one is the illuminating gas, and the other coke, one of the best combustibles known. Nevertheless, even now, in some gas-works, gas is made from oils or resins, or by passing a mixture of tar and steam, over coke heated to redness (gas of Sellfgue). Gas from oil or resin is brighter than that from coal; it owes this property to the presence of a much larger proportion of bicarburet of hydrogen, and other volatile carburets of hydrogen. Besides, it does not contain sulphuretted hydrogen, sulphuret of carbon, nor ammonia, which are ordinarily found in coal gas badly purified. The distillation of coal furnishes gas, the composition of which varies with the temperature to which the coal has been exposed. At the commencement of the distillation, the gas is very rich in bicarburet of hydrogen, and in consequence, very brilliant; the proportion of this gas diminishes as the operation progresses, and at the end the gas contains a considerable quantity of hydrogen and carbonic oxide, which are but slightly illuminating. The gas always contains com* The bituminous slate-marl of Autun affords, when distilled, about 10 to 20 per cent. of oily products, two-thirds of which consist of a light oil for the production of gas.-Knapp. ILLUMINATING GAS. 119 bustible vapors, which increase its illuminating power. These vapors, which do not condense entirely in the purifying and the conduit-pipes of the gas, are principally formed of different carburets of hydrogen. The sulphur, of which the greater part is contained in coal in the state of pyrites, transforms itself, during the distillation, into sulphydric acid. This acid is readily absorbed by the lime in the purifiers. The purity of the gas is tested with a paper impregnated with acetate of lead, which remains colorless when the gas is pure, but becomes black when the gas contains sulphydric acid. The sulphuret of carbon forms only in small proportion, and without doubt condenses with the liquid products; it is rarely met with in coal-gas. Illumination with coal-gas is at this time an object of great importance. We will briefly point out the details of this manufacture. This gas is produced when bituminous coal is heated in retorts. The gas-works use, in general, the coals called demigrasses in preference to those called grasses, which are difficult to distil, and besides give tar in great abundance. Poor (maigre) coals leave a coke which does not run together well, and give besides but little gas. The coal-basins of Anzin, Douchy, of Commentry and Mons, supply the gas-works of the north of France, as well as those of the upper and lower Seine. The vein of Mons resembles the cannel-coal of Lancashire, which is the most valuable coal in England for the manufacture of gas. The vein of Mons is inferior to cannel-coal, as to the quality and quantity of gas which it produces; but it leaves a coke less burnt, and more merchantable. The basin of St. Etienne furnishes coals of excellent quality for the manufacture of gas, which supply the works of the centre and east of France. 120 MET ALLOIDS. The city of Paris has now eight gas-works, which have been successively established since 1818. They had, at their organization, and also for pipes through the streets, a capital of more than 30 million francs. The pipes are laid through a route of more than 450,000 metres. They are made of cast-iron; for some years past, they have also used pipes of sheet-iron, covered with a thick coating of bitumen. The number of retorts used in the eight gas-works has been put down at a mean of 1200; their capacity varies from one hectolitre to one and a half, and rarely reaches two hectolitres; they are made of cast-iron, or of fire-clay. The earthen retorts present incontestible advantages over the others. They are less costly, more durable, and the coal distilled in them furnishes a superior gas to that made in metallic retorts. The coal intended for the manufacture of gas is first broken up, then introduced into the retorts, which are about half filled. This precaution is necessary to allow of the free development of the coke, the volume of which is about one and a third, or one and a half times that of the coal which produces it. When coal of a good quality is used, the operation generally lasts four hours. It may be estimated that, in a wellconducted distillation, 100 kilogrammes of coal will furnish 25 cubic metres of gas to the gasometer. This quantity of coal is about that ordinarily distilled. A single retort can thus yield, in 24 hours, 150 cubic metres of gas. The retorts are heated and kept to a cherry red, either with coke or tar. The distillation of one hectolitre of coal requires 75 litres of coke. The gaseous products are conducted through large subterranean pipes, communicating at different distances with cisterns in which are condensed ammonia, ammoniacal salts, water, tar, &c. ILLUMINATING GAS. 121 When the condensation is complete, the gas is received in the apparatus intended for its purification. The purifiers, the form of which is variable, are ordinarily cases of iron of two and a half to three cubic metres. The gas runs through three purifiers, and then passes out by a pipe. It is received in the gasometer for distribution to the different parts of the city. The mean pressure to which the apparatus is subjected is 30 lines. The pressure can be increased or diminished by giving more or less opening to the valves placed in the pipes of exit. The dimensions of the gasometers vary with the size of the gas-works. They are, however, seldom required of greater capacity than from 70,000 to 80,000 hectolitres. The quantity of gas consumed in Paris, in 1846, was estimated at 25 millions of cubic metres, which were produced by about 100,000 tons of coal. These 100,000 tons furnish 60,000 to 65,000 tons of coke. About one-third of this quantity was consumed for the distillation of the coal. 40,000 tons of coke were put in the market, and mostly applied to domestic uses. 85,000 gas-burners supply Paris for public and private use. The price per burner is six centimes per hour. Each one consumes, on an average, 120 litres of gas per hour, and makes a light equal to about one and a half times that of a Carcel lamp. The price of the cubic metre of gas sold is 43 centimes (1852). The ordinary burners are pierced with 20 holes of the size of a third of a millimeter; the height of the flame is 8 centimetres; that of the glass chimney ought not to exceed 20 centimeters. The quantity of ammoniacal salts resulting from the condensed waters, may be estimated at more than 100,000 kilogrammes. A more complete purification would increase this quantity still more. It was thought that the tar from gas-works could be applied to the same uses as asphalte and 11 122 M ETALLOIDS. bitumen; but it cannot be prevented softening at a moderate temperature, so that it is used almost exclusively for heating the retorts. The proportion of tar produced by the distillation of coals, varies with their quality; on an average, it amounts to 4 or 5 per cent. of the weight of the coal. Endeavors are made to keep down as much as possible the formation of tar, because it is formed at the expense of the quantity of gas, and its power of illumination. CYANOGEN, PRUSSIC ACID, ACND CYANURETS. The discovery of cyanogen, which is due to Gay-Lussac, is justly considered as one of those which has exercised the most influence on the progress of chemistry. Gay-Lussac proved that nitrogen and carbon could combine together to form a new body which he named cyanogen, and that this compound, in all its relations, acts as a simple body. Cyanogen, like chlorine, forms a hydracid which is cyanhydric, or prussic acid; it combines, also, with oxygen to form acids. It unites with the metals to form cyanides, which may be compared with chlorides. Cyanogen is formed when azotized organic substances are burnt with potash; thus fibrine, gelatine, skins, &c., heated with potash, produce cyanide of potassium. Cyanogen was obtained in a state of purity, by GayLussac, by heating cyanide of mercury, which, under the influence of heat, separates into mercury and cyanogen, which last is disengaged. Cyanogen is gaseous; its odor is penetrating, and irritates the eyes; it is combustible, and burns with a blue, characteristic flame. Combining with hydrogen, it forms cyanhydric acid, often called Prussic acid. Cyanhydric acid is one of the most poisonous bodies known; two or three drops being sufficient to produce death. Cyanogen unites with the metals to form cyanides, many ~ ~!-",-......'.'~.,~ ~!:i.~i!'~i',,.'~,~'~ ~~i'~~ ~ ~ ~ ~ ~i- ~.?.~..~,:!~i~~1ii~~i!,: ~, i,.~!,~ ~?i!? ~~~~~~~~~~~~~~~~~-i ~ ~"~'~~~~~~~ ~i~ ~ ~',~ ~" ~"~' "~"??? "11":':~,.1; ~~~i ~::~~~ ~i~~i-'.. U V liA.i CM j C it C I DO. 123 of which are used in the arts. We name the cyanide of potassium, which is used in electric gilding, to dissolve the cyanides of gold and silver. The ferro-cyanide of potassium, which may be considered a double cyanide of potassium and iron, is used in dying. It precipitates most metals from their solutions, and is used as a reagent in chemical laboratories. Prussian blue, which is cyanide of iron, is prepared by precipitating a salt of the peroxide of iron by the ferro-cyanide of potassium, and is much used as a blue coloring-matter. COMBINATION OF BORON WITIH OXYGEN. Boracic Acid. Boracic acid is a compound of boron and oxygen, which exists in nature. This body presents itself in lamellated, colorless crystals, without smell, of a slightly acid taste; it feebly reddens litmus. Boracic acid is found in Tuscany, in the small lakes, called lagunes. Small craters (soffioni), in the bottom of these lagunes, discharge vapor with boracic acid. This boracic acid dissolves in the waters of the lagunes, and is reduced to a proper degree of concentration by heat; after this, the solution is allowed to cool, and boracic acid crystallizes. The solutions of boracic acid are evaporated by using the heat resulting from the condensation of the vapors of the soffioni; this vapor is conducted through passages of stone, under the evaporating pans. The boracic acid crystallized, and still moist, is first placed in willow baskets, where it drains, and then in brick dryers, likewise heated by the vapors of the soffioni (Fig. 16). Boracic acid is also extracted from borate of soda dissolved in water; this salt is decomposed by a slight excess of sulphuric acid, and the boracic acid which precipitates is then purified by crystallization. Boracic acid, extracted 124 METALLOIDS. from borate of soda by sulphuric acid, generally retains a small quantity of this last acid. To purify it, it is washed with cold water, until the water no longer forms a precipitate with a salt of barytes, mixed with weak nitric acid. Uses. -Boracic acid is used in medicine under the name of sedative salt of Homberg; it is used for preparing borax (borate of soda). It enters into the composition of some glasses, that of paste, and of the glazing of common potteries. COMBINATIONS OF SILICIUM WITH OXYGEN. Silicic Acid, or Silex. This is one of the most wide-spread bodies in nature. It forms part of all primitive rocks, of clay, of soils of different formations, of the gangue of a great number of minerals, and of nearly all precious stones. It is met with in small quantities in the ashes of nearly all vegetables. Some waters contain silex in solution. Anhydrous Silicic Acid. Anhydrous silicic acid is white, tasteless, without smell, infusible in the fire of the forge, but, as shown by M. Gaudin, capable of being melted by the oxy-hydrogen blowpipe, and drawn out in very fine threads. Anhydrous silicic acid, after having been heated to redness, is completely insoluble in water, and the acids. Fluorhydric acid alone attacks it, and transforms it into water and fluoride of silicium. This last property is one of the most characteristic of silex and the silicates. Hydrogen, carbon, phosphorus, chlorine, and the metals, are without action on silex. Certain metals, particularly iron, reduce silex in presence of carbon, and form carbonic oxide, and a metallic siliciuret. Thus, when a mixture of oxide of iron and silex are melted ANHYIDIROUS SILICIC ACID. 125 in a clay (brasqu6) crucible, a residue of iron is obtained, in which the proportion of silicium amounts to 5 or 6 per cent. of the weight of the iron. Silicic acid is a very feeble acid, but, on account of its fixedness, it can expel from their combinations the most energetic acids; it is thus that, by heat, it decomposes the sulphates. When silex in powder is thrown on carbonate of soda, kept in a state of fusion, a brisk effervescence takes place in the mass, and carbonic acid is set free. Potash, soda, and barytes react, when hot, on silex free or combined. The two first bases form with silex, silicates, which are attacked by the acids; this property is made use of to render a great number of mineral substances soluble in the acids. There are found in nature numerous varieties of silex, bearing the names of opal, rock-crystal, agate, flint, millstones, tripoli, sand-stone, &c. Opal is silex, containing 10 per cent. of water. Rockcrystal, so called on account of its limpidity and transparency, is pure anhydrous silex; it crystallizes in hexagonal prisms, terminated by hexagonal pyramids. Rock-crystal, or iyalin quartz, becomes electric by friction, it strikes fire with steel, and is sufficiently hard to scratch glass, and even steel. Quartz is of different colors, When it is colored a clear yellow by the peroxide of iron, it takes the name of false topaz, and that of amethyst, when it is colored violet by the oxide of manganese. Agate quartz is concreted, and often presents strata of different colors, The white agates, or those of a pearly, translucid gray, are called chalcedony; cornelians are of a blood red, and waved, Gun-flints are met with in irregular tuberculous masses, of a conchoidal fracture, and great hardness. They are 11 * 126 METALS. silex, containing two per cent. of water, and one per cent. of alumine. They are almost always covered with a white layer, composed of disaggregated silex and carbonate of lime. Silex enters into the composition of most potteries. Tripoli is earthy silex, in very fine grains, united together by the force of adhesion, assisted by compression. Sandstone is a quartz sand, agglutinated by a calcareous or silicious cement. METALS. GENERALITIES ON THE METALS. THE metals are solid at the ordinary temperature, with the exception of mercury, which is liquid. Most metals possess a characteristic brilliancy, which they lose when brought to a state of minute division. Their powders, which are ordinarily black or gray, become brilliant again when rubbed with a hard body. Metals, taken in mass, are all opaque; but the light passes through them, if they are reduced into leaves of an extreme thinness. It is thus that gold-leaf appears green, when placed between the eye and the light. The ordinary color of metals is a grayish white; gold, however, is yellow, and copper of a peculiar red. They are, in general, without smell; however, tin, copper, iron, and lead, exhale a disagreeable odor, particularly when they are rubbed with the hand. Some metals have a peculiar and disagreeable taste, as iron and tin. The metals, with the exception of sodium and potassium, are heavier than water; hammering ordinarily increases their density. The hardness of metals is variable: some, as lead and tin, are very soft; others, as iron and antimony, are very hard. The presence of small quantities of carbon, arsenic, and phosphorus, increase in general their hardness. Metals PROPERTIES OF METALS. 127 possess the property of ductility, that is, of being elongated into wires by being passed through the drawing-plate. Malleability is the property they possess of being hammered or rolled out into thin sheets or leaves. Metals or their alloys, which have undergone the action of the hammer, the drawing-plate, or the roller, become nearly always hard and brittle; to continue to reduce them into wires and sheets, they must be annealed from time to time, and allowed to cool slowly. Metals have different degrees of ductility and malleability. We will here class the principal metals in the order of their ductility and malleability. ORDER OF DUCTILITY. ORDER OF MALLEABILITY. Gold, Gold, Silver, Silver, Platinum, Copper, Iron, Tin, Copper, Platinum, Zinc, Lead, Tin, Zinc, Lead. Iron. Malleability and ductility are in general increased by heat. Tenacity is the power which prevents their rupture. This property differs in different metals. The tenacity of metals is compared by finding out what weight will break wires of the same diameter, but of different metals. Metallic wires of two millimeters diameter break with the following weights: Iron................ 249'159 kilogrammes. Copper........... 137'399 " Platinum......... 124'000 " Silver........... 85'062 " Gold.............. 68'216 " Tin............ 24'200 Zinc................ 12'710 " 128 MET A LS. When metals are elastic and sonorous, they have these properties developed in proportion as they are harder. This remark appears to apply to the alloys; thus, bronze formed of copper and tin, is harder and more sonorous than either of these metals. Their fracture is to be considered in several different aspects, for it is often a means of distinguishing one from the other. Thus, the fracture is lamellated in bismuth and antimony, granulated in tin, &c. Metals can assume regular crystalline forms, which are in general octahedral, the cube or forms derived from them. Metals conduct heat and electricity the best of all simple bodies. Their fusibility is very variable. Some, like lead and tin, fuse much below a red heat; others, as platinum, rhodium, and iridium, do not melt except with the aid of powerful lenses, or the oxyhydrogen blow-pipe. The following table shows the order of fusibility of the principal metals - Mercury............. - 390 C. Manganese, between iron Potassium............ + 580 C. and cast iron Sodium............... 90~ C. Nickel, do. Tin................. 2300 C. Forged iron........... 2118~C. Bismuth.............. 2460 C. Palladium, Lead................ 312~ C. Molybdenum, Almost infusible, Cadmium............. 3600 C. Uranium, running together Zinc.................. 370~ C. Tungsten, only in the fire of Antimony............. 432~ C. Chromium, a powerful forge. Silver................ 1022~ C. Titanium, J Copper............... 10920 C. Cerium, Infusible in the fire of Gold................. 1102~ C. Osmium, the most powerful Cast iron......... 15870C. Iridium, forge; fusible with the Steel, between iron and Rhodium, oxyhydrogen. blowcast iron Platinum, pipe. CLASSIFICATION OF ME'TALS. 129 The action of Oxygen, of Atmospheric Air, and of Water, on the Metals. Some metals, as potassium and sodium, absorb oxygen at the ordinary temperature; but most metals are not oxidized except at a higher temperature. Others, as gold, platinum, palladium, rhodium, and iridium, do not absorb oxygen at any temperature. Dry air acts on the metals like oxygen, but with less energy; moist air oxidizes them more rapidly than dry air; it then forms oxides which are ordinarily hydrated and carbonated. Several metals decompose water at the ordinary temperature, as potassium and sodium; others, as iron, zinc, tin, antimony, &c., do not act on water, except at a temperature approaching redness. Some metals, as gold and platinum, do not exercise any action on water, even under the influence of a red heat. Acids sometimes bring about the decomposition of water by the metals; the oxygen of the water, in this case, unites with the metal, to form an oxide, which combines with the acid, while hydrogen is disengaged. Some acids, as nitric and concentrated sulphuric, will even give up a part of their oxygen to the metals. CLASSIFICATION OF METALS. The best classification of metals is that proposed by M. Thenard. We shall adopt it with the modifications introduced by M. Regnault, which, however, leave undisturbed the basis of the classification of M. Thenard. The metals are classed in six sections, according to their degree of affinity for oxygen. This affinity is shown, 1st. By the action which oxygen exercises on the metals. I 130 M E T A L S. 2d. By the action of heat on the oxides, and by the more or less easy reduction of these oxides. 3d. By the decomposition of water which the metals give rise to, either directly or in the presence of acids. First Section. - The metals of the first section absorb oxygen at a low temperature; their oxides resist the most elevated temperatures, and are decomposed with difficulty by bodies having a strong affinity for oxygen. They decompose water in the cold, disengaging hydrogen. The metals of this section are: potassium, sodium, lithium, barium, strontium, and calcium. Second Section.-The metals of the second section absorb oxygen at a higher temperature; their oxides are, in general, as difficult to reduce as the preceding. But these metals do not decompose water till between 1000 and 2000C., and sometimes only at a low red-heat. The metals of this section are: gluciniumr, aluminium, magnesium, zirconium, thorium, yttrium, cerium, lanthanium, didymium, manganese, uranium, pelopium, niobium, erbium, and terbium. Third Section.-The metals of this section do not absorb oxygen, but at a moderately-elevated temperature; their oxides, indecomposable by heat, are easily reduced by hydrogen, carbon, and carbonic oxide. These metals do not decompose water, except at a red heat, or at ordinary temperature in the presence of acids. The metals of this section are: iron, nickel, cobalt, zinc, cadmium, chromium, and vanadium. Fourth Section. - The metals of this section are distinguished from those of the preceding, by not decomposing water in the presence of acids, though they do so at a red heat. But as they have a great tendency to acidify, they decompose water in the presence of energetic bases like potash. The metals of this section are: tungsten, molybdenum, osmium, tantalum, titanium, tin, and antimony. Fifth Section. - The metals of the fifth section do not METALLIC OXIDES. 131 decompose watery vapor but slowly, and at an elevated temperature; their oxides are not reduced by heat. These metals are bismuth, lead, and copper. Sixth Section. - This section comprises the metals called noble, which do not decompose water; their oxides are reduced by heat. These metals are: mercury, silver, rhodium, palladium, rhuthenium, platinum, and gold. It may be remarked that metals of the first section form the most energetic bases. Those of the second give less energetic bases, and often acids. In the third are found, among oxides of the same metal, both bases and acids. The fourth section gives principally acids. The metals are sometimes divided into 1st. Alkaline Metals, which are: potassium, sodium, and lithium. 2d. Alkaline-earthy Metals, which are: calcium, barium, and strontium. 3d. Earthy Metals, which are: aluminium, magnesium, glucinium, zirconium, yttrium, erbium, terbium, thorium, pelopium, niobium, cerium, lanthanium, and didymium. 4th. Into Metals properly called, which are: manganese, iron, chromium, zinc, cadmium, cobalt, nickel, tin, titanium, antimony, bismuth, lead, copper, uranium, molybdenum, vanadium, tungsten, tantalum, mercury, silver, gold, platinum, osmium, iridium, rhodium, palladium, and rhuthenium. METALLIC OXIDES. This name is given to the binary compounds formed by the combination of a metal with oxygen. The oxides are divided into four classes, viz: 1st. Basic oxides; 2d. Acid oxides (oxacids, metallic acids); 3d. Indifferent oxides; 4th. Saline oxides. 132 METALS. Basic Oxides belonging to the metals of the first section, have the property of neutralizing acids, of turning green the syrup of violets, of restoring the blue color to litmus reddened by acids, and of changing to a reddish brown the yellow color of turmeric. The Oxacids possess acid properties, neutralize bases, form salts with them, and often redden litmus. The Indifferent Oxides are those which combine neither with acids nor bases. The bioxides of barium, of calcium, mzanganese, strontium, &c., are indifferent oxides. The Saline Oxides are those which result from the combination of two oxides of one metal, one acting as an acid, the other as a base. Action of Heat on the Oxides.-The oxides of the metals of the sixth section lose their oxygen, and are reduced to the metallic state, by the action of heat. None of the other oxides are completely reduced by heat; but certain metallic acids, as chromic acid, ferric acid, manganic and permanganic acids, plumbic acid, some peroxides, as those of manganese and copper, lose a part of their oxygen when they are heated. Moreover, the metallic oxides are nearly all fixed. Most of these melt only at a very high temperature. Action of the Pile.-All the oxides, with the exception of the earthy oxides, may be decomposed by the pile. When an oxide is placed in contact with the two poles of a pile, the metal reduced is soon seen to appear at the negative pole. When the metal will form an amalgam, the decomposition of the oxide is facilitated by using mercury. The oxide, slightly moistened, is formed into a cup, which is filled with mercury; this is placed on a metallic plate, which communicates with the positive pole of the pile, while the negative pole is plunged into the mercury; at the end of a short METALLIC OXIDES. 133 time an amalgam is obtained, which by distillation gives the metal which formed the oxide. Action of Oxygen. —Many oxides absorb oxygen when they are in contact with this gas or with air, either at the common temperature, or at a high temperature. Such are the protoxides of potassium, of sodium, of barium, of iron, of manganese, of tin, of copper, of lead, &c. Action of Jlydrogen.-Hydrogen, under the influence of heat, reduces the oxides of the four last sections; excepting, however, the oxide of manganese, and the oxide of chromium. IIydrogen brings to the state of protoxide, the peroxides of the two first sections, as well as the peroxide of manganese. Certain oxides, particularly those of the last section, are reduced by hydrogen at a somewhat elevated temperature. The oxides reduced by hydrogen always leave the metal pure; it is thus metals are often prepared in laboratories. Action of Carbon.-Carbon reduces the metallic oxides at a more or less elevated temperature, excepting the earthy oxides, or those of the second section, and the alkalineearthy oxides. Carbon, in its action on the oxides, produces either carbonic acid, or carbonic oxide, according to the proportion of carbon used, and the affinity of the metal for oxygen. If the oxide is easily reduced, like the oxides of copper and silver, carbonic acid is always obtained. If the reduction does not take place except at an elevated temperature, and if the carbon is in excess, carbonic oxide is produced. When the reduction is made at a temperature approaching redness, carbonic acid and carbonic oxide are both produced. Carbon is used in metallurgic operations, to extract metals from their oxides. The carbon produces, in burning, the heat necessary for the reduction and at the same time takes away the oxygen of the oxide, which it transforms 12 134 M ETALS. into carbonic oxide, or carbonic acid. Metals extracted from their oxides by carbon, ordinarily retain a small quantity of carbon. Iron, for example, obtained in blastfurnaces, may contain from two to six per cent. of carbon. Manganese and chromium, reduced by carbon in a clay crucible, also retain carbon. SALTS. Before Lavoisier, the name salt was applied indiscriminately to a certain number of bodies, whose composition and properties often did not present any analogy. It was enough that a body was solid, crystallizable, transparent, and soluble in water, to give it the name of salt. Lavoisier was the first to fix the true nature of salts, and gave the following definition of them: A salt is a body formed by the combination of an acid with a base, in which the properties both of the acid and the base are found to be more or less neutralized. At the time Lavoisier proposed this definition of salt, the hydracids were not known. It was thought that a salt resulted necessarily from the combination of a base with an oxacid, and ought to contain the elements of the acid and the base. Later, the existence of a new class of acids was demonstrated, the hydracids, which in uniting with the bases, form water and binary compounds.. Chemists then found themselves placed in the alternative of abandoning the definition given by Lavoisier, or of rejecting from the class of salts, bodies which, like marine salt, have all their general properties, but differ in composition. Berzelius proposed the name of haloid salts, to binary compounds resulting from the reaction of hydracids on bases. The chlorides, bromides, iodides, fluorides, cyanides, sulphurets, were considered by Berzelius as haloid salts. The salts formed by the SALTS. 135 oxacids, and which are called oxysalts, may unite together to form double salts; it is thus the sulphate of potassa combines with the sulphate of alumine, and constitutes alum. Phenomena of Saturation. When a base is made to act gradually on an acid, the properties of the acid and the base are seen gradually to disappear. There arrives a time when these two bodies have lost their characteristic taste, their action on turnsol (litmus), &c. The acid is then said to be saturated by the base. At first, the term neutral salt was given to the saline compounds in which the respective properties of acid and base were neutralized. But the expression neutral salt afterwards had another signification. The moment when the neutrality is perfect is recognised by the aid of coloring substances, which readily become modified under the influence of acids or bases. Thus the tincture of turnsol, syrup of violets, the solution of the coloring matter of Campeachy wood, turmeric, &e., may be used when neutralizing a solution,.by seizing upon the moment when the acid and the base shall have ceased to act on these colored reagents. As the tincture of turnsol is the colored reagent most frequently used to test the presence of acids and bases, it is necessary to know its composition. We borrow from M. Chevreul the following observations on this reagent. Blue turnsol ought to be regarded as a true salt, resulting fiom the combination of a base with an acid which is red. An acid reddens turnsol, because it isolates the red acid which exists in the tincture of turnsol. The sulphate of potash does not react upon turnsol, because sulphuric acid and potash have so strong a mutual affinity, that the coloring principles of the tincture of turnsol 136 METALS. cannot unite to the base or the acid, so as to form combinations of another color than that of the coloring principles in a state of purity. If there existed a coloring matter sufficiently energetic to remove the potash from the sulphuric acid, the sulphate of potash, in presence of this coloring matter, would have an acid reaction. General Properties of Salts. Salts are nearly all solid. Their color is variable, and depends, in general, on the nature of the base they contain. The alkalies, the earthy oxides, and some metallic oxides, form colorless salts, when the acids with which they are united are themselves colorless. Most of the metallic oxides, as those of copper, iron, cobalt, nickel, chromium, gold, platinum, &c., give colored salts. When the acid which enters into the composition of the salt is colored, as chromic and permanganic acid, the salt has a color, in general, like that of the acid. The taste of salts is often characteristic, and depends nearly always on the base. Thus the salts of magnesia are bitter; those of alumine, sweet and astringent; those of lead, sweet and styptic. It may happen, however, that the taste of the salts partakes of the nature of the acid. The sulphites, sulphurets, and cyanides, have a taste and properties affecting the senses and organs generally, which depend especially on the nature of the acids with which they were formed. Certain acids may modify, or even change completely, the taste of a base; thus, the citrate of magnesia has not the ordinary taste of the salts of magnesia. Action of Heat. Heat produces effects on salts which vary with the nature of the acid and that of the base. When a salt contains much water of crystallization, it readily fuses without losing this water, and thus presents the phenomenon of aqueous SALTS. 137 fusion. In continuing to heat it, the water of crystallization volatilizes, the salt returns to the solid state, and may enter into fusion a second time; it then undergoes igneous fuision. Some salts, submitted to the action of heat, give rise to a peculiar noise, which is called decrepitation. When sea-salt is thrown on live coals, it is thrown about on all sides, producing a series of slight detonations. For a long time, the decrepitation was attributed to the sudden expansion of the water contained in the crystals; but it is now demonstrated that the volatilization of the water is not the only cause of the decrepitation; for some salts decrepitate by heat, even after they have been dried a long time in vacuo, and thus lost the small quantity of water contained between their molecules. The decrepitation ought then to be attributed to an unequal distribution of heat between the molecules of the salt, which causes the rupture of the crystals. Some salts are rendered phosphorescent by heat, as the fluoride of calcium, some sulphurets, &c. Some salts and certain oxides, throw out a bright light when their temperature is gradually raised; they then display new properties, and are in general more difficult of solution. Action of Electricity on Salts. All salts are decomposed by the pile, when they are moist or dissolved. The acid goes to the positive pole, the base to the negative. It often happens that the base is decomposed, and the reduced metal goes to the negative pole, while the acid and the oxygen of the base go to the positive pole. Action of Metals on Saline Solutions. When we place in a saline solution, containing one of the metals of the four last sections, another metal, belonging 12* 138 METALS. also to one of these sections and having a greater affinity for oxygen than that which is in the solution, this metal, in general, substitutes itself for that of the salt, and precipitates it. Generally the precipitated metal attaches itself to the precipitating metal, with which it forms an element of the pile which causes the complete decomposition of the salt. The metal, in slowly depositing, sometimes assumes beautiful crystalline forms. The most remarkable crystallization, is that called arbor Saturni, which is made by placing a plate of zinc in a solution of the acetate of lead. This is done by pouring into a wide-necked bottle, water containing the 30th part of its weight of acetate of lead, first rendered acid by acetic acid. A piece of zinc, attached to a stopper by brass or copper wire, is introduced into the bottle. Shortly the zinc and the wires are coated with brilliant and long crystals of lead. The name of arbor _Diane is given to the crystallization which is obtained by the precipitation of nitrate of silver by mercury; the body which crystallizes is an amalgam of silver. Iygrometric action of the Air on Salts. — Those salts which attract moisture from the air, when in contact with it, and become liquid from being dissolved in the water which they absorb, are called deliquescent salts. All very soluble salts are deliquescent in an atmosphere saturated with moisture. There are, on the contrary, salts which give up to the air, either in whole or in part, their water of crystallization. These are called efflorescent salts. It may happen, however, that certain anhydrous and fused salts become efflorescent by absorbing water from the air. Thus the sulphate of soda, fused, absorbs water in contact with the air, and crumbles into powder; it remains in this form, because the hydrated sulphate of soda is not SALTS. 189 deliquescent. Some salts, as the sulphate of soda, lose in dry air all their water of crystallization, while others, as the carbonate of soda, always retain,a certain quantity, no matter how dry the air may be. Action of Water on Salts. — The solubility of salts in water is'very variable. Some salts, as the sulphate of barytes, the phosphate of lime, &c., are insoluble; other salts often require less than their weight of water to dissolve them. Anhydrous salts, in forming solid hydrates with water, in general develop heat, when placed in contact with this liquid. Salts which do not combine with water, or those which contain all their water of crystallization, produce, on the contrary, cold, when passing from the solid to the liquid state in contact with water; such are the chloride of potassium, nitrate of ammonia, and sulphate of magnesia. An absorption of heat takes place from the same cause in the first case; the heat developed by the combination is, nevertheless, often sufficiently intense to raise the temperature. The cold produced is great in proportion to the rapidity of the solution; so that dilute acids often are used instead of water, as they dissolve hydrated salts more quickly. A most intense cold is obtained by mixing hydrated salts with pulverized ice, or, better, with snow. This easily explains why ice or snow, in melting, absorb a considerable quantity of heat. All these mixtures are called freezing mixtures. Wlater is said to be saturated with a salt at a given temperature, when it is no longer able to dissolve the smallest quantity of the salt at this same temperature. A mother-water, which has deposited crystals on cooling, or even a solution agitated for a long time with an excess of the salt in powder, ought to be considered as saturated. A saturated solution of a salt may dissolve a new salt. 140 M E T A L S. Water saturated with nitrate of potash can still dissolve a considerable quantity of sea-salt, and even a certain quantity of a third, or even a fourth salt, provided the mutual action of these different salts does not produce other compounds which would precipitate. A saturated solution of a salt sometimes deposits some of this salt, when it dissolves a new salt. It is thus that water, charged with nitre, precipitates a part of this salt, when it is agitated with chloride of potassium. Many industrial operations, and some analytical processes, are founded on the property which water possesses, when saturated with a salt, to dissolve several other salts. Variations of temperature modify the solvent power of water. This generally increases with the temperature. There are, however, certain salts which are more soluble in cold than warm water. The air in general does not exercise any influence on the crystallization of salts. The sulphate and seleniate of soda do not crystallize when they are preserved from contact with the air; a bubble of air is sufficient to cause the sudden crystallization of these two salts. Certain saline solutions have also the singular property of remaining surcharged with an excess of salt for some time; such are the solutions of the nitrate of silver, and acetate of lead. When they are agitated, or a solid body is introduced, the solution sometimes thickens into a mass. The cause of this supersaturation is yet unexplained. Whatever process is used to crystallize saline solutions, the crystals which form retain some water. When the water is combined with the salt in definite proportions, it is called water of crystallization, or of combination. If the quantity of water retained by the crystals is small, and is not found in simple relation with that which enters into the composition of the salt, the name of water of interposition is given to it. POTASSIUM. 141 Exposure to the air, or a few instants in vacuo, or even pressure between the folds of unsized paper, will remove the water of interposition, which is not an integral part of the salt, and exists in it only in variable and very feeble proportions. Phenomena w/hich determine the Composition of Salts. We are indebted to Berthollet for valuable observations on the laws which determine the decomposition of salts. This chemist has proved that, when a salt is placed in presence of an acid, a base, or even another salt, a decomposition always takes place if, by the reaction either of the acid or the base, or the salt, on the elements of the first salt, new compounds would be formed, more volatile, less soluble, or more fusible, than those which were placed together. Thus, by consulting the physical properties of bodies which ought to form in the preceding mixtures, it will always be easy to foretell, in advance, the nature of the bodies which will be formed. When sulphuric acid is poured on marble (carbonate of lime), we might, relying on the laws of Berthollet, say in advance that there will be effervescence; that the sulphuric acid will expel the carbonic acid from its combination with lime, because carbonic acid is less fixed than sulphuric acid. What we have said of the influence of volatility is applicable to the influence of insolubility, or to that of fusibility. We shall now examine the principal metals and their compounds. POTASSIUM. Potassium was first isolated in 1807, by Humphrey Davy. This discovery, one of the most important which has been made in chemistry, has fixed the true nature of the alkalies and earths. The properties of potassium were then studied with the 142 METAL S. greatest care by Gay-Lussac and Thenard, who were the first to explain a practical process for preparing this metal..Properties.-Potassium is solid at the ordinary temperature, and possesses a metallic brilliancy. Melted in the oil of naptha, it is as white as silver; but when exposed to the air, it tarnishes rapidly, and becomes of a bluish-gray color. It is softer than wax, and may be moulded between the fingers. This experiment ought to be made under naptha, because potassium takes fire in the air, even at ordinary temperatures. The density of potassium, at 15~C., is 0O865; it is then lower than that of water. Potassium is, next to mercury, the most fusible of all metals; it fuses at 580 C., and is volatile at a red heat. It may be volatilized in a glass tube by means of a spirit-lamp; its vapor is green. This experiment ought to be made in an atmosphere of nitrogen, in order to prevent the oxidation of the metal. Potassium has a great affinity for oxygen. When its temperature is raised, when it is touched, for example, with a rod of red-hot iron, it burns with vivacity, and is transformed into oxide of potassium (potash). Potassium may be kept in perfectly dry oxygen, or atmospheric air, without change. It decomposes water at the common temperature, and deprives it of its oxygen. Thus, when a globule of potassium is thrown into a jar filled with water, it is seen to dance about rapidly, and become incandescent; it then combines with the oxygen of the water to form potash, which remains in solution while the hydrogen of the water is set free. The reaction of the potassium on the water developing a very high temperature, the hydrogen takes fire in contact with the air, and reproduces water. To show lke production of hydrogen in the preceding experiment, a small quantity of water is introduced into a tube filled with mercury, and a globule of potassium is made to pass up to it. As soon as the metal comes in contact 'Qi.~~~~~~~' ~~~~~~~~~:. i~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~ I.'i. POTAS S IU. 143 with the water, the reaction takes place; the hydrogen, in disengaging, depresses the column of mercury contained in the tube, and in a few moments the tube is filled with hydrogen. Preparation.-Davy isolated potassium by subjecting the hydrate of potassa to the action of a strong pile. He made a cavity in a piece of the hydrate, and filled it with mercury; he placed this upon a metallic plate which he made to communicate with the positive pole of a pile of 150 pairs, while the negative pole was plunged into the mercury. The hydrated potassa was decomposed under the influence of the electric current; the oxygen of the oxide, and the oxygen of the water, went to the positive pole, while the potassium and the hydrogen went to the negative pole. The potassium, finding the mercury at the negative pole, forms with it an amalgam; submitting this amalgam to distillation in a small glass retort, the mercury volatilizes, and the potassium remains in the retort in a state of purity. As this experiment only gives small quantities of potassium, this metal is always prepared by reducing the hydrate of potassa by iron, or by decomposing the carbonate of potassa by carbon. We shall first speak of the mode of preparing potassium by means of the hydrate of potassa and iron, which we owe to Gay-Lussac and Thenard. In this method, the hydrate of potassa, in fusion, is brought to react on the turnings of iron heated to redness, placed in a curved gun-barrel (Fig. 18). The hydrogen, proceeding from the decomposition of the water of the hydrate of potassa, is disengaged; the iron absorbs the oxygen of the water, and of the oxide of potassium, while the potassium set free volatilizes, and condenses in the receiver. The potassium should be taken from the receiver with an iron rod, and placed under liquid carburet of hydrogen, 144 METALS. which preserves it from oxidation. It is generally preserved in the rectified oil of petroleum. The process of M. Brunner consists in decomposing, in an iron vessel, the carbonate of potassa, by carbon, which completely reduces the potassa at a very high temperature, and transforms the carbonic acid of the carbonate into carbonic oxide. The potassium distils and condenses in a cool receiver containing the oil of naptha, Fig. 17. The retort is covered with a refractory lute; it is placed over two horizontal bars of iron, in a blast-furnace, furnished with a chimney having a strong draught; it is filled from above, first with charcoal, afterward with a mixture of charcoal and coke. The operation is commenced by heating strongly the iron retort, and the receiver should not be adapted until the vapors of potassium begin to disengage. Often the oil of naphtha, in the receiver, is fired, to absorb the oxygen of the air which it contains, and avoid the oxidation of the potassium. The potassium thus obtained is not pure, it always contains carbon. To purify it, it is first filtered in a cloth, under oil of naphtha heated; then it is distilled in an iron vessel, or a refractory glass retort, covered with an earthy lute. The vapors are condensed in the oil of naphtha. This operation lasts about three hours, and gives 30 to 40 grammes of potassium; it is more easy of execution than the preceding, but it gives a metal less pure. We shall speak now of the principal compounds of potassium, commencing with the protoxide of potassium, which is called potash. Potash. Hydrated potash is solid, white, caustic, very alkaline, and unctuous to the touch; it instantly alters the skin; in its contact with organic substances, it develops a peculiar odor, which is that of lye. 2. PO T AS II. 145 Potash acts on nearly all organic substances, and destroys them rapidly. It decomposes or dissolves a great number of animal substances, such as the skin, silk, &c.; it saponifies fat bodies. Potash fuses below a red heat, and then volatilizes, producing white vapors. It has a great affinity for water; exposed to the air, it attracts moisture, and falls into deliquescence. Potash dissolves alumine and silex, and attacks glass and porcelain; so that it can never be concentrated in vessels either of glass or porcelain. This operation ought to be done in capsules of silver. Preparation. -Potash is extracted from its carbonate: this salt may be obtained by different methods. The ashes of vegetables contain different salts of potash, and chiefly the carbonate. By treating the ashes with water, a great part of the potash is taken up, combined with carbonic, sulphuric, chlorohydric, and silicic acids, and the insoluble residue is thrown upon a cloth. It is principally composed of silex, and phosphate and carbonate of lime. (Fig. 20.) The liquors evaporated to dryness give a residue called calcined potash, or saline. To extract from this the carbonate of potash, the mass is dissolved in boiling water, and then evaporated to crystallization. The foreign salts are deposited, while the carbonate of potash remains in the mother-waters. The carbonate of potash having been once obtained, the potash is extracted from it by submitting it to the action of the hydrate of lime, which, according to the rules established by Berthollet, decompose the carbonate of potash, because the carbonate of lime is insoluble. The carbonate of lime is then formed, and the potash remains free. In order to decompose the carbonate of potash by lime, dissolve one part of carbonate of potash in ten or twelve parts of water; put this solution in an iron vessel, carry it to ebullition, and add to it a boiling liquor of hydrate of 13 K 146 METALS. lime. The carbonate of potash requires for its decomposition about its own weight of lime. The lime ought to be added so gradually as not to stop the ebullition, to facilitate the deposit of the carbonate of lime. To be sure that the carbonate of potash has been completely decomposed, a small quantity of the liquid is taken out and allowed to stand; to this nitric acid is added. If all the carbonate of potash has been decomposed, there will be no effervescence, and it ought no longer to precipitate lime-water. As soon as the decomposition of the carbonate is complete, the liquor is allowed to cool protected from the air; and when clear, it is decanted, and evaporated in a capsule of iron, or, better, of silver. When the hydrate of potash is melted, it is poured on plates of iron, or into cylindrical moulds, made in two parts, so that they can be separated one from the other, and the potash, when it has become solid, can be taken out. If the carbonate used in this preparation is pure, the potash which is extracted is equally so; at least it will only contain a small quantity of its carbonate, produced during the evaporation, at the expense of the carbonic acid of the air. If, however, the carbonate of potash of commerce is used, the potash always contains chlorides, sulphates, and carbonates. To purify it, alcohol is used; which only dissolves the hydrated potash, and precipitates the foreign salts. With this end, the impure potash is evaporated to the consistence of honey; to this is added a quantity of alcohol at 33~*, which represents about one-third of the original weight of the potash; the mixture is stirred and boiled for some minutes, and put into a bottle with a ground-stopper. The liquor, left to repose, divides into three layers: the lower of which is formed of sulphate of potash and anhy* See article Alcohol, in the Organic Chemistry. NITRATE OF POTA SH. 147 drous lime; above is a solution of sulphate, carbonate and chloride of potash; the superior is an alcoholic solution of potash. This last is decanted and distilled, so as to drive off about two-thirds of the alcohol it contains; it is then rapidly evaporated in a silver capsule. The potash thus prepared is called potash a' l'alco]ol; it is nearly pure, and ordinarily contains only traces of chloride. Uses of Potash. —Iydrated potash is a valuable reagent; it is used in the preparation of a great many oxides, and also to attack, by the dry method, the silicates, and to render them soluble in the acids. It is used in medicine as a caustic, from which it has received the name of caustic potash. It is also used in the manufacture of soaps, glass, &c. Natural State of Potash. - Potash exists largely in nature; it is always combined with acids; it is met with in nearly all rocks, and principally in feldspar. It is often found in considerable quantity in arable land and in clay; it saturates, in part, vegetable acids, and forms different organic salts, which, by calcination, produce carbonate of potash, which is found in their ashes. ZNitrate of Potash. Nitrate of potash, also known by the names of nitre, salt of nitre, saltpetre, and nitrate of potash, is white, without odor, of a taste at first pleasant, but soon sharp and bitter; its crystals are very friable; when they are kept some time in the hand, they break, giving rise to a slight noise. Nitre is unalterable in the atmosphere in ordinary circumstances; it does not deliquesce, except in an atmosphere almost saturated with moisture. It fuses at about 300~, and in cooling leaves a white compact mass, known by the name of crystal mineral. This mass pulverizes more readily than the crystals of nitre, which always have a certain degree of elasticity. 148 METALS. The nitrate of potash is hardly soluble in alcohol of 900 of the alcoholometer, and is altogether insoluble in absolute alcohol; it dissolves freely in water. The solubility of nitre, which increases considerably with the temperature, permits it to be purified with great facility, and to be freed, by crystallization, from the foreign salts which it may contain. A mixture of nitre and carbon burns with vivacity, when it is heated, or when it is touched with an incandescent body: the carbon changes into carbonic acid, at the expense of the oxygen of the nitric acid; a part of this carbonic acid is disengaged; another part remains united with the potash, and nitrogen is liberated. Sulphur reacts on nitrate of potash, under the influence of heat. This reaction always gives rise to a lively disengagement of heat. If the nitrate of potash is in excess, the sulphur is entirely changed into sulphuric acid, which remains combined with the potash. The name of detonating powder is given to a mixture of 3 parts of nitre, 2 of potash, and 1 of sulphur. When a few grammes of this powder are slowly heated in a spoon, the mass at first fuses, and then detonates with violence; the detonation is due to an instantaneous disengagement of gas. The flux of Baurmn is a mixture of 3 parts of nitre, 1 part of sulphur, and 1 part of saw-dust; it has the property of determining the fusion of different metals. It acts thus, not only on account of the elevated temperature which is produced by the reciprocal action of the bodies which it contains, but also because a part of the sulphur unites directly with the metals, and forms with them fusible sulphurets. Acids more fixed than nitric acid decompose nitre, under the influence of heat. Clay itself may bring about this decomposition. For a long time nitric acid was prepared by decomposing nitre with clay. NITRATE OF POTASH. 149 Natural State of Nitre. - Nitre abounds in nature; it is principally found in Egypt, in India, in America, and in Spain. In these countries, the nitre effloresces on the surface of the soil; it is collected with long brooms made of twigs, which are called houssoirs; hence the name saltpetre de hoztssage (swept saltpetre). In France there are many places which produce saltpetre; thus in the departments of Seine and Oise, and in la RocheGuyon, these efflorescences are found, which are very rich in saltpetre. This salt crystallizes on the walls of old buildings, in stables, and in places inhabited by animals; plaster from old buildings also contains saltpetre. Though France does not possess deposits rich in saltpetre, it can nevertheless furnish it in great quantities; thus, at the time of the most active wars, the annual production of saltpetre in France has been 1,900,000 kilog. Paris furnished seven-twentieths of it; Tourraine, two-twentieths; all the other provinces, ten-twentieths; the artificial nitreries, one-twentieth. In France they always extract nitre from the saltpetre materials, which, besides the nitre, contain nitrates of lime and magnesia. According to Gay-Lussac, the plaster saltpetres of Paris contain about 5 per cent. of nitrates. The soluble part of the saltpetre materials has in general the following composition: — Nitre,.. 25 Nitrate of lime,.. 3 Nitrate of magnesia,. 5 Sea-salt,. * 5 Other salts,.. 32 100 M. Kuhlmann has also found a notable quantity of nitrate of ammonia in them. 13* 150 METALS. The operation which transforms the nitrates of lime and magnesia into nitrate of potash, is called saturation of the liquors. The waters in which the saltpetre materials have been washed, containing the nitrates of lime and magnesia, are mixed with carbonate of potash, which forms insoluble carbonates of lime and magnesia, and soluble nitrate of potash. In this reaction, the nitrate of ammonia is also decomposed into nitrate of potash and carbonate of ammonia, which disengages during the concentration of the liquors. Washing the Saltpetre Materials.-In order readily to wash the saltpetre materials (old plaster and mortar), they begin by breaking them up, and placing them on hurdles or sieves, of willow; they are then mixed with ashes, or any other substance containing the carbonate of potash intended for the saturation of the liquors. The materials mixed and broken up, are placed in casks having one head out, and placed over a trough called a recipient. A quantity of water, equal in volume to one half of the solid matter, is poured in; it remains there ten hours; after this time, it is allowed to run off; the plaster retains the half. Suppose that 500 litres of water are poured over one cubic metre of plaster containing 40 kilogrammes of saltpetre; then 250 litres of water are drawn off, containing 20 kilogrammes of saltpetre. To draw off the remaining 20 kilogrammes, 250 litres of fresh water are poured in; and 250 litres run off, containing 10 kilogrammes of saltpetre. In continuing to add successively 250 litres of fresh water, each operation will give solutions less and less strong, but will reduce to one-half, the quantity of saltpetre retained by the solid materials. When this quantity becomes so small as not to cover the expense of the evaporation of water used the solution, the operation is stopped. a waters from the washing are not sufficiently rich in to be subjected to evaporation; they are made to NITRATE OF POTASII. 151 pass over new materials, following the same course as before. The waters are not evaporated unless they contain 14 kilogrammes of saltpetre per hectolitre. These last waters ought to mark about 12~ on the areometer; they are then evaporated. Boiling.-The evaporation of the saltpetre waters, called boiling, is made in large boilers of iron or copper. The liquors, in concentrating, deposit carbonate of lime, sulphate of lime, and animal matters; these deposits are called boues (impurities). During this concentration, there is a strong ammoniacal odor. Boues, or Impurities.-Boiling brings the impurities from the circumference of the boiler to the centre; so that by placing in it, at some distance from the bottom, a vessel into which the insoluble deposits will settle, they can be easily taken away. When the liquor arrives at a certain point of concentration, it drops crystals of chloride of potassium and sodium, which are not near so soluble as nitre in boiling water. These salts are taken off with skimmers. The deposit of chlorides generally takes place at the time the liquor marks 420 on the test liquor; the evaporation ought to be pushed to about 450. At this time, the evaporation is stopped; and it is known that it has been carried sufficiently far, if a drop of the solution solidifies when placed on a cold body. The liquor is then carried to the crystallizer, and as it crystallizes, it is agitated with wooden rakes. Thus is obtained a crude saltpetre, crystallized in small needles, which bear the name of saltpetre of the first boiling. This first operation is performed in France by the licensed manufacturers, who have not the right to refine saltpetre. Crude saltpetre is delivered in this state to the governmentshops; it contains about 25 per cent. of foreign matter. It is purified from this by the process of refining. Refining Saltpetre.-To refine saltpetre, the crude salt is 152 METALS. dissolved in its own weight of warm water. The liquor is clarified with bullock's blood, which forms a scum that is taken off with care, and then it is crystallized. This operation gives saltpetre of the second boiling. This is, however, not yet sufficiently pure for the manufacture of gunpowder. It ought to be subjected to a third crystallization, which gives saltpetre of the third boiling, the only one which can be employed to make powder. In the crystallization of saltpetre, the solution should be continually agitated till the moment of crystallization, in order that the salt may deposit in small crystals, which are purified more easily than large ones. To clear the crystals of saltpetre from the mother-water, charged with foreign salts, with which they are impregnated, they are treated with a water saturated with nitrate of potash, which deposites the chlorides and sulphates, &c., and leaves the nitre pure. Nitre is chiefly used for the preparation of gunpowder. We shall treat with some minuteness of this important fabrication. GUNPOWDER. Powder is an intimate mixture of nitre, sulphur, and carbon. It is distinguished into three principal kinds, gunpowder, sporting or hunting powder, and mining or blasting powder. Sporting powder, made in the government works, is composed as follows:Nitre,....78 Carbon,... 12 Sulphur,.. 10 Mining powder: - Nitre,. 62 Carbon,... 18 Sulphur,.. 20 GUNPOWDER. 153 The composition of gunpowder for war is nearly the same in all countries. It may be conceived that there must be but little difference in the composition of a substance which, like powder, has to fulfil certain invariable conditions. The following table shows the composition of the gunpowder used for war by different countries. Nitre. Charcoal. Sulphur. France, ) Prussia, 75'00. 1250. 1250 United States of America, England.............. 75'00..... 1500..... 1000 Russia....;.. 73'78..... 1359..... 12-63 Austria..... 7600..... 11'50..... 12-50 Spain................... 76-47..... 10'78..... 12'75 Switzerland (round powder) 76-00..... 14'00..... 10'00 Holland 7000.....00.......1600 14'00 Sweden.7500..... 9'00..... 1600 China................... 75'00..... 1440..... 9-60 Properties of Powder. -It is well known that powder easily takes fire under the influence of heat, and that it almost instantaneously develops a considerable volume of gas, which then acts as an energetic power. It is not only the proportions of bodies used in the composition of gunpowder which influence the volume of gas developed at the moment of combustion; the physical condition of powder also exercises a great influence on its projectile effects; it is necessary then carefully to point out the physical properties of powder. Powder ought to be so hard that the rubbing together and jarring, which the grains are subjected to during its transportation and in the manufacture of munitions, will not produlce dust in such quantities as would be hurtful to the rapidity of the inflammation of the powder. A good powder ought to be so hard that the grain will not break down under the fingers, nor soil the back of the hand. Before the grains have been glazed, the slightest rubbing 154 METALS. suffices to detach a considerable quantity of dust, which diminishes greatly the projectile power of the powder. The glazing takes off the asperities from the grains, and gives them a lustre and polish by hardening their surface: this however should not be too prolonged, because the grain would be less easy to inflame, and would not be homogeneous, especially if the glazing is done on a very moist grain. The graining will cause the qualities of powder to vary. In comparing grained powder with powder in mass, it will be seen that the first inflames almost instantly, because the flame penetrates the interstices between the grains, while the powder in mass burns slowly, and in layers. Powder in mass, fired in a gun, would hang fire. The size of grain ought to be adapted to the nature of the arms; in small arms, powder of fine grain is always used. The size of the grain varies with the kind of powder, and mode of manufacture. Density exercises great influence on the quality of powder. Dense powders inflame less easily than when they are light and porous, and give less waste in their transport. Powders which give off their gas too rapidly are called brisantes (bursting). They expend a great part of their effect on the sides of the gun, which they may burst without their projectile force being increased in any considerable proportion. In this respect, they have an analogy with fulminating powders. The causes which contribute to render powders too explosive, are the employment of too inflammable charcoal, the lightness of the grain, or want of pressure in the intimate mixture of the nitre, sulphur, and charcoal. A hard charcoal makes powder not sufficiently inflammable. To make a good article, there should be observed a certain relation between the condition of the charcoal, the density of the powder, and the size of the grain. GUNPOWDER. 155 The same mixture of nitre, sulphur, and charcoal, produces a powder of good quality, or one too explosive, according to the size, form, and density of the grain. It is by varying these different conditions, that powder of such a quality as may be required is obtained. ~ The best powder is that which is completely on fire before the projectile has left the gun, and of which the combustion takes place successively as the projectile is displaced. Powder detonates when it is inflamed by an electric spark, by a violent shock, by the contact of an ignited body, or by a heat of about 300~ C. The flame of alcohol, or that of hydrogen, is not always sufficient to inflame powder. Good powder ought to inflame rapidly, without leaving an appreciable residue, on a sheet of paper, which it ought not to set on fire. It was for a long time thought that iron was the only substance which, when struck against a very hard body, would inflame powder; but experience has shown that blows of copper against copper, and of iron against marble, will also cause it to detonate. The shock of a leaden ball, projected from a gun, inflames powder placed against lead, or even against wood. Different causes favor or retard the inflammation of powder. Moist powders always burn more slowly than those that are dry; angular powders inflame more rapidly than those which are round; unglazed powders are more inflammable than glazed powders. Powder does not inflame unless brought suddenly to a red heat; if it is subjected to the action of heat whose intensity increases progressively, the sulphur which it contains fuses and separates from the mass. The sulphur may even be distilled by heating powder in vacuo. From the fact of charcoal, on account of its porosity, having the property of absorbing moisture, powder, whatever be its quality, cannot be preserved perfectly dry even 156 METALS. in the best magazines. Water, penetrating the grain of the powder, causes its projectile effects to vary, also causes an efflorescence of the saltpetre on the grain, and destroys the intimate mixture and aggregation of its components. Powders made with pure saltpetre absorb moisture with a rapidity and quantity proportioned to the quantity of charcoal and dust they contain. The absorption of moisture is also greater in powders made with red coal than with black. In general, fine powder absorbs moisture more rapidly than when coarse. The solid residue from the combustion of powder forms in the gun a coating, which increases in thickness with each discharge. This is a very great inconvenience, especially for sportingpowder, or that for musketry, and does not permit of a brisk fire being kept up for a long time. The quantity of dirt or coating which powder produces, depends on the purity of its ingredients and their proportions; an excess of sulphur, or an incomplete trituration, are causes of dirt; the degree of dryness, and the rapidity of the combustion, also exercise a great influence. Powders with very large grains, or those which are moist, leave a good deal of crust. The products of the detonation of powder are solid and gaseous. The solid products are mostly carried off, or even volatilized, by the high temperature which results from the detonation of the powder. It has been attempted to ascertain the nature of the gases proceeding from the combustion of powder, by inflaming some in a small cartouch-box of copper, and then emptying it under a bell-glass filled with mercury; a gaseous mixture is thus obtained, formed principally of carbonic acid, nitrogen, and carbonic oxide. We should, however, GUNPOWDER. 157 observe that the gases which are thus produced cannot be compared to those produced in guns, because the conditions of combustion are not the same. Gay-Lussac analyzed the gases from powder, by dropping it, grain by grain, into an incandescent tube; he found that these gases were formed of 53 parts of carbonic acid, 5 of carbonic oxide, and 42 of nitrogen. The quantity of gas which powder produces in burning is subject to variations, the cause of which has not yet been explained. A litre of powder, weighing 900 grammes, gave to GayLussac 450 litres of gas, supposed at a pressure of 07T6m., and at the temperature of zero. The moment the explosion takes place, these gases are at a very elevated temperature, which dilates them considerably; it may be said that 1 volume of powder gives, in burning, at least 2000 volumes of gas. The temperature produced at the moment of explosion is very high; it may be estimated at more than 1200~C. It is sufficient to melt gold, pieces of money, and copper; it does not fuse platinum. In order that this elevation of temperature should have full effect, it is necessary that the combustion of the powder should take place rapidly, so that the heat may act on the gaseous mixture, dilate it, and thus augment its elastic force. We will now point out the conditions which the different substances entering into the composition of powder ought to present. Nitre. This body is always delivered to the French powder-works, by the government refineries, in a state of almost absolute purity; it ought never to contain more than three onethousandths of foreign bodies. 14 158 METALS. In France the nitre is used in small crystals; in England, fused nitre is preferred; this is more easily pulverized, and keeps better. Sulphu1r. In France, distilled sulphur is used, collected in masses in barrels, into which it is poured when liquid. Melted sulphur is preferred to the flour of sulphur, which is nevrer pure, containing a quantity of sulphuric and sulphurous acid. All the sulphur intended for the manufacture of the different species of powder, comes from the refinery of Marseilles. Charcoal. The choice and preparation of charcoal for the manufacture of powder are of great importance. All woods are not equally fitted for furnishing charcoal for powder. The wood of the black alder is almost exclusively used in France for the fabrication of charcoal intended for gunpowder for war and the chase. The woods of the black alder, the poplar, the white alder, the aspen, the linden, and the willow, are about equally fit for the fabrication of the charcoal to be used in the composition of mining-powder. When there is no choice, the charcoal of any of these different kinds of wood will answer for all sorts of powder. The shoots of the alder are cut when they have five or six years' growth at the most; their bark is taken off, as in burning it leaves more ashes than the sapwood, and shoots are chosen whose diameter varies from 00151n. to 0O030m. The charcoal is prepared either in trenches or tight vessels. In the carbonization in trenches, a trench is made 3 metres long and 12 decimetres deep, the sides of which are of brick; fagots of wood are introduced, weighing 15 kilo G U N P 0 W D E'R. 159 grammes, sustained by a longitudinal rest, leaving in the trench a vacuity, which is filled with branches. This is set fire to, which soon spreads through the mass, and causes the wood to settle down. As this takes place, the trench must be filled with new fagots. When there is no more flame from the combustion, the carbonization is considered complete; to stop the combustion, the trench is covered with wet woollen cloths, then with a quantity of earth. It should not be emptied for two or three days after the fire is extinguished; without this precaution, the charcoal, still hot, would inflame in the atmosphere. This process produces charcoal which should be sorted with the greatest care, to separate foreign bodies, and parts not well charred, which are called brul0ts. The yield of charcoal fit for powder is 20 per cent. The sides of the trenches, being of brick, rapidly burn out. Vessels of iron have been substituted for these, on which are placed sheet-iron covers, to smother the fire. The charcoal used for gunpowder is now prepared in this way. Carbonization in close vessels. - This is performed in cylinders of iron, like those used for the manufacture of gas for illumination. The volatile parts escape by a tube which leads into a chimney. This distillation is made at a temperature which never reaches redness. In preparing charcoal for choice powder, it ought to be conducted slowly; the distillation of wood, for the government powder, lasts about twelve hours. Thus is obtained red charcoal. It is a charcoal which contains much hydrogen and oxygen, and is rather roasted wood, or half-charred wood, than pure charcoal. Red charcoal does not contain above 70 to 72 per cent. of carbon. 100 parts of dry wood give about 40 parts of red charcoal (charbon roux). lifanefacture of Powder. —There are two principal 160 METALS. methods of manufacturing powder. 1st, by pestles or stampers; 2d, by millstones. The Process by Pestles. —This mode of trituration is effected by means of pestles in mortars of wood. These pestles are made of pieces of beech-wood, of the weight of about 20 kilogrammes, fitted with a metallic shoe at the end, weighing also about 20 kilogrammes. The mortars are hollowed out of a stout piece of oak wood. The pestles are moved by a hydraulic wheel, Fig. 25. The preparation of the composition commences by breaking up the charcoal; the sulphur is pulverised in the triturating vessels. To each mortar is added 1'25 kilogrammes of charcoal, 7'5 kilogrammes of saltpetre, and 1'25 kilogrammes of sulphur. These three materials are then submitted to a battage (stamping), which lasts about eleven hours, and is not interrupted except by the changes necessary to prevent the adherence of any of the material to the bottom of the mortar. After eight changes, the battage continues two hours more without interruption. The material should be moistened to get it into mass. On taking out the pestles, it is in the form of a moist cake, called galette. This, before being grained, is subjected to two preliminary operations, le guillaumage and l'essorage. The ingredients are then passed over a sieve, called guillaume, the holes of which are about 8 millimeters in diameter. The workman, by a backward and forward, or rotary motion, given to the sieve, gives to a disc of hard wood placed in the sieve, a circular movement; this disc of wood breaks up the cake. The essorage consists in exposing to the air the ingredients of the mass, which is the commencement of the drying that facilitates the graining. This gives to the powder a uniform grain. It is done by means of a disc of wood in sieves of skin, the holes of 25!"' P GUNPOW D E R. 161 which are 2'40 millimeters in diameter for cannon-powder, 1'50 millimeters for musket-powder. The dust, or fine grain, is taken out by passing it over still finer holes. After these different manipulations, the powder is subjected to desiccation, either in open air, spreading it on cloths, or in the drying-rooms, or large closed apartments, into the interior of which warm air is blown, which passes over the surface of the powder, taking up its moisture. This last operation producing a great deal of dust, the powder is afterwards cleared of the dust, and has then to be tried and barrelled up. Process of the Mills.-This process is used for sporting powder: - 1. The carbon and sulphur are pulverised in drums of wood containing 120 kilogrammes of bronze balls of a diameter of 6 to 7 millimeters. The charcoal is first introduced by charges of 21 kilogrammes, and triturated alone for eight to twelve hours. Fifteen kilogrammes of sulphur are then added, which is turned with the charcoal for four hours. 2. Six kilogrammes of this binary mixture, and 20 kilogrammes of saltpetre are placed with 60 kilogrammes of bronze balls, of about 5 millimeters in diameter, in a drum of leather. The trituration of the ternary mixture is performed in the space of twelve hours, at the speed of about 20 to 25 revolutions per minute. 3. On being taken out of this, the mixture is wet with one to two per cent. of water, and placed under millstones of iron rimmed with bronze, of the weight of 2500 kilogrammes, which are called light millstones, to distinguish them from the heavy wivtls onzes used at the Government works. (Fig. 26.) The basin in which the millstones turn is of wood, and is charged with 50 kilogrammes of mixture, which is triturated during two hours. In checking the progress of 14 * L 162 METALS. the mills at the end of the trituration, cakes (galettes), are formed, which contain two to three per cent. of moisture. 4. These are reduced into grains by means of a machine which gives motion to eight sieves or bolters, mounted on the same stand, and which grains about 80 kilogrammes of cake an hour. 5. The glazing is done in barrels of wood, divided into three or four compartments by transverse divisions, each of the capacity of 100 to 150 kilogrammes of powder. The grains, by rubbing on themselves and on the sides of the vessel, acquire, in the space of twenty-four hours, the hardness and polish required. 6. The drying is done by exposure to the sun, or in an artificial drying-house. CARBONATE OF POTASIH. This salt is often known in commerce by the name of vegetable alkali, salt of tartar, dulcified alkali, or simply potash. It is formed by the combination of carbonic acid with potash. It has a sharp, and slightly caustic taste; it is very soluble in water, and deliquescent; water, at the ordinary temperature, dissolves its own weight of it; it has an alkaline reaction. It is insoluble in alcohol, fusible at a red-heat, and indecomposable by heat. Carbon, at a red-heat, acts on carbonate of potash, and gives rise to potassium; it is upon this reaction is based the preparation of potassium by the process of M. Brunner. Lime, in presence of water, transforms the carbonate into hydrate of potash. Preparation.-Vegetables contain potash united to different organic acids, such as acetic, malic, oxalic, tartaric, &c. When these salts are calcined, they are decomposed into CARBONATE OF POTASH. 163 carbonate of potash, which is found in the ashes of the vegetable. The name of pdtash of commerce is given to the soluble part of ashes, evaporated to dryness. Carbonate of potash from the lye of ashes, is not pure; it is always mixed with different soluble salts, such as the sulphate, chloride, and silicate of potash. The quantity of real potash contained in the potash of commerce, varies with the constituents of the wood which is used for making the ashes. The purest potash is that made from the ash of the birch, and the least pure from that of the pine. One hundred kilogrammes of ashes oirdinarily yield about ten kilogrammes of soluble residue called saline. The salts which accompany the carbonate of potash being much less soluble than this last salt, the saline is often purified by treating it with its own weight of cold water, which dissolves the carbonate of potash, and leaves, for the most part, the other salts. The solution, evaporated to dryness, leaves carbonate of potash purer than saline. This is generally colored brown by organic matters; after it has been calcined in contact with the air, it becomes white, and is called pearlash. In commerce, potashes bear names which refer to their origin. They are known as potashes of America, Russia, Vosges and Treves, &c. Uses.-Neutral carbonate of potash is used in the manu facture of soft soaps, of crystal glass, and Prussian blue. It is also sometimes used to transform into nitrate of potash, the nitrates of lime and magnesia contained in the saltpetre materials. The neutral carbonate, submitted to the influence of an excess of carbonic acid, is changed into the bicarbonate, which is used in the treatment of gout and gravel. 164 METALS. SODIUM. Sodium has a great resemblance to potassium. This body was isolated by Sir H. Davy, by decomposing soda with the pile. A short time after, Gay-Lussac and Thenard showed that sodium could be obtained by the action of iron on soda, under the influence of an elevated temperature. Sodium is actually prepared by the process of.M. Brunner, by decomposing the carbonate of soda by carbon, with the apparatus described in treating of the preparation of potassium. Sodium is of a silver white, of a metallic brilliancy when recently cut; but it tarnishes in contact with the air. Its density is equal to 0'92. It fuses at 900 C., and volatilizes at a red-heat. Sodium is less volatile than potassium. It decomposes water like this last metal, at the ordinary temperature. When a piece of sodium is thrown upon water, hydrogen is disengaged; but the heat produced by the reaction of this metal on water, not being so great as by potassium, does not inflame the gas. If water is thickened by dissolving gum in it, to impede the movements of the metal, or if the sodium is thrown into a glass containing but a few drops of water,.there is less loss of heat, the metal becomes incandescent, and the hydrogen inflames. The other properties of sodium resemble those of potassium. There is also a great analogy between the protoxide of sodium (soda), and the protoxide of potassium (potash). CHLORIDE OF SODIUM. The chloride of sodium, often called sea salt, is white, colorless, and of a salty but agreeable taste. It is slightly soluble in alcohol. Sea salt is nearly as soluble at the ordinary temperature, as at the boiling-point of water which is CHLORIDE OF SODIUM. 165 saturated with it; thus, the boiling and saturated solution of chloride of sodium, only deposits small quantities of it on cooling. This property permits sea-salt to be separated from most other salts, and particularly from nitrate of potash, the solubility of which, in water, increases greatly with the temperature. In fact, in treating a mixture of sea-salt and nitrate of potash with boiling water, a great part of the nitrate deposits on cooling, while the sea-salt remains in solution in the water. Sea-salt crystallizes in cubes, or in aggregated forms produced by the symmetric arrangement of a number of small cubes; these crystals are anhydrous, and decrepitate strongly when heated to 200~ or 300~. They will keep in the atmosphere in dry weather. The chloride of sodium is fusible at a red-heat, and volatilizes at a still higher temperature, producing white fumes. This vaporization increases considerably in a current of gas. Sea-salt fused, crystallizes in cubes on cooling. In this state, it does not decrepitate when heated. Some oxides, and especially the oxide of lead, decompose sea-salt dissolved in water, producing a metallic chloride and caustic soda. This reaction takes place so readily, that it has been proposed in the arts to prepare soda by treating sea-salt with litharge; but then the soda always contains in solution a considerable quantity of oxide of lead. The process of Leblanc presenting, besides, incontestable advantages over all others, the use of litharge has been given up in the preparation of soda. When a mixture of sea-salt well dried, and silex, are heated together, no reaction takes place; but if a current of steam is made to pass over the mixture, there is formed a silicate of soda, and chlorhydric acid. It is on this reaction is based the use of sea-salt in glazing some kinds of potteries, such as stone-ware. A quantity of sea-salt is 166 METALS. thrown into the oven; this volatilizes, and meeting with the silex which exists in the paste of the pottery, and with steam, produces silicate of soda, which forms a vitreous coating on the surface of the pottery. Uscs. -Tho uses of sea-salt are numerous. It is used in the preparation of artificial soda, and of sulphate of soda; also in the glazing of potteries. It is used in preparing chlorhydric acid, in the manufacture of the decolorizing chlorides, and to produce chlorine. It is used for domestic purposes, and in agriculture, in large quantities. It is one of the most widely-spread salts in nature. It exists in large quantities in sea-water, in lakes, and salt springs. It is also deposited in abundant strata in the earth, and is then called mineral or rock salt. ROCK SALT. Rock salt is crystallized, and often in transparent masses of a milky white; it presents an easy cubic cleavage. It is sometimes met with in fibrous masses. It is generally colored grey by a small quantity of bitumen, and often has a reddish color, due to the presence of the oxide of iron. Rock salt is sometimes of great purity, as is the case with that of Wieliczka; but often is mixed with sulphate of lime, clay, &c. Some specimens of the salt of Wieliczka have a curious peculiarity, which has been examined by MM. Dumras and Rose. When this salt is placed in water, a series of decrepitations are heard, and there is a disengagement of gas, which sometimes appears to be pure protocarburet of hydrogen; and sometimes a mixture of this gas with hydrogen and carbonic oxide. It is probable that the gas is imprisoned under more or less pressure between the molecules of the salt, the layers of which it breaks as soon as they are ROCK SALT. 167 reduced in thickness by the action of water, and then produces a decrepitation. In France, attempts have been made to introduce pulverized rock salt into the market; but it has not yet been introduced into consumption, on account, no doubt, of the presence of foreign bodies which it always contains, and the slowness with which it dissolves in water. Rock salt has all the properties of common sea-salt. It is, however, attacked much more slowly than this last salt by sulphuric acid, and does not decrepitate with heat; in both these respects, it resembles melted sea-salt. Extraction of Chloride of Sodium. —Rock salt is extracted either in the solid state, by shafts and galleries, or from its solution. Extracted in the solid state, rock salt is exposed for sale in blocks, or after having been first broken. When it is not very pure, it is dissolved in water, and crystallized by evaporating the solution. To extract the sea-salt held in solution in salt-springs, the waters are first evaporated in the open air in places called drying-houses, so arranged as to offer great surface for evaporation. These drying-houses are composed of large sheds, in which are piled up bundles of thorny sticks. The salt water is brought to the top of the sheds in canals, which communicate with troughs leading to lateral gutters, which tumble the water over the fagots; this, in falling, is divided, and evaporates rapidly. The course of the water is changed to suit the direction of the wind, which has a great influence on the rapidity of the evaporation. The drying-houses are generally divided into two sections. The first receives the water from the spring; the second, the waters which have already circulated over the fagots. Pumps placed at intervals, and moved by water-wheels, raise the water from the lower reservoirs to the conduits, which pour it out on the fagots. 168 METALS. As the water concentrates, it deposits on the fagots sulphate of lime, generally mixed with carbonate of lime and oxide of iron. These deposits are removed from time to time. When the water is concentrated to a point corresponding to about 15 to 20 per cent. of salt, the evaporation is finished in boilers. The salt contained in sea-water is extracted by subjecting this water to a spontaneous evaporation in large reservoirs, called salt-pans. In cold countries, where the method of salt-pans cannot be used, the salt is extracted by exposing the sea-water to a very low temperature. The water divides itself into two parts; one solidifies first, and is nearly pure water; while the other remains liquid, and retains in solution all the soluble salts. In taking away, from time to time, the ice which forms, water is finally obtained highly charged with salt, which is then evaporated in boilers. BORATE OF SODA. This is called Borax in commerce. It exists in nature, and is found in Persia, India, and China. Native borax crystallizes in hexahedral prisms. These crystals are impure, and are often found mixed with a fat matter whose composition is unknown. It is sometimes known in commerce by the name of tinal. To purify tincal, it is treated with lime-water, which forms, with the fat matter, an insoluble compound; and the salt is crystallized in vessels of wood or lead. Borax is generally produced by uniting boracic acid directly with soda, which comes from Tuscany. The borate of soda is white, and of an alkaline taste and reaction; when it is heated, it melts, and forms, in cooling, a solid, transparent, and vitreous mass. CARBONATE OF SODA. 169 Uses. — Borax is used in soldering. When oxydizable metals are to be soldered, they are covered with borax, which, in melting, prevents oxidation, and dissolves traces of oxide which would prevent the soldering. Borax enters into the composition of many glasses. It is chiefly used in the fabrication of very fusible glasses, and in the glazing of some potteries. CARBONATE OF SODA. Carbonate of soda is a white salt, inodorous, of a sharp and slightly caustic taste, and of an alkaline reaction. It is very soluble in boiling water, and crystallizes in large prisms which are hydrated. Exposed to the air, it loses part of its water of crystallization, and effloresces. Carbonate of soda is decomposed at a red heat by steam, which disengages all its carbonic acid, and produces soda. Lime decomposes the carbonate of soda, by depriving it of its carbonic acid, and isolating the soda. Preparation of Carbonate of Soda. — For a long time, the carbonate of soda used in the arts was extracted either from marine plants, such as fuci and sea-weed, or from certain terrestrial plants, as salsola soda, or barilla, which grow on the sea-shore. These plants are burnt, and from their ashes is extracted, by lixiviation and evaporation, salts more or less rich in carbonate of soda, which are called sodas of sea-weed, of Alicant, of C'arthtagena, of Malaga, of Narbonne, &c. But little has been done at working these natural sodas, since Leblanc discovered a method of obtaining carbonate of soda artificially, by decomposing by chalk and carbon, the sulphate of soda, which is produced by treating sea-salt with sulphuric acid. This discovery is considered, justly, as one of the most 15 170 METALS. important which has ever been made in the industrial arts. The process of Leblanc, perfected by D'Arcet and Aufrye, is now exclusively used in the manufacture of carbonate of soda. Wle will describe it somewhat in detail. Into a reverberatory furnace of an elliptic form, with a bottom of large surface constructed of fire-brick, is introduced a mixture of 400 kilogrammes of anhydrous sulphate of soda,- 400 kilogrammes of dry chalk, and 135 to 140 kilogrammes of coal. These materials are stirred, from time to time, with an iron poker; they soften down at a red heat, and acquire by degrees a pasty consistence, giving off a great quantity of gas, which burns with a blue flame. After four or five hours of calcination, the semi-fluid mixture is stirred anew, brought to the front of the furnace with an iron rake, and put into a sort of wheel-barrow of thick sheet-iron, where it is left to cool. This product is called crude artificial soda. The above mixture usually yields'from 550 to 600 kilogrammes of crude soda, of 38 to 40~. Two workmen can manufacture many thousand kilogrammes per day. Crude soda is of a bluish grey; it is slightly porous; exposed to moist air, it deliquesces, and becomes friable. When first made, it is hard. It is then pulverized or broken up, and subjected to the action of warm water, which dissolves all the soluble parts which it contains. The sulphuret or oxysulphuret of calcium, and the excess of charcoal which it contains, are separated by decantation. The solution is evaporated in iron boilers. The carbonate of soda precipitates to the bottom of the vessel; it is then taken out with ladles, as it deposits, and placed to drain. The carbonate thus deposited, after its calcination in a reverberatory furnace, is often exposed for sale in this state. To perfect its purification, it is sometimes subjected to a new solution, and then evaporated to dryness. The product thus obtained is called, in commerce, sal soda. C A L C I U M- L I I E. 171 Uses.-The carbonate of soda is used in the manufacture of glass and soaps. These consume it in enormous quantities. It is also used in some operations of dyeing, and chiefly in the washing of threads and tissues. Crude soda, mixed with quick-lime and lye, furnishes a liquor used in the manufacture of Marseilles soap. The neutral carbonate, treated with carbonic acid, is transformed into the bicarbonate of soda. The bicarbonate is used in medicine in the preparation of pastilles of vichy, and in the treatment of calculous affections. CALCIUM. Calcium was isolated by Davy, with the pile. A small cupel of moist lime was filled with mercury; the positive pole was placed in communication with the metallic plate on which the cupel was placed, and the negative pole with the mercury: he thus produced an amalgam of calcium. This amalgam, distilled, gave pure calcium. Potassium will decompose lime at a high temperature, and free the calcium. Calcium is white, and of a metallic brilliancy; burnt in the air, it is transformed into lime. LIME. Lime is a base formed by the combination of oxygen with calcium, and is oxide of calcium. Lime was known to the ancients; it entered into the composition of the mortars used by them. This base is white, caustic, and very alkaline. It greens the syrup of violets. When it is plunged into water and taken out, as soon as the air contained in its pores escapes, it hydrates, giving rise to a disengagement of heat as high even as 3000 C., and produces a hissing noise, accompanied with thick vapors of water. 172 METALS. The high temperature developed by lime, with water, is sufficient to inflame gunpowder. Lime which is fallen, that is, reduced to powder by absorbing water, is often called slaked lime, to distinguish it from quick-lime, which is anhydrous lime. Lime suspended in water, constitutes milk of lime. When lime is of good quality, and is placed in water, it increases in volume; such lime called fat lime. Lime-water is much used in laboratories as a reagent. To prepare it, some lime is put into a bottle, which is then filled with water. The excess of lime deposits, and the lime-water remains clear; this first solution is never pure, mostly containing some potash which existed in the lime. To prepare a pure lime-water, the lime must be washed three or four times, before leaving it definitively in contact with water. Lime, in the anhydrous or hydrated state, absorbs carbonic acid from the air, and produces carbonate of lime; in carbonating, it becomes excessively hard, and transforms itself into a substance which has the composition of limestone, and often its hardness. It is from this important property that lime is used in making mortars. Natural State of Lime. Lime is never met with in nature in a free state; if this base was for a moment isolated, it would immediately absorb the water and carbonic acid contained in the air. It is found combined with carbonic acid, constituting all the varieties of carbonate of lime, known as chalk, marble, Iceland spar, &c. The shells of mollusca are almost entirely carbonate of lime. Lime, combined with sulphuric acid, constitutes plaster, a body widely spread in nature. This base enters into the composition of bones, and is there found in the state of phosphate and carbonate of lime. It is, besides, combined in nature in different proportions r~~~~~~~~~~ F ass if; -93 Id A2 LIMlE. 173 with silex. Most vegetables contain lime united with organic acids. Extraction of Lirne.-As lime is indecomposable by heat, those salts of lime are used for its preparation, whose acids can be driven out by a high heat. The nitrate of lime would do; but this salt not being abundant, the carbonate is always used. Pure carbonates furnish, by calcination, fat lime. If it is impure, or especially if it is clayey, it leaves a poor lime, and hardens by contact with water. When pure carbonate of lime is calcined, its decomposition is slow, and it requires a high heat. Different gases, such as the carbonic oxide, hydrogen, and especially steam, influence the decomposition of the carbonate of lime. When a porcelain tube is filled with pieces of lime, and exposed to a low red heat, no disengagement of carbonic acid gas is seen to take place; but if, as Gay-Lussac has shown, a stream of gas or steam is made to pass into the tube at this temperature, the decomposition of the salt soon takes place. The influence of steam on the decomrposition of carbonate of lime, has been long known to lime-burners, who know that a stone, while moist, decomposes much more easily than one which has been dried in the air. They often throw into their lime-kilns a small quantity of water, which, in vola. tilizing, hastens the decomposition of the liome-stone. Different kilns are used for burning lime. In some places, the kilns are made of the lime-stone itself, resembling in some respects that used in forests for making charcoal, This is covered with turf and earth, leaving within a central chimney, by which the gases of combustion escape. Sometimes, also, an intermittent kiln is used, the vault of which is made also of lime-stone. These are now replaced by continuous, perpetual, or draw kilns, which are formed of truncated and inverted cones, Fig. 23. Lime and coal are alternately introduced into 15' * 174 M E T ALS. these kilns, in the proportion in bulk of four parts of stone to one part of coal. These kilns are filled up to the level of the upper opening. The fire is lighted below by means of fagots, and communicates gradually; the carbonate of lime loses its carbonic acid, and the calcination is considered finished whenever the smoke ceases to come out. The lime is then taken out below, and fresh quantities of lime and coal are added above. Uses of Lime.-The uses of lime are numerous. Lime is the basis of mortars; it is used in tanning, and in the purification of illuminating gas, to absorb sulphydric and carbonic acids. It is also used in the preparation of potash and soda, to carry off the carbonic acid from the alkaline carbonates. Lime is used in the saponification of fat bodies intended for the manufacture of stearin-candles, and in the manufacture of sugars, in the operation called defecation. It is used in agriculture as a manure. When a soil is too clayey, a quantity of lime is mixed with it, which, by absorbing water and carbonic acid, slakes and makes the ground lighter, and assists vegetation. This addition of lime also restores the calcareous element which vegetation carries off with each crop. CHLORIDE OF LIME. The name chloride of lime, bleaching chloride, is given to a body prepared by treating the hydrate or milk of lime with chlorine. This body is white, amorphous, pulverulent, and of a strong odor recalling that of chlorine; water dissolves it freely. When treated with weak acids, it gives off chlorine. Chloride of lime acts on organic substances and decomposes them: it destroys coloring matters: its action however on these matters is slow when it contains an excess of lime. Thus the blue color of litmus and chloride of lime SULPHATE OF LIME. 175 may be mixed together, without the color being sensibly changed; but the intervention even of a very feeble acid displaces hypochlorous acid, and thus causes the decolorizing of the litmus. Preparation.-To prepare chloride of lime, a stream of chlorine is made to pass through milk of lime or into slaked lime. The mass mixed through water, and decanted or filtered, gives a concentrated solution of chloride of lime. Uses.-Chloride of lime is used in large quantities for bleaching cloths, and the paste for making paper, and in the manufacture of painted stuffs. The miasmas of hospitals and dissecting rooms may be destroyed by chloride of lime; this should always be used in small quantity, because chlorine in excess, mixed in the atmosphere, would exercise an unwholesome influence on the lungs. The name ecu de Javelle, is given to a liquor obtained by passing chlorine into potash. Eau de Javelle disengages chlorine under the influence of acids, and may be used as a decolorizer and disinfecter. Chloride of lime and eau de Javelle then act like Chlorine when treated with acids; in the arts they are preferred to chlorine, because their odor is less active, their action slow, successive, and continued, and finally, they are more easily transported. SULPHATE OF LIME, GYPSUM, PLASTER. Gypsum, or plaster, is sulphate of lime; this salt is crystallized in large transparent tables, which are ordinarily grouped infer de lance. When the crystals of sulphate of lime are opaque, they are called gypsum alabaster. This must not be confounded with real alabaster, which is carbonate of lime. Hydrated sulphate of lime is insipid, or of a slightly bitter taste; it is colorless, and indecomposable by heat. 176 METALS. This salt is equally soluble in hot or cold water, for at +100 to 1000 C., 1000 parts of water dissolve 3 parts of plaster. It is more easily dissolved in concentrated sulphuric acid; it is completely insoluble in alcohol mixed with water: it contains 20.9 per cent. of water: it is entirely dehydrated at a temperature below 200~, particularly in a current of gas. The hydrated sulphate of lime has the hardness of stone; after it is dehydrated, it becomes pulverulent and like flour. When sulphate of lime dehydrated is placed in contact with water, it hydrates anew, combines with 2 equivalents of water which heat had caused it to lose, and resumes its original hardness. On account of this property, sulphate of lime is used in building; when it is calcined, it dehydrates; when water is added, it takes up precisely the quantity necessary to restore it to its original hardness. The calcining of plaster is done in furnaces, the vault of which is made with the stones of the plaster itself. (Fig. 24.) The temperature ought not to be very high, for a temperature of 2000 C. is sufficient; after the calcining, it is reduced to powder in mills. Calcination at too high a temperature causes the plaster to undergo a fusion (sorte de fritte) which prevents its hydrating easily. Plaster once calcined ought to be protected from the moisture of the air, or it will by degrees hydrate, become dead, and then lose a part of its good quality. Plaster well prepared ought to produce heat when mixed with water. Often, indeed, its quality is tested by the temperature it produces in hydrating. Plaster often gives off sulphuretted hydrogen when mnixed with water, owing to the presence of a quantity of sulphuret of calcium produced by the action of carbon or carburetted gases on the sulphate.of lime; this sulphuret gives off traces of sulphuretted hydrogen under the influence of water and carbonic acid. SULPIIATE OF LIME. 177 Plaster, in solidifying, increases in volume; this property renders it eminently fit for moulding, for in dilating itself it takes the impression of the finest marks. The most esteemed plaster is that found in the vicinity of Paris. The hardness of plaster depends altogether on that of the hydrated sulphate of lime which it produces. This hardness is found to be, after mixing, what it was before calcining, that is, what it was in the stone itself. The sulphate of lime transforms itself into sulphuret of calcium, under the influence of organic matters in a state of decomposition. This sulphuret, afterwards decomposing by the action of carbonic acid, gives rise to the disengagement of sulphydric acid. In this way can be explained the presence of sulphydric acid in certain waters which originally contained sulphate of lime. MJ. Chevreul was the first to show that a like decomposition might take place in the soil in some large cities, which contain, like that of Paris, a large quantity of sulphate of lime. This sulphate, in changing into sulphuret, might become in time a cause of insalubrity. M. Chevreul proposed aerating the soil of large cities, so as to change the sulphurets found there into sulphates. Uses.-Sulphate of lime is used in buildings as a cement, as it presents the advantage that it solidifies in a few minutes. The name of stucco is given to plaster mixed with water containing glue, and sometimes gum in solution. Stucco is easily polished, and has the appearance of marble. Various colors may be given to it: often pieces of marble are put into the stucco before solidification, which are then polished with the stucco. Stucco does not resist moisture, but may be used in the interior of houses. Stucco with lime is a composition obtained by mixing lime M 178 M ETALS. with pulverized marble; in its composition it presents no analogy with that made with sulphate of lime. For some time, a peculiar cement has been made with plaster, called aluminated plaster. This body, which is prepared by calcining plaster.with alum, has, like plaster, the property of solidifying rapidly when slaked with water; but it becomes harder, and produces a mass which has both the hardness and semi-transparency of marble. It appears to resist the influence of moisture. M. Kuhlman has shown that plaster may be made very hard by subjecting it to the action of a solution of silicate of potash, which produces on the surface of the plaster a coating' of silicate of lime. Plaster is used in agriculture: it increases the growth of some plants, particularly leguminous plants. CARBONATE OF LIME. Carbonate of lime, from its numerous uses and its abundance in nature, is one of the most important salts. It presents itself in different states bearing the name of Iceland spar, aragonite, limestone, marble, chalk, alabaster, and lithographic stone. Carbonate of lime assumes variable forms, which, however, may all be reduced to two principal ones, which are the form of calcareous spar and that of aragonite. These two forms are incompatible. Properties.-Pure carbonate of lime is perfectly white; small traces of foreign matters are sufficient to color it. The color of marbles is generally ascribed to metallic oxides. The hardness of carbonate of lime varies greatly with the different varieties of this salt. Thus, it is well known that marble is harder than lime-stone, and the hardness of this latter is greater than that of chalk. This depends, probably, on the mode of formation of the carbonate of lime. It decomposes, at a red heat, into carbonic acid and lime. LIME. 179 It is on this property is founded the manufacture of lime; nevertheless, this decomposition does not take place when the carbonate is burned in a hermetically closed vessel. Hall showed that if chalk is introduced into a gun-barrel closed at both ends, the calcareous carbonate, instead of decomposing, fuses, and forms a body which has all the properties of marble. It was attempted, some years ago, to make artificial marble, by the fusion of an amorphous carbonate of lime. A manufactory was established at Paris, where white or variously-colored marbles were made by melting pure chalk, or chalk mixed with the metallic oxides. This enterprise fell through; but the problem of the manufacture of artificial marbles is nevertheless resolved. The experiment of Hall, too, has explained the presence of crystallized carbonate of lime in soils of igneous origin. Carbonate of lime is insoluble in water, so that it can easily be produced by treating a soluble carbonate with a salt of lime. When this double decomposition takes place at the common temperature, a crystalline precipitate is obtained, which, under the microscope, presents small crystals having the form of Iceland spar. If the decomposition is made hot, the crystals which are produced belong to the crystalline system of aragonite. Carbonic acid dissolves the carbonate of lime, and gives rise to a bicarbonate soluble in water. This property accounts for many natural phenomena. Whenever water, holding carbonic acid in solution, passes over calcareous deposits, it dissolves the carbonate, and gives rise to a bicarbonate of lime. As the bicarbonate is not very stable, and as it readily loses half of its carbonic acid, this salt may, in a great number of circumstances, form deposits of insoluble carbonate of lime. Neutral carbonate of lime, proceeding from the decomposition of the bicarbonate, produces — 180 METALS. 1. The calcareous deposits which often obstruct waterpipes, particularly about the joints. 2. The calcareous deposits in steam-boilers. 3. The crystalline deposits called stalactites and stalagmites. When water, holding carbonic acid in solution, passes over calcareous rocks, it dissolves, as we have said above, the carbonate of lime; then, on reaching the interior of caves, each drop evaporates, and deposits the neutral carbonate of lime. When these concretions form under the roofs of caves, they are called stalactites; if they are produced on the ground where the water falls, they are called stalagmites (Fig. 27). When the stalagmites have yellow and red zones, they form the oriental alabaster, which will often take a beautiful polish. 4. The calcareous incrustations are also produced by the carbonate of lime; when different objects, such as fruit, birds'-nests, &c., are exposed to the action of certain mineral waters holding bicarbonate of lime in solution, these objects become covered with calcareous incrustations, often known as petrifactions. The most celebrated incrusting springs are the Sprudel at Carlsbad, which produce a calcareous deposit, zoned, and of great fineness, which is used for objects of ornament. The waters of San Filippo, in Tuscany, and of St. Allyre, in Auvergne, are also examples. 5. The calcareous tufs or gravel stones, which are very abundant in some countries, and which are used for building stones, have also the same origin. Many towns in Italy are built with calcareous tufs riddled with small holes, and evidently produced by calcareous deposits formed from the decomposition of bicarbonate of lime. MARBLES. Marbles belong to two calcareous varieties: the saccharoid and compact varieties. The saccharoid carbonate of lime is formed of small crys - - ~~~~~~~~~~~~-~ r~r - - ~~~~~~~~~i' -4.::;it~- -rL6r.ii::: ~- -: -~i:'L:::-::,,,.:....,:.i~~ "'4,",