A TEXT-BOOK OF INORGANIC CHEMISTRY RICHTER. STANDARD TEXT-BOOKS. RICHTER'S CHEMISTRY OF THE CARBON COMPOUNDS, OR ORGANIC CHEMISTRY. New Edition. A text-book for students. By Prof. victor von Richter, University of Breslau. Second American from Sixth Revised and Enlarged German Edition. Authorized Translation. By Edgar F. Smith, m.a., ph.d., Professor of Chemistry, University of Pennsylvania, Phila. Illustrations. Cloth, $4.50 The present edition will be found to differ very considerably, in its arrangement and size, from the first. The introduction contains new and valuable additions upon analysis, the determination of molecular weights, recent theories on chemical structure, electric conductivity, etc. The section devoted to the carbohydrates has been entirely rewritten. The sections relating to the trimethylene, tetramethylene, and pentamethylene series,*the furfurane, pyrrol, and thiophene derivatives, have been greatly enlarged, while the subsequent chapters, devoted to the discussion of the aromatic compounds, are quite exhaustive in their treatment. Such eminent authorities as Profs. Ostwald, von Baeyer, and Emil Fischer have kindly super¬ vised the author's presentation of the material drawn firom their special fields of investigation. The characteristic features of the first edition have been retained, so that the work will continue to be available as a text-book for general class pmposes, useful and reliable as a guide in the preparation of organic compounds, and well arranged and satisfactory as a reference volume for the advanced student as well as for the practical chemist. SMITH'S ELECTRO-CHEMICAL ANALYSIS. a practical handbook. By Edgar F. Smith, Professor of Chemistry, University of Pennsylvania, Translator of Riehter's Chemistries, etc. 26 Illus¬ trations. i2mo. Cloth, $1.00 This handbook is designed to meet the wants of a large and growing class of students and chemists, who are desirous of becoming acquainted with this method of quantitative analysis, which is daily acquiring more importance. The author has devoted much time to this branch of analysis and has succeeded in making a book that is exceedingly clear, concise, and practical. SMITH & KELLER, EXPERIMENTS. Second Edition. arranged for students in general chemistry. By PrOF. EDGAR F. Smith, Translator of Richter's Chemistries, and Prof. H. F. Keller, Prof, of Chemistry in the Central High School, Philadelphia, Pa. Second Edition. Enlarged. With 37 Illustrations. Interleaved for Notes. 8vo. Cloth, «í/, .60 *^*This little work is designed as a guide for beginners in chemistry. The arrangement of the course is such as the authors have used with success in the instruction of their classes; its object is not to dispense with the supervision of an instructor, but rather to assist him. Although reference is made to Richter's " Inorganic Chemistry," any other text-book on the subject can be employed in its stead. The experiments have been collected from various sources, and no claim is made for originality. P. BLAKISTON, SON & CO., PUBLISHERS, PHILADELPHIA. A a 0 10 . 20 In H ! 11 i 1-Ii! i 11 I ! I [|[ B C 30 20 30 i!i Ml III iii Ml I nil I lililí MI nil!! E b ; 70 60 ) G 120. 130 h H Hl 150 150 170 II III III III 1111 il ! I i Ii I Mil INI i Mill 11 M i NIM MM M.;: i M il'.; I ;• - ii-M: . ; G 10 20 1 ! 1M 11 M 11 1 in 1 ni 1 1 1 M 1 i 1 h:œ 60 70 ÔO 9 0 100 110 120 130 140 150 160 170 II! M I lin M I mmImM I I Mill M 11 M II In M I M M i M Iii M M I II M I MM 11 M i Im M I M I 11 I M I ! 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Mil 1 II V: M| 1 II 1: ■ M H j, î G ; y 40 50 60 'i M M i Ml MJM-!, IM 1 M 40 vy' Oj 'v '■ i 50 '0--vi-.. 60 111 M 1 i 1 ! 1 il M M 1 M I M f • M ml-M: 80 910 100 110 120 130 140 150 ! 160 M I I M,j I h M^tlj 11 M I M M 11 M I h Illl II M II M 11 M M I MIM 1 M M I M M I M m! 11 Iii M M I M II i li M •'i!ÎM M; 11: M: -Ml ■ " -,.•1 ! /li il I hiilk^ " y 6 |c6|/3 • 'M'!lu il ..111 ÔO 9|0 100 110 120 Il i ; I i IJ i 1 II 111! M M i I n 11 M I I ! M II I M II i M M 111 M 11 130 140 II! M Hill Ml M 60 9j0 100 110 120 13i liiIrMMlilMihillilM.iillIllMiliMlhlMlItMlIllMl 120 130 140 150 160 ; 170 MltMllMMlMMllll!.lMM!MM!íMIÍMMllllll Uli i i 1 ■ a . ß O m-* o o 30 40 5 0 60 70i' ÔO ^ 80 100 110 120 130 140 150 IßO M M 1 M M 1 fin il III 1 Mill ! M 11 M 1 1 m 11 M II ! M IJ, ijüliMilii mImmIímiI MM MM! MM m lin M MM im im imlimlii iiliiMlimliMiliiiilmi mill m ! >3*■'"v 1 1 Kaa Lia Sró Rb Kay3 A TEXT-BOOK OF INORGANIC CHEMISTRY. BY PROF. VICTOR VON RICHTER, UNIVERSITY OF BRESLAU. AUTHORIZED TRANSLATION, BY EDGAR F. SMITH, PROFESSOR OF CHEMISTRY IN THE UNIVERSITY OF PENNSYLVANIA, PHILADELPHIA. FOURTH AMERICAN FROM THE SIXTH GERMAN EDITION. CAREFULLY REVISED AND CORRECTED. WITH EIGHTY-NINE ILLUSTRATIONS ON WOOD AND COLORED LITHOGRAPHIC PLATE OK SPECTRA. PHILADELPHIA: P. BLAKISTON, SON & CO., No. IOI2 Walnut Street. 1892. Copyright, 1887, »y P. Blakiston, Son & Cb. WM. F. FELL & CO., Eleotrotypeps ano Printers, 1220-24 SANSOM STREET, PHILADELPHIA. PREFACE TO THE FOURTH AMERICAN EDITION. In its present form this volume represents a translation of the sixth German edition. The inductive method adopted in previous editions has been continued. The changes that will be observed throughout the text consist mainly of corrections and additions made necessary by the most recent investigations. It may be added that as three years have passed since the publication of the German edition, the translator has taken the liberty of intro¬ ducing here and there such new matter as the present state of chemical science has seemed to demand. • • vu PREFACE to the THIRD AMERICAN EDITION. The present edition contains a rather extended section upon the thermal behavior of bodies, and throughout the work frequent occasion is taken to call attention to the dynamical side of chemi¬ cal reactions. The sections upon the pressure and condensation of gases, and that upon the dissociation phenomena, have also been considerably increased, while new facts relating to the ele¬ ments and their derivatives, to their preparation, etc., have been introduced in their proper places. The leading and characteristic features of the preceding editions remain unchanged. The present edition is a translation of the fourth German edition. Many parts of the work have been rewritten, and new matter incorporated. The features which recommended the previous edition have been preserved, and as now presented it hoped that the work may continue to be of service to the student «f chemical science. The translator would here express his obligations to Mr. Allen J. Smith, who has read the entire proof, revised the index and table of contents, as well as rendered other valuable assistance* to the SECOND AMERICAN EDITION. vm PREFACE to the FIRST AMERICAN EDITION. The success of Prof, von Richter's work abroad would indi¬ cate its possession of more than ordinary merit. This we believe true, inasmuch as, in presenting his subject to the student, the author has made it a point to bring out prominently the relations existing between fact and theory. These, as well known, are, in most text-books upon inorganic chemistry, considered apart, as if having little in common. The results attained by the latter method are generally unsatisfactory. The first course—that adopted by our author—to most minds would be the more rational. To have ex¬ periments accurately described and carefully performed, with a view of drawing conclusions from the same and proving the inti¬ mate connection between their results and the theories based upon them, is obviously preferable to their separate study, especially when they are treated in widely removed sections or chapters of the same book. Judging from the great demand for von Richter's work, occasioning the rapid appearance of three editions, the com¬ mon verdict would seem to be unanimously in favor of its inductive methods. In the third edition, of which the present is a translation, the Periodic System of the Elements, as announced by Mendelejeff and Lothar Meyer, is somewhat different, in the manner of de¬ velopment and presentation, from that appearing in the previous edi¬ tions. This was done to give more prominence to and make more general the interesting relations disclosed by it. Persons examining this system carefully will be surprised to discover what a valuable aid it really has been, and is yet, in chemical studies. Through it ix X PREFACE TO THE FIRST AMERICAN EDITION. we are continually arriving at new relations and facts, so that we cannot well hesitate any longer in adopting it into works of this character. It is, indeed, made the basis of the present volume. In accordance with it, some change in the treatment of the metals, ordinarily arbitrarily considered, has been made. A new feature of the work, and one essentially enlarging it, is the introduction of the thermo-chemical phenomena, briefly pre¬ sented in the individual groups of the elements and in separate chapters, together with the chemical afiinity relations and the law of periodicity. ''Hereby more importance is attributed to the principle of the greatest heat development than at present appears to belong to it, because it was desired, from didactic con¬ siderations, by the explanation of the few anomalies, to afford the student the incentive and opportunity of deductively obtaining the majority of facts from the thermal numbers, on the basis of a simple principle. To facilitate matters, there is appended to the volume a table containing the heat of formation of the most im¬ portant compounds of the metals." Trusting that the teachings of this work will receive a hearty welcome in this country, and that they will meet a want felt and often expressed by students and teachers, we submit the following translation of the same. TABLE OF CONTENTS. INTRODUCTION. Province of Chemistry, 17, Definition of Chemistry, 19. Chemical Elements, 19. Principle of Indestructibility of Matter, 21. Principle of Conservation of Enei^, 21. Chemical Energy, 22. Chemical Symbols and Formulas, 23. Conditions of Chemical Action, 26. Thermo-Chemical Phenomena, 26, Crystallography, 30. SPECIAL PART. Classification of the Elements, 39. Hydrogen, 40. Purifpng and Drying of Gases, 42. Apparatus for the Generation and Collection of Gases, 42. Condensation of Gases, 47. Critical Condition, 47. Group of Halogens, 49. Chlorine, 49. Bromine, 52. Iodine, 54. Fluorine, 55. General Charac¬ teristics of the Halogens, 56. Compounds of the Halogens with Hydrogen, 56. Hydrogen Chloride, 56. Acids, Bases, Salts, 59. Hydrogen Bromide, 61. Hydrogen Iodide, 62. Hydrogen Fluoride, 64. General Character¬ istics of the Hydrogen-Halogen Compounds, 66. Thermo-Chemical Deport¬ ment of the Halogens, 66. Compounds of the Halogens with Each Other, 69. Weight Proportions in the Union of the Elements, Law of Constant Proportions, Atomic Hypothesis, 6g. Density of Bodies in State of Gas, Volume Propor¬ tions in the Union of Gases, Atomic Molecular Theory, 73. Avogadro's Law, 77. Status Nascens, 79. Oxygen Group, 80. Oxygen, 81. Oxidation and Reduction, 84. Ozone, 85. Isomerism and Allotropy, 88. Compounds of Oxygen with Hydrogen, 89. Water, 89. Thermo-Chemical Deportment of Solutions, 93. Dissociation, 94. Quanti¬ tative Composition of Water, 97. Molecular Formula of Water, Atomic Weight of Oxygen, 98. Hydrogen Peroxide, loi. Sulphur, 105. Mole- CONTENTS. cules of the Elements, 107. Hydrogen-Sulphide, 108. Hydrogen Per- sulphide. III. Compounds of Sulphur with the Halogens, 112. Selenium, 113. Tellurium, 114. Summary of the Elements of the Oxygen Group, 114. Thermo-Chemistry of their Hydrogen Compounds, 115. Nitrogen Group, 116. Nitrogen, 1x6. Atmospheric Air, 118. Eudiometry, 121. Measuring Gases, 123. Diflusion of Gases, 124. Compounds of Nitrogen with Hydrogen, 125. Ammonia, 125. Ammonium Salts, 129. Atomic Weight of Nitrogen, 130. Hydroxylamine, 131. Diamide, 132. Hydrazoic Acid, 132. Compounds of Nitrogen with the Halogens, 133. Phosphorus, 134. Compounds of Phosphorus with Hydrogen, 137. Phosphonium Salts, 139. Compounds of Phosphorus with the Halogens, 140. Arsenic, 143. Arsine, 144. Compounds of Arsenic with the Halogens, 146. Antimony, 147, Stibine, 148. Compounds of Antimony with the Halogens, 149. Characteristics of the Elements of the Nitrogen Group, 150. Their Thermo-Chemical Deportment, 151. Carbon Group, 151. Carbon, 151. Carbon Compounds of Hydrogen, 153. Methane, 153. Atomic Weight of Carbon, 154. Ethane, 154. Ethylene, 155. Acetylene, 156. Natme of Flame, 156. Compounds of Carbon with the Halogens, 161. Silicon, 162. Hydrogen Silicide, 163. Compounds of Silicon with the Halogens, 164. Hydrogen Silico-Fluoride, 165. Atom and Molecule, 167. Determination of Molecular Value from Chemical Re¬ actions, 169. Valence of the Elements, Chemical Structure, 170. Equiva¬ lence, 171. Maximum Valence, 174. Oxygen Compounds of the Metalloids, 176. Oygen Compounds of the Halogens, 178. Oxides of Chlorine, 178. Hypochlorous Oxide, 179. Hypochlorous Acid, 179. Chlorine Trioxide, Chlorous Acid, Chlorine Tetroxide, 180. Chloric Acid, 181. Perchloric Acid, 182. Oxides of Bromine, 183. Oxides of Iodine, 183. Hydrates of the Acids, 184. Thermo-Chemistiy of the Oxygen Compounds of the Halogens, 185. Oxygen Compounds of the Elements of the Sulphur Group, 185. Oxygen Compounds of Sulphur, 187. Sulphm Dioxide, 187. Sulphurous Acid, 189. Hydrosulphurous Acid, 190. Sulphur Heptoxide, 190. Per- sulphates, 191. Sulphur Trioxide, 191. Sulphuric Acid, 193. Lead Chamber Process, 194. Pyrosulphuric Acid, 197. Sulphuric Acid Chlor-anhydrides, 199. Polythionic Acids, 200. Thiosulphmic Acid, 201. Oxygen Derivatives of Selenium and Tellurium, 203. Thermo-Chemistry of the Oxides and Acids of the Sulphur Group, 204. CONTENTS. Xlll Oxygen Compounds of the Elements of the Nitrogen Group, 204. Oxygen Derivatives of Nitrogen, 205. Nitric Acid, 206. Nitrogen Pentoxide, Nitrogen Trioxide, 208. Nitrous Acid, 209. Nitrogen Tetroxide, 210. Nitrosyl-sulphuric Acid, 211. Nitric Oxide, 213. Nitrous Oxide, 215. Hyponitrous Acid, 216. Thermo-Chemistry of the Oxides and Acids of Nitrogen, 216. Oxides of Phosphorus, 217. Hypophosphorous Acid, 218. Phosphorous Acid, 219. Phosphorus Trioxide, 219. Phosphoric Acid, 219. Pyrophosphoric Acid, 220. Metaphosphoric Acid, 221. Phosphorus Pent- oxide, 222. Chlor-anhydrides of the Acids of Phosphorus, 222. Phosphorus Compounds with Sulphur, 223. Oxides of Arsenic, 223. Arsenic Trioxide, 223. Arsenious Acid, Arsenic Acid, 224. Compounds of Arsenic with Sulphur, 225. Sulpho-Salts, 226. Oxygen Derivatives of Antimony, 226. Antimony Oxide, 226. Antimonio Acid, 227. Antimony Sulphides, 228. Thermo-Chemistry of the Acids of the Nitrogen Group, 229. Vanadium, Niobium, Tantalum, 229. Oxygen Compounds of the Elements of the Carbon Group, 230. Oxides of Carbon, 231. Carbon Dioxide, 231. Critical Pressure, 233. Carbon Monoxide, 235. Compounds of Carbon with Sulphur, 237. Cyanogen Compounds, 238. Thermo-Chemistry of the Carbon Compounds, 239. Oxygen Compoimds of Silicon, 239. Dialysis, 240. Crystalloids and Col¬ loids, 241. Silicates, 241. Titanium, Zirconium, Thorium, 241. Boron, 243. Boron Hydride, 244. Boron Fluoride, 244. Boric Acid, 245. The Periodic System of the Elements, 246. Periodicity of Chemical Valence, 252. Periodicity of Thermo-Chemical Phenomena, 254. Reduction of Metallic Oxides by Magnesium, 255. THE METALS. Physical Properties of the Metals, 257. Atomic Volumes, 258. Specific Heat, Atomic Heat, 260. Thermal Atomic Weights, 262. Isomorphism, 263. Chemical Properties of the Metals, 264. Alloys, 265. Halogen Compounds, 266. Oxides and Hydroxides, 266, Peroxides, 267. Salts, 268. Action of Metals on Salts and Acids, 270. Electrolysis of Salts, 271. Transposition of Salts, 275. Principle of the Greatest Heat Disengagement, 277, Group of the Alkali Metals, 277. Thermo-Chemistry of the Alkali Metals, 279. Potassium, 279. Potassium Hydride, 280. Dissociation, 280. Potassium Hydroxide, 281. Potassium Chloride, 282. Potassium Chlorate, 283. Potassium Hypochlorite, 284. xiv CONTENTS. Potassium Sulphate, 284. Potassimn Nitrate, 285. Gunpowder, 286. Potas¬ sium Carbonate, 286. Potassium Silicate, 287. Potassium Sulphides, 287. Recognition of the Potassium Compounds, 288. Rubidium, Caesium, 289. Sodium, 289. Sodium Hydroxide, 290. Soditun Chloride, 290. Sodium Sulphate, 292. Super-saturated Solutions, 293. Sodium Hyposulphite, 293. Sodium Carbonate, 294. Sodium Nitrate, Sodium Phosphates, 296. Borax, 297. Sodium Silicate, 298. Recognition of Sodium Compounds, 298. Lithium, 298. Ammonium Compounds, 299. Ammonium Chloride, Am- moniiun Nitrate, 300. Ammonimn Carbonate, 301. Ammonium Sulphide, 302. METALS OF GROUP II, 302. Group of the Alkaline Earths, 303. Calcium, 304. Calcium Oxide, 304. Cement, 305. Calcium Chloride, 306. Calcium Fluoride, Chloride of Lime, 306. Calcium Sulphate, 308. Calcium Phosphates, 308. Calcium Carbonate, 309. Glass, 310. Calcium Sulphides, 310. Strontium, 311. Barium, 312. Barium Oxide, 312. Barium Per¬ oxide, 312. Recognition of the Compounds of the Alkaline Earths, 313. Diammonium Compounds, 314. Magnesium Group, 314. M^nesium, 316. Magnesia, 316. Magnesium Chloride, 316. Magnesium Sulphate, 317, Magnesium Phosphates, 318. Magnesium Carbonate, 318. Recognition of Magnesium Compoimds, 319. Beryllium, 319. Zinc, 320. Zinc Oxide, 321. Zinc Sulphate, 321. Zinc Sulphide, 322. Cadmium, 322. Thermo-Chemistry of the Metals of Group II, 323. Mercury, 326. Mer- curous Compounds, 327. Mercuric Compounds, 329. Copper, Silver, Gold, 331. General Characteristics, 332. Forms of Combina¬ tion, 333. Copper, 334. Metallurgy of Copper, 334. Cuprous Compounds, 336, Cupric Compounds, 337. Copper Sulphate, 338. Alloys of Copper, 339. Silver, 340. Metallurgy, 340. Silver Oxide, 342. Molecular Formu¬ las, 342, Silver Chloride, 343. Photography, 344, Nitrate of Silver, 344. Silvering, 345. Gold, 346. Amous Compounds, 347. Auric Compounds, 347. METALS OF GROUP III, 348. Group of Earth Metals, 350. Aluminium, 350. Aluminium Chloride, 351. Aluminium Oxide, 352. Aluminates, 353. Alum, 356. Aluminium Silicates, 357. Porcelain, 357. Ultramarine, 357. Rare Earth Metals, 358. Scandium, 358. Yttrium, Lanthanum, Cerium, Ytterbium, 358. Didymium, Samarium, Holmium, Thulium, 359. CONTENTS. XV Gallium Group, 359. Gallium, 360. Indium, 360. Thallium, 361. Thallous Compoimds, 362. Thallic Compounds, 363. Germanium, 365. Germanous Compounds, 366. Germanic Compoimds, 366. Tin, 367. Stannous Compounds, 368. Stannic Compounds, 369. Stannates, 370. Sulpho-stannates, 370. Lead, 371. Lead Oxide, 372. Plumbic Acid, 373. Lead Sulphide, 374. Bismuth, 374. Bismuthic Acid, 375. Chromium Group, 376. Chromium, 377. Chromous Compounds, Chromic Compounds, 378, Chro¬ mates, 380. Chromic Acid Chlor-anhydrides, 383. Molybdenum, 385. Tungsten, 386. Uranium, 387. Manganese, 389. Forms of Combination, 389. Manganous Compounds, 390. Manganic Compounds, 391. The Acids of Manganese, 393. METALS OF GROUP VIII, 394. Iron Group, 396. Iron, 396. Metallurgy of Iron, 397. Ferrous Compoimds, 400. Ferric Compounds, 402. Ferric Acid Compounds, 402. Cyanogen Compounds, 403. Cobalt, 406. Cobalt-amine Compounds, 407. Cobalt-cyanogen Com¬ pounds, 408. Nickel, 408. Platinum Metals, 409. Ruthenium and Osmium, 411. Rhodium and Iridimn, 412. Palladium, 413. Platinum, 414. Spectrum Analysis, 416. Periodicity of the Spectrum Lines, 421. INDEX, 421. A TEXT-BOOK OF INORGANIC CHEMISTRY. INTRODUCTION. The study of Nature reveals an endless multitude of objects or bodies. That which forms the basis of the latter, strongly character¬ ized by extent and weight, we designate substance or matter. The investigation of the internal and external structure of bodies, their classification according to conformable or distinguishing character¬ istics, constitute the task of the descriptive sciences ; of mineralogy, of geology, of descriptive botany and zoology, of anatomy, etc. A closer scrutiny of natural objects discloses the fact that they in time succumb to many more or less serious alterations or changes. We observe that minerals form, crystallize, or disintegrate and crumble to pieces ; that plants and animals spring up, grow, and then fall into decay and decomposition. Such changes in the condition of bodies occurring with time are entitled phenomena. The investigation of these during their progression, the determination of the law's according to which they occur, the explanation of the causes under¬ lying them, form the task of the speculative sciences^ physics and chemistry—depending upon the nature of the phenomena. Like every other classifícation or definition, the division of the natural sciences into speculative and descriptive is not strictly correct. It does not completely cover the nature of the phenomena. We approximate the actual facts more closely by designating the natural sciences as general and special. The general sciences, mechanics, physics, and chemistry, occupy themselves with the study of the general properties and transformations of bodies, regardless of the external form, and deal chiefly with their substance only. Special branches—like botany and zoology— consider distinct classes of bodies, first in reference to their form (morphology, etc.), and afterward in relation to their transformations and alterations. Physi¬ ology of animals and plants, and geology, investigate the physical and chemical phenomena of particular classes of bodies, and are, therefore, speculative sciences. On the other hand, chemistry is also a descriptive science, inasmuch as it considers the external properties of chemical substances. 2 17 18 INORGANIC CHEMISTRY. Although no abrupt boundaries are presented in Nature, but gradual transitions and intermediate steps throughout, two tolerably distinct classes of phenomena may be observed. Some changes in the condition of bodies are only superficial (external), and are not accompanied by material alteration in substance. Thus heat con¬ verts water into steam, which upon subsequent cooling is again condensed to water, and at lower temperatures becomes ice. In these three conditions, the solid, liquid, and gaseous, the substance or the matter of water or ice is unchanged ; only the separation and the motion of the smallest particles—their states of aggregation—^are different. If we rub a glass rod with a piece of cloth, the glass acquires the property of attracting light objects, e. g., particles of paper. It becomes electrified. An iron rod allowed to remain sus¬ pended vertically for some time slowly acquires the power of attract¬ ing small pieces of iron. Through the earth's magnetism it has become magnetic. In both instances the glass and iron receive new properties ; in all other respects, in their external and internal form or condition, they have suffered no perceptible alteration ; the glass is glass, and the iron remains iron. All such changes in the condition of bodies, unaccompanied by any real alteration in substance, are known as physicalphenomena. Let us turn our attention now to the consideration of another class of phenomena. It is well known that ordinary iron under¬ goes a change, which we term rusting, i. e., it is transformed into a brown substance which is entirely different from iron. On mixing finely divided copper filings with flowers of sulphur (pulverulent sulphur) there results an apparently uniform, grayish-green powder. If this be examined, however, under a magnifying glass, we can very plainly distinguish the red metallic copper particles in it from the yellow of the sulphur ; by treating with water, the specifically lighter sulphur particles can easily be separated from those of the copper. Carbon disulphide will also dissolve out the sulphur par¬ ticles. Hence this powder represents nothing more than a mechan¬ ical mixture. If, however, this mixture be heated, e. g., in a glass test-tube, it will commence to glow, and on cooling, a black, fused mass remains, which differs in all respects from copper and sulphur, and even under the strongest microscope does not reveal the slightest trace of the latter, and elutriation with water or treating with carbon disulphide will not effect a separation of the ingredients. By the mutual action of sulphur and copper in presence of heat, a new body with entirely different properties has been produced, and is named copper sulphide. Mixtures of sulphur with iron or with other metals act in a similar manner ; and the resulting bodies are known as sulphides. Such mutual action of different bodies occurs not only under the INTRODUCTION. 19 influence of heat, but frequently at ordinary temperatures. If, e. g., mercury and sulphur are rubbed continuously in a mortar, there is produced a uniform, black compound, called mercury sulphide. The action of gaseous chlorine upon various metals is quite ener¬ getic. When finely divided antimony is shaken into a flask filled with yellow chlorine gas, flame is produced ; each antimony par¬ ticle burns in the chlorine with a bright white light. The pro¬ duct of this action of solid metallic antimony and gaseous yellow chlorine is a colorless, oily liquid, known as antimony chloride. Such occurrences, therefore, in which a complete and entire altera¬ tion takes place in the bodies entering the reaction, are termed chemical phenomena. Chemistry, then, is that department of natural science which occupies itself with the study of those phenomena in which an alteration of substance has occurred. In the previously described experiments we observed the phe¬ nomena of chemical combination; from two different bodies arose new homogeneous ones. The opposite may occur, consisting in the decomposition of compound bodies into two or more dissimilar ones. If red mercuric oxide be heated in a test-tube, it will dis¬ appear; a gas (oxygen) is liberated, which will inflame a mere spark on wood ; in addition, we find deposited upon the upper, cooler portions of the tube, globules of mercury. From this wc observe that on heating solid red mercuric oxide two different bodies arise ; gaseous oxygen and liquid mercury. We conclude, then, that mercuric oxide holds in itself, or consists of, two con¬ stituents—oxygen and mercury. This conclusion, arrived at by decomposition, or analysis, may be readily verified by combina¬ tion or synthesis. It is only necessary to heat mercury for some time, at a somewhat lower temperature than in the preceding ex¬ periment, in an atmosphere of oxygen, to have it absorb the latter and yield the compound we first used—red mercuric oxide. The direct decomposition of a compound body into its constituents by mere heat does not often happen. Generally, the cooperation of another substance is required, which will combine with one of the constituents and set the other free. In this manner we can, for example, effect the decomposition of the previously synthesized mercury sulphide, viz., by heating it with iron filings; the iron unites with the sulphur of the mercury sulphide, to form iron sulphide, while the mercury is set free. If, in a similar manner, natural objects be decomposed, bodies or substances are finally reached which have withstood all attempts to bring about their division into further constituents, and which cannot be formed by the union of others. Such substances are chemical elements ; they cannot be converted into each other, but constitute, as it were, the limit of chemical change. Their number. 20 INORGANIC CHEMISTRY. at present, is about 70 ; some have been only recently discovered. To them belong all the metals, of which iron, copper, lead, silver, and gold are examples. Other elements do not possess a metallic appearance, and are known as metalloids. It would be more correct to term them non-metals. To these belong sulphur, carbon, phos¬ phorus, oxygen, etc. The line between metals and non-metals is not very marked. When the elements unite with each other in smaller or larger numbers they produce the compound bodies known to us. Water is a compound of two gaseous elements—hydrogen and oxygen ; common salt consists of the metal sodium and the gas chlorine. The elements make up not only our own earth, but the heavenly bodies are composed of them ; at least as far as has been proved by spectrum analysis. THE PRINCIPLE OF THE INDESTRUCTIBILITY OF MATTER. If the quantities by weight of substance entering into a chemical change be determined, we notice that in all transpositions, in the decomposition of a compound into its constituents, and in the union of the elements to form compound bodies, loss in weight never occurs. The weight of the resulting compounds is invariably equal to the sum of the weights of the bodies entering the reaction. Well-known, general phenomena apparently contradict this scien¬ tific conclusion. We observe plants springing from a small germ and constantly acquiring weight and volume. This spontaneous increase of substance, however, is only seeming. Closer inspection proves conclusively that the growth of plants occurs only in conse¬ quence of the absorption of substance from the earth and atmos¬ phere. The opposite phenomenon is seen in the burning of com¬ bustible substances, where an apparent annihilation of matter takes place. But even in this, careful observation will discover that the combustion phenomena consist purely in a transformation of visible solid or liquid bodies into non-visible gases. Carbon and hydro¬ gen, the usual constituents of combustible substances, e.g.^ a candle, combine in their combustion with the oxygen of the air and yield gaseous products—the so-called carbon dioxide and water—which diffuse themselves in the atmosphere. If these products be collected, their weight will be found not less, but indeed greater, than that of the consumed body, and this is explained by the fact that in addition to the original weight they have had the oxygen of the air added. From what has been remarked we can conclude that in chemical transpositions loss in matter does not occur, nor is there a new INTRODUCTION. 21 creation of the same observed. Compounds are formed and dis¬ appear, because they are converted into new forms, but their sub¬ stance (matter), their weight, does not disappear, and is not produced anew. This fundamental truth is called the principle of the indestructibility of substance {matter). Lavoisier, in the eigh¬ teenth century, first established it by convincing experiments. Combined with the principle of the conservation of energy, it con¬ stitutes the firm foundation of all scientific knowledge. THE PRINCIPLE OF THE CONSERVATION OF ENERGY- CHEMICAL ENERGY. Causes underlie and influence all material phenomena. The final cause of phenomena we term force, accepting for the various sorts of phenomena a variety of forces. Some of these are attraction and repulsion, light, heat, electricity, cohesion, chemical affinity. These names, however, only represent kinds of phenomena, without explaining their true nature. Of the nature of some of these forces we know, positively, that they consist of various modes of motion of portions of matter. In the case of mechanical force it is obvious that it depends solely upon the motion of masses ; but other forces are nothing more than modes of motion. The phenomena of light are explained by the very rapid movements of the smallest particles, and these act upon the eye through the aid of a gaseous medium— ether. The phenomena of heat are due to the less rapid motion of weighable portions of matter which affect our sense of feeling. Accurate physical investigations have established that the different forces or modes of motion can never be destroyed, but only trans¬ ferred from one body to others, and changed from one kind to another. The movements or vibrations of one variety pass into those of another. For example, a discharged bullet is heated by coming in contact with any obstruction in its course; the visible motion of the entire mass in this instance is transformed into the invisible motions of the smallest particles, and appears as heat. The heat motions can, on the other hand, be again changed into mechanical motion (molecular motion), or into light, magnetism, or electricity. In all these transformations of the different modes of motion into one another, we observe a perfect equivalence of their quantity. If a mass motion, whose quantity is designated as mechanical work, can produce a certain degree of heat, so vice versa, the latter can perform the same mechanical work (the mechanical equivalent of heat, light, electricity). Upon this equivalence of transformation rests the principle of the conservation of force or energy, according to which the various forces or motions of matter can neither be anni- 22 INORGANIC CHEMISTRY. hilated nor produced anew. This principle, forming one of the most important corner-stones of natural science, was first sharply defined by the speculative observations of Dr. J. R. Mayer, of Heilbronn, in 1842, and since then has been repeatedly confirmed by experiment. The most recent advance in physics has led to the negation of the objective existence of all abstract physical forces. Not considering the phenomena of elec¬ tricity and those of chemical affinity—the reduction of which to forms of motion is clearly foreseen, and not to be doubted—the only remaining enigmatical force is that of attraction or gravity. To affirm the existence of gravity is nothing more than to give expression to the fact that bodies in space tend to approach each other. The supposition that the active cause of gravity existed within the bodies themselves, was long ago discarded by Newton as " absurdum ; " it is merely a mathematical fiction. The action of a body in aplace where it does not exist, with¬ out the aid of a medium, is inconceivable. The transference of the gravitation into material bodies, further contradicts the principle of conservation of energy, as gravity is neither transferred nor exhausted—whether it be through the approach of bodies, whereby the force always increases—or by planet movement, in which the centrifugal component is constantly overcome. Therefore, the active cause of gravity is not to be sought after in bodies themselves, but without them, and, indeed, in a substantial medium—ether—without the acceptance of which natural investigation cannot proceed. If we desire to make a preliminary presentation upon these relations, the follow¬ ing would be the simplest and most probable : Space is filled by the smallest pos¬ sible material particles, but as they are all alike, they do not possess gravity, and are found in constant transferable motion—e¿ker substance. By the congress of the smallest ether particles to mass-aggregates arise the chemical elementary atoms, which constitute material bodies—substance or matter. If, now, in addi¬ tion to this one mass-aggregate, a second appear in space, an effort to approach each other produced by the action (collision) of the disturbed ether surrounding them, will appear—they possess gravity. By these suppositions the obscure ideas upon potential energy and energy of place are removed. A much clearer and more distinct presentation and confirmation of these representations, especially as regards the nature of forces, may be found in A. Secchi's " Die Einheit der Naturkräfte." In the chemical union of bodies heat is almost invariably disen¬ gaged, and as it is a mode of motion, and as motion of one kind can only be derived from another, we must conclude that bodies acting chemically, especially the elements, do possess a peculiar kind of motion, which, in chemical union, is partially converted into heat motion (also into light and electricity). This special motion of matter is designated chemical energy or chemical tension. And in the chemical decomposition of a compound body into its constitu¬ ents, heat is absorbed, disappears as such, and is transformed into chemical energy. Thus, for instance, in the union of i kilogram of hydrogen with 8 kilograms of oxygen a quantity of heat is liber¬ ated which can perform a mechanical work equal to 34.462 X 423.5 INTRODUCTION. 23 = 14,629,000 kilogrammeters. In the decomposition, on the other hand, of 9 kilos of water into hydrogen and oxygen, the same force or quantity of heat is necessary. Therefore, in the liber¬ ated hydrogen and oxygen, the same quantity of force or motion must be contained in the form of chemical energy. Chemical energy is not only a quantitative phenomenon ; it also presents qualitative differences. Although all bodies, and particu¬ larly the elements, possess it, they do not disclose it in the same way in their action upon each other. Some unite or react readily with each other; others, on the contrary, with difficulty, or not at all. The reason for this deportment is to us entirely unknown, but it is in all probability due to the different form and mode of motion of the smallest particles of matter. We designate it by the phrase chemical affinity y and add that bodies capable of union have affinity for each other (are related), and that by union they satisfy their affinity. This expression is incorrectly chosen, because, generally, the bodies least alike chemically unite with each other most readily. CHEMICAL SYMBOLS AND FORMULAS. For simplicity and convenience the elements are represented by the first letters of their names, derived either from the Latin or Greek. Hydrogen is represented by the letter H, from the word hydrogenium ; nitrogen by N, from nitrogenium. When several elements happen to have the same letter there is added to the capital a second small letter; thus, Na represents natrium; Ni, nickel; Hg, mercury (hydrargyrum), etc. The subjoined table comprises all the elements known at present with certainty (69), together with their chemical symbols and atomic weights. The latter have been determined with more or less accuracy. 24 INORGANIC CHEMISTRY. Elements. Symbol. | Atomic Weight. Elements. Symbol. Atomic Weight. Aluminium, A1 27.08 Molybdenum, .... Mo 95-9 Antimony (Stibium), . Sb 120.3 Nitrogen, N 14.041 Arsenic, As 75 Nickel, Ni 59 Barium Ba 137 Niobium, Nb 94.2 Beryllium, Be 9.1 Osmium, Os 192 Bismuth, Bi 208 Oxygen, 0 16 Boron, Bromine B I I.Ol Palladium, Pd 106 Br 79-963 Phosphorus, P 31-03 Cadmium, Cd 112.1 Platinum, Pt 194.8 Caesium, Cs 132.9 Potassium (Kalium), . K 39 14 Calcium, Ca 40 Rhodium, Rh 103 Carbon, C 12 Rubidium, .... Rb 85.4 Cerium, Ce 140.2 Ruthenium, Ru 101.7 Chlorine, Cl 35-453 Samarium, Sa 150 Chromium, Cr 52.2 Scandium, Sc 44-1 Cobalt Co 59 Sulphur, S 32.06 Copper, ...... Cu 63-3 Selenium, Se 79.1 Didymium, Di 142.3 Silver (Argentum), . . Ag 107.66 Erbium, Er 166 Silicon, Si 28.4 Eluorine, Fl 19 Sodium (Natrium), . . Na 23.06 Gallium, Ga 699 Strontium, Sr 87-5 Germanium, .... Ge 72.3 Tantalum, Ta 183 Gold (Aurum), . . . Au 197.2 Tellurium, Te 125 Hydrogen, H 1.003 Thallium T1 204.1 Indium, In 113-7 Thorium, Th 232.4 Iodine, I 126.86 Tin (Stannum), . . . Sn 118.1 Iridium, Ir 193 2 Titanium, Ti 48.1 Iron (Ferrum), . . . Fe 56 Tungsten (Wolfram), . W 184 Lanthanum, La 138.5 Uranium, Ur 239-4 Lithium, Li 7-03 Vanadium, Yd 51-2 Lead (Plumbum), . . Pb 206.91 Ytterbium Yb 173-2 Magnesium, Mg 24.38 Yttrium, Y 88.7 Manganese, Mn 55 Zinc, Zn 65-5 Mercurv, Hg 200.4 Zirconium, Zr 90.7 Holmium and thulium (the element X of Soret) might also be mentioned, although the latest researches seem to indicate that they and also erbium, didymium and samarium are not simple substances, but rather mixtures of several elements. Compounds produced by the union of the elements are repre¬ sented by placing their corresponding symbols together and designat¬ ing these chemicalformulas. Common salt, a compound of sodium and chlorine, is represented by the formula NaCl \ mercuric oxide, a compound of mercury and oxygen, by HgO ; iron sulphide by FeS ; hypochlorous acid, a compound of chlorine, hydrogen and oxygen, by ClOH. These chemical formulas not only express the nature of the ele- INTRODUCTION. 25 ments, but also the relative proportions by weight, according to which they unite, compared with hydrogen as unity. Thus H represents i part by weight of hydrogen ; CI, 35.45 parts by weight of chlorine ; Na, 23.06 parts by weight of sodium. (See table, p. 24.) These numbers indicate the relative weights of the atoms consti¬ tuting the elements. If we seek to obtain a representation of the constitution of the elements and matter in general, two possibilities appear to be fore¬ most. Either the substance continuously fills space, or it consists of very small separate particles, chemical individuals, filling space, which are termed atoms. The latter idea alone corresponds to the present state of physical and chemical investigation, so that the atomic constitution of matter only has value at present. The inductive derivation and establishment of the atomic theory will be given subsequently (see page 69) ; here we will only state the following propositions: Each distinct element consists of similar atoms, of like size and similar weight, while atoms of different ele¬ ments possess a different weight. The absolute atomic weights are, at present, not determined with sufficient accuracy; the relative weights are referred to the hydrogen atom, which has the smallest weight, hence is made equal to i (H=. i). The chemical union of the atoms produces the smallest particles of compound bodies, termed molecules, physical individuals ; these are chemically divisible. By these premises chemical formulas acquire a very precise and evident importance. The formula NaCl represents the union of i atom of sodium (Na) with i atom of chlorine, and indicates that in it 23.06 parts, by weight, of sodium are combined with 35.45 parts of chlo¬ rine. If several atoms of an element are present in a compound, this is denoted by numbers which are attached to the symbol of the atom :— HCl Hp H3N CH^ Hydrochloric acid. Water. Ammonia. Methane. The formula of water (HjO) means that its molecule consists of 2 atoms of hydrogen (2 parts by weight) and i atom of oxygen (O = 16 parts by weight). The formula of sulphuric acid (H2SO4) indicates it to be a compound consisting of i atom of sulphur (32.06 parts), 4 atoms of oxygen (4 X 16 = 64 parts), and 2 atoms of hydrogen (2X1 = 2 parts), from which the com¬ position of the acid may be at once calculated into per cent., or into any desirable quantity by weight. Atomic Composition. In per cent. Sulphur, S = 32.06 32.69 Oxygen, = 64 65.26 Hydrogen, H2= 2 2.05 HjSO^ = 98.06 100 00 3 26 INORGANIC CHEMISTRY. The chemical union of bodies is shown by the sign -j-, and the resulting products are placed to the right, follo#ing the = sign :— HgS + Fe = reS4- Hg. By this equation of chemical transposition is meant that by the union of mercury sulphide (HgS) and iron (Fe), iron sulphide (FeS) and free mercury (Hg) are formed. The equation 2H +0 = HjO indicates that i molecule of water has béen formed by the union of 2 atoms of hydrogen with i atom of oxygen. At the same time such equations show the proportions by weight of the substances entering into and resulting from the reaction ; the weight of the acting substances is equal to that of those resulting. Therefore every chemical equation is at once an expression of the principle of the indestructibility of matter (substance). (See p. 21.) CONDITIONS OF CHEMICAL ACTION—THERMO-CHEMICAL PHENOMENA. The chemical changes sustained by matter, the decomposition of compounds into their constituents, or their conversion into other forms of like composition (into allotropie and isomeric modifica¬ tions, see ozone and phosphorus), are only effected through exter¬ nal agencies, e. g., by heat, electricity, or light, and also by mechanical action (compression, concussion). These agents are nothing more or less than modes of motion, and their action con¬ sists in imparting motion, which induces the chemical changes, or renders them possible. The mutual action of substances differing chemically—the formation of compound from simple bodies—occurs only when they are in intimate contact. It is only in physical changes that action at a distance is observed. The intimate con¬ tact requisite for chemical action with solids is generally not attaint- able by mere mechanical mixture. It is brought about by liquefying one or both of the solids, either by fusion or by solution in some solvent. In this way the physical cohesion of the molecules is diminished. The early chemists expressed these conditions by saying : Corpora non agunt nisi fluida. Liquid and gaseous bodies are therefore, eo ipso, adapted to chemical action. Besides intimate contact, some external physical action is neces¬ sary to effect the chemical union of different bodies, or at least to induce their reaction. The agents (heat, electricity, light) employed for the decomposition of compound substances will serve this pur¬ pose. We may explain this action by saying that the atoms chemi¬ cally combined with each other must often be loosened in order that they may react upon other atoms (see below). INTRODUCTION. 27 The real cause of the chemical transposition, of the union of the atoms with each other, is ascribed usually to a special attractive force, which may be called chemical affinity or attraction. We are, however, aware that all forces are referable to modes of motion : S^sides, the acceptance of chemical affinity is very uncertain and determined with difl&culty. At the ordinary temperatures hydro¬ gen and oxygen do not react ; there is apparently no affinity between them. But when they are raised to a temperature of 400°, or upon the passage of an electric spark, they unite to form water, and an explosion ensues. Again, a mixture of hydrogen and chlorine remains unaltered in the dark ; their union occurs very slowly in diffused light, but very rapidly and with explosion when exposed to sunlight or under the influence of an electric spark. The affinity or tendency to combine is consequently relative ; it is dependent on external conditions. Experiments have been made to ascertain how great this affinity or tendency might be ; the results were not concordant. The speed of reactions was studied with this in view. For the present, at least, it is therefore better not to attempt to define or estimate the magnitude of chemical affinity. Far better and more satisfactory results are afforded by the idea of energy of combination, and by the energy-content of the ele¬ ments and compounds (p. 22). Every chemical change is invariably accompanied by a change of energy—by the disengagement or absorption of heat (electricity, etc.). The customary chemical f3(|uations, such as are employed upon p. 26, represent merely the iüaterial side of a chemical reaction, the nature and quantities by weight of the reacting and resulting substances. But when, for example, hydrogen and oxygen unite to yield water, there is an ^companying dynamical change, a definite and considerable quan¬ tity of heat is disengaged—in this instance equaling 68,000 calories (p. 66). An equation showing the union of hydrogen with oxygen to form water—an equation which would include both weight reac¬ tions and those of energy—would read as follows :— 2H -|- O = HjO "h 68,000 Cal. In the union of hydrogen with chlorine, forming hydrogen chloride, we have a liberation of 22,000 calories :— H -f- CI =HC1 -j- 22,000 Cal. ; while the formation of hydrogen iodide from hydrogen and iodine is accompanied by an absorption of energy—an absorption of 7,000 calories :— H -|- I ==: HI — 7,000 Cal. 28 INORGANIC CHEMISTRY. When we desire to decompose water into hydrogen and oxygen we must restore all the energy which has escaped (68,000 caL). It is contained in the free elements as chemical (potential) energy. The decomposition of hydrogen iodide, on the contrary, would occur with a disengagement of energy (of 7,000 cal.). These alterations of energy, consisting in the liberation and absorption of heat (electricity, etc.), constitute an important char¬ acteristic of every chemical transposition, because they afford real explanations for the course of a reaction, the mode of formation, the stability, and entire character of the resulting compounds. Reactions are distinguished as exothermic and endothermic^ de¬ pending upon whether heat (energy) is liberated or absorbed. The union of hydrogen with oxygen to form water is an exo¬ thermic reaction (see above), and water is an exothermic compound. All such exothermic compounds possess less energy than their component elements; they are, consequently, more stable than these, and can only be converted into them by the addition of energy. The conditions of their formation are present in the components ; for this reason they may be designated direct com¬ pounds. Exothermic reactions frequently take place immediately upon contact of the components, but usually the reaction must be induced ; there must be some external impulse (by heat, electricity, etc.) given, which may be regarded as a liberation of the chemical energy, and is made manifest by the fact that the molecules of the compounds and of the elements must first be torn asunder into their component atoms. When the exothermic reaction has been induced in a part of the mixture, it will proceed farther of itself, corresponding with the quantity of heat liberated, and will advance with greater or less rapidity, and may even rise to an explosion, e. g., the formation of water or hydrogen chloride from their elements (p. 67). The endothermic reactions, e.g., production of hydrogen iodide from iodine and hydrogen, proceed with the absorption of heat (energy). Endothermic compounds, e.g., hydrogen iodide and nitrogen chloride, contain more energy than their components, and are, therefore, less stable than the latter, into which they pass with loss of energy. The conditions of their formation are not present in their components, but must be added from without, and on this account they may be designated as indirect compounds. Accordingly, endothermic compounds do not result by the imme¬ diate contact of their components. To effect their union there must be a constant addition of energy, otherwise the reaction will cease. The decomposition of compounds into their elements or com- INTRODUCTION. 29 ponents proceeds in a manner exactly opposite to that displayed in their formation from the same components. Exothermic com¬ pounds, possessing less energy than their components, can only be separated into the latter by adding to them the entire quantity of heat which was developed in their formation. The decomposition of an exothermic compound, like that of water into hydrogen and oxygen, is, consequently, an endoihermic decomposition. It proceeds in a manner similar to that of an endothermic compound. It requires the constant addition of energy, proceeds gradually, is never accompanied with explosion, and is limited by the opposite tendency of the components to unite (hydrogen and oxygen, to yield water). Compare the dissociation of water. The endothermic compounds, on the other hand, e.g.^ hydrogen iodide and nitrogen iodide, containing more energy than their components, separate easily and completely into the latter. An exothermic decomposition such as this proceeds in a manner similar to that of an exothermic compound. It only requires an external impulse when it will proceed alone, sometimes with explosion—just as hydrogen unites with oxygen and chlorine, with explosion. Endothermic compounds are, therefore, generally explosive bodies. Some, like chlorine oxide and nitrogen iodide, explode when touched or when warmed, others require a powerful blow. For example, nitric oxide, acetylene, and cyanogen decompose with explosion when a slight quantity of fulminating mercury is ignited in them (Berthelot). It follows, from what has been recorded above, that exothermic reactions, liberating heat, can proceed of themselves, while the endothermic always require additional energy. We may then deduce this principle, that when several bodies react with each other without the, addition of external energy, the reaction must proceed in the direction in which there is the greatest liberation of heat—or, that from a given system of bodies that one must result which contains the least energy, and in the formation of which the largest amount of heat has been developed. This is the principle of the greatest development of heat. This principle may be considered as a special case of the principle of the mechanical theory of heat resulting from the degradation (dissipation), or entropy of energy, according to which every form of energy tends to pass into heat-energy. It is expressed in the principle of mechanics, that an isolated system of bodies will be in equilibrium if all the potential is converted into kinetic energy. It is due to Berthelot that this principle, in its generality, has been applied to chemical phe¬ nomena and conñrmed. This fundamental principle of the greatest heat development in no manner proves that the chemical reactions, when in presence of 30 INORGANIC CHEMISTRY. an external energy (heat), must invariably proceed in the direction of the heat disengagement, nor does it prove that only those bodies are formed in which there is less energy—in other words, those that have given off the greatest amount of heat. Heat separates com¬ pounds into their components, and this opposes the tendency on the part of the elementary atoms to unite. Solvents occasionally act in a manner similar to heat. At a certain temperature every compound is split into its elements (see Dissociation of Water). At absolute zero (—273°) heat motion does not occur ; at this point, then, only exothermic reactions take place. The reaction is endothermic at the temperature at which a compound is fully decomposed. As the temperatures of decompo¬ sition are mostly very high (1000-3000°), it is evident, that at ordinary temper¬ atures most reactions belong to the exothermic class. We must not omit the circumstance, that chemical substances do not consist of free, but of atomic-aggregations (molecules), which, in chemical changes, must first be separated into atoms, and further, that the stability of compounds (their decomposition by heat or solvents) does not always correspond to their heat of forma¬ tion (p. 68). The principle of the greatest development of heat may, therefore, be formulated into the principle of maximum work- power (F. Braun). It is embraced in the following sentence : In chemical changes the tendency is toward the production of that body or aggregation of bodies, in whose formation there is the greatest development of heat. This principle is important in the representation and explanation of chemical changes, and will be applied in the several groups of elements which follow. CRYSTALLOGRAPHY. Chemistry occupies itself chiefiy with the investigation of the chemical alterations of bodies. Its subject is not the latter in them¬ selves, in their external properties, but only with reference to their material composition, and their genetic relations to other substances. The investigation of the physical properties of the non-organized bodies constitutes the province of Mineralogy, or, if the same is not limited to naturally occurring bodies, but includes also the innumerable substances which have been prepared artificially, it becomes the province of Inorganography. Pure chemistry considers the physical properties only as far as they serve for the external characterization and eventual recognition of the given substances and for the deduction of chemical laws. The most important physical properties,—the state of aggregation, the temperature of fusion and boiling, the specific gravity, capacity for heat, etc.— are partially treated in Physics, and in part will be considered later, in special cases. INTRODUCTION. 31 Here, therefore, the morphological characters of the solid bodies will receive only a brief consideration. The homogeneous solids exhibit either similar properties in all their parts, are amorphous^ or show differences in certain definite directions, giving rise to a crystalline appearance. The cause of this deportment lies in the arrangement of the smallest particles of substance of the molecules, which in the first instance is irregular, hence cannot cause differences in any direction ; while in the crys¬ talline structure the molecules are regularly grouped according to directions of varying density and coherence, which find expression in the cleavage and the optical and thermal behavior of bodies. A consequence of this regular arrangement is, in the case of undis¬ turbed formation, the external limitation of bodies by planes, edges, and angles, which represent the crystal form. The number and forms of these crystal elements are very numerous, since several thousands are known. It is, however, possible to reduce the num¬ berless varieties to a few classes or systems^ by comparing their modes of formation, and by referring their principal elements—the planes —to definite axes, i. e., to directions or lines, which are imagined to be so placed through the middle point of the crystals that their planes lie symmetrically with reference to them. In this manner, by distinguishing the various axis-intersections, crystallography arrives at the following six systems of crystallization :— 1. The Regular, or Tesseral System. 2. Quadratic, or Tetragonal System. 3. Rhombic System. 4. Hexagonal, or Rhombohedral System. 5. Monoclinic System. 6. Triclinic System. The position of any even plane in space is determined by three points of a system of coordinates, and therefore the position of a crystal face is also determined by its points of intersection with the three axes, or, by the distances from the centre of the axis at which the plane (by suitable expansion) cuts or intersects the three axes. The distances are termed the parameters of the plane. In the regular system all three axes are of equal length and equal value, for which reason they are designated by the same letter, a (Fig. i). If a plane intersect all three axes at equal distances (as in the octahedron) the parameter ratio is a: a: a; if at unequal distances, the ratio is a : a : ma, or a : ma : na ; in which case the parameter of the first axis is always made = i. If a plane lie parallel to an axis intersecting it at infinity, its parameter with reference to this axis = infinity (a : a : 00 a) ; if it be parallel to two axes, two para¬ meters are = infinity (a : 00 a ; 00 a). Hence these parameter ratios 32 INORGANIC CHEMISTRY. designate the position of a plane ; and as all the planes in simple crystal forms are similar, these symbols represent the entire simple form, i. - ings of points or molecules in space, i. e., from such groupings in which about each point the arrangement of the remaining points is the same. In this way it is possible to embrace the point positions which can occur in space into seven groups, corresponding to the seven crystallographic systems. The rhombohedral forms, which crystallography views as the hemihedral of the hexagonal, then appear as an independent system. I. Regular System.—The crystals of this system are similarly developed in three directions, and therefore have three similar axes, a, of the same length, which intersect each other at right angles (Fig. i). If we imagine the faces of the crystal arranged in such manner that they intersect the three axes at similar or dis¬ similar distances, we have the following seven simple forms for the various possible parameter ratios (page 31): I. The octahedron (Fig. 2), with the parameter ratio a : a : a; for which the abbreviated symbol O may be written. 2. The cube, or hexahedron (Fig. 3), with the symbol a : 00 a : 00 a or briefly 00 O 00. 3. The rhombic dodecahedron, twelve sides, a : a : 00 a or 00 O (Fig. 4). 4. The trigonal tris-octahedron (Fig. 5), a ; a ; INTRODUCTION. 33 ma or mO ; there are several of these, as m can equal f, 2, or 3. 5. The trapezohedron (Fig. 6), a : ma : ma or mOm, in which ra= 2 (in garnet) or m = 3 (in ammonium chloride). 6. The tet- Fig. a a cc a ct Fig. 3. 1 1 1 ooO oo 1 1 1 X y rahexahedron (Fig. 7), a : ma ; 00 a or mO 00, in which m = f, 2, or 3. 7. The hexoctahedron (Fig. 8), a : ma : na or mOn, where, e. g., in fluorite, m = 2 and n = 3. Fig. 4. Fig. 5. These seven simple forms may appear combined on one crystal, and we thus obtain crystal combinations. For example, Fig. q represents a combination of octahedron and cube (on alum) ; Fig. Fig. 7. Fig. 8. Fig. 9. ID, a combination of octahedron, cube, and dodecahedron (gale- nite). In addition to these simple forms, appearing with their full 34 INORGANIC CHEMISTRY. number of planes, hence termed holohedral, others are found having only the half or fourth of the possible faces, hemihedral ox Fig. to. Fig. Fig. 12. tetartohedral forms. We may suppose them produced by the enlargement of the symmetrically distributed half number of faces of the holohedral form, and therefore they are marked with the corresponding holohedral symbol divided by 2. Such hemihedral forms are: The tetrahedron (Fig. Ti), resulting from the octahedron, Fig. 13. Ç and the pentagonal dodecahedron m Q OS (Fig. c' 12) derived from the tetrahexahedron. ^ 2. Quadratic System.—Crystals of this system are developed at right angles in three directions, of which two are alike, the third unlike the others. Therefore, they possess two equal secondary axes, a, and an unequal, longer or shorter, principal axis, c (Fig. 13). The axis ratio, therefore, is a : a : c j the ratio of a : c is definite but not rational for every body of the system (see page 31), e.g.^ for tin I : 0.3857. The principal (ground) form* of this system is the quadratic pyramid a : a : c, with the symbol P, an obtuse pyramid (Fig. 14) if c be snialler, or an acute pyramid if c is greater than a. From these pyramids of the first order, in which the axes pass through the angles, we distinguish pyramids of the second order, in which the two secondary axes pass through the middle of the edges. The planes of the latter are parallel to one secondary axis, there¬ fore its symbol is a : 00a : c, or Poo . Fig. 15 represents a com¬ bination of a pyramid of the first and second orders. In addition to the above pyramids, others may occur upon the same crystal, which intersect the principal axis at the distance mc. In this case m is also a simple rational number, e.g., 3, 2, etc. The symbol of such secondary * The principal (ground) form is that to which the remaining forms of the same mineral may be most readily referred. INTRODUCTION. 35 pyramids of first order, is a : a : mc = mP : that of the second order, a : oo a : mo = mPoo . The coefficient of the principal axis is always written before P. If the pyramid planes intersect the principal axis at infinity, the quadratic prism results, and, indeed, either a prism of the first order, with the symbol a : a : oo c = Poo ; or a prism of second order, a : oo a : oo c = ooPoo. Fig. i6 represents a combination Fig. 14. FIG. 15. FIG. 16. of the quadratic prism with the quadratic pyramid, as seen on zircon and potassium phosphate. With the parameter ratio of the planes, a : ma ; nc, we get the ditetragonal pyramid nPm. Its corresponding ditetragonal prism has the symbol a : ma : oo c = oo Pm. Different hemihedral forms are possible in this system, of which P . may be mentioned the tetragonal sphenoid ± j' corresponding to the tetrahedron. Chalcopyrite, tin, potassium phosphate, etc., crys¬ tallize in this system. Fig. 17. Fig. 18. Fig. ig. 3. Hexagonal System.—The forms of this system, like those of the preceding, exhibit one peculiarly striking direction of development, and hence this is chosen as the direction of the prin¬ cipal axis c. They are distinguished from the four-sided forms of the quadratic system by their sixfold symmetry, which finds expres- 0 36 INORGANIC CHEMISTRY. sion in their equal secondary axes (Fig. 17), intersecting each other at 60°. The principal axis is at right angles to these. The axis ratio is a : a ( : a) c, and the ratio of a : c for every substance is definite, but not rational ; e.g., in quartz, i : i.ioo ; calcite, i : 0.8543, etc. The fundamental (ground) form of the system is the hexagonal pyramid a : a : (00 a) : c — P, from which is derived the hexagonal prisma :a: (00a): ooc = ooP; and, indeed, as in the quadratic system, there are pyramids and prisms of first and second order—- the latter with the symbol a : 2a : (2a) : c = P2, and a : 2a : (2a) : (xc = 00 P2. Further, other pyramids can occur, intersecting the principal axis, at the distance mc ; their symbol is mP and mPa. Fig. 18 represents the combination of pyramids and prisms found on apatite. With the common parameter ratio a : na : (ra) : mc (in which the parameter of the third secondary axis, ra, is always given as in all other hexagonal forms, by the parameters of the first two secondary axes), result the dihexagonal pyramid, mPn, and the dihexagonal prism, 00 Pn. Therhombohedra±^=± mR(Fig. 19) are the hemihedral forms of the pyramids mP, produced by the growth of the alternate faces. Another important hemihedral form is the scalenohedron derived from the dihexagonal pyramid. It is a remarkable fact that the hemihedral forms of the hexagonal system occur much more frequently in nature in numberless combinations (especially in calcite), and they are, therefore, sometimes treated as a separate system. 4. Rhombic System.—Three axes of unequal length, a. b. c., at right angles to each other. Any one, as c, is chosen as principal axis, and of the secondary axes, the shorter a is designated as the brachydiagonal, the longer b as macrodiagonal (Fig. 20). The INTRODUCTION. 37 Fig. 23. axis ratio a : b : c is definite for every substance, enings near their bottom, in which are glass tubes, are connected by a rubber tube. The bottle A is filled with granulated zinc, and B with dilute sulphuric acid. The cock R closes A. When this is opened the sulphuric acid flows firom B to A, to the zinc, and the evolution of gas commences. On closing the stop-cock the hydrogen presses the acid back, from A to B ; the evolution of gas ceases. By elevating and sinking the flasks the regulation can be hastened. Gasometers of various construction serve to collect and preserve gases. In Fig. 31 we have the ordinary gasometer of Pepys. It is constructed from sheet copper or zinc, and consists of two cylindrical vessels, the lower one closed, the upper open, communicating with each other by the two tubes a and b. The tubes c and e are only supports. To collect gases in this apparatus it must first be filled with water. To this end, pour water into the upper cylinder, and open a and e ; the water then flows through a, nearly reaching the bottom of the lower cylinder, while the air escapes through e. The side glass tube f allows the operator to observe the HYDROGEN. 43 height of the water-level. When the lower cylinder is filled with water close a and e (the last traces of air can be removed by opening b). To fill the gasometer with gas, remove the cover of the side tubulure d, and introduce the tube from which the gas is escaping. The latter rushes up into the cylinder, while the water flows out the tubulure. When the water is displaced by the gas, close d, after filling the upper cylinder, and then, if desired, open a, and the gas can be set free, either by c or b. In addition to the gasometer described, various other forms are employed ; gas¬ bags are very well adapted for preserving gases. Physical Properties.—Hydrogen is a colorless, odorless, and taste¬ less gas. It has a metallic character, and in accord with this it conducts heat and electricity better than all other gases. This may be proved by the following experiment: A current of electricity Fig. 30. is sent through a thin platinum spiral, and while the latter remains in the air or some other gas, it will glow, but in an atmosphere of hydrogen the spiral will not become luminous, or at once cease glowing. Of all gases hydrogen is the most difficult to liquefy, because its critical temperature is the lowest (about—180°) ; it must, therefore, be exposed to the most intense cold (p. 48). Pictet subjected it to a pressure of 650 atmospheres and a temperature of—150°. When the cock of the compression apparatus was opened (p. 48) the hydrogen escaped as a steel-blue liquid, which evaporated rapidly. Cailletet and Wroblew.sky, employing the compression method of Cailletet, which is that of the rapid expansion of a strongly compressed gas (p. 48), only succeeded in obtaining the hydrogen in the form of a gray mist. Olczewsky, however, main- 44 INORGANIC CHEMISTRY. Fig. 31. tains that compressed hydrogen, in a mixture with compressed oxygen, is a colorless liquid. Like all gases, coercible with difficulty, hydrogen is but slightly soluble in water, 100 volumes dissolving 1.9 volumes H at 0°—20° C. The coefficient of absorption of hydrogen by water is therefore 0.0193. 's the lightest of all gases. A cubic decimetre (= i litre) of hydrogen, according to Regnault's determinations, weighs 0.089567 grams at 0° and an atmospheric pressure of. 760 millimeters (at Paris), while a litre of air, under similar conditions, weighs 1.29318 grams. Hydrogen is, therefore, 14.43 times lighter than air. Its specific gravity, or, more correctly, its gas density com¬ pared with air as unity, is = 0.06928. Hydrogen has been selected as unit in the determination of the spe¬ cific gravity of gases, because it possesses the smallest den¬ sity, and because it is well adapted thereto for chemical Fig. 32. reasons. If the specific gravity of gases compared with H = i be represented by A, and the specific gravity compared with air = i by D, then A = D x 14.43 ^.nd JD = That hydrogen is lighter than air is shown by a balloon of col¬ lodion or gum filled with the former rising in the latter; this can also be seen in soap bubbles filled with hydrogen. In consequence of its levity, hydrogen may be collected in inverted vessels (opening turned down) by replacing the air, and can also be poured from one cylinder into another, as is represented in Fig. 32. The hydrogen fiows from the inclined cylinder into the one held vertically and HYDROGEN. 45 filled with air, which it expels. Such a separation of gases, based on their varying specific gravity, is only temporary, as they soon mingle with each other by diffusion. By virtue of its levity and mobility, which the kinetic gas-theory attributes to the great velocity of the gas particles, hydrogen penetrates porous bodies with ease, and diffuses through both animal and vegetable membranes, as well as through gutta-percha. Consult air, upon diffusion of gases. Metals, e. ^., iron, platinum, palladium, permit a free passage to hydrogen, when they are raised to a red heat. They are impene¬ trable to other gases. This behavior is in part probably dependent upon the chemical attraction of these metals for hydrogen. Chemical Properties.—Hydrogen is characterized by its ability to burn in the air, when it combines with the oxygen of the latter and forms water ; hence its name hydrogenium (from udwp, water, and yewdto, I produce). sa¬ lts flame is faint blue, and almost non-lumin¬ ous, but possesses a very high temperature. When a mixture of hydrogen and air is ignited a violent explosion ensues ; there¬ fore, before bringing a light in the vicinity of hydrogen disengaged in a vessel filled with air, allow the latter to fully escape, otherwise the vessel will be shattered to pieces by the explosion. As hydrogen itself is inflammable, it can¬ not sustain the combustion of other bodies which will burn in the air. If a burning candle be introduced into an inverted cyl¬ inder containing an atmosphere of the gas (Fig. 33) the latter will ignite at the mouth of the vessel, but the candle will be extinguished. Water is the product of the combustion of hydrogen in the air. It is a chemical compound containing hydrogen and oxygen. To render the formation of it visible, by the combustion of hydrogen, the flame of the latter is made to burn under a cold glass jar (Fig. 34). The sides of the latter are soon covered with moisture, which collects in drops of water. To avoid any deception the hydrogen is first conducted through sulphuric acid or a tube filled with calcium chloride, to absorb all moisture. The union of hydrogen with oxygen to form water occurs only at high temperatures: at a red heat, in contact with a flame, or by the passage of an electric spaik through the mixture. The com¬ bination can be effected at ordinary temperatures with the aid of platinum sponge; the latter consists of finely divided metal, obtained 46 INORGANIC CHEMISTRY. by the ignition of ammonio-platinum chloride (see Platinum). If a stream of hydrogen be directed upon a piece of freshly ignited platinum sponge, the gas will at once ignite. This is due to the Fig. 34. Fig. 35. power of the metal to condense hydrogen and oxygen upon its surface, and thereby increase their ability to unite. Upon this behavior depends the action of the so-called Dœbereiner Lamp (Fig. 35). This is a continuous hydrogen generator. The outer glass cylinder is filled with dilute sulphuric acid, into which projects the pear-shaped vessel b. This is open below and above, and is provided with a stop¬ cock, e, aifording communication with the air ; in it a piece of zinc is suspended by a wire. On opening the stop-cock the sulphuric acid presses from the outer cylinder a into b and meets the zinc—when the liberation of hydro¬ gen begins. The stop-cock directs the gas upon the support, y, in which is fixed some platinum sponge that effects the ignition. Upon again closing e the gas causes the acid to re¬ cede from the inner vessel, the zinc is freed of acid, and the hydrogen evolution ceases. The absorption of hydrogen by the metal palladium is very characteristic. As already known, water is so decom« posed by the electric current that hy¬ drogen separates at the electro-nega¬ tive pole and oxygen at the electro¬ positive. Now, if a piece of palladium, in sheet or wire form, be CONDENSATION OF GASES. 47 attached to the electro-negative pole, the disengagement of hydrogen does not occur, because it is absorbed by the palladium, in a quantity over nine hundred times the volume of the latter. Palladium also absorbs hydrogen when it is heated to ioo°. The palladium expands, becomes lighter in weight, but retains its metallic appearance. Its tenacity and power of conducting heat and electricity are but little impaired. The compound of palladium and hydrogen, PdjH, therefore, conducts itself like an alloy of two metals. From the specific gravity of the compound (according to Graham), the specific gravity of the condensed hydrogen is really found tobe 0.62 (water = i), and is, therefore, somewhat heavier than the metal lithium. The metals potassium and sodium absorb hydrogen when heated from 200° to 400°, forming alloys (NaaH and K2H) in which the density of hydrogen is again equal to 0.62. These facts prove the metallic character of hydrogen, which is also indicated by its ability to conduct heat and electricity. According to Pictet liquid hydrogen has a metallic appearance (p. 43). Later, we will observe that this element displays the character of a metal in its entire chemicaf deportment, and that it must be regarded as a gaseous metal at ordinary temperatures. CONDENSATION OF GASES. CRITICAL CONDITION. Until recently hydrogen and several other gases (oxygen, nitro¬ gen, carbonous oxide, methane, nitric oxide), were considered as non-condensable—permanent gases, inasmuch as all attempts to liquefy the same were failures, notwithstanding Natterer (1852) had employed a pressure of 3,600 atmospheres for this purpose. These negative results find their explanation in a general property of gases, first recognized by Andrews (1871), and called by him the critical condition of matter. There is a temperature common to all gases, above which they cannot be condensed—this is the critical tempera¬ ture. It was first observed with carbon dioxide (see this). Again, Caignard de la Tour (1822) showed that all liquids when heated above a certain temperature (the same critical temperature), would be transformed into gases (absolute boiling point of Mendelejeff), although they were subjected to intense pressure (in closed tubes). The pressure exerted by the gas at the critical temperature (at which it would immediately condense upon lowering the tempera¬ ture) is called the critical pressure \ the volume filled by the sub¬ stance at this time is the critical volume. An explanation of the existence of the critical condition is afforded by the consideration, that at the critical temperature and 48 INORGANIC CHEMISTRY. pressure the volume of the gas (or saturated vapor) is equal to the volume of an equal quantity, by weight, of the liquid—there exists no longer a difference between the gaseous and liquid condition. That the gases may be liquefied, we need not only a pressure, but also a definite temperature, and this must be lower than the critical. By this means Cailletetand Bietet (1879) succeeded in condensing nearly all the permanent gases. Bietet pursued the method of Faraday, who had condensed various gases in sealed tubes (see con¬ densation of chlorine (p. 51). The gases were generated in a powerful, iron retort (oxygen from potassium chlorate ; hydrogen from sodium formate) by the application of heat. They were then allowed to be compressed under their own pressure in a copper tube attached to the retort. Solid carbon dioxide surrounded the tube, and by its evaporation under the air pump its temperature was reduced to —140° C. On opening the cock of the copper tube, the liquefied gas escaped in a stream which rapidly evaporated. Cailletet employed a capillary glass tube, provided with a reservoir, and a pressure pump. The strongly compressed gas was cooled by opening a cock and permitting it to expand suddenly. In its expan¬ sion and in overcoming the external pressure it performs work and there follows an absorption of an appreciable quantity of heat, which is taken from the gas. This causes a partial liquefaction of the gas in the form of a dense cloud, or in small drops. It is better to liquefy the gas compressed in a glass tube by exter¬ nal cooling. This may be effected by vaporizing solid CO2, or liquid ethylene under an air pump, when the temperature of the latter will fall to —150° at a pressure of 10 mm. In this way Wroblewsky and Olczewsky, and also J. Dewar (1883) suc¬ ceeded in obtaining oxygen, carbon monoxide and nitrogen in the liquid form (static liquids). Lower temperatures, necessary for the liquefaction of hydrogen gas, may be obtained by the vaporiza¬ tion of liquid oxygen or nitrogen. The first of these boils under a pressure of 9 mm. at —211.4°, the second at —255° with a pressure of 4 mm. Temperatures lower than these can only be reached by the evaporation of liquid hydrogen. The critical temperatures (T) and critical pressures in atmos¬ pheres (B) of the gases condénsed with difficulty are as follows ;— T + 31° + 9-2° P Carbon dioxide, CO^ . . . . Ethylene, . . . Nitric Oxide, NO Marsh gas, CH^ . . Oxygen, O2 Carbon Monoxide, CO . . . Nitrogen, Nj Hydrogen, 73.6 atm 58 71 — 113° — 141° — 146® 55 50 35 33 180° (about) 99 (calculated). HALOGENS. 49 The lowest temperatures are ascertained by means of a hydrogen thermometer or a thermo-electric element composed of copper and German silver. The critical temperature, pressure and volume can be determined not only experimentally, but may also be deduced from the variations of the gases from the laws of Boyle and Gay-Lussac, by the following equation of van der Waals :— (P + ^) (v— b) = (l -f a) (i -f- b) (i -f at). From this formula the values of the critical volume (V), of the critical pressure (P), and the critical temperature are found to be V = 3b P = TO I + Í (approximately), in which a and b represent the constants of the above equation. HALOGEN GROUP. To this group belong chlorine, bromine, iodine, and fluorine. These elements show a similar chemical deportment. They are termed halogens or salt producers, because by their direct union with the metals salt-like derivatives result. I. CHLORINE. Cl — 35.453- Clj = 70.906. It does not occur free in nature, as it is endowed with strong affinity for the majority of the elements. Its most important derivative is sodium chloride, or rock salt, which is composed of chlorine and sodium. The Swedish chemist. Scheele, discovered chlorine in 1774. Its elementary character was first established by Gay-Lussac and Thénard in France (1809), and by Davy in England (1810). Preparation.—To obtain free chlorine, heat a mixture of black oxide of manganese (MnOj) and hydrochloric acid in a flask (Fig. 36), provided with a so-called Welter's safety-tube to equalize the gas pressure. The escaping gaseous chlorine is washed and freed from acid that is carried along mechanically by passing it through water in a three-necked Woulff's bottle, and then collecting it over water. The reaction which occurs above is indicated in the follow¬ ing equation :— MnOj -f- 4 HCl = MnCla + CI^ -f aHjO. The manganous chloride formed dissolves in the water. The evolution of the chlorine proceeds more regularly if a mixture of man¬ ganese oxide (5 parts), sodium chloride (4 parts) and sulphuric acid (12 parts diluted with 6 of water) is employed :— 5 50 INORGANIC CHEMISTRY. MnOj + 2NaCl + 2H2S04= MnSO^ + Na^SO^ + Clj + 2Hp. Manganese Sodium Sulphuric dioxide. chloride. acid. This reaction comprises two phases: First, the sodium chloride (NaCl) is decomposed by the sulphuric acid, yielding sodium sulphate and hydrochloric acid :— 2NaCl + UßO^ = NagSO^ -f 2HCI. The latter acid then acts, together with a new portion of sulphuric acid, upon the manganese dioxide :— MnOg + HjSO^ 4- 2HCI = MnSO^ + 2H2O + Cl^. The second method is more advantageous for laboratory purposes ; the firstp however, is preferred in practice, as it is cheaper. The resulting manganous chloride (MnClj) is converted by the Process of Weldon into manganese peroxide (see this). Technically, chlorine is also obtained by the Process of Deacon, by conducting HCl mixed with air over strongly ignited porous substances (bricks) saturated with metallic salts (copper sulphate). An excellent laboratory method for the preparation of chlorine consists in allow¬ ing dilute hydrochloric acid to act upon bleaching lime (see this). The latter is previously mixed with burnt gypsum pt.), and a little water added, when the mass can be formed into cubes or stout sticks which are introduced into a Kipp generator (p. 42) (Winkler, Berichte, 20, 184). As chlorine gas dissolves readily in cold water it is àdvisable to collect it over warm. It cannot be collected over mercury, as it readily combines with the latter. When perfectly dry chlorine is sought, conduct the liberated gas through WoulflPs bottles containing sulphuric acid, to absorb the moisture, then collect in an empty upright flask (compare Fig. 44, p. 63). As chlorine is so much heavier than air it will displace the latter. Physical Properties.—Chlorine isa yellowish-green gas (hence its name from yXu)pbossess; it does, however, very probably belong to that which is in the ad of forming,—of becoming free. We will learn, later, that many other elements, at the moment of their birth {in statu nascendi), act more energetically than when free; the cause for this will be explained hereafter. 3. BROMINE. Br = 79.963. Br, = 159.926. Bromine, the perfect analogue of chlorine, was discovered by Balard, in 1826. It occurs in the sea water as sodium bromide, accompanied by sodium chloride, but in much smaller quantity than the latter (especially in the water of the Dead Sea), and in many salt springs, as at Kreutznach and in Hall. When sea water or other salt water is evaporated, sodium chloride first separates ; in the mother-liquor, among other soluble salts, are found sodium and magnesium bromides. Bromine is found in greatest abundance in the upper layers of the rock-salt deposits of Stassfurth, near Mag¬ deburg, where it exists in the form of bromides together with other salts. At present, large quantities of bromine are obtained in America. The method of its preparation is similar to that employed under chlorine. A mixture of manganese dioxide and sodium bromide is warmed with sulphuric acid :— HALOGENS. 53 MnO, + 2NaBr + 2H,S04 = MnSO^ + Na,SO^ + Br, + 2HjO. The operation can be executed in the apparatus pictured in Fig. 38. This can also be used for many other distillations. The retort containing the mixture, is heated in a water-bath ; the tube B serves to cool the vapors which are condensed in the receiver C, surrounded by cold water or ice. When free chlorine is conducted into an aqueous solution of sodium bromide, bromine separates. Bromine is a heavy, reddish-brown liquid, with an exceedingly penetrating, chlorine-like odor (hence the name Bromine, from ßowfioqf stench). At —7.3° it crystallizes to a yellow-green, scaly mass, having a metallic lustre, and resembling iodine. Liquid bromine at o® has the specific gravity 3.18 (water = i) ; it is very volatile, forming dark-brown vapors at the ordinary temperature, and boils Fig. 38. at 63®, changing at the same time into a yellowish-brown vapor. Its density equals 79.96 (hydrogen = i), or 5.53 (air = i). Bromine is more soluble in water than chlorine. Cooled below 4® C., the hydrate (Bra + loHaO) crystallizes out : this is analogous to the chlorine hydrate. It is decomposed at moderate temperatures. Bromine dissolves with ease in alcohol, and especially in ether, chloroform and carbon disulphide. An aqueous solution of bromine sustains a decomposition similar to that of chlorine water on exposure to sunlight. In a chemical point of view, bromine is extremely like chlorine, combining directly with most metals to form bromides ; but it pos¬ sesses a weaker affinity than chlorine, and is liberated by the latter from its compounds :— KBr -f CI = KCl + Br. 54 INORGANIC CHEMISTRY. With hydrogen it only combines on warming, not in sunlight. Upon hydrocarbons it acts like chlorine, withdrawing hydrogen from them. Bromine water gives starch an orange color. 3. IODINE. I = 126.86. ij = 253.72. Iodine, as well as bromine, occurs in combination with sodium, in sea water and some mineral springs, especially at Hall, in Austria, and the Adelheit Spring in Bavaria. In these springs the iodine can easily be detected ; in the sea water it is, however, only present in such minute quantity that its separation, practically, is disadvantageous. Sea algae absorb it from the water, and these are then thrown by the tide on various coasts, where they are burned, yielding an ash (known as kelp in Scotland, as varec in Normandy) which is the principal source for the manufacture of iodine. It was in this ash that the element was accidentally discovered, in 1811 ; in 1815, it was investigated by Davy and Gay-Lussac, and its ele¬ mentary character established. To obtain the iodine, the ash is treated with water, the solution concentrated, and the sodium and magnesium iodides are further worked up. Lately, iodine has been obtained from the mother-liquors of the crude Chili saltpetre. It is set free from its compounds in the same manner as chlorine and bromine—by distillation with manganese dioxide and sulphuric acid. It is more convenient, however, to pass chlorine(orbetter,nitrous acid) through a solution of the salt, when all the iodine will separate :— KI -f CI = KCl -f I. The grayish-black powder thus liberated is collected on a filter, dried, and then sublimed. Iodine is a gray-black solid, subliming in large rhombic crystals, possessing strong metallic lustre. It has a peculiar odor, reminding one somewhat of that of chlorine ; it stains the skin brown, and is corrosive, although not as strongly so as bromine. Its specific gravity is 4.95. It fuses at 113° to a dark-brown liquid, and boils near 200°, passing at the same time into a dark-violet vapor (hence the name iodine, from violet-blue^. The vapor density of iodine equals 8.7 up to 600° C. (air = i) or 126.86 (H = l), corresponding to the molecular weight Ij = 253.72. Above 600® the vapor den¬ sity gradually diminishes, and about 1500® it is only half the original. This is explained by the gradual decomposition (see Dissociation of Water) of the normal diatomic molecule Ij into the free atoms T -|-1- In like manner the bromine mole¬ cules Br,, suffer a separation into the free atoms. The dissociation of bromine vapor (diluted with 11 vols, of nitrogen) commences about 1000® and is com¬ plete at 1600°. The vapor density of chlorine is still normal at 1200®, and it is only at 1400° that it sustains a slight diminution. Oxygen and nitrogen on the contrary ^how no alteration in their vapor density even at 1690®, and possess equal power of expansion (C. Langer and V. Meyer). Iodine is very slightly soluble in water, more readily in alcohol HALOGENS. 55 (Tinctura lodt)^ very easily in ether, chloroform and carbon disul- phide, the last two assuming a deep red-violet color in consequence. It affords a particularly beautiful crystallization, consisting of forms of the rhombic system, when it separates from a solution of glacial acetic acid. In chemical deportment iodine closely resembles bromine and chlorine ; it possesses, however, weaker affinities, and for this reason is liberated from its compounds by those elements. With the metals it usually combines only when warmed ; with hydrogen it does not combine directly, and it does not remove it from its carbon compounds. The deep blue color it imparts to starch is characteristic of iodine. On adding starch-paste to the solution of an iodide, and following this with a few drops of chlorine water, the paste will immediately be colored a dark blue by the separated iodine. This reaction serves to detect the smallest quantity of it. Iodine is largely employed in medicine, photography, and in the preparation of aniline colors. 4. FLUORINE. Fl = 19. (Flj = 38.) Fluorine is found chiefly in the minerals fluorite (CaFla) and cryolite (AlFl3.3NaFl). By proper treatment these minerals yield hydrofluoric acid. The latter dissolves potassium hydrogen fluoride. The resulting liquid is a good conductor of electricity, and by its electrolysis, in a U-tube of platinum, immersed during the decom¬ position in a bath of liquid methyl chloride, boiling at —23°, Moissan (1887) was enabled to isolate the element. Fluorine is a greenish-yellow gas with an odor resembling that of a mixture of hypochlorous acid and nitrogen dioxide. It attacks the mucous membrane very powerfully. Its effects are very per¬ sistent. It remains as a gas at —95°, and at the ordinary pressure of the atmosphere. In the dark it unites with hydrogen with great violence, and even combines with it at —23°. It decomposes water, producing hydrofluoric acid and liberating oxygen in the form of ozone. It does not react with chlorine and nitrogen, but combines with sulphur, bromine, iodine, phosphorus, arsenic, carbon, boron, and silicon with the production of flame. Sodium and potassium unite very energetically with fluorine at the ordinary temperature. Calcium burns in an atmosphere of the gas. Iron combines with fluorine with incandescence. Lead and mercury are attacked by it in the cold. Silver is slowly attacked at the ordinary temperature. Nascent fluorine rapidly attacks platinum at a tem¬ perature of —23°, while at 100° it appears to be without action {Ann. Chitn. et Fhys., 6th Series, 12, 472; 24, 224). Upon the basis of theoretical observations developed later the speciflc gravity of free fluorine is 19 (hydrogen =: i). 56 INORGANIC CHEMISTRY. These four similar elements, fluorine, chlorine, bromine, and iodine, exhibit gradual differences in their properties; and, what is remarkable, this gradation stands in direct relation to the speciflc gravity, of the elements in the state of gas or vapor. Fl Cl Br I Specific gravity, 19 35-453 19-9^3 126.86 With the increase of specific gravity occurs a simultaneous con¬ densation of matter, which expresses itself in the diminished vola¬ tility. Fluorine is a gas ; chlorine can readily be condensed to a liquid ; bromine is a liquid at ordinary temperatures, and iodine is a solid. Other physical properties, as seen in the following table, are also in accord with the preceding :— Fusing point Boiling point Specific gravity in liquid or solid condition Color Fluorine. Chlorine. Bromine. +1 —33° ••••••••• Green-yellow 1-33 Yellovir 3.18 Brown Iodine. + "3° -I-200® 4.97 Black Just such a gradation, as we have seen, is observed in the chem¬ ical affinities of these four elements for the metals and hydrogen ; fluorine is the most energetic, iodine the least. Therefore, each higher element is separated from its soluble metallic and hydrogen compounds by the lower. We shall discover, later, that in the affinity-energy of the halogens for oxygen and some other metalloids, the reverse is true. COMPOUNDS OF THE HALOGENS WITH HYDROGEN. With hydrogen the halogens form compounds of an acid nature, readily soluble in water. I. HYDROGEN CHLORIDE. HCl = 36.456. Density = 18.22. The direct union of chlorine with hydrogen takes place through the agency of heat, and by the action of direct sunlight or other chemically active rays; in diffused light the action is only gradual, and does not occur at all in the dark. On introducing a flame of hydrogen ignited in the air into a cylinder filled with cWorine (Fig. 39), it will continue to burn in the latter. The opposite, the com¬ bustion of chlorine in an atmosphere of hydrogen, may be shown easily by the following experiment (Fig. 40). An inverted cylinder hydrogen chloride. 57 is filled with hydrogen by displacement, the gas is ignited at the mouth, and a tube immediately introduced which will conduct dry chlorine into the cylinder. The burning hydrogen will inflame the chlorine, which will continue to burn in the former. From these experiments, we perceive that combustibility and combustion are only relative phenomena ; if hydrogen is combustible in chlorine (or air), so, inversely, is chlorine (or air) combustible in hydrogen. By the term combustion, in chemistry, is understood every chemical union of a body with a gas, which is accompanied by the phenom¬ enon of light. A mixture of equal volumes of chlorine and hydrogen explodes with very great violence under the conditions given above for the irnion of the gases. The product is gaseous hydrogen chloride. fig. 40. /ÜI Í fct The formation of the latter compound succeeds best by allowing sulphuric acid to act upon sodium chloride, when solid sodium sul¬ phate and hydrogen chloride gas will result :— 2NaCl + HjSO^ = NajSO^ -j- 2HCI. Pour over 5 parts sodium chloride, 9 parts sulphuric acid, somewhat diluted with water (2 parts), and warm the mixture gently in a flask, A (F^. 41). The escaping hydrogen chloride is conducted through a Woulff's bottle containing sul¬ phuric acid or through the cylinder B (filled with pumice stone saturated with sulphuric acid), intended to free it from all moisture, and afterward collected over mercury. Physical Properties.—Hydrogen chloride is a colorless gas, with a suffocating odor. In moist air it forms dense clouds. Its critical temperature is about -j- 52.3®, and the critical pressure 86 atmos- fig. 39. 58 INORGANIC CHEMISTRY. pheres, i. m aqueous hydrogen iodide, water forming at the same time. This corresponds to the thermo- chemical equation :— 2(HIAq) + O = HaO-liquid + 2I . . . (-f 41.9 Cal.) Iodine, on the other hand, does not act upon water. Oxygen cannot affect aqueous hydrogen chloride, but water may be decomposed gradually by chlorine into hydrochloric acid and oxygen (p. 52) according to the equation :— HjO-liquid 2CI = 2(HClAq) -j- O . . . ( -|- 10.3 Cal.) WEIGHT PROPORTIONS. 69 As far as water is concerned, bromine holds an intermediate position between chlorine and iodine (p. 66). To illustrate these interesting relations, let us study the formation of hydrogen Iodide by the action of iodine upon hydrogen sulphide (p. 63). Since the heat of formation of 2HT (gas) ( — 12.0) is less than that of HjS-gas (-{- A'S)> h is impossible for iodine to act upon gaseous hydrogen sulphide :— H^S + Ia = 2HI + S . . .( — 7.5 Cal.) In presence of water the result is different. Aqueous HI is formed ; its heat of formation (2 X ^3-^ Gal.) is greater than that of the aqueous H^S-water (9.2 Gal.) :— HjS Aq + 2l = 2HIAq + S . . . ( + 17.2 Gal.) The transposition is incomplete if the quantity of water is small. COMPOUNDS OF THE HALOGENS WITH EACH OTHER. These compounds, formed by the union of the halogens with each other, are very unstable, and it may be remarked here, that this is also true of most derivatives obtained from elements which are similar in chemical respects. When chlorine is conducted over dry iodine, the latter being in excess, mono-chlor-iodine results, and when the chlorine is in ex¬ cess, trichlor-iodine is formed Iodine Chloride—ICI—is a red crystalline mass, fusing at 24.7° C., and distilling a little above too° C. Water decomposes it easily, with formation of iodic acid, iodine, and hydrogen chloride. Iodine Trichloride—IClj—is formed upon mixing iodic acid with concentrated hydrochloric acid, and by the action of PCI5 upon I2O5. It crystallizes in long, yellow needles, and, when heated, suffers decomposition into ICI and chlorine (at ordinary pressure, the dissociation commences at 25° C.). It dissolves in a little water without alteration ; but large quantities cause partial decomposition, with formation of iodic acid. Iodine Bromide—IBr—obtained by the direct union of the elements, consists of iodine-like crystals, fusing at about 30°. Iodine Pentafluoride—IFI5—is produced by the action of iodine upon silver fluoride, and forms a colorless, strongly fuming liquid. WEIGHT PROPORTIONS IN THE UNION OF THE ELEMENTS. THE LAW OF CONSTANT PROPORTIONS. ATOMIC HYPOTHESIS. If in the halogen derivatives considered, as well as in all other chemical compounds, we determine the quantity of the elements (according to methods described in analytical chemistry), we will discover that they are always combined with each other in the same 70 INORGANIC CHEMISTRY. proportions by weight. In every chemical compound the proportions by weight of the constituents contained in it are invariably the same. Thus chemical analysis shows the following percentage composition for the halogen derivatives of hydrogen :— H = 5.0 2.7 H= 1.2 H= 0.8 Fl = 95.0 01=97.3 Br =98.8 I =99.2 HFl =100.0 HCl=100.0 HBr =100.0 HI = 100,0 Experience has shown that hydrogen, of all the elements, enters compounds in the least quantity, therefore its quantity is chosen as unity, and we calculate those weights of the elements which com¬ bine with one part by weight of H. In this manner we find the following proportions for the halogens :— H= I H= I H= I H= I Fl =19 Cl = 35.45 Br = 79.96 I = 126.86 HFl = 20 HCl = 36.45 HBr = 80.96 HI = 127.86 Experiments have also established the remarkable fact that the same proportions of the halogens by weight are also obtained by the union of the same with other elements. Thus 19 parts of Fl by weight combine with the following weights of the metals : 23.06 parts Na, 39.14 parts K, 32.7 parts Zn, 31.6 parts Cu, 100.2 parts Hg, and 35.45 parts CI, 79.96 parts Br, and 126.8 parts I combine with exactly the same quantities of these metals by weight. Let us take another example. On bringing copper into the solution of a mercuric salt the former dissolves, while Hg separates out j indeed, 31.6 parts Cu displace 100.2 parts Hg. If zinc be brought into the copper solution thus obtained, it will dissolve, while copper sepa¬ rates—and 32.7 parts of Zn separate 31.6 parts Cu. Furthermore, zinc displaces the hydrogen in acids; from all of them 32.7 parts Zn separate i part H. In all these reactions we observe the ele¬ ments appearing in the same quantities by weight. These remarkable facts are fully verified by experiments. Such facts may be formulated into a rule, and when a rule comprises a great number of facts—true for all and expressible in numbers— we designate it a law. The facts presented above find their ex¬ pression in the empirical law of constant proportions, first proposed by Dalton, and reading : The elements combine with each other in definite proportions by weight ; and the proportions by weight of two elements remain the same in their combinations with other elements. Causes underlie facts. The cause is first expressed in the form of a supposition or hypothesis, and when the latter includes a long ATOMS. 71 series of facts, if it is repeatedly substantiated by other phenomena and has acquired a high degree of probability, it is termed a theory. If an hypothesis completely satisfies all the observations to which it refers, it becomes a fact, for the further explanation of which a new hypothesis may be necessary. Conversely, something which long passed as a fact or a theory may be shown to be erroneous, if not any longer consistent with new observations. H3rpothesis and that which we designate a fact, are distinguished really by the different degree of probability only. If, for example, we make a sight observa¬ tion we assume the hypothesis that the same has been caused by an external pro¬ cess, of the reality of which (in distinction from subjective perceptions) we can only assure ourselves by repeated observations. The hypothesis of the revolving of the earth, which at first was only a suitable, improbable supposition, proposed for simplifying calculation, has become a fact. The combustion theory of Lavoi¬ sier met a like result. The same may be true with regard to the supposition of atoms—whether we comprehend them as material particles or as ether motion. The law of constant proportion finds its clearest explanation in the hypothesis of the existence of atoms. Grecian philosophers even conjectured that matter consisted of indivisible and very small par¬ ticles—atoms (from ¿1, privative, and rófioc;^ division). This a priori supposition was subsequently repeatedly announced, but Dalton (1804) first gave it an actual confirmation, in that he applied it to the law of constant proportions. According to the atomic view, matter consists of extremely small (although not indefinitely small) particles, atoms, which cannot be further divided, either mechani¬ cally or chemically. The atoms of different elements possess different weights ; all atoms, however, of one element have the same absolute weight and are like each other. By the aggregation of the elemen¬ tary atoms arise the smallest particles of compound bodies. Upon the basis of these representations, the law of constant proportions becomes very simple ; we can comprehend that the quantities of the constituents of a compound should be constant, and that the relative quantities of the elements by weight, must be the same in all their compounds, as they express the relative weights of the atoms. As yet only the relative atomic weights of the elements have been determined by chemical researches ; in these the hydrogen atoms, as they possess the least weight, have been taken as unity. Until now the knowledge of the absolute atomic weight, for chemical considerations, has been unessential. At the present time different physical phenomena permit fixing the absolute size of the atom with considerable approximate accuracy. Very different considerations lead to the same conclusion, that the atoms cannot be smaller than the fifty-millionth part of a millimeter. (Thomson.) We can determine the diameter of the molecules more accurately with the gases. With hydrogen Hj it has been found equal to the 4 —, with N, the 3 —, with Oj the 7— lo-millionth part of a millimetre. If we grant that in the preceding halogen-hydrogen compounds one atom of hydrogen is combined with every halogen atom, the conclusion follows, that the ratio found expresses the relative atomic 72 INORGANIC CHEMISTRY. weights of the halogens. This supposition, however, appears ques¬ tionable, in view of the more complicated proportions which occur in the union of some elements. Observation shows, to wit, that very frequently two elements unite with each other in not only one, but, indeed, several proportions. For example, 35.45 parts of chlorine combine not only with 31.6 parts copper and 100.2 parts mercury, but also with 63.3 parts copper and 200.4 parts mercury. One part, by weight, of hydrogen, combines with 8 parts of oxygen (more accurately 7.98) to form water, and with 16 parts oxygen (to form the so-called hydrogen peroxide) ; further, with 16 and 32 parts sulphur. Oxygen forms five different compounds with nitrogen according to the following proportions by weight :— Similar proportions are observed in the union of many other elements. Therefore, they combine with each according to several ratios by weight. As we have noticed in the examples given, the varying quantities of one of the elements (calculating upon the same quantity of the other element), bear a simple ratio to each other ; they are mostly multiples of the smallest quantity. These facts are enunciated in the Law of Multiple ProportionSy also proposed by Dalton (1807), which forms an essential amplification of the law of constant proportions. Based on the atomic hypothesis, these facts are explained by saying that the elements can not only unite with each other, atom for atom, but in variable quantities. This con¬ siderably complicates the problem of determining the relative atomic weights of the elements, as these are directly dependent upon the conceived number of atoms in a compound. If, for example, in water, one atom of hydrogen is combined with one atom of oxygen, the atomic weight of the latter would = 8 (regarding that of hydro¬ gen as i). It is just as likely that water consists of two atoms of H and O, or of one of H and two of O, etc.; in the first case the atomic weight of O would =16, in the latter, 4. Analytical results afford nothing positive for the solution of this difficulty. This was the condition in which the question relating to the magnitude of the atomic weights existed thirty years ago. To establish these correctly, various views were allowed to prevail, none, however, with positive foundation. The question can only be solved upon a new and accurate basis : the specific gravities of the chemical compounds in a gaseous or vapor form answer well for this purpose. Nitrogen. 14 parts. 14 parts. 14 parts. 14 parts. 14 parts. 8 parts = I X 8. 16 parts — 2 X 24 parts = 3X8. 32 parts = 4X8. 40 parts = 5X8. Oxygen. Nitrous oxide, . Nitric oxide. Nitrous anhydride. Nitrogen dioxide. Nitric anhydride. GAS VOLUME. 73 DENSITIES OF BODIES IN STATE OF GAS. VOLUME RATIO IN THE UNION OF GASES. ATOMIC MOLECULAR THEORY. The halogens, fluorine, chlorine, bromine, and iodine, unite with hydrogen in only one proportion. The supposition, therefore, that in the halogen-hydrogen compounds, i atom of H is combined with I atom of the halogen, is the simplest and most probable. Then their weight proportions, derived from analysis, directly ex¬ press their relative atomic weights. By comparing these atomic numbers (referring to H = i) with those expressing the density in a state of gas (also referred to H=: i) the astonishing result is seen that the two series are identical. Elements. Hydrogen, Fluorine, Chlorine, Bromine, Iodine, Density. Air = i 0.0692 (I-3I) 2.45 5-52 8.75 Density. hydrogen = i (19) 35-45 79.96 126.86 Atomic Weights. I 19 35-45 79.96 126.86 From this similarity of the atomic (combination) weights with the densities, follows the cogent conclusion that in equal volumes of these elementary gases there is contained an equal number of atoms. Indeed, if in one volume of hydrogen, for example, there are con¬ tained 1000 atoms of hydrogen, which equal 1000 weight units, and in a like volume of chlorine there are also present 1000 atoms of chlorine, which equal 1000 X 35-45 weight units, then it is evident that the relation between the atomic weights and that between the densities (the weights of like gas volumes) must be the same. 1000 X I looox 35.4 I vol. Hydrogen. i vol. Chlorine. These relations can be expressed by the following rule : The atomic weights of the halogen elements are proportional or equal to their densities^ if referred to the same unit. Yielding to a too hasty generalization, this was incorrectly followed for all elements. We arrive at a perfectly similar, but much more general conclu¬ sion, by the consideration of the physical properties of gases or 7 74 INORGANIC CHEMISTRY. vapors. The similar deportment of the same under pressure (law of Mariotte and Boyle), their similar expansibility by heat (law of Charles and Dalton, ordinarily the law of Gay-Lussac), only appear comprehensible by the following suppositions. The gases consist of small portions of matter, which are separated by equal distances', very great in proportion to the particles (the distances of the centres are equal and suffer equal alterations). It immediately follows from this, that equal numbers of particles are contained in equal volumes of all gases (under like temperature and pressure). The kinetic gas theory, based on the same supposition, explains the similar deportment of gases by the equal kinetic energy of the smallest gaseous particles. From the proposition, that in equal volumes an equal number of particles are present, it follows directly that their relative weights are proportional to the volume weights or gas densities, and that by the determination of the latter, the first are also given. In what ratio these smallest particles (called molecules) stand to the chem¬ ically smallest particles (atoms), remains undetermined, and can only be obtained by a comparison of the volume ratios according to which the bodies combine (p. 73). It is, however, even now seen that, at least in the case of compound bodies, the smallest gas particles must be sums of atoms, as the same consist of combina¬ tions of atoms. It follows, from the equality of the atomic weights and the den¬ sities, that the halogens must combine with hydrogen in equal vol¬ umes, since I part of H by weight combines with 35.45 parts of chlorine by weight, etc., and the weights of equal gas volumes stand in the same ratio. Further: i part H and 35.45 parts cnionne yield 36.45 parts HCl \ one volume of the latter weighs, however, 18.2 (H = I, p. 79) ; consequently, 36.45 parts HCl occupy 2 vol¬ umes. Therefore, equal volumes of H and CI yield a double volume of HCl, or, as ordinarily expressed, i volume H and i volume CI yield 2 volumes HCl. In a similar manner it may be deduced that I volume H and one volume Br vapor yield 2 volumes of HBr; that I volume H and i volume I vapor yield 2 volumes of HI. These conclusions are confirmed by the following experiments :— I. The concentrated aqueous solution of hydrochloric acid is decomposed by the action of the galvanic current, and the chlorine and hydrogen collected ; these gases separate at opposite poles. The electrolysis may be made in an ordinary voltameter (Fig. 46). Hofmann's apparatus is better adapted to this purpose (Fig. 47). Two glass cylinders, provided at the top with stop-cocks, are connected at the lower end with each other and with a funnel tube; the latter serves to fill the apparatus with liquid; and also, by further additions, to press out the gases GAS VOLUME. 75 collected in the tubes. The platinum electrodes are fused into the lower part of both tubes. In another form of Hofmann's apparatus (Fig. 48) the electrodes are introduced by means of caoutchouc corks. When the separating gases (in this case the chlorine) attack the platinum, carbon electrodes are substituted for •the latter. To electrolyze hydrogen chloride, fill the apparatus with concentrated hydro¬ chloric acid, which is mixed with ten volumes of a saturated salt solution ; close the upper cocks, and connect the electrodes with the poles of the battery. Gases separate in both tubes, and in equal volumes ; that separated at the positive pole may be proved to be chlorine ; the other combustible gas is hydrogen. This experiment shows that hydrogen chloride decomposes into equal volumes of chlorine and hydrogen. The opposite—the production of HCl by the union of equal volumes of H and CI—is shown in the next experiment. Fig. 47. 2. Fill a cylindrical glass tube, provided with stop-cocks at both ends (Fig. 49), with equal volumes of chlorine and hydrogen. This is most conveniently done by conducting the gaseous mixture obtained by the electrolysis of HCl into the dry tube. (The tube should be filled in the dark, as the gases combine in day¬ light.) When the tube is filled with the mixture, sunlight or magnesium light is brought to bear upon it, when chemical union ensues. On immersing the lower end of the tube into water, and opening the lower cock, the water will rapidly fill the tube, as the hydrogen chloride that was produced dissolves ; all the hydrogen and all the chlorine have disappeared. 3. A modification of this experiment teaches us another important fact which has reference to the ratio of the volume of the hydrogen chloride to the volumes of its constituents. If the tube filled with equal volumes of CI and H be opened under Hg, after the explosion, no diminution in volume will be detected, although the mixture of CI and H has been changed to hydrogen chloride. It follows 76 INORGANIC CHEMISTRY. from this that a mixture of equal volumes of CI and H affords the same volume of HCl, or, as ordinarily expressed, one volume of CI and one volume of H yield two volumes of hydrogen chloride. The following experiment confirms this conclusion : Into a bent tube (Fig. 50), filled with Hg. conduct dry HCl, and then introduce in the bend of the upper part a little piece of metallic sodium. On heating the latter with a lamp, the HQ is decomposed, the CI combines with the Na to form sodium chloride, while hydro¬ gen is set free. Upon measuring the residual hydrogen it will be found that its Fig. 48. FIG. 49. Volume is exactly the half of the volume of HCl originally introduced. In the same manner may be shown the fact that in two volumes of HBr and HI there is con¬ tained in each one volume of H. It follows further from the densities of bromine and iodine vapors, that the quantities of these elements in gas form combining with one volume of hydrogen also occupy one volume. Hence, one volume of hydrogen and one volume of bromine vapor yield two volumes of HBr, and one volume of hydrogen and one volume of todine vapor two volumes of HI. The volume ratios in the chemical union of gases were first investigated by Humboldt and Gay Lussac (1805-1808). The latter derived the two following ATOMS AND MOLECULES. 77 empiñcal laws by experiment : (i) Gases unite according to simple volume ratios ; (2) The volume of the resulting body bears a simple ratio to the volumes of the constituents. Comparing this fact announced by Guy-LussaC, that in the chemical union of gases simple volume ratios do occur, with that discovered by Dalton (p. 70), that the quantities by weight of the combining elements also bear a simple ratio, and granting the atomic constitution of matter, it follows that the number of smallest gas particles (molecules) contained in equal volumes of different gases must bear a simple ratio to each other : the simplest supposition, however, would be that this number of molecules in equal volumes of all gases is the same. These important conclusions were deduced by Avogadro in 1811, and by Ampère in 1814. As deduced on p. 74, and confirmed by the described experi¬ ments, the quantities of the halogen-hydrogen compounds by weight, expressed by the chemical formulas, HCl, HBr, HI, occupy a volume twice as large as one part by weight of H, or 35.45 parts CI, 79-9 parts bromine, 126.8 parts iodine. While the gas densi¬ ties of the elements are equal to their atomic weights (p. 73), those of the compound bodies amount consequently to half that ex¬ pressed by their formulas. From this it would follow that in equal volumes of compound bodies only half as mapy atoms or particles are present as in an equal volume of an elementary form of matter. In fact, one volume of H, containing n atoms of H, combines with one volume of chlorine, which, too, contains n atoms of CI. n parts HCl result, which fill two volumes ; therefore, there are only \ parts of HCl contained in one volume of HCl :— nH -f nCl = nHCl. I vol. I vol. 2 vols. This conclusion contradicts the general postulate (p. 74), derived from the physical properties, viz., that all gases, both simple and compound, contain the same number of gaseous particles in equal volumes. This contradiction, which for a long time prevented the adoption of the atomic volume theory in chemical science, is now easily solved by the following supposition of Avogadro, announced in 1811. It is necessary to distinguish two different kinds of par tides : molecules and atoms. The smallest discrete particles in gases are not atoms^ but molecules, which consist of several atoms. That the molecules of compounds consist of atoms, is obvious, since, indeed, the same represent aggregates of atoms ; but the ele¬ ments also form molecules in a free condition, which are composed of several, generally, of two atoms. The previously deduced rule (p. 73), that in equal gas volumes of the halogen elements there is contained an equal number of atoms, must be formulated somewhat as follows : In equal volumes of all gases is found an equal number of molecules (law of Avogadro). 78 INORGANIC CHEMISTRY. The process of the combination of hydrogen with chlorine (and the other halogens) must be conceived therefore to be somewhat like the following : i molecule of H, containing 2 atoms of H, acts upon i molecule of CI, also composed of two atoms of 01, and there result 2 molecules of HCl ;— + CI2 = 2HCI. We can now understand that hydrogen chloride contains just as many molecules in an equal volume as H and CI. This is apparent from the following representation ;— nH, + nClj - - nHCl nHCl I volume. I volume. 3 volumes. In a similar manner 2 volumes H (containing 2« molecules) give with I volume oxygen (containing n molecules), 2 volumes aqueous vapor j consequently, 2n molecules of water. In 2« molecules of the latter (HjO) there are contained 2« atoms of O ; therefore in n molecules of oxygen, 2« atoms of oxygen—or one oxygen molecule consists 0/2 atoms. + nOa yield nHjO nHjO I vol. 2 vols. 2 vols. In the same way it may be shown that the nitrogen molecule con¬ sists of 2 atoms of nitrogen (N2), the phosphorus molecule, of 4 atoms of phosphorus (PJ, etc., etc. This peculiar result, following from the law of Avogadro, that the molecules of the elements consist of several atoms, etc., is shown by many other circumstances founded on facts, for example, by the existence of the allotropie modifications of the elements (compare ozone), by the chemical reactions (compare hydrogen peroxide), and by the remarkable action of the elements in the moment of their liberation. ATOMS AND MOLECULES. 79 Upon p. 52 we said that the oxygen separated from water by chlorine acted much more energetically than free oxygen. Other elements, especially hydrogen, behave similarly in the moment of formation—in statu nascendi. As viewed by the atomic molecular theory, this may be very easily explained. The free elements ftheir molecules) are compounds of similar atoms whose chemical affinity has always been partially satisfied. In the moment of their separa¬ tion from compounds free atoms appear, which, before they combine to molecules, must act more energetically. All that has been developed in the preceding statements may be summarized in the following sentences : All bodies are composed of elementary atoms. The latter unite to produce the molecules of the simple and compound bodies. Molecules are the smallest dis¬ crete particles existing in a free state. The same number of mole¬ cules is contained in equal volumes of all gaseous and vapor-form- ing bodies. Therefore, the gas densities bear the same ratio to each other as the molecular weights. The density is generally com¬ pared with that of hydrogen = i, while the molecular weights are referred to H2 = 2 ; therefore, the gas densities (the specific gravi¬ ties of gases) of all bodies are one-half their molecular weights. The atomic weights are compared with H = i, therefore, the den¬ sities of the elements whose molecules consist of two atoms, are equal to the atomic weights :— Atoms. Molecules. Density. H = 1.003 2 1.003 Cl = 35-45 CI2 70.9 35-45 Br = 79.9 Brj 159.8 79-9 I = 126.8 I2 = 253.6 126.8 HQ 36.45 18.2 • HBr 80.9 40.45 HI 127.8 63-9 0 = 16 O2 =r 32 16 H2O 18 9 N == 14.04 N, =■ 28.08 14.04 Nn3 17.04 8.52 P = 31.03 P* = 124.12 62.06 PHg ' 3403 17.01 A simpler deduction, that the molecules of the elements consist of two or more atoms, is the following : We proceed from the law of Avogadro, that an equal number of molecules is contained in equal volumes of all gases or vapors. This law, or better hypothesis, cannot be pro,ven mathematically, as was attempted ; just 80 INORGANIC CHEMISTRY. as little as any other fundamental hypothesis *—but it possesses, as basis of the en¬ tire recent kinetic theory of gases, a high degree of probability. It necessarily fol¬ lows from this law that the molecular weights of all bodies are proportional to the gas densities. Referred to hydrogen as unit, the empirical gas densities of HCl = i8.2, of HBr = 40.4, of HI = 63.9, etc. Analysis shows, however, that 35.4 parts of CI are in union with I part H in HCl, 79.9 bromine in HBr, 126.8 iodine in HI. As the weight of one atom of H is made equal to i, and 35.4 parts of chlorine are combined with it, the weight of a molecule of hydrogen chloride, consisting of at least one atom of H and one atom of CI, must equal 36.4 ; it is, therefore, twice as much as its density, 18.2. Hence the molecular weights of all other bodies, as they bear the same ratio as the densities, must also be twice as large (referred to H as unit) as the latter. The hydrogen molecule is = 2, and consists of two atoms, as its atomic weight equals I. The chlorine molecule weighs 70.9 units, and consists of two atoms (CIg), if we suppose that the atomic weight = 35-45- Its atomic weight could, however, be only the half (or another sub- multiple) of 35.4; then its molecule would consist of four chlorine atoms (Cl^ = 70.9 when CI is made equal to 17.6), and the formula of hydrogen chloride would be HClg. From the densities of the elements in gas form we only ascertain their molecular weights. Their atomic weights are derived from the molecular weights of their compounds, as we regard the smallest quantity of the element which analysis discloses in the molecule of any compound as the atomic weight. Thus, in the molecule of any compound of chlorine there are never than 35.45 parts by weight of CI. That the maximum values thus derived have not been found too high, but correspond to the actual relative atomic weights, follows from the agreement of these numbers with the atomic numbers obtained from the spe¬ cific heat of the elements. The complete certainty of their correctness we reach by the law of periodicity, which is formed from these numbers. Taking one atom of hydrogen as the unit of weight and volume, then two parts by weight of H, or one molecule (Hj), would oc¬ cupy two volumes. We say, therefore, although incorrectly, that the molecules fill two volumes, and designate the molecular formulas double volume formulas. The volume of molecules and atoms is, however, unknown to us ; we only know that in equal gas volumes there is contained an equal number of molecules. These convincing suppositions and conclusions deduced from these actual relations, form the atomic molecular doctrine, which is the foundation of the chemistry of to-day. As this doctrine completely explains the quantitative phenomena arising in the action of the chemical elements upon each other, and as it has been repeatedly confirmed by entirely opposite phenomena, it is only proper and correct that it be designated a theory (p. 71). OXYGEN GROUP. In this group are included the elements oxygen, sulphur, selenium, and tellurium. They are perfectly analogous in their chemical deportment. They unite with two atoms of hydrogen. * A mathematical proof is only possible upon the basis of another, more general, quantitative hypothesis (or of an axiom), which in turn is not provable. OXYGEN. 81 I. OXYGEN. O = i6. O, = 32. Oxygen (oxygenium) is the most widely distributed element in nature. It is found free in the air; in combination it exists in water. It is an important constituent of most of the mineral and organic substances. It was discovered, almost simultaneously, by Priestley, in Eng¬ land, 1774, and Scheele, in Sweden, 1775. Lavoisier, in France, 1774-1781, first explained the important rôle attached to oxygen in processes of combustion, of respiration, and of oxidation. Preparation.—Heat red mercuric oxide, a compound of mercury with oxygen, in a small glass retort ; in this way the oxide is decom¬ posed into mercury and gaseous oxygen :— Hg0 = Hg + 0. The following method is commonly pursued in the chemical laboratory : Potassium chlorate, a compound of potassium, chlorine and oxygen, is heated in a glass retort (Fig. 51) or flask, and thus decomposed into solid potassium chloride and oxygen ;— KClOg^KQ-l-aO.* The evolution of the gas proceeds more regularly and requires a less elevated temperature if the pulverized chlorate be mixed with ferric oxide or manganese peroxide. The liberated oxygen is collected over water. McLeod [Jr. Chem. Soc., 55, 192) explains the mechanism of the action of manganese peroxide on potassium Perchlorate, when heated, as follows : ist, the peroxide acts on the chlorate producing potassium permanganate, chlorine, and oxy¬ gen, aMnOj 2KCIO3 = KjMnjOg -f- Clj -|- Oj. The permanganate is then decomposed by heat : KjMujOg = KjMnO^ -f- MnO, -I- Oj, and in the third stage the change is likely : KjMnO^ -1- CI, = 2KCI -j- MnO, O,. Very pure oxygen may also be obtained by heating potassium dichromate with sulphuric acid :— KjCrjOj -)- 411280^ = Cr2(SO^)3 -(- KjSO^ -f- 4H2O -|- 30* Besides these, many other methods may be employed for the preparation of the gas: e. g., the ignition of manganese and barium peroxides ; the decomposition of sulphuric acid at a high heat; the boiling of a solution of bleaching lime with a cobalt *The chemical equations used here and previously are only intended to repre¬ sent the manner of the reaction, and to express the accompanying relative quanti¬ ties by weight. It should not be forgotten that free iUoms do not exist, but that they always occur combined in molecules. Molecularly written the equation would be :— 2KC108 = 2KC1-H302. 82 INORGANIC CHEMISTRY, salt, etc. These methods, applied technically, will be considered more fully later. A very convenient laboratory method for the preparation of oxygen consists in allowing dilute hydrochloric acid to act upon a mixture of barium peroxide (2 parts) and manganese peroxide (i part). The gas is evolved at the ordinary tem¬ peratures. If the solid ingredients are mixed with gypsum and a little water the mass can he moulded into cubes, and the oxygen then he generated in a Kipp apparatus {Berichte^ 20, 1585). Properties.—Oxygen is a colorless, odorless, tasteless gas. Its density equals 16 (H = 1.003), 1.1060 (air = i). One litre of oxygen at o°C., and 760 mm. pressure, weighs 1.4330 grams (16 times more than one litre of hydrogen). It is only slightly soluble in water; 100 volumes of the latter dissolve 4.1 volumes of the gas at 0°, and 2.9 volumes at 15®. It is more readily dissolved by absolute alcohol (28 volumes in 100 volumes). Fig. 51. h The critical temperature of oxygen is —118°, and its critical pressure equals 50 atmospheres (p. 48). Liquid oxygen under a pressure of i atmosphere boils at —181°, and under 9 mm. pressure at —225°. Its specific gravity at —118° equals 0 65, at —139® it is 0.87, and 1.124 at —181°. Oxygen combines with all the elements excepting fiuorine. With most of them it unites directly, accompanied by the evolution of light and heat. The combustion of bodies which burn in the air depends on their union with oxygen, which is present in the same to the amount of 23 per cent. The phenomena of the respiration of animals are also infiuenced by the contact of the oxygen of the air—hence the earlier designations of oxygen as inflammable air, and vital air. In pure oxygen the phenomena of combustion pro¬ ceed more energetically. Ignited charcoal or an ignited sliver in¬ flames immediately in the gas, and burns with a bright light. This OXYGEN. 83 test serves for the recognition of pure oxygen. Sulphur and phos¬ phorus ignited in the air burn in it with an intense light (Fig. 52). Even iron is able to burn in the gas. To execute this experiment, take a steel watch spring, previously ignited, attach a match to the end, ignite the same, and then introduce the spring into a vessel filled with oxygen gas (Fig. 53). At once the match inflames and ignites the iron, which burns with an exceedingly intense light and emits sparks. (To protect the vessel from the fusing globules of iron oxide, cover the bottom with a layer of sand.) Iron will bum in any flame if a current of oxygen be conducted into the same. Oxygen combines with hydrogen to form water. The union oc¬ curs at a red heat, by the electric spark or by the action of plati¬ num sponge (p. 45). Hydrogen burns in oxygen with a flame j vice versa, oxygen must also burn in hydrogen ; this may be demonstrated in the same manner as indicated under hydrogen chloride (p. 56). A mixture of hydrogen and oxygen detonates violently; most strongly if the proportions are i volume of oxygen and 2 volumes of hydrogen ; such a mixture is known as oxy- hydrogen gas. The explosibility may be shown in a harmless way by the following experiment : Fill a narrow-necked flask of 4-6 ounces, over water, ^ with hydrogen, and yi oxygen ; close the opening with a cork, then wr^p the flask up in a towel, remove the cork and bring a flame near the opening. A violent explosion en¬ sues, generally with complete breaking of the flask. Fig. 52. Fig. 53. Fig. 54. 84 INORGANIC CHEMISTRY. The oxy-hydrogen flame is only faintly luminous ; it possesses, however, a very high temperature, answering, therefore, for the melting of substances which fuse with great difficulty, e. g.y plati¬ num. To get a continuous oxy-hydrogen flame, efflux tubes of peculiar construction are employed (Fig. 54); through the outer tube, Wy hydrogen is brought from a gasometer; oxygen is con¬ veyed through the inner Sy and the mixture ignited at a. Such a flame impinging on a piece of burnt lime makes the latter glow and emit an extremely bright light—Drummond's Lime Light. The union of oxygen with other substances, is termed oxidation. This term, as well as the name oxygenium (from and yewáo»), or acid producer, arises from the fact that acids are sometimes formed in oxidation. This the combustion experiments, previously mentioned, prove. If the vessels, for instance, in which carbon, sulphur, and phosphorus were burned, be shaken up with water, the latter will give an acid taste, and redden blue litmus paper. It was formerly thought that the formation of acids is always con¬ ditioned by oxygen. We have, however, already noticed that the haloid acids HCl, HBr, and HI, contain no oxygen. Some of the elements yield acids by their union with oxygen, or more cor¬ rectly oxideSy which form acids with water. Most of these are the metalloids. Thus the following corresponding acids are derived from the acid-forming oxides of sulphur and phosphorus :— SO3 -I- HjO = HjSO^ Sulphur Sulphuric trioxide. acid. P3O5 + HP =r 2HPO3 Phosphorus Metaphosphoric pentoxide. acid. With oxygen the metals usually yield oxides, which form hydrox¬ ides (hydrates) or bases with water ;— Kp + HjO = 2KOH Pot. Potas, oxide. hydroxide. CaO + HjO = Ca (OH), Calcium Calc. oxide. hydroxide. The salts are produced by the alternating action of acids and bases (see p. 60). Thirdly, there exist the so-called indifferent oxideSy which yield neither acids nor bases, with water, e. g., Np NO BaO, Nitrous Nitric Barium oxide. oxide. peroxide. OZONE. 85 Oxidation is not only induced by free oxygen or bodies rich in it, but frequently, also, by the halogens ; in the latter case the halo¬ gens first decompose the water with the elimination of oxygen, which then oxidizes further (compare p. 52). The opposite of oxidation, the removal of oxygen, is called reduction. Hydrogen (in statu nascendi), and substances giving it off easily (as HI), have a reducing action. Most of the metallic oxides are reduced at a red heat, by hydrogen, e.g. :— CuO + Hj = Cu + Hp. Copper oxide. Copper. OZONE, O3. Ozone, discovered in 1840, by Schönbein, is a peculiar modi¬ fication of oxygen, characterized by a remarkable odor and great ability to react, therefore it is called active oxygen. It is obtained from oxygen in various ways ; it is almost always produced when this gas is liberated, or when it takes part in a reaction ; thus, in the decomposition of peroxides by concentrated sulphuric acid, in the electrolysis of water (at the positive pole), in the slow oxidation of moist phosphorus, in the combustion of hydrocarbons, and in the action of the so-called silent discharge in an atmosphere of oxygen or air. In none of these instances is all the oxygen ever converted into ozone ; only a small portion—in most favorable conditions 5-6 per cent.—suffers this change. The following methods serve for the preparation of ozone 1. Bring several pieces of stick phosphorus into a spacious flask, cover them about half with water, and allow them to stand for some hours. Or conduct oxy¬ gen over pieces of phosphorus placed in a glass tube and moistened with water. Ozone is also formed abundantly when a potassium bichromate solution is substi¬ tuted for water. 2. Pass the electric spark from an electrical machine or a Ruhmkorff coil through air or oxygen. The silent discharge from a powerful induction current is better. For this purpose we can employ a Siemen's induction tube (Fig. 55), which con¬ sists of a glass tube covered without with tin foil, in the interior of which is a smaller tube coated upon its inner side. The oxygen circulates between the two tubes ; the two coatings are in connection with the induction spiral, or the poles of a Holtz electrical machine. The Berthelot induction tube is well adapted for ozonization. In it dilute sulphuric acid replaces the metallic coatings. 3. Gradually add barium peroxide in small portions (or potassium permangan¬ ate) to cold sulphuric acid :— BaOj -f H2SO4 = BaSO^ -f H^O -f O. The escaping oxygen is tolerably rich in ozone, and is collected over water. Ozone possesses a highly penetrating, chlorine-like odor (phos¬ phorus odor), which by prolonged respiration produces bad results. In a long layer, ozone shows a bluish color. If ozonized air be sub- 86 INORGANIC CHEMISTRY. jected to powerful pressure (150 atmospheres) at a very low tempera¬ ture, or if ozonized oxygen be conducted through a small tube cooled to —181° by boiling oxygen, the ozone will condense to a liquid with an indigo-blue color. Liquid ozone, if preserved in a sealed tube, passes into a blue gas, that can be again liquefied by chilling it with boiling ethylene. Ozone is rather stable at the ordinary tem¬ perature; when heated to 300° C., it reverts to ordinary oxygen. It is somewhat soluble in pure water ; the larger portion of it is, how¬ ever, converted by the water into oxygen, without formation of hydrogen peroxide. Unlike ordinary oxygen, ozone, especially in a moist state, oxidizes strongly at ordinary temperatures. Phosphorus, sulphur, and arsenic are converted into phosphoric, sulphuric, and arsenic acids ; ammonia is changed to nitrous and nitric acids ; sil¬ ver and lead are converted into the corresponding peroxides; therefore paper moistened with a lead salt is colored brown. Iodine is separated from potassium iodide by it :— 2KI -f HgO + O = 2KOH + 12- Fig. 55- It also oxidizes all organic substances, like caoutchouc ; therefore the apparatus used in its preparation must not be constructed of the latter. Solutions of dye stuffs, like indigo and litmus, are decolor¬ ized. Very characteristic for ozone is its ability to turn an alcoholic solution of guaiacum tincture blue. Liquid ozone boils at — 106°. If it be enclosed in a glass tube it changes to a blue gas, which may be recondensed by cooling it by means of boiling ethylene. For the detection of ozone the ordinary potassium, iodide starch paper (Schôn- bein) may be used. This is prepared by immersing white tissue paper in a starch solution mixed with potassium iodide. The iodine which the ozone liberates from the potassium iodide blues the starch paper. Thequantity of ozone may be approximately determined from the rapidity and the intensity of the coloration ; the reactive power is, however, very much influenced by aqueous vapor. Thai- QZONE. 87 lous hydroxide is a more reliable reagent for ozone than the potassium iodide paper. Guaiacum tincture and paper saturated with a lead acetate solution may also be used to detect ozone ; the first acquires a blue color, the second is browned. Other substances also blue potassium iodized starch and guaiacum, e. g., chlorine, bromine, nitrogen dioxide, etc., etc. To distinguish ozone from these, proceed as follows (Houzeau): Take two strips of violet litmus paper, one of which is saturated with KI, and expose it to the action of the gas ; when O3 is present potassium hydroxide will be formed from the KI, and color the violet litmus blue. The second paper servers to show the absence of ammonia. The preceding reactions of ozone are all produced by hydrogen peroxide, although less rapidly. The only test answering for the distinction of very slight quantities of ozone from hydrogen peroxide, is the blackening of a bright strip of silver by ozone. Ozone is formed from pure oxygen, and is nothing more than the latter condensed. The molecules consist of 3 atoms of O :— 3O2 yield 2O3. 3 vols, oxygen. 2 vols, ozone. This is proved by the following experiments : In ozonizing oxy¬ gen its volume diminishes ; upon heating (whereby ozone is again changed to oxygen), the original volume is reproduced; when ozonized oxygen is brought in contact with oil of turpentine or cinnamon, all the ozone is absorbed and the volume of the gas is diminished. Comparing this diminution, corresponding to the ozone volume, with the expansion which an equal volume of ozon¬ ized oxygen suffers after the application of heat, we will find that the first is twice as large as the latter ; this indicates that i volume of ozone yields volumes of oxygen. From this it follows that the specific gravity of ozone must be i ^ times greater than that of ordinary oxygen, and that if the molecule of O consists of 2 atoms, the molecule of ozone must contain 3 atoms. This conclusion is confirmed by the specific gravity of ozone derived experimentally from the velocity of diffusion. The density of ozone is found to be 24 (H = i) ; the molecular weight of it, therefore, is 24 X 2 = 48, a number almost equal to the trebled atomic weight of oxygen (3 X 16 = 48). The molecular formula of ozone is, there¬ fore, O3. A diminution in the volume of the gas does not occur in the action of ozone upon oxidizable bodies like KI and Hg, although all the ozone disappears. It would appear from this, that in oxi¬ dizing, ozone only acts with one atom of oxygen, which occupies the same volume as the ozone :— O3 + 2KT = O, -f K2O -k I2. I vol. I vol. As a consequence of this behavior, ozone is also called oxidized 88 INORGANIC CHEMISTRY. oxygen; i. e., free oxygen (O2), which has combined with an additional oxygen atom. Thermo-chemical Deportment,—Compared with ordinary oxygen, ozone is an endothermic compound. Heat is absorbed in its formation from oxygen (p. 28) : (02,0) = 03 .... ( — 32.4 Gal.) This explains why ozone is produced with so much difficulty, and why the addition of considerable energy is necessary. This may be applied directly in the form of heat or electricity (electric sparks, silent electric discharge), or it may be withdrawn from the heat of formation of other exothermic compounds which are produced at the same time, e. g., the formation of ozone by the oxidation of phosphorus to phosphorous acid. Being an endothermic derivative, we readily perceive why ozone is so unstable. It reverts to oxygen with the separation of the excessive energy. Herein is mani¬ fest the cause of the greater reactivity of ozone as compared vdth oxygen. All oxidations performed by ozone are more energetic, because there are 32.4 Gal. more set free in them than in oxidations with ordinary oxygen. We observe, therefore, that the elementary substance oxygen occurs in free condition in two different forms—allotropie modifica¬ tions—ordinary oxygen (O2) and ozone (O3). We will learn later that very frequently substances of the same elementary composition possess different physical and chemical properties ; such bodies are called isomerides and the phenomenon isomerism. The isomerism of the elements is known as allotropy ; this is accounted for (as in the case of oxygen and sulphur) by the different number of atoms in the molecule. The phenomena of isomerism constitute an important argument for the atomic constitution of matter. If in the chemical union of two bodies the particles of matter would entirely permeate and blend into each other, the existence of isomeric bodies would scarcely be comprehensible. We can therefore only suppose a co- stratification of the atoms, and must consider isomerism as only a varied arrangement of the same. Special allotropy verifies the con¬ clusion drawn from the gas density that the molecules of the ele¬ ments are composed of atoms. We have already seen that ozone is absorbed, not only by turpen¬ tine and cinnamon oil, but also by other ethereal oils. These bodies are, however, only very slowly oxidized ; the ozone is contained in them in a peculiar, combined condition. In this form it acts upon some bodies like free ozone; in other instances, the oxidizing action is only rendered possible by peculiar substances which carry the ozone. Spongy platinum, ferrous sulphate, and the blood cor- WATER. 89 puscles are examples of this class. Thus, old turpentine oil, con¬ taining absorbed ozone, only acts on paper saturated with starch and potassium iodide, if a few drops of a ferrous sulphate solution have been added to it. Since ozone is formed when electricity acts upon air, and, indeed, probably, in all oxidation and combustion processes ; and, further, potassium iodide starch paper is blued when exposed to the air ; it was believed that ozone was a constant constituent of atmospheric air (i-io milligrams in loo litres of air) ; according to recent in¬ vestigations it is, however, probable that the imagined ozone reac¬ tions are frequently produced by hydrogen peroxide, which is very similar to ozone in reaction (p. 87), and is almost constantly in the air (Schöne). Antozone, which was regarded as a third peculiar modification of oxygen, has been proved to be hydrogen peroxide. COMPOUNDS OF OXYGEN WITH HYDROGEN. I. WATER. H,0 = 18. Density = 9. Water, the product of the union of hydrogen with oxygen (p. 83), is produced in many chemical processes, e. g., in the formation of salts from bases and acids (p. 60). Cavendish was the first (1781) to confirm the formation of water by the combustion of hydrogen. Lavoisier first (1783) determined its quantitative composition. Later (1805) Gay-Lussac showed that it was produced by the union of two volumes of hydrogen with one volume of oxygen. Physical Properties.—It is obtained chemically pure by the dis¬ tillation of naturally occurring water, which always contains other matter dissolved in it. It appears in all three states of aggregation ; in the liquid, gaseous (steam), and solid (ice, snow). When water is cooled it contracts and attains its greatest density at -|- 3.7° C., the maximum contraction. The weight of, a cubic centimeter of such water is taken as the unit of weight (= i gram). By further cooling the water expands—the opposite of most other bodies; its volume becomes greater, while the specific gravity decreases. The following table gives the volume and specific gravity of water for different temperatures (according to Kopp) :— 8 90 INORGANIC CHEMISTRY. Temperature. Volume. Specific Gravity, 0° 1.00012 0.99988 2° I .OOCX)3 0.99997 4° 1.00000 1.00000 6° 1.00003 0.99997 8° 1.00011 0.99989 10° 1.00025 0-99975 12° 1.00044 0.99956 14° 1.00068 0.99932 16° 1.00097 0.99903 18° 1.00131 0.99869 20° 1.00169 0.99831 22° 1.00212 0.99789 24° 1.00259 0.99742 By cooling water solidifies to ice. The solidification-tem¬ perature of water, or more correctly the fusing point of ice, is taken as the zero of Celsius's and Réaumur's thermometric scales. We can, howe^^er, reduce still water considerably below the o° point without its freezing, while the fusing point of ice, like all other solid bodies, is constant (at a definite pressure). In the conversion of water into ice, a considerable expansion occurs; loo vols. HjO at o° yield 107 vols, ice at 0° ; the specific gravity of the latter is, therefore, 0.93. Ice crystallizes in hexagonal forms, as may be distinctly observed in snow-flakes. Different bodies require different quantities of heat to bring them to the same temperature. The heat capacity of water is greater than that of all other liquid or solid bodies. It is customary to take the quantity of heat necessary to raise one part by weight of HjO from 0° C., to 1° C., as the unit of heat, or calorie. In the pa.ssage of a liquid to the solid state heat is always set free, while, on the other hand, in the fusion of the solid heat is absorbed. The latent heat of water equals 79 calories; that means, that for the fusion of one part of ice by weight, a quantity of heat is required which is capable of raising one part H2O from 0° to 79° C. Water boils upon the application of heat, and is converted into steam. The boiling temperature, like that of all other liquids, depends on the pressure ; it is also influenced by the substances dis¬ solved in it, although the temperature of the vapors is constant (at a given pressure). The temperature of the steam escaping from water at the ordinary pressure of 760 mm. is= 100° of the thermo¬ metric scale of Celsius ( = 80° Réaumur). One volume of water, at 100° C., yields 1696 volumes of vapor WATER. 91 of the same temperature. The specific gravity of steam == = 9 (H = i), or = 0.623 (air = i). One litre of aqueous vapor weighs 0.8064 grams (at 0°). The critical temperature of water (or its absolute boiling tempera¬ ture, p. 47) ¡8 + 370°, and its critical pressure 195.5 atmospheres, i. or (Dixon) : CO 2OHH O, = CO2 -|- HjO HjOj. More careful investigations have, however, demonstrated that the presence of moisture is not necessarily essential in oxidations. Carbon monoxide and oxygen also combine to carbon dioxide when perfectly dry if the temperature be sufficiently high. Their union in the presence of moisture is due solely to the fact that the following transpositions, CO -4- HjO = CO, -f- H, and 2Hj -j- O, = 2H2O, take place more readily and at a lower temperature than the direct union of carbon monoxide with oxygen : 2CO O, = 2CO, (Lothar Mey«). HYDROGEN PEROXIDE. 103 Hydrogen peroxide, concentrated as much as possible under the air pump, is a colorless, syrupy liquid, with a specific gravity of 1.45, and does not solidify at—30° C. ; from very dilute solutions, pure water freezes out. It possesses a bitter, astringent taste, is miscible in all proportions with water, and vaporizes in vacuo. Very dilute aqueous solutions can be boiled without decomposing the peroxide ; a portion of it distils over with the water. In concentrated solutions, hydrogen peroxide is very unstable, and easily decomposed with liberation of oxygen ; in more dilute acidulated solutions it may be preserved longer. Decomposition occurs, even at ordinary temperatures ; by heating the point of ex¬ plosion can be reached. In consequence of this ready decomposi¬ tion, hydrogen peroxide oxidizes powerfully, since oxygen appears (p. 79) in statu nascendi. It converts selenium, chromium, and arsenic into their corresponding acids ; sulphides are changed to sulphates (PbS to PbSOi) ; from lead acetate solutions the peroxide is precipitated, but is again decolorized by the excess of peroxide. Organic dyestuifs are decolorized and decomposed. From hydro¬ gen sulphide, sulphur, from hydrogen chloride and iodide, chlorine and iodine are set free :— H2O2 + 2HI = 2H2O + Ij. Thus hydrogen peroxide acts in a manner analogous to ozone ; in both there exists a slightly bound atom of oxygen, which can readily be transferred to other bodies. Hydrogen peroxide acts very slowly upon a neutral potassium iodide solution, while ozone separates iodine at once j but if plati¬ num-black, ferrous sulphate, or blood corpuscles (see p. 88), be added to the solution, iodine immediately separates out, and colors added starch paste a deep blue. In all these cases the action of hydrogen peroxide is oxidizing. Some substances, on the other hand, are reduced by it, oxygen sepa¬ rating at the same time ; this is true of certain unstable oxides, per¬ oxides, and the highest oxidations of some metals, like MojOt, CrOg. Thus, argentic, mercuric, and gold oxides are reduced to a metallic state with an energetic evolution of oxygen :— ~1~ H2O2 = zAg -|- H2O O2. Lead peroxide is changed to lead oxide :— PbOj -1- H2O2 = PbO -f H2O -f O2. In the presence of acids, the solution of potassium permanganate is decolorized and changed to a manganous salt (p. 104). In the same way chromic acid and its salts are altered to chromic oxide :— zCrOs -f- 3H2O2 = CrjOj 3^8^ "f" 3®a* 104 INORGANIC CHEMISTRY. Ozone and hydrogen peroxide decompose themselves into water and oxygen :— O3 -f- HjOj = O3 -|- HjO -|- Oj. Chlorine in aqueous solution is oxidized to hypochlorous acid by hydrogen peroxide, Gig -j- H2O2 — 2HOCI, but again reduced by an excess of the latter :— ClOH + H2O2 = CIH + HjO+Oj. All these reactions are generally explained by supposing that the oxygen atoms (also those of other elements), possess a certain affinity for each other; this is saturated by their union to molecules. Those present in other compounds, and not firmly bound, therefore separate and unite with each other, and form free oxygen molecules—OO. The conclusion derived from the gas density, viz., that the molecules of the free elements consist of two or more atoms, is corroborated by these reactions. The readiness with which ozone and hydrogen peroxide react, is explained by their thermo-chemical behavior. Hydrogen peroxide, like ozone (p. 88), is an endothermic compound when compared with water, i. (HjS, Aq.) = 4*7 î (^2» S, Aq.) = ^.2,. It is because of this low heat of formation that the gas is pro¬ duced with such difficulty from its elements, and it is for this reason that it is so readily dissociated by heat into its elements. Its entire chemical deportment is also accounted for by its heat of formation (P- 115)- MOLECULAR FORMULA OF HYDROGEN SULPHIDE. ATOMIC WEIGHT OF SULPHUR. The analysis of hydrogen sulphide shows that it consists of one part hydrogen and sixteen (more accurately 16.03) parts sulphur. If the molecular formula of hydrogen sulphide were HS, the atomic weight of sulphur would be sixteen (compare p. 98). The great analogy of the sulphur compounds with those of oxygen (p. 108), permits us to accept formulas for the former similar to those of the latter. The molecular formula of hydrogen sulphide would, therefore, be HjS = 34.06, and the atomic weight of sulphur would equal 32.06. Hence the gas density of hydrogen sulphide must be ^ = 17 (H = i), or 1.177 (iùr= l); this is confirmed by direct experiment. Conversely, it follows from the gas density HYDROGEN PERSULPHIDE. Ill that the molecular weight of hydrogen sulphide = 34.06. Since the analysis of 34,06 parts of hydrogen sulphide shows the presence of two parts of hydrogen, the molecule of H^S contains two hydrogen atoms. It then follows that the 32.06 parts of sulphur combined with the latter, correspond to one atom of sulphur, because less than 32.06 parts of this element have never been found in the mole¬ cule of any compound in which sulphur occurs (p. 80). From the molecular formula HjS, we further conclude that the hydrogen con¬ tained in one volume of hydrogen sulphide would occupy in a free condition the same volume as the latter : — nHjS contains nHg. I vol. I vol. This conclusion is verified experimentally as follows : In a bent glass tube filled with mercury (Fig. 60), introduce dry hydrogen sulphide gas; then in the bent portion place a piece of tin, which is heated by a lamp. The sulphur of the HjS combines with the metal to form solid tin sulphide, while hydrogen is set free : its volume is exactly equal to the volume of the employed hy¬ drogen sulphide. The quantity of sulphur, 32 parts, in vapor form, at 1000° C., when the density is 32 (p. 107) combined with hydrogen (2 parts) would equal exactly half the volume of the hydrogen; at 500° C., however, when the vapor density is three times as great, it will equal volvune of the hydrogen. I volume HjS, therefore, consists of i volume H and volume sulphur vapor (at 500®), or as ordinarily expressed ; 2 volumes HgS consist of 2 volumes H and y volume sul¬ phur vapor. Written molecularly, we have :— At 500® C. : Sg -f- 6H2 = óHjS. I vol. 6 vols. 6 vols. At lOCO® C., however : Sj -J- 2H2 = 2H2S. I vol. 2 vols. 2 vols. 2. HYDROGEN PERSULPHIDE. Just as hydrogen peroxide HgOa is formed by (p. loi) the action of acids upon some peroxide, so may hydrogen persulphide be ob¬ tained from metallic persulphides. Calcium persulphide is most suitable, and when its aqueous solution is poured into dilute hydro¬ chloric acid, CaSj -f 2HCI = CaClj H2S2, a yellow, oily, disagreeable liquid, insoluble in water, separates. It decomposes gradually at medium temperatures, more rapidly on warming, into hydrogen sulphide and sulphur :— H2S2=H2S-f S. It is generally supposed that the resulting hydrogen persulphide is similar in constitution to the peroxide and consists of hydrogen disulphide, containing an ex¬ cess of dissolved sulphur. 112 INORGANIC CHEMISTRY. As the calcium persulphide used is a mixture of CaS,, CaS3, and CaS^, it is probable that the oily liquid is a mixture of HgSj, HgSj and H^Sg. We must at least conclude that HjSg is present in it, because it unites with strychnine to form a crystalline compound. COMPOUNDS OF SULPHUR WITH THE HALOGENS. Sulphur and chlorine unite to form three compounds : SCI2, SCI4, and S2CI2. It is only the latter that meets with any practical application. Sulphur Dichloride—SCI2—is produced when S2CI2 is saturated with chlorine in the cbld:— ^2^12 H~ ~ zSClj. The excess of chlorine is removed by conducting a stream of COj through it. Fig. 61. It is a dark red-colored liquid, with a specific gravity of 1.62; boils at 64° C., with partial decomposition into SjClg and Clj; the dissociation commences at ordinary temperatures. Sulphur Tetrachloride—SCh—only exists at temperatures below o® C. It is formed by saturating SClj with CI at —30° C., and readily decomposes into SClj and Clj; the dissociation commences at —20° C., and is complete at 6®. It 5delds crystalline compounds with some chlorides, e. g., SnCl^ AsClj, SbCI,. The most stable of the sulphur chlorides is Sulphur Mono-chloride—S2CI2—which is formed when chlo¬ rine is conducted over molten sulphur. (Fig. 61.) It distils over and condenses in the receiver È ; the product is redistilled, to obtain it pure. SELENIUM. 113 Sulphur mono-chloride is a reddish-yellow liquid with a sha^ odor, provoking tears, having a specific gravity of 1.68, and boil¬ ing at 139° C. Its vapor density equals 67 (H = i), corresponding to the molecular formula S2CI2 = 134.7» It fumes strongly in the air, and is decomposed by water into sulphur dioxide, sulphur and hydrochloric acid :— 2S2CI2 + 2H2O = SO2 + 4HCI + 3S. Sulphur mono-chloride dissolves sulphur readily and serves in the vulcanization of caoutchouc. Bromine forms analogous compounds with S. S2Br2 is a red liquid, boiling at i90°-2oo° C. When gently heated, iodine unites with S to form S2I2. 3. SELENIUM. Se = 79.1. Scj =158.2. (at 1400° C.). This element is not very abundant in nature, and is only found in small quantities, principally in certain iron pyrites (in Sweden and Bohemia). Upon roasting this ore of iron, for the prepara¬ tion of sulphuric acid, selenium settles out in the chimney dust or in the deposit of the lead chambers (compare Sulphuric Acid), and was found there by Berzelius, in the year 1817. Like sulphur, selenium forms different allotropie modifications. Amorphous selenium^ obtained by the reduction of selenium dioxide (SeOz) by means of sulphur dioxide (SO2), is a reddish-brown powder, soluble in carbon disulphide, and has a specific gravity of 4.26. Selenium crystallizes from carbon disulphide in brownish- red crystals. The solution of potassium selenide is brown-red, and when it is exposed to the air, black leaf-like crystals of selenium separate. These are isomorphous with sulphur. Upon suddenly cooling fused selenium it solidifies to an amorphous, glassy, black mass, which is soluble in carbon disulphide and has a specific gravity of 4.28. When selenium (amorphous) is heated to 97° C., its temperature suddenly rises above 200° C. ; it is converted into a crystalline, dark gray mass with a specific gravity of 4.8. It pos¬ sesses metallic lustre, conducts electricity, and is insoluble in carbon disulphide. The crystalline, insoluble modification is obtained by slowly cooling the molten selenium. Selenium melts at 217°, and boils at about 700°, passing into a dark yellow vapor. The vapor density diminishes regularly with increas¬ ing temperature (similar to sulphur), and becomes constant at 1400° C. It then equals 79 ; the molecular weight is, therefore, 158, i. erim'ent, it was discovered that ^ volume of the air had disappeared and combined with the mercury to form red mercuric oxide. Lavoisier now strongly ignited the resulting mercuric oxide, and obtained a volume of oxygen equal to that withdrawn from the air during the experiment. By mixing this with the residual volume of N the original volume of air was again recovered. Thus it was demonstrated that air consists of f volumes N and ^ volume O gas. The elementary character of nitrogen was first established by Lavoisier in 1787. It was called azote (from life, and a, priva¬ tive) by him. The symbol Az, derived from azote, is used in France and England for nitrogen. The name nitrogenium (from which the symbol N) was given to nitrogen because it was a constituent of saltpetre (nitrum). Lavoisier made use of the above experiment for another important deduction. As he determined the weight, both of the employed mercury and the resulting mercuric oxide, he discovered that the increase in weight was exactly equal to that Fig. 64. of the oxygen withdrawn from the air, and by heating the mercuric oxide the same weight of oxygen was again separated. Thus was it demonstrated that the process of oxidation was the union of two bodies (not a decomposition), and that the weight of a compound body equals the sum of the weights of its constituents; the principle of the indestructibility of matter. Quantitative Composition of Air.—Its composition is expressed by the quantity of oxygen and nitrogen contained in it, as its remaining admixtures are more or less accidental and variable. Boussingault and Dumas determined the accurate weight composition of the air by the following experiment : A large balloon, with a capacity of about 20 litres (Fig. 64), is connected with a porcelain tube, a b, filled with metallic copper. Balloon and tubes, closed by stop-cocks, THE ATMOSPHERE. 121 are previously emptied and weighed apart. The bent tubes, A, B, and C, contain KOH and sulphuric acid, and serve to free the air undergoing analysis from aqueous vapor, carbon dioxide, and other impurities. The porcelain tube, filled with copper, is heated to a red heat, and by carefully opening the stop-cocks u, r, and a slow current of air is allowed to enter the empty balloon V. The impu¬ rities are given up in the bent tubes, and all the oxygen absorbed by the ignited Cu, forming cupric oxide, so that only pure nitrogen enters V. Now close the cocks and weigh the balloon and por¬ celain tube. The increase in weight of the latter represents the quantity of oxygen in the air ; the increase in V the quantity of nitrogen. In this manner Dumas and Boussingault found that in loo parts by weight of air there are contained :— Nitrogen 76.99 parts by weight. Oxygen 23.01 " " " As we know the specific gravity of nitrogen (14.01) and of oxygen (15.96), we can readily calculate the volume composition of air from that in parts by weight. We thus discover: — Calculating upon these numbers, we obtain 14.415 (H = i,0 = 15.96) as the specific gravity of air. The volume composition of air may be directly found by means of the absorptiometer. The latter is a tube carefully graduated, and sealed at one end. This is filled with mercury, and air allowed to enter ; the volume of the latter is determined "by reading off the divisions on the tube. Now introduce into the tube, through the mercury, a platinum wire having a ball of phosphorus attached to the end (Fig. 65), (or a ball of coke saturated with an alkaline solution of pyrogallic acid). The phosphorus absorbs the oxygen of the air, and only nitrogen remains, the volume of which is read off by the graduation. The eudiometric method affords greater accuracy. It is dependent upon the combustion of the oxygen with hydrogen in a eudiometer. The latter is an absorptiometer, having two platinum wires fused in its upper end (Fig. 66). Air and hydrogen are introduced into the eudiometer, and the electric spark then passed through the wires (Fig. 67). All the oxygen in the air combines with a portion of the hydrogen to form water. On cooling, the aqueous vapor con- 100.00 " Oxygen. Nitrogen Air 100.00 " 20.78 parts by volume. 79.22 " " " II 122 INORGANIC CHEMISTRY. denses and a contraction in volume occurs. Assuming that we had taken loo volumes of air and 50 volumes of hydrogen, and that the residual volume of gas, after allowing for all corrections (p. 123), equalled 87.15; then of the original 150 volumes of mixed gas, 62.85 volumes disappeared in the formation of water. As the latter results from the union of I volume of oxygen and 2 volumes of hydrogen, the 100 volumes of air em¬ ployed in the analysis there¬ fore contained §-2^ = 20.95 volumes of oxygen. From this air consists of 79.05 volumes nitrogen. Fig. 66. h= 20.95 100.00 oxygen. aur. Fig. 67. Numerous analyses show that the composition of the air everywhere on the earth's sur¬ face is constant. The opinion was long held that variations in composition equaling 0.5 per cent, existed at various periods of the year and in various localities. The recent extended investigations of Kreusler and Hempel prove these assumptions to have been based upon erroneous THE ATMOSPHERE. 123 determinations. The oxygen content of the air varies constantly from 20-91 to 20-93 cent, by volume. Measuring Gases.—The volume of gases is influenced by pressure, temperature, and the moisture contained in them. The volume of dry gases, at 760 mm. baro¬ metric pressure and 0° C., is accepted as the normal volume. If a gas has been measured under any other conditions it must be reduced to the normal volume. According to the law of Boyle and Mariotte, the volumes of the gases are inversely proportional to the pressure ; therefore, if the volume of the gas at pressure h, has been found equal to V, its volume at 760 mm. equals According to Gay-Lussac's law, all gases expand in proportion to the tempera¬ ture. Their coefficient of expansion is = o.3665 ; i. e., one volume of gas at o® occupies at 1° the volume 1.003665. If represents the observed gas volume at Vo, however, its volume at 0°, then Ft ir * I -|- 0.003665./ and, considering the pressure, Vth F„ = . 760(1 -j- 0.003665./) Further, the gas volume is enlarged by moisture, as the tension of the aqueous vapor opposes the atmospheric pressure. The moisture may be removed by intro¬ ducing into the gas a ball of coke saturated with sulphuric acid, which dries it. It is more convenient, however, to make the correction of the gas volume in the following manner : Water is brought in contact with the gas to be measured, in order to perfectly saturate it with aqueous vapor ; the gas is then measured and its normal volume calculated by the above formula, after deducting from the ob¬ served pressure h the number of millimeters* corresponding to the tension of the aqueous vapor for the given temperature (p. 91). From the great constancy of its composition air was supposed to be a chemical compound, consisting of nitrogen and oxygen. This supposition is, however, opposed by the following circum¬ stances. All chemical compounds contain their constituents in atomic quantities, which is not the case with air. In the mixing of nitrogen and oxygen to form air there is neither disengagement nor absorption of heat, which is always observed in chemical com¬ pounds. Further, the air absorbed by water or other solvents possesses a composition different from the atmospheric ; this is due to the unequal solubilities of nitrogen and oxygen in water. The air expelled from water upon application of heat consists of 34.9 volumes of oxygen and 65.1 volumes nitrogen. (Bunsen.) These facts indicate that air is not a chemical compound, but a mechanical mixture of its two constituents. *Fo=Vt — Fo .0.00366./, consequently Vo Fo .0.00366./ = Ft, and F« (I -j- 0.00366./) = Ft, 124 INORGANIC CHEMISTRY. Fig. 68. The great constancy in composition of the air depends on the mutual diffusion of the gases. As the gas molecules possess a direct progressive movement, they distribute themselves, without limitation, into space, and intermingle regularly with each other. The velocity of the diffttsion of gases is approximately inversely pro¬ portional to the square root of their densities—the law of the diffusion of gases. The density of hydrogen = I ; the density of oxygen = i6; therefore, hydrogen diffuses 4 times more rapidly than oxygen. The unequal diffusion of gases may be perceived if they are allowed to pass through very narrow apertures, or through porous partitions. The following experiment very clearly illustrates this : In the open end of an unglazed clay cylinder (as used in galvanic elements) there is puttied a glass tube about one meter long, its open end terminating in a dish containing water (Fig. 68) ; the cylinder and tube are filled with air. Over the porous cylinder is placed a wider vessel filled with hydrogen. The latter presses almost four times faster into the cylinder than the air escapes from it; the air in the tube and cylinder is displaced and rises in the water in bubbles. When the separation of gas ceases, tube and cylinder are almost filled writh pure hydrogen. On remtiving the larger hydrogen vessel the gas will escape much more rapidly into the external air than the latter can enter the cylinder; the internal pressure will therefore be less than the external, and water ascends in the glass tube. In addition to N and 0,air con¬ stantly contains aqueous vapor and carbon dioxide in very small quantities. The presence of the former can readily be recognized by the fact that cold bodies are covered with dew in moist air. Its quantity depends on the tem¬ perature and corresponds to the vapor tension of water (see p. 91). I cubic meter of air perfectly saturated with aqueous vapor contains 22.5 grams water at 25° C. ; on cooling too® 17.1 grams of these separate as rain. Generally the air contains only 50-70 per cent, of the quantity of vapor necessary for complete saturation. The amount of moisture in it is either determined according to physical methods (hygrometer), or directly by weigh¬ ing. To this end a definite quantity of air is conducted through a AMMONIA. 125 tube filled with calcium chloride or sulphuric acid, and its increase in weight determined. To detect the carbon dioxide in the air, conduct a portion of the latter through solutions of barium or calcium hydroxides, and a turbidity will ensue. To determine its quantity, pass a defi¬ nite and previously dried amount of air through a weighed potas¬ sium hydrate tube, and ascertain the increase in weight of the latter. 10,000 parts of atmospheric air contain, ordinarily, from 2.9-3.0 parts carbon dioxide. Besides the four ingredients just mentioned, air usually contains small quantities of ozone, hydrogen peroxide, and ammonium salts (ammonium nitrite). Finally, air contains microscopic germs of lower organisms; they are generally found in the lower air strata, and their presence infiuences the processes of the decay and fermentation of organic substances. COMPOUNDS OF NITROGEN WITH HYDROGEN. AMMONIA. NH3 = 17.04. Density = 8.52. Ammonia occurs in the air in combination with some acids, in natural waters and in the earth, but always in small quantities. The formation of ammonia by the direct union of nitrogen and hydrogen occurs under the influence of the silent electric discharge. Its compounds are frequently produced under the most varying conditions. Thus ammonium nitrate is formed by the action of the electric spark upon moist air :— Nj 4- O + 2H2O = NH4NO3. Ammonium nitrate. Small quantities of ammonium nitrite result by the evaporation of water in the air :— Nj + 2H2O = NH^NOa. Ammonium nitrite. The same salt isr formed in every combustion in the air ; by the rusting of iron and in the electrolysis of water. The white vapors which moist phosphorus forms in the air, consist of ammonium nitrite. Further, ammonium salts are produced in the solution of many metals in nitric acid, in consequence of a reduction of the acid by the liberated hydrogen :— HNO3 + 4H, = 3H3O -1- NH3. Ammonia is produced in large quantities in the decomposition 126 INORGANIC CHEMISTRY. and dry distillation of nitrogenous organic substances. Even as late as the last century the bulk of the ammonium chloride (the most important salt technically), was obtained by the distillation of camel's dung (in Egypt in the oasis of Jupiter Ammon—hence the name Sa/ ammoniacum). In the preparation of illuminating gas by the distillation of coal, ammonia appears as a by-product and may be obtained by combining it with sulphuric or hydrochloric acid. This method is used almost exclusively at present for its production. Fig. 6g. To prepare ammonia heat a mixture of ammonium chloride and slaked lime in a glass or iron flask :— 2NH4CI + Ca(0H)2 = CaCla + aH^O + 2NH3. Ammonium Calcium chloride. hydroxide. The disengaged ammonia gas is collected over mercury, as it is readily soluble in water (Fig. 69). For perfect drying conduct it through a vessel fllled with burnt lime (CaO). Calcium chloride is not applicable for this purpose, as it combines with the gas. In consequence of its levity, arhmonia, like hydrogen, may be col¬ lected by displacing the air in inverted vessels. Physical Properties.—Ammonia is a colorless gas with a suffocating characteristic odor. Its density is 8.5 (H = i), or 0.589 (air = i). Under a pressure of 6.5 atmospheres (at 10° C.), or by cooling to AMMONIA. 127 730 volumes of ammonia. —40° C., it condenses to a colorless, mobile liquid with a specific gravity of 0.613 solidifies at —80°. Fig. •jo. Ammonia gas may be condensed, just like chlorine. Take ammonium silver chloride (AgC1.2NH3), obtained by conducting am¬ monia over silver chloride, and enclose it in a tube with a knee-shaped bend (Fig. 70). The limb containing the compound is now heated in a water-bath, while the other limb is cooled. The compound is decomposed into silver chloride and ammonia, which condenses in the cooled limb. Ammonia gas dissolves very readily in water, with the liberation of heat. One part of water at 0° and 760 mm. pressure absorbs 1050 volumes (= 0.877 parts by weight); at 15° When a long glass tube, closed at one end and filled with ammonia, has its open end placed in water, the latter rushes up into the tube as it would into a vacuum ; a piece of ice melts rap¬ idly in the gas. The aqueous solution possesses all the properties of the free gas, and is called Liquor atnmonii caustici. The greater the ammonia content the less will the specific gravity of the solution be. The solution saturated at 40° contains about 30 per cent. NH3, and has a specific gravity of 0.897. -A-H the gas escapes on the application of heat. When the condensed liquid ammonia evaporates it absorbs a great amount of heat, and answers, therefore, for the production artificially of cold and ice in Carrè's apparatus. The simplest form of the latter is represented in Fig. 71. The cylinder A is filled about half with a con¬ centrated aqueous ammonia solution, and is connected, by means of the tubes from with the conical vessel in the middle of which is the empty cylindrical space E. The entire internal space of A and F is hermetically shut off. A is heated upon a charcoal fire until the thermometer a, in it, indicates 130° C., while F is cooled with water. In this way the gaseous ammonia is expelled from the aqueous solution in A, passes through b, in which most of the water runs back, and condenses to a liquid in B, of the receiver F. The cylinder A is removed from the fire, cooled with water and the vessel D constructed of thin sheet-metal and filled with water, placed in the cavity which is surrounded with a poor conductor, e. g., felt. The ammonia condensed in B evaporates, and is reabsorbed by the water in A. By this evapo- non Fig. 71. 128 INORGANIC CHEMISTRY. ration a large quantity of heat, withdrawn from i^and its surroundings, becomes latent ; the water in D freezes. The method of Carré for the artificial production of ice has acquired great appli¬ cation in the arts ; recently, however, it has been more and more replaced by the tnethod of Windhausen. The latter depends upon the expansion of compressed air. Chemical Properiies.—A red heat and continued action of the electric spark decompose ammonia into nitrogen and hydrogen. On conducting ammonia gas over heated sodium or potassium, the nitrogen combines with these metals and hydrogen escapes :— NH3 + 3K = NK, + 3H. Ammonia will not burn in the air ; in oxygen, however, it burns with a yellow flame :— 2NH, + 30 = N3 + 3H3O; ammonium nitrite and nitrogen dioxide are formed simultaneously. When a mixture of ammonia and oxygen is ignited it burns with explosion. To show the combustion of NHg in O, proceed as follows : A glass tube, through which ammonia is conducted, is brought into a vessel with oxygen, bringing the opening of the latter near a flame at the moment of the introduction of the glass tube. In contact with oxygen, the ammonia gas ignites and continues to burn in it. The following experiment (of Kraut) shows the combustion of ammonia very conveniently. Place a somewhat concentrated ammo¬ nia solution in a beaker glass; heat over a lamp, until there is an abundant disengagement of gas, and then run in oxygen gas, by means of a tube dipping into the liquid. Upon approaching the mixture with a flame, it ignites ' ignition may be induced without - a flame, by sinking a glowing platinum spiral into the mixture (Fig. 72) ; we then have a number of slight explosions. The glass is filled at the same time with white vapors of ammonium nitrite (NHiNOj) ; later, when oxygen pre¬ dominates, red vapors of nitrogen dioxide (NOa) and nitroqs acid appear. AMMONIA. 129 If chlorine gas be conducted into the vessel with ammonia, it immediately ignites and continues to burn in the latter, with the production of white fumes of ammonium chloride (NH4CI). The chlorine combines with the hydrogen of the ammonia, with sepa¬ ration of nitrogen, and yields hydrochloric acid, which unites to form ammonium chloride with the excess of ammonia. NH3 + 3CI = 3HCI + N, and 3NH3 + 3HCI = 3NH^C1. Chlorine reacts similarly upon aqueous ammonia (p. 118). In gaseous form, as well as in solution, ammonia possesses strong basic properties ; it blues red litmus paper, neutralizes acids, form¬ ing salt-like compounds with them, which are very similar to the salts of the alkalies—sodium and potassium. The following illus¬ trates the similarity :— NH3 + HCl = NH^Cl KG! Ammonium Potassium chloride. chloride. (NHJ^SO, K3SO, Am. sulphate. Potassium sulphate. NH^SH KSK. Am. Potassium sulphydrate. sulphydrate. In these ammonia derivatives NH^ plays the rôle of the metal potassium. Hence the group (NH^) has been designated Ammonium and its compounds, ammonium salts. The latter, when acted on by strong bases, yield ammonia gas :— 2NH^CI + CaO = 2NH3 + CaClj + H^O. The metallic character of the ammonium group is confirmed by the existence of the ammonium amalgam, and likewise by its entire deportment in compounds. Therefore, the ammonium derivatives will be considered with the metals. Thermo-chemical Deportment.—The heat of formation of ammonia from hydrogen and nitrogen equals 11.8 Cal. When ammonia gas is dissolved in much water 8.8 Cal. are set free, so that the heat of formation of ammonia from its elements in dilute aqueous solution equals 20.6 Cal. :— (N.Hj — gas) = 11.8. (NHg, Aq) = 8.8. (N.Hj, Aq) = 20.6. Although an exothermic compound, ammonia is produced from its elements with difficulty, and in turn is rather easily dissociated into them. The rather great heat of solution of gaseous ammonia explains why ice will melt in the same (p. 127). 2NH3 4- H^SOi = NH3 + Hß = 130 INORGANIC CHEMISTRY. The explosibility of a mixture of ammonia and oxygen is accounted for by the following large heat disengagement :— 2NH3 + 30 = 3H2O + N, . . . . (+ 148,0 Cal.) (23.6 Cal.) (3 X 57-2 Cal.) The action of chlorine upon gaseous or aqueous ammonia is also very energetic :— NHg gas -j- 3CI = 3HCI gas + N . . . . (-j- 54-2 Cal.) (11.8 Cal.) (66 Cal.) NHj-dissolved -{- 3CI = 3HCl-dissolved 4- N . . . (-f- 97«3 Cal.) (20.6 Cal.) (3 X 39-3 Cal.) When there is an excess of ammonia the hydrochloric acid com¬ bines with it to form ammonium chloride (NH3 -j- HC1= NH4CI), and in doing this both the heat disengagement and the energy of reaction are raised still further. QUANTITATIVE COMPOSITION OF AMMONIA. ATOMIC WEIGHT OF NITROGEN. The quantitative analysis of ammonia shows that it consists of i part hydrogen and 4.67 parts nitrogen; hence we conclude that the atomic weight of N is a multiple of the last number (see p. 98) :— H = I 2H = 2 3H = 3 N = 4.67 N = 9.34 N = 14.04 NH = 5.67 NHj = 11.34 NH3 = 17.04 As the density of ammonia equals 8.5 (H = i), its molecular weight would almost = 17. In 17.04 parts of ammonia there are 3 parts, and, therefore, 3 atoms of hydrogen. That the 14.04 parts nitrogen united with them correspond to one atom of N is a con¬ sequence, as never less than 14.04 parts of N are present in the molec¬ ular weight of any nitrogen derivative. The density of nitrogen equals 14.04, and its molecular weight 28.08 ; therefore, the mole¬ cule of N consists of two atoms (Nj). This is also concluded from the volume ratios occurring in the formation of ammonia. (See below.) From the molecular formulas NH., and Nj it follows, further, that i volume N and 3 volumes H form 2 volumes ammonia gas, or that 2 volumes NH, decompose into 3 volumes Ha and i volume Na, cor¬ responding to the molecular equation :— Na -f 3Ha = 2NH,. 1 vol. 3 vols. 2 vols. The following experiments prove these conclusions :— I. Decompose an aqueous ammonia solution, mixed with sulphuric HYDROXYLAMINE. 131 acid to increase its power of conductivity, in a Hofmann's appa¬ ratus (Fig. 47), by the galvanic current. Hydrogen will separate at the negative and nitrogen at the positive pole ; the former will have three times the volume of the latter. 2. The electric (induction) sparks are permitted to strike through dry ammonia gas enclosed in a eudiometer, or the apparatus repre¬ sented in Fig. 58. In this way the ammonia is decomposed into nitrogen and hydrogen, whose volume is twice as large as that of the ammonia employed. That 3 vols. H are present in the mixture for every vol. N is easily shown by the volumetric method, by burning the H with oxygen (p. 121). The volume ratios in the formation of ammonia confirm the conclusion drawn from the density of nitrogen (see above), that the molecule of the latter consists of two atoms. In two volumes of ammonia there are 2« molecules of NHj, therefore 2« atoms of N. The nitrogen contained in these 2 volumes of NHg occupies i volume in a free condition, and this contains » molecules and therefore 2« atoms of N. HYDROXYLAMINE. NH3O = NHjOH. This compound, very analogous to ammonia, was discovered (by Lossen) in the reduction of ethyl nitrate by tin and hydrochloric acid. It is produced, too, by the action of tin upon dilute nitric acid, and by tin and hydrochloric acid upon all the oxygen compounds of nitrogen. In all these reactions it is the hydrogen eliminated by the tin which, in statu nascendi, reduces the nitric Rcid * HNO3 + 3H2 = H3NO -f 2H,0. To prepare hydroxy lamine treat ethyl nitrate (l20 gr.) with granulated tin (400 gr.) and hydrochloric acid (800-1000 c.c. of specific gravity 1.19, mixed with three times its volume of water) until solution is obtained. The strongly concentrated liquid is cooled and supersaturated with soda, the filtrate slightly acidulated with hydrochloric acid and then evaporated to dryness. Hot alcohol will extract hydroxylamine hydrochloride, NH3O.HCI, from the residue. Hydroxylamine hydrochloride is most easily formed by acting with hydrochloric acid upon fulminating mercury (see Organic Chemistry). Hydroxylamine is very similar to ammonia, and like it unites directly with the acids to form salts :— H3NO + HCl = HsNO.HQ. The hydrochloride in distinction to ammonium chloride is per¬ fectly insoluble in absolute alcohol. It passes into ammonium chloride when allowed to stand exposed to the air. On adding to the aqueous solution of the sulphate of hydroxyl¬ amine sufficient barium hydroxide to remove all the sulphuric acid, an aqueous solution is obtained, which, like the ammonia solution, possesses strong basic properties, and blues red litmus 132 INORGANIC CHEMISTRY. paper. The solution is, however, very unstable, and readily decom¬ poses into water, ammonia, and nitrogen :— 3NH3O = NH3 + 3H3O + N3. Upon the application of heat a portion of the hydroxylamine will be carried over undecomposed along with the steam, but most of it is broken up. The hydroxylamine solution manifests a reducing action ; it precipitates metallic silver from silver nitrate, white mer- curous chloride, HgCl, from mercuric chloride, HgClj, and cuprous oxide from cupric salts. Owing to its great similarity to ammonia and its various reac¬ tions, it is supposed that hydroxylamine represents ammonia in which I H is replaced by the hydroxyl group OH ; therefore the name hydroxylamine :— NH3O = NH2OH. Diamide or Hydrazine, NjH^ rzr.HjN.NHj, a compound of two amido-groups (NHj), was until recently only known in its numerous organic derivatives. Curtius has at last succeeded in isolating it. Its salts, e. g., NjH^.aHCl, result from the transposition of a complex diazofatty acid. Alkalies liberate the free diamide from the latter. It is a stable gas with peculiar odor. It is very similar to ammonia, dissolves easily in water, colors red litmus blue and combines with the acids (2 equiv.) to form salts. One of its characteristic reactions is the reduction of silver and copper salts. Its salts will be described under the metals as diamide or diammonium salts. Hydrazoic Acid, HN3, Azoimide, is an interesting derivative of hydrazine. It is formed under proper conditions by the action of nitrous acid upon diamide. Ammonia and nitrous acid yield nitrogen, and hydrazine monochloride and nitrites should form hydrazoic acid :— I. NH^Cl -f- NOjNa = N, -j- 2H,0 4- NaQ. Ammonium Sodium Chloride. Nitrite. 2. NHj.HCl N. I -f NOjNa = II >NH 2HjO -f NaQ. NH, N/ Hydrazine Sodium Chloride. Nitrite. The direct preparation of the acid from hydrazine is, however, difficult. An easy method for its preparation consists in acting upon hippuryl hydrazine (CjHgOjj.NH.NHj, with sodium nitrite and acetic acid. The product is boiled with acids or alkalies, when the resulting hydrazoic acid is> carried over with steam. At ^he ordinary temperature hydrazoic acid is a gas, producing dizziness, headache, and inflammation of the membrane of the nasal cavity. Its aqueous solution has a penetrating odor. Blue litmus, held over the liquid, is colored an intense red. Its vapors form dense clouds with ammonia. A 70 per cent, solu¬ tion of the acid dissolves iron, zinc, copper, aluminium and magnesium energeti¬ cally. In concentrated form it appears to dissolve gold and silver. Its metallic salts are very similar to the chlorides. The acid itself is distinguished from the haloid acids by its extremely "explosive character. Silver nitrate and mercurous nitrate precipitate it quantitatively from its solutions. Silver hydrazoide, AgN,, is extremely like silver chloride, but is not changed on exposure to the light {Berichte, 23, 3023). NITROGEN IODIDE. 133 COMPOUNDS OF NITROGEN WITH THE HALOGENS. NITROGEN CHLORIDE. NCla. As we have seen, nitrogen is liberated when chlorine acts upon an excess of ammonia (p. 129); when, however, the chlorine is in excess, it acts upon the previously formed ammonium chloride, to produce nitrogen chloride :— NH^Cl -j- 3CI3 = NCI3 -f 4HCI. For the preparation of a small quantity of nitrogen chloride, dip a flask filled with chlorine, open end down, into an aqueous ammo¬ nium chloride solution, warmed to 30°. The chlorine is absorbed, and heavy oil drops separate, which are best collected in a small leaden dish. Nitrogen chloride is an oily, yellow liquid, with a disagreeable odor ; its specific gravity equals 1.65. Of all chemical compounds this is the most dangerous, as it decomposes by the slightest con¬ tact with many substances, and frequently, too, without any percep¬ tible external cause. Its decomposition is accompanied by an extremely violent report. Aqueous ammonia gradually decomposes it into ammonium chloride and nitrogen :— NCI3 + 4NH3 = sNH^Cl -f N^. It is converted into ammonium chloride and free chlorine by concentrated hydrochloric acid :— NCI3 + 4HCI = NH4CI -b 3CI2. This reaction is directly opposed to that by which nitrogen chloride is formed. The formation and explosibility of nitrogen chloride may be illustrated in a harmless way as follows : Decompose a saturated ammonium chloride solution with the electric current. Nitrogen chloride rising in small drops from the liquid will separate at the positive pole. Upon covering the surface of the solution with a thin layer of turpentine oil, each drop will explode as it comes in contact with the latter. Nitrogen Iodide. Upon adding ammonium hydroxide, or a mixture of ammonium chloride and caustic soda, to a solution of iodine in aqueous potassium iodide, a brownish black powder sep¬ arates. Its composition closely approximates the formula, NIjH. Its formation by means of ammonium chloride and caustic soda is represented in the equation ;— 4I 4- NH^Cl 4- 3NaOH = NI^H 4- 3HjO 4- NaCl -f- 2NaI. When the conditions are slightly changed a very similar com- 134 INORGANIC CHEMISTRY. pound separates. Its formula is N2l3H3(= NH3 + Nig). Pro¬ tracted washing with water decomposes it into ammonia and nitrogen tri-iodide, NI3. Nitrogen Di-iodide and Nitrogen Tri-iodide, NHIjand NI3, are, when dry, very explosive. The explosibility may be shown without danger in the following manner : The precipitate is col¬ lected on a filter, washed with water, the filter opened out and torn into small pieces, which are then allowed to dry ; upon the slightest disturbance these pieces explode with a sharp report. Nitrogen iodide dissolves in dilute hydrochloric acid and decom¬ poses into ammonia and iodine chloride :— NHjI + HCl = NH3 -f ICI. Hydrogen sulphide and sulphurous acid convert it into ammonia and hydrogen iodide. The nitrogen iodide formed by digesting powdered iodine with ammonia water manifests properties that are slightly different from those of the ordinary iodide. It is only stable in the presence of ammonia. It sometimes explodes even when moist—if it be washed with water, or when acted upon by hydrochloric acid. Thermo-chemical Deportment.—Nitrogen chloride and iodide are both strongly endothermic compounds ; considerable heat is absorbed in their production from the elements :— (N.Cy = — 38.1 Cal. This being the case, they can only be obtained from their constituents by the addition of energy from without. Yet this formation does not occur, because the slightest external impulse occasions their decomposition. When they are formed from ammonium chloride by the action of chlorine (iodine) it is at the expense of the total heat of transposition, which continues positive :— NH^Cl + 6C1 = NCI3 + 4HCI Aq. . . . 4- (81.4 Cal.) (75.8 Cal.) (— 38.1 Cal.) (4 X 39-3 Cal.) Inasmuch as these compounds contain a great deal more energy than their elements, they exist in a very uncertain equilibrium, and may be readily decom¬ posed with explosion—they are very explosive (see p. 29). 2. PHOSPHORUS. P = 31.03. P4 = 124.12. Density = 6s.o6. This element does not occur free in nature, because of its very great affinity for oxygen. The phosphates, especially calcium phosphate, are widely distributed. By the disintegration of the minerals containing phosphates the latter pass into the soil, are ab¬ sorbed by plants, and remain in their ash. In the animal kingdom calcium phosphate occurs in the bones. PHOSPHORUS. 135 Brand and Kunkel, in Hamburg (1669), first obtained phospho¬ rus by the ignition of evaporated urine. In 1769, Scheele, in Sweden, showed that it could be obtained from bones. Its name is derived from its power of giving light in the dark—(paxTKpôpoç, i. e , light-bearer. To obtain phosphorus from bones the latter are burned, thereby destroying all organic admixtures and leaving bone ashes, which consist principally of tertiary calcium phosphate (P04)2Ca3 (see Phosphoric Acid). The ashes are digested with of their weight of sulphuric acid, when the tri-phosphate becomes primary calcium phosphate, and g)rpsum (cal. sulphate) is produced :— Ca3(POj2 + 2H2SO4 = CaH^(POj2 + 2CaS0^. Tertiary Primary Calcium calcium phosphate. calc. phosphate. sulphate. The gypsum, which dissolves with difficulty in water, is separated from the readily soluble primary phosphate by filtration; the solution is mixed with char¬ coal, evaporated in leaden pans, and the residue raised to a red heat. This expels water from the primary phosphate and the latter changes to calcium metaphos- phâtc ' CaH,(PO,)2 = Ca(P03)2 + 2H2O. Calcium metaphosphate. The ignited residue is then raised to a white heat, in retorts of infusible clay. The carbon partly reduces the metaphosphate to phosphorus, by forming carbon monoxide with oxygen, and half of the phosphorus contained in the metaphos¬ phate remains as calcium pyrophosphate :— 2Ca(P03)2 "1" 5^ = 2P 5^0 4" Ca2P20,j. Calcium pyrophosphate. The liberated phosphorus escapes in vapor form, and is collected and condensed under water in receivers of peculiar construction. T o remove mechanically admixed impurities the phosphorus is again distilled from retorts and fused under water ; it is then moulded into sticks. The crystalline or yellow phosphorus obtained by distillation is a waxy, transparent, slightly yellow-colored substance, with specific gravity of 1.83 at 10° C. At ordinary temperatures it is soft and tough; at 0° it becomes brittle. It fuses underwater at 44.4° and boils at 290° C. (278.3°). By the action of sunlight it be¬ comes yellow, and is coated with a non-transparent, reddish-white layer. Phosphorus is insoluble in water, slightly soluble in alcohol and ether, and very readily soluble in carbon disulphide. It crystallizes from the latter solution in forms of the isometric (rhom¬ bic dodecahedra) system. When exposed to moist air, it oxidizes to phosphorous acid (HjPOs) ; the white vapors which arise contain ammonium nitrite (NH4NO2), ozone and hydrogen peroxide. Its odor resembles that of ozone. In the air it phosphoresces at night. It does this also in other gases, but only in such as contain 136 INORGANIC CHEMISTRY. oxygen. It appears the phosphorescence is influenced by the formation and combustion of the self-inflammable phosphine, as all substances which destroy the latter, prevent and put an end to the former. It is noteworthy that in pure oxygen the oxidation of phosphorus begins at 27°. If the oxygen be diluted by re¬ moval over an air pump or by the addition of neutral gases, so that its quantity is not more than 40 per cent., the absorption will be very energetic at 20°, but cease entirely at 7°. Another modification—the red or amorphous phosphorus—^pos¬ sesses properties entirely different from the ordinary variety. It is a reddish-brown amorphous powder, of specific gravity 2.14; in¬ soluble in carbon disulphide, non-phosphorescent, does not alter in the air, and is, indeed, very stable. While ordinary phosphorus is very poisonous, this variety is perfectly harmless. It does not fuse at a red heat, even when subjected to strong pressure, and vaporizes very slowly (above 260°), although only partially, the vapors pass¬ ing over into ordinary phosphorus. To prepare the red variety, yellow phosphorus is heated for some minutes to 300°, in closed, air-tight iron vessels ; there is a partial conversion at 250°. The resulting mass is then treated with car¬ bon disulphide or sodium hydroxide, to withdraw the unaltered, ordinary phosphorus. If some iodine be added to the ordinary phosphorus, the change will occur below 200°. A third modification—metallic phosphorus—is formed if the amorphous variety be heated in a glass tube, free of air, to 530°. Microscopic needles then sublime into the upper, less heated, portiori of the tube. It is more easily obtained if phos- phorus is heated with lead, in a closed tube, to a red heat. The molten metal dissolves the phosphorus, and on cooling, the latter separates in black, metallic, shining crystals. Metallic phosphorus possesses the specific gravity 2.34, vaporizes with difficulty, and is less active than the amorphous variety. Two green lines characterize the spectrum of phosphorus. On conducting hydrogen over a small piece of phosphorus, heated in ß. glass tube, the escaping gas will burn with a bright green flame. When ordinary phosphorus is distilled with water, some passes over with the steam and, in the dark, phosphoresces. This procedure serves for the detection of phosphorus in poisoning by this substance. The density of P equals 62.06 (H = i), or 4.29 (air=: i); the molecular weight is, therefore, 124.12. As the atomic weight of P is 31.03, it follows that the molecule in the form of vapor consists of 4 atoms: Pi = 124.12 (31.03 X 4). We saw that the sulphur molecule at 500° consists of 6 atoms (Se), and at 900° of 2 atoms (82). Such a dissociation does not, however, occur with phos¬ phorus; even at 1040° its vapor density remains unaltered, although a partial dissociation does take place at a very intense heat. PHOSPHORUS. 137 When phosphorus is burned in oxygen or in air, it forms the pentoxide (P2O5). The ordinary variety inflames at 40°, and also by gentle friction ; the amorphous is not ignited below 260°. The first will burn with a bright flame even under water. To this end heat some pieces of P in a flask with water, until they fuse, and conduct a current of oxygen through the water. Phosphorus com¬ bines very energetically with CI, Br, and I at ordinary temperatures ; by throwing a small piece of it into a vessel containing dry chlorine gas it at once inflames. The red only reacts with the halogens after applying heat. With most of the metals phosphorus unites on warming, and precipitates some of them from solutions of their salts. From a silver nitrate solution, it precipitates silver and phosphorus- argentide (PAgj) ; this solution, therefore, answers as a counter- irritant in phosphorus burns. The difference in dèportment of the yellow and the amorphous phosphorus is fully accounted for by the circumstance, that when the amorphous is produced from the yellow there follows a considerable heat-disengagement :— P-yellow = P-amorphous -}- 19.2 Cal. Hence, the red variety contains much less energy than the yellow. In its union with other substances there will always be liberated 19.2 Cal. less, and the reac¬ tion consequently will proceed more sluggishly and with less energy. COMPOUNDS OF PHOSPHORUS WITH HYDROGEN. PH3, P2H4, P^Hg. The compounds of phosphorus with hydrogen can be prepared by the action of nascent hydrogen upon phosphorus, as, for example, on gently heating dilute sulphuric acid with zinc and phosphorus (p. 145.) The usual course is to heat yellow phosphorus with concentrated potassium or sodium hydroxide, when sponta¬ neously inflammable phosphine will escape and a salt of hypophos- phorous acid enter solution. The liberated gas mixed with air in a closed vessel explodes violently, hence to make it proceed as follows; Fill a glass flask almost full of aqueous KOH, add a few pieces of P, and heat over a lamp (Fig. 73). When the liberation of gas commences, and the air in the neck of the flask has been expelled, close the same with the cork of the delivery tube, the other end of which dips under warm water, to prevent any obstruction arising in it from phosphorus that may be carried over and solidify by cooling. Each bubble rising from the liquid inflames in the air, and forms white cloud-rings which ascend. The gas thus produced consists of gaseous phosphine (PHg) and hydrogen, with which is mixed a small quantity of a liquid sub¬ stance (P2H4), whose presence imparts the spontaneous inflamma- 12 138 INORGANIC CHEMISTRY. bility to the gas. On conducting the latter through a cooled tube the P2H4 is condensed to a liquid, and the escaping gas no longer inflames spontaneously. The liquid compound may be isolated in a similar manner if the gas is conducted through alcohol or ether, which will absorb the compound P2H4. Liquid Phosphine, PgH^, separated from the gas by cooling, is a colorless, strongly refracting liquid, insoluble in water, and boiling at 30®. It inflames spontaneously in the air, and burns with great brilliancy to phosphorus pentoxide and water. Its presence in combustible gases, such as hydrogen, marsh gas, and PH3, gives to them their spontaneous inflammability. In contact with some Fig. 73. compounds, like carbon and sulphur, and by the action of sunlight, it decomposes into gaseous and solid phosphine :—• SP2H4 = 6PH3 + P.H^. Solid Phosphine, P4H2,(?) is a yellow powder, inflamed at 160® or by a blow. It is produced in the decomposition of calcium phosphate by hydro¬ chloric acid. Gaseous Phosphine, PHg, may be formed, in addition to the manner previously described, by the action of water or hydro¬ chloric acid upon calcium phosphide ;— Ca-3P2 -f 6HC1 =3CaCla + 2PHg. PHOSPHORUS. 139 Further, by the ignition of phosphorous and hypophosphorous Hcids 4H3PO3 = PH3 + 3H3P0,. Phosphorous Phosphoric acid. acid. It is a colorless gas, with a disagreeable, garlic-like odor, and is somewhat soluble in alcohol. Its density is 17.01 (H = i), or 1.176 (air = i). When pure—freed of P2H4—it ignites at 100° C. Oxidizing agents convert it again into the spontaneously inflammable variety, owing to the production of PaHi- It is extremely poisonous. Phosphine is decomposed into phosphorus and hydrogen when it is heated, or if it is exposed to the action of the electric spark. When ignited in the air it burns with a brightly luminous flame, disseminating at the same time a white cloud of phosphorus pentoxide (P2O5). Heat and the electric current decompose PH3 into phosphorus and hydrogen. When mixed with chlorine it explodes violently, with production of phosphorus trichloride and hydrogen chloride :— PH3 + 3CI3 = PCI3 + 3HCI. Like ammonia, phosphine possesses faint alkaline properties, and combines with hydrogen iodide and bromide to yield compounds similar to ammonium chloride ;— PH3 -f HI = PHJ. It combines with HCl at —30° to —35°, or, at ordinary tem¬ peratures, under a pressure of 20 atmospheres. The group PH4, figuring in the rôle of a metal in these compounds, is analogous to ammonium (p. 129), and termed Phosphonium. Phosphonium Iodide^ PHJ. It is best prepared by the decom¬ position of phosphorus di-iodide (PI2), by a slight quantity of water, or by adding yellow phosphorus (10 parts), and, after some hours, iodine (2 parts), to a saturated solution of hydriodic acid (22 parts). The liquid becomes a solid mass, consisting of phos¬ phonium iodide and phosphorous acid. Phosphonium iodide sub¬ limes in colorless, shining, cube-like rhombohedra ; fumes in the air, and, with water, decomposes into PH3 and HI. When decom¬ posed by potassium hydroxide it yields pure hydrogen phosphide, which is not spontaneously inflammable :— PH4I -f KOH = KI -f PH3 + H,0. Phosphine is a feebly exothermic compound :— P-yellow -f 3H = PHj -j- 11.6 Cal. This explains why in its power to react it differs so little from the elements com¬ posing it. 140 INORGANIC CHEMISTRY. MOLECULAR FORMULA OF PHOSPHINE. ATOMIC WEIGHT OF PHOSPHORUS. The analysis of phosphine shows that it consists of I part hydrogen and 10.34 parts phosphorus. Were its molecular formula PH, the atomic weight of P would be 10.34. The great analogy of phosphine to NH3, and that of all the P com¬ pounds to those of N, argues, however, for the formula PH3. The atomic weight of P, therefore is 31.03 (= 3 x 10.34), and the molecular weight of the phos¬ phine is 34.03 :— PH3 = 34.03 This view is confirmed by the density which, according to its formula, must be 34^ _ ly.oi. Direct experiment confirms this. Further, from the formula PH3 it follows that 3 volumes of hydro¬ gen are present in 2 volumes of the gas :— 2PH3 contain 3H2, 2 vols. 3 vols. or in I volume there are volumes of hydrogen. On decom¬ posing the gas in a eudiometer by means of electric sparks, it will be found that the volume increases i ^ times ; the gas consists, then, of pure hydrogen, while phosphorus separates in a solid con¬ dition. As the phosphorus molecule in the gaseous condition is composed of 4 atoms, the phosphorus (62.06 parts) separated from 2 volumes of PHg, will fill volume when in the form of vapor ; hence in 2 volumes of PH3 there are present 3 volumes of H and ^ volume of phosphorus vapor. Or, written molecularly :— P, + 6H2 = 4PH3. I vol. 6 vols. 4 vols. COMPOUNDS OF PHOSPHORUS WITH THE HALOGENS. Phosphorus combines directly with the halogens to yield com¬ pounds of the forms PX3 and PX5, in which X indicates an hal¬ ogen atom. Phosphorus T richloride — Phosphorous Chloride—PClj. Conduct dry chlorine gas over phosphorus gently heated in the retort D (Fig. 74). The phosphorus ignites in the stream of gas, and distils over as trichloride, which is collected in the receiver and condensed. The product is purified by a second distillation. It is a colorless liquid, boiling at 74° C., and has a sharp, peculiar PHOSPHORUS PENTACHLORIDE. Ml odor. Its specific gravity equals 1.616 at 0°. It fumes strongly in the air, and is decomposed by moisture into phosphorous and hydro¬ chloric acids :— PCIg 4- 3HP = H3PO3 + 3HCI. The vapor density of the trichloride equals 68.6 (H = i), cor¬ responding to the molecular formula PCl3= 137.3. Phosphorus Pentachloride—Phosphoric Chloride—PCI5.— This is produced by the action of an excess of chlorine upon the liquid trichloride. It is a solid, crystalline, yellowish-white com¬ pound. It fumes strongly in the air and sublimes without fusion when heated. It at the same time sustains a partial decomposition into trichloride and chlorine. Fig. 74. At lower temperatures (in an atmosphere of chlorine) the vapor density of the pentachloride has been found to be 104.1, corresponding to the molecular formula PCI5 = 104.1. At increased temperatures the vapor density steadily diminishes, and a gradual decomposition occurs—dissociation (p. 94) of the mole¬ cules PCI5 into the molecules PCI3 and Clj. The dissociation is complete at 236°, and then equals the vapor density 52 ; i- í.» the vapor then fills a volume twice as large as at a lower temperature. The breaking up of PCI5 into PCI3 and CI, explaihs this :— I vol. PCI5 = I vol. PCI3 -|- I vol. CI,. That such a decomposition of the penta- into trichloride and chlorine does really occur, is proven, among other things, by the originally colorless vapor gradually assuming the yellow color of chlorine as the temperature rises. The decomposi¬ tion products—PCI3 and CI,—may be separated from each other by diffusion (p. 124). 142 INORGANIC CHEMISTRY. PCI5 acts very energetically with water, when it yields phosphoric acid (see this) and hydrochloric acid. With a little water it forms the oxychloride and hydrochloric acid :— PCI5 + Hp = PCI3O + 2HCI. Phosphorus Oxychloride—POCI3, is a colorless liquid, fum¬ ing strongly in the air, with a specific gravity, at 12°, of 1.7. It boils without decomposition at 110°. Its vapor density equals 76.7, corresponding to the molecular formula P0Cl3= 153.4. Water decomposes it into metaphosphoric and hydrochloric acids :— POCI3 + 2H2O = HPÖ3 + 3HCI. This oxychloride may be obtained by decomposing the penta- chloride with a little water (see above), or by letting it gradually deliquesce in moist air. The most practical method is the distil¬ lation of PCI5, with an excess of phosphorus pentoxide :— or with crystallized boric acid (5 pts. with i pt.) :— aPClg + 2BO3H3 = 3POCI3 + B3O3 + 6HC1. Its production on conducting ozonized air through phosphorous chloride is quite interesting :— PCI3 4-03 = Pa30 + O3. Potassium chlorate acts quite energetically upon phosphorous tri¬ chloride, with formation of the oxychloride (Dervin) :— 3PCI3 + CIO3K = 3PCI3O + KCl. Phosphorus Sulpho-chloride, PCI3S, is analogous to the preceding com¬ pound. It results from the union of the trichloride with sulphur when heated together to 130°, and also by the action of the pentachloride upon hydrogen sulphide or some metallic sulphides :— PCI5 + HjS = PCI3S -f 2HCI. It is a colorless liquid of sp. gr. 1.6 and boils at 124°. It fumes in the air and is decomposed by water according to the following equation :— PCI3S + 4H2O = P0iH3 + 3HCI + H,S. The bromine and iodine phosphorus compounds are perfectly analogous to the chlorine derivatives. They are obtained by uniting the constituents in the proportions by weight expressed by their for¬ mulas. As the union is exceedingly energetic, it is best to proceed as follows : Dissolve the phosphorus in carbon disulphide, gradually add the calculated amount of Br or I, and then distil off the vola¬ tile solvent. ARSENIC. 143 Phosphorus Tribromide, PBrj, is a colorless liquid, boiling at 175°, and having a specific gravity of 2.7. The pentabromide, formed hy the gradual addition of aBr to PBrg, is a yellow, crystalline substance, which fuses when heated, and breaks up into PBrg and Br^. Water decomposes both compounds, as it does the corresponding chlorides. Phosphorus oxy-bromide (POBrg) is a colorless crystalline mass, fusing at 45°, and boiling at 19S®. Phosphorus Chlor-bromide, PCigBrg, is produced by the union of PCI3 with Br, in the cold. It is a yellowish-red mass, which decomposes at 35° C., into PClj and Brj. Phosphorus Tri-Iodide—PI3, forms red crystals, fusing at 55° and distils, with partial decomposition, at a higher temperature. The so-called phosphorus iodide, PIj, or Pgl^ (corresponding to PgH^), crystallizes in beautiful orange-red needles or prisms, and fuses at 110°. Its vapor density at 265® and 90.7 mm. pressure equals 234, corresponding to the molecular weight Pgl^. A little water decomposes it into phosphorous acid, PHj and HI. The last two bodies then form phosphonium iodide, PH^I (p. 139). The recently discovered Phosphorus Pentafluoride—PFlg—is interesting. It results upon heating PCI3 or PClg with arsenic trifluoride, AsFlj ;— "i" It is a colorless gas that fumes in moist air and is decomposed by water into phosphoric acid and hydrogen fluoride. Its density is 63, corresponding to the molecular formula PFI5 = 126.03. If I*® liquefied at 16° under a pressure of 46 atmospheres, and solidifies when the pressure is removed. It is rather remarkable that although phosphorus pentaiodide could not be obtained, the stability of the compounds, PBrg, PClg, PFlg, gradually increases with the diminution of the atomic weight of the combined halogens. PFlg can be gasified without decomposition. Thermo-chemical Deportment.—^While the halogen derivatives of nitrogen (like those of oxygen) are strongly endothermic, are produced with the absorption of much heat, and are, in consequence, readily exploded (p. 134), those of phos¬ phorus are exothermic. The heat disengaged in the union of yellow phosphorus and chlorine (p. 137) corresponds to the following symbols:— (P,Cl3) = 75*3 > (Pjdsj) = Ï04-9 j (P,Cl3,0) = 142.6. In this we observe a transition to the halogen derivatives of the metals, all of which are exothermic. In accordance with this (just as with the metals) we find that the heat of formation of the bromides and iodides diminishes in regular suc¬ cession :— (P,Br3)=42.6; (P,l3) = 10.9. The great reactivity of all these derivatives with water is fully explained by the large amount of heat set free at the same time. 3. ARSENIC. As = 75. AS4 = 300. Vapor density, 150, Arsenic is a perfect analogue- of phosphorus, but possesses a somewhat metallic character. In its free state it is similar to metals. Arsenic is found free in nature, although it occurs more fre¬ quently in combination with sulphur (realgar, orpiment), with oxygen (arsenolite, AsjOs), and with metals (mispickel, FeSAs, cobaltite, CoAsS). To prepare it, heat mispickel with some iron, 144 INORGANIC CHEMISTRY. and free arsenic will sublime. Or, in the customary way of isolat¬ ing metals from their oxides, heat the trioxide (arsenolite) with charcoal :— AsjOj -f 3C=2 As -|- 3CO. Arsenic appears in two modifications. Crystallized arsenic is obtained by the sublimation of ordinary arsenic. It forms a gray- white, more or less metallic, crystalline mass, but may be changed into acute rhombic octahedra. Its specific gravity equals 5.7. It is brittle, and may be pulverized without difficulty. The amorphous variety is formed when arsenic is sublimed in a glass tube in a current of hydrogen, and also upon heating arsine. It is black, with little lustre, and possesses the sp. gr. 4.71. When heated to 360° it sets heat free and reverts to the crystalline variety. Away from air contact and at the ordinary pressure, arsenic vaporizes at a dark-red heat (about 450°) without previously fusing; it will, however, fuse if heated under great pressure in a sealed tube. Its vapor possesses a lemon-yellow color. The vapor density isi5o(H = i), the molecular weight, therefore, 300. As its atomic weight equals 75, it follows that the molecule in the state of gas, like that of phosphorus, consists of four atoms (AS4 = 300 = 4 X 75)- It is only at a yellow heat that the vapor density alters any. Arsenic does not change in dry air. When heated in the air it infiames at 180° and burns with a blue-colored flame, disseminating the garlic-like odor of arsenic tri-oxide (AsjOs). It combines directly with most elements. Powdered arsenic will inflame when projected into chlorine gas. It yields arsenides with the metals. It is remarkable that arsenic, belonging to the nitrogen group and generally forming compounds which in constitution are quite different from those of sulphur, should be analogous to the latter in its metallic combinations. Thus the sulphides and arsenides have similar formulas, are isomorphous, and in them sulphur and arsenic can mutually replace each other in atomic ratios, e.g. :— FeSj, FeAsg and Fe(SAs). COMPOUNDS OF ARSENIC WITH HYDROGEN. Arsine, AsH3= 78. Like nitrogen and phosphorus, arsenic furnishes a gaseous compound containing 3 atoms of hydrogen. It is obtained pure by the action of dilute sulphuric acid or hydro¬ chloric acid upon an alloy of zinc and arsenic:— ASgZng -f- 6HC1 = 3ZnClj -p zAsH,. ARSENIC WITH HYDROGEN. 145 It also results in the action of nascent H (zinc and sulphuric acid), upon many arsenic compounds, as, e. g., the tri-oxide:— ASjOj -j- 6Hj == zAsHg sHgO. Arsine is a colorless gas, of strong, garlicky odor, and extremely poisonous action ; it may be condensed to a liquid at —40°. Its density equals 38.9 (H=i), or 2.69 (air=:i). It burns with a bluish-white flame when ignited, and evolves white fumes of arsenic tri-oxide : — 2ASH3 4- 30j = AsgOg + 3H2O. It is decomposed at a dull red heat or by the electric spark into arsenic and hydrogen. On conducting the gas through a heated tube the arsenic deposits itself behind the heated part as a metallic coating {arsenic mirror^. On holding a cold object, e. g.j a piece of porcelain in the flame of the gas, the arsenic forms a black deposit (arsenic spots). In its chemical behavior arsine is very similar to PH3 ; its basic properties are very slight, and it does not furnish any derivatives with the halogens. According to analysis, arsine consists of i part by weight of hydrogen and 25 parts arsenic. If, because of its analogy to PHj, we ascribe to it the formula AsH„ then the atomic weight of arsenic would be 75 (3 X 25) and the molecular weight of AsHj would = 78. Hence the density must be 'j® = 39, which is confirmed by experiment. The formula, too, shows that 3 volumes of hydrogen are present in 2 volumes of AsHj ;— zAsHj contain 3H2. 2 vols. 3 vols. We can satisfy ourselves of this by decomposing the gas by electricity in a eudiometer (see p. 140). Marsh's Methodfor the Detection of Arsenic.—The detection of arsenic is very important, because of the poisonous nature of the element. The method of Marsh is based upon the formation and the characteristic properties of arsine. It is as follows : Hydrogen is generated in a flask « (Fig. 75 j, by the action of dilute sulphuric acid upon zinc, and a portion of the solution to be tested for arsenic is introduced through the funnel-tube. The liberated gas, a mixture of hydrogen and arsine, is dried in the calcium chloride tube c and escapes through the difficultly fusible glass tube d, which is contracted at several points. Upon igniting the escaping hydro¬ gen (after all the air has been previously expelled from the vessel, as otherwise oxy-hydrogen gas will be present) it will burn with a bluish-white flame, if arsenic is present, and at the same time dis¬ seminate a white vapor. The dark arsenic spots are obtained by holding a cold porcelain dish in the flame. If the tube d be heated *3 146 inorganic chemistry. (as shown in Fig. 75), an arsenic mirror will be formed upon the adjacent contraction. The slightest traces of arsenic may be detected by this method. Besides the ordinary arsine, AsHg, we might expect the existence of AsjH^ and As^Hj, corresponding to the liquid and solid phosphines (P2H4 íuid P4H2). The first is not known; its derivatives exist, and contain hydrocarbon groups instead of hydrogen. An example of this class is cacodyl, AsjiCHj)^ = (CHj),- fig. 75. as-as(ch3)2. Nitrogen affords similar compounds—(CHjljN-NH, and (CH3)NH-NH2, derived from diamide or hydrazine (N^H^ = HjN-NH,). The solid arsine, As^H2, is obtained by the action of nascent hydrogen up>on arsenic compounds in the presence of nitric acid. It forms a reddish-biown pow¬ der, which decomposes when heated. COMPOUNDS OF ARSENIC WITH THE HALOGENS. These are perfectly analogous to the corresponding phosphorus compounds, and are the result of the direct union of their constitu¬ ents, The iodide is the only known representative of the com¬ pounds with the formula AsX5(see p. 140). The metallic character of arsenic is shown by the fact that arsenic chloride, like other ANTIMONY. 147 metallic chlorides, may be obtained by the action of hydrochloric acid upon the oxide :— AS2O3 -j- 6HC1 = aAsClg -j- sHjO. Arsenic chloride is evolved when a solution of AsjOj is boiled with concentrated hydrochloric acid. Arsenic Trichloride—^AsClg. A colorless, oily liquid, fum¬ ing in the air, and having a specific gravity of 2.2. It solidifies at —30° and boils at 134°. Its vapor density equals 90.5 (H = i), corresponding to the molecular formula AsCls = 181.0. The chloride dissolves in a small quantity of water without change, while much water converts it into the oxide and hydrochloric acid :— 2ASCI3 -j- 3H2O = AsgOj -j- 6HC1. Arsenic Tribromide, AsBrg, is a white crystalline mass, fusing at 2Q°, and boiling at 220° C. The Tri-iodide, Aslg, forms red crystals; the Trifluoride, AsFlg, is a liquid, fuming strongly in the air. It results in the distillation of AsCb or As^Os with calcium fluoride and sulphuric acid. Arsenic pentaiodide, Aslj, melts at 70°, and is very soluble in water and alcohol. TTiertno-chemical Deportment.—The arsenic halogen derivatives are exothermic (see p. X43). Their heats of formation correspond to the symbols :— (As,Cl3) = 71.4; (ASjBrj) = 47.1 ; (As,^ = 12.6. This readily explains their chemical behavior. The following thermo-chemi- cal equation, based on the principle of the greatest heat disengagement (p. 67), indicates that arsenic trioxide is converted by gaseous or concentrated hydro¬ chloric acid (see above) into arsenic trichloride :— AsgOg-solid 6HQ-gas = 2ASCI3 -f- 3H2O. . . . (-j- 61.1 Cal.) 154.6 6X22.0 2X71-4 3X68.3 If, however, dilute hydrochloric acid be used the thermal value on the right side is negative :— AsjOj dissolved 6HC1 in solution = 2ASCI3 3H2O. . • • ( — 35* i Cal.) 147 6 X 39-3 2 X 71-4 3 X 68.3 The reaction consequently pursues an inverse direction. Dilute hydrochloric acid cannot convert the trioxide into the trichloride, but the latter, on the other hand, is fully changed into the former and hydrochloric acid by much water. The course of the reaction manifestly is dependent upon the concentration of the solu» tion and is influenced by the heat of solution (pp. 68 and 93). 4. ANTIMONY. Sb = 120.3. The metallic character exhibited by arsenic becomes more dis¬ tinct with antimony, which at the same time retains its complete 148 INORGANIC CHEMISTRY. analogy to the metalloidal elements, arsenic and phosphorus. Anti¬ mony is a perfect metal so far as its physical properties are con¬ cerned. It (Stibium) occurs in nature chiefly in union with sulphur, as stibnite, SbgSg, and with sulphur and metals in many ores. It is almost always accompanied by arsenic. To prepare antimony, stibnite is roasted in a furnace, i. e., heated with air access, whereby the sulphur burns, and antimony trioxide remains :— SbjSg -f" 9^ = Sb203 "1" The residual oxide is ignited with carbon, which reduces it to metal (general procedure for the separation of metals). Antimony may also be obtained by heating its sulphide with iron, which com¬ bines with the sulphur :— SbjSg + 3Fe = 2Sb + 3FeS. The resulting commercial crude antimony is further purified in the laboratory by fusing it with nitre, whereby the admixed arsenic, sulphur and lead are removed. Chemically pure antimony is obtained by reducing the pure oxide. It is a silver-white, and very brilliant metal, of leafy crystalline structure; specific gravity 6.715. Like arsenic it crystallizes in rhombohedra, is very brittle, and may be easily broken. It fuses at 430°, and distils at a white heat. It is not altered in the air at ordinary temperatures; but when heated it burns with a blue flame, yielding white vapors of antimonio oxide, Sb^Os. Like phosphorus and arsenic it combines directly with the halogens ; powdered anti¬ mony inflames in chlorine gas. It is insoluble in hydrochloric acid ; nitric acid oxidizes it to antimonio oxide. Hydrogen Antimonide—Stibine—(SbHg), is produced like arsine, and is very similar to the latter. It is always obtained mixed with hydrogen. It is a colorless gas of peculiar odor, and when ignited, burns with a greenish-white flame, disseminating white vapors of antimonio oxide. A red heat decomposes it into anti¬ mony and hydrogen. In Marsh's apparatus (Fig. 75, p. 146) it affords an antimony mirror and spots. The mirror is distinguished from that of arsenic by its black color, lack of lustre, its insolubility in a solution of sodium hypochlorite (NaClO), and by its slight volatility in a current of hydrogen. When a solution of antimony trichloride is decomposed by the electric current there is deposited on the cathode a metallic compound, which explodes when crushed, scratched with a knife, heated or exposed to the electric spark. White clouds of antimony chloride escape and pure metallic antimony remains. This explosive body was supposed to be an alloy of antimony with hydrogen. It, however, consists of metallic antimony, 4.8 to 8 per cent, of antimonious chloride, and some hydrogen chloride mechanically or chemically combined. The cause of the explosibility is unknown. ANTIMONY WITH THE HALOGENS. 149 COMPOUNDS OF ANTIMONY WITH THE HALOGENS. Antimonous Chloride—Trichloride—SbClg, results from the action of chlorine upon the metal or its sulphide ; better by the solution of the oxide or sulphide in strong hydrochloric acid :— Sb^Sg + 6HCI = aSbClg -f- sH^S. This solution is evaporated to dryness and the residue distilled. It is a colorless, crystalline, soft mass {^ButyrumAntimonii^^ fusing at 73° and boiling at 223°. Its vapor density equals 112.8 (H = i), corresponding to the molecular formula, SbCl3 = 225.7. In the air it attracts water and deliquesces. It dissolves unchanged in water acidified with hydrochloric acid. Much water decomposes it ; the solution becomes turbid and a white powder—powder of algaroth—^separates :— SbClg + HgO = SbOCl + 2HCI. The composition of this powder varies with the conditions under which it is formed, but generally corresponds to the formula aCSbOCO.SbjOa. Pure Antimony Oxychloride, SbOCl, ob¬ tained by heating SbClg with alcohol, occurs in colorless crystals and is further decomposed by water. While the metallic chlorides are not decomposed by water at ordinary temperatures, the ready decomposition of the halogen derivatives of antimony, shows that this element yet possesses a partial metalloidal character. Antimonio Chloride—Pentachloride—SbClj, results from the action of an excess of chlorine upon antimony or the trichloride. It is a yellowish liquid which fumes in the air, becomes crystalline when cold and fuses at —6°. Heat partially decomposes it, like PCI5, into SbCls and Clj :— SbClg = SbClg -f Clj. I vol. I vol. I vol. Water converts it into pyroantimonic acid (HiSbaOï), and hydro¬ chloric acid. Antimony Tribromide—SbBrg—is a white, crystalline sub¬ stance, fusing at 94° and distilling at 270°. The Tri-iodide, Sbis, is a red compound, crystallizing in three distinct forms. The Pentaiodide, Sbig, is dark brown in color and fuses about 78°. The heat of formation of antimony chloride equals :— (SbjClg) = We possess no knowledge concerning the other halogen derivatives in this direc¬ tion. The decomposition of SbClg by water is analogous to that of arsenic tri¬ chloride (p. 147). 150 INORGANIC CHEMISTRY. We must also include Bismuth—Bi = 208—in the group of nitrogen, phosphorus, arsenic, and antimony; it forms similarly constituted compounds, e. g.y BÍCI3, Bilg, BiOCl. Its metallic character, however, considerably exceeds its metalloidal. Thus, it does not unite with hydrogen, and the oxide (BijOa), similar in constitution to the acid-forming AS2O3, possesses only basic charac¬ ters. We will, therefore, consider bismuth and its derivatives with the metals. TABULATION OF THE ELEMENTS OF THE NITROGEN GROUP. The elements belonging here—nitrogen, phosphorus, arsenic, antimony, and bismuth—present similar graded differences in their physical and chemical properties, just like the elements of the chlorine and oxygen group, and this gradation is intimately con¬ nected with the atomic weights. As the latter increase the substance condenses, the fusibility and volatility decrease, and the metallic character becomes more prominent :— N. P. As. Sb. Atomic weight 14.041 31-03 75 120.3 Specific gravity • • • 1.8-2.1 4-7-5-7 67 Fusion point • • • 44° red-white heat • • • Vapor density 0.972 4.32 10.3 • • • Excepting bismuth, which is perfectly metallic in its nature, the elements of this group form gaseous compounds with three atoms of hydrogen. Ammonia (NH3) possesses strongly basic properties, and combines with all acids to yield ammonium salts ; phosphine (PH3) combines with HBr and HI to form salt-like compounds. ASH3 and SbHj no longer show basic properties. Arsenic and antimony, as well as the two preceding elements, combine with the hydrocarbons (jc-g-, CH3 and C2H5) and form compounds that are analogous in consti¬ tution and similar in character to the hydrides. These compounds [As (CH3)3 and Sb (CHs)3] will be described in Organic Chemistry ; they possess basic properties and yield salts corresponding to the ammonium salts. The oxygen derivatives of these elements exhibit a similar gradation. With increase of atomic weight, corresponding to the addition of metallic character, the oxides that form strong acids in the lower series acquire a more basic nature. CARBON. 151 The gradation is more manifest in the thermal relations of the group. How- ev», of the hydrides of the elements of this group only those of nitrogen and phosphorus have been investigated in this direction (pp. 129 and 137) :— (N,H,) = ii.9; (P,H3) = II.6. It is very probable that in the case of the higher members ( just as in the halo¬ gen and oxygen groups) the heat of formation will diminish successively in accordance with their increasing decomposability. Yet, the heat of formation of the halogen derivatives increases successively with rise of atomic weight and metallic character :— (N.Cy =- 38.1 ; (P.Clj) = 75.3; (As, CI3) = 71.4; (Sb,Cl,) = 91.4. We here observe plainly the transition from metalloidal to metallic character. CARBON GROUP. The two non-metals, carbon and silicon, and the metals, tin and germanium, belong to this group. These unite with four atoms of hydrogen or four of the halogens. I. CARBON. C = 12. Carbon occurs free in nature as the diamond and graphite. It constitutes the most important ingredient of all the so-called organic substances originating from the animal and vegetable kingdoms, and is especially contained in the fossilized products arising from the slow decomposition of vegetable matter—in turf, in brown coal, bituminous coal, and in anthracite. In combination with hydrogen it forms the so-called mineral oils—petroleum and asphaltum. It occurs, further, as carbon dioxide (CO2) in the air; and in the form of carbonates (marble, calcite, dolomite) comprises many minerals and entire rock formations. It is found in different allotropie modifications when free ; these may be referred to the three principal varieties—diamond, graphite and amorphous carbon. In all these forms it is a solid, even at the highest temperatures ; non-fusible and non-volatile. This deport¬ ment can only be explained by the supposition that its free molecules are composed of a large number of carbon atoms combined with each other. (See p. 107.) All the modifications of carbon are quite stable, but not very reactive. When burned all yield carbon dioxide. I. The diamond occurs in alluvial soils in certain districts (in India, Brazil, and South Africa) ; less frequently in micaceous schist. It has great lustre, strong 152 INORGANIC CHEMISTRY. power of refraction, and the greatest hardness of all substances. It crystallizes in forms of the regular system, that are mostly rhombic dodecahedra, rarely octa- hedra. Ordinarily, it is perfectly colorless and transparent; sometimes, however, it is colored by impurities. Its specific gravity equals 3.5. It does not soften any unless exposed to the most intense heat—between the poles of a powerful galvanic battery. It is then converted into a graphitic mass. When heated in oxygen gas it burns to carbon dioxide. It is scarcely attacked at all when acted upon by a mixture of nitric acid and potassium chlorate. 2. Graphite is characterized by its oxidation to graphitic acid when it is heated with a mixture of potassium chlorate and nitric acid. Like amorphous carbon, it is oxidized to mellitic acid by an alkaline solution of potassium permanganate, or when it is made the p)ositive electrode in the elecholysis of alkaline solutions. Native graphite is found in the oldest rock formations, and of especially good quality at Altai, in Siberia. It occurs, too, in considerable quantities at many places in the United States. It is occasionally found crystallized in six-sided forms, but usually as an amorphous, grayish-black, glistening, soft mass, used in the manufacture of lead pencils. The specific gravity is 2.25. It conducts heat and electricity well. When away from air-contact it is not altered even at the highest temperatures. It usually burns when heated in an atmosphere of oxygen, but with more difficulty than the diamond, forming carbon dioxide, and leaving about 2—5 p)er cent, of ash. To purify the poorer and more impure kinds of graphite, the latter is pulverized and heated with a mixture of KCIO3 and HjSO^ ; the product is washed with water, and the residue ignited (Brody's Graphite). Graphite may be obtained artificially by fusing amorphous carbon with iron ; when the latter cools, a p)oition of the ffissolved carbon separates in hexagonal shining leaflets. 3. Amorphous Carbon is produced by the carbonization of organic (containing carbon) substances, and is found in a fossilized state. Nitric acid and potassium chlorate convert it in the cold into brown substances soluble in water. The purest amorphous carbon is soot which is obtained by the imperfect combustion of resins and oils (like turp>entine) rich in carbon. Gas Carbon^ called metallic carbon, deptosits in the manufacture of gas in the retorts, and is very hard, p)ossessing metallic lustre, and conducting electricity well; hence its use in galvanic batteries. Cokfy resulting from the ignition of bituminous coal, forms a sintered mass, con¬ ducting heat and electricity well. Charcoal is very pxjrous, and can absorb many gases and vapwrs ; i volume of it condenses 90 volumes NH3, 55 volumes HjS, and 9 volumes Oj. At 100®, and under the air pump, the absorbed gases are again liberated. Charcoal will also take up many odorous substances and decaying matter ; hence is employed as a disinfectant. Animal Charcoal is obtained by the carbon¬ ization of animal matter (bones, blood, etc.), and possesses the power of removing many coloring substances from their solutions ; hence it serves in the laboratory and in commerce for the decolorization of dark solutions. All these varieties of carbon contain smaller or larger quantities of nitrogen, hydrogen, and mineral substances, which remain as ash after combustion. Hydro¬ chloric acid will withdraw almost all the mineral constituents. The fossil coal varieties, bituminous coal, lignite and turf, are the products of a p>eculiar, slow decay of wood fibre, which gradually separates oxygen and hydro¬ gen, and enriches itself in carbon. Fossil coal contains 90 per cent, and brown coal 70 per cent, of carbon. The fossil coal richest in carbon, the last product of the alteration, is anthracite. This has lost all its organic structure, and contains 96-98 per cent, of carbon. CARBON WITH HYDROGEN. 153 COMPOUNDS OF CARBON WITH HYDROGEN. With hydrogen, carbon forms an unlimited number of compounds, into which all other elements, especially oxygen and nitrogen, can enter. The derivatives of carbon have been termed organic com¬ pounds, because they were formerly obtained exclusively from vege¬ table and animal organisms, and the idea was entertained that they were produced by the influence of forces other than those forming the mineral substances. At present, most carbon derivatives are pre¬ pared artificially from the elements by simple synthetic methods ; we are aware that they do not differ essentially from mineral sub¬ stances. Hence the description of the carbon compounds must be arranged in the general system of chemical bodies. This, however, is not readily executed without sacrificing the review of a defunct system. The derivatives of carbon are so numerous, and possess so many peculiarities, that it appears necessary, from a practical stand¬ point, to treat them apart from the other compounds, in a separate portion of chemistry, which we, pursuing the old custom, term organic chemistry. We then designate the chemistry of all other bodies as Inorganic Chemistry. Only the simplest carbon compounds will be considered here. It is only under the influence of the electric arc that the direct union of carbon and hydrogen may be effected; the product is acetylene (CgHj). All other hydrocarbons are obtained indirectly in various ways. Methane—Marsh Gas—CH4.—This simplest hydrocarbon, con¬ taining but one atom of carbon, is formed in the decay of organic matter under water (in swamps and coal mines), and escapes in large quantities in many regions of the earth (thus at Baku, on the Caspian Sea). It may be obtained synthetically by conducting vapors of carbon disulphide and hydrogen sulphide over ignited copper filings :— CS2 -f 2H2S + 8Cu = 4CU2S -f CH^. For its preparation, heat a mixture of sodium acetate with sodium hydroxide :— C^HaNaOa + NaOH = CH^ + Na^COa. Methane is a colorless, odorless gas, insoluble in water ; it can be condensed by pressure and cold. Its density equals 8 (H = i) or 0.552 (air = i), corresponding to the molecular formula CH4=: 16. Methane, under great pressure and at low temperatures (below — 82°), is condensed to a colorless liquid, which boils at—164° under the ordinary pressure. Its specific gravity is 0.415 at the boiling point. When ignited it burns with a faintly luminous flame. It affords a violently explosive mixture (fire-damp 154 INORGANIC CHEMISTRY. of the miners) with two volumes of oxygen (or ten volumes of air) :— CH^ + 20j = COj + aHjO. I vol. 2 vols. I vol. 2 vols. MOLECULAR FORMULA OF METHANE. ATOMIC WEIGHT OF CARBON. The quantitative analysis of methane shows that for every i part of hydrogen in it there are 3 parts carbon. Were the formula CH (analogous to hydrochloric acid) then the atomic weight of carbon would be 3. If it corresponded to the formula of water (H^O) then carbon would equal 6, etc. (see p. 98) :— H = i 2H==2 3H = 3 4H = 4 C = 3 C = 6 C — g C = 12 CH = 4 CH, = 8 CHj = 12 CH^ = 16 In this case the analysis yields (as in former instances) no con¬ clusive answer. We derive the molecular weight of methane, according to Avogadro's law, from its density. The latter equals 8 (H = i), hence the molecular weight is 16. In 16 parts by weight of methane there are 4 parts by weight, hence 4 atoms, of hydrogen, and 12 parts carbon. The atomic weight of C is, then, presuming that only i atom of it is present in methane, 12;— 4 atoms hydrogen H^ = 4 I atom carbon C =12 Methane molecule CH^ = 16 That the atomic weight of carbon is really 12, is proven by the fact that of all its innumerable derivatives, not one contains less than 12 parts, by weight, of this element. It follows, with certainty, from the periodic system of elements (p. 80). From the formula CH4 it follows that in i volume of methane there are 2 volumes of hydrogen (CH^ contains 2H2). This is I vol. 2 vols. proved indirectly by the combustion of methane with oxygen in a eudiometer (see p. 121). Four atoms of hydrogen yield two mole¬ cules of H2O ; I atom of C yields i molecule of CO2. Hence the volume relation in the combustion of CH^ in oxygen is expressed by the equation :— CH^ + 2O2 = CO2 -f- 2H2O. I vol. 2 vols. I vol. 2 vols. In two volumes of aqueous vapor there are 2 volumes of hydro¬ gen ; hence in one volume of CH^, there are 2 volumes of Hj. The result of the eudiometric analysis confirms these conclusions. Ethane—CaHg—is formed when hydrogen m statu nascendi acts upon ethyl chloride :— ETHYLENE. 155 C,HgCl + 2H = C,H, + Ha. Or by the action of potassium or sodium upon methyl iodide :— 2CH3I + Na, = + 2NaI. This is a colorless gas, insoluble in water, and when ignited it burns with a feebly luminous flame. Its density equals 15 (H = i ) or 1.037 (air = i) corresponding to the molecular formula CjHe = 30. Besides methane (CH4) and ethane (CjHg) there exists a long series of hydrocarbons of the general formula CnHjn+j {e. g., CaHg, QHjo, C5H12, etc.), in which each member differs from the preceding and next following by i C and 2 H (CHj). Bodies belonging to such a series, greatly alike in their chemical behavior, are termed homologues. In addition to this series of saturated hydrocarbons others exist, with less hydrogen, and by the addition of the latter, pass into the saturated^ and may, therefore, be termed unsaturated. The first unsaturated series is composed according to the formula CnHjn, the second according to CnH2n_2> etc. The lowest member of the series CnHjn is ethylene (see Chemical Structure, p. 170). Ethylene—C2H4—is formed in the destructive distillation of wood, bituminous coal, and many carbon compounds, hence is con¬ tained in illuminating gas. It is most easily obtained by the action of sulphuric acid upon alcohol, whereby the acid withdraws HjO from the latter ;— qHgO — HjO = C2H4. Alcohol. Ethylene. It is a colorless gas, of weak, ethereal odor, and condenses at —110° to a liquid. Its density equals 14 (H= i) or 0.969 (air = i), corresponding to the molecular formula Its critical temperature lies at-f-9.2°, and its critical pressure is 58 atmospheres. Under a pressure of i atmosphere, liquid ethylene boils at —102°, and when in a vacuum at —150°. Because it does not solidify at this point, it is well adapted for the liquefaction of other gases (p. 48). It burns with a bright, luminous flame, decomposing first into marsh gas and free carbon : — C2H4=CH4 + C. The CH4 then burns and heats the particles of carbon in the flame to incandescence; these are then consumed to carbon di¬ oxide (CO2). The unsaturated compound, ethylene, unites directly with two atoms of chlorine and bromine :— C2H4-fa2 = C2H4a2. The resulting compounds, C2H4CI2 and C2H4Bra, are oily liquids; hence the name defiant gas, for ethylene. The lowest member of the second unsaturated series is C,H,. 156 INORGANIC CHEMISTRY. Acetylene—CaHj—is produced in the dry distillation of many carbon compounds, and is present in coal gas, to which it imparts a peculiar penetrating odor. Its density = i3(H=i) correspond¬ ing to the formula CaHj = 26. It combines directly with 2 and 4 atoms of chlorine and bromine. The three hydrocarbons considered above, methane (CH4), ethylene (C2H4), and in slight amount acetylene (C2H2), constitute, together with H and carbonous oxide (CO), ordinary illuminating gas, which is produced in the dry distillation of bituminous coal, lignite, or wood. The illuminating power is influenced by its quantity of ethylene and acetylene (and their homologues.) Fig. 76. THE NATURE OF FLAME. We are aware that every chemical union which occurs in a gaseous medium, and is accompanied by the evolution of light is designated combustion. We observe herein, that some bodies, like sulphur and phosphorus, yield a flame when burned in the air or in other gas ] such substances are converted into gases or vapors at the temperature of combustion. Carbon burns without a flame, be¬ comes incandescent, because it is non-volatile. The carbon com¬ pounds, wood, bituminous coal, and tallow, are, indeed, not volatile, but burn with a flame because under the influence of heat they develop combustible gases. Flame is, therefore, nothing more than a com¬ bustible gas heated to incandescence. We have observed, too, that the com¬ bustibility is only a relative phenomenon ; if hydrogen burns in oxygen and chlo¬ rine, oxygen and chlorine, conversely, will burn in hydrogen (p. 57). Illumi¬ nating gas burns in the air, therefore air (its oxygen) bums in the former. This may be demonstrated in the same manner as in the case of chlorine and hydrogen. The relative combustibility and the so-called return of the flame may be very plainly illustrated by means of the following contrivance : An ordi¬ nary lamp chimney (Fig. 76) is closed at its lower end with a cork, through which two tubes enter ; the narrow tube, somewhat contracted at its end,is connected with a gas stop-cock; the other tube, b (best a cork borer), is about 5 mm. wide, and communicates with the air. The gas issuing from the tube a is ignited, and the chimney is then dropped over the not too large flame; it continues to burn along quietly, as sufficient air enters throt^h THE NATURE OF FLAME. 157 the wide tube b. Upon increasing the supply of gas, the flame becomes larger, the globe fills with illuminating gas, while the air is crowded out. The gas flame is extinguished, and an air-flame appears upon the wider tube, as the entering air continues to burn, in the atmosphere of illuminating gas. The excess of the latter escaping from the upper portion of the globe may be ignited, and we then have a gas-flame above, while within the globe we have an air-flame. On again lessening the gas-flow the air-flame will distribute itself, extend to the exit of the tube a, and then the gas-flame will appear upon the latter, while the flame above the globe is extinguished. In this manner, we may repeat the return process of flames at will. That the air actually burns in the air-flame may be plainly proved if we introduce a small gas-flame from i, through the wide metallic tube b ; the little flame will continue to burn in the air-flame, but will be ex¬ tinguished if it be introduced higher up into the atmosphere of illuminating gas. Fig. 77. We say ordinarily that only those bodies are combustible which, because of their power to unite with oxygen, burn in an atmos¬ phere of this gas or in air. If we imagine, however, an atmos¬ phere of hydrogen, or illuminating gas, then bodies rich in oxygen must be combustible in these. In fact, nitrates, chlorates, etc., burn in an atmosphere of illuminating gas with the production of an oxygen flame. This may be demonstrated as follows: An Argand-lamp chimney (Fig. 77) is closed at its lower end by a cork, bearing a gas-conducting tube. The gas which escapes through the 158 INORGANIC CHEMISTRY. opening of the sheet covering, is ignited. Then the substance (potassium or barium chlorate, etc.) is introduced into the flame on an iron spoon provided with a long handle, heated to the tempera¬ ture of decomposition (disengagement of oxygen), and the spoon then plunged through the opening into the gas atmosphere. The sub¬ stance burns with a brilliant light, as the resulting oxygen flame is brightly colored by the vaporizing and reduced metallic salts. The brilliancy or luminosity of a flame is influenced by the nature of the substances contained in it, also by its temperature and den¬ sity. Incandescent gases shine very faintly per se ; this is esj^cúdly true when they are diluted. Thus hydrogen, ammonia and methane burn with a pale flame. Even sulphur bums in the air with a slightly luminous flame. If, on the contrary, sulphur or phosphine be permitted to burn in oxygen, or arsenic and antimony in chlorine gas, an intense display of light follows. This depend on the fact that the flame is not diluted by the nitrogen of the air, is therefore more condensed, develops a higher temperature, and the combustion products (SOj, P2O5, PCI3) or the evaporating sub¬ stances are not immediately gasiñed. That the density of the flame of gases exercises a great influence upon the luminosity is proved by the fact that hydrogen, compressed into a smaller space with oxygen, burns with intense light display. A slightly luminous flame may be rendered intense by introducing solid particles into it. For example, if hydrogen be passed through liquid chromium oxychloride (CrOaCla) it burns with a bright, luminous flame, because the volatile CrOjClj in 7®- it is changed by the oxygen of the flame into solid, ® non-volatile chromium oxide, CrjOg, whose particles I are heated to incandescence by the hydrogen flame. A The illuminating power of the various hydrocarbons and nX carbon compounds is similarly explained. Marsh gas, CH4, and ethane, CaHg, afford a pale flame, because ^1% they burn directly to aqueous vapor and carbon dioxide. Ethylene, on the contrary, burns with a bright, lumi- mBUÈ nous flame, because, by the temperature of combustion, it decomposes flrst into CH4, and carbon, whose par- tides glow in the flame. (See p. 155.) Let us consider the flame of an ordinary stearin candle : On approaching the wick with a flame the ■ stearin melts, is drawn up by the fibres and converted ■ 9H into gaseous hydrocarbons, which ignite, and by their chemical union with the oxygen of the air, produce the flame. The unaltered gases exist in the inner non-volatile zone a (Fig. 78) ; they cannot burn because of lack of air access. If the THE NATURE OF FLAME. 159 lower end of a thin glass tube be inserted here the gases will rise in it, and may be ignited at the upper end. There is a partial com¬ bustion of the gases in the middle, luminous part, /, g ; ethylene, CjH*, breaks up into CH^ and C ; the first burns completely, while the C is heated to a white heat, because there is not sufficient oxygen present for its combustion. The presence of carbon particles in the luminous part may be easily proved by placing a glass rod or a wire net in it ; it will at once be coated with soot. In the outer, very feebly luminous and almost invisible mantle, by Cy dy of the flame, which is completely surrounded by air, occurs the perfect combustion of all the carbon to carbon dioxide. A perfectly identical structure is possessed by the ordinary illumi¬ nating gas flame. By bringing as much air or oxygen into it as is necessary for the perfect combustion of all the carbon, none of the latter separates (see below), and there is produced a faintly lumi¬ nous but very hot flame. Upon this principle is based the con¬ struction of the Bunsen burner, the flame of which is employed in laboratories for heating and ignition. Fig. 79 represents a form of the same. The upper end, c, is screwed into the lower portion, and in the figure is only separated for the sake of explanation. The gas enters through the narrow opening, 0, from the side gas tube, and mingles with air in the tube Cy which enters through the openings of the ring, b. In this way we obtain a flame that is but faintly luminous, although affording an intense heat. On closing the openings in b the air is cut off, and the gas burns at the upper end of the tube /?, with a bright, strongly smoking flame. The non- luminous flame contains an excess of oxygen, and hence oxidizes— oxidizingßame. It is employed to effect oxidation reactions. The luminous flame, on the other hand, is reducing in its action, and is designated the reductíon flame, because the glowing carbon in it abstracts oxygen from many substances. The construction and application of the ordinary and the gas blowpipes depend upon the same occurrences ; they are, however, replaced at present by gas lamps. The non-luminosity of the Bunsen burner flame, due to addition of air, depends Dn a more complete combustion of the separated carbon or of the yet undecom- posed hydrocarbons. The flame, in consequence, is smaller, more intense, and the combustion extends itself even to the inner cone of the flame. It is more difficult to render the flame non-luminous by pure oxygen, because it is then not diluted by nitrogen, and is, therefore, much smaller, the temperature much higher, and the flame gases are more condensed. Another variety of non-luminosity of hydrocarbon flames is induced by the admixture of inactive gases, like nitrogen and carbon dioxide. By this means the flame is enlarged and the combustion, as in the luminous flame, takes place only in the outer cone. In consequence of the dilution there are present fewer com- 160 INORGANIC CHEMISTRY. bustible particles in an equal space, and these can be more completely consumed by the oxygen of the air, which enters more readily ; further, the temperature is lowered, and probably does not acquire the decomposition temperature of ethylene (CjH^) in the adjoining cone, which is being continually renewed. The simple extension of an illuminating flame upon a plate, will render it non-luminous, because then the air comes in contact with a larger flame surface. On heatii^ a gas made non-luminous by the admixture of nitrogen, and then letting it bum, its flame becomes luminous because the increased temperature can induce the decom¬ position of ethylene. In rendering flame non-luminous by carbon dioxide, we must also consider that the same is converted, by the particles of carbon, into carbon monoxide :— CO2 4- C = 2CO. Indeed, but a few per cent, of COj in a gas flame suffices to considerably diminish its luminosity, C2H4 4- CO2 = CH^ 4- 2CO, I vol. I voL z vol. 2 vols. while the presence of nitrogen is far less detrimental. Every substance requires a definite temperature for its ignition— temperature of ignition. When a substance is once ignited it gener¬ ally burns further, because additional particles are raised to the temperature of ignition by the heat of combustion. By rapid cooling {e. g., by the introduction of a piece of metal into a small flame) every flame may be extinguished. By holding a metallic net over the opening of a gas lamp, from which gas issues, and igniting the same above the wire (Fig. 80), the latter, being a good con- CARBON WITH THE HALOGENS. 161 ductor of heat, cools the flame so much that it is incapable of igniting the gas below the gauze. Upon this phenomenon depends the construction of Davy's safety lamp, which is used in coal mines, to avoid ignition of the fire-damp (Fig. 8i). It is an ordinary oil lamp surrounded and shut off from the air by a metallic wire gauze. On bringing a lighted lamp of this sort into an explosive mixture, or into a combustible gas (ounds to those of carbon, that its formula is SiCl^, the atomic weight of the silicon would be 28.4 :— Si = 28,4 Ch = 141.80 (4x35.45). SiCh — 170.2 This supposition is confirmed by the vapor density of the compound. This equals 84.7 (H = i), hence the molecular weight is 2 X 84.7 = 169.4. As the analysis shows that there are 141.8 parts chlorine in 170 2 parts of silicon chlo¬ ride, the atomic weight of the metal must be 28.4. Silicon Bromide—SiBr^—and Silicon Iodide—SÍI4—are formed in the same manner as the chloride. The first is a colorless liquid, of specific gravity 2.8, becoming solid at —12° and boiling at -|- 153°. The iodide forms colorless octahedra, fusing at 120°, and boiling at 290°. Like the chloride, both are decomposed by water. Besides these compounds, which may be viewed as hydrt^en silicide, in which all the hydrogen is replaced by halogens (see p. 161), others exist, in which o- The decomposition of its aqueous solution by alkalies is an argument in favor of this view, as is also its analogy with nitrogen dioxide, NO2, or nitrogen tetroxide, N2O4 (see this), the existence of both of which molecules has been proved. Chloric Acid—HCIO3, or CIO2.OH—is obtained by decompos¬ ing an aqueous solution of barium chlorate with sulphuric acid :— (aOs)2Bä 4- SO^H, = BaSO^ 4- 2HCIO,. Barium chlorate. Barium sulphate. 182 INORGANIC CHEMISTRY. The barium sulphate separates as a white, insoluble powder, and can then be filtered off from the aqueous solution of the acid. This is concentrated, under an air-pump, until the specific gravity becomes 1.28, and it then contains about 40 per cent, of chloric acid ; it is oily and, when heated to 40°, decomposes into chlorine, oxygen, and perchloric acid, HCIO4. The concentrated aqueous solution oxidizes strongly ; sulphur, phosphorus, alcohol, and paper, are inflamed by it. Hydrochloric acid eliminates chlorine from the acid and its salts :— HCIO, + sHCl = 3H2O The chlorates are produced, together with chlorides, by the action of chlorine, in the presence of heat, upon many bases (com¬ pare potassium chlorate) :— 6K0H + 3CI2 = 5KCI + KCIO, + 3H2O. Perchloric Acid—HCIO4, or CIO3.OH. This is the most stable of all the oxygen derivatives of chlorine. As previously stated, it is produced by the decomposition of chloric acid, but is more easily obtained from its salts. Upon heating potassium chlor¬ ate to fusion, oxygen escapes and potassium Perchlorate results :— 2KC10j = KCIO4 -f KCl + O2. Upon warming the Perchlorate with four parts sulphuric acid, per¬ chloric acid distils over :— 2a04K -H H2SO4 = K2SO4 + 2HCIO4. The pure acid is a mobile, colorless liquid, fuming strongly in the air ; its specific gravity is i. 78 at 15 °. It may be solidified by cold, meltsabout 15°, and boils at 110°. It cannot be preserved, since after a few days it decomposes with violent explosion. It also explodes in contact with phosphorus, paper, carbon, and other organic sub¬ stances. It produces painful wounds when brought in contact with the skin. It dissolves in water with hissing, and with one molecule of the solvent forms the crystalline hydrate HCIO4 -}- HjO, fusing at 50° ; the crystals fume in the air and gradually deliquesce. The second hydrate—HCIO4 2H2O—is a thick, oily liquid, resem¬ bling sulphuric acid, and boils unchanged at 208°. It may also be obtained by evaporating the aqueous solutions of perchloric and chloric acids. When the crystalline hydrate is distilled it breaks up into anhydrous perchloric acid and the second hydrate :— 2CIO4H.H2O = CIO4H + CIO4H.2HJO. HYPOBROMOUS ACID—IODIC ACID. 183 Bromine yields the following oxygen compounds :— HBrO Hypobromous acid. HBrOj Bromic acid. HBrO^ Perbromic acid. The corresponding anhydrides are not known. The acids are perfectly analogous to the corresponding chlorine compounds. Hypobromous Acid—HBrO—is formed when bromine water acts upon mercuric oxide ; the aqueous solution can be distilled in vacuoy and possesses all the properties of hypochlorous acid. Bromic Acid—BrOjH. Bromates are formed by the action of bromine, in the heat, upon the aqueous solution of the alkalies or of barium hydroxide; an aqueous solution of the acid can be obtained from the barium salt by decomposing the latter with sulphuric acid. A more practical method of getting the free acid is to let bromine act upon silver brómate or oxidize bromine with hypochlorous cicid * 5CI2O + Brj + HjO = zBrOaH + loCl. The aqueous solution may be concentrated in vacuo until its con¬ tent reaches 50.6 per cent. BrOsH, and then closely corresponds to the formula BrOgH + 7H2O. When heated it breaks up into bro¬ mine, oxygen and water. Perbromic Acid—BrO^H—is said to be formed in the action of bromine vapor upon perchloric acid :— ClO^H + Br = BrO^H + CI. and is perfectly similar to the latter. Iodine forms the following anhydrides and acids :— I2O5 HIO3 —Iodic Acid. ( 12O2) HIO^ — Periodic Acid. Iodic Acid—HIO3. Its salts (iodates) are formed in the same manner as those of chloric and bromic acids, by dissolving iodine in a hot solution of potassium or sodium hydroxide ;— 6K0H + 3I2 = sKI -f IKO, + 3H2O. The free acid can be obtained by the oxidation of iodine with strong nitric acid, or by means of chlorine ; further, by the action of iodine upon chloric or bromic acids, whereby the iodine directly eliminates the chlorine and bromine :— aHClOj + la = aHIOg + Clj. 184 INORGANIC CHEMISTRY. Upon evaporating the aqueous solution the free iodic acid crystal¬ lizes in colorless rhombic tablets of specific gravity 4.63. The solution possesses strong oxidizing properties. When iodic acid is heated to 170° it decomposes into water and iodic anhydride :— 2HI03 = l205 + H,0. It is decomposed, similarly to chloric acid, by hydrochloric acid :— 2IO3H 4- loHCl = I2 + SCI2 + 6H2O. Reagents, like HjS, SO2 and HI, reduce it to iodine. Periodic acid sustains similar decompositions. Iodic Anhydride—I2O5—is a white crystalline powder, which dissolves in water to form iodic acid. It decomposes at 300° into iodine and oxygen. Periodic Acid—HIO^—is produced by the action of iodine upon perchloric acid :— 2HCIO4 + I2 = 2HIO4 + Clj. Upon the evaporation of the aqueous solution, the acid crystallizes out with two molecules of water (HIO4, 2H2O—compare below). In the air, the crystals deliquesce, fuse at 130°, and at a higher temperature decompose into water and periodic anhydride, the latter at once breaking up into oxygen and iodic anhydride :— 2(HI0, + 2H3O) = I3O3 + O3 + 5H3O. The existence of the hydrates of periodic and perchloric acids, as well as of many others (see Sulphuric and Nitric acids), which we once regarded as molec¬ ular compounds (p. 175), is interpreted at present by the acceptance of hy- droxyl groups, directly combined with the element of higher valence :— VII ClO^H -}- HjO = C102(0H)3— trihydrate or tryhydric acid. CIO4H -f- aHjO — CIO (0H)5— pentahydrate or pentahydric acid. CIO4H 4- 3H2O = CI (OH),j— heptahydrate or heptahydric acid. The extreme hydrates, CI(OH)^ and IfOH)^, in which all seven affinities of the halogen atom are attached to hydroxyl groups, are not known, but probably exist in aqueous solution. As they give up water, and one atom of O becomes simultaneously united with two bonds to the halogen, they yield the lower hy¬ drates—even to the monohydrate CIO3OH. Perchloric acid continues mono¬ basic in the polyhydrates, since but one hydrogen atom is replaced by metals : ClOgHg 4- KOH = ClO.K 4- 3H2O. On the other hand, periodic acid (IO3OH) is not only monobasic, but as a penta¬ hydrate (I0(0H)5) can, like the polybasic acids, furnish also polymetallic salts as : OXYGEN COMPOUNDS OF THE HALOGENS. 185 VII f fOHI. VII f fOH), VII VII I0{ ¡0Na)3 I0{ ¡OA^ IO(O.Na)5 lO(OAg),. Salts also exist which are derived from coudensed polyiodic acids, as lOi /(OH), —Diperiodic acid, etc. IO( ^(OH), (Compare disulphuric, dichromic acid, etc.) The existence of such salts plainly indicates that the hydrates of acids must be looked upon as hydroxyl compounds, and that iodine and the halogens are, in fact, heptads in their highest combinations. The oxygen compounds of the halogens in some respects display a character exactly opposite to that of the hydrogen derivatives. While the affinity of the halogens for hydrogen diminishes with increasing atomic weight from Fl to I (see p. 66), the affinity for oxygen is the exact reverse. Fluorine is not capable of combining with oxygen ; the chlorine and bromine compounds are very unstable, and are generally not known in free condition ; the iodine derivatives, on the contrary, are the most stable. In accord with this is the fact that in the higher oxygen compounds chlorine and bromine are set free by iodine, while in the hydrogen and metallic compounds of the halogens the direct reverse is the case, viz., that iodine and bromine are replaced by chlorine. Further, the oxygen compounds exhibit the remarkable peculiarity that their stability increases with the addition of oxygen. The lowest acids, HCIO, HBrO, HCIO2, are very unstable, even in their salts ; they possess a very slight acid character, and are, too, separated from their salts by carbon dioxide. The most energetic and most stable are the highest acids, HCIO4, HBrOg, HIO3, in which the higher valence of the halogens appears. In the sulphur and nitrogen groups those oxides, in which the elements manifest their maximum valence, are the most stable (compare p. 174). The peculiar behavior of the oxygen compounds of the halogens, their variable stability and decomposition, and their modes of formation, find a clearer explanation in their thermo-chemical relations. All oxide compounds of chlorine and bromine are endothermic, t. e., heat is rendered latent in their production from the elements (compare p. 67). They do not result, therefore, by direct union of the elements ; further, they are not very stable, 16 186 INORGANIC CHEMISTRY. decompose readily with elfmination of oxygen, and then oxidize strongly. The heat, appearing in the formation of chlorine monoxide, and of the hypothetical pentoxides, CI2O5 and BrjO^ (in their production from the elements and solution in water), corresponds to the symbols: — l8.oj (ClgjOgjAq.) = 20.4 (Br2j05,Aq.) = — 43'5* In the formation of iodine pentoxide and iodic acid heat is liberated :— (^2j05) ~ ~t" 44-8; (IjOjjH) = 4- S7»S* This explains its stability in comparison with the chlorine and bromine compounds, and also the direct production of iodic acid by the oxidation of iodine. When the pentoxides are compared with each other, it is seen that the most heat is rendered latent in the formation of bromine pentoxide, BrjOs—the affinity of bromine for oxygen, consequently, is the lowest, that of iodine the greatest. This is also evident from the heat of formation of the acids, in dilute aqueous solution, or of the potassium salts in solid condition :— (Cl.OgjH.Aq.) = 23.9; (Br,03,H,Aq.) = 12.4; (I,03,H,Aq.) = 55.7; (C1,03,K) =94.6; (Br,03,K) = 87.6; (lA.K) = 128.4. We now understand why chlorine and bromine are separated from chloric and bromic acids by iodine, with formation of iodic acid, while broifiine does not act upon chloric acid. Later, we will observe that also in the groups of oxygen, nitrogen and carbon, the second members (sulphur, phosphorus and silicon) exhibit a greater heat-disengagement in their oxygen derivatives, than the immediately succeeding members (selenium, arsenic) ; the former belong to the second, and the latter to the third period of the natural system of the elements. OXYGEN COMPOUNDS OF THE ELEMENTS OF THE SULPHUR GROUP. The elements sulphur, selenium, and tellurium combine with two atoms of H, and also yield oxygen acids, which contain 2 H atoms :— H2S (SO2H2) H2SO3 H2S0^. In these acids i and 2 atoms of H can be replaced by metals ; hence they are dibasic. By the replacement of i atom of H we get the so-called acid or primary salts^ while the neutral or secondary salts are obtained by the replacement of both hydrogen atoms :— SO^KH SO4K2. Acid potassium sulphate. Neutral potassium sulphate. OXYGEN COMPOUNDS OF SULPHUR. 187 I. OXYGEN COMPOUNDS OP SULPHUR. (SO2H3) Hyposulphurous acid. 502 SO3H2 Sulphurous anhydride. Sulphurous acid. 503 SO4H2 Sulphuric anhydride. Sulphuric acid. In addition to these compounds there are others of more com¬ plicated nature. They will be studied later. The structure of the former may be expressed by the following formulas :— IV 0 = S = 0 Sulphur dioxide. VI o = s=o I Sulphur trioxide. IV .OH o = s( ^OH* Sulphurous acid. VI .OH 0=S< II ^OH O Sulphuric acid. Sulphur Dioxide, SO2, or sulphurous anhydride, is formed by burning sulphur or sulphides in the air : S -f- O2 = SO2. I vol. 1 vol. The combustion may also be effected by the action of metallic oxides (copper oxide, manganese peroxide) which give up their oxygen quite readily. It is most conveniently prepared for labora¬ tories by heating sulphuric acid with mercury or copper:— 2H2SO4 + Cu = CuSO^ + SO2 -h 2H2O. Copper sulphate. The acid is similarly decomposed by heating it with carbon :— 2S04H3 -1- C = 2SO2 + COj + 2H,0. By this method we get a mixture of carbon and sulphur dioxides, which are separated with difficulty. A more convenient method for preparing sulphur dioxide consists in allowing ordinary English sulphuric acid to act upon calcium sulphite, CaSOs. The latter is * The structure of sulphurous acid must probably be expressed by the formula, IV H — SOj — OH, according to which i atom of H is connected with sulphur, but the other is contained as hydroxyl. This s^pears from the carbon derivatives of sulphurous acid. Probably both structural cases exist in compounds as two isomeric series of neutral ethers of the acid are known. 188 INORGANIC CHEMISTRY. mixed with burnt gypsum (}i part) and water, then moulded into cubes, which can be introduced into a Kipp's generator, as in the preparation of oxygen (p. 8i). Owing to its solubility in water, sulphur dioxide must be collected over mercury. Sulphur dioxide is a colorless gas, with a suffocating odor. Its density is 32.03 (H — i), corresponding to the molecular formula SO, = 64.06. It condenses at —15°, or at ordinary temperatures under a pressure of two atmospheres, to a colorless liquid, of spe¬ cific gravity 1.45, which crystallizes at —76° and boils at —10°. Upon evaporation the liquid sulphur dioxide absorbs much heat ; so that if some of the liquid is poured upon mercury in a clay crucible, and the evaporation accelerated by blowing air upon it, the metal will solidify. Water dissolves 50 volumes of sulphur dioxide gas with liberation of heat. The gas is again set free upon application of heat. The solution shows all the chemical prop¬ erties of the free gas. Sulphur dioxide has great affinity for oxygen. The gases com¬ bine when dry ; if their mixture be conducted over feebly heated platinum sponge* sulphur trioxide results ;— 2SO2 -f O2 = 2SO3. 2 vols. 1 vol. In aqueous solution the dioxide slowly absorbs O from the air, and becomes sulphuric acid :— SO2 + H2O + O = H2SO4. Aqueous sulphur dioxide is converted more rapidly into sulphuric acid by the action of CI, Br, and I :— SO3TT2 + H2O + CI2 = SO^Hj + 2HCI. Here the decomposition of a molecule of water is effected in con¬ sequence of the affinity of the halogen for hydrogen and of sul¬ phurous acid for oxygen. On adding sulphurous acid to a dark- colored iodine solution the latter is decolorized. Similarly, sulphurous anhydride and its solution withdraw oxygen from many compounds rich in that element ; hence it deoxidizes strongly and passes over into sulphuric acid. Thus chromic acid is reduced to oxide, and the red solution of permanganic acid is decol¬ orized with formation of manganous salts. Many organic coloring substances, like those of flowers, are decolorized by it. This prop- * Instead of platinum sponge, platinized asbestos maybe applied; this is ob¬ tained by immersing asbestos in a platinic chloride solution, then in ammonium chloride, and afterward drying and igniting. SULPHUROUS ACID. 189 erty is what leads to its application in the bleaching of wools and silks, which are strongly attacked by the ordinary chlorine bleach¬ ing agents (p. 52). The dioxide may be deoxidized by strong reducing agents ; thus sulphur is separated out by HjS :— SO2 -j- 2H2S = 2H2O -}" 3^- If, however, both gases are strongly diluted by other neutral gases, the action is but very slow. A mixture of equal volumes of SO3 and CI2 unites in direct sunlight to thionyl chloride, SO2CI2- (p. 200). When sulphur dioxide actsup>on warmed phosphoric chloride, the products are phosphorus oxychloride, and the comp>ound SOClj :— SO 2 + PCI5 = FOC\ + SOCI3. Chlorthionyl—SOCI2—maybe viewed as sulphur dioxide in which one atom of O is replaced by two atoms of chlorine. It is a colorless liquid with a sharp odor, and boils at 78°. Water decomposes it into hydrogen chloride and sulphur¬ ous acid :— SOCI2 -f HjO = SOj + 2HCI. Sulphurous Acid—H2SO3—is not known in free condition, but is probably present in the aqueous solution of SOj. On cooling the concentrated solution to 0°, colorless cubical crystals sep¬ arate, which have the composition (SO2 -1- 15H2O) or (SO3H2 14H2O). If the aqueous solution is allowed to stand for some time, especially in sunlight, sulphur separates with formation of sulphuric acid :— 3S0a -f 2H2O = 2S0^H2 + S. Sulphurous acid is dibasic and forms two series of salts ; the pri¬ mary (KHSO3) and secondary (K2SO3). Sulphites.—These are obtained by saturating solutions of bases with SO2. When sulphurous acid is separated out from its salts by stronger acids it decomposes into its anhydride and water :— NaaSOs -f 2HCI = 2NaCl -f SO2 + H2O. Thermo-chemical Deportment.—Sulphur dioxide is a very powerful exother¬ mic compound. 71.0 Cal., are set free in its formation from sulphur (solid) and oxygen. When it dissolves in much water there is an additional disengagement of 7.7 Cal., so that the heat of formation of the hypothetical sulphurous acid in dilute aqueous solution (from sulphur, oxygen and water) equals 78.7 calories :— (SjOa) gas = 71.0; (SOjjAq.) = 7.7; (S,02,Aq.) = 78.7. In consequence of this great loss of energy the dioxide is a very stable com¬ pound; it is only at high temperatures that it sustains a partial separation into 190 INORGANIC CHEMISTRY. sulphur and oxygen. Its whole .chemical deportment fully corresponds to its thermo-chemical relations, e. g., its reduction by hydreien sulphide :— SO, -j" == S, 4" 2H,0-vapor. . . . Cal.) (71.0) (2 X 4-5) (2 X 58.0) For its behavior toward oxygen, see p. 189. Hydrosulphurous Acid.—11,80, or 8,04!!,. On adding zinc to the aqueous solution of sulphurous acid the metal dissolves without liberation of hydrogen. A yellow solution is obtained, which decolorizes indigo and litmus solutions ener¬ getically. These properties are due to the hydrosulphurous acid contained in the solution, formed there by the action of the H set free by the zinc upon a second molecule of 8O3H, :— HjSO, -|- Zn = SOjZn H,, and HjSO, + H, = 80,H, + H,0. The pure aqueous solution of the acid is obtained by the decomposition its salts. Its solution has an orange yellow color, reduces powerfully, bleaches and soon decomposes with separation of sulphur. The bleaching action of this lowest oxygen compound of sulphur reminds us of a similar behavior of the lower oxygen derivatives of chlorine and bromine. The salts are more stable than the acid. The sodium salt is obtained by the action of zinc filings upon a concentrated solution of primary sodium sulphite. Its composition is not established with certainty; it corresponds to either the formula 80, HNa or 8,04Na,. The salt solutions absorb oxygen very rapidly fi?om the air and change to sulphites. Two peculiar oxides of sulphur, resembling peroxides, are sulphur sesquioxide and sulphur beptoxide. Sulphur Sesquioxide—8,0,—is obtained by the solution of flowers of sul¬ phur in anhydrous sulphuric anhydride; it separates out in blue drops, which solidify to a mass resembling malachite. It decomposes gradually, more rapidly on warming, into 80, and sulphur. It is very violently broken up by water, with formation of sulphur, S02,804H2 and polythionic acids. It dissolves with a blue color in concentrated sulphuric acid. Sulphur Heptoxide—8,0,—is produced by the action of a silent electric dis¬ charge of great tension upon a mixture of 80, or 80, and oxygen. It sepandes, in oily drops, which solidify to a crystalline mass at 0°. Upon standing, but especially upon warming, it gradually decomposes into 80, and oxygen :— 8,0,= 280, + O. It fumes strongly in the air, and with water decomposes into sulphuric acid and oxygen :— 8,0, + 2H,0 = 28O4H, -f O. The solution of sulphur heptoxide in sulphuric acid is formed by the electrolysis of the latter acid diluted with water (40 per cent. 11,804). appears at the airáde SULPHUR TRIOXIDE. 191 together with oxygen and ozone (p. 94). The solution is also produced on the addition of aqueous hydrogen peroxide to strongly cooled and concentrated sul¬ phuric acid ; active oxygen is evolved simultaneously. The solution of the heptoxide in sulphuric acid shows oxidation reactions simi¬ lar to those of hydrogen peroxide (p. 104). It oxidizes ferrous sulphate to ferric oxide, gradually separates iodine from potassium iodide, and decolorizes the blue solution of indigo-sulphuric acid. It cannot, however, decolorize a potassium p>ermanganate solution, does not impart a blue coloration to a chromic acid solu¬ tion, nor does it act upon a titanic acid solution, and thus may be distinguished from hydrc^en peroxide (p. 104). If the aqueous sulphuric acid should contain more than 60 per cent. HjSO^ when electrolyzed, hydrogen peroxide will also result from the decomposition of the heptoxide first produced (p. 102). It is also formed by the gradual decomposition of sulphuric acid containing the heptoxide, SjOf, on standing. This will occur rapidly if the acid be concentrated beyond 70 per cent, by the addition of more acid. Persulphates.—By the electrolysis of potassium hydrogen sulphate in a divided cell, Marshall obtained potassium persulphate—KSO^. This salt forms tabular, apparently asymmetric crystals. It leaves secondary potassium sulphate on igni¬ tion, oxygen and sulphur trioxide being evolved. It is slightly decomposed by hot water. At the ordinary temperature its conversion into acid potassium sulphate with liberation of oxygen is slow. The pure aqueous solution of the salt is neutral to test paper. It readily oxidizes ferrous sulphate to ferric sulphate. It does not react immediately with potassium iodide; heat hastens the reaction. It bleaches litmus and turmeric slowly. It rapidly oxidizes a warm solution of potassium ferrocyanide to potassium ferricyanide. Alcohol is oxidized to aldeh^e by it. Strong nitric acid and sulphuric acid evolve oxygen as ozone from solid per¬ sulphates on the application of a gentle heat. Hydrochloric acid liberates chlorine. Ammonium persulphate—NH^SO^—has also been prepared. Its reactions are similar to those of the potassium salt. It is utilized in the preparation of other persulphates. The formation of pure persulphates in solid crystalline form conclusively demonstrates that persulphuric anhydride and persulpinuic acid are not peroxides CA- Ch. Soc.y 60, 771)« Sulphur Trioxide—SO3—or sulphuric anhydride, is produced, as previously described, by the union of SOj and oxygen, aided by platinum black ; or when SO2 and air are conducted over ignited oxide of iron (Wöhler). It can also be made by heating sodium or potassium pyrosulphate (p. 198). It is most conveniently obtained by heating fuming (Nordhausen) sulphuric acid (p. 198) ; the escaping white fumes are condensed in a chilled receiver. Sul¬ phur trioxide exists in two different (polymeric) modifications. In the one form obtained by cooling the vapors, there is produced a white, asbestos-like mass which, after fusion, crystallizes in long, colorless prisms; it melts at 16° and boils at about 46°. The vapor density agrees with the formula SO». By keeping it below 25° it passes into another so-called solid modification, which does not fuse until above 50°, and then reverts to the liquid variety. 192 INORGANIC CHEMISTRY. According to later investigations of Weber neither modification is the pure anhydride, as both contain water. He obtained the pure anhydride by subjecting the asbestos-like variety to repeated and careful distillations in a closed tube. It is a readily mobile liquid, of specific gravity 1.940 at 16®, but solidifies to long, transparent, needles, resembling saltpetre. The crystals fuse at 14.8® and boil at 46.2®. By the addition of a small quantity of moisture the transparent crystals pass into the asbestos-like needles of the ordinary anhydride. Sulphuric oxide fumes strongly in the air, and attracts moisture with avidity. When thrown on water it dissolves with hissing, to form sulphuric acid (SO3 -}- HjO = H2SO4). When the vapors are led through heated tubes they are decom¬ posed into SO2 and oxygen. Thermo-chemical Deportment.—When sulphur dioxide and oxygen combine to form liquid sulphur trioxide 32,1 Cal., are disengaged, so that its heat of formation from the elements is 103.2 Gal. :— (SOgjO) liquid = 32.1. (SjOj) liquid = 103.2. inasmuch as the heat of formation of the dioxide = 71.0 Cal. Generally, when the formation of several compounds is possible in a reaction, that one is produced which sets free the most heat. Thus carbon yields COj and not CO. But when sulphur burns it yields SO 2 almost exclusively and not SO3. This is due to the instability of the trioxide, which at a red heat decomposes at once into SO 2 and O. It is further necessary to bear in mind that in the formation of the more com¬ plex compounds, a great number of atoms or molecules are acting simultaneously, and therefore, at least in the cooperation of gases, the simpler compounds can be more readily produced. We now comprehend why reactions such as S2 + 3O2 = 2SO3, N2 -j- 3H2 = 2NH3 and C2 -j- 4H2 = 2CH^, do not usually take place. And another consideration is that all bodies havir^ a high heat of formation and decomposed at a high heat, must have their heat of formation removed or con¬ ducted away, in order that their production may be at all possible. This behavior explains many phenomena which apparently contradict the thermo-chemical principles, as for example, reactions generating little heat, taking place more readily than those having a greater heat-disengagement. The catalytic action of many metals, e.g., platinum, in the reaction, SOg -j- O = SO3, is very likely due to their conducting off the heat. Another case is the greater reactivity of many bodies where they form a gal¬ vanic chain, because in this way the chemical eneigy is removed, as electricity, e.g., the union of hydrogen and oxygen at ordinary temperatures due to the fiar- mation of a polarization current. On dissolving sulphur trioxide in much water to form aqueous sulphuric acid, 39.1 Cal. are disengaged. The production, therefore, of the aqueous acid from sulphur, oxygen, and water equals (including the heat of formation of SO3) 142.4 Qui ; (SO3, Aq.) = 39-1 ; (S, Og, Aq.) = 142.4- If we add to this the heat of formation of water (liquid)—68.3 Cal.,—the heat of formation of sulphuric acid ( HjSO^ = SO3 -f- HjO) from the elements in dilute aqueous solution will be :— (S, 0„ Hj, Aq.) = 210.7. SULPHURIC ACID. 193 The heat of solution of anhydrous sulphuric acid, in much water, equals 17.8 Cal. ; hence the heat of formation of anhydrous sulphuric acid from its elements is;— 210.7 — 17.8 = 192.9 : (SO4H2, Aq.) = 17.8. (S, O^, H2) = 192.9. For the relations of selenic and telluric acids, see p. 204. SULPHURIC ACID—H2SO4. This acid has long been known and is extensively applied in technology, etc. Besides the reactions already mentioned, it arises in the oxidation of sulphur by nitric acid. It was obtained formerly by heating ferrous sulphate (FeSOi) ; at present, however, it is almost exclusively manufactured in large quantities, after the so-called English lead chamber process. This method is based upon the conversion of SO2 into SO4H2. Sulphur or pyrite (FeS2), is roasted in ovens, and the disengaged SO2 immediately conducted, together with air, into a series of large leaden chambers in which it is frequently brought in contact with nitric acid and steam. By the combined action of these substances (sulphur dioxide, nitric acid, oxygen of the air and water) sulphuric acid is formed in the chambers and collects upon the floor of the same. The lead chamber process is very complicated, being influenced by the quantity of the reacting substances and the temperature, and as yet is not fully explained. It is most simply represented as follows : in the presence of water, the nitric acid oxidizes the SO 2 to sulphuric acid, and the former is reduced to nitric oxide or nitrogen dioxide :— 3SO2 -f 2HNO3 + 2H2O = 3H2SO4 + 2NO. The oxygen of the air (which entered the chambers simultaneously with the SO2 ) and the steam convert the NO again into nitric acid :— 2NO + 3O + H2O = 2HNO3, and this converts a fresh portion of SO 2 into sulphuric acid. Or, the nitric oxide forms N2O3 and NO2 by union with oxygen, and these in the presence of steam oxidize sulphur dioxide to sulphuric acid :— SO2 -f HoO + N2O3 = SO4H2 -f 2NO and SOj + H2O + NO2 = SO4H2 -f NO. The regenerated nitric oxide is again subjected to the same transformations. The chamber gases, according to the latest investigations of Lunge and Naeff, contain only nitrogen trioxide of the oxides of nitrogen, and in Lunge's opinion the lead chamber process depends chiefly upon the intermediate formation of nitrosylsul- phuric acid. At first nitrogen trioxide, sulphur dioxide, oxygen, and a little water react upon each other, and the product is free, solid nitrosylsulphuric (p. 211):— 2SO2 + N,0, -f O, -f HjO == 2S0,/g-N^. 17 194 INORGANIC CHEMISTRY. The excess of water in the lead chamber immediately converts this product into sulphuric acid and nitrogen trioxide :— + "A- The trioxide, regeuerated in this manner, acts again upon new portions of the sul¬ phur dioxide. According to Lunge, nitrosylsulphuric acid is formed upon the very ñrst action of the nitric acid introduced into the lead chamber :— SO, + NOj.OH = and it then reacts in the manner already indicated. In this way, apparently, one and the same quantity of nitric acid, by sufficient oxygen access and water. Fig, 84. changes an unlimited amount of SO, into sulphuric acid; the oxides of nitrogen act, as it were, as carriers of oxygen. As much as 2-3 per cent, of nitric acid is lost in practice. This is probably due to the further reduction of the nitrogen oxides to nitrous oxide (NjO) and nitrogen (which are not absorbed in the Gay- Lussac towers). The recent s'udies of Raschig point to the intermediate forma¬ tion of amidosulphonic acids in the lead chamber process, and the loss in nitrc^en is due to their conversion into ammonia. In practice, the active nitrogen oxides (N,03 and NO,) are carried along and withdrawn from the action by means of the escaping nitrogen and excess of air. To avoid any further loss of nitric acid by this means, the escaping brown gases SULPHURIC ACID. 195 are conducted throi^h the so-called Gay-Lussac tower. This is constructed from lead sheets, and filled with pieces of coke, over which concentrated sulphuric acid constantly trickles. The acid completely absorbs the nitrogen oxides NjOj, NOj and NO, with formation of nitrosylsulphuric acid (see p. 211). The nitrogen oxides can be regained from the acid—the so-called nitroso-acids—collected at the bottom of the tower, and made useful in the production of sulphuric acid in the chambers. This is effected at present, in the so-called Glover tower, which is constructed of lead plates and fire-proof bricks, and inserted between the sulphur ovens and lead chambers. In this the nitroso-acid (diluted with the previously obtained chamber acid) is allowed to run over fire-brick, while the hot gases of combustion from the sulphur ovens stream against it. This cools the hot gases to the required temperature (70-80°), water evaporates from the chamber acid, and, at the same time, the nitrogen oxides are set free (see p. 212), and carried into the lead chambers. Hence, the Glover tower serves, not only for complete utiliza¬ tion of the nitrogen oxides, but also for the concentration of the chamber acid. The chamber process may be illustrated by the following laboratory experi¬ ment: A large,, glass flask (Fig. 84) A replaces the lead chamber; in its neck are introduced, by means of a cork, several glass tubes, which serve to introduce the various gases. In a, SOj is developed by heating a mixture of HjSO^ and Hg or copper strips. The flask b contains some dilute nitric acid and copper turnings, from which NO is evolved. Water is boiled in c to afford steam. Air enters through d while the excess of gases escapes through e. By the meeting of NO with the air, red fumes of nitrogen dioxide (NOg) and nitrogen trioxide (NjOg) arise, and these in presence of water change the sulphur dioxide to sulphuric acid (p. 193). The regenerated nitric oxide 3rields NOg with the oxygen of the air, and converts another portion of SOg into sulphuric acid. In time aqueous sulphuric acid collects upon the bottom of the vessel. If, at first, only SOg, NO and air enter without the steam, we get (by aid of the moisture of the air) the compound SOj I(the so-called nitrosulphonic acid) which covers the walls of the vessel with a white crystalline sublimate (comp. p. 212). These crystals, known as lead-chamber crystals, are also formed in the technical manufacture of sulphuric acid, when an insufficient quantity of steam is conducted into the chambers. Water decomposes them into sulphuric acid and nitrogen oxides. The acid collecting in the chambers (chamber acid) possesses, when the operation has been properly conducted, the specific gravity 1.5 (50° according to Beaumè); it contains about 60 per cent. H2SO4 and 40 per cent. HjO. For concentration the chamber acid is first heated in open pans until the specific gravity reaches 1.72 (60° Beaumè). The lead vats are strongly attacked by further evaporation, hence the acid is finally heated in glass vessels, or, better, platinum retorts, until the residual liquid has acquired the specific gravity, 1.83 (65.5° Beaumè). It is now entered upon trade under the name crude sulphuric acid (^Acidum sulphuricum cruduni). It still contains about 8 per cent, water and traces of lead and arsenic. By further concentration we can obtain 95-96 per cent. HjSOi (extra concentrated acid). By the distillation of the crude English acid an aqueous solution at first distils over ()^ distillate), but at 330° we obtain almost pure 196 INORGANIC CHEMISTRY. H2SO4 {Acidum sulphuricum or destillatum). This has the specific gravity 1.854 at 0° or 1.842 at 12°, and contains about 1.5 per cent, water. On cooling this to —35° white crystals separate, which, after repeated recrystallization fuse at -(- 10.5° ; this is the anhydrous acid, H2SO4. The crystalline acid is more readily obtained by cooling the 96-98 per cent, sulphuric acid to o° or —10°, and then adding already-formed crystals. This is the manner in which the anhydrous acid is produced technically ; the crystals are separated from the liquid, hydrous acid by a centrifugal machine. Pure anhydrous sulphuric acid, H2SO4, has the specific gravity 1.8372 at 15°, and is, therefore, lighter than slightly hydrous acid. When this is heated, white fumes of SO3 escape at 40° ; the liquid begins to boil at 200°, and at 330° the acid, with 1.5 per cent. H2O, again distils over. From these data it is obvious that sulphuric acid, even at a gentle heat, sustains a partial decomposition (dissociation) into SO3 and HjO, which again unite to form sulphuric acid when they cool. The vapor density, corresponding to the molecular formula H2SO4 = 98, must be equal to 49 (®2®)* The vapor density of sulphuric acid has been found to be 36.0 (or 2.5-air = i) at 332® (near its boiling point). It diminishes at higher temperatures, and is 24.5 at 416°, where it is constant. This behavior is explained by the dissociation of the acid molecules, according to the equation :— SO.H^ : SO3 + H3O. I vol. I vol. I vol. The vapor density corresponding to the molecular formula IIjS04 = 98 must be 49 ; that of the mixture of the decomp>osition products, filling a volume twice as large, is = 24.5. Hence, the dissociation of the acid is complete at 416°, while it is only about 34 per cent, at 332® (p. 96). Concentrated sulphuric acid is a thick, oily liquid. On cooling, a sulphuric acid containing about 15 per cent, water to 0°, large six-sided prisms of the hydrate SO4H2 -f- HjG separate ; these fuse at 8.5°, and give up water at 250°. The second hydrate^ SO4H, -j-2H20, corresponding to the maximum contraction, has the specific gravity 1.63, and yields water at 195°. The concentrated acid possesses an extremely great affinity for water, and absorbs aqueous vapor energetically, hence is applied in the drying of gases and in desiccators. It unites with water with the evolution of considerable heat, and, for this reason, it is practically recommended, in mixing the acid, to pour the latter in a thin stream into the water, and not the reverse, as otherwise explosive phenomena occur. In mixing sulphuric acid with water, a contraction of the mixture takes place ; its maximum corresponds to the hydrate SO4H2 -f- 2H,0. The existence of the hydrates of sulphuric acid is explained, as in the case of periodic acid, by the supposition of hydroxyl groups ;— PYROSULPHURIC OR DISULPHURIC ACID. 197 SO4H, -}- 2H2O = S(OH)g Hexahydroxyl sulphuric acid. SO^H, 4- HjO = SO(OH)^ Tetra " SO^H, = S02(0H)j Normal sulphuric acid. The tetra- as well as the hexahydroxylsulphuric acid yield only salts of the normal dibasic acid, when they are acted upon by bases. Salts, in which several H atoms are replaced by metals, are not known. The affinity of sulphuric acid for water is so great that the former withdraws the hydrogen and oxygen from many substances, with the production of water. In addition to carbon, many organic compounds contain hydrogen and oxygen in the proportion in which these elements yield water. The withdrawal of H and O from such substances leaves the carbon. This explains the charring action of H2SO4 upon wood, sugar, and paper. When sulphuric acid acts upon alcohol (CzHgO), ethylene, C2H4 (p. 155), results. By conducting H2SO4 over red hot porous bodies, it is decom¬ posed into sulphur dioxide, water, and oxygen :— H2SO4 = SO2 -f H2O + O. This decomposition affords us a method for manufacturing oxygen technically ; the sulphur dioxide is absorbed by water and after¬ ward converted into H2SO4. When heated with S, P, C, and some metals (Hg, Cu), the acid is reduced to dioxide (p. 187). Nearly all the metals are dissolved by it, forming salts ; only lead, pla¬ tinum, and a few others are scarcely attacked at all. It is a very strong acid, and, when heated, expels most other acids from their salts; upon this depends its application in the manufacture of hydrochloric and nitric acids. The barium salt (BaS04) is charac¬ terized by its insolubility in water and acids ; therefore, sulphuric acid added to solutions of barium compounds produces a white pulverulent precipitate, which serves to detect small quantities of the acid. Pyrosulphuric, or Disulphuric Acid—HjSjOt.—On with¬ drawing one molecule of water from two molecules of the acid there results the compound SaOtHa, whose formation and structure may be represented by the following formula : — SO so gq/OH As this contains two hydroxyl groups it is a dibasic acid ; yet its manner of formation shows that it possesses an anhydride character. Later, we will observe that almost all polybasic acids, like phosphoric acid, PO(OH)s, silicic acid, SiO(OH)j, and chromic acid, CrO, 198 INORGANIC CHEMISTRY. (0H)2, are capable, by the condensation and the elimination of several molecules of water, of forming like derivatives, which bear the name Poly- or Pyro-acids. The disulphuric acid is contained in the so-called fuming or Nordhausen sulphuric acid \Acidum sulphuricum fumans\ which is obtained by heating dehydrated ferrous sulphate—^green vitriol (FeS04). It is a thick, oily, strongly-fuming liquid, of specific gravity 1.85-1.9. When it is cooled, large colorless crystals of H2S2O7 separate ; these fuse at 35°. Heat breaks it up into sul¬ phuric acid and sulphur trioxide, which volatilizes;— SjOijH, = SO^Hj -f- SO3. Conversely, disulphuric acid may be obtained by dissolving SOj in sulphuric acid :—SO4H2 + SO3 = S2O7H2. The production of fuming sulphuric acid also depends on this, as it may be regarded as a solution of SO3 (or S2O7H2) in excess of sulphuric acid. Technically, fuming sulphuric acid is obtained from pyrites (FeS2)—(at present only in Bohemia). The decomposition of the pyrites in the air affords ferrous sulphate and ferric oxide. The first can be dissolved out with water. The solution is evaporated, and the residue roasted in a reverberatory furnace, where¬ by the ferrous salt is changed to ferric salts. The latter are then distilled from earthen retorts, when sulphuric acid and the trioxide pass over and are collected in the receivers :— Fe2(S04)3 = FejOg 4- 3SO3. Fe i (^04)2 = Fe203 + SO3 + S0,H2. ^n(0H)2 The residue, consisting of red ferric oxide, finds application as colcothar {caput mortuum) in polishing and as a paint. Solid, crystalline pyrosulphuric acid has been recently introduced into the market as a substitute for the fuming liquid sulphuric acid. It is made by con¬ ducting the theoretical amount of sulphur trioxide into concentrated sulphuric acid (see above). Sulphur trioxide is prepared by two distinct methods at present. In Winkler's method sulphuric acid of 66° Baumè is first allowed to run into retorts raised to a red heat. The mixed gases, SOj, O and H2O (p. 195) resulting from this action, are freed from steam by passing through a coke tower through which there trickles conc. sulphuric acid. The dry mixture is then conducted over ignited platinized asbestos (p. 188) and the resulting SO3 collected in concentrated sulphuric acid. The more recent method of Wolter consists in producing sodium pyrosulphate by heating sodium sulphate with concentrated sulphuric acid :— SO^Na^ + S0^H2 = S207Na2 + H2O. Sodium Sodium sulphate. pyrosulphate. An intermediate product in this reaction is primary sodium sulphate—SO^NaH. Upon stronger heating sod. pyrosulphate is broken up into the neutral sulphate and sulphur trioxide ;— SjO^Na, = S04Na3 -f- SO5. SULPHURIC ACID CHLOR-ANHYDRIDES. 199 The sulphur trioxide which escapes is collected in sulphuric acid, and the rest dual sodium sulphate again converted into pyrosulphate. Sulphuric Acid Chlor-anhydrides.—Under the name of halogen anhydrides we understand the derivatives resulting from the replacement of OH in hydroxides by halogens. Conversely, the chlor-anhydrides, by the action of water, pass into the corresponding acids :— OH SOj{Q + 2H2O = SO^ -f 2HCI. \)H The ordinary method for the preparation of the chlor-anhydrides consists in permitting PClg to act on the acids. Sulphuric acid has two hydroxyl groups ; tlierefore it can furnish two chlor-anhydrides. The firsty SOji^Q^ — Sulphuryl Hydroxy-chloride or Chlorsulphonic Acid — results when l molecule of PClg acts upon i molecule of H^SO^ :— CI SO2 { oS + PClg = sol + POCI3 + HQ. \h The resulting POClj acts upon two additional molecules of SO^Hj, with forma¬ tion of metaphosphoric, hydrochloric, and chlorsulphonic acids. It is formed, too, by the direct union of SO3 with HCl. The most practical method for its formation consists in conducting chlorine gas through SO^Hj (15 parts),and gradually adding PCI3 (7 parts). Or, HCl gas is led into solid fuming sulphuric acid (SjO^Hj), long as absorption occurs, and then it is distilled (Otto). Chlorsulphonic acid is a colorless, strongly fuming liquid of specific gravity 1.776 at 18°, and boils at 155°. Its vapors possess the normal density at a tem¬ perature a little beyond the boiling point, but at 180° sustain dissociation, which is complete at 440°, and corresponds to the equation :— 2SO3CIH = SO3 + HjO + SO2 + CI2. The salt SOj results from tlie union of SO3 with KCl. The second chloranhydride, SOjClj, or Sulphuryl-chloride* forms when PClg acts upon SOg; by heating SO3HCI to 180°: 2S0gHCl = SOjClj -|- SO^Hj; and also by the direct union of SO, with CI3 in sunlight:— SO2 + CI, = SOjCl,. I vol. I vol. I vol. The most convenient method for its formation is to pass equal volumes of sulphur dioxide and chlorine gas into a capacious flask, containing some camphor, which causes the union of the gases to form sulphuryl chloride. A colorless, suffocating, * The group SO, combined with 2OH in H,SO^, is known as SulphuryL 200 INORGANIC CHEMISTRY. ^ongly faming liquid, of specific gravity 1.708 at 0°, results. It boils at Water decomposes it energetically into sulphuric and hydrochloric acids (also the case with chlorsulphonic acid). A little water will first change it to chlorsulphonic acid:— ®o»\a + "'O = ^'\oH + Its vapor density is normal at 184®; at higher temperatures it gradually dimiidshes, and at 44®** corresponds to the decomposition equation SOjQj = SO, + CI,. Pyrosulphuryl Chloride, SjOgCl,, is the chloranhydride of pyrosulphuric acid. It is obtained by several reactions, chiefly, however, by the action of PQg or PjOj upon chlorsulphonic acid :— 2SO,<^OH — SO,/ ^a. It is a thick liquid, fuming in the air; has a specific gravity = 1.858 at o®, and, when perfectly pure, boils at 153®. At 183-210® it shows a normal density; it is dissociated at higher temperatures, and this is complete at 440®, corresponding to the equation:— SjOjCl, SO3 + SO, + CI,. It dissolves gradually in water, without hissing, and decomposes into SO^H, and HCl; it at first yields chlorsulphonic acid with a httle water. Thionyl chimie—SOCl,—(p. 189) may be regarded as a chloranhydride of sulphurous acid. POLYTHIONIC ACIDS. By this name (from BtTov^ sulphur) are understood the complex acids of sulphur, containing two or more atoms of the latter. The following are known:— S,03H, — Thiosulphuric acid. SjOjH, — Dithionic acid. SgOgH, — Tri « " S^OgH, — Tetra " '♦ SgOgH, — Penta « " The general chemical character of these acids is represented simply and distinctly in the following structural formulas. We suppose that they contain one or two univalent groups, SOsH, or THIOSULPHURIC ACID. 201 VI — SOjOH, in which one affinity of sulphur is unsaturated. This group is known as the sulpho group; it is also present in organic sulpho-acids, and corresponds to the acid-forming carbon group, COOH, called carboxyl. From this group (written in another form) are derived the above-observed acids:— H.SO2.OH HO.SO2.OH ^^xsOjOH Sulphurous acid. Sulphuric acid. Disulphuric acid. The constitution of the polythionic acids is expressed by the following formulas :— HS so H c/'S03H ilb.bUsil \SO3H \SO3H Thiosnlphuric acid. Dithionic acid. Trithionic acid. c /SO3H o /SO3H ^2\S03H ^3\S03H Tetrathionic acid. Pentathionic acid. The last three acids may be viewed as derivatives of the hydrogen sulphides, SHj, S2H2, and S3H2, in which both H atoms are replaced by two univalent sulpho groups. In thiosnlphuric acid, only iH- atom is replaced by sulpho; the dithionic acid, on the other hand, results by the direct union of two sulpho groups, with their free affinities. / QJT Thiosulphuric Acid, H2S2O3 = S02.(^qjj, generally known as hyposulphurous acid, can be considered as sulphuric acid in which the oxygen of an hydroxyl group is replaced by sulphur. It is not known in a free condition, since as soon as it is liberated from its salts by stronger acids, it at once decomposes into SO2, S and H2O SjOgNaj -f 2Ha = 2NaCl SO, + S + HjO. Its salts, called hyposulphites^ are of practical importance (compare sodium hyposulphite). They are formed by the direct addition of sulphur to sulphites :— NajS03 S = NajS203 ; similar to the formation of sulphates by the addition of O to the sulphites. A very interesting formation of thiosulphuric acid is that of the action of iodine upon a mixture of sodium sulphite and sodium sulphide :— r NaSOj.ONa SO,.ONa -I-1, = I -f- 2NaI. NaSNa SNa Sodium hyposulphite. 1; 202 INORGANIC CHEMISTRY. Conversely sodium h)rposulphite is split up by sodium amalgam into SOsNaj and NajS. Dithionic Acid—HjSjOj—is only known in aqueous solution. When concen¬ trated in vacuo or when heated it decomposes into sulphuric acid and sulphur dioxide. Its manganese salt results from the action of sulphur dioxide unon MnOj suspended in water :— MnOj -f- 2SO, ■= MnSjOg. Barium hydroxide converts this into the barium salt, from which the free dithionic acid is obtained by means of sulphuric acid. It is very doubtful whether dithionic acid is produced by the action of an iodine solution upon primary sodium sulphite, NaSOj.OH (Spring). Trithionic Acid—HjSgOg—is not known in a free condition. Its salts are produced when an aqueous solution of primary potassium sulphite is digested with flowers of sulphur :— 6HKSO3 + 28 = aKjSjOe + + 3H3O. Separated from its salts by other acids it decomposes into HjSO^, SO2 and S. Its production by the action of iodine upon a mixture of sodium sulphite and hypo¬ sulphite is especially interesting :— NaSOo.ONa SOj.ONa NaS.SOj^'oNa + ^2 — + ^Nal. SOj.ONa Tetrathionic Acid—1128405. Its salts are produced when iodine acts upon solutions of the hyposulphites :— KS.SO3K S.SO3K + I2 = I + 2KI. S.SO3K Potassium, tetrathionate. KS.SO3K S.SO3K This reaction is applied in volumetric analysis for the quantitative determination of iodine, and such substances as separate iodine from potassium iodide (see sodium thiosulphate). The tetrathionic acid, separated from its salts by stronger acids, is very unstable, and when its aqueous solution is concentrated it decomposes into sulphuric acid, sulphur dioxide, and sulphur. An aqueous solution of the acid can be easily pre¬ pared by conducting hydrogen sulphide into aqueous sulphurous acid :— 4SO2 -j- 3lla^ — S^OgHj -j- 2H3O -|- 3S. If the solution be saturated with bases, neutral and acid tetrathionates result, e.g., S^OgBa 2H2O and (S40jH)2Zn. Pentathionic Acid, SgOçHg, is supposed to be produced together with the tetra¬ thionic acid when hydrogen sulphide acts upon aqueous sulphurous acid. Later researches, however, have proved that its supposed salts are identical with the acid tetrathionates (Curtius). Another method of production is supposed to be that in which sulphur chloride acts upon barium hyposulphite :— OXYGEN DERIVATIVES. 203 SjClj -f- Ba (8.803)2 Ba = 8 (8.80,>2 Ba + BaQ, + 8. It is, however, doubtful whether pentathionic acid is produced in this reaction. The polythionic acids are distinguished from sulphuric acid by the solubility of their barium salts. Amide Derivatives of the Sulphuric Acids.—DiflFerent nitrogenous com- px>unds, which must be regarded as derivatives of ammonia and hydroxylamine (p. 131), result upon introducing sulpho-groups, 8O2.OH, for amide hydrogen:— H,N.803H HO.NH.8O3H. Amido sulphonic acid, Hydroxylamine sulphonic or sulphaminic acid. acid. HN(803H), H0.N(803H), Imidosulphonic acid, Hydroxylamine disuiphonic or disulphimidic acid. acid. N(803H)3 Nitrilo-sulphonic acid. a. OXYGEN DERIVATIVES OF SELENIUM AND TELLURIUM. SeO, Se03H2 Selenium dioxide. Selenious acid. (SeO,) SeO^H, Selenium trioxide. Selenic acid. Selenium Dioxide—SeOj, or selenious anhydride, is produced when selenium burns in the air or in oxygen. It consists of long white needles, which sublime at about 320° without fusing. It dissolves readily in water, forming selenious acid, H2SeO,. The latter is also obtained by dissolving the metal in concentrated nitric acid. When the solution is evaporated it crystallizes in large, colorless prisms, which decompose, on heating, into the anhydride and water. Sulphurous oxide reduces selenious acid, with separa¬ tion of free selenium :— H28e03 + 280, 4- HjO = 2H2SO^ + 8e. Selenic Acid—H2Se04—is obtained by conducting chlorine gas into an aqueous solution of selenious acid :— H28e03 + H3O + CI2 = HjSeO^ -f 2HCI. The solution may be concentrated until it attains a specific gravity of 2.6 when it becomes an oily liquid, similar to sulphuric acid, and contains 95 per cent. H28e04. If the solution be heated above 280®, the acid breaks up into 8e02,0 and HjO. 8elenic anhydride is unknown. The salts of selenic acid are known as selenates, those of selenious acid as selenites- 204 INORGANIC CHEMISTRY. The derivatives of tellurium are analogous to those of selenium. The dioxide —TeOj—results when tellurium is burned, and forms a white cry^alline mass, fusing at a red heat and subliming. It is almost insoluble in water. Tellurous Acid—HjTeOg—^is produced when the metal is dissolved m con¬ centrated nitric acid. Water will precipitate it from such a solution in the form of a white amorphous powder. On warming it readily decomposes into TeOj and water. Telluric Acid—HjTeO^. Potassium tellurate is produced when tellurium or its dioxide is fused with saltpetre. The acid (telluric) is obtained from this salt by means of sulphuric acid, and crystallizes from its aqueous solution, in large colorless prisms, with 2 "molecules of HjO (HjTeO^-}- 2H2O), which are expelled at 100° C., and the acid remains in the form of a white powder. The latter is not very soluble in water, and manifests but a slight acid reaction. When care¬ fully heated, the acid breaks up into water, and the trioxide TeOg, which is a yellow mass insoluble in HjO, and by further application of heat decomposes into TeOj and oxygen. The affinity of the elements of the oxygen group for the haïtiens seems to increase with rise of atomic weight from oxygen to tellurium. This is just the reverse of what we observed for the halogens with hydrogen; OClj is very imstable and is formed with heat absorption; SCI, and SCl^ only exist at lower tempera¬ tures, while SeCl^, TeClj, and TeCl^ even exist as gases. The oxygen derivatives of sulphur, selenium, and tellurium appear to conduct themselves similarly. The thermo-chemical relations, however, show that the gradation is not complete, but that the heat of formation of the derivatives of sulphiu: is greater than that of those of selenium—as is the case in the chlorine group (p. 185). This may be seen in the heat of formation of the anhydrous dioxides:— (SjOj-gas) = 71.0. (Se,02)-solid = 57.7, and also in that of the acids from their elements and water (page 189) :— (S,02,Aq.) = 78.7 (Se,02,Aq.) = 56.7 (Te,0„Aq.) = 81.2 (S,03,Aq.) = 142 (Se,03,Aq.) = 77.2 (Te,03,Aq.) = 107.0 In all these compounds, consequently, the affinity of selenium to oxygen is the least, and this explains the reduction of selenious by sulphurous acid, as well as the slight stability of selenic acid. OXYGEN DERIVATIVES OF THE ELEMENTS OF THE NITROGEN GROUP. The halogens combine with one atom of hydrogen and also afford oxygen acids containing one atom of H. The elements of the sulphur group contain two atoms of H in the" hydrogen deriva¬ tives and oxygen acids. In accord with this we find that the elements of the N group combine with three atoms of H, and form acids which also contain three atoms of the same element:— HCl HgS PHj HCIO4 H2SO4 PO4H, Perchloric acid. Sulphuric acid. Phosphoric acid. HCIO3 HjSO, PO5H3 Chloric acid. Sulphurous acid. Phosphorous acid. OXYGEN DERIVATIVES. 205 The acids containing three atoms of H, designated normal or Ortho-Acids (as HgPOi, HgAsO^, HgAsOg) can yield monobasic acids by the removal of one molecule of water. Such derivatives, having one atom of H, are called meta-acids:— H3PO, HPO3 Orthophosphoric acid. Metaphosphoric acid. H3ASO3 HAsO, Orthoarsenious acid. Metaarsenious acid. These meta-acids of phosphorus and arsenic are less stable than the ortho-acids and pass into the latter by the absorption of water. The ortho-acids of N, on the other hand, are less stable and only exist in some salts. The ordinary acids and salts of N belong to the meta-series and contain one atom of H (or metal) :— (H3NOJ HNO3 Orthonitric acid. Ord. Nitric acid. (H3NO3) HNO3 Orthonitrous acid. Ord. Nitrous acid. The further exit of water affords the true anhydrides (p. 178). 1. OXYGEN DERIVATIVES OF NITROGEN. NP3 Nitrogen pentoxide. N3O3 Nitrogen trioxide. NO3H Nitric acid. NO^H N3O Nitrous acid. (N0H)3 Hyponitrous oxide. Hyponitrous acid. In addition to the above compounds we have : Nitrogen tetroxide (Nj04), the mixed anhydride of nitrous and nitric acids, and two oxides, nitrogen dioxide (NOj) and nitric oxide (NO), which do not yield acids. The following formulas express the structure of these compounds :— N = N Nitrogen. V 0,N — OH Nitric acid. Ill III N = N \/ O Hyponitrous oxide. Ill ON —OH Nitrous acid. Ill III ON — O — NO Nitrogen trioxide. Ill (ON - H)., Hyponitrous acid. OjN — o — NOj Nitrogen pentoxide. Ill ON- O — NO« Nitrous-Nitric anhydride. The salts of nitric acid are called nitrates ; those of nitrous, nitrites. 206 INORGANIC CHEMISTRY. NITRIC ACID.—HNOj. This acid occurs in nature only in the form of salts,—potassium, sodium, and calcium nitrates (compare these)—which have resulted from the decay of nitrogenous organic substances in the presence of strong bases (the alkalies). It is sometimes present in the air as ammonium salt. The free acid is formed in very slight quantity by conducting the electric sparks through moist air. To prepare nitric acid heat potassium or sodium nitrate with sulphuric acid, when the nitric acid will distil over and sodium sulphate remain :— aNaNO. + H^SO^ = Na^SO^ -f 2HNO3 and NaNOg + HgSO^ = HNaSO^ + HNO3. This process may be conducted in the distillation apparatus figured on page 53. The quantity, by weight, of sodium nitrate and sulphuric acid corresponding to the second equation must be employed, because with less acid a higher temperature is requisite to complete the reaction, and, in consequence, the nitric acid that is produced will be partially decomposed. Pure anhydrous nitric acid is a colorless liquid of specific gravity 1.54 at 0°, fumes in the air, and at—40° solidifies to a crystalline mass. At ordinary temperatures it undergoes a partial decomposi¬ tion (like H2SO4) into water, oxygen, and nitrogen dioxide, NOj, which dissolves in the acid, with a yellowish-brown cOlor; the colorless acid therefore becomes colored upon standing, and in sun¬ light soon turns yellow. At 86° the acid commences boiling and sustains a partial decomposition ; the first portions are colored yellow by the dissolved nitrogen dioxide, but subsequently, some aqueous acid distils over. Nitric acid is completely decomposed into nitrogen dioxide, oxygen, and water, when its vapors are con¬ ducted through red-hot tubes:— 2HNO3 = 2NO2 4- HjO 4- O. The acid mixes in all proportions with water. Upon distilling the dilute aqueous solution, only pure water passes over at first; the boiling temperature gradually rises, and at 121° a solution goes over, which contains 68 per cent. HNO3, and has a specific gravity of 1.414 at 15°. This is the ordinary concentrated nitric acid of trade. When this is distilled with 5 parts sulphuric acid, an almost anhydrous acid is obtained, which may be freed of NOa contained in it by conducting a stream of air through it. Generally, the anhydrides of acids distil at temperatures lower than the acids themselves (SO, is more volatile than H^SO^). The higher boiling-point of the aqueous nitric acid, in relation to the anhydrous, is probably explained by the NITRIC ACID. 207 fact that the hydrate HNO, -f- i. e., the normal nitric acid, NO^Hj = NO- (0H)3 (p. 205), is present in this solution. The liquid boiling at 121®, however, contains more water than corresponds to this hydrate (just as distilled sulphuric acid contains water), so that it can be regarded as a mixture of the trihydrate (NO(OH)j) and pentahydrate (N(0H)5). Nitric acid is a very powerful acid, oxidizing or dissolving almost all metals (gold and platinum excepted). Nearly all the metalloids, like sulphur, phosphorus, and carbon, are converted by it into their corresponding acids. It acts as a very strong oxidizing agent, destroying organic coloring substances, and decolorizes a solution of indigo very readily. In so doing the nitric acid itself is deoxi¬ dized to the lower oxidation products of nitrogen (NO and NOj). Some substances even reduce the acid to ammonia. Thus, for example, if zinc be brought into dilute nitric acid (5-6 per cent.) the metal will be dissolved without the liberation of hydrogen. The latter, in statu nascendi^ acts at once upon the excess of acid and reduces it to ammonia, which forms an ammonium salt with the acid ; hence, in solution, we have ammonium nitrate in addition to the zinc nitrate 2HNO3 4- Zn = Zn(N03)2 -|- and 2HNO, + 4H3 = NOjNH^ + 3H3O. If the aqueous nitric acid be less dilute (containing more than 10 per cent. NO3H) it will be reduced by zinc and other metals, not to ammonia, but to the nitrogen oxides, N2O, NO, N2O3, and N2O4. The more concentrated the acid, the higher will the oxides be. The reduction of nitric acid to ammonia by nascent hydrogen occurs more easily in alkaline solution. If an alkaline solution of nitrates be treated with zinc or aluminium filings, all the N of the nitric acid will be converted into ammonia :— HN03+4H2 = NH3+3H30. Hydroxylamine (p. 131) and ammonia are produced when nitric acid acts on tin. Nitric acid—NOjOH—and its hydrates, NO(OH)3 and N(0H)5 (see above), usually form salts of the form NOaMe, with i eq. of the metals ; these are called nitrates, and are all soluble in water. Red Fuming Nitric Acid{Acidum Nitricum Fumans^ is the name given a nitric acid containing much nitrogen dioxide in solution. It is obtained by the distillation of 2 molecules of KNOs with i molecule of sulphuric acid (p. 206), or better, by the distillation of commercial nitric acid with much sulphuric acid. It generally has the specific gravity i.5-1.54, and possesses greater oxidizing power than the colorless nitric acid. 208 INORGANIC CHEMISTRY. A mixture of i volume nitric acid and 3 volumes concentrated hydrochloric acid is known as aqua regia, as it is able to dissolve gold and platinum, which neither of the acids alone is capable of doing. The powerful oxidizing action of the mixture is due to the presence of free chlorine and the two chlorine derivatives (NOjCl and NOCI), which may be considered the chloranhydrides of nitric and nitrous acids. Nitrogen Pentoxide—N2O5—nitric anhydride, is produced when phosphoric anhydride acts on nitric acid :— 2HNO3 + P2O5 = N2O5 + 2HPO3 ; further, by conducting nitroxyl chloride over silver nitrate :— Ag0.N02 + NO2CI = NO' }^ + It forms colorless, rhombic prisms, fusing at 30° and boiling with partial decomposition at 47°. It is very unstable, decomposing readily into N2O4 and O, and sometimes exploding spontaneously. It yields HNO3 with water and evolves much heat by the union :— JJg2|0 + H30 = 2N030H. Nitroxyl Chloride—NOgCl—the chloranhydride of nitric acid, results by the union of NO, with chlorine, and according to the ordinary method of forming chloranhydrides (see p. 199), by the action of PCI5 or POCI3 upon nitric acid, or better, its silver salt :— 3N020Ag + POCI3 = PO(OAg)3 + 3NO3CI. Silver nitrate. Silver phosphate. Recent investigations make its existence questionable. Nitrosyl Chloride—NOCI—is produced by the union of NO (2 yolumesywith chlorine (1 volume) : 2NO -j- CI, = 2NOCI, and when phosphoric chloride, PCL, "s allowed to act upon liquid nitrogen tetroxide, N^O^, in the cold. The reddiw yellow vapors that escape, if cooled to —20°, condense to a red liquid of sp. gr. 1.416 at —12®, and boiling at It is a reddish-yellow gas which below o® condenses to a liquid. It forms nitrous and hydrochloric acids with water :— NOCI + H,0 = HNO, -f HCl. It may, therefore, be regarded as the chloranhydride of nitrous acid—NO.OH. Nitrogen Trioxide—NaOg—nitrous anhydride, is formed by the direct union of nitrogen oxide (4 volumes) with oxygen (i volume) at —18° ;— NITROUS ACID. 209 4NO -}- O2 — 2N2O3; 4 vols. I vol. 2 vols. by mixing liquid nitrogen tetroxide, N2O4, with a little cold water :— ^N0'}° + = No}o + =N0,.0H¡ by the introduction of nitric oxide into liquid nitrogen tetroxide;— N204 + 2N0= 2N2O3; and by conducting nitric oxide into anhydrous nitric acid :— 2NO3H + 4 NO = 3N2O3 + H2O. It is most easily obtained by the action of nitric acid upon arsenious acid, AsjOg. Nitrogen tetroxide is simultaneously pro¬ duced, from which it is readily separated by fractional distillation and condensation. Nitrogen trioxide condenses at —20° to a dark-blue liquid of 1.449 sp- gr. at 0°, and boils at decomposes when dis¬ tilled; its vapors consist of a mixture of the tetroxide and nitric oxide (N2O4 4" 2NO). Upon cooling they re-unite to the liquid nitrous anhydride. The latter is therefore only known in the liquid condition (Geuther). The trioxide mixed with a little cold water probably forms nitrous acid (HNO2) i more water, aided by heat, decomposes it into nitric acid and nitric oxide gas :— 3NO2H = HNO2 -1- 2NO + H2O. Nitrous Acid, HNO2, is not known in a free state. Its salts (the nitrites) are obtained by igniting the nitrates:— KNO3 = KNO2 -f O. The withdrawal of oxygen is rendered easier if oxidizable metals, lead, be added to the fusion. On adding sulphuric acid to the nitrites, brown vapors are disen¬ gaged ; these consist of NOj and NO. It may be that the nitrous acid, at first liberated, is broken up into water and the trioxide, which, as we have seen above, gradually decomposes into NOj and NO. Similar reddish-brown vapors are obtained if nitric acid be permitted to act upon starch or arsenious oxide (AsjOg). According to Lunge, if we employ nitric acid of i.30-1.35 sp. gr., nitrogen tri¬ oxide is produced almost exclusively, whereas in using the concentrated acid (1.4-1.5) we get a mixture of Ng08 and NjO^, and chiefly NO, with a little NjO,, if the acid be dilute. The nitrous acid that has separated out in the solution and its decomposition products—NO, and NO—are strong oxidizers, set¬ ting iodine free from the soluble iodides. In other cases, however, 18 210 INORGANIC CHEMISTRY. they exhibit a reducing action ; thus, e. g., the acidified red solutioa of pota^ium permanganate is decolorized by the addition of nitrites. In very dilute solution, the action proceeds according to the following equation :— SNOjH + zMnO^K -f sSO^H, = 5NO3H + SO^K, + zSO^Mn + 3H,0. This reaction serves for the determination of free nitrous acid, as well as for its salts (p. 212). Nitrogen Tetroxide—N,04 or nitrogen dioxide NO, (formerly called hyponitric acid^^ really constitutes two compounds. The former only exists at low temperatures ; when heated, it suffers a gradual decomposition into the simpler molecules NO, which recom- bine to N,04 upon cooling. We here meet the interesting case of dissociation, occurring even at the ordinary temperature. N,©* is colorless, while NO, is colored red-brown ; it appears, therefore, that the color gradually becomes darker as the temperature rises, and that it corresponds to the increasing dissociation of the complex molecules N,04. At ordinary temperatures, nitrogen tetroxide is a yellow liquid of sp. gr. 1.45 at 15°. When cooled to —20° it solidifies to a colorless crystalline mass, melting at —12°. In consequence of the disso¬ ciation that sets in at 0°, the liquid, at first colorless, becomes yellow, and the intensity in color grows with rising temperature. The liquid begins to boil at about 26°, and is converted into a yellowish- brown vapor that becomes dark as the temperature is increased. The theoretical v^por density of N2O4 (molecular weight = 91.86) equals 45-9, while that of NO, (45-9) = 22.9. The experimental vapor dena^ has been found to equal 38 at the point at which the liquid compound boils (26°) ; it may be calculated from this that, at this temperature, 344 per cent, of the N2O4 mole¬ cules are decomposed into NO, molecules. Hence we conclude that the dissoci¬ ation of the compound N2O4 commences already in the liquid state; this is con¬ firmed by the yellow coloration ïq)pearing at 0°. Sulphuric add, as we saw (p. 196), exhibits a similar dissociation in liquid condition. With rising tenq>eratnre the denáty of the vapor steadily diminishes, becomes constant finally at 150** and equals 22.9. Then all the molecules (N2O4) are decomposed into the simpler molecules NO, ; and the dark coloration of the vapors attaims hs maTÎmnm Nitrogen tetroxide is formed by the union of two volâmes of nitric oxide with one volume of oxygen :— 2NO 4- O, = N,04. 3 vols. X voL I roL We can get it more conveniently by heating dry lead nitrate, which decomposes according to the following equation ;— (NO,),Pb = PbO + 0 + 2NO,. The escaping vapors condense in the cooled receiver to liquid N,04. NITROSYLSULPHURIC ACID. 211 The varying molecular composition of nitrogen tetroxide at lower and higher temperatures manifests itself also in its chemical reactions. We saw that by the action of a little cold water, the tetroxide was decomposed into nitrogen trioxide and nitric acid (p. 209). With excess of cold water, and also with an aqueous solution of alkalies, it yields nitric and nitrous acids, that is, their salts NO, + H2O = NO,.OH + NO.OH. NO Both reactions plainly indicate that the liquid tetroxide repre¬ sents the mixed oxide of nitric and nitrous acids ; similarly, the compound CijO^ constitutes the mixed oxide of chloric and chlor¬ ous acids (p. 181). Warm water converts the tetroxide into the dioxide, NOj, which in turn yields nitric acid and nitric oxide:— 3NO2 + HjO = 2HNO3 + NO. The tetroxide and dioxide possess strong oxidizing properties j many substances burn in their vapors ; iodine is set free from the soluble metallic iodides by them. .O.NO. Nitrosylsulphuric Âcid, SO5NH = ^OH. This compound, termed nitrosulphonic acid, is an intermediate product in the manufacture of commercial sulphuric acid (see p. 194), and is quite important. It is employed in the analytical determination of the nitrogen oxides. It is produced by conduct¬ ing nitrogen trioxide and tetroxide into concentrated sulphuric acid;— ^0»\0H + + «'O SO _i_ so _|_ NO OH ^a\OH ^ NOj/^ ~~ ^ 2\0H ^ JNU,.Utt. Nitrogen monoxide—NO—is not absorbed by pure sulphuric acid, but will be if the same contains nitric acid :— O.NO / 3SO4H, + NOjH + 2NO = 3SO, + 2H,0. \ OH Further, the nitrosylsulphuric acid results from the combined action of sulphurous oxide, nitrogen tetroxide (or trioxide), oxygen, and a little water :— 212 INORGANIC CHEMISTRY. O.NO / 2SO2 + + O + HjO = 2SO2 \ OH. It is obtained most readily by conducting sulphur dioxide into strongly cooled anhydrous nitric acid :— O.NO / SO2 + NO3H = SO2 ; \ OH there results a thick magma, which may be dried upon porous earthen plates under the desiccator. Nitrosylsulphuric acid forms a leafy or granular crystalline, color¬ less mass (chamber-acid crystals, p. 195), which fuses about 73° and decomposes into its anhydride, sulphuric acid, and nitrogen trioxide. It deliquesces in moist air, and yields sulphuric and nitrous acids with water:— O.NO OH SO2 +H20 = S03 +NO.OH; \ \ OH OH the nitrous acid decomposes further into nitric acid and nitric oxide. Nitrosylsulphuric acid dissolves in concentrated sulphuric acid without any change; the solution, called niirosaactt/, is also produced in the sulphuñc acid manufacture in the Gay-Lussac tower, is very stable and may be distilled 'without decomposition. When diluted with water it remains unaltered at first, but when the specific gravity of the solution reaches i.55~i*5o ®)> then all the nitit^en oxides escape, especially on warming. When the nitroso acid is pwured into a large quantity of water, the nitrososulphuric acid breaks up into (like the pure acid, see above) sulphuric and nitrous acids, and the latter in part into NOjH and 2NO. Therefore, in titrating nitrous acid with permanganate of potassium (MnO^K. See p. 210), we get the results corresponding to nitrososulphuric acid, if the latter be poured into the permanganate (Lunge). All the nitrogen oxides and acids are separated as nitric oxide (NO) on shaking the nitroso acid with mercury—a procedure serving equally well for estimating tne amount of nitroso acid by means of the nitrometer. All nitrogen oxides are expelled from the nitroso acid by dilution with water and application of heaU (see above). This occurs more readily and completely (even upon concentration to 58® B. = 1.679 specific gravity) by the action of sulphur dioxide :— O.NO OH 2SO2 + 2H2O + SO2 = 2SO2 + 2NO. \ \ OH OH Upon this depends the denitrating action of the Glover tower—see p. 195. NITRIC OXIDE. 213 The anhydride of nitrosulphonic acid—SjNjO, = ^ SO^'o NO' produced upon heating the latter (together with SO^Hj and NjOg—see p. 212) beyond its point of fusion. It is obtained pure by saturating sulphur trioxide with nitric oxide: 380, 4-2NO =0 (SOg.O.NO), + SOg. It is a crystalline, colorless mass, fusing at 217®, and boiling without decomposition about 360®. Much water decomposes it, the same as nitrosylsulphuric acid. The chloranhydride of nitrosulphonic acid, SNO^Cl, is formed by the union of sulphur trioxide with nitrosylchloride : SOj-|-NOCI ^SOjCl (O.NO). It forms white leaflets, is decomposed by heat into its components, and with water breaks up into sulphuric, hydrochloric and nitrous acids. Nitric Oxide—NO. When different metals are dissolved in somewhat dilute nitric acid this oxide is formed, inasmuch as the hydrogen in statu nascendi reduces another portion of the acid. It is most conveniently obtained by pouring dilute nitric acid (specific gravity 1.2) upon copper filings:— 3Cu + 8HNO3 = sCufNOglg + 4HP + 2NO. A better procedure consists in gradually adding concentrated sulphuric acid to a saturated solution of saltpetre, covering thin copper turnings. The action begins in the cold. A colorless gas escapes, which, however, immediately forms brown vapors when it comes in con¬ tact with the air, as it unites with the oxygen of the latter to form NO,. Therefore, all the air must be expelled from the generating vessel hy NO, and the gas collected over water after the interior of the apparatus has become colorless. Nitric oxide is a colorless gas, of specific gravity 15 (H = i) or 1.038 (air = i). Its gas density remains unaltered at —100° (referred to air of the same temperature, V. Meyer). It is con¬ densed with difficulty. Its critical temperature is—93°, and its critical pressure 71 atmospheres (Olszewski). It is slightly soluble in water, but dissolves very readily in an aqueous solution of ferrous salts, imparting a reddish-brown color to the liquid j heat expels it from the same. Nitric oxide is readily soluble in nitric acid. As its solution becomes more con¬ centrated, it assumes a brown, yellow, green or blue color, as nitrogen trioxide is formed finally with anhydrous nitric acid :— 2HNO, -f 4NO = 3N,03 -f H,0. Potassium permanganate oxidizes it, like nitrous acid (p. 210), to nitric acid :— loNO -f fiMnO^K -f- 9SO4H, = 10NO3H -f- 3SO4K, + fiSO^Mn -f- 4H,0. 214 INORGANIC CHEMISTRY. As nitric oxide contains 53.3 per cent, oxygen, it is capable of sustaining the combustion of some substances, but to bring about the previous separation of oxygen from the nitrogen requires ener¬ getic reaction. Hence, phosphorus continues to bum in this gas, while a sulphur flame, developing only a slight heat, is extinguished ignited charcoal does the same, while a splinter, burning energetic¬ ally, will continue to do so, when introduced into the gas. On shak¬ ing a cylinder fllled with NO with a few drops of readily volatile carbon disulphide, and bringing a flame to the mouth of the vessel, the carbon disulphide vapors will quietly burn in the gas, giving a bright luminous flame, emitting strong actinic rays ; in this com¬ bustion, the C and S of the CS2 unite with the oxygen of the nitric oxide. Nitric oxide is a strongly endothermic compound (see p. 216), and is conse¬ quently explosive (p. 29). It maybe caused to explode by inflaming a little ful¬ minating mercury in it (Berthelot). On determining the quantity of heat disengaged in the combustion of phos* phorus, carbon or other substances in NO gas, it will be discovered that the same is greater (about 21,600 calories), than that which is developed by the combus¬ tion of these bodies in oxygen. This can only be explained upon the theory that less heat is necessary for the separation of NO into N and O than for the separa¬ tion of the molecules of combined oxygen atoms—an additional proof that the molecules of free oxygen (as of other elements) consist of atoms (p. 116). With oxygen or air, NO at once forms brown vapors of nitro¬ gen dioxide :— 2NO + 03 = 2NO,. 2 vols. X vol. 2 vols. With less oxygen, nitrogen trioxide is produced (p. 208). NO also combines with chlorine to nitrosylchloride, NOCI (p. 208), and the compound NOCI2, which is, as yet, little investigated ; with bromine, it yields NOBrg. At a red heat NO becomes NO, and N. With hydrogen and moderate heat it forms water and nitrogen : NO -j- H, = N + HjO ; a mixture of both gases burns with a white flame. On con¬ ducting NO and H together over platinum sponge, water and ammonia are produced ;— 2NO -f 5H2 = 2NH3 -f 2H3O. The volumetric analysis of ni¬ tric oxide gas may be easily executed as follows : Fill a bent glass tube (Fig. 85) over mercury with NO gas ; introduce into the same a piece of sodium NITROUS OXIDE. 215 and heat the latter with a lamp. The sodium combines with the oxygen, and free nitrogen separates ; the volume of the latter always equals half the volume of the nitric oxide gas employed ; this follows from the formula NO :— 2N0 = N, + O2. 3 vols. I vol. X vol. The molecular formula of the oxide is NO = 29.27 (even at —too®, see p. 210), as its vapor density is 14.98 (H = l). NO, NO, and chlorine dioxide, CIO, (p. 180), present an apparent anomaly as regards the common laws r^u- lating the valence of the elements. Ordinarily, the valence changes from an odd number to an odd, and from an even to an even number (p. 174). Nitrogen is usually pentavalent and trivalent ; in the cited compounds it appears di- and tetravalent. This abnormal behavior of N finds a partial explanation in the position it occupies in the periodic system of the elements. Nitrous Oxide—Hyponitrous Oxide—N,0—is formed when zinc or tin acts upon dilute nitric acid of sp. gr. i.i. It may be best obtained by heating ammonium nitrate, which at about 170® breaks up directly into water and nitrous oxide :— NH^NOg = N2O + 2H2O. This compound is a colorless gas, of sweetish taste and slight odor. Its density is 22.04 (H = 1), or 1.52 (air = i), correspond¬ ing to the molecular formula NgO = 44.08. In cold water it is tolerably soluble (i volume HjO dissolves at 0° 1.305 volumes N2O) ; therefore it must be collected over warm water or mercury. Cooled to —88°, or under a pressure of 30 atmospheres at 0°, it condenses to a colorless liquid of specific gravity 0.937. By evaporation of the liquid in the air its temperature diminishes to —100°, and it solidifies to a crystalline, snowy mass. If the äqueous nitrous oxide be evaporated under an air pump its temper¬ ature falls to —140° ; the lowest which has been attained. Although this oxide contains less oxygen than nitric oxide, it supports the combustion of many bodies more readily than the latter, because it is more easily decomposed into oxygen and nitrogen. A glimmer¬ ing chip inflames in it, as in oxygen ; phosphorus burns with a bright, luminous flame, while a sulphur flame is extinguished. The liquid nitrous oxide behaves like the gas ; an ignited coal thrown on its surface burns with a bright light. A mixture of equal volumes of nitrous oxide and hydrogen explodes like detonating gas, only less violently:— N,0 H, = N, + H,0. I vol. X vol. X vol. 216 INORGANIC CHEMISTRY. It resembles oxygen very much, but may be distinguished from it by not producing brown vapors with nitric oxide, as does the former. It is not capable of combining with oxygen. When it is conducted through a red-hot tube it is decomposed into nitrogen and oxygen. It has an exhilarating effect when inhaled in slight quantity, and is, therefore, termed laughing gas. Its volume composition may be determined in the same manner as with nitric oxide, viz.: by heating a definite volume of the gas with potassium. Then we learn that from a volume of NjO an equal volume of nitrogen will be separated—corresponding to the molec¬ ular formula :— N^O + K2 = N3 + K^O. I vol. I vol. Hyponitrous Acid—NOH. As nitrous and nitric acids correspond to nitro¬ gen trioxide and pentoxide, so hyponitrous oxide may be regarded as the anhy¬ dride of the recently discovered hyponitrous acid ;— N2O + H2O = 2NOH, although the latter is not formed by the hydration of N2O. On the contrary, it yields nitrous oxide when water is withdrawn from it by sulphuric acid. An aqueous solution of the acid may be obtained by decomposing the silver salt, AgNO, with hydrochloric acid. Dissolved in water it is colorless, reacts strongly acid and is tolerably stable. It liberates iodine from potassium iodide and reduces a permanganate solution. The silver salt is again precipitated by silver nitrate. The alkali salts of hyponitrous acid result from the action of sodium amalgam upon the nitrates or nitrites, or of freshly precipitated ferrous hydroxide upon sodium nitrite, and by the electrolysis of the nitrites. On neutralizing the solu¬ tion with acetic or nitric acid, and adding silver nitrate, the silver salt is obtained as a light-yellow amorphous powder, which decomposes gradually at lOO®, but with explosive violence at 110®. It is soluble in dilute sulphuric acid, and is reprecipitated upon neutralization with ammonium hydroxide. Concentrated sulphuric acid liberates NgO from it and the free acid. The latest researches make the composition of hyponitrous acid rather doubtful, since either the formula NgOjAg or N^OgAg^ may be ascribed to the silver salt The thermo-chemical relations of the oxygen derivatives of nitrogen give some clue to their chemical deportment. All nitrogen oxides are endothermic com¬ pounds, í.í., they are produced from their elements with heat absorption (com¬ pare p. 185) corresponding to the symbols:— (N2O)—— 18.3 (N,0)= —21.5 (N203)=—23.0 (N,02)= —2.0 (N2,05— gas) =—12.0. From this we observe that the oxides of nitrogen cannot be prepared from the elements without addition of energy, and their direct production has never been noticed. Proceeding from nitric oxide (NO), we see from the above numbeis that the formation of the higher oxides from it occurs with heat disengagement :— (2N0,0)=2O.I (N0,0) = I9.5 (2N02,01 = 2.8, OXYGEN COMPOUNDS OF PHOSPHORUS. 217 whereas heat is absorbed in the conversion of nitrous into nitric oxide : (NjOjO) = — 24.8. Hence nitric oxide combines readily with oxygen to yield NjOj and NOj, while nitrous oxide is not reactive. Conversely, the higher oxides are easily reduced to nitric oxide, and this explains their great ability to oxidize. The heat absorption of nitric oxide (N,0 = — 21.5) indicates, therefore, that the affinity of N to O is so much less than that of the oxygen atoms in the molecule (p. 214). Heat disengagement, on the contrary, occurs in the production of nitric acid from its elements :— (N,03,H —liquid) = 41.5 (N,03,H,Aq.) =49.0. This explains the relative stability of that acid, and the possibility of its pro¬ duction under different conditions. 3. OXYGEN COMPOUNDS OF PHOSPHORUS. PO3H3 Hypophosphorous acid. P2O3 PO3H3 Phosphorus Phosphorous trioxide. acid. P2O5 P0,H3 Phosphorus Orthophosphoric pentoxide. acid. The following anhydride acids (compare p. 204) are derived from orthophosphoric acid:— HPO3 — Metaphosphoric acid. H4P2O2 — Pyrophosphoric acid. The structure of these compounds is expressed by the following formulas :— V V V /OH H2PO —OH HP0CX„ PO—OH \0H Hypophosphorous Phosphorous Phosphoric acid. acid. acid. In hypophosphorous acid two atoms of hydrogen are directly combined with pentavalent phosphorus, while the third atom forms an hydroxy 1 group with oxygen. The latter is easily replaced by the action of bases, and, therefore, hypophosphorous acid is a mono¬ basic acid. Phosphorous acid contains one atom of H united to F and two hydroxyl groups j therefore, it is dibasic. Finally, phos¬ phoric acid has three hydroxyl groups, and forms three series of salts. By the elimination of one molecule of HjO from H5PO4, metaphosphoric acid results—an anhydride, which, at the same time, is a monobasic acid, as it contains one hydroxyl group ;— V PO3 — OH — Metaphosphoric acid. 19 218 INORGANIC CHEMISTRY. On removing one molecule of H2O from two molecules of H3PO4, pyro- or diphosphoric acid is formed (see p. 197) :— V /OH /OH PO—OH PO—OH \OH Q \Q V /OH "2^ — PO—OH PO—OH \OH \OH 2 Molecules Phosphoric acid. z Molecule Diphosphoric acid. Pyrophosphoric acid contains 4 hydroxyl groups, hence is tetra- basic. Finally, if from two molecules of phosphorous acid or phosphoric acid, all the H atoms be removed, in the form of water, two perfect anhydrides are obtained :— III in V V OP —O —PO and OjP —O —POj Phosphorous Phosphoric anhydride. anhydride. The salts of phosphoric acid are termed phosphates ; those of phosphorous acid, phosphites^ and of hypophosphorous acid, hypo- phosphites. Hypophosphorous Acid—HaPOj. Hydrogen phosphide es¬ capes when a concentrated solution of sodium or barium hydroxide is warmed with yellow phosphorus, leaving behind in solution a salt of hypophosphorous acid :— 4P + sNaOH 4- 3H2O = sH^PO.ONa +PH3. The free acid may be separated from the barium salt by means of sulphuric acid; the insoluble barium sulphate being filtered off from the aqueous solution of the acid, and the latter concentrated under the air-pump. Hypophosphorous acid is a colorless, thick liquid, with a strong acid reaction. Below o® it sometimes solidifies to large white leaflets, which fuse at Heat converts it, with much foaming, into hydrogen phosphide and phosphoric £LCld I 2P02H3 = PH3 +PO4H,. It absorbs oxygen readily, becoming phosphoric acid, hence acts as a powerful reducing agent. It reduces sulphuric acid to sulphur dioxide, and even to sulphur. It precipitates many of the metals from their solutions ; from copper sulphate it separates the hydride -~--CUjH3. PHOSPHOROUS OXIDE. 219 The acid is monobasic, HjPO.OH. Its salts dissolve readily in water, and absorb oxygen from the air, thus becoming phosphates. When heated in a dry condition, they set free the hydride of P, and are converted into pyrophosphates ; some also yield metallic phosphides. Phosphorous Acid—H3PO3—is formed at the same time with phosphoric acid in the slow oxidation of P in the air. The decom¬ position of the trichloride by water gives it more conveniently :— PCI3 + 3H2O = PO3H3 + 3HCI. By evaporating this solution under the air-pump the phosphorous acid becomes crystalline. The crystals are readily soluble in water, and deliquesce in the air. It fuses at 70°, and decomposes on further heating into PH3 and phosphoric acid : — 4P03H3 = PH3 +3P0,H3. In the air the acid absorbs oxygen, and changes to phosphoric acid. Hence, it is a strong reducing agent, and precipitates the free metals from many of their solutions. In the presence of water the halogens oxidize it to phosphoric acid. It is a dibasic acid, forming two series of salts, in which i and 2 atoms* of H are replaced by metals. In the air, the phosphites do not oxidize, except under the influence of oxidizing agents. When heated, they generally decompose into hydrogen and pyrophos¬ phates. Phosphorous Oxide—P2O3 or PiOg—is formed when phosphorus is burnt in a rapid current of air. It consists of feathery crystals, which melt at 22.5° to a clear, colorless, mobile liquid. It boils without decomposition at 173.1° in an atmosphere of nitrogen or carbon dioxide. Its density equals 7.7, corresponding to the for¬ mula P^Og. In this respect it is analogous to arsenious and antimo- nious oxides. It is stable up to about 200°. On long standing with water it slowly dissolves, and the resulting solution yields all the reactions characteristic of phosphorous acid (/r. Chem. Soc., 57. 545)- Phosphoric Acid—PO4H3—or Orthophosphoric acid, is pro¬ duced when the pentoxide is dissolved in hot water, and by the decomposition of the penta- or oxy-chloride (POCI3) by water (see * Therefore, the structural formula, HP0(0H)2 is assigned to this acid. There appears to exist another phosphorous acid, at least in compounds, to which the formula P(OH)j belongs. 220 INORGANIC CHEMISTRY. p. 142). It may be obtained by decomposing bone ash, (P04),Cas, with sulphuric acid, or, better, by oxidizing yellow phosphorus with nitric acid. The aqueous solution is evaporated to dryness in a platinum dish. The anhydrous acid consists of colorless, hard, pris¬ matic crystals, which in the air deliquesce to a thick, acid liquid. Phosphoric acid is tribasic, forming three series of salts, called acid (PO4H2K), neutral (PO4HK2), and basic (PO4K3). As this desig¬ nation does not entirely correspond with the behavior of the salts to litmus, it is more rational to term them primary^ secondary, and tertiary: or to speak of them according to the number of hydrogen atoms replaced by metals, as,, e. g., monopotassium phosphate (H2KPO4), dipotassium (K2HPO4) phosphate, and tripotassium (K3PO4) phosphate. The tertiary phosphates, excepting the salts of the alkalies, aye insoluble in water. With a silver nitrate (AgN03) solution,soluble phosphates give a yellow precipitate of tri-silver phosphate, P04Agj. Pyrophosphoric Acid—H4P2O7—(structure, p. 218) is formed by the continuous heating of orthophosphoric acid to 200°—300®, until a portion of it dissolved in ammonium hydroxide does not yield a yellow but pure white precipitate with silver nitrate. The sodium salt is easily obtained by heating di-sodium phosphate:— The acid presents a white crystalline appearance, and is readily soluble in water. When in solution, it slowly takes up water at ordinary temperatures, more rapidly when heated, and, like all anhydrides, passes into the corresponding acid—orthophosphpric acid. It is tetrabasic. Its salts are very stable, and are not altered by boiling with water; warmed with acids, they become salts of the ortho-acid. The soluble salts yield a white precipitate, Ag4P,0„ with silver nitrate. Hypophosphoric Acid—P20eH^.—While pyrophosphoric acid is an anhy¬ dride acid of phosphoric acid (p. 217), the so-called hypophosphoric may be viewed as a mixed anhydride of phosphoric and symmetrical phosphorous acids It is produced together with phosphorous and phosphoric acids, by the slow oxidation of moist phosphorus in the air. It is separated from these acids by means of its difficultly soluble sodium salt, PjOgNajHj 6HjO; by precipitating the solution of the latter with a soluble lead salt we get insoluble PjO^Pb,. Its silver salt is more easily obtained by oxidizing phosphorus in the presence of silver nitrate. The free acid from the lead or silver salt is rather stable in a dilute solution, and below 30° may be concentrated to a~ syrup. At higher temperatures. aNajHPO^ = Na^PjO, 4- H2O. PO(OH), >° P(0H)2. METAPHOSPHORIC ACID. 221 itiore readily in the presence of hydrochloric or sulphuric acid, the acid decom¬ poses into phosphoric and phosphorous acids. It is not a reducing agent, but is oxidized by potassium permanganate to phosphoric acid. Metaphosphoric Acid—HPO3 or POjOH—results upon heat¬ ing the ortho- or pyro-acid to 400°. It can be more conveniently obtained by dissolving the pentoxide in cold water:— PjOs 4- HjO = 2HPO3. It is a glassy, transparent mass {Acidum phosphoricum glaciale)y which fuses on heating, and volatilizes at higher temperatures, without suffering any change. It deliquesces in the air, and dis¬ solves with ease in water. (The commercial glassy phosphoric acid Fig. 86. contains sodium and magnesium phosphate, and dissolves with difficulty in water). The solution coagulates albumen; this is a characteristic method of distinguishing the meta- from the ortho- and pyro-acids. In aqueous solution, the acid changes gradually, by boiling rapidly, into the ortho-acid:— HPO3 + HjO = H3PO4. It is a monobasic acid. Its salts, the metaphosphates, are readily obtained by the ignition of the primary salts of the ortho-acid :■— NaHjPO^ = NaPO, -f- H^O. When the aqueous solutions of these salts are boiled, they are 222 INORGANIC CHEMISTRY. converted into the ortho-primary salts. With silver nitrate the soluble metaphosphates give a white precipitate, AgPO,. In addition to the ordinary salts of metaphosphoric acid, various modifications of the same exist ; these are derived from the polymeric meta-acids, HjPjOg, H3P3O,, H4P4O12, etc. These acids arise from the corresponding polyphosphoric adds, which are obtained by the union of n molecules of the ortho-acid, with the separation of « — i molecules of water (p. 197), just as the meta-acid is formed from the ortho-. They are all changed to primary ortho-phosphates by boiling their solutions. Phosphorus Pentoxide—P2O5, or Phosphoric anhydride—^is obtained by burning phosphorus in a current of oxygen or dry air. The following procedure serves for the preparation of it (Fig. 86) : A piece of P, placed in an iron dish attached to the glass tube a b, is burned in the glass balloon A. The necessary amount of air is drawn through the vessel by means of an aspirator. It is first passed through the bent tube containing pieces of pumice-stone, moistened with sulphuric acid, in order to dry it perfectly. After the phosphorus has been consumed, fresh pieces of it are introduced into the little dish through a b, and the upper end of the tube closed with a cork. The PjOj formed collects partly in A and partly in the receiver. Phosphorus pentoxide is a white, voluminous, flocculent mass. It attracts moisture energetically and deliquesces in the air. It dissolves in cold water with hissing and yields metaphosphoric acid. Owing to its great affinity for water it serves as an agent for drying gases, and also for the withdrawal of water from many substances. Chlor-Anhydrides of the Acids of Phosphorus.—The halogen derivatives of P, considered on page 140, may be viewed as the halc^en anhydrides of phos¬ phorous and phosphoric acids (p. 199).. The compounds PCI3, PBr,, and PI,, are derived from phosphorous acid, because they yield the latter acid with water :— PCI3 -f 3H2O =3 H3PO3 -}- 3HCI. The compounds POCl,, POBr,, are the halogen anhydrides of phosphoric acid :— poa, + 3H3O = po(oh)3 + 3HCI; while PCI5 and PB^ correspond to the normal hydroxide, PfOH),, which has not been obtained in a free condition. The compound PSCl, is analogous to the oxychloride POCl,. It is obtained by the action of PCI5 upon hydrogen sulphide and some metallic sulphides :— POj -f- HjS = PSCI3 -f 2Ha. This reaction is very similar to that occurring in the formation of phosphorus oxychloride. Phosphorus sulphochloride—VSC\^—is a colorless liquid, fumii^ in the air and boiling at 124®. Water decomposes it into phosphoric and hydro¬ chloric acids and hydrogen sulphide. OXYGEN DERIVATIVES OF ARSENIC. 223 COMPOUNDS OF PHOSPHORUS WITH SULPHUR. With sulphur, phosphorus affords a number of compounds which are obtained by direct fusion of P with S. As the union of ordi¬ nary P with S usually occurs with violent explosion, red phospho¬ rus should be employed in preparing these compounds. The compounds PjSg and P^Sj, analogous in constitution to PjOj and PjOs, are solid crystalline substances, melting at higher tem¬ peratures and subliming without decomposition ; P2S5 boils at 518°. Water changes them to hydrogen sulphide and the corresponding acids, phosphorous and phosphoric. They combine with metallic sulphides to form compounds (e.g., PS4K3) which possess a constitu¬ tion analogous to that of the salts of phosphoric acid (see sulpho- salts of arsenic). At ordinary temperatures, PjS and P^S are liquids, which inflame readily in the air. Besides the preceding, we have other phosphorus derivatives which contain N. These have been little studied, and at present offer little interest. Such com¬ pounds are PNjH fphospham), PNO, PNClj. The so-called amid derivatives, POCljNHj, POClfNHjlj and POfNHj), are produced by allowing ammonia to act upon POCl,. In these chlorine is replaced by the amido-group NH,. 3. OXYGEN DERIVATIVES OP ARSENIC. ASjO, ASO3H3 Arsenic trioxide. Arsenious acid. ASjOj AsO^TIj Arsenic pentoxide. Arsenic acid. Arsenic Trioxide, AsjOj, or Arsenious anhydride, occurs in nature as arsenic "bloom." It is produced by the burning of arsenic in oxygen or in the air, and by the oxidation of the metal with dilute nitric acid. It is obtained metallurgically on a large scale as a by-product in the roasting of ores containing arsenic. The trioxide thus formed volatilizes and is collected in walled cham¬ bers, in which it condenses in the form of a white powder (white arsenic, poison flour). To render it pure, it is again sublimed in iron cylinders, and obtained in the form of a transparent, amorphous glassy mass (arsenic glass), the specific gravity of which equals 3.78. Upon preservation, this variety gradually becomes non-transparent and porcelaneous, acquires a crystalline structure, and its specific gravity decreases to 3.69. Upon dissolving this oxide in hot hydro¬ chloric acid, it crystallizes, on cooling, in shining, regular octa- hedra. At the same time the interesting phenomenon is observed, that when the solution of the glassy variety crystallizes it phospho- 224 INORGANIC CHEMISTRY. resces strongly in the dark, while the porcelaneous does not exhibit this property. Arsenic trioxide crystallizes in similar forms of the regular system when its vapors are rapidly cooled, but upon cool¬ ing slowly, it assumes the shape of rhombic prisms ; therefore, it is dimorphous. When heated in the air, it sublimes above 218°, without fusing ; upon higher pressure, however (in sealed tubes), it melts to a liquid which solidifies to a glassy mass. The vapors of AsjO, have the vapor density 198 (H = l), which con¬ tinues constant to 1560°. Corresponding to formula AsjO, (= 198), the vapor density should be -î-fA = 99. The vapor density determined experimentally is just twice as great, therefore the gas molecules of the trioxide possess the double formula As^Oj. We have already noticed that the molecule of free arsenic con¬ sists of four atoms (As^) : this complex arsenic group, consequently, is retained ih the trioxide ; while in arsine (AsH,) and arsenious chloride (AsClj) the molecules contain but I atom of arsenic. The trioxide dissolves with difficulty in water ; the solution pos¬ sesses a sweetish, unpleasant metallic taste, exhibits but feeble acid reaction, and is extremely poisonous. The oxide is very soluble in acids, and probably forms salts with them ; at least, on boiling a solution of AsjÖj in strong hydrochloric acid, arsenious chloride, AsCl,, volatilizes. From this and its feeble acid nature we perceive an indication of the basic character of the trioxide corresponding to the already partially metallic nature of arsenic (see p. 150). Nascent hydrogen converts the trioxide into arsine (AsHj) ; but when heated with charcoal it is reduced to the metallic state. Upon heating AsjOg in a narrow glass tube with C, the reduced arsenic deposits as a metallic mirror on the sides. Oxidizing agents con¬ vert it into arsenic acid. Arsenious Acid—H3ASO3, corresponding to AsjOj, is not known in a free condition. It probably exists in the aqueous solu¬ tion, but the anhydride separates out upon evaporation. In its Salts {arsenites) it is tribasic and usually affords tertiary deriva¬ tives ;— AggAsOg, Mg3(As08),. The alkali salts, soluble in water, absorb oxygen from the air and serve as powerful reducing agents, they themselves becoming arsen- iates. Other salts exist which are derived from the meta-arsenious acid, HAsOj (p. 205). Arsenic Acid—HsAs04—is obtained by the oxidation of arsenic or its trioxide with concentrated nitric acid or by means of chlorine. Upon evaporating the solution rhombic crystals of the formula H3ASO4-f~ j^HjO separate out; these deliquesce on ex¬ posure. They melt at 100°, lose their water of crystallization and ARSENIC WITH SULPHUR. 225 yield orthoarsenic acid HjAsO*, which heated to 140-180® passes into pyroarsenic acid—H^AsjOt :— 2H3ASO4 = AsjO^H^ -f- HjO. At 200° this again loses water and becomes Meta-arsenic acid— HAsOs- With water the last two acids become ortho- again ; hence the latter is perfectly analogous to phosphoric acid. At a red heat the meta-arsenic acid loses all its water and becomes Arsenic Pentoxide—AS2O5, a white, glassy mass. Very strong igni¬ tion breaks this up into AsjOs and Oj ; in contact with water it gradually changes to arsenic acid. Orthoarsenic acid is readily soluble, and is a strong tribasic acid. Its salts—the arseniates—are very similar to the phosphates and are isomorphous with them. With the soluble salts silver nitrate gives a reddish-brown precipitate of tri-silver-arseniate, AgaAsO^. COMPOUNDS OF ARSENIC WITH SULPHUR. Like phosphorus, arsenic, upon fusion with sulphur, yields several compounds. The metallic nature of arsenic is seen in these deriva¬ tives, because they, açcording to the common method of forming the metallic sulphides, can be obtained by the action of hydrogen sulphide upon the oxygen derivatives of arsenic ;— AS2O3 +3H2S = AsjSj + 3H2O. Arsenic Trisulphide—AsjSj—is precipitated from solutions of arsenious acid or its salts by hydrogen sulphide, as a lemon-yel¬ low amorphous powder. It may also be obtained from arsenic acid solutions, but then it contains admixed S, as the acid is first reduced to arsenious acid and then precipitated :— AsjOj -|- 2H2S = AsjOg -f- 2Hj,0 -f- 2S. This compound is most readily prepared by fusing AsjOs with sulphur. It occurs as auHpigment in nature in the form of a bril¬ liant, leafy, crystalline mass of gold-yellow color, and the specific gravity 3.4. On fusing artificially prepared arsenic trisulphide it solidifies to a similar yellow mass, the specific gravity of which equals, however, 2.7. In water and acids the trisulphide is insolu¬ ble, but dissolves readily in ammonium hydroxide and the alkalies. Arsenic Pentasulphide—^AsjSs—separates as a bright yellow powder from the solution of sodium sulpharseniate, NajAsS* (see below), upon the addition of acids. The Arsenic Disulphide—AsjSa—also exists. It occurs in nature as Realgar^ forming beautiful, ruby-red crystals, of specific 226 INORGANIC CHEMISTRY. gravity 3.5. It is applied as a pigment. It is prepared artificially by fusing As with S. Arsenic Sulpho-Salts.—Owing to the similarity of sulphur to oxygen we may anticipate for arsenic (as also for other elements) the existence of sulphur acids corresponding to the oxygen acids, sulpharsenious acid, HjAsI^ and sulpharsenic acid, HjAsS^. However, these acids are unknown in a free stat^ although their salts, known as sulphur- or sulpho-salis, are found, and they corre¬ spond perfectly with the oxygen salts. Just as the latter arise by the union of metallic oxides with acid oxides, so the sulpho-salts are formed by the combination of alkaline sulphides with those sulphur derivatives corresponding to the add oxides (acid sulphides) :— AsjS, -f- 3K2S = 2K3AsSj Tripotassium sulpharsenite. AsjSg + 3K2S = aKgAsS^ Tripotassium sulpharseniate. For the preparation of these sulpho-salts, arsenic sulphide is dissolved in the aqueous solution of potassium or sodium sulphide, or hydrc^en sulphide is con¬ ducted through the alkaline solution of the oxygen salts :— KjAsO^ 4H2S = KgAsS^ + 4H2O. The sulpho-salts 01 the alkalies and ammonium are easily soluble in water, and when the solution is evaporated they generally separate in crystals. Acids decom¬ pose them, arsenic sulphide separating out and hydrogen sulpMde becoming free ;— 2K3ASS4 + 6HC1 = AS2S5 + 6KC1 + 3H2S. Antimony, carbon, tin, gold, platinum and some other metals form sulpho-salts similar to those of arsenic (and of phosphorus). 4. OXYGEN COMPOUNDS OF ANTIMONY. The oxygen derivatives of antimony are analogous in constitution to those of arsenic: SbjO., and SbjOs. The metallic nature of antimony, which we observed appearing in the halogen derivatives, shows itself quite distinctly here. The lowest oxygen compound does not possess acid, but basic properties almost solely ; it forms salts with acids only, hence is called Antimony oxide. The normal hydrate HjSbOg, corresponding to arsenious acid, HjAsOs, is not known. A hydrate, SbOjH or SbO.OH, analogous to meta-arseni- ous acid, does exist ; it deports itself like a base. The higher oxidation product, the pentoxide, Sb^Og, on the contrary, has an acid nature and yields salts with the bases. The hydrate, Sb04H3, or ortho-antimonic acid, and its salts, have not been obtained. The known salts are derived from pyro-antimonic acid, H^SbjOï, and meta-antimonic acid, HSbO,; these exist in a free condition. Antimony Oxide—SbjOa or Sb^O«—is obtained by burning ANTIMONIO ACID. 227 the metal in the air, or by oxidizing it with dilute HNOs- By sublimation it may be obtained in two different crystal systems, in regular octahedra and in rhombic prisms. Arsenic trioxide also crystallizes in the same forms; therefore the two compounds are isomorphous. On adding sodium carbonate to the solution of the trichloride a white precipitate of antimony hydrate or antimonious acid, HSbOj, separates out :— 2Sba3 + sNajCOj + HjO = 2SbO.OH + 6NaCl + 3CO3. The hydrate is changed to oxide by boiling. The latter and the hydrate dissolve in sodium and potassium hydroxide, and, very probably, form salts (NaSbOj) which decompose upon evaporating the solution. In this behavior the acid nature of antimony hydrate is also seen; therefore it has received the name of antimonious acid. The oxide forms salts with acids, which are derived either from the normal hydrate, HgSbOa, or from the hydrate, HSbO, = SbO- OH. In the salts of the first kind we have 3 hydrogen atoms of the hydrate replaced by acid radicals, or, what is the same, a triva¬ lent antimony atom displacing 3 hydrogen atoms of the acids :— SbOaiNO^), or (N03)3Sb. Antimony nitrate. In the second variety of antimony salts derived from the hydrate, SbO.OH, hydrogen is replaced by a monovalent acid residue, or the hydrogen of the acid is substituted by the monovalent group, SbO, known as antimony I :— SbO.O.NOj or NOg.SbO. Antimonyl nitrate. Of these salts may be mentioned the following :— Antimony Sulphate—(804)38^—which separates when a solution of the oxide in sulphuric acid is cooled. Antimonyl Sulphate—80^(8b0)2—is formed when antimony oxide is dis¬ solved in somewhat dilute sulphuric acid, and crystallizes in fine needles on cooling. Water decomposes both, forming basic salts; hence the basic nature of antimony oxide is slight. Antimonic Acid—HSbOs—or, more correctly, Meia-antimonic acidy is obtained upon warming antimony with concentrated nitric acid, and is a white powder, almost insoluble in water and in nitric acid, but reddens blue litmus paper. It is a weak monobasic acid, the salts of which are mostly insoluble. If antimony pentachloride be mixed with much water, Pyroan- timonic Acid, H4Sb20T, separates as a white powder. Its salts are produced by fusing antimonic acid or meta-antimoniates with potassium or sodium hydroxides :— 228 INORGANIC CHEMISTRY. 2KSb03 4- 2KOH = K^SbjO, + HjO. Hydrochloric acid precipitates pyroantimonic acid from the solu¬ tions of these salts. By gentle ignition the meta- and pyro-acids yield Antimony Pentoxide, SbjOj, a yellow, amorphous mass, soluble in hydro¬ chloric acid. By heating the oxygen compounds for some time with air access they are converted into the oxide, SbjOi, which can be viewed as antimonyl antimoniate (SbOj.SbO), or as a mixed anhydride | O. It is a white powder, becoming yellow when heated, and is non-volatile. COMPOUNDS OF ANTIMONY WITH SULPHUR. These are perfectly analogous to the S compounds of arsenic, and form sulphosalts with alkaline sulphides, corresponding to the oxygen salts. Acids precipitate antimony sulphide from the sul¬ phosalts. Antimony Trisulphide—SbgSs—is found in nature as stibnite, in radiating crystalline masses of dark-gray color and metallic lustre; specific gravity = 4.7. When heated it melts and sublimes. The artificial sulphide obtained by precipitating a solution of the oxide with hydrogen sulphide, is an amorphous red powder. When fused, it solidifies to a mass exactly like stibnite. The sulphide dissolves in concentrated HCl, upon warming, to form antimony trichloride. The compound, SbgSgO, occurring in nature as red stibnite, can be artificially prepared, and serves as a beautiful red color, under the name of antimony cinnabar. Kermes minerahy employed in medicine, is obtained by boiling antimony sulphide with a sodium carbonate solution, and is a mixture of SbjSj and SbgOg. Antimony Pentasulphide—SbjSj—or gold sulphur (julphuf^ auratuni) is precipitated by HgS from acid solutions of antimonio acid; it is more conveniently obtained by the precipitation of sodium sulphantimoniate, NajSbS*, with hydrochloric acid ;— aNajSbS^ + 6HC1 = Shß^ + ÓNaQ + sHjS. It is an orange-red powder, like the trisulphide ; it decomposes on being heated into SbgSj and Sg. It dissolves to antimony tri- chloride in strong hydrochloric acid, with separation of sulphur and hydrogen sulphide. Sodium Sulphantimoniate—NasSbS4 (Schlippe's salt), results from boiling pulverized SbgSg with sulphur and sodium hydroxide. Upon concentrating the solution it crystallizes in large yellow tetrahedra, VANADIUM—NIOBIUM—TANTALUM. 229 containing 9 molecules of HjO (SbSiNaj -f- çHjO) > exposed to the air it becomes covered with a brown layer of SbïSs. It serves principally for the preparation of the officinal gold sulphur. The affinity of the elements of the nitrogen group for hydrogen diminishes with increase of atomic weight, and corresponds to the addition of metallic character, while the affinity for chlorine, concluded from the thermo-chemical relations, generally increases (compare p. 150). However, the heat disengagement in the formation of ASCI3 is somewhat less than that in the case of PCI3, which would afford a partial explanation for the non-existence of the compounds AsX 5 (see p. 146). The slight affinity of arsenic is seen more distinctly in the oxygen com¬ pounds, because, as in the case of the halogen and oxygen groups (pp. 186 and 204), the arsenic corresponding to bromine and selenium— Br = 79.96 Se = 79.1 As = 75, shows a less heat disengagement in the formation of its compounds:— (N,04,H3,Aq.) = 117.4 (N3,05,Aq.) = 29.8 (P,0^,H3,Aq.) =305.3 (PaA) =363-8 (As,04,H3,Aq.) = 215.2 (As3,05) = 219.4 Phosphoric acid is, therefore, more stable and more energetic than nitric and arsenic acids; nitric acid oxidizes phosphorus, and arsenic to phosphoric and arsenic acids. The latter acid is readily reduced to arsenious acid. VANADIUM. NIOBIUM. TANTALUM. Vd = 51.2 Nb = 94.2 Ta = 183 The three rare elements, vanadium, niobium and tantalum, are closely related to the Phosphorus group. They yield derivatives very much like those of P; but possess a more metallic character, and do not combine with hydrogen. They exhibit many characteristics similar to those of chromium, iron and tungsten, with which they are frequently associated in their naturally occurring compounds (compare the Periodic System of the elements). Vanadium occurs in nature principally in the form of salts of vanadic acid (vanadium lead ore) and in some iron ores. It may be obtained free by igniting its chlorides in a current of hydrogen. It is a grayish-white, metallic, lustrous powder, of specific gravity 5.5. It is difficultly fusible, and does not change in the air. When heated, it bums to VdjOj. Vanadium Trichloride—^VdClj—forms red plates, which readily deliquesce in the air; it is not volatile. Vanadium Oxychloride—VdOCl3—results on heating a mixture of VdjOj and C in chlorine gas. It is a lemon-yellow liquid, of specific gravity 1.84, and boils at 120®. It fumes strongly in the air and decomposes with water (analogous to phosphoras oxychloride) into vanadic acid and hydrochloric acid. Its vapor density equals 86 (VdOClg = 173.2). Vanadium Oxide—VdjOg—is a black powder obtained by heating VdjOg in hydrogen. It combines with O, to form Vd^Og. Vanadium Pentoxide—Vd^Oj—or vanadic anhydride, is a brown mass obtained by fusing the naturally occurring vanadates with nitre, etc. It is soluble in the 230 INORGANIC CHEMISTRY. alkalies, and forms salts of Vanadic, HjVdO^, and Metavanadie acids, HVdO,, with the metals. All these compounds are similar in constitution to those of P. In addition to these, vanadium forms other compounds, constituted like those of sulphur and chromium. In this class belong VdClj (dichloride),the tetrachloride, VdCl^, vanadious oxide, VdO, vanadium dioxide, VdOj, and VdOCl,. The tetra¬ chloride, VdCl^, is a red-brown liquid, boiling at 154° ; its vapor density equals 96 (VdCl^ = 191.6). Niobium and tantalum are not known in a free state. The chlorides, NbClg and TaClj, are volatile and are decomposed by water. Niobium and tantalum unite with potassium fluoride, forming double salts, e. g-., 2KFI. NbFlj and aKFLTaFlj; also 2KFl.NbOFl3 and 2KFl.TaOFl3. When pot^ sium niobium fluoride, 2KFi.NbFlg, is heated with sodium, niobium hydride is formed. This is a grayish-black powder. If it is heated it burns to niobic anhydride, NbjOg, and water. The oxides, Nb^Oj, and Ta^Oj, form salts of niobic (HgNbO^) and tantalic (HjTaO^) acids with bases. OXYGEN DERIVATIVES OF THE ELEMENTS OF THE CARBON GROUP. Because of the analogy with the hydroxyl derivatives of the ele¬ ments of the first three groups, CIO3.OH, S02(0H).3, PO(OH)s, we may assume the existence of the following normal hydroxides, cor¬ responding to the halogen derivatives, CCh, SÍCI4, GeCU, and SnCl*, for the tetravalent elements—carbon, silicon, germanium, and tin :— IV IV IV C(OH)^, GeCOH)^. Normal Normal Normal Carbonic acid. Silicic acid. Germanic acid. These normal hydrates or acids have but little stability, and mostly exist only in some derivatives. By the separation of a mole¬ cule of water, they pass into CO3H3 SÍO3H3 Ge03H3 or CO(OH), SiO(OH), GeO(OH),. Carbonic acid. Silicic acid. Germanic acid. These hydroxyl derivatives deport themselves toward the normal just as the meta-acids of the elements of the N group do to the ortho-acids (see p. 204). They constitute the ordinary acids of the tetravalent elements, carbon, silicon, and tin, and as they contain 2 hydroxyl groups, are dibasic. Carbon is the lowest member of this group, with the least atomic weight. Among the elements of the other three groups correspond¬ ing to it are : nitrogen, oxygen, and fiuorine :— C = 12, N = 14.04, O = 16, Fl = 19. Fluorine and oxygen do not afford any oxygen acids. The normal acids of nitrogen, N(OH)6 and NO(OH)3, are very unstable, and pass readily into the meta-acids, NOj.OH and NO.OH. The nor¬ mal carbonic acid (C(0H)4) corresponds to this, but is not capable OXYGEN COMPOUNDS OF CARBON. 231 of existing. Indeed, the meta- or ordinary carbonic acid, H2CO3, is also very unstable and at once decomposes, when separated from its salts, into water and carbon dioxide, CO2. Even silicic, ger- manic, and stannic acids break up readily into water and their oxides :— CO, SiO, SnO, Carbon dioxide. Silicon dioxide. Stannic oxide. 1. OXYGEN COMPOUNDS OF CARBON. Carbon Dioxide—CO2—or carbonic anhydride (generally called carbonic acid). It is produced when carbon or its com¬ pounds are burned in air or oxygen. It is found free in the air (in 100 volumes, upon average, 0.05 volumes CO2), in many mineral springs (acid springs), and escapes in large quantities from the earth in many volcanic districts. It occurs in the liquid form, enclosed in the cavities of many crystalline minerals. It is pre¬ pared on a large scale by burning coke ; in the laboratory it may be most conveniently obtained by the decomposition of calcium carbonate (marble or chalk) with dilute hydrochloric acid ;— CaCOj + 2HCI = CaCl, + CO, + H.p. Calcium Calcium carbonate. chloride. Carbon dioxide is a colorless gas, of sweetish odor and taste. Its gas density equals 1.527 (air = i), or 22 (H=i), correspond¬ ing to the molecular formula, CO2 = 44. Owing to its weight, the gas may be collected by air displacement, and may be poured from one vessel into another filled with air. Faraday was the first to liquefy carbon dioxide by pressure. The apparatus of Thilorier and Natterer were employed to this end. At present liquid dioxide is brought into market enclosed in wrought-iron cylinders, and is applied quite regularly, e.g.y to compress cast steel. Carbon diox¬ ide can only be liquefied below + 30.9°—this is its critical tempera¬ ture (p. 47). Its tension (critical pressure) at this point equals 73.6 atmospheres. If liquid CO2, enclosed in some suitable vessel, be allowed to escape into the air by opening a stop-cock (ordinary pressure), it immediately solidifies (see below) to a white, snowy mass. This is because in the evaporation of a part of the liquid so much heat is withdrawn that the remainder becomes solid. Solid carbon dioxide is a very poor conductor of heat and vaporizes very slowly. Notwithstanding its low temperature it can be handled without serious result, because it is always surrounded by a gaseous layer, and is, consequently, not in immediate contact with the skin. If, however, it be pressed between the fingers it will produce pain¬ ful burns. 232 INORGANIC CHEMISTRY. The temperature of the solid carbon dioxide vaporizing in the air under ordinary temperature is about—78° (below—99° accord¬ ing to Faraday). When the solid dioxide is mixed with a little ether it forms a paste, and then conducts heat better, and is, therefore, well adapted as a cooling agent. In vacuo its temperature dimin¬ ishes to —140° C. Liquid carbon dioxide is readily obtained by filling the snowy mass into a thick tube, and then sealing the latter. It is a color¬ less, very mobile liquid. Its sp. gr. is 0.923 at 0°, 0.868 at 10°, and 0.782 at 20°. Its coefficient of expansion is, consequently, greater than that of the gases. Other gases behave similarly, but only such as are condensed under great pressure. If liquid carbon, dioxide, contained in a glass tube, be heated, it expands rapidly and suddenly passes, at the critical temperature -j- 30.9°, into gas. This behavior enables us to determine without difficulty whether the liquids contained (see above) in minerals are liquid carbon di¬ oxide. If liquid carbon dioxide, confined in a glass tube, be cooled by a mixture of solid, showy dioxide and ether (see above), it will solid¬ ify to a transparent ice-like mass, which will fuse at —65°. The tension of the solid or liquid dioxide, which at the same time indicates the pressure necessary for condensation at various temperatures, is given in the following table :— Temperature. Tension. Temperature. Tension. + 30.9° 73.6 atmos. — 21® 21.5 atmos. 20° 56.0 " -40° II.Q " 10® 450 " — 59-4° 4.6 « 0® 38-5 " — 70.6® 2.3 « — 78® 1.2 " At the temperature of fusion of the soIi4 dioxide (—65®) the tension equals about 3.5 atmospheres : the resulting liquid has this tension at this temperature. If the external pressure exerted upon it be less, it cannot exist as a liquid, but must immediately pass into the gaseous state. Herein we observe why the solid dioxide (under ordinary pressure) does not melt in the air, but vaporizes at once; and, further, it explains why the liquid dioxide, subjected to the ordinary atmospheric pressure, cannot continue in this state—why it either is gasified at once, or changed to the solid form. Many other fusible solids behave like the dioxide. If the tension of their vapors at the fusion temperature exceeds that of the external atmospheric pressure they do not melt in the air, because the resulting liquid is immediately trans¬ formed into vapor; they vaporize (sublime) directly, without previous fusion. Such bodies are, e. g., arsenic, arsenic trioxide, ASjOj, camphor, hexachlor- ethane CjClj, etc. They can only be fused under increased pressure (in sealed tubes). Again, all solids fusible in the air (under ordinary pressure) may be converted directly into gases by removing the external pressure. Thus iodine OXYGEN COMPOUNDS OF CARBON. 233 fuses at 114®, but sublimes in a vacuum without previous fusion. Mercuric chloride fuses at 265°, but not if the external pressure be less than 420 mm. Water melts at 0°, its tension at this temperature is 4.6 mm. If the external pres¬ sure be less (in vacuo), it will no longer melt, but vaporize at once. The pressure below which solids no longer melt has been called the critical pressure (Carnelley). It is, of course, understood that this is nothing more than the tension of the sub¬ stance at its point of fusion. Water dissolves its volume of the gas at 14° ; at 0° it takes up 1.79 vols. This proportion remains constant for every pressure, i. e.j at every pressure the same volume of gas is absorbed. As gases are condensed in proportion to the pressure, the quantity of absorbed gas is also proportional to the former. Hence i volume of HjO absorbs, at 14° and two atmospheres, 2 volumes, at 3 atmos¬ pheres 3 volumes, etc., of carbon dioxide—measured at ordinary pressure. The gas absorbed at higher pressure escapes with effer¬ vescence of the liquid when the pressure is diminished; upon this depends the sparkling of soda water and champagne, which are saturated with CO2 under high pressure. Every naturally occurring water, especially spring water, holds CO2 in solution, which imparts to it a refreshing taste. As the product of a complete combustion carbon dioxide is not combustible, and is unable to support the combustion of most bodies, a glimmering chip is immediately extinguished in it. In a similar manner it is non-respirable. Although it is not poisonous, in the true sense of the word, yet the admixture of a few per cent, of COj to the air makes it suffocating, as it retards the separation of the same gas from the lungs. It is decomposed by continued action of the electric sparks into carbonous oxide (CO) and oxygen; upon heating to 1300° it suffers a partial decomposition (dissociation) into CO and O. It is also decomposed when conducted over heated K or Na, with separation of carbon ; the potassium combines with oxygen to form potassium oxide :— CO2 -j- 2K2 = c 4" 2K2O, which forms potassium carbonate (KjCOj), with excess of CO2. Glowing carbon reduces the dioxide to the monoxide (p. 235). CO is analogously formed on conducting a mixture of the dioxide and hydrogen (equal volumes) through a tube heated to redness :— CO2 + H2 = CO 4- H2O. The composition of carbon dioxide is readily determined by burning a weighed quantity of pure carbon (diamond or graphite) in a current of oxygen, and ascertaining the weight of the result¬ ing gas. From the formula COj it follows that in one volume it 20 234 INORGANIC CHEMISTRY. contains an equal volume of O. We may satisfy ourselves of this by burning C in a definite volume of O ; after cooling, there is obtained an equal volume of carbon dioxide:— C -\- O^ — COj. I vol, 1 voU The experiment is most practically executed by aid of the appa¬ ratus of Hofmann pictured in Fig. 87. The spherical-shaped ex¬ pansion of the eudiometer limb of the U tube is closed by means of a glass stopper, through which two copper wires pass. The one wire bears a com¬ bustion spoon at its end, upon which lies the car¬ bon to be burned, while the other wire terminates in a thin piece of platinum, which is in contact with the carbon. For the performance of the experiment, the air is expelled from the globe limb by means of a rapid current of oxygen, the stopper placed in air-tight, the mercury level noted, and the copper wires connected with the poles of an induc¬ tion current from 3-4 Bunsen elements, which induces the burning of the carbon. As the volume of the enclosed gas is greatly expanded by the heat developed, it is advisable, in order to avoid the jumping out of the stopper, to previously re¬ duce the pressure of the gas about two-thirds, by running out mercury. A practical modification of this apparatus consists in having the usual cylin¬ drical eudiometer limb provided with two side tubes, in which carbon electrodes can be inserted. This same apparatus can also be employed for the illustration of the volume relations observed in the combustion of sulphur and other bodies. The Physiological Importance of Carbon Dioxide.—The gas is present in the atmosphere, and is inhaled by the plants. The chlorophyl grains in the green parts of the plant decompose carbon dioxide in sunlight, with a partial separation of oxygen ; by the mutai action of water and ammonia the innumerable carbon compounds peculiar to plants are formed from the residue. The respiration and life process of animals are essentially the reverse of the above. These absorb the oxygen of the air through the lungs ; and, influenced by the blood corpuscles, it oxidizes the substances present in the blood, and in this manner shapes the life process. The final products of the oxidation are carbon dioxide and water, which are exhaled. The absorption of O by animals, and its separation by plants, as also the reverse course of COj, are about alike, so that the quantities of O and COj in the air show no appreciable alteration. In dry condition, carbon dioxide, like all anhydrides, exhibits neither basic nor acid reaction. In aqueous solution it colors blue CARBON MONOXIDE. 235 litmus paper a faint red ; upon drying the latter the red disappears, in consequence of the evaporation of the carbon dioxide. We may then regard it as probable that free carbonic acid, H2CO3, is contained in the aqueous solution, but this readily decomposes into the dioxide, CO2, and water. The salts of carbonic acid are produced by the action of carbon dioxide upon the bases :— 2KOH -f CO2 = COgKj 4- HjO. Potassium carbonate. Carbon dioxide is, therefore, easily absorbed by potassium and sodium hydroxide. On conducting it through a solution of calcium or barium hydroxide, a white precipitate of barium or calcium carbonate, CaCOs, is produced. Carbonic acid is dibasic, forming primary (acid) and secondary (neutral) salts, CO3HK and CO3K2, called carbonates. As the acidity of carbonic acid is only slight, the secondary salts, obtained from strong bases, exhibit a basic reaction. Most acids expel the weak carbonic acid, from its salts, with decomposition into carbon dioxide and water. Carbon Monoxide—CO—is produced in the imperfect com¬ bustion of carbon by insufficient access of air, and when carbon dioxide is conducted over red-hot coals: CO2 C = 2CO. The forms of apparatus described, p. 234, serve for the demonstration of this volume relation. Carbon monoxide is, therefore, found in the generator gases produced by incomplete combustion. Zinc dust reacts like carbon :— COj 4- Zn = ZnO -|- CO. When carbon dioxide is conducted through a glass tube, containing zinc dust heated to a faint red heat, almost pure carbon monoxide escapes. A more con¬ venient procedure consists in heating pulverized magnesium carbonate and zinc dust in a glass retort, when CO, containing CO,, is eliminated; subsequently the former alone escapes. Pure monoxide is also formed upon heating zinc dust with chalk (in equivalent quantities) : Zn 4- CaCOg = ZnO 4- CaO CO. The monoxide is produced, further, by igniting carbon with different metallic oxides, e.g, zinc oxide : ZnO -j- C = Zn4- CO. Water is similarly decomposed. On conducting aqueous vapor over burning carbon there is produced a mixture of carbon monoxide and hydrogen :— C 4- HjO = CO 4- Hj. This gas mixture is known as water gas, and is applied technically. For preparation of carbon monoxide, oxalic acid is warmed with sulphuric acid : the latter withdraws water from the former, and the residue breaks up into carbon dioxide and monoxide:— CjO^H, =i= COj 4- CO 4- HjO. 236 INORGANIC CHEMISTRY. The disengaged mixture of gases is conducted through an aqueous solution of sodium hydroxide, by which the CO2 is absorbed, the monoxide passing through unaltered. Pure monoxide may be prepared by heating yellow prussiate of potassium (see Iron) with 9 parts H2SO4. The resulting gas is conducted through sodium hydroxide to remove from it traces of CO2 and SO2. Its specific gravity equals 0.969 (air = i) or 14 (H=i), corresponding to the molecular formula CO = 28. It is one of the gases which are condensed with difficulty. Its critical temperature is —141® and its critical pressure is 35 atmospheres. Liquid carbon monoxide solidifies below 100 mm. pressure at —270°, and at 4 mm. pressure indicates a temperature of —220° (Olszewski). It is almost insol¬ uble in water, but is readily dissolved by an arnmoniacal solution of cuprous chloride (CuCl) with which it forms a crystalline com¬ pound, but this decomposes when its solution is heated and CO Í5 again liberated. When ignited, it burns in the air, with a faintly luminous, beautiful blue flame, which distinguishes it from other combustible gases. With air or oxygen, it affords a very explosive mixture :— 2CO + O, = 2CO2. 3 vols. I vol. 3 vols. The union of carbon monoxide and oxygen takes place at very high temperatures; hence the burning flame of the gas is extin¬ guished upon cooling. A flame or the spark from a powerful induction coil is necessary to ignite a dry mixture of carbon mon¬ oxide and oxygen. When the two gases are moist they are more easily ignited and combustible. This is explained by the fact that carbon monoxide unites with the aqueous vapor and yields the dioxide and hydrogen (CO + H2O = CO2 + Hj), which then combines with the oxygen and forms water (p. 102). In consequence of its ready oxidation, it is capable of reducing most metallic oxides at a red heat ;— CuO + CO = Cu + CO, ... + (31.3 Cal.) (37.1 Cal.) (38.5 Cal.) (96.9 Cal ) Some noble metals are precipitated from solutions of their salts by CO even in the cold. Thus, palladium is thrown out from its chloride solution, and a piece of paper moistened with palladious chloride is blackened by it (delicate test for CO). As it easily oxidizes, carbon monoxide is reduced to carbon with difficulty. Burning bodies are extinguished by it. When it is heated with potassium it is decomposed with separation of carbon. When inhaled, it acts very poisonously, even in slight quantity, as it expels the oxygen of the blood. The carbon vapor, developed in heated ovens closed too soon, is carbon monoxide. As an unsaturated compound, this oxide, like ethylene, unites directly with COMPOUNDS OF CARBON WITH SULPHUR. 237 2 atoms of chlorine, to yield carbon oxychloride, ox phosgene gas, COCL:— CO + CI, = COCI,. X voL 1 vol. I vol. This is obtained by bringing together equal volumes of CO and CI, in direct sunlight, or, better, by conducting CO into SbClj. It is a colorless, suffocating gas, of specific gravity 49.4 (H = i), agreeing with the molecular formula COCI, = 98.9. Water decomposes it into hydrogen chloride and carbon dioxide :— COCI, 4- H2O = CO, + 2HCI. COMPOUNDS OF CARBON WITH SULPHUR. Carbon Disulphide—CS,—is formed, like the dioxide, by the direct union of carbon and sulphur; if vapors of the latter are led over ignited carbon the escaping disulphide vapors are condensed in a cooled receiver. It is a colorless, mobile liquid, of characteristic odor, solidifies at —116°, and refracts light strongly. Its specific gravity equals 1.29 at 0°. It is very volatile, boils at 47°, and burns with a blue flame, to carbon dioxide and sulphurous acid. When a mixture of carbon disulphide vapors and oxygen is ignited, a violent explosion ensues :— CS, + 3O, = CO, + 2SO,. I vol. 3 vols. . 1 vol. 2 vols. In nitrous oxide, the vapors burn with a bright, blinding flame. On blowing a strong current of air upon carbon disulphide in a porcelain capsule (which conducts heat poorly), so much heat is absorbed by the evaporation, that the residual liquid solidifies to a white, snow-like mass. Carbon disulphide is insoluble in water ; but mixes, in every proportion, with alcohol and ether. It dis¬ solves iodine with a violet-red color, and serves as an excellent solvent for sulphur, phosphorus, caoutchouc, and the fatty oils. On conducting the CS, vapors over heated zinc dust, all the sulphur unites with the zinc, forming zinc sulphide, while the carbon separates as soot :— CS, + 2Zn = 2ZnS + C . . . + (95.6 Cal.) (— 12.6 CaL) (2 X 41.S Cal.) Most metals react in a similar manner. Carbon disulphide may be viewed as the anhydride of sulpho- carbonic acid—H,CS8. The salts of this acid are obtained by the solution of CSa in alkaline sulphides (see Sulpho-salts, p. 226):— CS, -f- K,S = KjCS,. On adding hydrochloric acid to the solutions of these salts the sulphocarbonic acid separates as a reddish-brown oil. This decom¬ poses readily. 238 INORGANIC CHEMISTRY. The sulphur compound corresponding to CO is not known there exists, however, one containing both oxygen and sulphur— Carbon oxysulphide, COS. It is produced (in small quantity) when a mixture of sulphur vapors and carbon monoxide gas is passed through red-hot tubes and by heating carbon disulphide with sulphuric oxide :— CSj + 3SO3 = COS + 4S0a. It is most readily obtained from potassium sulphocyanide— CN.SK (see Organic Chemistry) by the action of dilute sulphuric acid. It is a colorless gas, with an odor reminding one of carbon dioxide and hydrogen sulphide. It is present in some sulphur springs. It is very readily inflammable and burns with a blue flame :— 2COS + 3O2 = 2CO2 + 2SO2. 2 vols. 3 vols. 2 vols. 2 vols. It is decomposed at a red heat into CO and sulphur. It is soluble in an equal volume of water, decomposing gradually into the dioxide and hydrogen sulphide :— COS + HjO r= CO2 + SH,. CYANOGEN COMPOUNDS. Of the innumerable compounds of C treated in organic chem¬ istry, we will here mention only those of cyanogen, as they are of importance in inorganic chemistry. Nitrogenous carbon compounds heated with potassium hydroxide yield potassium cyanide—CNK—which with iron forms the so- called yellow prussiate of potassium, El4Fe(CN)6. All the other cyanogen derivatives may be prepared from these two compounds. They all contain the group CN, called cyanogen. In it we have a trivalent nitrogen atom combined with a tetravalent carbon atom ; III IV the fourth affinity of the latter is not saturated ; N-^C— : it is similar, therefore, to the groups OH, NHj, CHg, and is a monova¬ lent radical. In chemical behavior the cyanogen group is very similar to the halogens chlorine and bromine; with the metals it forms metallic cyanides (KCN, AgCN) very similar to the haloid salts. Hydrogen cyanide is evolved when the cyanides are heated with sulphuric acid:— 2KCN + H2SO4 = K^SO^ + 2HCN. Hydrogen Cyanide, HCN, is a colorless, mobile liquid, of peculiar odor, and boiling at 27°. It is an acid, forming salts with metals and bases, and is known as Hydrocyanic or Prussic * Upon standing in sunlight CSj is said to break up into S and CS—a chestnut- brown powder of specific gravity 1.66. OXYGEN COMPOUNDS OF SILICON. 239 acid. Both it and its salts are very powerful poisons. If the CN group is separated from its salts it doubles itself, yielding dicyanogen or free cyanogen^ C2Na (N=C—C=N), because, like the other monovalent groups (as CHg, see p. 172), it cannot exist in a free condition. The heat occurring in the formation of the simplest carbon compounds (from amorphous carbon) above cited corresponds with the symbols :— (C,0) = 28.5. (C0,0) = 68.3. (CA) = 96.9. (C,0,S) = 1.4- (QSj) =—12.6. (C,N,H) = —28.3. (COjClj) = 4 3* (C02»^'l*) " 5-^* If an element combine with another according to multiple proportions, there usually occurs, in the union of the first atom, a greater disengagement of heat than with the following atom (compare nitrogen oxides, p. 216). The numbers above, on the contrary, show that the union of the second atom of oxygen with carbon (C0,0) sets free 68.3 calories ; that of the first atom (C,0), however, only 28.5 calories. This can only be explained by the fact that, for the vaporization and disaggregation of the solid carbon molecules, heat is necessary. If we assume that the direct union of the first atom, also disengaged 68.3 calories, it would follow from this that, in the dissociation of 12 parts carbon by weight into gaseous free atoms, 39.7 (= 68.3 — 28.5) calories were absorbed. This would explain the heat absorption in the production of CSg, CNH, CjHj, while otherwise heat is invariably disengaged in every direct chemical union. Comparing the elements of the carbon group with each other, we discover that the heat disengi^ement is greatest with the compounds of silicon :— C,C1J = 28.3. (Si,ClJ = 157-6. (Sn,Cl^l = 129.2. CA) — 96.9- (Si,02) = 219.0. (Sn,02,H20) = 133.5. From these numbers we observe that tin dioxide, but not that of silicon, can be reduced by carbon. 2. OXYGEN COMPOUNDS OF SILICON. Silicon Dioxide, SÍO2 {Silica)^ is widely distributed in nature as rock crystal, quartz, sand, etc. It is artificially obtained as a white, amorphous powder, of specific gravity 2.2, by the com¬ bustion of amorphous silicon, or by the ignition of the silicic acids. It only occurs in nature crystallized in figures of the hexagonal system, with the specific gravity, 2.6; these crystals are colorless, or colored by impurities. In the oxy-hydrogen flame it fuses to a transparent glass. Silicon dioxide is insoluble in water and all acids -, but is decom¬ posed by hydrofluoric acid with the formation of silicon fluoride (SÍFI4) and water (p. 165). Strong ignition with sodium or potas¬ sium reduces it to metallic silicon. The dioxide prepared artificially dissolves when boiled with potassium or sodium hydroxide ; some of the naturally occurring amorphous varieties are also soluble, but not the crystallized dioxide. By fusion with the hydroxides or 240 INORGANIC CHEMISTRY. carbonates of the alkalies all varieties of silica yield a glassy mass (water glass) soluble in water and containing silicates (K4SÍO4 or K2SÍO3). Upon the addition of hydrochloric acid to the aqueous solution of the potassium or sodium salt, a very voluminous, gela¬ tinous mass separates; this is probably normal silicic acid, H4SÍO4:— Na^SiO^ + 4HCI = 4NaCl + H^SiO^. It becomes a white amorphous powder having the composition H2SÍO3 when washed with water and dried in the air. The freshly precipitated acid is somewhat soluble in water, more readily in dilute hydrochloric acid. On adding a solution of sodium silicate to an excess of dilute hydrochloric acid the separated silicic acid remains dissolved. From the hydrochloric acid and sodium Fig. 88. chloride solution we can obtain a perfectly pure aqueous solution of silicic acid by dialysis by proceeding in the following manner : Pour the hydrochloric acid solution into a wide cylindrical vessel whose lower opening is covered with animal bladder or parchment paper, and then suspend the vessel (dialyser) in another contain¬ ing pure water. (Fig. 88.) Osmosis now sets in. The sodium chloride and hydrochloric acid pass through the parchment paper into the outer water, while on the other hand, water passes from the outer vessel into the dialyser; the parchment paper is not permeable to silicic acid. This alternate diffusion of the different particles occurs until the outer and inner liquids show the same quantity of diffusible substances. Upon introducing the dialyser into a fresh portion of water, the dialysis commences anew. Finally, after repeated renewal of the external water, the dialyser will contain a perfectly pure silicic acid solution, free from sodium chloride and hydrochloric acid. The solution may be concen¬ trated by evaporation; it then readily passes into a gelatinous mass. The same occurs instantaneously in dilute solutions if a trace of sodium carbonate be added or carbon dioxide be led into it. TITANIUM—ZIRCONIUM—THORIUM. 241 Like sodium cbloiide, all crystallizable soluble substances diffuse through jjarchment. These are known as crystalloids, to distinguish them from the non- difhisible colloids. To the latter belong gum, gelatine, albumen, starch, glue {KÓTíkcLy hence the name colloid), and especially most of the substances which occur chiefly in vegetable and animal organisms. Like silicic acid these colloids exist in liquid, soluble, and solid gelatinous condition. Many other substances (like ferric and aluminium oxides) which ordinarily are insoluble, can be brought into aqueous solution by dialysis. We have already seen that acids like sulphuric, phosphoric, and arsenic, are capable of forming anhydro- or poly-acids by the union of several molecules and the elimination of water. Silicic acid is particularly inclined to that kind of condensation. It forms a large number of poly-silicic acids, SÍ203(0H)2, SÍ304(0H)4, 81305(011)2, etc., derived from the normal and ordinary acid, according to the common formula:— wSi(0H)4 — «Hp. These poly-acids are not known free ; it appears, however, that many amorphous forms of silica occurring in nature, as agate, chalcedony, opal, which lose 5-15 per cent. HjO by ignition, represent such poly-acids. The natural silicates are the salts of such acids. The majority are derived from the acids : HaSiaOs, H4SÍ30a, H2SÍ3O7, HiSi^Og, and others. Only a few silicates are obtained from the normal acid, e. g., chrysolite—Mg2Si04. Corresponding to CSg is Silicon Disulphide, SÍS2, which may be made by heating amorphous silicon with sulphur, or by conducting sulphur vapors over an ignited mass of silica and carbon. It sublimes in shining, silky needles, which water changes to silicic acid and hydrogen sulphide. Germanium, Tin, and Lead also belong to the group of carbon and silicon. Their oxygen compounds, MeOj and Me (OH)*, are perfectly analogous to those of the latter elements. The metallic character is, however, more pronounced in germanium, tin, and lead, and becomes more noticeable as the atomic weight increases. Germanium and tin form lower oxides, GeO and SnO, which mani¬ fest a perfectly basic nature and unite with acids to form salts ; therefore germanium, tin, and lead will be discussed under the metals. TITANIUM. ZIRCONIUM. THORIUM. Ti = 48,1. Zr = 90.7. Th — 233.4. The same relation that vanadium, niobium, and tantalum show to the elements of the phosphorus group, is manifested by the three elements, titanium, zirco¬ nium, and thorium, for the silicon group. (See Periodic System, p. 249.) 21 242 INORGANIC CHEMISTRY. P= 3103 V= 51.2 Si = 28.4 Ti = 48.1 As = 75 Nb = 94.2 Ge = 72.3 Zr = 90.7 Sb = 120,3 T = 183 Sn =118.1 Th = 232.4. In all their deportment they strongly resemble tin ; they possess, however, a more metallic character in their derivatives. They are tetravalent, affording compounds of the form MeX^, in which X represents monovalent elements and groups; those of the form MeX^, corresponding to the stannous derivatives, are unknown. The hydroxides, Me(OH)^ and MeO(OH)2, have a stronger basic nature than stannic acid and form stable salts with acids ; the basicity increases successively with the atomic weights, in the order, Ti,Zr,Th. Corresponding to this, the acidity of the hydrates, i. e., their capability of exchanging H for metals, gradually diminishes. Thorium hydroxide, Th(0H)4, is not able to form metallic salts. TITANIUM. Ti = 48.1 This metal occurs in nature as titanium dioxide (rutile, anatase, brookite) and in titanates (perofskite, TiOgCa, menaccanite, FeTiOg). Free titanium is a gray, metallic powder, obtained by heating potassium titan-fluoride (KjTiFlg) with potassium. It burns when heated in the air, and decomposes water on boiling. It dissolves in dilute hydrochloric and sulphuric acids, with evolution of hydrogen. Titanium Chloride—TiCl^—is formed, like silicon chloride, by conducting chlorine over an ignited mixture of the dioxide and carbon. A colorless liquid, of specific gravity 1.76, fuming strongly in the air (with decomposition into hydrochloric and titanic acids), and boiling at 136°. The vapor density equals 95 (H = i). corresponding to the molecular formula TiCl^ = 190. It behaves like tin tetrachloride with water. A compound TijClj, analogous to CjCl^, is known. Titanium Fluoride—TiFl^—is not known in a free condition, but forms beautifully crystallized double salts, e.g., TiFl^ 2KFI, corresponding to the silico- fluorides (KgSiFlg). Titanic Acid—H^TiO^—separates as a white, amorphous powder, on adding ammonium hydroxide to the hydrochloric acid solution of the titanates. When dried over sulphuric acid it loses i molecule HgO and becomes TiO(OH)j. Titanic acid, like silicic and stannic acids, forms poly-acids. The hydrates dis¬ solve in alkalies and strong acids, to form salts. On igniting the hydroxides we get Titanium Dioxide—TiOg, which may be procured crystallized as rutile, brookite, and anatase. When ignited in a stream of hydrogen it changes to the oxide TigOg. Titanium dioxide is almost insoluble in the acids ; it is only dis¬ solved by hydrofluoric acid. It forms titanates upon fusion with the alkalies. The hydroxides, TiO^H^, TiOgHg, etc., conduct themselves as feeble bases with strong acids and afford salts with them (er, moistened with it, is colored red-brown after drying. On boiling the solution, boric acid escapes with the steam. An alcoholic solution of the acid burns with a green flame. These reactions afford a ready means for its detection. When heated to 100°, the acid loses i molecule of HjO, and passes into the anhydro- or meta-acid, HBO2, which at 140° is converted into Tetraboric acid—B4OTH2. When ignited, boric anhydride or Boron trioxidcy B2O3, is produced. This is a fusible, glassy mass, oí specific gravity 1.8, and is slightly volatile at a very high heat. Water dissolves the anhydride to boric acid. It is a very weak acid ; and can be expelled from its salts by most other acids. By fusion it removes the most acids from their salts, in consequence of the difficult volatility of its anhydride. Salts of normal boric acid, B(OH)5, are not known, while the ethers, B(O.CH8)3, are. The salts of metaboric acid, BO.OK, can be obtained crystallized, but they are very unstable. They are decomposed by carbon dioxide with production of salts of tetraboric acid :— 4NaB0, + COj = B40TNa, -f COjNa,. The latter, from which the ordinary borates are derived (see Borax), may be viewed as an anhydro-acid, produced by the union of 4 molecules of trihydric boric acid (compare p. 197) :— 4B(0H)3 —5H,0 = BAH2. 246 INORGANIC CHEMISTRY. On heating amorphous boron in a stream of N or ammonia, or by igniting a mixture of the trioxide and carbon in nitrogen gas, there is formed boron nitride^ BN. This is a white amorphous powder, which gives forth an extremely intense greenish-white light when heated in a flame. Boric acid and ammonia result when steam at 200° is conducted over the nitride :— BN + 3H,0 = B(0H)3 + NH3. PERIODIC SYSTEM OF THE ELEMENTS. In the preceding pages we have studied four groups of elements and their compounds with hydrogen, the halogens and oxygen. We have repeatedly directed attention to the remarkable relations of the elements of a single group, as well as to those of the various groups to each other, but they appear more manifest if viewed in the connection in which they present themselves in the periodic system of elements. The position which these elements occupy in this system determines their entire physical and chemical character to a marked degree. The system is based upon the grouping of the elements according to the magnitude of their atomic weights. For the longest time we have been cognizant of the remarkable relations existing between the atomic weights of analogous elements, but only recently has the law of periodicity underlying them been announced by Men- delejeif and Lothar Meyer, and, according to this, the properties of the elements and their compounds present themselves as periodic functions of the atomic weights^ Arranging the elements according to increasing atomic weight we observe that similar elements return after definite intervals. Thus they arrange themselves in several periods, consisting of the following horizontal series (for brevity the atomic weights are not attached to the symbols) :— I. Li Be B C N 0 Fl 2. Na Mg Al Si F S Cl K Ca Sc Ti V Cr Mn Fe Co Ni • Cu Zn Ga Ge As Se Br Rb Sr Y Zr NbMo Ru Rh Pd • Ag Cd In Sn Sb Te I Cs Ba (La Ce Di) — Yb— TaW Os Ir Pt i Au Kg T1 Pb Bi Th Ur The first two series, lithium (Li) to fluorine (Fl), and sodium (Na) to chlorine (CI), present two periods of seven members each, * John A. B. Newlands published similar views as early as 1865. He recDg- nized that every seventh element was an analogue of the first, and called these relations the law of octaves (see On the Discovery of the Periodic Law, by John A. B. Newlands, 1884). PERIODIC SYSTEM. 247 in which the corresponding (above and below) members exhibit a great but not complete analogy. Sodium resembles lithium ; mag¬ nesium, beryllium ; chlorine, fluorine, etc. Then follow two periods, consisting of 17 elements each ; potassium (K)to bromine (Br), and rubidium (Rb) to iodine (I). The series 5 and 6 are incomplete, and together probably constitute a period. In the 7th series there are as yet but two elements: thorium = 232 and uranium = 240. Thus result 3 great periods^ whose corresponding members exhibit an almost complete analogy : the elements K Rb Cs, Ca Sr Ba, Ga In Tl, As Sb Bi, etc., are so similar that they remind us of the homologous series of the carbon compounds (com¬ pare p. 155), and, therefore, can be designated as homologous elements. It is only in the third great period (series 5 and 6) that the middle members exhibit any variations. Now on comparing the three great periods with the two small ones, we discover that the first members are analogous to each other : K,Rb,Cs resemble Na and Li ; Ca,Sr,Ba resemble Mg and Be. Then the similarity gradually lessens, disappears apparently in the middle members, and only appears again toward the end of the periods : I and Br resemble chlorine and fluorine ; Te and Se, sulphur and oxygen; Bi, Sb, As, phosphorus and nitrogen, etc. The character or the function of the three great periods is therefore other than that of the two small periods. But in all five periods we can detect a gradual, regular alteration in the properties of the adjoining heterologous elements. This is particularly manifest in the measurable physical properties, all of which show a maximum or minimum in the middle of the periods (both of the great and small), as may be seen, for example, in the specific gravity of the elements in solid condition (compare further the atomic volumes, p. 258):— Na Mg Al Si P S Cl Specific gravity, 0.67 1.7 2.5 2.5 2,0 1.9 1.3 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br 0.86 1.6 3.8 — 5.56.8 7.2 ,7.9 8.5 8.8 8.8 7.1 5.9 5.5 5.6 4.6 2.9. These relations show themselves very clearly in a graphic repre¬ sentation, by making the atomic weights the abscissas, the numer¬ ical values of the properties the ordinates; then the individual periods represent segments of curves, which blend to a curve with alternating maxima and minima. The same regularity exhibits itself even in chemical properties, in the two small periods, especially in the valence of the elements in their compounds with hydrogen or the hydrocarbon groups CHsjCaHg, etc. (compare p. 176 and p. 251). The hydrogen 248 INORGANIC CHEMISTRY. valence rises and falls periodically with the condensation of the substance (corresponding to the specific gravity) ;— 1 II III IV HI II I NaR MgR, AIR, SiH^ PH, SH, CIH. On the other hand, the maximum valence of the elements in¬ creases successively in the salt-forming oxides (p. 174) :— I n HI IV V VI VII Na,0 MgO AljjO, SiO, Pp, SO, Cip,. The chemical valence expresses itself somewhat differently in the three great periods. In them we have a double periodicity ; thus, e. g.y with the salt-forming oxides : — I II III rv V VI VII KjO CaO SCjO, TiO, Vp^ CrO, Mnp, VI IV n FeO, CoO, NiO I II HI rv V VI VH Cup ZnO Ga^O, — Asp, SeO, Brp^. In consequence of this double periodicity, the first seven and the last seven members of the three great periods, with respect to their valence (and consequently also their compounds), resemble the seven members of the two small periods. To bring out this double periodicity and analogy, the first seven and last seven mem¬ bers of the great periods are divided into two series, and arranged under the corresponding seven members of the small periods. In this wày the three middle members of the great periods (which are found between the dotted lines of the table, p. 246) come to stand apart, as they have no analogues. In this manner arises the following table, in which the seven (or ten) vertical columns include analogous elements ;— Li Be B C N 0 Fl Na Mg Al Si P S CI K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br" Rb Sr Y Zr Nb Mo — Ru Rh Pd Ag Cd In Sn Sb Te I Cs Ba La (CeDi) — — — — Yb — Ta W — Os Ir Pt Au Hg T1 Pb Bi — — The Periodic System of the Elements. I Group. II Group. Ill Group. IV Group. V Group. VI Group. VII Group. VIII Group. H - COMTOUKDS. Highest salt- foming oxides. MjO MO MjOj MH^ MOj MHg MgOg MH, MOg MH MgO^ MOg MOj (M,H) MO Periods, ist Series. 1st H I Li 7 Be 9 B II C 12 N 14 0 16 Fl 19 2d 2d Na 23 Mg 24 Al 27 Si 28 P31 S 32 Cl 35-4 3d f 3d I 4th K39 Cu 63 Ca 40 Zn 65 Sc 44 Ga 70 Ti 48 Ge 72 y 51 As 75 Cr 52 Se 79 Mn 55 Br 79.7 Fe 56 Co 58 Ni 58 4th - r 5th 6th Rb 8s Ag 108 Sr87 Cd 112 Y 89 In 113 Zr 90 Sn 117 Nb94 Sb 120 Mo 96 Te 126 — 100 I 126.5 Ru 104 Rh 104 Pd 106 5th 7th 8th 9th loth Cs 133 Au 197 Ba 137 Hg 200 La 139 Yb 173 T1 204 (Ce 140 (Er 166) ? Pb 206 Di 142) Ta 182 Bi 210 W 184 (Sa 150)? Os(i95) Ir 193 Pt 195 — — — Th 232 — Ur 240 — § o Ö n cn S s bO CO 250 INORGANIC CHEMISTRY. In the preceding table (p. 249) we have presented the same grouping of the elements, together with their atomic weights, given in round numbers. In this arrangement we must always bring into consideration that the principé analogy (homology) of the three great periods finds expression in the three uninterrupted horizontal series (p. 246), and that the decomposition of the latter into two series each only corresponds to the secondary, double analogy with the small periods. It may be further remarked that, in the second small period the last three members, P, S and CI, show a complete homology with the corresponding members of the large periods—as is represented in the table. When the periodic grouping of the elements was first presented, some atomic weights, not sufficiently well established at that time, had to be more or less altered. Thus, the atomic weight of indium, formerly 75.6, was made 113.4, and that of uranium 240 (before 120). All such alterations have been proved to be established through recent investigations. Further, the atomic weight of tellurium (formerly determined to be 128) had to be less than that of iodine (126.5) î this, also, has been established by recent researches, which place it at 125 (p. 114). There is, therefore, little doubt that the atomic weight of osmium (found I95) will also prove to be somewhat less. This is only the more likely, since it has been shown that the atomic weight of iridium, which was formerly given as 197, is really 192.7. Hence, the periodic system offers a control for the numbers of the atomic weight, while formerly they appeared to be irregular, and, at the same time, accidental. Further, upon the basis of the periodic system, the existence of new, not yet known, elements may be ascertained, which correspond to unoccupied, free places or gaps in the table. In fact, three such gaps have been filled up by the dis¬ covery of gallium (Ga = 69.8), scandium (Sc = 44), and germanium (Ge = 72.3) ; their properties have shown themselves to be perfectly accordant with those deduced from the periodic system. At present, only the first homolc^ue of manganese (with atomic weight of about 100) is wanting. The series 5 and 6 are very incomplete; the elements, terbium, samarium, and erbium, little investigated as yet, will probably find positions in them. It may be, however, that the two series will together form a single period of somewhat varying character. Consult further Annalen der Chemie^ 8 Supplement Band, p. 133. The entire character of a given element is determined to a very high degree by the law of periodicity ; hence, all physical and chemical properties of the same are influenced by its position in the system. These relations we will examine more closely in the indi¬ vidual groups of the elements, and here confine ourselves to a notice of some general relations, and the connection of atomic weight with the chemical valence of the elements and the thermo-chemical phenomena. The relation of metalloids to metals is shown with g^eat clearness PERIODIC SYSTEM. 251 in the periodic system. The first members of all periods (on the left side) consist of electro-positive metals, forming the strongest bases, the alkalies—Cs, Rb, K, Na, Li, and metals of the alkaline earths—Ba, Sr, Ca, Mg, and Be. The basic character diminishes successively, in the following heterologous members, and gradually passes over into the electro-negative, acid-forming character of the metalloids, Fl, Cl, Br, I. Here is observed that, in the periods following each other, with higher atomic weights, the basic metallic character constantly exceeds the metalloidal. The first period comprises five metalloids (B, C, N, O, Fl), the second only four (Si, P, S, CI), the fourth and fifth periods each only three (or two) metalloids (As, Se, Br, and Sb, Te, I), which, at the same time, become less negative. With the metalloidal nature is combined the power of forming volatile hydTogen compounds. Similar volatile derivatrves are also afforded by the metalloids with the monovalent hydrocarbon groups (as CHg, C2H5, CgH^, etc.), which resemble hydrogen in many respects. Such metallo-organic compounds, in which the elements show the same valence as in the hydrogen compounds, are also produced by the metals adjacent to the metal¬ loids ;— IX III IV III IX I Mg(CH3)„Al(CH3),. Si(CH3L, P(CH3)3, S(CH3)3, CICH3. Their stability gradually diminishes with the increasing basic nature of the metals ; hence, in the three large periods, this power extends only to Zn, Cd, and Kg. In consequence of the opposite (metalloidal and metallic) char¬ acter of the two ends of the periods, there are in the table repre¬ senting the double periodicity of the great periods (pp. 248 and 249) two sub-groups each, with the seven vertical groups ; on the left with the more positive, basic, and on the right with the more negative, metalloidal elements. Thus in group VI, in addition to Ö and S (belonging to the small periods) stands the more basic sub-group Cr, Mo, W, and the metalloids Se and Te ; in group It stand the strong basic metals Ca, Sr, Ba, and the less basic heavy metals, Zn, Cd, Hg. The elements of group VIH form the gradual transition from the latter to the former. The fundamental deduction necessarily resulting from the law of periodicity is, that the various elementary atoms must be aggrega¬ tions or condensations of one and the same primordial substance, a necessary correlative postulate of the recognized unity of all forces. Then only can we comprehend that the properties of the elements 252 INORGANIC CHEMISTRY. are functions of the atomic weight. It was once believed that this primordial substance was hydrogen (h)^othesis of Prout), because it seemed that the numbers representing the atomic weights were all whole numbers (multiples of the hydrogen atom = i). The most accurate determinations, made with exceeding care by Stas, prove that this is not correct in all instances. It is, however, note¬ worthy that of the 18-20 elements whose atomic weights have been carefully established, ten (Li, K, Na, C, O, S, N, etc., — p. 24), so nearly approach whole numbers, that a complete coincidence is not unlikely. It is possible that these elements represent multiples of the hydrogen atom. We can only expect to arrive at the under¬ lying law when the atomic numbers of a majority of the elements have been determined with equal accuracy. Periodicity of Chemical Valence.—Group I of the table comprises the monovalent metals, group II the divalent. In group III are the trivalent metalloid, boron, and the trivalent metals Al, Sc, Y, and Ga, In, Tl. In the tetravalent carbon group the valence arrives at its maximum ; from here it gradually decreases with in¬ creasing atomic weight ; the nitrogen group is trivalent, the oxygen group is divalent, that of the halogens monovalent. This valence is derived from the compounds with hydrogen and hydrocarbons (compare p. 251), or where such do not exist, as in the case of boron and many metals, from the halogen compounds :— IV CH^ III NH3 II OH, FIH LiCl NaCl II BeClj MgCl, III BCI3 AICI3 IV CCl^ SÍC14 III NCI, PCI3 II OCl, SCI, I Fl, Cl,. The elements of the first four groups are not capable of yielding higher compounds with the halogens. On the other hand, as we have seen, the higher analogues of nitrogen and other metalloids can unite with a larger number of halogen atoms (see p. 171). The higher valence of these elements is more manifest in the more stable oxygen compounds. On bringing together the highest oxides of the seven groups capable of forming salts (salt-building oxides), we get this series :— 1 n III IV V VI vii LijO BeO B2O3 COj NjOj SO3 IjO,. The elements of the first four groups in their oxygen compounds exhibit, consequently) the same valence as in the compounds with hydrogen (or hydrocarbon radicals) and the halogens ; in the last PERIODIC SYSTEM. 253 three series, however, there is noticed a constant increase of valence for oxygen. Besides these highest oxides, remarkable for their greater stability, the elements of the last three groups afford lower oxides, returning in this manner to the hydrogen valence ;— III IV V ^2^3 •'2^5" II III SClj CljOj. CljO. PH, SHj CIH. The hydroxyl compounds of the elements of the 7 groups arc analogous to the oxides in constitution. They afford the following series, expressing the maximum valence (compare p. 174) :— I II III IV V rv VII Na(OH) Mg(0H)2 A1(0H)3 Si(0H)4 P(0H)5 SiOH)^ Cl(OH)7. The hydroxyl compounds of the elements of the first 4 groups exist in free condition, excepting that of carbon, C(OH)4, which is only represented in its derivatives. The strong basic character of the hydroxides of group I (NaOH) diminishes, step by step, in the succeeding groups, down to the weak acid hydrate, Si(OH)4. The hydrates of the last three groups are of acid nature, and mostly unstable or not known. By elimination of i, 2 and 3 molecules of HjO, they yield the ordinary highest acids ;— V VI VII P0(0H)3 S02(0H)2 CIO3OH. Phosphoric acid. Sulphuric acid. Perchloric acid. The non-saturated hydroxides behave in the same way :— III IV V P(0H)3 S(0H), C1(0H)3 II III S(OH)j CI(0H)3 Cl(OH). Sulphurous acid, S0(0H)2 (p. 189), is derived from the hydrate, S(0H)4 J chloric acid, ClOa-OH, from the hydrate, Cl(OH)5 ; and cblorous acid, CIO.OH, from the hydrate, Cl(OH)5 The hydrates, P(0H)3, S(OH)j and ClOH, are very unstable, and the first two appear to pass readily into HPO(OH)s and HSO.OH (compare p. 187). It has been shown already in the case of periodic, sulphuric, and 254 INORGANIC CHEMISTRY. nitric acids, how the so-called hydrates with water of crystallization (regarded as molecular compounds) are explained by the accept¬ ance of the existence of such hydroxyl derivatives. The same may be done for many salts with water of crystallization. Thus, we see, and in the following pages will find it more exten¬ sively developed, that the relations of valence of the elements have their complete expression in the periodic system, are regulated by it, and hence we must conclude that, in fact, the valence is not only a property attaching to the elements per se, but is influenced also by the nature of the combining elements ; the hydrogen valence is constant, the valence to oxygen and the halogens, on the con¬ trary, varies according to definite rules. Valence, therefore, is a relative function of the elements (p. 174). Periodicity of Thermo-Chemical Phenomena. We observed in the case of the elements of the chlorine and sulphur group, that, in their hydrogen com¬ pounds the heat liberation decreased successively with increasing atomic weight (pp. 66 and 116), while there is generally an increase in their oxygen derivatives (pp. 178 and 186). Similar relations exhibit themselves with the halogen, oxy¬ gen, and sulphur compounds of the metals, as will be more fully exemplified later with the individual groups. Here it is sufficient to call attention to the relations in the heterologous series. In the formation of the hydrogen compounds of the elements of the first two periods, the following quantities of heat are set free according to present data ;— (C.HJ (N,H3) (O.H^) (F1,H) 21.7 II.9 57.2 — (Si,H,) (P,H3) (S,H3) (Cl.H) 33.2 116 4.5 22.0 In the halogen compounds the heat modulus is more regular :— (Li,CI) (Be.ClJ (B,Cl3) (QCIJ (N,Cl3) (0,013) (F1,C1) 93.8 — 104 28.3 —38.1 —18 — (Na,a) (Mg,Cl3) (Al,Cl3) (Si,ClJ (P.Cl,) (S.Cl^) (€1,01) 97.7 151.0 160.9 157-6 75-3 — — With the bromides, it is less throughout, and the least with the iodides. Con¬ sequently, a maximum appears to lie in the middle of the periods. However, on calculating the thermal value, which corresponds to one equivalent of the elements (united with one equivalent of chlorine), we obtain numbers that diminish successively and correspond to the decrease of the basic metalhc character of the elements :— (Na,01) ^Mg.Cl^-^ ^Al.013^ ^Si,01^^ ^P.Cla^ ^S.Ol, ^ (0,01). 97-7 75-5 53-6 3^4 25^1 Perfectly similar relations are furnished by the oxides :— (Na3,0) (Mg,0) (A1„0,) (Si,03) {VM (SA) (QaA)- (100.2) 145.8 388.8 219 363 104 REDUCTION OF METALLIC OXIDES. 265 Calculated upon one equivalent, the thermal value is :— 50.1 72.9 64.6 54.7 36.3 17.3 — That the heat disengaged in sodium oxide is less than that in magnesium oxide, depends partly upon the solubility of the first, as this property is also to be included as a thermal function. A like diminution of the heat modulus is also seen with the heterologous elements in their compounds of similar type :— (Mn.Cla) (Fe, CI2) (Co, Cl^) (Ni, CI3) (Cu, Cl^) (Zn. CI j 111.9 82 76.4 74.5 51.6 97.2 (Mn, O) (Fe,0) (Co, O) (Ni, O) (Cu, O) (Zn, O) 94.7 68.2 63.4 60.8 37.1 85.4 The following series are also noteworthy (Ag,Cl) {Cd,C\^) {In, a,) (Sn, CIJ (Sb, CI3) (Te, CI3) 29.3 93.2 127.2 87 (Ag2,0) (Cd,0) (103,03) (Sn,03) (Sb3,03) (Te, O) 5.9 65.6 133.5 THE REDUCTION OF METALLIC OXIDES BY METALLIC MAGNESIUM. Winkler has made interesting observations upon the deportment of metallic magnesium toward the oxides, hydroxides, and carbonates of the metals when heated together and in the presence of hydrogen. His results are briefly as fol¬ lows :— Magnesium heated in contact with the oxides, hydroxides, and carbonates of the alkali metals causes their reduction to the metallic condition; csesium alone is excepted. The reducing action loses in intensity as the atomic weights of the metals rise. With the oxides of copper, silver, and gold the intensity of reduction increases as the atomic weight rises. The oxides and hydroxides of beryllium, magnesium, calcium, strontium, and barium are also reduced, but the metals are not volatilized. When the reduction occurs in an atmosphere of hydrogen, hydrides of these metals result. The pro¬ duction of the hydrides of beryllium and magnesium is not very complete. Cal¬ cium hydride is more readily produced and in considerable quantity, while the union of hydrogen with strontium and barium is more easily effected and the yield of product much greater. These hydrides are all earthy masses, without lustre. The oxides of zinc, cadmium, and mercury are reduced with much energy. The heat of reduction causes either a partial or entire volatilization of the correspond¬ ing metals. Magnesium heated together with boron trioxide and borax yields a magnesium boride of varying composition. Aluminium oxide is converted into a product that contains a monoxide of aluminium. Yttrium and lanthanum are both apparently reduced ; they also combine with hydrogen to form hydrides. The reduction of the oxides of gallium, indium, and thallium by magnesium is very violent ; it increases with increasing atomic weight. Magnesium has a decided preference for carbon, as is evident from its deport¬ ment toward carbon dioxide, carbon monoxide, and the carbonates. With silicon, 25(5 INORGANIC CHEMISTRY. magnesium at low temperatures forms magnesium silicide, while at an intense heat the products are amorphous silicon and magnesium oxide. Zirconium dioxide, heated together with magnesium in an atmosphere of hydro¬ gen, yields a product black in color, and it may be that a zirconium hydride does exist. Titanium dioxide is reduced to monoxide with the simultaneous production of magnesium titanate. The hydride is apparently not formed. Cerium dioxide is reduced to metallic cerium or to the sesquioxide. If the reaction occtirs in hydrogen cerium hydride is formed. Thorium dioxide is reduced to metal, and in the presence of hydrogen its hydride results. It would thus appear that all the members of the principal division of the quad¬ rivalent elements are capable of yielding hydrides, although it is not so evident with titanium. More recently Seubert and Schmidt have shown that metallic magnesium is capable of removing the chlorine from metallic chlorides in neutral aqueous solu¬ tion ; the exceptions, in this case, being the chlorides of the alkali and alkaline earth metals. At elevated temperatures, however, all chlorides are reduced by magnesium. THE METALS. Although there is no sharp line of demarcation between metals and non-metals, yet these two classes of bodies form a distinct contradiction in their entire deportment, as may be plainly seen in the periodic system of elements. In physical respects the character of metals is determined by their external appearance and by their ability to conduct heat and electricity ; chemically, it shows itself chiefly in the basicity of the oxygen compounds ; yet we see that with the increase of the number of the oxygen atoms, the basic character gradually diminishes and becomes acidic. PHYSICAL PROPERTIES OF THE METALS. At ordinary temperatures all the metals excepting mercury are solid, slightly volatile bodies. They are opaque, and only a few, like gold, permit the passage of light to a limited extent when beaten into thin leaflets. In compact mass they exhibit metallic lustre and mostly possess a whitish-gray color ; gold and copper are, however, brilliantly colored. In powder form almost all the metals are black. Most of them crystallize in the forms of the regular system ; only a few, showing a metalloidal character, are not regular. Thus antimony and bismuth crystallize in the hexago¬ nal system, and tin is quadratic. The speciflc gravities of the metals vary greatly—from 0.59 to 22.4 as seen from the following arrangement :— Lithium, Potassium, Sodium, Rubidium, Calcium, Magnesium, Aluminium, Barium, Germanium, 059 0.86 0.97 1-52 1.58 »•75 2.56 375 5-47 Arsenic, Antimony, Zinc, Tin, Iron, Cobalt, Cadmium, Copper, Bismuth, 5-67 6.7 7» 7-3 7-8 8.5 8.6 8.8 9.8 Silver, Lead, Palladium, Thallium, Mercury, Gold, Platinum, Iridium, Osmium, 10.5 11.4 »»•5 11.8 »3-59 »9-3 2Ï.5 22.4 22.4 In general the specific gravities of the metals, and also those of the metalloids, increase with the atomic weights ; they stand more especially in sharp periodic dependence with reference to the latter. 22 257 2 ">8 INORGANIC CHEMISTRY. The first members of all periods possess low specific gravities; the latter grow gradually until the middle of the period, when the maximum is attained, and then they again decrease (p. 247). These relations show themselves more fully if, instead of the specific gravity, we compare the specific volumes or atomic volumes ; i. c,, the quotients from the atomic weights (A) and specific gravities (d):- ^ = specific volume. These quotients express the relative volumes of the atoms (in solid or liquid state). Thus the atomic volume of lithium (tf-W) = 11.9, that of potassium = 45.4; /.° Cadmium, 315° Palladium, I5CX5® Lead, 334° Platinum, 1780° Zinc, 423° Iridium, 1950° Aluminium, 750° A greater volatility also corresponds with the greater fusibility. Mercury boils at 360° ; potassium and sodium about 440° ; cad¬ mium at 860° j zinc toward 1000°, and the difficultly fusible metals may also be volatilized by the galvanic current. All these physical properties bear a periodic dependence to the atomic weights, as will be more plainly indicated in the individual groups. 260 INORGANIC CHEMISTRY. SPECIFIC HEAT-ATOMIC HEAT. Of all physical properties of the elements, from a chemical standpoint, their heat capacity is the most important, as it can serve for the determination of the atomic weights. To heat one and the same quantity, by weight, of the different metals or sub¬ stances to one and the same temperature, would require very dif¬ ferent amounts of heat. This is evident from the following experiments. If we add to i kilogram of HjG at o° i kilogram of H2O at 100°, the temperature of the mixture of 2 kilograms of water is 50°. The quantity of heat necessary to raise i part, by weight, of H2O, 1°, is almost the same for all temperatures from 0-100°; this is designated the am/ or calorie. On bringing to I kilogram H2O at 0° i kilogram Hg at 100°, the temperature of the water and of the mercury after their compensation, will equal only 3.2°. Consequently, the mercury has cooled about 96.8° (from 100 to 3.2°), and given off 3.2 calories. The quanti¬ ties of heat, contained in equal parts, by weight, of water and mercury, therefore, areas 96.8 to 3.2, i.e., the specific heat of mercury (that of water being made = i) is = 0.0332. On comparing the specific heats of solid elements found in this way with their atomic weights, we discover that these are inversely proportional to the latter, and hence the product of the specific heat and atomic weight for all solid elements (few excepted) is a constant quantity. This fact was first discovered by Dulong and Petit (1819), and formulated in the following law: The solid ele¬ ments possess the same atomic heat. In the table on p. 261 are presented the specific heats of the elements in solid condition (as far as they have been determined). W represents the specific heat, A the atomic weight, and the product, IV y. A, the atomic heat. From the table, it is evident that the atomic heats of most of the elements lie between 5.0 and 6.8, and equal, upon an aver¬ age, 6.4. It is only in the case of a few elements that the atomic heat is somewhat less (S, P, Si, Al, Ge), or considerably less (C, B, Be), than the mean. They are such as have low atomic weight, a metalloidal character, and occupy the middle of the two small periods. These variations bear distinct periodic dependence to the atomic weights:— Li Be B C N 0 Fl 6.6 3-8 2.8 1.9 — — — Na Mg Al Si P S Cl 6.7 5-9 5-5 4.6 5-8 5-7 — SPECIFIC HEAT. 261 Elements. W 1 A wx A Hydrogen* 5,880 1 5-9 Lithium Li 0,941 7.6 6.6 Beryllium Be 0,408 9.1 3-8 Boron (amorphous) B 0,254 10.9 2.8 Graphite ) c 0,174 \l2 2.1 Diamond f 0,143 / 1-7 Sodium Na 0,293 23 6.7 Magnesium Mg 0,245 239 5-9 Aluminium Al 0,202 27-3 5-5 Silicon (cryst.) Si 0,165 28 4.6 Phosphorus (yellow) P 0,189 3* 5-9 Sulphur (rhombic) S 0,178 32 5-7 Potassium K 0,166 39 6-5 Calcium Ca 0,170 39-9 6.8 Chromium Cr 0,100 52.4 5-2 Manganese Mn 0,122 54-8 6.7 Iron Fe 0,112 55-9 6.3 Cobalt 0,107 58.6 6.3 Nickel Ni 0,108 58.6 6.4 Copper 0,093 63.2 5-9 Zinc 0,093 64.9 6.1 Gallium ....Ga 0,079 69.8 5-5 Germanium Ge 0,057 72.3 5-4 Arsenic (cryst.) 0,082 74-9 6.2 Selenium (cryst.) 0,080 78.9 6.4 Bromine (solid) Br 0,084 79-7 6.7 Zirconium Zr 0,066 90 6.0 Molybdenum Mo 0,072 95-8 6.9 Ruthenium Ru 0,061 103 6-3 Rhodium Rh 0,058 104 6.0 Palladium Pd 0,059 106.2 6-3 Silver Ag 0,056 107.6 6.0 Cadmium 0,054 111.9 6.0 Indium 0,057 "3-4 6.5 Tin 0,054 117.5 6.5 Antimony 0,052 119.6 6.2 Tellurium Te 0,047 *25 6.0 Iodine 0,054 126.5 6.8 Lanthanum 0,045 139 6.2 Cerium 0,045 140 6.2 Didymium 0,045 142 6.5 Tungsten W 0,033 184 6.1 Osmium 0,031 *95 6.2 Iridium Ir 0,032 192.5 6-3 Platinum Pt 0,032 *94-3 6.4 Gold 0,032 196.2 6.4 Mercury (solid) Hg 0,032 200 6.4 Thallium 0,033 203.6 6.8 Lead 0,031 206.4 6.5 Bismuth..., 0,030 207 6-5 Thorium 0,027 232 6.4 Uranium 0,027 240 66 *A8 Palladium hydride. 262 INORGANIC CHEMISTRY. The variations from the mean are in part explained by the fact that most of the elements, in their different modifications (crystal¬ line, amorphous, malleable), possess a somewhat different heat capacity, as observed with carbon. The influence of temperature is, however, more important. The figures in the table mostly in¬ dicate the heat capacities at medium temperatures. It was known before that these show a slight increase with the temperature, but it is only recently, that H. E. Weber has proved that the increase is very considerable for the elements, C, B and Si, which, at me¬ dium temperatures, possess a remarkably low atomic heat ; that, beyond a definite temperature, the atomic heat becomes tolerably constant, and then almost agrees with the law of Dulong and Petit. According to Nilson, beryllium shows a similar deportment :— W A WXA Diamond, graphite, above 600° 0.459 11.97 5-4 Boron, above 600° 0-5 10.9 5-5 Silicon, above 200° 0.203 28 5-6 Beryllium at 257° 0.58 9.1 5-2 It is probable that there is a definite temperature for all elements, at which their heat capacities can be compared with accuracy. From this close agreement of the found atomic heat of the metals with the mean, it follows, without doubt, that there does occur a regularity, and we must conclude that the slight variations, apart from the inaccuracy of the observations, are influenced by second¬ ary causes. Hence, the specific heat may serve for the derivation of the atomic weight of the elements ; the atomic weight is equal to the constant quantity, 6,4 divided by the found specific heat :— The atomic weights derived from the specific heat—the so-called thermal atomic weights—agree in almost all instances with those ob¬ tained from the vapor density of the free elements or their volatile compounds. Where no volatile compounds of an element are known, the specific heat is the only certain means of fixing the actual atomic weight. The equivalent weight, 37.8 (InCl), of indium is fixed with great accuracy by analysis ; it is, however, unknown whether the atomic weight is double or triple that quantity. The specific heat of indium is 0.0569, from which the atomic weight would be 07^5^-5 = 112.5—a number clqsely approaching the trebled ISOMORPHISM. 263 equivalent weight of indium, 113.4 (= 37.8 X 3)» From this it fol¬ lows that the true atomic weight of indium is 113.4 and that indium is trivalent (InClj). In their solid compounds the elements retain the specific heat attaching to them in their free, solid state; hence the molecular heat is nearly equal to the sum of the atomic heats of the elements constituting the molecules—law of Neuman and H. Kopp. Hence the atomic heat of elements not known in solid condition may be derived from the molecular heat of their compounds. In this manner the fol¬ lowing atomic heats are found: For nitrogen, 5.0; for chlorine, 5-9; for oxygen, 4; for fluorine, 5 ; for hydrogen, 2.3. In the free state the gaseous elements usually have a slighter atomic heat, as seen from the following table :— A w* AX W Oxjrgen 16 0.156 2-5 Hydrogen 1.003 2.405 24 Nitrogen 14.041 0.172 2.4 Chlorine 35-45 0.093 3-3 The law that the atoms in solid condition possess the same ther¬ mal capacity (^A. W. =A/ W.'), finds an interesting analogy and exemplification in the results derived from the kinetic gas theory, and in the proposition of Avogadro, that the molecules in gas con¬ dition, at like temperatures have a similar degree of motion (ü/i v. = M'. z/.), and that the latter experiences like increases. The mol¬ ecules are the smallest particles for gases, and the atoms the small¬ est parts of the solid, which possess the same heat energy. The velocity of their heat motion, both for the molecules and the atoms, is, therefore, greater the smaller their masses. ISOMORPHISM. As indicated in the preceding pages, the atomic weights of the elements may be derived directly from the heat capacity of solids, while from the gas density of the volatile compounds we get the molecular weights, and from the latter, indirectly, ascertain the atomic weights (compare p. 80). A third, although less general and certain means of determining the atomic and molecular weights is afforded by isomorphism. By this is understood the phenomenon observed by Mitscherlich (1819), that bodies chemi¬ cally similar possess the same or almost the same crystal form. An essential mark of isomorphous bodies is their ability to crystallize together—to form so-called isomorphous mixtures. Conversely * By constant volume. 264 INORGANIC CHEMISTRY. from the isomorphism of two compounds may be concluded an analogous chemical composition, an equal number of atoms in the molecule. This would lead us to accept as relative atomic weights, those quantities of the elements which replace each other in iso- morphous compounds. For example, the metals calcium, stron¬ tium and barium do not afford volatile derivatives. Their atomic weights could not be deduced from their thermal capacity, and it was the isomorphism of many of their compounds with those of magnesium that determined the same ; the quantities of these ele¬ ments, replacing 24 parts by weight of magnesium (i atom), were accepted as the true atomic weights. In the present state of chemistry we attach but secondary import¬ ance to isomorphism as a method of determining atomic weights. The phenomena of pleomorphism, according to which one and the same substance frequently possesses several crystalline forms, teach us that the latter are not only dependent upon the chemical mole¬ cules, but that these (according to yet unknown laws) may unite to more complicated crystal molecules. Hence, isomorphism affords a means for determining the molecular value of solid sub¬ stances. On the other hand we know of many cases where compounds chemically dissimilar possess a similar isomorphous crystalline form. Thus, dimorphous calcium carbonate (CaCOj), as calcite is iso¬ morphous with sodium nitrate (NOsNa), while as aragonite it is isomorphous with potassium nitrate (KNO3). Consequently, iso¬ morphism is only to be applied with care in chemical conclusions. Yet it is generally seen that bodies chemically similar have like crystalline forms, especially if the similarity of the elements be taken into consideration according to groups, as expressed in the periodic system. Thus, the isomorphism of the sodium com¬ pounds with the silver and cuprous derivatives, of the permanga¬ nates with the Perchlorates (CIO4K), of the chromâtes with the sul¬ phates (SO^Naj), confirms the relations presented in the periodic system. Details upon this will be noticed in the consideration of the individual groups. CHEMICAL PROPERTIES OF THE METALS. As a usual thing the metals combine, without difficulty, with the metalloids, and with them yield well-characterized compounds, the properties of which are essentially different from the elements com¬ posing them. The greater the chemical difference of two bodies (metals and non-metals, bases and acids) the more energetic, in general, is their tendency to unite, and the more different and CHEMICAL PROPERTIES OF THE METALS. 265 rtiore stable the resulting products. As we have seen, the anal¬ ogous metalloids (the groups of chlorine, of sulphur) form deriva¬ tives with each other that are not very characteristic. In the same manner when fused together the metals form indefinite metal-like compounds, known as alloys. Alloys, for the solid condition, are essentially the same as solu¬ tions for the liquid. Solutions and alloys constitute the transition from mechanical mixtures to the real chemical compounds. In both instances the constituents possess but a slight aifinity for each other, and, therefore, unite in almost all proportions to the so- called undetermined compounds (see p. 92). We, however, know that definite compounds frequently exist in solutions ; thus in an aqueous solution of sulphuric acid there is present the hydrate HJSO4.2H2O j in aqueous nitric acid the hydrate HNO3.H2O. And in the solutions of the salts crystallizing with water of crystalliza¬ tion there are definite compounds with water (Cu PO,.|NH, SO,/ H SO, {: '4 Pot. am. phosphate. Pot. copper sulphate. Pot. aluminium sulphate. CI Pb/ rNO, ^CO. Cr, Pb/ ' Mci ^Cl f 3 NO 3 '4 The halogen double salts are usually viewed as molecular com¬ pounds :— MgClj.KCl AuClg.KCl RC1,.2KC1. If, however, the fluorides of boron and silicon, BFlj.KFl, SiFl,. 2KFI, are derived from peculiarly constituted atomic acids, HBF1„ HjSiFle, then a peculiar union of atoms may be regarded as existing in the metallic double chlorides,* which are often very similar and isomorphous. * Consalt American Chemical Journal.^ 11, p. 291. 270 INORGANIC CHEMISTRY. ACTION OF METALS UPON SALTS AND ACIDS. We have seen that the metals by solution in acids are able to form salts. In this case the hydrogen is directly replaced by the metal and separated in free condition (providing in the moment of its formation it does not act upon the acids) :— Zn -j- SO4H2 = ZnS04 -j- Hg. The metals deport themselves in the same manner with the salts. Zinc introduced into a solution of copper sulphate is dissolved to sulphate and metallic copper deposits:— Zn -j- CUSO4 = ZnSO^ -j- fcu. Herein is shown the perfect analogy between acids and salts. In chemical nature hydrogen is a metal. Hence the acids may be viewed as hydrogen salts : hydrogen sulphate for sulphuric acid, hydrogen nitrate for nitric acid, etc. The similarity of salts and acids shows itself, too, in their acidity. All soluble salts of the metals, whose hydroxides are weak bases, exhibit acid reaction, and color blue litmus paper red. Only the salts of the strong basic metals, like potassium and calcium, show a neutral or basic reac¬ tion—providing the base is stronger than the acid. The displacement of metals from their salts by others, was formerly regarded as exclusively influenced by their electrical deportment. Indeed the more electro-positive, basic metals replace the electro-negative, less basic. In the following series each metal throws out from solution those preceding it : Au, Pt, Ag, Hg, Cu, Pb, Sn (Fe, Zn). Iron and zinc precipitate almost all the heavy metals from solutions of their salts. The strongly positive potas¬ sium is able to displace all other metals. This is very evident from the action of molten potassium upon the haloid salts—a reaction which frequently serves for the separation of the metals in a free condition :— AICI3 + 3K = Al 4- 3KCI. In its electrical deportment hydrogen stands near zinc ; like the latter, it must, therefore, displace all more negative metals. If this does not happen, the cause must be sought in its volatility ; in fact, we know that hydrogen, under pressure, is capable of sepa¬ rating gold, silver and some other metals from their salt solutions. Formerly great importance was attributed to the electrical behavior of the ele¬ ments, and all were arranged in an electro-chemical series, in which oxygen figured as the most negative and potassium as the most positive member— O + K. The opinion prevailed that the chemical affinity of the elements depended upon their electrical differences, and that chemical union occurred because the opposite electricities united—electro-chemical theory of Berzelius. Now, however, we know that in the expression of chemical affinity only secondary ELECTROLYSIS OF SALTS. 271 importance is attached to the electrical deportment of bodies. Although the affinity, in general, corresponds to the electrical difference, yet this does not always occur. Thus, the strongly negative chlorine expels bromine and iodine from their hydrogen, and nearly all their metallic compounds ; however, chlorine and bromine are conversely displaced by iodine from their oxygen compounds (CIO3H and CIO4H) (p. 186). Similarly, lead separates tin from its chloride, SnCl^, while, on the other hand, tin throws out lead from the solution of its oxides in alkalies. At present, it is established that the mutual deportment of the metals is depend¬ ent upon and regulated by their thermo-chemical relations. A metal displaces another from its oxygen salts, as well as from its oxides, sulphides, or halogen compounds, if the heat of formation of the resulting bodies is greater than that of those acting ; this agrees with the principle of greatest heat development. Thus, copper displaces silver from its sulphate, because the heat of formation of the cop¬ per sulphate (in aqueous solution) is about 33.5 calories greater than that of silver sulphate. Sulphuric acid dissolves most metals with liberation of hydrogen, because their heat of formation, (8,04,112) = 192.9 (S,04,H2,Aq.) = 210.7, is less than that of most of the sulphates. The heat of formation of lead sulphate (Pb,S,04) equals 213 5; therefore, lead would be dissolved by dilute sulphuric acid, did not the insolubility of lead sulphate in the dilute acid prevent it from so doing. Concentrated sulphuric acid, on the other hand, does dissolve lead, because lead sulphate is soluble in it. For the same reason, potassium dis¬ places almost all the other metals; on the other hand, potassium is separated by sodium amalgam, with formation of potassium amalgam, as the heat of formation of the latter is much greater than that of sodium amalgam, and therefore, in the equation, (K,C1) + (Na,Hg) = (Na,Cl) + (K,Hg), the thermal value upon the right side overbalances. (Berthollet.) Although the affinity relations dependent upon the quantity of heat frequently correspond with the electrical differences of the free elements, this is so influenced that the electro-motive energy is induced by the heat, and is proportional to the same (see p. 274). The heat of formation of the compounds constitutes the primary cause of their chemical transposition; it varies in the different compounds for the same element, and thus explains the opposing deportment of the elements. Chlorine displaces iodine in iodides, not because it is more strongly electro-nega¬ tive, but because the heat of formation of the chlorides is greater than that of the iodides. Conversely, chlorine is eliminated from chloric acid by iodine, because the heat of formation of the iodic acid is the greater (compare p. 186). HgSand I, are similarly transposed, in the presence of water, into HI and sulphur, while iodine is separated from concentrated hydriodic acid by boiling with sulphur (P- 67). ELECTROLYSIS OF SALTS. On subjecting a salt in a fused or dissolved condition to the action of an electric current, it is decomposed, so that the metal separates upon the negative pole and the acid group or halogen upon the positive ;— + — NaCl = Na -f CI. 272 INORGANIC CHEMISTRY. The oxygen salts behave in the same way ; the metal upon the neg¬ ative pole, the acid residue upon the positive :— + — CuSO^ = Cu -j- SO4. As the liberated acid residue cannot exist in a free condition, a sec¬ ondary reaction occurs, by which it generally, especially in the electrolysis of aqueous solutions, breaks up into oxygen and an acid oxide, which, with the water of the solution, again forms the acid :— SO4 + H2O = SO4H2 + O. Thus, in the electrolysis of salts, the metal and oxygen separate out—the first at the negative, the latter at the positive pole. That the decomposition, indeed, occurs in the manner indicated is con¬ firmed by the fact that the free acid arises at the positive pole. All neutral salts are similarly decomposed. If, however, the metal contained in the salt acts upon water when free, manifestly a secondary reaction must occur at the negative pole. The real elec¬ trolytic decomposition of potassium sulphate would then take place according to the following equation :— ~f- — SO4K2 = Kg (SO3 -|- O). The separated potassium decomposes the water with formation of potassium hydroxide and the disengagement of hydrogen :— K -f HOH = KOH + H. Therefore, hydrogen and potassium hydroxide occur as definite decomposition products, at the negative pole ; at the positive, how¬ ever, we have oxygen and sulphuric acid. On coloring the liquid exposed to the electrolysis with a little violet syrup, that part at the -f- pole will become red, owing to the acid formed, while that at the — pole will have a green color from the base. That the electrolytic decomposition of potassium sulphate and similar salts proceeds in the manner given, may be proved experimentally by using mercury as negative electrode ; then the separated potassium will combine with the mercury and form an amalgam, which will act gradually upon the water. It was formerly believed that the alkali salts were directly decomposed by elec¬ trolysis into metallic and acid oxides, which yielded the hydrates (KOH and SO4H2) with water; the appearance of H and O was attributed to the simulta¬ neous electrolytic decomposition of water (a view which was set aside by the behavior of the other salts). With this erroneous idea as a basis, all salts were held to be binary compounds of the metallic oxides (bases) with acid oxides (acids), e. g., KgO.SOg = KjSO^, KJO.N2O5 = 2KN0j — duaUstic theory of Ber- ELECTROLYSIS OF SALTS. 273 zelius. The acids and bases were also thought to be binary compounds of a metallic oxide or acid anhydride with water :— KjO.HaO = 2KOH. SO3.H2O = H2SO4. The acid oxides or anhydrides were termed acids and the true acids hydrates. Other compounds are decomposed in the same way as the salts. Thus, molten caustic potash, KOH, breaks up into K and OH ; the first separates in metallic form upon the negative pole (and gradually acts upon KOH with hydrogen disengagement), while at the positive pole water and oxygen appear—produced by decom¬ position of the hydrogen peroxide formed at first :— (0H)2 = HjO -h O. Acids sustain decompositions similar to those of salts. They are, indeed, nothing more than hydrogen salts. Hydrochloric acid, for example, breaks down into hydrogen and chlorine. Sulphuric acid, H2SO4, is decomposed in aqueous solution into the ions 2H and SO4 or H and SO4H. The anion SO4 (or SO4H) is immediately converted by the water into sulphuric acid and oxygen : SO4 -j- OH2 = SO4H2 + O. The final decomposition products are, there¬ fore, hydrogen and oxygen. Sulphur heptoxide, S2O7 ; 2SO4 = S2OT -1- O, is formed in minute quantities. It is, therefore, probable that the water is also decomposed in an analogous manner :— 2HOH = H2 + OjHj; the peroxide produced at first breaks up, however, for the most part, into water and oxygen. Considering the quantities which are deposited from various compounds by the same electric current, we will discover that a like number of valences are invariably dissolved in like time, i, e., equivalent quantities are separated according to the idea of the valence theory (p. 171). (The law of Faraday and Becquerel.) Thus in the simultaneous decomposition of hydrochloric acid, water and ammonia (pp. 74, 99, 130), equal volumes of hydrogen (= I part) are liberated, while at the positive pole i volume of chlorine (=35.45 parts), volume of oxygen (=8 parts) and Yl volume of nitrogen (= 4.67 parts) appear. The quantities decomposed by electrolysis, therefore, bear the ratio :— HCl, íísN 2 3 In the same way, equal quantities of chlorine are set free from all metallic chlorides (and other salts, as the chlorine atoms are alike in all), while the quantities of the precipitated metals agree with the values according to which they enter chemical action. The quantities of the different salts, decomposed by electrolysis, stand in the following relation :— 274 INORGANIC CHEMISTRY. AgNO , 9^, 9^, SbCl, Fe,Cle SnCl, HgCl, Hg,(N03)3 ^'2' 2 ' 2 (>' 4 ' ^ 2 Therefore, 31.8 parts Cu are deposited for the 35.4 parts CI in cupric chloride (Cu"Cl2), but from cuprous chloride (Cu'Cl) we obtain 63.6 parts Cu; from mercuric chloride (Hg^'Clj) we obtain 100.2 parts Hg, and from mercurous nitrate (Hg'NOa) 200 parts Hg, etc. The quantities of the metals existing in the different states of oxidation, and which are equivalent to each other, vary and correspond to their chemical affinity. Like valences are dis¬ solved in equal periods of time. As the quantity of heat liberated in the union of like valences (in KCl, CuCl, AgNOj, etc.), and that necessary for the decomposition are very different, and since Faraday's law calls for the solution of like valences in equal periods by the same current—the performance of a different amount of work—^there must occur, in consequence of the law of the conservation of energy (if the electric current does indeed do the electrolytic work), an unequal distribution of the energy of the current upon the different electrolytes. The manner of this distribution is not known. Faraday thought it probably took place in such a manner that the greater consumption of energy by the electrolytes was compensated by their resistance, which, consequently, is so much less. The relation of the electro-motive force of a galvanic cell to the chemical trans¬ position occurring within it is equally obscure. Disregarding the electric contact theory, the source of the former (with the principle of the conservation of eneigy as basis) can only be found in the loss of chemical energy (heat disengagement), which corresponds to the change taking place within the galvanic element- Fol¬ lowing the experiments of Joule and others, W. Thomson assumed that the energy, developed by a cell, was proportional or equal to (providing no sec¬ ondary actions occurred) the thermal value of the chemical reaction producing it. Thus, the effectiveness of a Daniell element (combination of zinc and dilute sulphuric acid, or zinc sulphate and copper in a solution of copper sulphate) de¬ pended on the replacement of Cu in CuSO^ by Zn—a reaction in which 50.1 calories are developed. Further, the action of a Bunsen cell (zinc and car¬ bon in a solution of potassium bichromate and sulphuric acid) was supposed to be due to the formation of zinc and chromium sulphates, in which instance 99.8 calo¬ ries are set free. In all such constant batteries J. Thomson and others concluded that the electro-motive force is equal or proportional to the energy developed in the chemical reaction. According to a later theory of Helmholtz, and the deter¬ minations made by F. Braun, chemical energy cannot be completely, but only partly, transformed into electric energy, just as heat cannot be completely changed to mechanical work. Hence, the electro-motive force of an element is usually less, but should be greater than the heat energy corresponding to the chemicsd transposition. In the latter cases, we must renounce any relation of the electro¬ motive force to the thermal value, in case these, as is probably true, cannot be explained by secondary chemical reactions which have gone unconsidered heretofore. Helmholtz asserts that the entire chemical enei^ is only completely converted into electric energy when the electro-motive force of the element is independent of the temperature. If the electro-motive force increases with the temperature, then heat is developed in addition to the chemical energy, and is withdrawn from the element. In such cases the electro-motive force is greater than the heat value of the chemical transposition. Jahn has confirmed these statements by very recent experiments. TRANSPOSITION OF SALTS. 275 When two salts in solution or fusion come together, a chemicál action will frequently occur. Berthollet endeavored (close of pre¬ ceding century) to explain the resulting phenomena by referring them to purely physical causes, and excluded every special chemical affinity. In the opinion of Berthollet, four salts always arise in the solu¬ tion of two. For example, on mixing solutions of copper sulphate and sodium chloride, there exist in solution copper sulphate, sodium sulphate, copper chloride, and sodium chloride :— 2CuS0^ -(- 4NaCl yield CUSO4 + Na^SO^ -1- CuClj -|-2NaCl. That copper chloride is really present in the solution together with the sulphate, follows, from the fact that the blue color of the latter acquires a greenish color, peculiar to the copper chloride, by the addition of sodium chloride; other phenomena are not notice¬ able at first. Suppose one of the four salts formed in the solution is insoluble or volatile, the reaction will occur somewhat differently. Upon adding barium chloride to the copper sulphate solution four salts will be formed at the beginning just as in the first case. The barium sulphate produced separates, however, in consequence of its insolubility, the equilibrium of the four salts will be disturbed, and new quantities of CuSO^ and BaCh act upon each other until the transposition is complete :— CuSO^ BaClj = BaSO^ -|- CuClj. The chemical transposition may, therefore, be explained by the in¬ solubility of the barium sulphate. On adding HCl, or soluble chlorides, to the solution of a silver salt all the silver is precipitated as chloride, because the latter is insoluble. Take another example. On adding sulphuric acid to a solution of potassium nitrate there is apparently no perceptible alteration. We may suppose that the four compounds, KNO3, K2SO4, H2SO4 and HNO3, are present in the solution. Upon warming the latter volatile nitric acid will evaporate, and, in proportion to its separa¬ tion, new quantities of potassium nitrate and hydrogen sulphate will act upon each other until the transposition is complete :— 2KNO3 -f H2SO4 = K2SO4 2HNO3. The decomposition of potassium nitrate by sulphuric acid occurs, therefore, in consequence of the volatility of the nitric acid. Sul¬ phuric acid decomposes sodium chloride in the cold, because hydro¬ gen chloride is volatile. Carbonates are even decomposed by very weak acids, because the carbonic acid, H^COs, at once separates gaseous carbon dioxide, COj. In many instances the chemical transpositions may be explained by such physical causes, and there is no doubt that an important 276 INORGANIC CHEMISTRY. rôle attaches to them. It is, however, not justifiable to ignore any special chemical aifinity between the various substances, as did Berthollet. Irrespective of all physical causes, the reactions are determined by chemical affinity. This is seen in the solutions of salts. Mix, e. g., ferric chloride with potassium acetate, and there is obtained a dark-red solution, in consequence of the formation of iron acetate. Although an insoluble salt is not pro¬ duced here, yet the rearrangement of the two salts, evident from the optical properties of the solution, is a perfect one ; only iron acetate and potassium chloride are present in the solution :— Fe^Cl, -f 6C2H3O2K = (QHjOJeFe, + ÓKQ. Pot. acetate. The transposition is determined by the strong affinity of potas¬ sium for chlorine and by the weak basic nature of the ferric oxide. If the difference between the affinities of the bases and salts is not so great, then four salts can exist in solution ; their quantity, how¬ ever, will be proportional to the different affinities and determined by the equilibrium of all the forces of attraction. Thus four salts are present in the previously mentioned solution of copper sulphate and sodium chloride, the quantities of copper chloride and sodium sulphate are, however, much greater than those of copper sul¬ phate and sodium chloride (proved by the optical properties of the solution), because the affinity of sulphuric acid for sodium is greater than the same for copper. The phenomena recorded above are the subject of controversy and many in¬ vestigations at the present time. Generally the chemical transpositions of salts with salts or with acids (hydrogen salts) and bases, are determined and governed by the tendency (Bestreben) toward the greatest heat disengagement. This is true, too, of other chemical reactions (the precipitation of metals from their salts by other metals, the action of metals upon acids and water, the alternating deport¬ ment of metalloids, etc.). In most cases the reaction corresponding to a chemical equation is easier and more complete, the more the sum of the heats of formation of the resulting bodies exceeds that of those reacting (the chemical energy of the first is less than that of the latter). The reverse of the reaction can only succeed by the consumption of energy and requires addition of heat or electricity. This is of practical importance (see Magnesium Chloride) in the review of the details of chemical reactions. In the action of salts and acids these relations are, how¬ ever, complicated because many reactions, especially in aqueous solution, occur accompanied by a direct absorption of heat. The aim of thermo-chemistiy, in ac¬ cord with the efforts of Berthollet, is to classify such reactions under the principle of the greatest heat disengagement, and declare its exceptions due to the influence of secondary causes. This would bring into consideration the formation of acid and double salts, the influence of the heat of solution, the decomposition of salts by solvents, etc. In this manner very many of the apparent exceptions have been satisfactorily accounted for. However, it is very evident, that the principle of greatest heat-development cannot have an unlimited value, and that all reactions are not of the class in which heat is set free, but that chemical transpositions can also proceed even when heat is absorbed. The decomposition (dissociation) of elements and compounds oppost^ their tendency to unite, and in such instances ALKALI METALS. 277 heat disappears as such, passing then into chemical energy. In most cases the greater heat formation of compounds argues for a greater stability, yet many com¬ pounds are known which have been produced with the liberation of much heat, but are themselves very unstable. In such cases the formation of the most stable derivative prevails, although it does possess the lowest heat of formation. The transpositions then follow with heat absorption in a direction directly opposite that presented by the principle of greatest heat disengagement. Thus, e.g., the heat of formation of water is greater than that of an equivalent quantity of hydro¬ gen chloride :— (H2,0) vapor = 57.2 Cal. 2(H,C1) = 44.0 Gal. Yet the latter is more stable than the former. The dissociation of water com¬ mences about 1000°, while that of HCl begins at 1500®. Consequently, with a mixture of hydrogen, chlorine and oxygen (in equivalent quantities), if the reac¬ tion be induced by application of heat or by the electric spark, the product will be hydrogen chloride exclusively and not water. Again, in a mixture of steam and chlorine, exposed to a temperature of about 1000®, a transposition will occur with heat absorption and result in the production of hydrogen chloride and oxy¬ gen. Heat is not the only agent which causes the decomposition of bodies ; many salts undergo changes in water and other solvents, quite analogous to that of dis¬ sociation, and the extent of these changes is not always limited to the heat of for¬ mation of the component elements. This accounts for the fact that very often transpositions occurring in dilute solutions are accompanied by the absorption of heat. It is therefore plain that the principle of the greatest heat development is sub¬ ject to a limitation, conditioned by the stability of the compounds (p. 30). GROUP OF THE ALKALI METALS. Potassium, 39-14 Lithium, 7.03 Rubidium, 85.4 Sodium, 23.06 Caesium, ^3^-9 (Ammonium). The metals of this group are decidedly the most pronounced in metallo-basic character, and this constitutes a visible contrast with the elements of the chlorine group, the most energetic among the non-metals. This contradictory character of both groups is seen, too, in their monovalence : in their combinations with each other, they saturate their affinity by single atoms. The more distinct the chemical character of two elements and the more unlike they are, the simpler and the more definite will the expressions of valence in general be between them. The alkali metals in physical and chemical properties exhibit great similarity. They oxidize readily in the air, decompose water violently, even in the cold, with the formation of strong basic hydroxides, which dissolve readily in water and are called alkalies (caustic potash, caustic soda),—hence the name alkali metal. They are not decomposed by ignition. Their chemical energy increases with increasing atomic weight (more correctly atomic volume, p. 258), sodium is more energetic than lithium, potassium more than 278 INORGANIC CHEMISTRY. sodium, and rubidium more than potassium. Caesium has not been studied in a free condition, but, judging from its compounds, it possesses a more basic character than rubidium. We saw in other analogous groups (of chlorine, oxygen, phosphorus, carbon), that the metalloidal, negative character diminishes, and the basic in¬ creases with the increasing atomic weight. The atomic weights increase simultaneously with the specific gravities ; but as the increase of the former is greater than that of the latter, the atomic volumes (the quotients p. 258), are always the greater. The increasing fusibility and volatility corre¬ spond to the increase of the atomic volumes ; rubidium distils a.t a red heat, while lithium only volatilizes with difficulty :— Li Na K Rb Cs Atomic weight 7-03 23.06 39-03 85.4 132.9 Specific gravity 0.59 0.97 0.86 1.52 1.85 Atomic volume 11.9 237 45.4 56.1 71.7 Fusion temperature 180° 95-6° 62.5® 38.5° 26.5® Although all the alkali metals exhibit a great similarity in their chemical deportment, we discover more marked relations between potassium, rubidium and caesium upon the one hand, and lithium and sodium on the other, which accords with their position in the periodic system of the elements. Especially is this noticed in the salts. The first three metals form difficultly soluble tartrates and chlorplatinates (see Platinum). Their carbonates deliquesce in the air, while those of sodium and lithium are stable under similar cir¬ cumstances ; the last is, indeed, rather insoluble in water. The phosphates deport themselves similarly ; lithium phosphate is very difficultly soluble. It must be remarked that the normal carbonates and phosphates of all other metals are insoluble. In lithium, then, which possesses the lowest atomic weight, it would seem the alkaline character has not yet reached expression, and it in many respects approaches the elements of the second group, especially magnesium, just as beryllium approaches aluminium ; this is indicated by the position of the elements in the table, p. 249. The elements of the two small periods are, indeed, similar, but not completely analogous, while the homology of the three great periods finds expression in K, Rb, Cs. See further p. 255. The affinity relations of the alkalies are expressed and explained by their thermo- chemical relations. Generally the heat liberation is greater as the atomic we^hts increase: thus, in the formation of the chlorides and hydroxides:— POTASSIUM. 279 (Li,a) (Na,Cl) (K,CI) (Na2.0) (K„0) 93-8 97-7 105.6 100.2 97.1 (Li,Cl,Aq.) (Na,Cl,Aq.) (K,Cl,Aq.) (Na2,0,Aq.) (K„0,Aq.) 102.2 915 lOI.I 155-2 164.5 (Li,0,n,Aq.) =117-4 (Na,0,H,Aq.) =111.8 (K,0,H,Aq.) = 116.4 (Na,0,H) = 102.0 (K,0,H) = 103.9 The position of lithium in the periodic system explains the varying deport¬ ment of its compounds, which frequently show a greater heat disengagement than those of sodium. Again, it is very probable that a constant increase in the heat modulus occurs with the true homologues of potassium—^rubidium and caesium. On the basis of the principle of the greatest evolution of heat, the numbers above would explain why sodium and lithium are displaced from their chlorides, etc., by potassium. It separates most other metals because the heat of formation of the potassium compounds is generally much greater (see p. 271). On compar¬ ing the heat of formation of water (HjO = 68.6 calories), we immediately per¬ ceive why it is so readily decomposed by the alkali metals. All metals, disengag¬ ing more than 68.6 calories in the formation of their oxides, MejO, or their hydroxides, MeOH, decompose water, and the energy will be greater, the greater the difference of heat. The insolubility of the oxides constitutes an obstacle to the action ; this, however, may be removed (see Aluminium) by addition of neutral solvents. Conversely, all oxides, affording less heat in their formation, are reduced by hydrogen. POTASSIUM. K - 39-14- In nature, potassium is found principally in silicates, viz. : feld¬ spar and mica. By the disintegration of these frequently occurring minerals, potassium passes into the soil, and is absorbed by plants ; the ashes of the latter consist chiefly of different potassium salts. The chloride and sulphate are also found in sea water, and in large deposits in Stassfurt, at Magdeburg, and in Galicia, where they were left by the evaporation of the water of inclosed seas. Metallic potassium was first obtained by Davy, in the year 1807, by the decomposition of the hydroxide, by means of a strong galvanic cur¬ rent. At present it is prepared by igniting an intimate mixture of carbon and potassium carbonate :— KjCOj + 2C = 2K -h 3CO. For demonstration purposes potassium can be prepared by heatii^ a mixture of dry potassium carbonate (l mol.) and magnesium powder (3 ats.) in a porcelain boat placed in a combustion tube in an atmosphere of hydrogen. The potassium volatilizes and forms a brilliant metallic mirror upon the tube. Such a mixture may be made by the carbonization of organic potassium salts, crude tartar. It is then ignited to white heat, in an iron retort, and the escaping potassium vapors collected in receivers of peculiar construction, filled with rock oil. The latter, an hydrocarbon, serves as the best means of preserving potassium, which would otherwise oxidize in the air, and decompose other liquids. In a fresh section, potassium shows a silver-white color and brilliant metallic lustre. At ordinary temperatures it is soft, like 280 INORGANIC CHEMISTRY. wax, and may be easily cut. It crystallizes in octahedra, and has a specific gravity = 0.86. It melts at 62.5° C., and when raised to a red heat, is converted into a greenish vapor. It oxidizes in the air, and becomes dull in color ; heated, it burns with a violet ñame. It decomposes water energetically, with formation of potassium hydroxide and the liberation of hydrogen. If a piece of the metal be thrown upon water, it will swim on the surface with a rotary motion ; so much heat is disengaged by the reaction that the gen¬ erated hydrogen and the potassium inflame. Finally, a slight explo¬ sion usually results, whereby pieces of potassium are tossed here and there; it is advisable, therefore, to execute the experiment in a tall beaker glass, covered with a glass plate. Potassium combines directly and very energetically with the halogens. On Conducting hydrogen over metallic potassium heated to 300®—4®, potas¬ sium hydride, KgH, results. This is a metallic, shining, brittle compound, which, upon stronger heating (410°), more readily in vacuo, is again decomposed. Exposed to air, it ignites spontaneously. The sodium hydride, Na^H, obtained in the same way, does not possess this latter property. The influence of heat and pressure in the formation and decomposition of these compounds is very noteworthy (p. 96). If, for example, potassium hydride be heated it melts, but otherwise remains unchanged. Above 200° (in a vacuum) it sustains a partial decomposition (dissociation), which gradually increases as the temperature rises. If the heating should take place in a closed vessel provided with a manometer, it will be observed that the decomposition at a given tempera¬ ture will continue until the liberated hydrogen has acquired a definite tension— until it exerts a definite pressure. For potassium hydride, this tension, at 330° C., equals 45 mm. The decomposition will then cease, but will proceed further at the same temperature if the hydrogen gas be removed, until the pressure of 45 mm. is again reached. In this manner a complete decomposition of the hydride may be effected at the temperature given above. If, however, the disengaged hydrogen is not removed, but be added to the completely or partially decomposed hydride, and the pressure be raised to 45 mm. (at the temperature 330°), the potassium hydride will be re formed. Consequently, both the decomposition and the formation of a body can follow, depending upon whether the external partial pressure be lowered or increased. This pressure is designated the tension of dissociation. Similar phenomena occur at higher temperatures, the correspond¬ ing pressure, of course^, increasing by regular steps. Appended are tensions of dissociation of potassium and sodium hydrides for different temperatures :—■ Temperature. Tension of Dissociation. KaH NajH 330® 45 mm. 28 mm. 350° 72 " 57 " 370° 122 " 100 " 390® 363 " 284 " 410® 736 " 594 " 430® HOC " 910 " POTASSIUM OXIDE—POTASSIUM HYDROXIDE. 281 The tension of dissociation is independent of the relative quantity of the disso¬ ciated body and of the space which the disengaged gas can occupy, whereas in solutions and absorptions (ammonia by charcoal) the pressure at one and the same temperature increases with the quantity of the absorbed gas. All exothermic compounds behave like potassium and sodium hydrides when they are decomposed into their components; if the pressure be raised above the tension of dissociation the components reunite—the compounds are re-formed. The decomposition of the endothermic compounds (potassium chlorate into chloride and oxygen) is quite different (pp. 29 and 94). It proceeds with heat dis¬ engagement (KCljOj =— 11.o Cal.), corresponding to the chemical affinities, and is only induced by application of external heat. It is independent of exter¬ nal pressure, and there is no reunion of the- decomposition products upon increas¬ ing the external pressure or upon lowering the temperature. These decomposi¬ tions are, therefore, quite different from the dissociation of exothermic compounds. We must also not omit mentioning the great analogy between the phenomena of dissociation and the vaporizing of liquids, which occurs in a similar manner. Like dissociated bodies, liquids exhibit at all temperatures a definite tension in vapor form ; the evaporation is conditioned by this. The evaporation will occur in a closed space by the lowering of the pressure, but if the latter be raised the vapors will condense. Consequently, we observe that the phenomena of dissociation reveal an intimate connection between the forces of chemical affinity and those of physical cohesion. Potassium forms three oxygen compounds, of which only the following yields corresponding salts. Potassium Oxide—KjO—results from the oxidation of thin pieces of metallic potassium in dry air, and by heating potassium peroxide with metallic potassium. It is a white powder, fusing at a red heat, and evaporating at higher temperatures. It unites with water, with evolution of much heat, and the formation of potassium hydroxide. When heated in a stream of hydrogen it yields the hydroxide, and metallic potas¬ sium is separated :— KjO -f H = KOH 4- K. This peculiar behavior is explained by the heat of formation of KOH (103.9 C.) being greater than that of KgO (97.1) ; hence the reaction occurs according to the preceding equation and heat is disengaged. Conversely, KOH cannot, therefore, be decomposed by potassium with the production of K2O (Beketoff ). Potassium peroxide^ KOj or K2O4, and potassium suboxide, K^O, are very unstable, and readily pass into potassium oxide. The first is formed together with potassium oxide, by the combustion of potassium in dry air or oxygen, and is a yellow mass. The suboxide has a violet color, due to the oxidation of potas¬ sium vapors. Potassium Hydroxide, or Caustic Potash—KOH—is obtained by the action of potassium or its oxide upon water. For its prepa¬ ration, potassium carbonate is decomposed by calcium hydroxide (slaked lime) ;— KjCOj -f- Ca(OH)j = CaCOj 4- 2KOH. 24 282 INORGANIC CHEMISTRY. The solution of i part potassium carbonate in 10-12 parts water is boiled with I part slaked lime in an iron p>ot, until a filtered portion does not effervesce when hydrochloric acid is added ; i. e., until there is no longer any qarbonic acid present. On standing awhile, the insoluble calcium carbonate subsides, and the liquid becomes clear. The solution of potassium hydroxide is then poured off, evapo¬ rated, the residue melted in a silver dish (which it does not attack), and poured into moulds. The caustic potash, prepared in this way, is not entirely pure, but contains potassium chloride and other salts. To obtain a product that is chemi- cally pure, fuse potassium nitrate with coppef filings, and treat the fusion with water. Potassium hydroxide forms a white, crystalline mass that fuses rather easily, and volatilizes undecomposed at a very high temper¬ ature. Exposed to the air it deliquesces, as it absorbs water and carbon dioxide and changes into carbonate. It is very soluble in alcohol, and especially in water. The solution possesses a strong alkaline reaction, saponifies the fats, and has a corrosive action upon the skin and organic tissues ; hence it cannot be filtered through paper. At low temperatures the hydrate KOH-f-zHjO crystallizes out from concentrated solutions. The haloid salts of potassium are obtained by the direct union of the halogens with potassium, and by the saturation of the hydroxide or carbonate with haloid acids. They are readily soluble in water, have a salty taste, and crystallize in cubes. When heated they melt, and are somewhat volatile. Potassium Chloride — KCl — occurs in Stassfurt in large deposits, as sylvite, and combined with magnesium chloride exists as carnallite (MgClg, KCl-f-ôHjO). The latter salt serves as the chief source for the preparation of potassium chloride, which meets with varied application in the arts, and also for the preparation of potassium carbonate. The chloride crystallizes in vitreous cubes, of specific gravity 1.84. It melts at 734°, and volatilizes at a strong red heaJ;. 100 parts water dissolve 30 parts of the salt at 0°, and 59 parts at 100°. Potassium Bromide—KBr—is generally obtained by warm¬ ing a solution of potassium hydroxide with bromine, when the brómate is also produced :— 6KOH -f sBrg = sKBr -f- KBrOg -j- 3H2O. The solution is evaporated to dryness, mixed with charcoal, and ignited, which reduces the brómate to bromide :— KBrOs -f- 3C = 3CO -h KBr. It is readily soluble in water and alcohol ; forms cubes of sp. gr. 2.4, and melts at 699°. POTASSIUM IODIDE—POTASSIUM CHLORATE. 283 Potassium Iodide—KI—may be prepared like the preceding. It is usually obtained according to the following method : Iodine and iron filings are rubbed together under water, and potassium carbonate added to the solution of the iron iodide ; this will pre¬ cipitate ferrous-ferric oxide ; carbon dioxide escapes, and potassium iodide will be found in the solution. It forms large white crystals, fuses at 634°, and is tolerably volatile. Its specific gravity equals 2.9. At medium temperatures it dissolves in 0.7 parts water and 40 parts of alcohol. The aqueous solution dissolves iodine in large quantity. Many metallic insoluble iodides dissolve in it without diflSculty, forming double iodides, Hgl2.2KI. The iodide is employed in medicine and photography. Potassium Pluoride—KFl—^is obtained by dissolving the carbonate in aqueous hydrofluoric acid. It crystallizes in cubes at ordinary temperatures, with aHjO, but above 35° does not contain water of crystallization. It is very soluble in water. The aqueous solution attacks glass. It is greatly inclined to combine with other fluorides: KFl.HFl; BFI3.KFI. On ad^ng hydrofluosilicic acid to the solution of potassium salts, a gelatinous precipitate of potassium silicofluoride is thrown down, which dissolves with difficulty in water. Potassium Cyanide—KCN.—This salt can be produced by saturating potassium hydroxide with hydrocyanic acid, and by heating yellow prussiate of potash (see Iron). It forms a white, easily fusible mass, which deliquesces in the air. The solution may be easily decomposed. It crystalli:2es in cubes, has an alkaline reaction, and smells like prussic acid, as this is set free by the car¬ bon dioxide of the air. By fusion potassium cyanide reduces many oxides, and hence is employed in reduction processes. It is just as poisonous as prussic acid. It is applied in many ways, especially in photography and for galvanic silvering and gilding. Potassium Chlorate—KClOj.—^The following reaction occurs when chlorine gas is conducted through a hot concentrated potassium hydroxide solution ;— 6K0H 4- 3Clj = SKQ + KOO, -f 3H2O. When the solution cools, the difficultly soluble potassium chlorate separates out. It is generally made, in trade, by the action of chlorine upon a mixture of calcium hydroxide and potassium chloride. The reaction occurs in two phases ; first, calcium chlor¬ ate is formed :— 6Ca(0H), 4- óClj = sCaClj -f- Ca(C103)j + óHjO; this then reacts with the potassium chloride :— CaCClOg), 4- 2KCI = 2KC10s + CaCl,. 284 INORGANIC CHEMISTRY. Potassium chlorate crystallizes from the hot solution in shining tables of the monoclinic system, which dissolve with difficulty in water (loo parts at the ordinary temperature dissolve 6 parts of the salt). Its taste is cooling and astringent. When heated it melts (at 359°) giving up a portion of its oxygen, and changes to the Perchlorate—KCIO^—which on further heating decomposes into oxygen and potassium chloride (see p. 182). As it gives up oxygen readily, it serves as a strong oxidizing agent. With hydrochloric acid it liberates chlorine :— KClOg + 6HC1 = KCl + 3H2O + 3CI2. Mixed with sulphur, or certain sulphides, it explodes on heating and when struck a sharp blow. The igniting material upon the so-called Swedish (parlor) matches consists of antimony sulphide and potassium chlorate ; when this is rubbed upon the friction sur¬ face coated with red phosphorus it ignites. Potassium Hypochlorite—KCIO—is formed when chlorine is permitted to act upon a cold solution of potassium hydroxide :— 2KOH + CI2 = KCl + KCIO + H2O. It only exists in aqueous solution ; when the latter is evapo¬ rated the salt is decomposed into chloride and chlorate :— 3CIOK = 2KCI -f CIO3K. The solution has an odor resembling that of chlorine, and bleaches strongly, especially upon the addition of acids. The bleaching solutions occurring in trade (Eau de Javelle) are pre¬ pared by the action of chlorine upon solutions of sodium and potas¬ sium carbonates ; they also contain free hypochlorous acid. The oxy-salts of bromine and iodine are perfectly analogous to those of chlo¬ rine. Potassium Brómate—KBrOg—and Potassium lodate—KIO3—are pre¬ pared by the action of bromine or iodine upon hot potassium hydroxide ; the second is also produced by the action of iodine upon potassium chlorate, when the chlorine is directly replaced (p. 186). If chlorine be passed through a hot solu¬ tion of potassium iodate in potassium hydroxide—the periodate of potassium, KIO4, arises ; it is difficultly soluble and when heated decomposes into O and KIO3, which then breaks up into potassium iodide and oxygen. Besides the normal periodates, KIO4, NaI04, other salts exist which are derived from the highest hydroxyl compound, I(0H),7, and its anhydro-deriva- tives (p. 185). These salts are very numerous, and are in part monoperiodates, I0(0H)5 and 102(011)3, and partly polyperiodates, produced by the conden¬ sation of several molecules of the highest hydroxides, e.g., I„0,(0H)a and Potassium Sulphate—K2SO4—is formed in the action of sul¬ phuric acid upon potassium chloride, and as a by-product in many technical operations. It crystallizes without water, in small rhombic POTASSIUM NITRATE. 285 prisms, having a bitter, salty taste, and dissolves in lo parts H2O of ordinary temperature. It is employed principally for the prep¬ aration of potassium carbonate, according to the method of Le Blanc. (See Soda.) The acid or primary salt—HKSO4—crystallizes in large rhombic tables, and is very readily soluble in water. It fuses about 200°, loses water, and is converted into potassium pyrosulphate—K2S2O7 (p. 198)—which at 600° yields K2SO4 and SOg. The salts of sulphurous acid—^the primary, SO3KH, and the secondary sul¬ phites, SO3K2—are produced when SOg comes in contact with a potassium car¬ bonate solution ; they are very soluble and crystallize with difficulty. The first salt shows an acid, the second an alkaline reaction. If sulphur dioxide be passed into a solution of potassium carbonate until effervescence ceases and then cooled, the pyrosulphite—KgSgOs—corresponding to the pyrosulphate, will crystallize out. Potassium Nitrate, Saltpetre, KNO3, does not occur anywhere in large quantities, but is widely distributed in the upper strata of the earth and is found as an efflorescence on the soil in some regions of the hot zone (in Egypt and East India). It is produced whenever nitrogenous organic substances decay in the presence of potassium carbonate—conditions which are present in almost every soil. The intentional introduction of these is the basis of the arti¬ ficial nitre production in the so-called saltpetre plantations. Ma¬ nures and various animal offals are mixed with wood ashes (potas¬ sium carbonate) and lime, arranged in porous layers, and submitted to the action of the air for two or three years, when nitrates are produced from the slow oxidation of the nitrogen. The heaps are then treated with water and potassium carbonate added to the solu¬ tion, which contains potassium, calcium and magnesium nitrates, to convert the last two salts into potassium nitrate :— Ca(N0,)2 + KgC03 = CaC03 2KNO3. The precipitate of calcium and magnesium carbonate is filtered off and the solution evaporated. The procedure was formerly employed universally in the manufacture of potassium nitrate. At present, however, almost all of it is obtained by the decomposition of the sodium salt, occurring in large deposits in Chili, by means of potassium carbonate or chloride :— NaN03 + KCl = NaCl -f- KNO3. Warm saturated solutions of sodium nitrate and potassium chlo¬ ride are mixed and boiled, when sodium chloride, being less soluble in hot water, will separate. On cooling the solution potassium nitrate, being less soluble in cold water, crystallizes out; sodium chloride is about equally soluble in hot and cold water, for which reason the portion not separated by boiling remains in solution. 286 INORGANIC CHEMISTRY. Potassium saltpetre crystallizes without water of crystallization in large six-sided rhombic prisms. It is far more soluble in hot than in cold water; loo parts of water dissolve 244 parts at 100°, but at 0° only 13 parts. It possesses a cooling taste, fuses at 338°, and decomposes, when further heated, into oxygen and potassium nitrite, KNO2. Heated with carbon it yields potassium carbonate :— 4KNO3 -f 5C =- 2K2CO3 + 3CO3 + 2N3. Its principal use is in the manufacture of gunpowder. This is a granular mix¬ ture of potassium nitrate, sulphur, and charcoal. The relative quantities of these constituents are somewhat different in the various kinds of powder (sporting, blasting, and cannon). Upon an average, the powder consists of 75 percent. KNO3, 12 percent, sulphur, and 13 per cent, carbon, which closely corresponds to the atomic composition 2KNO3 + S + 3C. When the powder bums, its decomposition is approximately expressed by the following equation :— 2KNO3 + S + 3C = K3S + 3CO3 + Nj. The effectiveness of the powder, therefore, depends upon the diseng^ement of carbon dioxide and nitrogen gas, the volume of which is almost 1000 times as great as that of the decomposed powder. Potassium Nitrite—KNO2—isobtained by fusing saltpetre with lead, which withdraws one atom of oxygen from the former. A white, fusible mass results ; this deliquesces in the air. The potassium salts of phosphoric acid: K3PO4, K2HPO4, and KH2PO4, meet with no practical application, they are readily soluble m water and crystallize poorly ; therefore, the sodium salts are generally used. The borates, BO2K and B4O2KJ + 5H2O (see Borax), crystallize with difficulty. Potassium Carbonate—K2CO3—ordinarily known as pot¬ ashes, is a principal ingredient of plant ashes. The field plants absorb potassium salts from the earth ; these are then transformed in them into salts of organic acids. When the plants are burned the organic acids are destroyed and potassium carbonate remains. The ashes are lixiviated with hot water, the filtrate evaporated and the residue ignited. The crude potashes thus obtained contain, besides the carbonate, also chloride, sulphate, and other salts. To purify them, treat with a little water, which will dissolve the easily soluble carbonate, leaving nearly all of the other ingredients behind. In this way we obtain pure potashes. This method of getting potashes from plant ashes was formerly pursued extensively in America, Hungary, and Russia ; it is not much used at present, because potassium carbonate is, upon the one hand, replaced by the cheaper sodium carbonate in practice; on the other hand, the immense deposits in Stassfurt and Galicia afford an inexhaustible supply of potassium salts. Considerable quantities of potassium SULPHUR COMPOUNDS. 287 carbonate, used at present almost entirely for the production of Bohemian or crystal glass, have been recently obtained from Stass- furt, according to the methods employed in the preparation of sodium carbonate from the chloride. (See Soda.) Chemically pure potassium carbonate is obtained most conveniently by the ignition of cream of tartar or by heating the primary carbonate. The commercial carbonate is a white, deliquescent powder melt¬ ing at 830°, and vaporizing at a red heat. It crystallizes from concentrated aqueous solutions with i molecules of water, in monoclinic prisms; at 100° it loses ^ molecule water. The solu¬ tion has a caustic taste and shows an alkaline reaction. When CO2 is conducted through the liquid it is absorbed and primary potassium carbonate is produced :— CO3K2 + H^O + CO2 = 2KHCO,. This salt, ordinarily called bi-carbonate^ crystallizes in mono- clinic prisms, free from water. It dissolves in 3-4 parts water and exhibits a neutral reaction. Heated to 80°, it decomposes into KjCOjjCOa ítiid water. The decomposition of the dry salt does not begin until about 110°, while the aqueous solution decomposes even on evaporation. Potassium Silicate, water-glass, does not possess a constant composition and cannot be obtained crystallized. It is produced by solution of silicic acid or amorphous silicon dioxide in potas¬ sium hydroxide, or by the fusion of silica with potassium hydroxide or carbonate. The concentrated solution dries, when exposed, to a glassy, afterward opaque mass, which, when reduced to a powder, will dissolve in boiling water. Potassium (and also sodium) water- glass has an extended application, especially in cotton printing, for the fixing of colors (stereochromy), in rendering combustible material fireproof, in soap boiling, etc. SULPHUR COMPOUNDS OP POTASSIUM. Potassium Hydrosulphide—KSH—is obtained when potas¬ sium hydroxide is saturated with hydrogen sulphide :— KOH 4- Hß = KSH + H^O. Evaporated in vacuo it crystallizes in colorless rhombohedra, of the formula 2KSH -f- H2O, which deliquesce in the air. At 200°, it loses its water of crystallization, and at a higher temperature fuses to a yellowish liquid» which solidifies to a reddish mass. Like the hydroxide, it has ap alkaline reaction. On adding an equivalent 288 INORGANIC CHEMISTRY. quantity of potassium hydroxide to the sulphydrate solution, we get potassium sulphide :— KSH 4- KOH = KjS + HjO. Potassium Sulphide—K2S—is usually obtained by fusing potassium sulphate with carbon :— K2SO4 + 2C KjS 4- 2CO2. When fused, it solidifies to a red crystalline mass. It crystallizes from concentrated aqueous solutions with 5 molecules of H2O, in colorless prisms, which deliquesce in the air. The solution absorbs oxygen from the latter, and is decomposed into potassium hyposul¬ phite and caustic potash :— 2K2S 4- HjO 4- 2O2 = K2S2O3 4- 2KOH. Potassium hydrosulphide and sulphide precipitate insoluble sul¬ phides from the solutions of many metallic salts. They are decom¬ posed by acids with liberation of hydrogen sulphide. When the aqueous solution of the sulphide is boiled with sulphur the poly sulphides y K2S3, K2S4, and K2S5, are formed, which after fusion solidify to yellowish-brown masses. The aqueous solutions of the polysulphides are decomposed by acids, with disengagement of H2S and separation of sulphur (milk of sulphur). The so-called liver of sulphur {Hepar sulphuris')^ a liver-brown mass, used in med¬ icine, is obtained by the fusion of potassium carbonate with sulphur, and consists of a mixture of potassium polysulphides with potassium sulphate. The aqueous solution of the potassium, as well as that of the sodium sulphide, dissolves some metallic sulphides and forms sulpho- salts with them (p. 226). When dry ammonia is conducted over heated potassium, potas- samide (NHjK) results. This is a dark-blue liquid which solidifies to a yellowish-brown mass. Water decomposes it into potassium hydroxide and ammonia. When potassamide is ignited away from the air, it loses ammonia, and leaves behind potassium triamide, NK3, a blackish compound which is spontaneously inflammable. Recognition of the Potassium Compounds.—Almost all the potassium compounds are easily soluble in water, with the ex¬ ception of a few, which, therefore, serve for the characterization and separation of potassium. Tartaric acid added to the solution of a potassium salt gives a crystalline precipitate of acid potassium tartrate. Platinic chloride (PtCh) produces in potassium solutions RUBIDIUM AND CAESIUM—SODIUM. 289 a yellow, crystalline precipitate of PtCli- 2KCI. Potassium com¬ pounds introduced into the flame of an alcohol or gas lamp impart to the same a violet coloration. The spectrum of the flame is char¬ acterized by two bright lines, one red and one violet (see Spectrum Analysis). RUBIDIUM AND C/ESIUM. Rb = 85.2. Cs == 132.7. Rubidium and Caesium are the perfect analogues of potassium (p. 278}. They were discovered by means of the spectroscope, by Bunsen and Kirchhoff, in i860. Although only occurring in small quantities, they are yet very widely distributed, and frequently accompany potassium in mineral springs, salt, and plant ashes. The mineral lepidolite contains 0.5 per cent, of rubidium; upward of 30 per cent, of caesium oxide is present in the very rare pollucite, a silicate of aluminium and caesium. The spectrum of rubidium is marked by two red and two violet lines ; caesium by two distinct blue lines ; hence, the names of these elements. Rubidium and caesium form double chlorides (PtCh. 2RbCl) with platinum chloride, and they are more insoluble than the double platinum salt of potassium, hence may answer for the separation of these elements from potassium. Rubid¬ ium and caesium may be obtained free by decomposing their fused chlorides with the electric current. Rubidium is also prepared by igniting its carbonate with charcoal. Metallic rubidium has a silver-white color, with a somewhat yellowish tinge; its vapor is greenish-blue. Metallic caesium has been obtained by the electrolysis of a mixture of caesium and barium cyanides. Electrodes of alumin¬ ium are employed for this purpose. Caesium is a silver white metal, of sp. gr. 1.85. It oxidizes quite readily and inflames in the air. It melts at 26.5° and boils at 270°. SODIUM. Na = 23.06. Sodium is widely distributed in nature, especially as chloride in sea water and as rock-salt ; and is also found in silicates. The metal was obtained in 1807, by Davy, by the action of a strong electric current upon fused sodium hydroxide. At present, like potassium, it is obtained upon a large scale by igniting a mixture of sodium carbonate and carbon in an iron retort :— NojCOg 2C = 2Na -f- 3CO. The liberated sodium vapors are condensed on flat iron receivers of peculiar construction, and the liquefied sodium collected under rock-oil. Sodium in external properties is very similar to potassium. It melts at 95.6°, distils at a red heat, and is converted into a color¬ less vapor, which burns with a bright yellow flame in the air. It oxidizes readily on exposure, and decomposes water even in the cold, although less energetically than potassium. A piece of sodium thrown upon water swims about upon the surface with a rotatory movement, the disengaged hydrogen, however, not igniting. If we prevent the motion, by confining the metal to one place, the heat 25 290 INORGANIC CHEMISTRY. liberated by the reaction attains the ignition temperature of hydro¬ gen, and a flame follows. Sodium Oxide,—NajO, and suboxide, Na^O, are very similar to tue corre¬ sponding potassium compounds; the peroxide, is somewhat different. It is ob¬ tained by burning sodium in a stream of oxygen. Its formula is Na202. When heated it absorbs iodine vapors, forming the compound, NagOIj (NagOj -{-1^ = Na20l2 O), soluble in water, but decomposed by acids into free iodine and sodium salt. This compound like some others, seems to indicate that sodium has several valences. When heated with hydrogen sodium oxide is decomposed with separation of metallic sodium and formation of sodium hydroxide. This is explained by the fact that the heat of formation of NaOH is greater than that of Na^O (p. 279). Carbon monoxide decomposes sodium monoxide in a similar manner, when the latter is heated to 290-310® :— 2Na20 -j- CO == Na2C03 -|- Na2. This reaction occurs because the heat of formation of sodium carbonate (271.2 C.) is greater than that of 2Na20(2 . 100.2 C.) and CO(30.i C.). Na20 combines with CO2 about 400° with production of light and yields C03Na2(Na20, COj = 74.1 C ). Sodium Hydroxide, Sodium Hydrate, or Caustic Soda, NaOH, like potassium hydroxide, is formed by boiling a solution of sodium carbonate with calcium hydroxide :— Na2C03 + Ca(0H)2 = CaC03 + 2NaOH. At present it is directly produced in the soda manufacture by adding a little more carbon to the fusion (see Soda), or by igniting sodium carbonate (Löwig) with ferric oxide, which aflbrds a com¬ pound of re203.Na20, decomposed by warm water into ferric oxide and sodium hydroxide. The sodium hydroxide which solidifles after fusion is a white, radiating, crystalline mass, and resembles caustic potash very much. It attracts water from the air, becomes moist, and coats itself by carbon dioxide absorption with a white layer of sodium carbonate (caustic potash deliquesces perfectly, because the resulting carbonate is also deliquescent). The aqueous solution, called sodium hy¬ droxide, resembles that of potassium. Crystals of NaOH -j- ß^HjÖ separate at 0° from the concentrated solution ; they melt at 6°. Sodium Chloride—NaCl—is abundant in nature. It is found almost everywhere in the earth and in natural waters ; in sea-water it averages 2.7-3.2 per cent. As rock-salt it forms large deposits in many districts—at Stassfurt and Wielizca in Galicia. In warm climates, on the coasts of the Mediterranean Sea, sodium chloride is gotten from the sea, according to the following procedure. At high tide, sea- SODIUM lODATE. 291 water is allowed to flow into wide, flat basins (salt gardens), in which it evaporates under the sun's heat; the working is limited, therefore, to summer time. After sufficient concentration, pure sodium chloride first separates, and this is collected by itself. Later, there crystallizes a mixture of sodium chloride and magnesium sulphate; Anally potassium chloride, magnesium chloride and some other salts appear (among them potassium iodide and bromide), the separation of which con¬ stitutes a particular industrial branch in sôme regions. In cold climates, as in Nor¬ way and at the White Sea, the cold of winter is employed for the production of salt. In the freezing of sea-water, as well as of other solutions, almost pure ice separates at first; the enriched sodium chloride solution is then concentrated in the usual way. The rock-salt is either mined in shafts, or, where the strata are not so large and are admixed with other varieties of rock, a lixiviation process is employed. Bor¬ ings are made in the earth and water run into them, or into any openings already formed. When the water has saturated itself with sodium chloride, it is pumped to the surface and the brine then further worked up. In many regions, especially in Reichenhall, in Bavaria, more or less saturated natural salt or brine springs flow from the earth. The concentration of the non-saturated brine occuis at first in the so-called " graduation " houses. These are long wooden frames fiTed with fagots, and on letting the salt water run upon them it will be distributed and evaporated by the fall; the concentrated brine collects in the basin below, and is then evaporated over a free fire. Sodium chloride crystallizes from water in transparent cubes, which arrange themselves by slow cooling into hollow, four-sided pyramids. It melts at 772° and volatilizes at a white heat. It is slightly more soluble in hot than in cold water; 100 parts at 0° dissolve 36 parts salt; at 100°, 39 parts. The saturated solution, therefore, contains 26 per cent, sodium chloride. The specific gravity of the crystals equals 2.13. If the saturated solution be cooled below —10°, large monoclinic tables (NaCl -f- 2H2O) sep¬ arate ; these lose water at 0° and become cubes. The ordinary sodium chloride usually contains a slight admixture of magnesium salts, in consequence of which it gradually deliquesces in the air ; the perfectly pure salt is not hygroscopic. When heated the crystals crackle, because of the escape of the mechanically enclosed water. Sodium Bromide and iodide crystallize at ordinary temperatures with 2 molecules of H jO, which they lose again at 30®; above 30® they separate in anhydrous cubes. Sodium bromide fuses at 708® and the iodide at 628® ; the former is difficultly soluble in alcohol and the latter is very soluble. Sodium Chlorate (NaClOg) and Perchlorate (NaClO^) are considerably more soluble in water than the corresponding potassium salts. Sodium lodate—NalOg— is obtained the same as the potassium salt, and crys¬ tallizes at ordinary temperatures with 3 molecules of HgO in silky needles. If chlorine gas be conducted through the warmed solution of sodium iodate in sodium hydroxide, the periodate 10 | |Q^^^2(seep. 184) crystallizes out on cooling. This becomes the normal salt (NalQ^ -f- 3H2O) when dissolved in nitric acid. 292 INORGANIC CHEMISTRY, Sodium Sulphate—NajSO^—crystallizes at ordinary tempera¬ tures with ID molecules of water of crystallization, and is then known as Glauber's salt (Sa¿ mirabile Glauberi). It occurs in many mineral waters, and in large deposits, with or without water of crystallization, in Spain. It is a by-product in the manufacture of sodium chloride from sea-water and brine. It is produced in large quantities by heating salt with sulphuric acid : — aNaCl -f H2SO4 = NagSO^ + 2HCI, and is used in making soda (sodium carbonate). Or it may be prepared by the method of Hargreaves, by conducting SO2, air and steam over strongly ignited sodium chloride :— 2NaCl + SO2 + O -h H2O = SO^Nag -f 2HCL More recently the sulphate has been obtained by a transposition of sodium chloride with magnesium sulphate at a winter temperature —a procedure which is prosecuted chiefly in Stassfurt, where immense quantities of magnesium sulphate exist : — 2NaCl -f- SO^Mg = MgCIa + SO^Nsl^. Sodium sulphate crystallizes at ordinary temperatures with 10 molecules of H2O, in large, colorless, monoclinic prisms, which crumble in the air and fall into a white powder. When the salt is heated to 33°, it fuses in its own water of crystallization ; by further increase of temperature it gradually loses this, becomes solid, and again fuses at a red heat. The solubility of Glauber's salt (NasSO^ -f- 10 H2O) shows the following interesting deportment : 100 parts of water dissolve, at 0°, 12 parts; at 18°, 48 parts; at 25°, 100 parts; at 30°, 200 parts ; at 33°, 327 parts of the hydrous salt. At the last temperature the solubility is greatest ; by further increase of heat it gradually diminishes ; at 50°, 100 parts water dissolve only 263 parts ; at 100°, 238 parts of the salt. While, ordinarily, the solubility increases with temperature, Glauber's salt exhibits a varying deportment. This is explained in that the hydrate, NajSO« -|- 10H2O, in aqueous solution, above the temperature of 33°, decomposes into water and the salt, NaaSO^ -f- HjO, which is less soluble in water. The decomposition does not occur at once, but only gradually, with increasing temperature, for which reason the quantity of the salt dissolved gradually grows less. Here we have an example of dissociation taking place in aqueous solution (p. 196). The solution, saturated at 33°, becomes turbid upon heating, and a portion of the dissolved salt separates in anhydrous, small, rhombic octahedra. The following interesting deportment in the solution of Glauber's salt may also be noticed. When the solution, saturated at 33®, is allowed to cool to the ordi- SODIUM HYPOSULPHITE. 293 nary temperature, and even lower, not the slightest separation of crystals occurs, although the salt is vastly more insoluble at lower temperatures than at 33°. Many other salts form supersaturated solutions, although they are less striking than that of Glauber's salt. The supersaturated solution of the latter may be agitated and twirled about without crystallization setting in. If, however, a glass rod, or some other solid body, be introduced into the solution, it will solidify suddenly to a crystalline mass. The particles of dust floating about in the air will have a like effect : therefore, to preserve the supersaturated solution, the vessel containing it should be kept well corked. By accurately made investigations, it has been deter¬ mined that the crystallization of the supersaturated Glauber's salt solution is only induced by contact with already formed crystals. These must then be present every¬ where in the atmosphere, because only solids which have been exposed to the air, and have not been carefully cleansed afterward, bring about the crystallization. Hence, the formation of a crystal of Glauber's salt is always dependent upon the previous existence of a similar crystal—^just as the production of cells is only caused by cells. In the crystallization of a supersaturated Glauber's salt solution, considerable heat is disengaged, and the mass increases in temperature. This is because the latent heat of all substances in the liquid condition is greater than in the solid. At 10°, occasionally, and of their own accord, transparent crystals, NagSO^ yHgO, separate from the supersaturated solutions. Exposed to the air and in contact with solid bodies, these crystals are changed to anhydrous sodium sulphate and Glau¬ ber's salt. This salt is employed in medicine as a purgative, and finds extended application in the manufacture of glass and the prepara¬ tion of soda. The primary or acid sodium sulphate — NaHSOi — is obtained by the action of sulphuric acid upon the neutral salt or upon sodium chloride :— NaCl -f H2S0^ = NaHSO^ + HCl. At ordinary temperatures, it crystallizes with one molecule of water, and is perfectly analogous to the potassium salt. The sodium salts of sulphurous acid are obtained by conducting sulphur dioxide into solutions of sodium hydroxide or carbonate. The secondary sulphite, NajSOg, crystallizes with 7 molecules of HjO at ordinary temperatures ; in the presence of sodium hydroxide, or by warming the solution, it separates in the anhydrous state. The primary sulphite—NaHSOg—gives up sulphur dioxide in the air, and is oxidized to sodium sulphate. Sodium Hyposulphite—NagSgOa—is prepared by boiling the aqueous solution of neutral sodium sulphite with flowers of sul¬ phur:— NagSOg + S = Na^SgOg. It is obtained as a by-product in the recovery of sulphur from the soda residues. It crystallizes with five molecules of HjO, in large monoclinic prisms, dissolves very readily in water, and is somewhat deliques¬ cent in the air. At 56°, it melts in its water of crystallization ; loses all water at 100°, and decomposes by further heating into 294 INORGANIC CHEMISTRY. NajSO^ and Na2S5. When the dry salt is heated in the air, the polysulphide burns with a blue flame. Acids decompose the aqueous solution with separation of sulphur and evolution of sulphur dioxide :— S^OgNaj + 2HCI = aNaCl + SOj + S + H,O. Like the sulphate, it readily affords supersaturated solutions. The hyposulphite is used as a reducing agent ; chlorine, bromine and iodine are converted by it into the corresponding halogen salts :— 2S203Na2 -j- L ~ -|- aNal. Sodium tetrathionate. An iodine solution is instantaneously decolorized by sodium hyposulphite. Chlorine behaves differently; sulphuric acid and sodium chloride are produced. Upon this reaction rests the appli¬ cation of sodium hyposulphite as an antichlor in chlorine bleaching, to remove the excess of the chlorine, which has a destructive action upon the fibre. In consequence of its property of dissolv¬ ing the halogen silver derivatives, it is employed in photography. Sodium Carbonate (Soda)—Na^COa.—This, technically, very important salt occurs frequently in nature. In some districts, as in Hungary and in Africa, it disintegrates from the soil, and occurs also in the so-called sodium seas (in Egypt, and upon the coast of the Caspian Sea). It is contained in the ashes of many sea-plants, chiefly the algae, etc. These assimilate the sodium salts of the earth, while the land-plants absorb the potassium salts, and for this reason contain potashes in their ash. The ash of the sea-plants, called varec in Normandy, kelp in England, formerly served as the principal material for the preparation of soda. At present it is, however, almost exclusively made in large quantities from sodium chloride, according to a method devised in 1787 by Leblanc. According to this method, the sodium chloride is converted, by warming with sulphuric acid, into sodium sulphate (p. 292). When the latter is dry, it is mixed with charcoal and chalk, and ignited in a reverberatory furnace. Two principal phases may be distinguished in this reaction. First, the carbon reduces the sodium sulphate :— Na^SO^ + 2C = Na,S -f-aCOj. The sodium sulphide then acts upon the calcium carbonate to form calcium sulphide and sodium carbonate :— Na^S -fCaCOs = CaS -fNa^CO,. At the same time, the high temperature converts a portion of the SODIUM CARBONATE. 295 calcium carbonate into calcium oxide and carbon dioxide, which is reduced by the ignited carbon to the monoxide ; the appearance of the latter, which burns with a bluish flame, indicates the end of the action. The chief products in the soda fusion are, then, sodium carbonate, calcium sulphide and oxide ; in addition, different other sulphur salts are formed in smaller quantity. The fusion is lixivi¬ ated with cold water; the sodium carbonate dissolves, and there remains behind an insoluble compound of calcium sulphide with oxide, CaO.aCaS, the sûi/a residue. By the evaporation of the solution and the ignition of the residue, we get the commercial or crude calcined sodoy containing different admixtures, among them sodium hydroxide. The latter is formed by the action of excess of carbon upon sodium carbonate :— NiLjCOg 4" C = NhjO -f- 2CO. By purposely adding more carbon to the fusion, sodium hydrox¬ ide may be obtained, together with the carbonate. To purify the crude soda it is recrystallized from water ; large, transparent crystals, NajCOg loHgO, crystallized soda, separate out ; the sodium hy¬ droxide remains dissolved. Considerable quantities of soda are obtained at present from cryolite, a com- p>ou,nd of aluminium fluoride and sodium fluoride (AlFl3,3NaFl), which occurs in great deposits in Iceland. The pulverized mineral is ignited with burned lime ; insoluble calcium fluoride and a very soluble compound of aluminium oxide with sodium oxide, called sodium aluminate (see Aluminium) are produced :— 2(AlFl3.3NáFl) + 6CaO = ÓCaFl^ + A]303.3Na20. The mass is treated with water and carbon dioxide conducted into the solution, which causes the precipitation of aluminium oxide, and sodium carbonate dis- solves AlA-SNap + 3H3O 4- 3CO3 = Al3(OH)6 + 3Na3C03. Latterly, a third procedure has appeared. It depends upon the double decom¬ position of a solution of sodium chloride with primary ammonium carbonate, by heat, under high pressure ;— NaCl + COgCNHJH = NaHCOg -f NH^. The primary sodium carbonate being rather insoluble in water, separates from solu¬ tion, and is converted by heat into the secondary carbonate. The ammonium chloride remains dissolved, and afterward is converted again into carbonate by aid of cal¬ cium carbonate. In this way, one and the same quantity of ammonium carbonate will suffice for the conversion of an indefinite quantity of sodium chloride into soda. The technical difficulties which at first opposed the extension of this process, so simple in chemical respects, are now entirely removed, and at present half of all the soda manufactured in Europe is made by this so-called ammonia process (Solvay-soda). The reason that the Leblanc method has not been entirely supplanted, is that until now theie was no means by which the chlorine of the sodium chloride could be converted into some available form. 296 INORGANIC CHEMISTRY. At ordinary temperatures sodium carbonate crystallizes with lo molecules of H20(Na2C03 -|- 10H2O) in large monoclinic crystals, which crumble upon exposure and become a white powder. It melts at 50° in its water of crystallization, and upon additional application of heat a pulverulent hydrate—NajCOs -)- 2H2O— separates, which in dry air his i molecule of H2O, and at 100° loses all of this. At 30^-50° rhombic prisms of the composition C03Na2 + 7H2O, crystallize from the aqueous solution. The anhy¬ drous salt absorbs water from the air but does not deliquesce. It melts at a red heat and volatilizes somewhat at a very high tempera¬ ture, 100 parts H2O dissolve 15 parts at 0°, and at 38°, 138 parts of the dry salt. At more elevated temperatures the solubility is less, owing, as in the case of the sulphate, to the formation of less soluble, lower hydrates. Sodium carbonate has a strong alkaline reaction ; acids liberate carbon dioxide from it. Primary Sodium Carbonate—ordinary Bicarbonate of Soda —Natrium Bicarbonicum—NaHCOa—is produced by the action of carbon dioxide upon the hydrous secondary carbonate :— NajCOj + CO2 -f- H2O — 2NaHC03. It crystallizes without water, in small monoclinic tables; it dissolves, however, at ordinary temperatures in lo-ii parts water, and possesses feeble alkaline reaction. By heating and boiling the solution it passes into the secondary carbonate with disengage¬ ment of carbon dioxide. The salt decomposes rapidly even under 100°. By rapid evaporation small monoclinic prisms of the so- called sodium sesquicarbonate—CgOgNa^ -f- 3H2O or Na4H2(C03)3 -|- 2H2O, separate; this also deposits in the sodium seas of Hungary and Egypt. It is called Trona or Urao, Sodium Nitrate—NaN03—Chili saltpetre, is found in im¬ mense deposits in Peru. It crystallizes in rhombohedra very simi¬ lar to cubes, hence designated cubic saltpetre. It fuses about 318°. in water it is somewhat more easily soluble than potassium salt¬ petre. In the air it attracts moisture, hence it is not adapted to the manufacture of gunpowder. In other respects it is perfectly similar to potassium nitrate. It is largely used in the manufacture of nitric acid, and especially in preparing potassium saltpetre (p. 285). Sodium Nitrite, NaNOg, is prepared like potassium nitrite (p. 286), by heating sodium nitrate with lead, iron, or graphite. It crystallizes more readily than potassium nitrite, and does not deliquesce in the air. It occurs in trade in small colorless crystals, containing from 93-98 per cent, of the pure salt. It is largely used in the dye industry for the preparation of the azo-compounds. Sodium Phosphates. The sodium salts of phosphoric acid are less soluble and crystallize better than those of potassium. The tri-sodium phosphate—NajPOi—is made by saturating i molecule of phosphoric acid with 3 molecules of NaOH, and crystallizes in six- SODIUM BORATE. 297 sided prisms with 12 molecules of H^O. It has a strong alkaline reaction, absorbs carbon dioxide from the air, and is converted into the secondary salt. Di-sodium Phosphate—NajHPOi—is the most stable of the sodium phosphates, and hence, is generally employed in labora¬ tories {Natrium phosphoricum^. It may be obtained by saturating phosphoric acid with sodium hydroxide to feeble alkaline reaction, or may be prepared on a large scale by decomposing bone ashes (tricalcium phosphate) with an equivalent amount of sulphuric acid. It crystallizes at ordinary temperatures with 12H2O in large monoclinic prisms which effloresce rapidly upon exposure. It separates from solutions with a temperature above 30° in non- efflorescing crystals containing 7 mois. HjO. It is soluble in 4-5 parts water, and shows a feeble alkaline reaction. The solution absorbs large quantities of carbon dioxide, without suffering any alteration. When heated the salt loses water, melts about 300° and becomes Sodium Pyrophosphate—Na4P207—which crystallizes with 10 molecules of HgO, and upon boiling with nitric acid passes into primary sodium phosphate. The primary or monosodium phosphate—NaH2P04, crystallizes with I molecule of H2O, and exhibits an aóid reaction. At 100° it loses its water of crystallization, and at 200° becomes Na2H2P207, disodium pyrophosphate^ which at 240° forms sodium metaphosphate —NaPO^:— H2Na2P207 = 2NaP03 ~i~ HjO. We get various modifications of the metaphosphate, according to the conditions of fusing and cooling ; they are probably polyme- rides, corresponding to the formulas Na2P206, NasPsOg, etc. Upon heating sodium metaphosphate with metallic oxides the latter dis¬ solve, and salts of orthophosphoric acid are formed, e. g. :— NaPOg -f CaO = NaCaPO^. In this manner, characteristic colored glasses (phosphorus beads) are obtained with various metals. In blowpipe analysis this beha?- vior serves for the detection of the respective metals. The salts of arsenic acid are perfectly analogous to those of phosphoric acid. Of the antimoniales may be mentioned the disodium pyroantimoniate, NajHjSbjOj -j- óHjO, which is insoluble in cold water, and is therefore pre¬ cipitated from the soluble sodium salts on the addition of dipotassium pyroan¬ timoniate. Sodium Borate. The normal salts of boric acid, B(OH)s, and metaboric acid, BO. OH (see p. 245), are not very stable. The ordinary alkaline borates are derived from tetraboric acid (H2B4O7), which results from the condensation of 4 molecules of the normal boric acid :— 4B(0H), — 5H2O H2B4O7. 298 INORGANIC CHEMISTRY. The most important of the salts is borax, which crystallizes at ordinary temperatures with lo molecules of H2O in large mono- clinic prisms, Na2B407 + loHjO. Borax occurs naturally in some lakes of Thibet, whence it was formerly imported under the name of tinkal. At present, it is prepared artificially by boiling or fusing boric acid with sodium carbonate. At ordinary temperatures, the crystals dissolve in 14 parts water; at 100° in one-half part; and the solution has a feeble alkaline reaction. When heated to 70® rhombohedra crystallize from the solution, and have the composi¬ tion NajBiOT -j- 5H2O, formerly known as octahedral borax. Both salts puif up when heated, lose water, and yield a white, porous mass {burned borax), which fuses at 561° to a transparent vitreous mass (Na2B407). In fusion this dissolves many metallic oxides, forming transparent glasses {borax beads), which frequently possess characteristic colors ; thus copper salts give a blue, chromic oxide, a green glass. Therefore, borax may be employed in blowpipe tests for the detection of certain metals. Upon this property of dissolving metallic oxides depends the application of borax for the fusion and soldering of metals. Sodium Silicate—sodium water glass—is analogous to the potassium salt, and is most readily obtained by fusing quartz with sodium sulphate and charcoal. The sulphur compounds of sodium are also analogous to those of potassium. Recognition of Sodium Compounds.—Almost all the sodium salts are easily soluble in water, sodium pyroantimoniate — H2Na2Sb207—excepted; this is precipitated from solutions of so¬ dium salts by potassium pyroantimoniate, and can serve for the detection of sodium. Sodium compounds, exposed in a colorless ñame, impart to the latter an intense yellow. The spectrum of the sodium ñame is characterized by a very bright yellow line, which, when more strongly magnified, splits into two lines. LITHIUM. Li = 7,03. Lithium only occurs in nature in small quantities, but is tolerably widely disseminated, and is found in some mineral springs and in the ashes of many plants, notably in that of tobacco and the beet. As a compound silicate, it occurs in lepidolite or lithia mica ; as phosphate (with iron and manganese) in triphylite. The metal is separated from the chloride by means of the electric AMMONIUM COMPOUNDS. 299 current, and is silver-white in color, decomposing water at ordinary temperatures. Its specific gravity is 0.59. It is the lightest of all the metals, and swims upon naphtha. It melts at 180°, and burns with an intense white light. The lithium salts are very similar to the salts of sodium, but closely approach those of magnesium (p. 278). Lithium Chloride—LiCl—crystallizes, at ordinary tempera¬ tures, in anhydrous, regular octahedra ; below 10°, however, it has two molecules of HjO, and deliquesces in the air. Lithium Phosphate—LÍ3PO4 -f- —and Lithium Car¬ bonate—LÍ2CO3—are difficultly soluble in water ; therefore they are precipitated from solutions of lithium salts by sodium phosphate or carbonate. By strong ignition the carbonate loses carbon diox¬ ide. So far as these two salts are concerned, lithium approaches the metals of the calcium group (p. 278). Its compounds color the ñame a beautiful red ; the spectrum shows an intense red line. AMMONIUM COMPOUNDS. upon page 129 we observed that ammonia combines directly with the acids to form salt-like compounds, which are analogous to the metallic salts, especially those of potassium. The monovalent group, NH4, playing the rôle of metal in these derivatives, is called ammonium, and the derivatives of ammonia, ammonium compounds. The metallic character of the group NH4 is confirmed by the exist¬ ence of ammonium amalgam, which, as regards its external appear¬ ance, is very similar to the sodium and potassium amalgams. Am¬ monium amalgam may be prepared by letting the electric current act upon ammonium chloride, NH4CI, viz., by immersing the neg¬ ative platinum electrode into a depression in the ammonium chlo¬ ride, which is filled with mercury. Then, as in the case of the decom¬ position of potassium or sodium chloride, a metal—ammonium— separates on the negative pole, and combines to an amalgam with mercury. The amalgam may also be obtained if sodium amalgam be covered with a concentrated solution of ammonium chloride :— (Hg -f Na) and NH4CI yield (Hg + NH^) and NaCI. Sodium amalgam. Ammonium amalgam. Ammonium amalgam forms a very voluminous mass with a metallic appearance. It is very unstable, and decomposes rapidly into mer¬ cury, ammonia, and hydrogen. On dissolving in water ammonia yields a strong alkaline solution ; no proofs, however, exist which would lead us to accept the exist¬ ence of ammonium hydroxide (NH4OH) in the solution. On the 300 INORGANIC CHEMISTRY. Other hand, there are organic derivatives of ammonium hydroxide, in which the hydrogen of the ammonium is replaced by hydro¬ carbon residues ; e. g., tetramethyl ammonium hydroxide— N(CH3)40H. These are thick liquids, of strong basic reaction and, in all respects, are very similar to potassium and sodium hydroxides. Ammonium Chloride—NH4CI—is sometimes found in vol¬ canic districts, and was formerly obtained by the dry distillation of camel's dung {Sa/ ammoniacun^. At present it is prepared almost exclusively by saturating the ammonia water from gas works with hydrochloric acid. The solution is evaporated to dryness, and the residue heated in iron vessels, when the ammonium chloride sub¬ limes as a compact, fibrous mass. It dissolves in 2.7 parts of cold, and in one part of boiling water, and crystallizes from the solution in small, feather-like, grouped octahedra or cubes, of sharp, salty taste. When heated, ammonium chloride sublimes without melt¬ ing ; at the same time a dissociation into NH3 and HCl is sustained, but these products recorabine again to ammonium chloride, on cooling. The dissociation is complete at 350°, and the vapor den¬ sity then equals 13.3 = i), corresponding to that of a mixture of similar molecules, of NH3 (8.5), and HCl (18.2). A like decom¬ position is sustained by the ammonium chloride when its solution is boiled ; ammonia escapes and the solution contains some free hydrochloric acid. Ammonium Sulphate—(NH4)2S04—is obtained by saturating the ammonia water from gas works with sulphuric acid. It crystal¬ lizes without water in rhombic prisms, and is soluble in two parts of cold and one part of hot water. It fuses at 140°, and by further heating decomposes into ammonia, nitrogen, water and ammonium sulphite. Ammonium Nitrate—NH4NO3—is isomorphous with potas¬ sium nitrate and deliquesces in the air. When heated it melts at 159°, and then decomposes (about 86°) into hyponitrous oxide and water (p. 215). Ammonium Nitrite—NH4NO2—is present in minute quanti¬ ties in the air, and results from the action of the electric spark upon the latter when moist, and also in the oxidation of phosphorus. It may be obtained by the saturation of aqueous ammonia with ni¬ trous acid—in a perfectly pure condition, by the decomposition of silver or lead nitrite by ammonium chloride. Heat decomposes it at 180° C., into nitrogen and water (p. 117). The decomposition of ammonium nitrite into water and nitrogen, and ammo¬ nium nitrate into hyponitrous oxide and water are both exothermic reactions, occurring with the disengagement of heat (p. 302) ;— AMMONIUM CARBONATE. 301 NH^NOj = 2H2O 4- NjO + 8.0 Cal. (88.1 Cal.) (2 X 57-2 Cal.) ( — 18.3.) NH4NO2 = 2H2O + Nj -f 49.5 Cal. (64.9.) (2 X 57-2 Cal.) Consequently, these decompositions are not dissociations, but proceed independ¬ ently of the pressure of the disengaged gas, and the components do not reunite to form their original compounds (p. 280). Ammonium Carbonate.—The neutral or secondary salt (NHjjCOa, separates as a crystalline powder, when ammonia gas is conducted through a concentrated solution of the so-called sesqui- carbonate. It parts with ammonia in the air and becomes the pri¬ mary ox acid NH4HCO3, which, when heated to 58°, decom¬ poses into carbon dioxide, ammonia, and water. The common, commercial, so-called sesquicarbonate of ammo- nium, (C02)2(NH3)3.H20, which can be regarded as a compound of primary ammonium carbonate with ammonium carbamate, CO3- (NH4) H -f- NH2CO2 NH4 (see Organic Chemistry), arises in the decay of many nitrogenous carbon compounds, e. g., the urine, and was formerly prepared by the dry distillation of bones, horn, and other animal substances. At present it is obtained by heating a mixture of ammonium chloride, or sulphate, with calcium car¬ bonate. It then sublimes as a white, transparent, hard mass, which gives off ammonia and çarbon dioxide in the air, falling into a white piowder of primary ammonium carbonate. The latter, obtained by the efflorescence of the first two salts, or by saturating ammonium hydroxide with carbon dioxide, is a white, odorless powder, more insoluble in water. In aqueous solution it gradually loses carbon dioxide and is changed to the secondary carbonate. Ammonium Phosphates.—The most important of these is the secondary ammonium sodium phosphate^ P04(NH4)NaH 4H2O, ordinarily termed salt of phosphorus. It is found in guano and decaying urine. It can be obtained by the crystallization of a mix¬ ture of di-sodium phosphate and ammonium chloride :— NagHPO^ -}- NH4CI = NH^NaHPO^ -f NaCl. It consists of large, transparent, monoclinic crystals. When heated it fuses, giving up water and ammonia and forms a transparent glass of sodium metaphosphate, NaP03 (p. 297). It serves in blow-pipe tests for the detection of various metals., The tertiary ammonium phosphate—(NE[4)3P04—separates upon mixing concentrated solutions of phosphoric acid and ammonia. Upon drying, it loses ammonia and passes into the secondary salt., (NH4)2HP04, which changes to the primary salt, P04(NH4)H„ when its solution is boiled. 302 INORGANIC CHEMISTRY. Ammonium Sulphide—(NH4)2S—results upon mixing i vol¬ ume of HjS with 2 volumes of NH3 at —18°. It is a white crystal¬ line mass, dissociating, at ordinary temperatures, into NH4HS and NHg. It is obtained in aqueous solution by the saturation of an ammonium hydrosulphide solution with ammonia ; and also seems to dissociate into its constituents. A higher heat (even at 45°) com¬ pletely dissociates ammonium sulphide :— (NH4)2S = 2NH3 +H2S. 2 vols. I vol. Ammonium Hydrosulphide—NH^SH—is produced upon conducting hydrogen sulphide into an alcoholic ammonia solution. It is dissociated when warmed ; this is complete at 45° :— NH4SH =NH3 + H2S. I vol. 1 vol. It is obtained in aqueous solution by saturating aqueous ammonia with hydrogen sulphide. At first the solution is colorless, but becomes yellow on standing in contact with the air, owing to the formation of ammonium polysulphides—(NH4)2Sn. The so-called yellow ammonium sulphide is more simply obtained by the solution of sulphur in the colorless hydrosulphide. Both solutions are often employed in laboratories for analytical purposes. Recognition of Ammonium Compounds.—All ammonium salts are volatile and decompose upon heating. The alkalies and other bases liberate ammonia from them, which is recognized by its odor and the blue color it imparts to red litmus paper. Platinum chloride produces a yellow crystalline precipitate of ammonio-pla- tinum chloride, PtCli.zNHiCl, in solutions of ammonium chloride. Tartaric acid precipitates primary ammonium tartrate. The heat of formation of the most important ammonium compounds corresponds to the following symbols :— (N,H3) = 11.9 (NH3Aq.) = 8.4 (N,H3,S.H) = 39.0 (N,H4,C1) = 75.8 (N,H4,Br) = 65.3 (N.H^,!) = 49.3 (NH4,Cl,Aq.)=—3.8 (NH4,Br,Aq.) =— 3.3 (NH^.I.Aq.) =—3.5 (N2,H3,S,04)= 282.1 (N,H4,N,03) = 88.1 (N.H^.N.Oa) = 64.9 NH3,HC1 = 4I.9 NH3,HBr = 45.o NH3,HI=4i.4 METALS OF THE SECOND GROUP. Ca 40 Sr 87.5 Ba 137. Be 9.10 Mg 24.38 2065.5 Cd 112.1 Hg 200.4. The second group of the periodic system (see table, p. 249) com¬ prises chiefly the divalent metals, which form compounds of the GROUP OF THE ALKALINE EARTHS. 303 type MeXa, and in their entire deportment exhibit many analogies. Their special relations and analogies are more closely regulated by the law of periodicity. Beryllium and magnesium belong to the small periods whose members are similar but do not show complete analogy. Beryllium exhibits many variations from magnesium, and in many properties approaches aluminium ; just as lithium attaches itself to magnesium (p. 278). The metals, calcium, strontium, and barium, constitute the second members of the three great periods, are among themselves perfectly homologous (p. 247), and in accord with their strong basic character, attach themselves to the alkali metals K, Rb, and Cs. Zinc, cadmium, and mercury, which cor¬ respond to them and constitute the second sub-group, really belong to the right, negative sides of the three great periods. They fall in with the heavy metals, are much less basic, and resemble the alkaline earth metals only in their combination types. In conse¬ quence of the double periodicity of the three great periods both sub-groups (Ca, Sr, Ba and Zn, Cd, Hg) exhibit many analogies to magnesium and beryllium. GROUP OF THE ALKALINE EARTHS. CALCIUM. STRONTIUM. BARIUM. Ca = 40. Sr = 87.5. Ba = 137, The metals of the group are termed alkaline earth metals, because their oxides attach themselves in their properties, on the one side to the oxides of the alkalies, upon the other to the real earths (alumina, etc.). They show the same gradation in properties as the elements of the potassium group, and, with respect to their atomic weights, bear the same relation to each other. With increase in atomic weight, their chemical energy and basicity become greater.* Barium decomposes water energetically, and oxidizes more readily than strontium and calcium. In accord with this, we find barium hydroxide a stronger base; it dissolves rather readily in water, does not decompose upon ignition, and absorbs carbon dioxide rapidly from the air. Barium carbonate is also very stable, fuses at a white heat, and only disengages a little carbon dioxide. Calcium hydroxide, on the other hand, dissolves with more difficulty in water, and when ignited, breaks up into water and calcium oxide ; the carbonate also yields carbon dioxide when similarly treated. In its entire character, strontium stands between barium and calcium. All these affinity relations find full expression in the heat of formation of the corresponding compounds. While the alkaline earth metals are similar to the alkalies in their * See p. 255. 304 INORGANIC CHEMISTRY. free condition and in their hydroxides, they essentially distinguish themselves from them by the insolubility of their carbonates and phosphates, and still more by their sulphates. Barium sulphate is almost insoluble in water and acids, while that of calcium dissolves in 400 parts water ; strontium sulphate occupies an intermediate position. The metals of this group do not form volatile compounds and their specific heats have not yet been determined. As the determination of the vapor densities of the elements or their volatile compounds, and the ascertainment of the specific heat of the metals, afford the only two direct means for the derivation of the true atomic weights, it was allowable to make the atomic weights of the calcium group I equal to their equivalent weights (Ca = 19.9, CaCl). But the great analogy of their compounds to those of the metals of the magnesium group, their isomor¬ phism for instance, argues with great probability for the divalence of the metals of the group, and that the present accepted double atomic weights are correct (compare p. 263). For calcium, this conclusion has already been confirmed by the experimental determination of its thermal capacity. CALCIUM. Ca = 4Q. Calcium belongs to the class of elements most widely distributed upon the earth's surface. As calcium carbonate (limestone, marble, chalk) and the sulphate (gypsum, alabaster), it represents immense deposits in all stratified formations. As phosphate, it constitutes phosphorite^ as fiuoride, fluorite^ both of which are abundant. As silicate, it is found in most of the oldest crystalline rocks. The metal is obtained by the electrolysis of the fused chloride; further, by heating calcium iodide with sodium, or calcium chloride with sodium and zinc. Although the affinity of calcium for oxygen is less than that of the alkalies, yet the oxide (also BaO and SrO) cannot be reduced to metal by ignition with carbon, iron, or sodium —due, probably to the non-fusibility of the oxide. Calcium is a yellow, shining metal, of specific gravity 1.55-1.6. In dry air it is tolerably stable, in moist it covers itself with a layer of hydroxide. It decomposes water with considerable energy. It fuses at a red heat, and in the air burns with a brilliant yellow light. Calcium Oxide—Lime—CaO—may be obtained pure by ignit¬ ing the nitrate or carbonate. It is prepared on a large scale by burning the ordinary limestone or marble (CaC03) in lime-kilns. It is a grayish-white mass, which does not fuse at the highest tem¬ peratures. The oxy-hydrogen ñame thrown upon a piece of lime CALCIUM HYDROXIDE—CALCIUM PEROXIDE. 305 causes it to emit an extremely intense white light (Drummond's Lime Light). In the air lime attracts moisture and CO2, becoming calcium carbonate ; burned lime unites with water with evolution of much heat, breaking up into a white voluminous powder of calcium hydroxide Ca(OH)2—slaked lime. When limestone contains large quantities of alumina, magnesium carbonate, or other constituents, the lime from it slakes with difficulty, and is known as poor lime, to distinguish it from pure fat or rich lime, which is readily converted into a powder with water. Calcium Hydroxide—Ca(OH)2—slaked lime—is a white, por¬ ous powder, forming a thick paste, milk of lime, with water. It dissolves with difficulty in cold water (i part in 760 parts), but with still more difficulty in warm water ; the solution saturated in the cold (lime water) becomes cloudy upon warming. It has a strong alkaline reaction. In the air it attracts carbon dioxide and forms calcium carbonate. At a red heat it decomposes into oxide and water. Slaked lime is employed in the preparation of ordinary mortar, a mixture of calcium hydroxide, water and sand. The hardening of the mortar in the air depends mainly upon the fact that the calcium hydroxide combines with the CO2 of the air to form the carbonate, and at the same time acts upon the silicic acid of the sand forming a calcium silicate, which, in time, imparts durability to the mortar. Hydraulic mortar, or cement, is produced by gently igniting a mixture of limestone or chalk with aluminium silicate (clay) and quartz powder. On stirring the powdered, burnt mass with water it soon hardens, and is not dissolved by water. Some naturally occurring limestones, containing upwards of 20 per cent, clay, yield hydraulic cements, without any admixtures after burning. Their composition varies, and also the process of their hardening j it however depends principally upon the formation of calcium and aluminium silicates. Calcium Peroxide—Ca02—is precipitated as a hydrate in crystalline leaflets, if lime water be added to a solution of barium peroxide in dilute hydrochloric acid ; it is very unstable. It con¬ tains 8H2O, which it gradually loses in dry air. The halogen derivatives of calcium, like those of other metals, are prepared by the solution of the oxide or carbonate in the haloid acids. They are formed by the direct union of calcium with the halogens ; calcium burns in the vapors of chlorine, bromine and 26 306 INORGANIC CHEMISTRY. iodine. Technically, calcium chloride is often obtained as a by¬ product, e. g., in the preparation of ammonia. Calcium Chloride—CaClj—crystallizes from aqueous solution with 6 molecules of H2O, in large, six-sided prisms, which deli¬ quesce in the air. In vacuo it loses 4 molecules of íi^O. When heated, it melts in its water of crystallization, loses water, but it is only after it has been exposed above 200° that it becomes anhy¬ drous; then it is a white, porous mass. The dry salt fuses at 719°, and solidifies to a crystalline mass, which attracts water ener¬ getically, and may be employed in the drying of gases and liquids. The dry calcium chloride also absorbs ammonia, forming the com¬ pound CaClz-SNHg. The crystallized hydrous salt dissolves in water with reduction of temperature ; by mixing with snow or ice the temperature is lowered to —48°. Upon fusing the dry chloride in air it will partially decompose into the oxide and hydrogen chloride. Calcium bromide and iodide are very similar to the chloride. Calcium Fluoride—CaFlj—occurs in nature as ßucrite^ in large cubes or octahedra, or even in compact masses. It is often discolored by impurities. It is found, in sparing quantities, in the ashes of plants, bones and the enamel of the teeth. A soluble fluoride added to the solution of calcium chloride throws down insoluble calcium fluoride as a white voluminous precipitate. The fluoride is perfectly insoluble in water, and is only decom¬ posed by strong acids. It fuses easily at a red heat, serving, there¬ fore, as a flux in the smelting of ores. When heated it phos¬ phoresces. Calcium Hypochlorite—Ca(C10)2—is not known in a pure condition. The so-called bleaching lime or chloride of lime, obtained by conducting, chlorine, at ordinary temperatures, over slaked lime, contains calcium hypochlorite as active principle. According to analogy to the action of chlorine upon potassium, or sodium hydroxide, the reaction in the case of calcium hydroxide may be expressed by the following equation :— 2Ca(OH)2 -f 2CI2 = Ca(OCl), -f CaCl, 2H,0. This would incline us to regard chloride of lime as a mixture of cal¬ cium hypochlorite, calcium chloride, and water. In accordance with the equation, the completely chlorinated chloride of lime must con¬ tain 48.9 percent, chlorine, which is never the case, because a portion of the calcium hydroxide invariably remains unaltered. Calcium chloride does not exist free in bleaching lime, because it is not with¬ drawn from the latter by alcohol, and nearly all the chlorine of the bleaching lime can be expelled by carbon dioxide. It is there- CALCIUM HYPOCHLORITE. 307 /Cl fore probable that the compound, Ca (Odling and Lunge), \OCl is present in bleaching lime :— Ca(OCl)j -f CaClj = 2Ca<^Q^,j Chloride of lime is a white, porous powder with an odor resem¬ bling that of chlorine. The aqueous solution has a strong alkaline reaction, and bleaches. It decomposes in the air as the carbon dioxide of the latter liberates hypochlorous acid. Even in closed vessels it gradually breaks up, with elimination of oxygen ; the de¬ composition is hastened by sunlight and heat, and may occur with explosion. Hence chloride of lime should be preserved in loosely closed vessels, in a cool dark place. Dilute hydrochloric or sulphuric acid will expel chlorine from chlo¬ ride of lime ; the quantity liberated is just twice that which the hypo¬ chlorite eventually found in it contains :— Ca(C10)2 4- 4HCI = CaCla + aH^O When sulphuric acid acts, the calcium chloride present partici¬ pates in the reaction :— Ca(C10)2 + CaCl^ -f aHaSO^ == aCaSO^ + aClg + aH^O. The application of chloride of lime for the production of chlorine in chlorine bleaching and disinfection is based on this deportment. The quantity of chlorine set free by acids from the chloride of lime represents its quantity of so-called active chlorine ; good chloride of lime should contain at least 25 per cent. Calcium chlorate and chloride are produced when the aqueous solution of chloride of lime is boiled :— 3Ca(C10)3 = (ClOsjaCa + 2C2ia^. On this is based the application of chloride of lime for the production of potas¬ sium chlorate (KCIO3) by a transposition of calcium chlorate with potassium chloride. When a small quantity of cobaltic oxide is added to the solution of bleaching lime, and heat applied, a regular stream of oxygen is disengaged; this is an ad¬ vantageous method of preparing oxygen. Other oxides, like those of manganese, copper, and iron, behave similarly. In this reaction there occurs, apparently, a contact action of the oxides. The reaction is explained, doubtless, in the same way as the action of hydrogen peroxide upon certain oxides (see p. 103). The feebly combined oxygen atom in cobaltic oxide unites with the oxygen of the cal¬ cium hypochlorite to form free oxygen :— Ca(C10)3 4- 2C02O3 = CaClj -f 2O3 4- 4C0O. Cobaltic Cobaltous oxide. oxide. 308 INORGANIC CHEMISTRY. The resulting cobaltous oxide is again converted by the chloride of lime into cobaltic oxide, which acts upon a fresh quantity of bleaching lime. On wanning bleaching lime with ammonia, the following decomposition occurs:— SCaOClj + 2NH3 = aCaClj + + 2N. This reaction can be used for the preparation of nitrogen. Calcium Sulphate—CaSO^—is very abundant in nature. In an anhydrous condition it forms the mineral anhydrite^ crystallizing in forms of the rhombic system. With two molecules of water it occurs as gypsum, in large monoclinic crystals or in granular, crys¬ talline masses (Alabaster, etc. ). It also separates as a fine crys¬ talline powder, CaSO^ -j- 2H2O, when soluble calcium salts are precipitated with sulphuric acid. Calcium sulphate dissolves with difficulty in water ; i part at average temperatures is soluble in 400 parts H2O. When heated to 110° gypsum loses all its water, and becomes burnt gypsum ; when this is pulverized and mixed with water, it forms a paste which hardens to a solid mass in a short time. The hardening is dependent upon the reunion of anhydrous calcium sulphate with 2 molecules of H2O. On this depends the use of burned gypsum for the production of moulds, figures, etc. In case gypsum has been too intensely heated (above 160°, dead- burnt gypsum) it will no longer harden with water ; the naturally occurring anhydrite behaves in the same manner. Calcium Nitrate—Ca(N03)2—is produced by the decay of the nitrogenous organic substances in the presence of lime, there¬ fore, it is frequently found as an efflorescence upon walls (in cattle stables). It crystallizes from water in monoclinic prisms, having four molecules of water ; the anhydrous salt deliquesces in the air. By the action of potassium carbonate or chloride, calcium nitrate may be transposed into potassium nitrate (p. 285). Calcium Phosphates. The tertiary phosphate—Ca3(P04)2— is found in slight quantities in most of the mountain rocks. In combination with calcium fluoride, it crystallizes as apatite. As phosphorite, it forms compact masses, more or less intimately mixed with other constituents, and occurs in immense deposits in Spain, France, Germany, and Russia. When these minerals disintegrate the calcium phosphate passes into the soil, and is absorbed by the plants. In the latter, it accumulates chiefly in the seeds and grains. In the animal kingdom, it is principally found in the bones, the ashes of which contain upwards of 85 per cent, calcium phosphate. Tertiary calcium phosphate is perfectly insoluble in water. If disodium phosphate be added to the aqueous solution of a calcium salt, and then ammonium hydroxide, it will separate as a gelatinous precipitate, which, after drying, becomes a white amorphous pow¬ der. It is very readily soluble in acids, even acetic. The secondary calcium phosphate—PO^CaH zHjO—is some¬ times present in guano, in the form of small, shining prisms, and CALCIUM CARBONATE. 309 separates as an amorphoos precipitate, if disodium phosphate be added to a solution of calcium chloride mixed with some acetic acid. When ignited, it passes into calcium pyrophosphate, PaOiCa,. primary phosphate—Ca(H2F04)2—is produced by the action of sulphuric or hydrochloric acid upon the first two phosphates. It is readily soluble in water, and deliquesces in the air. Heated to 200°, it decomposes into pyrophosphate, metaphosphoric acid and wâtcr * 2Ca(H¡,POd2 = Ca^PjO^ + 2HPO3 -}- sH^O. When intensely ignited, pure calcium metaphosphate remains (P- 135)- Calcium phosphate is present in all plants. Its presence in the soil is, therefore, an indispensable condition for its fertility. When there is a scarcity of phosphoric acid it must be added. To this end, bone meal and pulverized phosphorite were formerly employed. Since, however, the phosphoric acid is contained in these substances as tri-calcium phosphate, which is not easily absorbed by the plants, the primary phosphate is extensively employed at present as a fertilizer, or, better, the mixture resulting from the action of sulphuric acid upon the tertiary salt. Superphosphate is the name applied to the resulting mass. The Thomas slag, obtained in the dephosphorization of iron ores by the Gilchrist-Thomas process, constitutes a very valuable and important crude product from which to prepare calcium phosphate. Calcium Carbonate—CaCOs—is very widely distributed in nature. It crystallizes in two crystallographic systems, hence is dimorphous. In rhombic crystals, with the specific gravity 3.0, it forms aragonite. In hexagonal rhombohedra, with specific gravity 2.7, it occurs as calcite. Iceland spar, employed for optical pur¬ poses, is perfectly pure, transparent calcite. The common calcite, which constitutes immense mountain chains, is an amorphous or indistinct crystalline stratum, and is usually mixed with other constit¬ uents, as clay. When the limestone is granular and crystalline, it is termed marble. Dolomite also constitutes large layers, and is a compound of calcium and magnesium carbonates, with generally an excess of the former. Chalk is very pure amorphous calcium carbonate, consisting of the shells of microscopic sea animals. Cal¬ cium carbonate is, further, a regular constituent of all plants and animals; the shells of eggs, of mussels, even corals and pearls, consist chieñy of it. A soluble carbonate, added to the aqueous solution of a calcium salt, precipitates calcium carbonate as a white, amorphous powder, which soon becomes crystalline. In the cold, it assumes the form of calcite ; upon boiling the liquid, it generally changes into aragonite crystals. The carbonate is almost insoluble in pure water ; but dissolves somewhat in water containing carbon dioxide, as it very probably is changed to the primary carbonate—Ca(HC08)a. For this reason. 310 INORGANIC CHEMISTRY. we find calcium carbonate dissolved in all natural waters. When the solution stands exposed, or if it be heated, carbon dioxide escapes, and the secondary carbonate again separates out. The formation of lime scales, thermal tufts, stalactites, boiler scales and similar deposits are due to this. Calcium carbonate, like all car¬ bonates, is decomposed by acids, with evolution of carbon dioxide. At a red heat, it dissociates into CaO and COj. The change begins at 600°. The tension of dissociation is 85 mm. at 860°, and at 1040® 510-520 mm. For this reason calcium carbonate is not decom¬ posed when heated in a sealed tube. Calcium Silicate—CaSiOg—occurs as white, crystalline wol- lastonite. It is also a constituent of most natural silicates and of the artificial silicate fusion—^glass. Glass.—The silicates of potassium and sodium are readily fusible and soluble in water. The silicates of calcium and the other alkaline earths are insoluble, very difiScultly fusible, and generally crystallize when they cool. If, however, the two silicates be fused together, an amorphous, transparent mass, of average fusibility, results ; it is only slightly attacked by water and acids—^it is glass. To prepare the latter, a mixture of sand, lime, and soda, or potash, is heated to fusion in a muffle furnace. Instead of the carlx>nates of p)otassium and sodium a mix¬ ture of sulphates with charcoal can be employed ; the carbon reduces the sul¬ phates to sulphides, which form silicates when fused with silicon dioxide. The following are the varieties of glass :— Soda Glass—a mixture of sodium and calcium silicates—fuses readily, and is employed for window-panes and ordinary glass vessels. Potash or Bohemian GlasSf also called Crown Glass, consists of calcium and potassium silicates, is not very fusible, is harder and withstands the action of water and acids better than soda glass; it is, therefore, employed in the manufacture of chemical glassware. Glass Crystal or Flint Glass is composed of potassium and lead silicates. It is not as hard, fuses with tolerable readiness, refracts light strongly, and when pol¬ ished, acquires a clear lustre. On this account it is employed for optical purposes (for lenses, prisms) and is used in ornamental glassware. Strass y—a lead glass containing boron trioxide, is used to imitate precious stones. The opaque varie¬ ties of enamel consist of lead glass and in the fused glass are insoluble admixtures, as tin dioxide and calcium phosphate. Ordinary window glass is obtained by the fusion of rather impure materials; in consequence of the presence of ferrous oxide it is ordinarily colored green. To remove this coloration, manganese peroxide, MnOg, is added to the fusion. It oxidizes a portion of the ferrous to ferric oxide, the silicate of which is colored slightly yellow, while manganese forms a violet silicate. These colors, violet and green, almost neutralize each other as complementaries. The colored glasses con¬ tain silicates of colored metallic oxides; chromic and cupric oxides color green; cobaltic, blue; cuprous oxide, a ruby red, etc. The sulphur compounds of calcium are very much like those of the alkalies. Calcium Sulphide—CaS—is most readily obtained by heating the sulphate with carbon, and is a whitish-yellöw mass. When it is dissolved in water we get Calcium Hydrosulphide STRONTIUM. 311 —Ca(SH)8—^which decomposes on boiling the aqueous solution. When calcium oxide is ignited with sulphur in a closed crucible a yellowish-gray mass is obtained, which consists of calcium polysul- phides and sulphate. Milk of lime boiled with sulphur yields a deep, yellow solution of calcium polysulphides. When the solu¬ tions of the latter are acted upon by acids, finely divided sulphur —milk of sulphur—is precipitated and HaS set free. If the reverse occur, viz., the addition of a solution of polysulphides to an ex¬ cess of diliite acids, hydrogen persulphide will separate (p. ni). STRONTIUM. Sr = 87.5. This element is rather rare in nature, and is principally found in strontianite (strontium carbonate) and celestite (strontium sul¬ phate). Its compounds are very similar to those of calcium. The metal is obtained by the electrolysis of fused strontium chlo¬ ride. It has a brass-yellow color, and a specific gravity, 2.5. It oxidizes in the air and burns with a bright light when heated. It decomposes water at the ordinary temperatures. Of the compounds of strontium we may mention the following :— Strontium Oxide—SrO—is most readily obtained by igniting the nitrate. It unites with water, with strong evolution of heat, forming Strontium Hydroxide—Sr(OH)2—which is more read¬ ily soluble in water than calcium hydroxide. It crystallizes from aqueous solution with 8 or 9 molecules of HjO. When ignited it decomposes into SrO and H2O, but with more difficulty than cal¬ cium hydroxide. Strontium Chloride—SrCh -f- 6H2O—crystallizes from water in hexagonal tables, which deliquesce in the air; it is somewhat soluble in alcohol. Strontium Sulphate—SrSO*—is much more difficultly soluble in water than calcium sulphate, but is not as insoluble as barium sulphate. Strontium Nitrate—Sr(N03)2—is obtained by dissolving the carbonate in nitric acid, and is readily soluble in water, but insolu¬ ble in alcohol. It crystallizes from warm solutions in anhydrous octahedra, but from cold, with 4 molecules of HjO, in monoclinic prisms. Mixed with combustible substances it colors the flame a beautiful carmine red, and for this reason is employed in pyrotechny. Strontium Carbonate—SrCOs—is precipitated from aqueous solutions of strontium salts by soluble carbonates, as an amorphous, insoluble powder. When ignited it breaks up into SrO and CO2. This decomposition does not, however, occur so easily as with cal¬ cium carbonate. 312 INORGANIC CHEMISTRY. BARIUM. Ba = Ï37. Barium occurs in nature in large masses, as heavy spar (or barium sulphate), and as witherite (barium carbonate). All its compounds are distinguished by their high specific gravity, hence the name barium, from ßapuso,. NHj.HCl NH3CI NH3/ Diamide Diammonium- Diammonium Hydrochloride. Chloride. Sulphate. The similarity of hypothetical diammonium to the alkaline earth metals is shown in the insolubility of its sulphate. Diammonium Chloride^ Nj¡H^(HCl)2, hydrazine hydrochloride, is obtained from the sulphate by transposition with barium chloride. It is very readily soluble in cold water, consists of large isometric crystals, and melts with decomposition at about 200°. Diammonium Sulphate, NjH^.SO^Hj, hydrazine sulphate, dissolves with diflS- culty in cold water, and crystallizes in vitreous plates. It deflagrates when exposed to a strong heat. The powerful reducing property of the hydrazine compounds is characteristic oí them. They precipitate metallic silver from an ammoniacal silver nitrate solu- tion, and cuprous oxide and metallic copper from an alkaline copper solution (Fehling's solution). METALS OF THE MAGNESIUM GROUP. In this group are usually included beryllium, magnesium, zinc, and cadmium. However, these metals do not exhibit complete analogy, as is clearly seen in the periodic system (p. 302). Beryl¬ lium shows the greatest variations. It approaches aluminium, while magnesium resembles not only zinc and cadmium, but also the alkaline earth-metals, calcium, strontium, and barium. Its simi¬ larity to the latter is expressed by the basic nature of its oxide, whereas it resembles zinc and cadmium mainly in isomorphism of compounds. Beryllium and magnesium bear the same relation to Ca, Sr, and Ba, as lithium and sodium bear to the metals of the potassium group. MAGNESIUM GROUP. 315 The alkaline character of the alkaline earth-metals gradually diminishes from barium to calcium, and becomes almost nothing in magnesium and beryllium, which possess the lowest atomic weights (see p. 302). Magnesium and beryllium are scarcely capable of decomposing water, even at boiling temperatures. Their oxides and hydroxides are almost insoluble in it; the hydroxides decompose, on gentle ignition, into oxides and water. Their carbonates are very unstable ; their chlorides, too, suffer a partial decomposition into oxide and hydrogen chloride, even on drying. The solubility of the sulphates of magnesium and beryllium further distinguishes them from the metals of the alkaline earth group. The specific properties of beryllium and magnesium are main¬ tained in zinc and cadmium, which constitute a natural group with the former. Zinc and cadmium do not decompose water at a boil¬ ing heat ; their hydroxides are insoluble in it, and are not very stable j their carbonates and chlorides easily undergo decomposi¬ tion ; their sulphates are readily soluble in water. The similarity is further expressed by the isomorphism of most of their compounds. Thus, magnesium and zinc sulphates crystallize with 7 molecules of H2O, in perfectly similar forms. If the solution of a mixture of both salts be allowed to crystallize, we get crystals with variable quantities of zinc and magnesium : the formation of such isomor- phous mixtures in ad libitum proportions, is a characteristic indica¬ tion of the isomorphism of compounds chemically similar. The difference between beryllium and magnesium upon the one side, and zinc and cadmium on the other, is shown distinctly in their specific gravities. While the first two elements possess a low specific gravity (Be—2.1, Mg—1-75), zinc and cadmium (with specific gravities 7.2 and 8.6) belong to the so-called heavy metals (see p. 259). The difference in specific gravity determines, also, many differ¬ ences in chemical character. The light metals (especially the alkalies and alkaline earths) form rather unstable sulphides, readily soluble in water, while the sulphides of zinc and cadmium, like those of all heavy metals, are insoluble in water, and, usually, in acids ; in these respects, magnesium and beryllium behave like the alkalies, while zinc and cadmium are precipitated by hydrogen sulphide or alkaline sulphides as sulphides from solutions of their salts. Further, the oxides of the light metals are very stable, and are only reduced by carbon if they are readily fusible (like potassium and sodium ox¬ ides) ; the heavy metals, on the other hand, are easily separated from their oxides by carbon. Zinc and cadmium oxides are reduced by carbon, while those of magnesium and beryllium are not altered. See further, p. 255. All these affinity relations are more clearly expressed and explained in their thermo-chemical relations (p. 323). 316 INORGANIC CHEMISTRY. MAGNESIUM. Mg == 24.38. Magnesium is abundant in nature, and almost always accomimnies calcium in its compounds. As carbonate, it occurs in compact masses, as magnesite, etc. Dolomite, which forms entire moun¬ tains, is an isomorphous mixture of calcium and magnesium carbon¬ ates. Magnesium is also present in most of the natural silicates; its soluble salts are contained in almost all natural waters. Metallic magnesium may be obtained by the electrolysis of the chloride, or by heating the same with sodium. It was formerly prepared by heating the double chloride of magnesium and sod¬ ium with metallic sodium and fluorspar, the latter serving merely as a flux:— MgCl2.NaCl + 2Na = aNaCl + Mg. At present magnesium is obtained in large quantities by the method of Grätzel, which consists in electrolyzing the chloride, heated to fusion in crucibles. Magnesium is a brightly shining, almost silver-white metal, of specific gravity T.75. It is tenacious and ductile, and when heated may be converted into wire and rolled out into thin ribbons. It melts at about 800°, and distils at a bright-red heat. At ordinary temperatures, it scarcely oxidizes in the air ; it burns, when heated, with an extremely intense white light, owing to the glowing non¬ volatile magnesium oxide. Magnesium light is rich in chemically active rays, and, for this reason, it is employed for photographing in dark chambers. Its alloy with zinc is generally employed as a substitute for pure magnesium, as it burns with an equally bright light. Boiling water is very slowly decomposed by magnesium. It dissolves easily in dilute acids, forming salts ; the alkalies do not attack it. Magnesium Oxide—MgO—or magnesia, formed by the com¬ bustion of magnesium, is ordinarily obtained by the ignition of the hydroxide or the carbonate {magnesia ustd). It is a white, very voluminous, amorphous powder, which finds application in medi¬ cine. The feebly ignited magnesia combines with water, with slight generation of heat, to produce magnesium hydroxide. Magnesium Hydroxide—Mg(HO)2—is precipitated from so¬ lutions of magnesium salts by potassium or sodium hydroxide as a gelatinous mass. Dried at 100° it is a white amorphous powder. It is almost insoluble in water and alkalies ; moist litmus paper is, however, colored blue. Ammonium salts dissolve it quite easily, forming soluble double salts. Magnesium hydroxide attracts carbon dioxide from the air and forms magnesium carbonate. It yields the oxide and water when gently ignited. .Magnesium Chloride—MgClj—is present in traces in many mineral springs. It may be prepared by the solution of the car¬ bonate or oxide in hydrochloric add ; in large quantities it is ob¬ tained as a by-product in the technical production of potassium MAGNESIUM SULPHATE. 317 chloride. When its solution is evaporated the salt crystallizes out with six molecules of HjO in deliquescent crystals, isomorphous with calcium chloride. When these are heated they give up water, and there occurs at the same time a partial decomposition of the chloride into oxide and hydrogen chloride :— MgQa + H20 = MgO + 2HCI. As magnesium chloride is produced in large quantities in various technical pro¬ cesses, repeated efforts have been made to utilize the above reaction for the pre¬ paration of hydrochloric acid, by conducting steam over heated magnesium chloride. However, the thermal relations of the reaction indicate that this could only be accomplished with difficulty. From the chemical affinities coming into play, the reaction pursues an opposite course, as the magnesium oxide is readily decomposed by hydrochloric acid into magnesium chloride and water :— MgO + 2HCI = MgClg + HjO. This is because the heat of formation of MgClj (iS^-o C.) and steam (58.0) is greater than that of MgO (145.8) and 2HCI (44.0). The decomposition is even easier in the presence of water, as is evident from the thermal numbers. The reverse reaction is, therefore, endothermic, requires the addition of chemical energy in the form of heat, and like all similar reactions is incomplete. To get anhydrous magnesium chloride ammonium chloride is added to the solution of the former. The double salt, MgClj. NH4CI + 6H2O, is formed. When this is heated it first loses water, and at 460° throws off ammonium chloride, leaving anhydrous mag¬ nesium chloride. This is a le.afy, crystalline mass, which fuses at 708°, and distils undecomposed at a red heat ; it is very deliques¬ cent in the air. Double salts, similar to the above, are also formed from potas¬ sium and calcium chloride. The potassium double salt—MgCl,. KCl -j- 6H2O—occurs in considerable deposits as carnallite at Stassfurt. Magnesium Sulphate—MgSOi—is found in sea-water and in many mineral springs. With more or less water it is kieserite, which abounds extensively at Stassfurt. At ordinary temperatures it crystallizes with 7 molecules of H2O—MgS04 -j- 7H2O—in four- sided rhombic prisms, readily soluble in water (at 0° in 2 parts water). It has a bitter, salt-like taste, and serves as an aperient. It crystallizes with 6 molecules of HjO from solutions heated to 70°; at 0°, however, it has 12 molecules. When heated to 150° these hydrates lose all their water of crystallization, excepting one molecule y which escapes above 200°. One molecule of water, in magnesium sulphate is, therefore, more closely combined than the rest. Many other salts containing water deport themselves simi¬ larly. The more intimately combined water is termed Water of Constitution. Magnesium sulphate forms double salts with potassium and ammonium sulphates, which crystallize with 6 molecules of H^O in monoclinic prisms, e.g. :— 318 INORGANIC CHEMISTRY. MgSO^.KjSO^ + 6H2O. The sulphates of zinc and several other metals, e.g., iron, cobalt, and nickel, in their divalent forms, are very similar to magnesium sulphate. Their sulphates crystallize with 7 molecules of HgO, are isomorphous, and contain i molecule of intimately combined water. They form double salts with potassium and ammon¬ ium sulphates; these crystallize with 6H2O, and are isomorphous; e.g. :— ZnSO^ + 7H2O ZnS04.K2S04 + 6H2O. FeSO^ + 7H2O FeS04.K2S04 + 6H2O. The constitution of these double salts may be viewed in the same way as that of potassium-sodium sulphate, or of mixed salts of polybasic acids. We may sup¬ pose that in the given instance the divalent metal unites two molecules of sulphuric acid :— SO4 ^Mg + 6H2O. SO4 \k Magnesium Phosphates.—The tertiary phosphate (P04)2Mg9, accompanies the tertiary calcium phosphates in small quantities in bones and in plant ashes. The secondary phosphate, MgHPO^ -f" 7H2O, is precipitated from the soluble magnesium salts, by disodium phosphate (Na2HP04) as a salt dissolving with difficulty in water. If ammonium salts be present, the precipitated double salt will be magnesium ammonium phosphate, MgNH4P04 -{- 6H2O, insoluble in water. The latter is found in guano, forms in the decay of urine, and is sometimes the cause of the formation of calculi. The primary salt, H4Mg(P04)2, has not been obtained. The magnesium salts of arsenic acid, H3ASO4, are very similar to those of phosphoric acid. Magnesium-ammonium arseniate (MgNHiAsOi 6H2O) is likewise almost insoluble in water. Magnesium Carbonate, MgCOs, occurs in nature as magne¬ sium spar, crystallized in rhombohedra (isomorphous with calcite), in compact masses as magnesite. Combined with calcium carbon¬ ate, it forms dolomite, to which, when pure, is ascribed the formula, CaCOs.MgCOsj however, it usually contains an excess of calcium carbonate. On adding sodium or potassium carbonate to the aqueous solution of a magnesium salt, some carbon dioxide escapes, and a white precipitate forms, which consists of a mixture of mag¬ nesium carbonate and hydroxide. If the predipitate be dried at low temperature, we obtain a white, voluminous powder, whose com¬ position generally corresponds to the formula Mg(0H)8.4C03Mg -f- 4H2O. This salt is employed under the name Magnesia alba in medicine. If it be suspended in water, and carbon dioxide passed through it, the salt will dissolve, and upon standing exposed to the air, crystals of neutral carbonate, MgCOs -j- 3HjO, separate. When BERYLLIUM. 319 these are boiled with water they give up carbon dioxide and are again converted into the basic carbonate. The naturally occurring magnesite sustains no change when boiled, and it is only when it is heated above 300° that it decomposes into MgO and CO2. Magnesium carbonate yields isomorphous double salts, with potas¬ sium and ammonium carbonate; e. g., MgCOs-KjCOs -(- 4H2O. Of the silicates of magnesium, we may mention olivine (Mg2Si04)» serpentine (MgaSijOT -|- 2H2O), talc (SÍ50i4Mg4), sepiolite (SijOgMg, -1- 2H2O), or meerschaum. The mixed silicates of magnesium and calcium are very numerous ; to these belong asbestos, the augites and hornblendes. Recognition of Magnesium Compounds.—The fixed alka¬ line hydroxides precipitate magnesium hydroxide from magnesium salts ; the carbonates throw down basic magnesium carbonate. The precipitates are insoluble in pure water and the alkalies, but dissolve readily in solutions of ammonium salts. In the presence of the latter, neither the alkaline hydroxides nor carbonates cause precipi¬ tation. In presence of ammonium salts, disodium phosphate pre¬ cipitates magnesium-ammonium phosphate, MgNHiPO* + 6H2O, insoluble in water. BERYLLIUM. Be == 9.1. Among the metals of the second group beryllium occupies a position similar to that of lithium in the first group ; in both elements, which have the lowest atomic weight in their group, the specific group character is considerably diminished, or does not find expression. As lithium attaches itself in many respects to magnesium, so does beryllium approach aluminium. Like the latter, it is scarcely at all attacked by nitric acid, but dissolves easily in sodium or potassium hydroxide, with elimi¬ nation of hydrogen. Like aluminium oxide, that of beryllium dissolves in the alkalies, and is almost invariably accompanied by the former in its natural com¬ pounds. Beryllium sulphate, like that of aluminium, forms a difficultly soluble double salt with potassium sulphate. However, beryllium, in most of its com¬ pounds, stands nearer to magnesium than to aluminium. The determination of the vapor density of beryllium chloride (see below) has finally established the atomic weight and the valence of this element. Beryllium is not very abundant in nature and is found principally in beryl, a double silicate of aluminium and beryllium—Al3Bej(Si03)g. Emerald has the same composition, and is only colored green by a slight amount of chromium oxide. Metallic beryllium is obtained by the ignition of the chloride with sodium, and is a white ductile metal, of specific gravity 1.64. Its specific heat equals 0.4084; the atomic heat is, therefore, 3.8 (p. 261). It does not decompose water, even 320 INORGANIC CHEMISTRY. upon boiling. It does not oxidize in the air at ordinary temperatures. When finely divided it will burn in the air with a very bright light when heated. In a com¬ pact mass (like magnesium), it does not do this. It is readily dissolved by dilute hydrochloric and sulphuric acids ; also by potassium and sodium hydroxides. Beryllium Chloride—BeClj—^is obtained, like aluminium chloride, by the ignition of a mixture of beryllium oxide and carbon in a stream of chlorine. It sublimes in shining needles, which deliquesce in the air. Its vapor density equals 2.8 (air = i) or 40.3 (H = i) at 680-800° C., corresponding to the molecular formula BeClj = 79 9 (Nilson.) It crystallizes from aqueous solution with four molecules of HgO ; upon dr3dng it suffers a decomposition similar to that of mag¬ nesium chloride. The salts of beryllium have a sweet taste, hence it has been called glucinum. Ammonium hydroxide precipitates a white, gelatinous beryllium hydroxide, Be(OH)2, from solutions of the soluble salts. This dissolves readily in sodium and potassium hydroxide, but on boiling, separates again from solution. When heated, the hydroxide breaks up into water and beryllitim oxide, BeO, which is a white, amorphous powder, of specific gravity 3.06. Its sp>ecific heat equals o. 2471. Beryllium Sulphate—BeSO^—crystallizes from water at various temperatures, with four or seven molecules of H2O, of which one is rather closely combined. It crystallizes with magnesium sulphate in an isomorphous mixture. The double salt, SO^Be.SO^Kj + sHgO, does not dissolve readily in water ; in this respect it resembles the alums. ZINC. Zn = 65.5. The natural compounds of the heavy metals have generally a high specific gravity, frequently possess metallic lustre, usually occur in the older crystalline rocks in veins, and are termed ores. The most important zinc ores are the carbonate—ZnCOa—the silicate, and sphalerite or blende, ZnS. The principal sources of these ores are in Silesia, England, Belgium, Poland and the United States. To get the metal the carbonate or sulphide is converted into oxide by roasting in the air ; the product is then mixed with carbon and ignited in cylindrical clay tubes. In this manner the oxide is reduced :— ZnO C = Zn -f- CO» and the liberated zinc distilled off. The receivers contain the fused, compact zinc and a gray, pulverulent mass, called zinc dust, which consists of a mixture of zinc oxide with finely divided metal. This material is used in laboratories as a strong reducing agent. Metallic zinc has a bluish-white color, and exhibits rough, crys¬ talline fracture; its specific gravity equals 7-7.2. At ordinary temperatures it is brittle and can be pulverized; at 100-150° it is malleable and can be rolled into thin leaves and drawn out into wire. At 200° it becomes brittle again and may be easily broken. It fuses at 412° and distils about 1000°. It becomes coated with a thin layer of basic carbonate in moist ZINC HYDROXIDE—ZINC CARBONATE. 321 air. Heated in the air it burns to zinc oxide with a very intense, bluish-white light. Compact zinc will only decompose water at a red heat ; zinc dust, however, acts at ordinary temperatures. Zinc is readily soluble in dilute acids, and dissolves with liberation of hydrogen in potassium or sodium hydroxide, as well as in ammonia, when the solutions are boiled. Owing to its slight alteration in the air zinc meets with extensive application as sheet-zinc for coating statues and in architectural adornment, and in galvanizing sheet-iron. It also forms an import¬ ant constituent of many valuable alloys, such as brass and argentan (see these). Zinc Hydroxide—Zn(0H)2—is precipitated as a white amor¬ phous powder, from aqueous solution, by alkalies, and is soluble in excess of the reagent. When heated it decomposes into water and zinc oxide. Zinc Oxide—ZnO—is usually prepared by igniting the pre¬ cipitated basic carbonate, and, as zinc white, is employed as a stable white paint. The oxide obtained by burning the metal is a white, voluminous, flocculent mass, called flores Zinci or Lana philo' sophica. When zinc oxide is heated it acquires a yellow color, which disappears on cooling. Zinc oxide occurs in nature as zincite, colored by impurities. Zinc Chloride—ZnClg—anhydrous, is obtained by heating zinc in a stream of chlorine, by the evaporation of the solution of zinc in hydrochloric acid, and by the distillation of zinc sulphate with calcium chloride. It forms a white, deliquescent mass, which fuses when heated and distils about 680°. When the aqueous solu¬ tion of zinc chloride is evaporated it partially decomposes (like magnesium chloride) into zinc oxide and hydrochloric acid. When the concentrated zinc chloride is mixed with zinc oxide, a plastic mass is obtained, which hardens rapidly ; a mixture of magnesium chloride and oxide does the same. In both instances the harden¬ ing depends upon the formation of basic oxy-chlorides, e. g., ZnClOH. Zinc chloride forms deliquescent double salts with the alkaline chlorides, e. g., ZnClg-zKCl. With ammonia it yields various compounds, of which ZnClj-NHa is characterized by great stability. Zinc Sulphate—ZnSO^—is obtained by dissolving zinc in sulphuric acid. It is prepared upon a large scale by a gentle roast¬ ing of zinc blende (ZnS) ; the zinc sulphate is extracted by water. It crystallizes at ordinary temperatures from aqueous solutions with 7 molecules of HjO (zinc or white vitriol) in rhombic crystals, resem¬ bling those of magnesium sulphate. It affords double salts with the alkaline sulphates ; these contain 6 molecules of water (p. 308). Zinc Carbonate—ZnCOs—occurs native as smithsonite in hexagonal crystals, isomorphous with those of calcite. Sodium car« 322 INORGANIC CHEMISTRY. bonate precipitates basic carbonates of varying composition from solutions of zinc salts. Zinc Sulphide—ZnS—is zinc blende, usually colored brown by ferric oxide or other admixtures. Ammonium sulphide pre¬ cipitates it as a white compound, from zinc solutions. Although fused zinc reacts with difficulty with sulphur, zinc dust combines with the latter in powdered form quite readily, and if the mixture be heated or struck with a hammer the union is accompanied by an explosion. Zinc sulphide is insoluble in water, but is readily dissolved by dilute acids, excepting acetic ; therefore it may be precipitated by hydrogen sulphide from zinc acetate solutions. This reaction serves to separate zinc from other metals. Zinc Silicate—ZnjSiO^ + H2O—occurs in rhombic cr3^tals as calamine. CADMIUM. Cd = 112.1. Cadmium very often accompanies zinc in its ores. As much as 5 per cent, of this metal is present in the Silesian zinc ores ; it was first discovered in these in 1819. Being more volatile than zinc, in obtaining the latter it distils off first, and may be easily separated from the first portions of the distillate. It is a white, tenacious, and rather soft metal, of specific gravity 8.6. It fuses at 315°, and boils at 770°. It does not alter much in the air. Heated, it burns with the separation of a brown smoke of cadmium oxide. It dis¬ solves with difficulty in dilute hydrochloric and sulphuric acids, but readily in nitric. Zinc throws out the metal from solutions of the soluble cadmium salts. St. Claire Deville found the specific gravity of cadmium vapors (at 1040®) to be 3.9 (air = i) or 56 (H = l). Therefore, the molecular weight of cadmium is 112. Since the atomic weight of cadmium (determined from its specific heat) is also 112, it follows that the gas molecule of cadmium consists of but one ztom. We know that the molecules of other elements in the gaseous state are composed of two or more atoms (OjjNj.P^jSg). Cadmium, therefore, forms an exception to this rule. This is also true of mercury, and perhaps, too, of other divalent metals, such as zinc. , These relations remind us of the behavior of the hydro¬ carbon residues (radicals) ; while the divalent or tetravalent groups, e.g-.j ethylene C2H4 and acetylene C2H2, exist in free condition, the monovalent groups (as CHjjCN) cannot appear free, but double themselves, when separated from their compounds. Of the cadmium compounds may be mentioned ;— Cadmium Hydroxide—Cd(0H)2—is precipitated as a white powder, from the soluble cadmium salts, by the alkalies ; it is in¬ soluble in sodium and potassium hydroxides, but dissolves readily in ammonium hydroxide. Cadmium Oxide—CdO—is prepared by igniting the nitrate. It is a brownish-black powder, consisting of microscopic octahedra. CADMIUM. 323 Cadmous Hydroxide, CdOH, and its Oxide, CdjO, have been prepared. The first is a grayish white compound, while the second is composed of yellow translucent crystals. It is obtained on heating the hydroxide to a temperature at which sulphuric acid gives off dense white fumes. The hydroxide is a reducing agent, yielding hydrogen with hydrochloric acid and oxides of nitrogen with nitric acid. The oxide conducts itself similarly. These compounds are of interest, as they foreshadow the tendency toward the formation of lower oxides, so strongly shown by mercury ^Am. Chem. Jr., 12, 493). Cadmium Chloride—CdCl,—crystallizes from aqueous solu¬ tion, with two molecules of HjO, and may be dried without decom¬ position. The anhydrous salt melts at 541° and sublimes in scales. Cadmium Iodide—Cdia—is obtained by the direct action of iodine upon metallic cadmium in the presence of water. It crystal¬ lizes from the latter in hexagonal tables. It is used in photography. Cadmium Sulphate—CdSOi—crystallizes from water, not like the sulphates of zinc and magnesium, with 7 molecules of HjO, but with f HjO ; the crystals effloresce in the air. It, however, forms double salts with the sulphates of the alkali metals, e. g., CdSO^.KjSO^ 6H2O; these are perfectly analogous to those of zinc and magnesium, and isomorphous with them (p. 318). Cadmium Sulphide—CdS—occurs native as greenockite, in yellow hexagonal prisms. Hydrogen sulphide precipitates it from cadmium salt solutions as a yellow powder, insoluble in dilute acids. It is employed as a pigment. Almost all the alloys of cadmium have a low fusion temperature. Freshly prepared cadmium amalgam is a white plastic mass, which soon becomes hard. It is used in filling teeth. The chemical energy of cadmium is less than that of zinc ; this is evident from the fact that the former may be displaced from its salts by the latter. We saw that, with the elements of the groups of potassium and calcium, the chemical energy increases inversely with the increasing atomic weight ; caesium is more energetic than rubidium, barium more than calcium. It is worthy of remark that nearly all of the more electro-negative elements, belonging to the second sub¬ groups of the seven main groups of the periodic system, exhibit a diminution in chemical energy with rising atomic weight similar to that shown by the members of the magnesium group ; copi>er displaces silver ; phosphorus is more energetic than arsenic and antimony; sulphur more energetic than selenium and silver; chlorine sets free or displaces bromine and iodine. These relations of affinity find full expression in the thermo-chemical phenom¬ ena in which are clearly shown the double periodicity of the great periods and the relations of the two sub-groups, Ca, Sr, Ba and Zn, Cd, Hg, to magnesium. The basic character increasing from Mg to Ba corresponds to tbe increase in heat developed by the formation of their compounds, i.^., the chlorides, hydroxides, and sulphydrates. fMg.CU = 151.0. (Ca,Cl2) = 170.2. iSr,(X) = 184.5. (B^»Clj) — 194*5* (Mg.O.HjO) = 148.9. (Ca,0,Aq.) = 149.4. (Sr,0,Aq.) =157*7* (Ba,0,Aq.) == 158.2. Mg,S,Aq.) = Ca,S,Aq.) = 98.3. Sr,S,Aq.) = 106.6. (Ba,S,Aq.) = 107. i. 324 INORGANIC CHEMISTRY. That the increase with the hydroxides is so slight is explained probably by the decreasing solubility of the same from Ba to Mg, inasmuch as an evolution of heat (heat of precipitation) corresponds to the difficult solubility. The heat of forma¬ tion of the carbonates (from metallic oxides and carbon dioxide) must also be introduced here :— (CaOjCOj) = 42.5 (SrOjCOj) =r 53.2 (BaOjCOj) = 55.9. These seem to indicate that calcium carbonate is less stable and more easily decomposed than barium carbonate (p. 303). The series Mg, Zn, Cd, Hg departs itself differently. In this the heat disengage¬ ment becomes successively less and corresponds with the diminishing basicity (Mg,Cl2) = 151.0. (Mg,0) = 145-0. (Mg,S) = (Zn,Cl2) = 97.2. (Zn,0) = 86.4. (Zn,S). = 41.3- (Cd,Cl2) = 93-2. (Cd,0) = 66.4. (Cd,S) = 33.9. (Hg,c4) = 63.1. (Hg,0) = 30.6. (Hg,S) = 16.8. Comparing these numbers with the quantity of heat which is disengaged in the formation of aqueous hydrochloric acid (H,Cl,Aq. =: 39.3), we find explained the behavior of the metals toward this acid. All metals liberating a greater quan¬ tity of heat than 39.3 C. in the formation of their chlorides (calculated for I equivalent of metal) are in condition to decompose the dilute acid. Most of the metals belong to this class ; mercury, copper, silver, gold, lead, thallium, and some others, set free a less amount of heat, and hence are not able to decompose dilute hydrochloric acid (see p. 270). The slight quantity of heat developed in the formation of hydrc^en sulphide (S,H2 = 4.5) indicates that the same is readily decomposed by all the metals. In the same way, by adding the heat of solution (S,H2, Aq. = 9.2), we can easily ascertain which metals are precipitated by hydrogen sulphide from their chloride, etc. If in the thermo-chemical equation, (Me, Cl2.Aq.) + (S,H2Aq.) = (Me, S) -f 2(H, CI, Aq.), the sum of the heat developed upon the right side is greater than that upon the left, the reaction will occur (precipitation of metallic sulphides) ; in the oppo¬ site case the sulphide is decomposed by the dilute hydrochloric acid. The magnitude of the atomic weight of Mercury would place the latter in the group of zinc and cadmium. The relationship of these three heavy metals is observed in many similarities of the free ele¬ ments and of their compounds (p. 325). Occupying a similar position in the three great periods (p. 248) they are distinguished among the heterologous members in a physical point of view by their ready fusibility and volatility, which nearly reach a maximum in them. In the homologous series, Zn, Cd, Hg, these properties, like the specific gravities, increase with rising atomic weight (just as with the metals of the potassium group, p. 277) :— MERCURY. 325 Zn Cd Hg Atomic weight Fusing point 65-5 412° 940° 7-1 1I2.I 315'' 765° 8.6 200.4 —40° 360® 13.6 Boiling point Specific gravity The gradation in the heat of formation of their compounds (p. 324) clearly indicates that mercury must be arranged in a group with cadmium and zinc. Liike zinc and cadmium, it yields compounds of the form, HgXa, in which it appears divalent. These derivatives are, in many respects, similar to the corresponding compounds of zinc and cad¬ mium. Thus, mercuric sulphate affords double salts with the alka¬ line sulphates, which crystallize with six molecules of H2O (SO^Hg. SO4K2 4- 6H2O), and are isomorphous with the double sulphates of the magnesium group (p. 318). The similarity, however, limits itself to few compounds. Since the properties of each group sus¬ tain a slight change in virtue of the increasing atomic weight, we are not surprised to observe this to be very evident in the case of Hg (with the high atomic weight 200.4), especially as tbe middle (transition) member of the third great period is not known (p. 248). Mercury differs essentially from zinc and cadmium* in that, in addi¬ tion to the compounds of the form HgXj (mercuric compounds), it is also capable of yielding those of the form HgX (mercurous compounds), in which it seems to be monovalent. Here we meet an instance, frequently observed, in which one and the same metal (as with the most metalloids) is capable of forming compounds of two or more forms, which are to be referred to a different valence of the metal ; and it often happens that the derivatives of a metal, appearing in different forms or types, are frequently more essen¬ tially distinguished from one another than the compounds of differ¬ ent elements having the same type. Thus, the mercuric compounds (HgXa) are similar to those of zinc and cadmium, after the same I form, while the mercurous compounds (HgX) exhibit great resem¬ blance to the cuprous (CuX) and silver (AgX) compounds, consti¬ tuted according to a similar type. It shows that the similarity of the compounds is not only influ¬ enced by the nature of the metals, but frequently, to a marked * See p. 255. 326 INORGANIC CHEMISTRY. degree, by the forms or types according to which they are consti¬ tuted (p. 333). As viewed above, mercury in its ic compounds is a dyad, in the ous a monad. According to the theory of constant valence, the mercury atom in the ous com¬ pounds is, however, also divalent. We suppose that the molecules of the same are twice as large, and that in them two Hg atoms form a divalent group, as seen from the following :— Hg. Hg—CI Hg —NO3 Hg I >0 I I I )s Hg/ Hg —CI Hg—NO3 Hg/ Mercurous oxide. Chloride. Nitrate. Sulphide. This view is, however, not justified, inasmuch as the molecular weight of mer¬ curous chloride corresponds to the formula HgCl (p. 328). MERCURY. Hg = 200.4. Mercury {Hydrargyrum) occurs in nature principally as Cinnabar^ more rarely native in the form of little drops scattered through rocks. Its most important localities are Almadén in Spain, New Almadén in California, Idria in Illyria, Mexico, Peru, China, and Japan. The metallurgical separation of mercury is very simple. Cinnabar is roasted in reverberatory furnaces, whereby the sulphur burns to dioxide, and the mercury vapors are condensed in large chambers. Or, it is distilled with lime or iron from iron retorts. Commercial mercury usually contains a slight quantity of other metals dissolved in it. For its purification, it is poured in a thin stream into a deep layer of sulphuric or dilute nitric acid, by which the accompanying tin and lead are more easily dissolved than the mercury. The metal is finally distilled out of a small glass retort and pressed through chamois skin. This is the only metal which is liquid at ordinary temperatures. At 0° its specific gravity equals 13.59 ; it solidifies at —40°, and crystallizes in regular octahedra ; it evaporates somewhat at medium temperatures, and boils at 360°. Its vapors are very poisonous. The specific gravity of the vapor of mercury is 100.2 (H = i) or 6.91 (air = i). Therefore, the molecular weight of the metal is 200.4, as its atomic weight is also 200.4, the molecule, like that of cadmium, is composed of only one atom. At ordinary temperatures, mercury is not altered by exposure to the air; near the boiling point, however, it gradually oxidizes to red mercuric oxide. Hydrochloric and cold sulphuric acids do not act upon it ; hot sulphuric acid converts it into mercury sulphate, with MERCUROUS COMPOUNDS. 327 evolution of sulphur dioxide. Even dilute nitric acid will readily dissolve it. It combines with the halogens and sulphur at ordinary temperatures. Mercury dissolves almost all metals (not iron) forming amalgams. It unites with potassium and sodium upon gentle warming, with production of heat and light. When the quantity of potassium and sodium exceeds 3 per cent., the alloy is solid and crystalline ; by less amount it remains liquid. Tin amalgam is employed for coating mirrors. Mercury forms two series of compounds, mercurous and mercuric. The first are analogous to the cuprous, and have the form, HgX. In them mercury appears to be monovalent ; we, however, do not know whether their molecules are not to be expressed by the double formula HgaXa (p. 326). In many respects the ous compounds are similar to the cuprous and silver derivatives. The halogen com¬ pounds are insoluble, and darken on exposure to light. In the ic derivatives—HgXj—mercury is divalent, and is very much like zinc and cadmium. The ic compounds almost always form, if the substance reacting with the mercury is in excess ; when the opposite is the case, mercurous salts result. The ic deriv¬ atives, by the addition of mercury, pass into the ous^ Hg (N03)2 -Ç- Hg = Hg2(N03)2. Oxidizing agents convert the ous into the ic compounds; the latter are, on the other hand, con¬ verted by reducing substances into the first. The heat of formation of some of the mercuric compounds corresponds to the symbols :— (Hg,0) = 3o.6 (Hg,Cg=63.i (Hg.y = 34-3 (Hg,S) = 16.8. That of the corresponding mercurous salts :— (Hgj,0)=42.2 (Hg,Cl) = 4i.2 (Hg,I) = 24.2 (Hgj.S) = —. MERCUROUS COMPOUNDS. Mercurous Chloride—HgCl or Hg2Cl2—calomel, is an amorphous white precipitate, produced by the addition of hydro¬ chloric acid or soluble chlorides to the solution of mercurous salts. It is generally formed by the sublimation of HgCl2 with mercury ; or a mixture of HgSO^, mercury and sodium chloride is sublimed :— HgSO^ 4- 2NaCl + Hg = Na^SO^ -f Hg^Ch. It then forms a radiating, crystalline mass (quadratic prisms) of specific gravity 7.2. Calomel is insoluble in water, in alcohol, and 328 INORGANIC CHEMISTRY. dilute acids ; it gradually decomposes when exposed to the light, with separation of mercury. When heated, it sublimes without fusing. By the action of strong acids it is converted into mercuric salts and free mercury. When ammonium hydroxide is poured over calomel, it blackens (hence the name calomel, from xako¡í£Xa<¡)j and reacts according to the equation :— Hg^Cl^ + 2NH3 = NH^Cl + NH^Hg^Cl. The compound NH2Hg2Cl is viewed as ammonium chloride, in which 2 H are replaced by Hg2. The vapor density of calomel vapors at 440° is 117.6 (H = 1), the molecular weight, therefore, 235.2, and corresponds to the formula HgCl (235.2). It ap¬ pears, however, that its vapors consist of a mixture of mercury and mercuric chlor¬ ide. Such a mixture must have the same density as HgCl :— HgCl + HgCl = Hg -f HgCl2. 1 vol. I vol. I vol. I vol. The question, whether the mercurous compounds contain one or two atoms of mercury, whether, for example, the formula HgjCh or HgCl properly belongs to calomel, is, therefore, not decided by the direct determination of its vapor density. This difficulty has, however, been solved by our better knowledge of the phe¬ nomena of dissociation and especially of the tension of dissociation. Pressure obstructs the dissociation of a compound, even though it be heated in the vapor of one of the components, into which it might separate. The determination, therefore, of the vapor density of mercurous chloride in an atmosphere of mer¬ curic chloride has shown it to be 117.6, corresponding to the simple formula HgCl—under this condition a dissociation is impossible (Fileti). Mercurous Iodide—Hgl or Hg2l2—is prepared by rubbing together 8 parts of mercury with 5 parts I, or by precipitating mer¬ curous nitrate with potassium iodide. It is a greenish powder, in¬ soluble in water and alcohol. Light changes it to Hgl2 and Hg. Mercurous Oxide—Hg20—is black in color, and is formed by the action of potassium or sodium hydroxide upon mercurous salts. In the light or at 100°, it decomposes into HgO and Hg. Mercurous Nitrate—HgNOg or Hg2(N03)2—is produced by allowing somewhat dilute nitric acid to act upon excess of mercury in the cold. It crystallizes with i molecule ofHaO in large mono- clinic tables. It dissolves readily in water acidulated with nitric acid ; pure water decomposes it into the acid salt which passes into solution, and the basic salt—Hga^^^Q^, which separates as a yel¬ low powder. The nitric acid solution of mercurous nitrate oxidizes when ex¬ posed to the air, and gradually becomes mercuric nitrate; this may be MERCURIC COMPOUNDS. 329 prevented by adding metallic mercury to the solution, whereby the resultant ic salt is again changed to the ous state :— HgiNOj), + Hg = Hg,(N03),. Mercurous Sulphate—Hg2(S04)—results when an excess of mercury is heated gently with sulphuric acid ; it separates as a crystalline precipitate, difficultly soluble in water, if sulphuric acid be added toa mercurous nitrate solution. It fuses upon application of heat, and decomposes into SO2, O2, and Hg. Mercurous Sulphide—Hg2S—is precipitated by potassium hydrosulphide, as a black compound, from the dilute solution of mercurous nitrate. When gently warmed, it decomposes into HgS and mercury. MERCURIC COMPOUNDS. Mercuric Chloride—HgCb—Corrosive sublimate—is pro¬ duced when mercuric oxide is dissolved in HCl, or metallic mercury in aqua regia. It is obtained on a large scale by the sublimation of a mixture of mercuric sulphate with sodium chloride :— HgSO^ + 2NaCl = HgCla + Na^SO^. It crystallizes from water in fine rhombic prisms, and dissolves at medium temperatures in 15 parts, at 100°, in 2 parts water; it is still more soluble in alcohol. Its specific gravity is 5.4. It fuses at 265°, and boils about 293°. Its critical pressure is about 420 Mm. (p. 233). The vapor density is 135.6 (H = 1), corresponding to the molecular formula HgClj (= 271.3) ;— Hg + CI, = HgCl,. I vol. I vol. I vol. Reducing substances, like SO2 and SnCb, change it to insoluble mercurous chloride :— 2HgCl2 + SO, + 2H,0 = Hg^Ch + H^SO, + 2HCI. Stannous chloride first precipitates mercurous chloride : 2HgCl2 + SnClz = Hg2Cl2 -f- SnCl4, which is afterward reduced, by excess of the first, to metallic mercury ; HgjCIj + SnClj = 2Hg -f- SnCl4. Mercuric chloride is greatly inclined to form double salts with metallic chlorides, ^., HgClg.KCl -f- H2O. When ammonium hydroxide is added to its solution, a heavy white precipitate, called white precipitate, NHaHgCl, is thrown down. This compound is regarded as a derivative of ammonium chloride, in which two atoms of H are replaced by a divalent mercury atom, and it has been called Mercur-ammonium Chloride. It forms the compound 28 330 INORGANIC CHEMISTRY. NHjHgClNHiCl with ammonium chloride ; the structure of this is expressed by the formula :— H /NH3CI Similar mercur-ammonium derivatives are numerous. Mercuric Iodide—Hglg—is formed by the direct union of mer¬ cury with iodine. When potassium iodide is added to a solution of mercuric chloride, Hglj separates as a yellow precipitate, which im» mediately becomes red. Hgl2 is readily soluble in HgClj and KI solutions ; it crystallizes from alcohol in bright red quadratic octa- hedra. Upon warming Hglj to 150°, it suddenly becomes yellow, fuses and sublimes in yellow, shining, rhombic needles. On touch¬ ing these with some solid, they become red, with separation of heat, and are changed into an aggregate of quadratic octahedra. Mer¬ curic iodide is therefore dimorphous. Mercuric Oxide—HgO—is obtained by the prolonged heating of metallic mercury near the boiling point in the air, or by the ignition of mercurous or mercuric nitrate. It forms a red, crys¬ talline powder, of specific gravity 11.2. When sodium hydroxide is added to a solution of mercuric chloride, mercuric oxide separates as a yellow, amorphous precipitate. Both modifications become black when heated, but change to a yellowish-red on cooling. Mercuric oxide breaks up into mercury and oxygen about 400°. Mercuric oxide combines directly with ammonia, to form the compound aHgO.NHg, which explodes with violence when heated. Mercuric Nitrate—Hg(N03)2.—It is difficult to obtain this salt pure, because it is inclined to form basic compounds. A solu¬ tion of it may be made by dissolving mercury or mercuric oxide in an excess of hot nitric acid. On diluting the solution with watei the basic salt, Hg(N03)2.2Hg0 -|- HjO, separates, and this may be converted into pure mercuric oxide by boiling with water. Mercuric Sulphate—RgSO^—is produced by digesting mer¬ cury or its oxide with an excess of concentrated sulphuric acid. It forms a white, crystalline insoluble mass, which becomes yellow on heating. It yields the hydrate HgSO^ -f H2O with a little water, but much of the latter decomposes it into sulphuric acid and the yellow insoluble basic salt, IlgSO*. 2HgO ( Turpetum minérale, Tur- peth mineral). Mercuric sulphate forms double salts with the alkaline sulphates, e.g., HgSOi. K2SO4 + 6H2O ; these are isomorphous with the cor¬ responding double salts of the magnesium group (p. 318). Mercuric Sulphide—HgS—occurs in nature as cinnabar, in radiating crystalline masses, or in hexagonal prisms of red color. COPPER, SILVER AND GOLD. 331 It is obtained by rubbing together mercury and flowers of sulphur with water, or it is produced as a black amorphous mass by the pre¬ cipitation of a solution of a mercuric salt with hydrogen sulphide. If the black sulphide be heated with exclusion of air it sublimes as a dark red mass of radiating crystalline structure, and is perfectly similar to natural cinnabar. A similar conversion of the black modiflcation into the red is effected by continued heating of the same to 50® with a solution of potassium or ammonium sulphide. The red mercury sulphide thus obtained is employed as artificial cinnabar in painting. The mercury compounds can be readily recognized by the fol¬ lowing reactions. On fusion with dry sodium carbonate, mercury escapes, and (if the operation be executed in a small tube) con¬ denses upon the side in metallic drops. Tin, copper and zinc throw out metallic mercury from its solutions. If a pure piece of sheet copper be dipped into the same, mercury is deposited as a gray coating, which on being rubbed acquires a metallic lustre. The mercurous compounds are distinguished from the mercuric by their precipitation by hydrochloric acid. COPPER, SILVER, AND GOLD. Considering the magnitude of their atomic weights, copper, sil¬ ver, and gold, bear the same relation to the alkali group, especially to sodium, as zinc, cadmium, and mercury bear to magnesium :— Na == 23.06 Mg = 24.38 Cu = 63.3 Zn = 65.5 Ag = 107.938 Cd = 112.1 Au = 197.2 Hg = 200.4. They occupy an entirely analogous position in the three great periods of the periodic system of the elements (p. 249), and con¬ stitute the transition from the elements of group VIII, especially from nickel, palladium, and platinum, to the less basic elements of group II—zinc, cadmium, and mercury ;— Ni = 58.6 Cu = 63.3 Zu = 65.5 Pd = 106 Ag = 107.938 Cd z= 112.1 R = 194.8 Au = 197.2 Hg = 200.4. This intermediate position of the three elements about to be dis¬ cussed is clearly shown in their entire physical deportment. While the elements of group VIII, with the last members, Ni, Pd, and Pt, fuse with difficulty and do not volatilize, Cu, Ag, and Au, in pokit of 332 INORGANIC CHEMISTRY. fusion and volatility, constitute the transition to the readily fusible and volatile elements, Zn, Cd, and Hg. They take an interme¬ diate position, too, with reference to their coefficients of expansion, their atomic volumes, and other physical properties. It is note¬ worthy that the ability to conduct heat and electricity attains its maximum in Cu, Ag, and Au. Consult p. 255. Not only are the properties of the free elements determined by the position of the latter in the periodic system, but those of their derivatives, and especially such as depend upon the valence of the elements, are influenced to a marked degree by the above relation. In consequence of the double periodicity of the great periods, Cu, Ag, and Au attach themselves to group I, and especially to sodium, just as the elements immediately following, Zn, Cd, and Hg, arrange themselves with group II and magnesium. Hence we find Cu, Ag, I and Au, like Na, yielding compounds of the form MeX, in which they appear monovalent. Some of these are isomorphous; thus NaCl, CuCl, and AgCl crystallize in forms of the regular system. Silver sulphate, AgjSO^, is isomorphous with sodium sulphate, SO^Naa ; and the same is true of other salts of these two metals. Cu and Ag, like the alkalies, afford so-called sub- or quadrant oxides, Na^O, CU4O, AgiO. But we may say that the similarity of Cu, Ag, and Au to Na is confined to these few external properties. Just as the heavy metals, Zn, Cd, and Hg differ in many properties from the light metal magnesium (p. 314), so do the metals Cu, Ag, and Au, possessing a high specific gravity distinguish themselves in a still higher degree from the light metal sodium. They possess all the properties belonging to the heavy metals, which are mainly characterized by the insolubility of the oxides, sulphides, and many salts. This character which separates them from sodium is explained by the fact that they really belong to the three great periods, and are clas¬ sified with the alkali metals in but few properties. Gold, with the high atomic weight, 197.2, corresponds in this respect to mercury (p. 325), but is very variable. I In the compounds constituted according to the form MeX, in which Cu, Ag, and Au appear monovalent, they exhibit great simi¬ larity in respect to their physical and chemical properties. The chlor¬ ides, CuCl, AgCl, and AuCl, are colorless and insoluble in water ; soluble, however, in hydrochloric acid, ammonia, the alkaline hypo¬ sulphites, etc., and furnish perfectly similar double compounds. I While silver only enters compounds of the form AgX, copper and gold are capable of yielding another form ; copper forms, besides COPPER, SILYER, AND GOLD. 333 I II cuprous, CuX, also cupric, CuXa, derivatives, in which it appears to be divalent. The latter are much more stable than the former, and embrace the most usual copper salts. Gold, however, besides I _ III _ _ furnishing ous, AuX, compounds, has /V derivatives, AuXg, in which it appears trivalent. While Cu and Au, in their ous forms, are analogous to silver (and in less degree, Na), the cuprü derivatives show a great resemblance to the compounds of the metals of the magnesium group, and other metals in their divalent combinations. Thus, the sulphates of zinc, magnesium, cupric oxide (CuO), ferrous oxide (FeO), nickelous oxide (NiO), cobaltous oxide (CoO), and manganous oxide (MnO), are similarly constituted, resemble each other, are isomorphous, and form entirely analogous double salts (p. 318) with the alkaline sul- II phates. In the same way the carbonates (MeCOs), the chlorates II and bromates (MeClaOg + 6H2O) and others, are similarly consti¬ tuted and isomorphous. In its ic derivatives, gold exhibits some III similarity to the aluminium compounds (AIX3), to those of indium III (InXg) and other metals, in their trivalent combinations. Here we see, as already observed with mercury (p. 325), that ¿Ae similarity of the compounds of the metals is influenced by the similarity of forms or typesf according to which they are composed, i. e.y by the valence of the metals. If a metal form several series of compounds of dif¬ ferent types, each series is usually more or less similar to the com¬ pounds of other metals of like type. In this manner is shown the resemblance of the compounds of the following types :— Na^O AgjO QujO AugO TljO Sodium oxide. Silver oxide. Cuprous oxide. Aurous oxide. Thallous oxide. MgO ZnO CuO FeO HgO Magnesium oxide. Zinc oxide. Cupric oxide. Ferrous oxide. Mercuric oxide. AI2O3 Fe203 AU2O3 TI2O3 Aluminium oxide. Ferric oxide. Auric oxide. Thallic oxide. The character of their derivatives varying with the degree of combination or valence, becomes quite marked with chromium, manganese and iron^ as we shall later see. The heavy metals also exhibit a strong, positive basic character in their monovalent com¬ binations. Thus silver oxide (AgjO) and thallous oxide (TljO) are strong bases, forming neutral reacting salts with acids, and even cuprous and aurous oxides are more strongly basic than their higher forms of oxidation. The metalloidal character of the metals, and 334 INORGANIC CHEMISTRY. the acid nature of their oxides begin to appear in their trivalent com¬ binations. Thus in the hydroxy! derivatives of aluminium, indium, and gold, A1(0H)3, In(0H)3, Au(0H)3, hydrogen may be replaced by the alkalies just as in boric acid, B(OH)3. Their higher forms of oxidation show, like those of the metalloids, a pronounced acid-like character (as Pb02, PtOj, CrOs, PeOs) which is only lessened by a high atomic weight of the metal (as in PbOa and PtOj). The character of the compound is influenced in a less, if not an unimportant degree, by the position of the elements in the periodic system. Hence the properties of the metallic compounds are not only influenced by the nature of the metals, but to a high degree by the combination forms. These forms of the elements, and particularly those of the metals, are regulated, however, if not entirely, yet to a consider¬ able degree, by the periodic system, as previously observed (see p. 252). This connection of Cu, Ag, and Au in an analogous group, expresses itself, too, in the heat resulting from the formation of their compounds of the form MeX :— Consequently relations occur here perfectly similar to those of the elements of the zinc group (p. 323), and perfectly analogous conclusions may be deduced from them with respect to the affinity relations. Thus, for example, copper is able to decompose concentrated but not dilute hydrochloric acid. The heat of formation of some cupric compounds equals :— Native copper is found in large quantities in America, China, Japan, also in Sweden and in the Urals. It frequently occurs crys¬ tallized in cubes and octahedra. The most important and most widely distributed of its ores are : cuprite (CujO), malachite and azurite (basic carbonates), chalcocite (CuaS), and especially chal- copyrite or speckled copper ore (CuFeSj). Metallurgy of Copper.—The extraction of copper from its oxygen ores is very simple: metallic copper is melted out when the ores are ignited along with charcoal. The sulphur ores are more difficult to work. The divided material is first roasted in the air, by which means copper sulphide is partially converted into oxide. The mass is afterward ignited with sand, silica fluxes, and carbon, when iron sulphide is converted into oxide and passes into the slag. By several repe¬ titions of this process we get the so-called copper stone—a mixture of cupric sul- (Naj, S) 88.0. (CUj, S) 20.2. (Cu, O) = 37.1 (Cu, CI2) = 51.6 (Cu, CI2, Aq.) = 62.7 (Cu, S, OJ = 182. COPPER. Cu = 63.3. COPPER. 335 phide with oxide. This is repeatedly roasted and heated, and metallic copper obtained by the action of the cupric oxide upon the sulphide :— 2CuO + CuS = 3CU -f- SOj. The copper obtained in this way is fused again with charcoal, to free it from the oxide. To obtain chemically pure copper, the pure oxide is heated in a stream of hydrogen, or the solution of copper sulphate is decom¬ posed by electrolysis. Metallic copper possesses a characteristic red color, and transmits a green light in thin leaflets. It is rather soft and ductile, and possesses a specific gravity 8.9. It fuses about 1054°, and vapor¬ izes in the oxy-hydrogen flame. It remains unaltered in dry air ; in moist, it is gradually coated with a green layer of copper carbon¬ ate. When heated, it oxidizes to black cupric oxide. Copper is not changed by dilute hydrochloric or sulphuric acids ; if it be moistened with these, and exposed to the air, it absorbs oxygen, and gradually dissolves. It is similarly dissolved by ammonium hydroxide. Concentrated sulphuric acid converts it into copper sulphate, with evolution of sulphur dioxide. It dis¬ solves in dilute nitric acid in the cold, with evolution of nitric oxide. Zinc, iron and also phosphorus precipitate metallic copper from the aqueous solutions of its salts. Copper forms two series of compounds, known as cmçtous and cupr/V. In the compounds, copper is divalent : — CuO CuClj Cu(0H)2 SO4CU. These are more stable than the ous derivatives ; the ordinary copper salts belong to them. In many respects they resemble the com¬ pounds of other dyad metals, especially those of the magnesium group, and ous compounds of iron (FeO), manganese (MnO), cobalt and nickel (see p. 333). The cuprous compounds are, on the other hand, very unstable, absorb oxygen from the air, and pass into cupric derivatives. They show some similarity to the mercurous derivatives (p. 327), and possess an analogous composition : — CuQ Cul CujO CujS. Oxygen salts of cuprous oxide are not known. From the formulas given above, copper, like silver, is monovalent in its ous compounds. It is, however, questionable, whether these formulas express the real molecular values. It is ordinarily assumed that the cuprous derivatives, like those of mercury in its ous state (p. 328), correspond to the doubled formulas, and that 336 INORGANIC CHEMISTRY, the copper atom is divalent, and forms a divalent group composed of two copper atoms, as may be seen from the following formulas :— The vapor density of cuprous chloride corresponds to the formula CUjQ, (com¬ pare p. 342), and, therefore, rather favors the above opinion. Cuprous Oxide—CU2O—occurs as cuprite crystallized in regu¬ lar octahedra. It is obtained artificially by boiling a solution of copper sulphate and grape sugar with potassium hydroxide, when it separates as a crystalline, bright-red powder. It does not change in the air, and is readily soluble in ammonium hydroxide. The solution absorbs oxygen, and while forming cupric oxide acquires a blue color. By the action of sulphuric and other oxygen acids, it forms cupric salts, the half of the copper separating as metal :— The hydroxide, Cu2(OH)2, is precipitated by the alkalies as a yellow powder from hydrochloric acid solutions of CU2CI2. It oxidizes in the air to cupric hydroxide. Cuprous Chloride—CuCl or CU2CI2—^is produced by the com¬ bustion of metallic copper in chlorine gas (together with CuClj), upon conducting HCl over copper at a low, red heat, by boiling the solution of cupric chloride with copper (CUCI2 4" Cu = CU2CI2), and by the action of many reducing substances upon cupric chloride. It is most conveniently made by passing sulphur dioxide through a concentrated solution of copper sulphate and sodium chloride, when it separates as a white, shining powder, consisting of small tetrahedra. It fuses at 430°, and distils about 1000° ; its vapor density corresponds to the formula CujCla- In the air, it rapidly becomes green, owing to oxygen absorption, and the for- /C1 mation of basic cupric chloride, CIu^qjj. Cuprous chloride is readily soluble in concentrated hydrochloric acid and in ammonium hydroxide; both solutions possess the characteristic property of absorbing carbon monoxide. Cuprous Iodide—Cul or CU2I2—is precipitated from soluble cupric salts by potassium iodide :— CujO or I Cu^ Cuprous oxide. CUPROUS COMPOUNDS. CujO + SO^Hj = CuSO^ + Cu + HjO. CuSO^ 4- 2KI = Cul -f KjSO^ 4- I. CUPRIC COMPOUNDS. 337 By extracting the co-precipitated iodine by means of ether it is obtained as a gray powder, insoluble in acids. Cuprous Sulphide—CujS—occurs as chalcocite crystallized in rhombic forms. It is produced by burning copper in vapor of sulphur, and by heating cupric sulphide in a current of hydrogen ; after fusion it solidifies in crystals of the regular system. Com¬ bined with silver sulphide it constitutes the mineral stromeyerite, I s or CuaS.AgjS, isomorphous with chalcocite. Copper Hydride—CuH or CuaHj—belongs to the derivatives of monovalent copper. If a solution of copper sulphate be digested with hypophosphorous acid, the hydride separates as a yellow amorphous precipitate which soon acquires a brown color. At 60° it decomposes into copper and hydrogen. With hydrochloric acid it forms cuprous chloride ;— CuH -b HCl = CuCl + Hj. Copper suboxide, Cu^O, or quadrantoxide, corresponds to potassium suboxide. On adding an alkaline stannous chloride solution to one of copper sulphate there separates, at first, cupric hydroxide, which is further reduced to cuprous hydroxide, and then to suboxide. The latter is an olive-green powder, which oxidizes readily and is decomposed by HjSO^ into CuSO^ and 3CU. CUPRIC COMPOUNDS. The cupric salts, when hydrous, are generally colored blue or green ; they are colorless in the anhydrous condition. Cupric Hydroxide—Cu(OH)2—separates as a voluminous bluish precipitate when sodium or potassium hydroxide is added to soluble copper salts. When heated, even under water, it loses water, and is changed to black cupric oxide. Cupric Oxide—CuO—is usually obtained by the ignition of copper turnings in the air, or by heating cupric nitrate. It forms a black amorphous powder, which, at higher temperatures, settles together and acquires a metallic lustre. By heating with organic substances their carbon is converted into carbon dioxide, and the hydrogen into water, the cupric salt being reduced to metal ; upon this rests the application of cupric oxide in the analysis of such compounds. Copper oxide and hydroxide dissolve in ammonium hydroxide with dark blue color. The solution possesses the power of dissolv¬ ing wood fibre (cotton-wool, Ijnen, filter-paper, etc.)—Schweizer's reagent. Cupric Chloride—CuClj—is formed by the solution of cupric oxide or carbonate in hydrochloric acid. It crystallizes from 29 338 INORGANIC CHEMISTRY. aqueous solution, with 2 molecules of water, in bright green rhom¬ bic needles, and is readily soluble in water and alcohol. When heated, it parts with its water, becoming anhydrous chloride, which at a red heat is decomposed into chlorine and cuprous chloride. It yields beautifully crystallized double salts with potassium and ammonium chlorides. Cupric bromide is like the chloride; the iodide is not known, since in its formation it at once breaks up into cuprous iodide and iodine. Copper Sulphate—CUSO4 5H2O—cupric sulphate, copper vitriol—may be obtained by the solution of copper in concentrated sulphuric acid. It is produced on a large scale by roasting chalco- cite. It forms large blue crystals of the triclinic system, which effloresce somewhat upon exposure. At 100° the salt loses 4 mole¬ cules of water; the fifth separates above 200°. The anhydrous sulphate is colorless, absorbs water very energetically, and returns to the blue hydrous compound. Although copper sulphate only crystallizes with 5 molecules of H2O, it is capable, like the sulphates of the magnesium group, of forming double salts with potassium and ammonium sulphates, which crystallize with 6H2O, and are isomorphous with the double salts of the metals of the magnesium group. Copper sulphate is employed in electro-plating. When its solu¬ tion is decomposed by the galvanic current copper separates at the negative pole, and deposits in a regular layer upon the conducting objects connected with the electrode. Ammonium hydroxide added to a copper sulphate solution in suf¬ ficient quantity to dissolve the c-upric hydroxide produced at first, changes the color of the liquid to a dark blue. From this solution alcohol precipitates a dark-blue crystalline mass with the composi¬ tion CUSO4.4NH3 -|- HjO. Heated to 150° this compound loses water and 2 molecules of NH3, and becomes CUSO4.2NH3. It is supposed that these compounds are ammonium salts in which a part of the hydrogen is replaced by copper ; they have been designated cuprammonium compounds^ e. g. :— gO^/NH,\cu \NH,/ Cuprammonium sulphate. The other soluble copper salts afford similar compounds with ammonium hydroxide. Cupric Nitrate—Cu(N03)2—crystallizes with three or six mole¬ cules of water, has a dark-blue color and is readily soluble in water and alcohol. Heat converts it into cupric oxide. Copper Carbonates.—The neutral salt (CuCOj) is not known. COPPER ARSENITE—PHOSPHORUS BRONZE, 339 When sodium carbonate is added to a warm solution of a copper salt the basic carbonate, CuÇ03.Cu(0H)2, or Cu OH' separates as a green precipitate. It occurs in nature as malachite^ which is especially abundant in Siberia. Another basic salt— 2C03Cu.Cu(0H)2—is the beautiful blue azurite. Copper Arsenite—(As03)2Cu3—separates as a beautiful bright green precipitate, upon the addition of sodium arsenite to a copper solution. It was formerly employed as a pigment, under the name of Scheele's green, but at present, owing to its poisonous character, it has been replaced by other green colors (Guignet's green and aniline green), Cupric Sulphide—CuS—is a black compound, precipitated from copper solutions by hydrogen sulphide. It is insoluble in dilute acids. When moist, it slowly oxidizes in the air to cupric sulphate. Heated in a stream of hydrogen, it forms cuprous sulphide, CuaS. Alloys of Copper.—Pure copper is very ductile, and may be readily rolled, and drawn out into a fine wire. It cannot be well poured into moulds, because it contracts unequally upon cooling and does not fill out the moulds. For such purposes, alloys of copper are employed, which, in addition, possess other technically valuable properties. The most important copper alloys are ;— Brass, consisting of two or three parts copper and one part zinc. It has a yellow color, and is considerably harder than pure copper. Ordinarily, one to two per cent, of lead are added to the brass, which facilitates its working upon the turning-lathe. Tombac con¬ tains 15 per cent, zinc, and has a gold-like color. The alloy of i part zinc and 5.5 parts copper answers for the manufacture of spurious gold leaf. The alloys of copper with tin are called bronzes. Most of the modern bronzes also contain zinc and lead; those from Japan, gold and silver. The cannon bronze contains 90 per cent, copper and 10 per cent, tin ; bell metal has 20-25 P®"" cent, of tin. Argentan is an alloy of copper, zinc, and nickel (see latter). The German copper coins consist of 95 per cent. Cu, 4 per cent. Sn, and I per cent. Zn. Of the more recently introduced alloys of copper we may men¬ tion :— Phosphorus Bronze. This consists of 90 parts copper, 9 parts tin and 0.5-0.8 parts phosphorus. By the latter ingredient the bronze is increased in hardness, and its solidity and resistance to oxidation are also increased. It is employed in making machin¬ ery. 340 INORGANIC CHEMISTRY. Silicon bronze, containing silicon instead of pliosphorus, is char¬ acterized by great firmness and conductivity. It is used for tele¬ phone wires. Manganese bronze contains 70 per cent, copper and 30 per cent, manganese (it is the cupro-manganese of Létrange), and may be melted with copper and copper alloys, imparting to these solidity and great hardness. See p. 351 for aluminium bronze. Recognition of Copper Compounds.—Most copper com¬ pounds containing water have a blue or green color. With the excep¬ tion of copper sulphide they all dissolve in ammonium hydroxide, with a blue color. When a piece of pure iron is introduced into a copper solution, it becomes covered with a red layer of metallic copper. Volatile copper compounds tinge the ñame blue or green. The spectrum of such a ñame is characterized by several blue and green lines. SILVER. Ag = 107.938. Silver occurs native. Its most important ores are AgjS and various compounds with sulphur, arsenic, antimony, copper and other metals. ^ Of rarer occurrence are combinations with chlor¬ ine (hornsilver, AgCl), bromine and iodine. Slight quantities of silver sulphide are present in almost every galenite (PbS). The principal localities for silver ores are America (Chili, Mexico, Cali¬ fornia), Saxony (Freiberg), Hungary, the Altai and Nertschinsk. Metallurgy of Silver.—The separation of the metal from its ores is rather complicated and variously effected; its elaborate description belongs to the pro¬ vince of metallurgy. At present, the ores containing silver and copper are, in Saxony and the Hartz, roasted in a divided state and fused with slags rich in silicic acid. In this way, as with copper, there is obtained a copper stone consist¬ ing of iron, copper and silver sulphides. This is then oxidized in a furnace; from the resulting mixture of ferric and cupric oxides and silver sulphate (SO^Agj), the latter is extracted by water. The silver is precipitated from this aqueous solu¬ tion by copper. Formerly, in Saxony, the separation of the silver was executed according to the so-called amalgamation process. According to this the mixture of sulphides is roasted with sodium chloride, whereby silver chloride is produced. The divided material is then mixed with iron scraps and water in rotating vessels. The iron causes the precipitation of the metallic silver from its chloride :— 2AgCl -)- Fe = FeClj -j- 2Ag. To free the metal from various impurities it is dissolved in mercury and the liquid amalgam ignited; mercury distils off and silver remains. Owing to scar¬ city of combustible material, the conversion of silver ores into silver chloride is executed, in Mexico and Peru, by mixing the ores with sodium chloride and copper sulphate in the presence of water. In this way cuprous chloride is pro¬ duced, which is transposed, with silver sulphide, into silver chloride and cuprous sulphide :— SILVER. 341 2CuCl -|- AgjS = CU2S -j- 2AgCl. To get silver from galenite, proceed as follows : First, metallic lead is obtained. In this way all the silver in the ore passes into the lead and may be obtained with profit from the latter, even if it does not constitute more than per cent, of it. To this end, the metallic lead is fused and allowed to cool slowly ; pure lead first crystallizes out, which can be removed by sieves, while a readily fusible alloy of lead with more silver remains behind. This method of Pattison's is repeated until the residual liquid lead contains I per cent, of silver. The lead, rich in silver, is subjected to cupellation ; it is fused in a reverberatory furnace with air access. The bottom of the oven consists of some porous substance. In this pro¬ cess the lead is changed to readily fusible oxide, which partly flows out of side openings from the hearth, or is, in part, absorbed by the porous bed ; the unoxi- dized silver remains in the cupel in met^lic condition. The ordinarily occurring silver (work silver) is not pure, but invariably contains copper and traces of other metals in greater or less quantity. To prepare chemically pure metal, the work silver is dissolved in nitric acid, and from the solution of nitrates thus obtained, hydrochloric acid precipitates the silver as chloride :— AgNOs + HCl == AgCl + HNO 3. The latter is reduced by various methods ; either by fusion with sodium carbonate, or by the action of zinc or iron in the presence of water :— 2AgCl Zn = ZnClj 2Ag. Silver is a pure white, brilliant metal, of specific gravity 10.5. It is tolerably soft and very ductile, and can be drawn out to a fine wire. It crystallizes in regular octahedra. It fuses at 954°, and is converted into a greenish vapor in the oxy-hydrogen flame. Silver is not oxidized by oxygen ; by the action of ozone it is cov¬ ered with a very thin layer of silver peroxide. When in molten condition, silver absorbs 22 volumes of oxygen without combining chemically with it ; the absorbed gas escapes again when the metal cools. Silver is capable of existing in three allotropie forms, which have properties greatly different from those of ordinary silver. The first form is soluble in water and has a blue color. The second variety is insoluble, and somewhat resembles the first form. The third closely resembles gold in color and lustre. These allotropie varieties of silver are broadly distinguished from normal silver by color. They very likely are more active conditions of silver, common silver being a polymerized variety (Am./r. Science^ 37, 476). Silver unites directly with the halogens ; by the action of hydro¬ chloric acid it becomes coated with an insoluble layer of silver chloride. Boiled with strong sulphuric acid, it dissolves to sul¬ phate ï—*~ 2Ag + 2H,S04 = AgjSO^ -f SO, + 2H,0. 342 INORGANIC CHEMISTRY. The best solvent of silver is nitric acid, which even in a dilute state and unaided by heat, converts it into nitrate. As silver is rather soft, it is usually employed in the arts alloyed with copper, whereby it acquires a greater hardness. Most silver coins consist of 90 percent, silver and 10 per cent, copper; the English shillings contain 92.5 per cent, silver. Oxygen forms three compounds with silver, but only the oxide affords corresponding salts. Silver Oxide—AgzO—is thrown out of a silver nitrate solution by sodium or potassium hydroxide as a dark-brown, amorphous precipi¬ tate. It is somewhat soluble in water, and blues red litmus paper. In this, and in the neutral reaction of the nitrate, the strong, basic, alkaline nature of silver and its oxide exhibits itself ; the soluble salts of nearly all of the other heavy metals show an acid reaction. When heated to 250°, the oxide decomposes into metal and oxygen ; at 100° it is reduced by hydrogen. The hydrated oxide is not known ; the moist oxide reacts, however, very much like the hy¬ droxides. On dissolving precipitated silver oxide in ammonium hydroxide, black crystals (Ag20.2NH3) separate when the solution evaporates, and when dry these explode upon the slightest disturbance. (Ful¬ minating silver.) Silver Suboxide—Ag^O—corresponding to potassium suboxide, is produced by heating silver citrate in a current of hydrogen, and is a black, very unstable powder, which decomposes readily into silver oxide and silver, v. d. Pfordten has recently shown that this product is really a silver hydroxide and assigns the formula Ag4^Qjj to it. Silver Peroxide—AgO or Ag^Oj—is formed by passing ozone over silver or its oxide, or by the decomposition of the nitrate by the electric current. It con¬ sists of black, shining octahedra, and at 100° decomposes into AgjO and oxygen. The salt-like compounds of silver correspond to the oxide AgjO, and are all constituted according to the form AgX, hence are termed argentic. They are analogous to the cuprous and mercurous derivatives, and show a great resem- blance to the former in physical and chemical qualities. It would, therefore, be more correct to designate them argento/«. Compounds of the divalent form AgX,, are not known for silver. If, however, the mercurous and cuprous compounds are expressed by double formulas (p. 336) :— CaCl Cu\ HgCl IIg\ 1 I O and I I O CuCl Cu/ Hga Hg/ which view is supported by their chemical deportment, those of silver might be répresented by analogous formulas :— AgCl AgNO, IgCl Ag/ -igNO,. SILVER CHLORIDE. 343 Tht n the silver atom would be divalent and a complete parallelism would be established with copper. The chemical formulas of solid bodies do not generally designate their true molecular values as in the case of gases, but only their sim¬ plest atomic composition. It is very probable that even the simplest chemical com¬ pounds, i.g.f KCl and AgCl, consist in their solid condition (as crystal molecules, p. 107) of complex molecules corresponding to the formulas (KC1)b, (AgCl)ni. An argument supporting this view is afforded by the existence of different modi¬ fications of chloride and bromide of silver ; these differ from each other in their external properties, and in their different susceptibilities to light. The doubling of formulas, as shown above with CojClj, HgjCIj etc., is mainly due to the ten¬ dency to deduce ail the compounds of an element from a constant value, accord¬ ing to the doctrine of constant valence. This is, however, impossible (p. 175). According to present notions of valence, and as it is presented in the periodic system, compounds (MeCl, MeClj, MeClg, etc.) are constituted according to defi¬ nite forms or types that may materially determine their properties (p. 333). So far as the similarity of metallic compounds is concerned, it is of secondary im¬ portance whether the quantities corresponding to the simple formulas, in the solid or gaseous state, do unite to larger, complex molecules (compare HgCl, CU2CIJ — BCI3, A1(CH3)3 and AljClg, GaCl3 and GagClg—SnClj, SujCl^, PbClj, etc.). In case of the sesquioxides MjOj it is also immaterial whether they are derived from supposed trivalent elements (as AljOj, GujOg, lujOg), or from those that are tetra- valent (as Fe^Og, CrjO,, MnjOg). TTie same may be remarked of the metallic compounds McgO^ = (MeOO)2.Me (see Spinels). The use of simple or of double formulas for" the metallic compounds is there¬ fore of no special importance. Silver Chloride—AgCl—exists in nature as hornsilver. When hydrochloric acid is added to solutions of silver salts, a white, curdy precipitate separates; the same fuses at 451° to a yellow liquid, which solidifies to a horn-like mass. The chloride is insol¬ uble in dilute acids ; it dissolves somewhat in concentrated hydro¬ chloric acid and in sodium chloride, readily in ammonium hydrox¬ ide, potassium cyanide, and sodium hyposulphite. It crystallizes from ammoniacal solutions in large, regular octahedra. Dry silver chloride absorbs 10 per cent, of ammonia gas, forming a white compound—2AgC1.3NH3 — with it, which at 38° gives up its ammonia. Silver Bromide—^AgBr—is precipitated from silver salts by hydrobromic acid or soluble bromides. It has a bright yellow color, and dissolves with more difficulty than the chloride in am¬ monium hydroxide ; in other respects it is perfectly similar to the latter. Heated in chlorine gas it is converted into chloride. Silver Iodide—Agi—is distinguished from the chloride and bromide by its yellow color, and its insolubility in ammonia. Fused silver iodide at first solidifies in isometric crystals, which gradually change to hexagonal forms, but when the latter are heated to 146°, they suddenly revert to the isometric forms. It dissolves readily in hydriodic acid, to Agi.HI, which, upon evaporation of the solution, separates in shining scales. Heated in chlorine or 344 INORGANIC CHEMISTRY. bromine gas, it is converted into chloride or bromide ; conversely, chloride and bromide of silver are converted into silver iodide by the action of hydriodic acid. These opposite reactions are explained by the principle of the greatest evolu¬ tion of heat. Chlorine and bromine expel iodine from all iodides because the heat of formation of the latter is less than that of the bromides and chlorides (p. 271). Again, hydriodic acid (gaseous or in aqueous solution) converts silver chloride into the iodide according to the equation :— AgCl + HI = Agi -f HCl, because the heat modulus of the reaction is positive (for gaseous HI and HQ -f- 12.5 C.,for the solution 10.6 C. See the Table at close of book). Sunlight, and also other chemically active rays (magnesium light, phosphorus light) color silver chloride, bromide, and iodide, at first violet, then dark black, whereby they are probably converted into compounds of the form AgjX. In such an altered condition they are capable of fixing finely divided silver ; on this depends their application in photography. In photographic work a nep^ative is first prepared. A glass plate is covered with collodion (a solution of pyroxylin in an ethereal solution of alcohol) holding in solution halogen salts of calcium or cadmium. After the evaporation of the ether the glass plate is covered by a dry collodion layer containing the haloid salts. The plate is now immersed in a solution of silver nitrate, whereby haloid salts of silver are precipitated upon the surface. The plate thus prepared is ex¬ posed to light in the camera obscura, and, after the action, dipped into a solution of pyrogallic acid or ferrous sulphate. These reducing substances separate metallic silver in a finely divided state, which is precipitated upon the places where the light has acted. The plate is now introduced into a solution of piotassium cyanide, which dissolves the silver salts not affected by the light, while the metal¬ lic unaltered silver remains. The negative thus formed is covered at the places upon which the light shone, by a dark layer of silver, while the places corre¬ sponding to shadows of the received image are transparent. The copying of the glass negative on paper is executed in a similar manner. Silver Cyanide—AgCN—is precipitated from silver solutions by potassium cyanide or aqueous hydrocyanic acid, as a white, curdy mass, not affected by light. It dissolves readily in ammonium hydroxide and potassium cyanide, forming with the latter the crystal¬ line compound AgCN.KCN. The solution in potassium cyanide is employed in the electro silver-plating of metals. Silver Nitrate—AgNOs—is obtained by dissolving pure silver in somewhat dilute nitric acid, and crystallizes from its aqueous solution in large rhombic tables, isomorphous with potassium salt¬ petre. At ordinary temperatures it is soluble in one-half part water or in four parts alcohol, the solution having neutral reaction. In this respect it differs from the salts of almost all metals, which SILVER SULPHIDE. 345 react acid (p. 333). It fuses at 218°, and solidifies to a crystalline mass. When perfectly pure it is not affected by light, but it usually turns black in sunlight with separation of metallic silver. Organic substances also reduce it to metal. Silver nitrate is employed in the cauterization of wounds (Lunar caustic). By dissolving work silver in nitric acid a mixture of silver and copper nitrates is obtained. To separate the silver salt from such a mixture it is heated to redness, the copper thus converted into oxide and the unaltered silver nitrate extracted with water. Silver Nitrite—AgN02—is precipitated from concentrated silver nitrate solutions by potassium nitrite. It crystallizes in needles, dissolves with difficulty in water, and decomposes above 90°. Silver Sulphate—^Ag^SO^—^is obtained by the solution of silver in hot sul¬ phuric acid, and crystallizes in small rhombic prisms •which are difficultly soluble in water. It is isomorphous with anhydrous sodium sulphate. Silver Sulphite—AgjSOj—is precipitated as a white, curdy mass, if sulphurous acid be added to the solution of the nitrate. It blackens in the light and decom¬ poses at 100°. Silver Sulphide—Agg S—occurs in regular octahedra, as argen¬ tita. Hydrogen sulphide precipitates it as a black amorphous sul¬ phide from silver solutions. By careful ignition in the air it is ox¬ idized to silver sulphate. It is insoluble in water and ammonium hydroxide and dissolves with difficulty in nitric acid. Silvering.—When silver contains more than 15 per cent, copper it has a yel¬ lowish color. To impart a pure white color to objects made of such silver they are heated to redness with access of air. The copper is thus superficially oxidized, and may be removed by dilute sulphuric acid. The surface of pure silver is then polished. The silvering of metals and alloys (German silver, argentan) is executed in a dry or wet way. In the first, the objects to be silvered'are coated with liquid silver amalgam, with a brush, and then heated in an oven; the mercury is volatilized, and the silver surface then polished. At present, the galvanic process has almost completely superseded the other processes. It depends on the electrolysis of the solution of the double cyanide of silver and potassium, whereby the silver is thrown out upon the electro-negative pole and deposits upon the metallic surface in connection with that electrode. To silver glass, cover it with a mixture of an ammoniacal silver solution, with reducing organic substances like aldehyde, lactic, and tartaric acids. Under defi¬ nite conditions, the reduced silver deposits upon the glass as a regular metallic mirror. Recognition of Silver Compounds.—Hydrochloric acid throws down a white, curdy precipitate of silver chloride, which dissolves readily in ammonium hydroxide. Zinc, iron, copper, and mercury throw out metallic silver from solutions of silver salts, and from insoluble compounds, like the chloride. 346 INORGANIC CHEMISTRY. GOLD. Au = *97.2 Gold {aurutn) usually occurs in the native state, and is found dis¬ seminated in veins in some of the oldest rocks. Gold sands are formed by the breaking and disintegration of these. It is found, in slight quantity, in the sand of almost every river. Combined with tellurium it forms sylvanite, found in Transylvania and Cali¬ fornia. It is present in minute quantity in the most varieties of pyrites and in many lead ores. For the separation of the gold grains the sand or pulverized rocks are washed with running water, which removes the lighter particles and leaves the specifically heav¬ ier gold. Native gold almost invariably contains silver, copper, and various other metallic admixtures. To remove these, the gold is boiled with nitric or concentrated sulphuric acid. The removal of the silver by the latter acid is only complete if that metal predominates ; in the reverse case a portion of it will remain with the gold. There¬ fore, to separate pure gold from alloys poor in silver they must first be fused with about three-fourths their weight of the latter metal. Gold may be separated from copper and lead by cupellation (P- 340- Pure gold is rather soft (almost like lead) and has a sjiecific gravity 19.32. It is the most ductile of all metals, and may be drawn out into extremely fine wire and beaten into thin leaves, which transmit green light. About 1035° it melts to a greenish liquid. It is not altered by oxygen, even upon ignition ; acids do not attack it. It is only in a mixture of nitric and hydrochloric acids (aqua regia), which yields free chlorine, that it dissolves to gold chloride, AuCb- Free chlorine produces the same. Most metals, and many reduc¬ ing agents (ferrous sulphate, oxalic acid) precipitate gold from its solution as a dark-brown powder. As gold is very soft it wears away rapidly, and is, therefore, in its practical applications, usually alloyed with silver or copper, which have greater hardness. The alloys with copper have a red¬ dish color, those with silver are paler than pure gold. The Ger¬ man, French, and American gold coins contain 90 per cent, gold and 10 per cent, copper. A 14-karat gold is generally employed for ornamental objects ; this contains about 58.3 per cent, pure gold (24 karats representing pure gold). Gold, according to its atomic weight, belongs to the group of copper and silver ; and, upon the other hand, forms the transition from platinum to mercury. Its character is determined to a high degree by these double relations (p. 331). Like the other elements GOLD. 347 of high atomic weight, mercury, thallium, lead, and bismuth, belonging to the same series of the periodic system, it varies con¬ siderably in character from its lower analogues. Gold, like silver and copper, yields compounds of the form AuX —aurons, analogous to the cuprous and argentous. Besides, it has those of the form AuXg, auric derivatives, in which it is trivalent. These show the typical character of the trivalent combination form, which expresses itself in the acidity of the hydroxides (p. 333); auric hydroxide, Au(0H)3, unites almost solely with bases. On the other hand, they show many similarities to the highest com¬ bination forms of the metals with high atomic weight ; platinum (PtX^), mercury (HgXj), thallium (TIX3), and lead (PbX^) (p. 360). AUROUS COMPOUNDS. Aurous Chloride—AuCl—is produced by heating auric chlo¬ ride, AUCI3, to 180°, and forms a white powder insoluble in water. When ignited, it decomposes into gold and chlorine ; boiled with water it decomposes into the trichloride and gold. Aurous Iodide—Aul—separates as a yellow powder, if potas¬ sium iodide be added to a solution of auric chloride :— AUCI3 -{- 3KI = Aul + Ij + 3KCI. When heated it breaks up into gold and iodine. When auric oxide or sulphide is dissolved in potassium cyanide, large colorless prisms of the double cyanide, AuCN.KCN, crys¬ tallize out upon evaporation. The galvanic current and many metals precipitate gold from this compound ; hence it serves for electrolytic gilding, which, at present, has almost entirely superseded the gilding in the dry way (see p. 345). Aurous Oxide—Au,O—is formed by the action of potassium hydroxide upon aurous chloride. It is a dark violet powder which at 250° decomposes into gold and oxygen. It is changed to AuClj and gold by the action of hydrochloric acid. Only a few double salts of the oxygen derivatives of monovalent gold are known. AURIC COMPOUNDS. Auric Chloride—AuCl,—results by the solution of gold in aqua regia, and by the action of chlorine upon the metal. When the solution is evaporated the chloride is obtained as a reddish- brown, crystalline mass, which rapidly deliquesces in the air It dissolves readily in alcohol and ether. 348 INORGANIC CHEMISTRY. Gold chloride forms beautifully crystallized double salts with many metallic chlorides, e.g., AuCh-KCl 2)4^20 and AuClg.- NH4CI + H2O. When auric chloride is heated with magnesium oxide a brown precipitate is obtained, from which all the magnesia is removed by concentrated nitric acid, leaving Auric Oxide (AujOg). This is a brown powder which decomposes, near 250°^ into gold and oxygen. If the precipitate containing the magnesia be treated, not with concentrated, but with dilute nitric acid^ Auric Hydroxide—Au(OH)3—remains as a yellowish-red powder. Both the oxide and hydroxide are insoluble in water and acids ; they possess, however, acid properties, and dissolve in alkalies. Therefore the hydroxide is also called aun'e acid. Its salts, the 1 aurates, are constituted according to the formula MeAuOj, and are derived from the meta-acid, HAuOg =HO.AuO. Potassium Anrate—KAUO2-I-3H2O—crystallizes in bright yellow needles, from a potassium hydroxide solution of auric oxide. These are readily soluble in water ; the solution reacts alkaline. The cor¬ responding aurates are precipitated from this solution by many metallic salts, e.g. :— KAuOj + AgNOg = AgAuOj + KNOg. The precipitate produced by magnesia in a solution of auric chloride (see above) consists of magnesium aurate (Au02)2Mg. Oxygen salts of auric oxide are not known. Auric Sulphide—AugSg—is precipitated as a blackish-brown compound, from gold solutions, by hydrogen sulphide. It dissolves in alkaline sulphides with formation of sulpho-salts. Stannous chloride (SnClg) added to an auric chloride solution produces, under certain conditions, a purple-brown precipitate, purple of Cassius, which is employed in glass and porcelain painting. Alumina and magnesia yield similar purples, and it appears that their red coloration is due to finely divided metallic gold. On pouring ammonium hydroxide over auric oxide a brown com¬ pound is ^roá.\icQá—fulminating gold. When this is dried and heated or struck a blow, it explodes very violently. METALS OF GROUP III. The triatomic elements, affording derivatives mainly of the form MeXg, belong to group III of the periodic system (p. 248) :— Sc == 44.1 Y = 88.7 La = 138.5 Yb == 173.2 ß = 10.9 Al = 27.0 Ga = 69.9 In = 113.7 Ti = 204.1 METALS OF GROUP III. 349 These bear the same relations to each other as do the elements of group II (p. 303). Boron has the lowest atomic weight, and the basic, metallic character in it is reduced very much or does not appear at all. In its exclusively acidic hydroxide, B(0H)3, it approaches the metalloids, and is therefore treated with them (p* 243)' _ _ _ . , Aluminium is a perfect metal; its hydroxide, A1(0H)3, exhibits a predominating basic character, and yields salts with acids. Its relations to boron are like those of silicon to carbon, or of magne¬ sium to beryllium. The connection of aluminium and boron with the same group plainly shows itself in the entire character of the free elements, and in their compounds. Thus aluminium and boron are not dissolved by nitric acid, but by boiling alkalies :— Al + 3KOH = A1(0K)3 + 3H. There is only a gradual difference between their hydrates. Boron hydroxide, B(0H)3, not only acts as a feeble acid, but we also find that aluminium hydroxide manifests an acidic character, inasmuch as it is capable (p. 333) of forming metallic salts with strong bases (chiefly the alkalies) ; but owing to the higher atomic weight of aluminium the basic character exceeds the acidic. The similarity is also shown by the existence of perfectly analogous compounds ; thus, e. g.f the chlorides BCI3 and AICI3 can unite with PCI5 and POCI3. Scandium, yttrium, lanthanum and ytterbium attach themselves to aluminium as the first sub-group. These constitute the third members of the great periods, and hence exhibit a pronounced basic character. As light metals, they are very similar to alumin¬ ium in their compounds, so that they all are embraced in one group, which (corresponding to the earthy nature of their oxides) is designated the Group of Earth Metals. Cerium and didymium bear ä peculiar relation to lanthanum; their atomic weights are nearly alike and their properties very similar. Their apparently abnormal existence is explained by the fact that the 5th period (series 7 and 8), which is very incomplete, shows a somewhat vary¬ ing function in its intermediate members (p. 250). The metals, erbium, terbium, thulium, and samarium, of recent discovery and but little characterized, may probably also be included in the same period. "WíQ second sub-group is more distinctly characterized and accu¬ rately investigated ; it consists of the heavy metals, gallium, indium and thallium. These belong to the right side of the great periods, possess, therefore, a less basic character, and bear the same relation to each other as Zn, Cd and Hg.* * * Consult p. 255 upon the behavior of the oxides of this group when heated with metallic magnesium. 350 INORGANIC CHEMISTRY. Aluminium was formerly classed together with chromium, iron, manganese, cobalt and nickel, in one group, because they all afford sesquioxides, M^Oj, whose salts are very much alike. Another fact which was thought to give weight to this classiñcation was the existence of the similarly constituted alums:— (SOJaAlgSO^Kj -f 24Hj,0 (S04)3Fe2S04K2 + 24H,0. Potassium aluminium sulphate. Potassium iron alum. In its entire behavior, aluminium is, however, very essentially distinguished from the other metals here mentioned—by the acid nature of its hydrate, Al(OH)3 —^and by its inability to form higher or lower combination forms, while the others yield basic monoxides, MeO, and acid-forming trioxides (CrOj, FeOg, MnO,). Here, again, the similarity of the sesquioxide compounds, Me^Oj, like those of the monoxide derivatives, is to be regarded as mainly influenced by the similarity of the combination forms (p. 333). At present aluminium is assumed to be trivalent and this fact apparently con¬ tradicts the circumstance that not the simple formulas, AICI3, AlBrj, but the double ones, Al^Clg, AljBrg, fall to its halogen derivatives (the result of vapor density determinations). On the other hand, however, the so-called metallo- organic compounds of aluminium exist, whose molecules are constituted accord¬ ing to the formulas, Al(CH3)3,Al(C2Hg)3; these undoubtedly prove the trivalence of aluminium, because the compounds with hydrocarbon groups (like those with hydrogen) afford the surest guide for the deduction of the valence (p. 251), The existence of the molecules, AljClg, AljBrg, etc., does not prove anything against its being a triad, but must be explained by a polymerization of the simple chemical molecules, AICI3, AlBrj. At higher temperatures the vapor density of aluminium chloride corresponds to the formula AICI3. We find the same to be the case with arsenious oxide, As^O,, and antimony trioxide, Sb203, whose molecules in vapor form correspond to the doubled formulas, As^Og (= As203,As203) and SbgO,; and with stannous chloride, whose molecule in vapor form at low temperature is SugCl^, but higher it becomes SnCl,, or with gallium chloride that possesses the formulas, GaClg and Ga^Clg (p. 343). GROUP OF THE EARTH METALS. ALUMINIUM. Al = s/.oS. This is one of the most widely distributed elements. As oxide, it crystallizes as ruby, sapphire and cprundum ; less pure as emery. It is commonly found as aluminium silicate (clay, kaolin), and in combination with other silicates, as feldspar, mica, and also in most crystalline rocks. It occurs, too, united with fluorine and sodium, as cryolite, in large deposits, in Iceland. Metallic aluminium is obtained by igniting the chloride, or, better, the double chloride of sodium and aluminium with metallic sodium :— AlClj.NaCl -1- 3Na = Al -f- 4NaCI. The metal may also be prepared by electrical processes. In the Hall method cryolite, or an electrolyte of analogous composition, is electrolyzed, the bath being constantly kept saturated with alumina. The aluminium in the end separates at the negative pole. In the Cowles process the underlying principle is the inter¬ ruption of a powerful electric current, the formation of an immense ALUMINIUM. 351 arc, and the reduction in the arc of the aluminium oxide, by car¬ bon, in the presence of the metal. It is a sil ver-white metal of strong lustre, is very ductile, and may be drawn out into fine wire and beaten into thin leafiets. Its specific gravity is 2.583; it belongs, consequently, to the light metals and possesses, therefore, all the properties opposed to those of the heavy metals (see p. 314). It fuses at a red heat but will not vaporize. It changes very little in the air at ordinary tem¬ peratures, and even when heated. If, however, thin leaves be heated in a stream of oxygen, they will burn with a bright light. Nitric acid does not affect aluminium ; sulphuric acid only dissolves it on boiling, while it is readily soluble, even in the cold, in hydro¬ chloric acid. It dissolves in potassium and sodium hydroxide, with evolution of H, and forms aluminates :— Al + 3KOH = K3AIO3 + 3H. Owing to its stability in air and beautiful lustre, aluminium is sometimes employed for vessels and ornaments. The alloy of cop¬ per with 10-12 per cent, aluminium is distinguished by its great hardness and durability. It may be poured into moulds, and pos¬ sesses a gold-like color and lustre. Under the name of aluminium bronze, it is used for the composition of various articles, as watches, spoons, etc. Its firmness and elasticity render it suitable for physical instru¬ ments (arms of balances) and watch springs. Aluminium affords compounds of the form AIX3, or AlgX# (p. 350) exclusively. Its salts, soluble in water, have an acid reaction, and a sweet, astringent taste. The heat of formation of some of the aluminium compounds equals :— (Al^.CIg) = 321.8. (Al2,Bre) = 239.3. (Algjle) = »40.6. (Al2»Cl3,Aq.) = 475*5* (Al2,Bfg,Aq.) = 4^9-9* (Al2,l0,Aq.) = (Al3,03,3H20) = 388.8. The heat evolved in the formation of a quantity of aluminium hydroxide, cor¬ responding to one atom of oxygen, is 129.6; since that of water is far less (Hg.O = 69.0),it must be decomposed by aluminium, with liberation of hydro¬ gen (p. 279). If this does not transpire under ordinary conditions, the reason must be sought for in the insolubility of aluminium hydroxide. Indeed, the reaction occurs if aluminium chloride, or another salt, in which the aluminium oxide is soluble, be added to the water. Conversely, the high heat of formation of aluminium oxide explains why it is not reduced by carbon. Aluminium Chloride, AICI3, or AlaClg, is produced by the action of chlorine upon heated aluminium : also by heating a mix¬ ture of aluminium oxide and carbon in a current of chlorine :— AI2O3 + 3C -h 6C1 = AI2CI3 -f 3CO. 352 INORGANIC CHEMISTRY. Chlorine and carbon do not act separately upon the oxide ; by their mutual action, however, the reaction occurs in consequence of the affinity of carbon for oxygen, and of chlorine for aluminium. The oxides of boron and silicon show a similar deportment. Aluminium chloride may be obtained in white, hexagonal leaflets by sublimation. It sublimes readily, but will only fuse when sub¬ jected to high pressure. Deville and Troost (185 7) first determined the vapor density of aluminium chloride, bromide, and iodide at 440°, and found it to correspond to the formulas AbCle, AbBrg, and Alale- This led to the supposition that aluminium was quadrivalent in all its compounds. The most recent investigations of Nilson and Petterson have shown that at higher temperatures (above 700°) the vapor density of the chloride corresponds to the formula AICÍ3. Aluminium chloride absorbs moisture from the air, and deli¬ quesces. It crystallizes from a concentrated hydrochloric acid solution, with 6 molecules of water. On evaporating the aqueous solution, the chloride decomposes into aluminium oxide and hydro¬ gen chloride :— Alacie + 3H2O = AI2O3 + ÓHCl. It forms double chlorides with many metallic chlorides, viz. : AlClg.NaCl, AICI3.KCI. The solutions of these maybe evaporated to dryness without decomposition. It also unites with many halo¬ gen derivatives of the metalloids AICI3.PCI5, AlClj.POClj, AICI3.SCI4. Aluminium Bromide—Al2Brg-^is obtained like the chloride, and consists of shining leaflets which fuse at 90° and boil at 265-270°. Its vapor density is 267.4 (H = i), corresponding to the formula AlgBrg. It behaves like the chloride. Aluminium Iodide—AI2I6—is formed on heating aluminium filings with iodine. It is a white, crystalline mass, fusing at 185°, and boiling about 400°. It is best prepared by covering sheet aluminium with carbon disulphide, and then adding the calculated amount of iodine gradually, letting the whole stand for some time, and then distilling off" the CSg. The reaction occurring between aluminium iodide and oxygen is interesting. If the vapor of the former be mixed with the latter, and then brought in contact with a flame, or if acted upon by an electric spark, a violent detonation will ensue ; aluminium oxide and iodine T6Slllt Î Alale 3^ = AI2O3 61. This deportment is due to the great difference in the heats of formation of the aluminium oxide (about 380 C.), and the iodide (140.6 C.). The chloride and bromide are similarly decomposed, but with less violence. Aluminium Pluoride—AIFI3 or Al2Flg—obtained by conducting hydrogen fluoride over heated aluminium oxide or hydroxide, sublimes at a red heat in colorless rhombohedra. It is insoluble in water, unaltered by acids, and is very stable. It yields insoluble double fluorides with alkaline fluorides. The com¬ pound—AIFI3. 3NaFl—occurs in Greenland, in large deposits, as cryolite, and is employed in the soda manufacture (p. 295). Aluminium Oxide—AI2O3—is found crystallized in hexagonal prisms in nature, as ruby, sapphire, and corundum, colored by other admixtures. Impure corundum, containing aluminium and ALUMINIUM HYDROXIDES. 353 iron oxides, is called emery, and serves for polishing glass. The specific gravity of these minerals is 3.9; their hardness is only a little below that of the diamond. Artificial aluminium oxide may be obtained by igniting the hydroxide, and is a white amorphous powder, which fuses to a transparent glass in the oxy-hydrogen flame. A mixture of aluminium fluoride, and boron trioxide, heated to a white heat has the boron fluoride volatilized, and crys¬ tallized aluminium oxide remains :— AljFlj B2O3 = AljOj 2BFI3. The crystallized or strongly ignited aluminium oxide is almost insoluble in acids; to decompose it, it is fused with caustic alkalies or with primary potassium sulphate—HKSO4. Aluminium Hydroxides.—The normal hydroxide, A1(0H)3 or Al2(OH)6, occurs in nature as hydrargillite. The hydroxide, Al202(0H)2, is diaspore. Bauxite is a mixture of the hydroxide, Al20(0H)4, with ferric oxide. The normal hydroxide is artificially obtained as a white voluminous precipitate, by adding ammonium hydroxide or an alkaline carbonate (in latter case carbon dioxide escapes, p. 355) to a soluble aluminium salt. Freshly precipitated, it dissolves in acids and in potassium and sodium hydroxides. By long standing under water, or after drying, it is, without any alter¬ ation in composition, difficultly soluble in acids. When carefully heated, the normal hydroxide first passes into AlO.OH. The freshly precipitated hydroxide dissolves readily in a solution of aluminium chloride or acetate. On dialyzing (p. 240) this solution the aluminium salt or crystalloid diffuses, and in the dialyzer remains the pure aqueous solution of the hydroxide. This has a faint alkaline reaction and is coagulated by slight quantities of acid, alkalies and many salts; the soluble hydrate passes into the insoluble gelatinous modification. Gelatinous aluminium hydroxide possesses the property of precipitating many dyestuffs from their solutions, forming colored insoluble compounds (lakes) with them. On this is based the application of aluminium hydroxide as a mordant in dyeing. The acetate is generally used for this purpose. Goods saturated with this salt are heated with steam, which causes the decomposition of the wèak acetate ; acetic acid escapes, whilfe the separated aluminium hydroxide sets itself upon the fibre of the material. If the latter now be introduced into the solution of coloring matter the latter is fixed by the aluminium hydroxide upon the fibre. At present, sodium aluminate is employed instead of the acetate. Aluminium hydroxide has a feeble acid character, and can form salt-like compounds with strong bases. On carefully evaporating its solution in sodium or potassium hydroxidê, or upon addition of alcohol, white amorphous compounds of KA10„ NaA102 and (NaO)3Al are obtained. The potassium compound can be obtained in crystalline form. These derivatives, known as alumínales, are not very stable, and are even decomposed by carbon dioxide, with elimination of aluminium hydroxide ;— 2A102Na + COj + 3H2O = Al2(OH)3 + COjNa,. 30 354 INORGANIC CHEMISTRY. The aluminium hydroxide obtained in this manner, in distinction from that precipitated from acid aluminium solutions by the alka¬ lies, is not gelatinous, and is more difficultly soluble in acids, espe¬ cially acetic. It comprises the ordinary alumina of commerce. On adding calcium chloride, strontium chloride, or barium chloride to the solution of potassium or sodium aluminate, white insoluble aluminates are precipitated :— 2A102Na + CaClj = (A102)2Ca + 2NaCl. Similar aluminates frequently occur as crystallized minerals, in nature. Thus the spinels consist chiefly of magnesium aluminate, AlO O/^^' chrysoberyl is beryllium aluminate, aIO gahnite is zinc aluminate, ^jo Nearly all these minerals, commonly called spinels, crystallize in regular octahedrons, like the corresponding chromium com¬ pounds (see these) ; the exceptions are chrysoberyl, crystallizing in the rhombic system, and hausmannite, MnjOi, in the quadratic system. Technically, alumináis obtained from cryolite, bauxite and other minerals con¬ taining aluminium. The pulverized bauxite is heated with dry sodium carbonate in furnaces, and the resulting sodium aluminate extracted with water. From the clear solution carbon dioxide precipitates the hydroxide, while sodium carbonate remains dissolved, and is afterward recovered. The dried aluminium hydroxide occurs as a white powder in trade. The gelatinous, readily soluble (colloidal) aluminium hydroidde (see above) precipitated from acid solutions by alkalies, has lately been prepared upon a large scale, according to the method of Löwig, by treating the sodium aluminate solu¬ tion with milk of lime ; calcium aluminate precipitates, while sodium hydroxide remains in solution :— zAlO^Na + Ca(OH)2 = (A102)2Ca + zNaOH. The calcium aluminate is dissolved in hydrochloric acid ;— (A102)2Ca + 8HC1 = 2AICI3 + CaQ, + 4H2O, and to the solution now containing the alumina as chloride the corresponding amount of calcium aluminate added, and aluminium hydroxide is precipitated :— 2AICI3 -f- 3(A102)2Ca = 4AI2O3 + 3CaCl2. According to this procedure, the sodium hydroxide formed in the first reaction is obtained together with the alumina. On conducting carbon dioxide into a solu¬ tion of alkaline carbonates^ and adding a solution of an alkaline aluminate at the same time, white aluminium-alkali carbonates are precipitated :— AI2O8.K2O + 2C0gNaH = Al2O3.K2O.2CO2 -I- 2NaOH. The caustic alkali that is formed in this way is converted ^[ain into bicarbonate by carbon dioxide. In a dry state the precipitates are white, chalk-like masses ALUMINIUM SULPHATE. 355 which at 90® contain 5 molecules of water: Al2O3.K2O.2CO2 + S^jO. Their constitution may be expressed by the formula :— They dissolve readily in dilute acids, even acetic, with evolution of carbon diox¬ ide, and are suitable for the preparation of pure alumina mordants and antiseptic solutions (Löwig). The basic character of aluminium hydroxide exceeds the acid ; but it is so feeble that it is not capable of forming salts with weak acids, as carbon dioxide, sulphurous acid, and hydrogen sulphide. When sodium carbonate is added to solutions of aluminium salts, aluminium hydroxide is precipitated, while carbon dioxide is set free :— AI3CI3 + 3Na2C03 + 3H2O = Al2(OH)e -(- 6NaCl -f3C02. The alkaline sulphides behave similarly :— Allele + 3(NH4)2S + 6H2O = Al2(OH)3 + ÓNH^Cl + 3H2S. Aluminium Sulphate—Al2(S04)3—crystallizes from aqueous solution with 18 molecules of HjO in thin leaflets with pearly lustre. These dissolve readily in water ; when heated, they melt and lose all their water of crystallization. The sulphate is obtained by dis¬ solving the hydroxide in sulphuric acid, or by the decomposition of pure clay with the same acid ; the residual silicic acid is removed by ñltration, and the solution of the sulphate evaporated. When a quantity of ammonium hydroxide, insufficient for complete precipi¬ tation, is added to the sulphate, basic sulphates separate out. Salts similar to the latter are also found in nature ; thus, aluminite, used to prepare alum, has the composition :— Aluminium sulphate can combine with the alkaline sulphates and affords double salts, termed alumsy e. g.y potassium alum :— (S0J3A12.S04K2 -f 24H20or (S04)2A1K -f- 12H2O. Their constitution is expressed by the following formula : O ^^2 \Sof + 7H3O or (A10.0)2S02 + 9H2O. -f- 24H2O or In this compound the potassium may be replaced by sodium, am- 356 INORGANIC CHEMISTRY. monium, rubidium, caesium, and also by thallium. Iron, chromium and manganese afford like derivatives :— Fe2(SO^)3.K2SO^ -f- 24H2O Mn2(SO^)3. Na2(S0)^ 24H2O. Potassium iron alum. Sodium manganese alum. All these alums crystallize in regular octahedra or cubes, and can form isomorphous mixtures. The most important of them is Potassium Alutninium Sulphate ox ordinary alum, A1K(S04)2+ 12H2O. It crystallizes from water in large, transparent octahedra, soluble in 8 parts water of ordinary temperature, or in ^ part boiling water. The solution has an acid reaction and a sweetish, astringent taste. When placed over sul¬ phuric acid, alum loses 9 (or 18) molecules of HgO. When heated it melts in its water of crystallization, loses all the latter and becomes a white, voluminous mass—burnt alum. Upon adding a little sodium or potassium carbonate to a hot alum solution the hydroxide first produced dissolves, and when the liquid cools, the alum crystallizes out in cubes, as cubical alum. The addition of more sodium carbonate causes the precipitation of the basic salt— A1K(S04)2.A1(0H)3. Alunite, found in large quantities near Rome and Hungary, has a similar composition (S04)2(A10)3K -j- 3H2O. Commercial alum is obtained according to various methods : I. From alunite, by heating and extracting with hot water. In this way alum dissolves while the hydroxide remains; from such solutions the former crystallizes in combinations of the octahedron with cube faces—Roman alum. 2. The most common source of alum was formerly alum skate, a clay containing pyrite and peat. This is roasted and after moistening with water is exposed for a long time to the action of the air. By this means FeSj is converted into FeSO^ and free sulphuric acid, which, acting upon the clay, forms aluminium sulphate. The mass is extracted with water, potassium sulphate added, and the whole permitted to crystallize. 3. At present clay is treated directly with sulphuric acid, and to the solution of aluminium sulphate potassium or ammonium sulphate is added. 4. Bauxite and cryolite are admirable material for the preparation of alum. The working of cry¬ olite for alumina and soda is described ®n p. 295, and that of bauxite, p. 354. Ammonium Alum—(S04)2A1NH4 12H2O—cr5rstallizes, like potassium alum, in large crystals, and at present, owing to its cheapness, is applied almost exclusively for technical purposes. Sodium alum is much more soluble, and crystallizes with difficulty. As the alum employed in dyeing must contain no iron, we under¬ stand why this salt is not applicable. At present the alum is being more and more supplanted by aluminium sulphate and sodium aluminate in all practical operations, because these chemicals can be procured perfectly free from iron. ALUMINIUM. 357 Aluminium Phosphate—AIPO4 + 4H2O—is thrown out of aluminium salt solutions by sodium phosphate, as a white gelatinous precipitate ; this is readily soluble in acids, acetic excepted. Aluminium Silicates.—The most important of the aluminium double silicates, so widely distributed in nature, are : leucite^ (Si03)2AlK, albite or soda feldspar, SigOgAlNa, ordinary feldspar— orthoclase—AlKSigOg—and the various micas, which, with quartz, compose granite. When these disintegrate under the influence of water and the carbon dioxide of the air, alkaline silicates are dissolved and carried away by water, while the insoluble aluminium silicate, clay y remains. Perfectly pure clay is white, and is called kaolin^ or porcelain clay ; its composition mostly corresponds to the formula, Al2(Si03)3Al205H4, or SÍ2O9AI2H4. When clay is mixed with water a tough kneadable mass is obtained. By drying and burning, it be¬ comes compact and hard, and is the more ñre-proof, the purer the clay. On this depends the use of clay for the manufacture of earthenware, from the red brick to porcelain. To produce porcelain a very fine mixture of kaolin, feldspar and quartz is employed. On strong ignition, the feldspar fuses, fills the pores of the clay and thus furnishes a fused transparent mass—porcelain. When it is not so strongly ignited, it remains porous—faience—serving for finer clay vessels. To render these impervious to water, they are covered with glazing. This consists of various readily fusible silicates. Rough earthenware vessels are constructed from impure clay, and they are usually glazed by throwing salt into the ovens at the time of burning. The hot steam decomposes the salt into hydrochloric acid and sodium hydroxide, which forms an easily fusible silicate on the surface of the clay. Ultramarine.—The rare mineral Lapis lazuli, which was formerly employed as a very valuable blue color under the name of Ultra¬ marine, is a compound of aluminium sodium silicate with sodium polysulphides. At present ultramarine is prepared artiñcially, in large quantities, by heating a mixture of clay, dry soda (or sodium sulphate), sulphur and wood ashes, away from air. Green ultra¬ marine is the product. This is then washed with water, dried, mixed with powdered sulphur and gently heated with air contact until the desired blue color has appeared—blue ultramarine. The cause of the blue coloration is generally assumed to be due to the existence of a complicated sulphur compound, whose nature is not yet explained. On pouring hydrochloric acid over the blue pro¬ duct, the color disappears with liberation of sulphur and hydrogen sulphide—this would point to the existence of a polysulphide. Violet and red ultramarines are prepared at present by conducting dry hydrogen chloride gas and air over common ultramarine at 100-150°. 358 INORGANIC CHEMISTRY. RARE METALS. In some very rare minerals, like cerite, gadolinite, euxenite, and orthite, occur¬ ring principally in Sweden and Greenland, is found a series of metals which, in their entire deportment, closely resemble aluminium (p. 349). These are scan¬ dium, yttrium, cerium, lanthanum, didymium, ytterbium, and the more recent erbium, terbium, thulium, samarium, and holmium* (Sorel's X element). These generally form difficultly soluble oxalates, and are, therefore, precipitated from solution by oxalic acid. They also afford difficultly soluble sulphates and double sulphates, of which the potassium double salts are constituted, according to the formula, ^62(804)3.3X280^. The different decomposabiliiy of their nitrates upon application of heat affords an excellent means for their isolation and separation. Scandium, yttrium, lanthanum, cerium, and 3^terbium have been most accurately investigated. Their atomic weights are very approximately correct The most interesting of the group is scandium, atomic weight 44. It fills out the gap between calcium and titanium. It coincides in all its properties with those deduced theoretically from the periodic system by Mendelejeff for the.element ekaboron (compare Gallium). Other elements have been observed whose salt solutions possess the remarkable property of yielding absorption spectra. They are didymium, erbium, samarium (decipium), holmium, and thulium. Nilson, Kriiss, and other investigators regard them as complex mixtures, consisting of several elements that have not yet been isolated. Scandium—8c =: 44—contained in euxenite and gadolinite, has not yet been obtained in a free condition. Its oxide, 8C2O3, is obtained by igniting the hydroxide or nitrate, and is a white,infusible powder (like magnesia and oxide of beryl¬ lium). Its specific gravity equals 3.86; the specific heat 0.1530. hydroxide, 8c(OH)3, is precipitated as a gelatinous mass from its salts by the alkalies, and is insoluble in an excess of the latter. The nitrate crystallizes in little prisms, and is decomposed with difficultyby heat. The potassium double sulphate, 862(804)3. 3K2SO4, is soluble in warm water, but not in a solution of potassium sulphate. The chloride affords a characteristic spark spectrum. Yttrium—¥=89—has long been known in its compounds, but has never been investigated in a pure condition. Its chief source is gadolinite (upward of 35 per cent.). Its potassium double sulphate is soluble in a potassium sulphate solu¬ tion, and in this manner it can be readily separated from cerium, lanthanum, and didymium. Its nitrate is much more difficult to decompose than those of scandium and ytterbium. The chloride, YCI3 -j- 7H2O, forms large prisms, and gives a spark spectrum. Lanthanum—La= 138.2—separated from its chloride by electrolysis, resem¬ bles iron as regards color and lustre, oxidizes in the air, and burns in a flame with a bright light. Its specific gravity equals 6.16, the specific heat 0.0448. The hydroxide, La(0H)3, is precipitated as a gelatinous mass, and reacts alkaline. Its chloride shows a distinct spark spectrum. The metal forms a hydride (p. 255)- Cerium—Ce = 140—occurs in cerite (60 per cent.), and is also obtained by the electrolysis of the chloride. It is very similar to lanthanum, but at ordinary tem¬ peratures is more stable than the latter ; burns much more readily, so that broken- off particles of it inflame of their own accord. The specific gravity of the fused metal is 6.72, the specific heat 0.0448. Besides the salts of the sesquioxide, €6303, it forms some from the dioxide, CeOj. The first are colorless, while the latter are colored yellow or brown ; red eerie hydroxide, Ce(OH)4, is precipitated from the first, on the addition of hypochlorites. A little aqueous hydrofluoric acid will convert the eerie hydroxide into cerium tetraßuoride. These compounds THE GALLIUM GROUP. 359 indicate that cerium is tetravalent and that it probably belongs to the fourth group of the periodic system (p. 249). Compare also p. 255. Ytterbium—Yb=i73.—Its oxide^ Yb, O3, is obtained from the so-called erbium earth (from euxenite and gadolinite) by repeated partial heating of the mixed nitrates, whereby the scandium nitrate is the first to decompose. It is a white infusible powder, of specific gravity 9.17 ; its specific heat is 0,0646. The salts of ytterbium are colorless, and show no absorption spectrum. Terbium—Tr=l50.— Terbium oxide, Tr^Og, occurs in large amount in samarskite. It has an orange-yellow color, resembles the oxide of erbium, but does not show an absorption spectrum. THE GALLIUM GROUP. The three heavy metals, gallium, indium, and thallium, bear the same relations to aluminium that we see exhibited by Cu, Ag, and Au to sodium, Na, and Zn, Cd, and Hg to magnesium. Cu 63.3 Zn 65.5 Ga 69.9 Ge 72.3 As 75.0 Ag 107.938 Cd II2.I In 113.7 Sn 118.1 Sb 120.3 Au 197.2 Hg 200.4 TI 204.1 Pb 206.91 Bi 208. They constitute the corresponding members of the three ^eat periods; and as second sub-group attach themselves to aluminium, while cerite metals form the first, more basic group (p. 349). The entire character of the three elements under consideration is influ¬ enced by this position in the periodic system, because regular rela¬ tions appear in all directions, as may be observed, for example, in the specific gravities, fusing points, and other physical properties in the free metals :— Ga In T1 Atomic weight Specific gravity Fusing point 69.9 S-9 30® II3-7 7-4 176° 204-1 II.8 290° Being members of group III of the periodic system, Ga, In, and T1 yield compounds of the trivalent form, and these are analogous to those of aluminium in many respects. Consult p. 255. Thallium, like other elements with high atomic weight (Au, Hg, Pb), exhibits great variations from the group properties (p. 325). It yields, for example, not only derivatives of the form TiXg, but also those of TlX. If we include thallium as a member of the last great period (Pt, Au, Hg, TI, Pb, Bi), we will discover that, as in case of the other metals of this series, a remarkable regularity 360 INORGANIC CHExMISTRYi underlies all its forms of combinations—the highest as well as the lowest î""" PtCl, AuCl HgCl TlCl PbCl, BiQ, PtCl^ AuClg HgClj TlClg PbCl^ BiXg, I. GALLIUM. Ga = 69.9. Gallium was discovered in zinc blende from Pierrefitte, in 1875, by Lecoq de Boisbaudran, by means of the spectroscope. As early as the year 1870, Mende- lejeiF, taking the table of the periodic system devised by him as basis, predicted the existence of a metal (standing between aluminium and indium, with an atomic weight of nearly 69), which he named Eka aluminium. Its properties were nec¬ essarily deduced from its position in the periodic system. All the properties of gallium known at that time agreed with those of eka-aluminium, and it seemed very probable that this element, which had been theoretically established, was in reality gallium. This is now confirmed by the fact that the atomic weight, deter¬ mined by experiment, agrees with that deduced theoretically. As yet gallium has only been found in very small quantity, and is but imper¬ fectly investigated. It is characterized by a spectrum consisting of two violet lines. Separated by electrolysis from an ammoniacal solution of its sulphate,it is a white, hard metal, of specific gravity 5.9, with a fusing point 30°. It is only superficially oxidized in the air, not altered by water, and is not volatile up to a red heat. Like aluminium, it is scarcely attacked by nitric acid, but dissolves readily in hydro¬ chloric acid. Gallium Oxide—GagOg—is obtained by igniting the nitrate. It is a white mass which sublimes when heated in a current of hydrogen. The hydroxide— Ga(0H)3—is thrown out of solutions of its salts by the alkalies as a white flocculent precipitate, readily soluble in an excess of the precipitant, but rather difficultly soluble in ammonium hydroxide. Gallium Chlorid&—GaClg—is produced on heating gallium in a current of chlorine gas ; it forms colorless crystals that fuse at 75°, sublime about 60® and boil at 215-220®. Its vapor density at 440® corresponds to the formula GaClg, at 270° very closely to GagClg. The chloride fumes in the air^ like aluminium chloride, deliquesces and decomposes in the evaporation of its aqueous solution. Gallium Nitrate—Ga(N03)3, and Gallium Sulphate, Ga2(S04)3—are crystalline and very deliquescent. The latter forms a double salt with ammonium sulphate—similar to the alums;— (SOJgGag.SO^CNHJg -f 24H2O. Hydrogen sulphide only precipitates gallium from acetate solutions. a. INDIUM. In = 113.7. Owing to its resemblance to zinc, indium was regarded as a divalent metal, and its compounds composed according to the formula, InX g ; this fixed its atomic weight at 75.6. The specific heat, however, made the atomic weight one and a half times as large (p. 262). Hence it is trivalent and its derivatives are con¬ stituted according to the form, InXg. It belongs to the group of aluminium, and, in its derivatives, manifests some similarity to this metal. It was discovered, in 1863, by Reich and Richter, by the aid of sroctnim analy¬ sis. Its spectrum is characterized by a very bright indigo-blue Une, hence its THALLIUM. 361 name. It only occurs in very minute quantities in some zinc blendes from Freiberg and the Hartz. It is a silver-white, soft and tenacious metal, of specific gravity, 7.42. It fuses at 176° and distils at a white heat. At ordinary temperatures it is not altered in the air; heated, it bums with a blue flame to indium oxide. It is difficultly soluble in hydrochloric and sulphuric acids, but dissolves readily in nitric acid. Indium Chloride—InClg—results from the action of chlorine on metallic indium, or upon an ignited mixture of indium oxide, and carbon. It sublimes in white, shining leaflets, which deliquesce in the air. Its vapor density corre¬ sponds to the formula InClg. It does not decompose when its aqueous solution is evaporated. Indium Oxide—In203^is a yellow powder resulting from the ignition of the hydroxide. Indium Hydroxide—In(0H)3^is precipitated as a gelatinous mass, by alkalies, from indium solutions. It is soluble in sodium and potassium hy¬ droxides. Indium Nitrate—In(N03)3—crystallizes with three molecules of water, in white deliquescent needles. Indium Sulphate—In^(804^)3—remains on evaporating a solution of indium in sulphuric acid as a gelatinous mass, with three molecules of water. It forms an alum with ammonium sulphate. Indium Sulphide—In2S3—is precipitated by hydrogen sulphide as a yellow- colored compound from indium solutions. Indium also forms a dichloride and monochloride. Indium Dichloride—InClj—is produced when metallic indium is heated in a current of hydrogen chloride. It is a white crystalline mass, which on exposure to a more intense heat becomes a yellow liquid and sublimes. Its vapor density at I000®-I040° corresponds to the formula InClj. Water decomposes it at once into indium trichloride and metallic indium : 3lnCl2 = 2lnCl3 -f- In. Indium Monochloride—InCl—like gallium dichloride (p. 360), results when the dichloride is heated with metallic indium. It is a crystalline, reddish yellow mass, which has a reddish black color when fused. Its vapor density at 1100°- 1400° corresponds to the formula InCl. Water decomposes it into the trichloride and metallic indium : 3lnCl = InCIg -|- 2ln. 3. THALLIUM. T1 == 204.1. Thallium is rather widely distributed in nature, but in very small quantity. The very rare mineral crookesite contains 17 per cent, of the metal, together with copper, selenium and silver. It is often found with potassium in sylvite and carnallite, in mineral springs, and in some varieties of pyrite and zincblendes. When these pyrites are roasted for the production of sulphuric acid, according to the chamber process, the thallium deposits as soot in the chim¬ ney and in the chamber sludge, and was discovered in the latter, in 1863, almost simultaneously, by Crookesand Lamy, by means of the spectroscope. To get the thallium, the chimney-dust is boiled with water or sulphuric acid, and thallous chloride precipitated from the solution 31 362 INORGANIC CHEMISTRY. by hydrochloric acid. The chloride is then converted into sul¬ phate, and the metal separated from the latter by means of zinc or the electric current. Thallium is a white metal, as soft as sodium, and has the specific gravity 11.8. It fuses at 290°, and distils at a white heat. It oxidizes very rapidly in moist air. It does not decompose water at ordinary temperatures. It is, therefore, best preserved under water in a closed vessel. By air access it gradually dissolves in the water, forming thallium hydroxide and carbonate. Heated in the air it burns with a beautiful green ñame whose spec¬ trum shows a very intense green line, hence the name thallium, from líaAiló?, green. Thallium dissolves readily in sulphuric and nitric acids, but is only slightly attacked by hydrochloric acid, owing to the insolubility of thallous chloride. , Thallium forms two series of compounds : thallous—TlX and III thallic—TlXg. The first are very similar to the compounds of the alkalies (and also to those of silver). The solubility of the hydroxide and carbonate in water shows this ; their solutions have an alkaline reaction. Again, many thallous salts are isomorphous with those of potassium, and afford similar double salts. In the insolubility of its sulphur and halogen compounds, monovalent thallium approaches silver and lead. In its compounds of the form TIX3 thallium is trivalent, like aluminium, but otherwise shows scarcely any similarity to the latter. The heat of formation of some of the thallous compounds is :— (Tl2,0)= 42.2 (T1,C1) =48.5 (Tl,Br) = 41.2 (Tl.I) = 30 (TIAH) = 56.9 (Th,S,OJ = 210.9 (T1,N,03) = 58.1- The heat of solution of these compounds is negative. From the numbers cited above we can understand the deportment of thallium toward water and the acids. The heat of formation of the ic compounds in aqueous solution equals (Tl,Cl3,Aq.) = 89.0 (Tl,Br3,Aq.) = 56.1 (Tl.Ig.Aq.) = 10.5 (TljjOgjßHjO) = 86.9. THALLOUS COMPOUNDS. Thallous Oxide—TbO—is formed by the oxidation of thallium in the air, or by heating the hydroxide to 100°. It is a black pow¬ der which dissolves in water with formation of the hydroxide. Thallous Hydroxide—Tl(OH)—may be prepared by decom¬ posing thallium sulphate with an equivalent amount of barium hydroxide, and crystallizes with one molecule of water in yellowish prisms. It dissolves readily in water and alcohol, affording strong alkaline solutions. THALLIC COMPOUNDS. 363 Thallous Chloride—TlCl—is thrown down from solutions of thallous salts by hydrochloric acid as a white, curdy precipitate, which is very difficultly soluble in water. It separates in small crystals from the hot solution. It fuses at 427°, and boils about 715°. Like potassium chloride, it affords an insoluble salt with platinic chloride—PtCl4.2TlCl. Thallous bromide forms a white, and thallous iodide a yellow precipitate. Thallous Sulphate—TI2SO4—crystallizes in rhombic prisms, isomorphous with potassium sulphate. It dissolves in 20 parts of water at ordinary temperatures. It affords double salts with the sulphates of the metals of the magnesium group, of ferrous oxide, of cupric oxide, etc. (p. 318), e. g., MgSOi.TlaSO^ + 6H2O ; these are perfectly similar and analogous to the corresponding double salts of potassium and ammonium. It affords thallium alum with the sulphates of the sesquioxides of the iron group, e. g., A1T1(S04)2 + 12H2O ; these are similar to potassium alum—A1K(S04)2 + 12H2O. Thallous Carbonate—TI2CO3— is obtained from the oxide by the absorption of CO2; it crystallizes in needles, which dissolve at ordinary temperatures in 20 parts of water. The solution has an alkaline reaction. Thallous Sulphide—TI2S—is precipitated from thallous salts by hydrogen sulphide as a black compound, insoluble in water. THALLIC COMPOUNDS. Thallic Chloride—TICI3—is produced by the action of chlorine upon T1 or TlCl in water, and is very soluble in water. It decom¬ poses at 100° into TlCl and CI2. The alkalies precipitate from its solutions thallic hydroxide, TIO.OH, a brown powder, which, at 100°, passes into thallic oxide, TI2O3. Further heating decom¬ poses the latter into thallous oxide and oxygen. The oxide and hydroxide are soluble in hydrochloric, nitric, and sulphuric acids, forming T1(N03)3, TÍ2(S04)3, TICI3. On conducting chlorine through a solution of thallic hydroxide in potassium hydroxide, it assumes an intense violet color, due probably to the formation of the potassium salt of thallic acid, the composition of which is yet unknown. The thallium compounds are poisonous. They are employed in making thallium glass, which refracts light more strongly than lead glass. The spectrum of the thallium flame shows a very bright green line. 364 INORGANIC CHEMISTRY. METALS OF THE FOURTH GROUP. The elements of group IV in the periodic system (p. 249), Ti = 48.1 Zr — 90.7 (Ce = 140.2) Th = 232.4 C = i2,Si = 28.4 Ge= 72.3 Sn = n8.i Pb = 2^.91, show the same analogies that were observed with the members of group III (p. 348). Their character is, however, more non-metallic ; their derivatives are chiefly of the types MeXi and MeOj, of which the latter are acid (p. 267). The first two elements, carbon and silicon, with low atomic weights, belong to the two short periods and are true non-metals. Their oxides and hydroxides are acid in nature. The first more basic sub-group comprises titanium, zirconium (cerium), and thorium. They constitute the fourth members of the large periods. Their compounds are almost exclusively of the type MeOj, are very similar to the silicon derivatives, and are usually discussed with the non-metals after silicon (p. 241). The other sub-group consists of more electro-negative heavy metals: germanium, tin, and lead. These constitute the transition from tiie elements in group III, corresponding to them, to those of group IV : Ga 69.9 Ge 72.3 As 75 In 113.7 Sn 118.1 Sb 120.3 Tl 204.1 Pb 206.91 Bi 208. Their intermediate position accounts for their more non-metallic character. In this group, as in all other groups, it is noticed that as the atomic weight rises (from germanium to lead) there is a successive rise in metallo-basic character. All three members form dioxides, GeOj SnOj PbO^, which may be viewed as anhydrides of the acids GeOsHj SnOsHj PbOgHj. These are perfectly analogous to silicic acid, but their stability and acidity diminish as the atomic weights of their basal elements increase. Lead dioxide, PbOg, combines with bases (especially the alkalies), forming salts of plumbic acid, e.g., PbOaKa. These are not very stable ; water decomposes them into their components. Lead dioxide does not unite with acids to yield salts. When digested with sulphuric acid it liberates an atom of oxygen, and forms salts of lead monoxide, PbO. It yields chlorine with hydrochloric acid. In this respect lead dioxide resembles the peroxides, e.g.^ manganese peroxide, MnOa, and is commonly known as lead peroxide. However, the salts, PbOsMe,, and the organo-metallic compounds, such as Pb(CH3)4, argue in favor of quadrivalent lead, and make it perfectly analogous to tin (pp. 268,372). GERMANIUM. 365 The elements of this group yield monoxide derivatives, GeO SnO PbO. These are commonly known as ous compounds. They are basic and only form salts with acids. The basicity and stability of their derivatives increase as the atomic weights rise. The germanûus and stann¿7«j compounds are readily oxidized to derivatives of the dioxide type, while lead monoxide, PbO, is a strong base, and forms very stable salts. Compare p. 255 upon their behavior with metallic magnesium. t. GERMANIUM. Ge = 72 3. This element was discovered in 1886 by CI. Winkler, of Frei¬ berg. As early as 1871 Mendelejeff, with the periodic system as his basis, predicted the existence of an element with an atomic weight of about 73, which corresponded to the then existing gap between silicon and tin ; he called it ekasilicon (the first analogue of silicon). The perfect agreement of the essential properties of germanium with those of the theoretical ekasilicon constitutes a brilliant confirmation of the law of periodicity (p. 360). Winkler discovered germanium in the very rare mineral, argyro- dite. The latter is a double sulphide of germanium and silver, GeS2.3Ag2S. It is also present in minute quantities in euxenite (together with titanium and zirconium) (Kriiss). It may be sepa¬ rated from these minerals by fusing them with sulphur and soda. Sodium sulpho-germanate is then produced, and it is soluble in water (p. 367). To obtain free germanium, its dioxide is heated in a current of hydrogen or reduced with carbon. The product is a dark-gray powder, which melts at 900°, and upon solidifying readily crystal¬ lizes into beautiful, grayish white, metallic octahedra. Its specific gravity at 20° equals 5.469. Its specific heat was found equal to 0.0737 at 100°, and at 440°, 0.0757. Therefore, its atomic heat at 100° is 5.33 and at 440°, 5.47. It increases very slightly with rise of temperature and is a little less than the mean atomic heat (p. 262). Germanium is very stable in the air. When ignited it burns with the production of white vapors of germanium dioxide, Ge02. The metal (like silicon) is insoluble in hydrochloric acid. Nitric acid converts it (like tin) into the hydrate of the dioxide. It is soluble in alkalies upon fusion. If it be heated in the gas flame, germanium will not impart a color to the same, and it does not show a spectrum. The latter can only be produced by the action of the induction spark. Germanium, like tin, forms derivatives of the oxides GeO and Ge02 ; the first are called geim&nous compounds, the latter germamV, or derivatives of germanic acid. 366 INORGANIC CHEMISTRY. GERMANOUS COMPOUNDS. These are not very stable, and are readily oxidized to the higher form. Germanous Oxide—GeO—is formed when the hydroxide is ignited in a current of carbon dioxide. It is a grayish-black powder. Germanous Hydroxide—Ge(0H)2—is precipitated as a yellow-colored compound upon the addition of caustic alkali to the solution of the chloride. It is soluble in hydrochloric acid. Germanous Chloride—GeCh—has not been obtained pure. It is formed when hydrochloric acid gas acts upon heated germa¬ nous sulphide. Germanous Sulphide—GeS—is a reddish-brown precipitate produced by the action of hydrogen sulphide upon the solution of the dichloride. It may be obtained in grayish-black crystals by heating germanium sulphide in hydrogen gas. It is soluble in hot hydrochloric acid, forming the corresponding chloride. GERMANIC COMPOUNDS. Germanium Tetrachloride—GeCU—is formed by the direct union of germanium with an excess of chlorine. The metal, when gently heated, burns in an atmosphere of chlorine, with a bluish color. When in powder form it inflames at the ordinary tempera¬ ture. The tetrachloride is also produced if the sulphide, GeSa, be heated together with mercuric chloride. It is a colorless, mobile liquid, of specific gravity 1.887 ^8°. It boils at 86°. It fumes strongly in the air, and is decomposed by water into hydrochloric acid and germanium hydroxide, Ge(OH)4. Its vapor density,from 300-740°, equals 7.4 (air = i), or 106.5 (H = i), corresponding to the molecular formula, GeCU = 213.8. Germanium Chloroform—GeHClg—corresponding to ordi¬ nary chloroform, CHCI3 (see p. 161), is produced when metallic germanium is heated in a current of hydrochloric acid gas. It is a mobile liquid, boiling about 72°. Its vapor density approximates the molecular formula GeHClg. It becomes cloudy on exposure to the air, and colorless, oily drops of Germanium Oxychloride, GeOClj (?), separate. Germanium Bromide—GeBr^—is a strongly fuming liquid, which solidifies at 0° to a crystalline mass. Germanium Iodide—Gel^—results upon heating germanium chloride with potassium iodide, or more readily by conducting iodine vapor over heated and finely divided metal. It is an orange- colored solid, melting at 144°, and boiling at 400°. Germanium Dioxide—GeOj—germanic anhydride, is formed upon roasting the metal or the disulphide, or by treating the latter with nitric acid. It is a stable, white powder, of specific gravity TIN. 367 4.703 at 18°. It is slightly soluble in water (i part in 95 parts at 100°) and imparts to the latter an acid reaction. Germanic Hydroxide, GeCOH)«, or GeO(OH)2, Germanic Acid, is pro¬ duced by directly transposing the chloride with water. It has not been obtained perfectly pure, as it loses more or less water. Like silicic acid, it is wholly acid in its character, and only forms salts with bases. It is soluble in the hydroxides and carbonates of the alkalies, especially on fusion, while it is almost insoluble in acids. Germanic Sulphide—GeSa.—Concentrated hydrochloric acid or sulphuric acid will precipitate it from solutions of its sulpho- salts. It is also formed when hydrogen sulphide is conducted through strongly acidulated solutions of the oxide. It is a white, voluminous precipitate, insoluble in acids, but readily soluble in water. If the precipitate is washed with water it dissolves. It is reprecipitated by acids, especially if hydrogen sulphide be con¬ ducted through the solution. The sulphide dissolves readily in the fixed alkaline hydroxides and ammonia. It forms sulpho-sdXx.?, with the alkaline sulphides. These are perfectly analogous to the sul- pho-stannates. Argyrodite is an example of this class (p. 365). 2. TIN. Sn = 118.1. Tin occurs in nature principally as dioxide (Cassiterite—tin stone) in England (Cornwall), Saxony, and India. To prepare the metal the oxide is roasted, lixiviated, and heated in a furnaco' with charcoal :— SnOj 2C = Sn 4" 2CO. Thus obtained, it usually contains iron, arsenic, and other metals ; to purify it the metal is fused at a low temperature, when the pure tin fiows a\^y, leaving the other metals.. The tin obtained in the Indian isles (Malacca) is almost chemically pure, while that of England contains traces of arsenic and copper. Tin is an almost silver-white, strongly lustrous metal, with a specific gravity of 7.3. It possesses a crystalline structure; and when a rod of it is bent it emits a peculiar sound (tin cry), due to the friction of the crystals. Upon etching a smooth surface of tin with hydrochloric acid, its crystalline structure is recognized by the appearance of remarkable striations. At low temperatures perfectly pure compact tin passes gradually into an aggregate of small quad¬ ratic crystals. The metal is tolerably soft, and very ductile, and may be rolled out into thin leaves. It becomes brittle at 200°, and may then be powdered. It fuses at 228°, and distils at a white heat (about 1700®) ; it burns with an intense white light when heated in the air, and forms tin dioxide. It does not oxidize in the air at 368 INORGANIC CHEMISTRY. ordinary temperatures, and withstands the action of many bodies, hence is employed in tinning copper and iron vessels for household use. The most interesting of the tin alloys, besides bronze and soft solder, is britannia metal. It contains 9 parts tin and i part anti¬ mony, and frequently, also 2-3 per cent, zinc and i per cent, copper. Tin dissolves in hot hydrochloric acid, to stannous chloride, with evolution of hydrogen gas :— Sn + 2HCl SnClj 2H. Concentrated sulphuric acid, when heated, dissolves tin, with formation of stannous sulphate. Somewhat dilute nitric acid oxidizes it to metastannic acid; while anhydrous nitric acid, HNO3, does not change it. It dis¬ solves when boiled with potassium or sodium hydroxides, forming stannates :— Sn -f 2KOH -f Hp = SnOjK, -f 2H3. There are two series of compounds : the stannous, and stannic or compounds of stannic acid (p. 364). STANNOUS COMPOUNDS. Tin Dichloride—Stannous chloride^ SnCIa—results when tin dissolves in concentrated hydrochloric acid. When its solution is evaporated it crystallizes with two molecules of water (SnClj -|- zHjO), which it loses at 100° C. It is used in dyeing, as a mordant, under the name of Tin Salt. The anhydrous chloride, obtained by heating the metal in dry hydrochloric acid gas, fuses at 250° and distils without decomposition at 606°. Its vapor density at 600-700° agrees with the formula SnaCh ; at 900° with SnCh. Stannous chloride dissolves readily in water. Its solution is strongly reducing, and absorbs oxygen from air with the separation of basic stannous chloride :— 3SnCl2 4- O + HjO = SnCh. In the presence of hydrochloric acid, only stannic chloride is produced. Stannous chloride precipitates mercurous chloride and metallic mercury from solutions of mercuric chloride (p. 329). It unites with chlorine to form stannic chloride, and with many chlorides to yield double salts, e. g. ;— SnG2.2KCl and SnQj.2NHpi. Tin Monoxide—SnO, or Stannous oxide—is obtained by heat¬ ing its hydroxide, SnOaHj, in an atmosphere of carbon dioxide ; it is a blackish-brown powder, which burns when heated in the air, and becomes stannic oxide. Sodium carbonate added to a solution of stannous chloride precipitates white STANNIC COMPOUNDS. 369 Stannous Hydroxide—stanno-hydrate—Sn(0H)2 :— SnClj + COgNaj + HjO = Sn(0H)2 + 2NaCl + CO^, It is insoluble in ammonium hydroxide, but is readily dissolved by potassium hydroxide. Upon slow evaporation of the alkaline solution, dark crystals of SnO separate ; but, on boiling the solution, the hydrate decomposes into potassium stannate, KjSnOg, which remains dissolved, and metallic tin. The hydroxide affords salts by* its solution in acids. Stannous chloride—SnCh—and stanno-sulphate—SnS04—are formed when tin is warmed with concentrated hydrochloric or sulphuric acid. The sulphate separates in small, granular crystals, when its solution is evaporated. Tin Monosulphide—Stannous sulphide—SnS—is precipitated from stannous solutions by hydrogen sulphide, as a dark-brown amorphous precipitate. Obtained by fusing tin and sulphur to¬ gether, it is a lead-gray crystalline mass. It dissolves in concen¬ trated hydrochloric acid, with liberation of HjS, and forms stannous chloride. It is insoluble in alkaline sulphides, but, if sulphur be added and the solution boiled, it will dissolve as a sulpho-stannate (p. 370):— SnS + S -f- KgS = KjSnS,. STANNIC COMPOUNDS. Tin Tetrachloride—Stannic chloride—SnCb—is produced by the action of chlorine upon heated tin or stannous chloride—SnCh. It is a colorless liquid {Spiritus fumans Libavii), fuming strongly in moist air, of specific gravity 2.27, and boiling at 114° ; its va¬ por density equals 129 (H = 1), corresponding to the molecular formula, SnCl^ = 258.9. It attracts moisture from the air and is converted into a crystalline mass (Butter of Tin), SnCU -}- 3H2O, readily soluble in water. Boiling decomposes the solution into metastannic acid (HjSnOg) and hydrochloric acid :— SnCh -f 3H2O = HaSnOs -f- 4HCI. Stannic chloride possesses a salt-like nature, and combines with metallic chlorides to the so-called double salts, SnCl4.2KCl and SnC^. 2NH4CI ; the latter compound is known as//>/^ W/in calico printing. It also yields crystalline double salts with chlorides of the metalloids, e. g., SnCUPCls and SnCh.aSCb. Tin tetrachloride combines with hydrochloric acid, forming SnClßHj.öHjO, analo¬ gous to the chlorplatinate, PtClßHj.öHjO. It is formed when hydrochloric acid gas is conducted into a concentrated solution of tin tetrachloride in water. In the cold it solidifies to a leafy crystalline mass, melting at -|- 9®. Tin Bromide—SnBr^—forms a white, crystalline mass that melts at 30° and boils at 200°. It unites with hydrogen bromide, forming SnBrjHj.7H,0, crystal¬ lizing in yellow needles and plates. 370 INORGANIC CHEMISTRY. Tin Iodide—Snl^—consists of orange-red octahedra, fusing at 146® and boil¬ ing at 295°. It may be obtained by heating tin with iodine to 50°- Tin Fluoride—SnFl^—is only known in combination with metallic fluorides KjSnFlg), which are very similar to and generally isomorphous with the salts of hydrofluosilicic acid (SiFlgKj). Tin Dioxide—Stannic oxide—SnOj—is found in nature as tin stone, in quadratic crystals or thick brown masses, of specific gravity 6.8. It is prepared, artificially, by heating tin in the air, and it then forms a white amorphous powder. It may be obtained crystallized, by conducting vapors of the tetrachloride and water through a tube heated to redness. The dioxide is infusible, and not soluble in acids or alkalies. When fused with sodium and potassium hydroxide it yields stannates soluble in water. On adding ammonium hydroxide to the aqueous solution of tin tetrachloride or hydrochloric acid to the solution of potassium stan- nate (Sn03K2), a white precipitate of stannic acid will separate. This dissolves readily in concentrated nitric acid, hydrochloric acid and the alkalies. If preserved under water, or in vacuo, it becomes insoluble in acids and sodium, hydroxide. Both modifications appear to have the same composition, HzSnOg, and the cause of the isomerism is not yet explained. The insoluble modification is com¬ monly called metastannic acid. It is also obtained as a white pow¬ der by digesting tin with dilute nitric acid. On adding sodium hydroxide to the insoluble stannic acid it is converted into sodium metastannate, insoluble in sodium hydroxide, but readily dissolved by pure water. The salts of stannic oxide with acids, e.g., the sulphate, are not very stable, and washing with warm water decom¬ poses them. The metal salts of stannic acid are more stable. The most important of these is sodium stannate—NajSnOa 3H2O— which is employed in calico printing, under the name of preparing salts. It is produced upon a large scale by fusing tin stone with sodium hydroxide. On evaporating the solution, it crystallizes in large, transparent, hexagonal crystals, containing three molecules of water. Tin Disulphide—Stannic sulphide—8082—is precipitated as an amorphous, yellow powder by H2S from stannic solutions. If a mixture of tin-filings, sulphur, and ammonium chloride be heated it is obtained in form of a brilliant crystalline mass, consisting of gold- yellow scales. It is then called mosaic gold, and is applied in bronz¬ ing. Concentrated hydrochloric acid dissolves the precipitated disulphide, forming stannic chloride ; nitric acid converts it into metastannic acid. The sulphides and hydrosulphides of the alka¬ lies dissolve tin disulphide forming sulphostannates (see p. 226). Sodium sulphostannate, 8n83Na2 + 2H2O, crystallizes in colorless octahedra. Acids decompose the sulphostannates with the separa¬ tion of tin disulphide. LEA». 371 3. LEAD. Pb = 206.91. Lead {Plumburn) is found in nature principally as Galenite—PbS. The other more widely distributed lead ores are Cerussite—PbCOs— Crocoisite (PbCr04) and Wulfenite (MoO^Pb). Galenite is the chief source of lead ; the process of its separation is very simple. The galenite is first roasted in the air and then strongly ignited away from it. In the roasting, a portion of the lead sulphide is oxidized to oxide and sulphate :— PbS 4- 3O = PbO_+ SO2 and PbS + O4 = PbSO^. Upon ignition, these two substances react with the lead sulphide according to the following equations :— 2PbO 4- PbS = 3Pb 4- SO., and SO^Pb 4- PbS = 2Pb 4- 2SO2. Metallic lead has a bluish-white color, is very soft, and tolerably ductile. A freshly cut surface has a bright lustre, but on exposure to air becomes dull by oxidation. Its specific gravity is 11. 37- It fuses at 325°, and distils at a white heat (about 1700°). It burns to lead oxide when heated in the air. In contact with air and water lead oxidizes to lead hydroxide, Pb (0H)2, which is somewhat soluble in water. If, however, the water contain carbonic acid and mineral salts—even in slight quan¬ tity, as in natural waters—no lead goes into solution, but it is cov¬ ered with an insoluble layer of lead carbonate and sulphate. When much carbon dioxide is present the carbonate is somewhat soluble in the water. This behavior is very important for practical pur¬ poses, as lead pipes are frequently employed in conducting water. Sulphuric and hydrochloric acid have little effect on the metal, owing to the insolubility of its sulphate and chloride j yet, if the lead be in the form of powder, both acids will dissolve it. It forms lead nitrate with dilute nitric acid. Zinc, tin, and iron precipitate it, as metal, from its solution ; a strip of zinc immersed in a dilute solution of lead acetate is covered with an arborescent mass, con¬ sisting of shining crystalline leaflets (lead tree). Alloys.—^An alloy of equal parts lead and tin fuses at 186°, and is used for soldering (soft solder). An alloy of 4-5 parts of lead and I part of antimony is very hard, and answers for the manufac¬ ture of type (hard lead or type-metal). The usual lead compounds are constituted according to the type PbXa, and are called plumbic (p. 364). They show a slight similarity to the stannous derivatives. Many of the lead salts are isomorph- ous with those of barium ; the sulphates of both metals are insoluble. 372 INORGANIC CHEMISTRY. The heat of formation of some of the lead compounds equals :— (Pb,Cl2) == 82.7 (Pb,Br,) = 64.4 (Pb.T^) = 39-6 (Pb,0) = 50.3. (PbO.O) = 12.1 (Pb,S) = 20.4 (PbjSjO^) = 2i6.2 (Pb,N2,06) = 105.5. If we include the heat of solution with the above numbers, they will afford us an explanation for the deportment of lead toward acids, as well as for the various transpositions of its compounds. The heats of formation of the corresponding tin compounds are as follows (Sn.Clj) = 80.8 (Sn,Cl4) = 127 (Sn,0) = 68 (SnjOj) = 136. Compared with the lead derivatives, it is evident from these figures that tin will separate lead from the plumbates, while tin is usually precipitated from its salts by lead. Lead Oxide—PbO —is produced when lead is heated in air. After fusion it solidifies to a reddish-yellow mass of rhombic scales (litharge). When lead is carefully roasted, or the hydroxide or nitrate ignited, we obtain a yellow amorphous powder called massi¬ cot. Lead oxide has strong basic properties ; it absorbs carbon dioxide from the air, and imparts an alkaline reaction to water as it dissolves as hydroxide. Like other strong bases it saponifies fats. It dissolves in hot potassium hydroxide, and on cooling crystallizes from solution in rhombic prisms. Lead Hydroxide—Pb(OH)2.—Alkalies throw it out of lead solutions as a white, voluminous precipitate. It imparts an alkaline reaction to water, as it is somewhat soluble, and absorbs carbon dioxide with formation of lead carbonate. When heated to 130° it decomposes into lead oxide and water. If lead or the amorphous oxide be heated to 300-400°, for some time, in the air, it will absorb oxygen and be converted into a bright-red powder, called red lead^ or minium. Its composition corresponds to the formula, Pb304 ; it is considered a compound of PbO with lead peroxide (PbsO* = 2PbO + PbOj). When minium is treated with somewhat dilute nitric acid, lead nitrate passes into solution, while a dark-brown amorphous powder—lead peroxide, PbOa—remains. This oxide is more conveniently obtained by adding bleaching lime (and milk of lime) to a concentrated lead chloride solution (in calcium chloride) :— 2PbCl2 -I- Ca (0C1)2 + 2H2O = aPbOa + CaCIj + 4Ha. Or chlorine is conducted into a mixture of lead chloride (2 mois.) and calcium hydroxide (3 mois.) with water. Lead peroxide dissolves in cold hydrochloric acid to a reddish- yellow liquid containing lead tetrachloride^ PbCl^. The latter also results from the action of chlorine gas on a mixture of lead chloride with hydrochloric acid. The tetrachloride cannot be obtained free, because its solution readily decomposes into PbCl, and chlorine. Oxygen is disengaged when sulphuric acid acts upon it, and lead LEAD SULPHATE. 373 sulphate (PbS04) formed. When dry sulphur dioxide is con¬ ducted over it, glowing sets in and lead sulphate results ;— PbOj + SOa = PbSO^. When ignited PbOa breaks up into PbO and oxygen. As previously mentioned (p. 364), lead dioxide, like that of tin, has an acid nature. When warmed with potassium hydroxide, it dissolves, and on cooling, large crystals of potassium plumbate—KjPbOj -|- 3^2^—separate out ; these are perfectly analogous to potassium stannate—KjSnOg -)- 3H2O. An alkaline lead oxide solution added to a solution of potassium plumbate produces a yellow precipitate (PbgO^ -f- HgO), which loses water upon gentle warming, and is converted into red lead. Therefore, the latter must be considered as the lead salt of a normal plumbic acid, Pb(0H)4, which corresponds to stannic, Sn(OH)4, and silicic, Si(OH)4, acids:— PbgO^ = PbgPbO^. Another oxide—PbgOg—which is precipitated as a reddish-yellow powder on the addition of sodium hypochlorite to an alkaline lead solution, is very probably the lead salt of metaplumbic acid : PbgOg = PbPbOg. Nitric acid decomposes it into lead nitrate and peroxide. It dissolves in hydrochloric acid without liber¬ ation of chlorine ; this gas escapes, however, when the solution is heated, Lead Chloride—PbClj—separates as a white precipitate, when hydrochloric acid is added to the solution of a lead salt. It is almost insoluble in cold water ; from hot water, of which it requires 30 parts for solution, it crystallizes in white, shining needles. It melts about 500° and solidifies to a horn-like mass. It is volatile at a white heat ; its vapor density corresponds to the formula PbClg. Lead Iodide—Pbig—is thrown down as a yellow precipitate from lead solutions by potassium iodide; it crystallizes from a hot solution in shining, yellow leaflets, melting at 383°. Lead Nitrate—Pb(N03)2—obtained by the solution of lead in nitric acid, crystallizes in regular octahedra (isomorphous with barium nitrate) and dissolves in 8 parts water. It melts at a red heat, and is decomposed into PbO, NO2 and oxygen. When boiled with lead oxide and water, it is converted into the basic nitrate, which separates in white needles. Lead Sulphate—PbS04—occurs in nature as Anglesite^ in rhom¬ bic crystals, isomorphous with barium sulphate. It is precipitated from lead solutions as a white crystalline mass by sulphuric acid. It dissolves with difficulty in water, more readily in concentrated sul¬ phuric acid. When ignited with carbon, it is decomposed accord¬ ing to the following equation: — PbS04 -f 2C = PbS 2C0g. Lead Carbonate—PbCOs—occurs as Cerussite in nature. It is precipitated by ammonium carbonate from lead nitrate solutions. Potassium and sodium carbonates precipitate basic carbonates, the composition of which varies with the temperature and concentration 374 INORGANIC CHEMISTRY, of the solution. A similar basic salt, whose composition agrees best with the formula :— is prepared on a large scale by the action of carbon dioxide upon lead acetate. It bears the name of white lead. White lead was formerly manufactured by what is known as the Dutch process. Rolled lead sheets were moistened in earthenware pots, with acetic acid, and then covered with manure and permitted to stand undisturbed for some time. In this way, the action of the acetic acid and air upon the lead produced a basic acetate, which the COj, evolved from the decaying manure, converted into basic lead carbonate. At present it is prepared by dissolving litharge in acetic acid, and converting the resulting basic acetate into a car¬ bonate by conducting carbon dioxide into it. White lead is employed for the manufacture of white oil colors. As it is poisonous, and blackened by the hydrogen sulphide of the air (formation of lead sulphide), it is being replaced more and more by zinc white and permanent white (BaSO^). Lead Sulphide—PbS—occurs crystallized in metallic, shining cubes and octahedra. Hydrogen sulphide precipitates it as an amorphous black powder. It is insoluble in dilute acids. The lead compounds are very poisonous. The soluble salts have a sweetish, astringent taste. They are readily recognized by the following reactions : sulphuric acid precipitates white lead sulphate, which is blackened by hydrogen sulphide ; potassium iodide pre¬ cipitates yellow lead iodide. Bismuth constitutes a natural group with antimony, arsenic, phos¬ phorus and nitrogen, and, like these, affords compounds of the forms BÍX3 and BÍX5. We observed that, with increase of atomic weight, the metalloidal character of the lower members becomes more metallic (see p. 150) ; the acid nature of the oxides passes into a basic. Antimony oxide (SbjOs) is a base, while the higher oxide, SbîOs, represents an acid anhydride. In bismuth, the me¬ tallic nature attains its full value. This is manifest in its inability to unite with hydrogen. Bismuth trioxide is a base, and the pen- toxide possesses a very feeble acid character, yielding indefinite 2PbCO,.Pb(OH)2 = BISMUTH Bi = 2o8. BISMUTH. 375 compounds with the alkalies ; it behaves more like a metallic perox¬ ide, and in its properties exhibits great similarity to lead peroxide. Bismuth usually occurs native, and in combination with sulphur, as bismuthinite. To obtain the metal, the sulphide is roasted in the air, and the resulting oxide reduced with charcoal. Bismuth is a reddish-white metal, of specific gravity 9.9. It is brittle and may be easily pulverized. It crystallizes in rhombohe- dra. It fuses at 267° and distils at a white heat (about 1300°). It does not change in the air at ordinary temperatures. When heated it burns to bismuth oxide—BigOg. It is insoluble in hydrochloric acid, but dissolves in boiling sulphuric acid with formation of sul¬ phate of bismuth, and the evolution of sulphur dioxide. Nitric acid dissolves it readily in the cold. Water decomposes bismuth solutions in the same manner as those of antimony; insoluble basic salts are precipitated, but these are not dissolved by tartaric acid. Bismuthous Chloride—BiClg—arises from the action of chlo¬ rine upon heated bismuth, and by the solution of the metal in aqua regia. It is a soft white mass which fuses at 230° and boils about 435°. It deliquesces in the air. Water renders its solution turbid, a white, crystalline precipitate of Bismuth Oxychloride—BiOCl— separating at the same time :— BÍCI3 + H2O = BiOCl + 2HCI. The metalloidal character of bismuth is indicated by this reaction. The compounds BiBrg and Bilg are very similar to bismuth chloride. All three combine with many metallic haloid salts to form double halogen derivatives. Halogen derivatives of pentavalent bismuth are unknown. Bismuth Oxide—BijOg—prepared by burning bismuth or heating the nitrate, is a yellow powder, insoluble in water and the alkalies. Normal bismuth hydroxide—Bi(OH)g—is not known in a free state. Potassium hydroxide added to a bismuth solution pre¬ cipitates a white amorphous metahydrate—BiO.OH. Chlorine conducted through a concentrated potassium hydroxide solution in which bismuth oxide is suspended precipitates red bismuthic acid (BiOgH or BigH^OT), which when gently heated becomes BijOg, bismuthic oxide. Strong ignition converts the latter into BijOg and Oj ; hydrochloric acid dissolves it to bismuth chlo¬ ride, with liberation of chlorine. 376 INORGANIC CHEMISTRY. Bismuth Nitrate—Bi(N0s)3—is obtained by the solution of bismuth in nitric acid, and crystallizes with 5 molecules of HjO in large, transparent tables. In a little water it dissolves without any change ; much water renders it turbid, owing to the precipitation (NOs fNOa of white, curdy basic salts : Bi -< NO3 and Bi -( OH. The precipi- (OH (OH täte is employed in medicine under the name of Magisterium bismuthi (subm'irate). Bismuth Sulphate—BÍ2(S04)3—is formed when bismuth dis¬ solves in sulphuric acid. It crystallizes in delicate needles. Bis¬ muth Sulphide—BÍ2S3—occurring as bismuthinite, is thrown down as a black precipitate from bismuth solutions by hydrogen sulphide. Unlike antimony and arsenic sulphides, it does not form sulpho-salts. The alloys of bismuth are nearly all readily fusible. An alloy of 4 parts Bi, i part Cd, i part Sn and 2 parts Pb, fuses at 65° (Wood's metal). The alloy of 2 parts bismuth, i part lead and i part tin (Rose's metal) fuses at 94°. CHROMIUM GROUP. We observed that a group of more metallic analogous elements attached itself to the metalloidal elements, carbon, silicon and tin (p. 241) ; and further that there was an analogous group of metallic elements corresponding to the metalloidal group of phosphorus (p. 229). We now meet a group of metals, consisting of chromium, molybdenum, tungsten, and probably uranium, that bears a like relation to the elements of the sulphur group (see Periodic System of the Elements). The resemblance of these elements to sulphur and its analogues is plainly manifest in their highest oxygen deriv¬ atives (see also manganese). As the elements of the sulphur group in their highest oxygen compounds are hexavalent, so chromium, molybdenum, tungsten and uranium form acid oxides—CrOs, M0O3, WO3, UrOs. Many of the salts corresponding to these are very similar to and isomorphous with the salts of sulphuric acid. Sodium Chromate, like sodium sulphate, crystallizes with 10 mole¬ cules of water ; the potassium salts of both groups form isomor¬ phous mixtures; their magnesium salts, as well as that of tungstic acid, have the same constitution : — MgSO^ "h 7H2O and MgCrO^ -f- 7H2O. Corresponding to the acid oxides are the chlorine derivatives ;— SOjClj, CrOaClj, MoOaClj, M0OCI4, WOCI4 and WCl,, CHROMIUM. 377 which are perfectly analogous, so far as chemical deportment is concerned. Besides these highest oxides the elements of the sulphur group form the less acid oxides :— IV IV IV SOj, SeOj and TeO,. Of these the last approaches the bases. Their analogues in the chromium group : CrOa, MoOg, WO2, in which the elements appear tetr'avalent, possess an undetermined, neither acid nor basic, character. The most important basic oxide of chromium is its sesquioxide. This affords salts with the acids, and they are perfectly similar to those of the sesquioxides of iron (Fe203), manganese (Mn203), and aluminium (AI2O3) (p. 334). Since the vapor density of ferric chloride declares its formula to be Fe2Clg, we assume that in their sesquioxide compounds Cr, Mn, and Fe are tetravalent, and that these contain a hexavalent group consisting of fwo atoms of metal :— rv IV IV rv =Cr — Cr= ClgCr — CrClg, etc. We can, however, conceive these derivatives, as well as those of aluminium, aS derived from trivalent metallic atoms, and then make use of the simple, instead of the double formulas (CrXg and FeXg). (Compare pp. 343 and 350.) These facts are confirmed by the most recent determinations of the vapor density of alu¬ minium, ferric and chromic chlorides, which accords with the formulas AICI3, FeClj, and CrClg (p. 378). II Finally, compounds of chromium, CrX2, are known in which the metal figures as a dyad. These so-called chromons compounds are very much like (p. 333) the derivatives of the metals of the mag¬ nesium group, especially the ferrous salts (FeX2). They are very unstable, and are oxidized by the air into chromic compounds. Salts of molybdenum and tungsten, corresponding to the states of lowest oxidation, are not known, because these metals occur as hexads in most of their derivatives. Uranium, which has the high¬ est atomic weight of the group, shows some variations from its analogues; these are explained, as in similar cases, by its high atomic weight. I. CHROMIUM. Cr = 52.2. Chromium Í5 found principally as chromite in nature. This is a combination of chromic oxide with ferrous oxide—Cr203FeO—and occurs in North America, Sweden, Hungary, and in large quantities in the Urals. Crocoisite, or lead Chromate (PbCrOi), is not met with so frequently. Chromic iron is used almost exclusively for the preparation of all other chromium derivatives, as it is first converted into potassium Chromate (see this) by fusion with potassium car¬ bonate and nitrate. 32 378 INORGANIC CHEMISTRY. Metallic chromium may be isolated by the very strong ignition of the oxide with charcoal. It is more conveniently obtained by igniting a mixture of chromium chloride, potassium chloride or sodium chloride with zinc, in a closed crucible. The separated chromium dissolves in the molten zinc, and when the latter is dis¬ solved in nitric acid the chromium remains behind as a gray, metallic, crystalline powder, of specific gravity 6.8. It is very hard (cuts glass), and fuses with difficulty. When heated in the air it slowly oxidizes to chromic oxide ; ignited in oxygen it burns with a bright light. It dissolves readily in hydrochloric and warm dilute sulphuric acid, with elimination of hydrogen ; it is not altered by nitric acid. Three series of chromium compounds are known: chromons— CrXj, chromic—Cr2X6, and the derivatives of chromic äcid, called chromâtes. All chromium compounds are brightly colored, hence, the name chromium (from ypmiia^ color). CHROMOUS COMPOUNDS. These are very unstable, and by oxidation pass readily into ic compounds. Like ferrous salts, they are produced by the reduction of the higher oxides. The following may be mentioned : Chromous Chloride, CrCl2. This is obtained by heating chromic chloride, Cr2Cle, in a stream of hydrogen. It is a white crys¬ talline powder. It volatilizes without decomposition. At 1300-1600® the vapor density corresponds to a mixture of the molecules CrClj and CrjCi^. It dissolves in water with a blue color; the solution absorbs oxygen with avidiiy, and becomes green in color. The alkalies precipitate yellow chromous hydroxide, Cr^OH)^, from it. This is readily oxidized. When heated it parts with hydrogen and water and becomes chromic oxide : 2Cr(0H)2 = Cr^Oj -f- -f- HjO. CHROMIC COMPOUNDS. Chromic Chloride—CrjClg, like AbClg—is obtained by ignition of the oxide and charcoal in a current of chlorine. When raised to a red heat in this condition it sublimes in shining violet leaflets, which are transformed into chromic oxide by igpition in the air. Its vapor density at 1200-1300° corresponds to the formula CrClj; It vaporizes very slowly below 1000°. Pure chromic chloride only dissolves in water after long-continued boiling ; if, however, it con¬ tains traces of CrCb, it dissolves readily at ordinary temperatures. Green crystals of CrjCle + izH^O separate from the green solution on evaporation ; these deliquesce in the air. The same crystals may be obtained from solutions of chromic hydroxide, Cr2(OH)e, in hydrochloric acid. When they are dried intermediate oxychlor- ides, Cr2Cl4(OH)2, and CraClafOH)^, and at last Crj(OH)6 result. Chromic Hydroxide—Cr2(OH)6 or Cr(OH)s.—It is precipita¬ ted by ammonium hydroxide from chromic solutions as a voluminous CHROMIC OXIDE. 379 bright brown, hydrous mass. The green precipitates produced by sodium and potassium hydroxides, contain alkali that cannot be removed even by boiling water. They dissolve readily with an emerald green color, in an excess of KOH or NaOH (slightly in ammonia), but are reprecipitated upon boiling their solutions. When it is heated to 200° in a current of hydrogen, the product is the hydroxide, CrO.OH, which isa grayish-blue powder, insoluble in dilute hydrochloric acid. When chromium hydroxide is ignited, it becomes chromic oxide. Chromic Oxide—CrjOg—is a green, amorphous powder. It is also formed by the ignition of chromium trioxide ;— 2Cr03 = Cr203 + 3O, or of ammonium bichromate :— (NHJ^Cr^O, = Cr203 + 4H3O -f- N^. It may be obtained in black, hexagonal crystals, by conducting the vapors of the oxychloride through a tube heated to redness :— 2Cr02Cl2 = CfjOg -j- 2CI2 + O. Ignited chromic oxide is insoluble in acids. • When fused with sili¬ cates, it colors them emerald green, and serves, therefore, to color glass and porcelain. 6^«z¿7íí?/j^^^^«isabeautifully green-colored chromium hydroxide, which is applied as a paint. It is obtained by igniting a mixture of one part potassium bichromate with three parts boric acid ; after treating the mass with water, which dissolves potassium borate, there remains a green powder, the composition of which corresponds to the formula :— CrjOCOH)^ = Cr203.2H20. The predominating properties of chromic oxide are basic, as it readily affords salts with acids ; yet its basic nature, like that of all sequioxides, is but slight, so that it does not afford salts with weak acids (p. 380). In addition to all this it possesses a slightly acidic character, and metallic salts are derived from it, generally from the hydroxide, CrO.OH, which are analogous to thealuminates (p. 354). Salts, like (CrO.O)2Mg and (CrO.O)2Zn, can be obtained crystal¬ lized in regular octahedra by fusing chromic oxide with metallic oxides and boron trioxide (as flux). Chromic iron is such a salt ;— Cr203.FeO = (CrO.Olj.Fe. 380 INORGANIC CHEMISTRY. Chromium Sulphate—012(804)3—is obtained by dissolving the hydroxide in concentrated sulphuric acid. The solution, green at first, becomes violet on standing, and deposits a violet-colored crys¬ talline mass. This may be purified by solution in water and pre¬ cipitation by alcohol. This salt crystallizes from very dilute alco¬ hol in bluish-violet octahedra containing 15 molecules of water. If the aqueous solution of the violet salt be heated, it assumes a green color, because the salt breaks up into free acid and a basic salt which, upon evaporation, separates as a green amorphous mass, soluble in alcohol. When the green solution stands, it reverts to the violet of the neutral salt. The other chromic salts, the nitrate and the alum, behave in a similar manner. Chromium sulphate forms double salts with the alkaline sulphates —the chromium alums (p. 355). Potassium Chromium Alum—CrK(S04)2 -f 12H2O—rcrys- tallizes in large, dark violet octahedra. It may be most conve¬ niently prepared by acting upon a solution of potassium bichromate mixed with sulphuric acid, with sulphur dioxide :— + H^SO, -f 3SO2 = Cr^CSOJa-K^SO, -f H2O. At 80° the violet solution of the salt becomes green, and on evaporation affords an amorphous green mass. As chromium oxide possesses only a slightly basic nature, salts with weak acids, like CO2, S02,H2S (see Aluminium, p. 355) do not exist. Therefore, the alkaline carbonates and sulphides precipitate chromium hydroxide from solutions of chromium salts :— Cr2(504)3 "b 3^33003 -|- sHjO = Cr2(0H)g -j- 3Na2S04 -|- 3^0, and Cr^fSOJa + 3(NH4)2S + ÓH^G = Cr^fOH)^ -f 3(NH4)3S04 -f 3H2S. Ammonium sulphide produces a black precipitate—CrS—in solu¬ tions of chromous salts. DERIVATIVES OF CHROMIC ACID. In its highest oxygen derivative, CrOs, chromium possesses a com¬ plete metalloidal, acid-forming character. Chromic acid, HjCrOi, is perfectly analogous to sulphuric acid, H2SO4, but has not been obtained free, since when liberated from its salts it at once breaks up into the oxide and water :— HjCrO^ = CrOg -1- HjO. The chromâtes are often isomorphous with the sulphates (p. 376). Folychromates also exist, and are derived from polychromie acids CHROMIUM TRIOXIDE. 381 produced by the condensation of several molecules of the normal acid (see Disulphuric acid, p. 197) :— KjCrO^ K2Cr20^, K2Cr30j^Q, etc. Potassium Chromate. Potassium bichromate. Potassium trichromate. The constitution of these salts is expressed by the following for¬ mulas :— .OK Cr02. ^OK /OK CrO / \o Crû \OK CrO.^/ OK. The free polychromie acids are not known, because as soon as they are separated from their salts, they immediately break up into the acid oxide and water :— ^2^1-3010 = sCrOs + H2O. The polychromates are frequently, but incorrectly, called acid salts ; true acid or primary salts, in which only one H atom is re¬ placed by metal (CrO^KH), are unknown for chromic acid. The salts of normal chromic acid are mostly yellow-colored, while the polychromates are red. The latter are produced from the former by the action of acids :— 2K2Cr04 -f 2HNO3 = K2Cr207 -j- 2KNO3 + H2O. Conversely, by the action of the alkalies, the polychromates pass into the normal salts :— KaCrgOy + 2KOH = 2K2Cr04 + H2O. Their formation may also be as follows : The chromic acid liberated from its salts by stronger acids breaks up into water and the acid oxide, which combines with the excess of the normal Chromate :— CrO^Kj -1- Cr03= KgCtgO^. When there is an excess of acid the anhydride (CrOs) is set free. Chromium Trioxide—Chromic Acid Anhydride—CrOs.—It consists of long, red, rhombic needles or prisms, obtained by add¬ ing sulphuric acid to a concentrated potassium bichromate solution. The crystals deliquesce in the air and are readily soluble in water. When heated, they blacken, melt, and at about 250° decompose into chromic oxide and oxygen :— 2Cr03 = Cr303 + 3O. 382 INORGANIC CHEMISTRY. Chromium trioxide is a powerful oxidizing agent, and destroys organic matter ; hence its solution cannot be filtered through paper. When alcohol is poured on the crystals, detonation takes place, the alcohol burns, and green chromic oxide remains. By the action of acids, e. g., sulphuric, the trioxide deports itself like a peroxide ; oxygen escapes and a chromic salt results. When heated with con¬ centrated hydrochloric acid chlorine is evolved :— 2Cr03 + 12HCI = CraClg + 6H2O + 3CI2. Reducing substances, like sulphurous acid and hydrogen sulphide, convert chromic acid into oxide : — 2Cr03 + 3H,S = CroOg + 3H„0 + 3S. 2Cr03 + 3SO2 + 3H2O = Cr203 + 3SO,H2. Chromate of Potassium—KjCrO^—is obtained by adding potassium hydroxide to potassium bichromate. It forms yellow rhombic crystals, isomorphous with potassium sulphate {K2SO4) ; isomorphous mixtures crystallize out firom the solution of the two salts. Bichromate of Potassium—KjCrjOi—called acid potassium Chromate, is manufactured on a large scale, and bears the name red Chromate of potash in commerce. It is obtained by igniting pul¬ verized chromite, CrjOsFeO, with potashes and nitre, whereby potassium Chromate and ferric oxide are formed. The fusion is treated with water, and the resulting solution of potassium Chro¬ mate, KjCrOi, mixed with acetic or nitric acid (see p. 381), when potassium bichromate crystallizes from the concentrated solution. In practice the above method is advantageously replaced by the following: The pulverized chromite is ignited, together with lime, in furnaces allowing air access. Calcium Chromate (CaCrO^) (together with ferric oxide) is produced. This dis¬ solves in dilute sulphuric acid to bichromate, CaCrjO^ ; the latter is converted by potassium carbonate into potassium bichromate. Bichromate of potassium crystallizes in large, red, triclinic prisms, soluble at ordinary temperatures, in 10 parts of water. When heated, the salt fuses without change ; at a very high heat it decomposes into potassium Chromate, chromic oxide and oxy¬ gen 2K2Cr20^ = 2K2CrO¿ ^*"2^3 ~1~ ®3* When the salt is warmed with sulphuric acid, oxygen escapes and potassium chromium alum is produced :— + 4H2SO4 = Cr2(S0j3.K2S04 + 4H2O + 3O. This reaction answers for the preparation of perfectly pure oxygen. Further, the mixture is made use of in laboratories, as an oxidizing agent. CHROMYL CHLORIDE. 383 Chromate of Sodium—NajCrO^ -f- loHjO—forms deliques¬ cent crystals, and is analogous to Glauber's salt (NajSO* -f- loHjO). Barium and Strontium Chromates—BaCr04 and SrCrO^—are almost insoluble in water. Calcium Chromate—CaCrO^—dis¬ solves with difficulty in water, and crystallizes like gypsum with two molecules of water. The magnesium salt, MgCrO^-f- 7H2O, dis¬ solves readily and corresponds to Epsom Salt. The chromâtes of the heavy metals are insoluble in water, and are obtained by trans¬ position. Chromate of Lead—PbCrO^—is obtained by the precipitation of soluble lead salts with potassium Chromate. It is a yellow amorphous powder which serves as a yellow paint. When heated it melts undecomposed, and solidifies to a brown, radiating crystal¬ line mass. It oxidizes and easily decomposes all the carbon com¬ pounds, and is therefore used in their analysis. In nature lead Chromate exists as crocoisite. The oxide, CrOj, called peroxide, may be obtained by the careful ignition of chromium trioxide, and is, most likely, a salt-like compound: CrgOg.CrOg or CrO.CrOg. Its hydrate is precipitated from chromium solutions upon the addi¬ tion of potassium Chromate. On warming the peroxide with hydrochloric acid, chlorine is evolved. Chromic Acid Chloranhydrides.—Chromic acid forms chloranhydrides similar to those of sulphuric acid (p. 199)- Cor¬ responding to SO2CI2, we have chromyl chloride, Cr02Cl2 ; and f CI for the first sulphuric acid chloranhydride, SO2 -j qjj, is the salt. Chromyl Chloride—Cr02Cl2—Chromium oxychloride, is pro¬ duced by heating a mixture of potassium (or sodium) bichromate (or monochromate) and sodium chloride with sulphuric acid :— CrO^Nag + aNaCl -f 2SO4H2 = CrOgCla + 2S04Na2 + 2H2O, CrgOyNag + 4NaCl + 3SO4H2 = 2Cr02Cl2 + sSO^Nag + 3H2O. The water produced at the same time must be absorbed by the excess of sulphuric acid. To prepare chromyl chloride, first fuse salt (10 parts) with potassium dichromate (12 parts) or with potas¬ sium monochromate (17 parts). The yellowish-brown mass is broken into coarse pieces, placed into a retort provided with a CI OK- VI /CI VI /CI VI /OK Cr02\ci CrOg^oK CrOg^oK' and 384 INORGANIC CHEMISTRY. condenser, and anhydrous or slightly fuming sulphuric acid (30 parts) poured over them. When a gentle heat is applied chromyl chloride distils over and is purified by further distillation. It is a red, transparent liquid, of specific gravity 1.92 at 25°, and fumes strongly in the air. It boils at ii6°-ii8° ; its vapor density equals 77 (H =1), corresponding to the molecular formula Cr02Cl2 — 154. Chromyl chloride has a strong oxidizing action. With water it is decomposed according to the following equation :— CrOaCla + H2O = CrOg + 2HCI. /CI Chloro-chromic Acid, C!r02^Qj^ (see above), is only known in its salts. The potassium salt is formed by heating potas¬ sium bichromate (3 parts) for a short period, with concentrated hydrochloric acid (4 parts) :— Cr^O.K^ + 2HCI = + H^O. It crystallizes from the solution on cooling in flat, red prisms. Heated to 100° it gives up chlorine. It is decomposed by water into hydrochloric acid and potassium bichromate : — 2Cr02<^^K + ^2^ = + zHŒ Chromium Hexafluoride, CrFlg, corresponds to the trioxide. It may also be regarded as the fluoranhydride of normal chromic acid, Cr(0H)6. It is obtained by heating a mixture of lead Chro¬ mate and calcium fluoride with fuming sulphuric acid. The result¬ ing dark-red vapor may be condensed in a platinum or lead tube to a very volatile red liquid. It fumes strongly in the air, and with water decomposes into chromic acid and hydrofluoric acid :— CrFlg + 3H2O = CrOj -f 6HF1. It reacts very energetically with glass. The chromium compounds can be readily recognized by their color. The following reaction is very characteristic for chromic acid : On adding hydrogen peroxide to a solution of chromium trioxide, or the acidified solution of a Chromate, the red liquid is colored blue. The nature of the compound causing this colora¬ tion is not known ; it is usually assumed to be a higher oxide x)f chromium. On shaking the blue solution with ether, the latter withdraws the blue compound and is beautifully colored in conse¬ quence. The ethereal solution is somewhat more stable than the aqueous. Both are gradually decolorized, with liberation of oxygen. MOLYBDENUM. 385 3. MOLYBDENUM. Mo = 95.9. Molybdenum occurs rather rarely in nature ; usually as molybdenite (MoSj) and wulfenite (MoO^Pb). Free jnolybdenum is obtained as a silver-white metal, of specific gravity 8.6, by igniting the chlorides or oxides in a stream of hydrogen. It is very hard, and fuses at a higher temperature than platinum. When heated in the air it oxidizes to molybdenum trioxide. It is soluble in concentrated sul¬ phuric and nitric acids. It is also converted by the latter into insoluble M0O3. Like chromium, molybdenum affords compounds of the forms M0X2, MoX^ and MoXg; besides which derivatives are known in which it appears to act as a pentad and also as a triad. Molybdenum Dichloride—MoCl^—resulting from the trichloride, M0CI3, when heated in a stream of carbon dioxide (together with MoCl^), is a bright yellow, non-volatile powder. It is converted by potassium hydroxide into the hydrate, Mo(OH)2, a black powder. Molybdenum Trichloride—M0CI3 or Mo2Clg—^produced by gentle heating (at 250°) of M0CI5 in a current of H or CO2, is a reddish-brown powder, which, when strongly ignited, yields a dark-blue vapor. It dissolves with a beautiful blue color in concentrated sulphuric acid, upon heating, with an emerald green color. Potassium hydroxide converts it into the hydroxide, Mo(OH)3 or Moj (OH)g, which forms salts with acids. The ignition of the hydrate affords the black oxide, M02O3. Strong heating of the trichloride in a current of COj leaves MoClj and it sublimes. Molybdenum Tetrachloride—MoCl^— is a brown, crystalline powder, which appears to break up by evaporation into M0CI5 and M0CI3. It yields a hydrate with ammonium hydroxide, forming salts with acids. The brown solution of the salts readily assumes a blue color by oxidation in the air. The ignition of the hydrate leaves the dioxide, M0O2, which is converted by nitric acid into the tii- oxide, M0O3. Molybdenum disulphide, MoSg, is produced by the ignition of the trisulphide, M0S3, away from air. It is a shining black powder, which occurs native as molybdenite^ in hexagonal, graphite-like crystals, with a specific gravity of 4-5- Molybdenum Pentachloride—M0CI5—is prepared by heating MoSj or molybdenum in dry chlorine gas. It is a metallic, shining, black, crystalline mass, fusing at 194° and distilling at 268° ; its vapor density equals 136, corresponding to the molecular formula M0CI5 == 272.6. It fumes and deliquesces in the air, and dissolves in water with hissing. Its aqueous solution has a brown color. It dissolves in absolute alcohol and ether with a dark-green color. The hexachloridef MoClg, is not known, but the oxychlorides, MoOCl^, and M0O2CI2, are. The first results from the ignition of MoOj and carbon in a stream of chlorine, and is a green crystalline mass subliming under 100° and yielding a dark-red vapor. Bromine forms perfectly analogous compounds with molybdenum. Molybdenum Trioxide—M0O3—^results on roasting metallic molybdenum or the sulphide in the air. It is a white, amorphous mass, which turns yellow on heating ; it fuses at a red heat and then sublimes. It is insoluble in water and acids; but dissolves readily in the alkalies and ammonium hydroxide. When fused with the alkaline hydroxides or carbonates, salts are produced, partly derived from the normal acid, HjMoO., and partly from the polyacids, and correspond to the polychromates :— K2M0O4, K2Mo20y, K2Mo30io, Na2Mo40i3, K^Mo,0^^, etc. The ammonium salt—(NH4)2Mo04—is obtained by dissolving the trioxide in 33 386 INORGANIC CHEMISTRY. concentrated ammonium hydroxide. In the laboratoiy it serves as a reagent for phosphoric acid. Alcohol causes it to separate out of its solution in crystals ; upon evaporation, however, the salt (NH^)gM07O24 + 4H2O crystallizes out. Both salts are decomposed by heat, leaving molybdenum trioxide. Hydrochloric acid added to a concentrated solution of a molybdate precipitates molybdic acid—H2M0O4. It is a white, crystalline compound, readily dissolved by an excess of acid. Zinc (tin, stannous chloride, sulphur dioxide) added to this solution causes it to become blue and then green (formation of sesquioxide), in consequence of the formation of lower oxides (like MogOg = 2M0O3.M0O2), and finally brownish-red and yellow, when a suboxide, M05O7 = 2M0O3.M0O, is produced. The final product, by the action of tin and hydrochloric acid, is molybdenum sesquioxide, MojOj. Potassium permanganate converts all these lower oxides into molybdic acid. Molybdic acid can also form polyacids with phosphoric and arsenic acids, e.g.^ H3PO4.11M0O3. These complex phosphomolybdic acids are distinguished by the fact that they form salts insoluble in dilute acids with the metals of the potassium group, with ammonia and with organic bases. On adding a solution containing phosphoric (or arsenic) acid to the nitric acid solution of ammonium molybdate, there is produced a yellow crystalline precipitate of ammonium phospho-molyb¬ date—(!NH4)3P04. 11M0O3 -1- óHgO. This reaction serves for the detection and separation of phosphoric acid. Molybdenum Trisulphide—M0S3—is thrown down as a brown precipitate from acidulated molybdenum solutions by hydrogen sulphide. It dissolves in alkaline sulphides forming sulpho-salts. Ignited away from air it is converted into molybdenum disulphide, MoSg, which occurs native as molybdenite. In addition t(j these molybdenum compounds in which the element is hexava- lent there is a Molybdenum Tetrasulphide—M0S4. From this are derived persulpho- molybdic acid, MoSgHg and its salts, e. g.^ MoSgKj and M0S5KH. 3. TUNGSTEN. W= 184. Tungsten is found in nature in the tungstates : as wolframite, FeW04, scheelite, CaW04, and as stolzite, PbW04. The metal is obtained, like molybdenum, by the ignition of the oxides or chlo¬ rides in a stream of hydrogen, in the form of a black powder, or in steel-gray crys¬ talline leaflets, having a sp. gr. 19.1. It is very hard and difficultly fusible. It becomes trioxide when ignited in the air. Tungsten forms the following chlorides: WCl,, WCI4, WCI5 and WClg. The Dichloride—^WClg—arises by strong ignition of WClg and WCI4 in a cur¬ rent of carbon dioxide, and is a bright gray, non-volatile mass. The Tetrachloride—WCI4—obtained by gentle ignition of WClg and WClg, in a current of hydrogen or carbon dioxide, is grayish-brown and up>on sublima¬ tion decomposes into WCU and volatile WCI5. It forms a brown oxide (WO,) with water. The Pentachloride—WCI5—is obtained by the distillation of WClg in a cur¬ rent of hydrogen or carbon dioxide, and consists of shining, black, needle-like crystals. It fuses at 248° and boils at 275®, forming a dark-brown vapor, with the density 180 (WCI5 = 360.4). It affords an olive-green solution and a blue oxide, WgOg, with water. It dissolves with a deep blue color in carbon disulphide. Tungsten Hexachloride—WClg—is produced when the metal or a mixture of wolframite with carbon is heated in a current of chlorine. It forms a dark violet, crystalline mass, fusing at 275° and boiling at 346®. The vapor density URANIUM. 387 equals 198 (WClg = 395.8). It dissolves in carbon disulphide with a reddish- brown color ; it forms WO3 with water. The Oxychloride—WCl^O—consists of red crystals, fusing at 210° and boiling at 227° ; its vapor density equals 170 (WCl^O = 341). The Dioxychloride— WCljOj—sublimes in bright yellow, shining leaflets. Tungsten Trioxide—WO3—is thrown out of the hot solution of tungstates by nitric acid, as a yellow precipitate insoluble in acids, but dissolving readily in potassium and sodium hydroxides. Tungstic acid, WO(OH)^, is, however, pre¬ cipitated from the cold solution, but on standing over sulphuric acid it becomes / W02(0H)2 and at 100° passes into ditungstic acid, = W205(0H)2. When tungstic acid is reduced in hydrochloric acid solution by zinc it first becomes blue (formation of WgOj) and then brown, when the salt of the dioxide, WO2, is formed. Potassium permanganate oxidizes this to tungstic acid. The salts of tungstic acid are perfectly analogous to the molybdates and are derived from the normal acid or the polyacids. The normal sodium salt, Na2W04 -|- 2H2O, and the so-called meta-tungstate of sodium, Na2W^Oj3 + loHjO, are applied practically. Materials saturated with their solutions do not burst into a flame, but smoulder away slowly. The reduction of the tungstates (by fusion with tin, etc.) affords peculiar com¬ pounds, e. g., K2W3O9 or KgW^Ojg; these have various colors, possess metallic lustre, and are applied as tungsten bronzes. Tungstic acid also combines with phosphoric and arsenic acids, forming deriva¬ tives analogous to those of molybdic acid with the same acids. The metal is used in the manufacture of tungsten steel ; a slight quantity of it increases the hardness of, the latter very considerably. 4. URANIUM. Ur or U = 239.4. In nature it occurs chiefly as uraninite, a compound of uranic and uranous oxides—UO2.2UÜ3 = U3O8. This mineral has been discovered to contain as high as 3 % of nitrogen. The metal is obtained by heating uranous chloride with sodium. It has a steel-gray color and a specific gravity of 18.7. When heated in the air it burns to uranous-uranic oxide. Its specific heat equals 0,0267, and its atomic volume is therefore 6.6. It melts about 1500®. There are two series of uranium com¬ pounds. In the one, the metal is a tetrad UX¿; these uranous or urano-compounds are very unstable, and pass readily into the uranic or derivatives of hexavalent uranium. Uranous oxide is of a basic nature, and only forms salts with acids. The compounds of hexavalent uranium are called the uranic compounds. UO3 and U02(0H)2 have a predominant basic character, but are also capable of form¬ ing salts with bases which are called uranates. In the salts derived from acids, e. g., U02(N03)2 and U02S0^, the group UO2 plays the rôle of a metal ; it is called uranyl, and its salts are termed uranyl salts. They may also be regarded as basic salts. URANOUS COMPOUNDS. Uranous Chloride—UCl^—is obtained by heating metallic uranium in a stream of chlorine, or uranous oxide in hydrochloric acid. It consists of dark green octahedra with metallic lustre. It volatilizes at ä red heat, forming a red vapor, whose density agrees with the formula UCl^. It deliquesces in the air, and dis- 388 INORGANIC CHEMISTRY. solves with hissing in water. Uranous hydroxide remains when the solution is evaporated. Uranous Oxide—UOj—^is formed when the other oxides are heated in a cur¬ rent of hydrogen. It is a black powder, which becomes uranous-uranic oxide, UO2.2UO3, when heated in the air. Uranous oxide dissolves with a green color in hydrochloric and concentrated sulphuric acids. Uranous sulphate, U(S04)2 + SIIjO, consists of green crystals. From the salts the alkalies precipitate the voluminous, bright green uranous hydroxide, U(OH)^, which becomes brown on exposure. HEXAVALENT URANIUM COMPOUNDS. Uranium Hexachloride—UClg—has not been obtained, but the oxychloride, UOgClg (Uranyl chloride), exists ; it is obtained by heating UOj in dry chlorine gas, or by the evaporation of uranyl nitrate with hydrochloric acid. It is a yellow crystalline mass, deliquescing in the air. Uranic Oxide, UO3 or Uranyl oxide, UOj O—is a yellow powder, and is obtained by heating uranyl nitrate to 250°. When warmed with nitric acid it be¬ comes uranyl hydrate or uranic acid, U02(0H)2, which is also yellow-colored. Uranyl Nitrate—U02(N03)2—^results from the solution of uranous or uranic oxide, or more simply of uraninite in nitric acid. It crystallizes with six mole¬ cules of water, in large, greenish-yellow prisms, which are readily soluble in water and alcohol. On adding sulphuric acid to the solution, Uranyl sulphate —UO2SO4 -|- 6H2O—crystallizes out, on evaporation, in lemon-yellow neecQes. If sodium or potassium hydroxide be added to the solutions of uranyl salts, yellow precipitates of the uranates —U2O7K2 and U207Na2—are obtained. These are soluble in acids. In commerce the sodium salt is known as uranium yeUow, and is employed for the yellow coloration of glass (uranium glass) and porcelain. The uranates can be obtained in crystalline form, by igniting uranyl chloride with alkali chlorides in the presence of ammonium chloride. Zinc and sulphuric acid reduce uranic to uranous compounds. The so-called uranic-uranous oxide, which constitutes uraninite, and is formed by the ignition of the other oxides in the air, must be viewed as uranous VI IV uranate—2U03.UC)2 = (U02.02)2 U. Many uranium salts exhibit magnificent fluorescence. The oxide colors glass fluxes a beautiful greenish-yellow (uranium glass). Uranous oxide—UOj—imparts a beautiful black color to glass and porcelain. Besides these compounds, in which uranium appears to be tetravalent and hex- avalent, it also affords apentachloride, UCI5, like molybdenum and tungsten. The same results on conducting chlorine gas over a moderately heated mixture of car¬ bon with one of the uranium oxides. It consists of dark needles which, in direct light, are metallic green, but in transmitted, ruby red. It deliquesces in the air to a yellowish-green liquid ; upon heating it is dissociated into UCl^ and CI (at 120®- 235°). There is also a tetroxide, UO4, which, like the trioxide, UO3, yields salts with the bases. It corresponds to molybdenum tetrasulphide M0S4 (p. 386). MANGANESE. 389 MANGANESE. Ma = 54.8 (ss-o)- According to its atomic quantity, manganese bears the same re¬ lation to the elements of the chlorine group as chromium to the elements of the sulphur group. The relationship manifests itself distinctly in the higher states of oxidation. Permanganic oxide, MnjOj, and acid, HMnOi, are perfectly analogous to CI2O7 (or I2O7) and HCIO4. The permanganates and the Perchlorates are very similar, and for the most part are isomorphous. The man¬ ganese in them appears to be heptavalent, like the halogens in their highest state of oxidation. The similarity of manganese to the halogens is restricted to this one point of resemblance. In the rest of its derivatives, manganese shows great resemblance to the ele¬ ments standing in the same horizontal series of the periodic system, viz., with iron and chromium (p. 349). Like these two elements, it forms three series of compounds. 1. In the manganous derivatives—MnXj—the metal is divalent. These salts are the more stable, and comprise the most common manganese compounds. They resemble and are usually isomor¬ phous with the ous salts of iron and chromium, and the salts of metals of the magnesium group (p. 314). 2. The manganic compounds—MuaXg—are similar to and isomor¬ phous with the ferric, chromic and aluminium derivatives ; they are, however, less stable, and easily reduced to the manganous state. Their composition is due to the tetravalent nature of manganese (P- 377). 3. The derivatives of manganic acid—HaMnO^ = MnOg (OH)j, in which manganese is hexavalent—are analogous to those of ferric (HzFeOi) and chrbmic acid (HjCrOi), and, of course, to those of sulphuric acid (H2SO4). Consequently, in manganese we plainly observe how the similarity of the elements in their compounds is influenced by the valence (see p. 333). In its ous condition, manganese, like the elements of the magnesium group, has a rather strong basic character, which diminishes considerably in the ic state. Hexavalent manganese has a metalloidal acidic character, and, in manganic acid, approaches sulphur. By the further addition of oxygen, manganese finally (in permanganic acid) acquires the metalloidal character of the halo¬ gens. We have already noticed that many other metals, especially chromium andiron, exhibit a similar behavior. Osmium tetroxide, OSO4, wholly resembles the halogens. On-the other hand, the metalloidal and the weak basic metals acquire a strong basic, alkaline character, by the addition of hydrogen, or hydrocarbon groups (CHg, CjHj). The groups, NH^ (ammonium), P(CH5)4 (tetramethylphospho* 390 INORGANIC CHEMISTRY. nium), SiCjHg^g (triethylsulphine), Sn(C2H5)3 (tin triethyl), etc., are of metallic nature, because their hydroxides, P(CH3)^.OH, S(C2H5)j.OH, Sn(C2H5)3.0H, are perfectly similar to the hydroxides (KOH, NaOH) of the alkali metals. Manganese is widely distributed in nature. It is found native in meteorites. Its most important ores are pyrolusite, MnOa, haus- mannite, MnsO^, braunite, MnjOa, manganite, MnaOg.HaO, and rhodochroisite, MnCOa. Metallic manganese is obtained by igniting the oxides with char¬ coal. It has a grayish-white color, is very hard, and fuses with difficulty; specific gravity 7.2. It oxidizes readily in moist air. It decomposes water on boiling, and, when dissolved in acids, forms manganous salts. The heat of formation of the most important manganese compounds corre¬ sponds to the symbols :— MnjO.HjOl = 94.7 (MnjClj) = 111.9 (Mn,Cl2,Aq.) = 128.0. Mn,02,H20) = 116.2 (Mn,S,OJ = 249.8 (Mn,0^,K) = 194.8. MANGANOUS COMPOUNDS. Manganous Oxide—MnO—results from ignition of the car¬ bonate, with exclusion of air, and by heating all manganese oxides in hydrogen. It is a greenish, amorphous powder, which, in the air, readily oxidizes to MugOi. Manganous Hydroxide—Mn(0H)2—is a voluminous, red¬ dish-white precipitate, formed by the alkalies in manganous solu¬ tions. When exposed to the air, it oxidizes quickly to manganic hydroxide, Mn2(OH)6. Manganous salts usually have a pale, reddish color, and are formed by the solution of manganese or manganic oxides in acids. Manganous Chloride—MnCla—crystallizes with four molecules of water in reddish tables. On drying, it is decomposed with separation of hydrochloric acid. Anhydrous manganous chloride is prepared by igniting the double salt MnCla.aNHiCl + H2O (see Magnesium Chloride), or by heating manganese oxides in hydrochloric acid gas ; it is a crystalline, reddish mass, which deliquesces in the air. Manganous Sulphate—MnSO^—crystallizes below -f- 6° with 7 molecules of H2O (like magnesium and ferrous sulphates), and at ordinary temperatures with 5H2O (like copper sulphate) ; the last molecule of water does not escape until 200°. It forms double salts with the alkaline sulphates, e. g., MnSO^.KjSO* -f 6H,0. MANGANIC COMPOUNDS. 391 Manganous Carbonate—MnCOg—exists in nature as rhodo- chroisite, and is precipitated by alkaline carbonates from manganous solutions, as a white powder, which turns brown on exposure. Manganous Sulphide—MnS—is found in nature as alaban- dite or manganese blende. Alkaline sulphides precipitate a flesh- colored sulphide from manganese solutions. It becomes brown in the air. MANGANIC COMPOUNDS. Manganic Oxide—MnjOg, manganese sesquioxide—is a black powder produced by the ignition of the manganese oxides in a cur¬ rent of oxygen gas. It occurs as Braunite in dark-brown quadratic crystals. Manganic Hydroxide—MugCOHg) or Mn(OH)3, manganic hydrate—is precipitated by ammonium hydroxide from manganous solutions containing ammonium chloride as a dark-brown mass. It dissolves in cold hydrochloric acid to a dark-brown liquid, con¬ taining, in all probability, manganic chloride, MnClg or MnjClc. When this is heated it decomposes into MnCb and chlorine. Manganite, occurring in iron-black crystals, is the hydroxide, MnACOH)^ or MnO.OH. Manganous-manganic Oxide—Mn304 = MnO.MnaOa. It constitutes the mineral hausmannite, crystallized in dark-gray quad¬ ratic octahedra, and is obtained as a reddish-brown powder by the ignition of all other manganese oxides in the air. It reacts with hydrochloric acid, according to the equation :— MnjO^ -1- 8HC1 = sMnCb -f 4H2O + CI,. Since manganic oxide is quadratic in its crystallization, while all other sesqui- oxides (like corundum and hematite) are rhombohedral, and since the first is decomposed by dilute nitric and sulphuric acids into MnOj and a manganous salt, it has been generally supposed that manganic oxide is not a sesquioxide, but rather a compound of the dioxide with manganous oxide :— IV O MnO-.MnO == MnO(f ^Mn. Hausmannite is quadratic, whi^e other metallic oxides (the spinels, p. 354 and p. 379, and magnetite, FejO^) are isometric ; therefore the former is not consid¬ ered a compound of manganese sesquioxide and protoxide :— MnO.O. Mn„O..MnO = ✓Mn, Mn0.0/ 392 INORGANIC CHEMISTRY. but as manganous oxide and the dioxide ;— IV O.Mn> MnO,.2MnO ■= MnO(^^ ^O. ^O.Mn^ This is shown by its behavior toward dilute nitric and sulphuric acids, which decompose it into manganese dioxide, and two molecules of manganous oxide. Chrysoberyl, unlike other spinels, is trimetric, and other reactions clearly prove (chiefly their deportment with concentrated sulphuric acid) that manganic and mangano manganic oxides are to be regarded as sesquioxide derivatives. Manganic oxide, like the other sesquioxides, is a very feeble base, which does not afford salts with dilute or weak acids, and by sep¬ aration of oxygen reverts to the manganous condition. Its salts are very unstable. Manganic Sulphate—Mn2(S04)3—is obtained by the solution of manganic oxide, hydroxide, or, better, manganous-manganic oxide in concentrated sulphuric acid. When the last oxide is employed manganous sulphate also results. The best procedure is to heat the hydrate of manganese dioxide (see below) with concen¬ trated sulphuric acid to i68°, when the sulphate will separate as an amorphous, dark-green powder. It dissolves with a dark-red color in a little water. It forms alums, with potassium and ammonium sulphates—e. g., Mn2(S04)3. -|- 24H2O. Much water will decompose these with the separation of manganic hydroxide. Manganese Dioxide—Mn02—peroxide. This is the mineral pyrolusite, occurring in dark-gray radiating masses, or in almost black rhombic prisms, which possess metallic lustre. When gently heated it is converted into oxide, by strong ignition into manga- nous-manganic oxide:— 3Mn02 = MujO^ -f- 2O. It is used for making oxygen. Manganous oxide results at a white heat. Chlorine escapes when it is warmed with hydrochloric StCid * MnOg + 4HCI r= MnCl2 -f 2H2O -f O,. The dioxide may be obtained artificially by heating manganous nitrate to 150-160°. Its hydrates—Mn02.H20 and Mn02.2H20 —are produced on adding a hypochlorite to the solution of a man¬ ganous salt, or if chlorine be conducted through a solution of man¬ ganese containing sodium carbonate, or by adding KMnO^ to a boiling solution of a manganous salt. The precipitated dioxide dissolves in cold hydrochloric acid, without liberating chlorine, as MANGANIC COMPOUNDS. 393 MnCl^ is probably formed ; when heat is applied it breaks up into MnCla and Cb. This deportment would indicate that manganese is tetravalent in the dioxide. Manganese dioxide also unites with bases, affording the so-called manganites, e. g., MnjOsBa and Mn, O5K,. Manganese peroxide (also MnjOg and MngO^) serves chiefly for the manufac¬ ture of chlorine gas, and it is, therefore, important from a technical point to esti¬ mate the quantity of chlorine which a given dioxide of manganese is able to set free. This is done by boiling the oxide with hydrochloric acid, conducting the liber¬ ated chlorine into a potassium iodide solution, and determining the separated equivalent amount of iodine by means of sodium hyposulphite. Or the oxide is heated in a flask with oxalic and sulphuric acids, when the oxalic acid is oxidized to carbon dioxide, and from the quantity of this set free we can calculate the quantity of active or available oxygen in the manganese oxide. In the preparation of chlorine the manganese is found in the residue as man- ganous chloride. With the relatively high value of pyrolusite, it is important for trade that the peroxide be recovered from the residue. This regeneration is at present largely executed by the method proposed by Weldon, according to which the manganous chloride, containing excess of hydrochloric acid, is neutralized with lime, the clear liquid brought into a high iron cylinder (the oxidizer), milk of lime added and air forced in. The mixture becomes warm, and so-called cal¬ cium' manganite, MnOgCa = MnOg.CaO, is precipitated as a black mud :— MnClg 4" 2CaO O = MnOgCa CaClg, The calcium chloride solution is run off, and the residual calcium manganite employed for the preparation of chlorine, when it conducts itself as a mixture of MnOg -(-CaO. COMPOUNDS OF MANGANIC AND PERMANGANIC ACID. When oxygen compounds of manganese are heated in the air in contact with potassium hydroxide, or, better, with oxidizing sub¬ stances, like nitre or potassium chlorate, a dark-green amorphous mass is produced, which dissolves in cold water, with a dark-green color. When this solution is evaporated under the air-pump, dark- green metallic rhombic prisms of potassium manganate— EbMnOi-^crystallize out. This salt is isomorphous with potassium sulphate and Chromate. It suffers no change by solution in potas¬ sium or sodium hydroxide, but is decomposed by water, brown hy- drated manganese dioxide separating, and the green solution of the manganate changing into a dark-red solution of the permanganate, KMnO^: — SKgMnO^ + SHgO = 2KMiiÔ^ + MnOg.HgO -f 4KOH. A similar conversion of the green manganate into red perman¬ ganate occurs more rapidly under the influence of acids :— SKgMnO^ -f 4HNO3 = 2KMn04 + MnOg -f 4KN0g -f 2H2O. Owing to this ready alteration in color the solution of the man< ganate is called chameleon mineral. 394 INORGANIC CHEMISTRY. Potassium Permanganate—KMnO^—is best prepared by conducting CO2 into the raanganate solution until the green color has passed into a red. When the solution is concentrated the salt crystallizes in dark-red rhombic prisms isomorphous with potassium Perchlorate, KCIO4. It is soluble in twelve parts of water at ordi¬ nary temperatures. The permanganate solution is a strong oxidizing agent, converting lower oxygen compounds into higher, and in doing this it is reduced to a colorless manganous salt. When a permanganate solution is added to an acidulated ferrous solution, the former is decolorized, and there results a faintly yellow-colored solution of ferric and manganous salts:— 2KMn04 + loFeSO^ + SH^SO^ = aMnSO^ + 5Fe2(S04)3 + SH^O -f- K^SO^. Hence the solution of this salt serves for the volumetric estima¬ tion of ferrous salts. In the same manner, the permanganate oxidizes and destroys many organic substances, therefore its solution cannot be filtered through paper; it serves as a disinfectant. The permanganate is also reduced by hydrogen peroxide (p. 105); the reaction proceeds according to the following equa¬ tion :— Mn^OjKjO -|- 5H2OJ = 2MnO -(- K2O -j- 5II2O -f- 5^2' the formation of oxides requires the presence of acids (sulphuric acid) for the completion of the reaction. The remaining permanganates are similar to and isomorphous with the Perchlorates. The sodium salt is very soluble in water, and does not crystallize well. Very cold sulphuric acid added to dry permanganate causes the separation of Manganese Heptoxide—MujOt—a-n oily, dark- colored liquid. By careful warming it is converted into dark-violet vapors, which explode when heated rapidly. Manganese heptoxide has a violent oxidizing action ; paper, alcohol and other orgamic matter are inflamed by mere contact with it. METALS OF GROUP VIII. Of the known elements, those standing in the eighth column of the periodic system remain for consideration (p. 249) :— Fe =56 Co = 58.6 Ni = 58.6 Ru = ICI.7 Rh = 103.0 Pd = 106 Os = 192 Ir = 193.2 ft = 194.8 METALS OF GROUP VIII. 395 These elements are the middle members of the three great periods, and they have no analogues in the two short periods (pages 246, 248). As regards both atomic weights and physical and chemical deportment, these elements constitute a transition from the preced¬ ing members of the great periods (Mn and Cr, Mo, W) to the next following members (Cu, Ag, An, and Zn, Cd, Hg, p. 324). The elements standing side by side (heterologous) and belonging to the same periods are very similar in their physical properties, and show, e. g., very close specific gravities. They are, therefore, usually arranged in groups, and distinguished as, (i) the iron group (Fe, Co, Ni), with the specific gravity 7.8-8.6; (2) the group of the light platinum metals (Ru, Rh, Pd), with the specific gravity 11.8-12. i, and (3) the group of the heavy platinum metals (Os, Ir, Pt), with the specific gravity 21.1-22.4. On the other hand, the homologous elements (Fe, Ru, Os ; Co, Rh, Ir j and Ni, Pd, Pt) show a like similarity in their chemical properties, as do the other homologous groups, and therefore may be considered in such groups. This resemblance shows itself chiefly in their combination forms, and, of course, too, in the properties of the compounds (p. 333). We know that the metals of group VI (chromium, molybdenum, tungsten) and of group VII (manganese) form the highest oxides (MeOs and MejOY) having an acidic nature. In the adjacent elements of group VIII (iron, ruthenium, and osmium) we find salts;— FeO^Kj, RUO4K2, OSO^K2, derived from the unstable trioxides FeOs, RuOs and OsOg. This acid-forming function disappears in the following members, Co, Rh, Ir, and Ni, Pd, Pt ; their chemical valence diminishes rapidly and they attach themselves to Cu, Ag and Au. Consequently the whole physical and chemical deportment of the 9 elements about to be considered is governed by their position in the periodic system. As mentioned on pp. 247,252 and at other places, the valences of the elements in their highest salt-forming oxides present themselves as periodic functions of the atomic weights. A similar dependence is also seen in the lowest salt-forming oxides, and may be observed in the following tabulation of both classes of oxides of the middle members of the great periods :— V VI VII VI IV II IV VjOg CrOj MiijOj Fe03 C02O3 h CUO h m GeOj "» II II II n NiO I ZnO Ga,0, « V2O3 CrO MnO FeO CoO CujO ^ ® GeO 396 INORGANIC CHEMISTRY. NbPg M0O3 III n MoO — VI RUO3 II RuG nr RhjOa 11 RhO II PdO AgîO II CdO III InaOj IV SnO, II SnO VI Tb^O, WO3 III II (Ta^Og) (WO) — VI OS03 II OsO IV IrO, II IrO IV PtO^ II PtO III AU2O3 AujO II HgO HgP III ThOg TI3O IV PbO, II PbO. METALS OF THE IRON GROUP. The metals of this group, iron, cobalt and nickel, form a gradual transition from manganese to copper. Their magnetic properties distinguish them from the other elements. Iron forms three series of compounds after the forms, FeOg, FejOa and FeO. In its highest combinations iron has an acidic character, and the derivatives of ferric acid (HaFeO^) are perfectly similar to those of chromic and manganic acids (p. 389) ; they are, however, less stable than the latter. Their analogues with cobalt and nickel are unknown. The ferriV compounds—Fe^Xg—containing the hexavalent group VI Fea (p- 377)1 are much like the aluminium, chromic and manganic derivatives. They are generally isomorphous with them. They are characterized among iron salts by their relative stability. The highest oxides of cobalt are far less stable, and only a few double salts of this form are known, while the higher salts with nickel are unknown. Again, iron, cobalt, and nickel afford ous compounds, (FeXj, CoXjjNiXî) in which they appear to be dyads. They resemWe the compounds of chromium, manganese, and copper of the same form, and those of the magnesium metals. The ferrous salts are not as stable as the ferric ; they are readily oxidized to the latter. The cobaltous and nickelous compounds are quite stable, and in this respect these metals ally themselves with copper and zinc. I. IRON. Fe = 56. This metal, of such great practical importance, is very widely distributed in nature. It is found native on the earth's surface almost exclusively in meteorites ; it is, however, present in great masses in other worlds which (like the sun) are surrounded by an atmosphere of hydrogen. IRON. 397 The most important iron ores are ; magnetite (FcjO^), hematite (FcaOs), brown iron ore and limonite (hydrates of the oxide) and siderite (FeCOg). These ores constitute almost the sole material for the manufacture of iron ; the sulphur ores, like pyrite, are less adapted to this end. In commerce there are three varieties of iron : cast-iron, steel, and wrought-iron. Their chief chemical difference is in the vari¬ able quantity of carbon contained in them. Cast-iron contains 3-6 per cent, carbon, in part chemically combined, and in part mechanically mixed in the form of graphite. When molten cast-iron is cooled rapidly it yields the so-called white irofiy in which the greater portion of the carbon is chemically combined with the iron. It has a whitish color, exhibits a granular crystalline structure upon fracture, and is very hard and brittle. Its specific gravity is 7. It fuses to a pasty mass about 1200°, and is on this account not suited for castings. The chemically com¬ bined carbon in it can easily be removed by oxidation, and, there¬ fore, it is adapted for the manufacture of steel or wrought iron. When molten cast-iron is allowed to cool slowly, the greater part of the carbon in it separates in the form of small leaflets of graphite. The gray cast-iron produced in this way has a darker gray color, is not so hard and brittle, fuses more readily (about 1150°) than white cast-iron, and serves for the manufacture of castings. Neither variety can be forged or welded, on account of its brittleness. Steel contains 0.8-1.8 per cent, of carbon, all of which is chemically combined with the iron. It has a steel-gray color and a fine-grained structure; its specific gravity equals 7.6-8.0. It fuses with more difficulty (about 1400°) than cast-iron, but easier than wrought-iron. When molten steel is rapidly cooled, it be¬ comes very hard and brittle. In this process more carbon is chemically combined. If cooled slowly, it is soft and malleable, and may be forged and welded. Welding becomes more and more difficult with the addition of carbon. Wrought-iron contains the least amount of carbon, 0.2-0.6 per cent. It possesses a bright-gray color, has a specific gravity of 7.6, is rather soft and tough, and, at a red heat, may be readily forged, rolled, and welded. The rolled iron possesses a fibrous texture, while the forged is fine grained j the former is more com¬ pact and tenacious. Wrought-iron fuses at a bright white heat (1500°). Metallurgy of Iron.—The extraction of iron from its oxygen ores is based upon the reduction of the same by carbon at a red heat. In the oldest method, the ores were heated with carbon in wind furnaces ; in this way the excess of heat consumed the greater portion of the carbon, and the product was an iron poor in 398 INORGANIC CHEMISTRY. carbon, wrought-iron, a spongy mass, which was then forged under the hammer. The present methods were adopted since the beginning of the previous century. According to these cast-iron is first prepared from the ores, and this afterwards converted into steel or wrought-iron. The smelting of the ores is executed in large, walled blast furnaces, that permit the process to proceed without interruption. The furnaces are filled from openings above, with alternating layers of coal, broken ore and fluxes containing silica and lime ; the latter facilitate the melting together of the reduced iron. The air necessary for the process is blown into the contracted portion of the furnace by means of a blast engine. The combustion of the coal aflbrds carbon monoxide, which reduces the iron oxides to metal :— ^^2^3 "4" 3^^ = 2Fe -J- 3^C)a* As the reduced iron sinks in the furnace it comes in contact with the coal, takes up carbon and forms cast-iron, which fuses as it sinks lower and flows into the hearth of the furnace. Protracted and strong heating converts the chemically united carbon into the graphitic form, and thus accelerates the formation of the gray cast-iron. The earthy impurities of the ores combine with the fluxes to a readily fusible slag, which envelops the fused iron and protects it from oxidation. To convert the cast-iron thus produced into steel or wrought-iron, carbon must be withdrawn from it. In making the wrought-iron the cast-iron is fused in open hearths (refining process), or in reverberatory furnaces with air access, and the mass stirred thoroughly until it has become semi-pasty (puddling process). In this way almost all the carbon is burned to carbon monoxide and the other admixtures, like silicon, sulphur, and phosphorus, present in small quantities, are oxidized. The wrought-iron is then worked up by rolling, or under the iron hammers (bar- iron). Steel was formerly manufactured from wrought-iron (not cast-iron), by cement¬ ation. The iron bars, mixed with fine charcoal, were exposed to a red heat, when the iron took up carbon from the surface. The bars were then reforged, again heated with fine charcoal, and the process repeated until the mass became as homogeneous as possible (cementation steel). A more uniform steel is obtained if it be fused in crucibles (cast-steel). At present, steel is chiefly prepared directly from cast-iron, by the method invented by Bessemer, somewhere in 1850. It consists in blowing air, under high pressure, into the molten iron, until the necessary amount of carbon has been con¬ sumed (Bessemer steel). An iron rich in silicon (1.5-2% Si) is well adapted to the purpose, because by the simultaneous combustion of the silicon the temperature is considerably aug¬ mented. The operation is conducted in a pear-shaped vessel known as the con¬ verter. The air is blown in through openings in the bottom. The decarboniza- tion is only partial in this way. This is better accomplished by the English method, which removes the carbon so as to convert the mass into wrought iron, and then it is again carbonized by adding molten spiegeleisen. The Bessemer process is only adapted to crude iron containing as little sulphur and phosphorus as possible (at the highest, .05%?), because in this process this phosphorus is not consumed, but remains unaltered in the steel. By a slight, yet very essential alteration, Thomas and Gilchrist (1880) rendered it suitable for iron containing much phosphorus. Their process is now known as the " basic pro¬ cess," and consists in lining the converter with a basic lining material, composed of clay and silica, mixed with lime and magnesia. By contact with these sub¬ stances (bases) the iron is completely dephosphorized and the phosphorus changed to calcium phosphate. All the phosphorus contained in iron ores collects in the slag of the converter. IRON. 399 The latter contains as much as 15-20 % phosphoric acid (and may even be in¬ creased to 30% ), existing as calcium phosphate, and this may be applied as a fer¬ tilizer in agriculture. A third method bf making'steel consists in puddling the different varieties of iron together, or with iron ores. Martin steel is obtained by fusing cast-iron with wrought-iron. It is much used. Uchatius steel is prepared by fusing cast-iron together with some iron ore and pyrolusite. The various and more recent processes for manufacturing steel and wrought- iron, the knowledge that their difference is mainly in hardness, and that the so- called Bessemer steel is not tempered, have led to the introduction of a new divi¬ sion and nomenclature for these substances (which are difficultly fusible and mal¬ leable compared with cast-iron). We now distinguish :— (i) Weld iron as a non-fused, non-tempered mass, formerly wrought-iron ; (2) weld steel, not fused, tempered, formerly puddle steel; {3) ingot iron, fused, nc^ tempered, formerly Bessemer steel ; (4) ingot steel, has been fused and tempered. Ordinary iron, even the purest wire, always contains foreign ingredients, principally carbon and manganese, and minute quanti¬ ties of silicon, sulphur, phosphorus, nitrogen, nickel, cobalt, titan¬ ium and others. The quantity of manganese is purposely increased (to 30 per cent.), as by this means the iron acquires valuable tech¬ nical properties ; it becomes more compact and solid. When iron, containing carbon, is dissolved in hydrochloric acid the chemically combined carbon unites with hydrogen, forming hydrocarbons, while the mechanically admixed graphite remains behind. The whole quantity of carbon is determined by the solution of the iron in bro¬ mine water or cupric chloride, when all the carbon remains behind. To prepare chemically pure iron, heat the pure oxide or the oxal¬ ate in a current of hydrogen :— FeP3 + 3H2 = 2Fe + 3H20; the iron then remains as a fine black powder. If the reduction occurs at a red heat, the powder glows in the air, and burns (pyro- phoric iron). The strongly ignited powder is not inflammable. Iron obtained by the electrolysis of ferrous sulphate contains some hydrogen. Chemically pure iron has a grayish white color, is tolerably soft, and changes but slowly in the air. Its specific gravity is 7.78. It melts in an oxy-hydrogen flame at 1800°. Ordinary iron rusts rapidly in moist air, as it covers itself with a thin layer of ferric hydroxide. When ignited in the air it is coated with a layer of ferrous-ferric oxide (Fe304) which is readily detached. It burns with an intense light in oxygen. In contact with a magnet iron becomes magnetic ; steel alone retains the magnetism, while cast-iron and wrought-iron soon lose the property after the removal of the magnet. 400 INORGANIC CHEMISTRY. Iron decomposes water at a red heat, with the formation of fer¬ rous-ferric oxide, and the liberation of hydrogen :— 3Fe + 4H2O = FegO^ + 4H2. The metal dissolves without trouble in hydrochloric and sulphuric acids, with evolution of hydrogen ; the latter has a peculiar odor, due to hydrocarbons that are liberated at the same time. Iron dis¬ solves in nitric acid with separation of nitric oxide. On dipping iron into concentrated nitric acid, and then washing it with water, it is no longer soluble in the acid (passive iron) ; this phenomenon is probably due to the production of ferrous oxide upon its surface. Metallic iron unites with carbon monoxide, forming Fe(C0j5, ferro-pentacar- bonyl. This is a liquid boiling at 102.8® and solidifying below—21°. It isdecom- posed on exposure to the light, with the formation of yellow-colored flakes, which probably have the formula Fe2(C0),j {Chem. your. 59, 1090). FERROUS COMPOUNDS. These are prodiiced by the solution of iron in acids, and may also be obtained by the reduction of ferric salts :— Fe2Cl6 -f Zn = 2FeCl2 + ZnClj. In the hydrous state they are usually of a green color ; in the air they oxidize to ferric salts :— 2FeO -j- O = Fe203. Ferrous Chloride—FeClj—crystallizes from aqueous solutions in green monoclinic prisms, with four molecules of water. These deliquesce in the air and oxidize. When dried they sustain a par¬ tial decomposition. The anhydrous salt is formed by conducting hydrogen chloride over heated iron. It is a white mass, which fuses on application of heat and sublimes at a red heat in white, six-sided leaflets. Its vapor density at 1300-1500° corresponds to the formula FeClz, but it appears that at lower temperatures it is also possible for the molecules FeaCh to exist. It forms double salts with the alkaline chlorides, e. g. :— FeCh-zKCl -f 2H2O. Ferrous Iodide—Fela—-is obtained by warming iron with iodine and water. It crystallizes with four molecules of water. Ferrous Oxide—FeO—is a black powder, resulting from the reduction of ferric oxide by carbon monoxide. When warmed it oxidizes readily. Ferrous Hydroxide—Fe(OH)2—is thrown out of ferrous solutions by the alkalies, as a greenish-white precipitate. Exposed to the air, it oxidizes, becoming green at flrst, then red¬ dish-brown. It is somewhat soluble in water, and has an alkaline reaction. Ferrous Sulphate—FeSO*—crystallizes with 7 molecules of H2O in large, greenish, monoclinic prisms, and is generally called FERROUS SULPHIDE. 401 green vitriol. The crystals effloresce somewhat in dry air. They oxidize in moist air, and become coated with a brown layer of basic ferric sulphate. At ioo° they lose 6 molecules of HjO, and change to a white powder. The last molecule of water escapes at 300°. Therefore, ferrous sulphate behaves just like the sulphates of the metals of the magnesium group. Like them, it unites with alkaline sulphates to double sulphates, which contain six molecules of water, e.g.y SO^Fe.SO^Ka -j- óHjO. These are more stable than ferrous sulpliate, and oxidize very slowly in the air. Ferrous sulphate is obtained by dissolving iron in dilute sulphuric acid ; or from pyrites (FeSj). When the latter are roasted they lose one molecule of sulphur, and are converted into ferrous sulphide (FeS), which, in the presence of water, absorbs oxygen from the air, and is converted into sulphate, which may then be extracted by water. Iron vitriol has an extended practical application ; among other uses, it is employed in the preparation of ink, and in dyeing. When heated it decomposes according to the following equa¬ tion ;— 2FeS04 = FCgOg 4" SO3 -f- SOg. On this is based the production of fuming Nordhausen sulphuric acid (p. 198), and of colcothar. Ferrous Carbonate—FeCOg—exists in nature as siderite, crystallized in yellow-colored rhombohedra, isomorphous with cal- cite and smithsonite. Sodium carbonate added to ferrous solutions precipitates a white voluminous carbonate, which rapidly oxidizes in the air to ferric hydroxide. Ferrous carbonate is somewhat soluble in water containing carbon dioxide, hence present in many natural waters. Ferrous Phosphate—Fe3(P04)2 + 8H2O—occurs crystallized in bluish monoclinic prisms as Vivianite. Precipitated by sodium phosphate from ferrous solutions, it is a white amorphous powder, which oxidizes in the air. Ferrous Sulphide—FeS—is a dark-gray, metallic mass, ob¬ tained by fusing together iron and sulphur. It is made use of in laboratories for the preparation of hydrogen sulphide. If an intimate mixture of iron filings and sulphur be moistened with water, the union will occur even at ordinary temperatures. Black ferrous sulphide is precipitated from ferrous solutions by alkaline sulphides. When the moist sulphide is exposed to the air it oxidizes to ferrous sulphate. The alkaline sulphides also precipitate ferrous sulphide from ferric salts, but the latter first suffer reduc¬ tion ;— FCgCl, + (NH^)gS = 2FeCl2 + 2NH,a + S, and FeClj + (NHJgS = FeS -f 2NHp. 34 402 INORGANIC CHEMISTRY. FERRIC COMPOUNDS. Ferric Oxide—Sesquioxide of Iron—FejOa—exists in nature, in compact masses, as hematite, and as iron mica, in dark-gray metallic rhombic prisms. It may be prepared by heating the iron oxygen compounds in the air, and is obtained on a large scale by the igni¬ tion of green vitriol. It is then a dark-red powder {colcothar or caput mortuuni) used as a paint and for polishing glass. Ferric Hydroxide—Fe3(OH)6—is precipitated by alkalies from ferric solutions as a voluminous, reddish-brown mass. On boiling, it becomes more compact, gives up water, and is converted into the hydrate, Fe20(0H)4. Many iron ores, like bog-iron ore, FejO- (OH)^, pyrosiderite, Fe202(0H)2 (isomorphous with diaspore), and brown hematite Fe403(0H)6, are derived in a similar manner. Freshly precipitated ferric hydroxide is soluble in a solution of ferric chloride or acetate. When such a solution is subjected to dialysis, the iron salt diffuses, and there remains a pure aqueous solution of ferric hydroxide. All of the latter is precipitated as a jelly from such a solution upon the addition of a little alkali or acid. Ferrous-Ferric Oxide—FegO^—FeO-FcaOs—occurs in nature crystallized in black regular octahedra—magnetite. It is abundant in Sweden, Norway, and the Urals. It may be obtained artificially by conducting steam over ignited iron (p. 400). Magnetite con¬ stitutes the natural magnets. Ferric hydroxide, like other sesquioxides, is a feeble base, and does not yield salts with weak acids, like carbonic or sulphurous (P- 355)- Ferric salts arise by the solution of ferric oxide in acids, or by the oxidation of ferrous salts in the presence of free acids (best by chloric or nitric acids) :— aFeSO^ -f H^SO^ -f O = Fe^fSOJg -f H^O. They generally have a yellow-brown color, and are converted by reduction into ferrous salts :— Fe2a6 -f H^S = 2 FeCl, + 2HCI -f. S. Ferric Chloride—FejClg.—It is obtained in aqueous solution by conducting chlorine into a solution of ferrous chloride :— aFeClj + Clj = Fe^Clg. The hydrate—FejCle + 6H2O—remains upon evaporation. It is a yellow crystalline mass, readily soluble in water, alcohol, and ether. It is partially decomposed when heated ; hydrogen chloride escapes, and a mixture of chloride and oxide remains. Anhydrous ferric chloride is produced by heating iron in a cur¬ rent of chlorine gas ; it sublimes in brownish-green, metallic, shin¬ ing, six-sided prisms and scales, which deliquesce in the air. The specific gravity of their vapor at 440° closely approximates the for- IRON CYANIDES. 403 mula FcaCle (found 10.5 ; calculated 11.2) ; with rising temperature it diminishes successively and from 750-1050° corresponds to the simple formula FeCh. At still higher temperatures a partial decom¬ position occurs into ferrous chloride and chlorine (Grünewald and V. Meyer), Ferric Sulphate—Fe2(S04)3—is obtained by dissolving the oxide in sulphuric acid. When its solution is evaporated, it remains as a white mass, which gradually dissolves in water, with a reddish- brown color. It forms alums (p. 350) with alkaline sulphates, Fe2(SOJ3. K2SO4 -f 24H2O. Potassium iron alum. Ferric Phosphate—Fe2(P04)2—is a white precipitate, thrown out of ferric solutions by sodium phosphate. It is insoluble in water and acetic acid. Ferric Sulphide—FeSa—occurs in nature as pyrites, crystallized in yellow, metallic, shining, regular cubes or octahedra. It is employed in the manufacture of sulphuric acid and green vitriol. The artificial sulphide can be prepared in many ways. COMPOUNDS OF FERRIC ACID. On fusing iron filings with nitre, or by conducting chlorine into potassium hydroxide, in which ferric hydroxide is suspended. Potassium Ferrate, K2Fe04, is produced, and crystallizes from the alkaline solution in dark-red prisms. This salt is isomorphous with potassium Chromate and sulphate. It dissolves quite easily in water; but the dark-red liquid soon decomposes with separation of ferric hydroxide and oxygen. The free acid is not known, as it immediately breaks up when liberated from its salts. CYANOGEN DERIVATIVES OF IRON. Iron unites with the cyanogen group to form compounds which are very characteristic, and important in a commercial sense. When potassium cyanide is added to aqueous solutions of the ferrous or ferric salts, the cyanides, Fe(CN)2 and Fe2(CN)8, are thrown down as white precipitates, but decompose rapidly in the air. They dis¬ solve in an excess of potassium cyanide to form the double cyan¬ ides, Fe (CN)2«4KCN and Fe2(CN)6.6KCN. When acidsare added to these solutions the hydrogen compounds, H4FeCy6*(= FeCyj. 4HCy) and Fe6Cyi2H6(=: FcaCye-öHCy) separate. These are of acid nature, and form salts by exchanging their hydrogen for metals. The iron and the cyanogen group in these * The cyanogen group, CN, is usually designated by the letters Cy. 404 INORGANIC CHEMISTRY. salts and the free acids cannot be detected by the usual reagents (e.g., the iron is not precipitated by the alkalies). It is supposed that compound groups of peculiar structure are present in these double cyanides, and that they conduct themselves like the halogens. The group, FeCye, ia the ous compounds is called ferrocyanogen, that of FezCyia in the ic, ferricyanogen. The ferro- behave toward the ferri-compounds the same as the ferrous toward the ferric salts ; oxidizing agents convert the former into the latter, and reducing agents transform the latter into the former:— zFeCyeK, + CI^ = KeFe^Cy^^ + 2KCI and Fe^Cy^jKe + 2KOH + = aK.FeCy« zHp. Cobalt, manganese, chromium and the platinum metals afford similar cyanides. Potassium Ferrocyanide—Yellow Prussiate of Potash— K4FeCy6, is produced by the action of potassium cyanide upon iron compounds, or upon free iron (in which case the oxygen of the air or water takes part). It is prepared commercially by igniting carbonized nitrogenous animal matter (blood, horns, hoofs, leather offal, etc.) with potashes and iron. In this operation, the carbon and nitrogen of the organic matter combine with the potassium of the potashes to form potassium cyanide, while the sulphur present forms iron sulphide with the iron. (By means of alcohol, potassium cyanide can be extracted from the fusion.) Upon treating the fusion with water, the potas¬ sium and iron sulphide react upon each other, and ferrocyanide of potassium results and is purified by crystallization :— FeS + 6KCy = K^FeCyg + KjS. It crystallizes from water in large, yellow, monoclinic prisms, having three molecules of water, and soluble in 3-4 parts HjO. The crystals lose all their water at 100°, and are converted into a white powder. At a red heat the ferrocyanide breaks up into potassium cyanide, nitrogen, and iron carbide (FeCa). When the salt is warmed with dilute sulphuric acid, half of the cyanogen escapes as hydrogen cyanide ; concentrated sulphuric acid decom¬ poses it, according to the following equation :— K^FeCye -f- ÓH^SO^ + ÓHp = FeSO^ + + sSO^ÍNH^)^ + 6CO. When strong hydrochloric acid is added to a concentrated solu¬ tion of potassium ferrocyanide hydrogen ferrocyanide, H4FeCy6, separates as a white crystalline powder, which soon turns blue in the air. It has the nature of an acid. Its salts with the alkali and alkaline earth metals are very soluble in water. The sodium salt FERRICY ANIDES. 405 crystallizes with difficulty. The salts of the heavy metals are insol¬ uble, and are obtained by double decomposition. When potassium ferrocyanide is added to the solution of a ferric salt a dark-blue cyanide (FeCy6)3(Fe2)2, called Prussian Blue is precipitated :— SK.FeCye + aFe^Cl^ = (FeCy6)3(Fe2)2 + 12KCI. This is the ferric salt of hydroferrocyanic acid ; and if potassium or sodium hydroxide is poured over it, it is converted into ferrocy¬ anide of potassium and ferric hydroxide :— VI VI (FeCy6)3(Fe2)2 + 12KOH = aK^FeCy^ + 2Fe2(OH)3. Potassium ferrocyanide produces a reddish-brown precipitate of FeCyeCuj in copper solutions. Oxidizing agents convert the ferro- into potassium ferricya- nide—K«Fe2Cyi2—redprussiate of potash. This conversion is most conveniently effected by conducting chlorine into the solution of the yellow prussiate :— 2K4FeCy6 -f- CI2 = K6Fe2Cyi2 + 2KCI. The ferrocyanogen group, FeCye, is then changed to the ferri-, Fe2Cyh (p. 404)- The red prussiate crystallizes from water in red rhombic prisms. The free hydroferricyanic acid, H6Fe2Cyi2, is precipitated upon the addition of concentrated hydrochloric acid. It is rather unstable. With ferrous solutions potassium ferricyanide affords a dark-blue precipitate, Fe3Fe2Cyij, very similar to Prussian Blue, and called Turnbull 'j Blue :— KeF^zCyia + sFeSO^ = Fe2Cyi2Fe3 3KßO^. This blue is the ferrous salt of hydroferricyanic acid. Alkalies convert it into ferricyanide of potassium and ferrous hydroxide:— Fe3Cyi2Fe3 -f 6K0H = Fefy^K, -f sFeCOH)^.* Potassium ferricyanide does not cause precipitation in ferric solu¬ tions. Ferrocyanide yields Prussian blue, while it forms a bluish white precipitate in ferrous solutions. By these reactions, ferric salts may be readily distinguished from * According to recent investigations it appears that Turnbull's blue and Prus- VI sian blue possess the same composition (FeCyj)2 | S®®. The simpler relations are retained here. '• ® 406 INORGANIC CHEMISTRY. the ferrous. Potassium sulphocyanide (CNSK) produces a dark- red coloration in ferric solutions, while it leaves the ferrous un¬ altered. a. COBALT. Co = 58.6. Occurs in nature as smaltite (CoAsj) and cobaltite (CoAsa.CoS,). The metal is obtained by the ignition of cobaltous oxide with car¬ bon, or in a current of hydrogen. It has a reddish-white color and strong lustre, is very tenacious, and fuses with difficulty. Its specific gravity is 8.9. It is attracted by magnets, but to a less degree than iron. It -s not altered by the air or water. It is only slightly attacked by hydrochloric and sulphuric acids ; nitric acid dissolves it readily, forming cobaltous nitrate. The predominating compounds have the form C0X2, and are called cobaltous. They are very stable, and generally isomorphous with the ferrous salts. The hydrous cobaltous compounds have a reddish color, the anhydrous are blue. COBALTOUS COMPOUNDS. Cobaltous Chloride—CoCh—is obtained by the solution of cobaltous oxide in hydrochloric acid, and crystallizes with 6HjO in red monoclinic prisms. When heated, it loses water, and becomes anhydrous and blue in color. Characters made with this solu¬ tion upon paper are almost invisible, but when warmed they become distinct and blue (sympathetic ink). Cobaltous Hydroxide—Co(OH)2—is a reddish precipitate produced by the alkalies in hot, cobaltous solutions. When exposed to the air, it browns by oxidation. Basic salts are pre¬ cipitated from cold solutions. When heated out of air contact, the hydroxide passes into green cobaltous oxide^ CoO. Cobaltous Sulphate—SO4C0 -f- 7H2O—crystallizes in dark- red monoclinic prisms ; the hydrated sulphate, C0SO4 -j- 6H2O, separates from hot solutions. It is isomorphous with ferrous sul¬ phate, and yields double salts with alkaline sulphates. Cobaltous Nitrate—Co(N03)2 -j- 6H2O—forms red deliques¬ cent prisms. Cobaltous Sulphide—CoS—is a black precipitate, produced in neutral cobalt solutions by alkaline sulphides. It is insoluble in dilute acids. Cobalt Silicates.—When glass is fused with a cobalt com¬ pound it is colored a dark blue, and when reduced to a powder is used as a paint, under the name of smalt. COBALT. 407 Smalt is prepared commercially by fusing cobalt ores with potashes and quartz. The cobalt forms a silicate (smalt) with the SiOj and potassium, while the other metals accompanying it in its ores, such as Bi, As, and especially nickel, are thrown out as a speiss. This is called speiss-cobalt and serves for the preparation of nickel. On igniting cobalt oxide, C02O3, with alumina, a dark-blue mass is produced—cobalt ultramarine or ThenariTs Blue. When zinc oxide and cobalt oxide are ignited a green color—green cinnabar or Rinman^s Green—is obtained. COBALTIC COMPOUNDS. Cobaltic Oxide—C02O3—is left as a black powder on the igni¬ tion of cobaltous nitrate. It becomes cobaltous-cobaltic oxide, C03O4, at a red heat, and cobaltous oxide at a white heat. The hydroxide—Co2(OH)6—separates as a dark-brown powder, if chlor¬ ine be passed through an alkaline solution containing a cobaltous salt. A cobaltous salt is produced, and oxygen set free, when sul¬ phuric acid acts upon the oxide or the hydroxide. Chlorine is generated when it is heated with hydrochloric acid :— CojOg 6HC1 = aCoClg 3H2^ ~1~ Clg. The cobaltic hydroxide dissolves in dilute, cold hydrochloric acid, with scarcely any liberation of chlorine ; the solution prob¬ ably contains CojClg, which decomposes into 2C0CI2 and CI2 on evaporation. Cobaltous-Cobaltic Oxide — C02O3.C0O—corre¬ sponding to magnetite, FejO^, is formed upon the ignition of the oxygen cobalt derivatives, and is a black powder. Only a few salts of cobalt in the ic state are known. The most interesting of these is potassio-cobaltic nitrite. When potassium nitrite, KNO2, is added to a cobaltous solution acidified with acetic acid, nitrogen is set free, and in course of time Co2(N02)6.6KN02 -f- nHaO, the double salt, separates as a yellow crystalline powder. This reaction is very characteristic for cobalt, and serves to separate it from nickel. Ammonio-Cobalt Compounds.—Cobalt is capable of forming a series of peculiar compounds with ammonia, in which the metal appears inks highest state of oxidation ; the structure of these derivatives has not yet been explained. On adding ammonium hydroxide to a cobaltous chloride solution, the precipitate first lormed dissolves in the excess of the reagent, and when this liquid is permitted to stand exposed to the air, the color, which is brown at first, gradually passes into red. On adding concentrated hydrochloric acid to this solution, a brick-red, crystalline powder, of the composition—CojClj.ioNHj zHjO—called roseoco- 408 INORGANIC CHEMISTRY. baltic chloride—is precipitated. If, however, the red solution be boiled with hydrochloric acid, a red powder—■purpureocobaltic chloride, Co^C\^.lomi^—sepa¬ rates out. If the ammoniacal red solution contain much ammonium chloride, hydrochloric acid will precipitate a yellowish-brown compound — luteocobaltic chloride, CojClg.iaNHj, The other salts of cobalt, such as the sulphate and nitrate, yield similar com¬ pounds, e.g., Co2(N03)g.loNH3, roseocobaliic nitrate. Cyanogen Cobalt Compounds.—In solutions of cobaltous salts, potassium cyanide produces a bright brown precipitate of cobalto-cyanide Co(CN)2, soluble in an excess of the reagent. The solution absorbs oxygen from the air, and is converted mio potassium cobalticyanide, KgCo2(CN)j2, corresponding to potassium ferricyanide. When the solution is evaporated the cobalticyanide crystallizes in colorless rhombic prisms, very soluble in water. Sulphuric acid precipitates hydrogen cobalticyanide, HgCo2(CN)j2, from the concentrated solution. This acid crystallizes in needles. 3. NICKEL. Ni = 58.6. Nickel exists in native condition in meteorites ; its most impor¬ tant ores are Niccolite—NiAs—and Gersdorffite, NiSj.NiAsa (con¬ stituted like cobaltite). It is always accompanied in its ores by cobalt, and vice versâ, cobalt usually by nickel. The isolation of the latter from its ores and from speiss cobalt (p. 407) is very complicated. Nickel usually appears in commerce in cubical forms, which in addition to the chief ingredient always contain some cop¬ per, bismuth, and other metals. Chemically pure nickel is pro¬ cured by igniting the oxalate or carbonate in a current of hydrogen. Nickel is almost silver-white in color and is very lustrous, and very tenacious. Its specific gravity is 9.1, and that of the fused variety 8.8. It fuses at a somewhat lower temperature than iron, and like it is attracted by the magnet. It is not altered in the air \ it dis¬ solves with difficulty in hydrochloric and sulphuric acids, but readily in nitric acid. Nickel, like iron, also combines with carbon monoxide, yielding the liquid com¬ pound Ni(CO)4, boiling at 43°, and solidifying at —25° to needle-shaped crystals. At 60° the vapor of the liquid explodes with violence {Chem. Jour., 749). Nickel can be perfectly separated from cobalt by means of this derivative. Its derivatives are almost exclusively of the ous foftn NiXj ; nickelic oxide behaves like a peroxide, and does not afford corre¬ sponding salts. Nickelous Hydroxide—Ni(0H)2—is a bright-green precipi¬ tate produced by alkalies in nickelous solutions. It dissolves in ammonium hydroxide, with a blue color. When heated it passes into gray nickelous oxide, NiO. Nickelous Chloride—NiCla -{- 6H2O—consists of green, monociinic prisms. When heated they lose water and become yellow. PLATINUM METALS. 409 Nickelous Cyanide—Ni(CN)2—is precipitated by potassium cyanide as a green-colored mass from nickel solutions. It is solu¬ ble in excess of the precipitant. The double cyanide, NiCyj.aKCy -f H2O, crystallizes from the solution. This salt is readily decom¬ posed by acids. Cyanogen compounds of nickel, constituted like those of iron and cobalt, are not known. Nickelous Sulphate—NÍSO4 -{- 7H2O—appears in green, rhombic prisms, isomorphous with the sulphates of the magnesium group, and forms analogous double salts. Nickelous Sulphide—NiS—is precipitated, black in color, by alkaline sulphides from nickel solutions. Nickelic Oxide—NÍ2O3—and Hydroxide—NÍ2(OH)6—are perfectly similar to the corresponding cobalt salts ; when warmed with hydrochloric acid they liberate chlorine. Nickel is used for certain alloys. Argentan consists, ordinarily, of 50 per cent, copper, 25 per cent, nickel and 25 per cent. zinc. The German nickel coins consist of 75 per cent. Cu and 25 per cent. Ni. The alloy will be whiter in color and harder, and also receive a higher polish in proportion to the amount of nickel that it contains. At present, cast-iron ware is coated with a layer of nickel to prevent it from rusting and to impart to it a beautiful white surface. This is accomplished in an electrolytic manner, or by boiling the iron ware in a solution of zinc chloride and nickel sulphate. In the electrolytic method the solution of the double sulphate of nickel and ammonium is employed ; the positive electrode consists of a pure nickel plate, while the object to be coated is attached to the negative electrode. GROUP OF THE PLATINUM METALS. Besides platinum, this group comprises palladium, rhodium, ruthenium, osmium and iridium—the constant companions of the first in its ores. On page 395 we observed that these metals are divided into two groups ; the group of light platinum metals^ and the group of heavy platinum metals which have higher atomic weights and specific gravities :— Ru, 101.7 Rh, 103 Pd, 106 Os, 192 Ir, 193.2 Pt, 194.8 Sp.gr. " 12.26" 12.1 " 11.8 " 22.4 " 22.38 " 21.4 The relations of the metals of this group to each other are per¬ fectly similar to those of the iron group ; and they show in their physical and chemical properties a great resemblance to the corre¬ sponding members of the iron group. Osmium and ruthenium, like iron, have a gray color, fuse with difficulty and are readily 35 410 INORGANtC CHEMISTRY. oxidized in the air. Palladium and platinum, on the other hand, have an almost silver-white color like nickel, are more fusible, and are not oxidized by oxygen. In chemical respects osmium and ruthenium, like iron, also show a metalloidal nature, inasmuch as their highest oxygen compounds form acids. Their derivatives show a complete parallelism with those of iron :— II HI IV VI OsO OSjOg OsOj (OSO3) Osmous Osmic Osmium Osmic oxide. oxide. dioxide. trioxide. RuO RUjOg RuOj (RuOg) Ruthenous Ruthenic Ruthenium Ruthenium oxide. oxide. dioxide. trioxide. The acid oxides OsOs and RuOs are unknown, but the corre¬ sponding acids, HgOsOi (osmic acid) and HaRuO^ (ruthenic acid), and their salts have been obtained. Besides the derivatives already mentioned we find that osmium and ruthenium are capable of still higher oxidation, yielding OsO^, per-osmic oxide, and RuO^, per- ruthenic oxide—which is not the case with iron ; in these com¬ pounds the metals appear to be octads, yet these oxides do not afford corresponding acids or salts. Rhodium and iridium, like cobalt, do not yield acid-like deriva¬ tives. Their salts correspond to the forms:— II III IV RhO RhPa RhOj. Rhodous Rhodic Rhodium oxide. oxide. dioxide. The rhodic compounds are the more stable. Palladium and platinum, finally, are relatively of more basic nature, as their ous derivatives, PdXj and PtXj, are proportionally more stable than the ü forms, PdX^ and PtX^. Palladium also forms a lower oxide, palladium suboxide, PdjO, in which it ap¬ proaches silver. The platinum metals are found in nature almost exclusively in the so-called platinum ore, which usually occurs in small metallic grains in accumulated sands of a few regions (in California, Aus¬ tralia, the Island of Sumatra, and especially in the Urals). The platinum ore, like that of gold, is obtained by the elutriation of the platiniferous sand with water whereby the lighter particles are car¬ ried away. Platinum ofe usually contains 50-80 per cent, platinum, besides palladium (to 2 per cent,), iridium (to 7 percent.), osmium (i/^ per cent.), and ruthenium (i}4 per cent.), and different other metals, as gold, copper, and iron. RUTHENIUM AND OSMIUM. 411 The separation of the platinum metals is generally executed in the following manner : The gold is first removed by dilute aqua regia. Then the ore is treated with concentrated aqua regia, when platinum, palladium, rhodium, ruthenium, and a portion of iridium are dissolved. Metallic grains or leafiets, an alloy of osmium and iridium—platinum residues—remain. Ammonium chloride is then added to the solution and platinum and iridium precipitated as double salts. When the precipitate is ignited, a spongy mass of iridium-bearing platinum (platinum sponge) is obtained, which is applied directly in the manufacture of platinum vessels. The filtered solution from the insoluble chlorides contains palladium, rhodium, and ruthenium, which are thrown down as a metallic powder by iron ; their further separation is then effected in various ways. Formerly spongy platinum was employed almost exclusively for the manufacture of platinum objects ; it was pressed into moulds, then ignited and hammered out. Now the fusibility of Pt in the oxy-hydrogen flame is resorted to, and the fused metal run into moulds. Platinum containing both iridium and rhodium may be fused directly out of the platinum-ore by means of the oxy-hydrogen blow¬ pipe. The greater portion of the osmium and ruthenium is con¬ sumed in this operation. The presence of iridium and rhodium makes platinum harder and less readily attacked by many reagents. RUTHENIUM AND OSMIUM. Ru = 101.7. Os — 192. Ruthenium has a steel-gray color ; it is very hard, brittle, and difficultly fusible (about 1800°). When pulverized aud ignited in the air it oxidizes to RuO and RujOj. It is insoluble in acids, and only slowly dissolved by aqua regia. When fused with potassium hydroxide and nitrate, it forms potassium rutheniate, KjRuO,. Ruthenium heated in chlorine gas yields ruthenium aichloride^ RuClj, a black powder, insoluble in acids. The sesquichloride, Ru^Clj, is obtained by the solu¬ tion of Ru2(0H)g in hydrochloric acid, and is a yellow, crystalline mass, which deliquesces in the air. It yields crystalline double chlorides with potassium and ammonium chlorides, e.g.y Ru2Cle.4KCl. The tetrachloride, RuCl^, is only known in double salts. Ruthenious oxide, RuO, the sesquioxide, RUjG,, and dioxide, RuOj, are black powders, insoluble in acids, and are obtained when ruthenium is roasted in the air. The hydroxides, Ru2(0H), and Ru(OH)^, are produced by the action of the alkalies upon the corresponding chlorides, and are very readily soluble in acids. Ruthenic acid, HjRuO^, is not known in a free condition. Its potassium salt, 412 INORGANIC CHEMISTRY. KgRuO^, is formed by fusing the metal with potassium hydroxide and nitre. It dissolves in water with an orange-yellow color. When chlorine is conducted through the solution ruthenium tetroxide, RuO^, separates as a gold-yellow crys¬ talline mass. It fuses at 40°, boils about 100°, and )delds a yellow vapor, the odor of which is similar to nitrogen dioxide, NOj. It decomposes with explosion at 108®. Water breaks it up with formation of RujiOHlg. It dissolves to RuO^Kj in concentrated potassium hydroxide. When less chlorine is introduced into the solution of RuO^Kg, greenish-black crystals separate out, which are isomorphous with potassium permanganate, and appoar to be RuO^K. Osmium is very much like the preceding. It is not even fusible in the oxyhy- drogen flame ; it only sinters together. According to Violle it fuses at 2500®. Reduced to a fine powder it will burn when ignited in the air to OsO^. Nitric acid and aqua regia convert it into the same oxide. The compounds, OsQj and OsO, OSjClg and OsjOjjOsOg and OsCl^, are very similar to the corresponding ruthenium derivatives. By fusion with potassium hydroxide and nitre we get potassium osmate—KjOsO^—which crystallizes from aqueous solution with 2H,0 in dark-violet octahedra. The most stable and a very characteristic derivative of osmium is the tetroxide, OsO^, which is produced by igniting the metal in the air, or by the action of chlorine on osmium in the presence of water. It crystallizes in large colorless prisms, which fuse below 100® and distil at a somewhat higher temperature. It has a very sharp, piercing odor, similar to that of sulphur chlo¬ ride. Reducing and organic substances precipitate pulverulent osmium from it. This is the basis of its application in microscopy. OsO^ and RuO^ do not afford corresponding salts. RHODIUM AND IRIDIUM. Rh = 103. Ir = 193.2. These metals are lighter in color and are more easily fusible than ruthenium and osmium, (Iridium fuses at 1950®.) When pure they are not attacked by acids or aqua regia ; but dissolve in the latter when alloyed with platinum. Rhodium forms three oxides: RhO, RhjOg and RhOj, of which the second forms salts with acids. RhOg results when rhodium is heated with nitre. Of the chlorides only RhjClg is known. It results when the metal is heated in chlorine gas. It is a brownish -red mass. It forms readily crystallizing, red-col¬ ored double salts with alkaline chlorides. Iridium has perfectly analogous derivatives : IrO, Ir^Og, IrOj, and IrC3„ Ir,Clg, IrCl^. The sesquichloride, IrjClg, formed by heating Ir in chlorine, is an olive- green, crystalline mass, insoluble in water and acids. It affords double salts with the alkaline chlorides, e.g., IrjClg.óKCl + óHgO, which cr3rstallizes from water in green crystals. They are also produced by the action of SO, upwjn the double salts of IrCl^. Iridium Tetrachloride—^IrCl^—^is produced in the solution of iridium or its oxide in aqua regia, and remains, on evaporation, as a black mass, readily soluble in water (with red color). When alkaline chlorides are added to the solution double chlorides are precipitated, e. g., IrCl^.aNH^Cl, isomorphous with the dou¬ ble chlorides of platinum. When a solution of IrCl^ is boiled with KOH, Ir(OH)g will be precipitated. PALLADIUM. 413 PALLADIUM. Pd == 106. Palladium, in addition to occurring in platinum ores, is found alloyed with gold (Brazil), and in some selenium ores (Hartz) ; it has a silver-white color, and is somewhat more fusible (about 1500°) than platinum. When finely divided it dissolves in boiling concentrated hydrochloric, sulphuric, and nitric acids. When ignited in the air it at first becomes dull by oxidation, but at a higher temperature the surface again assumes a metallic appearance. Palladium absorbs hydrogen gas (occlusion) to a much greater extent than platinum or silver. Freshly ignited palladium leaf absorbs upwards of 370 volumes of hydrogen at ordinary tempera¬ tures, and about 650 volumes at 90-100° C. A greater absorption may be effected at ordinary temperatures in the following man¬ ner :— Water is decomposed by the electric current, palladium foil being used as a negative electrode. The liberated hydrogen is then taken up by the palladium (to 960 volumes) ; the metal expands (j^ its volume), becomes specifically lighter, but retains its metallic ap¬ pearance entire. According to the investigations of Debray, the compound PdgH is produced, which contains dissolved hydrogen, and deports itself similarly to an alloy (compare p. 47). Palladium charged with hydrogen usually remains unaltered in the air, and in a vacuum ; it, however, sometimes becomes heated in the air, as the hydrogen is oxidized to water. The same occurs when palla¬ dium hydride is heated to 100° ; in vacuo^ all the hydrogen escapes as gas. Palladium hydride is a strong reducing agent, like nascent hydrogen. Ferric salts are reduced to the ferrous state ; chlorine and iodine in aqueous solution are converted into hydrochloric and hydriodic acids. Palladium black absorbs hydrogen more energetically than the compact variety (at 100° upwards of 980 volumes). This substance is obtained by the reduction or electrolysis of palladic chloride. If palladium sponge be heated in the air until the white metallic color becomes black, in consequence of the superficial oxidation, it will absorb hydrogen very energetically at ordinary temperatures, and partially oxidize it to water. When palladium-sheet or sponge is introduced into the flame of a spirit-lamp, it is covered with smoke ; this is due to the fact that the metal withdraws the hydrogen of the hydrocarbons, and carbon is set free. There are two series of palladium compounds : the palladiouSy PdXa, and palladic^ PdX4. The first are well characterized and are distinguished by their stability. 414 INORGANIC CHEMISTRY. Palladious Chloride—PdClj—remains as a brown, deliques¬ cent mass, on evaporating the solution of palladium in aqua regia. It yields easily soluble crystalline double salts, with alkaline chlo¬ rides, e.g.^ PdCla.aKCl. Palladious Iodide—Pdlj—is precipitated from palladium solu¬ tions by potassium iodide as a black mass, insoluble in water. Palladious Oxide—PdO—is a black residue left upon careful ignition of the nitrate. It is difficultly soluble in acids. When heated, it loses oxygen, and forms palladium suboxide^ PdjO. When palladium is dissolved in sulphuric or nitric acids, the corresponding salts are produced. The sulphate, PdSO* + 2H2O, is composed of brown crystals, readily soluble in water. Much of the latter decomposes it. Pailadic Chloride—PdC^— is formed when the metal is dis¬ solved in aqua regia. It decomposes, on evaporation, into PdCl, and Clj. When potassium or ammonium chloride is added to its solution, red-colored, difficultly soluble double chlorides crystallize out ; they are analogous to the corresponding salts of platinum. PLATINUM. Pt = 194.8. The separation of platinum from the ore was described on pagfe 411. The metal has a grayish-white color, and a specific gravity of 21.4. It is very tough and malleable, and may be drawn out into very fine wire and rolled into foil. It becomes soft without melting at an intense heat. It fuses in the oxy-hydrogen flame (about 1770°—Violle), and is somewhat volatile. On fusion, it absorbs oxygen, which it gives up again on cooling (like silver). At ordinary temperatures, it also condenses oxygen upon its surface, especially when in a finely divided state, as platinum black or sponge. The first is obtained, if reducing substances, like zinc, be added to solutions of platinic chloride or upon boiling the solution with sugar and sodium carbonate; it absorbs as much as 250 volumes of oxygen. Platinum sponge is obtained by the ignition of PtCV 2NH4CI. The production of various reactions is due to this power of platinum to condense oxygen ; thus hydrogen will inflame in the air, if it be conducted upon platinum sponge; sulphur dioxide combines with O at 100° to form the trioxide. At a red heat pla¬ tinum permits free passage to hydrogen, while it is not permeable by oxygen and other gases (p. 265). PLATINUM. 415 Platinum is not attacked by acids ; it is only soluble in liquids generating free chlorine, e.g., aqua regia. In consequence of this resistance to acids, and its unalterability upon ignition, this metal answers as an undecomposable material for the production of chemical crucibles, dishes, wire, etc. The usual presence of iridium in ordinary platinum increases its durability. The alkaline hydroxides, sulphides, and cyanides attack it strongly at a red heat. It forms readily fusible alloys with phos¬ phorus, arsenic, and many heavy metals, especially lead, and many heavy metals are reduced from their salts by platinum. Therefore such substances must not be ignited in platinum crucibles, etc. Platinum, like palladium, affords platinous^ PtXj, and platinicy PtX4, derivatives ; in the first it is more basic, in the latter more acidic. Platinic Chloride—PtCl4—is obtained by the solution of pla¬ tinum in aqua regia, and when the liquid is evaporated with an excess of hydrochloric acid, hydrogen chloroplatinate, PtCh- 2HCI = HîPtClg, crystallizes with six molecules of water in brownish red, deliquescent prisms. It forms characteristic double chlorides, PtCli.zKCl, with ammonium and potassium chloride. These are difi&cultly soluble in water; hence, on mixing the solutions, they immediately separate out as a crystalline yellow powder. Ignition completely decomposes the ammonium salt, leaving spongy platinum. Platinum chloride affords similar insoluble double chlorides with those of rubidium, caesium, and thallium, while that with sodium, PtCleNaj + 6H2O, is very soluble in water. On adding NaOH to platinic chloride and then supersaturating with acetic acid, there separates a reddish-brown precipitate of plat¬ inic hydroxide, Pt(0H)4. This dissolves readily in acids (excepting acetic), with formation of salts. The oxygen salts, as Pt(S04)j, are very unstable. The hydroxide has also an acidic character (^platinic acid^, and dissolves in alkalies, yielding salts with them. These, also, result on fusing platinum with potassium and sodium hydroxide. The barium salt, Pt | -f 3H2O, is precipitated from platinic chloride, by barium hydroxide, as a yellow, crystalline compound. The acidic nature of its hydroxide allies platinum to gold. If hydrogen sulphide be conducted through platinic solu¬ tions, black platinum disulphide, PtS2, is precipitated ; it is soluble in alkaline sulphides, with formation of sulpho-salts. Platinous Chloride—PtCl,—is a green powder, insoluble in water, remaining after heating PtCl^ to 200°. It affords double salts with alkaline chlorides, e.g., PtClj.2NaCl. When digested with potassium hydroxide it yields the hydroxide, Pt(OH),. 416 INORGANIC CHEMISTRY. Cyanogen Compounds.—Like iron and cobalt, platinum af¬ fords double cyanides corresponding to the ferrocyanides. When platinous chloride is dissolved in potassium cyanide platinum potas¬ sium cyanide^ KjPtCyi -f- 3H2O, crystallizes on evaporation in large prisms exhibiting magnificent dichroism ; in transmitted light they are yellow and in refiected light blue. This salt must be viewed as the potassium compound of hydro platino-cyanic acid, HjPtCyi. When separated from its salts it crystallizes in gold-yellow needles. Its salts with heavy metals are obtained by double decomposition, and all show a beautiful play of colors. Platinum-Ammonium Compounds.—There is a whole series of these, which must be viewed as platinum bases and their salts. They are constituted according to the following empirical formulas :— Pt(NH3)2X2, Pt(NH3)3X„ Pt(NH3),X3, Pt(NH3),X„ in which X indicates various acid residues, or halogen atoms. They arise by the action of ammonium hydroxide upon platinous chloride. The bases are obtained by substituting hydroxyl groups for the acid radicals, e.g., Pt(NH3)4(OH)2. They resemble alkaline hydroxides in their chemical properties. The other platinum metals afford similar amine derivatives. The nature and chemical constitution of these interesting compounds is, however, not fully explained. SPECTRUM ANALYSIS.* We observed that various substances, when introduced into a non-luminous ñame, imparted to it a characteristic coloration. Thus, the sodium compounds color it yellow -, the potassium, vio¬ let ; thallium, green ; etc. The decomposition of the light thus obtained, and, indeed, of every light, by means of a prism and the study of the resulting spectrum form the basis of spectrum analy¬ sis, established in 1859 by Kirchhof! and Bunsen. Its important applications and universal use constitute one of the greatest of scientific achievements of all ages. As we well know, every substance, solid or liquid, heated to white heat (e. g., molten platinum ; lime heated in the oxyhydrogen flame ; the ordinary flame containing glowing particles of carbon ;) * A more exhaustive, concise, and distinct presentation of the spectrum phenom¬ ena may be found in Herman W. Vogel's " Practische Spectralanalyse irdischer Stoffe," 1888. SPECTRUM ANALYSIS. 417 emits rays of every refrangibility ; and hence furnishes a continuous spectrum^ which brings to view all the colors of the rainbow, from red to violet, if the light be conducted through a prism. Glowing gases and vapors, on the contrary, whose molecules can execute unobstructed oscillations, emit light of definite refrangibility, and, therefore, afford spectra, consisting of single, bright lines. Thus, the spectrum of the yellow sodium flame is recognized as composed of one very bright yellow line, which by increased magnifying power is shown to consist of two lines lying very near each' other. This reaction is so very delicate, that of a milligram of sodium may be detected by it. The violet potassium light affords Fig. 89. a spectrum, consisting of a red and a blue line. The crimson strontium light shows in the spectrum several distinct red lines and a blue line. (See the spectrum plate.) Each of these lines corre¬ sponds to a definite relative position in the spectrum. If substances affording different colors be introduced into a flame, the most intense color generally obscures the others; in the spec¬ trum, however, each individual substance shows its peculiar bright lines, which appear simultaneously or succeed each other, accord¬ ing to the volatility of the various substances. The spectrum apparatus or spectroscope, figured in Fig. 89, serves for the observation of the spectra. 418 INORGANIC CHEMISTRY. In the middle of the apparatus is a flint-glass prism P. At the further end of the tube A'v&q. movable vertical slit, in front of which is placed the light to be investigated. The entering light rays are directed by a collecting lens into the tube A, upon the prism, and the refracted rays (the spectrum) are observed by the telescope B. The tube C is employed to ascertain the relative position of the spectrum lines. This is provided at its outer end with a transparent horizontal scale. When a luminous flame is placed before the scale its divisions are reflected from the prism surface and thrown into the telescope B. We then see the spectrum to be studied and the scale divisions in B at the same time, and can readily determine the relative position of the lines of the spectrum. To study two spectra at the same time, and compare them, a three-sided, right- angled glass prism is attached in front, of one-half (the lower or upper) of the slit of the tube A ; this directs the rays of a light placed at the side (p, Fig. 89) through A upon the prism P. By means of B, two horizontal spectra will be observed, one above the other, and between are the bright divisions of the scale. Adjustment of the Spectroscope.—^To observe the spectra in the apparatus de¬ scribed, it is necessary to first adjust the same correctly. The tube A contains, besides the slit, also a lens (Collimator Lens), which serves to render parallel the bunch of rays proceeding from the slit; hence, the latter must be accurately placed in the focus of the lens. This is best accomplished as follows : The tele¬ scope (.5) is drawn out and adjusted for some distant object, that it may be adapted for the reception of parallel rays; it is then replaced in the stand, pointed toward the slit, illuminated by a sodium chloride flame, and the slit then removed far enough to appear perfectly distinct in the telescope. To have the spectrum lines sharply defined, the slit must be made quite narrow; for feebly luminous lines it must, however, be widened. The horizontal black lines that appear in the spectrum arise from dust particles adhering to the slit. The proper position of the tube with the slit with respect to the prism is usually fixed by the frame on which they rest. It must be so adjusted that the refracted rays pass through the prism as symmetrically as possible i.e., in the minimum of their deviation, otherwise the spectrum will be less distinct and (owing to unequal refraction) will appear distorted. The symmetrical passî^e of the spec¬ trum rays is approximately attained when the intermediate green rays pass through the prism symmetrically. Such a position must, therefore, be given the prism, with reference to the tube canying the sht, that the middle green rays (Ime E of the sun spectrum) pass through it in the minimum of their deviation. It is then only necessary to so arrange the telescope that the green rays lie in the middle of the field of vision. The determination of the position of the lines of the spectra is usually effected by means of a scale (see above). This arrangement is such (according to Bunsen) that the yellow sodium line coincides with the line 50 of the scale ; then the red potassium line (a) will lie at 17, and the violet (ß) at 152 (apparatus of Desaga). Since the refraction and dispersion of the rays are influenced by the quality of the glass of the prism, the scale indications of different forms of apparatus are not directly comparable. They must, therefore, be referred to an absolute measure. SPECTRUM ANALYSIS. 419 This is most conveniently attained by reduction to the sun spectrum, which may be rendered visible in the telescope at the same time by means of a comparison- prism attached to the tube carrying the slit. We then determine the dark lines of the sun with which the lines of the flame under investigation coincide. In accu¬ rate determinations, the spectrum lines are represented in wave lengths, according to the millionth of a millimeter. The apparatus described above is that usually employed in chemical laborato¬ ries. There are other forms adapted for special purposes—^for investigation under the microscope, for the observation of the sun and stars. In accurate observations, where clear, broad spectra are required, the light is permitted to pass through several (3-9) prisms, consisting of hollow glass filled with carbon disulphide, which refracts light very strongly. Diffraction spectra may be advantageously applied in many instances in the place of the prism spectra. The direct line {à vision directe) spectroscopes are very excellent for labora¬ tory purposes. With these the spjectra may be viewed in the direction in which the luminous objects really are; there are no deflections. This is accomplished by a combination of several prisms of crown and flint glass, whereby dispersion is attained, together with the simultaneous removal of deflection. To observe the spectra of metals, it is only necessary, in many instances, e. g., with the alkalies and alkaline earths, to introduce their volatile salts into a non-luminous alcohol or gas flame. A reduction of the metal usually takes place, and the spectra of the free metals themselves are obtained ; thus, for example, sodium chloride is decomposed in the flame first into HCl and NaOH, which is then reduced by the carbon of the glowing gases to metal¬ lic sodium, and colors the flame yellow. The compounds diflicult to decompose (as the barium salts), often afford independent spectra, differing from those of the free metal ; this is plainly recognizable in the copper compounds. Most metals, however, require a much higher temperature than that of the gas flame for their conversion into gases. To vaporize them and observe their spectra, the electric spark is made to pass from electrodes constructed of them. In this manner we can study all metals, even the most non-volatile, like gold, iron, and platinum. Their spectra are generally quite complicated, and exhibit a great number of single lines. Thus, over 450 lines have been established for iron. Instead of making the electrodes of the metals we wish to study, we can, according to Bunsen (Poggend, Ann., 155), employ carbon points saturated with solutions of the metallic salts. A Ruhmkorf induction apparatus, with a dip battery of 4 elements, and an attached Leyden jar, will suffice to produce the electric spark. Such spark spectra are often quite different from the flame spectra obtained in gas flame. A peculiar, very interesting and practical procedure for the production of spark spectra is due to Lecoq de Boisbaudran (Spectres lumineux, Paris, 1874). He allows the induction sparks to strike into the solution of the salt that is being examined. This is placed in a smafl reagent titibe, in the bottom of which is fused 420 INORGANIC CHEMISTRY. a platinum wire. Above the surface of the liquid is the second electrode, a plati¬ num wire connected with the positive pole of the induction spiral. The spectra of all the metals may be easily obtained by this contrivance, and, indeed, Lecoq de Boisbaudran discovered gallium by it. The spectra of the elementary gases may be obtained by passing electric sparks through them ; these will be variously colored. Hy¬ drogen shines with a red light, which is decomposed and gives a spectrum consisting of a bright red, a blue, and a green line. Ni¬ trogen shines with a violet light and affords a spectrum of many lines, chief of which are the violet. The spectra of gases may be more conveniently observed by aid of Geissler's tubes, which are filled with very dilute gases and the induction stream then passed through them. These methods give us a means of readily distinguishing the individual chemical elements, and even detecting them in traces. Since the year i860 various new elements, e. g.j caesium, rubidium, thallium, indium, scandium, gallium, and several others, not accu¬ rately studied, have been discovered by their aid. In addition to the direct, bright spectra just described there are yet ¿fark absorption spectra. If a white light giving an uninter¬ rupted spectrum be allowed to pass through different transparent bodies, the latter will absorb rays of definite refrangibility, allow¬ ing all others to pass. Therefore we observe the sun spectrum interrupted by dark lines or bands in the spectroscope. The solu¬ tions of didymium and erbium absorb certain rays, and show cor¬ responding dark lines in the spectrum. The gases deport them¬ selves similarly. White light that has traversed a broad, layer of air shows in the spectrum several dark lines peculiar to nitrogen, oxygen, and steam. This power of absorption is possessed to a much higher degree by all incandescent gases or vapors. If a white light, like the Drummond calcium light, be conducted through the yellow sodium ñame (through glowing sodium vapors), a dark line will appear in the rainbow spectrum of the white light, and its position will correspond exactly to that of the yellow sodium line ; the latter thus appears converted into a dark line. If white light be passed through the potassium ñame, two dark lines will be visible in the spectrum, corresponding to the red and blue lines of the potassium spectrum. Such spectra are designated the inverted spectra of the corresponding metals. The inverted spectra of all elements may be obtained in this way, and they correspond accurately to the direct bright spectra. The cause of these phenomena lies in the proposition deduced by Kirchhoff from the undulatory theory of light, that the ratio between the emissive and absorptive fower is the same for almost all bodies at like temperatures. According to this. SPECTRUM ANALYSIS. 421 incandescent gases only absorb rays of just the same refrangibility as those which they emit. For example, when bright white light is passed through the yellow sodium flame the yellow rays of the former are absorbed and retained, while all others pass on almost entirely unaltered. Therefore, in the rainbow spectrum of white light the yellow rays of definite refrangibility will be absent ; and if the other refracted rays of the white light are brighter than the yellow rays emitted from the sodium flame, the latter will be rela¬ tively darker ; a dark line will therefore make its appearance. These phenomena have presented a new and wide province to spectral analysis, inasmuch as they open up avenues for the inves¬ tigation of the chemical nature of the sun and other bodies. It is known that the bright rainbow sun spectrum is intersected by a number of dark lines which have been called the Frauenhofer lines from their discoverer. Kirchhoff has shown that these lines can be easily accounted for, after what has already been said, by the following hypothesis upon the nature of the sun : The latter consists of a solid or liquid luminous nucleus, surrounded by an atmosphere of incandescent gases and vapors. Then the uninter¬ rupted spectrum of the glowing nucleus must be intersected by the dark lines of the inverted spectra of those gases and vapors which occur in the sun's atmosphere. An accurate comparison of the Frauenhofer lines with the spectrum lines of the various elements, has revealed the fact that iron, sodium, magnesium, calcium, chro¬ mium, nickel, barium, copper, zinc, and hydrogen are present in the sun's atmosphere. Thus dark lines have been found in the sun's spectrum corresponding to all of the 450 lines of the iron spectrum. The inferences upon the chemical constitution of the sun possess as much and, indeed, a higher degree of probability than falls to many other deductions. The investigation of the sun's spectrum has cleared up many other changes occurring there, and this has led to a complete sun meteorology. All the fixed stars thus far investigated possess a constitution like that of the sun. They give spectra intersected by dark lines, and therefore consist of incandescent nuclei surrounded by gaseous atmospheres. The spectra of nebulae, however, only show bright lines ; hence these consist of uncondensed, incandes¬ cent masses of vapor. Periodicity of the Spectrum Lines.—Since all other properties of the ele¬ ments and their compounds have proved themselves to be periodic functions of the atomic weights, we may expect the same with reference to the spectrum phenomena. But few such regularities have occurred thus far. Most of the ele¬ ments, especially the metalloids and the difEcultly volatile metals, afford very 422 INORGANIC CHEMISTRY. complex spectra, which often vary at the same time with the temperature, so that spectra of ist, 2d, etc., order are distinguished for some of them. It appears that not all spectral lines are of the same importance, since it has been possible, in some instances, to refer the various lines of a spectrum to particular fundamental lines whose relations to each other are comparable with those existing between the harmonics and the primary tones. Thus the four lines of the hydrogen spectrum may be looked up>on as the harmonics of a single wave. Therefore, only individual lines are to be regarded in the comparison of spectra. This is clearly observed with the easily volatile metals belonging to the homol(^ous groups, K, Rb, Cs; Ca, Sr, Ba; Ga, In, Tl; whose lines, lying in the violet px>rtion of the spectrum, advance more towards the red end as the atomic weights increase; the wave lengths (in millionths of a millimeter) become successively greater with the increase of the latter or that of the atomic volumes :— K 39 Wave length, 404 Rb 83 " " 420.421 Cs 132 " « 456-459 Ca 40 Wave length, 422 Sr 87 " « 461 Ba 137 « •« 525-550 Ga 69 Wave length, 403.417 In 113 " " 410.450 Tl 209 " « 535. These relations are made apparent by the arrangement of the spectra in the sp)ectrum plate. A similar shifting of the spectral lines is observed with heterologous elements belonging to the same periods : K Ca, Rb Sr, Ba Cs, so that it seems possible to draw a conclusion upon the spectra of the succeeding elements. Indeed, the ele¬ ment scandium (45) succeeding calcium, shows intense violet lines of the wave length 425-440. eo CN Heat of Formation of the Most Important Compounds of the Metals {according to J. Thomsen). In usual state of aggregation (columns a), and in dilute aqueous solution (columns b). b a 101.2 KBr 95-3 96-5 NaBr 85-7 102.2 LiBr 196.3 BaBr, 169.4 195-3 SrBr, 157-7 187.6 CaBr, 141 3 186.9 MgBr, 128.0 MnBr. ... 237-7 AlBr, 129.6 112.8 ZnBr, 75-9 96.2 CdBr, 73-9 99-9 FeBr, 94-8 CoBr, ... 93 7 NiBr, • •• 127.7 FeBr, • •• 81.1 SnBr, • •• .39-3 TlBr 41.4 157-0 HBr 38-4 PbBr, 64-4 75-9 HgBr .34-1 .59-8 HgBr, .50-5 CuBr 249 ... CuBr, 32-5 o\ 10 SbBr, ... AgBr TlBr, 27 7 ... As Br, 47-1 ... HBr 8.4 89.0 AuBr —0.1 Au Br, 8.8 ... PtBr, 27.2 KCl NaCl LiCl BaClj SrClj CaCl, MgClj MnCl, AlCl, ZnCU CdClg FeCU CoClj NiCl, Feci, SnCl, HCl SnCL TlCl PbCU HgCr HgCl, CuCl CuCl, SbCl, BiCl, AgCI AsCl, TiCl, HOI AuCl AuCl, PtCl, 105.6 97.6 93.8 194.2 184-5 170.2 151-0 III.9 160.9 97-2 93-2 82.0 76.4 74.0 96.0 8q.8 127.2 48-5 82.7 41.2 63.1 32 8 5» 6 91.4 90.6 29-3 71-4 22.0 5-8 22.8 90, 85 91, 174. 173- 165. 165. 105, 204. 90. 74- 78. 72. 71- 28.3 54-4 20.4 56-1 5-0 KI Nal Li I Bal, Sri- Caí, Mgf, Mnl, All, Znl, Cdl, Fel, Col, Nil, Sni, Til Pbl, Hgf, Hgl Cul Cul, Agi Hi Til, Asl, Aul HI PtI, 8o.i 69.1 107.6 70-3 49.2 44-9 30.1 39-6 34-3 24.2 16.2 13-8 12.6 —5-5 —6.0 75-0 7°-3 76.1 144.0 143-3 135 3 134-6 75-7 159-3 60.5 43-9 47.6 42-5 41.4 Ï3-I 10.5 KOH NaOH LiOH BaO,H, SrOaH, CaOjH, MgO-rf, AIO3H3 MnO, H,0 ZnO, HjO HjO (liquid) FeO, H,0 FcjOo. sHjO CdO, H,0 CoO, HoO NiO, H,0 SnO, HjO SnO,, H,0 H (vapor) ASgC, T1,0 HgO CuO Cu,0 T1,03. 3H,0 PdO, H,0 Ag,0 AuoO,, sHjO 103.9 102.0 216.3 216.4 214.7 217.2 296.9 94-7 82.6 68, 68.2 191.1 65 6 63-4 60.8 68.1 133-5 58. 50-3 154-6 42.2 30-6 42.2 37-1 40.8 86.0 22.7 5-9 —13.2 116.4 III.8 117.4 226.6 226.1 217.8 147-0 39-2 Li,O K,0 Na,0 BaO SrO CaO (99.1)* (100.2) 130-4 130.9 ^3»-7 166.5 164-5 155-2 158.2 157-7 149 KSH NaSH LiSH BaSjH, SrH,Sa CaS.li, MnS, nH,0 ZnS, nHjO CdS, nHjO FeS, nHjO CoS, nH,0 NiS, nH,0 TUS PbS CuS Cu,S Hg^ (64)t (55) 46-3 41-5 33-9 23-7 21.7 19-3 21.6 20.4 10.o 20.2 16.8 5-3 4-5 65 60 66 124 123 "5 114 K,S Na^S LlgS BaS SrS CaS MgS A1,S, 102 (88 (99-o) (99 2) (92-0) (79 6) (124-4) 113-2 103.9 115-2 107, 106-6 98-3 (110) K,SO. Na,^S(34 LijSO, BaSO, 6SrS04 CaSO, MgSO, AljSjCJxj MnSO. ZnS04 CdSO, FeSO, C0SO4 NÍSO4 T1,S04 HaSO« PbSO^ H,S04 CuSOi AgjSO^ 344-S 328.5 334-1 337-5 320.8 319-9 302.2 249.8 230.0 221.1 221.0 216.2 192.9 182.5 167.3 337 329 340. 323- 878 263 248 231 235 230 229. 212. 198 102. KNO, NaN(5, LiNO, BaN,ôs SrN-O, CaN.O, MgN,Oe MnN,p, .sZnNjO, 4CdN,0, 8 FeN,Oe K,CO, Na,CO, BaCO, SrCO, CaCO, MnCO, CdCO, PbCO- Ag,CO, 279-5 271.0 281.3 279.6 269.2 209.2 179.9 168.2 121.3 286. 276. 6CoN,Oa NiN,0, HNO, PbNlO, T1N(3, HNO, CuN.O, HgNb, HgN,Oe AgNÖ, 119.4 111.2 111.6 : 225.7 219.8 : 203.2 : ioS-5 58.1 41.5 28.7 * According to Beketoff. t According to Sabatier. I N D KX. A. Absorptiometer, 121 Acetylene, 156 Acids, 59, 39, 84 haloid, 60 meta-, 205 normal, 205 Ortho-, 205 oxygen, 60 radicals, 178 Active oxygen, 85 chlorine, 307 Afl&nity, chemical, 23 Air, composition of, 120, 121, 123, 125 destroyed, 119 fire, 119 Algaroth, 149 Alkali metals, 277 Alkalies, 277 Alkaline earths, 303 Allotrophy, 88 Alloys, 92, 265 Alums, 355 Alum, ammonium, 356 burnt, 356 cubical, 356 ordinary, 356 Aluminates, 353 Aluminium, 350 bronze, 351 chloride, 351 hydroxides, 353 oxide, 352 silicates, 357 sulphate, 355 Amalgams, 265 Ammonia, 125 chemical properties of, 128 physical properties of, 126, 127 quantitative composition of, 130 Ammonio-cobalt compounds, 407 Ammonium, 129 amalgam, 299 carbonate, 3cx> chloride, 300 compouncÊ, 129, 299 nitrate, 300 nitrite, 300 phosphates, 301 sulphate, 300 sulphides, 302 Ampere, law of, 77 Anglesite, 373 Anhydrite, 308 Anhydrides of acids, 177 Animal charcoal, 152 Anthracite, 152 Antimony, 147 acids, 227 bromides, 149 butter of, 149 chlorides, 149 iodides, 149 mirror, 148 oxides, 226, 228 sulphides, 228 Antimonyl, 227 Apatite, 308 Apparatus of Carré, 127 Marsh, 145 Aqua regia, 208 Argentan, 339 Argyrodite, 367 Arsenic, 143, 144 acids, 224, 225 bloom, 223 bromide, 147 chloride, 147 fluoride, 147 iodides, 147 Marsh's test for, 145 mirror, 145 pentoxide, 225 spots, 145 sulphides, 225 sulpho-salts, 226 trioxide, 223 Arsenious acid, 224 Arsine, 144, 146 Atmosphere, 118 Atomic compounds, 175 hypothesis, 71 volume, 258 weight, 167 thermal, 262 Atomicity, 170 Atoms, 25, 167 Auric acid, 348 compounds, 347, 348 Aurons compounds, 347 Avogadro, law of, 77 Azote, 120 Azurite, 339 36 425 426 INDEX. B. Barium, 312 carbonate, 313 nitrate, 313 oxide, 312 peroxide, 312 sulphate, 313 Baryta water, 312 Bases, 39, 59, 84 Basicity of acids, 177, 217 Bauxite, 353 Beryllium, 319, 320 Bessemer steel, 398 Bismuth, 150, 374, 375 alloys, 376 nitrate, 376 oxide, 375 oxychloride, 375 subnitrate, 376 Bismuthic acid, 375 Bleaching, 52 lime, 306 Borax, 245, 298 Boric acid, 245 meta-, 245 Boro-fluoride, hydrogen, 245 Boron, 243 chloride, 244 fluoride, 244 hydride, 244 nitride, 245 trioxide, 245 Brass, 339 Britannia metal, 368 Bromic acid, 183 Bromine, $2 Bronzes, 339 Brown coal, C. Cadmium, 322 oxide, 322 sulphate, 323 sulphide, 323 Cadmous oxide, 323 Caesium, 289 Calamine, 322 Calcium, 304 carbonate, 309 chloride, 306 fluoride, 306 hydroxide, 305 hypochlorite, 306 oxide, 304 Calcium peroxide, 304 phosphates, 308, 309 silicate, 310 sulphate, 308 sulphides, 310, 311 Calomel, 327 Caloric, 66, 260 Caput mortuum, 402 Carbon, 151 amorphous, 152 chlorides, 161 dioxide, 231, 233, 234 liquid, 232 disulphide, 237 gas, 132 group, 151 monoxide, 235, 236 oxychloride, 237 oxysulphide, 238 Carbonates, 235 Carbonic acid, 231, 235 Carré's ice machine, 127 Camallite, 282, 317 Cassiterite, 367 Catalytic reactions, 104 Caustic potash, 281 soda, 290 Cement, 305 Cementation, steel, 398 Cerium, 358 Cerrusite, 373 Chalk, 309 Chamber acid, 195 Chameleon minerals, 393 Charcoal, 152 animal, 152 Chemical affinity, 23 elements, 19 energy, 22 equations, 26, formulas, 23 Chemical structure, 171, 173 Chemistry, defined, 19 Chloric acid, 181 Chloride of lime, 306 Chlorine, 49 dioxide, 180 hydrate, 5* preparation, 49 properties, 50, 51 tetroxide, 180 trioxide, 180 Chlorites, I to Chlorous acid, ito INDEX. 427 Chromates of potassium, 382 Chromic acid, 380 anhydride, 381 chloranhydrides, 383 compounds, 379 oxide, 379 Chromium, 377 group, 376, 377 hexafluoride, 384 Chromous compounds, 378 Chromyl chloride, 383 Cinnabar, 326, 330, 331 Coal, anthracite, 151 bituminous, 151 brown, 151 Cobalt, 406 ammonio-compounds, 407 cyanides, 408 Cobaltic compounds, 407 Cobaltous compounds, 406 Coke, 152 Colcothar, 402 Colloids, 238 Condensation of gases, 47 Constant proportions, law of, 70 Copper, 331, 334, 335 alloys, 339 carbonates, 338, 339 hydride, 337 sulphate, 338 Corrosive sublimate, 329 Corundmn, 352 Critical condition, 47 pressure, 47, 233 temperature, 47 volume, 47 Crystallography, 30 Crystalloids, 241 Cupric compounds, 337 Cuprous compounds, 336 Cyanogen, 238 D. Davy's lamp, 161 Density, gas, 73 Determined compounds, 92 Dialysis, 240 Diämide, 132 Diammonium compounds, 314 Diamond, 151 Dissociation, 96 Dithionic acid, 202 Doebereiner's lamp, 46 Drummond light, 81 E. Earth metals, 349, 350 Electro-chemical theory, 270 Electrolysis of salts, 271 Endothermic compoimds, 28 reactions, 28 Equivalence, 171 Equivalent weight, 171 Ethane, 154 Ethylene, 155 Eudiometric analysis, 121 Exothermic compounds, 28 reactions, 28 F. Ferric acid, 403 chloride, 402 compounds, 402 sulphide, 403 Ferropentacarbonyl, 400 Ferrous compounds, 400 sulphate, 400 sulphide, 401 Flame, nature of, 156, 157» ^S^» *6° oxidizing, 159 reducing, 159 Fluorine, 55 Fluorite, 306 Fulminating gold, 348 silver, 342 G. Gadolinite, 358 Gahnite, 354 Galenite, 371 Gallium, 360 compounds, 360 group» 359 Gas, laughing, 216 Gases, dihusion of, 124 drying and purifying, 42 measuring of, 123 Germanium, 166, 241, 365 Germanic compormds, 366, 367 Germanous compounds, 366 Glass, 310 Glauber's salt, 292 Glucinum, 320 Gold, 331, 346 chlorides, 347 fulminating, 348 oxides, 347, 348 Graphite, 151, 152 Greenockite, 323 428 INDEX. Guignet's green, 379 Gunpowder, 286 Gypsum, 308 H. Halogen compounds of metals, 266 hydrides, 56 Halogens, 49 thermo-chemical department of, 66 Haloid acids, 60 salts, 60 Heat, atomic, 260 of formation of compounds, 423 latent, 91 modulus, 66 of solution, 93 specific heat, 260 unit of, 90 Hepar, 288 Homsilver, 343 Hydraulic cement, 305 Hydrazine, 132 Hydrazoic acid, 132 Hydriodic acid, 62 Hydrobromic acid, 61 Hydrocarbons, saturated, 155 vmsaturated, 155 Hydrochloric acid, 58 Hydrocyanic acid, 238 Hydrofluoric acid, 65 Hydrosulphuric acid, 109 Hydrosulphurous acid, 190 Hydrogen, 40 antimonide, 148 bromide, 61 chloride, 56 fluoride, 64 iodide, 62 peroxide, lOI detection of, 104 thermo-chemistry of, 105 persulphide, iii preparation, 41 properties, 43, 45 selenide, 114 silicofluoride, 165 sulphide, io8, 109 telluride, 114 Hydroxyl, 105 Hydroxylamine, 131 Hypobromous acid, 183 Hypochlorous acid, 179 oxide, 179 Hyponitric acid, 210 Hypxjnitrous acid, 216 oxide, 215 Hypophosphoric acid, 220 Hypophosphorous acid, 218 Hyposulphites, 201 Hyposulphurous acid, 201 I. Indium, 360 compounds, 361 Iodic acid, 183 anhydride, 184 Iodine, 54 bromide, 69 chloride, 69 fluoride, 69 trichloride, 69 Iridium, 412 Iron, 396, 399 carbide, 404 cast, 397 group, 306 metallurgy, 397, 398 vitriol, 401 wrought, 397 Isomerism, 88 Isomorphism, 263 K. Kaolin, 350, 357 Kelp, 54 Kermes mineral, 228 Kieserite, 317 L. Lanthanum, 358 Lapis lazuli, 357 Laughing gas, 216 Lead, 241, 371 alloys, 371 carbonate, 373 chamber crystals, 195 chloride, 373 Chromate, 383 iodide, 373 nitrate, 373 oxides, 372 peroxide, 372 red, 372 sulphate, 373 sulphide, 374 tree, 371 white, 373 Lime, 304 chloride of, 306 light, 84 INDEX. 429 Litharge, 372 Lithium, 298 carénate, 299 chloride, 299 phosphate, 299 Lunar caustic, 345 M. Magisterium bismuthi, 376 Magnesia alba, 318 Magnesium, 316 carbonate, 318, 319 chloride, 316 oxide, 316 phosphates, 318 reducing action of, 255 sulphate, 317 Magnetite, 402 Malachite, 339 Manganese, 389 bronze, 340 heptoxide, 394 peroxide, 392 Manganic acid, 393 compounds, 391, 392 Manganites, 393 Manganous compounds, 390 sulphate, 390 Marsh gas, 153 test, 145 Massicot, 372 Maximum valence, 174 Mercuric compounds, 329, 330 Mercurous compounds, 327, 328 Mercury, 324, 326 Meta-acids, 205 phosphoric acid, 221 stannic acid, 370 Metalloids (non-metals), 20^ 39 Metals, 20, 39, 257 heavy, 259 light, 259 properties of, 257 rare, 358 Methane, 153, 154 Mineral water, 92 Molecular compounds, 175 weight determination, 168, 169 Molecules, 25, 167 Molybdenum, 385 compounds, 385, 386 Molybdic acid, 386 Mosaic gold, 370 Multiple proportions, 72 N. Nickel, 408 alloys, 409 oxides, 408, 409 plating, 409 sulphate, 409 Niobium, 230 Nitrates, 207 Nitric acid, 206, 207 fuming, 207 oxide, 213, 214 Nitrogen, 116, 117, 118 chloride, 133 iodides, 133, 134 oxides, 205 pentoxide, 208 tetroxide, 210 trioxide, 208 Nitroso-acid, 212 Nitrous acid, 209 oxide, 215 Nitrosyl chloride, 208 sulphuric acid, 211 Nitroxyl chloride, 208 Non-metals, 20 Nordhausen sulphuric acid, 198 O. Olefiant gas, 155 Organic chemistry, 153 Ortho-acids, 205 Osmium, 411,412 Oxidation, 84 Oxides, 84, 266 indifferent, 84 Oxygen, 81 atomic weight of, 98 group, 80 liquid, 82 preparation of, 8l properties of, 82 oxidized, 87 Ozone, 85 tests for, 86 P. Palladium, 413 hydride, 47, 413 Palladious compounds, 414 Passive iron, 400 Pattison's method, 341 Pentathionic acid, 202 Perbromic acid, 183 Perchloric acid, 182 hydrate of, 184 430 INDEX. Periodic acid, 184 hydrate of, 184 system, 246 Periodicity of chemical valence, 252 Permanent white, 313 Permanganic acid, 393 Peroxides, 267 Persulphates, 191 Phlogiston, 119 Phosgene gas, 237 Phospham, 223 Phosphates, 218 Phosphine, 138, 139 Phosphites, 218 Phosphonium, 139 Phosphoric acid, 219 hypo-, 220 meta-, 221 pyro-, 220 Phosphorite, 304 Phosphorous acid, 219 oxide, 219 Phosphorus, 134, 140 acids, 218 amorphous, 136 bromides, 143 bronze, 339 bums, 137 fluoride, 143 iodides, 143 metallic, 136 oxides, 217 oxychloride, 142 pentachloride, 141 pentoxide, 222 red, 136 salt of, 301 sulphochloride, 142 sulphur derivatives, 223 trichloride, 140 yellow, 135 Phot(^raphy, 344 Pink salt, 369 Platinum, 414 ammonio-compounds, 416 black, 414 metals, 409, 410, 411 sponge, 414 Platinic compounds, 416 Platinous compounds, 415 Plumbates, 373 Plumbic acid, 373 Polychromates, 380 Polysulphides, 288 Polythionic acids, 200 Porcelain, 357 Potashes, 286 Potassium, 279 borates, 286 bromide, 282 carbonate, 286 chlorate, 283 chloride, 282 chromâtes, 382 cyanide, 283 ferrocyanide, 404 ferricyanide, 405 fluoride, 283 hydroxide, 281 iodide, 283 nitrate, 285 nitrite, 286 oxides, 281 perbromate, 284 Perchlorate, 284 permanganate, 394 phosphates, 286 silicate, 287 sulphates, 284 sulphides, 287, 288 sulphites, 285 Preparing salts, 370 Prussian blue, 405 Prussic acid, 238 Purple of Cassius, 348 Pyrites, 403 I^ophosphoric acid, 220 P^osulphuric acid, 197 Q. Quadrant oxides, 268 Quartz, 239 Quicksilver, 326 R. Radical, 178 Realgar, 225 Reducing action of magnesium, 255 Reduction, 85 Residue, acid, 178 Rhodium, 412 Rose's metal, 376 Rubidium, 289 Ruby, 352 Ruthenium, 411 S. Safety lamp, i6l Saltpetre, 285, 296 INDEX. 431 Salt of phosphorus, 301 producers, 49 Salts, 59, 268 ammoiMum, 129 basic, 269 double, 269 haloid, 60 oxygeû, 60 Sapphhe, 352 Scandium, 358 Scheele's green, 339 Schlippe's salt, 228 Schweizer's'reagent, 337 Selenimn, 113 acids, 204 chlorides, II4 hydride, 114 oxide, 204 Silica, 239 Silicates, 241 Silicic acid, 240, 241 Silicon, 162 bromide, 164 bronze, 340 chloride, 163 chloroform, 164 dioxide, 239 disulphide, 241 fluoride, 165 hydride, 163 iodide, 164 Sflver, 331, 340, 341 bromide, 343 chloride, 343 coins, 342 cyanide, 344 iodide, 343 nitrate, 344 oxides, 342 plating, 344 sulphide, 345 Silvering, 345 Slags, 398 Slaked lime, 305 Smalt, 406, 407 Soda, caustic, 290 residue, 295 solving, 295 Sodium, 289 borate, 298, 299 bromide, 291 carbonate, 294, 296 chlorate, 291 chloride, 290 Sodium hydroxide, 290 hyposulphite, 293 iodate, 291 iodide, 291 nitrate, 296 nitrite, 296 oxides, 290 Perchlorate, 291 periodate, 291 phosphates, 296, ,297 silicate, 298 sulphate, 292, 293 sulphite, 293 Soft solder, 371 Solutions, 92 supersaturated, 293 thermo-chemistry, 93 Specific gravity of gases, 73 Specific volume, 258 Spectrum analysis, 416, 417 lines, periodicity of, 421 Spinels, 354 Stannic acid, 370 compounds, 369, 370 Stannous compounds, 368, 369 Status nascens, 52, 79, 103, 131, 168, 207 Steel, 397, 399 Stibine, 148 Strass, 310 Strontium, 311 carbonate, 311 nitrate, 311 oxide, 311 Structure, chemical, 171 Substitution, 161 Sulphates, 186 Sulphides, no Sulphites, 189 Sulphocarbonic acid, 237 Sulpho-group, 201 Sulpho-stannates, 370 Sulphur, 105, 106, 107 bromides, 113 chemical properties of, 108 chlorides, 112 dioxides, 187, 188 heptoxide, 190 iodides, 113 sesquioxide, 190 trioxide, 191, 192 Sulphuric acid, 193, 194, 195 amides, 203 chlor-anhydtides, 199 432 INDEX. Sulphuric acid, di-, 197 fuming, 198 hydrates, 196 Nordhausen, 198 Sulphurous acid, 189 Sulphiuyl, 199 Syloite, 282 S3nnpathetic ink, 406. T. Tantalum, 230 Tellurium, 114 acids, 204 bromides, II4 chlorides, II4 hydride, 114 Tension of vapors, 91 Terbium, 359 Tetrathionic acid, 202 Thallic acid, 363 compounds, 363 Thallium, 361 alum, 363 Thallous compounds, 362 Thenard's blue, 407 Thermo-chemistry of the elements, 26, 66, 88, 93, 94, 105, no, 115, 129, 134, 143, 147, 151, 161, 185, 189, 192, 204, 216, 229, 236, 237, 239, 254, 271, 277, 279, 290, 301, 302, 323» 324. 334. 362, 372, 390- Thionic acids, 200 Thionyl chloride, 189, 200 Thiosulphuric acid, 20i Thorium, 241, 243 Thulium, 24 Tin, 166, 241, 367 dichloride, 368 dioxide, 370 disulphide, 370 monoxide, 368 salt, 368 stone, 367 tetrachloride, 369 Titanium, 241, 242 Trithionic acid, 202 Tombac, 339 Trona, 296 Tungsten, 386 compounds, 387 Turpeth mineral, 330 Tumbull's blue, 405 Type metal, 371 U. Ultramarine, 357 Undetermined compounds, 92 Uranium, 387 Uranic compounds, 388 Uranous compounds, 387 Uranyl, 387 Urao, 296 V. Valence, 170, 174 periodicity of, 252 Vanadium, 229 Vapor density, 73 Varec, 54 Vitriol, copper, 338 green, 403 iron, 401 oil of (see sulphuric acid) Volume, atomic, 258 specific, 258 W. Water, 89 chemical properties of, 93 crystallization, 92 constitution, 317 dissociation of, 94 electrolysis of, 93 gas, 235 glass, 287 hard, 92 mineral, 92 natural, 91 physical properties of, 89 quantitative composition of, 97 soft, 92 White precipitate, 329 lead, 374 Wood's metal, 376 Y. Yellow prussiate of potash, 404 Ytterbium, 358 Yttrivun, 359 Z. Zinc, 320 blende, 322 dust, 320 oxide, 321 sulphate, 321 sulphide, 322 white, 321 Zircon, 242 Zirconium, 241, 242