7 i II ii ii , PP PLATE OF SPECTRA The numbers at top and bottom of plate give wave lengths for each spectrum, in hundredths of a micron (//). TEXTBOOK OF CHEMISTEY BY WILLIAM A. NOYES DIRECTOR OF THE CHEMICAL LABORATORY OF THE UNIVERSITY OF ILLINOIS NEW YORK HENRY HOLT AND COMPANY 1919 iv PREFACE time displaced by something very different. It is believed that the development of our knowledge during the last few years fully justifies this course. The theory of ionization has also been freely used, as the only means we have by which a large class of phenomena can be clearly presented and understood. It seems desirable to give some material which it is not possible to emphasize or teach thoroughly in a brief course and some things which are rather for reference than to be learned. To aid teachers and students in distinguishing such paragraphs, they are indicated by an asterisk. This device, which was, I think, first used by Professor Ostwald, is better adapted than the use of fine print to the need of teachers who may wish to make a different selection. Students are earnestly advised to read these paragraphs, even when they are not expected to acquire a full knowledge of them. I wish to express my very sincere thanks to the following members of the Chemical Staff of the University of Illinois, some of whom have met with me almost weekly for three years to read and criticize the successive chapters of the book. The criticisms have been very valuable and helpful. C. W. Balke, S. J. Bates, Edward Bartow, G. D. Beal, L. L. Burgess, E. S. Curtiss, C. G. Derick, Lambert Thorp, B. S. Hopkins, Helen Isham (Mattill), Grinnell Jones, C. G. MacArthur, Ellen S. McCarthy (Foley), D. F. McFarland, D. A. Maclnnes, C. F. Nelson, S. W. Parr, G. McP. Smith, E. K. Strachan, E. W. Washburn, H. C. P. Weber. I wish also to express my gratitude to Professor Edward W. Morley, Professor Julius Stieglitz of the University of Chi- cago, Professor J. Bishop Tingle of MacMaster University, and Mr. C. M. Wirick of the Crane Technical High School of Chicago, who have read the proofs and made many useful sug- gestions. I am also indebted to Professor K. B. Moore of the Bureau of Mines for some valuable criticisms, of the paragraph on Radioactivity. CONTENTS CHAPTER I INTRODUCTION The Nature of Scientific Knowledge, 1. Subdivisions of Science, 3. Physical Sciences, 4. Matter and Energy, 5. Conservation of Matter and Energy, 6. Pure Substances and Mixtures, 7. Preparation of Pure Sub- stances, 8. Elements and Compounds, 9. How Pure Substances are dis- tinguished from Mixtures. Law of Constant Proportion, 12. Inductive Reasoning, 13. Law of Combining Weights, 13. The Atomic Theory, 14. Selection of Atomic Weights, 16. Formulas, 16. Composition of Pure Substances, 17. Study of Chemistry, 18. CHAPTER II OXYGEN SYMBOL, O. ATOMIC WEIGHT, 16 Occurrence, 19. Preparation, 19. Collection and Storage of Gases, 22. Properties of Oxygen, 22. Oxygen and Acid Properties, 23. Combustion. Effect of Concentration on a Chemical Reaction, 24. Kindling Tempera- ture, 24. Heat of Combustion. Calorimeter, 25. The Nature of Chemical Energy, 27. Catalysis, 28. Chemical Affinity, 29. Nomenclature, 29. CHAPTER III LAWS OF GASES UNITS OF LENGTH, WEIGHT, VOLUME, TEMPERATURE, TIME AND ENERGY Unit of Length. Meter, 31. Unit of Weight. Gram, 31. Unit of Vol- ume. Liter, 31. Units of Time, 32. Unit of Temperature, 32. Units of Energy. Kilogram-meter. Erg, 32. Centimeter-gram-second System. Ab- solute Units, 33. Units of Mechanical Energy, 33. Unit of Power, 33. Units of Heat, 33. Electrical Units, 33. Chemical Energy, 34. Effect of Pressure on a Gas. Law of Boyle, 34. Corrections for Readings of the Barometer, 36. Effect of Temperature on a Gas. Law of Charles, 38. Absolute Temperatures, 39. Significance of the Absolute Zero, 40. Deter- mination of the Weight of a Liter of a Gas, 40. Graphical Representation of the Gas Laws, 42. Exercises, 43. v yi CONTENTS CHAPTER IV HYDROGEN SYMBOL, H. ATOMIC WEIGHT, 1.0078 Occurrence, 45. Radicals, 46. Salts, 47. Preparation of Hydrogen. 1, Electrolysis of Dilute Sulf uric Acid, 47. Electrolytes. Ions. Theory of Electrolysis, 48. 2, Preparation of Hydrogen from Iron and Steam, 48. Reversible Reactions, 50. 3, Decomposition of Water by Metals at Ordi- nary Temperatures, 50. Contrast between the Action of Iron and of Sodium on Water, 51. 4, Hydrogen from "Hydrone," 52. 5, Preparation of Hydrogen by the Action of Metals on Acids, 52. Apparatus for the Prepa- ration of Hydrogen, 53. Purification of Hydrogen, 54. Properties of Hydrogen, 55. Diffusion of Gases, 56. Kinetic Theory of Gases, 58. Chemical Properties of Hydrogen, 59. Dissociation, 59. The Oxyhydrogen Blowpipe, 61. Explosions. Catalysis, 62. Oxidation. Reduction, 63. Valence, 63. Heat of Combustion of Hydrogen, 65. CHAPTER V WATER, HYDROGEN PEROXIDE Analysis, Synthesis, 66. Qualitative Analysis and Synthesis of Water, 66. Quantitative Synthesis of Water by Volume, 66. Composition of Water by Weight, 68. The Unit for Atomic Weights, 68. Determination of the Composition of Water by the Use of Copper Oxide, 69. Determination of the Composition of Water by Weighing Oxygen and Hydrogen, 71. Proper- ties of Water, 72. Heat of Fusion and Vaporization, 74. Vapor Pressure of Water, 74. Equilibrium, 76. Effect of Water Vapor on the Pressure of a Gas, 76. Phases. Degrees of Freedom, 77. Water as a Solvent. Solutes, 79. Chemical Activity in Solutions. Metathesis, 81. Hydrates, Deliques- cence, Efflorescence, 81. Natural Waters, 82. Purification of Water, 83. Hydrogen Peroxide, 83. Properties and Uses of Hydrogen Peroxide, 86. Tests for Hydrogen Peroxide, 85. Structure of Hydrogen Peroxide, 86. Law of Multiple Proportion, 87. CHAPTER VI AVOGADRO'S LAW. SELECTION OF ATOMIC WEIGHTS. OZONE Gay Lussac's Law of Combining Volumes, 89. Avogadro's Law, 91. Selection of an Atomic Weight, 92. Molecules of the Elements, 93. Gram Molecular Volume, 94. Number of Molecules in one Cubic Centimeter of a Gas, 95. Allotropic Forms. Ozone, 97. Exercises, 99. CHAPTER VII CHLORINE SYMBOL, CL. ATOMIC WEIGHT, 35.46. FORMULA, C1 2 Occurrences of Chlorine, 100. Preparation of Chlorine. 1, By Electrol- ysis of Sodium Chloride, 100. 2, Preparation by Oxidation of Hydrochloric CONTENTS Vii Acid, 100. 3, Preparation of Chlorine by the Deacon Process, 102. 4, The Weldon Process for Chlorine, 103. Properties of Chlorine, 104. Chlorine and Water. Bleaching, 106. Chlorine Hydrate. Phases, 107. The Heat of Combination of Chlorine and of Oxygen with Other Elements, 108. Equi- librium in Chemical Reactions, 108. Principle of van't Hoff-Le Chatelier, 111. Effect of Water on Chlorides. lonization, 112. Effect of Water on Chlorides. Hydrolysis, 115. Exercises, 116. CHAPTER VIII HYDROCHLORIC ACID. OXIDES AND OXYACIDS OF CHLORINE Hydrochloric Acid, 118. Properties of Hydrochloric Acid, 119. 1, Reac- tion with Metals, 120. 2, Reaction with Hydroxides of Metals, 121. 3, Re- action with Oxides of Metals, 122. 4, Reaction with Oxidizing Agents, 122. Indicators, 122. Oxides and Oxygen Acids of Chlorine. Nomenclature, 123. Hypochlorous Acid. Hypochlorites, 124. Hypochlorous Anhydride, or Chlorine Monoxide, 126. Chlorous Acid and Chlorites, 127. Chloric Acid and Chlorates, 127. Chlorine Peroxide, 127. Perchlorates and Per- chloric Acid, 128. Structure of the Oxyacids of Chlorine, 130. The Atomic Weight of Chlorine, 130. CHAPTER IX CLASSIFICATION OF THE ELEMENTS. THE PERIODIC SYSTEM, 132 CHAPTER X THE HALOGEN FAMILY General Properties of the Halogens, 139. Compounds of the Halogens with Hydrogen and Oxygen, 138. Bromine, Br, 79.92. Occurrence. Prepa- ration, 140. Properties, 141. Hydrobromic Acid, 142. Sodium Hypo- bromite, 143. Iodine, I, 126.92. Occurrence, Preparation, 144. Properties of Iodine, 144. Hydriodic Acid, 145. Direct Combination of Hydrogen and Iodine. Reversible Reactions. Equilibrium, 146. Speed of Chemical Reac- tion, 148. Concentration and Speed of Reaction, 149. Calculation of the Relative Speed of Two Reactions from the Composition of an Equilibrium Mixture, 151. Effect of Removing One of the Reacting Substances. Dis- placement of the Equilibrium Point, 152. Heat of Formation of Hydriodic Acid, 152. Flourine, F, 19.0. Occurrence, 153. Preparation, 153. Prop- erties, 154. Etching Glass. Hydrofluoric Acid, 154. Metallic Elements of Group VII, 156. Exercises, 156. CHAPTER XI SULFUR, SELENIUM AND TELLURIUM Sulfur, S, 32.0. Occurrence, 160. Allotropic Forms of Sulfur, 162. Liquid Forms of Sulfur, 162. Gaseous Forms of Sulfur, 163. Properties viii CONTENTS and Uses of Sulfur, 163. Hydrogen Sulfide, 164. Solution of Hydrogen Sulfide. Henry's Law, 165. Sulfides. Groups of Analytical Chemistry, 166. Hydrosulfuric Acid. Strength of Acids, 167. Application of the Idea of Strength of Acids to Explain the Conduct of Sulfides, 168. Hydrogen Sul- fide as a Reducing Agent, 171. Sulfur Dioxide, 172. Sulfurous Acid, 174. Sulfites, 175. Sulfur Trioxide, 175. Sulf uric Acid, 177. The Electron Theory, 181. Sulf uric Acid as a Dehydrating Agent, 182. Sulfates. Di- basic Acids, 183. Normal, Standard and Formular Solutions, 183. Acid- imetry and Alkalimetry, 185. Pyrosulfates, 186. Hyposulfites, 186. Thiosulfates, 186. Persulfuric Acid, 187. Permonosulfuric Acid, 188. Polythionic Acids, 188. Compounds of Sulfur Containing Halogens, 188. Sulfur Monochloride, 188. - Chlorosulfonic Acid, 189. Sulfuryl Chloride, 189. Selenium, Se, 79.2, 189. Hydrogen Selenide, 190. Selenium Dioxide, 190. Selenic Acid, 190. Tellurium, Te, 127.5, 190. Atomic Weight of Tellurium, 190. General Properties of the Elements of the Sixth Group, 191. Crystals, 192. 1, The Isometric or Regular System, 193. 2, The Tetragonal System, 194. 3, The Rhombic System, 194. 4, The Hexagonal System, 194. 5, The Monoclinic System, 195. 6, The Triclinic System, 196. Exercises, 196. CHAPTER XII NITROGEN SYMBOL, H. ATOMIC WEIGHT, 14.01 Occurrence and Natural History of Nitrogen, 198. Preparation and Prop- erties of Nitrogen, 200. Ammonia, 201. Properties of Ammonia, 202. Aqua Ammonia, 203. Ice Machines, 204. Derivatives of Ammonia, 205. The Electron Theory, 206. Solutions in Liquid Ammonia, 207. The Volu- metric Composition of Ammonia, 208. Nitric Acid, 210. Hydrates of Nitric Acid, 211. Chemical Properties of Nitric Acid, 212. Aqua Regia, 213. Nitrosyl Chloride, 214. Oxides of Nitrogen, 214. Nitrous Oxide, 214. Nitric Oxide, 215. Nitrous Anhydride, 218. Nitrous Acid, 218. Nitrogen Peroxide and Nitrogen Tetroxide, 219. Nitrogen Pentoxide, 220. Other Compounds of Nitrogen, 220. Hyponitrous Acid, 221. Hydro xylamine, 221. Hydrazine, 222. Hydronitric Acid, or Azoimide, 223. Iodine Trinitride, 223. Nitrogen Trichloride, 224. Nitrogen Iodide, 225. Nitro Nitrogen Trichloride, 225. Endothermic Compounds, 225. Exercises, 225. CHAPTER XIII THE ATMOSPHERE. NOBLE GASES Determination of Oxygen, 227. Composition of Air, 227. Air is a Mix- ture, 228. Carbon Dioxide in the Air, 229. Ventilation, 230. Moisture, 231. Liquid Air. Critical Temperature, 232. Argon, A, 39.88, 235. Atomic Weight of Argon. Specific Heat of Gases, 236. Helium, He, 3.99, 237. Neon, Krypton, Xenon and Niton, 238. Exercises, 238. CONTENTS ix CHAPTER XIV PHOSPHORUS Phosphorus , P, 31.04. Occurrence, 240. Preparation of Phosphorus, 241. Allotropic Forms of Phosphorus, 241. Matches, 242. Phosphine, 243. Phosphonium Salts, 243. Phosphorus Trichloride, and Phosphorus Penta- chloride, 244. Hydrolysis of the Chlorides of Phosphorus, 245. Phosphorus Oxychloride, 245. Oxides of Phosphorus, 246. Acids of Phosphorus, 246. Basicity of the Acids of Phosphorus, 247. Hypophosphorous Acid, 248. Phosphorous Acid, 248. Orthophosphoric Acid, 248. lonization of Ortho- phosphoric Acid, 250. Decomposition of Primary and Secondary Salts of Orthophosphoric Acid, 252. Pyrophosphoric Acid, 253. Metaphosphoric Acid, 253. Hypophosphoric Acid, 254. Sulfides of Phosphorus, 254. Exercises, 255. CHAPTER XV ARSENIC, ANTIMONY AND BISMUTH Arsenic, As, 74.96. Occurrence, 256. Preparation and Properties of Arsenic, 257. Arsine, Marsh's Test, 257. Arsenic "Trioxide," 258. Arsenious Acid, 259. Arsenic Pentoxide and Arsenic Acid, 259. Arsenic Trichloride, 260. Sulfides of Arsenic, 260. Arsenic Disulfide, or Realgar, 260. Arsenic Trisulfide, or Orpiment, 260. Arsenic Pentasulfide, 260. Sulfarsenites and Sulfarsenates, 261. Colloidal Arsenic Trisulfide, 261. Antimony, Sb, 120.2. Occurrence and Preparation, 263. Properties, 264. Uses, 264. Stibine, 264. Oxides of Antimony, 265. Antimony Hydroxide. Antimonious Acid, 265. Tartaric Emetic, 266. Antimonic Acids, 267. Chlorides of Anti- mony, 267. Antimony Trichloride, 267. Antimony Tetrachloride, Hydro- tetrachloroantimonic Acid, 267. Antimony Peutachloride, 267. Metachlo- roantimonic Acid, 268. Antimony Trisulfide, 268. Antimony Pentasulfide, 268. Sulfantimonites and Sulfantimonates, 268. Bismuth, Bi, 208. Occur- rence, Properties, Uses, 268. Oxides of Bismuth, 269. Bismuth Chloride, 269. Bismuth Nitrate, 270. Bismuth Trisulfide, 270. Tables of Compounds of the Elements of the Fifth Group, 270. Vanadium, Columbium Tantalum, 271. Exercises, 272. CHAPTER XVI CARBON Carbon. Occurrence, 273. Diamonds, 274. Graphite, 276. Amorphous Carbon, 277. Lampblack, 277. Wood Charcoal, 277. Animal Charcoal and Bone Black, 278. Coke, 278. Gas Carbon. Carbon Electrodes, 279. Coal, 280. Chemical Properties of Carbon, 281. CHAPTER XVII HYDROCARBOUS, ILLUMINATING AND PRODUCER GAS. FLAME Marsh Gas or Methane, 286. Substitution, 287. The Davy Safety Lamp, 287. Homologues of Methane, 289. Petroleum, 289. Ethylene or X CONTENTS Ethene, 290. Unsaturated Compounds. Ethylene Chloride and Ethylene Bromide, 291. Acetylene, 292. Benzene, 294. Illuminating Gas, 295. Oil Gas, 296. Water Gas, 296. Producer Gas, 297. Blast-furnace Gas, 298. Luminous Flames, 299. Bunsen Burner, 300. Explosion Waves, 301. Temperature of Flames, 302. Blowpipe, 303. Reversed Flames, 304. Exercises, 305. CHAPTER XVIII OXIDES AND SULPHIDES OF CARBON. ASSIMILATION AND RESPIRATION. CYANIDES Carbon Dioxide, 306. Isothernials of Carbon Dioxide, 307. Density of Carbon Dioxide, 308. Aqueous Solutions of Carbon Dioxide, Carbonic Acid, 309. Carbonates and Bicarbonates. Hard Waters, 309. Carbon Monoxide, 311. The Cycle of Carbon in Nature, 312. Respiration Calorimeter, 313. Carbon Suboxide, 316. Carbon Oxychloride, or Phosgene (Carbonyl Chlo- ride), 316. Carbon Bisulfide, 317. Sulfocarbonates, 317. Sulfocarbonic Acid, 318. Carbon Oxysulfide, 318. Cyanides, 319. Hydrocyanic Acid, or Prussic Acid, 319. Potassium Cyanide, 319. Complex Cyanides, 319. Potassium Cyanate, 321. Potassium Thiocyanate, 321. Cyanogen, 322. Exercises, 322. CHAPTER XIX ALCOHOLS, ALDEHYDES, KETONES, ACIDS, FATS, CARBOHY- DRATES Structural Formulas, 323. l, Valence, 323. 2, Radicals, 323. 3, Substi- tution, 324. Alcohols, 324. Methyl Alcohol, 324. Ethyl Alcohol, 325. Phenol or Carbolic Acid, 326. Glycerol, 326. Aldehydes and Ketones, 327. Formaldehyde, 327. Benzaldehyde, 328. Acetone, 328. Acids, 328. Formic Acid, 329. Acetic Acid, 329. Oxalic Acid, 329. Lactic Acid, 330. Tartaric Acid, 330. Citric Acid, 330. Ammonium Ferric Citrate, 331. Benzoic Acid, 331. Palmitic, Stearic and Oleic Acids, Fats, 331. Soaps, 332. Carbohydrates, 332. Cane Sugar, or Saccharose, 333. Maltose, 334. Lactose, or Milk Sugar, 334. Glucose, 334. Fructose, 335. Starch, 335. Dextrin, 336. Pectose, Pectin, 337. Cellulose, 337. Paper, 337. Gun Cotton, Celluloid, Lacquers, 338. CHAPTER XX AMINES, DYES, ALKALOIDS, PROTEINS, ENZYMES, FOODS AND NUTRITION Methyl Amine, 339. Aniline, 340. Dyes, 340. Alizarin, 341. Indigo, 341. Mordants, 342. Alkaloids, 342. Nicotine, 342. Coniine, 342. Atropine, 343. Cocaine, 343. Morphine, 343. Quinine, 343. Strychnine, 343. Ptomaines, 343. Proteins, 343. Enzymes, 344. Toxins, Antitoxins, 344. Urea, 345. Nutrition, 345. CONTENTS X i CHAPTER XXI SILICON, BORON, GERMANIUM, TIN, LEAD, TITANIUM, ZIRCO- NIUM, CERIUM, THORIUM SILICON, Si, 28.3 Occurrence, 348. Preparation, 349. Hydrogen Silicide, 349. Silicon Carbide, Carborundum, 349. Silicon Fluoride, 350. Fluosilicic Acid, 350. Silicon Tetrachloride, 351. Silicon Hexaiodide, 351. Silicon Dioxide, or Silica, 351. Artificial Silicates, 352. Silicic Acids, 353. Natural Silicates, 355. Calculation of the Formula of a Mineral, 356. Dialysis, Semiperme- able Membranes, 357. Osmotic Pressure, 358. Germanium, 361. Tin and Lead, 361. Titanium, 362. Zirconium, 363. Cerium, 363. Thorium, 364. Welsbach Mantles, 364. Boron, 365. Preparation, Properties, 365. Boron, Trioxide, Borax Beads, 365. Boric Acid, 366. Other Acids of Boron, 366. Borax, 367. Sodium Perborate, 367. Other Compounds of Boron, 367. Exercises, 368. CHAPTER XXII METALLIC ELEMENTS. DIFFERENCES BETWEEN METALS AND NON-METALS. PREPARATION OF COMPOUNDS Metals and Non-metals, 369. Classification of the Metals, 370. Melting Points of the Elements, 372. Preparation of Chemical Compounds, 372. Effect of Volatility, 374. Effect of Insolubility, 376. Effect of a Common Ion. Solubility Product, 377. Formation of Complex Ions, 378. Degree of lonization, 379. Effect of Degree of lonization, Neutralization, 384. Hydrolysis, 385. Illustration of the Strength of Acids, 386. Use of Indi- cators, 387. Systematic Study of the Metals, 390. Metallurgy, 390. Ox- ides, 392. Hydroxides, 392. Solubility of Salts, 393. CHAPTER XXIII ALKALI METALS: LITHIUM, SODIUM General Properties of the Alkali Metals, 395. Lithium, 395. Lithium Urate, 396. Atomic Weight of Lithium, Law of Diilong and Petit, 396. The Quantum Theory, 398. Sodium, 398. Metallurgy, Properties, 399. The Alkali Industry, 400. Sodium Hydroxide, 401. Sodium Oxide, 404. Sodium Peroxide, 404. Sodium Chloride, 404. Sodium Sulfate, Glauber's Salt, 406. Acid Sodium Sulfate or Sodium Bisulf ate, 408. -Sodium Sulfite, 408. Acid Sodium Sulfite, or Sodium Bisulfite, 408. Sodium Hyposulfite, 408. Sodium Thiosulfate, 408. Sodium Tetrathionate, 409. Sodium Sul- fide, 409. Sodium Hydrosulfide, 409. Sodium Nitrate, 410. Sodium Ni- trite, 410. Sodamide, 410. Sodium Trinitride, 410. Disodium Phosphate, 410. Sodium Carbonate or Sal Soda (Washing Soda). The Leblanc Soda xii CONTENTS Process, 411. Sodium Bicarbonate or Baking Soda, The Ammonia Soda Pro- cess, 412. Sodium Silicate, or Soluble Glass, 413. Sodium Tetraborate, or Borax, 413. CHAPTER XXIV ALKALI METALS: POTASSIUM, AMMONIUM, RUBIDIUM, CAESIUM. THE SPECTROSCOPE Potassium. Occurrence, 414. Metallic Potassium, 415. Potassium Oxide, 415. Potassium Hydroxide, 415. Potassium Chloride, 416. Potassium Chlorate, 416. Potassium Perchlorate, 416. Potassium Iodide, 417. Potas- sium Polyiodides, 417. Potassium Sulfates, 417. Acid Potassium Sulfate, or Potassium Bisulfate, 417. Potassium Nitrate, or Saltpeter, 417. Gun- powder, 418. Potassium Nitrite, 419. Potassium Carbonate, 419. Potas- sium Bicarbonate, or Saleratus, 420. Potassium Cyanide, 420. Ammonium, 420. Ammonium Hydroxide, 420. Ammonium Chloride, 421. Ammonium Sulfide, 421. Ammonium Hydrosulfide, 421. Ammonium Sulfate, 422. Ammonium Nitrate, 422. Ammonium Nitrite, 423. Ammonium Sodium Hydrogen Phosphate, 423. Ammonium Carbonate, 423. Ammonium Bi- carbonate, 423. Ammonium Chloroplatinate, 423. Rubidium and Caesium, 424. Spectrum Analysis, 424. CHAPTER XXV THE ALTERNATE METALS OF GROUP I. COPPER, SILVER, GOLD. PHOTOGRAPHY Copper. Occurrence, 428. Metallurgy, 428 Electrolytic Refining of Cop- per, 429. Properties of Copper, 430. Alloys of Copper, 431. Copper Hy- droxide, 431. Cupric Oxide, 432. Cuprous Oxide, 432. Cupric Chloride, 432. Cuprous Chloride, 432. Cuprous Iodide, 433. Cupric Sulfide, 433. Copper Sulfate, or Blue Vitriol, 433. Vitriols, 434. Cupric Nitrate, 434. Ammoniocupric Sulfate, 434. Cuprous Cyanide, 434. Precipitation of Cop- per by Iron, Electromotive Series, 435. Faraday's Law, 438. Silver, 439. Metallurgy, 439. Pattison's Process, 439. Cupellation, Assaying, 440. Parke's Process, 440. Amalgamation Process, 441. Other Processes for the Recovery of Silver, 441. Properties of Silver, Alloys, 442. Silver Plating, 442. Silver Oxide, 442. Silver Peroxide, 443. Silver Nitrate, 444. Silver Nitrite, 444. Silver Sulfate, 444. Silver Chloride, Silver Bromide, Silver Iodide, 444. Photography, 444. Gold, 445. Metallurgy, 446. Cyanide Process, 446. Properties of Gold, 448. Alloys of Gold, 448. Oxides of Gold, 448. -Gold Hydroxide, 448. Chlorides of Gold, 450. Exercises, 450. CHAPTER XXVI GROUP II. ALKALI-EARTH METALS: BERYLLIUM, CALCIUM, STRONTIUM, BARIUM, RADIUM Beryllium, 451. Calcium. Occurrence, 452. Preparation, Properties, 452. -Calcium Hydride, 452. Calcium Oxide, 452. Dissociation of Calcium CONTENTS xiii Carbonate and the Phase Rule, 453. Mortar, 454. Cement, 454. Cal. cium Chloride, 455. Chloride of Lime, 455. Calcium Chlorate, 456. Cal- cium Flouride, 456. Calcium Sulfide, 456. Acid Calcium Sulfite, 457. Calcium Sulfate, Plaster of Paris, 457. Plaster of Paris and the Phase Rule, 458. Calcium Nitrate, 460. Calcium Phosphates, 460. Solubility of Cal- cium Phosphates, 461. Calcium Carbide, 462. Calcium Cyanamide, 462. Calcium Carbonate, 463. Hard Waters, 463. Determination of Free and Combined Carbonic Acid in Natural Waters, 464. Calcium Acetate, 465. Calcium Oxalate, 465. Calcium Silicate, 466. Glass, 466. Strontium. Oc- currence, 467. Strontium Hydroxide, 468. Strontium Nitrate, 468. Ba- rium. Occurrence, 468. Barium Oxide, 468. Barium Peroxide, 469. Barium Hydroxide, 470. Barium Chloride, 470. Barium Nitrate, 470. Barium Sulfide, 470. Barium Sulfate, 470. Flame Colors for Calcium, Strontium and Barium, 471. Radium, 471. Disintegration of Atoms, 472. Nature of the Radiations from Radioactive Substances, 473. The Life of an Element, 474. Other Radioactive Elements, 475. Chemical Action of the Rays, 475. Radiochemistry in Relation to Geology and Medicine, 475. Exercises, 476. CHAPTER XXVII ALTERNATE METALS OF GROUP II. MAGNESIUM, ZINC, CADMIUM AND MERCURY Magnesium, 478. Preparation, Properties, 478. Magnesium Oxide, 479. Magnesium Hydroxide, 479. Magnesium Chloride, 480. Magnesium Am- monium Chloride, 480. Magnesium Sulfate, 480. Magnesium Sulfide, 480. Magnesium Ammonium Phosphate, 480. Zinc. Occurrence, 481. Metal- lurgy, 481. Uses. Galvanized Iron, 481. Sherardized Iron, 482. Zinc Oxide, 482. Zinc Hydroxide, 483. Zinc Chloride, 483. Zinc Sulfate, or White Vitriol, 483. Zinc Sulfide, 483. Cadmium, 483. Cadmium Hydrox- ide, 484. Cadmium Sulfate, 484. Cadmium Sulfide, 484. Mercury, Hg, 200.6. Occurrence. Metallurgy, 484. Properties and Uses, 485. Amalgams, 486. Compounds of Mercury, 488. Mercurous Oxide, 488. Mercuric Oxide, 488. Mercuric Sulfide, 489. Mercurous Chloride, or Calamel, 489. Mercuric Chloride, or Corrosive Sublimates, 489. Mercuric Iodide, 490. Mercurous Nitrate, 490. Mercuric Nitrate, 490. Mercuric Cyanide, 490. Mercuric Fulminate, 491. lonization of Compounds of Cadmium and Mercury, 491. Solubility of the Sulfides of Group II, 491. Conduct of Solution of Mag- nesium, Zinc and Cadmium Salts towards Ammonium Hydroxide, 491. Ammono-mercuric Compounds, 492. Nessler's Reagent, 492. Exercises, 493. CHAPTER XXVIII METALS OF GROUP III. ALUMINIUM FAMILY. RARE EARTH METALS Aluminium, 494. Metallurgy, 495. Properties of Aluminium, 497. Alloys, 497. Goldschmidt's Thermite Process, 497. Aluminium Chlo- ride, 498. xiv CONTENTS Aluminium Fluoride, 499. Aluminium Hydroxide, 499. Aluminium Oxide, 499. Aluminium Sulfate, 500. Alums, 500. Brick, Earthenware, Porcelain, 501. Ultramarine, 502. The Rare Earths, 502. Scandium, 503. Yttrium, 503. Lanthanum, 503. Ytterbium, 504. Praseodymium and Neo- dymium, 504. Samarium, 505. Europium, Gadolinium, Terbium, 505. Holmium, 505. Dysprosium, 505. Erbium, 505. Thulium, 506. Lute- cium, 506. Gallium, 506. Indium, 506. Thallium, 507. Exercises, 507. CHAPTER XXIX TIN AND LEAD Tin. Occurrence, Metallurgy, 508. Uses of Tin, Alloys, Tin Plate, 509. Compounds of Tin, 510. Stannous Oxide, 510. Stannous Chloride, 510. Stannous Sulfide, 510. Stannic Oxide, 510. Stannic Acids, 511. Stannic Acid, 511. Metastannic Acid, 512. Parastannic Acid, 512. Stannic Chlo- ride, 512. Stannic Sulfide, 512. Firep roofing of Cotton Goods, 513. Lead. Occurrence, Metallurgy, 513. Properties and Uses of Lead, Alloys, 514. Oxides of Lead, 515. Lead Monoxide, or Litharge, 515. Storage Batteries, 516. Lead Sulfide, 518. Lead Chloride, 518. Lead Tetrachloride, 518. Lead Sulfate, 519. Lead Nitrate, 519. Lead Acetate, or Sugar of Lead, 519. Basic Lead Acetates, 519. Lead Carbonate, 519. Basic Lead Carbonate, or White Lead, 520. CHAPTER XXX VANADIUM AND CHROMIUM GROUPS Group V. Vanadium, 522. Columbiurn (or Niobium), 523. Tantalum, 523. Group VI. Chromium, 524. Metallurgy, Uses, 524. Chromous Chloride, 525. Chromic Oxide, 525. Chromic Hydroxide, 525. Chromic Chloride, 525. Hydrates of Chromic Chloride, 525. Potassium Chromium Sulfate, or "Chrome Alum," 527. Potassium Chromate, 527. Potassium Dichromate, or Pyrochromate, 527. Lead Chromate, or Chrome Yellow, 527. Barium Chromate, 528. Chromium Trioxide, or Chromic Anhydride, 528. Chromyl Chloride, 528. Molybdenum, 528. Molybdium Trioxide, or Molyb- dic Anhydride, 528. Compounds of Molybdenum, 529. Molybdic Anhydride, 529. Tungsten, 530. Compounds of Tungsten, 531. Phosphotungstic Acid, 531. Uranium, 531. CHAPTER XXXI MANGANESE Group VII. Manganese, 533. Occurrence, Properties, 533. Compounds of Manganese, 534. Manganous Manganic Oxide, 534. Manganous Hydrox- ide, 535. Manganous Chloride, 535. Manganous Sulfate, 535. Manganous Sulfide, 535. Manganese Dioxide, or Black Oxide of Manganese, 535. Manganates, 536. Permanganates, 537. Potassium Permanganate, 537. Manganese Heptoxide, or Permanganic Anhydride, 538. CONTENTS xv CHAPTER XXXII IRON, COBALT, NICKEL Group VIII. Iron, 539. Occurrence of Iron, 540. Metallurgy and Iron, 540. Pig Iron, Cast Iron, 543. Wrought Iron, 544. Cementation Steel, Cast Steel, 545. Bessemer Steel, 547. Open Hearth, or Siemens-Martin Process, 548. Alloy Steels, 552. Compounds of Iron, 552. Potassium Fer- rate, 553. Ferrous Chloride, 553. Ferrous Hydroxide, 553. Ferrous Oxide, 553. Ferrous Sulfate, Green Vitriol, or Copperas, 554. Ferrous Carbonate, 554. Ferrous Chloride and Nitric Oxide, 554. Ferric Chloride, 554. Ferric Hydroxide, 555. Dialyzed Iron, 555. Ferric. Oxide, 555. Ferric Sulfate, 556. Magnetic Oxide of Iron, 556. Ferrous Sulfide, 556. Ferric Sulfide, 556. Iron Bisulfide, or Iron Pyrites, 556. Ferric Thiocyanate, 556. Co- balt, 557. Compounds of Cobalt. Oxides, 557. Cobaltous Hydroxide, 557. Cobaltous Chloride, 557. Cobalt Sulfide, 558. Cobalt Nitrate, 558. Cobalt Glass, 558. Potassium Cobaltocyanide, 558. Potassium Cobalticyanide, 558. Potassium Cobaltinitrite, 558. Cobalt Ammines, 559. Nickel, 559. Compounds of Nickel, 560. Nickel Dimethylglyoximine, 560. Nickel Car- bonyl, 561. CHAPTER XXXIII THE PLATINUM METALS Ruthenium, 563. Rhodium, 563. Palladium, 563. Osmium, 564. Iridium, 565. Platinum, 565. Platinous Chloride, 565. Chloroplatinic Acid, 565. Platinic Chloride, 566. Platinum Bisulfide, 566. INDEX . 567 A TEXTBOOK OF CHEMISTRY CHAPTER I INTRODUCTION The Nature of Scientific Knowledge. The phenomena pre- sented to our senses are so complex and varied that a complete description of each of them is impossible. It is the purpose of science to classify these phenomena and to discuss relation- ships between them which are frequently repeated and separate them from those relationships which are not repeated and which are to be considered as more or less accidental. When a given relationship between two or more phenomena is found always to exist, we conclude that the relationship is necessary or inherent in the nature of things, and it is called a law. The original thought conveyed by the name was, doubtless, that the material universe acts as it does because it was commanded to do so by some higher power, but the word has come to signify simply a statement of some constantly recurring rela- tionship between phenomena. To illustrate : we find that if we let go of an object held in the hand it will fall. The first time this is observed it is simply a fact of experience or observation. But a further examination of the relations involved leads to the general statement that any body which is not supported will fall. We may, perhaps, call this a law, but it is still a comparatively imperfect statement of the existing relations. A further study teaches us that all bodies fall, in a vacuum, at the same rate, irrespective of their size or weight and that the velocity of a falling body varies as the time during which it has fallen. These may be called the empirical laws of gravitation, that is, 1 2 i V A tSB00K OF CHEMISTRY the /laws derived m?m ^experiment and observation. For the class of phenomena to which they apply such laws have, prac- tically, as great a certainty as any individual fact which we ob- serve. A further study of the matter and especially of the motions of the earth and planets and stars has led to the more complete generalization that all bodies attract each other directly as the product of their masses and inversely as the squares of their distances. This is Newton's law of universal gravitation. When once stated, it is seen that the empirical laws stated above follow from it. A natural law, when discovered, usually requires further study and development in two directions. It calls, in the first place, for a careful examination of the logical consequences which follow from the law, and it furnishes the means of predicting in count- less cases just what will occur in given conditions where the law applies. Thus the law of gravitation is assumed, instinctively, every time that we move. It enables us to predict, accurately, the motion of a pendulum or projectile and to calculate the exact relative positions of the sun, moon and planets for a hundred or a thousand years to come. The law lies at the very basis of all of those calculations of the engineer by which he deter- mines the necessary strength for the parts of a bridge or a truss. Many other illustrations might be given of its practical impor- tance and usefulness. In the second place, a natural law suggests the need of some additional explanation. An inquiry in this direction may lead to some more fundamental and general law, as when the laws of falling bodies were found to be only special cases of the universal law of gravitation, or it may lead to the confines of our present knowledge at a point where no further progress can be made without the use of speculation or hypothesis. Following our illustration, it appears to most minds almost or quite incon- ceivable that one body should act upon another at a distance without some medium, and there have been many speculations with regard to some medium which may be the cause of gravita- tion. Several hypotheses with regard to the mechanism of the INTRODUCTION 3 action of such a medium have been proposed, but these specula- tions have not met with any notable success. With regard to such speculation there are, at present, among scientific men, and notably among chemists, two somewhat dis- tinct schools or classes. One of these takes the ground that the number of possible explanations of those parts of the universe which lie beyond our knowledge is so great that it is hopeless to attempt to find the true explanation. While admitting the value of hypotheses in stimulating and directing research, this school claims that such hypotheses can never give us any real knowledge of matters which are beyond the cognizance of our senses, and that all genuine scientific advance consists in giving a fuller description of things about which we can gain direct, positive knowledge. The other school points out that, while there are many things in the universe which must always remain beyond the possibility of direct knowledge, we can accumulate so much evidence with regard to these that it may be possible, ultimately, to give to our theories with regard to them a high degree of probability. The danger of the first attitude of mind is that the investigator will be content with a full description of phenomena and will fail to discover relations which can be under- stood only by a knowledge of matters about which we can secure only indirect evidence. The danger of the other point of view is that it may lead one to overestimate the amount of knowledge which has been acquired about unseen things and to spend time in useless speculations which would be better spent in acquiring new facts. Whichever view is accepted, the science of our time includes a knowledge of a very great number of facts, of the natural laws which express the relations between these' facts and of the theories which are our best present explanation of the laws. Subdivisions of Science. There is, properly speaking, only one science, which includes all classified, systematic knowledge. The amount of such knowledge has become so great, however, that it is customary to subdivide it into a number of parts, each of which is called a science. It should always be remembered A TEXTBOOK OF CHEMISTRY that the boundaries between various divisions are more or less arbitrary and that very many facts belong about equally to two or more of the sciences. It is also very important to understand that no one can make much progress in any one science without considerable knowledge of several others. The more important subdivisions of science are the following : abstract sciences, which deal, primarily, with forms of abstract reasoning mathematics and logic; physical sciences, dealing with the phenomena of matter and energy apart from life physics, chemistry, astronomy, mineralogy, geology; biological sciences, dealing with the phenomena of living bodies bacteri- ology, botany, zoology, paleontology; psychological sciences, dealing with the phenomena of mind and of society psychology, language, history, social science, political economy, ethics. Mathematics Abstract sciences Logic Physics Chemistry Astronomy Mineralogy . Geology Bacteriology Botany Zoology Paleontology Psychology Language History Social science Political economy .Ethics Physical Sciences. The two fundamental physical sciences are physics and chemistry. The other three mentioned, astron- omy, mineralogy and geology, are concerned with the applica- tion of the laws of physics and chemistry in studying particular bodies or substances and so are more special in their nature. Physical sciences Biological sciences Psychological sciences INTRODUCTION 5 Roughly speaking, physics treats of energy and chemistry of matter. Thus chemistry tells us of the properties and composi- tion of substances, as of water, of iron or of sulfur, of the action of substances on each other and of the changes in composition which they undergo in a great variety of circumstances. Phys- ics, on the other hand, deals with the varieties of energy, as mechanical energy, sound, heat, light, electricity, and with the transformations of each of these into other forms. As we can have no knowledge of matter except through the energy which it possesses and the effect of that energy on our senses in one way or another, and since the changes in energy which result when substances act on each other are often of great importance, the chemist can make little progress in his study without a considerable knowledge of physics. And as we have no knowledge of energy apart from matter, the physicist finds some knowledge of chemistry desirable. This interrelation be- tween the sciences is so close, also, that there is a large domain which is common to both and of which it is scarcely worth while to ask whether it belongs to chemistry or to physics. Matter and Energy. The two most fundamental concepts of physical science are matter and energy. Matter may be defined as anything which has mass, or, in its relation to the earth or to other bodies, weight. 1 Putting the same thought in quite different words, matter is anything which requires energy to set it in mo- 1 This definition of matter, which is based, of course, on Newton's conception of inertia that no body can move, if at rest, or change the direction or velocity of its motion, if moving, unless it is acted on by some external force is not entirely satisfactory. It has been shown that the mass of a body is changed by a change in its velocity, though the change is inappreciable until the velocity approaches that of light. (D. F. Comstock, J. Amer. Chem. Soc., 30, 683). In such a case it is usually better to retain the older, simple definition, frankly recognizing that it is imperfect. It is character- istic of a scholastic rather than a scientific attitude of mind to be much troubled because a definition is imperfect or incomplete. The scientific worker sees that our knowledge is incomplete in every direction and is constantly developing. The definitions are simply a means of conveying this incomplete knowledge to others and of at- tempting to discover the most fundamental conceptions of science. We succeed best in both directions by keeping these definitions simple. 6 A TEXTBOOK OF CHEMISTRY tion or to change its rate of motion. Energy is usually defined, on the basis of etymology, as anything which can do work. A more satisfactory definition is that energy is anything which may set matter in motion or change its rate of motion. The most important forms of energy are mechanical energy, sound, heat, light, electrical energy and chemical energy. Conservation of Matter and Energy. A superficial observation of many phenomena in nature appears to show that under some conditions, especially in burning, matter is destroyed, and that under other conditions, as in the rusting of iron, matter in- creases in weight. It is comparatively easy to show, however, that a part of the air, which has weight, takes part in these pro- cesses and that while a candle, for instance, seems to be destroyed in burning, water and carbon dioxide are formed by its combus- tion ; and if these are absorbed by soda lime and weighed, the sum of their weights is very considerably greater than the weight of the candle which has been burned. Still more careful experi- ments will show that the products of combustion weigh exactly the same as the weight of the candle and the weight of the por- tion of the air (oxygen) with which it has combined. The question whether there is any change in the weight of matter during a chemical reaction is so fundamental that one chemist (Landolt) has considered it worth while to give ten years of most careful and painstaking work to its study. His con- clusion is that in the cases which he studied no change so great as the one millionth part of the weight of the substances which reacted with each other occurred. We say, therefore, that no method is known by which we can create, or destroy, matter. This is known as the law of conservation, or indestructibility of matter. If we place a wheel with vanes in a can of water and wind around its axle a cord tied to a weight in such a manner that as the weight falls the wheel will revolve and stir the water, we shall find that the temperature of the water will rise. If the experi- ment is carefully performed, it will be found that a weight of one kilogram falling 427 meters will raise the temperature of a INTRODUCTION 7 kilogram of water one degree. On the other hand, if the steam from a boiler is caused to drive the piston of a steam engine which is pumping water or doing other work, it is found that a part of the heat of the steam disappears and that the heat which can no longer be found in the exhaust steam or anywhere about the engine corresponds accurately to the amount of work which the engine performs, and that the ratio is exactly the same as that found between the falling weight and the rise in tempera- ture of water stirred by the paddle wheel. The energy of the engine may be used to drive a dynamo which will furnish an elec- tric current ; the electric current may be used to decompose a chemical compound, giving substances which contain more chem- ical energy than the compound ; and these substances, in turn, may be recombined, giving out heat in the process. Each form of energy which we know may be transformed into some other, and there is always an exact relation between the quantity of energy of one kind which disappears and the quantity of other kinds of energy which takes its place. This is the law of the conservation of energy. It might be called the law of the inde- structibility of energy. Pure Substances and Mixtures. As has been stated, chemistry deals, primarily, with the properties and composition of sub- stances. 1 If we examine certain substances, such as pure water or gold, we find that they are alike throughout their whole mass, or that they are homogeneous. We find that every sample of such a substance which we examine melts or freezes at exactly the same temperature ; and that if it boils without decomposition, it will always boil at the same temperature under atmospheric pressure. We find, too, that if the substance is a liquid or gas, the density or specific gravity is always the same under the same condition of temperature and pressure. In spite of the large 1 The distinction between the words body and substance should be carefully observed. Body always refers to some definite, con- crete thing, as a heavenly body, speaking of the sun or a star, a body of ore, etc. Substance, on the other hand, refers to some par- ticular kind of matter, as water or gold. A given piece of gold might be called a body, but gold, in general, is a substance. 8 A TEXTBOOK OF CHEMISTRY number of substances which are known to exist (more than one hundred thousand) it is possible to identify many of these with practical certainty by the examination of a comparatively small number of their properties. For instance no other substance has the same freezing point, boiling point and density as water. Contrasted with pure substances, as water or gold, most sub- stances which we meet in daily experiences are mixtures. For example, if we take a cereal, as wheat, and powder it, as is done in the milling process, a portion will pass through fine bolting cloth, while another portion, the bran, will not. If the portion which passes the bolting cloth is warmed gently, it loses weight, and it can readily be shown that the loss is almost wholly due to the escape of water, which may be condensed and identified by its freezing point and boiling point. From the dry flour ether will dissolve an oil or fat which will be left behind on evaporating the ether. If the portion which remains is kneaded between the fingers in a stream of running water, a fine white powder, consist- ing mainly of starch, will be washed away, while a residue, called gluten, which consists largely of proteins, will remain. The processes described show clearly that the cereal is a very complex mixture of many different substances, but of the substances sepa- rated only the water and starch can be considered as even approx- imately pure substances. A more careful examination of the bran or oil or gluten will show that each of these is still a mixture. Preparation of Pure Substances. A large part of the work which must be done in the study of chemistry consists in the separation and characterization of pure substances. The most common means used for this purpose are treatment of mixtures with solvents, crystallization and distillation. Thus if we have a mixture of sugar and sand, the sugar may be easily separated by dissolving it in water and pouring off or filtering the solution from the sand. From a brine which contains other substances in solution along with salt, the salt may be obtained nearly pure by evaporating it till the salt separates in crystals. For some reason particles of the same kind separate from a solution on evaporation or, frequently, on cooling a hot solution, in definite, INTRODUCTION geometrical forms called crystals, and when they separate in this manner they usually exclude other substances which may be present. By repeated distillation of a mixture of alcohol and water, collecting the lower boiling portions by themselves each time, nearly pure alcohol can be separated from water. When a volatile substance contains a nonvolatile one in solution, the separation by distillation is much easier and , , more complete. Elements and Compounds. If the red oxide of mercury is heated in a small tube, metallic mercury will distill away, while a glowing splinter held at the mouth of the tube will burst into flame. The heat causes the decomposition of the oxide of mercury into mercury and a gas which supports com- bustion better than air, and which is called oxygen. An electric current passed between two strips of platinum immersed in a solution of sulfuric acid in water in the apparatus -^ shown in Fig. 1 will cause the separation of two gases, oxygen and hydrogen. As it can be shown that the amount of sulfuric acid remains unchanged, it is evident that the gases are formed by the decomposition of the water; and this view can be confirmed by burning the mixture of oxygen and hydrogen and regenerat- ing the water. While oxide of mercury can be decomposed into mercury and oxygen, and water may be decomposed into oxygen and hydrogen, no one has thus far succeeded in decom- posing mercury or oxygen or hydrogen. Substances like these, which it has not been found possible to decompose, are called elements. 1 Substances which can be separated into two or 1 This definition is not wholly satisfactory, since it has been found that radium, which has all of the other properties of an ele- ment, decomposes spontaneously into helium and a whole series of other elements. It seems best, however, to retain the simple definition, but also it is best to consider radium as an element. J Fig. 1 10 A TEXTBOOK OF CHEMISTRY more parts, neither of which can be converted into the other, or which can be prepared by the union of two or more elements, are called compounds. Elements sometimes exist in two or three different forms, but these may always be converted each into the other. Only about eighty elements have been positively identified. The names of these, together with their symbols and atomic weights, are given in the following table : ATOMIC WEIGHT 27.1 120.2 39.88 74.96 137.37 208.0 11.0 79.92 112.40 132.81 40.07 12.00 140.25 35.46 52.0 58.97 93.5 63.57 162.5 167.7 152.0 19.0 157.3 69.9 72.5 9.1 197.2 3.99 163.5 1.008 114.8 126.92 193.1 1 Also called niobium, Nb. 2 Often given as beryllium, Be. SYMBOL Aluminium . Al Antimony . . Sb Argon . . . A Arsenic . . As Barium . . . Ba Bismuth . Bi Boron . . . B Bromine . Br Cadmium . . Cd Caesium . Cs Calcium . Ca Carbon . . . C Cerium . . . Ce Chlorine . Cl Chromium . . Cr Cobalt . . . Co Columbium 1 . Cb Copper . . . Cu Dysprosium Dy Erbium . . . Er Europium . . Eu Fluorine . F Gadolinium . Gd Gallium . Ga Germanium . Ge Glucinum 2 . . Gl Gold . . . . Au Helium . . . He Holmium . . Ho Hydrogen . . H Indium . . . In Iodine . . . I Iridium . . . Ir ATOMIC SYMBOL WEIGHT Iron . ... Fe 55.84 Krypton . . Kr 82.92 Lanthanum La 139.0 Lead . . . Pb 207.10 Lithium . . Li 6.94 Lutecium . . Lu 174.0 Magnesium Mg 24.32 Manganese Mn 54.93 Mercury , . Hg 200.6 Molybdenum . Mo 96.0 Neodymium Nd 144.3 Neon . . . Ne 20.2 Nickel . . . Ni 58.68 Niton . .- . Nt 222.4 Nitrogen . . . N 14.01 Osmium . . Os 190.9 Oxygen . . . 16.00 Palladium . . Pd 106.7 Phosphorus P 31.04 Platinum . . Pt 195.2 Potassium K 39.10 Praseodymium Pr 140.6 Radium . . . Ra 226.4 Rhodium Rh 102.9 Rubidium . . Rb 85.45 Ruthenium Ru 101.7 Samarium . Sa 150.4 Scandium . . Sc 44.1 Selenium . . Se 79.2 Silicon . . . Si 28.3 Silver . . . Ag 107.88 Sodium . . Na 23.00 Strontium . . Sr 87.63 INTRODUCTION 11 ATOMIC ATOMIC SYMBOL WEIGHT SYMBOL WEIGHT Sulfur . . . S 32.07 Uranium . U 238.5 Tantalum . . Ta 181.5 Vanadium . . V 51.0 Tellurium . . Te 127.5 Xenon . . . Xe 130.2 Terbium . Tb 159.2 Ytterbium Thallium . . Tl 204.0 (Neoytter- Thorium . Th 232.4 bium) . . Yb 172.0 Thulium . . Tm 168.5 Yttrium . Y 89.0 Tin . . . . Sn 119.0 Zinc . . . . Zn 65.37 Titanium . . Ti 48.1 Zirconium . . Zr 90.6 Tungsten . . W 184.0 The symbols are either the first letter or the first letter together with some other characteristic letter of the name of the element. With few exceptions symbols are derived from the English names and the symbols readily suggest the names. The exceptions are : Antimony, Sb, Stibium Potassium, K, Kalium Gold, Iron, Lead, Mercury, Au, Aurum Fe, Ferrum Pb, Plumbum Hg, Hydrargyrum Silver, Ag, Argentum Sodium, Na, Natrium Tin, Sn, Stannum Tungsten, W, Wolfram For all of these except the last the symbols are derived from the Latin names. The elements vary greatly in their relative abundance. Of that portion of the earth which we are able to examine it is estimated that oxygen forms nearly one half of the total weight and sili- con one fourth. The percentage amounts of the twelve most common elements in the surface of the earth to a depth of ten miles, including the ocean and the atmosphere, are estimated as follows : l PER CENT PER CENT Oxygen, 49.78 Potassium, 2.28 Silicon, 26.08 Magnesium, 2.24 Aluminium, 7.34 Hydrogen, 0.95 . Iron, 4.11 Titanium, 0.37 Calcium, 3.19 Chlorine, 0.21 Sodium, 2.33 Carbon, 0.19 99.07 1 F. W. Clarke, Data of Geochemistry, p. 32. 12 A TEXTBOOK OF CHEMISTRY Some elements which form only a very small part of the whole are very important, especially nitrogen, phosphorus and several of the metals which are not included in the above table. How Pure Substances are distinguished from Mixtures. Law of Constant Proportion. A very large part of our knowledge of chemistry depends on the preparation of pure substances and on the determination of the properties and composition of these. It is, therefore, important to understand how we may distinguish between pure substances and mixtures. The first characteristic of a pure substance is that it must be homogeneous so long as it exists in one state of aggregation, that is, so long as it is all solid, all liquid or all gaseous. Second, it must have a constant melting point and boiling point, if it melts and boils without de- composition, and the specific gravity or density and other physi- cal properties must be invariable under the same conditions. 1 Third, a pure substance must always show the same conduct toward any other substance which may dissolve it or act upon it chemically, provided that the conditions are the same. A very careful examination of a large number of substances which have the characteristics just given in the highest degree has demonstrated that such substances are absolutely constant in composition. This is the law of constant proportion and may be stated thus : A pure substance always contains the same ele- ments in the same proportion by weight. Thus pure water always contains hydrogen and oxygen in the proportion of 1 to 7.94 parts by weight. This law has been tested by a large amount of most careful and painstaking work, and the more careful the work has been the more accurately has the law been found true, so that we may consider it as one of the most absolutely perfect laws of nature. Since a very large number of substances which fulfill the first three requirements of a pure substance are in- variable in composition, this constancy of composition is con- sidered as a fourth characteristic of a pure substance. It is a 1 The density of some solids and especially of metals may vary slightly according to the treatment to which they have been sub- jected. INTRODUCTION 13 characteristic of very great importance and one which is fre- quently used to determine whether a given substance is pure or not. Inductive Reasoning. It may seem at first that the use of constancy of composition as a means of determining whether a substance is pure or not is due to reasoning in a circle, or, as it is commonly called, is " begging the question." We say first that a pure substance has a constant composition and then that because a substance has a constant composition it is pure. The criticism would be justified if constancy of composition were the only characteristic applied to decide whether a substance is pure or not. But the first three characteristics mentioned above are the ones which will appeal to any one as being dictated by common sense. When we find that a very great number of substances having these characteristics are also constant in com- position, we come to the conclusion that there is some inherent, necessary connection between this fourth characteristic and the other three, and that when a given substance does not have this characteristic it probably lacks some of the other three as well. Such a conclusion is said to be reached by inductive reasoning. The truth of such a conclusion can never be absolutely proved any more than we can prove that the sun will rise to-morrow morning. But we 1 may reach practical certainty by means of such conclusions and may properly use them as the basis for further reasoning. Law of Combining Weights. If we select a series of com- pounds in such a manner that each compound has an element contained in the preceding and another contained in the follow- ing compound, it will be found that whenever the same element recurs the proportion of the element which combines with other elements will always be the same or some exact multiple or submultiple of the first proportion. This will be more clear from the following series of compounds * l 1 Here and elsewhere whole numbers are used for greater sim- plicity. The exact values will be found in the table of atomic weights, p. 10. 14 A TEXTBOOK OF CHEMISTRY Water Cuprous Cupric Hydrogen Hydrochloric Oxide Sulfide Sulfide Acid H:O 0:Cu Cu :S S:H H:C1 1:8 8 : 63.6 63.6 :32 32:2 2:71 Ferrous Ferrous Sulfur Sodium Sodium Sodium Chloride Oxide Dioxide Sulfide Chloride Chlorate Cl:Fe Fe:O 0:S S:Na Na:Cl Na : Cl : O 71:56 56:16 16:16 16:23 23 : 35.5 23 : 35.5 : 48 In this series of compounds hydrogen has been chosen as the starting point and has been given a value of 1. If oxygen had been chosen and had been given a value of 100, as was at one time proposed, the other numbers would all be different but exactly the same ratios between the different numbers for the same element would be found throughout the series. It is seen that the values for hydrogen are 1 and 2, for oxygen 8, 16, 32, and 48, for sulfur 32 and 16, for chlorine 71 and 35.5, the larger numbers for each element being in every case exact multiples of the smallest number for the element. This table might be extended to include all pure substances which have been analyzed. The law of combining weights stated above may be expressed more briefly as follows : A number may be selected for each element which represents the proportion of the element which enters into combination with other elements. The Atomic Theory. The laws of constant proportion and of combining weights find a very satisfactory explanation in the atomic theory, which was proposed by Dalton at the beginning of the nineteenth century. According to this theory the chemi- cal elements are composed of very small particles or atoms, the atoms of the same element being all alike in properties and in weight, while the atoms of different elements are different. If we suppose further that compounds are always formed by the union of atoms of different elements, it is evident that the ratio between the weights of the elements in a compound must be the same as the ratio between the weights of the atoms composing the smallest particle of the compound. Thus if the smallest particle (molecule) of water which can exist contains two atoms INTRODUCTION 15 of hydrogen united to one atom of oxygen and an atom of oxygen weighs 16 times as much as an atom of hydrogen, any quantity of water, whether large or small, must contain hydrogen and oxygen in the proportion of two to sixteen. If, for instance, 1000 atoms of oxygen could be mixed with 2001 atoms of hydro- gen, after combination had taken place one atom of hydrogen would be left uncombined. In this way the theory explains very satisfactorily the law of constant proportion. It explains equally well the law of combining weights, for these combining weights must be directly connected with the relative weights of the atoms of the elements. In accordance with the atomic theory we may select some ele- ment as our unit for atomic weights, and by determining the amounts of other elements which combine with a given weight of this element and the number of atoms of each element in the compounds formed, we can determine the weights of the atoms of the other elements as compared with the weight of an atom of the element taken as a unit. Thus if we take hydrogen as our unit and find that hydrochloric acid contains one part of hydro- gen to 35.5 parts of chlorine, and can show, further, that a mole- cule of hydrochloric acid contains one atom of chlorine and one atom of hydrogen (p. 92), the atom of chlorine must be 35.5 times as heavy as the atom of hydrogen and we say that the atomic weight of chlorine is 35.5. Or if we find that water con- tains 8 parts of oxygen for one of hydrogen 1 and a molecule of water contains one atom of oxygen and two atoms of hydrogen, the atom of oxygen must be 16 times as heavy as the atom of hydrogen and we say that the atomic weight of oxygen is 16. The atomic theory, which could be considered* as scarcely more than a doubtful hypothesis when it wag first proposed by Dalton, became the central, guiding principle in the development of the science of chemistry during the nineteenth century ; and evidence in its favor has been accumulated from very many dif- ferent and independent directions, so that, now, the actual existence of atoms and molecules can scarcely be doubted. 1 The exact composition of water will be considered later. 16 A TEXTBOOK OF CHEMISTRY We even have a half dozen different ways of estimating the actual weight of an atom and the estimates agree fairly well. These estimates give the number of molecules in a cubic centimeter of air under standard conditions as about 2.71 X 10 19 or nearly thirty million million millions (Millikan. See also Rutherford, Presidential Address before Section A of the British Association at the Winnipeg meeting). Sir William Thomson (known later as Lord Kelvin) once used the illustration that if a drop of water could be magnified to the size of the earth the molecules would be larger than small shot and smaller than cricket balls. This is something the same sort of an estimate as if we were to say that a certain animal is the size of a dog. Our knowledge of the space filled by a molecule is now much more accurate. Selection of Atomic Weights. In the series of compounds used to illustrate the law of combining weights, the combining weights of oxygen are 8, 16 and 48. If the table were extended, the values 4 and 32 might be found in other compounds, and almost any multiple of 8. It is evident that if we start with hydrogen and give it an atomic weight of 1 (see, however, p. 72), only one of these various combining weights can be the true atomic weight of oxygen. Since the atoms and molecules are so small as to be beyond the possibility of direct observation, it seemed for a long time impossible to select the true atomic weight from among the various possible combining weights. Dal ton thought it most natural to suppose that the molecule of water contains one atom of hydrogen and one atom of oxygen and on this basis the atomic weight of oxygen would be 8 instead of 16. The reasons for considering that the true atomic weight of oxygen is 16 and thte methods used in selecting what are believed to be true atomic weights will be considered later (p. 92). Formulas. The atomic weights selected for the elements used to illustrate the law of combining weights are: 1 H = 1, Cu = 63.6, S = 32, Cl = 35.5; Fe = 56; Na = 23. If we 1 These values are rounded off. The accepted values are : H = 1.008, O = 16.00, Cu = 63.57, S = 32.07, Cl = 35.46, Fe = 55.84, Na = 23.00. INTRODUCTION 17 express the composition of water, cuprous oxide and cupric sulfide in such a manner as to avoid the use of fractions of atomic weights, the ratios for these compounds become : Water, H : O = 2:16 Cuprous oxide, O : Cu = 16 : 127.2 Cupric sulfide, Cu : S = 63.6 : 32 In accordance with the atomic theory it follows from these ratios that a molecule of water contains two atoms of hydrogen for each atom of oxygen, that a molecule of cuprous oxide con- tains two atoms of copper for one of oxygen and cupric sulfide contains the same number of atoms of sulfur as of copper in its molecule. It has been found very convenient to express these relations by using the symbol of each element to stand for one atom of the element and so to write formulas for compounds, using numerical subscripts to designate the number of atoms of each element contained in a molecule of the compound. The formulas for the compounds are: H 2 O, Cu 2 O, CuS. Since a formula is always based on the proportion by weight of each element contained in the compound, it tells us not only how many atoms of each element are contained in a molecule of the compound, but it also tells us the exact composition of the com- pound by weight. Thus the formula H 2 SO 4 , for sulfuric acid, means that a molecule of sulfuric acid contains two atoms of hydrogen, one atom of sulfur, and four atoms of oxygen ; but it also means that the acid is composed of 2 parts by weight of hydrogen, 32 parts of sulfur and 64 parts of oxygen. Strictly speaking, the formulas given should be written H 2 Oi, Cu 2 O b CuiSi and H 2 SiO 4 , but by common consent the subscript 1 is always understood when no subscript is given. What are the formulas of the other compounds mentioned on p. 14? Composition of Pure Substances. From what has been stated we may derive a fifth characteristic of a pure substance. The composition of a pure substance can always be expressed by exact multiples of the atomic weights of the elements compos- 18 A TEXTBOOK OF CHEMISTRY ing it. This may be considered as still another way of stating the law of combining weights. The law has been tested by the analysis of thousands of compounds, and, like the law of constant proportion, it is one of the perfect laws from which no deviation has been discovered. Study of Chemistry. To obtain a knowledge of the elements of chemistry it is necessary to become acquainted with a large number of facts about the substances with which the science deals, but it is still more important to understand the relations connecting these facts with each other and the fundamental laws and theories by which the facts are grouped together and explained. Success in the study depends especially on the abil- ity to learn new facts in their relation to those which have al- ready been acquired and on the cultivation of a logical as dis- tinguished from an arbitrary memory. Formulas, especially, should be derived, whenever possible, from the formulas of other compounds of the same elements, and not learned individually, except in the earliest portion of the study. Chemical equations should be written on the basis of a knowledge of the reacting substances and of the products of the reaction, and should never be learned by brute memory. In the systematic treatment of the subject the more common elements will be considered first and under each element the compounds of that element with each of those previously studied will be mentioned so far as this is desirable. See p. 132. CHAPTER II OXYGEN SYMBOL, O. ATOMIC WEIGHT, 16. Occurrence. Oxygen is the most abundant and one of the most important of all the elements. It forms about one fifth of the volume of the air, eight ninths of the weight of water and nearly one half the weight of the mineral substances which com- pose the crust of the earth. Oxygen is found in all living bodies and is a constituent of a larger number of compounds than any other element except carbon. Preparation. 1. When metallic mercury is heated at the right temperature in contact with the air, it is slowly converted B G Fig. 2 into a bright red compound called oxide of mercury. The French chemist, Lavoisier, carried out the experiment in the apparatus shown in the figure, and proved that after the mercury had been heated several weeks the air no longer decreased in volume and he concluded that this was because the oxygen of the air had 19 20 A TEXTBOOK OF CHEMISTRY all been removed by combination with the mercury. He then collected the oxide of mercury and heated it to a higher tempera- ture till it was all decomposed into mercury and gaseous oxygen. The volume of oxygen was the same as the decrease in volume of the air during the heating in contact with the mercury. The quantitative relation between the mercury, oxygen and oxide of mercury may be very briefly expressed by means of the symbols for the elements, as follows : Hg + O HgO Mercury Oxygen Oxide of Mercury HgO Hg + O Oxide of Mercury Mercury Oxygen Since the atomic weight of mercury is 200 and the atomic weight of oxygen is 16, the first equation means that 200 parts by weight of mercury combine with 16 parts by weight of oxy- gen to form 216 parts of oxide of mercury ; and the second equa- tion means that 216 parts of oxide of mercury decompose into 200 parts of mercury and 16 parts of oxygen. It will be noticed that the symbols of two elements placed side by side represent a compound, while a symbol by itself represents a free element. 2. The portion of air which was not absorbed by the mercury was chiefly nitrogen and formed about four fifths of its volume. Liquid nitrogen boils at 194, while liquid oxygen boils at 182.5. If liquid air is allowed to boil, the nitrogen goes off, chiefly, at first, and the gas which comes off toward the end is nearly pure oxygen. In this way the oxygen and nitrogen may be separated very much as alcohol and water are separated by distillation. Oxygen prepared in this manner is compressed into strong steel cylinders for medicinal and other uses. 3. When potassium chlorate is heated, it melts and begins to decompose slowly into potassium chloride and oxygen. Potassium chlorate has the composition represented by the formula KC1O 3 . OXYGEN 21 Potassium, K = 39.1 parts or 31.90 per cent Chlorine, Cl = 35.46 parts or 28.93 per cent Oxygen, 3 O = 48. parts or 39.17 per cent Total 122.56 parts or 100. per cent The decomposition may be represented by the equation : KC1O 3 = KC1 + 3O Potassium Potassium Oxygen Chlorate Chloride If the potassium chlorate is mixed with one fourth of its weight of finely powdered manganese dioxide, MnO 2 , the de- composition will begin at a much lower temperature and pro- ceed more rapidly than when the potassium chlorate is heated alone. If the residue in the retort, after the decomposition is complete, is treated with water, the potassium chloride will dissolve, while the manganese dioxide will remain undissolved and may be readily separated from the solution by filtration. If the manganese dioxide is examined, it will be found that it has not changed in composition or amount. We may, there- fore, write the equation : KC10 3 + Mn0 2 = KC1 + 3O + MnO 2 Potassium Manganese Potassium Oxygen Manganese Chlorate Dioxide Chloride Dioxide 4. When fused sodium peroxide containing a very little copper oxide 1 is dissolved in water, it gives sodium hydroxide and oxygen : Na 2 2 + H 2 O = 2NaOH + O Sodium Water Sodium Oxygen * Peroxide Hydroxide 2 As the copper oxide is left unchanged at the end and as the reaction will take place, though more slowly, in its absence, we may omit it in writing the equation. 1 The substance is known commercially as " oxone." 2 The student should notice the connection between the name and the composition of sodium hydroxide. Many similar compounds containing oxygen and hydrogen are called hydroxides. 22 A TEXTBOOK OF CHEMISTRY Collection and Storage of Gases. Gases which are only slightly soluble in water and which are not required in a high state of purity are conveniently collected and stored in a gasometer of the form shown in Fig. 3. If the stopcocks A and B are opened and the cap C screwed on, water placed in the cup above will run into the body of the gas- ometer till it is filled with water and all of the air is expelled. Now on closing the stop- cocks the gasometer will still remain filled with water after removing the cap at C. By means of a tube inserted through C gas may be in- troduced and will fill the gasometer, displacing the water, which will flow out of C by the side of the tube delivering the gas. After filling the gasometer and replacing the cap at C, on opening the stopcock A, water will run into the body of the gasometer and the gas may be drawn off as desired through B. Properties of Oxygen. Oxygen is a colorless, odorless and tasteless gas. 1 The weight of one liter at and under a pressure of 760 millimeters of mercury (about the average atmospheric pressure at sea level) is 1.429 grams. 2 Under the same conditions of temperature and pressure it is about one tenth heavier than the same volume of air. The most striking property of oxygen is the vigor with which it supports combustion. All substances which burn in air burn much more rapidly and vigorously in oxygen. A splin- ter of wood having a live coal on the end will burst into flame, 1 These statements refer, of course, to the ordinary form of the element. Ozone, which is another form of oxygen, is colored and has a strong odor (p. 98). 2 At 45 latitude. It is slightly less at lower latitudes because the pressure of 760 mm. of mercury is less and the density of the gas is less. At the latitude of New York the weight of one liter of oxygen is 1.42845 grams. If, however, the reading of the barometer is cor- rected for latitude and altitude, the weight of one liter of oxygen is 1.429 grams at any place. Fig. 3 OXYGEN 23 if thrust into the gas. A piece of charcoal, barely ignited, will glow intensely and be surrounded by a pale blue flame, scarcely visible in the intense light of the glowing mass. The product of the combustion is carbon di- oxide, a colorless gas. Sulfur burns with a brilliant blue flame, giving sulfur dioxide, also a colorless gas. Phosphorus burns with an intense white light, giving a white, solid com- pound, phosphorus pentoxide. A coil of iron wire or a steel watch spring to which is at- tached a string that has been dipped in melted paraffin, may be set on fire and will burn in oxygen (Fig. 4), throwing off brilliant sparks and forming white-hot, molten globules of the magnetic oxide of iron, which will drop off from time to time. Fig. 4 The equations which represent the quantitative relations in these experiments are : = CO 2 Carbon Dioxide = SO 2 Sulfur Dioxide = P 2 5 Phosphorus Pentoxide = Fe 3 4 Magnetic Oxide of Iron Oxygen and Acid Properties. If sulfur dioxide or phosphorus pentoxide is dissolved in water, the solution obtained will have a sour taste, and acid properties. Many other compounds of nonmetallic elements with oxygen combine with water in a similar manner to form acids, and it is because of this that the name oxygen, meaning " acid former," was first given to the element. When the name was given, it was supposed that all acids contain oxygen, but it was discovered later that this is not the case. c 4- 2O Carbon Oxygen s + 20 Sulfur Oxygen 2P + 5O Phosphorus Oxygen 3Fe + 40 Iron Oxygen 24 A TEXTBOOK OF CHEMISTRY Combustion. Effect of Concentration on a Chemical Reac- tion. The similarity between ordinary combustion and the burning of substances in oxygen is apparent. A more careful study of the matter has shown that exactly the same compounds are formed when charcoal, sulfur or phosphorus burn in the air as are formed when they burn in oxygen, and even in the case of iron, the magnetic oxide formed by burning the steel watch spring has just the same composition as that of the scale formed when white-hot iron is exposed to the air. The burning of iron in air can also be shown by sprinkling fine iron filings through a flame. About four fifths of the air by volume is nitrogen. This does not combine with the burning substances, and by its presence it moderates the action of the oxygen, partly by diluting it, partly because it must be heated to the same temperature as the other substances, and this absorbs a large part of the heat of the reaction and so lowers the temperature to which the burning substance is heated. These facts illustrate two prin- ciples of almost universal application in chemistry : first, that the speed of a chemi- cal reaction is increased by increasing the concentration of one of the reacting sub- stances, here the concentration of the oxygen; and, second, that the speed of a reaction is affected by the temperature and is greater at high temperatures than at low ones. We shall find later that the first of these principles can be stated in the form of an accurate, quantitative law, but the phenomena of ordi- nary combustion are not well suited for a quantitative study of this kind. Kindling Temperature. If we place on an iron plate smail pieces of phosphorus, sulfur and charcoal, it will be found on warming the plate that phosphorus takes fire at a quite low temperature, the sulfur at a moderate heat, while the charcoal Fig. 5 CALORIMETER 25 will not burn till the plate is nearly red-hot. The temperature at which combination of a substance with the oxygen of the air is sufficiently rapid so that it takes fire is called the kindling temperature. The temperature rises rapidly from the heat of the reaction as soon as the substance is kindled. Well-known appli- cations of the gradations of kindling temperature are the old- fashioned sulfur match and the methods commonly used in kin- dling a fire. Kindling temperature is not a satisfactory measure of the affinity of a substance for oxygen nor is it closely con- nected with the heat generated on combination with oxygen. Heat of Combustion. Calorimeter. When substances burn, a part of the chemical energy of the burning substance and of the chemical energy of the oxygen is converted into heat. The amount of energy transformed into heat when one gram of the substance burns, or, for most scientific pur- poses, the energy obtained from one gram atom or gram molecule of the substance, is called its heat of com- bustion. By gram atom is meant as many grams of the substance as there are units in the atomic weight as 12 grams of carbon or 31 grams of phosphorus. By gram molecule is meant as many grams of the sub- stance as there are units in the mo- lecular weight, as 44 grams of carbon dioxide (12 grams of carbon -f- 32 grams of oxygen). The heat of combustion is de- termined in an instrument called a calorimeter, the most accurate form being known as a bomb- calorimeter because the combustion is carried Fig. out in an atmosphere of compressed oxygen in a strong, tightly closed bomb, which is immersed in water. The substance is 26 A TEXTBOOK OF CHEMISTRY placed in a small crucible within the bomb and is ignited by means of a fine iron wire, which is heated for a moment by an electric current. The weight of the substance, the weight of water surrounding the bomb and the temperature of the water before and after the substance is burned are accurately de- termined. There are, of course, many other details about the apparatus and manipulation which need not be described here. (See Atwater, J. Am. Chem. Soc. 25, 659, and Richards and Jesse, ibid. 32, 268.) The amount of heat required to raise the temperature of one gram of water one degree at 15 is called a calorie (see p. 33). This is often called the small calorie and designated by the abbreviation cal. to distinguish it from the large Calorie, which is the amount of heat required to raise the temperature of a kilogram of water one degree and which is designated by the ab- breviation Cal. * The corresponding unit of the English system is the British Thermal Unit (B. T. U.), the heat required to raise the tempera- ture of a pound of water 1 F. Since 1 Kg. = 2.204 Ib. and 1 F. = | of 1 C., 1 Cal. = 3.968 B. T. U. Heat of combustion ex- pressed in calories per kilogram, however, is reduced to British Thermal Units per pound by multiplying by - . From the re- sults of determinations with the calorimeter it is possible to cal- culate how many grams of water can be raised one degree in tem- perature by burning one gram or one gram atom of a substance and this will be the heat of combustion of the substance in calories. 1 The heats of combustion for the substances mentioned in this chapter are : 1 Since the amount of heat required to raise the temperature of one gram of water one degree varies slightly with the tempera- ture, it is necessary in accurate work to define the temperature at which the experiment is performed. A mean temperature of 15 is most often used. On account of the variability of the calorie it has been proposed to use the joule as a unit. The small calorie is equal to 4.182 joules at 15. See G. N. Lewis, Journal of the American Chemical Society, 35, 4 (1913). CHEMICAL ENERGY 27 For one gram of carbon (charcoal), 8080 calories For one gram of sulfur (rhombic), 2190 calories For one gram of phosphorus (yellow), 5970 calories For one gram of iron (to Fe 3 O 4 ), 1616 calories For one gram of mercury (to HgO), 105 calories For one gram atom of carbon (charcoal), 96,960 calories 1 For one gram atom of sulfur (rhombic), 70,180 calories For one gram atom of phosphorus (yellow), 185,000 calories For one gram atom of iron (to Fe 3 O 4 ), 90,200 calories For one gram atom of mercury (to HgO), 21,000 calories The Nature of Chemical Energy. The amount of energy liberated by a burning substance is very large. If it were possible completely to transform the energy liberated by burning a pound of good coal into mechanical energy, it would lift a ton weight over 4000 feet. It is a very good engine which will con- vert ten per cent of the energy of the coal burned under its boiler into useful work, but, in spite of the fact that more than ninety per cent of the energy of the coal is dissipated and lost, the total energy is so enormous that the steam engine is commercially economical. A very natural question which arises here is; What is the source of the energy which suddenly appears as heat when par- ticles of carbon and oxygen unite to form carbon dioxide ? Is there some motion within the particles of oxygen and carbon which is transformed into heat when they unite? or do the particles collide when their mutual attraction brings them to- gether, somewhat as a meteor collides with the earth? For these questions there are, at present, no answers, and specula- tions about them are of very little value till some one can dis- cover some sort of experimental evidence bearing upon them. 1 In joules these are : For one gram atom of carbon (charcoal) to CO 2 , 405,700 joules For one gram atom of phosphorus, to PzO&, 773,000 joules For one gram atom of sulfur, to SO 2 , 293,000 joules For one gram atom of iron, to Fe 3 O 4 , 3 77,000 joules For one gram atom of mercury, to HgO, 87,800 joules 28 A TEXTBOOK OF CHEMISTRY It is well, however, to recognize how imperfect and fragmentary our knowledge is and that there are hundreds of questions like these for which we have no answer. It is also well, at times, to ask such questions and consider whether there is any tangible method of attacking the problem, for, while the explanation seems beyond our grasp, at present, many similar problems which would have seemed beyond the possibility of a solution one hun- dred years ago have been solved. Catalysis. It has been pointed out that when manganese dioxide is mixed with potassium chlorate the latter decomposes at a lower temperature or more rapidly than when the chlorate is heated by itself, but that the manganese dioxide is left un- changed in the end. When a substance acts in this manner by its mere presence, causing a reaction or decomposition to take place at a lower temperature or more rapidly, it is called a catalytic agent and the process is called catalysis. * A study of this particular case makes it seem probable that the oxygen is at first transferred from the potassium chlorate to the manganese dioxide and that the compound of the manganese dioxide with the oxygen decomposes at a lower temperature than the potassium chlorate. Such an explanation seems, at first, paradoxical, for it seems to imply that manganese dioxide has a greater affinity for oxygen than potassium and chlorine have and so can take the oxygen away from the potassium chlorate, while in the resulting compound the affinity of the manganese dioxide for oxygen seems to be less than that of potassium and chlorine for oxygen, because the decomposition of the manganese compound occurs at a lower temperature than that required for the decomposition of potassium chlorate. A partial explanation of this seeming paradox is found in the fact that very many substances act upon each other chemically at a much lower temperature than that at which either decom- poses into its elements. It is also true that the stability of a compound is not an accurate measure of the affinity between the elements of which it is composed. The affinity between the elements of a compound which decomposes at 200 is not neces- CHEMICAL AFFINITY. NOMENCLATURE 29 sarily less than that between the elements of a compound which decomposes at 400. Still further, it is by no means always true that when an element is transferred from one compound to another its affinity for the element with which it combines is greater than that for the element which it leaves. These ques- tions will be considered further later ; but it is well, at the outset, to avoid certain misconceptions which are very liable to arise . because the facts of chemistry are often so very different from what our first and most natural idea of chemical affinity would lead us to expect. Chemical Affinity. The term affinity has been used in the pre- ceding paragraph and seems to call for some definition. The word is generally used in a rather indefinite way to designate that attraction between elements which causes them to unite to form compounds. * The real nature of chemical affinity is not known. This is another of those questions, like the cause of the heat generated when elements combine, which waits for an answer. Doubtless the two questions are intimately connected. But while we do not know its real nature, we can learn a great deal about the con- ditions under which chemical affinity acts. Thus it seems al- ways, in accordance with the laws of constant proportion and of combining weights, to be exerted between definite quantities of the elements. We shall find, too, that there are certain ways in which we can give to chemical affinity an accurate, mathemati- cal definition and measure it quantitatively. It seems natural to think of chemical affinity as a force similar to the force of gravity or to the force of electrical attraction. It may be that it is closely connected with one or both of these. Nomenclature. The compounds of oxygen have been called, in this chapter, oxides. This is an application of a system of naming substances which is used for all compounds consisting of two elements. As compounds of oxygen are called oxides, compounds of sulfur are called sulfides, compounds of chlorine, chlorides, etc. In order to give more definite names, prefixes derived from the Greek numerals are used. CO is called carbon 30 A TEXTBOOK OF CHEMISTRY monoxide; CO 2 , carbon dioxide; HgCl 2 , mercury dichloride; P 2 O 3 , phosphorus trioxide ; SO 3 , sulfur trioxide ; CCU, carbon tetrachloride ; P2O 5 , phosphorus pentoxide. In addition to these names, which tell how many atomic weights of the ele- ment are contained in the molecular weight of the substance, the prefix per is used to name compounds containing more oxy- gen than some other oxide of the same element. Thus sodium peroxide, Na 2 O 2 , contains more oxygen than the other oxide of sodium, Na 2 O, the prefix per meaning more or beyond. Still another method of naming oxides is to add the suffixes -ous and -ic to the name of the metal or other element which is combined with oxygen. The ending -ic is used for the compound containing the larger proportion of oxygen. Thus Hg 2 O is called mercurous oxide; HgO, mercuric oxide; FeO, ferrous oxide; Fe 2 O 3 , ferric oxide. The choice among these three methods of naming oxides and other compounds is more or less arbitrary and conventional. CHAPTER III LAWS OF GASES UNITS OF LENGTH, WEIGHT, VOLUME, TEMPERATURE, TIME AND ENERGY. Unit of Length. Meter. The meter was originally intended to be one ten-millionth of the distance from the equator to the pole of the earth, measured on the surface. The measurements by means of which the first meter was prepared were inaccurate, however, and the real meter is the distance, measured at the freezing point of water, between two marks on a bar of platinum- iridium kept at the International Bureau of Weights and Meas- ures at Sevres, France. The meter is divided into tenths, hun- dredths and thousandths, called decimeters, centimeters and millimeters. Its most common multiple is the kilometer, 1000 meters. Unit of Weight. Gram. The gram was intended to be the weight of one cubic centimeter of water at its maximum density, 4 centigrade. Here, again, the first measurements were not quite accurate and the real kilogram (1000 grams) is the weight, in a vacuum, of a mass of platinum-iridium kept at the Inter- national Bureau. The most common division of the gram is one milligram, the thousandth of a gram. Unit of Volume. Liter. The liter was intended to have a volume of one cubic decimeter. Because volumes can be most accurately compared by weighing the water which fills them, the real liter is the volume occupied by one kilogram of water, weighed in a vacuum at 4 C. The cubic centimeter is defined, conventionally, as one one- thousandth of a liter. 1 1 The actual weight of water contained in a cube whose edge is one centimeter is 0.999982 g. according to the best measurements. Because the edge of a cube of water weighing one gram is not ex- 31 32 A TEXTBOOK OF CHEMISTRY Units of Time. The units of time used in chemistry are the second, minute, hour, day and year. These are all fixed by means of astronomical observations with the aid of accurate clocks. Unit of Temperature. The freezing point of water under at- mospheric pressure has been selected as zero for the ordinary centigrade scale of temperature, and the boiling point of water under atmospheric pressure as 100. For the International scale, the interval between the two points is divided into one hundred equal parts by measuring the increase in pressure of hydrogen gas, at constant volume, the initial pressure being that of a column of mercury one meter high at 0. Absolute tempera- tures will be considered later. It is well to notice that the unit of temperature is a unit of intensity and not of quantity. In this respect it corresponds to the height to which a weight is raised in mechanical energy or to the volt in electricity. Units of Energy. Kilogram-meter; Erg. The simplest unit of energy is the kilogram-meter, the energy required to lift one kilogram to a height of one meter. Since the force of gravity varies with the latitude and altitude, another unit, which is independent of these, is often used. This is the erg, and is twice the energy of one gram 1 moving with a velocity of one centimeter a second. Or it may be defined as the energy required to im- part to one gram a velocity of one centimeter per second or to increase its velocity by one centimeter per second. One joule is 10,000,000 ergs. actly one centimeter in length, some persons prefer to call the con- ventional cubic centimeter a mimliter. The suggestion has not, however, been generally accepted. 1 This is more often stated as the mass of one gram, but since weights are always accurately determined by the balance, one gram determined by weighing is just as "absolute" a quantity of matter as the mass of one gram. For the same reason chemists are justi- fied in speaking of atomic weights instead of atomic masses. It is well to remember, however, that weight is in its accurate, scientific use, a measure of a force and not a measure of a quantity of matter. In the common everyday use of the word we use it for a quantity of matter. UNITS 33 Centimeter-gram-second System. Absolute units. Since in accordance with the law of conservation of energy every form of energy bears an exact, quantitative relation to every other, any quantity of energy which we can measure may be expressed in terms of the velocity of a moving mass. The units necessary for such a purpose are a unit of length, a unit of mass and a unit of time. Physicists have agreed upon the centimeter, gram and second as fundamental units and have developed a system of " absolute units " in which all forms of energy are measured by reference to these. These units are called absolute because they are independent of the force of gravity. Units of Mechanical Energy. The absolute unit for mechan- ical energy is the erg, which has been defined above. The most common unit used by engineers is the kilogram-meter (or the foot-pound in the English system). At 45 latitude and sea level the kilogram-meter is 98,066,700 ergs. Unit of Power. Power is the rate of production of energy. One horse power is 4600 kilogram-meters or 33,000 foot-pounds per minute. Units of Heat. The calorie is the heat required to raise the temperature of a, gram of water one degree (p. 26). It varies slightly with the temperature and for accurate work the tem- perature must be specified usually a temperature at 15 l is taken as the standard. As an absolute unit, independent of the temperature of the water the joule has been suggested. One calorie at 15 is equal to 4.187 joules. Electrical Units. The primary electrical units are the wit, ohm and ampere. These are so related that an electromotive 1 If the calorie at 15 is taken as one, the values of the calorie at other temperatures are as follows : 10 1.0016 15 1.0000 20 0.9991 25 0.9988 30 0.9989 These are the mean of the values of Georg lanke, Ann. Tables of Physical Constants for 1910, and of Bausfield, Phil. Trans. 211, A, 199 (1911). 34 A TEXTBOOK OF CHEMISTRY force (E. M. F.) of one volt acting through a resistance of one ohm gives a current of one ampere or : Amperes = E ' M ' F ' R(in ohms) The unit of electrical power is the watt, a current of one ampere flowing under a difference of potential of one volt. It is equiva- lent to 10,000,000 ergs or one joule per second. The kilowatt is, of course, 1000 watts and is the most common measure for electrical service in lighting, heating, running of motors and the like. The watt is one of the " absolute " units. An electrical horse power is 746 watts, and is equivalent, of course, to 33,000 foot pounds per minute. Chemical Energy. By chemical energy is meant the energy which appears as mechanical energy, heat, light, sound or elec- tricity when two or more elements unite, 1 or when an element is changed from one form to another, as ozone to oxygen. It is usually expressed in terms of heat units or electrical units. It must always refer to some definite chemical action which takes place and can never refer to the total energy contained in an ele- ment or compound, as we have no means of measuring this. Effect of Pressure on a Gas. Law of Boyle. When the pres- sure applied to a gas is doubled, the volume is reduced to one half ; or when the pressure is reduced to one half, the volume be- comes twice as great. Another method of stating this property of gases in a general way is to say that the volume of a gas varies inversely as the pressure. Or, mathematically : V : V : : P f : P, or VP = V'P' = Constant, where V and V are two volumes of the same quantity of gas and P and P f are the corresponding pressures. This is known as Boyle's law. It is not an accurate law, as the law of constant 1 Heat is absorbed when some substances unite, and in such cases the energy of the compound is considered as negative or less than nothing in comparison with that of the elements from which it is formed, but the idea that energy can be really negative seems absurd. LAWS OF GASES 35 proportion is, but is sufficiently accurate for use in all ordinary cases. * The extent of the deviation from the law for several gases is shown in the following table : TABLE Volumes filled at by two liters of each gas when the pressure is increased from one atmosphere to two atmospheres. Hydrogen 1.0006 liters Nitrogen 0.9996 liters Carbon monoxide 0.9995 liters Oxygen 0.9991 liters Nitric oxide 0.9989 liters Carbon dioxide 0.9931 liters Nitrous oxide 0.9924 liters Hydrochloric acid 0.9919 liters Ammonia 0.9845 liters Sulfur dioxide 0.9739 liters Those gases which are liquefied most easily depart farthest from the law, and all gases except hydrogen and helium are com- pressed more than they should be under moderate pressures. For a pressure of many atmospheres a point is reached where all gases which do not liquefy are compressed less than they should be in accordance with the law. According to the kinetic theory (p. 58) the greater compressibility under moderate pressure is caused by the attraction of the molecules for each other the same forces which cause the gas to liquefy at low tempera- tures or under pressure. The point of too little compressibility is reached when the molecules are brought so close together that the molecules themselves fill a considerable fraction of the total space. The law may be easily illustrated by taking a gas measuring tube, graduated in cubic centimeters, filling it partly full of mercury and immersing the mouth in a deep, narrow jar con- taining mercury. It is evident that if the tube is raised or lowered till the top of the mercury within the tube is exactly level with the surface of the mercury on the outside, the pressure 36 A TEXTBOOK OF CHEMISTRY of the gas within the tube will be the same as that shown by a barometer in the same room. 1 If, now, the tube is raised, the volume of the gas will be seen to in- crease, and for any given position the pressure of the gas must be equal to the reading of the barome- ter less the height of the mercury in the tube above that in the jar. By reading the volumes in two different positions of the tube and determining the corresponding pressures the data for a verification of the law may be easily obtained. For practical uses it is convenient to select some standard pressure to which the volume of a gas may be referred. The pressure universally used by chemists for this purpose is the pressure of a column of mercury 760 mm. high at 0, at 45 latitude and at sea level. This is, approximately, the average pressure of the air at sea level and is called a pres- sure of one atmosphere. Other pressures are most easily determined by measuring, directly or indi- rectly, the height of the column of mercury which will balance the elastic pressure of the gas. * Corrections for Readings of the Barometer. In accurate work, when the barometer is read at some other temperature than a cor- rection must be subtracted, owing to the fact that the column of mercury is lighter as the metal expands with rise of temperature. The correction in millimeters at temperatures from 5 to 35 is : Fig. 7 TEMPEBATURE CORRECTION FOR BAROMETER CORRECTION FOR BAROMETER DEGREES WITH GLASS SCALE WITH BRASS SCALE 5 0.7 0.6 10 1.3 1.2 15 2.0 1.9 20 2.6 2.5 25 3.3 3.1 30 4.0 3.7 35 4.7 4.3 1 For the sake of simplicity, the lowering of the meniscus of the mercury in the tube by capillary action is disregarded. LAWS OF GASES 37 If the pressure is less than 760 mm. the correction will be less in pro- portion. Thus at 730 mm. the correction for a glass scale is 2.5 mm. at 20 instead of 2.6 mm. The corrections for latitude and altitude are usually less important. Corrections of barometer for latitude, to be added for latitudes less than 45 or subtracted for latitudes greater than 45 : LATITUDE CORRECTION LATITUDE 1.97 90 5 1.94 .85 10 1.85 80 15 1.70 75 20 1.51 70 25 1.27 65 30 0.98 60 35 0.67 55 40 0.34 50 45 0.00 45 Correction for altitude, to be added. ALTITUDE CORRECTION BAROMETER READING 300 meters 0.04 720 600 meters 0.08 700 900 meters 0.12 680 1200 meters 0.16 660 1500 meters 0.19 640 2000 meters 0.24 630 If the corrections for latitude and altitude are applied to the barometer readings, the weight of one liter of the gas at 45 latitude may be properly used in calculating the weight of a quantity of gas measured at any other latitude or altitude. A problem which often presents itself in dealing with gases is the calculation of the volume which a quantity of gas, that has been measured at some other pressure than that of one at- mosphere, would assume if it were brought to atmospheric pres- 38 A TEXTBOOK OF CHEMISTRY sure. Such problems are most easily solved by putting the mathematical expression given above into the following form : VP' P f V =-^~ r V at 760 mm. = V ^~ The student is advised most earnestly that this formula should not be committed to memory. Instead of this it should only be remembered that when the volume at one pressure is given and the volume at another pressure is desired, the first volume is to be multiplied by a fraction in which one pressure is the numera- tor and the other pressure the denominator. A consideration of the fact that an increase in pressure will cause a decrease in the volume will at once indicate which pressure is to be taken as the numerator of the fraction. The proper method of using the formula is emphasized because in the study of chemistry it is of the greatest importance to cultivate the ability to reason quickly from one point to another and to acquire a knowledge of the subject by a rational process rather than by mere memory. Effect of Temperature on a Gas. Law of Charles. When the temperature of a gas is increased one degree while the pres- sure remains constant, the volume will increase ^^ (or 0.003663) of its volume at O . 1 This will be most easily understood with the aid of the accompanying diagram, which gives the volume x This law, while sufficiently accurate for ordinary calculations, is only approximate, the deviations from it being of somewhat the same order of magnitude as the deviations from the law of Boyle. The coefficients of expansion of some of the more common gases as deter- mined by the increase of pressure at constant volume are : Air 0.003666 or 1/272.8 Argon 0.003668 or 1/272.6 Oxygen 0.003674 or 1/272.2 Helium 0.003663 or 1/273.0 Nitrogen 0.003668 or 1/272.6 Carbon mon- Nitric oxide 0.003676 or 1/272.0 oxide 0.003667 or 1/272.7 Hydrogen 0.003663 or 1/273.0 Carbon dioxide 0.003698 or 1/270.4 Sulfur dioxide 0.003845 or 1/260.1 As with the law of Boyle, those gases which are easily liquefied vary most from the rate of expansion for a " perfect " gas. ABSOLUTE TEMPERATURES 39 373 C 283 273 173 C 73 C O c TEMPER- ATURE 100 10 --283cc. O c -100 -- 173 cc. -200 -\- 73 cc. -273 Fig. 8 VOLUME 373 cc. 273 ec. which 273 cubic centimeters of a gas at would assume at other temperatures. Only hydrogen or helium would obey the law at atmospheric pressure over the range of volumes given in the diagram. All other gases are liquid or solid at - 200. Absolute Temperatures. On the left side of the diagram is given a series of numbers which are called absolute temperatures. A little examination of the dia- gram will show that these tem- peratures are based on the thought that if we could find a gas which does not liquefy and which continued to con- tract at the same rate at very low temperatures it would dis- appear at -273. If we take this point as the starting point for the "absolute" scale of temperature it is evident that the freezing point of water will be at 273 absolute and the boiling point 373. Any temperature may be readily converted to the absolute scale by adding to it, algebraically v 273. The absolute scale of temperature enables us to give a very simple statement of the law of Charles, viz. : The volume of a gas varies directly as the absolute temperature. This becomes, mathematically : V : V : : T : T', or F = j- If, as is customary in dealing with gases, we wish to find the volume which a gas, which has been measured at some other temperature than 0, would assume if cooled or warmed to zero, 40 A TEXTBOOK OF CHEMISTRY 273 the formula may be written, VQ= V'. This formula should be used rationally, not by rote (see p. 38), and may be combined with the formula for pressures for practical uses. If the volume of a gas is known at one temperature and pressure, its volume at some other temperature and pressure may be calculated by multi- plying by two fractions one of which involves the two pressures and the other the two absolute temperatures. Significance of the Absolute Zero. The absolute scale of temperature may be treated merely as a mathematical conven- ience in dealing with problems of gases and of thermodynamics ; but the question naturally arises whether the absolute zero has any further, real meaning. Is it, in reality, as the name indi- cates, a point of absolute cold at which all phenomena of tempera- ture begin and below which it is impossible to go ? Many differ- ent phenomena seem to indicate that the absolute zero is actual and not merely a mathematical fiction. It can be no mere acci- dent that hundreds of gases and vapors obey the law of Charles so closely ; and the further the study of the physical and chemical properties of gases is carried, the more clear does it become that the law is intimately connected with some of the most funda- mental properties of matter. From the side of experiment, also, every recent attempt to reach very low temperatures has pointed to 273 as a point which can never be passed. The lowest point thus far reached is that of helium boiling under a pressure of 10 millimeters and is estimated as 270, or 3 absolute. (Kamerlingh Onnes, Chemical Abstracts, 1908, p. 2752.) Determination of the Weight of a Liter of a Gas. The weight of a unit volume of any gas under standard conditions is one of its most important properties, not only for the purpose of cal- culating the weight of a gas when we know its volume, but be- cause of relations between these weights for different gases, upon which one of the most important laws of chemistry is based (p. 89). The unit volume usually chosen is the liter and the standard conditions are a temperature of zero and a pressure of 760 mm. of mercury. WEIGHT OF GASES 41 If a bulb 1 is connected with a manometer and evacuated by means of an air pump, by reading the manometer and tempera- ture, we can, if we know the capacity of the bulb, calculate the volume which the air remaining in the bulb would fill at and 760 mm. If we weigh the bulb and then fill it with some gas at To oirpurrfi Fig. 9 atmospheric pressure (to be determined by reading the barom- eter) and weigh it again, the difference between the two weights will evidently be the weight of the gas which has entered, while the volume of the air which was left in the bulb plus that of the gas which has entered can be readily calculated for stand- ard conditions as before. 2 The difference between this calcu- lated volume and the corrected volume of the air which re- mained in the bulb will be the volume, under standard con- ditions, of the gas which was admitted. From this and the weight it is easy to calculate the weight of one liter of the gas under standard conditions. 1 A capacity of 125 to 150 cc. is suitable for a lecture or labora- tory experiment. The volume may be determined by weighing the bulb empty and then filled with water, but the bulb must be thoroughly dried by warming it and evacuating it repeatedly before it is used for the determination. 2 This assumes, of course, Dalton's law of partial pressures, that when two gases which do not act on each other are mixed, each exerts the same pressure as if it filled the whole space alone, and the total pressure is the sum of the pressures exerted by each gas. 42 A TEXTBOOK OF CHEMISTRY Graphical Representation of the Gas Laws. It is often useful in studying physical and chemical phenomena to use a method of graphical representation which is illustrated in Figs. 10 and 11. In Fig. 10 distances from the line OX represent pressures, while distances from the line OY represent volumes. If we X V 5 4 1 s P 2 1 ( Jb i \ \ \ \ a \ ^ ^-^. "^^ o r ' ) 1 2 3 4 5 PY PRESSURES Fig. 10 start with a unit volume of a gas under unit pressure, represented by the point a, as the pressure increases the volume will decrease along the line ac, while as the pressure decreases the volume will increase along the lines ab and PV, the product of pressure and volume must always remain constant. The geometrical curve which satisfies these conditions is a hyperbola. Charles's law may be represented in a similar way by Fig. 11. LAWS OF GASES 43 Here the relation between volumes and absolute temperatures is represented by a straight line, but all gases liquefy before the 500 400 300 200 100 V O 100 c 200 300 400 ABSOLUTE TEMPERATURES 500 Fig. 11 absolute zero is reached, and the line can never be continued, experimentally, to the origin. 1 . EXERCISES 1. A quantity of gas fills a volume of 175 cc. at 20 and under a pressure of 735 mm. What will be its volume under standard con- ditions (0 and 760 mm.) ? 2. A flask having a capacity of 3.5 liters is filled with oxygen at 25 and 770 mm. What weight of oxygen does it contain ? (See p. 22.) 3. What volume will 33. 5 cc. of a gas measured at 18 and 715 mm. assume at 25 and 731 mm. ? 4. A cylindrical gasometer has a diameter of 30 cm. and height of 60 cm. What weight of oxygen will be required to fill it at 22 and 745 mm. ? 5. A bulb having a capacity of 127.2 cc. was exhausted till the manometer showed a pressure of 35 mm. while the temperature was 1 The line Y is called the axis of abscissas and any line parallel to it and perpendicular to OX is called an abscissa, while OX is the axis of ordinates and any line parallel to this and perpendicular to Or is called an ordinate. O is called the origin. 44 A TEXTBOOK OF CHEMISTRY 23. After weighing, it was filled at atmospheric pressure with a gas. The reading of the barometer was 751 mm. and the temperature 23, as before. The increase in weight was 0.1382 gram. What is the weight under standard conditions of one liter of the gas which was used ? (Ans. 1.2507.) 6. A sample of bituminous coal has the following composition : Carbon 75.00 per cent Hydrogen 5.25 per cent Oxygen 10.00 per cent Ash, nitrogen, etc. 9.75 per cent 100.00 per cent What is the heat of combustion of the coal in calories per kilo- gram and in B. T. U. per pound, assuming the heat of combustion of one gram of carbon as 8080 calories, one gram of hydrogen as 34,179 calories (burned to liquid water), and that 1.25 per cent of the hydrogen is combined with the oxygen and contributes nothing to the heat of combustion ? It is assumed further that the combination between the carbon and hydrogen is of such a nature that these ele- ments give the same amount of heat when the coal is burned as they would give if they were in the free state. CHAPTER IV HYDROGEN SYMBOL, H. ATOMIC WEIGHT, 1.0078. Occurrence. Although the quantity of hydrogen in the world is very much smaller than the quantity of oxygen, it is very widely diffused, especially in the form of its most common com- pound, water. It forms a little more than one ninth of the weight of water and is present both as water and as a constitu- ent of all of the most important compounds found in vegetables and animals. Hydrogen is an essential element, also, in the large class of substances called acids. A minute quantity, pos- sibly 0.001 per cent, or one part in 100,000, is found free in the air. There is some evidence that at very high altitudes the at- mosphere consists almost exclusively of hydrogen. Acids. In order to understand one of the most convenient methods for the preparation of hydrogen in the laboratory, it is necessary to know something of the properties of the important class of substances called acids. It has been shown that when sulfur is burned in the air, sulfur dioxide, SO 2 , is formed. A small amount of the sulfur usually, or perhaps always, combines with more oxygen to form sulfur trioxide, 80s, and by means of suitable apparatus and a catalytic agent, nearly all of the sulfur can be converted into this compound (see p. 175). When sulfur trioxide is dissolved in water it combines with it, giving sulfuric acid : SO 3 + H 2 O = H 2 S0 4 -Sulfur Water Sulfuric Trioxide Acid It will be recalled that this sort of combination between oxides of nonmetallic elements and water gave to chemistry the name 45 46 A TEXTBOOK OF CHEMISTRY of oxygen. Sulfuric acid, when pure, is a heavy liquid of an oily consistency, sometimes called oil of vitriol. By the action of sulfuric acid on common salt we can obtain hydrochloric acid, HC1, a gas which dissolves easily in water and which is ordinarily used in the form of its solution. By the action of sulfuric acid on saltpeter, nitric acid, HNOs, is formed. This is a liquid, which is usually diluted with water for use. The most common acid of ordinary experience is acetic acid, HC2HaO2, the acid of vinegar. This is the acid from which we have all learned to associate the word acid with the sour taste which is characteristic of all moderately strong acids. Radicals. An examination of the formulas of the acids named above shows that each of them contains hydrogen, but a still more important characteristic of these and of all other acids is that in a great variety of reactions this hydrogen may be re- placed by other elements and especially by metals. The follow- ing are illustrations of such replacement : H 2 SO 4 + Zn ZnSO 4 4 2H Sulfuric Zinc Zinc Sulfate Hydrogen Acid HC1 + Na NaCl 4 H Hydrochloric Sodium Sodium Chloride Acid (Common Salt) HNO 3 + NaOH = NaN0 3 h H 2 Nitric Acid Sodium Sodium Water Hydroxide Nitrate HC 2 H 3 O 2 + NaOH = NaC 2 H 3 2 - h HOH Acetic Sodium Sodium Water Acid Hydroxide Acetate In each of these reactions one or two atoms of hydrogen are replaced by an atom 1 of some metal, while all of the rest of the 1 One of the most common mistakes of beginners in such cases is to say "one or two parts of hydrogen are replaced by one part of the metal." The distinction between one part and one atomic weight (or in accordance with the atomic theory, one atom) of an element ought always to be kept clear. HYDROGEN 47 acid passes into the new compound without any change in com- position. A group of atoms, which remain in combination in this way when they pass from one compound to another, is called a radical. Thus SO 4 is the radical of sulfuric acid ; NO 3 , of nitric acid ; C 2 H 3 O 2 , of acetic acid. Salts. The compounds formed by the replacement of the hydrogen of an acid by a metal are called salts. These are so intimately connected with the acids in their composition that it is natural to use for them names which are derived from the names of the acids. How this is done is clear from the illustra- tions given. The name of the metal of the salt is placed first and this is followed by a word in which the -ic of the acid is changed to -ate. Sulfuric acid gives sulfates; nitric acid, nitrates ; acetic acid, acetates. The name o common salt, sodium chloride, seems to be an exception, but this is because, as a binary compound, it belongs to the class of substances which take names ending in -ide (p. 29). Additional principles which are used in naming acids and salts will be considered later. Preparation of Hydrogen. 1. Electrolysis of Dilute Sulfuric Acid. If an electrical current is passed between two strips of platinum (called electrodes) which are immersed in dilute sul- furic acid, bubbles of gas will rise from the electrodes ; and if an apparatus is so arranged (p. 9) that these can be collected, it will be found that the gas rising from the negative electrode (cathode) is hydrogen, while that from the positive electrode (anode) is oxygen. The volume of the hydrogen will be almost exactly twice that of the oxygen. Since, as we shall find later, the hydrogen and oxygen are lib- erated in the same proportion in which they combine to form water, this experiment is often spoken of as a decomposition of water by electricity and in a certain sense this is correct. That the sulfuric acid is more than a merely passive agent in what takes place is evident, however, first, because pure water is nearly a nonconductor for electricity ; and, second, because if we examine the liquid in the U-tube by appropriate means, we shall find that the hydrogen atoms of the sulfuric acid are being 48" A TEXTBOOK OF CHEMISTRY transferred through the liquid toward the cathode as the current passes, while the radicals of the sulfuric acid, the SO4 group of atoms, are transferred toward the anode. In other words, elec- trolysis is not merely something which takes place at the two electrodes, but it is always accompanied by a transfer of material through the whole of the space between ; and while hydrogen is carried in one direction, it is the sulfate radical and not oxygen, which is carried the other way. Electrolytes. Ions. Theory of electrolysis. Any substance which carries the electric current in this way is called an elec- trolyte. The most satisfactory theory which has been proposed to explain the facts which have just been given is that electro- lytes in solution are more or less completely separated into parts which are charged with positive or negative electricity. Ac- cording to this theory sulfuric acid separates partly into hydro- gen atoms with a positive charge of electricity and the sulfate radical with two negative charges. This is indicated by the symbols H + , H + , SO 4 . When the positive and negative elec- trodes are dipped in the dilute acid, the positively charged hy- drogen atoms are attracted by the negative cathode and move toward it, while the negatively charged sulfate radicals are re- pelled by the cathode and attracted by the positive anode. This causes the motion of the hydrogen atoms in one direction and the motion of the sulfate radicals in the other, through the solution. This motion constitutes the current of electricity in an electrolyte. At the cathode the hydrogen atoms lose their positive charge and at once appear as hydrogen gas. At the anode the action is more complicated, but the final result i& that oxygen gas is liberated. The charged atoms or groups are called ions. The positive ion is called the cation, the neg- ative ion, the anion. The decomposition of an electrolyte by an electric current is called electrolysis. 2. Preparation of Hydrogen from Iron and Steam. If steam is passed over red-hot iron contained in an iron tube (Fig. 12), a part of it gives up its oxygen to the iron and hydrogen is liberated. If the compound which remains in the tube is ex- amined, it is found to have the same composition as the magnetic HYDROGEN 49 oxide of iron, Fe 3 O 4 , formed when iron burns in oxygen. The equation is not quite so simple as those which have been given Fig. 12 before. In order to arrive at the true equation the formulas of the substances used and the products obtained should be written first : Fe + H 2 O -> Fe 3 O 4 + H Iron Water Magnetic Hydrogen Oxide of Iron On examining the above it is seen that 4 atoms of oxygen will be required to form one molecule of the magnetic oxide of iron, hence we must have 4 molecules of water in the first member of the equation to furnish these. The 4 molecules of water will give 8 atoms of hydrogen and 3 atoms of iron will also be re- quired to form the magnetic oxide. Putting all together we have: 3 Fe + 4 H 2 O = Fe 3 O 4 + 8 H It would, doubtless, be easier for a beginner to learn this last equation outright than to learn how to derive it in the manner indicated, but things which are a mere matter of memory are likely to be evanescent, while a rational process like the above can be reproduced at will. It is very important in studying chemistry to reduce those portions which are remembered as distinguished from those portions which are derived by a logical 50 A TEXTBOOK OF CHEMISTRY process just as far as possible. At the same time many simple, fundamental facts, as here the composition of the magnetic oxide of iron, must be remembered and used over and over again. Reversible Reactions. It was stated above that a part only of the steam is decomposed by the iron. If we reverse the con- ditions and pass hydrogen over magnetic oxide of iron, part of the hydrogen will be converted into water and metallic iron will be obtained. The earlier and most natural idea of chemical affinity was that when three elements are present those two would unite which had the strongest affinity for each other. If this were true, either the hydrogen would be able to take the oxygen from the iron or the iron should be able to take it from the hydrogen. We see from the experiments described that this simple idea is not correct, but that either element can take the oxygen from the other. While there is a certain sense in which iron has a stronger affinity for oxygen than hydrogen has, the direction of the reaction depends on the quantities of the sub- stances present as well as upon their relative affinities. If steam is used and the hydrogen is constantly removed, the tend- ency is to form magnetic oxide of iron and hydrogen. If hydro- gen is used and the steam is constantly removed, the tendency is to form metallic iron and water. It is often convenient to express such reversible reactions as follows : 3 Fe + 4 H 2 O ^ Fe 3 O 4 + 8 H 3. Decomposition of Water by Metals at Ordinary Tempera- tures. Potassium and sodium have a much stronger affinity for oxygen than iron has, and partly for this reason, partly, perhaps, for other reasons which we do not fully understand, these metals will decompose water and liberate hydrogen at ordinary temperatures. If potassium is thrown on water, the heat of the reaction is great enough to cause the hydrogen to ignite. It burns with a violet flame, the color being given to it by the potassium. Sodium when thrown on water usually rolls over the surface in a globule, evolving hydrogen, which does not take fire, but if thrown on a piece of filter paper lying on the HYDROGEN 51 water so that the globule remains at one spot, the hydrogen will catch fire and burn with the yellow flame characteristic of sodium. If a small piece of sodium is wrapped in paper and thrust quickly under the mouth of a jar which has been filled with water and inverted with the mouth under water, the sodium will act on the water as before and the hydrogen may be collected and examined. If the water in which the potassium or sodium has dissolved in these experiments is examined, it will be found to have a soapy feel and disagreeable, acrid taste. It will also turn the color of red litmus paper blue. If the solution is evaporated in a dish of platinum or of some material which is not affected by it and under such conditions that it cannot absorb carbon dioxide from the air, a white solid will be obtained, which will have the composition represented by the formula KOH or NaOH. These substances absorb and retain water so greedily that it is necessary to heat them nearly to redness before the last of the water is expelled. They are called, in accordance with their composition, potassium hydroxide or sodium hydroxide. The equations are : K + H 2 O = KOH + H Potassium Water Potassium Hydrogen Hydroxide Fig. 13 Na + H 2 = Sodium NaOH + Sodium Hydroxide H or, Na + HOH = NaOH Hydrogen Hydroxide + Contrast between the Action of Iron and of Sodium on Water. The last form expresses a little more clearly that the metal has replaced only one of the two atoms of hydrogen in each molecule 52 A TEXTBOOK OF CHEMISTRY of water. The action is seen to be quite different from that of iron on steam. This is closely connected with the amount of chemical energy changed to heat in each reaction. In the reac- tion, Na + H 2 O = NaOH + H, 43,450 calories are liberated for each gram atom of hydrogen set free, while in the reaction, 3 Fe + 4 H 2 O = Fe 3 O 4 + 8 H, if it could be carried out at 100, only 4160 calories would be given for each gram atom of hydro- gen liberated. In general, those reactions in which large amounts of chemical energy are changed to heat take place most easily. We must, however, guard against the impression that this is a universal law. The ease with which a reaction takes place is by no means proportional to the heat generated. Other factors are involved, and some of these are, at present, but little understood. 4. Hydrogen from " Hydrone." The action of water on so- dium is too violent for use as a laboratory method of preparing hydrogen in quantity. If, however, the sodium is alloyed with lead, the action is moderated, and such an alloy containing about 35 per cent of sodium is sold under the name of " hydrone." By means of it very pure hydrogen can be easily prepared. 5. Preparation of Hydrogen by the Action of Metals on Acids. If a strip of zinc and one of platinum, copper or lead are dipped in dilute sulfuric or hydrochloric acid while the strips are con- nected by means of a wire, an electrical current will pass through the wire while bubbles of hydrogen will be seen to collect and rise from the surface of the platinum, copper or lead. If the liquid between the two metallic plates is examined, as in the electrolysis of dilute sulfuric acid, it will be found that the hydrogen travels through the liquid toward the platinum, while the sulfate radical or the chlorine travels toward the zinc. At the surface of the zinc, the sulfate radical or the chlorine combines with the zinc, forming zinc sulfate, ZnSO4, or zinc chloride, ZnC^. If pieces of chemically pure zinc are placed in dilute hydrochloric or sul- furic acid, there will be almost no action at all, while commercial zinc will dissolve rapidly. After the action of the acid on the commercial zinc has continued for a short time it will be seen HYDROGEN 53 that the surface is dark, and a closer examination will show that it is covered with lead and other impurities found in the zinc. When we consider these facts along with the experiment with the strips of platinum and zinc, we reach the conclusion that the ac- tion of the acid on the zinc requires some catalytic agent like the platinum or lead before it can be very rapid, and that the phenomenon of the solution of the zinc is partly electrical, being accompanied by electrical currents between the particles of lead and zinc in the impure zinc. If we disregard the catalytic agent, the process may be expressed by the equations : Zn + H 2 SO 4 = ZnSO 4 + 2H Zinc Sulfuric Zinc Hydrogen Acid Sulfate Zn + 2HC1 = ZnCl 2 + 2H Hydrochloric Zinc Acid Chloride It is to be noticed that one atom of zinc replaces two atoms of hydrogen in each case and that when an acid is used which has only one atom of hydrogen in its molecule, two molecules of the acid are required for the reaction. If iron is substituted for zinc, these reactions become : Fe + H 2 SO 4 = FeSO 4 + 2H Iron Sulfuric Ferrous Hydrogen Acid Sulfate Fe + 2HC1 = FeCl 2 + 2H Ferrous Chloride Apparatus for the Preparation of Hydrogen. In the labora- tory, small quantities of hydrogen may be generated in the simple apparatus shown in Fig. 14. Zinc and some water are placed in the generating flask, and dilute sulfuric acid 1 is added in portions through the thistle tube. 1 As sulfuric acid is heavier than water (sp. gr. 1.84) and much heat is generated on its dilution, it should always be poured slowly into water and should never be diluted by pouring water upon the acid. Pouring water on concentrated sulfuric acid may cause an explosion. 54 A TEXTBOOK OF CHEMISTRY A more convenient apparatus for the preparation of larger amounts, or when it is desired to have the gas always ready for use, is the Kipp generator (Fig. 15). The zinc is placed in the middle bulb and the dilute acid is poured in through the upper bulb, which communicates with the lower one through the tube A. When the stopcock B is opened, the acid rises and comes in contact with the zinc in the middle bulb and the gen- eration of hydrogen begins. Whenever the stopcock is closed the hydrogen generated forces the acid away from the zinc and the action ceases as soon as the acid moistening the surface of the zinc is exhausted. The Fig. 14 generator is not altogether satisfactory because the spent acid containing zinc sulfate is mixed with that which has not been used, diluting it and causing the action to become very slow before the acid has been exhausted. A more suitable form of apparatus for generating large quantities of hydrogen is described on p. 165. Purification of Hydrogen. The hydrogen prepared by any of the methods described is impure. Spray from the generating liquids may be removed by passing the gas through a tube filled with cot- ton wool. Moisture, or water vapor may be removed by means of calcium chloride, 1 contained in a Fig. 15 1 One liter of a gas dried with cal- cium chloride retains 1.0 mg. of water at 15, 1.5 mg. at 20, 2.5 mg. at. 25, 3.3 mg. at 30. When dried with concentrated sulfunc acid the amount of water retained by one liter of the gas is only 0.002 mg. at 15 to 19. When dried with phosphorus pentoxide HYDROGEN 55 tube such as shown in Fig. 16, or more perfectly by means of pumice stone or glass beads moistened with concentrated sul- furic acid or by phosphorus pentoxide. Hydrogen sulfide 1 and some other impurities, especially some of those which give an unpleasant odor to the gas, may be removed by passing it through a wash bottle containing a solution of potassium permanganate, but a small amount of oxygen will be introduced into the gas (V. Meyer, and Recklinghausen Ber. 29, 2550). Oxygen may be removed by passing the gas through a hot tube containing platinized quartz, which will cause the oxygen to combine with some of the hydrogen. This should be done, of course, be- fore the gas is dried. Nitrogen from the air cannot be removed and when pure hydrogen is re- quired very great care is necessary to prevent its entrance (Cooke ji- jg and Richards, Am. Chem. J. 10, 81; Morley, ibid. 17, 267; Noyes, J. Am. Chem. Soc. 30, 1724). Properties of Hydrogen. Hydrogen is a colorless, tasteless and odorless gas. One liter weighs at and 760 mm. pressure 0.08987 gram. As the weight of a liter of air is 1.293 grams, air is Fig. 17 0.08987 times heavier than hydrogen, or approxi- mately 14 J times. Oxygen is ! = 15.90 times heavier than hydrogen, or approximately 16 times. the amount of moisture retained by one liter of the gas is less than 0.00002 mg. (Morley). 1 Hydrogen sulfide may be removed to better advantage by passing the gas through a wash bottle or serpentine tube containing lead oxide dissolved in a solution of potassium hydroxide. 56 A TEXTBOOK OF CHEMISTRY Fig. 18 At a very low temperature hydrogen condenses to a liquid which boils under atmospheric pressure at 252.5 or 20.5, ab- solute. The liquid has a density of only 0.07 gram per cubic centimeter or one gram fills a volume of about 14 cc. If the liquid is made to boil by reduc- ing the pressure, it grows still colder and at - 260, or 13, ab- solute, what remains freezes to a solid. The vapor pressure of the solid is 58 mm. The lightness of hydrogen as compared with air may be easily shown by pouring it upward through the air, by showing that it may be collected in an inverted jar or beaker while it will not remain in one which is placed upright, and by filling soap bubbles or toy balloons with the gas. The use of the gas for filling balloons is well known. * What is the lifting power of one cubic meter of pure hydrogen under standard conditions? What will be the lifting power of one cubic meter at a mile above sea level when the pressure is 620 mm. and the temperature 20 ? It is con- venient to remember that as the volume of a gas varies inversely with the pressure and directly with the absolute temperature, the weight of a given volume must vary directly as the pressure and inversely as the absolute temperature. Diffusion of Gases. If two cylinders are filled, one with hydrogen and the other with air, and placed with their mouths together but with the hydrogen above (Fig. 19), it will be found after a comparatively short time that the gas in each cylinder will explode, if ignited, with the whistling sound char- acteristic of a mixture of air and hydrogen fired in an open cylinder or test tube. Although the air is fourteen and a half Fig DIFFUSION 57 times heavier than the hydrogen, it makes its way quite rapidly up into the hydrogen above it and the hydrogen passes down into the heavier air below. This property of mixing with each other is true of all gases without exception and is called diffusion. While there are many liquids which do not dissolve in each other or which dissolve only to a limited extent and such liquids separate into layers in accordance with their specific gravities, gases, which differ much more than liquids in their densities, always mix when brought in contact and when once mixed will sep- arate to only a very slight extent (p. 45) in accordance with their densities. Very closely related to the diffusion of gases which are in contact with each other is the diffusion of gases through a wall full of fine openings, which separates them. If a cylinder of porous porcelain with openings so fine that pressure will cause a gas to pass through them only very slowly is fitted with a rubber stopper and con- nected with a bulb and bent tube filled with water as shown in Fig. 20, on bring- ing a beaker filled with hydrogen over the cylinder the pressure within will sud- denly increase and force water out of the tube in a jet, showing that hydrogen is Fig. 20 passing through the walls of the cylinder to the interior. When the stream of water ceases, if the beaker is removed, the movement of the water in the tube will show that diffu- sion of the hydrogen outward is taking place. By appro- priate experiments it is possible to show that some air passes out through the porous wall while the hydrogen is passing in. A careful study of the phenomena by Graham has shown that gases pass through a porous wall of this sort, at a rate which varies inversely as the square root of the density. Oxygen, which is about sixteen times as heavy as hydrogen, will pass the wall only one fourth as fast. 58 A TEXTBOOK OF CHEMISTRY Kinetic Theory of Gases. When water is converted into steam at 100 the volume at atmospheric pressure is increased more than 1600 times. This fact and many of the other proper- ties of gases makes it seem highly probable that the space be- tween the molecules of a gas is very large in comparison with the size of the molecules themselves. A very satisfactory explana- tion of the law of diffusion of gases, given in the last paragraph, and of the fact that a gas expands at once to fill any empty space, however large, which it is allowed to enter, is found in the kinetic theory of gases. According to this theory the molecules of gases are moving constantly at a comparatively high velocity, and whenever they meet each other they rebound according to the laws of elastic bodies. According to these laws when two elastic bodies meet each other squarely each rebounds with the energy of the other. If the bodies meet at an angle, the inter- change of energies will be only partial, and the effect of this constant interchange must be to give to all molecules of the same weight approximately the same average energy and hence the same average velocity. If molecules of different weights are mixed, however, the interchange of energies must give a greater velocity to lighter molecules. As the energy of a moving body varies as the square of its velocity, the velocities of mole- cules of different weights must vary inversely as the square roots of their weights, if their energies are the same, since the pressure of a gas is due to the impacts of its molecules on the walls of the containing vessel and the pressure does not change when two gases having different densities are mixed. As an illustration we may take oxygen and hydrogen. The molecule of oxygen is 16 times as heavy as that of hydrogen (p. 95). If the two gases are mixed, the molecules of hydrogen must, through frequent collisions with molecules of oxygen, soon have the same average energy as the latter, and to do this must have, on the average, four times the velocity of the molecules of oxygen. What is true of the mixed gases must be true also of the gases when separate. Accordingly, if we have a porous wall with fine openings separating the two gases, as the hydro- DISSOCIATION 59 gen molecules have four times the velocity of the oxygen mole- cules and there are the same number in equal volumes, four hydrogen molecules will hit the openings and pass into them while one oxygen molecule does so. This ratio of four to one is title same as the ratio of the square roots of the densities, Vl6 : Vl = 4:1, the law for diffusion through a porous wall given above. Chemical Properties of Hydrogen. If hydrogen is brought to the air, through a glass tube drawn to a narrow opening, and lighted, it will burn for an instant with a pale blue, almost in- visible flame, but the color quickly changes to yellow from par- ticles of sodium or its compounds, which are volatilized from the glass by the heat. From a platinum jet the pure gas burns with a flame almost or quite invisible in daylight. The product formed is water, as may be shown, roughly, by holding a cold glass over the flame, or, more accurately, by burning the gas for some time, condensing the water formed and determining its freezing point and boiling point. The hydrogen and air used should, of course, be carefully dried. Mixtures of air and hydrogen explode when ignited, as the flame travels through such a mixture with a very high velocity, and both the steam formed and the nitrogen of the air are heated to a high temperature and expanded greatly by the heat of com- bustion. Mixtures of oxygen and hydrogen explode still more violently. Hydrogen does not support the combustion of substances which burn in oxygen (Fig. 21). It will support the combustion of oxygen or chlo- rine. Dissociation. .If steam is passed through a tube of porous porcelain A, Fig. 22, which is inclosed in a larger tube B of glazed porce- lain through which is passed a current of carbon dioxide, introduced through C, while the whole is heated to a very high temperature (2000, perhaps) the steam 60 A TEXTBOOK OF CHEMISTRY will be partly decomposed into oxygen and hydrogen, and, since the hydrogen diffuses through the porous porcelain four times as fast as the oxygen, more oxygen than enough to com- bine with the hydrogen which remains will stay in the inner tube while more hydrogen will pass through than enough to combine with the oxygen which passes through. If the cooled Fig. 22 gases which are delivered at the ends of the tube are mixed and the carbon dioxide is absorbed by a solution of sodium hy- droxide, it will be found that the gas which remains consists of a mixture of two volumes of hydrogen with one volume of oxygen. In this way Deville showed that water can be de- composed into oxygen and hydrogen by heat alone, and that the reaction between hydrogen and oxygen is reversible : A decomposition of this kind, when the products of decomposi- tion recombine on cooling or on a reversal of the process which caused the decomposition, is called a dissociation. Probably all compounds would be decomposed into their elements at a suffi- ciently high temperature, but it is only in those cases where the elements recombine, on cooling, to form the same compound, that the decomposition is called a dissociation. * By other methods the per cent of dissociation of water has been determined up to 2300, absolute. From the results the dissociation at still higher temperatures may be calculated OXYHYDROGEN BLOWPIPE 61 approximately. 1 The results are as follows, at atmospheric pressure : ABSOLUTE TEMPERATURE TEMPERATURE CENTI- GRADE PER CENT OP DISSOCIATION T r 1000 727 0.00003 1500 1227 0.022 2000 1727 0.59 2500 2227 3.98 2773 2500 8.12 3273 3000 20.0 4000 3727 40.5 The Oxyhydrogen Blowpipe. If oxygen and hydrogen are brought together in such a manner that they burn as they come in contact, an almost colorless flame having a very high tempera- Fig. 23 ture is produced. If steam could be heated to very high tem- peratures by the expenditure of the same quantity of energy, pro- 1 By the formula, 2 P (100 log - i.oo.iu -^r - 1000) - 0.685. 10- 7 (T 2 - 1000 2 ), in which P is the pressure in atmospheres, T, the absolute temperature, and x, the fraction dissociated. (Nernst, Theoretische Chemie, 6 te Aufl., p. 681.) 62 A TEXTBOOK OF CHEMISTRY portionally, as that required to heat it to 1000, the heat of com- bustion of oxygen and hydrogen (p. 65) is great enough to heat the steam produced to 10,000, at least. A little consideration of the dissociation of water at high temperatures shows us that such an extreme temperature cannot be reached, for it is evident that oxygen and hydrogen cannot combine to produce heat at a temperature at which water largely decomposes into its ele- ments. Since there is a dissociation of 40' per cent at 3700, it is doubtful if even that temperature can be obtained. The tem- perature of an electric arc between carbon points is estimated at about 3600. The temperature of an open-hearth steel fur- nace is only 1500 to 1700. Iron wire will take fire in the oxyhydrogen flame and burns brilliantly. Platinum melts easily (1755) and the flame has long been used in working with this metal. A piece of lime held in the flame glows intensely, giving the light known variously as the oxyhydrogen, lime or Drummond light. It will be noticed that any solid substance, which does not volatilize too easily, gives an intense light in the flame, while the oxyhydrogen flame alone is almost nonluminous. Explosions. Catalysis. A mixture of oxygen and hydrogen may remain in a glass tube indefinitely without combining to an extent that can be measured. It is not till a temperature of 300 is reached that the gases combine rapidly enough so that the rate of combination can be measured after some days or weeks. At a temperature a little above 500, the combination is fast enough so that the heat of combination raises the mix- ture to a higher temperature, which still further accelerates the combination, and an explosion results. It is characteristic of most explosions caused by chemical action that the reaction causing the explosion is accelerated enormously by the heat of the reaction. If the mixture of oxygen and hydrogen is brought into contact with platinum in the spongy form or with platinized asbestos 1 the reaction is hastened and the gases will 1 Prepared by moistening asbestos with a strong solution of chloroplatinic acid and heating it. OXIDATION. VALENCE 63 usually take fire and burn. This is another illustration of ca- talysis and recalls the effect of other metals on the solution of zinc in acids. It may be that the two phenomena are closely related. Oxidation. Reduction. If a piece of copper is held over a flame so that it is heated quite hot while exposed to the air, it will be oxidized, the surface changing to black copper oxide. On holding the hot, oxidized copper in an atmosphere of hydrogen, the black oxide will be quickly changed back to metallic copper. This process is called reduction, and hydrogen is called a reducing agent. It will be seen from the above that oxidation and reduc- tion are opposite processes. The two words are often used in chemistry in a much more general sense. The addition of other elements than oxygen is frequently called oxidation, and the re- moval of other elements, or the substitution of hydrogen for other elements, or even the addition of hydrogen to a com- pound, may be called a reduction. Valence. It may have been noticed that when sodium and zinc act on water or on hydrochloric acid, one atom of the sodium replaces one atom of hydrogen, while one atom of zinc replaces two atoms. This characteristic of metals may be made more striking by selecting three metals whose atomic weights are close together. If 23 milligrams of sodium, 24 milligrams of magne- sium and 27 milligrams of aluminium are allowed to act on hy- drochloric acid in such a way that the hydrogen generated is collected in separate tubes, 1 it will be found that the sodium will give about 11 cc. of hydrogen, 2 the magnesium 22 cc. and the aluminium 33 cc. The property of the metals illustrated here is called valence. A metal which replaces one atom of hydrogen 1 The sodium may be placed in a short lead tube, 3 mm. in diam- eter and closed at one end. The mouth should be closed with a little cotton till ready for use. The magnesium and aluminium may be weighed in small watch glasses about the size of the mouths of the test tubes to be placed over them. The aluminium must be etched with a solution of sodium or potassium hydroxide and after- wards washed and dried before use. 2 One cc. of hydrogen weighs 0.09 mg. 0.09 X 11.1 = 1 nig. of hydrogen from 23 mg. of sodium. 64 A TEXTBOOK OF CHEMISTRY for one atom of the metal is called univalent; one which re- places two atoms is called bivalent; three, trivalent; four, quadrivalent. On a somewhat different basis, which will be discussed later (p. 156), when other meanings of valence are considered, some elements may be quinquivalent, sexivalent, septivalent or even octovalent. The valence of the metals is Fig. 24 such an important characteristic and a knowledge of it is so useful in the writing of formulas that it seems best to present it from the standpoint of replacement here. Some metals have two or more kinds of valence in different compounds. Thus, in the sense in which valence is used here, iron is bivalent in ferrous chloride, FeCl 2 , and ferrous sulfate, FeSO4, but trivalent in ferric chloride, FeCla, and ferric sulfate, Fe2(SO4)s. This last formula illustrates the value of a knowledge of valence in writing formulas. (What is the formula of alu- minium sulfate? What are the formulas of stannous sulfate and of stannic sulfate if tin is bivalent in stannous compounds and quadrivalent in stannic compounds ?) From the point of view of the atomic theory the facts by means of which the valences of elements are determined point very strongly to differences in the powers of atoms to combine with other atoms. Thus in the compounds NaCl, MgCl 2 and HYDROGEN 65 " Aids it seems that an atom of sodium can hold one atom of chlo- rine in combination, an atom of magnesium can hold two and an atom of aluminium, three. This property of valence is often expressed by means of such formulas as the following : Cl /Cl Na Cl, Mg< , Alf-Cl, H O H X C1 X C1 The lines are intended to represent lines of force holding the atoms in combination, and the number of lines proceeding from the symbol of an element indicates its valence. We must distinguish sharply between the intensity of the force holding two atoms together and the valence of the atoms. Thus the valence of sodium in Na Cl is the same as that of hydrogen in H Cl, but the force which holds the sodium and chlorine together is much greater than that which holds hydrogen and chlorine together. Heat of Combustion of Hydrogen. The heat generated when 2.015 grams of hydrogen combine with 16 grams of oxygen and the water formed is condensed to the liquid state at 18 is 68,414 small calories, or 34,179 calories for one gram of hydro- gen (Thomsen). This value is often used in calculating the heat of combustion of coal. If 2.015 grams of hydrogen combine with 16 grams of oxygen at 100 and the water formed remains as steam, the heat of combination is only 58,000 calories, at constant pressure, or 28,970 calories for one gram. This is, of course, the maximum amount of heat which can be obtained by burning hydrogen un- der practical conditions. CHAPTER V WATER, HYDROGEN PEROXIDE Analysis. Synthesis. The two methods by which the com- position of a substance is determined are by analysis, the sepa- ration of the substance into the elements of which it is com- posed, and by synthesis, the putting together of the elements to form the compound. In analysis it is comparatively seldom that the elements are separated in the free state. Thus, in order to determine the amount of hydrogen in an organic compound, such as sugar, the substance is burned and the water formed is collected and weighed. Knowing what per cent of hydrogen is contained in water, it is easy to calculate the amount of hydro- gen contained in the compound. Either an analysis or a synthe- sis may be qualitative, giving simply the elements which are present, or quantitative, giving the quantity or per cent of each element. Qualitative Analysis and Synthesis of Water. The prepara- tion of hydrogen by passing steam over red-hot iron is a qualita- tive analysis of water. In order to make the analysis complete it would be necessary to show that hydrogen and the magnetic oxide of iron are the only products of the action, that the density and properties of the hydrogen are the same as those of hydrogen prepared in other ways and that the density and properties of the magnetic oxide formed are the same as those of the oxide formed by burning iron in oxygen. The experiment showing that water is formed when dry hy- drogen burns in air or in oxygen is a qualitative synthesis of water. Quantitative Synthesis of Water by Volume. The determina- tion of the composition of water by volume may be made in an instrument called a eudiometer, a tube graduated usually to 66 COMPOSITION OF WATER 67 tenths of a cubic centimeter by means of fine lines etched on the surface. For the synthesis of water two platinum wires must be sealed in, near the closed end. A capacity of 50 cc. is suitable for the experiment to be described. Such a eudiometer is carefully dried and filled with mercury and 10-12 cc. of dry oxygen introduced. The volume of the gas is then accurately measured and the height of the mercury in the eudi- ometer above the mercury in the reservoir, the tem- perature and the reading of the barometer are noted. From these measurements the volume of the oxygen under standard conditions is calculated. Enough dry hydrogen to give a total volume of 25-28 cc. is then introduced and these measurements repeated. The hydrogen should be in excess, but, as the pres- sure is greater, the total volume need not be three times the vol- ume of the oxygen. The eudiom- eter is then clamped firmly, with a piece of sheet rubber placed under its mouth, and the mix- ture exploded by passing an elec- tric spark between the platinum points by means of an induction coil. After cooling, the volume of hydrogen remaining is meas- ured as before, except that the water formed by the explosion remains partly as vapor in the Steam =18 =20 .-22 =24 -23 =32 =36 Fig. 25 hydrogen and the pressure of the hydrogen is to be found by subtracting from the reading of the barometer the height of the mercury in the eudiometer plus the pressure of vapor of water for the temperature which is read (see p. 75). After correction of the three volumes of gas to standard conditions the proportion by volume in which the gases have united may be calculated. 68 A TEXTBOOK OF CHEMISTRY By placing a tube over the eudiometer (Fig. 26) and passing steam through it, the water formed in the experiment may be converted into steam and the volume of the excess of hydrogen plus the steam determined and from this the volume of the steam calculated. * The results of Morley's exceedingly careful experiments (Amer. J. Sci. 41, 220 and 276 ; Chem. News, 63, 218) show that when oxygen and hydrogen are measured in tubes the ratio of the volumes which combine is, O : H = 1 : 2.0002. Curiously enough Scott (Phil. Trans. 184, A, 543 (1893)) has found that when the gases are measured in globes the ratio is O : H = 1 : 2.0025. It seems, therefore, that the same quantity of gas may fill a different volume when measured in a tube from what it does when measured in a globe, but no one has proved this by direct experiment. The volume of the steam is very nearly the same as the volume of the hydrogen which goes to form it. We can express the re- lation by the following diagram : II 1 vol. oxygen 2 vols. hydrogen 2 vols. steam Composition of Water by Weight. From the composition of water by volume and the weights of one liter of each gas we may calculate the composition by weight. This gives : O : H = 1.429 : 2.0025 X 0.08987 or 16:2 X 1.0075 The Unit for Atomic Weights. The quantity of hydrogen combining with 16 parts of oxygen is given because an atomic weight of 16 has been assigned to oxygen, somewhat arbitrarily, as a basis for comparison with all other atomic weights. Hy- drogen with an atomic weight of one was originally chosen as the unit for atomic weights. For 70 or 80 years it was sup- posed, on the basis of inaccurate determinations of the composi- tion of water, that the atomic weight of oxygen was, on that COMPOSITION OF WATER 69 basis, almost exactly 16 (or 8). When the composition of water was finally determined more accurately, it was decided by the majority of chemists that it is better to make oxygen, with an atomic weight of 16, the basis for all other atomic weights, rather than to make such large changes as would be necessary in many of the common atomic weights, if hydrogen were retained as the unit. A few chemists, however, still prefer hydrogen as the unit. Determination of the Composition of Water by the Use of Copper Oxide. The first moderately accurate determination of the composition of water was made by the Swedish chemist, Berzelius, in 1819. He weighed a quantity of copper oxide in a glass bulb, heated it, passed dry hydrogen through the bulb, and collected and weighed the water formed. The copper oxide was reduced to metallic copper, and the loss of weight gave the weight of oxygen which had been converted into water. The difference between the weight of the water collected and the weight of oxygen taken from the copper oxide gave the weight of the hydrogen. A number of years later (1842) a French chemist, Dumas, carried out an elaborate series of experiments by the same method, with the apparatus shown in Fig. 27. The hydrogen was generated in the large bottle and passed through a series of tubes to purify and dry it. It was then passed through the bulb containing the copper oxide and the water formed was collected, partly in a bulb, and partly in dry- ing tubes. The atomic weight of hydrogen calculated from the results of 19 experiments with this apparatus is 1.0025. 1 For some unknown reason the result is too low by about one part in 200. About 50 years later the copper oxide method was modified by the author, who used the apparatus shown in Fig. 28. After placing some copper oxide in the bulb A and ex- 1 The student is not, of course, expected to remember these various values. A very brief description of 4 out of some 16 deter- minations of the composition of water is given as an illustration of the amount of labor which has been expended on the determination oi atomic weights and also to illustrate how successive workers attain, sometimes, a closer approximation to the truth. 70 A TEXTBOOK OP CHEMISTRY COMPOSITION OF WATER 71 hausting it with a good mercury air pump the apparatus was weighed. It was then connected at C with an apparatus fur- nishing pure hydrogen, the bulb A was heated in an air bath and the tube B was cooled. The stopcock E permitted the hydrogen to pass out at first through D so that no air should enter the bulb from the connecting tubes. On admitting hydrogen to the bulb it was converted into water by the copper oxide and the water was condensed in B. After from one to two grams of hydro- Fig. 28 gen had been converted into water in this way the stopcock was closed and the apparatus cooled and weighed. The gain in weight was the weight of the hydrogen which had entered. The apparatus was then connected with a tube into which the water formed could be driven by warming B and A. The loss in weight of the apparatus gave the weight of the oxygen which had been taken from the copper oxide. The water was also collected and weighed. Twenty-four determinations, partly by this method, partly by another which need not be described here, gave 1.00787 as the atomic weight of hydrogen. Determination of the Composition of Water by weighing Oxygen and Hydrogen. For more than twelve years Professor Morley worked at Cleveland on the composition of water by volume, on the determination of the weights of oxygen and hy- drogen gases and finally on the composition of water by weight. The weights of one liter of hydrogen and of oxygen which have been given are the values determined in this long, classical in- 72 A TEXTBOOK OF CHEMISTRY vestigation. The composition of water by weight was also de- termined in the apparatus shown in Fig. 29. Into this apparatus were brought hydrogen from a tube containing metallic palla- dium, in which it had been absorbed, and oxygen from globes in which it had been weighed. The water formed was frozen in the bottom of the apparatus and was weighed at the end of the experiment. From twelve experiments the atomic weight of hydrogen is calculated as 1.00762. The atomic weight of hydrogen which is now used (J. Am. Chem. Soc. 31, 1) is 1.0078. This is probably not in error by so much as one part in 5000. For ordinary calculations the value is rounded off to 1.008, or, frequently, to 1.01. Properties of Water. Pure water appears, ordinarily, to be colorless and transparent, but light transmitted through a layer of water some meters in thickness has a blue color. Water is the only substance for which we have three names according as it is in the form of a solid, liquid or gas. Water has a maximum density at 4. If either cooled or heated from that temperature, it expands. For this reason in the fall and winter large bodies of water cool by convection, that is by the sinking of the cooler, heavier water on the surface, till a temperature of 4 is reached. On further cooling the water grows lighter again and the colder, lighter water floats on the surface, protecting the warmer water beneath from further rapid cooling. The ice which finally forms has a density of only 0.92 (accurately 0.91674) and continues to float on the surface. The density of water at dif- ferent temperatures is given in the following table : l 1 This table is useful especially for determining the capacity of burettes or flasks by weighing the water which they contain. Thus it is seen that the water which will fill 1 cc. at 20 weighs 0.99823 gram. This is, however, when weighed in a vacuum. If weighed with brass weights in air, the apparent weight will be 0.00105 gram. Fig. 29 PROPERTIES OF WATER 73 TEMPERATURE DENSITY TEMPERATURE DENSITY 0.99987 20 0.99823 1 0.99993 21 0.99802 2 0.99997 22 0.99780 3 0.99999 23 0.99756 4 1.00000 24 0.99732' 5 0.99999 6 0.99997 25 0.99707 7 0.99993 30 0.99567 8 0.99988 35 0.99406 9 0.99981 40 0.99224 45 0.99024 10 0.99973 50 0.98807 11 0.99963 55 0.98573 12 0.99952 60 0.98324 13 0.99940 65 0.98059 14 0.99927 70 0.97781 75 0.97489 15 0.99913 80 0.97183 16 0.99897 ' 85 0.96865 17 0.99880 90 0.96534 18 0.99862 95 0.96182 19 0.99843 100 0.95838 less, or 0.99718 gram. The following table, which gives the appar- ent weight of one liter of water weighed with brass weights in air and the corresponding correction to volume, is still more conven- ient. Apparent weight of one liter of water in air and corrections to be added to the apparent weight of one liter of water to find the true volume in cubic centimeters. TEMPERA- TURE GRAMS CORRECTION TEMPERA- TURE GRAMS CORRECTION 15 998.05 2.07 CC. 23 996.53 3.40 CC. 16 997.90 2.20 cc. 24 996.29 3.61 cc. 17 997.74 2.34 cc. 25 996.04 3.83 cc. 18 997.56 2.49 cc. 26 995.79 4.06 cc. 19 997.38 2.65 cc. 27 995.52 4.31 cc. 20 997.18 2.82 cc. 28 995.24 4.56 cc. 21 996.97 3.00 cc. 29 994.96 4.82 cc. 22 996.76 3.19 cc. 30 994.66 5.08 cc. 74 A TEXTBOOK OF CHEMISTRY Heat of Fusion and Vaporization. If heat is applied to a kilogram of ice at the freezing point, 0, it will absorb 79 large calories in melting ; that is, if the same amount of heat is applied to a kilogram of ice and a kilogram of water, both of them at 0, when the ice is melted and still at the other water will be at a temperature of 79. If a kilogram of water at the boiling point, 100, is heated till it is all converted into steam at the same temperature and at atmospheric pressure it will absorb 536 large calories, or enough to warm 5.36 kilograms of water from the freezing point to the boiling point. The steam formed will fill a space of about 1700 liters, while the liquid water fills only one liter. The heat which is absorbed in the melting of ice and vaporiza- tion of water was formerly called latent heat because it seems to disappear, but this expression is not as fre- quently used now as it was some years ago. The heat of vaporization varies with the tem- perature, being greater at lower and less at higher temperatures. Vapor Pressure of Water. If two dry tubes, about 800 mm. long, are filled with mercury and inverted, the mercury will fall until the weight of the mercury in the tubes is the same as the weight of a column of air of the same cross section and reaching to the top of the atmosphere. In other words, the mercury will stand at the same height as the mercury in a barometer. If a drop of water is introduced into one of the tubes, the mer- cury in the tube will fall and remain at a position lower than that in the dry tube. If the tube containing the water is warmed, the mercury will fall further, if cooled, it will rise higher and for each temperature there will be a definite difference between the heights of the mercury in the two tubes. It is evident that this must be caused by the fact that a part of the water in the tube Fig. 30 VAPOR PRESSURE OF WATER 75 is converted into a vapor or gas and exerts a pressure on the mercury, partially balancing the pressure of the air. This pres- sure is called the pressure of water vapor, or sometimes, and less correctly, the aqueous tension. It is given for different tempera- tures in the following table : VAPOR PRESSURE OF ICE AND WATER TEMPERA- TUBES PRESSURE IN MILLIMETERS OF MERCURY TEMPERA- TURES PRESSURE IN MM. TEMPERA- TURES PRESSURE IN MM. -10 2.0 27 26.5 100.5 773.7 - 5 3.0 28 28.1 101.0 787.6 2 3.9 29 29.8 - 1 4.2 30 31.6 4.6 31 33.4 1 4.9 32 35.4 2 5.3 33 37.4 3 5.7 34 39.6 PRESSURE IN 4 6.1 35 Af\O 41.9 er er A ATMOSPHERES 5 6 6.5 7.0 40 50 55.0 92.2 111.7 1.5 7 7.5 60 149.2 120.6 2 8 8.0 70 233.8 127.8 2.5 9 8.6 80 355.5 133.9 3 10 9.2 90 526.0 144.0 4 11 9.8 95 634.0 159.2 6 12 10.5 96 657.7 170.8 8 13 11.2 97 682.1 180.3 10 14 11.9 98 707.3 188.4 12 15 12.7 99 733.2 195.5 14 16 13.6 99.1 735.9 201.9 16 17 14.5 99.2 738.5 207.7 18 18 15.4 99.3 741.2 213.0 20 19 16.4 99.4 743.9 '224.7 25 20 17.4 99.5 746.5 " 21 18.5 99.6 749.2 22 19.7 99.7 751.9 23 20.9 99.8 754.6 24 22.2 99.9 757.3 25 23.5 100.0 760.0 26 25.0 100.1 762.7 100.2 765.5 100.3 768.2 100.4 770.9 76 A TEXTBOOK OF CHEMISTRY Equilibrium. So long as water remains in the liquid form in the barometer tube described in the last paragraph, the volume of the tube above will have no effect on the vapor pressure of the water. This relation, which is very important, will be clearer from Fig. 31. Suppose that the cylinder contains water in the bottom with vapor of water above it and that it is fitted with an air-tight piston. If the piston is raised, the vapor, which acts like any other gas in this respect, will immediately expand and fill this additional space, and the pressure will be momentarily lowered. Immediately, however, some of the water will evaporate, and this evaporation will continue till the pressure is the same as before, provided that the water is kept at the same temperature. If the piston is pressed down, the reverse operation will occur. The vapor will be momentarily compressed and the pressure increased, but some of the vapor will immediately con- dense to water and in this way the pressure will fall to its original value. When two or more forms of a sub- stance are related in this way, they are said to be in equilibrium, that is, they are so balanced against each other that any change in temperature or pressure will cause a partial conversion of one form into the other. It is not necessary to suppose that the same molecules Fig. 31 of water are always in the form of vapor in such a case. On the contrary, it seems more probable that some of the water particles are all of the time leaving the liquid and pass- ing into the vapor and that molecules of vapor are constantly leaving the vapor and passing back into the liquid. When the two are in equilibrium, just as many molecules must pass in ohe direction as in the other within a given time. Effect of Water Vapor on the Pressure of a Gas. If a small glass bulb filled with water (Fig. 32) is placed in a bottle filled with dry air and closed with a stopper bearing a manometer to show the pressure within the bottle, on breaking the bulb it will be seen that the pressure within the bottle increases and after some time the increase will be almost exactly equal to the PHASES 77 vapor pressure of water at the temperature of the experiment. It is evident from this that as the water evaporates the vapor diffuses into the air above just as any other gas would do, and as it does so it adds its pressure to that of the air, in accordance with the law of partial pressures, that each gas in a mixture exerts the same pressure as though it were present alone (Dalton's law of partial pressures). The experiment suggests the proper method of finding the vol- ume which a quantity of dry gas would fill under standard condi- tions, when the gas has been meas- ured in a moist condition. The actual pressure exerted by the gas is less than the apparent pressure by the pressure exerted by the vapor of water at the given tem- perature. For this reason in the experiment described on p. 67 the direction was given to subtract from the reading of the barometer both the height of the mercury in the eudiometer and the pressure of water vapor, when the corrected volume of the moist hydrogen was to be found. Phases. Degrees of Freedom. Water may exist in the three forms of ice, water and vapor. These three" forms of the same substance are called phases. As long as only one phase is present we may change either the temperature or pressure or both at will, and in order to know the condition of the phase we must know both the temperature and the pressure. The system is said to have two degrees of freedom and the system is called divariant. When two phases are present, any change in the temperature will cause a corresponding change in the pressure, and as long as Fig. 32 78 A TEXTBOOK OF CHEMISTRY the temperature is fixed the pressure cannot be changed without the disappearance of one of the phases. Or, as long as the pres- sure is fixed, the temperature cannot be changed. Thus for water and vapor at a given temperature an increase of the volume does not decrease the pressure, but, instead of this, causes some of the water to change to vapor ; and it is only when the liquid phase disappears that a further increase in the volume causes a decrease in the pressure. In the same way, if water and ice are present, an increase in the pressure will cause some of the ice to melt and the- temperature will fall, and for every pressure there will be a corresponding temperature at which ice and water can exist together. To know the condition of such a system of two phases we need give only one factor. If we know the temperature, the pressure is fixed ; or if we know the pressure, the temperature is fixed by the properties of the substance. Such a system has only one degree of freedom and is called uni- variant. When the three phases, water, ice and vapor, are present, it is impossible to change either the temperature or the pressure without the disappearance of one of the phases. Such a system has no degree of freedom and is called invariant. The relations between temperature and pressure for the three phases of water may be seen clearly from the diagram (Fig. 33). If water and vapor are present, the relation between temperature and pressure is fixed by the line OA. For water and ice the re- lation is fixed by the line OC, from which it is apparent that an increase in the pressure lowers the melting point, though very slowly. The relation for ice and vapor is fixed by the line OB. The line OB' gives the relation for vapor and supercooled water, the vapor pressure being slightly greater than that of ice at the same temperature. The line OB f represents a condition of un- stable equilibrium, and if a little ice is introduced, the pressure will fall or the temperature will rise to the line OB, if all of the water freezes, or both pressure and temperature will rise to the point O, which is the invariant point, called also the triple point. SOLUTIONS 79 * The temperature of the system at the triple point will be H- 0.0073, since the of our scale is determined by freezing water under atmospheric pressure and a pressure of one atmosphere lowers the freezing point 0.0073. The vapor pressure of water at the freezing point is only 4.6 mm. Water as a Solvent. Solutes. If salt or sugar is placed in water, it quickly disappears and a homogeneous liquid is obtained, which we call a solution. Any substance which passes into solu- tion in this manner is called a solute. The liquid in which the |4. 6mm. Ice Vapor -10 * O.O073 TEMPERATURE Fig. 33 -HO solute dissolves is called a solvent. Solutes may be either solids, liquids or gases, and they vary very greatly in their degrees of solubility. Thus one liter of water will dissolve at 20 670 grams of sugar or 358 grams of common salt, but it will dissolve only 0.00153 gram of silver chloride. No satisfactory reason for such differences has been discovered, though many empirical relations between the composition of substances and their solubility are known. Some substances, as alcohol or sulfuric acid, dissolve in water in all proportions, but others will dissolve only up to a definite limit. When a solution can remain in contact with the solute 80 A TEXTBOOK OF CHEMISTRY without taking up any more, it is said to be saturated. The solid, liquid or gaseous phase of the pure solute is then in equilibrium with the solution very much as vapor of water is in equilibrium with liquid water. The solubility of salts usually increases with the temperature, but there are some- exceptions, and the rate of increase varies very greatly, as will be apparent from the accompany- ing diagram (Fig. 240 1 1 1 1 1 T? rr~i 1 , 34). If a warm satu- rated solution of a salt which is more soluble in warm than in cold water allowed to cool tO O O O QO O Ci O n ^ O 120 3 100 a s out of contact with the solid phase, a supersaturated solu- tion may usually be obtained. The in- troduction of a little of the solid will start the separation of the solid phase, and after a short time the solution will assume the normal, saturated condition. In a similar manner still water may be cooled below its freezing point or a vapor may be cooled or compressed below the point at which a part would ordinarily exist in the liquid form. Such a system is always unstable, somewhat after the analogy of a pyramid standing on its apex, and can only occur in the absence of the solid or liquid phase which should normally be present. 20 40 60 80 100 120 140 160 Temperature Fig. 34 SOLUTIONS 81 Chemical Activity in Solutions. Metathesis. If common salt, NaCl, and silver nitrate, AgNO 3 , are powdered and mixed to- gether, there will be no apparent action ; but if each is dissolved in water and the solutions mixed, there will be formed immedi- ately a white precipitate of silver chloride ; and if the solution is filtered from the precipitate and the filtrate evaporated, crystals of sodium nitrate may be obtained : NaCl + AgNO 3 = AgCl + NaNO 3 Sodium Chloride Silver Silver Sodium (common salt) Nitrate Chloride Nitrate Hundreds of illustrations of similar reactions which do not occur readily between the solid substances but which take place easily in solutions might be given. This reactivity of substances in solution is evidently in part because they are brought into inti- mate contact, since no combination can take place except be- tween substances which are touching each other. But this does not appear to be the only reason. In very many cases when clear, sharp-cut reactions occur, each compound separates, as here, into the metal and an acid radical. If we subject these same compounds in solution to the influence of an electrical current, the metal will travel toward the cathode through the solution while the acid radical will travel toward the anode. It seems, therefore, that solution in some way loosens the combina- tion between the ions so that they can very readily enter into new combinations. A reaction of the sort just considered is called a double decom- position or metathesis. Hydrates, Deliquescence, Efflorescence. Many salts when they separate in crystals from solution do so in combination with a definite quantity of water. Thus crystals of copper sulfate have the composition CuSO 4 .5 H 2 O ; crystals of sodium sulfate, the composition Na 2 SO 4 .10H 2 O. Such compounds usually decompose rather easily into water and the anhydrous salt, and the water is spoken of as water of hydration, and the compounds 82 A TEXTBOOK OF CHEMISTRY are called hydrates. 1 If the hydrate is placed above the mercury in a barometer tube it will decompose, giving off water vapor till a definite vapor pressure is reached. This vapor pressure will vary greatly for different hydrates and will increase with the tem- perature as the vapor pressure of water does. Thus the vapor pressure of the hydrate CuSO 4 .5 H 2 O is 12.5 mm. at 30, the vapor pressure of CaCl2.H 2 O is only 3.1 mm., the vapor pressure of Na2SO4.10 H2O is 25.3 mm. and the vapor pressure of pure water is 31.6 mm. If an anhydrous salt like calcium chloride is exposed to air containing enough moisture so that the pressure of the water vapor in it exceeds 3.1 mm. at 30, water will be absorbed and the hydrate will be formed. In this case even a concentrated solution of calcium chloride has so low a vapor pressure that it is exceeded by that of the moisture in ordinary air. Accordingly calcium chloride when exposed to the air absorbs moisture and finally dissolves in the water absorbed. Such a salt is said to be deliquescent. On the other hand, if the hydrate, Na 2 SO 4 .10 H 2 O, is ex- posed to air in which the vapor pressure of the water which it contains is less than 25.3 mm. at 30 the salt will decompose and lose water to the air. As it does so it will fall to a fine powder or flour. Salts of this type are called efflorescent. Natural Waters. The water which is found in nature is never pure, the purest being rain water falling in the open country after it has been raining for some time, or water obtained by melting the ice from a pure, fresh-water lake. Even such water 1 The term water of crystallization, which is still used by many authors, is not as appropriate. During the first half of the nine- teenth century sodium hydroxide, NaOH, and calcium hydroxide, Ca(OH) 2 , were called sodium hydrate and calcium hydrate, and their formulas were written in a form which with modern atomic weights would be Na 2 O.H 2 O and CaO.H ? O. This older use of the word hydrate, which still clings to the literature of pharmacy, has inter- fered somewhat with the introduction of the word in the sense in which it is here defined. It is fair to say, too, that the distinction between hydrates and hydroxides is more or less arbitrary, as some hydroxides lose water more easily than some hydrates, and in many cases the water of hydration cannot be removed without decomposition of the rest of the salt. NATURAL WATERS 83 contains air in solution and a little carbonic acid from the carbon dioxide of the air. On falling upon the ground rain water begins at once to take up various substances, partly in suspension, partly in solution. Calcium or magnesium carbonate and calcium sulfate cause the water to become hard (pp. 310, 311) and injure it seriously for use in steam boilers or in laundries. If the water is mixed with sewage, it frequently becomes contaminated with bacteria which produce disease. The diseases of typhoid fever and of cholera, especially, are frequently transmitted in this way. During an epidemic of cholera in the cities of Hamburg' and Altona, Germany, the people in the houses on one side of a certain street used a contaminated water which caused very many cases of the disease. Across the street, water from the same source, but after passing a system of public filters, was used and there were very few cases. In Chicago" before the opening of the drainage canal there were 170 deaths from typhoid fever per year for each 100,000 people. After the drainage canal carried away the sewage which had formerly gone into Lake Michigan and contaminated the water supply of the city, the death rate from typhoid fell to 16 per 100,000. Purification of Water. Water may be purified most com- pletely by distillation, though special precautions are required to get rid of ammonia, carbon dioxide and other volatile impuri- ties. The bacteria in water may be killed almost completely by boiling the water for a short time, and this should always be done when it is necessary to use a suspected water for drinking or the preparation of food. The bacteria may also be almost completely removed by filtration, either on a large scale on beds of sand, or through filters of fine-grained stone or unglazed por- celain. Charcoal filters, which were formerly used, are rarely effective. Waters may also be sterilized by treatment with ozone, with ultra-violet light, or with bleaching powder. Hydrogen Peroxide. A second compound of hydrogen with oxygen, called hydrogen peroxide and having the formula H 2 O2, can be prepared by the action of acids on peroxides of univalent or bivalent metals : 84 A TEXTBOOK OF CHEMISTRY Na 2 O 2 + 2 HC1 = 2 NaCl + H 2 O 2 Sodium Peroxide Hydrogen Peroxide Ba0 2 -f H 2 SO 4 = BaSO 4 Barium Peroxide Barium Sulfate H 2 O Barium sulfate is almost insoluble in water, and if the barium peroxide and sulfuric acid are used in equiv- alent amounts a solution may be obtained which contains practically nothing except hydrogen peroxide and water. Such a solu- tion may be concentrated in a vacuum des- iccator (Fig. 35) over sulfuric acid, the vapor pressure of water being much greater than that of hydrogen peroxide, or it may be distilled under the low pressure obtained by means of a good air pump or filter pump. Water will distil away first, and finally the hydrogen peroxide will distil, almost pure. Hydrogen peroxide is a heavy liquid which decomposes slowly at ordinary temperatures Fig. 35 into water and oxygen : H 2 O 2 = H 2 + O This reaction is accompanied by the evolution of 46,200 cal- ories of heat (by the decomposition of 68 grams of hydrogen peroxide), and, as is usual with reactions evolving heat and lib- erating a gas, may become explosive. The reaction is catalyzed by many substances, especially by finely divided platinum and other metals, and it is quite dangerous to bring concentrated solutions of hydrogen peroxide into contact with organic sub- stances. Hydrogen peroxide is a powerful oxidizing agent, giving oxygen readily to many substances. This property, also, is closely re- lated with the fact that heat is evolved when it decomposes. If a solution of hydrogen peroxide is added to the black pre- HYDROGEN PEROXIDE 85 cipitate of lead sulfide suspended in water, the substance is oxi- dized to white lead sulfate : PbS + 4 H 2 O 2 = PbSO 4 + 4 H 2 O Hydrogen peroxide also acts on silver oxide and on a good many other substances as a reducing agent, but in all such cases oxy- gen is liberated. We may explain this by considering that oxygen gas is a compound of oxygen with itself l and that the oxy- gen of the hydrogen peroxide oxidizes the oxygen of the silver oxide to free oxygen. Or it may be that the oxygen of the silver oxide oxidizes the hydrogen of the hydrogen peroxide to water, leaving the two oxygen atoms of the peroxide combined together as free oxygen, O 2 2 : H 2 2 + A g2 = H 2 + 2 + 2Ag The reduction of potassium permanganate, KMn0 4 , in acid solution is frequently used for the quantitative deter- mination of hydrogen peroxide. The equation representing this reaction is most easily written in two parts, as indicated below. The first part of the reaction is based on the facts that potassium is univalent and manganese bivalent when combined with acid radicals and that, for this reason, when two molecules of potassium permanganate are acted on by sulfuric acid in the presence of some substance which can take up oxygen, five atoms of oxygen become available for the oxidation of that other sub- stance. These five atoms of oxygen are inclosed in brackets to indicate that the decomposition does not take place except in the presence of something with which this oxygen can combine : 2 KMn0 4 + 3 H 2 SO 4 = K 2 SO 4 + 2 MnSO 4 + 3 H 2 O + (5 O) (5 O) + 5 H 2 O 2 = 5 H 2 O + 5 O 2 2 KMnO 4 + 3 H 2 SO 4 + 5 H 2 O 2 = K 2 SO 4 + 2 MnSO 4 + 8 H 2 O + 5 O 2 1 Other reasons for such a view will be given later (p. 93). 2 The fact that the reaction H 2 O 2 = H 2 O + O is monomolecular points strongly to the latter explanation. 86 A TEXTBOOK OF CHEMISTRY The last equation is obtained by combining the other two algebraically, eliminating the five atoms of oxygen which appear on opposite sides. * Properties and Uses of Hydrogen Peroxide. ~ Hydrogen peroxide has a specific gravity of 1.4584 at 0. It boils at 69.2 under a pressure of 26 mm. or at 84-85 under a pressure of 68 mm. A dilute solution is fairly stable when pure, the stabil- ity being increased by the presence of a small amount of a mineral acid. The stability is decreased by alkalies and by many other substances, and the decomposition is also hastened by exposure to light. Hydrogen peroxide is used in medicine as a bactericide and for the diagnosis of pus, which causes its rapid decomposition with evolution of oxygen. The medicinal solution commonly used is known as a 10-volume solution, meaning that it evolves ten times its volume of oxygen when it decomposes, or 20 times its volume when treated with an oxidizing agent, one half of the oxygen coming from the latter, as explained above. Hydrogen peroxide is also used to bleach hair, silk and wool, being much more suitable than chlorine or hypochlorites (p. 127) for this purpose. A solution suitable for this purpose may be obtained by dissolving sodium peroxide in cold water and adding dilute sulfuric acid. * Tests for Hydrogen Peroxide. With potassium dichromate in an acid solution hydrogen peroxide gives a beautiful blue compound which is soluble in ether. The composition of the compound is not positively known. Its formation may be used either as a test for hydrogen peroxide or as a test for chro- mium. Another valuable test is the yellow color given with solutions containing titanium. * Structure of Hydrogen Peroxide. Two formulas have been proposed to represent the structure of hydrogen peroxide : H O O H and "> = O LAW OF MULTIPLE PROPORTION 87 The first represents both oxygen atoms as bivalent, the second represents one of the oxygen atoms as quadrivalent. The evi- dence in favor of the first formula is : 1. There are a number of reactions in which hydrogen peroxide seems to be formed by the reduction of oxygen, which has the formula O 2 (p. 93) : O=O + H-H = H O O H 2. Hydrogen peroxide does not seem to be formed by the oxi- dation of water, as it should be if the second formula were true. 3. When the two hydrogen atoms of hydrogen peroxide are replaced by the ethyl group, C 2 H5, diethyl peroxide, (2115)202, is formed and this gives ethyl alcohol, C2H 5 O H, by re- duction. This is easily explained by the first formula : OjjHs o o G 2 ri5 -p 12 == 2 C^HS o M If diethyl peroxide had a formula corresponding to the second one for hydrogen peroxide, it should give ethyl ether, C 2 H 5 O C 2 H 5 by reduction (Baeyer and Villiger, Ber. 33, 3387 (1900)) : C 2 H 5 Law of Multiple Proportion. In water and hydrogen peroxide the same elements, hydrogen and oxygen, combine to form two different compounds. In water one part of hydrogen combines with eight parts of oxygen (or 2 with 16), while in hydrogen per- oxide one part of hydrogen combines with sixteen parts of oxy- gen (or 2 with 32). Nitrogen and oxygen form a whole series of compounds represented by the formulas and composition : N:O Nitrous oxide, N 2 O 28 : 16 Nitric oxide, NO 14 : 16 Nitrogen trioxide, N 2 O 3 28 : 48 Nitrogen tetroxide, N 2 O 4 28 : 64 Nitrogen pentoxide, N 2 O 5 28 : 80 88 A TEXTBOOK OF CHEMISTRY In every such case, if we consider some fixed amount of one oj the elements (it makes no difference which one] the amounts of the other combining with this fixed amount will bear a simple ratio to each other. This is known as the Law of multiple proportions. The discovery of this law led Dalton to propose the atomic theory of the constitution of matter. A little consideration of the law shows that it follows, necessarily, from the law of com- bining weights (p. 13), and that the law of combining weights is more comprehensive and important. Two compounds of iron contain : Ferrous oxide, 77.73 per cent of iron 22.27 per cent of oxygen 100. Ferric oxide, 69.94 per cent of iron 30.06 per cent of oxygen To67~ Show that these proportions are in accordance with the law of multiple proportion. This illustration helps us to understand why the law was not discovered before the nineteenth century. CHAPTER VI AVOGADRO'S LAW. SELECTION OF ATOMIC WEIGHTS. OZONE Gay Lussac's Law of Combining Volumes. It has been shown that hydrogen and oxygen unite very nearly in the proportion of two to one by volume and that the volume of steam formed is very nearly the same as the volume of hydrogen which it contains. A study of many other gaseous elements and their compounds has shown that in every case there is a simple ratio between the volumes of gases which combine with each other and also between those volumes and the volume of the product, if that is a gas. This is known as Gay Lussac's law of combining volumes, and it is true of all elements or compounds which can be converted into gases without decomposition, as well as of substances which are gases at ordinary temperatures. It has been shown to be true for thousands of compounds. The law may be illustrated by the following diagrams : 36.5 grains 36.5 grams Hydrocloric Hydrocloric acid acid 90 A TEXTBOOK OF CHEMISTRY 2 grams Hydrogen 2 grams Hydrogen 2 grams Hydrogen 17 grams Ammonia 17 grams Ammonia 2 grams 2 grams Hydrogen Hydrogen 2 grams 2 grams Hydrogen Hydrogen 2 grams 2 grams Hydrogen Hydrogen 34 grams 34 grams Phosphine Phosphine 34 grams 34 grams Phosphine Phosphine The volumes have been chosen for these illustrations in such a way that the unit volume always contains one gram, or a whole number of grams, of hydrogen. In every case the weight of this unit volume bears a simple relation to the atomic weights of the ele- ments which it contains. This volume of any gaseous element or compound always contains one gram-atom or a whole number of gram-atoms of each element contained in the gas. This is not accidental, but follows of necessity from the two laws : (V\ that AVOGADRO'S LAW 91 the composition of every compound can be expressed by multi- ples of the atomic weights of the elements which compose it ; and (2) that there is always a simple ratio between the volumes of gases which combine and also between those volumes and the volume of the product, if that is a gas. A further examination of the illustrations shows that the unit volume of each compound contains one gram molecule of the compound. This result does not follow of necessity from the two laws just given. It depends on the values which we give to the atomic weights. Thus if we were to call the atomic weight of oxygen 8 and the formula of water HO, as was done by Dalton, the gram molecular weight of water would be 9 and the unit volume which has been chosen would contain two gram mole- cules of water. As has been stated, we have chosen as the unit volume for our illustration that volume which, in several com- pounds, contains one gram of hydrogen, of course because hy- drogen is our (approximate) unit for atomic weights. No com- pound of oxygen is known which contains less than 16 grams in this unit volume. Avogadro's Law. The selection of 16 instead of 8 as the atomic weight of oxygen is based on a hypothesis proposed by Avogadro, an Italian chemist, in 1811. This hypothesis is that all gases, under the same conditions of temperature and pressure, contain equal numbers of molecules in equal volumes. Some of the many facts which support this hypothesis so fully that it may now be considered as an established law are : 1. While solids and liquids vary greatly in their rate of expan- sion or contraction for changes of pressure or of temperature, all gases expand and contract alike. This points very strongly to a similarity in their structure. 2. When the law of combining weights is combined with the law of combining volumes it follows, of necessity, that there must be a simple ratio between the numbers of molecules in equal volumes of different gases. But if the ratio is one of simple whole numbers, it seems highly improbable that it is not one of equality. Thus it would seem very improbable that the number 92 A TEXTBOOK OF CHEMISTRY of molecules in a given volume of steam is exactly twice the number in the same volume of hydrochloric acid, as we should have to suppose if the molecular weight of water is 9 while that of hydrochloric acid is 36.5. 3. The kinetic theory of gases, which accounts so well for their properties, leads directly to Avogadro's law on the basis of the laws of the collision of elastic bodies (p. 58). 4. The atomic weights selected on the basis of Avogadro's law have made possible the classification of the elements in groups according to their atomic weights. This classification is known as the Periodic System (p. 132) and furnishes very strong, independent evidence that the atomic weights selected are in reality the true relative weights of the atoms. Selection of an Atomic Weight. According to Avogadro's law, equal volumes of different gases, under the same conditions of temperature and pressure, must have weights proportional to the weights of the molecules of the gases. If the molecule of one gas is twice as heavy as that of another, one liter of the first gas must weigh twice as much as a liter of the second. Accordingly if we can take as our unit volume the volume filled by a gram molecule of some compound which contains only one atom of hydrogen in its molecule, this unit volume will contain one gram molecule of every other gaseous element or compound. We can- not, of course, look at the molecules of different compounds to discover which one contains only a single atom of hydrogen ; but if we weigh the same volume of hydrochloric acid, steam and ammonia, we find that in a given volume of steam there is twice as much hydrogen as there is in the same volume of hydrochloric acid, and that in the same volume of ammonia there is three times as much hydrogen. As no compound has ever been found which contains less hydrogen than hydrochloric acid does in the unit volume, it seems pretty safe to conclude that there is only one atom of hydrogen in a molecule of this compound. As one gram of hydrogen combines with 35.5 grams of chlorine, 1 the volume 1 Approximate values are used for convenience, as always. The true values are 1.0078 grams of hydrogen for 35.46 grams of chlorine. SELECTION OF ATOMIC WEIG-HTS 93 filled by 36.5 grams of hydrochloric acid will contain one gram molecule, and it must contain one gram molecule of every other gas. In the discussion above it is pointed out that those compounds of hydrogen which contain the smallest amount of hydrogen in the unit volume probably contain only one atom of hydrogen in their molecules. Reasoning in the same way, we may find for each element those compounds which contain the smallest amount of the element in the unit volume, and it is probable that these compounds contain only one atom of the element in the molecule. Thus steam contains 16 grams of oxygen in the unit volume ; and as no compound of oxygen containing a smaller amount is known, we conclude that there is only one atom of oxygen in a molecule of steam and that the atomic weight of oxygen is 16. The atomic weights of the other elements which form gaseous compounds, or compounds which can be converted into gases without decomposition, have been selected in the same way. Molecules of the Elements. It was natural for Dalton when he proposed the atomic theory to think of the atom as the small- est particle of an element in the free state, and it did not occur to him that atoms of the same kind could combine. A refer- ence to the diagrams, however, shows that there is twice as much oxygen in the unit volume of oxygen gas as in the unit volume of steam. According to Avogadro's law it follows that oxygen gas contains two atoms of oxygen in each molecule. We may reach the same conclusion by another process. If we let each square below represent 1000 molecules, it is clear that 2000 molecules of steam are formed from 1000 molecules of oxygen, and, as each molecule of steam must contain at least one atom 94 A TEXTBOOK OF CHEMISTRY of oxygen, the 1000 molecules of oxygen must consist of 2000 atoms, or each molecule contains two atoms. It will be seen from the diagrams that four atoms combine to form a molecule of phosphorus vapor, while a molecule of mercury vapor contains only a single atom. In the latter case atom and molecule are identical. In general the atoms of non- metallic elements combine to form molecules of the element in the free state, but the atoms of the metals show little tendency to combine in this way. For a probable explanation of this re- markable difference see J. J. Thompson, The Corpuscular Theory of Matter, p. 120 ; and H. N. McCoy, J. Am. Chem. Soc. 33, 273. Gram Molecular Volume. The unit volume, which will con- tain one gram molecule of any gas, is best calculated from the weight of a liter of oxygen, since oxygen is the basis for atomic weights (p. 68). One gram molecule of oxygen, C>2, contains 32 grams, and, as a liter of oxygen weighs 1.429 grams, the gram 32 __ molecular volume for oxygen will be = 22.4 liters at and 760 mm. At the same temperature and pressure one gram molecule of any other gas will fill the same volume, 22.4 liters. This statement and the law of Avogadro are subject to limita- tions similar to those which apply to the laws of Boyle and Charles. Just as most gases, and especially those which are easily liquefied, contract too much when compressed from one to two atmospheres pressure and also contract too much when cooled from 100 to 0, so almost all gases are heavier than they should be in accordance with the law of Avogadro. As the volume increases under diminished pressure, however, gases ap- proach the condition of an "ideal " gas, and at low pressures the densities of gases agree very closely indeed with the law. The amount of the deviation from the law at atmospheric pressure and the agreement under low pressures will be obvious from the following table : DENSITY OF GASES 95 DENSITY OF GASES l NAME FOBMULA WEIGHT OF ONE LITER WEIGHT OF 22.4 LITERS AT AND 760 MM. DENSITY AT Low PRESSURE = 32 MOLECU- LAR WEIGHT Oxygen . . . 2 1.429 32.00 32.00 32. Hydrogen . . H 2 0.08987 1.997 2.01 2.016 Nitrogen . . . N 2 1.2507 28.02 28.01 28.02 Carbon monoxide CO 1.2504 28.01 28.00 28.00 Nitric oxide . . NO 1.3402 30.02 30.01 30.01 Argon .... Ar 1.7808 39.89 39.88 39.88 Carbon dioxide . C0 2 1.9768 44.28 44.01 44.00 Nitrous oxide N 2 1.9777 44.30 44.03 44.02 Hydrochloric acid HC1 1.6398 36.73 36.47 36.47 Ammonia . . NH 3 0.7708 17.27 17.01 17.03 Sulfur dioxide . SO 2 2.9266 65.56 64.07 64.07 Air 1.2928 28.96 The table shows that at low pressures the deviations from Avo- gadro's law scarcely exceed the experimental errors of the deter- minations. If the pressure and volume of a gas are corrected for the volumes occupied by the molecules and for their attractions for each other, corrections which can be determined experimentally (van der Waals) , the law of Avogadro also becomes almost rigor- ously exact. The weight of 22.4 liters of air furnishes a very simple method of calculating the density of any gas as compared with air. For approximate estimates the weight may be taken as 29 grams, and the molecular weight of any compound divided by 29 will be approximately its density as compared ' with that of air. Thus it will be seen that hydrogen is 14 J times lighter than air, while oxygen is 1.1 times as heavy and carbon dioxide one and one half times as heavy. * Number of Molecules in one Cubic Centimeter of a Gas. Avogadro's law was established with a high degree of probability before any means of estimating the number of molecules in a 1 Guye, J. Am. Chem. Soc. 30, 155 (1908). 96 A TEXTBOOK OF CHEMISTRY given volume of a gas was discovered. It is of some interest to know, however, that several different methods of estimating this number are now known and that the results obtained by different methods are in fair agreement. Two of these methods may be given here, in outline. It has been shown that the element radium slowly decomposes and that as it does so it shoots out atoms of helium with a tremen- dous velocity. The volume of helium given off by a gram of radium has been measured and is 0.46 cubic millimeter in a day. When an atom of helium shot out by radium hits a screen of zinc sulfide, it produces a flash of light ; and as the helium atoms are sent out in all directions equally, by placing a screen of zinc sulfide back of a small opening which is at a known distance from a weighed amount of radium and counting the flashes, it is possible to estimate the number of atoms of helium which pass the opening in a given time and so the number of atoms in a cubic centimeter of the gas. Rutherford has estimated the number of molecules in one cubic centimeter of helium, in this way, as 2.56 X 10 19 (or 25,600,000,000,000,000,000). (Report of the Winnipeg Meeting of the British Association, 1909, p. 377.) When very minute particles suspended in water are observed with a microscope, it is seen that they are never at rest, but move about constantly in a wholly irregular manner. This was first noticed by the English botanist Brown in 1827 and is called the Brownian movement. It has long been considered as an evi- dence that water and other liquids are composed of molecules which are in rapid motion. Perrin has succeeded in showing (1909) that in an emulsion of gamboge the minute particles are distributed as they should be if they are considered as very large molecules, and he has connected this distribution with the kinetic theory in such a manner as to furnish an estimate of the number of molecules in a cubic centimeter of a gas, which he gives as 3.15 X 10 19 . A number of other methods of estimating the same quantity give results of the same general value. As the different methods are quite independent of each other, we can OZONE 97 have considerable confidence that the values are approximately correct. 1 Allotropic Forms. Ozone. When oxygen is subjected to the action of electrical waves obtained by connecting the tinfoil at A and D (Fig. 36), with the poles of an induction coil, from 5 to 8 per cent of the gas is converted into ozone. If the mixture * Q 1 r^A- (( w\ ^ 1 1 B % =:i D : I ^J. 1 === ^ fa 1 112 Fig. 36 of oxygen and ozone is partly liquefied by passing it into a flask immersed in liquid air, a dark blue liquid is obtained. Oxygen boils at 182.5, while ozone boils at 119, and a mixture of oxygen and ozone containing 84 per cent of ozone has been ob- tained by allowing the oxygen to boil away from such a mixture. (Ladenburg, Ber. 31, 2508, 2830 (1898).) Ozone has a strong odor, which is noticed in the neighborhood of electrical machines. The weight of the gram molecular volume (22.4 liters) of the gas is 48 grams. Hence the formula is O 3 . Ozone is evidently formed in accordance with the equa- tion : The reaction is reversible and the ozone formed is unstable, decomposing slowly at ordinary temperatures, rapidly and completely at 250-300. The decomposition is accompanied by considerable evolution of heat, and so liquid ozone may easily give violent explosions. One gram molecule, 48 grams, gives 1 Professor R. A. Millikan of the University of Chicago has recently determined the number by a new method which is, appar- ently, much more accurate than any previously used. He gives the value 2.71 x 10 19 . 98 A TEXTBOOK OF CHEMISTRY out 29,400 calories in decomposing, and the properties of ozone are intimately connected with the fact that it contains so much more chemical energy than ordinary oxygen. It is a vigorous oxidizing agent and attacks metallic silver and many other sub- stances which are not affected by ordinary oxygen. Ozone may be detected by its action on moist potassium iodide starch paper : O 3 -f 2 KI + H 2 O = 2 KOH + I 2 -f O 2 The liberated iodine gives a deep blue color with the starch. A small amount of ozone is formed when a piece of clean phos- phorus, half covered with water, is allowed to stand for a short time in a bottle rilled with air. The ozone may be detected by the odor and by the potassium iodide starch paper. Ozone has been used to a limited extent for sterilizing drinking water, but this use has not proved very successful. Bleaching powder is cheaper and much more efficient. Ozone is, however, a powerful germicide, and it may be that the ozone generated in thunderstorms plays a beneficent part in nature. In ordinary oxygen and ozone we have two different forms of the same element, and it is evident that in this case we may ob- tain another substance from an element without adding any- thing to it. The formulas of the two forms of oxygen give us a partial explanation of the difference between the two substances. Both forms are really compounds in which oxygen is combined with itself. We can represent this graphically by the formulas : O O=O and /\ or O = O=O o o The last formula represents one of the oxygen atoms as quadrivalent. Several elements beside oxygen exist in two or more forms. Such forms are called allotropic. OZONE 99 EXERCISES 1. What will be the volume of a gram molecule of oxygen at 20 and 750 mm. ? 2. How many grams of sodium peroxide will be required to give 22.4 liters of oxygen under standard conditions ? 3. How many grams of potassium chlorate will be required to give 22.4 liters of oxygen ? 4. How many liters of carbon dioxide will be formed by burning 6 grams of carbon ? 5. How many liters of carbon monoxide ? How many liters of oxygen will be required for the reaction in each case ? 6. How many grams of iron will be required to give 22.4 liters of hydrogen, if the iron is dissolved in hydrochloric acid ? How many grams, if dissolved in sulfuric acid-? How many grams, if used to decompose steam ? 7. How much heat could be obtained, if two grams hydrogen could be burned in ozone ? 8. How many grams of sodium are required to give 22.4 liters of hydrogen ? 9. How many milligrams of aluminium are required to give 22.4 cc. of hydrogen ? CHAPTER VII CHLORINE SYMBOL, Cl. ATOMIC WEIGHT, 35.46. FORMULA, C1 2 . Occurrence of Chlorine. Chlorine is not found free in nature, chiefly because of its strong affinity for almost every other ele- ment and especially for the metals. Its most important com- pound is common salt, sodium chloride, NaCl. This is found in large amounts in sea water, in strong brines from artesian wells in very many places and in enormous beds of rock salt, some- times a hundred feet in thickness. Calcium chloride and mag- nesium chloride are also found in sea water and in many of the brines. Silver chloride is sometimes an important ore of silver. Preparation of Chlorine. 1. By Electrolysis of Sodium Chlo- ride. If an electric current is passed through a solution of common salt, NaCl, using a carbon anode and mercury as a cathode, the negative chloride ions, Cl~, will be discharged at the anode and the chlorine will be evolved as a gas. The sodium will dissolve in the mercury cathode, and by appropriate means it may be caused to react with water, giving sodium hydroxide and hydrogen. This process is now used extensively in the manu- facture of chlorine and caustic soda or sodium hydroxide. It is known as the Castner-Kellner process and will be considered further in connection with sodium hydroxide (p. 402). 2. Preparation by Oxidation of Hydrochloric Acid. In the preparation of oxygen, compounds (mercuric oxide or potassium chlorate) are selected which decompose with liberation of the gas when heated. In the preparation of hydrogen, compounds (water or hydrochloric or sulfuric acid) are selected from which the hydrogen can be displaced by another element. Both methods may be used for the preparation of chlorine. 100 CHLORINE 101 In all of the methods practically used except electrolysis, the chlorine of hydrochloric acid is displaced by oxygen, one atom of the bivalent oxygen displacing two atoms of chlorine : -H- ^i \ /-\ -H- \/^ I /"<i /~<] TJ f^l I" ^ TT S^* T ^"^ ^l 1 d ll/ In several methods compounds of manganese are used, and these methods depend upon the fact that while manganese forms compounds with oxygen in which it is quadrivalent, or in which it even shows a higher valence, the compounds of the element with chlorine which contain more than two atoms of chlorine for one of manganese are very unstable. The compounds with .0 oxygen are manganese dioxide, Mn^ , and potassium perman- ganate, K O Mn^O. The stable compound with chlorine % is Mn<g- If a concentrated solution of hydrochloric acid is poured on manganese dioxide, there is formed, at first, a dark brown solu- tion which probably contains some manganese tetrachloride : Mn0 2 + 4HC1 = MnCl 4 + 2 H 2 O Manganese Hydrochloric Manganese Dioxide Acid Tetrachloride This equation represents the reaction as a metathesis in which two compounds each separate into two parts, and then one part of each combines with one part of the other compound. A very large majority of chemical reactions belong to this class, and the principle is very useful in writing and understanding chemical equations. In the present case the manganese tetrachloride is so unstable that only indirect evidence of its existence has been obtained. When the solution is warmed gently, it is de- composed completely into manganous chloride and chlorine : MnCl 4 = MnCl 2 + C1 8 102 A TEXTBOOK ' OF CHEMISTRY The two reactions, which go on simultaneously, may be ex- pressed in one equation, thus : Mn0 2 + 4 HC1 = MnCl 2 + C1 2 -f 2 H 2 O If concentrated hydrochloric acid is al- lowed to drop on potassium permanganate, KMnC>4, which is a powerful oxidizing agent, the hydrochloric acid is oxidized and chlorine is liberated. A part of the chlorine, of course, remains combined with the potassium and manganese. In writing the equation we notice that we must have 8 molecules of hydrochloric acid to furnish the hydrogen to combine with the 4 atoms of oxygen in one mole- cule of the potassium permanganate, and we write, at first : KMnO 4 + 8 HC1 = KC1 + MnCl 2 + 4 H 2 O + 5 Cl But this form of the equation gives an odd number of atoms of chlorine, and as the chlorine is liberated in the molecular form it is necessary to double the equation to represent the substances actually formed : 2 KMnO 4 + 16 HC1 = 2 KC1 + 2 MnCl 2 + 8 H 2 O + 5 C1 2 3. Preparation of Chlorine by the Deacon Process. In the methods described in the last paragraph for the preparation of chlorine the element is liberated from hydrochloric acid by means of an expensive oxidizing agent. Henry Deacon of Eng- land devised a process a good many years ago by which the ex- pensive oxidizing agent is replaced by the oxygen of the air. If a mixture of hydrochloric acid and oxygen, or air, is heated, the reversible reaction represented by the equation : 4 HC1 + O 2 ^t 2 H 2 O + 2 C1 2 Fig. 37 CHLORINE 103 takes place, but at a temperature of 350 to 400 the reaction is too slow to be commercially possible, as time is a very important element in all technical processes. The reaction might be has- tened by using a higher temperature, but at a high temperature the equilibrium between the four substances is displaced in such a way that less chlorine is formed and more of the hydrochloric acid passes through the heated apparatus unchanged (p. 110). By the use of a catalyzer, however, the reaction at lower tem- peratures can be hastened so far as to become technically pos- sible. Several different catalyzers may be used, but Deacon found that copper chloride, in the form obtained by saturating pumice with a solution of the salt and drying, is most suitable. With this catalyzer the reaction is sufficiently rapid so that, at 345, 80 per cent of the hydrochloric acid can be oxidized to chlorine. The process does not, however, compete successfully with other processes for the manufacture of chlorine. * 4. The Weldon Process for Chlorine. The oxygen of the air is also used, indirectly, as the oxidizing agent for the hydro- chloric acid in the Weldon process. When calcium hydroxide (in the form of milk of lime) is added to the solution of manganese chloride obtained in the preparation of chlorine by means of manganese dioxide, a precipitate of manganese hydroxide is formed : /Cl /O H /OH /Cl Mn< +Ca< =Mn< + Ca< X C1 X O H X OH X C1 or MnCl 2 + Ca(OH) 2 = Mn(OH) 2 + CaCl 2 The manganese hydroxide, if exposed to the air, takes up OH oxygen and forms hydrated manganese dioxide, O = Mn( X OH (or MnO 2 .H 2 O). The oxidation is hastened, practically, by passing air and steam into the mixture. The addition of the oxygen to the manganese causes the hydrogen of the hydroxyl groups to become acid in character (see p. 206), and the com- 104 A TEXTBOOK OF CHEMISTRY pound reacts with more of the calcium hydroxide to form cal- cium manganite, CaMnOa : /O H /OH /O v O=Mr< + Ca< = O=Mr< >Ca + 2 H 2 O X H \)H XX Calcium Manganite The calcium manganite is insoluble and may be easily sepa- rated from the solution of calcium chloride. On treatment with hydrochloric acid it acts in the same way as a mixture of man- ganese dioxide and lime would. CaMnO 3 + 6 HC1 = CaCl 2 + MnCl 2 + C1 2 + 3 H 2 O The solution of manganese chloride may, of course, be treated with milk of lime, air and steam and the cycle repeated indefi- nitely. This process was used for many years in manufacturing chlorine and bleaching powder on a large scale, but it has now been replaced almost entirely by the electrolytic processes, which are simpler and more direct. Properties of Chlorine. Chlorine is a greenish yellow gas about two and one half times (, see p. 95) as heavy as air. It has a characteristic odor and, even when diluted with a large volume of air, attacks the nose and lungs strongly, producing the effect of a severe cold. The best antidote is to breathe, at once, the vapors of strong alcohol. A larger quantity of the gas acts as a violent poison and may produce fatal effects. Water dissolves about twice its volume of the gas and for laboratory experiments it is usually collected by displacement of air, in upright jars. It is less soluble in a concentrated solu- tion of salt. Chlorine may be condensed to a liquid by cold or by pressure. The liquid boils at 33.6 under atmospheric pressure and freezes at 102. The vapor pressure at is 3.66 atmospheres. Chlorine is a very active element and forms compounds with all of the elements except fluorine and those of the argon family. With many of the elements it will combine directly and rapidly CHLORINE 105 at ordinary temperatures and with others it combines at a much lower temperature than does oxygen. Chlorine and hydrogen combine too slowly for the rate to be measured, if the mixture of the gases is kept in the dark, but the mixture will explode at a much lower temperature than mix- tures of oxygen and hydrogen. If the mixture is exposed to diffused daylight, the elements combine slowly at ordinary temperatures, and if exposed to bright sunlight or to the light from burning magnesium, the combination is so rapid as to pro- duce a violent explosion. The effect is produced by light and not by heat, and the rays of light at the violet end of the spec- trum are especially effective, just as the same light rays are most effective in producing changes in photographic plates. Evi- dently the light vibrations set up or increase some kind of vibra- tion within the molecules or atoms of chlorine, which makes these more active, but we can form only a very vague idea of the mechanism of the action. Chlorine containing a minute amount of moisture will attack almost all of the metals, even at ordinary temperatures, and more rapidly on gentle warming. A strip of copper, if warmed and held in the gas, takes fire and burns to cuprous chloride, Cu2Cl 2 , which melts and runs from the end of the piece. Dutch metal, or false gold leaf, burns with a flash. It contains copper and zinc and gives cuprous chloride, Cu2Cl2, and zinc chloride, ZnCl2- Powdered antimony burns with brilliant flashes and gives antimony pentachloride, SbCls. Phosphorus takes fire in chlorine, burning to the liquid phosphorus trichloride, PCla, if the phosphorus is in excess, or to the solid phosphorus penta- chloride, PCls, if the chlorine is in excess. If chlorine gas is carefully- dried with phosphorus pentoxide, it will not act on copper, iron or other metals. This may be shown by passing the dry chlorine into a flask containing dry Dutch metal. The metal will remain perfectly bright, but the introduction of the slightest trace of moisture will cause the im- mediate combination of the chlorine with the leaf. The reason for this catalytic effect of the water is not understood. Because 106 A TEXTBOOK OF CHEMISTRY of this property, dry, liquid chlorine is kept and sold in strong steel cylinders. If a piece of tissue paper, which has been dipped in warm tur- pentine, CioHie, is thrust into a jar of chlorine, the turpentine will usually take fire and burn with a very smoky, red flame, giving hydrochloric acid and carbon : Ci H 16 + 8 C1 2 = 10 C + 16 HC1 Chlorine and Water. Bleaching. A solution of chlorine in water has the same greenish yellow color which is characteristic of the gas. Such a solution apparently contains most of the chlorine as such, but it has been shown (Jakowin, Z. physik. Chem. 29, 613) that a small amount of the chlorine reacts with the water : Cl Cl + H O H ^ HC1 + Cl O H Hypochlorous Acid The reaction is reversible, with the equilibrium far toward the side giving chlorine and water, and very little hypochlorous acid is present in the solution. If the solution is exposed to light, however, the hypochlorous acid decomposes in two ways, giving either chloric and hydrochloric acids or oxygen and hy- drochloric acid : 2 HC1O + HC1O = 2 HC1 + HC1O 3 Chloric Acid 2 HC1O = 2 HC1 + O 2 It is very noticeable that the compounds of chlorine with oxy- gen become more stable as the amount of oxygen in them in- creases. As the two reactions progress, the color of the chlorine gradually disappears. If dry litmus paper or a dry piece of colored calico is placed in dry chlorine, the color is affected very slowly, if at all, but if the paper or cloth is moistened, the color will be bleached very quickly. The effect of the water may be in part similar to the action of moisture in causing chlorine to combine with metals, PHASES 107 but the chlorine also reacts with the water giving hypochlorous acid, HC1O, and this oxidizes and destroys the coloring matter. * Chlorine Hydrate. Phases. If chlorine is passed into water which is cooled to 0, a crystalline compound, chlorine hydrate, C1 2 .8 H 2 O, separates. If this is warmed under atmospheric pressure, it decomposes at 9.6 ; but if the pressure of the chlo- rine is increased, it may exist at higher temperatures, or if the pressure is lessened, it will decompose at a lower temperature. We have, in this case, a system of three phases, solid, liquid and gas, which can exist over a range of several degrees of tempera- ture. This is true of any other system containing two com- ponents. Such a system with two phases may still have two degrees of freedom (see p. 77) freedom to change in tempera- ture and freedom to change in pressure. The addition of a second component, chlorine, increases the number of degrees of freedom for a given number of phases. If the system is cooled to 0.24, ice will separate from the solution of chlorine, as well as chlorine hydrate. When this occurs, there will be four phases present, liquid, ice, chlorine hydrate and vapor or gas. The pres- sure will also be fixed at 244 mm. No change in either tempera- ture or pressure can occur without the disappearance of one of the phases. With two components and four phases, there is no freedom. A further study of cases in which there are two or more components leads to the conclusion that the number of phases and number of degrees of freedom together are equal to the number of components increased by two, or : p-f p = C + 2 P = number of phases F = degrees of freedom C = number of components This is the celebrated "Phase rule," which was discovered by Willard Gibbs of Yale University. It applies equally well to the formation and decomposition of compounds, as of chlorine hydrate above, and to changes of state, as from ice to water and vapor. It is applicable only when the changes in state are 108 A TEXTBOOK OF CHEMISTRY reversible and is important only when the equilibrium between the different phases is reached within a measurable time. Faraday first prepared liquid chlorine by warming chlorine hydrate in a bent tube of the form shown in Fig. 38. By immersing the closed, empty end, B, in a freezing mixture while the chlorine hydrate in the end, A, was Fig. 38 warmed, the gas liberated by the de- composition of the hydrate exerted enough pressure to cause a part to liquefy in the cooled end. The Heat of Combination of Chlorine and of Oxygen with Other Elements. The following are the heats of combination of several elements with equivalent amounts of oxygen and chlorine : H 2 + O = H 2 O ( vapor) + 58,000 calories H 2 + C1 2 = 2 HC1 + 44,000 calories 2 Na + O = Na 2 O + 100,000 calories 2 Na + C1 2 = 2 NaCl + 195,000 calories Zn + O - ZnO + 85,300 calories Zn + C1 2 = ZnCl 2 + 97,200 calories Cu + O = CuO + 37,200 calories Cu + C1 2 = CuCl 2 + 51,600 calories P 2 -f 5 O = P 2 O 5 + 370,000 calories . P 2 + 5 C1 2 = 2 PC1 5 + 210,000 calories The heat of combination with chlorine is sometimes greater, sometimes less, than the heat of combination with oxygen. In general, the heat of combination of chlorine seems to be greater than that of oxygen in combining with metals and less than that of oxygen in combining with nonmetals. Equilibrium in Chemical Reactions. If a mixture of four volumes of hydrochloric acid with one volume of oxygen is passed slowly through a tube containing cuprous chloride at 345 , EQUILIBRIUM 109 four fifths of the hydrochloric acid will be oxidized, giving chlo- rine and water in accordance with the reversible reaction : 4 HC1 + 2 = 2 C1 2 + 2 H 2 If a mixture of equal volumes of chlorine and steam is passed slowly through the tube at the same temperature, one fifth of the chlorine will be converted into hydrochloric acid. In other words a mixture containing 4 volumes or 4 molecules of HC1 1 volume or 1 molecule of O 2 8 volumes or 8 molecules of C1 2 8 volumes or 8 molecules of H 2 O will be in equilibrium and will not change its composition when heated for a long time at 345. We do not suppose that chemical action ceases, but rather that, in a given time, just as many atoms 20 Vols. HC1 1 5 Vols. O 2 I [ 4 Vols. HC1 1 Vol. O 2 8 Vols. Cl, ( 8 Vols. H 2 O Fig. 39 of chlorine unite with hydrogen to form hydrochloric acid as there are atoms of chlorine separated from molecules of hydro- chloric acid. In this way the total number of molecules of each 10 Vols. C1 2 10 Vols. H 2 O [ 4 Vols. HC1 J 1 Vol. O 2 8 Vols. C1 2 8 Vols. H 2 Fig. 40 110 A TEXTBOOK OF CHEMISTRY of the four substances will remain unchanged after equilibrium is reached, but any given atom may frequently change its state of combination. In s"uch a case we may think of two opposing forces, one of which drives the reaction to the right and the other drives it to the left, and that these forces are in equilibrium. The force driving the reaction toward the right, when we start with hydrochloric acid and oxygen, must be much stronger than the force driving the reaction toward the left, when we start with chlorine and steam. This is probably due, in part, to the greater affinity of oxygen for hydrogen, as indicated by the heat of com- bination given in the last paragraph, but it is also connected with the change in volume which occurs in the reaction and with other factors which are less clearly understood. The reaction proceeds toward the right with the evolution of heat : H 2 + O = H 2 O -f 58,000 calories H 2 + C1 2 = 2 HC1 + 44,000 calories Hence 2 HC1 + O = H 2 O + C1 2 + 14,000 calories since the sum of the reactions : H 2 + Cl a = 2 HC1 and 2 HC1 + O = H 2 O + C1 2 must give the same amount of heat as the reaction H 2 + O = H 2 O, because the chlorine is in the same condition at the end as at the beginning. Whenever a reversible reaction proceeds with evolution of heat, a higher temperature always shifts the equilibrium in the direction to cause a smaller evolution of heat. In other words, the application of heat always helps the side of a reversible reac- tion in which heat is absorbed and retards that side of a reaction in which heat is given out. In accordance with this we find that the mixture in equilibrium at 384 contains : EQUILIBRIUM 111 4 volumes or 4 molecules of HC1 1 volume or 1 molecule of C>2 6 volumes or 6 molecules of C1 2 6 volumes or 6 molecules of H^O This means that while four fifths of the hydrochloric acid can be oxidized to chlorine by the Deacon process at 345, only Ijiree fourths of it can be oxidized at 384. It is this very unfavor- able effect of an increase in the temperature which makes it necessary to use a catalyzer and work at as low a temperature as possible. * Principle of van't Hoff-Le Chatelier. As has been stated above, an increase in temperature displaces any equilibrium in the direction in which heat is absorbed and an increase in pressure displaces any equilibrium in the direction in which the volume decreases. These are special cases of the principle of van't Hoff-Le Chatelier, which is that every force applied to a system which is in equilibrium causes a change which tends to resist the force that is applied. Thus, if pressure is applied to ice, a small amount of the ice will melt, but in melting it will absorb heat, the temperature will fall and this will tend to stop the melting. Or, if dry air is blown over the surface of water, it will cause the water to evaporate; but as it evap- orates, the temperature will fall and the lower vapor pressure will tend to stop the evaporation. In accordance with the law, if hydrochloric acid and oxygen are brought together at 345, heat will be evolved as they react, and this will tend to stop the reaction. On the other hand, if steam and chlorine are brought together at 345, heat will be absorbed as they react, and this will tend to increase the re- action. It is necessary to distinguish between the effect of an in- crease in temperature to increase the speed of a reaction, which seems to be universal, and the tendency to cause a reversal of those reactions in which heat is evolved. The speed of the combination of oxygen and hydrogen increases rapidly with the temperature and becomes explosive at a very mod- 112 A TEXTBOOK OF CHEMISTRY erate heat. The reverse reaction by which water dissociates into oxygen and hydrogen increases, however, with the tem- perature in accordance with the theorem of van't Hoff-Le Chatelier. If we could understand fully the mechanism of all physical and chemical processes, it seems likely that we should find that this* principle has its foundation in Newton's law that for every action there is an equal and opposite reaction. Effect of Water on Chlorides. lonization. When hydro- chloric acid or such chlorides of the metals as sodium chloride, NaCl, zinc chloride, ZnCl 2 , or copper chloride, CuCl 2 , dissolve in water, several different lines of evidence indicate that these compounds separate more or less completely into chloride ions, Cl~~, bearing a negative charge of electricity and hydrogen, H + , or metallic ions, Na + , Zn ++ , or Cu ++ , bearing a positive charge, or, if the atom is bivalent or trivalent, positive charges of electricity. The evidence for this view is, in part, as fol- lows : , 1. It is found that if substances which are not electrolytes are dissolved in water, the freezing point is lowered in direct propor- tion to the amount dissolved and in inverse proportion to the molecular weight of the solute. An aqueous solution of alcohol (C2H 6 O, molecular weight, 46) or of sugar (C^B^On, molec- ular weight, 342) is almost as poor a conductor as pure water. A solution containing 46 milligrams of alcohol dissolved in 10 cc. of water will freeze at 0.184. A solution of sugar containing 342 milligrams in 10 cc. of water will freeze at 0.188. One of these solutions must contain the same number of molecules as the other, and it is evident that the lowering of the freezing point is proportional to the number of molecules of the solute in a given volume of the solvent. But if we dissolve 58.5 milligrams of salt (NaCl, molecular weight, 58.5) in 10 cc. of water, the solution will freeze at 0.349. According to the law just stated, the freezing point indicates nearly twice as many molecules as there should be. The simplest explanation of this fact is that the sodium chloride separates into sodium (Na + ) and chloride (Cl~) IONIZATION 113 ions in the solution and that these, so far as this law is con- cerned, act as independent molecules. In a similar manner, a solution containing 36.5 milligrams of hydrochloric acid (HC1, molecular weight, 36.5) in 10 cc. of water freezes at 0.355, indicating that it is largely separated into hydrogen (H~) and chloride (Cl~) ions. 2. If charged strips of metal, as, for instance, pieces of platinum connected with the poles of an electric battery, are dipped in a solution of hydrochloric acid, the chlorine atoms are attracted by the positive electrode (anode) and the hydrogen atoms are attracted by the negative electrode (cathode) and there is a motion of the ions throughout the solution. This is called the migration of the ions. The rate of migration varies for different ions and with the same current it is quite different, the hydrogen ions moving nearly five times as fast as the chloride ions. This effect can be shown by passing an electrical current from a silver anode to a platinum cathode through hydro- chloric acid in the U-tube (Fig. 41). A silver anode is used because it will combine quantitatively with the chlorine which is liberated, and remove it from the solution. The solution on this side contains : Cathode At first, j^\- 73 mg. HC1 = 2 milligram atoms of H 2 milligram atoms of Cl At end, 66.5 mg. HC1 = 1.82 milligram atoms of H 1.82 milligram atoms of Cl 1 milligram atom of H is liberated Amount of hydrogen transferred across the line CD = 0.82 mg. atoms. Anode 4 V i g y y \ / ^^4-^ D Fig. 41 The solution on this side t contains : At first, 73 mg. HC1 = 2 milligram atoms of H 2 milligram atoms of Cl At end, 43 mg. HC1 = 1.18 milligram atoms of H 1.18 milligram atoms of Cl 1 milligram atom of Cl combines with anode Amount of chlorine transferred across the line CD = 0.18 mg. atoms. 114 A TEXTBOOK OF CHEMISTRY At the beginning of the experiment the concentration of the hydrochloric acid is the same in both arms of the tube ; but after decomposing a part of the hydrochloric acid by passing the cur- rent for some time, it will be found that while the concentration of the acid has decreased on both sides, the amount of acid on the cathode side has become much greater than that on the anode side. Since the number of hydrogen atoms liberated at the cathode must be exactly the same as the number of chlorine atoms which combine with the silver anode, the greater amount of acid on the cathode side must be due to the fact that the hydrogen ions migrate faster toward the cathode than the chloride ions migrate toward the anode. This will be clear from an examination of the figure and the accompanying statement f Fig. 42 about the composition of the solution at the beginning and end of the experiment. 3. If a solution of sodium iodide is subjected to a powerful centrifugal force, the heavier iodide ions may be separated to a slight extent from the sodium ions (Tolman, J. Am. Chem. Soc. 33, 121). Similar experiments were tried with hydriodic IONIZATION 115 acid, lithium iodide and potassium iodide. By whirling solu- tions of these substances in the apparatus shown diagrammati- cally in Fig. 42 the heavy iodide ions were thrown toward the outside, giving a negative electrical charge to the solution at that end of the tube and a positive charge at the inner end. All substances which are electrolytes are supposed to separate more or less into ions in aqueous solutions. Thus sodium nitrate, NaNOs, separates into sodium ions (Na + ) and nitrate ions (NOs~) ; sulfuric acid, H 2 SO4, may separate into two hydrogen ions (H + , H + ), and the sulfate ion (SO 4 ), or it may separate, in part, only into a single hydrogen ion (H + ) and the acid sul- fate ion (HSO4 + ). When solutions containing electrolytes are mixed, the reactions which occur are, in most cases, a simple exchange of ions, and the groups of atoms which form the ions remain unbroken. Thus in the reaction : AgNO 3 + HC1 = AgCl + HNO 3 the nitrate ion (NOa~) passes from one compound to the other without any change. Such reactions are always reversible, and the equilibrium is frequently displaced to one side or the other because one of the compounds is volatile or difficultly soluble. Thus it has been pointed out that in the reaction : NaCl + H 2 S0 4 ^ HC1 + NaHSO 4 if concentrated sulfuric acid is added to salt the equilibrium will be far to the right because the hydrochloric acid is a gas and escapes ; while if a concentrated solution of 'hydrochloric acid is added to a concentrated solution of acid sodium sulfate, the equilibrium may be carried to the left, because sodium chloride precipitates. Effect of Water on Chlorides. Hydrolysis. When water is brought in contact with a chloride of a nonmetallic element, the effect is very different. The chloride reacts with the water as though the water were composed of two parts, hydrogen, H, and hydroxyl, OH. 116 A TEXTBOOK OF CHEMISTRY C\ HOH ,OR OH C1 + HOH = 3 HC1 + P^OH or O=P^OH C1 HOH X OH X H Phosphorous Acid. 01 ,OH /OH O + 5 HOH = 5 HC1 + P^-OH = P^OH + H 2 O. yci Y OH \ OH \C1 \OH \OH Phosphoric Acid This sort of double decomposition with water is called hy- drolysis. The hydroxyl compounds which are formed are acids. In most reactions between these acids and other compounds, the oxygen is held by the nonmetallic element, while the hy- drogen may be easily replaced by metals. Thus with sodium hydroxide we have the reaction : H 3 PO 4 + 3 NaOH = Na 3 PO 4 + 3 HOH The division between the metallic and nonmetallic elements in the conduct of the chlorides is not a sharp one. While sodium or potassium chlorides are only ionized in solution and on evaporation of the water the ions recombine without any loss of hydrochloric acid, and phosphorus pentachloride is completely decomposed by water and on evaporation of the solution the hy- drochloric acid will escape entirely, leaving phosphoric acid, there are many other chlorides, like ferric chloride and aluminium chloride, which partly ionize and partly hydrolyze in solution. EXERCISES 1. If a mixture of salt and manganese dioxide is treated with sulfuric acid, the products will be manganous sulfate, sodium sulfate, water and chlorine. Write the equation. 2. How much hydrochloric acid will be required to give 10 liters of chlorine by the Deacon process, assuming that 80 per cent is oxidized to chlorine ? How much hydrochloric acid will be required by the second stage of the Weldon process ? CHLORINE 117 3. What per cent of the chlorine in hydrochloric acid is liberated when the acid acts on manganese dioxide ? What per cent when it acts on potassium permanganate? What .per cent when it acts on calcium manganite ? 4. What is the weight of chlorine absorbed by one liter of water ? 5. Assuming that air contains 21 per cent of oxygen (by volume), how many volumes of hydrochloric acid should be mixed with 100 vol- umes of air for the Deacon process ? What per cent of free chlorine will the resulting mixture contain, after the reaction, if there is an oxidiza- tion of 80 per cent ? CHAPTER VIII HYDROCHLORIC ACID. OXIDES AND OXYACIDS OF CHLORINE Hydrochloric Acid. The explosive combustion of hydrogen and chlorine under the influence of light has been mentioned. Hydrogen may be burned in a jar of chlorine or chlorine may be burned in hydrogen, Figs. 43 and 44. In each case hydrochloric acid is formed. It may be pre- pared more easily by pouring a mixture of 9 parts (by weight) of concentrated sulfuric acid with 2 parts of water on common salt, Fig. 43 NaCL Concentrated acid might be used, but the mixture with salt froths badly, while the slightly diluted acid does not froth : NaCl + H 2 SO 4 : NaHSO 4 + HC1 Sodium Acid Sodium Chloride Sulfate The compound NaHSO 4 is called acid sodium sulfate because it still contains an acid hydrogen atom which can be replaced by a metal. Thus if more salt is added and the mixture is warmed, the reaction : NaHSO 4 + NaCl ^ Na 2 SO 4 + HC1 will occur. These reactions are reversible and would be very far from com- plete in either direction if all of the substances remained mixed together. But as soon as the salt and sulfuric acid are mixed, hydrochloric acid begins to escape as a gas. When this occurs, 118 HYDROCHLORIC ACID 119 the acid which has gone can no longer have any effect in driving the reaction in the oppo'site direction, and a new quantity of the sulfuric acid will act on the salt. In this way, if the mixture is warmed, the reaction may finally be made practically complete. The equilibrium is displaced in the direction toward the forma- tion of the product which is continually removed from the mixture. That the reaction is reversible may be easily shown by adding a concentrated solution of hydrochloric acid to a strong solu- tion of acid sodium sulfate. A copious precipitate of sodium chloride will be formed : HC1 + NaH3O 4 ^ NaCl + H 2 SO 4 In this case, as the sodium chloride precipitates it can no longer act on the substances remaining in solution, and the equilibrium is displaced toward the formation of the compound which is precipitated. Not many years ago it was quite common to say that sulfuric acid is stronger than hydrochloric acid and so expels hydrochloric acid from its salts. The experiments described show that either acid may expel the other, and that the direction in which the reaction goes depends on the volatility or insolu- bility of the compounds formed and on the relative amounts of the reacting substances as well as upon the relative affinities of the chlorine and of the sulfate radical for the metal, and that the first three factors are frequently more important than the last. Properties of Hydrochloric Acid. Hydrochloric acid is a colorless gas which may be condensed to a liquid by cold and pressure. The liquid boils at - 83.7 and freezes at - 110. What is the weight of 22.4 liters of the gas ? What is the density as compared with air ? Water at will absorb 503 volumes of the gas. If the solu- tion is boiled, more hydrochloric acid than water escapes at first, and the temperature gradually rises till a boiling point of 110 is reached. After that, the portion which distils over and that which remains behind will have the same composition, contain- 120 A TEXTBOOK OF CHEMISTRY ing 20.2 per cent of the acid. If an acid which contains less than 20.2 per cent is boiled, the boiling point will be below 110 and the portion distilling over will contain less acid than that which remains. If the distillation is continued, the temperature will gradually rise to 110, and after that a mixture of constant composition (20.2 per cent) will distill as before. The ratio between the volumes of hydrochloric acid and of the hydrogen which it contains may be demonstrated roughly by filling a dry tube with the gas, pouring in a few cubic centi- meters of liquid sodium amalgam, inserting a rubber stopper quickly and shaking vigorously. On opening the tube with the mouth below the surface of water in a beaker the water will rise and fill the tube one half full. Does the experiment demonstrate that hy- drochloric acid is composed of equal volumes of hydrogen and chlorine? What would be the result if a similar experiment could be tried with steam and all of the hydrogen of the steam were replaced by the metal ? The composition of hydrochloric acid by volume may be demonstrated by the electrolysis of a strong solution of the acid, using carbon electrodes. The volumes of hydrogen and of chlorine liberated at the two electrodes will be nearly equal, if a moderately strong cur- rent is used and the current is continued till the solution around the anode is saturated with chlorine. A suitable ap- paratus is shown in Fig. 45 . See Brown- lee, J. Am. Chem. Soc. 29, 237. For most laboratory purposes to which hydrochloric acid is applied the solution in water is used. The most important chemical properties are : 1. Reaction with Metals. With many metals the hydrogen Fig. 45 HYDROCHLORIC ACID 121 is displaced by the metal and chlorides are formed, which dis- solve in the water. Thus, sodium, magnesium, zinc, iron, aluminium and tin give sodium chloride, NaCl, magnesium chlo- ride, MgCl 2 , zinc chloride, ZnCl 2 , ferrous chloride, FeCl 2 , alumin- ium chloride, A1C1 3 , and stannous chloride, SnCl 2 . What are the reactions for the formation of these chlorides ? It is worthy of notice that metals like iron and tin, which form two chlorides, give the lower chloride when the metals are dissolved in a solu- tion of hydrochloric acid. 2. Reaction with Hydroxides of Metals. Hydrochloric acid reacts with hydroxides of the metals, forming chlorides and water : HC1 + NaOH = NaCl + HOH 2 HC1 + Fe(OH) 2 = FeCl 2 + 2 HOH Ferrous Ferrous Hydroxide Chloride 3 HC1 + Fe(OH) 3 = FeCl 3 + 3 HOH Ferric Ferric Hydroxide Chloride 4 HC1 + Sn(OH) 4 = SnCl 4 + 4 HOH Stannic Stannic Hydroxide Chloride In these reactions the separation of the metallic hydroxide is between the metal and hydroxyl, just as the separation of the acid is between the hydrogen and chlorine. Compounds which react in this manner are called bases, the presence of a hydroxyl group, OH, which separates easily, being characteristic of a base, as the presence of hydrogen which separates easily is character- istic of an acid. Since hydrogen and hydro*xyl have a strong affinity for each other and separate only to a trifling extent in solutions or in pure water, bases and acids neutralize each other by the union of the hydrogen of the acid with the hydroxyl of the base. The compound formed by the union of the metal with the chlorine or with the acid radical is called a salt. In each case, for the formation of a normal salt there must be as many hy- droxyl groups in the base as there are hydrogen atoms in the acick 122 A TEXTBOOK OF CHEMISTRY 3. Reaction with Oxides of Metals. Some oxides of metals also react with hydrochloric acid to form salts and water : ZnO + 2 HC1 = ZnCl 2 + H 2 O 4. Reaction with Oxidizing Agents. With oxidizing agents hydrochloric acid is oxidized to water and chlorine, the oxidizing agent being at the same time reduced. In such reactions chlorides which contain two atoms of chlorine are in the same degree of oxidation as those which contain one atom of oxygen, since two atoms of chlorine replace one atom of oxygen in com- bination with hydrogen. Thus manganous oxide, MnO, is in the same state of oxidation as manganous chloride, MnCl2, and ferric chloride, FeCl 3 , corresponds in oxidation to either ferric oxide, Fe 2 O 3 , or ferric hydroxide, Fe(OH) 3 , while man- ganese dioxide, MnO 2 , is in a higher state of oxidation than man- ganese chloride, MnCl 2 . The reactions between hydrochloric acid and manganese diox- ide, MnO 2 , potassium permanganate, KMnC>4, and calcium man- ganite, CaMnO 3 , have been given. Similar reactions take place with lead dioxide, PbO 2 , and red lead, Pb 3 O 4 , which are re- duced to lead chloride, PbCl 2 ; also with potassium dichromate, K 2 Cr 2 O7, the chromium being reduced to chromic chloride, CrCl 3 , while the potassium, which is univalent, does not change its state of oxidation. These equations should be written by the student as an aid to an understanding of reactions of this type and also to give practice in writing equations correctly by devel- oping them from a knowledge of the compounds formed instead of as a matter of memory. Why does not barium peroxide, BaO 2 , give chlorine when treated with hydrochloric acid ? Indicators. A number of organic compounds are known which have one color in an acid solution, that is in a solution containing hydrogen ions, H + , and another color in an alkaline solution, that is, in a solution containing hydroxide ions, OH~. More strictly speaking, such compounds are, in reality, each of them, two different compounds so related that hydrogen ions OXIDES AND OXYACIDS OF CHLORINE 123 will change the first into the second, and hydroxide ions will change the second into the first. Thus litmus is a red compound in an acid solution, and a blue compound in an alkaline solution. Some of the common indicators are : NAME COLOR IN ACID SOLUTIONS COLOR IN ALKALINE SOLUTIONS Litmus Red Blue Phenolphthalei n Colorless Red Methyl orange Rose red Yellow Methyl red Red Yellow Congo red Red Blue Oxides and Oxygen Acids of Chlorine. Nomenclature Chlorine forms three oxides and four acids containing oxygen : , ~ (Chlorine monoxide or] TT^,^ i -j C1 2 O j . \ HC1O hypochlorous acid, [ hypochlorous anhydride j r^in /^ui j- -j f HC1O 2 chlorous acid. C1O 2 - Chlorine dioxide - .__ . . , [ HClUs chloric acid C^OT Perchloric anhydride HC1O4 perchloric acid The names of these acids should be learned carefully as an illustration of the principles used in naming acids. For chlorous and chloric acids the endings correspond to those which are used for oxides and chlorides (p. 29). The prefix hypo- means under, and the prefix per- means above or beyond. The endings and prefixes refer to the relative amounts of oxygen for different acids of the same element. The relations for other elements are not always so simple. Thus the acids of sulphur are : Sulfurous acid, H 2 SOs Sulfuric acid, H 2 SO 4 Persulfuric acid, H 2 S 2 O 8 [HSO 4 ] 2 The persulfuric acid is in a higher state of oxidation than sulfuric acid because it contains less hydrogen, not because it contains more oxygen in proportion to the sulfur. 124 A TEXTBOOK OF CHEMISTRY The salts of the acids are named by changing the -OILS of the acid to -ite for the salt, and the -ic of the acid to -ate for the salt. Hypochlorous acid, HC1O, gives potassium hypochlorite, KC1O Chlorous acid, HC1O 2 , gives potassium chlorite, KC1O 2 Chloric acid, HC1O 3 , gives potassium chlorate, KC1O 3 Perchloric acid, HC1O 4 , gives potassium perchlorate, KC1O 4 Hypochlorous Acid. Hypochlorites. When chlorine is dis- solved in water, it has been pointed out that a small amount of hypochlorous acid is formed by the reversible reaction : i C1 2 + HOH ^ HC1 + HC1O The equilibrium in this reaction is very far toward the left, but if a base is added to the solution, the two acids will be neutral- HC1 + KOH = KC1 + H 2 HC1O + KOH = KC1O + H 2 O The neutralization of the acids causes a displacement of the equilibrium toward the right side of the first equation and the reaction goes on to completion. By adding the three equations together and eliminating water, the result can be expressed in the single equation : 2 KOH + C1 2 = KC1 + KOC1 + H 2 O If slaked lime (calcium hydroxide, Ca(OH) 2 ) is used, a mixture of calcium chloride, CaCl 2 , and calcium hypochlorite, Ca(OCl) 2 , ,C\ or a calcium chloride-hypochlorite, Ca<^ , is formed. This X)C1 is called bleaching powder. /Cl Ca(OH) 2 + C1 2 = Ca< + H 2 O OCl Calcium Chloride-hypochlorite HYPOCHLORITES 125 The hypochlorites give up their oxygen readily to other sub- stances and so are powerful oxidizing agents. The action is much more vigorous in a faintly acid than in an akaline solution, however, because hypochlorous acid, HC1O, gives up its oxygen much more easily than a hypochlorite does. For this reason bleaching powder is applied to the bleaching of cotton or linen cloth by dipping the cloth first in a solution of the bleaching powder and then in very dilute acid, which liberates the hypo- chlorous acid. Not only may a hypochlorite be used to oxidize other sub- stances, but if a neutral or faintly acid solution of a hypochlorite is boiled, one portion is oxidized to a chlorate while another por- tion is reduced to a chloride : KC1O + 2 KC1O = KC1O 3 + 2 KC1 Potassium Chlorate If a small amount of a cobalt salt, as cobalt nitrate, Co(NOs)2, is added to a solution of a hypochlorite, the oxygen of the hypo- chlorite is liberated in the free state. The cobalt is oxidized to cobalt dioxide, CoO2, which then acts as a catalyzer, as man- ganese dioxide acts on potassium chlorate : ,C\ 2 Ca< + CoO 2 = 2 CaCl 2 + O 2 + CoO 2 X OC1 Hypochlorous acid is a very weak acid. While in the reaction of ionization : HCI ^t H + + cr . which occurs when hydrochloric acid is dissolved in water, the equilibrium is far to the right in moderately dilute solutions, for hypochlorous acid the corresponding reaction : has the equilibrium very far to the left. In other words hydro- chloric acid gives a large proportion of hydrogen ions in dilute solutions, while hypochlorous acid gives only a very small pro- 126 A TEXTBOOK OF CHEMISTRY portion of such ions. This fact may be used to obtain a solu- tion of hypochlorous acid. If hydrochloric acid is added to a solution of a hypochlorite, hypochlorous acid will be formed in accordance with the reaction : K + + cio- + H + + cr = HCIO + K + + cr From such a solution hypochlorous acid and water pass over together, on distillation, and this is the easiest method of getting a solution of hypochlorous acid. An excess of hydrochloric acid must be avoided, however, as this would cause the reaction : H + + Cl- + HCIO ^ H 2 O + C1 2 to occur, in which the equilibrium is far to the right. Hypochlorous acid can be obtained only in dilute solutions. Concentrated solutions decompose in accordance with the re- actions already given : 2 HCIO = 2 HC1 + O 2 HCIO + 2 HCIO = HC10 3 + 2 HC1 HC1 + HCIO = H 2 + C1 2 * Hypochlorous Anhydride or Chlorine Monoxide. Chlorine monoxide, C1 2 O, is formed when chlorine is passed through a tube containing cold, dry mercuric oxide, the mercury being con- verted into an oxychloride : 2 HgO + 2 C1 2 = HgO.HgCl, + C1 2 O The oxide of mercury used must be obtained by precipitation, and washed and dried at 300 400, as the crystalline oxide does not react readily enough. Chlorine monoxide may be con- densed to a liquid which boils at about 5. Either the liquid or the gas explodes violently on slight provocation, giving chlo- rine and oxygen : 2 C1 2 O = 2 C1 2 + O 2 In this case the affinity between atoms of the same kind is greater, apparently, than that between chlorine and oxygen in chlorine monoxide. Curiously enough, when chlorine combines CHLORATES 127 with a larger amount of oxygen, the compound is much more stable. * Chlorous Acid and Chlorites. When sodium peroxide is added to a solution of chlorine peroxide, C1O 2 , sodium chlorite is formed : 2 C1O 2 + Na 2 O 2 = 2 NaClO 2 + O 2 The chlorites are bleaching agents, similar to the hypochlo- rites, and are even more unstable. Free chlorous acid has not been prepared, even in solution. Chloric Acid and Chlorates. When a faintly acid solution of a hypochlorite is warmed, one portion oxidizes another to a chlorate : KC1O + 2 KC10 = KC10 3 + 2 KC1 Practically, if chlorine is passed into a solution of potassium hydroxide or into milk of lime, Ca(OH) 2 , till there is a slight excess and the solution becomes warm from the heat evolved by the reaction, a solution of potassium chloride and potassium chlorate, or of calcium chloride and calcium chlorate will be ob- tained. The student should write the equations and notice what portion of the potassium is converted into potassium chlo- rate. Why is it more economical to prepare calcium chlorate first and then obtain potassium chlorate by adding potassium chloride to the solution ? What must be the relative solubilities of calcium chlorate and potassium chlorate for such a method to be successful ? Chloric acid is stable only in solution and cannot be separated as a pure compound. * Chlorine Dioxide. When concentrated sulfuric acid is added to a chlorate the chloric acid liberated decomposes at once into perchloric acid, chlorine peroxide and water : KC1O 3 + H 2 S0 4 = HC10 3 + KHS0 4 3 HC10 3 = HC10 4 + 2 C10 2 + H 2 O Chlorine dioxide is a heavy, yellow gas, which is easily soluble in water. It seems to be even more unstable than chlo- 128 A TEXTBOOK OF CHEMISTRY rine monoxide and explodes violently if warmed or brought into contact with organic matter. These properties may be illus- trated by mixing some sugar and potassium chlorate and adding a drop of concentrated sulfuric acid. The chlorine dioxide will react with the sugar and ignite the mixture. The solution of chlorine peroxide in water gives with a base a mixture of chlorite and chlorate : 2 C1O 2 + 2 KOH = KC1O 2 + KC10 3 + H 2 O Chlorine dioxide may, therefore, be considered as an anhy- dride of both chloric and chlorous acids. Perchlorates and Perchloric Acid. When potassium chlorate is heated to its melting point, it partly decomposes into potassium chloride and oxygen, but a part is oxidized to potassium per- chlorate, while another part is reduced to potassium chloride : 3 KC1O 3 + KC1O 3 = 3 KC1O 4 + KC1 Potassium Perchlorate This illustrates, again, the fact that the compounds of chlorine with oxygen become more and more stable as more oxygen is taken up. This is true of the acids as well as of the salts. Per- chloric acid is the only oxyacid of chlorine which can be ob- tained as a pure compound, free from water. Perchloric an- hydride, C^OT, is also the most stable of the oxides of chlorine. It has been pointed out that concentrated sulfuric acid expels hydrochloric acid from salt chiefly because hydrochloric acid is a gas and escapes from the mixture ; also that concentrated hydrochloric acid will precipitate salt from a solution of acid sodium sulfate, NaHSO 4 , leaving sulfuric acid in solution, chiefly because the salt is nearly insoluble in concentrated hydrochloric acid. Both principles may be used to prepare perchloric acid. If concentrated hydrochloric acid is poured over some sodium perchlorate (30 cc. for 20 grams of perchlorate), the reversible reaction : NaC10 4 + HC1 = NaCl + HC1O 4 PERCHLORIC ACID 129 will proceed till 95 per cent of the sodium separates as sodium chloride. This may be removed by filtering on an asbestos filter and washing the salt with a small amount of concentrated hy- drochloric acid. The filtrate 1 contains a little salt with hy- drochloric and perchloric acids. The highest boiling point of an aqueous solution of hydrochloric acid is 110, while the boil- ing point of the hydrated perchloric acid is 203. On heating the mixture, therefore, the hydrochloric acid escapes and finally the reversible reaction : NaCl + HC10 4 = NaClO 4 + HC1 is carried to completion, leaving only perchloric acid containing a small amount of sodium perchlorate. A pure hydrated per- chloric acid, containing about 28 per cent of water, may be ob- tained from the mixture by distilling under diminished pressure. This hydrated acid boils with some decomposition at 203, under atmospheric pressure. Anhydrous perchloric acid may be ob- tained by distilling a mixture of potassium perchlorate and con- centrated sulfuric acid under diminished pressure. The anhy- drous acid is far less stable than the hydrated acid. This is possibly because the hydrated acid has the structure : /0-H : %o ^O * Perchloric Anhydride, C1 2 O 7 , may be prepared by adding perchloric acid to phosphoric anhydride cooled to 10 and dis- tilling after some time : 2 HC1O 4 + P 2 O 5 = C1 2 7 + 2 HPO 3 Metaphosphoric Acid Perchloric anhydride is a colorless, oily liquid which boils at 82. 1 The portion of a solution which has passed through a filter. 130 A TEXTBOOK OF CHEMISTRY Structure of the Oxyacids of Chlorine. If chlorine is univa- lent in the oxyacids of chlorine, the structure of these acids would be represented by the formulas : H O Cl Hypochlorous acid H O O Cl Chlorous acid H O O O Cl Chloric acid H O O O O Cl Perchloric acid There is very little evidence that oxygen atoms can unite in this manner to form chains that are stable, and the instability of hydrogen peroxide, H O O H, makes these formulas seem very improbable. There is also a good deal of evidence to show that chlorine and other related elements may have a valence as high as seven in some of their compounds. The fol- lowing structures are, therefore, considered much more probable : ^O H O Cl H O C1=O H O Cl^O H O ClO These compounds illustrate very clearly the effect of oxygen in giving an acid character to hydrogen compounds. Hypo- chlorous acid is a very weak acid and separates almost as easily into chlorine and hydroxyl, H O , as it does into hydrogen and O Cl. Perchloric acid on the contrary, is a strong, stable acid, especially in the hydrated form. For a possible explanation see p. 206. * The Atomic Weight of Chlorine. As oxygen is the basis for atomic weights, it would be most natural to determine the atomic weight of chlorine by determining the composition of one of the oxides of chlorine. But these oxides are so unstable that they cannot be prepared in a condition of sufficient purity for such a purpose. The composition of hydrochloric acid has been de- termined accurately, however, in two ways ; first, by combining a weighed amount of hydrogen with chlorine, which was weighed in the liquid form; second, by passing a weighed amount of ATOMIC WEIGHT OF CHLORINE 131 hydrogen over potassium chloroplatinate, K 2 PtCl 6 , which was reduced to potassium chloride and metallic platinum. The loss in weight gave the weight of the chlorine, while the hydro- chloric acid formed was also collected and weighed. The aver- age for the ratio between hydrogen and chlorine by the two methods is, H : Cl = 1 : 35.189. As the atomic weight of hy- drogen is 1.0078 (p. 72), the atomic weight of chlorine is 35.189 X 1.0078 = 35.463 The atomic weight has also been very accurately determined by an entirely different method. By dissolving lithium chloride, LiCl, in a solution of perchloric acid and evaporating the water, it was converted into lithium perchlorate, LiClO 4 . The ratio of the increase in weight to the weight of the lithium chloride was : 4 : LiCl = 1.50968 : 1 = 64 : 42.393 This gives the molecular weight of LiCl as 42.393. Next, the amount of silver required to combine with the chlo- rine of the lithium chloride was determined by dissolving a weighed amount of silver in nitric acid and adding the solution of silver nitrate, AgNOs, to the solution of lithium chloride. This gave the ratio : LiCl : Ag = 0.39299 : 1 = 42.393 : 107.871 This gives the atomic weight of silver, Ag = 107.871. The amount of silver chloride which could be obtained from a given weight of lithium chloride was also determined. This gave : LiCl : AgCl = 0.295786 : 1 = 42.293 : 143.325 Subtracting the atomic weight of silver from the molecular weight of silver chloride we have : 143.325 - 107.871 = 35.454 which is the atomic weight of chlorine. It will be seen that this value agrees cjosely with that given by the other methods, and it does not seem likely that the value 35.46, which is given in the atomic weight tables, can be far wrong. CHAPTER IX CLASSIFICATION OF THE ELEMENTS. THE PERIODIC SYSTEM THE three elements which have been studied, oxygen, hydrogen and chlorine, differ very greatly from each other, but as we pass on to other elements, it will be found that several of these have very marked resemblances to chlorine, while others have proper- ties which recall those of oxygen, though the resemblances are not so close. The elements fall into a number of more or less well-defined families or groups, and a knowledge of these groups is of great assistance in acquiring a knowledge of the elements and their compounds. The most satisfactory classification is the one known as the Periodic System, which is based on the atomic weights and is found in the accompanying tables. The elements are arranged in the order of their atomic weights, with a few exceptions, which will be referred to below. Hydrogen does not seem to fall into the classification and is omitted in the first table. The valence of hydrogen would put it in Group I, while its amphoteric (p. 206) character in HOH would relate it to aluminium or silicon. Beginning with helium, the first seven elements after helium, He, 4; Li, 6.94; Be, 9; B, 11 ; C, 12 ; N, 14; O, 16; F, 19, pass from lithium, which is strongly metallic, to fluorine, which is very strongly nonmetallic. The first five form oxygen compounds as follows : L^O, BeO, B 2 O3, CO2, N2Os. The last four form compounds with hydrogen, CH 4 , NH 3 , OH 2 , FH. In the second row we find that the ninth element, neon, resembles helium, sodium resembles lithium, magnesium re- sembles beryllium, and so on to chlorine, which resembles fluorine. Compounds with oxygen and hydrogen are : 132 CLASSIFICATION OF THE ELEMENTS 133 Na 2 O, MgO, A1 2 3 , SiO 2 , P 2 (V, SO 3 , C1 2 O 7 SiH 4 , PH 3 , SH 2 , C1H In the third row the oxides are : K 2 O, CaO, Sc 2 O 3 , TiO 2 , V 2 O 5 , CrO 3 , Mn 2 O 7 , but the last four elements do not, as in the first two rows, form compounds with hydrogen. At the end of the row are three elements, iron, cobalt and nickel, which resemble manganese in some of their properties. Similar groups of three elements are found after the fifth and seventh rows of elements. Beginning with the third row, the elements of alternate rows resemble each other much more closely than those of the successive rows, and each pair of rows taken together is spoken of as a long period .to distinguish these from the short periods of the first two rows. It will be noticed that the highest valence of the elements toward oxygen increases from "left to right, from one to seven : Bf M N^O x,0 C\fO fl "\ "\ J --V ^ x^ -. =O )0, Be=0, V), Cf , >0, S=0, >0. Li X B< ^O N^O % Cl%0 ^O ?0 The valence toward hydrogen, however, decreases from the center to the right from four to one : H I /H /H H C H, Nf-H, O< , F-H. T \H H H If a line is drawn in the table between beryllium and boron downward to the right and between tellurium and tungsten (W), all of the elements below and to the left of the line will be found to be metallic, except those of the argon family, while in the third, fifth and seventh rows the elements to the right of the line are also metallic. The remaining elements are nonmetallic. These 1 The true formula is P 4 Oi . 134 A TEXTBOOK OF CHEMISTRY 6g fa iri EM O V s*=? / K " CO S>2*^ a -O pL, ^ -a S 2: JH^ -H-P "^ GO & -as 02 t^ *** CO -.8 SPjj GO CLASSIFICATION OF THE ELEMENTS 135 S3 o S s t- o o J2 is 1C T-1 a ag w-s ll 136 A TEXTBOOK OF CHEMISTRY facts and also the short and long periods are better shown in the form of the table on p. 135. In this second table the melting points of the elements are also given and it is very clear that these increase to a maximum in the fourth to the sixth groups of the longer periods (horizontal rows) and fall off on either side. This periodicity of the melting points has proved useful in indi- cating those elements which are suitable for the filaments of in- candescent lights. There is also a periodic relation between the atomic weights of the elements and their atomic volumes. The atomic volume may be defined as the volume occupied by one gram-atom of the element. Thus, the specific gravity of potassium is 0.862 and the 39 1 atomic volume is ' = 45.4. The specific gravity of silver is 10.492 and the atomic volume is ^^ = W.2S. The peri- odic relation between atomic weights and atomic volumes is clearly shown in Fig. 46. When the Periodic Table was first proposed by Mendeleef, scandium, germanium and several other elements, which have been discovered since then, were unknown, and he predicted the discovery of these elements and pointed out some of their prop- erties. When these elements were discovered a few years later, the fulfillment of this prophecy helped very much toward the acceptance of the table among chemists. Still later the table caused Professor Ramsay, after the discovery of argon and helium, to search diligently for the other elements of the Zero Group which were predicted by the table. His search proved successful and resulted within a few years in the discovery of neon, krypton, xenon and niton. The group is called the zero group because the elements of the group do not combine with other elements and their valence is considered to be zero. The atomic weight which has been accepted for tellurium is greater than it should be in accordance with the properties of the element, which place it in the sulfur family. This has led a number of investigators to examine the compounds of the ele- CLASSIFICATION OF THE ELEMENTS 137 OIHOIV 138 A TEXTBOOK OF CHEMISTRY ment very carefully and to redetermine the atomic weight by different methods. In three other cases (A and K ; Co and Ni ; Pr and Nd) the relative positions of the elements do not correspond to the prop- erties, and all of these elements have been very carefully studied for this reason. No one has been able to show, however, that the accepted atomic weights for these eight elements are wrong, and we are forced to the conclusion that the same factors which cause the differences between successive atomic weights to be irregular have, in these cases, displaced the elements from what seem to be their normal places. The rare elements given in the footnote on p. 134 are not easily placed in the ordinary forms of the Periodic Table ; Werner and others have proposed arrangements which include these elements on the general principle that as the short periods, He F and Ne Cl are followed by longer periods, A Br and Kr I, these periods are, in turn, followed by periods containing each a still larger number of elements. It is clear from what has been said that the Periodic System is not only useful as a convenient means of classifying the ele- ments and for didactic purposes, but that it has also proved a powerful stimulus to chemical research. The relations between the atomic weights of the elements and their properties which are brought out in the Periodic System constantly suggest that the elements must have some common origin and that the atoms are complex aggregates built up in some way from simpler parts. Such an idea has received very strong support from the phenomena connected with radium and other radioactive elements (p. 471). CHAPTER X THE HALOGEN FAMILY General Properties of the Halogens. The four elements of the halogen family are : x Fluorine, F, 19 Chlorine, Cl, 35.5 Bromine, Br, 80 Iodine, I, 127 The elements of the halogen group are the most strongly non- metallic of all the elements. They are also called negative be- cause in the electrolysis of their compounds they are attracted toward the anode or positive electrode. In contrast with these and other nonmetallic elements, metals are called positive, the most strongly positive or metallic elements being those of the alkali group, to which sodium and potassium belong. The name halogen means " salt-former," and is given to these elements be- cause they combine directly with metals to form salts, sodium chloride or common salt, NaCl, being the most important ex- ample. The most common salts containing other nonmetallic elements are those which also contain oxygen, as sodium sulfate, Na 2 SO4, or potassium nitrate, KNO 3 . Compounds of the Halogens with Hydrogen and Oxygen. The compounds of the halogens with hydroge'n and with hydro- gen and oxygen are acids and have the following formulas : H 2 F 2 orHF HC1 HBr HI HC10 HBrO HIO HC10 2 HC10 3 HBrO 3 HIO 3 HC1O 4 HI0 4 1 In this and other similar tables approximate values are given for the atomic weights in order that the student may learn the rela- 139 140 A TEXTBOOK OF CHEMISTRY The elements of the group are univalent in combining with hydrogen or with positive elements and frequently, also, in com- bining with nonmetallic elements. In combining with oxygen or with oxygen and hydrogen the valence seems to vary from one in hypochlorous acid, H O Cl, to seven in perchloric acid, H O Cl^O, the odd numbers of valences being most common. Fluorine is the most strongly nonmetallic or negative element of the group, or, indeed, of all of the elements. It will displace any other element of the group from combination with hydrogen or a metal. In a similar manner chlorine will displace bromine or iodine, and bromine will displace iodine. This is prob- ably due to the same properties which cause hydrofluoric acid to be the most stable and hydriodic acid to be the least stable of 'the compounds of these elements with hydrogen. The resemblances between chlorine, bromine and iodine are much closer than the resemblances between these elements and fluorine. For this reason bromine and iodine are considered first. In studying these elements the properties of chlorine and its compounds should be constantly recalled and the re- semblances emphasized. Bromine, Br, 79.92. Occurrence, Preparation. In most cases where large quantities of chlorides are found in nature smaller amounts of bromides are found associated with them. In this way bromides are found especially in sea water and in the brines from which salt is obtained by evaporation and crystallization. Some of the American brines in Michigan are rich in bromine, and the bromine is obtained from these by subjecting them to electrolysis till all of the bromine is liberated, with a small amount of chlorine. As bromine boils at 59 and has a molec- ular weight of 160, on boiling the liquid the bromine will pass off with a comparatively small amount of water. (What con- tions among the atomic weights more easily. Accurate values are given on p. 10. BROMINE 141 nection has the last fact with the molecular weight of bromine ?) The chlorine may be removed by mixing the impure bromine with a solution of potassium bromide and distilling : 2 KBr + C1 2 = 2 KC1 + Br 2 Bromine may also be prepared by warming a mixture of potas- sium bromide, manganese dioxide and sulfuric acid. The prod- ucts are potassium sulfate, manganese sulfate, bromine and water. What is the equation for the reaction? Properties. Bromine is a heavy, very dark colored liquid, which gives off reddish brown vapors at ordinary temperatures. It has a strong, disagreeable odor, the name having been given to it for this reason, from /Spw/Aos, a stench. It is also an irritant poison. As with chlorine, the best antidote is to breathe the vapor of strong alcohol. If the liquid touches the skin, it produces a severe wound, which it is very difficult to heal. Although bromine vapor is much heavier than air (how many times heavier ?), if a little of the liquid is placed in the bottom of a tall cylinder, the vapor will diffuse rapidly upward through the air in the cylinder. How can this be explained by the kinetic theory ? Bromine combines directly with both metals and nonmetals, forming compounds which are, in almost all cases, very similar to the corresponding chlorides both in formulas and in properties. At 228 the volume of bromine vapor which would fill 22.4 liters at and 760 mm. 1 weighs about 160 grams, but at 1570 the gram molecular volume weighs only a little over 100 grams. Fig. 47 1 Supposing that the vapor could be cooled to this temperature at a pressure of 760 mm. without its condensing to a liquid. 142 A TEXTBOOK OF CHEMISTRY This indicates that at high temperatures bromine is largely dissociated into molecules which contain only a single atom. At lower temperatures the formula of bromine is evidently Br 2 . Bromine melts at 7 and boils at 59. It has a specific gravity of 3.1883 at or of 2.9483 at 59. It forms a hydrate which probably has the composition Br 2 + 8 H 2 O, though the analyses of the compound do not agree very well with the for- mula. Potassium and sodium bromides are used in medicine as sedatives, the latter by preference, because the bromide ion seems to be the constituent which produces the desired effect, while the potassium ion is much more irritant than the sodium ion when taken in moderate quantities. Silver bromide is used in pho- tography, especially in the preparation of " dry plates." Many compounds of bromine are used in the manufacture of coal- tar dyes. Bromine has also been used to a limited extent as a disinfectant. Hydrobromic Acid. From the method used in preparing hy- drochloric acid we should expect to get hydrobromic acid by the action of sulfuric acid on sodium bromide or potassium bromide : KBr + H 2 SO 4 : KHSO 4 + HBr This reaction takes place when the substances are mixed, but the hydrobromic acid gas which escapes will be colored brown, indicating the presence of free bromine. Sulfur dioxide, SO 2 , is also found in the vapor : H 2 SO 4 + 2 HBr ^ H 2 SO 3 + Br 2 + H 2 O Sulfurous Acid The sulfurous acid is unstable and decomposes into sulfur dioxide and water : H 2 SO 3 ^1 H 2 O + SO 2 This is evidently because, owing to the comparatively weak affinity between bromine and hydrogen, hydrobromic acid acts as a reducing agent toward sulfuric acid. BROMINE 143 To obtain hydrobromic acid free from bromine a mixture of hydrogen and bromine vapor may be passed through a tube con- taining a red-hot spiral of platinum wire. Another method is to drop bromine into a mixture of red phos- phorus and water and pass the hydrobromic acid gas through a tube containing red phosphorus and glass wool moistened with a strong solution of hydrobromic acid. The method depends on the hydrolysis of phosphorus tribromide by water (p. 115). 2P +3Br 2 = 2PBr 3 PBr 3 + 3 HOH = P(OH) 3 + 3 HBr Phosphorous Acid Hydrobromic acid is a colorless gas, which fumes strongly in the air owing to its condensation with the moisture of the air to form a concen- trated solution, which has a much lower vapor pressure than that of water. Water dissolves the acid even more readily than it dissolves hydrochloric acid. The constant boiling mixture of hy- drobromic acid and water boils at 125 and contains 47.7 per cent of hydrobromic acid; while the cor- responding mixture of hydrochloric acid and water boils at 110 and contains only 20.24 per cent of hydrochloric acid. The density of the hydrobromic acid solution is also considerably greater for a given per cent of acid. Sodium Hypobromite, NaBrO, is obtained by dissolving bro- mine in a cold solution of sodium hydroxide, or, better, by draw- ing the vapor of bromine through the solution with a current Fig. 48 144 A TEXTBOOK OF CHEMISTRY of air. (See hypochlorites.) If the solution is warmed, the hypobromite is changed to the bromate, NaBrO 3 : 2 NaBrO + NaBrO = 2 NaBr + NaBrO 3 Iodine, I, 126.92. Occurrence, Preparation. While a minute quantity of iodine is found in sea water and in almost all brines, the amount is too small for the practical preparation of the ele- ment. Many seaweeds, however, absorb a small amount of iodine from the sea water. The ash from these weeds is called kelp and contains a small amount of iodides. From these the iodine may be liberated by chlorine or by sulfuric acid and manganese dioxide. (If sodium iodide is the compound pres- ent, what will be the equation for the required action ?) Iodine is also found as sodium iodate, NaIO 3 , in the crude sodium nitrate from Chile and Peru. The crude nitrate (" caliche ") contains about 0.2 per cent of this compound and most of the iodine of commerce comes from this source. * Iodine is found in the thyroid gland, and its presence seems to be physiologically important. The diseases of goiter and cretinism are, apparently, connected with a deficiency of iodine. Properties of Iodine. Iodine is obtained in the form of black, crystalline scales which melt at 114.2. The liquid boils at 184.3, but gives off a beautiful violet vapor at much lower temperatures. The weight of the vapor indicates that the for- mula is I 2 at temperatures not far above the boiling point, but even at 700 the molecules dissociate appreciably into single atoms, just as the bromine and chlorine molecules dissociate at much higher temperatures. The stability of iodides is much less than that of bromides or chlorides, and the stability of the iodine molecule is also much less than that of the bromine mole- cule. In general the affinity of the nonmetallic elements toward metallic elements decreases with increasing atomic weight. Iodine dissolves very slightly in pure water. It dissolves more easily in alcohol, giving a brown solution, called tincture l of 1 The name tincture is given in pharmacy to a solution in alcohol of some substance or of the active constituents of some plant. IODINE 145 iodine, which is used in medicine. Iodine also dissolves in a so- lution of potassium iodide. There is evidence that in solution it forms an unstable compound, KIs, but the dilution of a solution having this composition causes the precipitation of a part of the iodine. In chloroform, carbon bisulphide and other solvents with which it does not combine, iodine forms violet solutions. Iodine gives with starch emulsion, in the presence of hy- driodic acid or an iodide, a deep blue color, which is very charac- teristic and which is used as a test for free iodine or for starch. From iodides the iodine must be liberated by some oxidizing agent, best by nitrous acid for the detection of minute quantities of the element. Chlorine may be used, but an excess oxidizes the iodine to iodic acid and destroys the color. The color of the starch iodide is also destroyed by heat, but returns, in part, on cooling the solution. Hydriodic Acid. It has already been pointed out that the affinity between hydrogen and iodine is much less than that between hydrogen and chlorine or bromine. If slightly diluted sulfuric acid is poured on potassium iodide, some hydriodic acid is liberated, but the larger part of the acid acts upon more of the sulfuric acid, reducing it to sulfur dioxide or even to hydrogen sulfide, H 2 S. The student should write the equations for the three reactions involved in the last statement and compare with the somewhat similar conduct of hydrobromic acid. Hydriodic acid is best prepared by melting together in a distilling bulb 1 part of red phosphorus with 20 parts of iodine, forming a mixture of phosphorus triodide, Pis, and iodine. 1 When the mixture is cold, a stopper bearing a* separatory funnel containing 4 parts of water is fitted to the neck of the bulb, and the side tube is connected with a small U-tube containing a very little water to wash the gas and retain nearly all of the iodine 1 Method of Lothar Meyer slightly modified, Ber. 20, 3381. The usual direction, which gives enough phosphorus to form PIa gives rise to the formation of phosphonium iodide, PH 4 I, and this may stop the exit tubes or contaminate the product. For the same reason iodine cannot well be removed from the gas by moistened red phosphorus, as directed by some authors. 146 A TEXTBOOK OF CHEMISTRY which passes over. If a solution of hydriodic acid in water is desired, the delivery tube should not dip beneath the surface of the water which is to absorb the gas. When all is ready, the water is dropped slowly on the mixture of phosphorus triodide and iodine. After all has been added the last of the hydriodic acid is driven over by warming. The equation is : PI 3 + I 2 + 4 H 2 O = H 3 P0 4 + 5 HI Phosphoric Acid The solution of hydriodic acid in water has properties similar to those of the corresponding solutions of hydrochloric and hydrobro- mic acids. The solu- tion of constant boiling point under atmos- pheric pressure boils at 127, has a specific gravity of 1.70 and contains 57 per cent of the acid. The hy- driodic acid is slowly oxidized with libera- tion of iodine on ex- posure of the solution to the air, and for this reason the aqueous Fig. 49 acid is almost always colored red or brown. Direct Combination of Hydrogen and Iodine. Reversible re- actions. Equilibrium. Mixtures of equal volumes of chlorine and hydrogen or of bromine vapor and hydrogen combine com- pletely when heated to the temperature of combination, and neither hydrochloric acid nor hydrobromic acid dissociates ap- preciably unless heated to a quite high temperature. 1 If a mix- 1 Haber (Thermodynaimk technischer Gasreaktionen, 1905, S. 95) calculates the dissociation of hydrobromic acid as only 0.15 per cent at 727. EQUILIBRIUM 147 ture of equal volumes of hydrogen and of iodine vapor is heated, however, there is no temperature at which the combination to form hydriodic acid will be complete, even if the mixture is heated for an indefinitely long time. On the other hand, if hydriodic acid is heated, it decomposes slowly, even at quite low temperatures, but never completely, no matter how long it is heated, unless the temperature is very high indeed. If we start with two sealed glass tubes, one containing one part by weight of hydrogen with 127 parts of iodine and the other containing hydriodic acid, and heat both tubes at the same tem- perature till the composition no longer changes, it will be found that each tube contains, on cooling, a mixture of hydrogen, iodine and hydriodic acid but that the composition of the mix- ture in the two tubes is identical. This result is most easily explained by supposing that we have here a reversible reaction : H 2 + I 2 ^ HI + HI and that when equilibrium is reached the reaction does not stop but continues in such a way that just as many molecules of hy- driodic acid are formed in a minute as are decomposed in the same time. The composition of the mixture which is in equilibrium varies with the temperature, as will be seen from the following table : COMPOSITION OF THE EQUILIBRIUM MIXTURE OF HYDROGEN, IODINE AND HYDRIODIC ACID TEMPERATURE PROPORTION OP HYDRIODIC ACID PROPORTION OF HYDROGEN AND IODINE 283 0.8213 0.1787 328 0.8115 0.1885 374 0.7990 0.2010 393 0.7942 0.2058 427 0.7843 0.2157 508 0.7592 0.2408 148 A TEXTBOOK OF CHEMISTRY It is evident from this table that when equilibrium is reached, the combination of hydrogen and iodine is more nearly complete at low temperatures than at high ones. (In accordance with the principle of Le Chatelier (p. Ill) is the combination accompanied by the evolution or by the absorption of heat ?) Speed of Chemical Reactions. If hydriodic acid is heated at a given temperature, the decomposition seems to proceed more and more slowly as the reaction goes on, as will be seen from the following table : RATE OF DECOMPOSITION OF HYDRIODIC ACID AT 374 Ol TOTAL TIME OF HEATING FRACTION DECOMPOSED FRACTION DECOMPOSED IN 1 MINUTE 360 minutes 0.0715 0.00020 720 minutes 0.1267 0.00015 1080 minutes 0.1596 0.00009 1440 minutes 0.1715 0.00003 A little consideration of these results leads us to the conclusion that the decreasing rate is due to two causes : first, because the amount of hydriodic acid in the mixture is constantly decreasing, and second, because the hydrogen and iodine which result from the dissociation are recombining. If we begin with the mixture of hydrogen and iodine, the com- bination appears to be rapid at first, but soon decreases in its rate as the amounts of hydrogen and iodine grow less and as the hydriodic acid which is formed begins to decompose. In order to get the real rate for the decomposition of pure hydriodic acid or for the combination of pure hydrogen and iodine, we might measure the rate for the first infinitesimal fraction of a minute, if that were possible. For the decomposition at 374 the deter- mination of the composition of the mixture at the end of six hours gives a rate for the decomposition of 0.0002, or 1/5000 1 Bodenstein, Z. physik. Chem, 29, 295. SPEED OF REACTIONS 149 part of the whole in one minute, and this is comparatively close to the rate of decomposition for pure hydriodic acid, since at the end of the first six hours, only about 1/15 of the whole has been decomposed. Concentration and Speed of Reaction. In order to calculate the true rate more accurately, it is necessary to use the law of the relation between the concentration of reacting substances and the speed of reaction, 1 which has been based on a careful study of many different reactions which take place slowly enough so that the rate can be measured. This is that the speed of any reaction at a constant temperature (i.e., the part of the whole which will react in unit time) is equal to the product of the concentrations of each reacting substance multiplied by a force which is characteristic of the given reaction. The force which causes the reaction, and which is given a numerical value in relation to the speed of the reaction by this law, is spoken of in most textbooks as chemical affinity, but it is evidently complex, depending on the attraction between the atoms which unite, the attraction between atoms which separate, the temperature, the presence of catalytic agents and perhaps on many other factors. 2 The temperature, especially, has a large effect, such that the speed of a reaction is usually doubled for an increase of 10. For hydriodic acid the law may be given the following expres- sion : Let C-g_ t = Concentration of hydrogen, Ci 2 = Concentration of iodine, CHI = Concentration of hydriodic acid, FI = Force driving the reaction, H 2 + I 2 ^ HI + HI, to the right, 1 Often called, less correctly, the "Law of Mass Action." 2 In the case of hydriodic acid the nature of the force is wholly changed by the action of light and the reaction becomes unimolec- ular: HI = H + I. Bodenstein, Z. physik. Chem. 22, 23. 150 A TEXTBOOK OF CHEMISTRY F 2 = Force driving the reaction to the left, Si = Speed of formation of hydriodic acid, 82 = Speed of decomposition, all at a given temperature Then : CW 2 x Cj 2 x FI = Si CHI X CHI X FZ = Sz l By means of these formulas it is possible to calculate from the results obtained by heating mixtures of hydrogen and iodine, or by heating hydriodic acid for different lengths of time, the rate of combination or of decomposition for the pure substances at unit concentration. The methods of calculation are compli- cated and need not be given here. The results are as follows : RATE OF FORMATION AND DECOMPOSITION OF HYDRIODIC ACID TEMPER- ATURE -Si FRACTION OF WHOLE FORMED IN ONE MINUTE & FRACTION OF WHOLE DECOMPOSED IN ONE MINUTE Si 52 302 0.000353 0.00000326 1:108 374 0.0140 0.000221 1: 63 427 0.172 0.0031 1: 55 It is clear from this table that the equilibrium mixtures as given on p. 147 contain much more hydriodic acid than hydrogen and iodine because the combination takes place much more rapidly than the decomposition. It is also seen that the speed of each reaction increases rapidly with the temperature and that the speed of decomposition increases more rapidly than the speed of combination. From this it follows that combination is ac- 1 This depends on the fact that the decomposition results from the action of two molecules of hydriodic acid on each other with the formation of molecules of hydrogen and iodine. Such a reaction is called bimolecular. The action of light seems to cause the direct separation of the atoms of hydrogen and iodine from each other and the reaction in that case is unimolecular. See footnote, p. 149. SPEED OF REACTIONS 151 companied by an evolution of heat, in accordance with the principle of Le Chatelier (see below and also pp. Ill and 201). Calculation of the Relative Speed of Two Reactions from the Composition of an Equilibrium Mixture. If we know the nature of a reversible reaction and the composition of the equilib- rium mixture, it is possible to calculate the relative speeds of the opposing reactions, at unit concentration. From the formulas given : C\s /~Y \.s If O Hj s^ Ij ^ 1 * CHI X CHI X F 2 = 02 From these equations, FI = Si and F* = <S 2 when the concen- tration is 1 for each of the reacting substances. At 374 the composition of the equilibrium mixture is very near to : 80 per cent of HI, 10 per cent of H2, and 10 per cent of I 2 , by volume. From this : CHI = 0.8 C Ha = 0.1 Ci, = 0.1 Let S 3 be the speed of combination and 84 the speed of decomposi- tion, at equilibrium. At equilibrium <S 3 must equal (84. Then, CH, X C Ia X Fi = S, 0.1 X 0.1 X Fi = S 3 CHI X CHI X F2 = $4 0.8 X 0.8 X F 2 = S 4 Since <S 3 = S 4 0.1 X 0.1 X F! = 0.8 X 0.8 X F 2 F and - = 64. It will be seen that this result agrees closely, as it should, with the result obtained by the direct measurement of the speeds of the two reactions (p. 150). The result may 152 A TEXTBOOK OF CHEMISTRY also be checked by the following calculation, putting FI = Si and FZ = 82 and using the values for Si and $2 given on p. 150 : C H2 X Ci 2 X Si = S, 0.1 X 0.1 X 0.0140 = S 3 = 0.000140 CHI X CHI X 02 = $4 0.8 X 0.8 X 0.00022 = S 4 = 0.000140 This means, of course, that after equilibrium is reached at 374, the dissociation and recombination still continue at the rate of about 1/7000 part of the whole each minute. Effect of Removing one of the Reacting Substances. Displacement of the Equi- librium Point. If a tube containing hy- . _~ driodic acid is heated in such a way that one end of the tube is kept cool, the iodine which results from the dissociation will partly condense and the concentration of the iodine, Ci 2 , will be diminished in the equation : C H2 X C l2 X Fi = S, which gives the rate of recombination. Under these conditions it is evident that the decomposition must go much farther than usual before S 3 = 84- In other words, the removal of one of the constituents of a reversible reaction always displaces the equilib- rium to the side on which the constituent removed appears. This effect has been noticed ' before in the reaction between iron and steam and in that between salt and sulfuric acid. If the end of the tube could be kept cold enough so that the vapor pressure of iodine in it would be reduced to zero, 83 would finally become zero, and the decomposition of the hydriodic acid would continue till it was complete. Heat of Formation of Hydriodic Acid. The heats of formation of the compounds of the halogens with hydrogen are as follows : H 2 + F 2 = 2 X 38,000 calories H 2 + C1 2 = 2 X 21,800 calories H 2 + Br 2 = 2 X 8300 calories H a H- I 2 = 2 X 96 calories FLUORINE 153 The heat energy liberated during the combination grows less with increasing atomic weight and becomes very small in the case of iodine. Bodenstein calculates (Z. physik. Chem. 29, 313) the heat of combination of hydrogen and iodine in gaseous form as follows At 510, 2 X 2222 calories At 290, 2 X 943 calories At 20, 2 X 96 calories Fluorine, F, 19.0. Occurrence. The chlorides, bromides and iodides of four of the most common metals, calcium, magnesium, sodium and potassium, are all easily soluble in water, and these three halogens are found chiefly associated with these metals, and especially with sodium, in the ocean and in brines. Fluorine, on the other hand, combines with calcium to form an almost insolu- ble compound, calcium fluoride, CaF 2 , and for this reason can never be found in more than very small amounts in natural waters, which practically always contain calcium. Fluorine is found chiefly as calcium fluoride, CaF 2 , in the mineral fluorite. Cryolite, a double fluoride of aluminium and sodium, NasAlF 6 (or AlF 3 .3NaF), found in Greenland, and apatite, a double phos- phate and fluoride, or chloride, of calcium, Ca 6 (PO 4 )3F or Ca 5 (PO 4 )3Cl, 1 found in Canada and elsewhere, are important as sources of aluminium and phosphorus rather than as sources of fluorine. Preparation. When we consider that the heat of combination of hydrogen and fluorine is nearly twice that for the combination of hydrogen and chlorine (p. 152), and remember that chlorine can take a part of the hydrogen away from oxygen, we may be led to expect that free fluorine cannot exist in the presence of water. It was not till this came to be clearly understood that Moissan succeeded in obtaining the free element in 1886. He did this by electrolyzing a solution of potassium fluoride, KF, in anhydrous hydrofluoric acid. He used a U-tube of platinum 1 Note the relation between this formula, the formula of phosphoric acid, H 3 PO 4 , and the valences of calcium and of fluorine or chlorine. 154 A TEXTBOOK OF CHEMISTRY at first, but showed later that a tube of copper is only very slightly attacked by the fluorine, if the temperature is kept at 23, or below, by a freezing mixture. In the electrolysis, fluorine goes toward the anode and is liberated there while potassium and hydrogen go toward the cathode, but only hydro- gen is liberated, because hydrogen ions are discharged at a much lower potential than potassium ions. Properties. Fluorine is a greenish yellow gas, less deeply colored than chlorine. The weight of a gram molecular volume is 38 grams, showing that the formula is F2. Fluorine is the most active of the nonmetallic elements, as is to be expected from its unique position in the periodic system. It combines directly and vigorously with nearly all elements, both metals and nonmetals, except with oxygen. Many elements, as iodine, phosphorus, arsenic, carbon as charcoal or lampblack, silicon, potassium and sodium take fire and burn in the gas, forming fluorides. Fluorine will also displace nearly all other non- metallic elements from their compounds. If led into water, it gives hydrofluoric acid and oxygen, rich in ozone : 3H 2 O + 3F 2 = O 3 + 6HF Ozone Hydrofluoric Acid Etching Glass. Hydrofluoric Acid may be easily prepared by warming a fluoride with concentrated sulfuric acid : CaF 2 + H 2 SO 4 = CaSO 4 + 2 HF Hydrofluoric acid is a gas which may be condensed to a liquid much more- easily than the other halogen acids. The anhydrous liquid boils at 19.4. The liquid mixes with water in all propor- tions, the concentrated solution fuming in the air in the same manner as concentrated solutions of the other halogen acids. The gaseous acid is very poisonous and the concentrated or anhy- drous acid causes painful wounds, which are very difficult to heal. The most interesting property of the acid is its action on sili- FLUORINE 155 cates and especially on glass, which is a complex silicate of calcium and sodium or other metals. When hydrofluoric acid comes in contact with glass, the fluorine combines both with the silicon and with the metals of the glass : CaSiO 3 + 6 HF = SiF 4 + CaF 2 + 3 H 2 O Calcium Silicon Silicate Tetrafluoride The reaction may be looked on as a displacement of oxygen by fluorine, two atoms of fluorine displacing one atom of oxygen in accordance with the valences of the two elements. Silicon tetrafluoride, SiF 4 , is a gas and escapes. By covering a glass object with beeswax, which is not affected by hydrofluoric acid, and exposing it to the action of the gas, after drawing lines or figures through the wax so as to expose part of the surface of the glass, it is possible to etch the exposed surface and obtain per- manent markings of any form that is desired. Graduation marks on thermometers, burettes, eudiometers, etc., are made in this way. The best results are obtained by exposing the glass to the anhydrous gas for some hours. Commercial hydrofluoric acid is kept in lead bottles, which are only slightly attacked. The pure acid must be kept in platinum or in bottles made of ceresin, a mineral wax with a higher melting point than that of paraffin. The constant boiling solution boils at 120 and contains 35 per cent of the acid. Hydrofluoric acid, unlike the other halogen acids, forms both acid and neutral salts. Thus it forms with potassium, acid potassium fluoride, KHF 2 , as well as the neutral, or normal, fluo- ride, KF. The formation of these acid salts seems to be closely related to the abnormal density of the gas and indicates that the true formula of the acid in solution or at low temperatures is probably H 2 F 2 instead of HF. The weight of a gram molecular volume of gas varies from 51.2 grams at 26 to 20.6 grams at 88. At the lower temperature the gas is evidently more complex than H 2 F 2 , for which the gram-molecular-volume would weigh 40 grams. 156 A TEXTBOOK OF CHEMISTRY Metallic Elements of Group VII. Manganese stands between chlorine and bromine in the seventh group of the Periodic System when the system is given its simplest form (p. 134). It re- sembles chlorine in the dioxide, MnO 2 , which corresponds to chlorine dioxide, C1O 2 , and in permanganic acid, HMnQj, corresponding to perchloric acid. But in most of its properties manganese is metallic, and it will be considered further later (p. 533). The Periodic System indicates the possibility of three or four other elements in the seventh group with atomic weights greater than that of bromine, but no such elements have been dis- covered. EXERCISES 1. Write the equations for sixteen reactions between the following acids and bases, giving normal salts : Hydrochloric acid, HC1 ; perchloric acid, HC1O 4 ; sulfuric acid, H2SO4; phosphoric acid, H 3 PO 4 ; sodium hydroxide, NaOH; ferrous hydroxide, Fe(OH) 2 ; ferric hydroxide, Fe(OH) 3 ; stannic hydroxide, Sn(OH) 4 . 2. Write the equations for the reactions between the following salts and sulfuric acid : sodium chloride, NaCl ; calcium chloride, CaCl 2 ; sodium perchlorate, NaClO 4 ; aluminium chloride, A1C1 3 . 3. Write the equations for the reactions between hydrochoric acid and the following oxidizing agents. Notice the changes in valence : MnO 2 -* MnCl 2 KMnO 4 -> KC1 and MnCl 2 HC10 -> HC1 Pb 3 4 -> PbCl 2 K 2 Cr 2 7 ^KClandCrCl 3 KC1O 3 ->KC1 4. Write the equation for the reaction between potassium iodide, manganese dioxide and sulfuric acid, giving K 2 SO 4 and MnSO 4 . 5. Write the equation for the reaction between calcium bromide, CaBr 2 , potassium permanganate, KMnO 4 and sulfuric acid, giving calcium sulfate and the other products to be expected. These reactions are introduced here to give the student facility in writing equations on the basis of the valence of the elements. The fundamental conception of valence is that each atom has the power of holding directly in combination a definite, small number of other atoms. Thus, when we write the graphical formula H Cl, the thought which it is intended to convey is that a hydrogen or THE HALOGEN FAMILY 157 3, chlorine atom holds directly to only a single other atom in the / H compound, hydrochloric acid. In water, H O H or (X , the H oxygen atom holds directly to two other atoms. In some sense we may think that an oxygen atom has two points of attachment for /H other atoms. In ammonia, N^-H or H N\ , in accordance with \H X H the same theory, each nitrogen atom holds directly to three hydrogen atoms. In the series of oxides of the Periodic System the elements of the zero group do not combine with other elements at all, and these elements are considered to have a valence of zero. The elements of the first group are univalent, and one bivalent oxygen atom can hold two atoms of these elements as in Na2O or Na O Na. The bivalent atoms of the second group can hold bivalent oxygen atoms, atom for atom, as in Mg=O. In the third group, where the elements are trivalent, if we consider an atom of such an element as combined with one atom of oxygen, one valence of the first element will remain unsatisfied, thus, B . If a second atom of oxygen is added, one valence of this will be unattached, B\Q- On adding a second atom of the trivalent element and a third atom of oxygen, all of the valences will be balanced In the next group a quadrivalent atom can balance two bi- ~ valent oxygen atoms, Cx' . The same principles may be easily V extended to the compounds, N2O 5 , SO 3 and C1 2 O 7 . When the oxides are those of nonmetallic elements, they will, in most cases, combine with water to form acids. In this case one valence of one oxygen atom separates from one nonmetallic atom, and the hydro- gen, H, of the water attaches itself to the oxygen, while the hydroxyl, OH, of the water attaches itself to the nonmetallic atom : H -X X H-0 158 A TEXTBOOK OF CHEMISTRY In the reactions between acids and bases the same principles of bal- ancing valences are to be applied, the only difference being that the valences of the metal on the one hand are to be balanced against the valences of the acid groups on the other. Since the hydroxyl group O H is univalent, the number of hydroxyl groups in the base gives the valence of the metal of the base, while the number of replaceable hydrogen atoms gives the valence of the acid group. Thus iron is bi- valent in ferrous hydroxide, Fe(OH) 2 , and trivalent in ferric hydroxide, Fe(OH) 3 , 1 while the sulfate group, SO 4 , of sulfuric acid is bivalent and the phosphate group, PO 4 , of phosphoric acid, H 3 PO 4 , is trivalent. By representing the valences with lines, it is a simple matter to balance the valences of a metal against the valences of an acid radical and so determine the correct formula of a salt. Thus for ferric sulfate the formula must be : or for ferric phosphate, Fe=PO 4 . A little practice of this sort will soon enable a student to write correct formulas, such as Fe2(SO 4 )s or FePO 4 , without the use of the lines to indicate valences. If these prin- ciples are once understood, one needs to remember only the formula of a single salt of any metal with some well-known acid in order to be able to write the formulas of the normal salts of the metal with a hundred or more acids whose formulas are known. In reactions which involve oxidation and reduction it often happens that the valence of some element changes. In a reduction, oxygen or some other element is removed without being replaced, or hydrogen is added, and to do this hydrogen must usually be furnished from some source, and the element combined with this hydrogen is often liberated in the free state. Thus in the reaction between manganese dioxide and hydrochloric acid quadrivalent manganese changes to the bivalent form. The extra oxygen atom is balanced by hydrogen from the hydrochloric acid and the chlorine of the latter is liberated. The manganese dioxide is reduced, the hydrochloric acid is oxidized: O Cl :}jMn<:+2H 2 o+ci 2 H-Cl 1 For the sake of simplicity the possibility of such doubled formulas as Fe 2 (OH) 4 and Fe 2 (OH) 6 is not presented here. THE HALOGEN FAMILY 159 In a similar way if hydrochloric acid acts on potassium permanganate, KMnO 4 , only three chlorine atoms are taken by the potassium and manganese, and the hydrochloric acid which furnishes these will give only three of the eight hydrogen atoms necessary to balance the four oxygen atoms of the permanganate molecule. To balance the remainder of the oxygen atoms, five more hydrogen atoms will be required. This gives us the reaction : KMnO 4 + 3 HC1 = KC1 + MnCl 2 + 4 H 2 O + 5 Cl + 5HC1 If we wish to take account of the fact that free chlorine has the formula C1 2 , the equation must, of course, be doubled, giving : 2 KMn0 4 + 16 HC1 = 2 KC1 + 2 MnCl 2 + 8 H 2 O + 5 C1 2 Which substance is reduced and which is oxidized in this reaction ? CHAPTER XI SULFUR, SELENIUM AND TELLURIUM THE nonmetallic elements of Groups VI and VII of the peri- odic system are : O .... 16 F .... 19 S .... 32 Cl . . . . 35.5 Se .... 78 Br .... 80 Te . . . . 127.6 I .... 127 Sulfur, S, 32.0. Occurrence. Oxygen is found free in nature, partly because of its great abundance, forming, as it does, one half of that portion of the earth which we can examine directly, partly, probably, because of its unique relationship to carbon and the growth of plants (p. 312). Sulfur is also found free, partly because it is a comparatively abundant element and partly because it is easily liberated from hydrogen sulfide and other sulfides by the action of oxygen and some compounds of oxygen. Free sulfur is found in large quantities in Sicily and in Louisiana. Until about 1903 the sulfur mines of Sicily held, for a long time, a practical monopoly of the sulfur markets of the world, almost the only competition coming from the sulfur obtained by the Chance process (p. 457) as a by-product in the manufacture of sodium carbonate. The sulfur in Sicily is mixed with other minerals, from which it is separated by piling up the mixture and setting fire to the sulfur in such a way that the heat from burning a part of the sulfur melts the rest and the latter runs out and is collected. The process is, of course, a wasteful one as a pound of coal would give nearly as much heat as four pounds of sulfur (p. 27). The crude sulfur is re- fined by distillation. If the vapors are condensed in cold cham- bers, the sulfur takes the form of flowers of sulfur, just as the freezing of water vapor gives snow. If the condensing room is 160 SULFUR 161 hot water above the melting point of sulfur, the liquid sulfur which collects on the bottom is run into molds and forms the roll brimstone of commerce. The extensive deposits of sulfur in Louisiana are below a layer of quicksand, and for a long time after they were discovered no practical method of working the deposits was known. The difficulty was finally solved by a process invented by Mr. Frasch of New York'. Three concentric iron pipes are sunk to the level of the sulfur, and hot water under pressure is forced down be- tween the two outer pipes, the pressure of the water being great enough so that the boiling point is raised above the melting point of the sulfur, 114.5. The hot water melts the sulfur, which rises in the second tube, the end of which is brought below the sur- face of the melted sulfur. To bring the sulfur to the surface, air is forced down through the cen- tral tube, the sulfur and com- pressed air rising together between the central and second tubes. By this process the production of sulfur in the United States was increased from 3500 tons in 1900 to 265,000 tons in 1911. The world's production of sulfur in 1909 was 818,000 tons. Sulfur is also found in nature combined with metals as metallic sulfides and with metals and oxygen as sulfates. The most important sulfides are lead sulfide, or galena, PbS, zinc sul- fide, or sphalerite, ZnS, iron sulfide, or pyrite, FeS 2 , and an iron-copper sulfide, copper pyrites, CuFeS 2 . The most important sulfates are calcium sulfate, or gypsum, CaSO4.2 H 2 O, and barium Fig. 51 162 A TEXTBOOK OF CHEMISTRY sulfate, or barite, BaSO 4 . Of these, only iron pyrites is used primarily as a source of sulfur, for the manufacture of sulfuric acid. The other sulfides are used primarily for the metal which they contain, but sulfuric acid is sometimes made from them as a by-product. Allotropic Forms of Sulfur. Sulfur may exist in three well- defined solid forms, in two liquid forms, which correspond closely to two of the solid forms, and in three gaseous forms. The solid forms are : 1. Rhombic Sulfur. Light yellow crystals, most often in the form of rhombic pyramids (p. 194), found in nature and formed by crystallization from carbon disulfide, in which sulfur is easily soluble. The specific gravity is 2.06 and the melting point 114.5. This is the most dense and most stable form at ordinary temperatures, and the other forms change to this form more or less rapidly at temperatures below 96. 2. Monoclinic Sulfur. When melted sulfur is allowed to cool slowly, it crystallizes in long, transparent needles of the monoclinic system (p. 195). These have a specific gravity of 1.96 and melt at 119. This form of sulfur is stable only at tem- peratures between 96 and 119. At lower temperatures it changes more or less quickly to the rhombic form. The outer form of the needles is retained, but they become opaque and then consist of microscopic crystals of the rhombic form. 3. Amorphous, Insoluble Sulfur. When sulfur which is heated above 160 is cooled quickly with care that it does not come in contact with crystals of sulfur, which would cause a rapid transformation to the crystalline form, it assumes a soft, plastic form, which hardens to a solid mass after some hours or days. If this hardened mass is treated with carbon disulfide, it will be found to be mostly insoluble and the insoluble portion is amorphous, i.e. it has no crystalline structure. The liquid forms of sulfur are : 1. Mobile Liquid Sulfur (S A ). Between the melting point (114.5 or 119) of either form of sulfur and 160 it forms a mobile, pale yellow liquid. SULFUR 163 2. Viscous Liquid Sulfur (S^). When heated to 160, sulfur suddenly becomes dark colored and so viscous that a test tube containing it may be inverted without its running out. If heated to a higher temperature, the liquid becomes gradually somewhat more mobile and finally boils at 444.7 . 1 The boiling point is frequently used to fix a point on the scale of ther- mometers and pyrometers. The gaseous forms of sulfur are : 1. Sg. When sulfur is converted into a vapor at 250, under low pressure, the weight of a gram molecular volume is nearly 256 grams, indicating that there are eight atoms in one molecule and that the formula is Sg. 2. S 2 . Even at the boiling point (444.7), the weight of a gram molecular volume of sulfur vapor is considerably less than 256 grams and it was formerly supposed that the formula at tempera- tures a little higher than this was 85. A more careful study of the matter has demonstrated that the formula Sg is the true one at low temperatures and that the heavy molecules dissociate as the temperature rises until, at 800, the formula becomes S 2 , the weight of a gram molecular volume at that temperature being 64 grams. It is still somewhat uncertain whether the larger molecules dissociate directly into molecules of 82 or whether intermediate molecules of 84 or 85 are found. 2 3. S. When sulfur vapor is heated to a very high tempera- ture (2000), it dissociates still further until the gram molecular volume weighs only 32 grams and the formula becomes 8. We may suppose that at high temperatures the collisions between molecules become more and more violent until, at last, the affinity between the atoms can no longer withstand the disrup- tive effect of the collisions. Properties and Uses of Sulfur. Sulfur burns readily in air or oxygen, forming sulfur dioxide, SO2, with usually a small amount of the trioxide, SOs. The volume of the sulfur dioxide 1 Bulletin of the Bureau of Standards, Vol. 7. pp. 3 and 129. 2 See Premier and Schupp, Z. physik. Chem. 68, 144 (1909), and Stafford, ibid. 77, 66 (1911). 164 A TEXTBOOK OF CHEMISTRY is almost the same as the volume of the oxygen from which it is formed. (How does this follow from Avogadro's law and the formulas of oxygen and sulfur dioxide ?) Sulfur combines with most metals when heated with them, forming sulfides. The combination with iron to ferrous sulfide, FeS, and with copper to cuprous sulfide, Cu 2 S, is attended with considerable evolution of heat. Sulfur is burned to sulfur dioxide for the manufacture of sul- furic acid, for use in bleaching straw goods, for the " sulfuring " of fruit in the process of drying, to prevent darkening and the growth of harmful organisms. Sulfur is also used in the manu- facture of carbon disulfide, of gunpowder and of india rubber. It is used directly or in a lime-sulfur wash for application to vines, fruit trees, etc., to prevent the growth of fungi or other harmful organisms. Hydrogen Sulfide, H 2 S, is found in many natural waters, the so-called sulfur waters. It is formed by the decomposition of organic matter containing sulfur and is one cause, though by no means the only reason, for the disagreeable odor of decayed eggs and sewage. Hydrogen sulfide is formed when hydrogen is passed over sulfur heated to its boiling point, as can be shown by passing the gas, subsequently, through a solution of lead nitrate, in which it will produce a black precipitate of lead sulfide : Pb(N0 3 ) 2 + H 2 S = PbS + 2 HNO 3 Lead Lead Nitrate Sulfide The reaction is reversible : 2 H 2 -f S 2 ^ 2 H 2 S as can be shown by passing hydrogen sulfide through a hot glass tube, in which a ring of sulfur will be deposited beyond the point that is heated. Hydrogen sulfide is prepared in the laboratory by the action of hydrochloric or sulfuric acid on ferrous sulfide. HYDROGEN SULFIDE 165 FeS + 2 HC1 = FeCl 2 + H 2 S Ferrous Chloride FeS + H 2 SO 4 = FeSO 4 Ferrous Sulfate H 2 S For the preparation of the gas on a small scale the apparatus used for the preparation of hydrogen is suitable. For the use of a laboratory the Parsons apparatus (J. Am. Chem. Soc. 25, 233) is better, because the acid remains in con- tact with the ferrous sulfide till the action is complete. Hydrochloric acid is more satisfactory than sulfuric acid for such a generator, be- cause ferrous sulfate is less soluble than fer- rous chloride and some- times crystallizes in the tube through which the spent acid escapes (Fig. 52). Hydrogen sulfide is a colorless gas with a very disagreeable odor. It is quite poisonous, if breathed in more than small amount. It may be condensed to a liquid, which boils at 62 and frozen to a solid, which melts at 85. Fig. 52 Solution of Hydrogen Sulfide. Henry's Law. One volume of water absorbs, or dissolves, 4.4 volumes of hydrogen sulfide at 166 A TEXTBOOK OF CHEMISTRY 0, 3.7 volumes, at 10 and 3.1 volumes at 20. The volume of the gas dissolved is, between quite wide limits, independent of the pressure. Since the weight, or amount of the gas in a given volume, is proportional to the pressure, it follows that the amount of the gas dissolved varies directly with the pressure. This is known as Henry's Law (discovered in 1803). It applies to partial pressures also. Thus if a gaseous mixture contains 10 per cent by volume of hydrogen sulfide, the amount dissolved from such a mixture at 20 will be only 0.31 volume. One hun- dred cubic centimeters of water in contact with pure oxygen dis- solve 4.9 cc. of the gas at ; in contact with nitrogen 100 cc. dissolve 2.35 cc. of nitrogen. One hundred cubic centimeters of water in contact with air will contain, therefore, 4.9 X 0.21 = 1.04 cc. of oxygen and 2.35 X 0.78 = 1.83 cc. of nitrogen. The law does not hold for gases which are very easily soluble in water, such as hydrochloric acid or ammonia. In accordance with Henry's law, water containing hydrogen sulfide loses the gas rapidly on exposure to the air, in which the partial pressure of the gas is, of course, zero. In addition to this the oxygen absorbed by the water reacts with the hydrogen sulfide, liberating sulfur : 2 H 2 S + O 2 = 2 H 2 O + 2 S The action is similar to the liberation of chlorine from chlorides by fluorine, but is far less rapid. Sulfides. Groups of Analytical Chemistry. When hydrogen sulfide is passed into a neutral or slightly acid solution containing salts of certain metals, such as arsenic, mercury and lead, the metal is precipitated as a sulfide because the sulfides of these metals are extremely insoluble : 2 AsCl 3 + 3 H 2 S = As 2 S 3 + 6 HCl HgCl 2 + H 2 S = HgS + 2 HCl Pb(N0 3 ) 2 + H 2 S = PbS + 2 HN0 3 If hydrogen sulfide is passed into an alkaline solution contain- ing the salts of some other metals, such as iron, zinc and manga- nese, which are not precipitated from acid solutions, these metals, STRENGTH OF ACIDS 167 whose sulfides are also very insoluble, but more soluble than those of the metals of the first group, are precipitated also as sulfides. FeSO 4 + H 2 S + 2 NaOH (or Na 2 S) = FeS + Na 2 SO 4 + 2 H 2 O ZnS0 4 + Na 2 S = ZnS + Na 2 SO 4 MnS0 4 + Na 2 S = MnS + Na 2 SO 4 A part, but not all, of the metals of the first group are precipi- tated from alkaline as well as from acid solutions. The reason for the exceptions need not be discussed here. (See p. 261.) A third class of metals form salts which are not precipitated from acid, neutral or alkaline solutions. The conduct of solutions of metals toward hydrogen sulfide, as just outlined, is the basis for the separation of metals into three fundamental groups for the purposes of analytical chemistry. Hydrosulfuric Acid. Strength of Acids. A solution of hydro- gen sulfide in water will redden blue litmus paper and will neutral- ize a solution of sodium hydroxide, or in other words a certain amount of a solution of a base must be added before the hy- droxide will turn the litmus blue. These are the properties of an acid, and hydrogen sulfide is sometimes very properly called hydrosulfuric acid, just as hydrogen chloride is called hydrochloric acid. We have seen that a solution containing a milligram molecule of hydrochloric acid in 10 cc. of water freezes at 0.355, while a solution of alcohol containing a milligram molecule in 10 cc. freezes at 0.184, and the difference was explained by supposing that the hydrochloric acid separates largely into hydrogen (H + ) and chloride (Cl~) ions. A solution of hydrogen sulfide which contains one milligram molecule in 10 cc. freezes at 0.196. 36.5 mg. HC1 in 10 cc. of H 2 O freezes at - 0.355 46 mg. C 2 H 6 O in 10 cc. of H 2 O freezes at - 0.184 34 mg. H 2 S in 10 cc. of H 2 O freezes at - 0.196 This indicates that such a solution of hydrogen sulfide con- tains comparatively few hydrogen ions. This conclusion is 168 A TEXTBOOK OF CHEMISTRY confirmed by the electrical conductivity of the solution. The solution of hydrochloric acid referred to is a very much better conductor (nearly 2000 times) of electricity than the solution of hydrogen sulfide. During a long period in the history of chemistry acids were spoken of as strong or weak according to whether they could expel other acids from their salts or not. Thus sulfuric acid was thought to be stronger than hydrochloric or nitric acid because it would expel these acids from salt or saltpeter. We have seen that such a view can no longer be held (p. 119) and that all such reactions are reversible. There is another sense, however, in which some acids are strong while others are weak, and the basis for a true distinction of this kind has just been indicated. A strong acid is one which separates largely into hydrogen ions and negative ions in an aqueous solution. A weak acid is one that separates to only a comparatively small degree into hydrogen ions and negative ions in solution. In this sense hydrochloric acid is one of the strongest of the acids, sulfuric acid is weaker but still a very strong acid, hydrofluoric acid is much weaker, acetic acid is still weaker and hydrosulfuric acid, H 2 S, is very weak indeed. Acids like hydrosulfuric acid which contain two hydrogen atoms may ionize in either of two ways : or In the case of weak acids the ionization probably takes place almost exclusively in the first form. It is worthy of notice that the halogen acids (hydrochloric acid, etc.), which contain only one hydrogen atom in the molecule, ionize very completely in moderately dilute solutions, while hydrogen sulfide, with its two hydrogen atoms, ionizes to only a slight extent. If oxygen is added, however, as in sulfuric acid, H^SCX, the ionization be- comes large, though it does not equal that of hydrochloric acid. Application of the Idea of Strength of Acids to explain the STRENGTH OF ACIDS 169 Conduct of Sulfides. 1 Practically all of the ordinary reactions in aqueous solutions are reversible. A reversible reaction leads to a stable condition only when the reaction has reached a point where it proceeds just as fast in one direction as in the other. In the reversible reactions : Pb ++ + 2 NOr + H + + HS- ^ Pb+ + 2 NO 3 ~ + H + and X SH b+ X +HS- ^ PbS + H 2 S 2 SH the reactions will proceed toward the right as long as lead sul- fide (PbS) is formed and separates from the solution. The equilibrium finally reached must depend upon whether there are enough lead ions (Pb ++ ), lead hydrosulfide (Pb+ ) ions and X SH hydrosulfide ions (HS~) in a given volume of the solution to form more lead sulfide than can remain in solution. As lead sulfide is very insoluble, only a very few hydrosulfide ions can remain in a solution containing lead ions. The ionization of hydrogen sulfide : takes place to a very limited extent even in pure water, the equilibrium in the reaction of ionization being very far to the left. If we add to a solution of hydrogen sulfide a strong acid, as hydrochloric acid, which gives a large number of hydrogen ions, the hydrosulfide ion (HS~) will meet hydrogen ions more frequently than before and will combine with them to form hydrogen sulfide. This must shift the equilibrium to the left and cause an increase in the unionized hydrogen sulfide and a 1 The student should read this paragraph, but it may be well to leave its careful study till review or a later period. (See pp. 379-386.) 2 This may involve the further ionization, Pb+ ^ Pb+ + H+ X SH \S- but the ion Pb+ if capable of existence at all, would immediately become PbS. Most authors are accustomed to write the reaction : Pb++ + 2 NOr + 2 H+ + S = PbS + 2 NO 8 ~ + 2 H+, but the form given above seems more probable. 170 A TEXTBOOK OF CHEMISTRY decrease in the number of hydrosulfide ions 1 (HS~). The presence of a moderate amount of hydrochloric acid in a solu- tion containing a lead salt will, therefore, so far decrease the con- centration of the hydrosulfide ions that lead sulfide can no longer be precipitated. It will be seen from what has just been said that the distinction between the first and second classes of metals (p. 166) in qualita- tive analysis depends on our definition of a <k slightly acid " solution. Such metals as lead, cadmium and zinc might belong to the first class or the second according to the concentration of the hydrogen ions present. The addition of an alkali, as sodium hydroxide (NaOH), to a solution of hydrogen sulfide has an effect opposite to the addition of an acid. The base gives hydroxide (OH~) ions, which combine with the hydrogen ions to form water. This displaces the equilibrium for the ionization of hydrogen sulfide in the oppo- site direction and results in a large increase in the number of hydrosulfide ions (HS~~). Under these conditions the sulfides of iron, zinc and some other metals, which are too soluble to form at all in acid solutions, will form and be precipitated. When hydrogen sulfide is passed into a solution containing a hydroxide of a metal of the third class, a hydrosulfide, which ionizes to a large extent in dilute solutions, is formed : Na + + OH- + H + + HS- = Na + + HS~ + H 2 O If a second, equal amount of sodium hydroxide is added and the solution is evaporated to dryness, sodium sulfide, Na 2 S, may be obtained : Na + + HS- + Na + + OH~ ^ Na 2 S + H 2 O 1 This may be stated mathematically as follows : Ce+ X CHS- X Fi = Si C H? s XFi ' = S 2 Since Si = S 2 at equilibrium and S 2 must be constant for a given quantity of hydrogen sulfide and Fi is also constant, the product of CH + X CHS~~ must be constant for any given concentration of hydrogen sulfide. Any increase in the number of hydrogen ions must, therefore, be accompanied by a corresponding decrease in the number of hydrosulfide ions. {REDUCTION BY HYDROGEN SULFIDE 171 If the sodium sulfide is dissolved in water, the ionization of water approaches so near in degree to that of hydrogen sulfide that the sodium sulfide is largely hydrolyzed : Na 2 S + H + + OH- = 2 Na + + HS~ + OH~ The presence of hydroxide ions in such a solution is indicated by the alkaline reaction of the solution, as shown by litmus or other test papers. Hydrogen Sulfide as a Reducing Agent. Hydrogen sulfide readily gives up its hydrogen to chlorine, bromine or iodine. It also gives up hydrogen to a great variety of compounds, re- ducing them. The following are typical illustrations : H 2 S + I 2 = 2 HI + S This reaction furnishes an excellent method of preparing a solu- tion of hydriodic acid, by suspending iodine in water and passing hydrogen sulfide into the mixture. K 2 Cr 2 7 + 8 HC1 + 3 H 2 S - 2 KC1 + 2 CrCl 3 + 7 H 2 O + 3 S Potassium Chromic Dichromate Chloride In writing the equation for this reaction, notice that the formulas of the chlorides determine the number of molecules of hydro- chloric acid required. Comparing the number of molecules of hydrochloric acid with the number of atoms of oxygen in the dichromate it is seen that after water has been formed from the hydrogen of the hydrochloric acid three atoms of oxygen remain. These will oxidize the hydrogen of three molecules of hydrogen sulfide. Fe 2 (S0 4 ) 3 + H 2 S = 2 FeSO 4 + H 2 SO 4 + S Ferric Ferrous . Sulfate Sulfate Here the hydrogen of the hydrogen sulfide takes the sulfate radical (SO 4 ) from the ferric sulfate, and the iron is reduced from the ferric to the ferrous state. Another method of writing such equations, which is preferred by some teachers, is based on the principle of positive and nega- tive valences. According to this principle : 172 A TEXTBOOK OF CHEMISTRY 1. The algebraic sum of the valences of any compound is zero. The valence of a free element is also zero. 2. Oxygen in compounds has a negative valence of 2. 3. Hydrogen in compounds has a positive valence of 1. 4. When the valence of one element changes, the valence of some other element or elements must change by the same amount in the opposite direction. In applying these principles to the present case, the equation is first written in the following form : K 2 Cr 2 O 7 + HC1 + H 2 S -+ 2KC1 + 2 CrCl 3 + S + H 2 O On inspection it is seen that the sum of the valences of the two potassium and two chromium atoms on the left is + 14, while on the right the sum of the valences of the same four atoms is only + 8, a loss of 6 positive valences. To balance this the valence of the sulfur atom changes from 2 in hydrogen sul- fide, H 2 S, to in free sulfur, S. It is obvious, at once, that to balance the changes in the chromium we must have three molecules of hydrogen sulfide. To furnish the chlorine for the chlorides there must be 8 molecules of hydrochloric acid, HC1. The equation becomes, therefore : K 2 Cr 2 O 7 + 8 HC1 + 3 H 2 S = 2 KC1 + 2 CrCl 3 + 7 H 2 O+3 S Similar reactions occur between hydrogen sulfide and chlorine, hydrogen sulfide with sulfuric acid and potassium permanganate, KMnO4, or hydrogen sulfide and ferric chloride, FeCls. The reaction between lead sulfide, PbS, and nitric acid, HNO 3 , giving lead nitrate, Pb(NO 3 ) 2 , nitric oxide, NO, sulfur and water, is also closely related to these, the positive, bivalent lead atom taking the place of the two positive hydrogen atoms in hydrogen sulfide. The student is advised to write the equa- tions for these reactions by use of both of the methods suggested above. Sulfur Dioxide is formed when sulfur is burned in the air, also when iron pyrites, FeS 2 , is burned, the latter method of preparation being used largely in the manufacture of sulfuric acid. SULFUR DIOXIDE 173 Sulfur dioxide may be prepared in the laboratory by the reduc- tion of concentrated sulfuric acid with copper or other sub- stances. If copper is used, copper sulfate is formed. The equation may be written as follows : Cu + H 2 SO 4 = [CuO] + SO 2 + H 2 O [CuO] + H 2 SO 4 = CuSO 4 + H 2 Combining, Cu + 2 H 2 SO 4 = CuSO 4 + SO 2 + 2 H 2 O The first two equations are written as an aid to the writing of the last. The [CuO] is placed in brackets to indicate that it is not a final product of the reaction. It may or may not be formed as an intermediate product. Another method of writing would be to represent hydrogen [2 H] as an intermediate product. The most convenient laboratory method for the preparation of sulfur dioxide is to drop concentrated sulfuric acid into a 40 per cent solution of acid sodium sulfite, NaHSOa : NaHSO 3 + H 2 SO 4 ^ NaHSO 4 + H 2 SO 3 Sulfurous Acid Sulfurous acid, H 2 SOs, decomposes very easily into sulfur dioxide and water, the sulfur dioxide escaping as a gas. This has the same effect on the equilibrium of the first reaction as if the sulfurous acid itself were volatile and escaped from the mix- ture. Sulfur dioxide is a colorless gas with a suffocating odor, famil- iar in the burning of sulfur-tipped matches. It may be con- densed to a liquid in a tube surrounded with" a freezing mixture and boils at 10. It freezes at a very low temperature and melts at - 73. Sulfur dioxide is used to bleach straw, wool and silk. The latter, especially, are injured by the action of chlorine, so that it cannot be used. The sulfur dioxide seems to combine with the coloring matter to form colorless compounds, or, in some cases, to reduce the colored compound to a colorless one. Exposure 174 A TEXTBOOK OF CHEMISTRY to the air and light frequently restores the color, as in the case of straw hats. Sulfur dioxide and sulfites are powerful germicides. Its use as a disinfectant, however, has been almost entirely replaced by formaldehyde, which is even more effective and does not injure fabrics or metallic articles, as sulfur dioxide does. The injury to fabrics may be either through its bleaching effect or because it is slowly oxidized by the action of air and moisture to sul- furic acid, which is corrosive. Sulfur dioxide is still used exten- sively in " sulfuring " fruit to destroy the organisms which cause darkening and injury during the drying. Sulfurous Acid. At ordinary temperatures water dissolves about 50 times its volume of sulfur dioxide. The solution red- dens litmus and neutralizes bases, showing that the sulfur dioxide combines with the water and forms an acid. The sodium salt obtained by neutralizing the acid is sodium sulfite, Na 2 SO3, and from the formula of this and other salts it is assumed that the formula of the sulfurous acid in such a solution is H 2 SOa. The structure of the acid is probably (k /H CT \)H It is very unstable, one hydroxyl group and one hydrogen atom separating very easily from the molecule. The solution smells strongly of sulfur dioxide, and all of the gas can be expelled by boiling the solution. Sulfurous acid is a comparatively weak acid. In a solution containing 0.05 gram molecule (-$ mol) about 20 per cent of its hydrogen is ionized, 1 while in a corresponding solution of sulfuric acid 60 per cent of the hydrogen is ionized. Sulfurous acid is a powerful reducing agent. It is oxidized to sulfuric acid 1 This is on the supposition that all of the sulfur dioxide has combined with water to form sulfurous acid. It is probable that some of the sulfur dioxide exists as such in the solution and that the ionization of the sulfurous acid really present is considerably greater than appears from these figures. SULFUR TRIOXIDE 175 by potassium permanganate, KMnO 4 , potassium dichromate, K 2 Cr 2 C>7, chlorine, bromine or ferric salts. Sulfites. Sulfurous acid forms both acid and normal salts, the salts of sodium being acid sodium sulfite, NaHSO 3 , and nor- mal sodium sulfite, Na 2 SO 3 . The calcium salts are CaH 2 (SO 3 ) 2 and CaSO 3 . These salts are prepared, commercially, by burn- ing sulfur in air and passing the mixture of sulfur dioxide and nitrogen through a solution of sodium carbonate or sodium hydroxide for the sodium salts, or through milk of lime (Ca(OH) 2 ) for the calcium salts. The acid sodium salt has been used as an addition to wine or cider to stop fermentation. The acid cal- cium salt is used in the purification of wood pulp for the manu- facture of paper. Sulfur Trioxide. Some heat is evolved when sulfur dioxide combines with oxygen to form the trioxide, SO 3 , but the speed of the reaction between the two is too slow to be measured at ordinary temperatures. At temperatures where the speed of the reaction of combination becomes sufficiently rapid to become a practicable method of preparation, the dissociation of sulfur trioxide into sulfur dioxide and oxygen becomes very large and renders this method of preparation from the substances alone impracticable. The reversible reaction : 2 SO 2 + O 2 ^ 2 SO 3 has its point of equilibrium shifted toward the left as the tem- perature rises, in accordance with the principle of van't Hoff- LeChatelier (p. 111). As early as 1831 it was discovered that the reaction is greatly accelerated by the presence of platinum, but it was nearly 70 years before the details for the application of this principle were so far worked out as to render the manufacture on a large scale possible, 1 so long does it often take to convert a scientific discovery into commercial success. The chief diffi- culties to be overcome were, first, that arsenic and other sub- 1 For very interesting historical details see Knietsch, Ber. 34, 4069 (1901). 176 A TEXTBOOK OF CHEMISTRY stances in the gases obtained by roasting pyrites " poison " the platinum and render it ineffective for the catalysis, and second, that the platinum catalyzes the dissociation of sulfur trioxide as well as the combination of sulfur dioxide and oxygen, and the temperature for rapid combination lies very close to a temperature at which the dissociation is large and so the com- bination becomes incomplete. These difficulties have been overcome by a careful purification of the sulfur dioxide as it comes from the pyrites burners and by a careful regulation of the temperature as the gases pass over the " contact mass." The platinum of the " contact mass" is disseminated in a very finely divided condition over asbestos or some other material which gives it a large surface in proportion to its weight. In the laboratory, on a small scale, sulfur trioxide can be read- ily prepared by passing dry sulfur dioxide and oxygen through a gently warmed tube containing platinized asbestos (Fig. 53). Fig. 53 It may be obtained still more easily by warming " fuming " sulfuric acid, which is a mixture of sulfuric acid, H 2 SO4, pyro- sulfuric acid, H 2 S 2 O 7 , and sulfur trioxide. Sulfur trioxide is a clear, volatile liquid which solidifies at a low temperature. It melts at 14.8 and boils at 46. In the presence of a trace of moisture a little sulfuric acid, H 2 SO4, or pyrosulfuric acid, H 2 S 2 O 7 , is formed and this acts as a catalytic agent causing sulfur trioxide to polymerize, forming the com- pound S 2 Oe, which crystallizes in white, asbestos-like needles. As it is extremely difficult to exclude moisture completely, this polymeric form is usually obtained instead of the true trioxide. On warming it gives a vapor, which consists of the true trioxide. SULFURIC ACID 177 Sulfur trioxide hisses like a hot iron when thrown into water, owing to the heat generated when it combines with water to form sulfuric acid. It fumes strongly in the air, forming minute drops of sulfuric acid, which settle only very slowly and are not readily absorbed by water. Curiously enough these minute drops are easily absorbed by concentrated sulfuric acid and this is used for the purpose in the manufacture of sulfuric acid by the contact process. Sulfuric Acid. The contact process for the preparation of sulfur trioxide has, thus far, been used almost exclusively for the manufacture of a very concentrated or a " fuming " sulfuric acid. It has been pointed out that the direct combination of sulfur dioxide and oxygen is too slow to be commercially possible as a method of manufacture, and that platinum is used to catalyze, or hasten, the reaction. Another catalytic agent, not so sensitive to impurities in the gases, or to temperature changes, and which acts rapidly at ordinary temperatures, has been used for a long time in what is called the " chamber process " for the manufac- ture of sulfuric acid. In this process large chambers lined with sheet lead, which is only slightly attacked by dilute sulfuric acid, are employed. Into these chambers are introduced : 1 . Sulfur dioxide from burning sulfur or iron pyrites : 2 FeS 2 +11O = Fe 2 O 3 + 4 SO 2 2. Nitric acid from Chile saltpeter, NaNO 3 , and sulfuric acid : NaNO 3 + H 2 SO 4 = NaHSO 4 + HNQ 3 3. Air, to furnish oxygen. 4. Water as steam or spray. The first reaction consists in the oxidation of the sulfur dioxide to sulfuric acid by the nitric acid : 3 SO 2 + 2 HNO 3 + 2 H 2 O = 3 H 2 SO 4 + 2 NO 1 1 This equation should not be learned by rote, but should be written on the following considerations : 1. When nitric acid is reduced to nitric oxide, two molecules give 3 atoms of available oxygen. 2. Each atom of oxygen will oxidize one molecule of sulfur dioxide. 178 A TEXTBOOK OF CHEMISTRY If it were necessary to stop here and the nitric oxide were lost, sulfuric acid would be, comparatively, an expensive substance on account of the limited supply and relatively high price of sodium nitrate. But nitric oxide combines almost instantly with the oxygen of the air to form nitrogen dioxide, NC>2 : 2 NO + O 2 = 2 NO 2 Nitrogen dioxide, in turn, can oxidize a new quantity of sulfur dioxide : S0 2 + NO 2 + H 2 O = H 2 SO 4 + NO It is pretty certain that the mechanism of the reaction is more complicated than is indicated by these equations, but the equa- tions given indicate clearly the fundamental facts on which the action depends. These are : first, that nitric oxide combines, practically instantaneously, with oxygen ; second that nitrogen dioxide can, directly or indirectly and very quickly, give its oxygen to the sulfur dioxide and water, converting these to sulfuric acid. Commercially, the whole process depends on the speed with which these actions occur. The theory of the lead chamber process which has received most acceptance is that of Lunge, who supposes the process to consist in the formation and decomposition of nitrosyl sulfuric add: NO OH a mixture of nitric oxide, NO, and nitrogen dioxide, NO2 (equivalent to nitrous anhydride, N 2 Os), being the effective 2 SO 2 + N 2 O 3 + O 2 + H 2 O = 2 S0 2 < X OH X) NO /OH 2 SO 2 < + H 2 O = 2 S0 2 < + N 2 O 3 X OH X OH Nitrosyl sulfuric acid is a definite, crystalline compound, which is formed in the chambers when the supply of water is SULFURIC ACID 179 insufficient, but it exists only as an intermediate product, if at all, in the normal manufacture. Nitrous anhydride, also, can exist only momentarily, if at all, as it decomposes at once into nitric oxide and nitrogen dioxide at the temperature of the chamber. For a further discussion of the subject see Trautz, Z. physik. Chem. 47, 513 ; Wentzki, Z. angew. Chem. 23, 1907 ; Raschig, ibid. 23. 2241, 24, 160; Ber. ibid. 23, 2250. If it were possible to lead into the chamber pure oxygen and sulfur dioxide, a small amount of nitric acid would convert an indefinitely large amount of sulfur dioxide into sulfuric acid. Since, however, air containing only 21 percent of oxygen mixed with 79 per cent of nitrogen (and argon) must be used, there must Glover Pyrites Tower Burners Leaden Chambers Fig. 54 G ay -L us sac Tower be a constant escape of nitrogen, carrying with it nitric oxide or nitrogen dioxide at the further end of the chamber or set of chambers. To recover these the gases are passed through a tower, known as the Gay-Lussac tower, Fig. 54, in which they are exposed to a large surface of concentrated sulfuric acid running down over broken coke or a series of earthenware plates. The strong acid absorbs the oxides of nitrogen, forming nitrosyl sulfuric acid, SO 2 (OH) (ONO). This nitrated acid is then forced by compressed air to the top of another tower, called the Glover tower, B. Here it is mixed with some of the more dilute 180 A TEXTBOOK OF CHEMISTRY acid from the chamber and a little nitric acid to replace the un- avoidable loss. The mixture runs down over broken coke and comes in contact with sulfur dioxide coming from the pyrites burners, C. This causes the denitrification of the acid : 2 SO 2 (OH)(ONO) + SO 2 + 2 H 2 O = 3 H 2 SO 4 + 2 NO The nitric oxide is, of course, carried back into the first chamber. When these towers are used, only from 25 to 40 pounds of sodium nitrate are required for the manufacture of a ton of sulfuric acid. Without them, two or three times as much is required. , The acid from the chambers has a specific gravity of 1.53 to 1.62, and contains 62-70 per cent of the pure acid. It is usually concentrated to about 79 per cent by evaporation in lead pans. At this point the acid begins to attack the lead more strongly and the concentration is completed to a specific gravity of 1.83-1.84 and 93 to 95 per cent, in glass, platinum or iron, the last metal being only slightly attacked by the concentrated acid, although it dissolves easily in the dilute acid. If the concen- trated acid is distilled, an acid of constant composition contain- ing about 98.5 per cent of the pure acid finally passes over at 338. The density of the vapor proves that the process is not ordinary boiling, but consists in the dissociation of sulfuric acid to sulfur trioxide and water and that the two recombine on cooling: H 2 S0 4 :S0 3 + H 2 The specific gravity of pure, 100 per cent sulfuric acid is slightly less than that of a 96 to 99 per cent acid, the difference being so small that the concentration of the acid cannot be de- termined satisfactorily by means of the density. When sulfuric acid is mixed with water, considerable heat is evolved, and the volume of the diluted acid is considerably less than the sum of the volumes of the acid and water which are mixed. There is a chemical combination between the acid and water, giving a compound which probably contains four or six hydroxyl (OH) groups : ELECTRON THEORY 181 /OH , H /OH H H VOH X OH It is noticeable that while sulfurous acid, S\ , loses O v OH of X OH water- easily, sulfuric acid, ^S v , dissociates at a much O^ X OH higher temperature and also has a strong tendency to take up more water. Along with this strong attraction of the sulfur atom for hydroxyl groups, which seems to be so closely con- nected with the addition of another oxygen atom, is the fact that sulfuric acid is a much stronger acid than sulfurous acid. Thus in a " tenth normal " l solution of sulfuric acid about 60 per cent of the hydrogen is ionized, while in a tenth normal solu- tion of sulphurous acid, H 2 SO 3 , only about 20 per cent is ionized. The Electron Theory. The facts which have just been given may be explained, in part, by the electron theory, which has been developed rapidly during the last few years. The electron 2 may be defined as an atom of negative electricity. When by itself and in rapid motion its mass is approximately one seventeen- hundredth of the mass of a hydrogen atom. It is supposed that atoms of the elements are composed, in part, of electrons and that they may either gain or lose these. If an atom gains an electron, it becomes negatively charged ; while if it loses one, it becomes positively charged. In hydrogen sulfide, H2S, it is supposed that each hydrogen atom has lost an electron which has been trans- 1 A solution containing one tenth of a gram atom of replaceable hydrogen or one twentieth of a gram molecule of sulfuric acid in one liter. 2 Professor J. J. Thompson uses the name "corpuscle" instead of electron. The evidence of the existence of electrons is very positive. 182 A TEXTBOOK OF CHEMISTRY ferred to the sulfur. The positive hydrogen atoms are then held by the negative sulfur atom. In sulfur dioxide and sulfur triox- ide, however, the sulfur atom is supposed to lose two electrons to each oxygen atom, and acquires either four or six positive charges. When water is brought in contact with sulfur dioxide, O = S = O, or sulfur trioxide, O = S^ , it separates into a pos- \) itive hydrogen atom and a negative hydroxyl group. The former adds itself to the negative oxygen or sulfur, while the latter unites H with the positive sulfur. This gives x+S+^T and H + -~cr +Nr For the reverse reaction to occur both hydrogen and hydroxyl must separate from the compound, and it seems probable that the negative hydroxyl groups will be held much more strongly by the sulfur atom with six positive charges than by the one which has only the equivalent of four. On the other hand, the positive hydrogen might be expected to separate more easily in the ionic form from the sulfuric acid on account of the indirect repulsion of the strongly positive sulfur atom. The electron theory is too recent for chemists to form a very positive opinion as to its value, but it is, at least, worthy of careful consid- eration, and it will be referred to repeatedly in the following pages. Sulfuric Acid as a Dehydrating Agent. On account of its affinity for water, sulfuric acid is an excellent drying agent for all gases which do not react with it chemically. It is much more efficient than calcium chloride. It will also take the elements of water from many such substances as wood or sugar, which contain oxygen and hydrogen. So much heat is liberated when sulfuric acid is mixed with water that the action may become explosive unless care is used. The concentrated acid should always be poured into water with which it is to be diluted instead of pouring water upon the acid. Why? SULFATES. NORMAL SOLUTIONS 183 Sulfates. Dibasic Acids. Either one or both of the hydrogen atoms of sulfuric acid may be replaced by a metal, giving acid and normal salts, as acid sodium sulfate, NaHSO 4 , and normal sodium sulfate, Na2SO 4 . Acids having this property are called dibasic. An acid like phosphoric acid, H 3 PO4, which forms three salts with sodium, NaH 2 PO 4 , Na 2 HPO 4 and Na 3 PO 4 , is called tribasic. The basicity depends, however, not on the number of hydrogen atoms in one molecule of an acid, but on the number of replaceable hydrogen atoms. Thus acetic acid, C 2 H 4 O 2 , is monobasic because only one of /ts^i^tirogen atoms can be replaced ; and phosphorous acid, HsPOa^pjdibasic because only two of the hydrogen atoms can be replaced. As with other strong acids, the normal sulfates of the metals of the sodium and calcium families are neutral in reaction, while the acid sulfates of all metals are strongly acid. The sul- fates of the metals of the calcium family, calcium, strontium, barium and radium, are difficultly soluble in water, the solubility decreasing with increasing atomic weight. Barium sulfate re- quires about 400,000 parts of water for its solution, while radium sulfate is still more insoluble. Lead sulfate, also, is almost insoluble, but all other sulfates which are not decomposed by water are soluble. In general, the salts of strong acids are solu- ble in water, and this fact is probably connected with the high degree of ionization of both acids and salts. No explanation has been offered for the exceptions to this general rule. The rule is useful because we have to learn only a short list of insolu- ble salts for these strong acids and can then assume that all other salts are soluble. Normal, Standard and Formular l Solutions. We have fre- quently found it convenient to use the gram molecule of sub- stances as a unit in dealing with them. This unit is often called, for the sake of brevity, one mol. In working with acids and 1 The designation "molar" (or molal) is often used, but "for- mular," if followed by the formula of the substance, is more definite. Thus a formular solution of ferric chloride, FeCl 3 , is definite, while a molar solution of ferric chloride might refer to either FeCl 3 or Fe 2 Cl 6 . 184 A TEXTBOOK OF CHEMISTRY bases it is often convenient to use a gram equivalent instead of a gram molecule, as the unit. The gram equivalent of an acid or base is that quantity which is equivalent to or will neutralize one gram molecule of a monobasic acid or of a monacid base. A solution containing one gram equivalent of an acid or base in one liter (or one milligram equivalent in one cubic centimeter) is said to be normal. Thus a normal solution of hydrochloric acid would contain 36.47 grams of the acid, HC1, in one liter ; but a normal solution of sulfuric acid would contain, not a gram molecule (98.08 grams), but a gram equivalent (49.04 grams) of sulfuric acid, H 2 SO4, in one liter. A normal solution of sodium hydroxide, NaOH, would contain 40.01 grams in one liter ; but a normal solution of calcium hydroxide, Ca(OH) 2 , if it could be prepared, would contain only 37.04 grams per liter. The advan- tage of the system is that one cubic centimeter of any normal solution will exactly neutralize or be exactly equivalent to one milligram equivalent of any acid or base. The name " normal " is also frequently applied to solutions of salts and ofcother sub- stances, but such a use is liable to lead to conf usi^B and it is better to call such solutions " standard " or^pKnular," a standard solution being simply one whose-^roncentration is known and a formular solution one which contains one formular weight in one liter. The formula on which a formular solution is based should always be given. If the term " normal " is applied to solutions of other sub- stances than acids and bases, one liter of the solution should always contain an amount of the substance which is equivalent to one gram atom of hydrogen in the reaction for which it is used. A normal solution of potassium chloride, KC1, or of silver nitrate, AgNOs, will contain one gram molecule of these compounds in a liter, but a solution of calcium chloride, CaCl2, will contain only one half of a gram molecule. A solution of potassium permanganate, KMnC>4, will contain only one fifth of a gram molecule, if to be used in an acid solution, and one third of a gram molecule, if to be used in an alkaline solution, because one gram molecule of the compound will oxidize five gram atoms ACIDIMETRY 185 of hydrogen in the first case and only three in the second. A normal solution of ferrous sulfate, FeSO 4 , will contain one gram molecule of the compound in a liter, because it requires only one half of a gram atom of oxygen to oxidize it. Acidimetry and Alkalimetry. If a very small quantity of an indicator (p. 122) is added to a solution of hydrochloric acid or any other strong acid, on adding a solution of sodium hydroxide or some other strong base, the change in color of the indicator will show very sharply when the acid has been ex- actly neutralized by the base. If we have a normal solution of hydrochloric acid (containing 36.47 milligrams of the acid, HC1, in 1 cc., as defined above), it is easy by measuring this from a burette (Fig. 55) exactly to neutralize a solution containing^ strong base. The num- ber of cubi(centimenters of the acid used will give at once the numbers of milligram equivalents of the base which were present in the solution neutralized. Thus one cubic centi- r D : . f j < 1 I* r Fig. 55 meter of normal hydrochloric acid will exactly neutralize 40.01 milligrams of sodium hydroxide, NaOH, 56.11 milligrams of potassium hydroxide, KOH, or 37.04 milligrams of calcium hydroxide, Ca(OH) 2 . In the same manner, by means of a normal solution of potassium hydroxide, KOH (containing 56.11 milligrams in 1 cc.), the number of milligram equivalents of any strong acid contained in a given solution can be readily determined. The process of making such determinations is called acidimetry or alkalimetry. The choice of an indicator and th*e application of the process to some cases involving weak acids and bases will be discussed in a later chapter (p. 387). 186 A TEXTBOOK OF CHEMISTRY Pyrosulfates. When acid sodium sulfate is heated, it loses water and is converted into a salt which is called, for this reason, sodium pyrosulfate : 2 NaHSO 4 = Na 2 S 2 7 + H 2 O Sodium Pyrosulfate A solution of sulfur trioxide in sulfuric acid doubtless always contains pyrosulfuric acid, H 2 S 2 O 7 , and pure pyrosulfuric acid is a definite compound which melts at 35, but it is very unstable, dissociating easily into sulfur trioxide and sulfuric acid. In solution the pyrosulfates take up water and pass back into acid sulfates. Hyposulfites. When zinc is dissolved in a solution of sulfurous acid, H 2 SO 3 , the acid is reduced and zinc hyposulfite is formed : Zn + 2 H 2 SO 3 = ZnS 2 O 4 + 2 H 2 O Zinc Hyposulfite Zinc hyposulfite is a salt of an unstable acid, hyposulfurous acid, H 2 S 2 O4, which is not known in the free state. The salts are very quickly oxidized to sulfites in the air and are powerful reducing agents. Sodium hyposulfite, Na 2 S 2 O 4 , is manufactured for use in the reduction of indigo to indigo white (p. 341). Hyposulfurous acid and the hyposulfites must not be con- fused with thiosulfuric acid and thiosulfates which were formerly called by the same name (see the next paragraph). Some authors prefer to call the acid hydrosulfurous acid to avoid possible confusion. Neither acid corresponds to the formula (H 2 SO 2 ), which we should logically expect for a hyposulfurous acid. Thiosulfates. A solution of sodium sulfite, Na 2 SO 3 , will dissolve sulfur, and there may then be crystallized from the solu- tion a salt called sodium thiosulfate, Na 2 S 2 O 3 .5 H 2 O. The change is similar to the oxidation of sodium sulfite to* sodium sulfate : Na 2 SO 3 + O = Na 2 SO 4 Na 2 SO 3 + S = Na 2 S 2 O 3 PERSULFURIC ACID 187 one atom of sulfur taking the place of an atom of oxygen, and the name thiosulfatc (from Greek Odov, sulfur) is given to the salt for this reason. The salt has been long known and was originally called sodium hyposulfite, a name which still clings to it among druggists and photographers. It is extensively used in photography as a solvent for silver chloride or bromide in " fixing " pictures. If a solution of a thiosulfate is acidified, the thiosulfuric acid at first liberated decomposes with the liberation of sulfur : H 2 S 2 3 = S0 2 + H 2 + S Iodine converts sodium thiosulfate into sodium tetrathionate : 2 Na 2 S 2 O 3 + I 2 = Na 2 S 4 6 + 2 Nal Sodium Tetrathionate This reaction is much used in connection with standard iodine solutions, in volumetric analysis. Persulfuric Acid. When a solution of acid potassium sulfate, KHSO4, is electrolyzed with a high current density, that is, with a current strong in comparison with the surface, at the anode, as the anions, HSO 4 ~, are discharged they combine, in part, with other anions of the same kind to form persulfuric acid, H O SO 2 O O SO 2 OH or H 2 S 2 O 8 . The persulfuric acid then reacts with some of the acid potassium sulfate present to form potassium persulfate, which is rather difficultly soluble: 2 KHSO 4 + H 2 S 2 O 8 = K 2 S 2 8 +.2 H 2 SO 4 Persulfuric acid is also formed when hydrogen peroxide, H 2 O 2 , is added to concentrated sulfuric acid : H O SO 2 iO-H H| O O !H H Oj SO 2 OH or 2 H 2 SO 4 + H 2 O 2 = H 2 S 2 O 8 + 2 H 2 O Persulfuric acid and the persulfates are used as oxidizing agents. 188 A TEXTBOOK OF CHEMISTRY * Permonosulfuric Acid. When a solution of persulfuric acid is diluted and allowed to stand, it changes to permonosulfuric acid: H O SO 2 O O SO 2 -OH + HOH = H O S0 2 O OH + H 2 S0 4 or H 2 S 2 O 8 + H 2 O = H 2 SO 5 + H 2 SO 4 The solution was formerly known as Caro's acid and is some- times used as an oxidizing agent for organic compounds. For instance, it will oxidize aniline to nitrobenzene. * Polythionic Acids. A series of acids having from two to six atoms of sulfur in a molecule has been obtained. These are : Dithionic acid H 2 S 2 Oe Trithionic acid H^SaOe Tetrathionic acid H 2 S 4 O 6 . (See thiosulfuric acid, above.) Pentathionic acid H 2 S 5 O 6 Hexathionic acid H 2 S 6 O 6 These acids need not be considered in detail here. Compounds of Sulfur containing Halogens. Quite a num- ber of such compounds are known. All of them except sulfur hexafluoride, SF 6 , are hydrolyzed by water, giving hydrochloric acid or a halogen acid and some acid of sulfur. Thus sulfuryl chloride, SO 2 C1 2 , gives : HOH /OH QfV/ HOH + = SO 2 < + 2 HC1 X)H * Sulfur Monochloride, S 2 C1 2 , is a clear, amber-colored liquid formed by passing chlorine over heated sulfur. It boils at 138 and is hydrolyzed by water to hydrochloric acid, thiosulfuric acid and sulfur. Its specific gravity is 1.7055. Because of the strong affinities of sulfur for oxygen and of chlorine for metals, the chlorides of a number of metals, which it is difficult to prepare from the oxides otherwise, may be ob- tained by passing a mixture of sulfur monochloride and chlorine over the heated oxides. (E. F. Smith.) It is also used in the manufacture of India rubber, SELENIUM 189 * Chlorosulfonic Acid 1 is easily prepared by passing hydro- chloric acid gas through warm, fuming sulfuric acid. Ov O^ X C1 >S=0 + HC1 = lS< CT CT X)H It boils at 152-153 and is easily hydrolyzed by water. Its specific gravity is 1.766 at 18. * Sulfuryl Chloride, SO 2 C1 2 , may be prepared by the union of sulfur dioxide and chlorine, or, more easily, by boiling chloro- sulfonic acid with mercuric sulfate, which acts as a catalytic agent, causing it to decompose in accordance with the equation : 2 SO 2 OHC1 = H 2 SO 4 + SO 2 C1 2 Sulfuryl chloride is sometimes called, less correctly, the chlo- ride of sulfuric acid. Similar acid chlorides may be formed by replacing the hydroxyl of other acids with chlorine. Acid chlo- rides are hydrolyzed by water, giving the acid from which they are derived and hydrochloric acid. Sulfuryl chloride boils at 69.1 and has a specific gravity of 1. 6674 at ^- 2 . Selenium, Se, 79.2, is found in small amounts as selenides of metals, associated, usually, with sulfides of these same metals. When such sulfides are used for the manufacture of sulfuric acid, selenium is sometimes found in the dust flues of the pyrites burners and in the slime on the bottom of lead chambers. Both selenium and tellurium are found in considerable quantities in the slimes from electrolytic copper refining. It occurs in several allotropic forms, the red variety obtained by crystallization from carbon bisulfide and a gray metallic form obtained by melting either of the other forms being the best defined. An amorphous form is also known. The metallic form melts at 217 and has a specific gravity of 4.8. Selenium boils at 680. 1 Acids containing the group S0 2 OH are called sulfonic acids. 2 This means the specific gravity at 20 referred to water at 4. 190 A TEXTBOOK OF CHEMISTRY The metallic form of selenium conducts electricity. Its con- ductivity is very greatly affected by changes of temperature or by exposure to light, and several important applications of this property have been invented, one of the most important being in stellar photometry. Hydrogen Selenide, H 2 Se, may be prepared by the action of hydrochloric acid on ferrous selenide, FeSe. It is very poison- ous, and the odor is more unpleasant than that of hydrogen sul- fide. Compare the series, water, hydrogen sulfide, hydrogen selenide, in this respect. Berzelius, one of the early workers with hydrogen selenide, reports that after breathing a single bubble of the gas he so far lost the sense of smell for several hours that he could not distinguish the odor of strong ammonia. Selenium Dioxide, SeO 2 , is a white solid prepared by burning selenium in a current of oxygen. It gives selenious acid, H 2 SeO 3 , on solution in hot water. From a solution of selenious acid sulfur dioxide precipitates selenium as a red powder : H 2 SeO 3 + 2 SO 2 + H 2 O = Se + 2 H 2 SO 4 Selenic acid, H 2 SeO 4 , is formed by the action of bromine on silver selenite, the silver bromide formed separating as a precipi- tate * Ag 2 SeO 3 + Br 2 + H 2 O = H 2 SeO 4 + 2 AgBr Selenic acid loses oxygen easily and is a strong oxidizing agent. Tellurium, Te, 127.5, is found in combination with gold, silver, copper and bismuth. It is a white, metallic-looking solid, which melts at 452, boils at 1400 and has a specific gravity of 6.44. Its most interesting compounds are hydrogen telluride, H 2 Te, tellurium dioxide, TeO 2 , tellurium trioxide, TeO 3 , tellurous acid, H 2 TeO 3 , and telluric acid, H 2 TeO 4 . Atomic Weight of Tellurium. Very many determinations of the atomic weight of tellurium have given values about 127.5, decidedly higher than the atomic weight of iodine, 126.9. The properties of tellurium and especially the formulas of its com- pounds indicate that it should precede iodine in the Periodic System, and this has led to many attempts to determine whether TELLURIUM. GROUP VI 191 the material used for the atomic weight determinations has been pure, or whether, possibly, it may have contained some other element from which it is unusually difficult to separate a pure tellurium compound. Some of these attempts to discover a method of preparing tellurium of greater purity and lower atomic weight have seemed, for a time, to be successful ; but none of these lower results for the atomic weight has been confirmed by other workers, and it seems pretty certain that the atomic weight of tellurium is greater than that of iodine. See Browning and Flint, Z. anorg. Chem. 64, 104, 112, and 68, 251 ; and Harcourt and Baker, J. Chem. Soc. 100, 1311. General Properties of the Elements of the Sixth Group. Just as chlorine, bromine and iodine are much more closely related in their properties than fluorine is related to them, oxygen stands somewhat by itself in the sixth group, while the relationships between sulfur, selenium and tellurium are comparatively close. The halogens have a valence of one in their compounds with hydrogen, as in HF, HC1, etc., and a maximum valence of seven in their compounds with oxygen, as in C^O?, HCIO^H O ClOa), HIO4, etc. The elements of the sulfur family have a valence of two in their compounds with hydrogen, as in H 2 O, H 2 S, etc., and a maximum valence of six toward oxygen, as in SOs, [2804, H-CX ,0 >SC ,H 2 TeO 4 , etc. H- (X ^O There is a similar gradation of physical properties in the two groups : fluorine and chlorine are gases, bromine a liquid and iodine a solid, with increasing depth of color as the atomic weight increases. In the same way, oxygen is colorless (ozone is blue), sulfur is a light yellow solid, selenium is dark red, and tellurium is opaque and has many of the properties of a metal. Indeed, if it were not for its position in the Periodic System and the re- semblance between the formulas of its compounds and those of selenium and sulfur, tellurium would be classed as a metal, or, at least, as a half metal. But it will be seen that in the succes- sive groups the metallic properties become more and more 192 A TEXTBOOK OF CHEMISTRY marked with increasing atomic weight. Thus arsenic, antimony and bismuth, of the fifth group, are usually classed as metals, though all of them are brittle. Tin and lead, of the fourth group, are clearly metals and are malleable, though deficient in tenacity. The most typical compounds of the sixth group are the follow- mg: H 2 O H 2 S H 2 Se H 2 Te O 3 SO 2 SeO 2 TeO 2 SO 3 TeO 3 H 2 S0 3 H 2 Se0 3 H 2 Te0 3 H 2 SO 4 H 2 SeO 4 H 2 TeO 4 As in the halogen family, the chemical activity, in general, decreases with increasing atomic weight. Hydrogen sulfide dissociates at a much lower temperature than water does, and sulfur dioxide will take oxygen from selenious acid, reducing it to free selenium. As manganese forms compounds which resemble some of the compounds of chlorine, there are four metals of the sixth group, chromium, molybdenum, tungsten (symbol W, from wolfram) and uranium, which form oxides and salts of acids similar to the oxides and acids of sulfur. The oxides are : CrO 3 , MoO 3 , WO 3 , UO 3 ; and the corresponding sodium salts of the acids are : Na 2 CrO4, Na 2 MoO 4 , Na 2 WO 4 . Uranium forms a compound, UO 2 (OH) 2 , similar in composition to sulfuric acid, but it is a base rather than an acid, another illustration of the fact that an increase in the atomic weight, within a given group, increases the metallic properties of the element. Crystals. When substances solidify from the molten condi- tion or when they separate on the evaporation of a solution, molecules of the same kind frequently arrange themselves in definite, geometrical relations to each other, forming solids bounded by plane faces, which are called crystals. This prop- erty has already been mentioned as an important means for preparing pure substances. It is also a very important and char- acteristic property of individual substances, and the shapes of the crystals of different compounds offer such an infinite variety CRYSTALS 193 that they may frequently be used as a very positive means of identification. In spite of the large number of crystalline forms all crystals may be classified in six systems. These systems are most easily defined by referring each to axes, which are used in much the same manner as the coordinates of analytical geometry, to de- fine the structure of the crystal and the relation between the planes bounding its surface. These systems are : 1. The Isometric or Regular System. Three equal axes at right angles. Some of the simplest forms of this system are the Fig. 56 Fig. 57 Fig. 58 cube (Fig. 56), octahedron (Fig. 57), rhombic dodecahedron (Fig. 58) and the tetrahedron (Fig. 59). The last has only half of the Fig. 59 Fig. 60 faces of the octahedron and is called a hemihedral form. Com- binations of two or more forms are also common. Figure 60 is a combination form called a tetrahexahedron. 194 A TEXTBOOK OF CHEMISTRY 2. The Tetragonal System. Three axes at right angles, two of them, only, being equal. The tetragonal pyramid (Fig. 61) and the square prism (Fig. 62) are common forms. Figure 63 shows a combination of the two. Fig. 61 Fig. 62 3. The Rhombic System. Three axes at right angles but of unequal length. Rectangular and rhombic prisms (Figs. 64 and 65) and pyramids are illustrations. '1 ^ ~1 1 f^L 1 m 'r. Fig. 64 Fig. 65 4. The Hexagonal System. Three axes in the same plane at an angle of 60 with each other and a fourth at right angles to CRYSTALS 195 the plane of the other three. Common forms are the hexagonal pyramid (Fig. 66) and prism (Fig. 67) and the rhombic hexahe- dron (Fig. 68), which has only half of the faces of the pyramid. Fig. 66 Fig. 67 5. The Monoclinic System. Two axes at right angles and a third at right angles to one and inclined to the other, the three axes being unequal (Fig. 69). Fig. 68 196 A TEXTBOOK OF CHEMISTRY Fig. 70 6. The Triclinic System. Three unequal axes, all inclined (Fig. 70). Crystals rarely exhibit the complete geometrical forms which are the ideal to which they are referred. They always have an internal structure characteris- tic of these forms, however, and this can often be detected by their optical properties. Thus for the regular system light travels with the same velocity in all directions through the crystal. For other, less symmetrical forms, the velocity is different in different directions and this causes double refraction, polarization of light and other phe- nomena frequently used for the identification of the crystal form. The angles between the faces of crystals are also accurately fixed by the system to which they belong and the properties of the individual substance, and the measurement of these angles is used for purposes of identification. Such substances as sulfur, which crystallize in two different forms, are called dimorphous. Different substances which crystallize in the same form are called isomorphous. A crystal of a substance should grow if placed in a supersaturated solution of a substance with which it is isomorphous. EXERCISES 1. One liter of water at absorbs 4.37 volumes of hydrogen sul- fide. What part of a gram molecule of the gas does the solution con- tain ? What would be the theoretical depression of the freezing point if the compound were completely ionized to HS~ and H + ? If it were not at all ionized ? 2. How many liters of air containing 21 per cent of oxygen will be required to burn one liter of hydrogen sulfide to water and sulfur ? How many liters would be required to burn it to sulfur dioxide and water ? SULFUR, SELENIUM AND TELLURIUM 197 3. How many cubic feet of air at (1 cu. ft. = 28.315 liters) will be required to burn enough pyrites to make one ton of chamber acid of 70 per cent? (1 ton = 907.18 kilograms.) How many cubic feet of air must be introduced into the chamber to convert the sulfur dioxide to sulf uric acid ? 4. Solve the same problem, substituting the metric ton and cubic meters for ton and cubic feet. How many kilograms of water must be introduced in the chamber ? 5. What is the percentage increase in the volume of the air which must be used if the temperature is 25 instead of ? 6. How many grams of ferric oxide will be obtained by burning one kilogram of iron pyrites ? CHAPTER XII NITROGEN SYMBOL, N. ATOMIC WEIGHT, 14.01. Occurrence and Natural History of Nitrogen. Fluorine and oxygen, the first elements of the seventh and sixth groups, are very active ; and while oxygen is found free in the air, this seems to be more on account of its abundance and because nearly all of the other elements in nature are already combined with oxygen than because of any lack of activity. Nitrogen, the first element of the fifth group, in very striking contrast to oxygen and fluorine, is found chiefly in the free state in the atmosphere because it does not readily combine with any of the other elements avail- able in the earth. Until 1894 it was supposed that the gas remaining when oxy- gen, carbon dioxide and moisture were removed from air was pure nitrogen. Rayleigh and Ramsay showed at that time that the residue left after the removal of these substances still con- tained about 1.2 per cent of gases, chiefly argon (p. 235), which are even more inert than nitrogen. Nitrogen forms, however, nearly 78 per cent of the volume of dry air ; and as the air above a square centimeter of the earth's surface weighs about one kilogram or that above a square meter weighs more than ten tons, it is evident that the amount of nitrogen in the atmosphere is very large, though it is small in comparison with the amounts of those elements which make up the bulk of the solid crust of the earth. In combination, nitrogen is found in all living organisms as an essential constituent, the four most important elements in organic matter being carbon, nitrogen, hydrogen and oxygen. But while plants can obtain the carbon for their growth from the carbon dioxide of the air and the oxygen and hydrogen from the moisture of the soil, very few, if any, of the higher forms of plant life can use the nitrogen of the air directly. The 198 NITROGEN 199 larger part of the nitrogen which is essential for the growth of crops must be supplied in the form of compounds which result from the decay of animal or vegetable substances, or from the combination of oxygen and nitrogen in the air through electrical agencies. There are, however, a few plants, especially clover, alfalfa and other leguminous plants, which are able to assimilate the nitrogen t>f the air with the aid of bacteria which grow in nodules on their roots. The decay of organic matter containing nitrogen is always caused by the growth of bacteria. In the absence of air, the conditions of decomposition lead to the reduction of the nitrogen to ammonia. In the presence of air, as in a well-aerated soil, the nitrifying bacteria, which are usually present, will convert the nitrogen to nitric acid, HNOs, which generally finds enough potassium, calcium or sodium present to form saltpeter, KNOs, calcium nitrate, Ca(NO 3 )2 or sodium nitrate, NaNOa. All plants can readily assimilate the nitrogen of the nitrates and so it finds its way back into the organic compounds of the plant life. As nitrogen in a readily available form is essential to the growth of wheat, corn and other crops, sodium nitrate, the cheapest of the commercial nitrates, and ammonium sulfate, (NH 4 )2SO 4 , also a comparatively cheap nitrogen compound, are often used as fertilizing materials. The course of nitrogen in nature is sh<5wn by the following diagram : l Leguminous plants with the help of bacteria Atmospheric elftntriftitv Plants ' Atmosphe Nitrogei _: "%^ Nitrates fing i Nitrif Organic compounds of Nitrogen Denitrifying bacteria Ni' bac k trif: teris Decay and ani or dis, of plant mal tissues tillation Mitr ying [ Ammonia See Abegg, Handbuchderanorg. Chemie, Bd. 3, Abth. 3, S. 215. 200 A TEXTBOOK OF CHEMISTRY Preparation and Properties of Nitrogen. Nitrogen which is pure with the exception of about 1.2 per cent of the inert gases of the argon family may be prepared by burning phosphorus in air or by passing air over heated copper turnings, which will take up the oxygen. A very convenient method for the preparation of considerable quantities of nitrogen is to pass a mixture of air and hydrogen through a tube containing hot copper oxide. By keeping the hydrogen slightly in excess of the amount necessary to combine with the oxygen, a part of the copper oxide will be reduced to metallic copper ; and it is then easy to regulate the currents of the two gases so that both copper and copper oxide will be present in the tube. The hy- drogen must, of course, be mixed with the air at the point where the two gases come in contact with the copper. (Why ?) To prepare nitrogen free from argon some compound of nitro- gen must be used, ammonium nitrite, NH 4 NO 2 , being most suitable. When a solution of the salt is warmed, it decomposes to water and nitrogen : NH 4 NO 2 = N 2 + 2 H 2 O Instead of ammonium nitrite a mixture of sodium nitrite and ammonium chloride may be used, the following reversible reac- tion occurring first : NaNO 2 + NH 4 C1 ^ NH 4 NO 2 + NaCl Nitrogen is a colorless, odorless gas which condenses to a liquid that boils at 196 and freezes at 210. It is very inert and does not combine with any element at ordinary tem- peratures, except under the influence of microorganisms, as referred to above. At the high temperatures of the electric discharge it combines with oxygen to form nitric oxide, NO, and this fact has been used recently as a basis for the commercial manufacture of nitrates (see below). As the supply of sodium nitrate in Chile, the only large supply now known, will be ex- hausted within a comparatively few years, it seems certain that this manufacture is destined to be very important. AMMONIA 201 At moderate temperatures and under high pressure nitrogen and hydrogen combine to form ammonia, NH 3 , but the reaction is very slow without some catalytic agent. The combination is exothermic and hence the equilibrium, is shifted toward the decomposition of the ammonia at high temperatures. As the volume decreases as the gases combine, pressure shifts the equilibrium to the right. Principle of van't Hoff-Le Chatelier, p. 111. (See Haber, Z. Elektrochem. 16, 244.) A careful study of the conditions best suited for the reaction has proved so encouraging that the Badische Anilin Soda Fabrik in Germany is preparing for the manufacture of synthetic ammonia on a large scale. The best catalyzers for the reaction seem to be metallic osmium or uranium. (See J. Ind. and Eng. Chem. 5, 328 (1913).) Several metals, especially lithium, magnesium or calcium, com- bine with nitrogen to form nitrides, at high temperatures : ,Li 6Li or Li 2Li 3 N Lithium Nitride = Mg 3 N 2 Magnesium Nitride Nitrogen will not support combustion nor burn. Ammonia. When organic matter containing' nitrogen decom- poses with exclusion of air, either under the influence of bacteria or of heat, the nitrogen is converted partly into ammonia, NH 3 . In this way ammonia is always found in sewage or in piles of manure. It is also formed in the destructive distillation of bi- tuminous coal for the manufacture of illuminating gas. The aqueous portion of the liquid distillate from the coal furnishes, at present, the chief source of the ammonia of commerce. These 202 A TEXTBOOK OF CHEMISTRY ammoniacal gas liquors are mixed with slaked lime and distilled, the lime retaining sulfur and other impurities. The distillate is mixed with hydrochloric or sulfuric acid and evaporated to obtain ammonium chloride, NH 4 C1, (NH 3 + HC1), or ammo- nium sulfate, (NH 4 ) 2 SO 4 , (2 NH 3 + H 2 SO 4 ). From these salts the ammonia may be liberated by any strong base, as sodium hydroxide or calcium hydroxide : (NH 4 ) 2 SO 4 + Ca(OH) 2 ^CaSO 4 + 2 NH 4 OH ^ 2 NH 3 + 2 H 2 O Ammonia may also be obtained by hydrolyzing a nitride with water : Li 3 N + 3 HOH = NH 3 + 3 LiOH For laboratory or lecture purposes ammonia gas is most easily obtained by boiling a strong solution known as aqua ammonia and passing the gas through a cylinder filled with quicklime to dry it. Properties of Ammonia. Ammonia is a colorless gas with a very pungent odor. It is very easily soluble in water, and hence must be collected by displacement of air (should the mouth of the bottle point up or down ?) or over mercury. Water at absorbs about 1000 times its volume of the gas, but gives off a large part of it on warming gently and all of it on boiling. The density of the solution is less than that of water, a 28 per cent solution having a specific gravity of 0.90. Ammonia combines directly with acids to form ammonium salts, in which the hydrogen of the acid combines with the am- monia to form the ammonium group, NH 4 , a radical which in its compounds possesses properties very closely resembling the properties of potassium or sodium : NH 3 + HC1 = NH 4 C1 Ammonium Chloride 2 NH 3 + H 2 S0 4 = (NH 4 ) 2 SO 4 Ammonium Sulfate .1 These compounds are most satisfactorily explained by sup- posing that nitrogen is trivalent when combined exclusively with AMMONIA 203 hydrogen or with positive groups but may become quinquivalent when one of the groups or atoms is negative : Any acid may take up as many molecules of ammonia as it has of replaceable hydrogen atoms. Thus a monobasic acid, as nitric acid, HNOa, may combine with one molecule of am- monia, forming ammonium nitrate, NH^NOs, or a tribasic acid, as H 3 PO 4 , may combine with three molecules. The formation of an ammonium salt may be very prettily illustrated by filling two cylinders with ammonia and hydro- chloric acid gas respectively. On bringing the mouths of the cylinders together the gases will combine to form solid am- monium chloride. Aqua Ammonia. The solubility of ammonia in water has already been mentioned. It dissolves, in part, without chemical change, as is shown by the strong odor of the solution due to the escape of the gas, but it partly combines with the water, which easily separates into H + and OH~, forming ammonium hydroxide : H \OH This ionizes to form ammonium, NH4 + , and hydroxide, OH~, ions, but the ionization is small in comparison with that of strong bases. A tenth normal solution of sodium hydroxide is ionized to the extent of about 84 per cent, while a tenth normal solution of ammonium hydroxide (or ammonia) shows only 1.3 per cent of ionization, if we assume that all of the ammonia in the solution has combined with water to form ammonium 204 A TEXTBOOK OF CHEMISTRY hydroxide. A normal solution, on the same basis, shows an ionization of only 0.3 to 0.4 per cent. How may the presence of hydroxide ions in a solution of ammonium hydroxide be demonstrated ? It is probable, however, that a large part of the ammonia exists as such in the solution. 1 In other words, the solution contains ammonia, NH 3 , as well as ammonium hydroxide, NH 4 OH, and ammonium, NH 4 + , and hydroxide, OH~, ions. The practical effect, that ammonium hydroxide is a weak base because its solution contains few hydroxide ions, is the same whether we suppose this to be because ammonium hydroxide is only slightly ionized or whether it is because the ammonium hydroxide is largely dissociated into ammonia and water. Ice Machines. Ammonia may be readily condensed to a liquid either by pressure (4.19 atmospheres at 0) or by cold ( 33 at 760 mm.). The heat of vaporization of the liquid is 330 calories per kilogram (at 33). This high value is intimately connected with the low molecular weight and also with the fact that liquid ammonia is, like water, a highly " asso- ciated " liquid, that is, consists of polymerized molecules such as (NH 3 ) 2 or (NH 3 ) 3 . Much heat is absorbed in the vaporiza- tion both because of the large volume of the vapor in proportion to its weight and because the polymerized molecules must be broken up. The high heat of vaporization is utilized in ice machines. The principle of one form of these machines is illus- trated in the diagram (Fig. 71). Liquefied anhydrous ammonia is allowed to evaporate in the coil A and the escaping gas is com- pressed by the pump B and condenses to a liquid in the coil C. From this coil the liquid ammonia is returned through the 1 Moore, J., Chem. Soc. 91, 1382 (1907), calculates on the basis of the partition of ammonia between water and chloroform at dif- ferent temperatures that from 30 to 40 per cent of the ammonia exists as ammonium hydroxide, NH 4 OH, at 20. This does not seem to be consistent, however, with the relative ionization constants of trimethyl amine, (CH 3 ) 3 N, and tetramethyl ammonium hydroxide, (CH 3 ) 4 NOH, which indicate strongly that the former forms only a small amount of trimethyl ammonium hydroxide, (CH 3 ) 3 NHOH, in aqueous solutions. Further experimental evidence on this ques- tion seems highly desirable. DERIVATIVES OF AMMONIA 205 regulating valve D to the coil A. The coil C is surrounded by cold water to absorb the heat evolved as the ammonia condenses. The coil A is surrounded with brine, either a solution of salt or of calcium chloride. The cold brine may be circulated by means of pumps through coils of pipe in refrigerator rooms, or cases of distilled water may be immersed in the brine, to be frozen. In another form of machine, which was in earlier use, the pressure to condense the ammonia was obtained by heating a Fig. 71 concentrated aqueous solution. Afterwards the weakened solution was cooled, and as it reabsorbed the gas the evapora- tion of the liquefied gas caused the refrigeration. * Derivatives of Ammonia. One or more atoms of hydrogen in ammonia may be replaced by a metal, giving such compounds as sodium amide, NaNH 2 , or by radicals, especially by organic radicals, giving such compounds as methyl amine, CHaNH^, phenyl amine or aniline, CeHsNH^, acetamide, C2H 3 ONH2, and phthalimide, C 6 H NH. If the group re'placing hydrogen is a hydrocarbon radical, as methyl, CH 3 , or phenyl, C 6 H 6 , the compound is called an amine. The amines combine with acids, as ammonia does, to form such salts as methyl ammonium chloride, CH 3 NH 3 C1 (or CH 3 NH 2 .HC1) and phenyl ammonium chloride, C 6 H 5 NH 3 C1. For this reason the amines are often called organic bases, but just as ammonia is a base only when it 206 A TEXTBOOK OF CHEMISTRY has combined with water to form ammonium hydroxide, NH 4 OH, the amines are true bases only when combined with water. As ammonium hydroxide dissociates to ammonia and water even in solution, such hydroxides as methyl ammonium hydroxide, CH 3 NH 3 OH, dissociate readily into water and the original amine and can exist as pure compounds only at very low tempera- tures, if at all. All four of the hydrogen atoms in the ammo- nium group, NH 4 , of ammonium hydroxide may be replaced by hydrocarbon radicals, however, and some compounds formed in this manner no longer dissociate when their solutions are evapo- rated. Thus a white crystalline mass, doubtless consisting of tetramethyl ammonium hydroxide, (CH 3 ) 4 NOH, separates on evaporating a solution of this compound. The preparation of this and several other similar compounds has been one of the reasons for believing that solutions of ammonia in water contain ammonium hydroxide, NH 4 OH. If the group replacing a hydrogen atom in ammonia is an acid radical, the compound is called an amide. Thus the com- pound containing the acetyl group, C 2 H 3 O, is called acetamide, C2H 3 ONH 2 . In aqueous solutions the amides are usually am- photeric; that is, they have both very weak acid and very weak basic properties. When two hydrogen atoms of ammonia have been replaced by a bivalent acid radical, the compound is / co \ called an imide, as phthalimide, CeH^ />NH. In the XXX imides the hydrogen can be replaced by metals forming C \ well-defined salts, as potassium phthalimide, C 6 H 4 ^ ">NK or C 6 H 4 <( ^>N. The Electron Theory. A comparison of the compounds of nitrogen on the basis of the electron theory (p. 181) is suggestive. In nitric acid, H O N^ , the nitrogen atom gives five ELECTRON THEORY 207 electrons to the oxygen atoms, forming a compound which i^ 0= readily ionizes to hydrogen, H + , and nitrate, O N+>^ +XT ions because of the strongly positive nitrogen atom, which holds the oxygen of the hydroxyl firmly but repels its hydrogen. In / H ammonia, H N<^ , the nitrogen atom receives three elec- trons from the hydrogen atoms, becoming negative. It can receive a fourth electron from another hydrogen atom only in case it also gives up one electron to the oxygen of a hydroxyl group or to chlorine or some other negative atom or group. In H\ /H + the ammonium hydroxide, H ? N+\ , which results, H +/ X ~O-H+ in the first case, the negative nitrogen atom no longer holds the negative oxygen of the hydroxyl group strongly, and so the com- pound may ionize to ammonium, NHU+, and hydroxide, OH~, ions. If one of the hydrogen atoms of ammonia is replaced by a nega- tive radical, as acetyl, CzH-sQ, giving the compound acetamide, C 2 H 3 Ov , the nitrogen atom no longer takes up hydrogen H^N W and hydroxyl readily to form a base, or the elements of an acid to form a salt. This seems to be because the presence of the negative acetyl group, C 2 H 3 O, so far reduces the positive character of the group C 2 H 3 (X H that it cannot form W the positive ion of a salt. * Solutions in Liquid Ammonia. Anhydrous ammonia may be condensed to a liquid which boils at 33.5. This liquid ammonia dissolves many substances, and the conductivity of the solutions indicates that some of these ionize in the ammonia as acids, bases and salts ionize in solutions in water. While a very large portion of our study of chemistry deals with reactions 208 A TEXTBOOK OF CHEMISTRY in aqueous solutions, there are closely parallel phenomena in ammoniacal solutions. In such solutions we may consider ammonia as consisting of H and NH 2 just as we think of water as consisting of H and OH. As in aqueous solutions derivatives of water ionize to form hydrogen, H + , or hydroxide, OH~, ions according to the nature of the radical, so in ammonia, com- pounds which are derivatives of ammonia may ionize to form hy- drogen, H + , or amide, NH 2 ~, ions. Thus acetamide, C 2 H 3 ONH 2 , ionizes to C 2 H 3 ONH- and H+ in solution in ammonia and is to be considered as an acid in such a solution. Sodium amide, NaNH 2 , on the other hand, ionizes to sodium, Na + , and amide, NH 2 ~, ions and is to be considered as a base. Curiously enough the latter will cause phenolphthalein to turn red in the ammonia solu- tion just as hydroxide ions cause it to turn red in aqueous solutions. Neutralization in such a solution must con- sist in the union of hydrogen, H + , and amide, NH 2 ~, ions to form ammonia. See Franklin and his coworkers, Am. Chem. J. 20, 820 and 826 ; 21, 8 ; 23, 277 ; 28, 83 ; 47, 285; J. Am. Chem. Soc., 26, 499; 27, 192, 820. The Volumetric Composition of Ammonia. The ratio between nitrogen and hydrogen in ammonia may be demonstrated by filling a tube, Fig. 72, with chlorine gas, allowing a small amount of concentrated aqua am- monia to enter it and following this with some dilute sulfuric acid. On allowing water to enter the tube till the gas remaining in it is at atmospheric pressure, the Fig. 72 tube will be found to be one third full of nitrogen. Under the conditions of the experiment the tube full of chlorine, which we know is capable of combining with its own volume of hydrogen, takes this amount of hydrogen from the ammonia and liberates the equivalent amount of nitrogen. In other words, one volume of nitrogen is combined with three volumes of hydrogen. As the gram molecular volume of ammonia weighs 17 grams, the complete reaction is : COMPOSITION OF AMMONIA 209 NH 3 NH 3 The primary reaction between ammonia and chlorine is : 3 NH 3 + 6 C1 2 = NC1 3 + N 2 + 9 HC1 This is followed by the reaction : NC1 3 + NH 3 = N 2 + 3 HC1 % which is favored in an acid solution. 1 The final result is the same as though the simple reaction 2 NH 3 + 3 C1 2 = 6 HC1 + N 2 took place. As in this reaction three molecules of chlorine, C1 2 , liberate one molecule of nitrogen, N 2 , it is evident that in accord- ance with Avogadro's law a tube full of chlorine should liberate from ammonia one third of a tube full of nitrogen. The composition of ammonia by volume may also be shown by another, quite different, experiment. If a small amount of ammonia gas is introduced into the apparatus shown in Fig. 73, and electric sparks are passed between the platinum wires, which pass through the walls of the tube, the ammonia will be decom- posed into hydrogen and nitrogen. The reaction is reversible, but with the equilibrium very far toward the side of decomposi- tion at the temperature of the electric discharge : 2 NH 3 N 2 + 3 H 2 As two molecules of ammonia give one molecule of nitrogen and three molecules of hydrogen, the volume of the gas would be doubled when the decomposition was complete. 1 J. Am. Chem. Soc. 23, 460. 210 A TEXTBOOK OF CHEMISTRY On the other hand, if a mixture of one volume of nitrogen with three volumes of hydrogen is placed in the same apparatus and a little dilute sulfuric acid is introduced, on passing electric sparks through the mixture as before, the nitrogen and hydrogen will slowly combine, and, as the ammonia formed will be absorbed by the sulfuric acid, the combination may be carried to comple- tion in spite of the unfavorable character of the equilibrium. Nitric Acid. The two compounds, ammonia and nitric acid, are to be considered as the fundamental ones for nitrogen. All other com- pounds of the element tend to return to one or the other of these, or their salts, or else to de- compose with the liberation of free nitrogen. And all compounds of nitrogen prepared in the laboratory, except those derived from organic materials, are prepared directly or indirectly from ammonia or nitric acid. It seems best, therefore, to speak of nitric acid next, though such an order of treatment is practical rather than logical. The formation of nitrates in the soil by the action of nitrifying organisms has been referred to. The present commercial source for nitrates is almost exclusively the sodium nitrate, NaNOs, or Chile saltpeter, found in enormous beds in Chile, in South America. From this nitric acid is prepared by a process similar to that for the preparation of hydrochloric acid. Nitric acid is a stronger acid than sulfuric, but if sodium nitrate is mixed with sul- furic acid and the mixture heated, the equilibrium of the reaction, NaN0 3 + H 2 SO 4 ^ NaHSO 4 + HNO 3 is displaced to the right as the nitric acid distills away from the mixture, the boiling point of nitric acid being much lower than that of the sulfuric acid. Nitric acid is not, however, very stable, Fig. 73 NITRIC ACID 211 and part of it decomposes, forming oxygen, water and nitrogen peroxide, NO 2 , when the distillation is under atmospheric pres- sure. To avoid this the operation is sometimes carried out at a lower temperature by reducing the pressure. Pure nitric acid is a colorless liquid, which boils at 86 and has a specific gravity of 1.52. The addition of water causes a rise in the boiling point, the acid of maximum boiling point, 120, containing 66 to 70 per cent of pure acid. Nitric acid is a strong acid, the tenth normal solution being ionized to the extent of 92 per cent, while a tenth normal solution of hydrochloric acid is 91 per cent ionized. * Hydrates of Nitric Acid. The addition of increasing amounts of nitric acid to water lowers the freezing point till an acid con- taining 32.8 per cent of nitric acid freezes at 43. Further addition of nitric acid causes the freezing point to rise and fall C S 10 *S \ 1 ~ 2 E r,. or* \ \ \ / \ TEMPERATURE Ol i A i I c C \ j \ V \ / \ 1 ) 10 20 30 40 50 60 70 80 90 100% PER CENT. OF NITRIC ACID. HNO 3 . Fig. 74 alternately as shown in Fig. 74, pure nitric acid freezing at 41.2. The two maximum freezing points shown in the figure correspond to acids containing 53.84 and 77.77 per cent of 212 A TEXTBOOK OF CHEMISTRY nitric acid respectively. In accordance with the principle that a pure liquid freezes at a higher temperature than when it con- tains some dissolved substance, these melting points indicate definite compounds. An acid containing 53.84 per cent of nitric acid corresponds to the formula HNO 3 .3 H 2 O and one contain- ing 77.77 per cent, to the formula HNO 3 .H 2 O. The freezing point curve demonstrates very clearly the existence of these two hydrates. 1 Chemical Properties of Nitric Acid. As an acid, nitric acid has the usual properties, forming salts as a monobasic acid with practically all metals. These may be prepared by the action of the acid on the metal or on a hydroxide, oxide or carbonate of the metal : NaOH + HN0 3 = NaNO 3 + HOH Ca(OH) 2 + 2 HNO 3 = Ca(NO 3 ) 2 + 2 H 2 O ZnO + 2 HNO 3 = Zn(NO 3 ) 2 + H 2 O The most important special properties of nitric acid depend on the ease with which it gives up oxygen to a great variety of substances. It is a powerful oxidizing agent, the concentrated or the anhydrous acid, HNO 3 , showing this property in a more marked degree than the dilute acid. Indeed the pure acid de- composes spontaneously into water, oxygen and nitrogen perox- ide, NO 2 , on distillation or on exposure to light : 2 HNO 3 = 2 NO 2 + O + H 2 O If pure nitric acid is boiled in a test tube having a plug of wool or feathers in its mouth, the latter will catch fire and burn. Ignited charcoal will also continue to burn beneath the surface of the liquid. When nitric acid acts upon a metal, hydrogen is rarely, if ever, liberated. Instead of this the nitric acid is reduced, either by the metal directly, forming an oxide of the metal, or by the hy- drogen displaced by the metal. Such hydrogen is often called 1 Kiister u. Kremann, Z. anorg. Chem. 41, I (1904). AQUA REGIA . 213 nascent" 1 hydrogen. The two ways of explaining the action may be illustrated as follows : 3 Cu + 2 HNO 3 = (3 CuO) + H 2 O + 2 NO (3 CuO) + 6 HN0 3 = 3 Cu(NO 3 ) 2 + 3 H 2 O 3 Cu + 8 HNO 3 = 3 Cu(NO 3 ) 2 + 4 H 2 O + 2 NO Or 3 Cu + 6 HN0 3 = 3 Cu(NO 3 ) 2 +(6 H) (6 H) + 2 HNO 3 = 4 H 2 + 2 NO 3 Cu + 8 HNO 3 =3 Cu(NO 3 ) 2 + 2 NO + 4 H 2 O The final reaction is the same whichever explanation of the mechanism of the reaction is adopted. Nitric acid may be reduced less or more by other metals, or by copper when the concentration of the nitric acid is different, giving the whole series of oxides, NO 2 , N 2 O 3 , NO and N 2 O, and even ammonia, NH 3 . In general a concentrated acid gives the higher oxides, while a more dilute acid or a metal which has a high heat of oxidation gives the lower oxides or ammonia. If nitric acid is added to a test tube containing zinc and sulf uric acid, the evolution of hydrogen may nearly cease and ammonium sulf ate will be formed. The addition of an excess of sodium hydroxide to the solution will then liberate ammonia, which may be recognized by its odor and effect on litmus paper. What is the series of reactions involved in the experiment ? Aqua Regia. Neither nitric acid or hydrochloric acid alone will dissolve the so-called noble metals, gold or platinum, but a mixture of the two will dissolve them readily. Such a mixture is called aqua regia 2 because of this property. The nitric acid acts upon hydrochloric acid in the same manner as other oxidiz- ing agents, liberating chlorine, and this attacks the gold or plati- num, forming soluble chlorides. Because it contains chlorine and oxides of nitrogen, aqua regia 1 Nascent means in the state of being born. The theory is that hydrogen atoms when first liberated are more active than the same atoms when they have combined with others to form hydrogen molecules, H 2 . The activity is dependent largely on the metal used, however, metals which dissolve with a large evolution of heat giving the most active hydrogen. 2 Royal water. 214 A TEXTBOOK OF CHEMISTRY is a powerful oxidizing agent and is frequently used for that pur- pose, especially to oxidize the sulfur of sulfides to sulfuric acid, as in the determination of sulfur in iron. Nitrosyl chloride, NOC1, a volatile compound with a very disagreeable odor, is also formed in the mixture of nitric and hydrochloric acids. This is hydrolyzed by water in the same manner as other nonmetallic chlorides : NOC1 + HOH = NOOH (or HNO 2 ) + HC1 Nitrosyl chloride is also a strong oxidizing agent. Oxides of Nitrogen. There are six oxides of nitrogen, but as two of these, nitrogen peroxide, NO2, and nitrogen tetroxide, N 2 O 4 , have the same percentage composition and change each into the other on merely changing the temperature, they are frequently spoken of as a single substance and the name nitrogen peroxide is applied to both. The oxides are : Nitrous oxide N 2 O Nitric oxide NO Nitrous anhydride N 2 O 3 Nitrogen dioxide NO 2 Nitrogen tetroxide N 2 O 4 Nitric anhydride N 2 O 5 Nitrous oxide, N 2 O, is most easily prepared by heating am- monium nitrate; nitric oxide, NO, by the reduction of nitric acid with metallic copper ; nitrous anhydride, N 2 O 3 , by reducing nitric acid with arsenious oxide, As 2 Os ; nitrogen dioxide, NO 2 , by heating lead nitrate, Pb(NO 3 ) 2 ; and nitric anhydride, N 2 O 5 , by dehydrating nitric acid with phosphoric anhydride. Nitric oxide is also formed by the direct union of the elements in an electric arc, and it unites with oxygen spontaneously to form nitrous anhydride, N 2 O 3 , and nitrogen dioxide, NO2. Nitrous Oxide, N 2 O. When ammonium nitrate, NH 4 NO 3 , is heated, the hydrogen of one part of the molecule combines with oxygen from another part to form water, while the two nitrogen atoms remain combined with the other oxygen atom : NH 4 NO 3 = N 2 O + 2 H 2 O NITRIC OXIDE 215 The reaction is exothermic, that is, it takes place with evolu- tion of heat and is liable to become explosive if the temperature is raised too high or if too large a quantity of the salt is heated at once. Nitrous oxide is a colorless gas with a sweetish odor and taste. Water at 20 absorbs about two thirds of its volume of the gas. It supports combustion. A glowing splinter will inflame in the gas somewhat as it does in oxygen, and phosphorus burns in it with an intense light. It will not support life. When inhaled, nitrous oxide sometimes causes hysterical laugh- ing, and it is called for that reason laughing gas. In larger amounts it produces insensibility and is used for this purpose in minor surgical operations, especially for the extraction of teeth. The structure of nitrous oxide is probably that represented by N the formula, II }O. This accounts best for the ease with which W it gives up oxygen and reverts to free nitrogen. It is the only oxide of nitrogen in which two atoms of nitrogen are supposed to be directly united. Nitric Oxide, NO, is formed when copper dissolves in dilute nitric acid of specific gravity 1.2. If a stronger acid is used, some nitrogen dioxide, NO2, will be formed ; while if the acid is much more dilute, nitrous oxide, N2O, and nitrogen, N2, will be formed along with the nitric oxide. The mechanism of the reaction has already been discussed. Nitric oxide, NO, is also formed by the direct union of nitrogen and oxygen in an electric arc. For instance,' if electric sparks from an induction coil are passed for some time between terminals in a large globe, the air in the globe will gradually become reddish brown in color from the formation of nitric oxide, which combines with more oxygen to form nitrogen dioxide, NO 2 . There. is evi- dence that the combination is caused by the high temperature of the arc and not by the electricity as such. The combination of nitrogen and oxygen is an endothermic 216 A TEXTBOOK OF CHEMISTRY reaction ; that is, heat is absorbed as it proceeds. It is also a reversible reaction : with the equilibrium very far to the left, so far, indeed, that the amount of nitric oxide formed from the elements is very small, even at high temperatures. The per cent of nitric oxide formed in air when the reaction comes to equilibrium is as follows : 1 ABSOLUTE TEMPERATURE .. PER CENT OF NO CALCULATED OBSERVED 1811 0.35 0.37 2195 0.98 0.97 2675 2.37 2.23 3200 4.43 About 5. With the equilibrium so far on the side toward its decomposi- tion, it seems at first thought that nitric oxide ought not to exist at all at ordinary temperatures, and it could not except for the fact that the speed of the formation or decomposition is very slow. Thus it has been shown that for the formation of half of the amount corresponding to a condition of equilibrium, 80 years would be required at a temperature of 725 and one and a fourth days at 1225. At 1825 it takes only 5 seconds. These facts are important in determining the heat conditions for the preparation of nitric oxide from the air as the first step in manufacturing nitrates. Evidently the highest possible tem- perature must be secured, and when the nitric oxide has been formed at that temperature, it must be cooled as quickly as possible through the temperatures at which formation and de- composition are both rapid and where the point of equilibrium lies farther toward the side of decomposition. Nitric oxide is a colorless gas which may be condensed to a liquid that boils at -153.6 and freezes at -167. Is the gas heavier or lighter than air ? 1 Nernst, Z. anorgf. Chem. 49, 213. NITRIC OXIDE 217 Nitric oxide, quite unlike nitrous oxide, will extinguish a glowing coal or a candle which is thrust into it, or even a piece of phosphorus which is barely ignited. If a little hotter, phos- phorus burns brilliantly in the gas, and a mixture of the vapor of carbon bisulfide with the gas will burn with a bright blue flash, which is rich in the light rays that affect a photographic plate. Nitric oxide combines directly with oxygen to form nitrous anhydride, N 2 O 3 , nitrogen dioxide, NO 2 , or nitrogen tetroxide, N 2 O4, or a mixture of the three, according to the temperature, and the relative volumes of the gases. If a strong base, such as potassium hydroxide, KOH, is present when nitric oxide and oxygen are brought together, a nitrite, KNO 2 , is formed, even though the oxygen is in excess. In what proportion by volume must nitric oxide and oxygen be brought together to form nitrous anhydride ? In what proportion to form nitrogen dioxide ? The structural formula of nitric oxide is usually written N = O, representing the nitrogen as bivalent, a valence which is very unusual for the element and which it is not known to have in any other compound. This unusual structure seems to be closely connected with the tendency of nitric oxide to combine with more oxygen. * Nitric oxide is formed by the reduction of nitric acid by ferrous sulfate, FeSO4, the latter being oxidized to ferric sulfate, Fe 2 (SO 4 ) 3 , if sulfuric acid is present : 2 HNO 3 = 2 NO + (3 O) + H 2 O 2 FeSO 4 + H 2 SO 4 + (O) = Fe 2 (SO 4 ) 3 + H 2 O Combining the equations, 6 FeSO 4 + 3 H 2 SO 4 + 2 HNO 3 = 3 Fe(SO 4 ) 3 + 3 H 2 O + 2 NO The nitric oxide is absorbed by a solution of ferrous sulfate with the formation of an unstable compound, FeSO 4 .NO (see Manchot and Zechentmayer, Ann. 350, 368), which gives to the solution a dark brown or black color and which is often used for the qualitative detection of nitric acid. The reduction 218 A TEXTBOOK OF CHEMISTRY of nitric acid to nitric oxide by ferrous chloride, FeCl2, and hydro- chloric acid is also used for its quantitative determination. Nitrous anhydride, N 2 O 3 , may be prepared by the union of oxygen with nitric oxide, by the action of a dilute acid on sodium nitrite, NaNO 2 , or by the reduction of nitric acid, especially by arsenious oxide, As 2 O 3 : 4 NO + O 2 = 2 N 2 O 3 2 NaNO 2 + H 2 SO 4 = Na 2 SO 4 + 2 HNO 2 Nitrous Acid 2 HNO 2 = N 2 O 3 + H 2 O 2 HNO 3 + As 2 O 3 + 2 H 2 O = 2 H 3 AsO 4 + N 2 O, Arsenic Acid At ordinary temperatures a gas having the composition of nitrous anhydride has a density which indicates that the com- pound dissociates into nitric oxide, NO, nitrogen dioxide, NC>2, and nitrogen tetroxide, N 2 O 4 : 2 NO 2 ^ N 2 O 4 (See below.) At a low temperature the gases recombine, in part, and con- dense to a dark blue or green liquid, but even at 90 there is still some dissociation (Ramsay). Nitrous Acid. If the mixture of gases spoken of in the last paragraph is dissolved in cold water, nitrous acid, HNO 2 , is formed : NO + NO 2 + H 2 O ^ 2 HN0 2 The acid is very unstable and can exist only in dilute solutions. Salts of nitrous acid are most easily prepared by reducing a ni- trate, as, for instance, sodium nitrate, with lead or copper : NaNO 3 + Pb = NaNO 2 + PbO Sodium Nitrite NITROGEN DIOXIDE 219 Nitrogen Dioxide, NO 2 , and Nitrogen Tetroxide, N 2 O4. When nitric oxide is mixed with air or oxygen, the colorless gas changes to a reddish brown color and is converted, at ordinary temperatures, into a mixture of nitrogen dioxide, NO 2 , and nitrogen tetroxide, N 2 O 4 . The molecular weight of nitrogen dioxide is 46, that of nitrogen tetroxide is 92. As the gram molecular volume of the gas weighs about 80 grams at 10, 69 grams at 64 and 46 grams at 150, it is evident that at the lower temperature it consists chiefly of nitrogen tetroxide, N 2 O 4 , but that this dissociates into nitrogen dioxide, NO 2 , as the temperature rises : At still higher temperatures nitrogen dioxide dissociates into nitric oxide, NO, and oxygen. The mixture of the two gases may be easily condensed to a reddish brown liquid which boils at 25. At lower temperatures the liquid becomes lighter colored and solidifies to colorless crys- tals at 10.5. From this it would seem that nitrogen tetrox- ide is colorless, while the dioxide is colored. The structure of the two compounds is probably : . and ^N O N=O CT The tetroxide, \N O N=O, may be considered as partly nitrous anhydride, O=N O N=O, and partly nitric anhydride, ^N O N^ In accordance with this view CT it gives both nitric and nitrous acid when dissolved in cold water : 220 A TEXTBOOK OF CHEMISTRY V / i N O N=O ,O # = H O NC +H O N=0 H O H ^O Nitric Acid Nitrous Acid In warm water nitrogen dioxide forms nitric acid and nitric oxide : ,/>O NO 2 + HOH = H O Nf + (H) 2(H)+ NO 2 = NO + H 2 O or 3 NO 2 + H 2 O = 2 HNO 3 + NO As the nitric oxide on coming to the air immediately combines with oxygen to form the dioxide, these reactions furnish a means for converting the nitric oxide obtained from atmospheric nitro- gen in the electric arc into nitric acid. Nitrogen Pentoxide is formed when nitric acid is mixed with phosphorus pentoxide : 2HNO 3 + P 2 O 6 = 2HPO 3 + N 2 O 6 Metaphosphoric Acid It forms colorless crystals which melt at 29.5. It decomposes easily into nitrogen dioxide and oxygen. With water it gives nitric acid. Other Compounds of Nitrogen. The halogens are univalent toward hydrogen and form only one compound, each, with that element. Hydrochloric acid, HC1, is typical for the group. Oxygen, which is bivalent, forms two compounds, with hydro- gen H O H and H O O H. Nitrogen, which is trivalent, forms at least five compounds with hydrogen. ,^ Ammonia, NHa, or Nr H \H H \ Hydrazine, 1 N 2 H4, or W H 1 Nitrogen is called "azote" (deprived of life) in French, and many names of nitrogen compounds are derived from the root " az " or "azo." HYDROXYLAMINE 221 Hydronitric acid, N 3 H, or H v H\ , , H-^N W N N Ammonium trinitride, N4EU, or H-N N^ X N H H AH >N N^- Hydrazine trinitride, N 5 H 5 , or >N N^- -- N W X H We shall find later that carbon, which is quadrivalent, forms many hundred of compounds with hydrogen. Nitrogen also forms a third oxygen acid, hyponitrous acid, H2N2O2, and a derivative of ammonia, hydroxylamine, NH 2 OH. * Hyponitrous Acid, H 2 N 2 O 2 . From its formula, nitrous oxide, N 2 O, might be looked upon as the anhydride of hyponi- trous acid, but it will not combine with water to form the acid nor will it combine with bases to form salts of the acid. The method of preparation which indicates most clearly the structure of the acid is by the interaction of hydroxylamine and nitrous acid: --------- . ' H H O N<^ +0=N OH = H O N=N O H + H 2 O >H ! The pure acid can be obtained in crystalline form, but is very explosive. The lowering of the freezing point of an aqueous solution of the acid (see p. 112) shows that the formula is H 2 N 2 O 2 , and not the simpler formula, HNO. In solutions of either alkalies or acids it decomposes slowly into water and nitrous oxide, N 2 O. * Hydroxylamine, NH 2 OH, may be prepared by the electro- lytic reduction of nitric acid, using an amalgamated lead cathode and in the presence of sulfuric acid, which combines with the hydroxylamine to form hydroxylammonium sulfate : l 1 Tafel, Z. anorg. Chem. 81, 289. 222 A TEXTBOOK OF CHEMISTRY ,0 K H O NC + 6 H = H O N< + 2 H 2 O H 2 NH 2 OH + H 2 SO 4 = (NH 3 OH) 2 SO 4 Hydroxylammonium Sulfate Pure hydroxylamine may be obtained in white scales or needles which melt at 33. It dissolves easily in water and the solution seems to contain some hydroxylammonium hydroxide, /OK NH 3 ^ , which is a much weaker base than ammonium hy- X)H droxide, NH 4 OH. Hydroxylamine is much used in organic chemistry for the preparation of derivatives of aldehydes and ketones, called oximes. * Hydrazine, N 2 H 4 . By means of a series of reactions a com- pound called bisdiazoacetic acid may be prepared. When this is hydrolyzed by warming with hydrochloric acid, it gives hydra- zine hydrochloride and oxalic acid. X \ H0 2 C CH< >CHC0 2 H + 2 HC1 + 4 H 2 O X N = N X Bisdiazoacetic Acid CO 2 H = 2H 2 N NH 2 HCl + 2 | Hydrazine CO 2 H Hydrochloride Oxalic Acid Pure hydrazine is a colorless liquid, which solidifies at a low temperature and melts at 1.4. In solution it forms a base, which seems to be H 2 N NH 3 OH, similar to ammo- nium hydroxide, NH 4 OH, and hydroxylammonium hydroxide, NH 3 OH.OH. With acids it forms salts by direct addition, taking up either one or two molecules of a monobasic acid, the chlorides being H 2 N NH 3 C1 and C1H 3 N NH 3 C1. The second molecule of the acid is held only feebly, however. In aqueous solution hydrazine gives a tenfold weaker base than ammonium hydroxide. HYDRONITRIC ACID 223 Hydrazine and its derivatives, especially phenylhydrazine, C 6 H 6 NHNH 2 , are much used in organic chemistry for the preparation of derivatives of aldehydes, ketones and sugars. /N * Hydronitric Acid or Azoimide, HN^ || . By passing am- \XT monia over heated metallic sodium a compound called sod- amide, NaNH 2 , may be prepared. When nitrous oxide, N2O, is passed over this at a temperature of 190, sodium trinitride, NaN 3 , is formed : /H /N /N Na N< +O< || = Na N< || +H 2 O X H X N X N By dissolving the sodium trinitride in water, adding sulfuric acid, and distilling, the hydronitric acid passes over with some water, from which it can be separated partly by redistillation and finally by treating with calcium chloride : 2 NaN 3 + H 2 SO 4 = Na 2 SO 4 + 2 HN 3 The pure acid freezes at 80 and boils at + 37. It is very explosive, as are also many of its salts, especially the salts with silver, AgN 3 , and copper, Cu(N 3 ) 2 . The silver and copper salts are very difficultly soluble in water. In aqueous solution the free acid is only slightly ionized, its strength in this respect resem- bling that of acetic acid. It is, however, a much stronger acid than hydrogen sulfide, H 2 S, or carbonic acid, H 2 CO 3 . Iodine Trinitride. By treating silver trinitride with iodine the silver may be replaced by iodine : , Ag-N< X N Agl + I- N Iodine trinitride is a colorless solid with a penetrating odor, resembling that of cyanogen iodide. It is hydrolyzed by water to hydronitric and hypoiodous acids : 224 A TEXTBOOK OF CHEMISTRY I N< || + H.OH = H N< || +1 O H X N N N This reaction is remarkable because the iodine conducts itself as the positive part of the molecule, combining with the negative hydroxyl, while the hydrazoic group, N 3 , combines with the posi- tive hydrogen. Probably for the same reason silver trinitride will not react with iodine trinitride to give the molecule N 6 , N N<^ || , because both the silver and iodine are positive in the two compounds. Nitrogen Trichloride. When a dilute solution of ammonia reacts with chlorine, one fourth of the chlorine combines with nitrogen to form nitrogen trichloride : 3 NH 3 + 6 Cla = NC1 3 + 9 HC1 + N 2 9 HC1 + 9 NH 3 = 9 NH 4 C1 Nitrogen trichloride is also formed when hypochlorous acid acts on ammonium chloride : H Cl N^-H + 3 HO.C1 ^ 3 H.OH + N^-C1 X H X C1 The last reaction seems to be reversible, and these reactions indi- cate that the chlorine of nitrogen trichloride is positive, just as the iodine of iodine trinitride is. Because of this, one molecule of nitrogen trichloride is equivalent to six atoms (or three mole- cules, 3 C1 2 ) of available chlorine in oxidizing power : 2 NC1 3 + 3 As 2 O 3 + 15 H 2 O = 6 H 3 AsO 4 + 2 NH 3 + 6 HC1 6 C1 2 + 3 As 2 O 3 + 15 H 2 O = 6 H 3 AsO 4 + 12 HC1 Nitrogen chloride is a volatile oil, which is very explosive and dangerous. Dulong, who discovered it, lost an eye and three fingers while working with it, and both Faraday and Davy were ENDOTHERMIC COMPOUNDS 225 injured while experimenting with the substance. It is soluble in benzene and may be handled more safely in such a solution. Nitro Nitrogen Trichloride. Some evidence has been obtained recently, which points to the existence of a nitro nitrogen tri- chloride in which the nitrogen is positive and the chlorine nega- tive. See J. Am. Chem. Soc. 35, 767 (1913). * Nitrogen Iodide, N 2 H 3 l3, is formed when a strong solution of ammonia is poured over powdered iodine. When dried in the air it forms a black powder which explodes violently at the lightest touch. Endothermic Compounds. Troost has given the heat of forma- tion of nitrogen chloride from its elements as 38,477 calories. The explosive character of nitrogen chloride and nitrogen iodide is evidently closely connected with the fact that they are en- dothermic compounds, that is, compounds which are formed from their elements with absorption of heat and which, conversely, decompose into their elements with evolution of heat. The absorption of heat in their formation indicates that the atoms have little affinity for each other in the compounds and so may easily separate, and the heat generated by the decomposition, when it once begins, raises the mixture rapidly to a higher temperature and this in turn hastens the reaction till it becomes explosive. The explosion is further caused, of course, by the formation of a large volume of gas at a high temperature from a small volume of a liquid or solid. EXERCISES 1. The specific gravity of mercury at is 13.6. What is the weight of air above one square centimeter of the earth's surface at sea level ? 2. Sketch an apparatus suitable for the preparation of nitrogen by the use of hydrogen and copper oxide. 3. What is the weight of a gram molecular volume of air if it contains 21 per cent of oxygen, 0.9 per cent of argon and 78.1 per cent of nitro- gen ? What is the weight of one liter ? 4. If a gram molecular volume of the mixture of nitrogen peroxide and nitrogen tetroxide weighs 69 grams at 64, what is the per cent of each compound present ? 226 A TEXTBOOK OF CHEMISTRY 5. How many grams of nitric acid will be required to dissolve 5 grams of copper ? How many liters of nitric oxide will be formed ? 6. What weight of ammonia will be contained in one liter of aqua ammonia having a specific gravity of 0.90 and containing 28 per cent of ammonia ? How much water ? How many grams of ammonia would a liter of water take up in forming a solution of specific gravity 0.90 ? How many liters of the gas ? 7. How much salt, NaCl, will be required to furnish the hydrochloric acid necessary to neutralize 17 grams of ammonia? How much sul- f uric acid ? 8. If an isomeric nitrogen trichloride, NC1 3 , could be prepared in which the chlorine atoms were negative, what would be the products of its hydrolysis? CHAPTER XIII THE ATMOSPHERE. NOBLE GASES Determination of Oxgyen. The first determination of the per cent of oxygen in the atmosphere was made by Lavoisier, who heated a measured quantity of air with mercury as long as the oxygen continued to be absorbed (p. 19). He determined roughly the contraction which occurred and also showed that by heating the oxide of mercury formed he obtained a volume of oxygen closely agreeing with this contrac- tion. His results showed that approximately one fifth of the volume of air is oxygen. The determination may be made more accurately by measuring a volume of air in a eudiometer, exposing it to the action of a stick of phos- phorus, which is slowly oxidized to phospho- rous acid, H 3 PO 3 , and measuring the volume of the gases which remain (Fig. 75). Or a measured volume of air may be mixed with a little more than two fifths of its volume of hydrogen and the mixture exploded. One third of the contraction in volume (or, more accurately , see p. 68) will be the volume o.OOZL of the oxygen in the air taken. Composition of Air. Very many careful analyses of air have shown that when samples Fig. 75 are taken out of doors under usual conditions the composition of dry air agrees very closely with the fol- lowing : 227 228 A TEXTBOOK OF CHEMISTRY BY VOLUME BY WEIGHT Oxvsren 20 95 per cent 23.15 per cent Carbon dioxide 0.03 per cent 0.05 per cent Argon 94 per cent 1.3 per cent Nitrogen 78.08 per cent 75.5 per cent 100 100 The percentage of oxygen varies slightly, but is very rarely less than 20.9 or more than 21.0 per cent, out of doors. Air is a Mixture. Although the composition of air is nearly constant, it is believed that the oxygen and nitrogen which it contains are merely mixed together and not chemically com- bined, for the following reasons : 1. No heat is generated when oxygen, nitrogen, argon and carbon dioxide are mixed, and yet the mixture has all of the properties of air. 2. Liquefied air when it boils gives at first a gas richer in nitrogen and at last nearly pure oxygen. A compound when it boils away without decomposition has the same composition from first to last. 3. When water is shaken with air, it dissolves the oxygen and nitrogen in proportion to the solubility and partial pressure of each, and as oxygen is more soluble than nitrogen, the mixture of gases obtained by boiling the water contains about 35 per cent of oxygen, while air contains only 21 per cent. This is the conduct to be expected of a mixture rather than that of a com- pound. 4. A gram molecular volume of air weighs 28.95 grams. The simplest formula which would give a composition approximating that of air is N4O. The gram molecular volume of a compound of this formula would weigh 72 grams. On the other hand, a liter of air weighs 1.2928 grams, which is almost exactly what it should weigh if it is a mixture of gases in the proportion which has been given, as will be seen from the following calculation : THE ATMOSPHERE 229 WEIGHT or ONE LITER] Oxygen 1.429 X 0.2095 = 0.2994 Carbon dioxide 1.9768 X 0.0003 = 0.0006 Argon 1.7828 X 0.0094 = 0.0168 Nitrogen 1.2507 X 0.7808 = 0.9765 1.2933 Carbon Dioxide in the Air. The carbon dioxide in the air comes from four principal sources : 1. From the breath of men and animals, the carbon of food which is eaten being mostly converted into carbon dioxide by the oxidation processes which take place in the body. 2. From the burning of compounds containing carbon, such as wood, coal, oil or natural gas. 3. From the decay of animal and vegetable substances under the influence of bacteria. 4. From volcanoes and other sub- terranean sources. While the amount of carbon dioxide in the air is very small in comparison with the amounts of oxygen and nitrogen, the total amount is very great. Thus it is estimated that 1,300,000,000 tons of coal are burned annually, and this gives nearly three times its weight of carbon dioxide, but this would increase the amount in the air by only one six-hundredth part. 1 The carbon dioxide of the air is also a very important factor in the economy of nature, as it furnishes practically all of the carbon for the growth of plants. In a sense it is the constituent of the atmos- phere which is most vitally important for the life of plants, as oxygen is the constituent necessary for the life of animals. In another sense it may be said to be even more important, since it furnishes the most important constituent for the growth of plants, while oxygen furnishes to animals only the means with which to consume food which is secured from some other source. In utilizing carbon dioxide plants decompose it and exhale oxygen to the air. In this way the growth of plants prevents the accu- mulation of carbon dioxide in the atmosphere. The process of 1 A. Krogh, quoted byj F. W. Clarke in Data of Geochemistry, p. 42. See also Science, 1911, p. 757. 230 A TEXTBOOK OF CHEMISTRY reducing carbon dioxide to the compounds of carbon synthe- sized by the plant is endothermic, of course, and the necessary energy is furnished by the sunlight. When we burn coal in our furnaces, we make use of the energy of sunlight which was stored by growing plants millions of years ago. It has even been sug- gested that all of the oxygen of the air came originally from car- bon dioxide through this process of plant growth (Lord Kelvin). Besides the equilibrium maintained by the balance between the evolution of carbon dioxide from the sources named and the absorption of the gas by growing plants, the ocean plays a very important part in maintaining a constant amount in the air during long periods. As carbon dioxide dissolves in water in proportion to the partial pressure of the gas (p. 165), any increase in the amount of carbon dioxide in the air would be followed very quickly by an increase in the amount in the ocean, while any decrease would be replaced from the storehouse in the ocean. Since the amount of " free " carbon dioxide in the ocean to a depth of 5 kilometers is nearly fifty times 1 the amount in the air above it, and about three fourths of the earth's surface is covered by the ocean, the importance of the store contained in the ocean is obvious. Ventilation. Shortly after the discovery of the composition of the air by Priestly and Lavoisier, a method was devised for analyzing air by mixing it with nitric oxide, to combine with the oxygen, and then absorbing the nitrogen peroxide formed by means of a solution of potassium hydroxide. It will be readily understood that such a method requires very great care to secure accurate results and the early determinations led the observers to think that there was a considerable fluctuation in the amount of oxygen present and that this fluctuation caused the difference between good and bad air. But Cavendish was able to use even the nitric oxide method so accurately that he very soon showed that the variation in the composition of the atmosphere must be between very narrow limits, and this result has been confirmed 1 Calculated on the basis of 45 milligrams per liter of sea water. VENTILATION 231 by later observers. It was then discovered that carbon dioxide in mines and in wells or caves frequently killed persons exposed to its action, and for many years it was supposed that this gas acts as a positive poison and is the chief cause of danger in poorly ventilated rooms. This fallacy and also the opinion that carbon dioxide will accumulate near the floor of a room because the gas is one and a half times as heavy as air were spread so widely in semipopular literature and became so firmly fixed in the minds of many people that it has proved very difficult to correct these errors. It has been shown that the amount of carbon dioxide present in the air of even badly ventilated rooms is practically never great enough to cause any injury to human beings. It has been found very difficult to demonstrate clearly just what sub- stances cause the ill effects which follow from poor ventilation, and some recent authorities have spoken doubtfully of the standards for ventilation which have been proposed. There seems to be little doubt, however, that lack of ventilation in factories, offices and dwellings is a frequent cause of disease. It is also very well established that abundance of fresh air secured by life out of doors, both by night and day, combined with a nourishing diet, furnish the best hope of recovery from incipient tuberculosis. While exhaled carbon dioxide is not in itself harmful, it fur- nishes the best means of determining whether a room occupied by people is properly ventilated or not. The amount of the gas should not exceed 0.07 per cent by volume. To secure this amount of ventilation 55,000 liters or 2000 cubic feet of fresh air will be required each hour for each person in a room. 1 Moisture. Natural air always contains a certain amount of water vapor, but this is subject to very great variations, depen- dent on the temperature and the conditions to which the air has been subjected. The pressure of the water vapor can never much exceed the normal vapor pressure of water for the given 1 Roscoe and Schorlenmer, Treatise on Chemistry, I, 589. Another authority recommends 85,000 liters per hour. Stewart, Manual of Physiology, p. 244. 232 A TEXTBOOK OF CHEMISTRY temperature, and, indeed, can only exceed that when in a state of unstable equilibrium such that the introduction of suitable nuclei to form points of condensation will at once cause the for- mation of a cloud. When the pressure of the water vapor in the air corresponds to the normal pressure of water vapor for the given temperature, the air is said to be saturated, but such a condition does not usually obtain close to the earth's surface. At a height of a few hundred or thousand feet, however, owing partly to the mixing of warm, nearly saturated air with colder air currents, partly to the lowering of the temperature, which results from the adiabatic cooling of air as it expands in rising, saturation and condensation to clouds and rain take place. The amount of moisture in air may be determined : 1. By aspirating a known volume through weighed bulbs containing concentrated sulfuric acid. 2. By determining the dew point, that is, the temperature at which the air will deposit moisture on a cooled, polished, metallic surface or the temperature at which moisture will just disappear from such a surface. 3. By com- paring the temperature of the air with the temperature of a thermometer whose bulb is covered with moist cotton over which air is blown. Tables have been prepared giving the humidity corresponding to the difference observed. In general the humidity in rooms which are heated is too low for healthfulness and should be supplemented by artificial means. In many factories, especially in those for spinning and weaving, the degree of humidity is of vital importance to the success of the operations. Liquid Air. Critical Temperature. After it had been shown early in the nineteenth century that such gases as chlorine, ammonia, carbon dioxide, and sulfur dioxide could be liquefied, many attempts were made to liquefy oxygen and nitrogen, or air. These gases were subjected to pressures of several hundreds of atmospheres, but it was always found that the gas continued to fill completely and uniformly any space left to it, while if liquefied it should have separated into a liquid and a gaseous portion. Finally, in 1869, Andrews showed that carbon dioxide, CRITICAL TEMPERATURE 233 which can be liquefied under a pressure of 38.5 atmospheres at 0, or 71 atmospheres at 30, cannot be liquefied, even under pressures very much greater than this, at temperatures above 31. If a thick-walled, sealed glass tube containing liquid carbon diox- ide is warmed gently, at a temperature of 31.35 the liquid in the lower part of the tube will suddenly disappear and the gas will now fill the tube uniformly. The pressure may be increased or decreased, but as long as the temperature is above 31.35 no pres- sure either high or low can be found at which the carbon dioxide will separate into a liquid phase and a vapor phase. Below this temperature carbon dioxide will be partly liquid and partly gas, provided the pressure is equal to the vapor pressure of the liquid at the given temperature and the volume filled by the substance is large enough to allow a part to assume the vapor phase. The temperature above which a gas cannot be liquefied is called the critical temperature. Andrews' experiment made it seem very probable that the fail- ures to liquefy air were due to the fact that the critical tempera- tures of oxygen and nitrogen are much below ordinary tempera- ture. Following this suggestion, Cailletet in Paris and Pictet in Geneva (1877), working independently, both succeeded in liquefying oxygen by the use of cold and pressure combined with the cooling effect produced by the expansion of the highly com- pressed gas. Some years later it was shown by Joule and Thomson (Lord Kelvin) that a moderately compressed gas scarcely changes its temperature on expanding into a vacuum for instance, if air compressed to 20 atmospheres is allowed to expand into a vacu- ous receptacle, both receptacles being surroun'ded by water, the temperature scarcely changes, though for all gases except hydro- gen and helium, there is a slight cooling effect. This cooling effect increases for higher pressures, or when a gas is so far compressed that it no longer obeys Boyle's law (p. 35). It would seem that the attraction between the molecules of the gas has a greater effect as the molecules are brought closer together, causing the gas to contract more than it should in accordance with the law. 234 A TEXTBOOK OF CHEMISTRY When the gas expands from such a condition work must be done in overcoming this attraction between the molecules, and the expansion is accompanied by a cooling effect. On the basis of these facts Linde, Hampson and others have devised machines by means of which air can be readily liquefied in large quantities. _ In these machines air is compressed to 150-200 atmospheres and is then allowed to expand to atmospheric pressure in such a manner that the expanded and cooled air passes back over the tube in which the air is expanding. In the Hampson machine the air expands through a copper tube of about three millimeters in internal diameter and one hundred and thirty meters in length. This is wound in a spiral to secure compactness and the ex- panded air is compelled to follow the course of the spiral backwards, Fig. 76. By these machines a portion of the air is soon cooled to the point of liquefaction and the liquid air collects in a receptacle placed beneath the end of the spiral. The carbon dioxide must be removed from the air which is to be liquefied, by passing it through a large apparatus filled with slaked lime, and the moisture must also be removed by calcium chloride or some drying agent, as otherwise these would condense in solid form and stop up the tube through which the air expands. For the liquefaction of hydrogen the compressed gas must be cooled by liquid air, as it is only at low temperatures that hydrogen depart^ sufficiently from Boyle's law so that it can be liquefied by this method. Liquid nitrogen boils at 194, liquid oxygen at 182.5. Liquid air will contain, therefore, a larger proportion of oxygen than ordinary air, and by a sort of fractional distillation it is easy Fig. 76 LIQUID AIR. ARGON 235 to obtain from it a gas which contains from 75 to 95 per cent of oxygen. Such a gas is used for medicinal (e.g. in pneumonia) and some technical purposes. The method is also used to obtain nearly pure nitrogen and is now the most important industrial method for the preparation of both oxygen and nitrogen. For experimental purposes liquid air is kept in Dewar flasks (Fig. 77), double- walled flasks having the space between the two walls evacuated to prevent loss of heat by convection currents. The inner bulb is often silvered to cause it to reflect radiant Fig. 77 heat which reaches it from outside. Argon, A, 39.88. In 1785 Cavendish described an experiment in which he mixed air with an excess of oxygen, passed electric sparks through the mixture, and absorbed the oxides of nitrogen formed by a solution of potassium hydroxide. He then absorbed the rest of the oxygen by means of " liver of sulfur " and re- ported that the gas remaining unabsorbed was not more than T ffr of the original volume of the air. The real significance of this remarkable experiment was not understood for more than a century. During the eighties and nineties of the last century Lord Rayleigh spent a great deal of time in determining very accu- rately the density of the elementary gases, oxygen, hydrogen and nitrogen. In the course of his work he prepared what he sup- posed to be nitrogen by removing oxygen and all other known substances from the air. He also prepared nitrogen by the decomposition of ammonia. To his surprise a liter of the nitro- gen obtained from the air weighed about 6 milligrams more than a liter of nitrogen prepared from ammonia. Lord Rayleigh is a physicist, and he called in the assistance of a chemist, Sir William Ramsay, to solve the problem which was presented. Within a short time, in 1894, the two succeeded in preparing argon, partly by a repetition of the Cavendish experiment with 236 A TEXTBOOK OF CHEMISTRY modern appliances, partly by removing the nitrogen of the air by passing it over heated magnesium, with which the nitrogen combined. Argon was not only a new element, but it belongs to a wholly new class of elements, now called the Zero group of the Periodic System, or the noble gases. The most remarkable property of these elements is that none of them enters into chemical combina- tion with other elements or with itself the valence of the group is zero. Argon may be condensed to a liquid, which freezes at 188 and boils at 186.1. The gram molecular volume weighs 39.9 grams. From this the molecular weight is 39.88. Atomic Weight of Argon. Specific Heat of Gases. The fact that argon will not combine with any other element would, of itself, lead us to expect that the molecule of argon consists of a single atom and that the formula of the gas is A and the atomic weight 39.88. Another, wholly independent, line of evidence points to the same conclusion. The specific heat of a gas may be determined either while the volume of the gas remains constant or while the pressure remains constant. It is evident that the specific heat must be greater at constant pressure than at con- stant volume because at constant pressure the gas must expand as it grows warm and do work as it expands against the pressure of the atmosphere. It can be shown that, on the basis of the fundamental assumptions of the kinetic theory of gases, in any gas in which the energy required to increase the temperature of the gas is all used in increasing the average velocity of the mole- cules, the ratio of the specific heats must be : Specific heat at constant pressure _ 1.67 Specific heat at constant volume 1 On the other hand, if a part of the energy is used in causing the atoms within these molecules to vibrate more violently, the numerator of the fraction expressing the ratio between the two kinds of specific heat will be smaller, since both kinds of specific heat will be greater and an addition to both the numera- HELIUM 237 tor and denominator of any fraction causes it to approach unity. The ratio between the two specific heats can be calculated from the velocity of sound in the gas. The two kinds of specific heat have also been determined directly for air and some other gases. It has been found that the ratio of the specific heats for mercury vapor, for argon and for .some other gases is very close to 1.67/1 and it is believed that all of these gases are monatomic. The ratio of the specific heats for nitrogen is 1.41/1 ; for carbon dioxide it is 1.305/1 ; for ethylene, 1.26/1 ; and in general the numerator becomes smaller as the molecule is more complex. This seems to mean that in gases with complex molecules a con- siderable part of the energy used in heating the gas is absorbed in doing internal work in the molecules, that is, in causing their atoms to vibrate more and more rapidly. It will be readily seen that this conclusion gives a simple explanation of the fact that complex molecules are generally unstable at high temperatures. Helium, He, 3.99. In 1868 Lockyer observed some bright lines in the spectrum of the corona of the sun, which did not correspond to the lines of any element then known. He called the element which gives these lines helium (from ^Aios, the sun), and he had, in reality, discovered a new element, which for nearly thirty years was known to exist only in the sun, 90,000,- 000 miles away. Shortly after the discovery of argon it was recalled that Dr. Hillebrand of the U. S. Geological Survey had obtained a gas from the mineral uraninite. Ramsay, on further examination of the gases obtained from cleveite, a variety of uraninite, found in them not only a small amount of argon, but also a gas which gave the same spectral lines which had been observed in the light of the sun's corona, and he soon separated helium from the mixture. Helium is only twice as heavy as hydrogen and has the lowest boiling point of any known substance (unless we call the electron an element). It boils at 268.5 or at 4.5 absolute. Helium has acquired a very extraordinary interest, also, from the discovery that it is formed by the decomposition of radium. In spite of this method of 238 A TEXTBOOK OF CHEMISTRY formation, radium cannot be considered as a compound of helium, and no one has been able to induce helium to combine with any other element. Helium is found in all gases issuing from the earth. It is doubtless derived from radium and other radioactive elements. It has been suggested that the reason why only a very minute quantity of helium is found in the atmosphere is because, owing to the lightness of the helium atoms, their kinetic velocity is such that they may fly away from the earth into space. Neon, Krypton, Xenon, and Niton. The following partial table of atomic weights taken from the periodic system indicates that there should be three or four other elements belonging to the same family as helium and argon : He 4 Li 7 F 19 Na 23 Cl 35.5 A 39.9 K 39 Br 80 Rb 85.4 I 127 Cs 133 A systematic search for these elements soon led Ramsay to the discovery of neon (Ne = 20.2), krypton (Kr = 82.92) and xenon (X = 130.2) as constituents of the air, each of them pres- ent, however, in only very small amounts. Several years later it was shown that an evanescent element formed by the disinte- gration of radium belongs to this series. The density of the gas has been determined only approximately because of the minute quantity which it is possible to obtain. From this determina- tion the atomic weight is about 222.4. The element is called niton. It disintegrates spontaneously and very rapidly, one half of it disappearing in a little less than four days. EXERCISES 1. An adult eats food containing about 300 grams of carbon daily. If this is exhaled as carbon dioxide, CO 2 , at a temperature of 37, how many liters of the gas are exhaled per hour ? 2. If a person breathes 20 times per minute, 500 cc. of air being ex- haled at each respiration and the exhaled air contains 4 per cent of carbon dioxide, how many liters of the gas are exhaled per hour ? THE ATMOSPHERE. NOBLE GASES 239 3. How does the volume of carbon dioxide from a gas jet burning 3 cubic feet of gas per hour and giving an equal volume of carbon dioxide compare with that exhaled by an adult ? 4. If outside air contains 0.03 per cent of carbon dioxide, how often must the air in a room 5 meters square and 3 meters high be changed in order that the amount of carbon dioxide may not exceed 0.07 per cent when two persons, each breathing out carbon dioxide at the rate of 20 liters per hour, are present ? CHAPTER XIV PHOSPHORUS THE atomic weights of the nonmetallic elements of the fifth, sixth, seventh and zero groups of the periodic system and of the semimetallic elements of the fifth group are, in round numbers : FIFTH GROUP SIXTH GROUP SEVENTH GROUP ZERO GROUP He 4 N 14 O 16 F 19 Ne20 P 31 S 32 Cl 35.5 A 40 As 75" Se 78 Br80 Kr83 Sbl20 Te 127.6 I 127 Xe 130 Bi 208 Nt222 Phosphorus, P, 31.04. Occurrence. Phosphorus, the second element of the fifth group, is a very important element both for vegetable and animal life. It is an essential mineral constituent in soils for the growth of plants, and it is also an important ele- ment in the protoplasm of the cells and in the bones of animals. The ash left when the organic matter is burned out of bones con- sists very largely of calcium phosphate, Ca 3 (PO 4 )2- The same compound is found mixed with other substances in extensive de- posits of "phosphate rock" in North and South Carolina, Georgia, Florida and Tennessee. These deposits are extensively mined for use in applying to soils which are deficient in phos- phorus. Phosphorus is also found in the mineral apatite, Ca5(PO4)sF or CasCPO^aCl, which has already been mentioned in connection with fluorine. Phosphorus compounds are found in almost all iron ores, lessening their value when present in 240 PHOSPHORUS 241 more than very small amounts, because of the injurious effect of the phosphorus on the iron made from such ores. Preparation of Phosphorus. When a. mixture of sand (sili- con dioxide, SiO 2 ), calcium phosphate, Ca 3 (PO 4 ) 2 , and charcoal or coke, C, is heated to a very high temperature in an electric furnace, calcium silicate, CaSiOa, phosphorus, P4, and carbon monoxide, CO, are produced : 2 Ca 3 (PO 4 ) 2 + 6 SiO 2 + 10 C = 6 CaSiO 3 4- P 4 + 10 CO * The phosphorus distills from the retort in which the mixture is heated and is condensed and collected under water. This electrical furnace method for manufacturing phosphorus has dis- placed older, more complicated methods, in comparatively recent times. Allotropic Forms of Phosphorus. The phosphorus obtained as described is a waxlike solid which usually has a slight yellow color and this form is called " ordinary " or " yellow " phos- phorus. When pure it melts at 44.5 and can be readily melted and cast into sticks under water. Its specific gravity is 1.8232 at 20. It boils at 290. In the gaseous form a gram molecular volume weighs about 124 grams, from which the formula must beP 4 . Ordinary phosphorus glows with a pale light when exposed to moist air. It may be distilled with steam, and a very minute quantity may be detected in a dark room by the use of these properties. The word phosphorescence recalls, of course, the luminous quality of the element. If ordinary phosphorus is heated to 240-250 in a closed vessel, it is gradually, though not quite completely, transformed into the allo tropic variety called red phosphorus. This was formerly called amorphous phosphorus, but it may be crystallized from solution in melted lead. When pure and free from yellow 1 In writing this equation notice that two molecules of calcium phosphate, Ca 3 (PO 4 ) 2 , are required to give one molecule of phos- phorus, P 4 . The rest of the equation follows logically from the formulas of the product formed. 242 A TEXTBOOK OF CHEMISTRY phosphorus its specific gravity is 2.34. It is not poisonous, while yellow phosphorous is very poisonous indeed. Yellow phosphorus dissolves readily in carbon disulfide, red phosphorus does not. Yellow phosphorus must be kept away from the air and is usually kept under water because of the very low kindling temperature. Red phosphorus takes fire at a much higher tem- perature and may be kept in open bottles. When heated to a high temperature, red phosphorus distills and goes back to the yellow form, but at lower temperatures the vapor pressure of red phosphorus is much lower than that of the yellow variety. The molecular weight of red phosphorus has not been deter- mined. Matches. The methods of obtaining fire in use before the nine- teenth century were difficult of application and people often sent to their neighbors even at some distance for coals rather than to take the trouble of starting a new fire. Phosphorus was dis- covered, it is true, in 1669 by Brandt, an alchemist of Hamburg, but it was not till 1827 that use was made of its low kindling temperature for the preparation of matches. For the first matches using phosphorus the match sticks were dipped in melted sulfur and then in a mixture of phosphorus and glue or some other adhesive substance. When dry a slight friction raises the phosphorus to its kindling temperature and this, as it burns, sets fire to the sulfur, which, in turn, ignites the wood. In the later manufacture the sulfur was replaced by other com- bustible substances which do not give an objectionable odor, and the kindling power of the phosphorus was reenforced by potas- sium chlorate, red lead or other oxidizing compounds. Ordinary phosphorus is extremely poisonous, however, and gives off enough vapor at ordinary temperatures so that, unless extraordinary pains are taken to ventilate the factories, the workmen often suffer from a very painful and fatal disease, which causes necro- sis of the jaw. Partly for this reason and partly to avoid the danger of accidental fires, most European countries have for- bidden the sale or even the manufacture of matches containing ordinary phosphorus. The " safety " matches used in these PHOSPHINE 243 countries have on their heads, usually, a mixture of antimony trisulfide, potassium chlorate and glue, and they are ignited on a prepared surface of red phosphorus, glue and a sulfide of antimony. In comparatively recent times it has been discovered that tetraphosphorus trisulfide, P4S 3 , may be substituted for yellow phosphorus in ordinary matches. As it does not give off poison- ous vapors, this sulfide of phosphorus ought soon to entirely displace the ordinary phosphorus for this manufacture. A law passed by Congress in 1912 will prevent the further use of ordi- nary phosphorus for matches in the United States. Phosphine, PH 3 . When yellow phosphorus is warmed with a strong solution of sodium hydroxide, it is oxidized to sodium hypophosphite, NaH^PC^. At the same time some of the hy- drogen of the water or of the sodium hydroxide combines with more of the phosphorus to form phosphine, PH 3 . In preparing the gas a rather small flask should be used, and it is well to add to the contents of the flask, before warming, a few drops of ether, which will expel the air and prevent a possible explosion. The phosphine prepared in this manner contains some hydrogen and some of the liquid hydrogen phosphide, P2H 4 , which corres- ponds in composition to hydrazine, N 2 H4. This liquid is volatile and takes fire spontaneously on exposure to the air. For this reason, although the kindling temperature of phosphine is about 150, the phosphine prepared as described takes fire at once as it comes to the air. Bubbles of the gas explode as they come to the surface of the water, forming a cloud of phosphoric acid, H 3 PO 4 , which gives beautiful vortex rings in still air (Fig. 78). Phosphine may be condensed to a colorless liquid, which boils at 86.2 and solidifies at lower temperatures to crystals which melt at - 133. Phosphonium Salts. Phosphine combines with acids to form phosphonium salts, as ammonia forms ammonium salts. The most stable and best known of these salts is phosphonium iodide, PH 4 I, which may be prepared by the direct union of phosphine, PH 3 , and hydriodic acid, HI. It forms white crystals, which 244 A TEXTBOOK OF CHEMISTRY sublime at 80. The salt is hydrolyzed by water and phosphine escapes from the solution : PH 4 I + HOH = PH 4 OH + HI PH 4 OH = PH 3 + HOH Evidently the phosphonium group, PH 4 , is very unstable, even in the presence of hydrogen ions. It is sometimes stated that phosphine is a much weaker base than ammonia. Correctly speaking neither is a base, and the true base, phosphonium hydroxide, PH 4 OH, is extremely un- stable, if it exists at all. We shall find that arsine, AsH 3 , and stibine, SbH 3 , do not com- bine with acids. In the series NH 3 , PH 3 , AsH 3 , SbH 3 , not only does the tendency to com- bine with a fourth hydrogen atom be* come less and less, but the compounds themselves are less and less stable, stibine, SbH 3 , decomposing at ordinary temperatures, especially in the presence of metallic antimony. Phosphorus Trichloride, PC1 3 , and Phosphorus Pentachloride, PCls, are easily prepared by the direct union of chlorine and phosphorus. The trichloride is a liquid which boils at 76. The pentachloride is a white solid which melts in a sealed tube at 148. Its vapor pressure, however, is 760 mm. at 140. In other words its melting point is higher than its boiling point and it sublimes without melting when heated under atmospheric pressure. Fig. 78 CHLORIDES OF PHOSPHORUS 245 The molecular weight of phosphorus pentachloride, PC1 5 , is 208.5 ; but a gram molecular volume of the gas at 182 weighs 147 grams, while at 300 it weighs only 105.7 grams, only a little more than one half the weight of a gram molecule. This indi- cates that the pentachloride dissociates into phosphorus trichlo- ride and chlorine, the dissociation being nearly complete at 300. This gives twice as many molecules as there are in the original pentachloride and one gram molecule of the pentachloride gives two gram molecular volumes of gas : PC1 5 ^ PC1 3 + C1 2 Hydrolysis of the Chlorides of Phosphorus. The chlorides of phosphorus are decomposed, or hydrolyzed by water in the same manner as most other chlorides of nonmetallic elements : PC1 3 + 3H.OH = H 3 P0 3 + 3HC1 Phosphorous Acid PC1 6 + 4H.OH = H 3 PO 4 + 5HC1 Phosphoric Acid Phosphorus Oxychloride, POC1 3 . When phosphorus penta- chloride is treated with a small amount of water or with almost any compound containing the hydroxyl group, OH, it is changed to phosphorus oxychloride, POC1 3 , while the two chlorine atoms which are lost combine with the two atoms which were united to the oxygen : cl x / .P^CI + HOH = c/ x ci x ci . an Ethyl Alcohol Ethyl Chloride a a a ;p^a + c 2 H 4 o 2 = Cl' X C1 (orC 2 H 3 O.OH) Acetic Acid 246 A TEXTBOOK OF CHEMISTRY Phosphorus oxychloride may also be prepared by oxidizing phosphorus trichloride with potassium chlorate. It is a color- less liquid which boils at 107.2. It has a very unpleasant odor and the vapor attacks the eyes strongly. It is, of course, hy- drolyzed by water to phosphoric and hydrochloric acids. It is the chloride of phosphoric acid, H 3 PC>4, in the same sense that sulfuryl chloride, SC^Ck, is the chloride of sulfuric acid, H2S04 (p. 189). Oxides of Phosphorus. When phosphorus is burned with an insufficient supply of air, a mixture of two oxides, phosphorus " trioxide," P^e, and phosphorus " pentoxide," P4Oio, is formed. The names were given long before determinations of the density of the vapors of these compounds showed that they have the formulas given. The names refer, of course, to the simple formu- las P 2 O 3 and P2O 6 . The trioxide, PA, is a solid which melts at 22.5 and boils at 173.1. The pentoxide is a solid which may be sublimed at a high temperature but which gives off almost no vapor at ordinary temperatures. It has a very strong affinity for water and is the most perfect drying agent for gases which we have. If 10,000 liters of air are passed through a comparatively small tube filled with the pentoxide, no moisture which can be determined remains in the gas, while the vapor of the pentoxide which is carried away by the gas weighs only one milligram (Mor- ley, Am. J. Sci. 34, 199 (1887) ; J. Am. Chem. Soc.26, 1171 (1904). Another oxide of phosphorus, ?2O4, called phosphorus tetrox- ide, is known, but is of little interest except as corresponding to nitrogen tetroxide, N2O4. Acids of Phosphorus. While the acids of nitrogen, nitrou acid, HNO 2 , and nitric acic^ HNO 3 , are formed by the addition of one molecule of water to the anhydrides, N 2 O 3 and N 2 O 5 , the normal acids of phosphorus corresponding to these are formed by the addition of three molecules of water to the cor- responding anhydrides (using the simpler formulas). Two other acids, metaphosphoric acid, HPO 3 , and pyrophosphoric acid, H 4 P 2 O 7 , are also derived from phosphorus pentoxide (phosphoric anhydride), P 2 O5. As the addition of water is considered as ACIDS OF PHOSPHORUS 247 neither an oxidation nor a reduction, the three acids derived from the pentoxide are all called " phosphoric " acids and are distinguished by prefixes. These relations will be clearer from the following table : (P 2 O 1 .3H 2 O)= 2 H 3 PO 2 Hypophosphorous acid (P 2 O 3 .3 H 2 O) = 2 H 3 PO 3 Phosphorous acid (P 2 O 6 .3 H 2 O) = 2 H 3 PO 4 Orthophosphoric acid ' (P 2 O 5 .2 H 2 O) = H 4 P 2 O 7 Pyrophosphoric acid (P 2 O 5 .H 2 O) = 2 HPO 3 Metaphosphoric acid Basicity of the Acids of Phosphorus. The formulas of hypo- phosphorous acid, H 3 PO 2 , phosphorous acid, H 3 PO 3 , and ortho- phosphoric acid, H 3 PO 4 , might lead us to expect each of these acids to be tribasic. It is found, however, that only one atom of hydrogen in hypophosphorous acid can be replaced by metals and only two of the hydrogen atoms in phosphorous acid. In 1 other words hypophosphorous acid is monobasic, phosphorous acid dibasic, and orthophosphoric acid tribasic. The normal sodium salts are : I Sodium hypophosphite, NaH 2 PO 2 Sodium phosphite, Na 2 HPO 3 Sodium orthophosphate, Na 3 PO 4 These and other facts, which cannot be given here, make it probable that the structure of these acids is correctly represented by the following formulas : Hypophosphorous acid H- X)-H H-0 6 Phosphorous acid /P / W H-0 O i Orthophosphoric acid /P\ KO' 1 This oxide is given in many of the books, but its existence is extremely doubtful. 248 A TEXTBOOK OF CHEMISTRY According to .these formulas only the hydrogen atoms which are united to oxygen are acid in character. Also the oxidation of the lower acids consists in the introduction of an oxygen atom between a hydrogen atom and the phosphorus. According to the electron theory the valence of the phosphorus is negative toward the hydrogen atoms and positive towards the oxygen. Oxidation, in such a case, consists in the change of a negative valence to a positive one. Hypophosphorous Acid, H 3 PO 2 . The sodium salt of hypo- phosphorous acid, NaH 2 PO2, is formed when phosphorus is warmed with a solution of sodium hydroxide, phosphine, PH 3 , being evolved at the same time. The acid is monobasic. It is a powerful reducing agent. Some hypophosphites are used in medicine. Phosphorous Acid, H 3 PO 3 , is formed with phosphoric and hypophosphoric acids, when ordinary phosphorus is allowed to oxidize slowly in moist air, but it is extremely difficult to separate the mixture into its components. The pure acid may be pre- pared by the hydrolysis of phosphorus trichloride. It is a bi- basic acid, the two sodium salts being monosodium phosphite, NaH 2 PO 3 , and disodium phosphite, Na 2 HP0 3 . Phosphorous acid is also a powerful reducing agent. Orthophosphoric Acid, H 3 PC>4, is formed when phosphoric anhydride, P4Oio, is dissolved in hot water. It is also formed when solutions of pyrophosphoric acid, H4P 2 O7, or metaphos- phoric acid, HPO 3 , are boiled, especially if some strong acid, as nitric acid or hydrochloric acid, is present to catalyze the reac- tion, which is to be considered as a hydrolysis : OH OH + HOH Pyrophosphoric Acid Orthophosphoric Acid (2 mols) PHOSPHORIC ACIDS 249 HOH = O= X OH Metaphosphoric Acid Pure orthophosphoric acid forms clear, rhombic crystals, which melt at about 40. These crystals dissolve in a small amount of water, forming a heavy, sirupy liquid somewhat resembling concentrated sulfuric acid in appearance. An impure solution of phosphoric acid was formerly prepared on a large scale, as a step in the manufacture of phosphorus, by treating bone ash with dilute sulfuric acid and filtering the solu- tion from the calcium sulfate, which is only slightly soluble in water : Ca 3 (PO 4 )2 + 3 H 2 SO 4 = 2 H 3 PO 4 + 3 CaSO 4 Tricalcium Calcium Phosphate Sulfate Orthophosphoric acid forms three classes of salts, in which one, two or three atoms of hydrogen are replaced in each mole- cule of the acid. These are called primary, secondary and ter- tiary, or, more often, mono-, di- and tri-metallic salts. The following are the names of the sodium and calcium salts : Monosodium phosphate, NaH 2 PO 4 (primary) Monocalcium phosphate, Ca(H 2 PO 4 ) 2 (primary) Disodium phosphate, Na 2 HPO 4 (secondary) Dicalcium phosphate, CaHPO 4 (secondary) Trisodium phosphate, Na 3 PO 4 (tertiary) Tricalcium phosphate, Cas(PO 4 ) 2 (tertiary) Orthophosphoric acid is much the most important of the acids of phosphorus, being the acid into which all of the others tend to pass either by oxidation or hydrolysis. Apart from its oc- currence in organic compounds, phosphorus is found almost exclusively in the form of orthophosphates, and these phosphates are an indispensable constituent of arable soils. 250 A TEXTBOOK OF CHEMISTRY lonization of Orthophosphoric Acid. Orthophosphoric acid is an acid of only moderate strength. A T V formular (i.e. con- taining one gram molecule in 10 liters of water) solution contains, of course, three times as many replaceable hydrogen atoms as a tenth normal solution of hydrochloric acid, but it contains only one third as many hydrogen ions. This means that in the ioniza- tion reaction : the equilibrium is comparatively far to the left, even in quite dilute solutions. Even in very dilute solutions the second and third hydrogen atoms ionize to only a very slight extent. This may be either because the three hydrogen atoms in Orthophos- phoric acid are different or because after the removal of the one hydrogen atom the negative ion, H 2 PO 4 ~, holds the remaining hydrogen atoms too strongly for them to separate easily as ions. The second explanation seems more probable. If the first hy- drogen atom is completely neutralized by the addition of a base : H + + H 2 P0 4 - + Na + + OH- ^ Na + + H 2 PO 4 - + H 2 O the dihydrogen phosphate ion, H 2 PO 4 ~, will ionize to a slight extent : H 2 P0 4 ~ ^. but the solution is only faintly acid, and if more sodium hydroxide is added to 'such a solution, the accumulation of the monohy- drogen phosphate ions, HPO 4 = , shifts the equilibrium to the left. This is because sodium salts of weak acids are always much more completely ionized than the corresponding acids. This shifting of the equilibrium prevents much formation of new hydrogen ions, as those which are present are removed by com- bination with the hydroxide ions of the sodium hydroxide. Before all of the second hydrogen atoms of the phosphoric acid have been neutralized, the tendency of the monohydrogen phos- phate ions, HPO 4 = , to combine with hydrogen ions will become so strong that even the hydrogen ions of water will combine with them, leaving an excess of hydroxide ions in the solution. Such a solution must, of course, react alkaline. From this conduct IONIZATION OF PHOSPHORIC ACID 251 of phosphoric acid it is evident that if we attempt to titrate a solution of phosphoric acid by adding sodium hydroxide, instead of the sharp change which occurs in titrating hydrochloric or sulfuric acid, there will be a gradual change from a solution con- taining a slight excess of hydrogen ions, H + , to one containing a slight excess of hydroxide ions, OH~. Two things result from these properties of solutions containing salts of phosphoric acid : first, unless a very sensitive indicator is chosen, that is, one in which the change in color is produced by a very slight change in the ratio between the hydrogen and hydroxide ions present, the end point of the titration will be indefinite ; and, second, since most indicators change color, not when the number of hydrogen, H + , and hydroxide, OH~, are equal, but when there is an excess of one or the other, and this excess differs for different indicators, the end point in titrating phosphoric acid will depend on the indicator chosen. (See p. 387.) Thus, if methyl orange or cochineal is used, the end point in fairly dilute solutions will be found when the solution corresponds very nearly to the composition NaH 2 PO 4 . With phenolphthalein, on the other hand, the end will correspond nearly to the composition Na 2 HPO 4 . With litmus the end lies between the two. If alizarine green is used, the change in color occurs when the composition of the solution is very nearly represented by the formula Na 3 PO 4 . It is well, also to consider the conduct of disodium phosphate from a somewhat different point of view, which, however, follows logically from what has been said. If the salt, which crystallizes with the composition Na 2 HPO 4 .12 H 2 O, is dissolved in water, we should expect the formation of the ions,- Na + + Na + + HPO 4 = . But, as has been stated, in the presence of many of the mono- hydrogen phosphate ions, HPO 4 = , these have a tendency to form, with the hydrogen ions of the water, dihydrogen phosphate ions, H 2 PO 4 ~, because the latter ionize only to a slight extent. This results in the presence of an excess of hydroxide ions in the solution, which will react alkaline toward indicators that are sensitive to a slight excess of hydroxide ions. 252 A TEXTBOOK OF CHEMISTRY Na + + Na + + HPO 4 = + H + Water in Ionic Form ^ Na + + Na + + H 2 PO 4 This sort of hydrolysis occurs with all salts of strong bases, as sodium hydroxide or potassium hydroxide, with weak acids or with acids whose second or third hydrogen atoms undergo slight ionization. Trisodium phosphate, Na 3 PO4, will, of course, be much more completely hydrolyzed : Na + + Na + + Na + + PO 4 ^ + H + + OH~ ^t 3 Na + + HP0 4 = + OH- The only tertiary salts of orthophosphoric acid which are sol- uble in water are those of the alkali metals, sodium, potassium, etc. All other tertiary or normal phosphates are insoluble. Many of the primary and secondary phosphates are either insoluble or are decomposed by water into phosphates which approach the tertiary phosphates in composition, and either phosphoric acid or more acid phosphates, which dissolve in an excess of the acid. Decomposition of Primary and Secondary Salts of Ortho- phosphoric Acid. Salts of orthophosphoric acid which contain hydrogen decompose on heating, losing all of their hydrogen as water and leaving salts of metaphosphoric or pyrophosphoric acid. As ammonium salts dissociate on heating, these give the same products as if they contained hydrogen in place of ammo- nium, NH 4 . Monosodium phosphate, NaH 2 PO 4 , gives sodium metaphosphate, NaPO 3 ; and sodium ammonium phosphate, NaNH 4 HPO 4 , gives the same compound. Disodium phosphate, Na 2 HPO 4 , gives sodium pyrophosphate, Na 4 P 2 O 7 ; and am- monium magnesium phosphate, NH 4 MgPO 4 , gives magnesium pyrophosphate, Mg 2 P 2 O 7 . Magnesium diammonium phosphate, Mg(NH 4 ) 4 (PO 4 ) 2 , or Mg [(NH 4 ) 2 PO 4 ] 2 , gives magnesium meta- phosphate, Mg(PO 3 ) 2 . PHOSPHORIC ACIDS 253 Pyrophosphoric Acid, H4P 2 O7. If orthophosphoric acid is heated carefully at 250 it loses water and is changed to pyro- phosphoric acid : 2H 3 PO 4 -H 2 = H 4 P 2 O 7 The acid may be dissolved in cold water, giving a solution which differs in its properties from those of a solution of orthophos- phoric acid. Especially, after neutralization it gives with silver nitrate, AgNOs, a white precipitate of silver pyrophosphate, Ag4P2O7, while orthophosphoric acid, or orthophosphates, will give a yellow precipitate of trisilver phosphate, AgsPC^. Sodium pyrophosphate is easily prepared by heating disodium phosphate, Na 2 HPO4. Metaphosphoric Acid, HPO 3 , is formed when phosphoric anhydride, P4Oio, is dissolved in cold water or when either ortho- phosphoric acid or pyrophosphoric acid is heated to a high tem- perature. It differs from the other two phosphoric acids in that its neutralized solution gives with silver nitrate a white pre- cipitate of silver metaphosphate, AgPO 3 , instead of the yellow precipitate of trisilver phosphate, Ag 3 PO 4 , given by ortho- phosphoric acid and the white precipitate of silver pyrophos- phate, Ag 4 P 2 O 7 , given by pyrophosphoric acid. Metaphos- phoric acid also precipitates a solution of albumin, as of the white of an egg, which has been acidified with acetic acid, while ortho- and pyrophosphoric acids or their salts do not do this. Sodium metaphosphate may be prepared by heating either monosodium phosphate, NaH 2 PO4, or sodium ammonium phos- phate, NaNH 4 HPO 4 . This last salt is called microcosmic salt and is used in blowpipe analysis. When this salt is heated in a loop of platinum wire, it melts to a clear bead of sodium meta- phosphate, NaPO 3 , which will dissolve the oxides of many of the metals, forming double salts of orthophosphoric acid : NaPO 3 + CuO = NaCuPO 4 Sodium Copper Orthophosphate 254 A TEXTBOOK OF CHEMISTRY The metaphosphate may be considered here as, in a certain sense, an acid anhydride which with oxides forms normal salts of orthosphoric acid. The copper sodium phosphate is blue, and the colors given to the microcosmic bead by different metallic oxides serve as a means for their identification. * Metaphosphoric acid may be vaporized at a high tempera- ture and the vapor has the formula (HPOs^. A study of the salts and of the properties of solutions of the acid prepared in different ways has shown that several polymeric forms of the acid exist, that is, forms having the same composition but dif- ferent molecular weights. The salts of these various forms are called dimetaphosphates, M 2 P 2 Oe, trimetaphosphates, M 3 P 3 O 9 , tetrametaphosphates, M 4 P 4 Oi 2 , etc. In these formulas " M " is used to represent any univalent metal. * Hypophosphoric Acid, H 2 PO 3 , is one of the products formed by the slow oxidation of phosphorus in moist air. From the solution obtained in this manner the rather difficultly soluble acid sodium salt, NaHPO 3 , is precipitated by a concentrated solution of sodium acetate, NaC 2 H 3 O 2 . From its formula we should expect that phosphorus tetroxide, P 2 O 4 , would be the anhydride of hypophosphoric acid, but, curiously enough, when phosphorus textroxide is dissolved in water, a mixture of phosphorous and orthophosphoric acids is formed : P 2 O 4 + 3 H 2 O = H 3 PO 3 + H 3 P0 4 The true anhydride of hypophosphoric acid (PO 2 ?) has not been prepared. * Sulfides of Phosphorus. Four sulfides of phosphorus have been prepared, tetraphosphorus trisulfide, P 4 S 3 , tetraphosphorus heptasulfide, P 4 S 7 , triphosphorus hexasulfide, P 3 S 6 , and diphos- phorus pentasulfide, P 2 S 5 . The last is usually called phosphorus pentasulfide. It melts at 274-276 and boils at 530. It has been used in chemical laboratories frequently to obtain a nearly constant, high temperature. Tetraphosphorus trisulfide, P 4 S 3 , melts at 165-166 and boils PHOSPHORUS 255 at 225-235 under a pressure of 10 mm. As it takes fire with slight friction and as its vapors are either nonpoisonous, or, in any case, far less poisonous than those of ordinary phosphorus, it is likely to replace the latter entirely for the manufacture of matches. EXERCISES 1. Write the equation for the reaction between phosphorus and a solution of sodium hydroxide, giving hydrogen and sodium hypophos- phite. 2. Write the equation for the reaction between phosphorus and sodium hydroxide in solution, giving phosphine and sodium hypophos- phite. 3. Write the equation for the reaction giving liquid hydrogen phos- phide, P2Hj, and sodium hypophosphite. 4. What is the distinction between a substance which sublimes and one which boils ? Under what conditions does water sublime ? 5. What percent of phosphorus pentachloride is dissociated when its gram molecular volume weighs 156.4 grams ? What per cent when it weighs 130 grams ? 6. If on heating phosphorous and hypophosphorous acids the prod- ucts formed are phosphine, metaphosphoric acid and water, what are the equations representing the decomposition of these acids? Are these decompositions consistent with the structural formulas which have been proposed for these acids ? 7. If the structure of phosphorous acid were correctly represented by X) H the formula P\-O H, how ought it to decompose on heating ? X 0-H 8. Metaphosphoric acid volatilizes at a very much higher tempera- ture than sulfuric acid. What will be the effect of heating a mixture of sodium sulf ate and metaphosphoric acid ? 9. How much iodine and water will be required to convert 10 grams of phosphorus into orthophosphoric acid if the reaction is quantitative ? 10. How many liters of air will be required to burn 10 grams of phosphorus to the pentoxide ? CHAPTER XV ARSENIC, ANTIMONY AND BISMUTH IT has been pointed out that with increasing atomic weights the elements of the nonmetallic groups of the Periodic System become more metallic in character. This is especially evident in the fifth group. Arsenic is metallic in its appearance, opaque and like steel on its surface when not tarnished. It is, however, brittle, and its chloride, AsCl 3 , is hydrolyzed by water, resembling the chlorides of the nonmetals rather than those of the metals. Arsenic forms no salts with sulfuric, nitric or other acids. Anti- mony and bismuth are still more metallic in their appearance and bismuth is malleable to a slight extent. Their chlorides are hydrolyzed by water, at first, only to the oxychlorides, SbOCl and BiOCl. Both of them form normal nitrates, Sb(NO 3 ) 3 , and Bi(NO 3 ) 3 , and sulfates, Sb 2 (SO 4 ) 3 , Bi(SO 4 ) 3 , though these are hydrolyzed to basic salts or even to the hydroxides or oxides by water. Arsenic, As, 74.96. Occurrence. Arsenic is found in the free state in nature, but occurs chiefly combined with sulfur, either alone, as in the disulfide, realgar, As 2 S 2 , or tiie trisulfide, orpiment, As 2 S 3 , or, much more frequently, with the sulfides of other metals, the most common compound of this kind being arseno- pyrite, or mispickel, FeAsS. Iron pyrites and copper pyrites almost invariably contain arsenic, often in considerable quan- tities. From the former the arsenic finds its way into commer- cial sulfuric acid and from that into a great variety of chemical products. From the copper pyrites the arsenic escapes along with the sulfur dioxide in the process of roasting, no less than twenty-five tons a day of arsenic trioxide escaping from a single smelting furnace in Montana (J. Am. Chem. Soc. 29, 993 (1907). 256 ARSENIC 257 Preparation and Properties of Arsenic. Metallic arsenic is usually prepared by heating arsenopyrite, FeAsS, the arsenic subliming and leaving ferrous sulfide, FeS, behind. Prepared in this manner it forms a dark gray, brittle mass. Fragments heated in a closed tube or before the blowpipe on charcoal, so that the tarnished surface is removed, appear like steel. When deposited on a glass or porcelain surface (Marsh's test), arsenic is brown or black according to the thickness of the deposit, usually showing brown at the edges where the deposit is thin, while antimony is a more sooty black. Here, again, we have an increase in metallic properties with increasing atomic weight, opacity being one of the most marked properties of metals. The specific gravity of gray arsenic is 5.73. The formula of its vapor at 560-670 is As 4 , at 1700, As 2 . The melting point of arsenic is higher than its boiling point, hence it sublimes with- out melting when heated on charcoal or in a tube closed at one end, a property which distinguishes it easily from antimony. Metallic arsenic is sometimes used for poisonous fly papers. Three one-hundredths of a per cent of arsenic lowers the con- ductivity of copper 14 per cent and injures it seriously, especially for electrical use. Arsine, AsH 3 . Marsh's Test. When almost any soluble compound of arsenic is added to a flask in which hydrogen is being generated from zinc and sulfuric or hy- drochloric acid, the arsenic is reduced to arsine. If the hydrogen containing arsine is con- veyed through a hard glass tube, narrowed at one point (Fig. 79), and Fig. 79 the tube is heated just back of the constriction with a Bunsen flame, the arsine is decomposed and metallic arsenic is deposited as a brown or black mirror on the glass. As small a quantity of arsenic as ytfW f a milligram can be seen in this form, 258 A TEXTBOOK OF CHEMISTRY and the process has been long used, under the name of Marsh's test, for the detection and estimation of small quantities of arsenic, especially in cases of criminal poisoning, or for the examination of wall papers or articles of food. One of the first requisites in making the test is, of course, that the zinc, sulfuric acid and other materials used should be entirely free from arsenic. Commercial zinc and commercial sulfuric acid almost invariably contain the element. It is necessary, also, to distinguish the mirror from that of antimony, which may be obtained in the same manner. When the amount of arsenic is considerable, it imparts to the burning hydrogen flame a pale blue color, and arsenic is deposited on a piece of porcelain held in the flame, very much as soot is deposited from a candle flame. Arsine may be condensed to a liquid, which boils at 55. It is very poisonous. Some years ago a chemist in Chile was fatally poisoned while working with it. Arsine does not show any tendency to combine with acids, as ammonia and phosphine do. Arsenic " Trioxide," As 4 O 6 . When arsenic is heated in the air, it burns to arsenic " trioxide," frequently called white arsenic. The simpler formula, As2Os, is commonly used for the compound, but the density of its vapor corresponds to the formula As^e. It crystallizes in octahedra which are highly characteristic, and the microscopic identification of the crystals is one of the' most important means of demonstrating the pres- ence of arsenic. Arsenic trioxide is one of the most common compounds of the element and is frequently used as a ratsbane and has often been used for criminal poisoning. The fatal dose for an adult is from 0.06 to 0.18 gram (one to three grains), but it seems possible to accustom the organism to its use, and the so-called arsenic eaters may sometimes take four times that amount without apparent injury. The best antidote is freshly precipitated ferric hydroxide, Fe(OH)3, or a colloidal solution of ferric hydroxide. ARSENIC 259 Crystallized arsenic trioxide dissolves in 50 parts of water at 25. The amorphous form is somewhat more soluble. The solution reacts faintly acid, and forms salts with bases, but on evaporation it deposits the trioxide. Arsenious Acid. As has just been stated, arsenious acid resembles sulfurous and nitrous acids in that it exists only in solution and decomposes easily into its anhydride and water. Salts of acids derived from this anhydride are known, however. Among these may be mentioned silver orthoarsenite, Ag 3 AsO 3 , and monopotassium diarsenite, KHAS2O4. The last seems to be derived from a diarsenious acid, H2As2O 4 , which would cor- respond to a doubled nitrous acid, (HNO2)2- Paris green is a double salt of copper with acetic and arsenious acids, Cu(C 2 H 3 2 ) 2 .Cu 3 (As0 3 ) 2 . Arsenic Pentoxide, As2O5, and Arsenic Acid, H 3 AsO 4 . When arsenic trioxide is warmed with nitric acid, it is oxidized to arsenic acid : 2 HNO 3 + As 2 O 3 + 2 H 2 = 2 H 3 AsO 4 + NO + NO 2 From a concentrated solution the acid crystallizes with one molecule of water, H 3 AsO 4 .H 2 O. At 140-180 this hydrate loses water and gives pyroarsenic acid, H 4 As2C>7, and at 200 the latter loses more water and gives metarsenic acid, HAsO 3 . At a higher temperature metarsenic acid loses more water and arsenic pentoxide, As2Os, remains. This cannot be volatilized without decomposition, and its true molecular weight is not known. Salts of pyroarsenic and metarsenic acids may also be prepared by heating secondary and primary salts of arsenic acid, but the acids are not known in solution, as they are hydro- lyzed by water at once to orthoarsenic acid. Trisilver arsenate, Ag 3 AsO 4 , is reddish brown and insoluble ; and the white, crystal- line magnesium ammonium arsenate, MgNH 4 AsO4, is also insoluble, closely resembling the corresponding phosphate. Arsenic acid is an oxidizing agent in concentrated solution, liberating chlorine from hydrochloric acid, while it is itself reduced to arsenious acid or oxide. But the action is reversible, 260 A TEXTBOOK OF CHEMISTRY and the reverse effect will occur in dilute solutions, chlorine oxidizing arsenious acid to arsenic acid : H 3 As0 4 + 2 HC1 ^ H 3 AsO 3 + C1 2 In a neutral or faintly alkaline solution the equilibrium is so far toward the formation of the arsenate that the oxidation by iodine, even, is practically quantitative and is used for the standardization of iodine solutions : Na 3 AsO 3 + 1 2 + 2 NaHCO 3 = Na 3 AsO 4 + 2 Nal + H 2 O + 2 CO 2 Arsenic Trichloride, AsCl 3 , may be prepared by the direct union of arsenic and chlorine. It is a colorless liquid, which boils at 130. It is almost completely hydrolyzed by water to hydrochloric acid and arsenious oxide or acid. Some arsenic trichloride is still present in the solution, however, as a part of the arsenic passes over on distilling the solution, while arsenic does not escape on distilling a solution of arsenious oxide. Sulfides of Arsenic. There are four sulfides of arsenic, As 2 S 2 , As 2 S 3 , As 2 $5 and As4S 3 . The last was prepared rather recently. Arsenic Bisulfide, or Realgar, As 2 S 2 , is found in nature and may be prepared by melting a mixture of arsenic and sulfur. It forms a red, crystalline mass which becomes lighter colored when powdered, and was formerly used by artists in painting. Arsenic Trisulfide, or Orpiment, As 2 S 3 , is also found in nature, and is prepared artificially by melting sulfur and arsenic mixed in the proper proportion. When prepared in this way it forms a yellow, crystalline mass and the powder is used as a pigment, especially by artists. From acid solutions of arsenic trioxide the trisulfide is precipitated in an amorphous form. It is one of the most insoluble sul fides known and is scarcely attacked by the most concentrated hydrochloric acid. It is, however, dissolved in the presence of oxidizing agents, as by nitric acid, aqua regia or potassium chlorate and hydrochloric acid. Arsenic Pentasulfide, As 2 S 5 , is precipitated from a solution of arsenic acid, H 3 AsC>4, containing hydrochloric acid, apparently through the substitution of sulfur for oxygen, giving the series ARSENIC 261 of acids, H 3 AsSO 3 , H 3 AsS 2 O 2 , H 3 AsS 3 O and H 3 AsS 4 . The last then dissociates into hydrogen sulfide and arsenic penta- sulfide. (MacCay, J. Am. Chem. Soc. 24, 661 (1902) ; Z. anorg. Chem. 29, 36 (1901). It may also be prepared by melting a mixture of the elements. Sulfarsenites and Sulfarsenates. Arsenic trisulfide, As 2 S 3 , and arsenic pentasulfide, As 2 Ss, dissolve easily in solutions of ammonium sulfide, (NH4) 2 S, or sodium sulfide, Na 2 S, giving solutions of sulfarsenites and sulfarsenates 3 (NH 4 ) 2 S + As 2 S 3 = 2 (NH 4 ) 3 AsS 3 Ammonium Sulfarsenite 3 Na 2 S + As 2 S 5 = 2 Na 3 AsS 4 Sodium Sulfarsenate These compounds may be considered as arsenites and arsenates in which the oxygen has been replaced by sulfur. Antimony forms similar compounds, but bismuth does not form them in this manner another illustration of the fact that bismuth is more distinctly metallic and does not show the same tendency as arsenic and antimony to form acid radicals. The formation of these compounds is used in analytical chemistry to separate arsenic and antimony from metals which are more metallic in character and which do not form similar compounds. From solutions of the sulfarsenites or sulfarsenates, acids precipitate the arsenic as the trisulfide, As 2 S 3 , or the pentasul- fide, As 2 S5. Colloidal Arsenic Trisulfide. It has been pointed out that arsenic trisulfide is one of the most insoluble compounds known. It requires at least two million parts of water to dissolve one part of the sulfide. In spite of this, however, hydrogen sulfide gives no precipitate with a solution of arsenic trioxide in pure water. A study of the properties of the solution obtained in this manner indicates that the interaction between the trioxide and hydrogen sulfide is practically complete : As 2 O 3 + 3 H 2 S = As 2 S 3 + 3 H 2 O 262 A TEXTBOOK OF CHEMISTRY The solution has the properties of a typical colloidal "solu- tion " a " solution " in which a substance, which under other conditions is insoluble and separates as a precipitate, remains in suspension. Such solutions will pass through ordinary filters un- changed and under an ordinary microscope they appear to be homogeneous. The freezing points and boiling points of such solutions are practically identical with the freezing point and boil- ing point of the pure solvent in this case water. This indicates that colloids are not in the ordinary molecular condition. The ultramicroscope reveals in many colloidal solutions the presence of aggregates which have a diameter of from 6 to 60 //./x. 1 Under the influence of a considerable electrical potential, colloidal arsenic trisulfide moves slowly toward the anode, indicating that the aggregates carry negative charges. In the case of some colloidal solutions the movement is toward the cathode. We may distinguish, therefore, negative colloids, as arsenic tri- sulfide, and positive colloids, as colloidal silver. The addition of an electrolyte to a colloidal solution will usually cause its precipi- tation. In general an electrolyte with bivalent ions, as barium chloride, BaCl 2 , is more effective than one with univalent ions, as sodium chloride, but the effect is dependent also on the degree of ionization of the electrolyte and it seems to be the cation (e.g. Ba ++ ) which is effective in precipitating a negative colloid, and the anion (e.g. Cl~ or SO4 ) which precipitates a positive colloid. The cation (or anion) is retained by the precipitate and cannot be washed away, though it may be displaced by another ion of the same sign. These facts are, at present, best understood in the light of the following theory. In the colloidal solution aggregates of a substance which is usually insoluble are formed around negative or positive ions, forming in the first case negative, in the second case positive, colloids. These aggregates are very much larger than ordinary molecules, but they are prevented from falling to 1 /* stands for one micron, y^ of a millimeter. /*/* stands for TTyW of a micron or one millionth of a millimeter. The wave length of sodium light is approximately 0.6 /*. ANTIMONY 263 the bottom of the solution, partly because they are still very small, but more because on account of their electrical charges they are prevented from cohering with other similar aggregates to form larger particles and also because there must always be in the solution, to balance them electrically, other, ordinary, ions with charges of the opposite signs. If these aggregates were to separate from the solution as a precipitate, the solution would be electrically positive and the precipitate negative in the case of arsenious sulfide. When an electrolyte, as barium chloride, is added to such a solution, the positive barium ions, Ba ++ , combine with the negative ions of the colloidal arsenious sulfide, forming neutral aggregates which can then cohere to larger aggregates and form an ordinary precipitate. At the same time the negative chloride ions, Cl~, balance the positive ions of the solution, usually hydrogen ions, H + , and the solution remains electrically neutral although the colloid has separated from it. It has been shown that in such a case the solution remains acid in exact proportion to the amount of barium carried down by the precipitate, and it has already been pointed out above that the barium cannot be removed from the latter by washing. A knowledge of the conditions which govern the formation and precipitation of colloids is often of very great importance in analytical chemistry. The phenomena of colloidal solutions also play a very important part in the digestion and assimilation of food and in the life processes of both plants and animals. Antimony, Sb, 120.2. Occurrence and Preparation. Small quantities of antimony are found free in nature, but the element occurs chiefly in the mineral stibnite, antimony trisulfide, Sb 2 S 3 . When this is heated in the air, the sulfur burns away as sulfur dioxide, SO 2 , and the antimony remains as the tetroxide, Sb2C>4. The process is called roasting and is a very common method of treating ores which contain sulfides of the metals. The crude oxide is then reduced by heating it with coke or charcoal and suitable substances to form a fusible slag with the impurities of the ore : sb 2 O 4 + 2 C = 2 Sb + 2 CO 2 264 A TEXTBOOK OF CHEMISTRY Properties. Antimony is a silver-white, brittle, crystalline metal. The specific gravity is 6.52. It melts at 630 and boils at 1300. A curious form of the element known as explosive antimony can be prepared by the electrolysis of a solution of antimony chloride, SbCla, in hydrochloric acid. It has a specific gravity of only 5.78. When rubbed in a mortar or when the dry sub- stance is heated to 200, it explodes violently, with an appearance of light and heat, being converted into ordinary antimony. The transformation is accompanied by the evolution of about 20 calories per gram. When antimony is heated in the air, on charcoal, it melts easily, differing in this respect from arsenic, which sublimes without melting. It burns slowly, giving vapors of antimony trioxide, Sb 2 O 3 . It does not dissolve in hydrochloric acid, but is easily converted into a mixture of insoluble oxides by nitric acid. Antimony and tin are the only metals which are acted upon by nitric acid in this manner, giving insoluble oxides or acids. Uses. Metallic antimony is a constituent of many important alloys, especially of type metal (lead, tin and antimony), stereo- type metal (lead, tin, antimony and bismuth), britannia metal (tin and antimony) and antifriction metals (lead, antimony and copper with a little bismuth) used for bearings in machinery. In type metal it gives hardness to the alloy and also causes it to expand slightly as it solidifies in the mold, giving clear-cut type. Stibine, SbH 3 , is formed when a soluble compound of antimony is introduced into a hydrogen generator containing zinc and hydrochloric or sulfuric acid. It resembles arsine closely, but gives a somewhat more sooty spot on porcelain or in a glass tube by Marsh's test. It is decomposed into antimony and hydrogen at a lower temperature than arsine, the decomposition taking place slowly at ordinary temperatures, especially in the presence of metallic antimony. For the methods of distinguishing between the deposits of arsenic and antimony, works on ana- lytical chemistry should be consulted. ANTIMONY 265 Oxides of Antimony. Antimony forms three oxides: anti- mony trioxide, Sb 4 O 6 (or Sb 2 O 3 ), formed by burning antimony in the air or by heating the hydroxide, Sb(OH) 3 ; antimony tetroxide, Sb2O4, formed when either the pentoxide or the trioxide is heated with free access of air ; and antimony pentox- ide, Sb 2 O 5 , obtained by repeated evaporation of metallic anti- mony or one of the lower oxides with nitric acid and finally heating the residue to 300. At a higher temperature it is decomposed into the tetroxide and oxygen. The trioxide is the only oxide which can be converted into a vapor without decompo- sition, and so is the only one for which we really know the molec- ular weight and true formula, Sb^e. It is altogether probable that the tetroxide and pentoxide have more complex formulas than those given, and, indeed, it is quite possible that the solid trioxide has a higher molecular weight and more complex formula than that of its vapor. For inorganic compounds, in general, it is more convenient to use the simplest formulas which express the composition in whole atomic weights. For this reason chem- ists continue to use the formulas P 2 O 3 , P 2 O5, Sb 2 O 3 , etc., and the corresponding names, for compounds whose molecules are known to be more complex. When questions of structure are con- sidered, however, it is important to remember that the molecules are more complex than these formulas indicate. Antimony hydroxide, Sb(OH) 3 or H 3 SbO 3 , Antimonious Acid, may be prepared by the precipitation of a solution of tartar emetic (see below) with dilute sulfuric acid. As is to be expected from the position of antimony in the Periodic System, it is amphoteric in character, that is, both an acid and a base. In a solution of a strong base it gives up hydrogen and forms a salt in which it furnishes the acid radical : H 3 Sb0 3 + NaOH = NaH 2 SbO 3 .H 2 O Sodium Antimonite In a solution of a strong acid, on the other hand, it gives up its hydroxyl and forms salts in which the antimony is the metallic element : 2 Sb(O H) 3 + 3 H 2 SO 4 = Sb 2 (SO 4 ) 3 + 6 H 2 O 266 A TEXTBOOK OF CHEMISTRY In further agreement with this character, the hydroxide or acid loses both hydrogen and hydroxyl (OH) easily even when in contact with water, going over into the oxide, Sb 2 O 3 : /io-H< JH-OJ Sb^-O-JHJ JH-t-O-Sb \O-!H H-oi As is to be expected, also, both classes of salts are hydrolyzed by water. The salts of the alkalies react strongly alkaline in solution, while the salts in which antimony forms the metallic part are mostly decomposed by water with the precipitation of a basic salt and liberation of the free acid. In these basic salts instead of the group SbO H, which might be expected, the group Sb^ , antimonyl, formed from this by the loss of hydro- gen and hydroxyl, is often present : Sb 2 (SO 4 ) 3 + 2 HOH = (SbO) 2 SO 4 + 2 H 2 SO 4 Antimony! Sulfate Tartaric Emetic, KSbOC 4 H 4 O 6 . One of the most interesting and important basic salts of antimony is tartar emetic, or potas- sium antimonyl tartrate. Cream of tartar, or acid potassium tartrate, KHC 4 H 4 O 6 , is the acid potassium salt of tartaric acid, H 2 C 4 H 4 O 6 , an acid found in the juice of grapes. When antimony trioxide, Sb 2 O 3 , is boiled with a solution of cream tartar, it dis- solves, forming tartar emetic : Sb 2 O 3 + 2 KHC 4 H 4 6 = 2 KSbOC 4 H 4 O 6 + H 2 O /o In the tartar emetic the univalent antimonyl group, Sbv , replaces hydrogen .as though it were a univalent metal, very much as the ammonium group, NH 4 , replaces hydrogen in the ANTIMONY 267 formation of ammonium salts. Tartar emetic dissolves easily in water. It is sometimes used as an emetic. Antimonic Acids. Three antimonic acids are known, cor- responding to the phosphoric acids of similar formulas : metanti- monic acid, HSbOa, pyroantimonic add, H4Sb2C>7, and orthoanti- monic acid, H 3 SbO 4 . Very few salts of the last are known. Chlorides of Antimony. Antimony forms three chlorides. Antimony trichloride, SbCla, can be prepared by dissolving the trioxide, Sb2Oa, or the trisulfide, Sb2Sa, in concentrated hydro- chloric acid, evaporating the solution and distilling the residue. It is a solid which melts at 73.2 and boils at 223. It is decom- posed by water with the precipitation of the oxychloride, which may be called antimony 1 chloride, SbOCl. From the method of preparing the trichloride it is evident that the decomposition by water is a reversible reaction, the direction of which depends on the concentration of the reacting substances. * Antimony tetrachloride, SbCU, and Hydrotetrachloroanti- monic acid, H 2 SbCl 6 . When solutions of antimony trichloride, SbCls, and antimony pent-achloride, SbCls, in hydrochloric acid are mixed, a dark brown solution is formed. The depth of color increases on warming and decreases on cooling, indicat- ing that the tetrachloride is formed with an absorption of heat, since its formation is promoted by an increase of temperature (p * l SbCl 3 + SbCl 5 = 2 SbCl 4 , The tetrachloride has not been prepared in pure condition, but double salts with other metals such as the caesium tetrachloro- antimonate, Cs2SbCl6, have been prepared. This corresponds to an acid, H 2 SbCle, which probably exists in the dark brown solution referred to above and which may be called hydrotetra- chloroantimonic acid. * Antimony Pentachloride, SbCls, is formed when antimony is burned in chlorine. It is a white solid at low temperatures, but melts at 6 and boils at 140. At the latter temperature it dissociates, partly, into the trichloride and chlorine, exactly as phosphorus pentachloride, PCls, does. 268 A TEXTBOOK OF CHEMISTRY * Metachloroantimonic Acid, HSbCle.4^ H 2 O. If chlorine is led into a concentrated solution of antimony trichloride in hydrochloric acid till the solution becomes colorless or light yellow and the solution is evaporated in a current of hydro- chloric acid to prevent hydrolysis, very hygroscopic crystals of metachloroantimonic acid can be obtained. This may be con- sidered as metantimonic acid, HSbO 3 , in which the three oxygen atoms have been replaced by six chlorine atoms. The freezing point of the solution indicates that the compound separates into the ions H^and SbCl~6- The solution gives a precipitate with silver nitrate only after some time, indicating that very few chloride ions, Cl~, are present. Very many salts of this acid have been prepared, among which the following may be mentioned : Potassium metachloroantimonate .... KSbCle.H 2 O Calcium metachloroantimonate .... Ca(SbCl6) 2 .9 H 2 O Aluminium metachloroantimonate . . . Al(SbCl 6 )3.15 H 2 O Antimony Trisulfide, Sb 2 S3, is obtained as an orange-red pre- cipitate when hydrogen sulfide is passed into an acid solution of a salt of antimony. It also occurs in nature as the black mineral, stibnite. It dissolves readily in concentrated hydrochloric acid, differing very markedly from arsenic trisulfide in this regard. Antimony Pentasulfide, Sb 2 S 5 , is best obtained by the decom- position of sodium sulfantimonate with hydrochloric acid : 2 Na 3 SbS 4 + 6 HC1 = 6 NaCl + Sb 2 S 5 + 3 H 2 S Sulfantimonites, M 3 SbS 3 , and Sulfantimonates, M 3 SbS4, may be obtained in the same manner as the corresponding sulfar- senites and sulfarsenates (p. 261). The alkali salts are soluble in water, hence the sulfides of antimony dissolve in solutions of sodium sulfide or ammonium sulfide and may be separated in this way from the sulfides of elements which are very decidedly metallic in character. Bismuth, Bi, 208. Occurrence, Properties, Uses. Bismuth is less abundant in nature than arsenic or antimony, as is usually, though not invariably, the case with elements of high atomic BISMUTH 269 weights. It is found mostly in the free state, but is found also as the sulfide, Bi 2 S 3 , both alone and with other sulfides, espe- cially with lead sulfide. It is also found as the oxide, Bi 2 O3. It is obtained commercially as a by-product in the electrolytic refining of lead. The specific gravity of the distilled metal is 9.78. The melting point is 271. As bismuth expands on solidi- fying, the melting point is lowered by pressure (Principle of van't Hoff-Le Chatelier, p. 111). Its boiling point is below 1700, but is not accurately known. Bismuth is used in a variety of alloys, usually because it lowers their melting points and renders them more suitable for specific purposes. It is used in this way in stereotype metal to give an alloy which can be cast in a papier-mache mold without injuring it, in Wood's metal (Bi, 4 parts, Pb, 2 parts, Sn, 1 part, Cd, 1 part), which melts at 60.5, and is used for heating baths in chemical laboratories, and for many other easily fusible alloys used as safety plugs in steam boilers and in automatic sprinklers for protection against fire, also for safety fuses in electrical work. The addition of a little bismuth has been found an advantage in Babbitt metal and in other antifriction metals used for bearings in machinery. Oxides of Bismuth. Bismuth forms two well-characterized oxides, BiO and Bi 2 O 3 . It also forms one or more higher oxides, called peroxides, for which we should expect the formulas Bi 2 O4 or Bi 2 O5. These higher oxides do not seem to have been obtained in a state of purity, probably because of the ease with which they and their hydrates lose water and oxygen. A mixture of these oxides or hydrates containing some sodium is prepared and has been called, without good reason, sodium bismuthate. It oxidizes manganese compounds to permanganic acid, HMnC>4, in nitric acid solutions and is used for that purpose in the de- tection and quantitative determination of manganese. Appar- ently no pure bismuthic acid or salt of bismuthic acid has been prepared. Bismuth Chloride, Bids, is formed by the direct union of bismuth and chlorine or by the solution of bismuth trioxide, 270 A TEXTBOOK OF CHEMISTRY Bi 2 O 3 , in concentrated hydrochloric acid. It melts at 225-230 and boils at 427-429. It dissolves in moderately strong hydrochloric acid to a clear solution, but the addition of water causes the precipitation of bismuth oxychloride, or bismuthyl chloride, BiOCl, which is extremely insoluble. Bismuth Nitrate, Bi(NO 3 )3.5 H 2 O, can be prepared by dis- solving either bismuth trioxide, Bi 2 O 3 , or metallic bismuth in an excess of nitric acid and evaporating to crystallization. When bismuth nitrate is treated with water, it is hydrolyzed with the formation of a mixture of basic nitrates, which varies in compo- sition according to the method by which it is prepared. The simplest of these compounds are Bi(OH)2NO3 and BiONO 3 : OH Bi (NO,). + 2 HOH = Bi OH + 2 HNO 3 X NO 3 n , H 2 O /.Pn ,0 i(-!OHl = Bif + \ L N6V N0 3 The mixture of basic nitrates is called in many medical works " bismuth subnitrate," an antiquated name which does not correspond to modern scientific usage. It is used in medicine and also as a slightly antiseptic face powder. Bismuth Trisulfide, Bi 2 S 3 , separates as a black or brownish black precipitate when hydrogen sulfide is passed into a solution of a soluble bismuth salt. Owing to the more metallic character of bismuth, it does not dissolve appreciably in solutions of sodium sulfide or ammonium sulfide as the sulfides of arsenic and antimony do. It dissolves easily in warm nitric acid, form- ing bismuth nitrate, Bi(NO 3 ) 3 . The following tables of the more important compounds of the elements of the fifth group will be of service in reviewing and comparing these. Compounds which correspond for different elements are selected, especially, for the table. Many other compounds are, of course, known. ARSENIC, ANTIMONY AND BISMUTH 271 N, 14 N 2 NO N 2 3 NO 2 , N 2 O 4 N 2 5 NC1 P, 31 PC1 PC1 6 Oxides As, 75 Sb, 120 AsCl SbCl Bi, 208 Bi 2 O 3 PA As 4 O 6 Sb 4 O 6 P<0 10 As 2 O 5 Sb 2 O 5 Bi 2 O 5 ? Chlorides BiCl H 2 N 2 O 2 HN0 2 HNO 3 Acids H 3 P0 2 H 3 P0 3 H 3 As0 3 H 2 P0 3 HP0 3 HAsO 3 H 4 P 2 O 7 H 4 As 2 O 7 H 3 P0 4 H 3 As0 4 H 3 SbO 3 (Bi(OH) 8 ) HSb0 3 HBi0 3 ? Salts of Sulfur Acids Na 3 PS 3 Na 3 AsS 3 M 4 P 2 S 7 Na 3 AsS 4 Na 3 SbS 3 NaBiS NH 3 N 2 H 4 N 3 H PH 3 P 2 H 4 Hydrides AsH 3 SbH 3 Vanadium (V, 51.0), Columbium (or Niobium) (Cb, 93.5), Tantalum (Ta, 181.5). These elements, which are found in the fifth group in the alternate rows of the Periodic System, are 272 A TEXTBOOK OF CHEMISTRY more decidedly metallic in their properties, corresponding to their positions in the system ; and while they show many analogies with the elements described in this chapter, they will be reserved for a later consideration (p. 522). EXERCISES 1. What are the reactions in Marsh's test, if arsenious oxide is used ? What, if arsenic acid is used ? 2. If an arsenic mirror on porcelain is warmed with ammonium sulfide, it dissolves and the addition of hydrochloric acid gives a lemon- yellow precipitate which does not dissolve in concentrated hydrochloric acid. Write the equations. 3. The antimony mirror conducts itself in a similar manner, but gives an orange precipitate with dilute hydrochloric acid which dissolves in concentrated hydrochloric acid. Write the equations. 4. What volume of air will be required for the complete combustion of one volume of arsine ? 5. What weight of arsenious oxide will be required to give one pound (453 grams) of Paris green ? 6. What will be formed by the ignition of magnesium ammonium arsenate ? 7. The specific gravity of fused arsenic trisulfide is 2.76. If a particle of the colloidal sulfide has this specific gravity and is one micron in diameter, how many such particles would there be in one gram of the sulfide? Assuming that it takes 1.5 X 10 21 molecules of arsenious sul- fide (As 2 S 3 ) to weigh one gram, how many molecules would there be in such a particle of colloidal arsenic trisulfide ? 8. Assuming the formula Bi 2 O 5 for bismuth peroxide, what is the equation for the reaction between this compound and manganese ni- trate, Mg(NO 3 ) 2 , in the presence of dilute nitric acid ? CHAPTER XVI CARBON Carbon, C, 12. Occurrence. Although carbon forms only about one five-hundredth part of that portion of the earth which we can examine, it is in many respects the most important of all of the elements. It forms an indispensable element in all living organisms, both animal and vegetable, and, indeed, it seems to be the peculiar properties of carbon rather than those of any other element, which make life, as we know it, possible. In addition to this preeminent role in living bodies, carbon is the principal constituent in all substances used for fuel and is the element by means of which iron is reduced from its ores. The unique character of carbon is suggested by its position in the Periodic System. It stands in the first row midway between the most strongly nonmetallic element fluorine and one of the alkali metals, lithium. Corresponding to this position its valence is four, but it combines both with the positive hy- drogen and with the negative fluorine and chlorine. Even in its elementary forms, it is nonmetallic, transparent and a nonconductor of electricity in the diamond, but approaches the metals in being opaque and a fairly good conductor of elec- tricity in graphite. This double character seems to be closely connected with the power which carbon atoms have to combine with each other as well as with other elements. The compounds of the element are bewildering in their variety and complexity. More than 100,000 such compounds have been prepared and analyzed, and some thousands of new compounds are discovered every year. On account of their number and many peculiarities, which distinguish them from the compounds of other elements, but also because of their importance and because so much time has been devoted to their study, the compounds of carbon are 273 274 A TEXTBOOK OF CHEMISTRY usually considered separately as a special subject, called organic chemistry. The study of these compounds has proved so important in its relation to the problems of general chemistry, however, that no textbook of inorganic chemistry is complete without a descrip- tion of some of them. Diamonds. A diamond of the first quality, weighing 0.2 gram, when properly cut, is worth approximately $100, while a kilo- gram of carbon in the form of coal or coke is worth less than one cent. This fact has been a constant challenge to chemists ever since the composition of diamonds was discovered. It was not, however, till near the close of the nineteenth century that even microscopic diamonds were prepared artificially ; and even now the theoretical conditions for their preparation are not fully understood and no one has succeeded in making diamonds large enough to be of commercial value. In 1892 Friedel discovered that a meteor which had fallen in Canon Diablo, Texas, contained carbon. By a careful examina- tion, Moissan, in Paris, demonstrated the presence of microscopic diamonds in the material. As he stated it afterwards, nature had been caught in the act of making diamonds. The high specific gravity of the diamond in comparison with graphite seems to have suggested to Moissan that diamonds were probably formed under conditions of high pressure. This and the dis- covery of diamonds in the meteorite suggested to him the follow- ing experiment. A mass of iron was heated to a very high tem- perature, 3000-3500, in an electric furnace and some pure sugar charcoal was dissolved in the hot iron. The mass was then suddenly thrust into water. This caused the exterior surface to solidify while the interior was still fluid and intensely hot. As the interior cooled, some of the dissolved carbon sepa- rated in crystalline form, and, as iron containing carbon ex- pands on solidifying, the interior portions were subjected to an enormous pressure. The lowering of the melting point of the iron by the pressure may also have had something to do with the success of the experiment. When cold, the iron was dissolved DIAMONDS 275 in acids, and silicon, graphite and other substances were removed by oxidation, solution, and finally by treatment with a liquid having a specific gravity greater than that of graphite and less than that of the diamond. A few minute crystals were dis- covered which were heavier than this liquid. By carefully rubbing one of the crystals against the surface of a ruby, it was shown that the surface of the latter was scratched. Only the diamond and carborundum are known to be harder than the ruby, and the specific gravity of carborundum is less than that of the artificial diamonds. After securing enough of the crystals to weigh a few milligrams, they were burned in a current of oxygen and the carbon dioxide formed was absorbed and weighed. Twelve parts by weight of the crystals gave 44 parts by weight of carbon dioxide. The crystals consisted, therefore, of pure carbon and were in reality diamonds. Until the beginning of the eighteenth century diamonds had been found only in India. In 1727 they were discovered in Brazil, at the beginning of the nineteenth century in the Ural mountains, and in 1867 at Kimberley in South Africa. The Kimberley mines now furnish most of the diamonds for the world's market. The product of the mines is valued at about $15,000,000 annually. The diamond crystallizes in octahedra and cubes of the iso- metric system. Its specific gravity is 3.5. It is the hardest known substance, and is not dissolved or oxidized by any known liquid or gas at ordinary temperatures. Its index of refraction is extraordinarily high, being 2.417 for sodium light. This and the high dispersive power give to diamonds, which are cut so as to accentuate these properties, the ability to reflect and re- fract light in such a manner as to produce brilliant colors. For cutting glass, an edge produced by cleavage must be used. For diamond drills to be used in boring in rocks in such a manner that a solid core can be removed, inferior black diamonds, called carbonado, set into the ends of tubes, are employed.- When heated to a high temperature with exclusion of air, the diamond is changed to graphite. When heated in oxygen, carbon 276 A TEXTBOOK OF CHEMISTRY dioxide begins to be formed at 720 and the diamond takes fire and burns at 800-850. Graphite, a second crystalline form of carbon, is found in nature, especially in Ceylon, Siberia, England and Canada. It may be prepared artificially by crystallizing carbon from cast iron, from one to two per cent of graphite being left behind when gray cast iron is dissolved in acids. It is also formed when any form of carbon is heated to a very high temperature in an electric furnace. The presence of silicon, aluminium, iron and other elements seems to assist in the transformation by the interme- diate formation of carbides. In the Acheson process, which has acquired considerable technical importance, impure carbon, as coke or anthracite, is found more suitable than pure carbon. Graphite crystallizes in six-sided leaflets of the monoclinic system. It is soft and has a gray, metallic luster and gives a metallic streak. Its specific gravity when pure is 2.255. In oxygen it begins to give carbon dioxide at 570 and takes fire at 690. At very high temperatures it seems to be the most stable form of carbon, into which all other forms tend to pass. The total energy of graphite at ordinary temperatures is, how- ever, greater than that of the diamond at ordinary temperatures, as is evident from the following table of heats of combustion, as determined by Berthelot : 12 grams of diamond give 94,310 small calories. 12 grams of graphite give 94,810 small calories. 12 grams of amorphous carbon give 97,650 small calories. From this table it appears that if 12 grams of graphite could be changed to diamond at ordinary temperatures, 500 small calories would be evolved. Graphite is used for " lead " pencils, as a lubricant, especially for surfaces of wood and where bearings are subjected to a high tem- perature, and in making crucibles for use at very high tempera- tures for melting and casting steel and difficultly fusible alloys. For the last purpose it must be mixed with some fire clay, which binds the particles of graphite together and also protects CARBON 277 the graphite from burning by giving an incombustible surface. Graphite is also used as a lubricant in a colloidal solution in oil (Acheson). It is used for stove polish, to protect iron from rusting. Amorphous Carbon. When almost any compound of carbon is heated, it will decompose with the separation of charcoal or car- bon. It is, however, extremely difficult to obtain perfectly pure carbon in this manner. The purest carbon is obtained by heat- ing sugar, which contains only carbon, hydrogen and oxygen, to a high temperature, but it seems doubtful whether the last traces of hydrogen can be expelled without the use of a temperature which would convert the amorphous carbon partly into graphite. The hydrogen may be almost completely removed, however, by heating the charcoal in a current of chlorine at 1000. There seems to be no form of amorphous carbon which can be properly spoken of as a definite chemical individual, as the density and kin- dling temperatures and conductivity for electricity vary gradually from a form which has a density of 1.45 and a kindling tempera- ture of 300, up to forms which approach closely to graphite in their properties. This fact is doubtless intimately connected with the almost infinite variety of ways in which carbon atoms unite with each other in the compounds of carbon. A great variety of impure forms of amorphous carbon are known. All of these contain at least some hydrogen and most of them contain oxygen and other elements. Lampblack is the soot deposited from substances rich in carbon, such as naphthalene, rosin, petroleum, etc., burning with a smoky flame. It is used as a pigment, especially in printers' ink. The insoluble and indelible quality of carbon makes such ink even more permanent than the paper on which it is printed. Wood Charcoal is manufactured by piling wood in heaps, covering it with sod and setting fire to it in such a manner that only a portion of the wood burns while the heat converts the re- mainder into charcoal. The process is wasteful and has been largely replaced by methods of charring in retorts or chambers so arranged that the wood tar, wood alcohol and acetic acid, 278 A TEXTBOOK OF CHEMISTRY which are formed by the decomposition of the wood, may be recovered. Combustible gases, which are also formed, are util- ized to heat the retorts or chambers. Charcoal is used by tinners, in small charcoal furnaces, for filtering alcohol to purify it and for the manufacture of a high grade of iron. It was formerly used in very large quantities for this last purpose, but has been almost entirely displaced by coke and coal. Charcoal, because of the infusible character of all forms of car- bon, retains the original structure of the wood and contains an immense number of pores of microscopic size. Apparently for this reason freshly ignited charcoal will absorb many times its volume of gases, especially of those gases which are easily lique- fied, such as ammonia, or hydrogen sulfide. This phenomenon is called adsorption, and seems to be due to the condensation of the gas on the very large surface offered by the porous charcoal. Charcoal cooled by liquid air is a very efficient means for ab- sorbing residual gases and producing a high vacuum. Charcoal was formerly much used in domestic water filters, but it has been found that such filters are only very temporarily effective. Animal Charcoal and Bone Black are obtained by charring the refuse of slaughterhouses and bones. They show the prop- erty of absorption, especially for coloring matters, in a very high degree, and are used in the removal of color from sirups and other liquids, as in the purification of sugar. They are also frequently used for the purification of organic compounds in chemical laboratories. For such use they should be purified by treatment with acids to remove calcium phosphate and other mineral matters which they contain. Coke bears very much the same relation to bituminous coal that charcoal does to wood. It is still manufactured in America, chiefly in the so-called " beehive " ovens hemispherical chambers built of brick, 12 feet in diameter and seven and one- half feet high. These are charged with coal while still hot from a previous charge, and the volatile matter given off from the coal takes fire and burns within the oven, over the surface of the CARBON 279 coal, furnishing the heat necessary to convert the coal into coke. An opening on one side of the oven supplies air and the products of combustion escape through a circular opening at the top. This wasteful method is being slowly replaced by methods of coking in retorts, by the use of which it is possible to recover the tar, ammonia and combustible gases. The gas produced is formed in excess of what is necessary for heating the retorts, and may be used in part for other purposes, while in the beehive ovens nothing is saved except the coke. Coke is used chiefly in blast furnaces for the production of cast iron. It is used in some other metallurgical processes and to a limited extent as a domestic fuel. Gas Carbon. Carbon Electrodes. On the walls of the retorts used in the manufacture of illuminating gas, carbon is deposited from the decomposition of carbon compounds in the volatile matter given off by the coal. As a result of the prolonged heating it assumes a semicrystalline, dense form, approaching graphite in its properties. It has a density of 1.9-2.0, a very high kindling temperature, and is a fairly good conductor of electricity. This gas carbon, or frequently, also, anthracite, petroleum coke, or some other form of amorphous carbon, is mixed with a little coal tar or some petroleum product for a binding material and molded into various forms for use as electrodes in the electroylsis of sodium chloride, aluminium com- pounds or other substances, for use in electrical furnaces and for the carbon electrodes of arc lights. The mixture is subjected to a hydraulic pressure of 500 atmospheres to render it as dense as possible, and is then heated to a temperature of 1200-! 400 for 24-48 hours, till all volatile compounds have been expelled and the carbon has become dense and hard and a good electrical conductor. Such electrodes are scarcely attacked by the chlo- rine evolved in the electrolysis of a solution of sodium chloride and are scarcely affected when heated, out of contact with air, to any temperature below that of the electric arc. Their resistance to oxidizing agents in electroylsis maybe further increased by con- verting them partly or wholly into graphite in an electric furnace. 280 A TEXTBOOK OF CHEMISTRY Coal. Very much the same process which occurs rapidly when wood is heated seems to have gone on slowly through some hundreds of thousands or millions of years with vast quantities of woody material accumulated during certain periods of geolog- ical time and afterwards covered with thick layers of clay and other materials, beneath the surface of ancient oceans. Wood consists chiefly of carbon, hydrogen, oxygen and nitrogen, with small quantities of mineral matter. The transformation to coal has been occasioned by the gradual loss of oxygen and some of the hydrogen, in such a manner that the per cent of carbon gradually increases. The per cent of the oxygen decreases until that element nearly disappears in anthracite. The per cent of hydrogen decreases only slightly till the last stage the trans- formation to anthracite is reached. It seems probable that this last transformation occurred at a more elevated temperature. The changes in composition which have taken place during these transformations are apparent in the following table: CHANGES OF WOOD MATERIAL DURING GEOLOGICAL TIME PERCENTAGE COMPOSITION EXCLUSIVE OP MOISTURE PERCENT- PERCENT- CALORIFIC VALUE ; MATERIAL AND ASH AGE OF AGE OF CALORIES Car- Hydro- Oxy- Nitro- ASH MOISTURE PER KIL- OGRAM bon gen gen gen Wood Oak . 50.35 6.04 43.52 0.09 0.37 20.00 2 3696 Peat .... 59.70 5.70 33.04 1.56 11.84 14.24 2 3979 Brown Lignite, North Dakota 74.88 4.99 19.12 1.01 9.35 35.38 3846 Black Lignite, Colorado . . 76.83 5.34 16.29 1.54 5.99 18.68 5635 Bituminous, Illinois . . 83.42 5.29 9.52 1.77 11.28 8.50 6542 Semibitumin- ous, West Vir- ginia Poca- hontas . . . 91.50 4.38 3.07 1.05 6.55 3.67 7939 Anthracite . . 93.76 2.72 3.11 0.41 10.80 2.18 7216 Charcoal . . . 84.11 1.53 14.36 2.50 6626 Coke .... 95.47 0.67 2.82 1.04 14.80 6768 1 Table prepared by Professor S. W. Parr. 2 Air dry. CARBON 281 The properties of the coals of different kinds follow from their composition. Peat, lignites, and bituminous coals increase progressively in calorific value as the amounts of moisture and oxygen decrease. The oxygen in these coals may be consid- ered as combined with either carbon or hydrogen and lessens by so much the amount of these elements which can evolve heat by combustion. Bituminous coals may equal or even exceed an- thracite coals in calorific value because a pound of hydrogen gives by its combustion more than three times as much heat as a pound of carbon. Such coals, however, give off volatile products which burn with a smoky flame, and hence require much greater care in use to secure effective combustion. Three classes of bituminous coals are distinguished : coking coals, which sinter together when heated, giving a hard, coherent coke ; noncoking coals, which do not sinter, or sinter imperfectly, giving a friable coke ; and cannel coals, coals of a peculiar, homo- geneous structure and conchoida.l fracture, which burn with a brilliant flame like that of a candle. These last coals are used in the manufacture of illuminating gas. The difference between coking and noncoking coals seems to be occasioned by the presence or absence of some compound whose character is little understood and which does not seem to be closely connected with the percentage composition of the coal. Chemical Properties of Carbon. The most remarkable prop- erty of carbon is the extreme slowness with which it reacts at ordinary temperatures with elements for which it has a very strong affinity at high temperatures. This is especially true in its relation to oxygen. Elementary carbon in either of its three forms may remain in contact with air for centuries without any apparent effect, although at very high temperatures there seems to be almost no element from which carbon will not take away oxygen. Practically all organic compounds must be considered as in a state of unstable equilibrium in the presence of oxygen, for we have only to heat them to. their kindling temperature when they will burn with very considerable evolution of heat. On this property depends the use of carbon and its compounds 282 A TEXTBOOK OF CHEMISTRY for fuel, for the reduction of iron ores and for other metallurgical operations. On this property, too, depends the existence of the almost infinite variety of compounds which form the material basis of the world of life compounds showing all possible gradations in their content of energy and the existence of which would be impossible if carbon passed quickly, as most other elements do, to the most stable forms of combination. In combination with other elements carbon is almost always quadrivalent. Methane, CH 4 , and carbon dioxide, CO 2 , may be considered as the most typical compounds. Carbon is bivalent in only a very few compounds and exclusively in combination with atoms or groups of a negative character, as in carbon monoxide, C=O, hydrocyanic acid, H N=C, and fulminic acid, H O N=C. There is some evidence that it may be tri- valent in very unusual combinations, but it is then extraordi- narily reactive. CHAPTER XVII HYDROCARBONS. ILLUMINATING AND PRODUCER GAS. FLAME CARBON combines with hydrogen to form many hundreds of compounds, called hydrocarbons. These compounds may be classified in a number of series in accordance with their com- position. The following table illustrates the relations which have been found between the formulas of successive hydrocar- bons in any series and between the hydrocarbons of different series. Each series is named from its first member. In the table only one series for a given general formula is given, but, as will be seen below, for each general formula, except the first, two or more series are possible. The series of the formula C w H 2n _ 4 is omitted from the table because the lower members of this series are relatively unimportant. MARSH GAS SERIES C w H 2n+2 ETHYLENE SERIES C w H 2n ACETYLENE SERIES C n H 2n-2 BENZENE SERIES CH 2w - 6 Methane CH 4 L- Ethane C 2 H 6 Ethene C 2 H 4 Acetylene C 2 H 2 Propane C 3 H 8 Propene C 3 H 6 Propine C 3 H 4 Butane C 4 Hio Butene C 4 H 8 Butine C 4 H 6 Pentane C 5 Hi 2 Pentene C 5 Hi Pentine CsHg Hexane C 6 Hi 4 Hexene C 6 H 12 Hexine C 6 Hi Benzene C 6 H 6 Heptane C 7 H 16 Heptene C 7 Hi 4 Heptine C 7 Hi 2 Toluene CyHg Octane CgHig Octene C 8 Hi 6 Octine C 8 Hi 4 Xylene C 8 Hi The existence of these compounds may be explained very simply on the hypothesis that carbon is quadrivalent and that carbon atoms unite readily with each other. This hypothesis 283 284 A TEXTBOOK OF CHEMISTRY gives us the following formulas for the first three members of the Marsh gas series : H H H I I I H C C C H I I I H H H Propane For the fourth member of the series the theory suggests two formulas : H i H H i i H C H i 1 1 H C C- 1 | H 1 H Methane 1 1 H H Ethane H H H H C C C H H H H H H Normal Butane H C H Boiling point, +1 | H Isobutane Boiling point, -11.5 These two hydrocarbons have been prepared by methods which leave no doubt as to the structure of each. For the series C TC H 2/l the theory suggests that we may have compounds in which carbon atoms are doubly united and also compounds in which there is a ring of carbon atoms. Thus we may have : H H H H \X H III C H C C = C H and Hv X\ /H H H/ H Propylene (Propene) Cyclopropane Boiling point, -37 Boiling point, -35 Both of these compounds are known, and the structure has been established by a study of the methods of preparation and of their conduct toward various reagents. HYDROCARBONS 285 For the series C n H 2n _2 there are four possibilities : one triple union, as in acetylene, H C = C H ; two double unions, as in butadiene, CH 2 =CH CH=CH 2 ; cyclic compounds with /CH 2 CH one double union, as cylopentene, CH 2 II ; and compounds \CH 2 CH with two cycles, or rings, as dekahydronaphthalene, CH 2 CH 2 CH CH 2 CH 2 The illustrations given would seem to include all of the types of combination possible for carbon and hydrogen atoms, since quadruple unions between carbon atoms would be impossible for atoms which are united to any other atoms. There are, however, certain other relations which seem to depend on the arrangement of the atoms in space. Some of the combinations of these forms give properties which would not be expected from the formulas of the compounds. This is especially true of the benzene series, in all of the com- pounds of which there is a ring of six carbon atoms, each of which is united to one hydrogen atom or to some other univalent atom or group. The simplest formula of benzene is that proposed byKekule: H H C/ ^C H II I H C jj H H and many of the properties and reactions of the hydrocarbon are satisfactorily represented by this formula, but for other proper- ties it does not give a satisfactory account. A further considera- tion of this and similar questions is impossible here. 286 A TEXTBOOK OF CHEMISTRY Marsh Gas or Methane, CH4. When decaying leaves in the bottom of a pond are stirred, bubbles of a combustible gas con- sisting largely of marsh gas or methane, CH 4 , rise to the surface. The same gas escapes from seams of coal, doubtless having been formed in a similar manner. In coal mines it is called fire damp. A combustible gas consisting very largely of methane is often found stored in large quantities in porous sandstones or lime- stones lying beneath an impervious layer of shale so situated as to form a large inverted reservoir. The gas is usually under strong hydrostatic pressure from water beneath. When such a reservoir is pierced by boring from above, the gas escapes through the opening, and is known as natural gas. When almost any kind of organic matter is heated, methane is one of the products of decomposition, hence it is always a constituent of illuminat- ing gas made by heating coal, oil or wood. At a white heat carbon will unite directly with hydrogen to form methane : C + 2 H 2 ^ CH 4 The equilibrium of the reaction is, however, very far on the side toward the decomposition of methane into carbon and hy- drogen. Water gas (p. 296) usually contains a very small amount of methane, which is probably formed by the direct union of the elements. In the laboratory methane is most easily prepared on a small scale by heating a mixture of sodium acetate, NaC 2 H3O 2 , and soda lime, which is a mixture of sodium hydroxide, NaOH, and slaked lime, Ca(OH) 2 . The slaked lime is added to render the mixture infusible : NaC 2 H 3 O 2 + NaOH = Na 2 CO 3 + CH 4 Sodium Carbonate Methane is the lightest gaseous compound known. It may be condensed to a liquid, which boils at 164. It is a compara- tively stable compound, and its kindling temperature is higher than that of hydrogen or than that of most other hydrocarbons. SUBSTITUTION 287 Mixtures of the gas with oxygen or with air explode violently when ignited. It burns from a jet with a blue flame, which gives very little light. Substitution. When a mixture of methane and chlorine is exposed to the sunlight, a double decomposition occurs in which one atom of the chlorine molecule combines with an atom of hydrogen while the other combines with the carbon, apparently taking the place of the hydrogen. This process, which occurs in a great variety of reactions of organic compounds, is called substitution : H H H C H + Cl Cl = H C Cl + H Cl H Methyl Chloride The process may be continued till all of the hydrogen has been replaced by chlorine : CH 3 C1 + C1 2 = HC1 + CH 2 C1 2 Methylene Chloride CH 2 C1 2 + C1 2 = HC1 + CHC1 3 Chloroform CHC1 3 + Cl a = HC1 + CC1 4 Carbon Tetrachloride Practically, a mixture of the four products is obtained by this process, so that these reactions have only a theoretical interest. The Davy Safety Lamp. The explosive character of mixtures of methane and air has been mentioned. Early in the nineteenth century the frequent explosions of fire damp in coal mines in Eng- land, often causing the death of miners, led to a request of Sir Hum- phrey Davy that he should investigate the matter and endeavor to suggest a remedy. He found, as the result of his investigation, that mixtures containing one volume of methane with more than six and less than fourteen volumes of air would explode when 288 A TEXTBOOK OF CHEMISTRY Fig. 80 ignited. 1 Outside of these limits explosions do not so readily occur. He also found that the kindling temperature of such mixtures is comparatively high, requir- ing contact with a surface heated nearly or quite to dull redness. 2 This prop- erty can be easily illustrated for illu- minating gas by pressing a cold piece of wire gauze down over the flame of a Bunsen burner (Fig. 80). Unburned gas from the center of the flame will pass through the gauze and the mixture of gas and air above the gauze will not take fire until the latter becomes nearly red-hot. On the basis of this fact Sir Hum- phrey Davy invented the Davy Safety Lamp, which has the flame of the lamp completely surrounded by wire gauze (Fig. 81). The lamp must, of course, be lighted and closed before the miner enters the mine. The lamps are usually so constructed that they can be locked. Sometimes a lock is used which can only be opened with a strong electromagnet, so that it will be impossible for the miner to open the lamp in the mine. While the danger of explosions is greatly lessened, it is not en- tirely removed by the use of the lamp. If much fire damp is present, a cap of flame ap- pears inside, above the flame of the lamp ; and this might, some- times, heat the wire gauze to the kindling temperature of the mixture. In blasting, too, the sudden vibration from the blast 1 Later investigations have shown that these results are only a very rough approximation. One authority states that air contain- ing 2 per cent of methane may be dangerous. 2 V. Meyer, many years later, found that the kindling temperature of mixtures of methane and oxygen in glass vessels is 650-680. Ber. 26, 2429. Fig. 81 PETROLEUM 289 may carry the flame through so quickly that it is not cooled below the kindling temperature by the gauze. Mixtures of very fine dust containing organic matter with air may explode in the same manner as fire damp. In this way destructive explosions have occurred with coal dust in dry mines, with flour in flour mills, and with similar dust in other factories. Homologues of Methane. On examining the formulas of the hydrocarbons given in the table it will be found that the succes- sive members of any series differ by one carbon and two hydrogen atoms. The reason for this is apparent from the structural formulas of methane, . ethane, propane, etc. Any hydrocarbon in such a series is called a homologue of the lower members of the series and the series is called a homologous series. Petroleum is found in great underground reservoirs somewhat similar to those containing natural gas. Oil fields have been found widely distributed in America, especially in Pennsylvania, Ohio, Indiana, Illinois, Kansas, Texas, California and Canada. A large field is found in the Caucasus, and doubtless very many undiscovered fields exist in other parts of the world. Petroleum consists chiefly of a very complex mixture of hydrocarbons. The petroleum from different localities differs very considerably in the nature of the hydrocarbons which it contains and also in the amount of the compounds of sulfur which are present. The Pennsylvania petroleum consists largely of homologues of me- thane. California and Caucasus petroleum contain compounds of the cyclic series. Crude petroleum is often used as a fuel. Its calorific value is about one half greater than that of the best quality of coal. Petroleum is refined chiefly by fractional- distillation. It is also treated with concentrated sulfuric acid to remove compounds with a disagreeable odor or objectionable properties. Com- pounds of sulfur are removed by boiling it with copper oxide. The principal product is usually kerosene, used as a burning oil in lamps. The low-boiling products are called petroleum ether and ligroin in chemical laboratories, or, commercially, gasoline, benzene and naphtha, partly in accordance with the boil- 290 A TEXTBOOK OF CHEMISTRY ing point, but chiefly according to the use made of the material. The vapors of the low boiling products form dangerously explo- sive mixtures with air. Kerosene should not give enough vapor to explode at any temperature below 65 (150 F.), and this is the legal flashing point in most states. Products boiling at a higher temperature than kerosene are used as lubricating oils. Solid products, called paraffin, are made into candles and are used for covering jellies and for many other purposes. A semisolid product, called vaseline, 1 is prepared for medicinal use. All of the products are very complex mix- tures of hydrocarbons. Ethylene or Ethene, C 2 H 4 . When ordinary alcohol, C 2 H 5 OH, and concentrated sulfuric acid are mixed in such proportion that the mixture boils at 140, ethyl ether, (C 2 H 5 )2O, and water distill over on heating. If more sulfuric acid is used (6 parts of sulfuric acid to 1 of alcohol by weight) so that the mixture distills or de- H \ / H composes at 170 - 180, ethylene, >C = C< , is formed. In W \H both cases we may consider that the sulfuric acid removes water from the alcohol, but the mechanism of the reaction is more com- plicated than such a statement indicates. 2 C 2 H 5 OH - H 2 O = C 2 H 5 O C 2 H 5 Ethyl Ether C 2 H 5 OH - H 2 O = C 2 H 4 Ethylene Ethylene is a colorless gas, very slightly lighter than air. It has a sweetish odor and burns with a bright, luminous flame. The difference between methane and ethylene in this regard seems to depend on the fact that methane is quite stable, even at comparatively high temperatures, and does not readily decom- pose with the separation of carbon, while at the temperature of the flame ethylene decomposes, partly, into methane and carbon : 1 Vaseline is a proprietary name used by the Chesebrough Manu- facturing Company. The name used in the Pharmacopoea and by other manufacturers is petrolatum. UNSATURATED COMPOUNDS 291 f~ C The carbon which is liberated temporarily assumes the solid form and, being raised to a white heat by the flame, makes it luminous. In spite of this instability, ethylene is formed when any hydrocarbon of the methane series is heated to a high temperature or, indeed, when almost any organic com- pound is heated. For this reason it is always present in illuminating gas prepared by heating coal or oil and is one of the most important constituents of the gas, because of the lumi- nous quality of its flame. It is formed to some extent even from methane, although it is less stable than methane and the reaction : 2 CH 4 = C 2 H 4 + 2 H 2 , is endothermic. It would almost seem that this ability of carbon to enter into many reactions in which heat is absorbed from sur- rounding objects is one of the most important characteristics of the element. It is doubtless intimately connected with those properties of the carbon atom which cause the rate of many of its reactions to be so slow (p. 281). The result of this seems to be that the speed of a reaction in one direction or another often has more effect in determining the direction of the reaction than the heat evolved or absorbed. Unsaturated Compounds. Ethylene Chloride and Ethylene Bromide. Ethylene combines directly with chlorine to form ethylene chloride, C2H4C12, and with bromine to form ethylene bromide, C 2 H 4 Br 2 . The process is known as addition. These compounds may be called, also, dichloroethane and dibromo- ethane and are to be considered as substitution products of ethane, C 2 H 6 . They illustrate the tendency of compounds having double or triple unions between carbon atoms to take up other elements and pass back into compounds which are derivatives of the hydrocarbons of the methane series. For this reason the hydrocarbons of the ethylene and acetylene series are called unsaturated, while the hydrocarbons of the methane 292 A TEXTBOOK OF CHEMISTRY series and their derivatives are called saturated. This conduct H H has led some chemists to prefer the formula H C C H for H H ethylene instead of the usual formula, H C = C H. Which- ever formula is true, it is evident that carbon atoms which are spoken of as doubly united are not more firmly held together than by a single union. The reverse of this seems to be true. The double union is a point of especial reactivity. Acetylene. When an electric arc is formed between carbon points in an atmosphere of hydrogen, some acetylene, C2H2, is formed. Acetylene is an endo thermic compound and de- composes into carbon and hydrogen with evolution of heat : C 2 H 2 = 2C + H 2 + 53,000 small calories for 26 grams of acetylene. The formation of acetylene is often used as an illustration of the fact that a high temperature is favorable to the formation of endothermic compounds. Acetylene is formed by the incomplete combustion of ethylene or of carbon compounds generally and so is found among the gases coming from a Bunsen burner burning at the base. The unpleasant odor of these gases is not, however, due to the acety- lene. The presence of the acetylene can be shown by driving the gases through an ammoniacal solution of cuprous chloride, with which the acetylene gives a precipitate of copper carbide, often incorrectly called copper acetylide : Cu 2 Cl 2 + 2 NH 3 + C 2 H 2 = Cu 2 C 2 + 2 NH 4 C1 Copper Carbide The formation of copper carbide in this way indicates that acetylene has some of the properties of an acid. The electri- cal conductivity of solutions of acetylene in water also indicates that it is an acid, but an extremely weak one, so that its salts ACETYLENE 293 are hydrolyzed by water. On this fact depends its prepara- tion for commercial uses from calcium carbide and water : CaC 2 + 2 HOH = Ca(OH) 2 + C 2 H 2 Calcium Carbide Acetylene is a colorless and odorless gas, which may be con- densed to a liquid or solid by cold or pressure. The boiling point is 83.6 and the melting point a little higher, 81.5. Under atmospheric pressure, therefore, it sublimes without melting. Acetylene is unsaturated in the same sense as ethylene and may combine with four atoms of bromine to form acetylene tetrabromide or tetrabromoethane, C 2 H2Br4. Acetylene dissolves readily in acetone, one volume of the liquid dissolving 25 volumes of the gas at 15 under atmos- pheric pressure and 300 volumes under a pressure of 12 atmos- pheres. As a solution containing 100 volumes of the gas for one of acetone is not explosive (see below) such a solution can be used to advantage for illuminating purposes. Acetylene burns from an ordinary jet with a smoky flame, due to the separation of carbon. When burned from a suitable burner it gives a very brilliant white light, developing from twelve to fifteen times as much light as can be obtained with the same volume of good illuminating gas burned with an ordi- nary burner. The intense light is due partly to the ease with which acetylene decomposes into carbon and hydrogen, but doubtless also to the heat developed by the decomposition, which aids in raising the temperature of the particles of carbon to a white heat. It is estimated that while only about 2 per cent of the energy of illuminating gas is actually effective as light, about 10 per cent of the energy of burning acetylene may appear as light. At 20 acetylene may be condensed to a liquid under a pres- sure of 41 atmospheres. This is considerably less than the vapor pressure of liquid carbon dioxide at the same tempera- 294 A TEXTBOOK OF CHEMISTRY ture, and when a process had been invented by means of which the manufacture of calcium carbide in an electric furnace be- came commercially possible it was thought that liquid acetylene could be very conveniently used in the liquid form, condensed in strong steel cylinders. Soon after the first attempts in this direc- tion were made, however, some unexpected and, at first, unac- countable explosions occurred. An investigation of the matter soon showed that liquid acetylene, or, indeed, gaseous acetylene, even under a pressure of somewhat less than two atmospheres, may be exploded by a glowing wire or by a fulminating cap. The explosion is due, of course, to the fact that acetylene de- composes into carbon and hydrogen with considerable evolution fheat: C 2 H 2 = 2C + H 2 Although the volume of the hydrogen is the same as that of the acetylene the heat of decomposition increases the pressure and is sufficient to cause the decomposition, when once started, to proceed explosively from one part of the liquid or compressed gas to another. It has been pointed out above that solutions of acetylene in acetone are not explosive, if the concentration is not carried too far. Such solutions are now extensively used, especially for automobile and motorcycle lights. At slightly elevated temperatures acetylene polymerizes easily, that is, it .combines with itself to form more complex compounds, some of which are liquid or solid at ordinary tem- peratures. When acetylene is generated by dropping water on calcium carbide the heat evolved by the reaction may cause a very considerable loss by polymerization. For this reason those forms of generators in which the carbide is dropped into the water are most suitable. Benzene, C 6 H 6 . In the manufacture of illuminating gas by the distillation of bituminous coal and also in the manufacture of coke by the methods in which retorts are used and the by- products are saved, large quantities of coal tar are produced. This is an extremely complex mixture from which many valuable products are obtained, chiefly by fractional distillation. Among ILLUMINATING GAS 295 these products are benzene and its homologues and some other hydrocarbons, especially naphthalene and anthracene. These hydrocarbons are very extensively used in the manufacture of the coal-tar dyes and of many other compounds which are used in medicine and in the arts. Benzene may also be formed by the polymerization of acetylene : 3 C 2 H 2 = It is a colorless liquid which melts at 5.4 and boils at 80.2. Illuminating Gas. The manufacture of illuminating gas by heating bituminous coal, especially cannel coal, in earthen- ware or iron retorts, has been repeatedly referred to. The gases obtained in this manner consist chiefly of a mixture of hydrogen, methane, carbon monoxide, carbon dioxide and hydrogen sulfide, with a small per cent of the so-called " heavy hydro- carbons." These last consist of ethylene, C2H4, vapor of benzene, CeH 6 , a little acetylene, C2H2, and small amounts of many other gases and vapors. In burning the gas only the heavy hydrocarbons decompose to an appreciable extent with separation of carbon. On this account these hydrocarbons are sometimes designated as illuminants. The composition of the gas varies greatly both with the tem- perature and the duration of the heating of the coal. A high temperature favors the decomposition of the hydrocarbons and increases the per cent of hydrogen, reducing the illuminating power of the gas, but it also greatly increases the volume of the gas obtained from a given weight of coal. The gas which es- capes from the coal during the first part of the heating is also much richer in the heavy hydrocarbons than that given later. These facts are easily understood from the instability of the hydrocarbons, at high temperatures. The hydrogen sulfide in the gas must be removed by passing it through boxes or chambers having a series of shelves covered with slaked lime, Ca(OH) 2 , or more often by passing it through boxes containing moist ferric hydroxide, Fe(OH)a: 2 Fe(OH) 3 + 3 H 2 S = 2 FeS + S + 3 H 2 O 296 A TEXTBOOK OF CHEMISTRY The ferrous sulfide passes back into a mixture of ferric hydrox- ide and sulfur on exposure to the air and so may be used re- peatedly : 2 FeS -f 3 O + 3 H 2 = 2 Fe(OH) 3 + 2 S The illuminating power of the gas is determined by com- parison with the light of a standard spermaceti candle which burns 120 grains per hour. For comparison, the gas is burned at the rate of 5 cubic feet per hour. Gas of good quality should give from 18 to 22 candle power. The light which can be obtained from a given quantity of gas may be greatly increased by burning it from a Bunsen burner under a mantle composed of oxides of thorium and cerium the Welsbach light, named after the inventor. The oxides not only furnish the solid substance heated to a high temperature, which is necessary in almost all forms of practical illumination, but they also catalyze the reaction of combustion, localizing the latter in immediate contact with the solid particles and so greatly increasing the temperature to which these are raised. (See p. 364.) With the inverted Welsbach burner, gas may give more than ten times as much light as could be obtained with an ordinary flat flame. Oil Gas. When petroleum is heated to a high temperature the hydrocarbons which it contains are decomposed with the formation of carbon, hydrogen, methane and other hydrocar- bons, some of which are gaseous at ordinary temperatures. If the temperature is high enough and especially with the aid of certain catalyzers, the final products are hydrogen and carbon. The process has been proposed as a method to obtain hydrogen for filling balloons. At lower temperatures it is possible to obtain a gas very rich in the heavy hydrocarbons, and a gas, called Pintsch-gas, is manufactured in this way and compressed in steel cylinders for use in lighting railway coaches and for similar purposes. Water Gas. When steam is passed over incandescent coal or coke a mixture of carbon monoxide, CO, and hydrogen, H 2 , is formed : C + H 2 O = CO + H 2 PRODUCER GAS 297 The heat of combustion of the carbon monoxide and hydrogen is very much greater than the heat of combustion of the carbon or coke. In other words the reaction is endothermic in a very high degree. The mass of incandescent coke cools very rapidly as the reaction proceeds. Practically, the reaction is carried out intermittently. The coke, contained in a large chamber, is brought to a white heat by burning a part of it in a blast of air, while the products of combustion are allowed to escape or are utilized as a fuel gas. The blast of air is then shut off, steam is turned on and the mixture of carbon monoxide and hydrogen, called " water gas " is collected for use. After a few minutes the mass cools below the temperature of rapid reaction. The steam is then shut off and the heating process repeated. By this method a gas having about one half of the heating value of a good illuminating gas can be manufactured very rapidly and cheaply. It is not suitable for use as an il- luminating gas, since hydrogen burns with a colorless flame and carbon monoxide with a blue flame which gives very little light. It may be enriched, however, by the addition of oil gas and in that form is used as illuminating gas in many cities of the United States. The most serious objection to its use is the very poisonous character of the carbon monoxide which it contains. Not only is this dangerous because of the acci- dental escape of gas from an open stopcock, but in the winter time gas may escape from leaks in pipes and may travel for some distance beneath the frozen ground till it finds an outlet in a cellar. The odor characteristic of illuminating gas is re- moved by passage through earth, still further increasing the danger. For these reasons some states forbid the manufac- ture of a gas containing more than a stated, small per cent of carbon monoxide. Producer Gas. The water-gas process is a wasteful one from the point of view of the per cent of the energy of the coke finally obtained in the water gas, largely because of the thick layer of fuel which must be used. Under such conditions the carbon is burned during the heating stage only to carbon monoxide, 298 A TEXTBOOK OF CHEMISTRY CO, the carbon dioxide which is formed in the lower part of the chamber being reduced to carbon monoxide above : CO 2 + C = 2CO The heats of combustion involved are as follows : Amorphous C + O 2 = CO 2 + 97,650 small calories CO -f O = CO 2 + 68,200 small callories Hence, amorphous C -f- O = CO -f- 29,450 small calories. It is evident from this that less than one third of the heat energy of coke is utilized when it is burned only to carbon monoxide. These heat relations, which are so unfavorable to the economy of the water-gas process may be utilized for the production of a low-grade fuel gas, commonly called " pro- ducer gas," which may, under favorable circumstances, retain from 80 to 85 per cent of the original heat energy of the coal or fuel employed in its manufacture. A chamber containing a thick bed of fuel has a blast of moist air forced through it in such a manner that a gas consisting chiefly of carbon monoxide and nitrogen, with a little hydrogen is obtained. The heat energy of such a gas can be much more perfectly utilized than that of solid fuel for many metallurgical operations, for the melt- ing of glass, for use in specially constructed gas engines and for many other purposes. Blast-furnace Gas. The reduction of oxides of iron by hydro- gen has been spoken of as a reversible reaction. The same is true when carbon monoxide is the reducing agent : Fe 2 O 3 + 3 CO ^t 2Fe -f 3 CO 2 Whichever reaction is used, the equilibrium of the reaction is so far to the left that the process can be successful only in the presence of a very large excess of hydrogen or carbon monox- ide. For this reason the gases escaping from the top of a blast furnace (p. 541) are of much the same nature as producer gas, with the advantage that the oxygen of the carbon monox- ide which they contain comes partly from the iron ore and so ILLUMINATING AND PRODUCER GAS 299 the per cent of nitrogen may be lower. This gas has long been utilized for heating the blast, generating steam, etc. During recent years it is coming into extensive use in gas engines. The following table illustrates the composition of the various kinds of gas which have been mentioned in this chapter : ENRICHED PRO- BLAST COAL OIL WATER DUCER FURNACE GAS GAS GAS GAS GAS Carbon dioxide, COo 1.1 3.0 1.5 11.4 Carbon monoxide, CO . . . 7.2 26.1 23.5 28.6 Hydrogen H2 49.0 [14.6] l 32.1 6.0 2.7 Methane CH 4 . . 345 388 19.8 3.0 0.2 "Heavy Hydrocarbons" . . 5.0 45.5 16.6 Nitrogen 3.2 1.1 2.4 66.0 57.1 Candle Power 17.5 65.0 25.0 1 Ethane, C 2 H 6 . Luminous Flames. It was formerly supposed that carbon separates in a flame because the hydrogen is more easily burned than the carbon. A study of the equilibrium between the gases present in a flame has shown, however, that this is not the case. It is only the carbon which results from the dissociation of hydrocarbons and which is momentarily heated to a very high temperature, which gives the luminous quality to the flame. In a gas flame burning from a round opening we may distinguish clearly three parts : 1. The interior cool portion of unburned gas. The head of a match inserted to this point enflames slowly or not at all. 2. A mantle of partial combustion and of dissociation of hydrocarbons, the luminous part of the flame. 3. A blue mantle surrounding the whole and especially notice- able at the base, where complete combustion to carbon dioxide and water occurs. The same parts may be seen in a candle flame, the gas or va- pors in the interior being formed by the heat of the flame acting on the materials of the candle. 300 A TEXTBOOK OF CHEMISTRY Bunsen Burner. A luminous flame may be made nonlumi- nous or slightly luminous by either of two methods. The flame may be diluted by an indifferent gas till it is cooled below the temperature of rapid dissociation for the hydrocarbons present, or oxygen may be introduced in such quantity that the carbon burns at once to carbon monoxide. Both effects are present in the Bunsen burner. The introduction of air at the base of the burner causes the combustion to take place in two stages. In the inner cone, Fig. 84, p. 303, the constituents of the gas are partially burned, giving a mixture of nitrogen, hydrogen, car- bon monoxide, carbon dioxide and water vapor. The last four substances are present above this cone in accordance with the equilibrium of the reversible reaction : At a temperature of about 850-900 the four compounds will be present in equal amounts, by volume, if the three ele- ments are present in the proportion given in the equation. Stated otherwise, at this temperature the reducing power of carbon monoxide is equal to that of hydrogen and the oxidiz- ing power of water vapor is the same as that of carbon dioxide. At lower temperatures the equilibrium is displaced to the right because the carbon monoxide becomes the stronger reducing agent, while at higher temperatures the equilibrium is displaced to the left. It is in accordance with this equilibrium that water gas prepared at a high temperature contains very little carbon dioxide. This is of considerable importance for an illuminating gas, because a small per cent of carbon dioxide greatly reduces the illuminating power. At the outer mantle of the Bunsen flame the carbon monoxide and hydrogen burn to carbon dioxide and water. By the arrangement shown in Fig. 82 it is possible to separate the two zones of the Bunsen flame from each other in such a way that the gases between the two zones may be drawn out and analyzed. By this method it has been shown that the equilibrium agrees with that of the water-gas reaction given above. ILLUMINATING AND PRODUCER GAS 301 In carrying out the experiment it is well to make the tube of the Bunsen burner, shown at the base of the figure, about 60 cm. long, to secure thorough mixture of the gas and air, and the proportion between the two must be properly regulated. If a spray of a solution containing a lithium and a copper salt is introduced into the gas, the lower cone will be colored red from the lithium and the upper cone green from /\ the copper, because lithium is oxidized by even the ' * small amount of oxygen in the water, carbon monoxide and carbon dioxide of the lower cone, while the copper is not oxidized till the freer oxygen supply of the upper cone is reached. Smithells and Ingle, J. Chem. Soc. 61, 204 (1902) ; Smithells, Phil. Mag. [5] 39, 132 (1895). When the proportion of air entering a Bunsen burner at the base is increased, the inner cone grows shorter until a point is reached where the flame " snaps back " and burns at the base. This is because the inner cone is to be looked upon as a stationary explosion wave where the velocity of the current of gases upwards is equal to the velocity of the combustion downwards. An increase in the proportion of oxygen increases this combustion velocity till it exceeds the velocity of the current of gas. Explosion Waves. When a mixture of gases is ex- ploded the combustion is not, of course, instantaneous, but proceeds with a definite velocity from the point of * ' ignition. This phenomenon has been carefully studied, partly by measurements of the velocity by stationary flames, as suggested above (Michelson, Wied. Ann. 37, 19 (1889)), partly by measuring the pressures developed (Michael and Le Chatelier, Ann. des Mines [8] 4, 379, 599 (1883)), partly by photographing the flame by its own light (Dixon, Ber. 38, 2426 (1905), Phil. Trans. 184, 97 (1893); 200, 315 (1903)). It has been found that there are two distinct stages in the explosion. At first the combustion proceeds with only a moderate velocity for a mixture of carbon monoxide and oxygen from 20 to 91 cm. per 302 A TEXTBOOK OF CHEMISTRY second, according to the composition of the mixture. The ex- pansion of the burning mixture by the heat compresses the gas in front of the explosion wave. This adiabatic 1 compression raises the temperature of the gas and, according to the prin- ciples of thermodynamics, in cases where the final volume is diminished by the explosion, lowers the explosion temperature. If the apparatus in which the explosion occurs is large enough, a point will finally be reached at which the compression raises the temperature of the mixture in front, to its explosion temper- ature. When this point is reached, the velocity suddenly changes and the flame proceeds with the velocity of sound. The pressure of the compression wave may become ten times that of the original mixture and it is estimated that temperatures of 5000-6000 may be produced. The velocity of sound at 5000 would be about 1400 meters per second and the velocities actually measured somewhat exceed this. This result is of considerable importance in connection with the design of gas or gasoline engines, where the form and dimensions of the explosion cylinder should be chosen so as to avoid the severe shock which comes with the second type of explosion waves. It also explains the shattering of a long eudiometer at the end farthest from the point of ignition. A glass tube of moderate thickness will readily withstand the explosion of a mixture of oxygen and hydrogen, if the dimensions are such as to avoid the explosion wave of the second type. Temperature of Flames. The temperature of the flame of a Bunsen burner varies from about 300 in the center, near the mouth of the burner, where combustion has not begun, to about 1550 in the portion between the inner cone and the outside of the flame. These temperatures are shown in detail in Fig. 84. In the Meker burner, Fig. 83, by widening the top of the burner and giving it a considerable number of fairly heavy metallic partitions, the inner cone is divided into a number of 1 Adiabatic means without escape of heat. Here the compres- sion is so rapid that the heat which results from the compression cannot escape. ILLUMINATING AND PRODUCER GAS 303 small and very short parts. This brings the high temperature of the upper part of the Bunsen flame down close to the mouth of the burner, concentrates the flame and gives it a more uni- form and somewhat higher temperature. The temperatures given in the figures are, of course, the tem- peratures of the flame when no substance radiating heat is 1540 1670* Fig. 83 Fig. 84 present. A platinum or porcelain crucible placed in the flame will be at a much lower temperature. A 20-gram platinum crucible placed 1 cm. above the Meker burner, with ordinary gas, will usually have a temperature of 900-950. Blowpipe. By means of a blowpipe (Fig. 85) the flame of a candle or of a Bunsen burner may be conveniently used for 304 A TEXTBOOK OF CHEMISTRY Fig. 85 heating or for oxidizing or reducing substances supported on charcoal or contained in beads of borax or of sodium meta- phosphate (from microcosmic salt). The interior of the flame, especially if it retains a slightly luminous character, will have a strong reducing effect, while at a point just beyond the tip of the flame, where substances are heated and at the same time can receive oxygen from the air, an oxidizing effect will be produced. These effects can be readily shown with litharge, PbO, and metallic lead. Similar effects can be ob- tained in the Bunsen flame. The tip of the inner cone is reducing, while the outer edge of the flame is oxidizing. Reversed Flames. Under ordinary conditions a flame con- sists of a combustible gas surrounded by air or oxygen with which it is combining. Because the oxygen of the air is available without expense further than the apparatus necessary to utilize it, we have become accustomed to speak and think of the combust- ible gas as the source of the energy which we use. We can, however, conceive of a world where the at- mosphere should consist of meth- ane or some other gas or gases which we call combustible. In such a universe we might obtain energy by preparing and burning oxygen. Such a condition can be illustrated with the apparatus shown in Fig. 86. When the illuminating gas entering through one tube is in excess, oxygen entering through the other will, if ignited, burn in the atmosphere of illuminating gas, while the ex- cess of the gas will burn at the end of the lamp chimney above. Gas Oxygett Fig. 86 ILLUMINATING AND PRODUCER GAS 305 EXERCISES 1. In what proportion must oxygen be mixed with the following gases, or vapors considered as gases, for their complete combustion? Methane, ethane, ethylene, acetylene, butane, benzene, gasoline if it has the average composition of heptane. 2. In what proportion must the same gases or vapors be mixed with air for their combustion ? 3. The heats of combustion of hydrogen, amorphous carbon and of some carbon compounds are : Hydrogen, H 2 + O 2 = H 2 O + 68,400 small calories. Amorphous carbon, C + O 2 = CO 2 + 97,650 small calories. Methane, CH 4 + 2 O 2 = CO 2 + 2 H 2 O + 214,000 small calories. Ethylene, C 2 H 4 + 3 O 2 = 2 CO 2 + 2 H 2 O + 333,350 small calories. Acetylene, C 2 H 2 + 2| O 2 = 2 CO 2 + H 2 O + 313,800 small calories. Benzene (vapor), C 6 H 6 + 7J O 2 = 6 CO 2 + 3 H 2 O + 799,350 small calories. These values are all for a final condition of liquid water at 18. If the following reactions could occur at ordinary temperatures, how much heat would be evolved or absorbed by each ? C + 2 H 2 = CH 4 C 2 H 4 = C + CH 4 C 2 H 2 = 2 C + H 2 C 2 H 4 = 2 C + 2 H 2 3 C 2 H 2 = C 6 H 6 C + H 2 = CO + H 2 C + 2 H 2 O = CO 2 + 2 H 2 CHAPTER XVIII OXIDES AND SULFIDES OF CARBON. ASSIMILATION AND RESPIRATION. CYANIDES. Carbon Dioxide. The sources of carbon dioxide in the air and the maintenance of a small constant amount, which fur- nishes carbon for the growth of plants, have been referred to in a previous chapter. For laboratory uses carbon dioxide is pre- pared by the decomposition of carbonates. Two properties of carbonic acid lead to the decomposition of carbonates when treated with almost any of the common acids. It is an exceed- ingly weak acid, the ionization: H 2 CO 3 ^H + -f-HCO 3 - taking place to only a very trifling degree, even in dilute solu- tions, and it is also very unstable, dissociating so readily to carbon dioxide and water : /O H 0=C 0=C=0 + H O H that the acid can exist only in solution, or possibly at very low temperatures. The carbonates usually employed in the laboratory are cal- cium carbonate, CaCO 3 , and sodium bicarbonate, NaHCO 3 . Carbon dioxide is also readily prepared by heating sodium bicarbonate, NaHCO 3 , magnesium carbonate, MgCO 3 , or cal- cium carbonate, CaCO 3 . The first dissociates to normal sodium carbonate, Na 2 CO 3 , carbon dioxide and water at a compara- tively low temperature. Magnesium carbonate requires a some- what higher temperature and calcium carbonate must be heated to 812, or very bright redness, before the dissociation pressure of the carbon dioxide is equal to atmospheric pressure. 306 CARBON DIOXIDE 307 Carbon dioxide is a colorless gas with a slightly sour taste and odor. It may be condensed to a liquid by pressure or to a solid by cold. The solid melts at 56.4 at a pressure of 5.1 atmospheres, while its vapor pressure is one atmosphere at 79. The gas can be liquefied, therefore, only under pressures greater than five atmospheres. If liquid carbon dioxide is allowed to escape from a cylinder in which it is kept under pressure, the evaporation of a portion of the liquid will cool the remainder below its freezing point and the solid carbon dioxide, which has a temperature of 79, can be collected in a small sack of closely woven cloth placed over the nozzle of the cylinder. Mixtures of the solid with alcohol, ether or acetone are very effective for securing low temperatures. The mixture with acetone boils at 88 and a temperature of 110 can be ob- tained by forcing air through it. Isothermals of Carbon Dioxide.. The critical temperature of carbon dioxide is 31.5. In the accompanying diagram, Fig. 87, VOLUMES. Fig. 87 308 A TEXTBOOK OF CHEMISTRY pressures are represented by the ordinates, volumes by the abscissas and temperatures by isothermal lines, which show the relation between pressure and volume for a given quantity of carbon dioxide. Such a diagram shows clearly all of the most important relations between pressure, volume and temperature for a typical gas. In the region of A where the pressures are moderate and the temperatures are very considerably above the temperature of liquefaction, for these pressures, the volumes are very nearly inversely proportional to the pressures, accord- ing to the law of Boyle, and the distance between the isother- mals either parallel to the axis of temperatures or parallel to the axis of pressures is closely proportional to the absolute tem- perature in accordance with the law of Charles. The higher the temperature and the lower the pressure, the more nearly are these laws accurate. In the region B there are two phases, liquid and vapor, and as the total volume is independent of the pressure, the isothermals are parallel to the axis of volumes. From this point the region where the gas laws are valid is reached either by an increase in temperature or a decrease in pressure. The latter can occur only when the liquid phase has disappeared. In the region C the isothermals are above the critical temperature and, while there is a flattening of the curves, showing some tendency to liquefy, there is no part parallel to the axis of volumes. In the region D the volume is less than that occupied by the liquid even at a considerably lower tem- perature, but it cannot be said that the substance is liquid, since no change of pressure, while the temperature is constant, will cause the separation into two phases. The coefficient of compressibility, however, approaches that of liquids. Density of Carbon Dioxide. The weight of a gram molecular volume of carbon dioxide is about 44, while that of air is about 28.9. Hence the gas is a little more than one half heavier than air. When the gas escapes somewhat rapidly from crevices in the earth, as is sometimes the case in wells and caves or mines, it may accumulate in sufficient amount to suffocate men or animals. In the Grotto del Cano in Italy the gas accumulates CARBONIC ACID 309 near the floor of the cave in such a manner that a dog entering the cave is suffocated, while a man with his head higher up may escape injury. The presence of a suffocating mixture in a well can be detected by lowering a candle into it, though a candle may be extinguished where respiration for some time is still possible. The conduct of carbon dioxide in wells and caves has given a popular impression that the gas will accumulate near the floor of a poorly ventilated room, but owing to the rapid diffusion of gases and because the conditions of breathing and flames cause an immediate mixture with the air, no appreciable accumulation of this sort can occur. Aqueous Solutions of Carbon Dioxide, Carbonic Acid. At ordinary temperatures water dissolves approximately its own volume of carbon dioxide (1.02 volume at 15, 0.88 volume at 20). In accordance with Henry's law the weight of the gas absorbed is very nearly proportional to the pressure and since, by the law of Boyle, the weight of a given volume of a gas is proportional to the pressure, water will absorb its own volume of the gas either at low or high pressures. By saturating water with the gas under high pressures a liquid is obtained from which carbon dioxide escapes with effervescence on relieving the pressure. Some mineral waters of this kind, especially Apolli- naris water, and Congress water from a spring in Saratoga, are found in nature, and similar waters are prepared artificially either for use directly or as the basis of the so-called " soda water." The carbonic acid, H^COs, formed by the union of the carbon dioxide with the water imparts to it a sightly sour taste. It also gives an acid reaction, which may be shown by the reddening of litmus or the discharge of the color of a faintly ilkaline solution of phenol phthalein. The carbon dioxide is ixpelled by boiling and the acid reaction disappears. Carbonates and Bicarbonates. Hard Waters. The acid character of a solution of carbon dioxide is also shown by the brmation of carbonates or bicarbonates (acid carbonates) with >ases : 310 A TEXTBOOK OF CHEMISTRY 2 NaOH + H 2 CO 3 = Na 2 CO 3 + 2 HOH Sodium Carbonate Ca(OH) 2 + H 2 C0 3 = CaCO 3 + 2 HOH NaOH + H 2 CO 3 = NaHCO 3 + HOH Sodium Bicarbonate The ionization of the bicarbonate ion : HC0 3 - ^ H + + C0 3 = is so very slight that very few carbonate ions can exist in aqueous solutions. Normal carbonates are therefore hydrolyzed by water, and their solutions have an alkaline reaction : Na 2 CO 3 ^ 2Na + + CO 3 = CO 3 = + HOH ^ HCO-T + OH- or Na 2 CO 3 + HOH ^ Na + + Na + + HCO-T + OH~ Bicarbonates may also be formed by the action of carbonic acid on carbonates : CaCO 3 + H 2 C0 3 = Ca(HC0 3 ) 2 The carbonates of calcium and magnesium are only very slightly soluble in water (CaCO 3 in 77,000 parts), but the acid carbonates, or bicarbonates, are much more easily soluble. Natural waters, which always contain carbonic acid, partly absorbed from the air but chiefly formed by the oxidation of the organic matter of the soils under the influence of bacteria, take up calcium carbonate and magnesium carbonate from the soils in the form of bicarbonates. Such waters are called " hard " waters. When boiled, owing to the ease with which the bicar- bonates dissociate into carbon dioxide, water and the normal carbonate, such waters give a deposit of calcium and magnesium carbonates, which forms the scale in teakettles and steam boilers. When the hardness is due to bicarbonate alone, it is called " temporary hardness " and can be largely removed by boiling the water: Ca(HCO 3 ) 2 = CaCO 3 + H 2 O + CO 2 CARBON MONOXIDE 311 Calcium sulfate, CaSO 4 , which is also present in many natural waters, is precipitated only when the water is concentrated or when it is heated to a high temperature under pressure. A water containing calcium sulfate is said to be "permanently hard." The precipitation at high temperatures is due to the decreased solubility of calcium sulfate under these conditions, and the scale formed is particularly adherent and objectionable. Such waters may be softened by the addition of an alkaline carbonate, phosphate, fluoride or borate, any one of which will precipitate the calcium as a nearly insoluble compound : CaSO 4 + Na 2 CO 3 = Na 2 SO 4 + CaCO 3 3 CaSO 4 + 2 Na 3 PO 4 = Ca 3 (PO 4 ) 2 + 3 Na 2 SO 4 CaS0 4 + 2 NaF = CaF 2 + Na 2 SO 4 CaSO 4 + Na 2 B 4 O 7 = CaB 4 O 7 + Na 2 SO 4 Carbon Monoxide. The formation of carbon monoxide in the manufacture of water gas and of fuel gas has already been discussed. The most familiar occurrence of the gas is probably in the burning of hard coal, where the carbon monoxide, formed by the reduction of carbon dioxide in the interior of the mass of coal, burns at the surface with a blue flame. Carbon monoxide is most easily prepared in the laboratory by the decomposition of oxalic acid, H 2 C 2 O 4 . Concentrated sulfuric acid assists in the decomposition, as a dehydrating agent : H 2 C 2 O 4 = H 2 O + CO 2 + CO The carbon dioxide must be removed from the mixture by means of soda lime or by passing the gases through a wash bottle containing of sodium or potassium hydroxide. Carbon monoxide is a colorless and odorless gas, which may be condensed to a liquid that boils at 190 and to a solid, which melts at 199. It burns with a characteristic blue flame, and mixtures of it with air or oxygen explode violently. Curiously enough, however, a perfectly dry mixture of carbon monoxide and oxygen will not explode, and the gao dried with phosphorus pentoxide will not burn in air or in oxygen which has been dried 312 A TEXTBOOK OF CHEMISTRY with the same agent. This recalls the fact that dry chlorine does not act on iron or copper. Many similar facts for which there is at present no very satisfactory explanation are known. The effect of a minute trace of moisture in promoting reactions suggests some connection with the effect of water in promoting reactions between electrolytes, but the two cases seem to be very radically different. One chemist has gone so far as to say that probably no reaction can occur between two substances which are absolutely pure but it is doubtful if such a generalization is justified. Carbon monoxide is very poisonous. Air containing one tenth of one per cent is distinctly dangerous if breathed for any length of time, and a smaller quantity, constantly present, would un- doubtedly cause chronic poisoning. It seems to combine with the hemoglobin of the blood and to alter it in such a way that the hemoglobin is no longer able to perform its proper function of combining with oxygen in the lungs and giving it up again for the oxidation of other substances in the tissues of the body. It acts as a cumulative poison, and recovery from its effects is often very slow. Carbon monoxide may be considered as the anhydride of formic acid, HCO 2 H (or H 2 CO 2 ), which might also be called carbonous acid. At 200-230, best under pressure, it combines directly with sodium hydroxide to form sodium formate : CO + NaOH = HCO 2 Na The Cycle of Carbon in Nature. The carbon dioxide of the air furnishes the great storehouse from which carbon finds its way into living bodies through the growth of plants. These are able to use the energy of the sunlight for the reduction of carbon dioxide, possibly first to formaldehyde, CH 2 O, which then unites with itself to form starch, (CeHuAs)^ The nitrogen which is necessary to form proteins and other compounds and the other elements required for the growth of plants, especially compounds of phosphorus, potassium, sodium, calcium, magnesium, silicon and iron, must be furnished by the soil. The compounds which RESPIRATION CALORIMETER 313 are built up by the growth of plants are ultimately broken down and their carbon is returned to the air as carbon dioxide by one of three processes : 1 . Wood or other vegetable material may be burned directly. 2. Vegetable substances may be used as food either by men or animals. After digestion and assimilation the carbon is sooner or later oxidized to carbon dioxide and is re- turned to the air, mostly through the lungs. 3. Vegetable or animal substances exposed to the action of bacteria decay, and the carbon is converted to carbon dioxide. Respiration Calorimeter. From the point of view of considera- tion of the transformations of energy, plants in their growth, by the reduction of carbon dioxide and the formation of com- pounds in which the carbon is combined with hydrogen and nitrogen as well as oxygen, store up energy received from the sun, in the form of combustible or edible carbon compounds. By burning and by the use of steam engines or other heat en- gines this stored chemical energy is transformed into heat energy or mechanical energy for practical use. When used as food by men or animals, the energy is also transformed both into heat and into muscular and mechanical energy. The amount of energy available in different kinds of food can be accurately determined by burning samples of them in a calorimeter (p. 25). In order to follow the transformations of energy which occur in the human body, the respiration calorimeter was devised by Professor W. O. Atwater with the aid of Dr. E. B. Rosa, and has been further developed by Benedict, Langworthy, and others. Figure 88 gives a cross section of one form of calorimeter which shows the general construction of the apparatus ; figure 89 shows an exterior view and some of the accessory apparatus for another form. The calorimeter consists of a small room in which a man may remain for experiments, which sometimes last for several days. The chamber is provided with tubes for the entrance and exit of the air necessary for respiration and also with pipes through which water can be circulated to take up the heat gen- erated by the body. The amount of food taken, any change in weight of the body which occurs, the amount of carbon dioxide 314 A TEXTBOOK OF CHEMISTRY Scale: 1 Meter Fig. 88. Horizontal cross-section of chair calorimeter, showing cross-section of copper wall at A, zinc wall at B, hair-felt at E, and asbestos outer wall at F ; also cross-section of all upright channels in the steel construction. At the right is the location of the ingoing and outgoing water and the thermometers. At C is shown the food aperture, and D is a gasket separating the two parts. The ingoing and outgoing air pipes are shown at the right inside the copper wall. The telephone is shown at the left, and in the center of the drawing is the chair with its foot- rest, G. In dotted line is shown the opening where the man enters. RESPIRATION CALORIMETER 315 evolved and the quantity of heat given out by the body are all carefully measured. The experiments involve, of course, many other details which cannot be given here. The investigations carried out with the calorimeter have demonstrated very Fig. 89 clearly that the amount of heat generated when food is digested and oxidized in the human body is the same, within the limits of error of the experiments, as the amount of heat generated by burning the same quantity of food in a calorimeter. In some of the experiments the man in the calorimeter ex- pended muscular energy in driving machinery, using the pedals of a bicycle for the purpose. It was found that during this period the amount of heat given oft 7 by his body was largely increased. Thus, while at rest his body, in one experiment, gave out 112 large calories per hour. When performing mechanical labor, on the other hand, 339 calories per hour were evolved, while the work done was 12,800 kilogrammeters, equivalent to 49 calories. It will be seen from this that only about 18 per cent 316 A TEXTBOOK OF CHEMISTRY of the additional expenditure of 276. calories required for the muscular work was actually converted into mechanical work. The human body, therefore, resembles a steam engine in dissi- pating as heat a large proportion of the energy required for its operation as a machine. It could not be shown by the respiration calorimeter that an appreciable amount of energy is required for mental work. In 22 experiments with persons performing mental work, as, for instance, in writing an examination or making arithmetical cal- culations, only one-half of one per cent more heat was given out from the body than that given out when the same individual was at rest. See Year Book of the United States Department of Agriculture for 1910, p. 307. The experiments seem to justify the conclusions that the compounds found in a living organism obey exactly the same chemical and physical laws as the same compounds outside of of the body. It cannot, however, be said that the difference between living and dead matter is simply a difference in physical and chemical properties. * Carbon Suboxide, C 3 O 2 , is formed when malonic acid, CH 2 (CO2H) 2 , is treated with phosphorus pentoxide, P 2 O5 : OH HH OH I \/ I O=C C C=0-* O=C=C=C=O + 2H 2 O Carbon Suboxide Carbon suboxide is a colorless liquid at a low temperature. It boils at 7 and has a very strong tendency to polymerize to a dark red substance, probably a mixture of compounds, having the same composition but evidently a much higher molecular weight. * Carbon Oxychloride, or Phosgene (Carbonyl Chloride) COC12. When a mixture of equal volumes of carbon monoxide and chlo- rine are exposed to the action of sunlight, the two gases unite to form carbonyl chloride. It is often called phosgene from this method of preparation (from <f>u<s, light and yewxw, to pro- duce). It is a colorless gas which may be easily condensed to CARBON BISULFIDE 317 a liquid that boils at 8.2. Carbonyl chloride is related to car- bonic acid just as sulfuryl chloride, SO 2 C1 2 , is to sulfuric acid. 0-H OH cl / C = -^ 0=0 ; OH Carbonic Acid As the chloride of an acid, it is hydrolyzed by water : COC1 2 + 2 H 2 -+ 2 HC1 + H 2 C0 3 -> CO 2 + H 2 O With ammonia it gives the amide of carbonic acid, urea : v - y \~ v ^ y X C1 ^0 OH Cl Carbonyl Chloride Sulfuric Acid Sulphuryl Chloride + 4 NH 3 = CO + 2 NH 4 C1 C1 X NH 2 Urea Carbon Bisulfide. When sulfur vapor is passed over heated charcoal, the two elements combine to form carbon bisulfide, CS 2 , which is now manufactured in electric furnaces. It is a volatile, inflammable liquid, which boils at 46.25. It takes fire so easily that the vapors ignite when a glass rod, which has been warmed gently, is held over a dish containing the liquid. As a mixture of the vapor with air is explosive, this property makes extreme care necessary in factories where the substance is used. Carbon bisulfide is a good solvent for fats and india rubber, also for ordinary phosphorus, sulfur and iodine. It is used in vul- canizing india rubber, in the preparation of rubber cements and in extracting grease from wool. It is poisonous and is sometimes used to kill rats and ground squirrels. * Sulfocarbonates. Carbon bisulfide dissolves in solutions of alkaline sulfides, forming sulfocarbonates exactly as carbon dioxide forms carbonates with alkaline hydroxides : 2 KOH + CO 2 = K 2 C0 3 + H 2 O Potassium Carbonate K 2 S -|- CS 2 = K 2 CSs 318 A TEXTBOOK OF CHEMISTRY or 2 KHS + CS 2 = K 2 CS 3 + H 2 S Potassium Sulfocarbonate Potassium sulfocarbonate is a yellow salt, easily soluble in water. It is used to destroy the phylloxera which sometimes cause great damage to grapevines. * Sulfocarbonic Acid, H 2 CS 3 , separates as an oily liquid when hydrochloric acid is added to a solution of a sulfocarbonate. It soon decomposes into carbon disulfide and hydrogen sulfide. K 2 CS 3 + 2 HC1 = 2 KC1 + H 2 CS 3 -> H 2 S + CS 2 The use of sulfocarbonates as germicides depends on a similar decomposition by means of the carbon dioxide of the air. * Carbon Oxysulfide, COS, is a compound intermediate be- tween carbon dioxide, CO 2 , and carbon bisulfide, CS 2 . It may be formed by the direct union of carbon monoxide, CO, and sul- fur, but it is best prepared by the decomposition of a thiocyanate by means of an acid : KCNS + HC1 = KC1 + HCNS Potassium Thiocyanic Thiocyanate Acid H N4=CS = NH 3 + O=C=S H 2 =fO or KCNS + 2 HC1 +H 2 O = COS + KC1 + NH 4 C1 Carbon oxysulfide is a colorless, odorless gas which may be condensed to a liquid that boils at 46.5, a boiling point inter- mediate between those of carbon dioxide and carbon bisulfide, but much nearer to the boiling point of the former. It burns in air to carbon dioxide and sulfur dioxide. It is hydrolyzed by water to thiocarbonic acid, which then decomposes into carbon dioxide and hydrogen sulfide : /SH /H O=C=S -f HOH -> O=C -> O=C=O + S< \ Thiocarbonic Acid OXIDES AND SULFIDES OF CARBON 319 Cyanides. Potassium cyanide, KCN, may be formed by the direct union of the three elements at a high temperature, as when nitrogen is passed over a mixture of carbon and potassium car- bonate at a white heat : K 2 CO 3 + 4 C + N 2 = 2 KCN + 3 CO If organic matter containing nitrogen is heated with potassium carbonate and iron, a somewhat similar reaction occurs at a com- paratively low temperature, with the formation of a double cyanide of ferrous iron and potassium, called potassium fer- rocyanide, K 4 FeC 6 N 6 (or Fe(CN) 2 .4 KCN). Cyanides are formed in a similar manner in the manufacture of illuminating gas and pass over into the ammoniacal gas liquors, from which they are recovered commercially. Hydrocyanic Acid or Prussic Acid, HCN, may be obtained by distilling a solution of potassium ferrocyanide or potassium cya- nide with dilute sulfuric acid. It is a liquid at low tempera- tures, which boils at 26.5. Hydrocyanic acid and its salts are among the most violent poisons known and should always be handled with extreme care. The acid has a characteristic odor, but some persons seem to be unable to perceive it , and such individuals need to be especially careful in handling the acid or cyanides. Hydrocyanic acid is so weak an acid that it is liberated from its salts by the carbon dioxide of the air, and the simple cyanides have the odor characteristic of the acid. A dilute solution of hydrocyanic acid is sometimes used in medicine. Potassium Cyanide, KCN, or rather -a, mixture of potassium and sodium cyanides, is manufactured commercially by heating potassium ferrocyanide with metallic sodium : K 4 FeC 6 N 6 + 2 Na = 4 KCN + 2 NaCN + Fe The salt is used for many purposes in the laboratory, es- pecially in the preparation of organic compounds. It is also used for the extraction of gold from its ores and in silver plating. Complex Cyanides. The cyanides of the heavy metals, such as silver, iron, zinc, etc., are most of them insoluble in water, but 320 A TEXTBOOK OF CHEMISTRY many of them will dissolve in a solution of potassium cyanide, ICCN. In the solutions obtained in this manner the atoms of silver or iron no longer conduct themselves as ions. 1 They will give no precipitate with reagents which precipitate them from their ordinary salts, as sodium chloride for silver or ammonium sulfide for iron. When an electric current passes through such ?. solution, the silver or iron travels with the cyanide ion toward the anode, while only the potassium goes toward the cathode. Finally, if the solutions are evaporated, definite crystalline compounds, potassium argenticyanide, KAgC2N2, and potas- sium ferrocyanide, K^FeCeNe.S H 2 O, can be obtained. These facts indicate that such solutions contain complex argenticyanide (AgC2N2~) and ferrocyanide (FeCeNe") ions which hold together in solution as the elements of the sulfate (SO4 = ) or nitrate (NO 3 ~) ions do. In further agreement with this interpretation, such solutions give with various solutions precipitates containing the characteristic complex group. A considerable number of fer- rocyanides, for instance, may be obtained in this way. Ferric salts give with potassium cyanide, potassium ferricy- anide, K 3 FeC 6 N 6 (or FeC 3 N 3 .3 KCN), a salt which forms red crystals, while the crystals of the ferrocyanide are yellow. The ferrocyanide may be easily oxidized to the ferricyanide and the latter can be reduced to the ferrocyanide : 2 K 4 FeC 6 N 6 + O + H 2 O = 2 K 3 FeC 6 N 6 + 2 KOH K 3 FeC 6 N 6 + H + KOH = K 4 FeC 6 N 6 + H 2 O Potassium ferrocyanide gives with ferric salts a blue precipi- tate of Prussian blue : 3 K4Fe n C 6 N 6 + 4 Fe m Cl 3 = Fe 4 III (Fe II C 6 N 6 ) 3 + 12 KC1 Prussian Blue Potassium ferricyanide gives a similar deep blue precipitate with ferrous salts, but this is of variable composition and usually contains potassium. 1 A very small number of ions of these metals are doubtless present. In some solutions of this type hydrogen sulfide will give a precipitate, owing to the extreme insolubility of the sulfide formed. COMPLEX CYANIDES . 321 The ferric ferrocyanide is decomposed by potassium hydroxide with separation of ferric hydroxide while potassium ferrocyanide passes into solution. When a smaller amount of ferric chloride is added to a solution of potassium ferrocyanide a deep blue solution containing potas- sium ferric ferrocyanide is obtained : K4Fe n C 6 N 6 + Fe m Cl 3 = KFe m Fe n C 6 N 6 + 3 KC1 Potassium Ferric Ferrocyanide In this solution a ferrous salt gives a precipitate of Turnbull's blue : 2 KFe m Fe 11 C 6 N 6 -f Fe 11 C1 2 = Fe 11 (Fe m Fe 11 C 6 N 6 ) 2 + 2 KC1 Turnbull's Blue The empirical formula of Prussian blue is FeyCigNis, while that of Turnbull's blue is FesC^N^. The empirical formula of ferrous ferricyanide is also FesCi^N^. Both Prussian blue and Turnbull's blue are used as blue pig- ments and for " blueing " for laundry purposes. A solution of potassium argenticyanide, KAgC 2 N 2 is used in manufacturing silver-plated ware. While the silver is trans- ferred toward the anode through the solution, it is also deposited on the cathode, and the fact that there are very few silver ions in the solution for some reason causes a smooth, coherent de- posit, while the silver deposited from a solution of silver nitrate usually assumes a crystalline form. From 'the silver anode, which is used in the electrolysis, silver passes into solution, re- generating the potassium argenticyanide which is decomposed at the cathode. * Potassium Cyanate, KCNO, may be prepared by heating potassium cyanide with lead oxide. It is very poisonous : KCN + PbO = KCNO + Pb * Potassium Thiocyanate, KCNS, is formed when a mixture of potassium cyanide and sulfur is heated. Potassium thio- cyanate and ammonium thiocyanate, NH 4 CNS, are used in testing solutions for the presence of iron in the ferric form because 322 . A TEXTBOOK OF CHEMISTRY of the intense red color of solutions containing ferric thiocyanate, Fe(CNS) 3 . Cyanogen, C 2 N 2 . When mercuric cyanide, HgC 2 N 2 , is heated, it decomposes into mercury and cyanogen, C 2 N 2 , just as mercuric oxidede composes into mercury and oxygen. Cyanogen is a colorless, poisonous gas, which burns with a characteristic pink flame. EXERCISES 1. Carbon dioxide melts at 56.4 under a pressure of 5.1 atmos- pheres. Under a pressure of 515 atmospheres (500 kg. per sq. cm.) it melts at 47.4. Is the density of solid carbon dioxide greater or less than that of the liquid when both are present together ? (See principle of van't Hoff-Le Chatelier, p. 111.) 2. Air contains, normally, about 3 parts in 10,000 of carbon dioxide. In accordance with Henry's law what weight of the gas will be absorbed by one liter of water in contact with such air at 20. (See p. 309 for the solubility of the gas in water at this temperature.) 3. What volume of normal sodium hydroxide will be neutralized with formation of sodium bicarbonate, NaHCOs, by one liter of water which has been in contact with ordinary air ? 4. An effervescent drink is sometimes prepared by mixing two solu- tions containing sodium bicarbonate and cream of tartar (HKC^Oe). What is the equation for the reaction ? 5. A water contains 0.130 g. per liter of calcium sulfate, CaSO4. How may grams of crystallized sodium carbonate (Na 2 CO 3 .10 H 2 O) per cubic meter will be required to soften the water ? How many grams of sodium fluoride ? 6. How many grams of potassium cyanide will be required to reduce 15 g. of stannic oxide, SnO 2 , to metallic tin ? CHAPTER XIX ALCOHOLS, ALDEHYDES, KETONES, ACIDS, FATS, CARBO- HYDRATES THE most important classes of compounds of carbon with oxygen and hydrogen are given in the heading of this chapter. The number of such compounds is very large and even a super- ficial knowledge of them can be gained only by a study of their structure, that is, by learning the arrangement of the atoms within their molecules. This is especially true because many different compounds having the same composition are known. Thus there are no less than seventy-five compounds having the formula C 7 Hi 4 O 2 . Compounds like these which have the same composition but different properties are called isomers. The empirical formula of such a compound will give very little infor- mation about its properties, but a structural formula often shows to a chemist, at once, many important relationships. Structural formulas are established mainly on the basis of three very simple principles : 1. Valence. The valence of the elements. Carbon is almost always quadrivalent, oxygen is bivalent and hydrogen univalent. Nitrogen may be trivalent or quinquivalent. An example of the use of the principle of valence in explaining and predicting the structure of the hydrocarbons has been given (p. 284). 2. Radicals. Groups of atoms called radicals hold together and retain their order of arrangement in passing from one com- pound to another. Thus in the reaction represented by the equation : C 2 H 6 O + HI = C 2 H 5 I + HOH Ethyl Ethyl Alcohol - Iodide the group, or radical, ethyl, C 2 H 5 , is supposed to retain the same arrangement of its atoms in ethyl iodide that it has in ethyl 323 324 A TEXTBOOK OF CHEMISTRY alcohol and on the basis of this reaction we usually write the formula of ethyl alcohol as C 2 H 5 OH. On the other hand methyl ether, which has the same composi- tion as ethyl alcohol, reacts with hydriodic acid thus : C 2 H 6 O + HI = CH 3 I + CH 3 OH Methyl Methyl Methyl Ether Iodide Alcohol On the basis of this reaction we give to methyl ether the formula CH 3 O CH 3 . Expanding these formulas, using the principle of valence, we have : H H H H H C C O H and H C O C H A Ethyl Alcohol Methyl Ether 3. Substitution. When an atom or group leaves a compound and another atom or group enters it, the group which enters takes the place which was occupied by the atom or group which has left. In the reaction between methyl ether and hydriodic acid given above the group CH 3 O leaves the compound and an iodine atom takes its place. Neither of the last two principles is universal in its application and there is some uncertainty as to the valence of the elements in some compounds, but, in spite of this, it has been possible to determine the structure of very many carbon compounds with practical certainty. It is impracticable within the limited space which seems suitable for this book to give the basis for the struc- ture assigned to the compounds mentioned, and in many cases only empirical formulas will be used. Alcohols. Alcohols may be defined as hydrocarbons in which one or more hydrogen atoms have been replaced by hydroxyl, OH. Methyl Alcohol, CH G OH, is obtained along with acetic acid and acetone by the destructive distillation of wood, and in an impure form it is often called wood spirit or wood alcohol. It is now used chiefly as an addition to ordinary alcohol to denaturize ALCOHOLS 325 it. (See below.) It is very much more poisonous than ordi- nary alcohol and fatalities have, often occurred from drinking it or from breathing its vapor when varnishes made with it were used in a confined space. Blindness sometimes results from drinking or breathing it. Methyl alcohol boils at 66 and has a specific gravity of 0.7931 at 15.6. Ethyl Alcohol, C 2 H 5 OH, is formed by the fermentation with yeast of liquids containing either ordinary cane sugar, as the juice of grapes or apples, or diluted sirups from the manufacture of sugar, or maltose or glucose, sugars formed by the action of malt on the starch of Indian corn or potatoes : H 2 O = 4 C 2 H 5 OH + 4 CO 2 Maltose Alcohol Only dilute alcohol can be obtained by fermentation and strong alcohol is prepared by fractional distillation, the boiling point of alcohol being 78.3. Absolute alcohol, or alcohol free from water, is prepared by allowing concentrated alcohol to stand with quicklime, CaO, which combines with the last portions of the water. The action of the lime may be hastened by warming the mixture. The specific gravity of absolute alcohol is 0.7933 at 15, referred to water at 4. The most common beverages containing alcohol are beer (3.5-7.5 per cent of alcohol by volume), wine (9-17 per cent), cider (3.5-7.5 per cent), brandy, from the distillation of wine (44-55 per cent), whisky, from the distillation of alcoholic liquids made from grains (46-55 per cent) and rum (30-50 per cent). The per cents by weight are approximately four fifths of these. Alcohol is extensively used for burning, as a solvent, in mak- ing varnishes and in making pharmaceutical extracts, ether and other products. " Denatured alcohol " is an alcohol to which some substance has been added to render it unsuitable for use as a beverage. Wood alcohol, gasoline and bone oil are the most common additions. Such an alcohol is sold free of tax and may be used 326 A TEXTBOOK OF CHEMISTRY for burning, for the manufacture of varnishes and for many other purposes. The substances added render it poisonous and unfit for drinking or for any medicinal use. Phenol or Carbolic Acid, C 6 H 5 OH, is obtained from coal tar and can be made artificially from benzene, C 6 H 6 . The pure compound is solid at ordinary temperatures but it liquefies on the addition of a small amount of water. It has been exten- sively used as a disinfectant and was the substance first used in antiseptic surgery. It is effective only when applied in a solution. The boiling point is high and the vapor is never sufficiently concentrated to be of service. Pure phenol cor- rodes the tissues and is a violent poison. Phenol is the chief active constituent of the " coal-tar dips " used in the care of sheep. CH 2 OH Glycerol, 1 CHOH , or C 3 H 5 (OH) 3 , is an alcohol containing CH 2 OH three hydroxyl groups. It is obtained as a by-product in the manufacture of soap from fats (p. 332). When glycerol is treated with nitric acid under proper conditions, it reacts with it, forming glyceryl nitrate (nitroglycerin) : C 3 H 5 (OH) 3 + 3 HNO 3 = C 3 H 5 (NO 3 ) 3 + 3 HOH Glyceryl Nitrate Nitroglycerin has a very much larger amount of chemical energy than the same elements combined in the form of carbon dioxide, water and free nitrogen, and under the influence of a denotating cap of fulminate of mercury the rearrangement of the atoms to the more stable, gaseous forms takes place so suddenly as to cause a violent explosion. It is used chiefly for blasting pur- poses. The liquid form is sometimes used directly, but usually 1 Glycerol is commonly known as "glycerin." The name glycerol is used by many writers and is to be preferred because the ending "ol" is used for names of alcohols in general. ALDEHYDES AND KETONES 327 it is absorbed in sawdust or some other porous material and is then known as dynamite. The conduct of nitroglycerin illustrates in an unusually striking manner the fact that very many carbon compounds, perhaps even the majority of them, are in a condition of unstable equilibrium. The same elements may be combined in some other form which contains less chemical energy. This seems to be because the atoms do not separate and recombine so easily as is usual with the compounds of other elements. This slow- ness of reaction velocities and the fact that there are enormous differences in the velocity of different possible reactions, so that the compounds which are formed seldom represent the lowest content of chemical energy, are certainly very important factors in the formation of the extraordinary number of carbon com- pounds. Aldehydes and Ketones. By the oxidation of an alcohol /H containing the group C-H , an aldehyde may be ob- \O H tained. For this the characteristic group is C<f , com- X H bined with a single carbon atom. In a similar manner an alcohol containing the group /C<f gives a ketone, which X)H has the group yC=O combined with two carbon atoms. Formaldehyde, H C^ (or H 2 CO), is easily prepared by the oxidation of methyl alcohol. A solution containing about 40 per cent of formaldehyde is known, commercially, as for- malin and is much used as a disinfectant. For disinfecting rooms it must be disseminated in the form of a spray, or vaporized. It has the great advantage over sulfur dioxide, which was for- merly used for the same purpose, that it does not cause the dark- ening of metallic objects or injury to fabrics. It is a very power- 328 A TEXTBOOK OF CHEMISTRY ful germicide and has been sometimes used as a preservative in milk and other articles of food. Such a use is very objec- tionable and is forbidden by law. Formaldehyde is a powerful poison and may produce painful wounds on the skin. s Benzaldehyde, C 6 H 5 C< , the chief constituent of oil of \H bitter almonds and used as a flavoring extract, is one of the most common aldehydes. Another aldehyde is Citral, CH 3 v /CH 3 .0 \C=CH-CH 2 CH 2 -C=CH-C<f , CH/ X H obtained from the oil of lemons and used as the starting point for the manufacture of ionone, an artificial substitute for the odor and the flavor of violets. CH 3 x Acetone, ^>C=O, is formed along with methyl alcohol CH 3 / and acetic acid by the destructive distillation of wood. It is also obtained by distilling calcium acetate or barium acetate. It is used in the manufacture of chloroform and as a solvent for acetylene. Acetone dissolves fats and is also used in the manu- facture of varnishes. Acids. The further oxidation of an aldehyde gives an acid, for which the characteristic group is called carboxyl, C OH The change from an alcohol to an acid is somewhat similar to /H the change from hydroxyl amine, N(-H , to nitrous acid, \OH . In both cases the replacement of the two hydro- H gen atoms by an oxygen atom gives acid properties to the hydro- gen of the hydroxyl group. Very many acids are found among natural products and in articles of food. ACIDS 329 Formic Acid, H C<T , was first obtained from ants and X H is, in part at least, the cause of the irritation from the sting of the bee. It is now manufactured on a large scale by absorbing carbon monoxide with sodium hydroxide at a temperature of 190-220 : NaOH + CO = HCOONa It is extensively used in making dyes and for other manu- facturing purposes as a substitute for acetic acid. > Acetic Acid, CH 3 Cf (or HC 2 H 3 O 2 ), is the acid of X)H vinegar. It is formed by the action of Bacterium aceti or "mother of vinegar " on alcohol, either slowly, as with wine or cider stored in casks which are partly open to the air, or more rapidly by allowing an alcoholic liquid to trickle over wood shavings which have been inoculated with the bacteria. Good vinegar contains 4 per cent of acetic acid. Acetic acid is also formed by the destructive distillation of wood, and the crude acid ob- tained in this way is sometimes called pyroligneous acid. This crude acid, containing phenol and other substances which ree'n- force the antiseptic properties of the acetic acid, is sold under the name of " liquid smoke " as a preservative for meat. c JNOH Oxalic Acid, Q (or H 2 C2O 4 ), is found in sheep sorrel and C/ X OH some other plants, but for technical purposes it is manufactured artifically by heating sodium formate with a catalyzer : CO 2 Na 2 HCO 2 Na = | +H 2 Sodium Formate CO 2 Na Sodium Oxalate 330 A TEXTBOOK OF CHEMISTRY Oxalic acid is a comparatively strong acid and is quite poisonous. It is used in calico printing and dyeing and in the bleaching of flax, straw and leather. The decomposition of oxalic acid into carbon dioxide, carbon monoxide and water when heated with concentrated sulfuric acid has been given (p. 311). The ammonium salt, (NH4) 2 C2O4.H 2 O, is used to precipitate calcium because calcium oxalate, CaC 2 O 4 , is only very slightly soluble. An acid potassium salt called potassium tetraoxalate, KH 3 (C 2 O4) 2 .2 H 2 O, derived from the doubled molecule, is sometimes used as a standard for alkalimetry. Lactic Acid, CH 3 CHOHCO 2 H (or HC 3 H 6 O 3 ), is formed by the fermentation of milk sugar and is found in sour milk. It is both an alcohol and an acid, as is apparent from its structural formula. CHOH C0 2 H Tartaric Acid, | (or H 2 C 4 H 4 O 6 ), is found in CHOH C0 2 H grape juice. When the sugar of the juice ferments, forming alcohol, the acid potassium salt, KHC4H 4 O 6 , which is only slightly soluble in dilute alcohol, separates in a crude form called argoL From this the pure salt, known as cream of tartar, and the free acid are prepared. Cream of tartar is used in cooking and sometimes to aid in the formation of jelly from fruit juices. Mixed with sodium bicarbonate, NaHCO 3 , and flour or starch, it is used in making the more expensive kinds of baking powders. Potassium antimonyl tartrate, KSbOC 4 H 4 O 6 , is called tartar emetic and is sometimes used as an emetic. CH 2 CO 2 H Citric Acid, C(OH)CO 2 H (or HaCeHsOy), is found in the juice CH 2 CO 2 H of lemons and is the constituent which gives the sharp, sour taste to the fruit. A cold, concentrated solution of ammonium citrate, FATS 331 will dissolve some of the acid calcium phos- phates which are insoluble in water, while it will not dissolve tricalcium phosphate, Ca 3 (PO 4 )2. Its action in this respect is supposed to resemble the action of water containing carbon dioxide in the soil. Such a solution is used in the analysis of commercial fertilizers to distinguish between phosphates which are supposed to be readily available for the growth of crops and those which are only slowly available. Ammonium Ferric Citrate, (NH^HFeCCeHsOy^, is used in the preparation of " blue-print " paper. The paper is moistened with a solution containing a mixture of ammonium ferric citrate and potassium ferricyanide. When dry it is exposed to the action of light under a " negative." The two salts are easily soluble in water, but in the light the ferric salt or the ferricyanide is reduced by the citric acid. In either case, a blue, insoluble compound is formed (p. 320). On washing the paper with water the soluble salts are removed, while the portions exposed to the light remain a permanent blue. Benzoic Acid, C6H 5 CO 2 H, is found in cranberries and in some other fruits. It is also manufactured by the oxidation of toluene, CeHsCHa, from coal tar. Even in dilute solutions it destroys or prevents the growth of bacteria and it has been much used for this purpose. At the present time there is some difference of opinion among authorities as to whether small amounts of the acid taken with food have an injurious effect or not. Its use as a substitute for cleanliness or to disguise inferior products is, of course, condemned by every one. Palmitic, Stearic and Oleic Acids, Fats. The natural fats, such as lard, tallow, butter, olive oil, cottonseed oil, linseed oil, and the oils found in nuts and cereals all contain compounds in which the hydrogen of three molecules of a monobasic or- ganic acid has been replaced by the trivalent radical glyceryl, CsHs. The most common of these fatty acids are palmitic acid, HCi 6 H3iO 2 , stearic acid, HCigHssC^, and oleic acid, HCisHsaC^. The corresponding compounds found in fats are palmitin, CsH^CieHs 102)3, stearin, CsH^CisH^^a, and 332 A TEXTBOOK OF CHEMISTRY olein, C 3 H 5 (Ci8H 33 O 2 ) 3 . All three of these are found in lard. Stearin is found especially in tallow, palmitin in palm oil, olein in olive oil. The substance known commercially as stearin is in reality an impure stearic acid. It is used in laundries and for the manufacture of candles. Stearic and palmitic acids are solid at ordinary temperatures. Oleic acid is a liquid. Soaps. When fats are heated with a concentrated solution of sodium hydroxide, they are decomposed, forming a sodium salt of the fatty acid, and glycerol : C 3 H5(C 18 H350 2 ) 3 + 3 NaOH = 3 NaCi 8 H 3 5O 2 + C 3 H 5 (OH) 3 Stearin Sodium Stearate This process is called saponification and the sodium salts form the chief constituents of the ordinary soaps. The sodium salts formed by the saponification are separated from the aqueous solu- tion containing the glycerol by the addition of a concentrated solution of salt, in which the salts of the fatty acids are nearly insoluble. From the aqueous solution the glycerol is recovered by evaporation and distillation under diminished pressure. The action of soap depends on the fact that water containing soap in solution readily forms an emulsion with oily or greasy substances and so aids in their removal from fabrics or from the skin. The calcium salts of the fatty acids are mostly so insoluble in water that they cannot act in this way ; and when soap is used with a water containing calcium salts in solution it can have little effect till enough of the soap has been used to pre- cipitate all of the calcium as calcium stearate, Ca(Ci 8 H 35 O 2 ) 2 , or in the form of similar compounds. The separation of insoluble calcium salts when soap is used with hard waters is, of course, a familiar experience. Carbohydrates. A large class of organic substances found in plants is made up of compounds which contain oxygen and hydrogen in the same proportion as in water. On account of this composition these compounds are called carbohydrates, CARBOHYDRATES 333 but they do not contain oxygen and hydrogen in the form of water. Many of these compounds contain molecules with six, twelve, eighteen, or some multiple of six carbon atoms. The most important of the carbohydrates are the sugars, dextrins, starch and cellulose. Cane Sugar or Saccharose, Ci 2 H 2 2On, is found in the juice of the sugar cane, in sugar beets, in the sap of the maple tree and in almost all sweet fruits and vegetables. From the sugar cane the juice is obtained by pressing the cane between heavy steel rollers. The juice is concentrated under diminished pressure, to avoid the decomposition which would occur if the solution were boiled in the open air at atmospheric pressure. The partly concentrated solution is filtered through boneblack or animal charcoal to remove coloring matters, and is then evap- orated further till sufficiently concentrated so that the sugar will crystallize on cooling. The crystals are separated from the colored sirup by means of' rapidly rotating centrifugal strainers. The manufacture of sugar from beets differs in many important details, but the general principles are the same. For maple sugar the juice is simply evaporated till the sirup is sufficiently concentrated to crystallize on cooling, the other substances present being of such a character as to give the sugar a desirable flavor. If purified by the methods described above, maple sugar would not differ from the sugar from sugar cane or sugar beets. Cane sugar crystallizes in well-formed monoclinic crystals, seen especially in rock candy. It melts at 160, and if heated for a short time at that temperature or a little higher it is partly decomposed, giving a dark brown substance of indefinite com- position called caramel. A solution of cane sugar turns the plane of polarization of a ray of polarized light, which passes through it, to the right. The degree of rotation is almost exactly proportional to the concentration of the solution, and the measurement of the rota- tion in specially constructed polarimeters called saccharimeters is very much used as a basis for the control of operations in 334 A TEXTBOOK OF CHEMISTRY sugar factories and for the collection of duty on sugar at ports of entry. When warmed for a short time with a dilute acid, cane sugar is hydrolyzed to a mixture of glucose and fructose : Ci 2 H 22 Qn + H 2 O = C 6 Hi 2 6 + C 6 Hi 2 O 6 Glucose Fructose Glucose rotates the plane of polarized light to the right but fructose rotates the plane in a greater degree to the left, and the mixture has a levo-rotation. It is called for this reason invert sugar. The ease with which the hydrolysis takes place is a source of very considerable loss in the manufacture of sugar. A similar hydrolysis often occurs in fruit juices and in honey. Maltose, Ci 2 H 22 Ou, is formed along with maltodextrin by the action of the enzyme diastase (p. 344) on starch. Its formation is a very important step in the manufacture of alcohol from corn and other grains (p. 325) . Maltose is hydrolyzed to glucose by the action of dilute acids. Lactose, or Milk Sugar, Ci 2 H 22 On, is found in milk and can be obtained from whey as a by-product in the manufacture of cheese. It forms an important constituent of milk, as a food, and is used by preference, rather than cane sugar, as an addition to cows' milk for feeding infants. Glucose, CeH^Oe, is formed together with fructose by the hydrolysis of cane sugar. It is also formed by the hydrolysis of starch with dilute acids : (C 6 HioO 5 ) n + n H 2 O = n C 6 Hi 2 O 6 Starch Glucose For the commercial manufacture dilute sulfuric acid is usually employed because on the subsequent addition of cal- cium carbonate the acid can be almost completely removed as the difficultly soluble calcium sulfate. Glucose, when pure, has about three fifths the sweetening power of the same weight of cane sugar. It is used in the manufacture of " corn sirup," in fruit preserves and in cheap grades of candy. The popular impression that glucose is harmful as an article of diet seems to have no experimental basis. Glucose is dextrorotatory and CARBOHYDRATES 335 was formerly called dextrose, a name still used by some au- thors. In the disease called diabetes sugar and starch of the food which is eaten are changed to glucose and eliminated in the urine instead of being assimilated as they should be. The glucose can be detected by means of Fehling's solution, 1 a solu- tion containing potassium sodium tartrate, KNaC 4 H 4 O 6 , copper sulfate, CuSO4, and sodium hydroxide. Glucose reduces the copper of such a solution to cuprous oxide, Cu2O, which sepa- rates as a red precipitate when the mixture is boiled. Fructose, CeH^Oe, is the second constituent of invert sugar (see above). It is levorotatory and was formerly often called levulose, but, since a second, exactly similar compound, which is dextrorotatory, is known, the designation fructose is preferred. Both glucose and fructose may be fermented to alcohol and carbon dioxide by the action of yeast. Starch, (CeHioOs)^, is found in the form of granules (Fig. 90) which differ very markedly in their organized structure but which, so far as is known, are identical in their chemical com- position. It is an important constituent of potatoes, Indian corn, rice, wheat, tapioca and many other cereals and vegetables. When the flour of a cereal is kneaded in a current of water, the fine starch granules float away, while most of the other con- stituents remain behind. The residue consists chiefly of ni- trogenous substances and is called gluten. From the water carry- ing the starch granules in suspension the latter will settle out on standing or on allowing the water to flow slowly over tables very slightly inclined. The practical manufacture of starch in- volves, of course, many other details which need not be given here. Starch is the most important non-nitrogenous constituent of foods. The granules are covered with a thin coating which 1 Fehling's solution may be prepared by mixing equal volumes of two solutions containing: 1. 34.65 grams of copper sulfate (CuSO 4 .5 H 2 O) in 500 cc. of water; 2. 173 grams of Rochelle salt (KNaC^Oe.HaO) and 50 grams of sodium hydroxide in 500 cc. of water. The copper of 1 cc. of the mixed solution will be precipitated by about 0.005 gram of glucose. 336 A TEXTBOOK OF CHEMISTRY interferes with their digestion, and one of the most important effects produced by baking bread and cooking cereals and vege- tables is the bursting of the granules by the combined effect of heat and moisture. In the process of digestion starch seems to Fig. 90. A, potato starch ; (x!60). B, rice starch ; C, wheat starch After Allen. be hydrolyzed to glucose, which is then used to form a part of larger molecules of compounds found in the tissues and fluids of the body. These compounds, in turn, are evidently easily avail- able in the animal economy for the production of heat and energy by their oxidation to carbon dioxide and water. Dextrin. If starch is moistened with very dilute nitric acid and heated for some time at 120, it is converted into a soluble compound or mixture of compounds called dextrin. Other forms of dextrin may be obtained by the use of hydrochloric acid, by heat alone at a somewhat higher temperature, or by the action of the diastase of malt. The dextrins are more or less soluble in water and are used for the preparation of some kinds CARBOHYDRATES 337 of mucilage and for the backs of postage stamps and of gummed labels. Pectose, Pectin. Fruits of nearly all kinds, especially when not fully ripe, contain a substance called pectose, which is in- soluble in water but which when boiled with water in the presence of the fruit acids is decomposed or hydrolyzed with the formation of a soluble compound called pectin. Pectin forms a jelly with sugar in a slightly acid solution. The best conditions require enough sugar to give a solution boiling at about 103 and having a specific gravity while hot of 1.27 to 1.29. The fruit juice should contain not less than 0.5 to 0.7 per cent of an organic acid, calculated as tartaric acid. Usually an amount of sugar about one half to three fourths of the volume of the fruit juice is used. The boiling must not be long continued after separating the juice from the fruit, as this seems to- destroy the pectin (see N. E. Goldthwaite, J. Ind. and Eng. Chem. 1, 333 and 2, 457 ; also Principles of Jelly Making, Univ. of 111. Bulletin, Vol. 9, No. 36 (1912)). Pectin may be precipitated from fruit juices, which have been extracted by cooking, by means of alcohol, but its com- position and its properties as a definite compound have not been established. Cellulose. The fiber of wood, flax, cotton, the outer coatings of cereals and many similar materials consist largely of an in- soluble substance having approximately the same composition as starch, (CeHioOs)^. As coal and peat were formed fronj woody material, cellulose must be considered as the principal original constituent of all of our fuels except petroleum and natural gas. In hay and alfalfa it forms a very important con- stituent of the food of herbivorous animals. It also furnishes the basis for the manufacture of paper. The best grades of filter paper are nearly pure cellulose. Paper. The cheapest grades of paper are made from straw, the better grades from wood and the best from flax, or from cotton or linen rags. Many other fibrous materials may also be used. The materials are first treated with a variety of 338 A TEXTBOOK OF CHEMISTRY chemicals to disintegrate them, bleach them and remove color- ing matters and other substances which are objectionable. This finally produces a thin, uniform pulp which can be spread out evenly to form the sheet of paper. In the best kinds of paper the fibers must remain as long and as strong as possible. It is also necessary to remove substances which turn brown on exposure to the light and which cause the paper to darken with age. The glazed surface of paper, necessary to prevent the absorp- tion and spreading of ink, is obtained by the application of rosin, aluminium sulfate and other substances as sizes. Gun Cotton, Celluloid, Lacquers. When cotton is digested with a mixture of concentrated nitric and sulf uric acids a variety of compounds called nitro-celluloses are formed, differing with the concentration of the acids used, the duration of the treatment and the physical condition of the material. The most highly nitrated product has the composition C^HnO^NOs^, and is called hexanitrocellulose. These products are powerful ex- plosives and are used in torpedoes and also as the principal constituent of smokeless powder. Less highly nitrated forms dissolve in a mixture of alcohol and ether, forming collodion. The evaporation of the solvent leaves the material as a thin, elastic film which is used to hold the silver compound for the wet plates in photography. It is also sometimes used to protect wounds from the access of bacteria. Similar products dissolved in amyl acetate form excellent lacquers. Mixed with or dissolved in camphor they form celluloid. CHAPTER XX AMINES, DYES, ALKALOIDS, PROTEINS, ENZYMES, FOODS AND NUTRITION IF one or more of the hydrogen atoms of ammonia are replaced by organic radicals, a great variety of compounds called amines are formed. These compounds combine directly with acids to form salts and are often called organic bases, but it should be remembered that the true bases are related to the amines in the same way that ammonium hydroxide, NKUOH, is related to ammonia, NH 3 . The " strength " of these bases varies greatly according to the nature of the radicals which they con- tain. Thus methyl amine, CH 3 NH 2 , forms in aqueous solution a much stronger base (CH 3 NH 3 OH) than ammonium hydroxide, while aniline, CeHsNH^, gives a very much weaker base. Both, however, form well-crystallized, definite salts, as methyl am- monium chloride, CH 3 NH 3 C1, or aniline hydrochloride, C 6 H 5 NH 3 C1. Methyl Amine, CH 3 NH 2 . Ammonia combines directly with methyl iodide, CH 3 I, to form methyl ammonium iodide, CH,I + NH, CH 3V , H I Methyl Ammonium Iodide Ciisv B = H-^NO When methyl ammonium iodide is warmed with a concentrated solution of sodium hydroxide, methyl amine escapes as a gas, exactly as ammonia escapes when ammonium chloride is treated in the same way : CH 3 NH 3 I + NaOH = Nal + CH 3 NH 2 + H 2 O 339 340 A TEXTBOOK, OF CHEMISTRY Methyl amine is a gas with a disagreeable odor resembling that of herring brine. It resembles ammonia in its general properties, but is more easily combustible. Aniline. By treating benzene, C 6 H 6 , with nitric acid a com- pound called nitrobenzene, C 6 H 5 NO 2 , can be prepared. When this compound is reduced by tin and hydrochloric acid, or, commercially, by iron and acetic acid, aniline, C 6 H 5 NH 2 , is ob- tained. When pure, aniline is a colorless oil which boils at 184. It is made on a large scale for use in the manufacture of a great variety of dyes and for the preparation of several valuable com- pounds used in medicine, especially of acetanilide (antifebrin), yCO CH CeH 5 NHC 2 H 3 O, and antipyrine CeH 5 N<; II \N(CH 3 )-C CH 3 . ,NHC 2 H 3 O Phenacetine, CeH^ , may also be considered as a XXiH, derivative of aniline. Dyes. Till the middle of the nineteenth century all of the dyes used for coloring fabrics were either inorganic compounds or were natural, vegetable or animal products. Vegetable products were chiefly used, but the number of those available was comparatively small, the two of greatest importance being, probably, indigo and alizarin, or Turkey red. In 1856 Sir William Perkin, in the course of some experiments which he tried in the hope of obtaining quinine from aniline, discovered a beautiful violet compound, mauve, which can be manufactured by the oxidation of aniline. During the next few years he estab- lished the manufacture of mauve on a commercial basis. This proved to be the starting point for a great industry for the prep- aration of thousands of different colors by synthetic processes. Some of the artificial dyes are identical with those obtained from vegetable sources. Many others rival these in brilliancy, in being insoluble or " fast " when the fabrics dyed with them are washed and in resisting the action of light. Others are not so good as some of the natural dyes in these last particulars. Colors of almost every conceivable shade are now available. (DYES 341 Alizarin, Ci4H 6 O 2 (OH)2. Shortly after the discovery of mauve, Graebe and Liebermann, two German chemists, showed that by distilling alizarin, the coloring matter of Turkey red, with zinc dust, it is reduced to anthracene, a hydrocarbon found in coal tar : C 14 H 6 O 2 (OH) 2 + 5 Zn + H 2 O = Ci 4 H 10 + 5 ZnO Alizarin . Anthracene Soon after this, methods were developed for the manufacture of alizarin from anthracene, and in a very few years the artificial product displaced the natural dye and the raising of madder root, from which the dye had been obtained, practically ceased. Indigo, Ci6HioO 2 N 2 . This dye, which has been extensively used for many centuries, has been obtained until recently almost exclusively from a plant growing in India. In 1881 Professor Baeyer in Munich discovered a method of making indigo arti- ficially. The process was complicated, however, and the original material used, toluene, was too expensive to allow of the profit- able manufacture of the dye. Twenty years of continuous study in University laboratories and in factories were required before a successful process was developed. One process used starts with naphthalene, Ci H 8 , and acetic acid as the original materials. In 1901 it was so far developed that the Badische Soda-Anilin Fabrik had been willing to spend $4,500,000 in pre- paring for the manufacture on a large scale. Since then the amount of the synthetic indigo produced has increased each year, and it seems likely that it will ultimately displace the natural product. Indigo is extremely insoluble in almost all solvents which do not change it chemically, and its value depends very largely on this property, which makes it a " fast " color. In order to fix it on the fiber it is dissolved by reducing it with ferrous hydroxide in the presence of calcium hydroxide or, of recent years, with an alkaline solution of sodium hyposulfite, Na 2 S 2 O4. The in- digo white formed by the reduction is a weak acid and forms a soluble salt with the calcium or sodium. Fabrics which are to 342 A TEXTBOOK OF CHEMISTRY be dyed are dipped in the alkaline solution. On exposure to the air the indigo white is oxidized back to indigo, which remains firmly attached to the fiber : Ci 6 Hi 2 N 2 + 2 Fe(OH) 2 + 2 H 2 O Indigo = H 2 C 16 H 10 2 N 2 + 2 Fe(OH) 3 Indigo White H 2 Ci 6 Hi O 2 N 2 + Ca(OH) 2 = CaCi 6 Hi O 2 N 2 + 2 H 2 O Soluble Calcium Salt of Indigo White Mordants. Some dyes, as indigo, are so insoluble that if once formed in contact with the fiber of a fabric, they will not dissolve and they produce a " fast " color. Other dyes seem to combine with fibers directly to form insoluble compounds. Such dyes are called " substantive " dyes because they are in- dependent of the use of other substance required to fix them on the fiber. Substantive dyes which may be used for silk or wool are much more common than those for cotton or linen. Other dyes, which are called " adjective " require a mordant, with which they form an insoluble compound, to fix them. The most common mordants are aluminium acetate, ferric acetate, potassium dichromate and tannic acid. Alkaloids. There is a considerable number of nitrogenous compounds, found in plants, which combine with acids to form crystalline salts in the same way that the amines do. Some of these are comparatively simple amines, but most of them are complex in their structure. Many of them have some very marked physiological action as poisons or as medicines. Many which are poisonous are used as medicines in small doses. Nicotine, CioHi4N 2 , is a colorless oil found in tobacco, which contains from 2 to 8 per cent of the alkaloid. It is very poison- ous. Coniine, C 8 Hi 7 N, the alkaloid of hemlock, is also a liquid. It is historically interesting as the active principle of the fatal draught taken by Socrates. ALKALOIDS. PROTEINS 343 Atropine, Ci 7 H 2 3O 3 N, is found in Atropa belladonna. It is used to dilate the pupil of the eye and is an active poison. Cocaine, Ci7H 2 iO 4 N, is found in coca leaves. It is used to produce local anaesthesia. A careful study of cocaine has shown that it is a derivative of benzoic acid and the group derived from that acid is chiefly effective in giving to it its valuable qualities. On the basis of this discovery other alka- loids having, in part, a similar structure have been prepared. Some of these retain the anaesthetic effect of cocaine and are less poisonous. Morphine, CnHigOaN.H^O, is the most important alkaloid of opium and is the chief constituent which gives to laudanum and paregoric their poisonous and sedative qualities. Paregoric also contains camphor and aromatic oils which may have as much effect as the morphine. Opium is obtained from the poppy. Quinine, C 2 oH24O 2 N2, is obtained from Peruvian bark. It is a specific in malarial fevers. Strychnine, C2iH22O2N2, is found in Strychnos nux vomica. It is a violent poison, producing convulsions. A dose of 0.06 gram is considered fatal. In small doses it is a powerful stimulant. Ptomaines. In the putrefaction of fish or meat under the influence of bacteria and sometimes in the putrefaction of vege- table substances, a variety of basic compounds called ptomaines is formed. Some of these are poisonous, and illness from ptomaine poisoning often results from eating spoiled meats, especially fish. A few of them give color reactions similar to those given by the vegetable alkaloids. Their presence often greatly increases the difficulty of identifying alkaloids in toxical analysis. Proteins. The most important compounds in both animal and vegetable organisms seem to be the proteins. These con- tain carbon, hydrogen, oxygen and nitrogen, usually sulfur and sometimes phosphorus or iron. The albumen of the white of an egg is a nearly pure, typical protein. Proteins form the larger part of muscular fiber, of the casein of milk and of the gluten of 344 A TEXTBOOK OF CHEMISTRY flour. The molecular weight of the proteins contained in the substances just mentioned is very high as much as 15,000, at least. When foods containing proteins are eaten, during the process of digestion they are hydrolyzed, under the influence of the enzymes (see below) of the digestive fluids, giving less com- plex proteins called albumoses, and by further hydrolysis amino acids. These pass into solution and so into the circula- tion of the blood and from them and from other portions of the food the organism reconstructs the proteins which enter into the tissues and fluids of the body, replacing those proteins which are constantly being oxidized to furnish heat and energy for the organism. In part, they are oxidized directly without being transformed into tissues. Enzymes. The larger portion of the foods which are eaten are insoluble in water and in a form which could not be directly assimilated. In the course of the digestive tract is found a series of substances, called enzymes, which act upon the food catalytically, hydrolyzing it or changing it so that it becomes soluble, and emulsifying the fats. The most important of these enzymes are the ptyalin of the saliva, which changes starch to sugar, pepsin of the gastric juice of the stomach, which, with the aid of hydrochloric acid, normally present, changes the proteins to albumoses and renders them soluble, and trypsin from the pancreas. The fluids of the digestive tract are alternately alkaline and acid. Many other enzymes are known. One of the first to be dis- covered was the diastase, which is formed during the germina- tion of barley and which forms the active constituent of malt. It transforms the starch of grains into maltose and dextrin in the manufacture of alcohol. Yeast cells contain an enzyme, zymase, which transforms glucose, fructose or maltose to alcohol. Toxins, Antitoxins. In snake venom and in a few plants, especially in the castor bean, substances are found which seem to resemble the enzymes, but which produce disastrous, poison- ous effects upon the animal organism. These are called toxins. Toxins seem also to be formed by the action of bacteria in cer- FOODS AND NUTRITION 345 tain diseases and doubtless their formation is often a chief factor in the progress of the disease. It has been discovered that animals subjected to the effect of such a toxin develop a sub- stance which appears to combine with it and render it harmless. By use of this principle it has been possible to prepare antitoxins which are powerful agents in combating these diseases. Urea, CON2H4. About 85 per cent of the nitrogen taken as food is eliminated from the human body in the form of urea. This may be considered as formed by the union of carbon dioxide and ammonia, followed by the loss of water : /NH 2 /OH /NH 2 Cf + 2NH 3 = cf =CO +H 2 NH 2 Urea is also formed by rearrangement when a solution of ammonium cyanate is evaporated : /NH 2 NH 4 O C=N - C^=O Ammonium Cyanate NH 2 Urea This transformation of ammonium cyanate into urea, which was discovered by Wb'hler in 1828, was the first synthesis of an " organic " compound from inorganic materials and it was the beginning of a long series of syntheses which have demonstrated that many compounds prepared in the laboratory are identical with the same compounds found in animals and vegetables. Nutrition. An adult man weighing about 70 kilograms, when on an average, mixed diet, eliminates from his system 16 to 18 grams of nitrogen per day. This is equivalent to the consump- tion of 100 to 112 grams of digestible protein, which could be obtained from 3 to 3j liters of milk, 1150 to 1250 grams of white bread, 600 to 750 grams of fresh fish, or 500 to 560 grams of lean beef. From the experiments with the respiration calorim- eter (p. 313) it seems that in a room at 20 the body of an 346 A TEXTBOOK OF CHEMISTRY adult weighing 70 kilograms loses about 2200 calories in 24 hours. The protein referred to above would give by its oxidation in the body, only 420 to 475 calories. 1 To furnish the balance of energy required by the body, about 325 grams of carbohydrates (starch, sugar, etc.) and 50 grams of fat would be required. Theoretically it does not matter whether the extra energy is supplied by carbohydrates or by fat. Practically, both are usually taken somewhat in the proportions given. When engaged in muscular labor for 8 hours a day, about 20 grams of nitrogen are eliminated and about 1800 addition calories are given out as heat and mechanical energy. This would re- quire a total of about 125 grams of protein, 625 grams of carbo- hydrates and 100 grams of fat. In the respiration calorimeter, where the man drove a stationary bicycle arranged to measure the work performed, about 20 per cent of the energy of the extra food required was converted into mechanical energy. The remainder was dissipated as heat given out from his body. This indicates that the human body, considered as a machine, is somewhat more efficient than the best steam engines. But the food required as the source of energy is, of course, much more expensive than coal. A large number of dietary studies have given results which indicate that the average American diet for an adult man weighing 70 kilograms is approximately that given in the table on opposite page. The consideration of the amount of protein and of heat energy required by the body, while undoubtedly of great importance in deciding upon the character of the diet which should be selected, takes account of only a few of the factors which ought to be considered. About many of these factors our knowledge is still very imperfect. It is claimed by some .writers that the human body may be maintained in a state of health with the 1 In the bomb calorimeter 1 gram of protein gives about 6.55 large calories. When taken as a food a part of the protein is elim- inated in the form of uric acid, creatinine and other compounds, which may still be burned with evolution of heat, hence the heat of oxidation of protein in the body is only about 4.2 calories per gram. FOODS AND NUTRITION 347 PROPORTIONS OF NUTRIENTS FURNISHED BY DIFFERENT FOOD MATERIALS IN THE AVERAGE AMERICAN DIETARY PRO- TEIN FAT CARBO- HYDRATES Animal foods : Total meats % 160 % 29 7 % 588 % Fish 1 8 3 ^ 1 Effffs 2 1 4 1 2Q -^feft Dairy products Unclassified animal foods . . 18.4 0.2 38.5 10.0 0.2 25.7 0.2 3.6 0.3 Vegetable foods : Total cereals 306 43 Q 1 61 8 Sugar, molasses, etc .... Legumes, tubers, vegetables . Fruits, including nuts . . . Unclassified vegetable foods . 5.4 20.3 4.4 0.5 61.2 8.7 0.6 1.0 0.7 17.6 12.0 4.3 Miscellaneous food materials . 0.3 0.2 0.6 0.4 100.0 100.0 100.0 100.0 use of a very much smaller amount of proteins, indeed with about 60 per cent of that given above, and experiments have been tried which tend to support this point of view. It has been shown, too, that the proteins from different sources differ very much in the character of the amino acids formed by their hydrolysis. Inasmuch as certain amino acids in definite quan- tities are required to restore the wasted tissues of the body, it seems certain that some of the proteins are much more suitable than others for use as food. But inquiry along these lines is recent and has not proceeded far enough to lead to final con- clusions. Finally, there are many of the inorganic elements which are absolutely essential to the health of the organism, such as sodium, chlorine, iron, calcium, silicon, sulfur, phosphorus and even iodine. CHAPTER XXI SILICON, BORON, GERMANIUM, TIN, LEAD, TITANIUM, ZIRCONIUM, CERIUM, THORIUM THE atomic weights of the nonmetallic elements, in round numbers, arranged in the order of the groups of the Periodic System, are as follows. The most closely related metallic ele- ments are also given. All of the nonmetallic elements are above and to the right of the dotted line. B 11 C 12 N 14 16 F 19 He Ne 4 20 Al 27! Si 28 P 31 s 32 Cl 35.5 A 40 Ga 70 Ge 72j As 75 Se 79 Br 80 Kr 83 In 114 Sn 118 Sb 120 JTe 127.5 I 127 Xe 130 Tl 204 Pb 207 Bi 208 '1 Ni 222 Silicon. Si, 28.3. Occurrence. Silicon is the most widely distributed and abundant element after oxygen. It forms about one fourth of that portion of the earth which we can examine. It is the characteristic element of minerals almost as much as carbon is the characteristic element of living matter though there are many minerals which do not contain sili- con, and the number of silicon compounds is very much smaller than that of the carbon compounds. Silicon is found in nature in the form of the dioxide, SiO 2 , in rock crystal, the purest form of quartz, amethyst, jasper and agate, flint, or chalcedony. Silicon dioxide also forms the prin- cipal constituent of the sandstones and of ordinary sand. A hydrated dioxide, containing varying amounts of water, is called opal. Silicon is found in a great variety of silicates, all of which may be considered as salts of silicic acids, of which silicon dioxide is the common anhydride. Among these may be mentioned orthoclase, one of the feldspars, KAlSisOs, mica, KH2A1 3(8104) 3, 348 SILICON CARBIDE 349 kaolin, H 2 Al 2 (SiO4)2.H 2 O, a chief constituent of clay, asbestos, a variety of amphibole, Ca 3 Mg3(SiO 4 )3, talc or soapstone (alberene), H 2 Mg 3 Si4Oi 2 , serpentine (meerschaum), H 4 Mg 3 Si 2 O 9 , garnet, Ca 3 Fe 2 (SiO 4 )3, or Ca 3 Al 2 (SiO 4 ) 3 , topaz, Ali 2 Si 6 O 25 Fio, tourmaline, Al 4 B 6 Oi 5 .4 H 2 NaAl 3 (SiO 4 ) 3 , and beryl, Be 3 Al 2 (SiO 3 ) 6 . Only a few ores of common metals are silicates, the most impor- tant being calamine, Zn 2 SiO 4 . Preparation. Silicon is never found free in nature. It is most easily prepared by heating a mixture of fine sand with magnesium : SiO 2 + 2 Mg = 2 MgO + Si The silicon obtained in this way is an amorphous brown powder insoluble in water and acids, except in a mixture of hydrofluoric and nitric acids. It dissolves in a solution of sodium hydroxide with evolution of hydrogen : Si + 2 NaOH + H 2 O = Na 2 SiO 3 + 4 H Sodium Silicate Silicon may be crystallized from its solution in melted zinc and then forms brilliant needles having a metallic luster. It is now made in electric furnaces for use in the steel industry. Hydrogen Silicide, SiH|. If fine sand is heated with four atoms of magnesium for each molecule of the silicon dioxide, the silicon combines with the magnesium to form magnesium silicide, Mg 2 Si. When this is treated with a dilute acid, hydrogen sili- cide is formed : Mg 2 Si + 4 HC1 = 2 MgCl 2 + SiH 4 Magnesium Hydrogen Silicide Silicide Hydrogen silicide is a colorless gas which takes fire spontane- ously on coming to the air and burns to water and silicon dioxide. It may be condensed to a liquid, which boils at a very low tem- perature. Silicon Carbide. Carborundum, SiC. By heating silicon dioxide with coke in an electric furnace it may be reduced and 350 A TEXTBOOK OF CHEMISTRY the silicon and carbon unite to form silicon carbide or carborun- dum, which, when pure, crystallizes in beautiful, colorless needles : SiO 2 + 3 C = SiC + 2 CO Carborundum is the hardest substance known except the car- bide of boron and the diamond, and it is manufactured for use as an abrasive. For this purpose it has partly displaced corundum, Al 2 Os, which is used under the name of emery. Carborundum is also used as a refractory material in the con- struction of furnaces and to remove gases from steel. Silicon Fluoride, SiF 4 . The formation of silicon fluoride in the etching of glass by hydrofluoric acid has been referred to. The compound is most easily prepared by warming a mixture of sand, SiO 2 , calcium fluoride, CaF 2 , and concentrated sulfuric acid : 2 CaF 2 + SiO 2 + 2 H 2 SO 4 = 2 CaSO 4 + SiF 4 + 2 H 2 O Silicon fluoride is a gas, but may be condensed to a solid which melts at 97 and has a vapor pressure of 760 mm. at 90. Fluosilicic Acid, H 2 SiFe. Silicon fluoride is hydrolyzed by water, as would be expected of a halogen compound of a non- metallic element, but the hydrofluoric acid formed combines with some of the undecomposed silicon fluoride to form a com- plex acid, fluosilicic acid : SiF 4 + 4 HOH = Si(OH) 4 + 4 HF Silicon Silicic Fluoride Acid 2 HF + SiF 4 = H 2 SiF 6 Fluosilicic Acid The potassium salt of fluosilicic acid, K 2 SiF 6 , is difficultly soluble. The barium salt, BaSiF 6 , is also extremely insoluble, even in dilute acids. It is the only salt of barium likely to be mistaken for barium sulfate, BaSO 4 , when barium chloride, BaCl 2 , is used to test for sulfates in a dilute acid solution. SILICON DIOXIDE 351 Silicon Tetrachloride, SiCU, is formed when chlorine is passed over heated silicon. It was formerly prepared by heating a mix- ture of silicon dioxide and charcoal in a current of chlorine : Si0 2 + 2 Cl a + 2 C = SiCl 4 + 2 CO This method of preparation, which was formerly much used to obtain chlorides of elements, such as silicon, aluminium and chromium, which cannot be reduced from their oxides by carbon at any moderate temperature, has ceased to be of practical impertance since other methods have been developed for the preparation of the free elements. The process depends on the simultaneous use of chlorine and carbon to cause the separation of the silicon and oxygen. Silicon tetrachloride is a liquid which boils at 56.9. It is hydrolyzed by .water to silicic acid, Si(OH) 4 , and hydro- chloric acid. Silicon Hexaiodide, Si 2 l6, is formed when silicon tetraiodide, SiI 4 , is heated to 290-300 with powdered silver. It crystallizes from carbon bisulfide in colorless prisms. It is hydrolyzed by water to silicooxalic acid : Si 2 I 6 + 4 HOH = H 2 Si 2 O 4 + 6 HI Silicooxalic Acid These compounds and some others, which have been prepared, show that silicon atoms may combine together as carbon atoms do, but only a few such compounds are known and most of these are comparatively unstable. Silicooxalic acid, H 2 Si 2 O 4 , is so unstable as to be explosive. It seems evident that stable com- plex molecules containing silicon are formed only when the atoms are held together through the agency of some other element, such as oxygen. (See below under Silicic Acids.) Silicon Dioxide or Silica, SiO 2 . In addition to the forms of occurrence already given, infusorial earth or " Kieselguhr," a porous material composed of the skeletons of infusoria, may be mentioned. It is used to absorb nitroglycerin in the manu- facture of dynamite and as a packing material for bottles con- 352 A TEXTBOOK OF CHEMISTRY taining bromine or other corrosive chemicals. It is also used in sapolio and in other scouring soaps and scouring materials. Clear specimens of rock crystal are sometimes used in the preparation of prisms and lenses which are more transparent than ordinary glass to ultra-violet light. Flint and pure sands and sandstones are used in the manufacture of glass. Silicon dioxide is found in nature in two forms. Rock crystal or quartz, which is much the more common, crystallizes in forms of the hexagonal system and has a specific gravity of 2.6. Quartz can be formed only at temperatures below 900. Above that temperature it changes to tridymite, which crystallizes in the rhombic system and has a specific gravity of 2.28. Quartz melts at about 1750, but softens and becomes plastic at 1600 or below. By means of the electric furnace or the oxyhydrogen flame it can be melted and fashioned into tubes, crucibles, beakers, flasks, thermometers and other laboratory utensils. It has the advantage over glass that it can be heated to very high temperatures without melting and also that it has such a small coefficient of expansion for a change of temperature that it does not crack when subjected to sudden heating or cooling. Partly for the same reason thermometers made from it show no depression of the zero point after use at high tempera- tures, as is common with glass thermometers. Fused quartz has a specific gravity of only 2.20, nearly the same as that of tridymite. Fused tridymite would doubtless be a more correct name than fused quartz. Artificial Silicates. When silicon dioxide is heated with so- dium carbonate or with the oxide or carbonate of almost any metal, a silicate is formed : Na 2 CO 3 + SiO 2 = Na 2 SiO 3 + CO 2 Sodium Sodium Carbonate Silicate CaO + SiO 2 = CaSiO 3 The industrial importance of reactions of this character will be seen when it is stated that similar reactions are used for the SILICIC ACIDS 353 manufacture of glass, for the formation of fusible slags in the manufacture of iron and in other metallurgical operations and in the manufacture of Portland cement. The silicates of sodium and potassium are soluble in water, and the former, especially, is called water glass and is used as an addition to laundry soaps, for preserving eggs and for the fire- proofing of wood and fabrics. Almost all other silicates and mixed silicates are insoluble or nearly insoluble in water. Silicic Acids. If concentrated hydrochloric acid is added quickly to a solution of sodium silicate, the silicate seems to be completely decomposed, but the silicic acid formed does not pre- cipitate. If such an acidified solution is dialyzed by placing it in a parchment sack suspended in water (Fig. 91), the sodium chloride will diffuse through the walls of the sack, and by repeat- edly changing the water on the outside the chloride can be al- most completely removed, leaving Fig. 91 a colloidal solution of silicic acid. Colloidal solutions may also be prepared by the hydrolysis of esters of silicic acid such as the methyl ester, (CH 3 ) 4 SiO4. According to the method of preparation silicic acid may be either a negative or a positive colloid. When it is a negative colloid it retains some anion, such as the chloride ion, Cl~, which cannot be removed by dialysis or washing. When it is a positive colloid it retains some cation as the sodium ion, Na + , in the same way. Solutions of the first class are precipitated by solutions containing a bivalent cation, such as barium chloride, BaCl 2 , while those of the second class are precipitated by solutions containing a bivalent anion, such as potassium sulfate, K2SO4 (see p. 261). Colloidal silicic acid, which is present in arable soils, retains cations in a form which cannot be washed out by the rain, 354 A TEXTBOOK OF CHEMISTRY probably owing to the relations which have just been given, and potassium seems to be held much more strongly than magnesium, calcium or sodium. This is doubtless of great importance in its relation to the fertility of the soil, since potassium is one of the most essential elements for the growth of crops. If hydrochloric acid is added to the solution of sodium silicate drop by drop, instead of suddenly, the silicic acid separates as a gelatinous precipitate. If this is dried on the water bath or at a slightly higher temperature, or if the colloidal solution obtained by the sudden acidification is evaporated and the residue dried, the silicic acid becomes almost completely insoluble in acid solutions. Such a process is much used in the quantitative analysis of silicates. The gelatinous precipitate obtained by precipitation has very nearly the composition Si(OH) 4 (Norton, J. Am. Chem. Soc. 19, 832 (1897)), but there is considerable doubt whether it is a definite compound, since it loses water very easily and its vapor pressure is scarcely, if at all, different from that of pure water. If dried, however, it loses the last portions of water with difficulty and only when heated to bright redness. Opal must be con- sidered as a mixture of silicic acids, (SiO 2 .H 2 O), but is variable in composition, and it cannot be said that the existence of any definite compound having the composition of a silicic acid has been established. Silicon dioxide, SiO 2 , resembles carbon dioxide, CO 2 , in this respect, but with the difference that while carbonic acid, H 2 CO 3 , dissociates directly into carbon dioxide and water the hydrates of silicon dioxide lose water gradually, forming hydrates containing less and less water, that are united together to form complex molecules, perhaps somewhat as follows : =0 -> H Q SiO O SiO OH OH H The final product of the dehydration of the silicic acid would, according to this view, consist, not of simple molecules of silicon dioxide but of complex molecules (SiO 2 ) ra in which the molecules NATURAL SILICATES 355 are held together through the agency of oxygen. The large number of complex silicates which are known furnish a strong support for this view. The extremely high melting point and boiling point of silica, especially as compared with carbon di- oxide, also indicate that its molecule is complex and that it does not have the simple formula SiO 2 . Natural Silicates. Carbon forms a great number of acids in which the complexity is due to the union of carbon atoms with each other and with varying numbers of other atoms and groups, and it forms only a single acid for which carbon dioxide is the -0\ /O- anhydride. Such a grouping as O=C O C=O seems to be extremely unstable. As has just been pointed out, how- ever, molecules containing silicon atoms united by oxygen seem to be stable, and a great variety of minerals are known which may be considered as salts of more or less complex silicic acids, all of which are derived from the same -anhydride, silica, SiO2. The hypothetical acids from which these natural silicates are derived are best classified according to the number of atoms of silicon contained in one molecule of the acid. The first two are given names similar to the names of the acids of phosphorus, arsenic and antimony : Orthosilicic Acid H 4 SiO 4 Metasilicic Acid H 2 SiO 3 Disilicic Acid H 6 Si 2 O7 Trisilicic Acids H 4 Si 3 O 8 and H 8 Si 3 Oi The following minerals may be given as illustrations of the salts of the above acids. ^ , _ Calamine, Zn 2 SiO 4 .H 2 O Orthosmcates ; Derivatives of ^^ H 2 Al 2 (SiO 4 ),H 2 O H4bl 4 Garnet, Ca 3 Fe 2 (SiO 4 )3 Metasilicates ; Derivatives of H 2 Si0 3 Amphibole Hornblende I CaMg 3 (SiO 3 ) 4 Asbestos J Talc or soaps tone, Mg 3 H2(SiO 3 ) Beryl, Be 3 Al 2 (SiO 3 ) 6 356 A TEXTBOOK OF CHEMISTRY Disilicates ; Derivatives of H 6 Si 2 O 7 Serpentine, I^MgaSiaOc, (or Mg 3 Si 2 O 7 .2 H 2 O) Trisilicates ; . , Derivatives of ( Orthoclase, KAlSi 3 O 8 H 4 Si 3 8 iAIbite, NaAlSi 3 O 8 Derivatives of H 8 Si 3 Oi Meerschaum, H 4 Mg 2 Si 3 Oi * Calculation of the Formula of a Mineral. The following analysis of a garnet may be taken as typical of the analysis of a silicate : PER CENT PER CENT OP OXYGEN RATIOS Silica, SiO2 40.45 19.67 4.05 2.60 6.90 5.78 20.79 21.57 9.26 1.21 0.81, 1.53 1.65 8.32 21.57 11.28 11.50 2 or 6 lor 3 1 or 3 Alumina A^Oa Ferric oxide, Fe2Os Chromic oxide, Cr 2 Oa .... Ferrous oxide, FeO Calcium oxide, CaO .... Magnesium Oxide, MgO . . . 100.24 In this and in similar analyses of minerals and rocks it is customary to calculate the results as though the elements were present as oxides. This custom was originally based on Lavoisier's system of nomenclature, according to which salts were considered as compounds of oxides of the metals with oxides of the nonmetals, and it was continued during the first half of the nineteenth century largely because of the older electrochemi- cal theory, which seemed to give a satisfactory theoretical basis for the old nomenclature. This method of calculation is still retained because many of the silicates can be prepared by the direct union of the oxides of the metals with silica, but also for the practical reason that when a silicate analysis is calculated DIALYSIS, SEMIPERMEABLE MEMBRANES 357 in this manner the failure of the percentages found to add up to 100 shows at once that there is an error in the analysis or that some element has been overlooked. In calculating a formula for a mineral it is convenient to make use of the principle that the amounts of oxygen in the different oxides must be in a simple ratio to each other. Thus if we write calcium silicate CaOSiO 2 , the ratio of the oxygen in the calcium oxide (" lime ") must be to that in the silicon dioxide (silica) as 1:2. An examination of the quantities of oxygen in the second column of figures above shows, however, no simple ratio between the amounts of oxygen in the various oxides. It is not till we put together the oxygen of the oxides of ferrous iron, calcium and magnesium and that of the oxides of ferric iron, chromium and aluminium that we obtain numbers which form simple ratios. When we do this we obtain, approximately, the ratios R"O : R 2 /// O 3 : SiO 2 = 3:1:3, and may write the gen- eral formula of garnet, 3 R"O.R 2 ' X 'O 3 .3 SiO 2 or R 3 "R'"(SiO 4 ) 3 , in which R" stands for ferrous iron, calcium or magnesium and R'" stands for ferric iron, chromium or aluminium. If this formula is written in the ordinary way, it becomes 3 CaO.Al 2 O 3 .3 SiO 2 or Ca 3 Al 2 (SiO 4 ) 3 , and we see that garnet is a derivative of orthosilicic acid, H 4 SiO 4 . It is fair to say, however, that the published analyses of many of the complex silicates agree only very roughly with the for- mulas which have been assigned to the minerals. The condi- tions under which such minerals have been formed in nature have evidently favored the formation both of complex, isomorphous mixtures and of solid solutions of variable composition. Dialysis, Semipermeable Membranes. Colloidal silicic acid may be separated from sodium chloride by dialysis, with the use of parchment or parchment paper. Graham, who first carefully studied phenomena of this kind, distinguished between crystalloids and colloids in regard to this effect of animal mem- branes. As he used these names they imply that crystalline compounds, such as sodium chloride, most salts and ordinary acids will pass through the parchment, while silicic acid, albu- 358 A TEXTBOOK OF CHEMISTRY men and other substances which do not crystallize, or which crystallize with difficulty, will not pass through. While this dis- tinction still has considerable force, in a general way, a fuller knowledge of the subject of colloids has led to a definition of them (p. 261) which is based on quite other properties than their relation to separating membranes, and the term crystalloid is now little used. A membrane or septum which allows one substance to pass through it while it prevents the passage of another is called semipermeable. A piece of parchment will allow water or salt in solution to pass, but is nearly impervious to silicic acid or albumen. It will allow cane sugar to pass, but more slowly than salt. If a precipitate of copper ferrocyanide, C^FeCeNe, is formed inside of the wall of a porous porcelain cup, by placing a copper sulfate, CuSO-i, solution within and a solution of potassium ferrocyanide, TQFeCeNe, on the outside, the gelatinous membrane, formed at the point in the wall where the two solutions come together is per- meable to water but may be made wholly impervious to cane sugar. A thin sheet of India rubber is readily permeable to pyridine but nearly impervious to glucose in solution in the pyridine. Metal- lic palladium is permeable to hydrogen but impermeable to ni- trogen and most other gases. The mechanism of the permeability is probably different in different cases. Palladium dissolves hydrogen very much as water dissolves carbon dioxide, but will give it up again to a vacuum or to any space containing no hydrogen, although it may contain another gas. There is some difference of opinion about the permeability of copper ferrocyanide. Some authors think the pores of the membrane are so fine that molecules of water can pass through them while the larger molecules of sugar are stopped. Others think that the membrane dissolves water on one side and gives it out on the other, as the palladium dissolves hydrogen and gives it up. Osmotic Pressure. The passage of a liquid or solution through a membrane in the manner which has been described is called osmosis. If a concentrated solution of copper nitrate is OSMOTIC PRESSURE 359 placed in the parchment sack, Fig. 91, it will be seen very soon that the solution inside of the sack is at a higher level than the water on the outside, indicating a greater pressure on the side of the solution and showing that some water has passed through the membrane into the solution. A pressure developed in this manner is called osmotic pres- sure. The pressures developed with the parchment sack will be small, partly because the salt solution as well as the water passes through the membrane, which does not form a perfect septum, and partly because any considerable difference of pres- sure between the two sides would burst the parchment. By the arrangement shown in Fig. 92 it is possible to measure an osmotic pressure of many atmospheres. A membrane of copper ferrocyanide, C^FeCeNe, is first prepared, as described in the last paragraph, within the walls of the porous porcelain cup, z. The tube, m, is filled with mercury, leaving air in the graduated capillary tube. The rest of the apparatus is then completely filled with a solution of sugar or of some other sub- Fig. 92 stance which is to be examined, and the cup, z, is placed in distilled water. The volume of air in the capillary tube will soon begin to diminish, and the contraction will continue till a pressure on the solution is developed which will just prevent the 360 A TEXTBOOK OF CHEMISTRY further passage of water through the membrane ; the system is then in equilibrium. This hydrostatic pressure on the solution, necessary to prevent the passage of the solvent through a semi- permeable membrane into a solution is the osmotic pressure of the solution (van't Hoff). It is evident that if equilibrium is reached when the volume of the air in the tube is one half the original volume, the pressure of the air must be two atmospheres and the osmotic pressure must be the difference between this pressure and the pressure of the air on the water outside, which would be one atmosphere. If the air in the tube is reduced to one third its original volume, the osmotic pressure must be two atmospheres. * The osmotic pressures developed in this manner are very con- siderable. A tenth-formular solution of cane sugar, C^H^On (containing 34.2 grams in a liter or 3.42 per cent), will give an osmotic pressure at of about 2.24 atmospheres, while a formu- lar solution of alcohol, C2H 5 OH (containing 4.6 per cent by weight), would give a pressure of more than 20 atmospheres. The pressures are very nearly proportional to the absolute tempera- ture and, in dilute solutions, nearly the same as though the dis- solved substance (solute) were in the form of a gas in the same volume and at the same temperature. It will be seen, at once, that a measurement of the osmotic pressure may be used to determine the molecular weight of a compound exactly as the density of a gas or vapor is used for this purpose (p. 94). It has been shown, also, by processes of reason- ing which it would carry us too far to give here, that there is a necessary connection between the osmotic pressure of a solution and its vapor pressure, boiling point and freezing point (van't Hoff). The use of the freezing point of solutions to determine molecular weights (p. 112) is intimately connected with these relations. A mental picture of what may be the cause of osmotic pressure can be formed by considering the effect of a septum of palladium on hydrogen gas. If a bulb of palladium containing nitrogen under atmospheric pressure is placed in an atmosphere of hydro- GERMANIUM 361 gen, also under atmospheric pressure, hydrogen will be absorbed by the palladium on the outside and given off on the inside until the pressure of the hydrogen on the two sides of the wall of palladium is the same. The pressure on the inside of the bulb will be two atmospheres, one atmosphere due to nitrogen and one atmosphere due to hydrogen. This is because nitrogen gas is discontinuous and offers no resistance to the escape of mole- cules of hydrogen from the surface of the palladium. When inside, the nitrogen and hydrogen each exert their normal pres- sure on the wall as gases. In a similar manner, if we have a dilute solution of sugar on one side of a membrane of copper ferrocyanide and pure water on the other side, the water will pass through till the pressure due to the water is the same on both sides. When equilibrium is reached the pressure on the side of the sugar solution will be greater than that on the side of the pure water by the amount of the pressure due to the sugar. It has been shown experi- mentally that this pressure is very nearly the same as though the solute existed as a gas in the same space. It is clear from what has been said that the kinetic theory (p. 58) must apply in a modified form to liquids as well as gases. We shall find later (p. 397) that it applies to solids also, as shown by the specific heats of the elements. * Germanium, Ge, 72.5. In 1886 Clemens Winkler found, after repeated analyses of a mineral called argyrodite, that the sum of the elements found was always 6 to 7 per cent less than 100. This led him to a further careful study of the material and to the discovery of germanium, an element intermediate in its properties between silicon and tin, but on the whole resem- bling the latter much more closely than the former. It is a brittle, grayish white metal, which melts at 958. It forms the compounds GeO, GeS and GeCl2, in which it is bivalent and GeH 4 , GeHCl 3 , GeCl 4 , GeO 2 , GeS 2 and K 2 GeF 6 , all of which correspond to similar compounds of quadrivalent silicon. Argyrodite has approximately the composition 4 Ag 2 S.GeS2. Tin, Sn, 119, and Lead, Pb, 207.1, are very distinctly metallic 362 A TEXTBOOK OF CHEMISTRY in their properties in the free state and also in the formation of salts in which they are the metallic elements. They give, how- ever, the oxides SnO 2 , stannic oxide, which is the anhydride of stannic (H 2 SnO 3 ), and metastannic acids, and PbO 2 , lead dioxide, which is the anhydride of plumbic acid, H 4 PbO4. These metals and their compounds will be considered more in detail later. The elements of the fourth group of the periodic system, which are found in the alternate rows, show a closer resemblance to silicon and germanium than is shown by tin and lead. It seems, on the whole, therefore, better to give an account of them here rather than among the metallic elements. * Titanium, Ti, 48. 1 . This element is found in the mineral rutile TiO 2 , in many iron ores, especially in the magnetic iron ore, and in small amounts in almost all rocks. Ordinary soils contain, on the average, more than half a per cent of titanium oxide and it is probably a constituent of the ash of most plants and animals. The free element has apparently never been obtained pure, largely because of its strong affinity for nitrogen and carbon. The purest specimens which have been obtained have a very high melting point, 1790, suggesting the possibility of using the ele- ment for the filaments of electric lamps, but no one has succeeded in applying it to this purpose. Titanium forms compounds in which the element is bivalent (TiCl 2 , TiO, TiS), trivalent (TiCl 3 , Ti 2 O 3 , Ti 2 S 3 , Ti 2 (SO 4 ) 3 , TiN) and quadrivalent (TiCl 4 , Ti0 2 , TiS 2 , K 2 TiF 6 ). Titanium tetrafluoride, TiF 4 , is a white powder which boils at 284, but it is much more easily decomposed by water than silicon tetrafluoride so that on treatment of a mixture of silicon dioxide, SiO 2 , and titanium dioxide, TiO 2 , with a mixture of hydrofluoric and sulfuric acids, the silicon may be expelled as the tetrafluoride on evaporation, while the titanium remains be- hind as the oxide. Acid solutions containing titanium give a deep yellow color with hydrogen peroxide, due to the formation of pertitanic acid. This is used for the detection and estimation of titanium and may also be used for the detection of hydrogen peroxide. Hydrofluoric acid interferes with the reaction. CERIUM 363 Some compounds of titanium are used as mordants in dyeing wool, cotton and leather. * Zirconium, Zr, 90.6, is found most frequently in the mineral zircon, ZrSiO4. It is also found in considerable quantities in Brazil in the form of the dioxide, ZrO 2 . The element exists in two forms. The amorphous form is a black powder, resem- bling carbon. The crystalline form is brittle, melts at about 1700 and has a specific gravity of 6.4. Zirconium hydride, ZrH2, is a black powder. The chloride, ZrCU, is hydrolyzed by water. Zirconates, which resemble some of the silicates, are prepared by fusing zircon dioxide, ZrO 2 , with metallic oxides, carbonates or chlorides. Crystalline silicozirconates have also been prepared. Zirconium dioxide glows very intensely in the oxyhydrogen flame, giving an even better light than lime, and the light is used for intensive illumination in spectroscopy and microphotography. The oxide is used with yttrium oxide as the chief constituent of the Nernst lamp. The mass conducts electricity well only at high temperatures and must be heated in some way to start the passage of the current. It was at one time used as the incan- descent material for the Welsbach light, but that use has been abandoned. There seems to be some possibility of using zircon tor some of its compounds in the filaments of electric lamps. Fused zirconium dioxide resembles fused quartz and cru- cibles for laboratory use have been made from it. Cerium, Ce, 140.25. In Brazil and in North and South Carolina there are found large deposits of a heavy sand, called monazite sand, consisting of a complex mixture of heavy, insoluble minerals which have been sorted out from other materials, partly by the disintegration of rocks in which they were originally disseminated, partly by a process of washing away the lighter minerals in operations which doubtless occurred during very long periods of geological time. These sands con- tain a great variety of minerals and are of considerable commer- cial value because of the thorium and cerium in them. The ce- rium seems to be mostly present as cerium phosphate, CePO 4 . Metallic cerium is an iron-gray, malleable metal, which burns 364 A TEXTBOOK OF CHEMISTRY easily to cerium dioxide, CeO2, in the air. Its specific gravity is 6.73. An alloy with iron throws off incandescent, burning particles on striking with a piece of steel. These will ignite gas or alcohol and the alloy is used for gas lighters and for similar purposes. * Cerium forms three oxides, Ce 2 O 3 , CeO 2 and CeOs. The dioxide, CeO 2 , forms 1 per cent of the material used for the Welsbach mantles. (See below.) Cerous sulfate, Ce2(SC>4)3, forms a difficultly soluble double salt with sodium sulfate, Ce2(SO4)s.Na2SO4.2 H 2 O, which is used in separating cerium from other minerals. The eerie sulfate, Ce(SO4) 2 .4 H 2 O, in which the cerium is quadrivalent is readily hydrolyzed to basic sulf ates . Thorium, Th, 232.4, is found in monazite sand, probably as the dioxide, ThO 2 , and this forms the practical source for thorium compounds, which have become very important for use in Wels- bach mantles. Thorium is also found in the mineral thorianite, an isomorphous mixture of thorium dioxide, ThO 2 , and uranium oxide, UO 2 . Thorianite is found in Ceylon. The mineral has proved of unusual interest as the source of radiothorium, one of the strongly radioactive elements. Apparently pure metallic thorium has not been prepared. From the properties of the impure metal it probably has a spe- cific gravity of about 12.2 and melts at 1700 or above. Thorium forms many salts in which it is quadrivalent, of which the sulfate, Th(SO 4 ) 2 .9 H 2 O, and the nitrate, Th(NO 3 ) 4 .12 H 2 O, may be considered as typical. Welsbach Mantles. Compounds of thorium and cerium are now important articles of commerce for the manufacture of the mantles used for the Welsbach light. In the preparation of these mantles a web of cotton or, more recently and better, of artificial silk is dipped in a solution containing nitrates of cerium and thorium. Careful experiments have shown that the best results are obtained with a mantle containing 99 per cent of thorium dioxide, ThO 2 , and 1 per cent of cerium dioxide, CeO 2 . After drying, the cotton or artificial silk is burned out and the BORON 365 skeleton of oxides is dipped in collodion and dried. The elastic film of nitrocellulose left by the evaporation of the collodion holds the fragile skeleton of oxides together for transportation. A study of the spectra of the light from mantles made of pure thorium oxide, pure cerium oxide and of mixtures of the two has shown that thorium oxide radiates comparatively little light of the wave lengths of the visible spectrum, apparently because it is nearly transparent. Cerium oxide, on the other hand, radiates so much energy of the wave length of the ultra red rays that it lowers the temperature of the flame too far to produce a satisfactory light. In the mixture, the thorium dioxide, which has a very low specific heat and slight power of emission, is heated to a high temperature (1500-1600), and at this tempera- ture the cerium oxide has a greatly increased power of radiation for rays of the wave lengths of the visible spectrum. Small amounts of chromium, platinum, manganese or uranium may produce effects similar to those produced by the cerium, but none of these give as successful a mixture as that of the oxides of thorium and cerium. Boron, B, 11. In some places in Tuscany, in northern Italy, steam which is charged with boric acid, HaBOs, escapes from subterranean sources. The boric acid is obtained by bringing the steam into contact with water, which absorbs the acid and by evaporating the solution till it crystallizes on cooling. Borax, Na 2 B 4 O 7 .10 H 2 O, is found in many natural waters and especially in Borax Lake, in California. A calcium salt, called colemanite, Ca 2 B 6 On.5 H 2 O, is also found in veins in California and is used for the manufacture of borax. Preparation, Properties. Amorphous boron may be prepared by heating the trioxide, B 2 O 3 , or dehydrated borax, Na 2 B 4 O 7 , with magnesium. It forms a brownish gray powder. Boron may be obtained in a crystalline form by crystallizing it from metallic aluminium. The crystals are extremely hard, approach- ing the hardness of the diamond. Boron melts at 2200-2500. Boron Trioxide, B 2 O 3 , Borax Beads. The trioxide or boric anhydride, B 2 Oa, is formed by heating either of the oxygen acids 366 A TEXTBOOK OF CHEMISTRY of boron to a high temperature. It is volatile only at a white heat and on this account it will decompose the salts of almost all other acids, when heated with them to a high temperature. It does this in spite of the fact that in solution boric acid is one of the weakest acids known so weak, in fact, that its solution does not taste sour and does not redden litmus. When fused with salts or with the oxides of metals, boric anhydride combines with the oxides to form borates, just as silicon dioxide combines with oxides to form silicates. Borax glass, Na 2 B 4 O 7 , which may be considered as boric anhydride that has been only partly neutralized by sodium oxide (Na 2 O.2 B 2 O 3 ), decomposes salts or combines with oxides of metals in the same manner, and almost all of the borates of metals prepared in this way remain dissolved in the anhydride or combined with it, giving a mixture of borates, which solidifies to a clear glass on cooling. Many of these glasses have characteristic colors and so the borax bead is much used for the detection of metals in mineralogy and qualitative analysis. Some of the colorless borates may be fused with silicates to form clear, transparent glasses, which are much less soluble than ordinary glass. Such borosilicates are now used in the manu- facture of glassware for chemical laboratories. Boric Acid, HsBOs. If sulfuric acid or hydrochloric acid is added to a moderately concentrated solution of borax, Na2B 4 O7.10 H 2 O, in an amount equivalent to the sodium present, orthoboric acid, H 3 BO 3 , crystallizes from the solution on cooling. It is usually called simply boric acid and is a very weak acid, having no sour taste and none of the corrosive qualities of ordinary acids. It dissolves in about 25 parts of water at 15. The solution has powerful germicidal properties and is often used for an eyewash. Both boric acid and borax have been used as food preservatives, but such a use is con- demned because the substances produce poisonous effects when taken in considerable quantities. Other Acids of Boron. At 100 boric acid loses one molecule of water and is changed to metaboric acid, HBO 2 , and at 140 OTHER COMPOUNDS OF BORON 367 it loses still more water and is converted into pyroboric acid, H 2 B 4 7 . Borax, Na2B4O7.10 H 2 O, is usually considered as a salt of pyroboric acid, H 2 B 4 O7, but it may with equal propriety be called an acid salt of boric acid and the formula written Na 2 Hio(BO 3 )4.5 H 2 O. When heated it loses water and finally melts to a clear glass consisting of sodium pyroborate, Na 2 B 4 O7. The conduct of this glass toward metallic oxides and its use in blowpipe analysis have been mentioned above. Because of these properties it is used as a flux in assaying and in metallurgy. It is sometimes sprinkled over iron which is to be welded. In this use it dissolves the oxide of iron on the surface, forming an easily fusible glass which is pushed out from between the two surfaces when the joint is hammered, leaving clean surfaces of iron to unite in the weld. Borax is used for laundry purposes, having the properties of a mild alkali. Its use as an antiseptic in milk and other foods is forbidden in most countries. Sodium Perborate, NaBO 3 .4 H 2 O, has recently come into use in laundries because it combines the properties of a mild alkali with the bleaching properties of hydrogen peroxide. Its struc- ture, from the methods of preparation, must be closely related to that of hydrogen peroxide, H O O H, and is probably Na O O B=0. Other Compounds of Boron. Boron forms a nitride, BN, a chloride, BC1 3 , and a sulfide, B 2 S 3 , all of which are hydrolyzed by water. The fluoride, BF 3 , gives boric acid and fluoboric acid by hydrolysis, resembling silicon fluoride, SiF 4 , in this respect : BF 3 + 3 HOH = H 3 BO 3 + 3 HF HF + BF 3 = HBF 4 Fluoboric Acid If boric acid or a borate is warmed with alcohol and con- centrated sulfuric acid, an ester of boric acid, (C2H5) 3 BO 3 , is 368 A TEXTBOOK OF CHEMISTRY formed. This is volatile and gives a green color to the alcohol flame, as it burns. If a piece of turmeric paper is dipped in a solution of boric acid or of a borate which has been made faintly acid with hydro- chloric acid, the paper assumes a very characteristic red color on drying. EXERCISES 1. How many grams of magnesium should be mixed with 3 grams of sand to form silicon ? How many grams to form magnesium silicide ? 2. In the first case above how many cubic centimeters of hydro- chloric acid, sp. gr. 1.10, containing 20 per cent of the acid will be required to dissolve the magnesium oxide formed ? 3. How many grams of fluorspar, of concentrated sulfuric acid and of sand will be required to prepare one liter of a 10 per cent solution of fluosilicic acid, assuming that one third of the materials used fails to react? The specific gravity of such a solution is 1.083. 4. What is the formula of a mineral having the following com- position ? PER CENT Silica, SiO 2 > . . 66.21 Alumina, A1 2 O 3 19.16 Potassium oxide, K 2 O 7.38 Sodium oxide, Na 2 O 7.25 100.00 5. How much sodium carbonate will be required to convert 60 grams of quartz into sodium metasilicate ? How many grams of calcium carbonate to convert it into calcium metasilicate ? How many grams of each to form the orthosilicates ? How many grams to form a salt of trisilicic acid, HiSisOs ? 6. What are the equations for the reactions for the preparation of the ethyl ester of boric acid, (C 2 H 5 )3BO3, from borax ? What compounds will be formed on burning the ester ? 7. When boron nitride is heated to 200 in a current of steam, pyro- boric acid and ammonia are formed. What is the equation for the reaction ? 8. How much crystallized boric acid can be obtained from 100 grams of borax ? CHAPTER XXII METALLIC ELEMENTS. DIFFERENCES BETWEEN METALS AND NON-METALS. PREPARATION OF COMPOUNDS Metals and Non-metals. The general, physical differences between the metals and non-metals are familiar to every one. Iron, copper and gold may be taken as typical of the first class, sulfur, phosphorus and carbon of the second. The metals are opaque, except in excessively thin layers, are malleable and ductile and good conductors of heat and of electricity. The non-metals are transparent or translucent, brittle, poor conductors of heat and nonconductors or poor conductors of electricity. It is true that these differences, which are so marked in the elements spoken of above as typical of the two classes, are not equally marked in all cases and that there is a gradation in these proper- ties such that some elements stand in a borderland between the two, but the physical properties mentioned are clearly those of metals on the one hand and of non-metals on the other. In chemical properties the metals and non-metals show equally strong contrasts. In these properties the most typical metals may be considered as sodium and potassium and the most typical non-metals as fluorine and chlorine. The chlorides and even the hydroxides of the former ionize in solution in such a manner that the metal becomes the positive ion. The chlorides of the nonmetallic elements, on the other hand, are hydrolyzed by water with the formation of hydrochloric acid and an acid containing the non-metal. The compounds of the non-metals with oxygen and hydrogen ionize in solution with the formation of hydrogen ions and some complex group containing the non- metallic element. As is the case with the physical properties, however, the chemical properties of the elements show continu- ous gradations from those of the metals to those of the non- 369 370 A TEXTBOOK OF CHEMISTRY metals. Some of the chlorides of the metals are hydrolyzed by water and some halides of elements which are nonmetallic in most of their properties may be formed in the presence of water. Further than this, some elements as they combine with more and more oxygen may pass from metallic to distinctly nonmetallic properties, while nitrogen or even sulfur, when combined with hydrogen or hydrocarbon radicals, may form radicals, such as NH4 or (0113)38, whose hydroxides are bases. The most typical of the metallic elements are univalent toward chlorine and toward oxygen and form no compounds of any kind in which it is certain that they have a higher valence. The most typical nonmetallic elements on the other hand show a varying valence, especially toward oxygen. The development of the electron theory makes it seem possible, or perhaps we may even say, probable, that both the physical and chemical properties which constitute the differences between metals and non-metals are very largely occasioned by the dif- ferences in the conduct of the atoms of the elements toward electrons. The passage of electricity through a metallic con- ductor consists in the flow of a stream of electrons, and metals are supposed to be good conductors because the electrons pass easily from one atom to another throughout the mass. For almost the same reason the atom of the metal becomes a posi- tive ion in solution because it easily loses one or more electrons the electrons being units of negative electricity. Metals are also good conductors of heat because the rapidly moving electrons transfer energy from one atom to another throughout the mass, and they are opaque because the electrons and atoms absorb and diffuse the light vibrations falling on the surface. Classification of the Metals. As in the case of the non-metals, the periodic system furnishes the most satisfactory classification for the metals. The relationships brought out in this way are not always so close as might be desired, and there are even sharper contrasts between the alternate metals of the first group than those which have been noticed among the elements of the sixth and seventh groups. METALLIC ELEMENTS 371 . S 3 H *3 1 s .3 oj a h- 1 .g 1 3 Oi g- P "2 S3 .3 S O H HH 3 II 1 1 O Carbon Family 's s C'5 3 .3 - 111 1 1 1 i i .1 I a * *"J3 ^_ Rutheniun Palladium Osmium, I Platinum d 'o "S S | 1 I ctf O , g 11 1 I ill 1 cu ft <U fl MENTS GROU ^(2 J| 1 jl jil I l| M 5 il i a aS 1^7 1 j^ > ^^ 14 . fill 5 <y s.s 2 S J ^ d J5 s p^ So 1 i 3.3 S 2 > 1 M II S c IH 1 ^ ^ II' S 1 II "i^J b ^ 2 -5 II ^ a u | -S s .1 ! 1 O 02 III 1 1 P IS i- Chromium Family Chromium Molybdenum Tungsten Uranium ^D if -I c i 3 3 If 1 >* |l; >5 S >i^H '3 0) 1 1 M ofe U C/3 1 ' co O .2 1| s s ill S 1 <t> i, 1 11 S IS o S 1 ^c^Pn PH O & > !>&H > fe H L O P * r ^* SERO GROU: lo d 3 ' d 3 a 9 Q ! M ii 1 & b S 1 1 1? o &1 S^ 2 g J'a .-s ^ 1 >> c -d 1 I G .2 2 g iPji ^1 < pq 372 A TEXTBOOK OF CHEMISTRY All of the elements sufficiently well known so that the atomic weights are given in the International Table of Atomic Weights are included in the table, with the exception of hydrogen. The nonmetallic elements are inclosed in brackets. An inspection of the table shows that even if the noble gases are counted as nonmetallic elements more than three fourths of the elements known are metallic in character. In spite of this, the chemistry of the metallic elements offers less variety than that of the non-metals largely because the atoms of the metals show less tendency to combine with each other, or indeed with other elements to form complex radicals. Metals usually sep- arate by themselves as positive ions in solution, while the non- metals more often form complex ions, such as NOs, SO4 or ClOs. Melting Points of the Elements. The following table gives all melting points of the elements which are known with some degree of accuracy. The melting points of the elements whose names are printed in capitals have been very carefully deter- mined, and are used as standard temperatures for calibrating thermometers and pyrometers. As nearly as may be, all values, in particular the standard points, have been reduced to a common scale, the thermodynamic scale. Preparation of Chemical Compounds. It is clear from many illustrations which have been given in the preceeding pages that many chemical reactions result in an equilibrium between reacting compounds such that action seems to cease only because two opposing reactions proceed with equal velocity in opposite directions. From a theoretical standpoint we are probably justified in considering all reactions as reversible, but in some of them the velocities of the reactions in opposite directions differ so greatly that the equilibrium lies very far on one side and the reaction is practically complete. Thus in the very simple case of the reaction between hydrogen and oxygen : 2 H 2 + O 2 ^ 2 H 2 O at 1000 when equilibrium is reached, only 3 parts in 10,000,000 of the oxygen and hydrogen will remain in the form of the free MELTING POINTS OF THE ELEMENTS 373 OO O5TH !> OS 05 O O .8 s 3 oc^ 1 -H 1 -ioooTHO ^l (N <M (M <M (M T-H ,-H ^H rH 111111111111 PQ d O ;-' sgs >* a^g-gpi] -- x = H Mx6a s jn:i 1 A 374 A TEXTBOOK OF CHEMISTRY gases (p. 61). At ordinary temperatures the velocity of com- bination and dissociation becomes so slow, in this case, that either a mixture of oxygen and hydrogen or water may remain apparently unchanged for an indefinite length of time, but we know that such a condition of apparent equilibrium, dependent on the slow velocity of a reaction, is a state of instability. Such conditions are, however, very common, and it has been pointed out that the varying velocity of different possible reactions is of very great practical importance in organic compounds. A very large number of the compounds of the metals are electrolytes, and many of these can be prepared in aqueous solu- tions. The velocity of ionization reactions is so great for all ordinary electrolytes and the interaction between ions in solu- tion is so rapid that equilibrium is reached almost instantane- ously. The amount of each substance present in this equilib- rium is often influenced by some property of one of the substances which removes it from interaction with the others. Such prop- erties are, especially, volatility, solubility, degree of ionization and the formation of complex ions. Effect of Volatility. If concentrated sulfuric acid is dropped into a concentrated solution of salt, there is, at first, no apparent action. In the solution the following reactions will be almost instantly in a state of equilibrium : NaCl Na + + CT H + + cr ^ HCI Na + + HS0 4 - ^ NaHS0 4 It is to be noticed that the reaction as usually given NaCl + H 2 S0 4 = HCI + NaHSO 4 is merely the final result of intermediate ionization reactions and probably occurs, at first, in very insignificant amounts. As more sulfuric acid is added, a point will be reached when the solution becomes saturated with hydrochloric acid. The addition of a further amount of sulfuric acid must now result VOLATILE COMPOUNDS 375 in the escape of hydrochloric acid in the gaseous form, as mole- cules of the compound, HC1. This will disturb the equilibrium of the reaction : H + + cr ^ HCI in such a manner that more of the hydrogen and chloride ions will unite to form un-ionized hydrochloric acid, HCI. This, in turn, will cause the formation of an increased number of sodium and chloride ions from the salt, NaCl, and of hydrogen and hy- drosulfate ions, HSC>4~, from the sulfuric acid. Of course, the more concentrated the original solution of salt is, the more hy- drochloric acid can be obtained in gaseous form, and if we start with a saturated solution and a large excess of solid salt, the latter will pass into solution as the reaction proceeds and nearly all of the chlorine may escape as hydrochloric acid. This result occurs because the hydrochloric acid is a gas and escapes from the mixture as a gaseous phase and in spite of the fact that the ionization reactions of hydrochloric and of sulfuric acids result in the formation of more of the sulfuric than of the hydrochloric acid when the two acids are present in equivalent amounts. When a solution of hydrochloric acid is dropped into a solu- tion of sodium carbonate the following reactions occur : HCI ^ H + + Cl- Na 2 C0 3 ^ Na + + Na + + CO 3 = CO 3 = + H^ ^ HCOr H + + HCOr ^ H 2 C0 3 H 2 C0 3 ; H 2 + CO 2 Na + + Cl~ ^ NaCl Na + + HCOr ^ NaHC0 3 In the third of these reactions the equilibrium is very far in- deed toward the formation of the hydrocarbonate ion, HCO 3 ~, and very few hydrogen ions, H + , can remain in the solution so long as any carbonate ions are present. Since many more hydro- gen ions separate from carbonic acid, H 2 CO 3 , when that is present, than separate from the hydrocarbonate ion, HCO 3 ~, carbonate 376 A TEXTBOOK OF CHEMISTRY ions, COs", and carbonic acid, H 2 CO3, cannot exist to any large extent in the same solution. It will be seen from the above that carbonic acid cannot be formed in the solution in suffi- cient amount for the rapid escape of carbon dioxide l until all of the carbonate ions present have been converted to hydrocarbon- ate ions. In other words, enough hydrochloric acid must be added to complete the reaction : Na 2 CO 3 + HC1 = NaHCO 3 + NaCl before carbonic acid can be formed in sufficient amount for the rapid escape of carbon dioxide. The addition of a small addi- tional amount of acid, however, causes the fourth and fifth reactions to occur and the escape of carbon dioxide begins. From this point on the gas will escape in almost exact proportion to the acid added, since the equilibrium is far to the right in the first, fourth and fifth reactions. It is clear from the above discussion that the formation of a volatile product, which escapes as a gaseous phase, has the same effect on the course of a reaction when it is a secondary product, formed by dissociation, as when it is formed directly by the ionization reactions, as was the case with hydrochloric acid. Effect of Insolubility. When a compound separates as a solid phase from a solution, the effect upon the equilibrium is exactly analogous to the effect of volatilit^. Thus in the reactions which occur on mixing a solution of salt with a solution of silver nitrate : NaCl ^ Na + + Cl~ AgNO 3 ^ Ag + + NO 3 Ag + + Cl- ^ AgCl 1 A solution of sodium bicarbonate, NaHCO 3 , contains some car- bonic acid, H 2 CO 3 , formed by the reaction : Na++ HCO 3 - + H+ + OH- = H 2 CO 3 + Na++ OH~ . Such a solution will lose carbon dioxide on boiling or on exposure to the air. SOLUBILITY PRODUCT 377 the fact that silver chloride is only very slightly soluble causes it to separate from the solution as a precipitate and shifts the equilibrium of the first three reactions toward the formation of this compound. The removal of a substance from the mixture as a solid phase has exactly the same effect upon the various equilibria involved as the removal of a volatile product. Effect of a Common Ion. Solubility Product. When a sub- stance having one of the ions of a difficultly soluble uni-uni- valent salt, that is, a salt such as silver chloride, AgCl, in which both the metallic and acid radicals are univalent, is added to a saturated solution of the salt, some of the difficultly soluble salt will usually be precipitated. Thus the addition of a few drops of a solution of silver nitrate to a saturated solution of silver chloride will cause the separation of some silver chloride from the solution and the addition of hydrochloric acid will also cause precipitation in a saturated solution of silver chloride. The precipitation depends on the following reactions : HC1 ^ H+ + Cl- In the first reaction the quantities of the three substances, silver chloride, AgCl, silver ion, Ag + , and chloride ion, Cl~, must be present in definite amounts in a saturated solution of silver chloride at a given temperature. At 18 the amount of silver chloride in the three forms dissolved by a liter of water is 1.6 milligrams of which about 0.4 of a milligram is chlorine. If, now, hydrochloric acid is added, the number of chloride ions must be largely increased, and this will cause the formation of more un-ipnized silver chloride. As the solution is already saturated with that compound, the silver chloride formed will be precipitated. Under such conditions as these, when a solid, difficultly soluble uni-univalent salt is in equilibrium with a solution con- taining a slight excess of one of its ions, the equation : C Ag + X*C C1 - = a constant 378 A TEXTBOOK OF CHEMISTRY has been found to be true. C Ag + arid C a - are the concentrations of the ions of the salt in the solution. 1 If, for instance, we add enough hydrochloric acid so that the solution is y^Vs" normal for hydrochloric acid, one liter will con- tain 35.5 + 0.4 = 35.9 milligrams of chloride ions, Cl~ , or nearly ten times as much as before, since at these dilutions both the silver chloride and hydrochloric acid will be almost completely ionized. Under these conditions, in order that the product C Ag + X CGI- may remain constant the quantity of the silver ions must become -^ as large as before. Since the silver chlo- ride is almost completely ionized at these concentrations, it will be seen that the addition of this small quantity of hydrochloric acid will reduce the quantity of silver in the solution to about one tenth of the original amount. The importance of these re- lations for the precipitation of difficultly soluble salts in quantita- tive analysis will be easily seen. This principle is often called the constancy of the solubility product, and may be stated as follows : In any dilute aqueous solution saturated with a slightly soluble uni-univalent salt the product of the concentrations of the ions of this salt is constant at a given temperature. In some cases the addition of a salt having a common ion will cause precipitation in a saturated solution containing a uni- bivalent salt, such as lead chloride, PbCl 2 , or in one containing a bi-bivalent salt, such as calcium oxalate, CaC 2 O 4 , but there are many other cases in which such precipitation does not occur. No general rules governing the conduct of such salts can be given in the present state of our knowledge. In some cases this failure of precipitation may be due to the formation of intermediate ions, such as PbCl~. In others it is probably caused by the for- mation of more or less stable complex ions. Formation of Complex Ions. If solutions of silver nitrate, AgNOs, and potassium cyanide, KCN, are mixed in equivalent 1 This relation can be derived theoretically from the ionization reaction : _,, ^ . AgCl ^ Ag+ + Cl- if it is assumed that in very dilute solutions the osmotic pressures of the ions are proportional to their concentrations. COMPLEX IONS. IONIZATION 379 proportions, nearly all of the silver ions and nearly all of the cyanide ions will be removed from the solution as a precipitate, in accordance with the equation : (AgN0 3 + KCN ^ AgCN + KNO 3 Silver Cyanide If, however, more of the potassium cyanide solution is added, the silver cyanide will dissolve. This seems to be in direct con- tradiction to the principle of the solubility product given above. A more careful examination shows that the apparent contradic- tion is due to the fact that the solution no longer contains an appreciable number of silver ions, Ag + . This can be shown in three ways : 1. Sodium chloride, NaCl, will cause no precipitate of silver chloride, AgCl, to form in the solution. 2. If an electric current is passed through the solution, the silver in the solution is carried toward the anode, not toward the cathode. This shows that the silver atoms form part of complex ions which carry negative charges instead of existing in the solu- tion as silver ions. If the silver were in the form of the ion, Ag + , the electric current would carry it toward the cathode, or negative pole. 3. By evaporating the solution a definite compound, potassium silver cyanide, KAgC 2 N 2 , may be crystallized from it. Trans- ference experiments have shown that the ions of this compound are K+ and AgC 2 N 2 ~. (See p. 320.) Many other complex ions are known which are very different in their properties from the ions which unite to form them. It is clear from this illustration that the principle of the solu- bility product depends on the character of the ions which are actually present and not on the amount of a given element which may be present in the solution. Degree of lonization. If two electrodes in circuit with a battery and ammeter are placed at the ends of a narrow, rectangu- lar cell filled with distilled water (Fig. 93) and a concentrated solution of potassium chloride is put in the bottom of the cell, 380 A TEXTBOOK OF CHEMISTRY the ammeter will indicate the passage of an electric current. If, now, the solution is stirred so that the potassium chloride is . + uniformly mixed with I, r the water, it will be ^^_J| n/ y seen that the current increases. Since the amount of potassium chloride in the cell is not changed by mixing Fig. 93 the solution with the water above it, it is evident that a given amount of the salt is more effective in conveying the current in a dilute solution than in a concentrated one. This is explained by the kinetic theory and theory of ionization by supposing, first, that only the ions, K + and Cl~, take part in the conduc- tivity of the solution, and second, that in the ionization reaction : the equilibrium is displaced to the right by dilution because the potassium and chloride ions meet each other to unite less fre- quently in the dilute solution, while the tendency to separate into ions is about the same in one solution as in the other. If the conductivity of solutions of potassium chloride at increasing dilution is measured in this way, the results given in the follow- ing table are obtained. Of course, for the dilute solutions it would be necessary to calculate from the conductivity of a relatively small quantity of the solution what the conductivity of the whole quantity of potassium chloride would be. CONDUCTANCE OF SOLUTIONS OF POTASSIUM CHLORIDE N Potassium Chloride (74.5 grams in 1 liter ) 98.28 mhos 1 0.76 N/10 Potassium Chloride (74.5 grams in 10 liters) 111.97 mhos 0.865 N/100 Potassium Chloride (74.5 grams in 100 liters) 122.37 mhos 0.945 N/1000 Potassium Chloride (74.5 grams in 1000 liters) 127.27 mhos 0.983 N/10000 Potassium Chloride (74.5 grams in 10000 liters) 129.00 mhos 0.996 N/oo Potassium Chloride (74.5 grams in oo liters) 129.5 mhos 1.000 1 The conductance in reciprocal ohms is the reciprocal of the resist- ance in ohms. The unit for conductance is one mho. DEGREE OF IONIZATION 381 Under A is given the conductance in reciprocal ohms * of one gram molecule of potassium chloride in the volumes of solution stated, when placed between two electrodes one centimeter apart and sufficiently large to contain the whole solution between them. For a normal solution the electrodes might be 25 X 40 cm., giving a surface of 1000 sq. cm. It will be seen that the values for the conductances with in- creasing dilution form a converging series from which a value of 129.5 for infinite dilution can be calculated. In accordance with the theory that the electric current is carried only by the ions and that at infinite dilution the compound is completely ionized, it is easy to calculate the degree of ionization by dividing the conductance for any given concentration by the conductance at infinite dilution. Thus the fraction of a normal solution of 98.28 potassium chloride in the form of ions is ' = 0.76. The values given in the last column of the table above have been calculated in this manner. In a previous chapter it was pointed out (p. 112) that a solu- tion containing 46 grams of alcohol in 10 liters of water freezes at 0.184. One containing 74.5 grams of potassium chloride in 10 liters freezes at 0.343. It would seem from this that if the potassium chloride were completely ionized, the depression of the freezing point would be twice that of the alcohol, or 0.368. This would correspond to an increase of 0.184 in the depression of the freezing point. Comparing this with the increased de- pression which is observed (0.343 0.184) we have = 0.864 as the fraction ionized. This is in very close agreement with the results found by the conductivity method. Cane sugar is hydrolyzed by dilute acids to a mixture of glu- cose and fructose (p. 334). It is found that a given amount of hydrochloric or nitric acid causes a much more rapid hydrolysis than an equivalent amount of acetic acid. If we assume that the rate of hydrolysis is proportional to the number of hydrogen ions present, it is possible to calculate from series of experiments 382 A TEXTBOOK OF CHEMISTRY with different acids the relative ionization of the acids. The results of such experiments are, again, in general agreement with the results obtained by the conductivity and freezing-point methods. The three methods all indicate that there are very great differ- ences in the degree of ionization of different compounds. The fact that the results obtained by three methods so radically different are in approximate agreement is very strong evidence that the three effects have a common basis, and no satisfactory theory other than that of ionization has been proposed to account for the phenomena observed. The lack of complete agreement indicates that some factors which are not yet entirely understood modify the effects, just as the mutual attraction of the molecules of gases prevent them from obeying exactly the laws of Avo- gadro and of Boyle. The following table gives the degrees of ionization of a number of common substances. In the measurement and calculation of these values some factors which have not been discussed are involved, but the fundamental principle used is a comparison of the conductivity of the solution for which the degree of ionization is given with the conductivity of the same substance when it is completely ionized. FRACTIONS IONIZED As the mobilities of the H + and OH~ ions change with the ion con- centrations, the conductance ratio ( A/ Aoo ) does not correctly represent the fractions ionized for strong acids and bases. Recent investigations have shown that these substances are ionized to about the same extent as salts of the same type, and values for them are so given in the table. With salts of acids and bases having bivalent ions, the ionization rela- tions are complicated by the presence of intermediate and complex ions, for instance HgCl 2 gives besides Hg ++ and Cl~, the ions HgCl~ and HgCl 4 = . The figures given must therefore be regarded only as relative measures of the tendencies of these substances to form ions. DEGREE OF IONIZATION 383 TENTH HUNDREDTH SUBSTANCE IONS FORMULA SOLUTIONS FORMULA SOLUTIONS V = 10 V = 100 Nitric Acid .... H+, N0 3 - 0.86 0.94 Hydrochloric Acid . H+, Cl- 0.86 0.94 Hydrobromic Acid . . H +, Br~ 0.86 0.94 Hydroiodic Acid . . H+, I- 0.86 0.94 Chloric Acid .... H+, C10 3 - 0.86 0.94 Perchloric Acid . . . H+, C1O 4 - 0.86 0.94 Permanganic Acid . . H+, MnO 4 - 0.86 0.94 Tartaric Acid H+, HC 4 H 4 6 - 0.098 0.31 Acetic Acid .... H+, C 2 H 3 2 - 0.013 0.043 Hydrocyanic Acid . . H+, CN- 0.00008 0.00026 Boric Acid .... H+, H 2 B0 3 - 0.00008 0.00026 Hydrosulfuric Acid . . H+, HS- 0.00095 0.00031 Sulfurous Acid . . . H+, HS0 3 - 0.50 0.70 Carbonic Acid H+, HCO-r 0.0017 0.0055 Phosphoric Acid . . H+, H 2 PO 4 ~ 0.28 0.64 Phenol . . . H+ , C 6 H 5 O- 0.00011 0.00036 Sulfuric Acid .... i (H+, H+, S0 4 =) 0.608 0.832 Oxalic Acid .... i (H+, H+, C 2 4 =) 0.17 0.398 Sodium Hydroxide . . Na+, OH- 0.86 0.94 Potassium Hydroxide . K+, OH- 0.86 0.94 Tetra methyl Ammo- nium Hydroxide . . N(CH 3 ) 4 +, OH- 0.85 0.94 Barium Hydroxide . !(Ba++,OH-,OH-) 0.76 0.88 Ammonium Hydroxide 1 NH 4 +, OH- 0.013 0.042 Water 2 H+, OH~ 0.0000001 0.0000001 1 This gives the fraction of the total ammonia in the solution which is in the form of ammonium, NH 4 +, and hydroxide, OH~, ions. Actually much of the ammonia is present as NH 3 and the proportion of real NH 4 OH ionized is much larger than that given in the table. 2 The value for water gives the fraction of a mol of water in one liter, which is in the form of ions at 25. For comparison with the other values in the table the values under V 10 must be divided by 10 and those under V 100 must be divided by 100. Thus one liter of tenth-normal hydrocyanic acid contains 0.000008 mol in the form of ions and one liter of hundredth-normal acid contains 0.0000026 mol in that form, while one liter of pure water contains 0.0000001 mol. 384 A TEXTBOOK OF CHEMISTRY FRACTIONS IONIZED Continued SUBSTANCE IONS TENTH FORMULA SOLUTIONS V = 10 HUNDREDTH FORMULA SOLUTIONS V = 100 Sodium Chloride Na+, Cl- 0.852 0.935 Potassium Chloride ... K+, Cl- 0.855 0.941 Ammonium Chloride . NH 4 +, Cl- 0.852 0.936 Sodium Nitrate . . . Na+, NO 3 ~ 0.832 0.932 Potassium Nitrate . K+, NO 3 - 0.824 0.935 Silver Nitrate . Ag+, N0 3 - 0.816 0.931 Potassium Chlorate K+, C1O 3 - 0.824 0.933 Sodium Acetate . . Na+, C 2 H 3 O 2 ~ 0.795 0.914 Potassium Cyanide K+, CN- 0.84 0.93 Sodium Bicarbonate . Na+, HCO 3 ~ 0.84 0.93 Potassium Sulfate . . i (K+, K+, S0 4 =) 0.724 0.870 Sodium Sulfate . . . i(Na+,Na+,SO 4 =) 0.704 0.857 Normal Sodium Car- i (Na+, Na+, CO 3 =) 0.71 0.86 Calcium Sulfate . . . i (Ca++, S0 4 =) 0.625 Zinc Chloride . . . i (Zn++,Cl-,Cl-) 0.71 0.86 Zinc Sulfate . . . . i (Zn++, S0 4 =) 0.405 0.633 Copper Sulfate . . . } (Cu++, S0 4 =) 0.396 0.629 Mercuric Chloride . . J(Hg++,Cl-,Cl-) 0.01 0.03 (This table was prepared by Dr. D. A. Maclnnes.) Effect of Degree of lonization. Neutralization. When a volatile product escapes from a mixture, or when a solid substance is precipitated, the equilibria of the reactions which lead to the formation of such compounds are shifted in such a way as to promote their formation. From the table which has just been given it is evident that certain ions cannot exist in any number in solutions which contain certain other ions. There can be very few hydrogen ions, H + , in solutions containing hydroxide, OH~, hydrosulfide, HS~, or hydrocarbonate, HCOs", ions, or in the presence of the ions of any of the weak acids. If the solutions of two substances giving hydrogen and hydroxide ions are mixed in equivalent amounts, the two ions unite to form water and HYDROLYSIS 385 the resulting solution is neutral, if both of the compounds have the same degree of ionization : HCI ; H + + cr NaOH ^ Na + + OH~ H + + OH- ^ H 2 O The equilibrium of the last reaction is so far toward the forma- tion of water that both hydrogen and hydroxyl ions and also practically all of the hydrochloric acid and sodium hydroxide disappear from the solution. This is the ordinary reaction of neutralization. A solution is neutral when the number of hy- drogen, H + , and hydroxide, OH~, ions is equal. An illustration of the effect of adding hydrochloric acid to a solution containing carbonate ions was given above (p. 375). Hydrolysis. It will be seen from the table that the ionization of carbonic acid, ^COs, to hydrogen, H + , and hydrocarbonate, HCOs", ions is very slight. The further ionization of hydrocar- bonate ions to hydrogen, H + , and carbonate, CO 3 = , ions must be almost vanishingly small in the presence of an excess of hy- drogen ions. A solution of sodium carbonate, Na2CO 3 , which may ionize, at first, as follows : Na 2 CO 3 ^ Na + + Na + + CO 3 = immediately gives with the ions of water COr + H + + OH- ^ HCOr + OH- This formation of hydrocarbonate ions, HCO*-, reduces the number of hydrogen ions and the solution must react alkaline because of excess of hydroxide ions, OH", present. Similar re- actions, due to the ions of the water and called hydrolysis, occur with the alkali metal salts of all of the very weak acids and especially with the salts of weak dibasic and tribasic acids, such as sulfides, borates and phosphates. The hydroxides of ferric iron, Fe(OH) 3 , aluminium, A1(OH) 3 , chromium, Cr(OH)s, and of many other elements are so insoluble that it is impossible to determine the degree of their ionization, 386 A TEXTBOOK OF CHEMISTRY and there is reason to think that the ionization of the hydroxyl in basic compounds such as FeCl 2 OH, which probably exist in so- lution, is very trifling. In a solution of ferric chloride, therefore, such reactions as the following occur : FeCl 2 + -f OH" + H + = FeCl 2 OH + H + Basic Ferric Chloride The resulting solution, which contains an excess of hydrogen ions, will react acid. These facts may also be stated in the form : Weak acids cannot neutralize strong bases and weak bases cannot neutralize strong acids completely, when mixed in equivalent proportions, because of the hydrolysis of the salts formed. It is chiefly because of these relations that acids and bases which are but slightly ionized in their solutions are called " weak." The degree of ionization furnishes the only satisfactory basis for classifying acids or bases as " weak " or " strong." Illustration of the Strength of Acids. If solutions of potassium iodide (KI, 6 grams per liter) and potassium bromate (KBrO 3 , 1 gram per liter) are mixed in equal proportions, no reaction occurs between them, but on the addition of an acid, iodine is liberated and colors the solution yellow or brown : 6 HI + HBrO 3 = HBr + 3 I 2 + 3 H 2 O If the acid is dilute, the reaction is sufficiently slow so that its progress can be noted by a slow change in the color. The rate of the change is proportional to the number of hydrogen ions present. If equal volumes of the iodide-bromate solution are placed in three glasses and there is added to these, respectively, 5 cc. of tenth normal hydrochloric acid (3.65 g. HC1 per liter), 5 cc. of tenth normal oxalic acid (5.3 g. H 2 C 2 O 4 .2 H 2 O per liter) and 5 cc. of tenth-normal acetic acid (6.0 g. HC 2 H 3 O 2 per liter), a very rapid change in color will occur in the first solution, a much slower change in the second and a very slow change, indeed, in the third. On the other hand, if 5 cc. of each acid are placed in three beakers and 5 cc. of tenth-normal sodium hy- USE OF INDICATORS 387 droxide are added to each, the three solutions, containing sodium chloride, NaCl, sodium oxalate, Na 2 C 2 O 4 , and sodium acetate, NaC 2 H 3 O 2 , respectively, will be neutral toward phenol phthalein or litmus. The experiment shows that while the three acids differ very greatly in " strength " as shown by a reaction which depends on the number of hydrogen ions, H + , they give at a given dilution, they are nearly equal in their power of neutral- ization, which depends for these acids and the indicator chosen on the number of hydrogen ions which can be obtained from them by complete ionization. It should be remembered, how- ever, that sodium oxalate and sodium acetate are not, strictly speaking neutral salts. The reason for this lack of neutrality will be clearer after a study of the following paragraph. Use of Indicators. It has been stated (p. 122) that indicators are colored substances which exist in two forms, one of which is stable in the presence of hydrogen, H + , ions while the other is stable in the presence of hydroxide, OH~~, ions. A neutral solution is one in which the numbers of hydrogen and hydroxide ions are equal. It has been pointed out in the table on p. 383 that the number of mols of water ionized in one liter is 0.0000001, or 10~ 7 at 25. Since in the ionization reaction we must have C H + X COH~ = constant so long as the solution consists mostly of water, it follows that in dilute solutions C H + X COH- = 10~ 7 X JO' 7 = 10~ 14 If an acid is added to water in sufficient amount to increase the concentration of the hydrogen ions from 10~ 7 to 10~ 6 , the concentration of the hydroxide ions must fall to 10~ 8 . Only 0.0000009 gram-mol of hydrogen ions would be required to produce such a change in a liter of water, and this would be given by about 0.001 cc. of tenth-normal hydrochloric acid, while 0.01 cc. would increase the hydrogen ion concentration to ID'*. 388 A TEXTBOOK OF CHEMISTRY t-i 1 \ -o cc S t 1 } 1 t jb 1 1 t Ii o > t t o 1 1 \ 1 1 1 t t t o \ t \ 1 4 a t t 1 > 1 8 1 { * * 1 \ -o oc t t II 1 1 ? "1 o \ k 1 i Greenish Yellow. i t t 1 1 oj cf 1 8 \ 1 1 t t t S 1 > } |i 1 1 "1 1 i-H 4 1 ' DC t t t S 3 i rr g, B ii i^ o rH $ S 1 1 > t t t 11 A| 5s 11 -ti S ii T S 2 o ii i Jl t t 11 t t 5 ^ i 1 ^ 1 o i"S 11 o t t i 1 t | ^ CO -o t 1 f t t t t t t t o t t t rH o Jl |1 1 i 1 1 i 1 i I i 1 i i 1 a i True Acidity =C H + True Alkalinity =C j f i o Sodium-alizarinsulfo T3 < Guiac Tincture Phenolphthalein Thymolphthalein Tropaolin O Trinitrobenzene Benzopurpurin Methyl Red Congo Red Litmus j USE OF INDICATORS 389 Indicators are weak acids or bases, and the change in color does not usually occur in an exactly neutral solution. The table on p. 388 gives the concentrations of hydrogen and hydroxide ions at which the change in color occurs for some of the more common indicators. It would seem, at first thought, that only an indicator which changes exactly at the neutral point would be suitable, but this is by no means true. In titrating a strong acid with a strong base a few hundredths of a cubic centimeter of tenth-normal acid or alkali will carry the concentration of hydrogen or hy- droxide ions so far to one side of the neutral point that any of the indicators for which the acidity is between 10~ 6 and 10~ 9 will give a sharp end point. In titrating a weak acid, such as acetic acid, HC 2 H 3 O 2 , with a strong base, such as potassium hydroxide, KOH, as the neutral point is approached the potassium acetate, KC 2 HaO 2 , formed is much more highly ionized than the acetic acid and the acetate ions carry the ionization : HC 2 H 3 O 2 H + + C 2 H 3 O 2 - far to the left. The number of hydrogen ions then becomes very small, while a considerable amount of acetic acid is still unneutralized. Under such conditions the change in color for methyl orange or methyl red may appear before the acid is completely neutralized and the end reaction will not be sharp, i.e. the change in color will appear gradually during the addition, sometimes, of a cubic centimeter or more of the alkali. But if phenol phthalein is used, the change in color will not occur till the true neutral point is passed and then a very slight excess of alkali will carry the concentration of the hydroxide ions far beyond the neutral point. The end will be sharp and will cor- respond closely to the exact neutralization of the acid. When a weak base such as ammonium hydroxide, NH 4 OH, is titrated with a strong acid, as hydrochloric acid, HC1, the conditions are reversed and such indicators as methyl red, 390 A TEXTBOOK OF CHEMISTRY methyl orange or cochineal, which change color in a faintly acid solution, are most suitable. With very weak acids or bases the hydrolysis of the salts formed may carry the acidity or alkalinity of the normal salt far to one side of the true neutral point. An illustration of the use of indicators in such a case will be given later (p. 464). Systematic Study of the Metals. In discussing the metals it is natural to consider first, as with the non-metals, their prepa- ration and properties. After this their compounds may be dis- cussed in the same order which has been followed in the consid- eration of the nonmetallic elements : oxides and hydroxides ; chlorides, hypochlorites, chlorates, bromides, iodides, fluorides ; sulfides, sulfites, sulfates ; nitrites, nitrates, phosphates, arsen- ates ; carbonates, bicarbonates, salts of organic acids, cyanides ; silicates ; borates. Most of the metals form nearly all of these classes of compounds, but only those compounds which are of some particular scientific, historical or industrial importance can be mentioned in this book. Before taking up the individual metals it seems desirable to give some general statements with regard to the preparation or metallurgy of the metals and about the preparation and proper- ties of the various classes of compounds. Metallurgy. The first metals used by man were those which are found free in nature, such as copper, silver and gold. The use of these metals and their alloys marks the beginning of the " Bronze Age," which reaches back into prehistoric times very late and recent when we consider the long history of the race, but early in a period when mankind became organized in society and differentiated sharply from the animal world. The discovery of methods of reducing iron from its ores by means of wind furnaces probably occurred during the period of history which has been recorded in inscriptions, but still so early that no definite record is to be found. Inscriptions in Egypt show that iron was made there at least 3000 years ago. From that period till the middle of the nineteenth century all methods for industrial metallurgy depended on the use of charcoal, coal or METALLURGY 391 other fuels containing carbon. The most important of these methods depended on the direct reduction of oxides of the metals by means of these fuels, while a few depended on the roasting of a sulfide (mercury, p. 485) or on the roasting of a sulfide followed by reduction through the interaction of an oxide with a sulfide or sulfate (copper, p. 427, and lead, p. 513). The first preparation of aluminium by Deville in 1854 in sufficient amount to demonstrate its valuable properties led to a strong desire to secure cheap sodium for use in its production. This resulted in the development of the reduction of sodium carbonate to metallic sodium by means of carbon and the use of the latter for the preparation of aluminium from its chloride, Aids. But this method of preparing aluminium did not attain any considerable industrial importance. Electrolytic methods were used for the deposition of thin films of copper and other metals as early as 1836, but electrical methods for preparing and refining metals could not be used on a large scale till the development of dynamos during the last quarter of the nineteenth century made the production of relatively cheap electrical energy possible. The first industrial use of an electric furnace seems to have been its application to the manufacture of aluminium bronze by the Cowles Brothers of Cleveland, Ohio, in 1884. Their labora- tory experiments with the method began in 1882. The Hall method for the electrolysis of aluminium oxide dissolved in cryolite soon rendered the Cowles furnace industrially worth- less for this particular purpose, but electric furnaces are now extensively used for metallurgical processes and for many other forms of chemical manufacture. The relatively cheap electrolytic manufacture of aluminium has not only given an abundant supply for use in the metallic form, but has led to the development of processes for the pro- duction of chromium and other metals by heating their oxides with aluminium. It will be seen from this brief sketch of the historical develop- ment of metallurgy that many important processes are of very 392 A TEXTBOOK OF CHEMISTRY recent origin. It is only a very few years since several metals which now have very important industrial uses, were scarcely more than scientific curiosities. Further rapid development along such lines is to be confidently expected. Oxides. All metals, without exception, may be combined with oxygen and nearly all metals combine with the element directly at suitable temperatures. Even the so-called " noble " metals, gold, silver and platinum combine with oxygen under some conditions, but the oxides are very unstable, are easily decomposed by heat alone and are reduced by hydrogen at ordinary temperatures. A few metals, especially sodium, potassium and barium, com- bine with oxygen to form peroxides, in which two oxygen atoms are united, as in Ba<Q | , but in nearly all of the metallic oxides Nj the oxygen seems to be united only with the metal, and the valence of the metal is apparent from the formula of the oxide. Nearly all nitrates, carbonates and hydroxides are decom- posed by heat with the formation of oxides : 2Pb(NO 3 ) 2 = 2PbO + 4NO 2 + O 2 CaCO 3 = CaO + CO 2 Cu(OH) 2 = CuO + H 2 O It is doubtful if any oxide dissolves appreciably in water as an oxide, and those oxides which do not combine with water to form hydroxides are practically insoluble. Hydroxides. The alkali metals, sodium, potassium, etc., and the alkali earth metals, calcium, strontium and barium decompose water with the formation of hydroxides at ordinary temperatures : Na + HOH = NaOH + H Ca + 2HOH = Ca(OH) 2 + 2H Magnesium decomposes water at 100. Other metals, as zinc and iron, which decompose water at higher temperatures, form oxides instead of hydroxides, though it is, of course, possible SOLUBILITY OF SALTS 393 that the latter are formed at first and immediately decom- posed. Iron is converted by the combined action of water and air into iron rust, a combined oxide and hydroxide having the composi- tion of the mineral limonite, Fe 2 O 3 .2Fe(OH) 3 . Practically all hydroxides of the metals except those of the alkali and alkali-earth metals are insoluble in water. For this reason the hydroxides of nearly all other metals are precipitated from solutions of their salts by solutions of sodium or potassium hydroxide. In a few cases, especially those of silver, cuprous copper, mercurous and mercuric mercury, these hydroxides precipitate an oxide instead of the hydroxide, doubtless because the hydroxides of these metals are unstable : AgNO 3 + NaOH = [AgOH] + NaNO 3 [2AgOH]=Ag 2 + H 2 Solubility of Salts. Practically all salts of the alkali metals (lithium, sodium, potassium, ammonium) are soluble in water, the only important exceptions being sodium pyroantimonate, Na2H 2 Sb 2 O7.6 H 2 O, potassium and ammonium chloroplatinates, K 2 PtCl 6 , and (NH 4 ) 2 PtCl 6 , 1 potassium perchlorate, KC1O 4 , and potassium cobaltini trite, K 3 Co(NO 2 )6, or potassium silver cobaltinitrite, K 2 AgCo(NO 2 )e. Some of these are, however, more soluble than those salts of other metals which are usually counted as insoluble. There is probably no salt which is wholly insoluble in water. Nearly all salts of the strong monobasic and bibasic acids are also soluble. This includes chlorides, bromides and iodides, fluorides, chlorates and perchlorates, sulfites and sulfates, ni- trites and nitrates. The most important exceptions are the chlorides, bromides and iodides of silver, cuprous copper, mer- curous mercury and lead, AgCl, AgBr, Agl, Cu 2 Cl 2 , Cu 2 I 2 , Hg 2 Cl 2 , Hg 2 I 2 , PbCl 2 (slightly soluble), PbBr 2 , and PbI 2 , mer- 1 Rubidium and caBsium chloroplatinates, Rb2PtCl and Cs 2 PtCl are still less insoluble. 394 A TEXTBOOK OF CHEMISTRY curie iodide, HgI 2 , calcium fluoride, CaF 2 , barium sulfite, BaSO 3 , and calcium, strontium, barium and radium sulfates, CaSO4 (slightly soluble), SrSO 4 , BaSO 4 , RaSO 4 . Normal salts of phosphoric, H 3 PO 4 , arsenious, H 3 AsO 3 , arsenic, H 3 AsO 4 , carbonic, H 2 CO 3 , silicic, H 2 SiO 3 , etc., and boric, H 3 BO 3 , acids, with the exception of those of the alkalies, are insoluble. Sulfides other than those of the alkalies are either insoluble in water or are hydrolyzed with the formation of a hydrosulfide (as Ca(SH) 2 ), and a hydroxide, or of hydrogen sulfide, H 2 S, and an insoluble hydroxide, such as A1(OH) 3 . CHAPTER XXIII ALKALI METALS : LITHIUM, SODIUM General Properties of the Alkali Metals. The alkali metals are univalent elements which combine with hydroxyl to form the strongest bases, hydroxides which are easily soluble in water and which are largely ionized in solutions of moderate concen- trations. They are the most active of the metallic elements,, decomposing water rapidly at ordinary temperatures and tar- nishing almost instantly in ordinary air, owing to the formation of a film of hydroxide. Their affinity for the halogens is also so strong that sodium and potassium have often been used to decompose halides for the preparation of metals and other elements. As has been stated in the last chapter, nearly all salts of the alkali metals are soluble in water. Normal salts of weak acids, such as the sulfides, carbonates, cyanides, phosphates, silicates and borates are hydrolyzed by water, and their solutions have a strongly alkaline reaction. The metals of the group have a low specific gravity, lithium, sodium and potassium being lighter than water. Their melting points range from 186 for lithium to 26.5 for caesium and the boiling points, from above a red heat for lithium and 742 for sodium to 270 for caesium. Lithium, Li, 6.94, is usually considered as one of the rarer elements and of comparatively little importance. It is found in a number of silicates and in small amount in practically all natural waters. The metal has a specific gravity of only 0.51. Hydrogen and helium are the only elements which are lighter than lithium when in the solid or liquid state. The metal com- bines with hydrogen to form the hydride, LiH, and with nitrogen 395 396 A TEXTBOOK OF CHEMISTRY to form the nitride, Li 3 N. It has been used in the separation of argon from the .atmosphere because of its strong affinity for nitrogen. Lithium carbonate, I^COs, may be decomposed to lithium oxide, Li2O, and carbon dioxide by heating it in a current of hydrogen. Both the carbonate and the phosphate, LisPO-i, are difficultly soluble in water. In these properties lithium re- sembles magnesium, the second element of the second group, rather than the other alkali metals. Beryllium, the first ele- ment of the second group, approaches aluminium in its proper- ties in a similar manner. Lithium Urate, LiC5H 3 O 3 N 4 , is soluble in water, and this fact led physicians to the belief that the administration of lithium carbonate or of natural waters containing lithium would be beneficial to patients suffering from rheumatism or gout and they have been much employed as remedies in those diseases. A more careful study has shown that these compounds are worth- less for such a purpose, but the ingestion of large quantities of water with the lithium compounds probably exerts a beneficial effect. Lithium compounds impart to the Bunsen flame a brilliant red color and give a spectrum of two red lines, one of which is very bright. Atomic Weight of Lithium. Law of Dulong and Petit. It has been pointed out (p. 92) that the most satisfactory method of selecting the true atomic weight of an element consists in finding the weight of the element contained in a gram-molecular volume (22.4 liters at and 760 mm.) of that gaseous compound which contains the smallest quantity of the element in this volume. But lithium forms no compound whose weight in the gaseous form has been determined, and a considerable number of other elements form no compounds which can be converted into gases without decomposition. The atomic weights of such elements must, of course, be selected in a different manner. For this purpose the law of Dulong and Petit, discovered in 1819, has been useful. These chemists found that the quantity LAW OF DULONG AND PETIT 397 of heat required to raise the temperature of one gram-atom of an element one degree is approximately 6.6 calories. If this quantity of heat is applied to 7 grams of lithium or to 65 grams of zinc or to 200 grams of mercury, it will, in each case, raise the temperature one degree. The law is also frequently stated that the specific heat of an element multiplied by its atomic weight is a constant quantity. The following table will make this clear : ELEMENT SPECIFIC HEAT ATOMIC WEIGHT SP. HT. X AT. WT. Lithium 0.94 7. 66 Graphite (at 11) . . Graphite (at 977) .... Silicon 0.16 0.467 16 12. 12. 28.4 1.9 5.6 4.5 Calcium Zinc 0.17 . 0093 40. 654 6.8 6.1 0.084 80. 6.7 0.033 200. 6.7 0.03 207. 6.2 It will be seen from the table that graphite and silicon depart rather widely from the law, though the former approaches it more closely at high temperatures. All of the metallic elements and all elements having atomic weights above 40 conform ap- proximately to the law. The law is at best, however, only approximate and is of service only in selecting between rather widely divergent possible values for an atomic weight. Thus the atomic weight of calcium might be 20, 40 or 60, according as the formula of the chloride is CaCl, CaCl 2 or CaCl 3 . But of these three values only an atomic weight of 40 agrees with the law. The laws of Avogadro and of Dulong and Petit have usually been considered as independent and wholly unrelated. 1 A 1 See, however, G. N. Lewis, J. Am. Chem. Soc. 29, 1165 and 1516 (1907). 398 A TEXTBOOK OF CHEMISTRY little consideration, however, shows us that if we accept the kinetic-molecular theory, this is not the case. At foundation Avogadro's law depends on the fact that molecules of different weights exchange energies, when in collision with each other or with the walls of the containing vessel at a given tempera- ture, in such a manner that the average value of J mv 2 (m = mass, v = velocity) is constant and is independent of the weight of the molecule. The law of Dulong and Petit must depend on a similar property of the atoms of the elements in the solid or liquid state. The Quantum Theory. Quite recently a new theory of molecular, atomic and radiant energy, called the quantum theory, has been developed by Plank, Einstein, Nernst, Sackur, Debye, Sommerfeld and others. The theory supposes that there are in the atoms of the elements, or associated with them, resonators or oscillators of such a nature that they can emit energy only in definite, unit quantities. The resonators may be atoms, ions, or electrons ; i.e. they may be particles with or without electrical charges. The theory seems to give a satisfactory explanation of the low and variable specific heats of some of the elements, of some photo-electric effects which were previously hard to under- stand and of a variety of other phenomena. From the character of the men who are working on the theory and the striking results already attained it seems likely to be developed very rapidly in the near future. Sodium, Na, 23. Sodium chloride, NaCl, or common salt, forms about 75 per cent of the residue left by the evaporation of sea water. It is also found in enormous beds of rock salt in Germany, Louisiana, Kansas, Utah and elsewhere and in strong brines found by boring deep wells in very many places. Sodium is taken up by seaweeds in their growth very much as potassium is taken up by land plants, and the ashes of sea- weeds contain considerable quantities of sodium carbonate. Deposits of sodium sesquicarbonate, NaHCO 3 .Na 2 CO3.2H 2 O, called trona, of sufficient extent to be of industrial importance are found in Egypt and in Venezuela. The occurrence of SODIUM 399 borax, Na2B4O7.10H2O, in lakes in California has been men- tioned; also that of sodium nitrate, NaNOs. These are of value, primarily, for the boron and nitrogen which they contain. Sodium is a constituent of practically all silicious rocks. Metallurgy. Properties. Metallic sodium and potassium were first prepared by Sir Humphry Davy in London in 1807 by the electrolysis of moist sodium and potassium hydroxides. The discovery awakened very great interest, both because the metals showed very striking properties, quite different from those of any metals hitherto known, and because it indicated very clearly that many other earthy substances contain ele- ments which could not at that time be prepared in the free state. Within a few years the new metals proved effective agents for the liberation of a number of other elements. Sodium may also be prepared by heating a mixture of sodium carbonate and carbon : Na 2 CO 3 + 2C = 3CO + 2Na The sodium, which boils at 742, distills from the mixture and is collected in iron condensers. During comparatively recent years metallic sodium is prepared commercially by various electrolytic methods, from the hydroxide, the nitrate or the chloride. Sodium is a silver white metal which tarnishes instantly on exposure to moist air. In dry air at 300 to 400 it is oxidized to sodium peroxide, Na 2 O 2 . It is kept in sealed cans or under kerosene to protect it from the action of the air. It melts at 97.5 and boils at 742, giving a dark green vapor. The specific gravity of the solid is 0.97. When a small piece of sodium is thrown on water the heat of the reaction causes it to melt and the globule of metal rolls rapidly over the surface of the water without taking fire or igniting the hydrogen, differing in this respect from potassium. If the metal is thrown on a piece of filter paper floating on the water, the heat is concentrated and the hydrogen takes fire and burns with a yellow flame. In both cases, of course, sodium hydroxide is formed. 400 A TEXTBOOK OF CHEMISTRY Sodium is manufactured in considerable amounts for use in preparing sodium peroxide, Na 2 O 2 , for the preparation of a mixture of potassium and sodium cyanides from potassium ferro- cyanide (p. 319), and for use in the synthesis of a variety of organic compounds. The Alkali Industry. Sodium is an essential constituent of common soap, of glass, of salsoda, or washing soda, and of baking soda. As common salt, NaCl, is much cheaper than any other compound of sodium, it now furnishes the basis for the preparation of all of these substances, but it is necessary to prepare from it, at first, one of the sodium carbonates or sodium hydroxide. Till the close of the eighteenth century an impure sodium carbonate from Egypt and the ash of seaweeds were used as the sources of sodium carbonate and sodium hydroxide for the manufacture of hard soap, while potassium carbonate from wood ashes was also extensively used for the manufacture of soft soap. During the disturbed commercial relations which followed the French Revolution, the foreign supplies of sodium carbonate were cut off and all available potassium compounds were needed for the manufacture of gunpowder. This caused* the French government to offer a prize for a method of manu- facturing sodium carbonate from salt. The prize was awarded to Leblanc and his process was used for a short time in France, but could not there compete with the sodium carbonate from other sources when commercial relations with other countries were again established. Leblanc himself did not secure any permanent advantage from his invention and died in a poor- house. About twenty years later Musgrave, in England, took up the process again and succeeded in making it a commercial success. It held the field of alkali manufacture almost without competition for fifty years. About 1860 it had to meet the competition of the ammonia-soda process, the principles of which had been discovered in 1838, but which was first put into successful operation by Solvay more than twenty years later. From then till the close of the nineteenth century the Leblanc process continually lost ground in competition with tke SODIUM HYDROXIDE 401 Solvay manufacture, maintaining a precarious existence only by the most careful conservation of the by-products, hydrochloric acid or chlorine and sulfur. In 1900 only two large Leblanc fac- tories remained in the world, one in England and one in Germany. The most serious difficulty with the ammonia-soda process is that the chlorine of the salt is converted into calcium chloride or magnesium chloride, practically worthless products. In the closing years of the nineteenth century, with the stimulus of comparatively cheap electrical energy, many electrolytic pro- cesses were developed which give both chlorine and sodium hydroxide directly from salt. It seems probable that these processes will eventually displace the Solvay process, at least for the production of caustic alkali. Sodium Hydroxide. So long as trona from Egypt or the ashes of sea plants were used, sodium hydroxide was prepared by treating a solution of these with 'slaked lime. A more pure sodium hydroxide was prepared in the same way from the sodium carbonate of the Leblanc or ammonia-soda processes : Na 2 CO 3 + Ca(OH) 2 = 2 NaOH + CaCO 3 The reaction depends, of course, on the insolubility of the calcium carbonate. In very concentrated solutions the reaction may be reversed because calcium hydroxide, Ca(OH)2, is also difficultly soluble, and with a high concentration of hydroxide ions, OH~, the solubility product for that substance may be exceeded even in a solution containing so insoluble a salt as calcium carbonate. As has been stated above, sodium hydroxide is now prepared on a large scale by electrolysis. Many different patents have been issued for such processes and it is probably too soon to decide which forms are likely to prove permanently suitable. In some forms a diaphragm, usually of asbestos, is used to sepa- rate the anode space from the cathode. The anode must be of carbon or of platinum or platinum-iridium and the anode space is inclosed so that the chlorine liberated may be collected and utilized. The cathode is usually of iron. Chlorine is liberated 402 A TEXTBOOK OF CHEMISTRY at the anode, while hydrogen is liberated at the cathode and sodium, Na + , and hydroxide, OH~, ions remain in solution, the hydrogen, of course, coming from the water, though the transfer of ions through the solution is mainly that of sodium and chloride ions. The points aimed at are to secure as high a concentration of hydroxide, OH~, and as low a concentration of chloride, Cl~, ions as possible in the cathode space and the reverse of this around the anode. To this end the salt solution is continuously introduced at the anode while the hydroxide solution is removed from the cathode. The presence of hy- droxide at the anode leads to the formation of hypochlorite and loss of current. The hydroxide solution may be concentrated till nearly all of the salt, NaCl, remaining in it separates, leav- ing a solution in which nearly all of the sodium is in the form of hydroxide. This solution is then evaporated till the water has been expelled, which requires a comparatively high temperature. The sodium hydroxide obtained in this way is sufficiently pure for the manufacture of soap and for many other industrial uses. The Castner-Kellner apparatus gives an almost pure solution of sodium hydroxide directly. It consists of a slate box divided into three com- partments by two partitions, which fit only loosely in grooves in the bottom of the box (Fig. 94). Mer- .JL Fig. 94 cury placed on the bottom of the box seals these, giving a contin- uous metallic layer for the three compartments, but prevents a dilute solution of sodium hydroxide placed in the central com- partment from mixing with the brine placed in the two side compartments. Graphite anodes are placed in the two side compartments and an iron cathode in the central one. Chlo- rine is evolved from the anodes and is, of course, collected and used for the manufacture of ' bleaching powder or for some other purpose. The mercury in the two side compartments is SODIUM HYDROXIDE 403 negative as compared with the graphite anodes and the sodium liberated at its surface combines with it to form a liquid sodium amalgam. By a slight tilting motion the amalgam is caused to flow alternately to one side or the other and so is brought into the central compartment. Here it is positive with refer- ence to the more negative cathode and the hydroxide ions brought to its surface by the current combine with the sodium of the amalgam to form sodium hydroxide, while the hydrogen ions of the water are discharged and liberated as free hydrogen, H 2 , at the surface of the iron cathode. The hydroxide solu- tion is kept at a constant concentration by introducing water at one side and removing some of the solution at the other. Salt is added from time to time to the side compartments. Sodium hydroxide is a white solid, which melts at a red heat* It deliquesces on exposure to the air, but the solution soon ab- sorbs carbon dioxide and then evaporates, leaving a residue of sodium carbonate. Sodium hydroxide is used in the manufacture of ordinary hard soaps and may be used for the preparation of many of the salts of sodium. The specific gravity of solutions of different concentrations is as follows : DENSITY OF SOLUTIONS OF SODIUM HYDROXIDE SPECIFIC GRAVITY PER CENT NaOH GRAMS OP NaOH IN 100 CO. 1.0555 5 5.277 1.1111 10 11.111 1.1665 15 17.497 1.2219 20 24.438 1.2771 25 31.928 1.3312 30 39.936 1.3838 35 48.433 1.4343 40 57.372 1.4828 45 66.726 1.5303 50 76.515 404 A TEXTBOOK OF CHEMISTRY Sodium hydroxide dissolves in water with the evolution of considerable heat, and it will be seen from the table that the addition of a small amount of sodium hydroxide causes the volume of the water to diminish. Sodium Oxide, Na2O. With the exception of lithium oxide, Li2O, the oxides of the alkali metals cannot be prepared by heating the hydroxides or carbonates. In this respect they differ from all other metallic oxides. Sodium oxide may be prepared by heating sodium hydroxide, NaOH, with metallic sodium. It combines with water to form the hydroxide, but has, at present, no practical importance. Sodium Peroxide, Na 2 O 2 , is prepared by heating metallic sodium to 300^400 in dry air. The sodium is placed in shallow aluminium trays, which are passed slowly through long ovens one way, while air passes in the opposite direction. In this way the pure sodium comes at first in contact with air which has been deprived of most of its oxygen and a too vigorous action is avoided, while the action is finally completed at the other end with fresh air. The use of fused sodium peroxide containing a little copper oxide, under the name of " oxone," for the preparation of oxygen has been mentioned (p. 21). It may be converted by cold, moist air into a hydrate, Na 2 O 2 .H 2 O, which can be dis- solved in water with little decomposition. It is hydrolyzed, how- ever, to sodium hydroxide, NaOH, and hydrogen peroxide, H 2 O 2 . On treatment with cold, dilute acids sodium peroxide gives a solution of hydrogen peroxide, H2O 2 , which is used to bleach silk, wool, hair and other substances which would be affected injuriously by chlorine. Sodium peroxide is also a very valuable oxidizing agent for many laboratory uses. Sodium Chloride. Salt is sometimes obtained by direct mining, but rock salt is rarely sufficiently pure for direct use, and it offers especial difficulties in mining, owing to its effect in dulling tools used to cut it and* because blasting does not loosen it up satisfactorily. It is found better to prepare an opening SODIUM CHLORIDE 405 in the bed of salt and allow water to stand in contact with it till a saturated solution is obtained, many of the impurities present remaining undissolved and settling to the bottom. The solution is then pumped out and evaporated to crystallize the salt. As salt is nearly as soluble in cold as in hot water, it cannot be crystallized practically by cooling a hot solution. In some places, especially at Syracuse, New York, and in Michi- gan, saturated brines are obtained directly from artesian wells. For the evaporation of the brines triple-effect evaporators are used to advantage. The principle of these is shown in the accompanying diagram (Fig. 95). The brine in pan A is heated directly, or by superheated steam in coils or in a false bottom, Fig. 95 and it boils under atmospheric pressure. The steam from this pan passes under B, in which a pressure of perhaps 550 mm. is maintained so that the condensation of the steam from the first pan beneath it will cause the brine which it contains to boil. The steam from this will, in turn, cause the brine in C to boil under a pressure of 300 mm. By such an arrangement a given weight of coal will evaporate nearly three times as much water as it would if used directly in the usual manner. 1 Commercial salt contains small quantities of various impuri- ties. The most objectionable, perhaps, is magnesium chloride, 1 It may be remarked, incidentally, that when this process is used for the preparation of distilled water, from waters containing relatively small amounts of solids in solution and the successive differences of pressure may be much less, as many as ten boilers may be used in series. For a description of the Yaryan Evapora- tor, which uses a modification of this system, see J. Soc. Chem. Ind. 14, 112 (1895). 406 A TEXTBOOK OF CHEMISTRY which makes it hygroscopic or even deliquescent in moist air. Pure sodium chloride can be obtained by precipitating a solu- tion of salt with concentrated hydrochloric acid. Sodium chloride crystallizes in cubes. It melts at 820 and may be volatilized at a high temperature. The crystals usually decrepitate on heating, owing to water inclosed in them. Salt is an essential constituent of human diet, furnishing chlorine for the hydrochloric acid of the gastric juice. Sodium Sulfate. Glauber's Salt, Na 2 SO 4 .10 H 2 O. By heat- ing salt with the theoretical amount of sulfuric acid it may be converted almost quantitatively into anhydrous sodium sulfate, Na 2 SO4. The operation is carried out on a large scale as the first step in the Leblanc soda process. The anhydrous sulfate is also used in the manufacture of glass. The crystallized hydrate, Na 2 SO 4 .10 H 2 O, has a solubility in water which increases very rapidly with rising temperature till the transition point, 32.383, is reached. If crystals of the hydrate are heated above this temperature, they are transformed into a mixture of anhydrous sodium sulfate and a saturated solution of the latter, which is less soluble than the crystallized hydrate at temperatures above this point. The transition is accompanied by an absorption of heat in very much the same manner as the melting of ice and may be used as an accurate, fixed point for the correction of thermometers. (Richards, Z. physik. Chem. 43, 465.) * The transition point is a quadruple point in the nomenclature of the phase rule (p. 107), the four phases being water vapor at a pressure of 30.8 mm., the hydrate, Na 2 SO 4 .10H 2 O, the anhydrous salt, Na 2 SO 4 , and the saturated solution. As there are four phases and only two components, sodium sulfate and water, the system is invariant and there can be no change in temperature or pressure without the disappearance of one of the phases. 1 1 Practically, the transition point is determined in contact with air at atmospheric pressure and the effect of pressure is not con- sidered. In the presence of the vapor phase only, the temperature SODIUM SULFATE 407 The solubility of sodium sulfate is shown in the diagram, Fig. 96. The concentrations are given for the anhydrous salt through- out. Below the transition point the hydrate separates on evap- orating or cooling the solution, though supersaturated solutions Solubility Grams of Na 2 SO 4 in 100 grams of water 1-1 o o o o o o c ' S o o ta \ \ o eg X. / c a --*J o M C / Fig. 96 of the transition point would be slightly different, just as the true transition point for water, ice, vapor is 0.0076 higher than the melt- ing point of ice under atmospheric pressure, which is used as the zero point for thermometers. 408 A TEXTBOOK OF CHEMISTRY are easily formed (p. 80). Above the transition point the anhy- drous salt separates on heating or on evaporation. Sodium sulfate is one of the chief active constituents of Hunyadi water and of some other similar medicinal waters. Acid Sodium Sulfate or Sodium Bisulfate, NaHSO4, is formed as the first step in the preparation of the sulfate from salt, or in the manufacture of nitric acid from sodium nitrate, NaNO 3 o At about 300 it loses water and is converted into sodium pyro- sulfate, Na 2 S2O7. At a still higher temperature it is decom- posed into sodium sulfate and sulfur trioxide, SO 3 . Sodium pyrosulfate is used in analytical chemistry to dissolve aluminium oxide, A1 2 O 3 , ferric oxide, Fe 2 O 3 , and titanium oxide, TiO2, after these have been brought to a difficultly soluble form by ignition. (Hillebrand, Analysis of Silicate and Carbonate Rocks, Bulletin 422, U. S. Geol. Survey, p. 105.) Sodium Sulfite, Na 2 SO 3 .H 2 O, is prepared by burning sulfur and passing the sulfur dioxide formed through a solution of sodium carbonate. It is used as a reducing agent in photog- raphy and for the preparation of sodium thiosulfate. Acid Sodium Sulfite, or Sodium Bisulfite, NaHSO 3 , is pre- pared by passing sulfur dioxide in excess into a solution of sodium carbonate. It is sometimes used as an addition to cider to stop fermentation. A 40 per cent solution is very conveniently used for the preparation of sulfur dioxide in the laboratory. Sodium Hyposulfite, Na 2 S2O 4 . A solution of this salt is prepared by the action of zinc on a solution of sodium bisulfite, NaHSO 3 , containing an excess of sulfurous acidjHaSOs. 2 NaHSO 3 + H 2 S0 3 + Zn = Na 2 S 2 O 4 + ZnSO 3 + 2 H 2 O Sodium hyposulfite is a powerful reducing agent and is used especially to reduce indigo to indigo white (p. 341). Sodium Thiosulfate, Na2S 2 O 3 .5 H 2 O, is prepared by dissolv- ing sulfur in a solution of sodium sulfite. It dissolves the halides of silver (AgCl, AgBr and Agl) and is used to fix photographic pictures (p. 445). It is usually called by photographers and pharmacists " sodium hyposulfite." It has also been used in SODIUM SULFIDE 409 extraction of silver from its ores in the so-called " hyposulfite- lixiviation " processes. The anhydrous thiosulfate, Na 2 S 2 O 3 , can be obtained by dry- ing the crystals at a moderate temperature. If the dry salt is heated, it decomposes into sodium sulfate, sodium sulfide, and sulfur or sodium polysulfide. If a solution of sodium thiosulfate is warmed with copper sulfate, cuprous sulfide, Cu 2 S, and sulfur are precipitated, while sodium sulfate remains in solution. These reactions show that the compound retains the reducing properties of the sulfites and also the properties, in part, of a sulfide. This recalls the method of preparation and agrees well with the formula Na Sv ^O /Sx' , which is assigned to the compound. Na CK ^O Sodium Tetrathionate, Na 2 S4O 6 .H 2 O, is formed by the action of iodine on a solution of sodium thiosulfate : 2 Na 2 S 2 O 3 + I 2 = Na 2 S 2 O 6 + 2 Nal The reaction is quantitative and is much used in volumetric analysis. Sodium Sulfide, Na 2 S, may be prepared by passing hydrogen sulfide into a solution of sodium hydroxide in the requisite amount and evaporating the solution to dryness with exclusion of air. It is hydrolyzed by water, giving a strongly alkaline solution. If a mixture of sodium carbonate and sulfur is heated, or if any metallic sulfide is heated with sodium carbonate on charcoal or if any metallic sulfate is heated with sodium carbonate in the reducing flame (p. 304) on charcoal, a sulfide is formed. If this is moistened with water on a silver coin a black spot of silver sulfide, Ag 2 S, will be formed. The reaction is used as a test for sulfur in any form of inorganic combination. Sodium Hydrosulfide, NaSH, is formed when hydrogen sulfide is passed into a solution of sodium hydroxide in twice the amount necessary to form the sulfide. It loses hydrogen sulfide on evap- 410 A TEXTBOOK OF CHEMISTRY oration of the solution with exclusion of air, in the same way that sodium bicarbonate, NaHCOs, loses carbon dioxide. Sodium Nitrate, NaNO 3 , is found in immense beds in Chile, South America. It has been the chief source from which nitric acid and saltpeter, KNO 3 , have been prepared and large quan- tities have also been used for fertilizers, to furnish the nitrogen necessary for the growth of crops. It is estimated that the supply from Chile will be exhausted in a comparatively few years, but there seems now a good probability that the manu- facture of nitrates from the nitrogen of the air will soon be in a position to supply its place. Sodium nitrate crystallizes without water of crystallization in rhombohedra. It melts at 316. It is hygroscopic and for that reason cannot be used in place of potassium nitrate for the manufacture of ordinary gunpowder, though its low molecular weight makes it, otherwise, more suitable. Sodium Nitrite, NaNO 2 , is prepared by heating sodium nitrate with metallic lead. It is very easily soluble, but crystallizes well. It is used in laboratories and in factories for the prepara- tion of diazonium compounds for the manufacture of dyestuffs and other important compounds. Sodamide, NaNH 2 , is prepared by passing ammonia over me- tallic sodium at 300-350. (See Dennis and Browne, J. Am. Ch. Soc. 26, 587.) It dissolves in liquid ammonia, ionizing to Na + and NH 2 ~ (p. 207). It is hydrolyzed by water to sodium hydroxide and ammonia. It has recently become important for the preparation of indigo. * Sodium Trinitride, NaN 3 , is formed by the action of nitrous oxide, N 2 O, on sodamide (p. 223). Disodium Phosphate, Na 2 HPO 4 .12 H 2 O, is the best known and most important of the phosphates of sodium. It is isomor- phous with the corresponding arsenate, Na 2 HAsQ4-12 H 2 O, and usually contains some of that salt derived from the impure phosphorus or from the sulfuric acid used in the preparation of phosphoric acid. The salt is sometimes used in medicine as a mild cathartic. SODIUM CARBONATE 411 Sodium Carbonate or Sal soda (Washing Soda), Na 2 CO 3 .10 H 2 O. The Leblanc Soda Process. The Leblanc soda process is carried out in three operations. 1. Salt is treated with sulfuric acid on the hearth of a furnace which can be heated to complete the reaction and expel all of the hydrochloric acid. The latter is conveyed through a tower filled with coke over which water is trickling and the aqueous hydrochloric acid obtained is sold or used for the preparation of chlorine. In the early years of manufacture the acid was al- lowed to escape and produced disastrous effects upon vegetation in the neighborhood of the works. This led to stringent legis- lation forbidding the escape of the acid. Later, the recovery of the hydrochloric acid proved profitable, and this has been an important factor in preventing the complete abandonment of the process : 12 NaCl + H 2 S0 4 = Na 2 SO 4 + 2 HC1 2. The sodium sulfate is mixed with charcoal, or coal, and limestone, CaCO 3 , and heated to fusion in the " black ash fur- nace." The mass melts, the sodium sulfate is reduced to sodium sulfide, Na 2 S, and the latter reacts with the calcium carbonate to form calcium sulfide, CaS, and sodium carbonate : Na 2 S0 4 + 2 C = Na 2 S + 2 CO 2 Na 2 S + CaCO 3 = CaS + Na 2 CO 3 3. The " black ash," when cold, is leached with water, which dissolves the sodium carbonate and leaves most of the calcium sulfide undissolved. There is a tendency for the calcium sulfide to hydrolyze to calcium hydrosulfide, Ca(SH) 2 , and calcium hydroxide, Ca(OH) 2 , but this takes place slowly and is repressed in the strongly alkaline solution, so that a fairly complete separa- tion is obtained. The sodium carbonate is obtained from the solution by evaporation and crystallization. If the evaporation and crystallization take place above 35.2, the monohydrate, Na 2 CO 3 .H 2 O, crystallizes from the solution. If the salt is crys- tallized below that temperature, by evaporation or by cooling 412 A TEXTBOOK OF CHEMISTRY the solution, the dekahydrate, Na 2 CO 3 .10 H 2 O, is formed. The temperature given, 35.2, is, of course, the transition point from the dekahydrate to the monohydrate. The Leblanc process, which was so important during the nineteenth century, seems likely to be completely displaced by other processes in the near future. Crystallized sodium carbonate is known as sal soda, or washing soda, and is used for laundry purposes, for the manufacture of soap and for softening water. Anhydrous, or calcined, sodium carbonate is used in the manufacture of glass and for the pre- paration of other compounds of sodium. Sodium carbonate is hydrolyzed by water to sodium bicar- bonate and sodium hydroxide, owing to the very trifling ioniza- tion of the hydrocarbonate ion, HCO 3 ~ : Na + + Na + + CO 3 = + H + + OH~ = Na + + Na + + HCO 3 ~ + OH~ Sodium Bicarbonate, or Baking Soda, NaHCO 3 . The Ammo- nia-soda Process. About 1860 Solvay succeeded in putting this process, which had been discovered many years before, into a successful form. It has been found more economical than the Leblanc process and by the close of the nineteenth century it had very largely displaced that method of manufacture. A strong brine is first treated with ammonia till it contains one molecule for each molecule of salt. A tower about 20 meters high and having a series of shelves is filled with the brine and carbon dioxide is forced into the tower at the bottom. Am- monium bicarbonate, NH 4 HCO 3 , is at first formed, but as sodium bicarbonate, NaHCO 3 , is the least soluble of the various com- binations of ions possible, it separates and is deposited in layers on the shelves : NH 3 + H 2 O + CO 2 = NH 4 HCO 3 NH 4 HC0 3 + NaCl = NH 4 C1 + NaHCO 3 The sodium bicarbonate, after removal from the shelves and washing in centrifugals with a little water, is almost chemically SODIUM SILICATE 413 pure. Sodium carbonate can be easily obtained from this by calcining it at a comparatively low temperature : 2 NaHCO 3 = Na 2 CO 3 + CO 2 + H 2 O A large part of the bicarbonate is converted to carbonate for the market by this process. The carbon dioxide obtained is, of course, returned to use in the first stage of the manufacture. Ammonia is much more valuable than sodium carbonate, at the present time, and the ammonia-soda process depends, economically, upon the recovery of the ammonia. This is ef- fected by treating the solution of ammonium chloride with slaked lime, Ca(OH) 2 , and distilling : 2 NH 4 C1 + Ca(OH) 2 = CaCl 2 -f 2 NH 3 In well-conducted factories not more than 5 kilos of ammonia are lost in preparing 1000 kilos of sodium carbonate. Sometimes magnesium oxide, MgO, is used in place of calcium hydroxide, as it is possible to recover chlorine or hydrochloric ' acid from the magnesium chloride, but the fact that the chlorine is left by the process in the form of a comparatively worthless compound may lead, ultimately, to its abandonment. Sodium bicarbonate is used as " baking soda " in cooking, to furnish carbon dioxide for lightening bread or cake. It is used with sour milk or, frequently, mixed with cream of tartar, tar- taric acid, alum or acid calcium phosphate in the various baking powders. It is also used as a mild alkali in medicine. Sodium Silicate or Soluble Glass, Na 2 SiO 3 , is made by fusing sodium carbonate and sand in the proper proportion. It is hydrolyzed by water and the solution reacts strongly alkaline, but the silicic acid remains in colloidal solution. The solution is used to fireproof wood and fabrics, covering them with a thin, glassy coating, which renders them much less inflammable. Sodium Tetraborate or Borax, Na 2 B4O7.10 H 2 O, has been con- sidered in connection with boron (p. 367). CHAPTER XXIV ALKALI METALS (Continued) : POTASSIUM, AMMONIUM, RUBIDIUM, CAESIUM : THE SPECTROSCOPE Potassium, K, 39.10. Occurrence. Many of the natural sili- cates and especially the potash feldspar, orthoclase, KAlSisOg, contain potassium. These feldspars form an essential constitu- ent of granite and in the disintegration of granites and other rocks through the action of water during very long periods of geological time the sodium and potassium have been partly re- moved and carried away to the ocean. Owing to the selective action of the colloidal silicates remaining in the beds of clay and soils formed by the disintegration of the. rocks, more potassium than sodium has been retained, and the element forms a constitu- ent of very great importance in all arable lands. From the soil, potassium is taken up by all plants in their growth and when vegetable material is burned, the ash is almost invariably alkaline from the presence in it of potassium carbonate. For- merly wood ashes were the most important source of potassium compounds. The potassium carbonate was obtained from the ashes by leaching them with water, and the " lye " prepared in this way was used for the domestic preparation of soft soap. The latter is a concentrated solution of potassium salts of the organic acids of ordinary fats, such as lard or tallow. These salts are deliquescent and cannot be readily brought to a solid form, as is the case with the sodium salts, which form ordinary hard soap. While beds of common salt are found in many different parts of the world and strong brines are quite common, large deposits of potassium chloride (sylvite, KC1) and of magnesium potassium chloride (carnallite, MgCl 2 .KC1.6.H 2 O) have thus far been 414 POTASSIUM 415 found only in Germany, and especially at Stassfurt. These deposits now furnish the larger portion of the potassium com- pounds used in the world, and especially they furnish the potas- sium required to maintain the fertility of the soil for raising tobacco, cotton and other crops. Considerable amounts . of potassium compounds are found in the seaweeds of the Pacific coast, and there is some hope that they may be profitably ex- tracted from that source. It is also possible to prepare potas- sium chloride on a large scale from feldspathic rocks, which are abundant in some parts of the United States. Metallic Potassium was prepared first by Sir Humphry Davy in 1807 at about the same time that he discovered metallic so- dium. It may also be prepared by reducing potassium carbon- ate with carbon and by the electrolysis of fused potassium chlo- ride. Potassium is a silver-white metal which tarnishes instantly in moist air and takes fire when thrown on water. The hydrogen which is liberated burns with the characteristic violet flame of potassium. It has a specific gravity of 0.8621 at 20. It melts at 62.3 and boils at about 760. Potassium Oxide, K 2 O, has been prepared by the partial oxidation of metallic potassium in dry air followed by distilling away the excess of metal. It combines energetically with water to form potassium hydroxide, KOH. Potassium Hydroxide was formerly prepared by treating a solution of potassium carbonate with slaked lime : K 2 CO, + Ca(OH) 2 = 2 KOH + CaCO 3 The reaction depends, of course, on the relative insolubility of calcium carbonate. Potassium hydroxide is now prepared commercially from potassium chloride by electrolysis, with the Castner-Kellner and other forms of apparatus (p. 402). Potassium hydroxide is a white, deliquescent solid. The solu- tion in water has a soapy feel and attacks the skin strongly if allowed to remain in contact with it. The ordinary solid potas- sium hydroxide of the laboratory contains 15 to 20 per cent of 416 A TEXTBOOK OF CHEMISTRY water. Sodium hydroxide, on the contrary, is usually almost anhydrous. Solutions of potassium hydroxide attack glass less than those of sodium hydroxide and do not give a precipitate of the carbon- ate as easily as the latter. For these reasons such solutions are usually employed in organic analysis and often in gas analysis, for the absorption of carbon dioxide. Potassium Chloride, KC1. The occurrence of potassium chlo- ride and of carnal lite in Germany has been mentioned. It crys- tallizes in cubes and melts at about 750. It is easily soluble in water. The solution is often used as a standard for electrical conductivity. The crude salt is extensively employed x in fertilizers and a purer form is used in the manufacture of saltpeter, KNOs. Potassium Chlorate, KClOs, may be prepared by saturating a warm solution of potassium hydroxide with chlorine (p. 127). Practically, milk of lime, Ca(*OH)2, is saturated with chlorine, forming calcium chloride and calcium chlorate, and to this solu- tion potassium chloride is added, causing the separation of potassium chlorate, which is not very easily soluble. The commercial reason for such a procedure is apparent. Potas- sium chlorate is also made by the electrolysis of a solution of potassium chloride under such conditions that the chlorine and potassium hydroxide formed react with each other. The final result may be expressed by the equation : KC1 + 3 H 2 O = KC1O 3 + 3 H 2 Potassium chlorate is used in the preparation of oxygen, in medicine and in the manufacture of matches. Potassium Perchlorate, KC1O 4 , is formed when potassium chlorate is heated slightly above its melting point : 4 KC10 3 = 3 KC10 4 + KC1 It is impossible to avoid some decomposition of the chlorate or perchlorate with evolution of oxygen, but with care a consid- erable portion of the chlorate may be converted into the per- POTASSIUM SALTS 417 chlorate. The latter is much less soluble than the chloride or chlorate and may be purified by crystallization from hot water. Potassium Iodide, KI, may be prepared by dissolving iodine in a solution of potassium hydroxide. The iodate, KIO 3 , formed at the same time, may be decomposed by heating the mixture alone, or, better, with charcoal or some other reducing agent. The commercial salt often contains a little iodate, which is very objectionable for many laboratory uses. The salt is used in medicine and for the preparation of photographic plates. Potassium Polyiodides. Solutions of potassium iodide dis- solve iodine readily, forming unstable polyiodides, the one having the composition KI 3 being probably present in solutions of mod- erate concentration. These solutions dissociate easily into potas- sium iodide and iodine and react in the same manner as free iodine toward reducing agents. For this reason such solutions are used in volumetric analysis. The reaction of such a solution with sodium thiosulfate has been given (p. 187). Potassium Sulfate, K 2 SO 4 , crystallizes without water of crys- tallization. It melts at 1080. It forms double salts with magnesium and calcium, which are found in the potash deposits in Germany and are an important source of potassium com- pounds. Acid Potassium Sulfate, or Potassium Bisulfate, KHSO4, can be prepared by heating a mixture of potassium sulfate and sul- furic acid in molecular proportions. When heated gently it is converted into potassium pyrosulfate, 1^28207, with loss of water. At a higher temperature the latter loses sulfur trioxide and goes back to the normal sulfate. Potassium pyrosulfate is often used in the laboratory as a solvent for aluminium oxide, A^Os, ferric oxide, Fe2Oa, titanium oxide, TiO 2 , and other difficultly soluble substances. Sodium pyrosulfate is, however, more suitable for this purpose. (Hille- brand, Analysis of Silicate and Carbonate Rocks, p. 105.) Potassium Nitrate or Saltpeter, KNO 3 . After the introduc- tion of gunpowder into Europe, about 1300, and especially after 418 A TEXTBOOK OF CHEMISTRY it came into general use in warfare in the sixteenth century, the preparation of pure saltpeter for use in its manufacture became continually more important. Until the nineteenth century saltpeter was obtained almost exclusively from natural sources, where it had been formed by the decay of organic matter con- taining nitrogen and potassium, in the presence of nitrifying bacteria. For a long time considerable supplies of saltpeter have been obtained from India, where it is formed in this way. Calcium nitrate, Ca(NOs)2, which is formed in a similar manner, sometimes occurs as an efflorescence on the walls of stables or in cellars. By interaction with potassium carbonate from wood ashes saltpeter is readily obtained. During the French Revolu- tion saltpeter was often prepared in this manner. During the War of 1812 the United States depended largely on saltpeter from the Mammoth Cave, Kentucky. After the discovery of Chili saltpeter, NaNO 3 , in South Amer- ica and of potassium chloride at Stassfurt, the manufacture of saltpeter from these salts was developed. The preparation de- pends on the fact that sodium chloride is about equally soluble in hot or cold water, while potassium nitrate dissolves in one half of its weight of water at 87, but only 25 parts of the salt will dissolve in 100 parts of water at 15. If potassium chloride is added in molecular proportions to a concentrated, hot solution of sodium nitrate, it will pass into solution, and sodium chloride, the least soluble of the four salts (NaNO 3 , KC1, NaCl, KNO 3 ) present, will separate. From the mother liquors potassium ni- trate will separate on cooling, since that is the least soluble con- stituent in the cool solution. For the manufacture of gunpowder a salt entirely free from chloride must be prepared by recrys- tallization and centrifugal drainage. Potassium nitrate crystallizes in rhombic prisms. It melts at 339. It is used in the manufacture of gunpowder and in the curing of meats, especially of salt beef. It imparts to the meat a desirable reddish color. Taken in considerable quantities it is a poison. Gunpowder is a mixture of about 75 parts of saltpeter, 13 parts POTASSIUM SALTS 419 of charcoal and 12 parts of sulfur. This corresponds very nearly to the equation : 2 KN0 3 + 3 C -f S = K 2 S + N 2 + 3 CO 2 The explosion depends on the fact that the oxygen for burning the carbon is contained in the mixture and on the large volume of the nitrogen and carbon dioxide formed in the reaction. The heat of combustion also raises these gases to a high temperature, increasing the force of the explosion. In the burning of the gunpowder the grains burn from the surface inward, and the speed of the combustion is closely connected with the size of the grains. For use in large ordnance, hexagonal blocks an inch or more in diameter are used to secure slower combustion and allow time for the heavy shot to gain momentum before the full force of the explosive is developed. The failure of guncotton when it was first tried in firearms was partly due to the too rapid burning of the material, which caused the guns in which it was used to burst. This difficulty was finally overcome by giving the " smokeless powder " made from guncotton a dense form, somewhat resembling that of ordinary gunpowder. Potassium Nitrite, KNO 2 , is prepared by the reduction of the nitrate with lead, iron or sometimes with charcoal or sulfur. It is very easily soluble in water and is used in the laboratory as a reagent for cobalt, with which it forms a difficultly soluble com- plex salt, Co(N0 2 ) 3 .3 KN0 2 or K 3 Co(NO 2 ) 6 . The forma- tion of the same salt or of a similar salt containing silver, K 2 AgCo(NO 2 )6, may also be used as a test for detecting potassium. Potassium Carbonate, K 2 CO 3 , was formerly obtained by leach- ing wood ashes. The salt is also obtained in the scouring of wool and from the residual sirups of the beet sugar manufacture after the sugar of the sirups has been converted into alcohol. It is also made from potassium chloride by processes similar to those used in manufacturing sodium carbonate. Potassium carbonate is a deliquescent salt, differing in this respect very markedly from sodium carbonate. The anhydrous 420 A TEXTBOOK OF CHEMISTRY salt melts at about 890. It is used in making soft soap and hard glass. Potassium Bicarbonate or Saleratus, KHCOs, is easily pre- pared by passing carbon dioxide into a concentrated solution of potassium carbonate. It was formerly used in cooking, but has been entirely displaced by the cheaper and more suitable sodium bicarbonate, NaHCO 3 . Potassium bicarbonate dis- solves in about three parts of water, being much more easily soluble than the sodium salt. The solution is nearly neutral to phenolphthalein, but loses carbon dioxide and becomes alka- line on boiling. Potassium Cyanide, KCN, can be prepared by heating potas- sium ferrocyanide : K4FeC 6 N 6 = 4 KCN + Fe + 2 C + &* By heating the salt with metallic sodium a mixture of potas- sium and sodium cyanides is obtained and all of the cyanogen of the original salt can be saved. This mixture may be used for nearly all purposes to which potassium cyanide is applied and especially for the extraction of gold from its ores. Ammonium, NH4, is a group which so closely resembles potas- sium in the salts which it forms with acid radicals that it seems desirable to speak of these salts at this point. It has not been found possible to separate ammonium, NH4, by itself, but if a solution of ammonium chloride is poured on some sodium amalgam the reaction represented by the equation : NH 4 C1 + Na(Hg) = NaCl + NH 4 (Hg) takes place and it can be shown by the electrical properties and by the effect in reducing metals from their salts that a small amount of ammonium amalgam is formed. The substance is, however, extremely unstable and decomposes rapidly into am- monia, NHs, hydrogen and mercury. Ammonium Hydroxide, NH 4 OH. Solutions of ammonia in water seem to involve the following equilibria : NH 3 + H 2 ^ NH 4 OH ; NH 4 + + GET AMMONIUM SALTS 421 The ratio between these various substances has not been deter- mined with any degree of certainty. It is known, however, that the concentration of the hydroxide ions is small in comparison with the concentration in a solution of sodium or potassium hydroxide of equivalent molecular concentration. In other words, ammonium hydroxide is a comparatively weak base. Ammonium Chloride, NH 4 C1, is prepared by neutralizing the aqueous gas liquors, obtained in the manufacture of coal gas or coke, with hydrochloric acid. It is partially purified by subli- mation. When heated to about 350 under ordinary conditions, it is converted into a gas which consists of a mixture of ammonia and hydrochloric acid, the weight of a gram molecular volume of the gas being only about 26.8 grams instead of 53.5 grams, as would be expected from the formula. This abnormal density was at one time used as an argument against Avogadro's hy- pothesis. It has been shown by diffusion experiments, however, that the gas is a mixture, as the ammonia, which is the lighter of the two constituents, diffuses away more quickly than the hydrochloric acid. Finally, many years after the theory of dis- sociation had been universally accepted as the correct explana- tion of the abnormal density, it was shown that very carefully dried ammonium chloride may be converted into a vapor without dissociation, and that a gram molecular volume of this vapor weighs about 53.5 grams. Ammonium Sulfide, (NH^S, is prepared by passing hydrogen sulfide, H 2 S, into a solution of ammonia. The most convenient method is to take a known quantity of a 10 per cent solution of ammonia (sp. gr. 0.96) and pass into it the hydrogen sulfide generated by the action of one fifth of its volume of concentrated sulfuric acid upon an excess of ferrous sulfide, FeS, contained in a bottle holding ten volumes of water for one volume of the acid. The gas should be generated rapidly and well washed with water before entering the ammonia. Ammonium Hydrosulfide, NH 4 SH. If hydrogen sulfide in excess is passed into a solution of ammonia, the hydrosulfide is formed. When a solution of either ammonium sulfide or of 422 A TEXTBOOK OF CHEMISTRY the hydrosulfide is exposed to the air, especially if exposed also to light, the sulfides are oxidized, sulfur separates and am- monium hydroxide is regenerated. The principal action is parallel to that of air on hydrogen sulfide water : H 2 S + O = H 2 O + S (NH 4 ) 2 S + O + H 2 O = 2 NH 4 OH + S The sulfur liberated in this manner will, for a time, dis- solve in the ammonium sulfide, forming poly sulfides, (NH 4 )2S 2 , (NH^A, etc. These polysulfides give to the solution a yellow color. After the oxidation has gone beyond a certain point, the separated sulfur no longer finds any ammonium sulfide with which to combine and begins to separate in the free state. The solution is then no longer fit for use as a reagent. Ammonium sulfide is used in the laboratory to precipitate those metals whose sulfides are too soluble for precipitation in acid solution but sufficiently insoluble for precipitation in the presence of a base. The polysulfide is used to convert stannous sulfide, SnS, into stannic sulfide, SnS 2 , and to dissolve the latter in separating it from lead, bismuth and other metals. Either may be used to dissolve arsenious sulfide, As 2 S 3 , or antimony sulfide, Sb 2 S 3 (p. 261). Ammonium Sulfate, (NH 4 ) 2 SO 4 , is prepared in a crude form by neutralizing the ammoniacal liquors of the gas works with sulfuric acid. As sulfuric acid is the cheapest of the commercial acids, this salt is often prepared to put the ammonia into suit- able form for transportation or for use in fertilizers. Ammonia can, of course, be readily regenerated from it by treatment with lime, CaO. Ammonium Nitrate, NH 4 NO 3 , may be prepared by neutraliz- ing nitric acid with ammonia or ammonium carbonate and evaporating the solution to crystallization. It is easily soluble m water. At 166 it melts and decomposes into nitrous oxide, N 2 O, and water. The decomposition is exothermic, and if the AMMONIUM SALTS 423 . temperature is too high or if a large amount of the salt is heated at once, the reaction may become explosive : NH 4 N0 3 = N 2 O + 2 H 2 + 29,500 cal. The salt is also used as a very important constituent of modern explosives. Ammonium Nitrite, NH 4 NO 2 , is a deliquescent, very unstable salt, which decomposes easily into nitrogen and water. Ammonium Sodium Hydrogen Phosphate, or microcosmic salt, NaNH 4 HPO 4 .4H 2 O, is used in blowpipe analysis to furnish a bead of sodium metaphosphate, NaPO 3 , which, when hot, will dissolve many metallic oxides, giving characteristic colors. The bead does not dissolve silica, SiO 2 . Ammonium Carbonate, (NH 4 ) 2 CO 3 . A salt known commer- cially as " ammonium carbonate "" is prepared by heating a mixture of calcium carbonate and ammonium sulfate. It con- sists of a mixture of ammonium bicarbonate, NH 4 HCO 3 , and ammonium carbamate, NH 4 O C NH 2 . The second of these salts may be prepared, also, by the direct union of carbon dioxide and ammonia. By dissolving the commercial carbonate in water and adding ammonia, the bicarbonate, NH 4 HCO 3 , is changed to the normal carbonate, (NH 4 ) 2 CO 3 . The carbamate is also soon hydrolyzed to the normal carbonate : NH 4 O C NH 2 +H 2 0= NH 4 O C^O NH 4 A solution prepared in this way is used as the ordinary labora- tory reagent. Ammonium Bicarbonate, NH 4 HCO 3 , is prepared by passing carbon dioxide into a solution of ammonia as one of the opera- tions of the ammonia-soda process (p. 412). Ammonium Chloroplatinate, (NH 4 ) 2 PtCl 6 , is a very difficultly soluble salt closely resembling the corresponding potassium salt. Similar compounds are formed from many amines, compounds in which one or more hydrogen atoms of ammonia have been re- placed by organic radicals. 424 A TEXTBOOK OF CHEMISTRY Rubidium, Rb, 85.45, and Caesium, Cs, 132.81. In 1860 and 1861 Bunsen and Kirchoff in applying their newly discovered method of spectrum analysis to the study of a mineral water from Durkheim found some spectral lines which did not corre- spond to those of any known element. In order to obtain enough material to study the compounds of the new elements, 40 tons of the water were evaporated and the compounds of rubidium and caesium were extracted from the residues. Rubidium was named from the Latin word rubidus, meaning red, and caesium from the Latin ccesiw, the blue of the sky, because of the red lines in the spectrum of the former and the blue lines in the spectrum of the latter. Rubidium is found to the amount of about 0.025 per cent in carnallite (KMgCl 3 .6H 2 O) and as a million and a half tons of this mineral are worked over annually for the preparation of potassium compounds it would be possible to obtain very considerable quantities of the element. Metallic rubidium melts at 38.5, caesium at 26.5, the lowest melting point of any metal except mercury. Their compounds resemble those of potassium in their general properties. The chloroplatinates, Rb 2 PtCl 6 , and Cs 2 PtCl 6 , and the alums, RbAl(SO 4 ) 2 .12 H 2 O and CsAl(SO 4 ) 2 .12 H 2 O, are less soluble than the corresponding potassium compounds and are used in sep- arating the elements. A chloroiodide of caesium, CsCl 2 I, is also especially useful for the preparation of pure caesium compounds. Spectrum Analysis. Early in the nineteenth century Fraun- hofer pointed out that when a solar spectrum is produced in such a manner that the colors are sharply separated from each other, the spectrum is crossed by a series of dark lines. During the years from 1820 to 1860 several different observers noticed the characteristic colors imparted to flames by different elements and the spectra of bright lines given by these colored flames and also those given by metals which are vaporized by the electric spark. It was not till 1860, however, that all of these phenomena were brought into clear relationship by the classical researches of Kirchoff and Bunsen, which culminated in the discovery of rubidium and caesium. THE SPECTROSCOPE 425 The simplest form of spectroscope is shown in Fig. 97. The prism A is of glass having a high dispersive power. A narrow slit at B is illuminated by the flame or light which is to be exam- ined. Between the slit and the prism is placed a lens at (7, which renders the rays of light from the slit parallel before they B Fig. 97 reach the prism. The prism is set in such a manner that the angle of incidence on the first face is the same as the angle of emergence from the second face, as this gives the purest spectrum. The light is examined by means of the telescope D. A scale placed at E, illuminated by a light placed before it and whose image is reflected from the surface of the prism, serves to locate the position of the lines. Such a flame as that of acetylene gives a continuous spec- trum, indicating that molecules of solid carbon in the white flame are vibrating in all possible periods required to give white light. If a Bunsen flame is placed before the slit and some com- pound of sodium, as sodium chloride, is introduced, the flame assumes a brilliant yellow color, and with a single prism spec- troscope the spectrum consists of a single, bright yellow line. 426 A TEXTBOOK OF CHEMISTRY A spectroscope having several prisms, or a spectroscope using a metallic mirror (" grating ") ruled with many thousands of equidistant lines and which gives a diffraction spectrum, will separate the line into two lines situated close together. The wave lengths of the lines are 0.5896 and 0.5890 microns, the micron being the thousandth part of a millimeter. The physi- cal significance of these lines seems to be that under the con- ditions of the flame either the sodium atoms as a whole or, more probably, portions of the sodium atoms or electrons within or around them vibrate at a definite rate, which is independent of the temperature. This rate is almost inconceivably rapid. The velocity of light is about 300,000 kilometers per second. This is equal to 3 X 10 14 microns, and since the wave length of the sodium light is only 0.59 micron, the number of vibrations o vx 1Q14 per second must be approximately = 5 X 10 14 per u.oy second. There are two dark lines in the solar spectrum which coincide exactly with the bright yellow lines of the sodium spectrum. This is explained by supposing the interior of the sun to be an incandescent mass which gives out light vibrations of all wave lengths corresponding to the visible spectrum. The photo- sphere of the sun, on the other hand, consists of a gaseous en- velope or atmosphere containing many different elements, among these sodium. The sodium atoms, if they have the power of producing light waves in the ether by their vibrations, must also be able to absorb waves of the same length from the ether, exactly as a tuning fork is set in vibration by sound waves of its own pitch, while waves of a different pitch do not affect it. The sodium atoms in the photosphere, therefore, absorb the waves of their own particular rate ; and while they give the energy ab- sorbed back again to the ether, they dissipate the energy by spreading it in all directions instead of allowing it to pass on toward the observer. The result is that the portion of the spectrum corresponding to the sodium vibrations will be rela- tively dark. By means of this principle it has been possible to THE SPECTROSCOPE 427 show that more than thirty elements found on the earth are found also in the sun. One of these elements (helium), indeed, was discovered in the sun before it was found on the earth. For the purpose of comparing spectra it is convenient to place a right-angled prism before the slit of the spectroscope in such a manner as to cover one half of it. This may be made to reflect the light from a second flame into the slit in such a way that the spectrum from one flame will occupy the upper half of the field of vision while the spectrum from the other flame will occupy the lower half. In this manner the coincidence of lines in the two spectra may be readily observed. In another form, known as the direct vision spectroscope, a series of prisms of different kinds of glass are so combined that one kind of glass counterbalances the mean refractive index of the other, while the dispersive effects . are not counterbalanced. The effect is exactly the reverse of that in an achromatic lens. Such spectroscopes are especially suitable for the detection of the alkali and alkali-earth metals in qualitative analysis. By means of the spectroscope it is possible to detect 3000000 milligram of soqjium. Only the methods used in studying radioactive sub- stances are more sensitive than this. To obtain the spectra of iron, copper and other metals, which are volatile only at high temperatures, electric sparks from a Rumkhorf coil are passed between terminals of the metal, or, in some cases, between platinum wires, one of which is in a small cup containing a solution of a salt of the metal. The spectra of gases are observed in Pliicker tubes, which have a narrow portion through which the electric discharge is passed. CHAPTER XXV THE ALTERNATE METALS OF GROUP I. COPPER, SILVER, GOLD. PHOTOGRAPHY COPPER, silver and gold, which alternate with potassium, rubid- ium and caesium in Group I of the Periodic System, are in almost the greatest possible contrast with those metals. The alkali metals are light. They melt and volatilize at compara- tively low temperatures and they react violently with water at ordinary temperatures. Copper, silver and gold are heavy metals, all three melt between 960 and 1083 and none of them decomposes water, even at high temperatures. They are also the best conductors of electricity that we have. As they do not decompose water, all three of these metals are found free in nature. All three have been known and used since very early times. Copper, Cu, 63.57. Occurrence. Copper is found free in nature, especially in the Lake Superior region. It is found as copper pyrites, or chalcopyrite, CuFeS 2 , a mineral closely resembling iron pyrites in superficial appearance but having a different crystalline form and usually showing blue, red or green colors, owing to superficial changes. Bornite, CuaFeSa, chalcocite, cuprous sulfide, Cu2S, and malachite, /O Cu OH CuCO 3 .Cu(OH) 2 , or C<X a Das i c carbonate, are \O-Cu OH the other most important minerals containing copper. Metallurgy. The most common ores of copper contain either the sulfide, Cu 2 S, or copper pyrites, CuFeS2. If the ore is poor in copper, it is sometimes concentrated by crushing it and washing away a part of the lighter minerals with water. The concentrated ore is then roasted in a furnace with the 428 COPPER 429 addition of sand, if enough silica is not already present. The iron is partly oxidized to ferrous oxide, FeO, which combines with the silica, SiO 2 , to form a fusible silicate or slag, Fe 2 SiO 4 . The cuprous sulfide and some of the ferrous sulfide melt and sink to the bottom of the furnace beneath the slag, which is much lighter. The mixture of sulfides obtained in this way is drawn off and is known as copper matte. Similar operations are now often carried out in a blast furnace somewhat similar to that used in the manufacture of pig iron (p. 543). In the older processes the ore was usually carried through a long series of complicated operations for the purpose of securing a compara- tively pure matte and reducing the latter to metallic copper. In the United States these processes have been very much shortened by the Use of a modification of the Bessemer converter (p. 547) for the reduction of the matte. The molten matte from the roasting furnace is poured directly into the converter, where it is subjected to a blast of air mixed with fine sand or silica. The sulfur of the ferrous sulfide and cuprous sulfide is burned out, the heat of the combustion maintaining the tem- perature of the converter. The ferrous oxide combines with the sand to form a slag of ferrous silicate, while the copper melts and may be cast into plates. The principal reactions may be expressed as follows : f 2FeS + 3O 2 = 2FeO + 2SO 2 1st btage | 2FeO+ SiO 2 = Fe 2 SiO 4 r Cu 2 S + 2O 2 = 2CuO + SO 2 2d Stage 1 2 CuO + Cu 2 S = 4 Cu + SO 2 I or Cu 2 S + O 2 = 2 Cu + SO 2 Electrolytic Refining of Copper. The copper obtained by means of the Bessemer converter or by any of the more com- plicated furnace methods usually contains a small amount of gold and silver and larger quantities of arsenic, lead and other metals which render it unfit for most industrial uses and es- pecially for use in electrical conductors. Nearly all of the copper is now refined electrolytically. The plates of crude 430 A TEXTBOOK OF CHEMISTRY copper are suspended, upright, in a long tank (Fig. 98) filled with a solution of copper sulfate. By connecting the two end plates with the poles of a dynamo the current flowing through the system will cause each plate to become negative on one side and positive on the other. On the positive side the copper will Fig. 98 pass into solution as cupric ion, Cu ++ . Gold, silver, bismuth and some other impurities in the copper fail to dissolve and are collected as " slimes " and treated for the recovery of silver and gold. On the negative side of each plate nearly pure copper will be deposited. When all of the original copper has been dissolved the plates of pure copper are removed and new plates of crude copper put in their place. The production of copper in the United States in 1910 was 500,000 tons, worth $137,000,000. Properties of Copper. Copper is red as ordinarily seen by reflected light. In very thin films it transmits green light. It melts at 1083 (1063 in air, because of oxidation), boils at 2310 and has a density of 8.93. It is the best conductor of electricity of the cheaper metals. Its conductance is very much impaired, however, by the presence of small amounts of other metals. Arsenic is especially harmful, 0.03 per cent lowering the con- ductance by 14 per cent. A table of conductances for the more common metals will be found at the close of this chapter. When exposed to the weather, copper slowly covers itself ALLOYS OF COPPER 431 with a green coating of basic carbonate of the composition of malachite, CuCO 3 .Cu(OH) 2 . When heated in the air a coating of the black oxide, CuO, is formed, which comes off in scales on cooling. In the absence of air, copper is not affected by hydrochloric or dilute sulfuric acid. It dissolves in nitric acid as copper nitrate, Cu(NO 3 )2, with evolution of nitric oxide, NO, nitrous oxide, N 2 O, nitrogen peroxide, NO 2 , or a mixture of these according to the concentration of the acid and the tem- perature. Hot, concentrated sulfuric acid dissolves it as cupric sulfate, CuSO 4 , while sulfur dioxide is evolved. Copper is used for electrical conductors, for the sheathing of ships and for much of the apparatus used in the fermentation industries. Alloys of Copper. Copper is used in a great variety of alloys, the two most important being brass and bronze. Brass con- tains 60 to 70 per cent of copper and 40 to 30 per cent of zinc, with usually small amounts of lead, tin and iron. Bronze is primarily an alloy of copper and tin, usually with 80 to 90 per cent of copper ; it contains, in most cases, small amounts of zinc and lead. Bronze is used for bells and statuary and was for- merly also used for cannon. Phosphor bronze contains from 0.25 to 2.5 per cent of phosphorus, which makes it hard and suitable for bearings in machinery. It has recently been dis- covered that the addition of a very small amount of copper to iron or steel greatly increases its resistance to corrosion. Copper Hydroxide, Cu(OH) 2 , is precipitated as a blue, amor- phous compound when a solution of sodium hydroxide is added to a solution of copper sulfate or any similar salt. If the solu- tion is heated to boiling, the hydroxide is decomposed into the black cupric oxide and water : Cu(OH) 2 = CuO + H 2 O The instability of copper hydroxide is in very marked con- trast with the stability of the hydroxides of the alkali metals, but the hydroxides of silver and mercury are still less stable and have never been prepared as pure compounds. 432 A TEXTBOOK OF CHEMISTRY Cupric Oxide, CuO, may be prepared, as just described, by the decomposition of the hydroxide, or by the decomposition of the nitrate, but neither method will give a pure compound. The copper oxide which is used for organic analysis is usually prepared by calcining copper wire in the air for a long time. Copper oxide ..obtained by heating the nitrate retains nitrogen obstinately and is not suitable for use in the determination of nitrogen in organic compounds. Cuprous Oxide, Cu 2 O, is formed as a bright red precipitate by the action of reducing agents and especially by the action of hydroxylamine, hydrazine, glucose, fructose or lactose on a solution of copper sulfate, sodium hydroxide and potassium sodium tartrate (p. 335). Cupric Chloride, CuCl 2 .2H 2 O, is a bright green salt obtained by dissolving cupric oxide in hydrochloric acid. The double salt, CuCl 2 .2KC1.2 H 2 O, is used to dissolve iron or steel without evolution of gas for the determination of carbon. Copper chloride dissolves in a small amount of water to a green solution similar in color to the crystallized salt. On dilution the color changes to the blue color characteristic of all solutions containing the cupric ion, Cu ++ . If not too dilute, concentrated hydrochloric acid will restore the green color, because the excess of chloride ions causes a decrease in the ionization of the cupric chloride. Cuprous Chloride, Cu 2 Cl 2 , may be easily prepared by digesting a solution of cupric chloride, CuCl 2 , in concentrated hydro- chloric acid with copper turnings : CuCl 2 + Cu = Cu 2 Cl 2 Cuprous chloride is very difficultly soluble in water but dis- solves rather easily in concentrated hydrochloric acid, owing to the formation of chlorocuprous acid, HCu 2 Cl 3 . On dilution this is dissociated into its components and cuprous chloride is precipitated. The insolubility of cuprous chloride, Cu 2 Cl 2 , is parallel to that of silver chloride, AgCl [Ag 2 Cl 2 ], and mercurous chloride, Hg 2 Cl 2 . COPPER SALTS 433 Cuprous chloride dissolves easily in ammonia, also. It is supposed that a complex molecule of the form, [Cu(NH 3 )n]Cl, is formed, but the exact composition of this molecule has not been determined. Solutions of cuprous chloride either in hydrochloric acid or in ammonia absorb carbon monoxide and are used for that purpose in gas analysis. Sodium hydroxide decomposes cuprous chloride with the for- mation of cuprous oxide. Cuprous hydroxide, CuOH, has not been prepared as a definite compound, and there is some reason to suppose that it could exist only at low temperatures. * Cuprous Iodide, Cu 2 I 2 . If potassium iodide is added to an acid solution of a cupric salt, cuprous iodide is precipitated and iodine is liberated : 2 Cu(NO 3 ) 2 + 4 KI = Cu 2 I 2 + 4 KNO 3 + I 2 The reaction is quantitative, and by titrating the iodine with a solution of sodium thiosulfate the amount of copper present may be determined. Cuprous iodide is white and nearly in- soluble in water. Cupric Sulfide, CuS, is formed as a 'black precipitate on pass- ing hydrogen sulfide into an acid solution of a cupric salt. It is very insoluble in water or dilute acids, but dissolves readily in warm nitric acid, the sulfur separating mostly in the free state. On ignition in a current of hydrogen it is converted into cuprous sulfide, Cu 2 S, which has the same composition as the mineral chalcocite. Copper Sulfate, or Blue Vitriol, CuSO 4 .5H 2 O, forms deep blue crystals of the triclinic system. At 120-140 these lose four fifths of their water, leaving the hydrate, CuSO 4 .H 2 O. This /\ ^ has, very probably, the structure Cu< )>S^-OH. At 240 XX \OH the last molecule of water may be expelled, leaving the anhydrous salt, CuSO 4 , as a white powder. Copper sulfate is the most common salt of copper. It is used in the electrolytic refining of copper, in electroplating and electrotyping, as a mordant in 434 A TEXTBOOK OF CHEMISTRY dyeing and in the gravity cells formerly much used for tele- graphic purposes. The anhydrous salt is sometimes used for drying alcohol, as it is very hygroscopic. Vitriols. Vitriol is a very old name given to sulfates because many of them have a glassy appearance. It is rarely used, now, except for copper sulfate, CuSO 4 .5H 2 O, or blue vitriol, ferrous sulfate, FeSO 4 .7 H 2 O, or green vitriol, zinc sulfate, ZnSO 4 .7H 2 O, or white vitriol, and for sulfuric acid, oil of vitriol. The last name is derived from an old method of preparing the acid by distilling green vitriol. * Cupric Nitrate, Cu(NO 3 ) 2 .6H 2 O, is formed in blue tabular crystals when solutions of the salt are crystallized at tempera- tures below 24.5. Above that temperature the hydrate Cu(NO 3 ) 2 .3H 2 O is obtained. On heating moderately the salt loses water, and at a higher temperature it is decomposed into copper oxide, CuO, oxygen and nitrogen dioxide, NO 2 . Ammoniocupric Sulfate, CuSO 4 .H 2 O.4NH 3 . When ammonia is added to a solution of copper sulfate, a very intense blue color is produced. Alcohol precipitates from such a solution ammoniocupric sulfate. This may be considered as the hydrate of copper sulfate in which four molecules of water have been replaced by four molecules of ammonia.. Similar compounds are formed with other salts of copper, and the reaction may be used for the detection of minute quantities of the element. The formula, [Cu(NH 3 ) 4 ]SO4.H 2 O, used by Werner and others brings out more clearly the intimate relation between the ammonia and the copper. *Cuprous Cyanide, Cu 2 (CN) 2 . Solutions of copper salts react with a hot solution of potassium cyanide in very much the same manner as with potassium iodide : 2 CuS0 4 + 4 KCN = Cu 2 (CN) 2 + 2 K 2 SO 4 + C 2 N 2 If the compounds are used in the proportions given in the equation, the cuprous cyanide separates as a white precipitate. If an excess of the potassium cyanide is used, a complex salt, K 3 Cu(CN) 4 , which is easily soluble, is formed. This salt gives ELECTROMOTIVE SERIES 435 a solution containing so few copper ions that no precipitate is formed by hydrogen sulfide. This property is sometimes used to separate copper from cadmium, as the latter is precipitated as cadmium sulfide, CdS, under similar conditions. Precipitation of Copper by Iron, Electromotive Series. If an iron nail is dipped in a solution of copper sulfate or copper chloride, metallic copper will be deposited and iron will pass into solution. We may formulate the reaction thus : Cu ++ + Fe = Fe ++ + Cu The copper ion, Cu ++ , loses its electrical charge, giving it to the iron, while the copper assumes the metallic form. If a strip of copper and one of iron are suspended in any ordinary electrolyte, an electrical current will pass from the copper to the iron on connecting the two metals by means of a wire, while within the electrolyte the positive ions will travel toward the copper and the negative ions toward the iron. As before, the iron will pass into solution and copper will be deposited on the copper strip, if the electrolyte is copper sulfate. On the basis of similar experiments a table of the metals may be arranged in such a manner that each metal in the series will be positive toward all of the metals on one side of it and negative toward those on the other side. The simplest method of looking at these phenomena is to consider that each metal in contact with a solution has a cer- tain solution-pressure which tends to cause the metal to pass into solution. We may suppose that a few atoms do actually leave the piece of metal and pass into solution as ions, but this would give the solution a positive charge and leave the metal negative ; and unless some means is provided for the escape of the electrical charges the electromotive force set up will very quickly balance the solution pressure and the process will cease. If the metal is dipped in a solution of one of its salts, the ions of the metal in solution will partly or wholly counterbalance the solution pressure of the metal, the quantitative effect depending on the concentration of the ions and the nature of the 436 A TEXTBOOK OF CHEMISTRY metal. For many metals this effect may exceed the solution pressure of the metal and the latter will assume a positive poten- tial with reference to the solution through the deposit and dis- charge of some metal ions on the plate. The differences in potential between the various metals and normal solutions of their ions are given in the following table : 1 ABSOLUTE POTENTIAL OP ELEMENTS IN CONTACT WITH NORMAL SOLUTIONS OF THEIR SALTS. ELECTROMOTIVE SERIES ELEMENT ABSOLUTE POTENTIAL Electropositive End Li - 2.740 K - 2.644 Na - 2.434 Ba (-2.6) Sr (-2.6) Ca (-2.4) Mg . . . . - 1.31 Al -1.04? Mn - 0.84 Zn - 0.52 S - 0.31 Fe - 0.19 Cd - 0.16 Te - 0.08 Co.. -0.05 ELEMENT Ni Pb Sn (H As Cu (bivalent) . Bi Cu (univalent) , Sb Hg (univalent) Pd Ag Pt Au F Cl Br I O Electronegative ABSOLUTE POTENTIAL + 0.02 ? + 0.12 + 0.14 + 0.24) + 0.53 ? + 0.58 + 0.63 ? + 0.71 + 0.71 ? + 0.99 + 1.03? + 1.04 + 1.10? + 1.7? (+2.1) (+ 1-59) (+ 1.32) (+ 0.78) (+0.65) End The values in parenthesis have not been measured directly but were calculated from thermochemical data. Elements which assume a high negative potential in contact with solu- tions of their salts are called electropositive because the ions Wilh. Palmaer, Nernst's Festschrift, p. 336 (1907). ELECTROMOTIVE SERIES 437 which separate are strongly positive. Elements which assume a positive potential are called relatively electronegative. This table may be used in two ways : first, any metal of the table may be precipitated by any other metal which has a lower absolute potential and it will precipitate any metal with a higher potential thus copper will precipitate silver, but it will be precipitated by lead or iron ; second, the values may be used to calculate the electromotive force of a galvanic jcell in which two of the metals are used. For an accurate calculation it is necessary to take account of the concentration of the solution in contact with each electrode, and when two solutions are used, the difference in potential between the solutions must also be considered. The common gravity cell (Fig. 99) consists of a copper plate A, in con- tact with a solution of copper sul- fate, and a zinc plate B, in contact with a solution of zinc sulfate. The electromotive force of the battery is approximately the difference be- tween the absolute potentials of copper and zinc, or + 0.58 (0.52) = 1.10 volts. The electromotive force of the Weston cells, consisting of mercury in contact with mercurous sulfate and cadmium in contact with cadmium sulfate, would be for normal solutions of each + 0.99 - (-0.16) = + 1.15. The Clark cell, which has zinc in place of the cadmium, would have for normal solutions an electromotive force of + 0.99 (0.52) .= + 1.51. In both cases, however, the electropositive metal (cadmium or zinc) is in contact with a concentrated or saturated solution of its salt, and the mercurous sulfate is only very slightly soluble. This lessens the difference of potential, and the actual value for the Clark cell is 1.434 Fig. 99 438 A TEXTBOOK OF CHEMISTRY volts, and for the Weston cell, in which a saturated solution of cadmium sulfate is used, it is 1.019 volts. Faraday's Law. If the same electrical current is passed through a series of cells containing electrolytes which are de- composed by the current, it is found that the amounts of the elements which separate at the electrodes will be proportional to the equivalents of the elements, i.e. to the atomic weights of the elements divided by their valences. In a series of cells con- taining solutions of the following compounds : AgNO 3 CuCl CuCl 2 SnCl 2 SnCl 4 H O Ag O Cu Cl Cu Cl Sn Cl Sn Cl Ig8g 108g8g 63.6 g 35.5 g 31.8 g 35.5 g 59 g 35.5 g 29.5 g 35.5 g the same current which causes the liberation of one gram of hydrogen will liberate 8 grams of oxygen, 35.5 grams of chlorine, 108 grams of silver, 63.6 grams of copper from cuprous chloride, 31.8 grams of copper from cupric chloride, 59 grams of tin from stannous chloride and 29.5 grams of tin from stannic chloride. This relation was discovered by Faraday in 1834 and is known as "Faraday's Law." The rational explanation of the law is that each univalent ion is associated with a unit charge of elec- tricity, each bivalent ion with twice this unit charge and a quadrivalent ion with four times this amount. In accordance with the electron theory this unit charge is the charge carried by an electron, and the fact that it does not seem possible to account for Faraday's law without assuming that there is a definite, unit charge, is one of the most important facts sup- porting the theory. It should not be overlooked that the amount of energy re- quired to decompose a gram equivalent of different electrolytes is not the same. The difference in potential between the two electrodes varies from cell to cell ; and the energy consumed in the cell depends on two factors the quantity of electricity passing through the cell and the fall in potential from one elec- trode to the other. The latter depends on the absolute poten- tials of the substances separating at the two electrodes and on SILVER 439 the concentration of the electrolyte exactly as in the galvanic cells discussed in the preceding paragraph. An electrolytic cell may be considered as a galvanic battery in which the direction of the current has been reversed by the application of an external electromotive force. Silver, Ag, 107.88. Copper, silver, gold and all metals which are more electronegative than hydrogen in the electromotive series are found sometimes in the free state in nature. The only other metals found in the free state are iron, cobalt and nickel, and possibly some others which have come to the earth in the form of meteorites and have not had time since their arrival to become completely oxidized. Silver is also found as the sulfide, Ag 2 S, either alone, or more often associated with other sulfides and especially with lead sulfide or galena, PbS. It is also found as the chloride in the mineral cerargyrite, or horn silver, AgCl. Metallurgy. In the electrolytic refining of copper a chloride is added to the electrolyte, and this causes the silver to separate as the insoluble chloride, Ag;Cl, with the slimes from which the silver and gold are recovered. The metallic lead obtained from galena, PbS, always contains some silver. This is recovered either by Pattison's process of crystallization or by Parke's process of extraction with zinc. * Pattinson's Process depends on the principle that a solution melts at a lower temperature than the solvent. If lead contain- ing 'silver is melted in an iron pot, on cooling, crystals of nearly pure lead separate at first, leaving a solution or alloy of silver and lead which is richer in silver than the original metal. The crystals of lead are skimmed out and transferred to another pot on one side while the richer alloy is transferred to a pot on the opposite side. By repeating the operation several times, nearly pure lead is obtained at one end of the series of pots and an alloy comparatively rich in silver at the other end. This rich alloy is then heated in a furnace with free access of air till the lead is oxidized to litharge, PbO, leaving very nearly pure silver behind. While this process is no longer used, it is of 440 A TEXTBOOK OF CHEMISTRY historical interest and also interesting because of the principles of crystallization involved. Cupellation, Assaying. The process of oxidizing lead contain- ing silver is often carried out on a small scale in a muffle furnace (Fig. 100) on a small cup of porous bone ash, called a cupel, which absorbs the litharge. Small amounts of other metals are oxi- dized with the lead and absorbed by the cupel so that an almost pure silver bead remains. This can be weighed as a means of determining the amount of silver in lead bullion. The gold and sil- ver of an ore may be concentrated in a lead button which can then be cupelled. Such determinations, MUFFLE FURNACE Fig. 100 or assays, fur- nish the basis for commercial trans- actions with ores and bullion. Parke's Proc- ess for the ex- traction of silver from lead is exactly analogous to the extraction of a substance from water by means of ether. When zinc and SILVER 441 lead are melted together, on allowing the mixture to stand for a few minutes in the melted condition, nearly all of the zinc will rise to the top in the form of a solution of lead in zinc, containing only a small quantity of lead. The solution of zinc in lead below will contain only a small amount of zinc. The silver and gold, however, are much more soluble in zinc than in lead, and a comparatively small amount of zinc will carry nearly all of the silver to the top with it. If much silver is present, it may be desirable to repeat the process. The alloy of zinc and silver, which is skimmed from the top, may be placed in a retort and the zinc distilled away. The residual lead is then removed by cupellation. * Amalgamation Process. From ores containing little lead or copper, silver was formerly often separated by an amalgamation process. The ore was pulverized and intimately mixed with metallic mercury and water. The mercury dissolved, or amal- gamated with the silver, if that was in the free state. If present as the chloride, the mercury reduced it to the metallic state : AgCl + Hg = Ag + HgCl Mercurous Chloride If the silver was present as the sulfide, it might be reduced by adding iron turnings, or the ore was roasted with salt before it was amalgamated : 2 NaCl + Ag 2 S = Na 2 S + 2 AgCl After the amalgamation was complete the lighter materials of the ore were washed away from the amalgam in a current of water, the amalgam was collected, and the larger portion of the mercury removed by squeezing it through chamois skin, which allows mercury but not the solid amalgam to pass. The latter was heated in a retort to remove the rest of the mercury. Other Processes for the Recovery of Silver. The recovery of silver in the electrolytic refining of copper has been referred to above. The cyanide process (p. 446) has largely displaced all others for obtaining both silver and gold. 442 A TEXTBOOK OF CHEMISTRY The production of silver in the United States in 1910 was 57,000,000 troy ounces, worth $31,000,000. Properties of Silver. Alloys. Silver is a soft, white metal, extremely malleable and ductile. It is the best conductor of electricity known. (See table at the end of this chapter.) It is also a very good conductor of heat, the two properties being closely parallel in most cases, a fact which is readily explained by the electron theory. Silver is too soft for satisfactory use in the free state and is usually alloyed with copper to harden it for the manufacture of coins or of table ware. " Sterling " silver, used in British coins, contains 1\ per cent of copper. The coins of the United States contain 10 per cent of copper. Silver does not tarnish in moist or dry air, but it is easily blackened by hydrogen sulfide or by a solution of a sulfide. Silver dissolves easily in nitric acid or in hot concentrated sulfuric acid, with evolution of nitric oxide, NO, in the former case and of sulfur dioxide in the latter. It is not attacked by hydrochloric acid, partly because of the insolubility of the chloride. Silver Plating. When silver is deposited on a cathode from a solution of silver nitrate, or of some other silver salt which is ionized largely in solution, it assumes a crystalline form and gives a surface which is not suitable for plated ware. If a solu- tion of potassium silver cyanide, KAg(CN) 2 , is used in place of one of the ordinary salts, a smooth, coherent deposit can be obtained. In the electrolysis the object to be plated is made the cathode while a plate of pure silver is used as an anode. The object to be plated requires very careful cleaning, as the slightest film of grease will cause the deposited silver to flake off. Silver Oxide, Ag2O, is formed when a solution of sodium hydroxide is added to a solution of silver nitrate. It may be supposed that silver hydroxide, AgOH, is formed at first, but that this immediately decomposes into silver oxide and water. The decomposition seems to represent an equilibrium, however, which lies far on the side toward the formation of the oxide, SILVER 443 since the oxide dissolves slightly in water and the solution has an alkaline reaction. Also, silver oxide cannot be completely freed from water below the temperature at which it dissociates rapidly into silver and oxygen. The solution of silver hydroxide is ionized to the extent of 70 per cent, indicating that the com- pound is a comparatively strong base. Silver oxide is easily decomposed by heat, but this reaction, also, is reversible : 2 ^0^4 Ag + Q 2 As in all cases of an equilibrium between a gaseous and a solid phase, the equilibrium depends only on the temperature and the pressure of the gaseous component and not upon the relative amounts of the two phases, because the reaction can occur only at the surface between the gas and the .solid, and not throughout the mass of either phase. At 302 the dissociation pressure of silver oxide is 20.5 atmospheres. If the pressure is decreased, more of the oxide will decompose until all is decom- posed or the pressure rises to 20.5 atmospheres again. If the pressure is increased, oxygen will combine with the silver till all is converted into the oxide or till the pressure falls to 20.5 atmospheres. At 325 the dissociation pressure is 32 atmos- pheres ; at 445 it is 207 atmospheres. The dissociation pres- sure of silver oxide is 0.2 of an atmosphere at 121. As 0.2 of an atmosphere is the partial pressure of oxygen in the air, it follows that at temperatures above 121, silver oxide will de- compose completely if exposed to the air, while at temperatures below that metallic silver would change to silver oxide. At that temperature, however, both the oxidation and its decom- position occur very slowly indeed, if no catalyzer is present (Lewis, J. Am. Ch. Soc. 28, 139 (1906)). Molten silver absorbs oxygen, which it gives out again, in part, as the metal solidifies. It seems probable that the absorbed oxygen is not in chemical combination with the silver. * Silver Peroxide, Ag 2 O 2 , is formed as a brown or black coat- ing by the action of ozone on silver. It is also deposited on the anode as a black, crystalline compound in the electrolysis of 444 A TEXTBOOK OF CHEMISTRY silver nitrate with an electrical pressure of about 15 volts and the use of a diaphragm of porous porcelain. Silver Nitrate, AgNOs. This is easily prepared by dissolving silver in nitric acid and evaporating the solution till the salt will crystallize on cooling. It forms tabular crystals of the rhombic system. It is easily soluble in water and is used as the starting point for the preparation of most of the other com- pounds of silver. It melts at 208.6 and is sometimes cast in small sticks for use as a cauterizing and germicidal agent in medicine. In connection with this use an old name, lunar caustic, is still employed a name which comes to us from the alchemists, who recognized a symbolical relation between silver and the moon (Latin, luna). * Silver Nitrite, AgNO 2 , is a white, difficultly soluble salt easily prepared by precipitating a solution of silver nitrate with a moderately concentrated solution of sodium nitrite. It is used in water analyses for the preparation of standard solu- tions of nitrites. * Silver Sulfate, Ag 2 SO4, is prepared by warming silver with concentrated sulfuric acid. The salt is only moderately soluble (1 : 200 in cold water). It is sometimes used to remove chlorine, bromine or iodine from solutions when the presence of a nitrate is to be avoided. Silver Chloride, AgCl, Silver Bromide, AgBr, and Silver Iodide, Agl, are almost insoluble compounds which separate as curdy precipitates on adding one of the halogen acids or a halide to a solution of silver nitrate. The chloride is white ; the bromide, yellowish white ; and the iodide, yellow. All three of the salts are sensitive to light and lose chlorine, bromine or iodine, slowly in diffused light, rapidly in sunlight or in artificial light which contains rays of short wave lengths those of the violet end of the spectrum or beyond. The action seems to be closely related to the effect of sunlight or the magnesium light in causing the combination of hydrogen and chlorine (p. 105). Photography. Photographic " dry " plates are prepared by spreading a thin film of an emulsion of silver bromide in a solu- PHOTOGRAPHY. GOLD 445 tion of gelatin over a plate of glass or of transparent celluloid and allowing the film to dry in a dark room. When such a film has been exposed to the image formed by the lens of a camera no change in the appearance of the film can be noticed, but if the exposed plate is placed in a solution of hydroquinone, pyro- gallol or some other reducing agent used as a " developer/' those portions of the silver bromide which were struck by the light will be reduced to metallic silver, which appears dark and opaque, while the portions under the dark parts of the image of the camera will not be affected. When the lights and shadows have been brought out sufficiently by the developer, the plate is placed in a solution of sodium thiosulfate, Na2S2Os, which dis- solves silver bromide readily, while the metallic silver is not affected. If the unchanged silver bromide were not removed, it would darken later on exposure to light, and the picture would be destroyed. The process of removing the unchanged silver bromide is called " fixing " the picture. The picture obtained is dark where the object photographed was light and light where the object was dark. For this reason it is called a " nega- tive." In order to obtain a " positive," which will reproduce the lights and shadows of the object, the negative is placed over a paper which has been coated with a film of silver chloride or bromide and is exposed to the light for " printing." The light passing through the light portions of -the negative darkens the silver salt beneath, while other portions of the salt are protected by the opaque portions of the negative. The picture on the paper must be fixed by means of sodium thiosulfate as is done with the negative. It may also be " toned " by immersion in a solution of gold trichloride, AuCls, or of chloroplatinic acid, H 2 PtCle, which will cause the silver to be replaced by gold or platinum. The gold gives a reddish tone, the platinum, a steel- gray color. Gold, Au., 197.2, is found almost always in the free state in nature, very rarely in large nuggets weighing several pounds, usually in small grains mixed with sand or gravel or dissemi- 446 A TEXTBOOK OF CHEMISTRY nated through quartzite, granite, pyrite and other rocks and minerals. In Colorado and in some other localities it is some- times found combined with or alloyed with tellurium in a mineral having approximately the composition AuTe 2 . Al- though gold is almost always found mixed with very large quantities of worthless or comparatively worthless minerals, it is very widely diffused and a careful examination reveals traces of gold in almost any rock, soil, clay or other natural, inorganic mixture which is tested. It is claimed that sea water contains from 0.03 to 0.06 gram of gold in a ton (Liversidge, see Chem. Centralblatt, 1905, II, 648). This would be worth from two to four cents, but the total amount of gold in the ocean is, of course, very large. Metallurgy. Gold has a specific gravity of 19.26, while that of ordinary minerals averages about 2.6. Such minerals may, therefore, be separated from the gold by " washing " in a current of water, which carries the lighter minerals away, leaving the gold. For rich sands or gravels the process may sometimes be carried out by hand in a pan an operation which has given the well-known expression " pan out." On a large scale, in " hydraulic mining," masses of sand and gravel are washed away by powerful streams of water, the material running through sluiceways in which are placed crosspieces to retain the gold. Metallic mercury is usually placed back of the cross- pieces, or riffles. This amalgamates with the gold and retains it. Dredging is also used to get gold-bearing material into sluices. Cyanide Process. A few years ago large quantities of gold were obtained by a " chlorination process," in which the gold was dissolved by chlorine obtained from bleaching powder and sulfuric acid. This process has been almost completely replaced by the cyanide process. The cyanide process has also displaced smelting processes in some cases. In the presence of air to furnish oxygen, gold dissolves in a solution of potassium cyanide as potassium aurous cyanide, KAu(CN)2. Metallic silver and silver chloride will also dissolve. GOLD 447 4 Au + 8 KCN + O 2 + 2 H 2 O = 4 KAu(CN) 2 + 4 KOH 4 Ag + 8 KCN + 2 + 2 H 2 O = 4 KAg(CN) 2 + 4 KOH The necessity of air in the solution can be shown by shaking gold leaf with a solution of potassium cyanide which is free from oxygen or through which a current of hydrogen is passed. It will not dissolve, while it dissolves readily on passing a current of air. Potassium permanganate, potassium chromate, sodium peroxide, nitrobenzene or some other oxidizing agent is sometimes used to assist in the solution of the gold. The gold is precipi- tated from the solution by means of zinc : 2 KAu(CN) 2 + Zn = K 2 Zn(CN) 4 + 2 Au The solution containing potassium zinc cyanide may be used for a new lot of ore. Native gold frequently contains silver, and gold is often obtained along with silver from silver or copper ores. From such alloys the silver can be removed by solution in dilute nitric acid or in concentrated sulfuric acid, provided not more than 35 per cent of gold is present. Alloys which are richer in gold than this are melted with enough silver to reduce the amount of gold to one third or one fourth of the whole. The separation with concentrated sulfuric acid can be carried out in cast-iron kettles and is cheaper than that with nitric acid. It also has the advantage that any platinum which is present remains with the gold, while small amounts of platinum dissolve with the silver in nitric acid. The separation of gold from silver by solution of the silver is called, technically, " parting." The production of gold in the United States in 1910 was 4,657,018 troy ounces worth $96,269,100. The coining value of an ounce of gold is $20.67183. The average yearly production of gold in the world during the first half of the nineteenth cen- tury was only about 800,000 troy ounces. The present annual production in the United States alone is nearly six times that and the annual production in the world is many times that of 70 or 80 years ago. 448 A TEXTBOOK OF CHEMISTRY Properties of Gold. Gold is yellow by reflected light but transmits the complementary color, green, through very thin films, as through gold leaf held between two glass plates. It is the most ductile and malleable of all the metals and can be drawn into exceedingly fine wires and beaten into very thin leaves. Its specific gravity is about 19.3, varying considerably with the method of preparation and treatment. It melts at 1063. " Its electrical conductivity is two thirds that of silver. Gold is insoluble in any one of the common acids, alone, but it dissolves readily in aqua regia, a mixture of three volumes of hydrochloric acid with one of nitric. Its solubility in chlorine water has been mentioned above. It dissolves also in selenic acid. Alloys of Gold. Gold is a very soft metal and is alloyed with other metals, especially with silver and copper for use in jewelry and coins. British gold coins are 22 carats fine, i.e. ff pure gold. In reference to gold the term " carat " is used to designate the number of parts in a total of 24 which consist of pure gold. Thus 18-carat gold is yf or J pure gold. American gold coins are made on a decimal basis " 900-fine," i.e. 900 parts in 1000 are pure gold. The American eagle contains 900 parts of gold and 100 parts of copper. * Oxides of Gold. Three oxides of gold have been described : gold monoxide, Au2O; gold dioxide, AuO; and gold trioxide, Au2Oa. There seems to be some doubt whether the second of these has been prepared as a definite compound. Each oxide is decomposed at a comparatively low temperature into gold and oxygen. M * Gold Hydroxide, Au O H, may be prepared by adding magnesium carbonate, MgCOs, to a solution of chloroauric acid and dissolving the excess with dilute nitric acid : HAuCl 4 + 2 MgCO 3 = HAuO 2 + 2 MgCl 2 + 2 CO 2 Gold hydroxide is a yellowish brown precipitate having the properties of both a base and an acid. It dissolves in hydro- ELECTRICAL CONDUCTANCES 449 SPECIFIC CONDUCTANCE AND RESISTANCE OF COMMON METALS* TEMPER- ATURE SPECIFIC CONDUCTANCE SPECIFIC RESISTANCE RESISTANCE IN OHMS OF WIRE 1 M. LONG 1 MM. IN DIAMETER Aluminium . . 35.6 X 10 4 2.81 X 10~ 6 0.036 Bismuth . . . 18 0.82 X 10 4 125. X 10~ 6 1.58 Cadmium . . 14.6 X 10 4 6.85 X 10" 6 0.087 Cobalt . . . 20 10.3 X 10 4 9.7 X 10~ 6 0.123 Copper . . . 25 58.6 X 10 4 1.71 X HT 6 0.022 Gold .... 47.5 X 10 4 2.10 X 10~ 6 0.027 Iron (electrolytic) 8.27 X 10 4 12.1 X 10~ 6 0.154 Iron (steel, 1%C) 18 5.02 X 10 4 19.9 X 10' 6 0.254 Lead .... 5.14 X 10 4 19.4 X 10~ 6 0.247 Magnesium . . 23.0 X 10 4 4.35 X 10' 6 0.055 Mercury . . . 1.063 X 10 4 94.07 X 10~ 6 1.197 Molybdenum 25 17.9 X 10 4 5.6 X HT 6 0.071 Nickel . . . 11.1 X 10 4 9.1 X HT 6 0.116 Palladium . . 9.47 X 10 4 10.6 X 10~ 6 0.135 Platinum . . . 20 10.7 X 10 4 9.3 X 10- 6 0.118 Potassium . . 18 14.9 X 10 4 6.7 X 10- 6 0.086 Silicon . . . 1.68 X 10 4 59.5 X 10~ 6 0.76 Silver .... 25 61.74 X 10 4 1.62 X HT 6 0.020 Sodium . . . 18 20.8 X 10 4 4.8 X 10~ 6 0.061 Tin .... 18 8.3 X 10 4 12.3 X 10~ 6 0.169 Tungsten (fila- ment) . . . 20 17.9 X 10 4 5.6 X 1(T 6 0.071 Zinc .... 16 16.6 X 10 4 6.0 X lO' 6 0.076 1 Specific conductance is the current in amperes which would pass between opposite faces of a cubic centimeter of the metal under a potential difference of one volt. The specific resistance is the reciprocal of the specific conductance. If the factor 10~ 6 is omitted, the specific resistance is given in microhms, i.e. in millionths of an ohm. The resistance of a wire 1 m. long and 1 mm. in diameter is found by multiplying the specific resistance by 10000 X 3.1416 450 A TEXTBOOK OF CHEMISTRY chloric acid with regeneration of chloroauric acid. It also dis- solves in a solution of potassium hydroxide and from the solution potassium aurate, KAuO 2 .3 H 2 O, may be crystallized. Chlorides of Gold. Corresponding to the three oxides there are three chlorides of gold : gold monochloride, AuCl ; gold dichloride, AuCl 2 ; and gold trichloride, AuCls. The last of these combines with hydrochloric acid to form chloroauric acid, HAuCU, and this last compound is formed when gold is dissolved in aqua regia. It is deposited by evaporating and cooling such a solution, in the form of yellow needles having the composition HAuCl4.4 H2O. It is a monobasic acid, which forms well- defined, crystalline salts with many of the metals and especially, also, with many organic bases. The potassium salt, KAuCl 4 .2H 2 O, the calcium salt, Ca(AuCl 4 ) 2 .6 H 2 O, the ammonium salt, 2 NH 4 AuCl4.5 H 2 O, and the strychnine salt, C2iH 22 N 2 O 2 .HAuCl4, are typical of such compounds. EXERCISES 1. Write the equation representing the action of nitric acid on copper ; also for the action of nitric acid on cupric sulfide. Free sulfur and nitric oxide are formed in the latter case. 2. Write the equations for the action of nitric acid and of sulfuric acid on silver. 3. How many cubic centimeters of nitric acid of specific gravity 1.42, containing 70 per cent of the pure acid, will be required to dissolve a pound (453 grams) of silver ? 4. How many grams of concentrated sulfuric acid of 98 per cent will be required to dissolve a pound of silver, taking account only of that which enters into the reaction ? 5. Write the equation for the action of potassium iodide on copper sulfate; also for the reaction between sodium thiosulfate and iodine. How many milligrams of crystallized sodium thiosulfate, Na 2 S 2 O 3 .5 H 2 O, are equivalent to 63.6 milligrams of copper in these reactions ? How many milligrams are equivalent to one milligram of copper ? 6. Write the equation representing the solution of gold in selenic acid, H 2 SeO4. The gold dissolves in the trivalent form, while some of the acid is reduced to selenious acid. CHAPTER XXVI GROUP II, ALKALI-EARTH METALS: BERYLLIUM, CALCIUM, STRONTIUM, BARIUM, RADIUM THE elements calcium, strontium and barium are called alkali- earth metals, a somewhat indefinite designation coming down to us from the time of the alchemists and referring to the fact that the hydroxides of these metals are strong bases. The elements beryllium and radium are not included under the designation. All of the elements of Group II are bivalent. Calcium, stron- tium, and barium decompose water at ordinary temperature, though much less rapidly than the alkali metals. Calcium hydroxide is difficultly soluble in water. Strontium and barium hydroxides are somewhat more soluble. All three carbonates are nearly insoluble in water, but dissolve as bicarbonates in water containing carbon dioxide. Barium sulfate is one of the most insoluble salts, while calcium sulfate is slightly soluble (1 part in 500 of water). * Beryllium, 1 Be, 9.1. The mineral beryl is a silicate of beryl- lium and aluminium, having the composition BesA^SieOig (or 3 BeO.Al 203.6 SiCy. As a precious stone, different forms of the mineral are known as emerald and aquamarine. The free element is silver-white and has a specific gravity of 1.9. Both as an element and in its compounds beryllium resembles mag- nesium and aluminium rather than calcium. The hydroxide, Be (OH) 2 , is nearly insoluble in water but dissolves both in acids and in alkalies, exhibiting in this way both acid and basic properties. The chloride, BeCl 2 .4 H 2 O, sulfate, BeSO 4 .4 H 2 O, 1 Sometimes called glucinum, Gl, a name which has some claim to priority. The name beryllium has the advantage of referring to the most important mineral containing the element. 451 452 A TEXTBOOK OF CHEMISTRY and nitrate, Be(NO 3 ) 2 .3 H 2 O, are soluble in water. The car- bonate is unstable and readily loses carbon dioxide, forming a basic carbonate. Calcium, Ca, 40.07. Occurrence. Calcium is one of the most abundant and most important of the elements. It is found as the carbonate, CaCO 3 , in the minerals calcite and aragonite, in marble and, less pure, in limestones ; as the sulfate in the mineral gypsum, CaSO4.2H 2 O; as the phosphate, Ca 3 (PO4) 2 , in bone ash and mineral phosphates and in the mineral apatite, Ca5(PO4)3F ; and as the fluoride in the mineral fluorite, CaF2. Preparation, Properties. Metallic calcium can be prepared most easily by the electrolysis of the fused chloride (E. F. Smith and Goodwin, J. Am. Chem. Soc. 25, 873 (1903)). It is a white, crystalline metal which decomposes water readily at ordinary temperatures, with the formation of the hydroxide, Ca(OH)2. It burns in air, forming both the oxide, CaO, and the nitride, Ca 3 N 2 . The latter is hydrolyzed by water with the formation of ammonia : Ca 3 N 2 + 6 HOH = 3 Ca(OH) 2 + 2 NH 3 Calcium melts at 800. * Calcium Hydride, CaH 2 , is a white, crystalline compound which can be prepared by the direct union of the elements. Calcium Oxide, CaO, or Lime is manufactured on a large scale by heating the carbonate in lime kilns. In the older forms a mixture of limestone and fuel was placed in the kiln. The fuel was then set on fire and allowed to burn till it was all con- sumed. In the newer forms a mixture of limestone and coal is charged into the kiln at the top, while the " burnt " lime is removed at the bottom, without stopping the process. The dissociation of the carbonate : CaCO 3 ^ CaO + CO 2 . is a reversible reaction. As there is a gaseous constituent, the system has a characteristic pressure for each temperature. This may be determined by heating calcium carbonate in a plat- inum bulb connected with a manometer. The pressure of the ALKALI-EARTH METALS: CALCIUM 453 carbon dioxide which will be in equilibrium with a mixture of calcium oxide, CaO, and calcium carbonate is as follows : Temperature 600 700 800 898 950 Pressure of CO 2 2.35 mm. 25.3 mm. 168 mm. 760 mm. 1490 mm. It will be seen from the table that the temperature of the lime kiln must be above 900 for the carbon dioxide to escape rapidly from the interior of a piece of limestone. When calcium car- bonate is heated in an open crucible in the laboratory under conditions such that the carbon dioxide is constantly displaced by air, the partial pressure of the carbon dioxide may become very low and complete decomposition could be secured at a tempera- ture of 750 or below. In the transformation to lime, pieces of limestone retain their shape but shrink somewhat in size. If water is added to the lime, it combines with it, evolving a very considerable amount of heat 15,540 small calories per gram molecule. At the same time the calcium hydroxide, Ca(OH) 2 , which is formed, swells and falls to a loose powder. The process is called " slaking." If lime is exposed to the air, it slowly absorbs water and falls to a powder, and the hydroxide also absorbs carbon dioxide and is converted back to the carbonate. Because of the latter change, " air-slaked lime " is usually worthless for the preparation of mortar. * Dissociation of Calcium Carbonate and the Phase Rule. In the system produced by the partial dissociation of calcium carbonate there are three phases (CaCO 3 , CaO and C(>2, and only two components (CaO and CO 2 ), the calcium carbonate formed by the union of the other two not being considered as a separate component. According to the phase rule (p. 107) a system hav- ing two components and three phases has only one degree of free- dom and is univariant, just as the system water vapor water has one component and two phases and is also univariant. Accordingly, in the system calcium carbonate calcium oxide carbon dioxide, if the temperature changes, the pressure must change also ; and for every temperature there is a fixed 454 A TEXTBOOK OF CHEMISTRY dissociation pressure, just as for every temperature there is a fixed vapor pressure for the system water vapor water. Mortar is prepared by mixing slaked lime with water and sharp sand, which has not been rounded by long action of waves. Where used between bricks the moisture is partly absorbed by the bricks and partly dries out in the air. This is followed by the action of the carbon dioxide of the air, which slowly changes the hydroxide to calcium carbonate : Ca(OH) 2 + CO 2 = CaCO 3 + H 2 O The carbonate crystallizes as it forms and adheres strongly to the particles of sand, binding them together. In plastered rooms the liberation of moisture by the reaction keeps the air of the rooms moist for some days or weeks. Cement. For the manufacture of cement a clayey limestone is sometimes used, but, as limestones having the proper composi- tion are rare, artificial mixtures of finely powdered limestone and a fine clay or shale, rich in silica, are usually employed. The slag from blast furnaces is also extensively used. The materials are finely ground and fed into the top of a long, slanting, slowly rotating, tubular furnace which is heated by means of powdered coal blown in at the lower end. A temperature of 1500-! 600 is reached in the hottest parts of the furnace. The carbon dioxide is completely expelled and a " clinker " composed of calcium silicate and calcium aluminate with an excess of calcium oxide is formed. This clinker is finely ground and mixed with a small amount of plaster of Paris. It should contain only a very small per cent of magnesium. The composition of the finished cement is as follows : Loss on ignition 0-2 per cent Silica, SiO 2 15-20 per cent Alumina, A1 2 O 3 3-8 per cent Ferric oxide, Fe 2 Os 3-6 per cent Lime, CaO 58-64 per cent Magnesium, MgO 0-4 per cent Potash and soda, K 2 O, Na 2 O 0-2 per cent Sulfur trioxide, SOs 0-2 per cent CALCIUM CHLORIDE 455 When the cement is mixed with water it slowly combines with it, forming partly crystals of calcium hydroxide, Ca(OH) 2 , partly, probably, hydrated silicates and aluminates of calcium which " set " to a firm, strong mass. Sand and sometimes other materials are usually added to increase the volume. As the set- ting results from the action of water alone, it will take place under water and hence the material is often called hydraulic cement. Calcium Chloride, CaCl2, may be prepared by dissolving cal- cium carbonate in hydrochloric acid, evaporating the solution and drying the residue, finally at a temperature of 260 or above. Prepared in this manner it forms a porous, extremely hygro- scopic and deliquescent mass, which is much used for drying gases. The anhydrous salt melts at about 800. With water it forms a series of hydrates. The one which is stable at ordinary temperatures and which may be crystallized by cooling very concentrated solutions is CaCl2.6 H 2 O. Considerable heat is evolved when the anhydrous chloride dissolves in water, but a mixture of the hydrate with a little less than its weight of snow will give a lowering of the temperature to 55. Even the hydrate, CaCl 2 .H 2 O, has an appreciable vapor pres- sure at ordinary temperatures, and gases cannot be fully dried by its use. At 15 it leaves 1.0 milligram and at 30 it leaves 3.3 milligrams of water in liter of a gas which has been passed over it. Calcium chloride is obtained in large quantities as a by-product in the ammonia-soda process (p. 412). Because of the low freezing point, solutions of calcium chloride are used in refrigerat- ing machines to surround the cans in which water is frozen to artificial ice. They have also been used for sprinkling roads. /OC1 Chloride of Lime," Ca^ , is prepared commercially by X C1 spreading slaked lime on the floor of a room and filling the room with chlorine gas : OC1 Ca(OH) 2 + C1 2 = Ca/ ' + H 2 X C1 456 A TEXTBOOK OF CHEMISTRY Treatment with a dilute acid liberates hydrochloric and hypo- chlorous acids and these, in turn, react to form free chlorine (p. 126). The compound is used for bleaching (p. 125) and for sterilizing water (p. 83). Good bleaching powder should contain 35-37 per cent of " available " chlorine. It deteriorates, partly by loss of hypochlorous acid through the action of the carbon dioxide of the air, partly by a slow transformation of the hypochlorite to the chlorate. * Calcium Chlorate, Ca(ClOs)2, is prepared by passing chlorine into " milk of lime," a mixture of calcium hydroxide and water. The hypochlorite formed at first changes to chlorate by auto- oxidation. The solution is used to mix with a solution of potas- sium chloride, KC1, for the preparation of potassium chlorate, KC1O 3 . (Why?) Calcium Fluoride, CaF 2 , is found as the mineral fluorite, which crystallizes in cubes or octahedra. It is nearly insoluble in water, and for this reason natural waters never contain more than minute traces of fluorine. Calcium fluoride melts at 1330. It is used as a flux in casting iron and in other metallurgical opera- tions and is the source from which hydrofluoric acid and all other compounds of fluorine are obtained. Calcium Sulfide, CaS, is formed by the reduction of calcium sulfate, CaSO 4 , by heating it with coal or charcoal in the Leblanc Soda Process (p. 411). It does not seem to be appreciably soluble in water, but is slowly hydrolyzed to the hydrosulfide, Ca(SH) 2 , which dissolves : 2 CaS + 2 HOH = Ca(SH) 2 + Ca(OH) 2 The hydrolysis is sufficiently slow in an alkaline solution so that the sodium carbonate of the Leblanc process may be dis- solved, leaving nearly all of the calcium sulfide as an insoluble residue. In the earlier manufacture the calcium sulfide was discarded as a waste product, but the slow hydrolysis of the material on exposure to the weather led to the contamination of streams in England and resulted in stringent legislation requiring the PLASTER OF PARIS 457 manufacturers to take care of their waste products in such a manner that they should not become a nuisance to others. This and the competition of the ammonia-soda process, com- pelling the manufacturer to practice every possible economy, led to the invention of the Chance process for the recovery of the sulfur. By exposing the moist calcium sulfide to the action of carbon dioxide, the calcium is converted to the carbonate and hydrogen sulfide is liberated : Ca(SH) 2 + H 2 C0 3 ^ CaC0 3 + 2 H 2 S The fact that the ionization constant of carbonic acid is con- siderably greater than that of hydrogen sulfide and also the insolubility of the calcium carbonate, both aid in shifting the equilibrium of the reaction to the right. By burning the hydrogen sulfide with a limited supply of air it is possible to convert the sulfur almost completely to the free state. * Acid Calcium Sulfite, CaH 2 (SO 3 )2, is prepared by burning sulfur and passing the sulfur dioxide, SO 2 , formed into milk of lime. The solution is used to dissolve and remove lignin from the fiber in the manufacture of paper from wood. Calcium Sulfate, CaSO 4 .2 H 2 O, Plaster of Paris. The min- eral gypsum, which has the composition CaSO 4 .2 H 2 O, is found in nature in clear, transparent, monoclinic crystals called selenite and in a white, opaque form called alabaster, also in large quan- tities in a form suitable for making plaster of Paris. Calcium sulfate is also sometimes found in nature in the anhydrous form in the mineral anhydrite. If gypsum is treated for some time at 130-160, it loses three fourths of the water which it contains and is converted into a compound having the composition 2 CaSO 4 .H 2 O and known commercially as plaster of Paris. When this is mixed with a small amount of water to a creamy consistency, it can be filled into a mold ; but after standing a short time part of the water unites chemically with the salt to form gypsum, CaSO 4 .2 H 2 O, which " sets " to a solid mass. If the plaster is heated to too high a temperature or for too long a 458 A TEXTBOOK OF CHEMISTRY time in driving out the water, it becomes " dead burnt " and will then combine with water only very slowly, and it is worthless for the ordinary uses of plaster of Paris. It seems probable that ordinary plaster of Paris retains a small amount of the unchanged dihydrate and that the molecules of this furnish the starting point for the crystallization in setting. Calcium sulfate is difficultly soluble, about two grams dissolving in a liter of water. It is much less soluble in alcohol. * Plaster of Paris and the Phase Rule. In the preceding paragraph two hydrates of calcium sulfate, CaSO4.2 H 2 O and 2 CaSO 4 .H 2 O, have been mentioned and also a natural anhy- dride, the mineral anhydrite. Another anhydride, which is more easily soluble than the natural anhydrite, is also known and is called the " soluble anhydrite." In speaking of the vapor pressure of hydrates (p. 82) it has been implied that each hydrate has a characteristic vapor pressure. This is strictly true only in case the loss of water leads to the formation of only one compound, either a lower hydrate or an anhydride. The majority of salts lose water in this manner; but gypsum, CaSO4.2 H2O, may lose water in such a manner as to form either of the three substances, natural anhydrite, soluble anhydrite or plaster of Paris, 2 CaSO 4 .H 2 O. The vapor pressure of the gypsum will depend on which of the three substances is formed, as is seen in the table on the opposite page. If these values are plotted and the vapor pressure curves extended, it is found (Fig. 101) that the curve for the system CaSO 4 .2 H 2 O, CaSO 4 (natural anhydrite) cuts the curve for water vapor at 66 and the other two curves cut it at 89 and 107. As there are only two components (H 2 O and CaSO 4 ), each of these temperatures represents a quadruple point (p. 78) where the sysem is invariant. The four phases at 66 are: vapor, solution, CaSO 4 .2 H 2 O and CaSO 4 (natural anhydrite). Any change in temperature or in pressure will cause the disap- pearance of one of the phases. If the pressure is decreased, water will evaporate till only the natural anhydrite is left. If the pressure is increased, the vapor phase will disappear. If the CALCIUM SULFATE: PHASE RULE 459 VAPOR PRESSURE IN MILLIMETERS OF MERCURY FOR SYSTEMS CONTAINING GYPSUM t PURE WATER SYSTEM CaSO4.2H z O AND NATURAL ANHYDRITE SYSTEM CaSO 4 .2 H 2 O AND SOLUBLE ANHYDRITE SYSTEM CaSO4.2 H 2 O AND PLASTER OF PARIS 2 CaS04.H 2 O 15 12.7 8.43 7 4.21 20 17.4 12.2 10.7 6.24 30 31.5 24 19.4 12.7 40 54.9 45.4 34 26.3 50 149 143 108 91.4 65 187 140 122 70 233 185 161 80 355 314 272 90 526 446 100 760 711 105 906 888 110 1075 1000 m m 900, 800 700 600 500 400 300 60 70 80 90 100 Fig. 101 110 460 A TEXTBOOK OF CHEMISTRY temperature is increased, all of the gypsum, CaSC>4.2 H 2 O, will be changed to anhydrite because the vapor pressure of gypsum would be greater than that of anhydrite above this temperature. The change might, it is true, take a long time, but there would be no stable equilibrium till the change was com- plete. If the temperature was lowered, all of the anhydrite would be converted into gypsum, since the vapor pressure of water is greater than that of gypsum at temperatures below 66. A consideration of the curves will enable one to predict whether gypsum or an anhydrite will crystallize from a salt solution, if we know the vapor pressure of the latter. Thus a solution of sodium chloride has a vapor pressure greater than 12.2 mm. at 20. From such a solution gypsum will crystallize. If mag- nesium chloride or calcium chloride is added till the vapor pres- sure is less than 12.2 mm. at 20, anhydrite will crystallize from the solution, because if gypsum and such a solution were placed side by side in a confined space water would escape from the gypsum and condense in the solution. These predictions of the theory agree with facts observed by geologists about the condi- tions under which gypsum and anhydrite are found in nature. * Calcium Nitrate, Ca(NO 3 )2-4 H 2 O, is readily prepared by dissolving the carbonate in nitric acid. It is so manufactured for fertilizers and other uses on a considerable scale in Norway, by absorbing the oxides of nitrogen, formed by the electric arc from air, in milk of lime. The anhydrous salt is sometimes used to dry oxides of nitrogen or other gases for which calcium chloride cannot well be used. Calcium Phosphates. Normal calcium phosphate, Ca 3 (PO 4 )2, is an important constituent of bones. It is also found as a min- eral phosphate in deposits in the southeastern part of the United States from North Carolina and Tennessee to Florida. Both bones and the mineral phosphates are extensively used as fertilizers, phosphorus being an element which is essential for the growth of crops and which is found in only limited amounts in some soils. In order to render the phosphate more easily soluble and available for the growth of plants, the powdered CALCIUM PHOSPHATES 461 mineral phosphate is often treated with sulfuric acid to convert it into monocalcium phosphate, CaH 4 (PO 4 ) 2 . The mixture of calcium sulfate and acid calcium phosphate is designated, com- mercially, as a " superphosphate." A slightly diluted acid is used and this "superphosphate " contains both salts as hydrates : Ca 3 (PO 4 ) 2 + 2 H 2 SO 4 + 6 H 2 O = Ca(H 2 PO 4 ) 2 .2 H 2 O + 2 (CaSO 4 .2 H 2 O) * Solubility of Calcium Phosphates. Even the monocalcium phosphate, Ca(H2PO 4 ) 2 , is only slightly soluble, about 5 grams dissolving in a liter of water. If a larger amount of the salt is added, the reaction : 2 H 2 PO 4 - ^ HPO 4 ~ + H 3 PO 4 causes the number of monohydrophosphate ions, HPO 4 , and of calcium ions in the solution to exceed the solubility product for the reaction, and there will be formed a precipitate of the very difficultly soluble dicalcium phosphate, CaHPO 4 , while the solution will contain more phosphoric acid than will correspond to the mono- calcium phosphate, Ca(H 2 PO 4 ) 2 . In the valuation of fertilizers it is customary to distinguish three forms of phosphoric acid, " water-soluble phosphoric acid," " citrate-soluble phosphoric acid " (phosphoric acid insoluble in water but soluble in a neutral solution of ammonium citrate of sp. gr. 1.09) and " insoluble phosphoric acid." The last is supposed to be in the form of tricalcium phosphate, Caa(PO 4 ) 2 . From what has been said above it is evident that more phosphoric acid will pass into solution if a dicalcium phosphate is treated with successive small portions of water than if it is treated at once with a large quan- tity of water. The amount of the " citrate soluble phosphoric acid " also depends on the exact conditions of the determination. Both determinations are to be considered as in a considerable measure conventional, and it is quite certain that they do not furnish an accurate measure of the availability of the phosphorus 462 A TEXTBOOK OF CHEMISTRY for the growth of plants. Finely ground phosphate rock also furnishes phosphorus which can be slowly absorbed by plants, and it is at least a question whether the acid phosphate is any better than the raw ground phosphate for maintaining the fertility of land during a series of years. The addition of even a weak acid, such as acetic acid, to cal- cium phosphate will cause it to pass into solution. This is because phosphoric acid is so weak an acid that in the presence of even comparatively few hydrogen ions, H + , the phosphoric acid must be almost completely either in the form of the un-ion- ized acid, H 3 PO4, or of the dihydrogen phosphate ions, H 2 PO4-. Under these conditions there can not be enough of the mono- hydrogen phosphate ions present to reach the solubility product : Ca ++ X HPO 4 ~ = K which would cause the precipitation of the dicalcium phosphate. As has been explained above, it is this precipitation which causes the apparent difficult solubility of monocalcium phos- phate. Calcium Carbide, CaC2. By heating a mixture of lime, CaO, and coke in a revolving electric furnace, calcium carbide is formed : 3 C = CaC ; 2 + 2 CO The carbide is easily hydrolyzed by water with the formation of calcium hydroxide and acetylene, C 2 H 2 (p. 293), and is manu- factured chiefly for that use. Calcium Cyanamide, CaCN 2 . By heating calcium carbide in a current of nitrogen at 1000 it is transformed into calcium cyanamide, Ca=N C=N : CaC 2 + 2 N = CaCN 2 + C Calcium cyanamide is hydrolyzed by water to ammonia and calcium carbonate : CaCN 2 + 3 H 2 = CaCO 3 + 2 NH 3 As these reactions furnish a means of transforming the nitrogen of the air into a form which is available for plant growth, calcium. HARD WATERS 463 cyanamide is now manufactured in considerable quantities for use as a fertilizer. For this use it is often called " lime-nitro- gen," or in German " Kalk-Stickstoff." Calcium Carbonate, CaCO 3 , is the most abundant compound of calcium. The various forms have been already mentioned, also its conduct when heated. Hard Waters. One liter of pure water at ordinary tempera- tures will dissolve only about 12 milligrams of calcium carbonate, CaCOs, but water saturated with carbon dioxide will dissolve nearly 100 times as much, or more than a gram in one liter. The solution contains the acid carbonate, Ca(HCO3)2, frequently called the bicarbonate. As natural waters always contain some carbon dioxide absorbed from the air and usually acquire much more from decaying vegetable matter in the soil, all such waters which have come in contact with a soil containing calcium car- bonate hold more or less of the calcium bicarbonate in solution. The properties of such a solution can be easily illustrated by passing carbon dioxide through a solution of limewater, slightly diluted. Calcium carbonate will be precipitated at first, but on continuing the current of the gas it will pass again into solution. On boiling the solution the acid carbonate dissociates into calcium carbonate, carbon dioxide and water. The carbon dioxide escapes with the steam and the calcium carbonate is precipitated. Natural waters of this type are said to have " temporary hardness," since the hardness is nearly all removed by boiling. The designation " hardness " refers to the effect of hard water in precipitating an insoluble calcium salt when soap is used with it. The water continues to have a harsh feel- ing to the skin until enough soap has been used to complete the precipitation of the calcium and magnesium salts which are present. Waters containing calcium sulfate in solution will not deposit the salt on short boiling and are said to have " permanent hardness." Such waters deposit the sulfate on concentration of the water, as is done in a steam boiler, and the decreased solubility of the calcium sulfate at a high temperature increases 464 A TEXTBOOK OF CHEMISTRY the amount of scale formed. The scale from such a water is especially coherent and troublesome. If milk of lime, Ca(OH)2, is added to water containing calcium bicarbonate in just the right proportion, nearly all of the calcium carbonate in the water and also the calcium of the milk of lime which is added will be precipitated. This is known as Clark's process of softening water : Ca(HCO 3 ) 2 + Ca(OH) 2 = 2 CaCO 3 + 2 H 2 O To remove the calcium of calcium sulfate, sodium carbonate, trisodium phosphate, sodium fluoride or some other salt which will precipitate the calcium must be used. * Determination of Free and Combined Carbonic Acid in Natural Waters. It has been shown (p. 389) that phenol- phthalein is a suitable indicator for weak acids ; and methyl red, for weak bases. If the acid is very weak indeed, the hydrolysis of the normal salt may cause the end point with phenolphthalein to appear before the acid is fully neutralized. In the case of carbonic acid, H 2 COs, when the point correspond- ing to the formation of NaHCOs is passed, the hydrolysis of the sodium carbonate, Na + + Na + + CO 3 = + H + = Na + + Na + + HC(V + OH~ causes the concentration of hydroxide ions to exceed the end point for phenolphthalein. The amount of alkali required to give a pink color with phenolphthalein will, therefore, indicate the amount of free carbonic acid, H 2 COs, in the solution. On the other hand, if a strong acid is added to a solution of a carbonate or bicarbonate, so long as any of either remains in solution the concentration of the hydrogen ions cannot exceed the concentration in a solution of carbonic acid, H 2 CO3. When enough acid has been added to decompose all of the carbonates and bicarbonates present, any further addition of acid will carry the concentration of the hydrogen ions past the end point for methyl orange or methyl red. This makes it possible to CALCIUM OXALATE 465 determine the amount of free carbonic acid and also of carbonates and bicarbonates present, by titrating first with alkali, using phenolphthalein as an indicator and then with an acid, using methyl orange or methyl red. * Calcium Acetate, Ca(C 2X1302)2, is prepared commercially by neutralizing the distillate obtained by the destructive distilla- tion of wood. It is used for the manufacture of glacial acetic acid, HC 2 H 3 O 2 , and of acetone, CH 3 COCH 3 . Calcium Oxalate, CaC2O4.H2O. When a solution of ammo- nium oxalate, (NH4)2C2O4, is added to a solution containing a soluble salt of calcium, calcium oxalate is precipitated as a fine, crystalline powder. One liter of water dissolves only 5.6 milligrams of the salt, and it is still less soluble in a solution con- taining ammonium oxalate. For this reason it is often used for the quantitative determination of calcium. Its value for this purpose is greatly enhanced by the fact that magnesium oxalate, MgC 2 O4, is much more easily soluble 300 milligrams in one liter of water. In strong acids, as hydrochloric or nitric acid, calcium oxalate dissolves. Oxalic acid is only moderately ionized in solutions of medium concentration 50 per cent to H + and HC 2 O 4 ~ in tenth normal solution. The ionization to H + , H + and 204 must be very much less. As the presence of the hydrogen ions of a highly ionized acid, such as hydrochloric acid, shifts the equilibrium of the reaction : H 2 C 2 4 ^ H + + H + + C 2 4 - to the left, the concentration of the oxalate ions, C 2 O4 , in such a solution cannot be great enough for the solubility product C 2 O 4 X Ca ++ to reach the point of precipitation for calcium oxalate. The salt will, therefore, dissolve in such a solution. In a solution of a weak acid such as acetic acid, HC 2 H 3 O 2 , the number of hydrogen ions is so small that they produce only a slight effect on the ionization of oxalic acid, which is a very much stronger acid than acetic acid. Calcium oxalate may be precipitated, for this reason, from solutions containing acetic 466 A TEXTBOOK OF CHEMISTRY acid, though the salt is more soluble in dilute acetic acid than in pure water. Owing to the extremely low ionization constant f or : H 2 PO 4 -+ ^ H + + HPO 4 " calcium phosphate is not precipitated from solutions containing even the very weak acid, acetic acid. The addition of ammo- nium oxalate to a solution of calcium phosphate in dilute acetic acid will, accordingly, cause the precipitation of nearly all of the calcium in the form of the oxalate, and even a solution of calcium sulfate will precipitate oxalic acid, but not phosphoric acid, from a solution of an oxalate containing acetic acid. Calcium oxalate loses carbon monoxide and is converted into calcium carbonate by gentle ignition : CaC 2 O 4 = CaCO 3 + CO Calcium Silicate, CaSiO 3 , is found in nature as the mineral wollastonite. It may be prepared by fusing a mixture of quartz or sand with calcium carbonate : CaCO 3 + Si0 2 = CaSiO 3 + CO 2 Calcium silicate is a constituent of a very large proportion of the most common natural silicates, such as pyroxene, amphibole, garnet and the zeolites. Glass. By melting together a mixture of calcium carbonate, CaCO 3 , sodium carbonate, Na 2 CO 3 , and a pure quartz sand, SiO 2 , in proper proportions, a silicate of calcium and sodium is obtained, which does not readily crystallize on cooling, but which passes through a stage in which it becomes more and more viscous and finally solidifies to a transparent, homogeneous mass. The sodium may be partly or completely replaced by potassium, the calcium may be replaced by lead, and part of the silica by boric anhydride, B 2 O 3 , giving glasses suitable for special uses. In all of these the glass is to be considered as a complex mixture of silicates in the form of an extremely viscous, supercooled liquid. The value of glass depends largely upon the fact that as an amorphous, supercooled liquid it is still viscous but sufficiently GLASS. STRONTIUM 467 plastic so that it can be worked over a considerable range of temperature. Some of the more important varieties of glass are as follows : window glass, plate glass and the glass of ordinary table ware are usually a silicate of calcium and sodium. The finer grades of such glass are often called crown glass. Flint glass is a silicate of lead and sodium or potassium prepared by melting together litharge, PbO, potassium carbonate, K 2 CO 3 , and silica, SiO 2 . The name comes from the former use of crushed flints for the silica. Flint glass has a higher index of refraction but also a relatively greater dispersive power than crown glass, and the two varieties of glass are used together for achromatic lenses and for direct vision spectroscopes. It melts easily. Strass or paste is a heavy lead glass with a high index of refraction, used in making imitations of diamonds and other precious stones. Bohemian glass is a silicate of potassium and calcium having a high melting point. It is also less soluble in water than ordinary glass and is used for combustion tubing and for beakers and flasks used in the laboratory. It is often called hard glass. It has been largely replaced by various borosilicate glasses. " Jena " glass, " Resistanz " glass and " Non-sol " glass are borosilicates containing a little zinc. They are much less sol- uble in water than the ordinary glasses, and some of them soften at much higher temperatures. These properties render them suitable for special uses in chemical laboratories, especially for combustion tubing, for beakers and flasks for use in quantitative analysis and for test tubes for bacteriological cultures. Durax glass is a variety especially resistant to alkaline solutions. For the manufacture of thermometers several borosilicate glasses are made which give a much smaller depression of the zero point after use at high temperatures than is the case with thermometers made from ordinary glass. Some of these glasses also give thermometers which correspond closely with the hydrogen thermometer at high temperatures. Strontium, Sr, 87.63. Occurrence. As the most important natural compounds of calcium are the carbonate and sulfate, so 468 A TEXTBOOK OF CHEMISTRY strontium and barium are found chiefly as carbonates and sulfates. The sulfates are much less soluble than calcium sul- fate. Strontium carbonate, SrCOs, is called strontianite, and strontium sulfate, SrSO4, celestite. The latter name is given because the mineral is often of a light blue color, but the pure sulfate is white. Water dissolves about 20 times as much cal- cium sulfate as it does of strontium sulfate. * Strontium Hydroxide, Sr(OH) 2 .8 H 2 O, is more soluble than calcium hydroxide and is much more easily soluble in hot than in cold water. It forms a difficultly soluble compound with cane sugar, Ci2H22On, and is sometimes used to recover sugar from the molasses of the beet sugar manufacture. Strontium Nitrate, Sr(NO3) 2 , is used, mixed with sulfur, charcoal and potassium chlorate or nitrate for red lights in fire- works. Barium, Ba, 137.37. Occurrence. Barium is found in nature as the carbonate, BaCOs, called witherite, and the sulfate, BaSC>4, called barite. The latter is more common. Barium Oxide, BaO. Barium carbonate is very much more stable than calcium carbonate. The dissociation pressure of calcium carbonate reaches atmospheric pressure at 898, but that of barium carbonate is equal to one atmosphere at 1350 (Finket- stein, Ber. 39, 1588 (1906)). Even at this temperature the decomposition, at first, appears to give a basic carbonate, prob- ably BaCO 3 .BaO, and a temperature of 1450 is required for the decomposition of this compound at atmospheric pressure. It is impracticable, therefore, to prepare barium oxide by the direct decomposition of the carbonate. If the carbonate is mixed with charcoal, however, the carbon dioxide formed by the dissociation of the barium carbonate will at once react with the carbon to form carbon monoxide. As a result of the two equilibria: BaCO 3 ^ BaO + CO 2 co 2 + c : 2 co the constant removal of the carbon dioxide by means of the second reaction makes it possible to prepare barium oxide at a ALKALI-EARTH METALS: BARIUM 469 temperature at which the dissociation pressure is small and the preparation is carried out, technically, by this method. Barium oxide may also be prepared by the decomposition of the nitrate, Ba(NOa) 2 , which occurs at a much lower temperature than that for the decomposition of the carbonate. Barium oxide combines directly with water to form the hydrox- ide, Ba(OH) 2 , the reaction being accompanied by considerable evolution of heat. It is used, chiefly, for the preparation of the peroxide, BaO 2 . It is also used in the laboratory as a powerful dehydrating agent, as in the preparation of absolute alcohol. Barium Peroxide, BaO 2 . The dissociation pressure for the reaction, 2 BaO + O 2 ^ 2 BaO 2 , is : Temperature 525 670 735 775 790 Pressure in mm. 20 80 260 510 670 When we remember that the partial pressure of oxygen in the air is 21 per cent of 760 mm., or 160 mm., it is evident that if air is passed over barium oxide at a temperature of 670, oxygen will be absorbed with the formation of the peroxide. On the other hand, if the peroxide is heated to about 800 under atmos- pheric pressure, the peroxide will be decomposed with evolution of oxygen. Again, if barium oxide at 800 is subjected to the action of air under a pressure of about five atmospheres, so that the partial pressure of the oxygen is a little more than one atmosphere, oxygen will be absorbed, and on lowering the pres- sure to atmospheric pressure oxygen will be evolved without any change in temperature. Both methods have been used for the technical preparation of oxygen. A little water vapor must be present to catalyze the reaction, and the carbon dioxide of the air must be carefully removed. Barium peroxide is also prepared and used for the manu- facture of hydrogen peroxide, H 2 O 2 . It forms a hydrate, BaO 2 .8 H 2 O, which is difficultly soluble and which is formed on treating the anhydrous barium peroxide with water. It is also formed by precipitating an ice-cold solution of barium chloride, BaCl 2 , with a cold solution of sodium peroxide, Na 2 O 2 . 470 A TEXTBOOK OF CHEMISTRY Barium Hydroxide, Ba(OH)2.8 H^O, is much more easily soluble in water than calcium or strontium hydroxides. It is also much more easily soluble in hot than in cold water and can be easily recrystallized from hot water. It gives a precipitate of barium carbonate, BaCOs, with carbon dioxide, and is a very sensitive reagent for the qualitative detection or quantitative determination of that gas. It is used to prepare an alkali solution which is free from carbonate, but such a solution must, of course, be carefully protected from the carbon dioxide of the air. Barium Chloride, BaCl2.2 H^O, is the most common sol- uble salt of barium. It is used especially for the detection and quantitative determination of sulfates. * Barium Nitrate, Ba(NOa)2, is sometimes used for the detec- tion of sulfates in solution containing silver or other metals which form insoluble chlorides. * Barium Sulfide, BaS, is formed "by hefting barium sulfate, the most plentiful natural source of barium compounds, with charcoal : g^ + 4 c = BaS + 4 CO Barium sulfide dissolves in hydrochloric acid with the forma- tion of the chloride or in nitric acid with the formation of the nitrate, methods which were formerly used for the preparation of these salts from the insoluble sulfate. Barium Sulfate, BaSO 4 . The mineral barite is sufficiently abundant to form a cheap source for barium compounds. It is also sometimes used in a finely pulverized form as an adulterant for white lead. It is crystalline and the particles are much more transparent than those of white lead, so that paints containing it have less covering power than those made from pure lead compounds. It is used in many mixed paints, especially in " lithopone," which is a mixture of barium sulfate, BaSO 4 , and zinc sulfide, ZnS, obtained by precipitating a solution of zinc sulfate with barium sulfide, BaS. Lithopone is not blackened by hydrogen sulfide and is more suitable than white lead for places where that gas is liable to be present. RADIUM 471 Barium sulfate requires about 400,000 parts of water for its solution and it is only slightly more soluble in dilute acids. It is very much used for this reason for the detection and estima- tion of sulfates. It has a very marked tendency to form solid solutions or mixed crystals with other sulfates, barium chloride or other compounds which may be present. The presence of these foreign substances may lead to rather serious errors in the determination of sulfates, unless great care is used in following proper methods of manipulation. Barium sulfate is appreciably soluble in solutions containing considerable hydrochloric acid and dissolves rather easily in concentrated sulfuric acid. Flame Colors for Calcium, Strontium, and Barium. Calcium compounds impart a brick red color to the Bunsen flame, stron- tium compounds a bright red and barium compounds a green color. The spectra show bright lines and bands which are easily distinguished, even f^fh the simplest forms of spectroscopes. Radium (Ra, 226.4). About 1878 Sir William Crookes dis- covered that in the discharge of electricity through a highly rarefied gas rays are shot out at right angles to the surface of the cathode and produce a beautiful green fluorescence at the point where they strike the containing tube. About twenty years later it was shown that these rays consist of electrons traveling with velocities approaching that of light. In 1895 Rontgen, partly by accident, discovered that rays, afterwards called Rontgen rays, emanate from the glass at the point of fluores- cence. These rays penetrate paper, wood and some other ob- jects, which are opaque to ordinary light, and affect a photo- graphic plate placed behind such screens. The rays are, however, intercepted by metals or substances containing compounds of the metals, as, for instance, by bones. The opacity of various sub- stances is closely proportional to their density. Very shortly after, in 1896, Becquerel, in Paris, discovered that minerals con- taining uranium have the property of emitting penetrating radi- ations resembling the Rontgen rays in their effect on a photo- graphic plate protected from ordinary light rays by a screen of 472 A TEXTBOOK OF CHEMISTRY black paper. Further study by Madame and Monsieur Curie led to the discovery in uranium ores of a new element, radium. This belongs to the calcium-barium family, and forms a sulfate which is much less soluble than barium sulfate and which is made use of in separating radium from other elements. Radium and its compounds exhibit the following remarkable properties : 1. It affects a photographic plate through black paper and will cause the fluorescence of zinc sulfide and some other com- pounds exposed to the action of its rays. 2. It causes the ionization of air, that is, the separation of the molecules of the air into positive and negative, charged particles, causing the air to become a conductor of electricity. A gold-leaf electroscope, which will remain charged for a long time in ordi- nary air, is rapidly discharged and the leaves fall together in air which has been exposed to the action of radium. The rate of discharge furnishes a quite accurate measure of the quantity of radioactive substances present, and the measurement of this rate is the most important method used in the study of radioactive elements. 3. Radium continually evolves heat. One gram of the ele- ment gives out 132 small calories per hour. This phenomenon is independent of the temperature or of the form in which the element is combined or of any other conditions which can be controlled. Disintegration of Atoms. In 1902-1903 Professor Rutherford, then at McGill University in Montreal, published a series of papers in which he proposed the hypothesis that the atoms of radioactive elements disintegrate more or less rapidly, breaking down with the formation of other elements. In the disintegra- tion, portions of the atom are shot out from it with tremendous velocity. Some sort of potential or kinetic energy within the atom is liberated in this manner and manifests itself as heat energy, and this explains the heat evolved by radioactive ele- ments. Incidentally this makes it probable that all atoms are complex in their internal structure and are storehouses of immense quantities of energy. The particles shot out by the RADIUM: DISINTEGRATION THEORY 473 disintegrating atoms seem to tear apart molecules which they strike, separating them into charged particles or ions, and in this manner air or other gases which are exposed to the action of radioactive substances become conductors of electricity as has been mentioned above. Rutherford's disintegration theory was very strongly sup- ported when Soddy, who began work with Rutherford, demon- strated, while working with Sir William Ramsay in London, that helium is one of the disintegration products of radium. The theory is now accepted, at least as a working hypothesis, by all investigators in this field. A few chemists, on account of the disintegration of radium, have contended that it is not properly called an element but should be classed as a chemical compound. But radium finds its place in the periodic system and has all of the other properties which usually characterize an element. Moreover, there is good reason to believe that radium is, in turn, a disintegration product of uranium. If radium is a compound, uranium is a compound also. It seems better, therefore, to revise our definition of an element and accept the notion that the atoms of some elements, and possibly of all, may disintegrate with the formation of other atoms. The relation between an atom and its disintegration products is evidently very different from the relation between a compound and the elements of which it is composed. Nature of the Radiations from Radioactive Substances. Four kinds of rays have been distinguished as emanating from radioactive substances. The first kind, called -rays, have been identified as atoms of helium, carrying a double positive charge. As they are charged particles moving with a high velocity, they are slightly deflected by a strong magnetic field. The /8-rays appear to be identical with the cathode rays of the Crookes tube, i.e. they are electrons moving with varying velocities, the velocity sometimes approaching that of light. As the charge is negative, they are deflected by a magnetic field in a direction opposite to that of the a-rays ; and as the mass is very much smaller in proportion to the charge, the deflection is much 474 A TEXTBOOK OF CHEMISTRY greater. The y-rays seem to be like the X-rays or Rontgen rays and are probably of the nature of ether waves. They will penetrate a very much greater thickness of metal than the /?-rays and these in turn are much more penetrating than the a-rays. Some of the radioactive elements give out all three kinds of rays, others give only one kind and still others two kinds. The S-rays are electrons moving much more slowly than those which form the /?-rays. They may be given out from the sur- face of radioactive material, or may be emitted from any sub- stance traversed by a-rays, whether solid or gaseous. The nature of the rays furnishes one of the most important means of identifying different elements. The Life of an Element. In accordance with the disintegra- tion theory some of the atoms of radium are constantly decom- posing into helium and another substance which was at first called radium emanation, but which has been characterized by Sir William Ramsay as an element of the argon family, the gas niton. This has an atomic weight of 222.4, the difference in weight between an atom of radium and an atom of niton being the weight of an atom of helium. The rate of the decomposition of radium has been determined by measuring the amount of helium given in a number of weeks or months by a given weight of radium. The result, of the measurement was that one half of a given quantity of radium would disintegrate in 1760 years. The rate of disintegration of niton, on the other hand, is so rapid that one half of a given quantity of the element will disintegrate in 3.8 days. The rate of decay is measured in this case by measuring the rate at which the radioactive effect on the elec- trometer decreases. The disintegration of uranium to form the first of a series of 3 elements which are supposed to stand between uranium (at. wt. 238.5) and radium (at. wt. 226.4) takes place so slowlv that one half of a given quantity of the element would decompose in 6,000,000,000 years. These periods, which are called the " half-life periods " of the elements, form one of the most important characteristics of the radioactive elements. It is evident from the periods which have been given for uranium, RADIOCHEMISTRY 475 radium, and niton that the amount of radium which can exist at a given time must be very small in comparison with the quan- tity of uranium in the world, and that the amount of niton must always be very small in proportion to the amount of the radium from which it is generated. For this reason Sir William Ramsay was compelled to establish the density and other properties of niton by working with a few cubic millimeters of the gas. He weighed the gas with a microbalance, designed for the purpose, with which it was possible to weigh innAnro f a milligram. Other Radioactive Elements. By the study of the rate of dis- integration of elements formed by the decomposition of others a considerable number of radioactive elements have been identified. They form three well-defined series. The uranium series, begin- ning with uranium (at. wt. 238.5) and probably closing with lead (at. wt. 207.5). There are twelve elements between, the best characterized being ionium, radium and niton. The thorium series commences with thorium (at. wt. 232.4) and prob- ably closes with bismuth (at. wt. 208), with ten elements between. The actinium series begins with actinium, an element of unknown atomic weight, and closes with an unknown inactive element, with nine elements between. The half -life of these elements varies from six billions of years for uranium and six hundred millions of years for thorium to a few seconds for some of the elements derived from thorium and actinium. Chemical Action of the Rays. If radium or niton is left in contact with water, the rays which they emit cause dissociation of some of the water into oxygen and hydrogen. Some hydro- gen peroxide is also formed. A glass tube containing a com- pound of radium soon assumes a violet or brown color. The tremendous velocity with which a-particles are expelled from radium or niton gives a unique and powerful form of energy, and it has even been thought that atoms of other elements may be broken into pieces by this means (Ramsay). Radiochemistry in Relation to Geology and Medicine. Radio- active elements are very widely diffused in the rocks of the earth ; and while the proportion of such elements is very small, it has 476 A TEXTBOOK OF CHEMISTRY been shown that the total amount present in the crust of the earth is sufficient to account for the increasing temperature which is observed in deep wells and mines and in tunnels. Indeed it would seem that the proportion of radioactive elements must be smaller at very great depths than it is near the surface. This discovery has thrown very grave doubts on estimates formerly made of the life of the earth, which were based on the supposition that the earth has cooled down from a molten condi- tion. The rays emitted from radium and other radioactive elements are fatal to bacteria. They also may produce severe burns somewhat resembling sunburn. They have been used with some success in the treatment of cancer, ulcers, lupus, etc. Some mineral waters in which chemical analysis has formerly shown no peculiar curative substances have been found to be radioactive, and it seems possible that beneficial results may be obtained by the use of such waters. EXERCISES 1. How many grams of water will be required, theoretically, to convert a pound (453 g.) of plaster of Paris into gypsum ? 2. How much lime can be obtained from a kilogram of marble ? 3. The heat of formation : CO 2 ->CaCO 3 is 42,520 small calories per gram molecule. How much coal having the same heat of combustion as carbon would be required, theoretically, to prepare one kilogram of lime ? See p. 27 for the heat of combustion of carbon. 4. The heat of formation : BaO + CO 2 -> BaCO 3 is 62,220 small calories per gram molecule. How much coal having the same heat of combustion as carbon will be required to give one kilogram of barium oxide on the basis of the reaction : BaCO 3 + C = BaO + 2 CO ? ALKALI-EARTH METALS 477 5. One hundred cubic centimeters of a natural water require 6 cc. of tenth normal sodium hydroxide to give a pink color with phenol- phthalein and the solution then requires 10 cc. of tenth normal hydro- chloric acid to give a red color with methyl red. How much carbonic acid and how much bicarbonate, calculated as calcium bicarbonate, are present in one liter of the water ? 6. How much lime (CaO) will be required to furnish enough calcium hydroxide to soften one U. S. gallon (3.785 liters) of the water just referred to ? CHAPTER XXVII ALTERNATE METALS OF GROUP II. MAGNESIUM, ZINC, CADMIUM AND MERCURY IN its occurrence and in the properties of its most common compounds, magnesium resembles calcium rather than zinc, but in the metallic form the resemblance to zinc is more marked. Magnesium decomposes water slowly at the boiling point, while zinc and cadmium decompose it readily at a higher temperature. Mercury in many of its properties seems to be more closely related to copper and silver than to cadmium, and some authors formerly placed it in the first group. It does not decompose water at any temperature and it forms compounds in which it is apparently univalent as well as those in which it is bivalent. Magnesium, zinc and cadmium are always bivalent. Magnesium, Mg, 24.32, is found as the carbonate, magnesite, MgCOs, as a double carbonate of calcium and magnesium, dolomite, CaCOs.MgCOs, as the double chloride with potassium, carnallite, KCl.MgCl 2 .6 H 2 O, and as a principal constituent of many silicates, especially talc or soapstone, MgsH^SiOs)^ serpentine, Mg 3 Si 2 O 7 .2 H 2 O, and meerschaum, Mg 2 Si 3 O 8 .2 H 2 O. The sulfate, Epsom salts, MgSO 4 .7 H 2 O, and the chloride, MgCl 2 .6 H2O, are also found in many natural waters. Preparation, Properties. Metallic magnesium is obtained by the electrolysis of fused carnallite, MgCl 2 .KCl, from which the water of hydration has been expelled by heat. It is a silver-white, very light metal, having a specific gravity of only 1.75, slightly lower than that of beryllium. It melts at 651 and boils at about 1100. Magnesium wire or ribbon burns in the air with a very intense white light that is particularly rich in the more refrangible rays, which affect the photographic plate. 478 GROUP II: MAGNESIUM 479 For this reason powdered magnesium is the effective constituent of flash-light powders. It is estimated that 10 per cent of the energy of burning magnesium appears as light, a very much larger per cent than is secured by any ordinary illuminant. The temperature of burning magnesium is not, 'however, very high only about 1340. Magnesium tarnishes only slightly, if at all, in dry air and very slowly in moist air, so that it can be kept indefinitely without special precautions. Magnesium is used in the laboratory as a powerful reducing agent for the preparation of silicon and boron. It is also used in a great variety of syntheses of organic compounds. Magnesium Oxide, MgO, is most easily prepared by heating magnesium carbonate, MgCO 3 , which decomposes at a much lower temperature than calcium carbonate. It is a light, white powder often called magnesia usta, or burnt magnesia. It is infusible in any ordinary furnace, but may be volatilized in the electric furnace. It is used for crucibles and for some forms of apparatus which must withstand extremely high temperatures. It is used as a basic lining for metallurgical furnaces, especially for the basic process for steel (p. 548). Magnesium Hydroxide, Mg(OH) 2 , is obtained as a white precipitate on the addition of sodium hydroxide, NaOH, or barium hydroxide, Ba(OH) 2 , to a solution containing almost any soluble salt of magnesium. It is much less soluble than calcium hydrox- ide, dissolving in about 6500 parts of water. In spite of this difficult solubility, however, it is not precipitated in the presence of ammonium salts. This is due probably to two reasons : (1) in the presence of an ammonium salt the ionization, NH 4 OH ^ NH 4 + + OH- is repressed by the ammonium ions, NH4 + , of the salt, and the concentration of the hydroxide ions, OH", is low in such a solu- tion; and (2) because magnesium hydroxide as a unibivalent com- pound very probably forms intermediate ions, MgOH + + OH", which interfere with the application of the ordinary law of the solubility product (p. 377). 480 A TEXTBOOK OF CHEMISTRY Magnesium hydroxide is easily decomposed by heat into magnesium oxide and water. Magnesium Chloride, MgCl 2 .6H 2 O, crystallizes from a con- centrated solution of the salt. It is very easily soluble in water. When "an attempt is made to drive out the water of the salt by heating it, both water and hydrochloric acid escape, and a mixture of variable composition, containing chiefly mag- nesium oxide, finally remains : MgCl 2 + H 2 O = MgO + 2 HC1 The process has been used to a limited extent as a basis for the preparation of hydrochloric acid and chlorine. Magnesium Ammonium Chloride, MgNH 4 Cl3.6 H 2 O. This salt is easily prepared by crystallizing from water a mixture of equimolecular amounts of magnesium chloride and ammonium chloride. The water of hydration may be expelled with very little loss of hydrochloric acid, and on heating the anhydrous salt to a slightly higher temperature the ammonium chloride dissociates and escapes, leaving anhydrous magnesium chloride behind. Magnesium Sulfate, MgSO 4 .7 H 2 O, or Epsom Salts, is found in some mineral waters used for their medicinal properties, es- pecially in Hunyadi water, in which it is associated with sodium sulfate, Na 2 SO 4 . * Magnesium Sulfide, MgS, may be prepared by heating a mixture of magnesium and sulfur. It is decomposed by water, giving magnesium hydroxide and hydrogen sulfide : MgS + 2 HOH = Mg(OH) 2 + H 2 S Both the insolubility of the hydroxide and the volatility of the hydrogen sulfide contribute to cause the reaction to go to com- pletion. * Magnesium Ammonium Phosphate, MgNH 4 PO 4 .6 H 2 O, is a difficultly soluble salt which is formed when solutions contain- ing magnesium, ammonium and a soluble phosphate are brought together. It is used in analytical chemistry for the determina- tion of both magnesium and phosphoric acid. A precipitate GROUP II: ZINC 481 having the exact composition represented by the formula can be obtained only by securing exactly the right conditions as regards the concentration of the various solutions employed. When heated, the compound decomposes quantitatively with the formation of magnesium pyrophosphate, Mg2P2O7. Zinc, Zn, 65.37. Occurrence. Zinc is found in nature as the sulfide, sphalerite, ZnS, the carbonate, smithsonite, ZnCOs, the silicates, willemite, Zn2SiO4, and calamine, H^ZnSiO-i, and in the mineral franklinite, which consists of oxides of iron, zinc and manganese. Franklinite is an important ore in New Jersey and is used as a source of zinc, manganese and iron. Metallurgy. The sulfide of zinc is converted into the oxide by roasting it, that is by heating it in a furnace with free access of air : ZnS + 3 O = ZnO + SO 2 The oxide, obtained in this way or by heating the carbonate, is mixed with coal and heated to a high temperature in an earth- enware retort having a receiver luted to it with clay. The zinc oxide is reduced, and the zinc, which boils at 925, distils over. Zinc melts at 419.4. Impure zinc dissolves easily in hydrochloric or in dilute sul- furic acid, with evolution of hydrogen. Pure zinc, which can be obtained by distillation in a vacuum, is attacked very slowly or not at all by these acids, but dissolves readily in contact with platinum. It has been pointed out that these facts indicate that the solution is always associated with electrical phenomena. Impure zinc which is covered by a thin film of amalgam, giving it a homogeneous surface, is also not attacked by the dilute acids, and such amalgamated zinc is used in electric batteries. Zinc has a specific gravity of 6.9. Uses. Galvanized Iron. Metallic zinc is used chiefly in brass, an alloy of the metal with about twice its weight of copper, and as a coating for iron to protect it from rusting. It is also a con- stituent of many of the bronzes, and is used as the metal which is dissolved or corroded in most forms of primary electrical batter- 482 A TEXTBOOK OF CHEMISTRY ies, especially in the gravity cell used in telegraphy and in the so-called " dry " batteries. " Galvanized " iron is prepared by dipping carefully cleaned sheet iron or other iron or steel articles in melted zinc. The value of the coating depends on two properties : first, zinc is electropositive with reference to iron and when the two metals are in contact with each other and also in contact with an elec- trolyte the zinc is attacked and the iron is protected ; second, the action upon the zinc causes the formation of a very thin, coherent coating of zinc oxide or hydroxide, which is practically insoluble in water and protects the zinc from further action. A small amount of zinc passes into solution, however, and this may be increased very considerably in the presence of even weak acids. As zinc salts are poisonous, pails or dishes of gal- vanized iron are not suitable for culinary use. The coating of zinc has been sometimes applied to the iron by an electrolytic method, and the term " galvanized iron " came from this method of manufacture. Sherardized Iron. A new process for coating iron with zinc has been developed by Sherard Cowper-Cowles. The articles to be coated are heated with zinc dust in iron drums at 500-600 for thirty minutes to several hours, according to the thickness of the coating desired. The process is somewhat analogous to the manufacture of cementation steel. See Johnson and Woolrich, Trans. Am. Electrochem. Soc. 21, 561 (1912). Cowper-Cowles, Electrochem. and Met. Ind. 6, 189 (1908). Zinc Oxide, ZnO, is prepared by burning the vapors of metallic zinc. It is a white powder and gives with linseed oil an excellent pigment, which has the advantage of not being blackened by hydrogen sulfide because zinc sulfide is also white. Zinc oxide is often used with phosphoric acid for a cement in dental work. The two substances combine to form a basic zinc phosphate which sets to a hard mass that adheres strongly to the surfaces with which it is in contact. Zinc oxide is yellow when hot, but turns white again on cooling, a property used for the detection of zinc compounds in blowpipe reactions. GROUP II: ZINC, CADMIUM 483 Zinc Hydroxide, Zn(OH) 2 , forms as a white precipitate on adding a soluble hydroxide to a solution of a zinc salt. It dis- solves in an excess of sodium hydroxide or potassium hydroxide, forming sodium zincate, Na2ZnC>2, or potassium zincate, K^ZnC^. In forming these compounds zinc hydroxide seems to act as an acid, which might be called zincic acid, and the formula might be written H2ZnO2. Toward acids, however, zinc hydroxide con- ducts itself as a true hydroxide or base. Compounds which exhibit a dual nature of this sort, acting in some conditions as acids and in others as bases, are said to be amphoteric. Zinc Chloride, ZnCl 2 . An aqueous solution of this salt is easily prepared by dissolving metallic zinc or zinc oxide in hydro- chloric acid. Unlike magnesium chloride, the solution loses only a small amount of hydrochloric acid when heated to a high temperature to expel the water. The pure, anhydrous salt melts at 290-297 and boils at 730. When boiled in an iron tube, it furnishes an easy means of securing a constant, rather high temperature and has been found useful for this purpose. Zinc chloride is used in the treatment of wooden ties to prevent decay. Zinc Sulfate, ZnSO 4 .7 H 2 O, or White Vitriol. Anhydrous zinc sulfate, ZnSO 4 , can be prepared by roasting the sulfide, ZnS, at a moderate temperature. At a higher temperature the sul- fide roasts to the oxide and sulfur dioxide. The hydrate, ZnSO 4 .7 H 2 O, forms rhombic crystals and is easily soluble. Zinc Sulfide, ZnS, forms as a white precipitate when hydrogen sulfide is passed into a neutral or alkaline solution of a zinc salt. The precipitate is formed even in slightly acid solutions, and care- ful attention must be paid to the amount and character of the acid present if a separation from other metals is desired. In the presence of sulfuric acid which is weaker than fifth normal (about 1 per cent) the sulfide will be precipitated. Cadmium (Cd, 112.40). Many zinc ores contain a small amount of cadmium. As the boiling point of cadmium (785) is considerably lower than that of zinc (925), the former distills over first in the preparation of zinc, and by collecting these por- 484 A TEXTBOOK OF CHEMISTRY tions and subjecting them to fractional distillation, nearly pure cadmium can be prepared. Metallic cadmium closely resembles zinc in appearance and in many of its properties. It melts at 320.9, boils at 785 and has a specific gravity of 8.65. It was formerly used in amalgams for filling teeth, but other amalgams are now considered more suitable. It is a constituent of Wood's metal and of the easily fusible alloys used for safety fuses in electrical circuits and for automatic sprinklers used for protection against fire. * Cadmium Hydroxide, Cd(OH) 2 , is easily obtained as a white precipitate. It dissolves easily in acids but does not dissolve in solutions of sodium or potassium hydroxides, as zinc hydroxide does. It is decomposed when heated, giving cadmium oxide, CdO, as a brown powder. * Cadmium Sulfate, 3 CdSO 4 .8 H 2 O, is an easily soluble salt used in the Weston standard cells, which are the most satisfac- tory primary standard for the measurement of electromotive force. Cadmium Sulfide, CdS, forms as a yellow precipitate in solu- tions of cadmium salts which do not contain too much free acid or too much of salts which interfere with the precipitation. From a solution containing both zinc and cadmium in which sulf uric acid is present and the concentration of the hydrogen ion is between fifth normal and twice normal and other interfering salts or acids are absent, the precipitation of the cadmium is practically complete, while only a small amount of the zinc will come down. For a complete separation, however, the cadmium sulfide must be dissolved and reprecipitated. In hydrochloric acid stronger than 0.3 normal cadmium sulfide is not completely precipitated. Cadmium sulfide dissolves readily in boiling, dilute sulfuric acid, but is insoluble in a solution of potassium cyanide, KCN. Mercury, Hg, 200.6. Occurrence. Metallurgy. From the positions of the elements of Group II in the electromotive series (p. 436) mercury is the only element of the group which could appear in the free state in nature. It is occasionally found GROUP II: MERCURY 485 in small globules disseminated in porous rocks. Mercury occurs chiefly, however, in the form of the native sulfide, cinnabar, HgS, a brilliant red mineral, when pure. When ores containing cinna- bar are roasted by heating in a current of air, the sulfur burns to sulfur dioxide while the mercury distills and is condensed in long flues where the vapors must be very thoroughly cooled to prevent loss: HgS + O 2 = Hg + SO 2 Mercury may also be obtained by mixing the sulfide with lime or with iron and distilling : 2 CaO + 2 HgS = 2 CaS + 2Hg+ O 2 Fe + HgS = FeS + Hg Mercury can be purified by allowing it to fall in very minute globules, through a chamois skin tied over the end of a glass funnel, into dilute nitric acid contained in a tube 2 meters long. The tube is drawn out and bent upward at the bottom so that a short column of mercury in the overflow tube balances the col- umn of nitric acid. The nitric acid dissolves zinc, arsenic, lead and nearly all of the other metals likely to be present. Mercury may also be separated from nearly all other metals by distillation under diminished pressure. If a very little air is allowed to pass through the mercury by means of a very fine, hairlike, capillary tube, troublesome bumping of the mercury can be avoided and zinc and some other metals are oxidized, giving purer mercury than if the distillation is carried out in the absence of air (Hulett). Properties and Uses. Mercury is a heavy, mobile liquid, with O 9O a density, at ^ = 13.5956 or at ^- = 13.5463. It freezes at 38.70. Its freezing point on a Fahrenheit thermometer is 37.7. It will be noticed that the scales of the Centigrade and Fahrenheit thermometers approach very closely together at this temperature. Mercury boils at 357. Its critical tem- perature is about 1275, and its critical pressure is calculated as about 675 atmospheres, an extraordinarily high value. (Menzies, 486 A TEXTBOOK OF CHEMISTRY J. Am. Chem. Soc., Sept., 1913. Konigsberger, Chem. Ztg. 13d, 1321 (1913) .) Mercury oxidizes slowly to red mercuric oxide, HgO, when heated to its boiling point in the air. (See Lavoisier's experiment, p. 19.) Mercury dissolves very slowly to mercur- ous nitrate, HgNOs, in dilute nitric acid, with evolution of nitric oxide, NO. It is insoluble in hydrochloric acid, but is converted into mercurous sulfate, Hg2SO4, by hot, concentrated sulfuric acid, with evolution of sulfur dioxide. Mercury is used in the amalgamation processes for the re- covery of gold and silver (p. 441), in making thermometers, for barometers and manometers, in mercury air pumps and in the collection and measurement of gases. Its advantages over all other substances used in thermometers are, especially, that it does not wet or attack the glass, that it is liquid over a wide range of temperature, including the common range of air tem- peratures, and that its rate of expansion is very uniform. Its coefficient of expansion between and 100 is so nearly constant that a mercury thermometer graduated in equal degrees does not differ from the standard hydrogen scale by more than 0.2 at any point between these temperatures. Mercury thermom- eters cannot, of course, be used at temperatures below 39 and ordinary thermometers cannot be used above 300 indeed, the thread of an ordinary thermometer will usually break before that temperature is reached. By filling the space above the mercury with nitrogen under pressure, however, thermometers graduated to 460 are made, and by filling the space with carbon dioxide the range has been carried to 550 or above. There is likely to be a large zero-point correction for such thermometers, and they must be carefully treated, if accurate results are required. The stem correction is also large, unless the whole thermometer is immersed in the sub- stance whose temperature is to be measured. Amalgams. The alloys of mercury are called amalgams. Many metals, such as sodium, potassium, copper, silver, gold, zinc, cadmium, tin and lead, dissolve in or alloy with mercury in all proportions or give amalgams having a wide range in their GROUP II: AMALGAMS 487 composition. Other metals, as iron and platinum, dissolve in mercury to only a trifling extent or not at all. The amalgams of gold and silver are used to separate these metals from large masses of other substances mixed with them in their ores. Sodium amalgam is often used as a reducing agent, especially for organic compounds. Zinc amalgam and other amalgams may be used in the same way. An amalgam of tin was formerly used for the backs of mirrors, but has been replaced by a thin film of metallic silver in modern mirrors. An amalgam with silver and other metals is used for filling teeth. In many cases mercury combines with metals to form definite compounds. Such compounds are most easily identified by a sou 300 250 8 2 H H 150 3 | 100 H P 50 -50 Per Cent.) Mercury ) Per Cent. ) II \ ' / \ 7 \\ \\ -^^. ^-^ ---^. ^^^^ / 1 ^ - -^ X 1 1 10 20 30 40 50 60 70 80 90 100 LOO 90 80 70 60 50 40 30 2O 10 Fig. 102 study of the freezing point curve of amalgams of varying composi- tion. In Fig. 102 the ordinates give the melting points and the abscissas give the composition of a series of amalgams of mercury with sodium. It will be seen from the figure that the addition 488 A TEXTBOOK OF CHEMISTRY of mercury to sodium lowers its melting point until a minimum is reached for an amalgam containing about 40 per cent of mer- cury and 60 per cent of sodium. This minimum is called a eutectic point. An amalgam of this composition melts at 21. Further addition of mercury raises the melting point till a maxi- mum is reached for an amalgam containing 5.4 percent of sodium and 94.6 per cent of mercury, which melts at 346. Further addition of mercury lowers the melting point till this would, undoubtedly, fall below the melting point of pure mercury. A compound having the formula NaHg2 would contain 5.43 per cent of sodium and 94.57 per cent of mercury. Evidently the amalgam of the highest melting point is a compound of this formula. The addition of either sodium or mercury to this compound lowers its melting point just as the addition of salt or any soluble substance lowers the melting point of ice. Changes in the direction of the curve at other points indicate that other compounds of mercury and sodium are present in some of the amalgams, but these details are not shown in the figure. See Kurnakow, Z. anorg. Chem. 23, 443 (1900). Compounds of Mercury. Mercury forms mercurous com- pounds, such as Hg 2 O, Hg 2 Cl2, Hg 2 SO 4 , in which it appears uni- valent, but in which it is probably really bivalent as expressed Hg-Cl by the graphical formula, | . It also forms mercuric com- Hg-Cl pounds, such as HgO, HgCl 2 , HgSO^ in which it is clearly bi- valent. In the formation of these two classes of compounds and also in its conduct toward nitric, sulfuric and hydrochloric acids mercury resembles copper rather than zinc or cadmium, and it has sometimes been classified under the first group of the Periodic System in place of gold. Mercurous Oxide, Hg 2 O, is formed as a black precipitate when a solution of sodium hydroxide is added to a solution of mercurous nitrate, HgNO 3 , or when calomel, Hg 2 Cl 2 , is digested with a solution of sodium hydroxide. Mercuric Oxide, HgO, is formed slowly as a heavy red crys- GROUP II: MERCURY 489 talline powder when mercury is heated to its boiling point in the air. It is obtained more easily by heating the nitrate. A yellow precipitate having the same composition is formed on adding an alkali to a solution of a mercuric salt. Mercuric Sulfide, HgS. The mineral cinnabar, HgS, is a bright red compound. When hydrogen sulfide is passed into a solution of a mercuric salt a black mercuric sulfide of exactly the same composition is precipitated. By subliming the black sulfide, or by warming it with a solution of sodium sulfide, it can be converted into the red variety. The red form is used under the name of vermilion as a brilliant red pigment. When applied to iron or zinc, however, it is decomposed with libera- tion of metallic mercury. Mercurous Chloride, or Calomel, Hg 2 Cl2, is prepared by sub- liming a mixture of mercuric chloride, HgCl2, and mercury, or a mixture of mercuric sulfate, HgSC>4, salt and mercury. The crude product usually contains a little mercuric chloride, which is removed by treatment with alcohol, in which the mercurous chloride is insoluble while the mercuric chloride is easily soluble. The gram molecular volume of the vapor of mercurous chloride weighs about 236 grams, corresponding to the formula HgCl, but it has been shown that the vapor really consists of a mixture of mercuric chloride and mercury (HgCl2 + Hg) (Alex. Smith, J. Am. Chem. Soc. 32 1541 (1910)). From this it seems probable that the true formula of mercurous chloride is Hg 2 Cl 2 . Calomel is used as a medicine. It is now usually administered in very small doses and mixed with sodium bicarbonate, NaHCO 3 , to render it less soluble in the acid gastric juice. In former times the careless administration of large doses sometimes caused sali- vation and other serious injuries to patients. Mercurous chloride is formed as a white precipitate on adding hydrochloric acid or a soluble chloride to a solution of mercurous nitrate or of some other soluble mercurous salt. Mercuric Chloride or Corrosive Sublimate, HgCl2, is prepared by subliming a mixture of mercuric sulfate, HgS(>4, and salt, 490 A TEXTBOOK OF CHEMISTRY NaCl. It is a white, crystalline salt, which melts at 265 and boils at 307. It is soluble in about 14 parts of cold water and more easily soluble in alcohol. It also dissolves in ether. When taken internally it is very poisonous. The best antidote is the white of an egg, with which it forms an insoluble compound. The solution in alcohol is sometimes used as a poison for insects. A dilute solution (usually 1 : 1000) is much used as an antiseptic in surgery. From such a solution the mercury does not seem to be absorbed from a wound or through the skin. Mercuric Iodide, Hgl2, is precipitated as a scarlet powder on adding a solution of potassium iodide, KI, to a solution of mercuric chloride. The precipitate dissolves in an excess of the potassium iodide, forming the complex salt, K 2 HgI 4 . Sodium hydroxide produces no precipitate in such a solution, evidently because it contains only a very small number of mercuric ions, Hg ++ . An alkaline solution prepared in this manner is used under the name of Nessler's solution as an extremely sensitive reagent for ammonia. Mercurous Nitrate, HgNO 3 or Hg 2 (NO 3 ) 2 , is formed by the solution of mercury in cold, dilute nitric acid. It is hydrolyzed by water, giving a basic nitrate, Hg 2 (OH)NO 3 , hence to secure a clear solution a little nitric acid must be added to carry the reversible reaction : Hg 2 (NO 3 ) 2 + HOH ^ Hg 2 (OH)NO 3 + HNO 3 to the left. To counteract the oxidation to mercuric nitrate, Hg(NOs) 2 , by the oxygen of the air, some metallic mercury must be kept in contact with the solution : 2 Hg 2 (NO 3 ) 2 + 4 HNO 3 + O 2 = 4 Hg(NO 3 ) 2 + 2 H 2 O Hg(N0 3 ) 2 + Hg = Hg 2 (N0 3 ) 2 * Mercuric Nitrate, Hg(NO 3 ) 2 .8 H 2 O, is obtained by dissolv- ing mercury in warm, concentrated nitric acid. * Mercuric Cyanide, Hg(CN) 2 , can be prepared by dissolving precipitated mercuric oxide in a solution of hydrocyanic acid, HCN. It decomposes into mercury and cyanogen, C 2 N 2 , when heated. MAGNESIUM, ZINC, CADMIUM AND MERCURY 491 * Mercuric Fulminate, Hg(ONC) 2 , is used in cartridges and percussion caps for firearms and in detonating caps for firing dynamite and nitroglycerin. It is hydrolyzed by hydrochloric acid and water to hydroxyl- amine hydrochloride, NH 2 OH.HC1, and formic acid, HCO 2 H. Hg< +4HC1 + 4H 2 ^C = HgCl 2 + 2 H-0 NH 2 .HC1 + 2 O= lonization of Compounds of Cadmium and Mercury. For some reason, not understood, the chlorides and sulfates of cad- mium and mercuric mercury ionize to a much smaller degree than the corresponding salts of most other metals. Solubility of the Sulfides of Group II. The sulfides of the metals of the first division of Group II, CaS, SrS and BaS, are hydrolyzed by water, forming hydroxides and soluble hydro- sulfides : 2 CaS + 2 HOH = Ca(OH) 2 + Ca(SH) 2 Magnesium sulfide, MgS, gives with water magnesium hy- droxide, Mg(OH) 2 , and hydrogen sulfide. Zinc sulfide, ZnS, is not affected by water, but dissolves in strong acids, if not too dilute. It is almost insoluble in such a weak acid as acetic acid. Cadmium sulfide, CdS, dissolves readily, especially on warming, in moderately concentrated, strong acids, especially in nitric acid (5 per cent) or sulfuric acid (15 per cent). Mercuric sulfide, HgS does not dissolve, even in boiling nitric acid, but dissolves easily in aqua regia. These relations furnish a ready means of separating magnesium, zinc, cadmium and mercury from each other and cause them to be classified in three different groups for analytical purposes. Conduct of Solutions of Magnesium, Zinc and Cadmium Salts toward Ammonium Hydroxide. Ammonium hydroxide gives no precipitate with salts of these metals in solutions con- 492 A TEXT BOOK OF CHEMISTRY taining ammonium chloride. The zinc and cadmium salts form complex compounds, which are soluble, such as Zn(NHs)4SO4 and Cd(NHs)4Cl2. These resemble the corresponding com- pounds of copper but are colorless. For Mg see p. 479. Ammono-mercuric Compounds. When a solution of an alka- line hydroxide is added to a solution of a mercurous or mercuric salt, mercurous oxide, Hg2O, or mercuric oxide, HgO, is precipi- tated, as has been stated. If ammonium hydroxide, NH^OH, is added to such a solution, however, compounds of a wholly differ- ent type, called ammonobasic mercuric compounds, are precipi- tated. These may be considered as formed by the ammonolysis of mercuric salts by a process which is closely analogous to the formation of an ordinary basic (aquobasic) salt by hydrolysis. Thus the partial hydrolysis of mercuric chloride may give in solution : HgCl 2 + H.OH ^ H-O HgCl + HC1 Aquobasic Mercuric Chloride In the presence of ammonia, by an exactly analogous reaction, we should have : HgCl 2 + H.NH 2 ^ H 2 N HgCl + HC1 Ammonobasic Mercuric Chloride or Hg(N0 3 ) 2 + H.NH 2 ^ H 2 N HgNO 3 + HNO 3 Ammonobasic Mercuric Nitrate The hydrochloric or nitric acid would, of course, unite with the excess of ammonia present to form ammonium chloride, NH 4 C1, or ammonium nitrate, NH^.NOs. Salts of many other metals undergo ammonolysis in solutions in anhydrous ammo- nia, but the ammonobasic compounds which are formed are decomposed by water in almost all cases, while the ammono- basic mercuric compounds are stable in the presence of water, either because of their extreme insolubility or because of some specific affinity between mercury and nitrogen. (See E. C. Franklin, J. Am. Chem. Soc. 29, 35 (1907) ; Am. Chem. J. 47, 363 (1912) ). MAGNESIUM, ZINC, CADMIUM AND MERCURY 493 Mercurous salts react with ammonia as though they were mixtures of a mercuric salt with mercury : Hg 2 Cl 2 + 2 H.NH 2 ^ H 2 N.HgCl + Hg + NH 4 C1. The metallic mercury colors the precipitate formed from mercurous salts black. If ammonobasic mercuric chloride is dissolved in a solution of ammonium chloride in anhydrous ammonia, the ammonolysis may be reversed exactly as the hydrolysis of a salt may be reversed by hydrochloric acid : H 2 N Hg Cl + NH 4 C1 ^ HgCl 2 + 2 NH 3 From such a solution a compound of the formula HgCl 2 .2 NHa, which contains ammonia of crystallization and is closely analo- gous to the hydrates of other salts (p. 82), may be crystallized. Nessler's Reagent. Mercuric iodide, HgI 2 , which is almost wholly insoluble in water, dissolves easily in a solution of potas- sium iodide, owing to the formation of a complex salt, potassium mercuric iodide, K 2 HgI 4 . In such a solution sodium hydroxide will give no precipitate ; but if ammonia or an ammonium salt is added to the alkaline solution, a precipitate of ammonobasic- aquobasic-mercuric iodide, HO Hg NH Hg I, is formed. In very dilute solutions of ammonia the solution, which is called " Nessler's reagent," produces a brown coloration which is used for the detection and quantitative estimation of ammonia. EXERCISES 1. How much crystallized hydrate of magnesium chloride and how much ammonium chloride will be required to furnish one pound (453 grams) of anhydrous magnesium chloride ? 2. How many liters of carbon dioxide at 20 and 760 nun. will be given by heating 84 grains of magnesium carbonate ? 3. How much dolomite would be required to give a pound of Epsom salts ? 4. What volume of gases will 0.284 gram of mercuric fulminate give by its explosion, supposing the temperature of the gases to be 546 ? CHAPTER XXVIII METALS OF GROUP III. ALUMINIUM FAMILY. RARE EARTH METALS WHILE all of the elements of both divisions of Group II are fully metallic in character and all except radium are comparatively common and their compounds well known, the first element, boron, of Group III, is decidedly nonmetallic, and aluminium is the only metal of the group which can be considered very common. Aluminium, Al, 27.1, is found in a great variety of natural silicates, especially in the feldspar and mica of the granites and similar rocks, which are still abundant and which must have been much more common in early geologic time. By the pro- longed action of water and the forces of nature such rocks have been slowly disintegrated. The potassium and sodium of the minerals, have been partly, though by no means completely, dissolved and removed, and a hydrated silicate of aluminium, mixed with fragments of quartz and of partially decomposed minerals, has been left in an extremely fine state of division. This material, after transportation for some distance by water, has been deposited by sedimentation and has formed immense beds of shales, clays and soils. Such clays and shales may be considered as ores of aluminium, though they contain only from 15 to 30 per cent of the metal. Pure kaolin, the mineral basis of clay, has the formula Al 2 Si2O7.2H 2 O. Aluminium also occurs as the oxide, A1 2 C>3, in crystals known as ruby and sapphire, which are used as gems, and in a massive, very hard form, called emery and used as an abrasive. The mineralogical name of the oxide is corundum. Bauxite, the hydrate, A1 2 O 3 .2 H 2 O, usually containing a considerable amount of the hydrate of ferric oxide, 2 Fe 2 O 3 .3 H 2 O, is the chief source from which aluminium oxide 494 ALUMINIUM 495 is prepared for the manufacture of the metal. Cryolite, NaaAlF 6 , is a soft, easily fusible mineral found in large quantities in Green- land, but, so far as known, nowhere else, except as a rare mineral. Metallurgy. Aluminium was first prepared by the German chemist Wb'hler in 1828 by the action of potassium on aluminium chloride, but he obtained only a very small quantity in the form of a gray powder. Twenty-six years later Sainte-Claire-Deville exhibited in Paris a quantity of the metal, which he obtained by the action of sodium on the chloride, and the element aroused a great deal of interest as " silver from clay." The desire of ob- taining the metal in larger quantities led to the development of cheaper and better methods for the manufacture of sodium, but as the valence of sodium is one while that of aluminium is three, and the atomic weights are not far different, it must always take about three pounds of sodium to give one pound of aluminium, and the metal manufactured by that process was never put on the market at a price below $10 to $12 a pound. In 1885 Professor Mabery of Cleveland, at a meeting of the American Association for the Advancement of Science held in Ann Arbor, gave an account of a new electric furnace devised in 1882, by the Cowles brothers, for the production of aluminium bronze. They had discovered that, at the high temperature of the electric arc, aluminium oxide, A^Os, can be reduced by carbon to the metallic form, and that if copper is present the alloy, aluminium bronze, can be obtained. This seems to have been the first application of the electric furnace to an industrial pro- cess. Its use for this particular purpose was short-lived. A few years later another American, C. M. Hall, discovered that aluminium oxide dissolves easily in melted cryolite, and that if an electric current is passed from a carbon anode through the molten mass contained in an iron pot, aluminium is deposited in the bottom of the pot, while oxygen liberated at the anode combines with the carbon and escapes as carbon dioxide (Fig. 103). Other materials are now used wholly or in part in place of the cryolite, but the principles used in the process are not changed. The aluminium oxide is obtained by heating bauxite 496 A TEXTBOOK OF CHEMISTRY with carbon in an electric furnace. The iron, silicon and other elements in the bauxite are reduced by this process and may be separated from the fused, pure aluminium oxide. * Thus far the alu- minium oxide used for the production of alu- minium has not been prepared, commercially, from clay, but has usually been made from bauxite. Very recently Alfred H. Cowles has Fig. 103 developed a process by which clay mixed with salt and charcoal is heated in a current of air and steam, giving the reaction : Al 2 Si 2 O 7 + 4 NaCl + 2 H 2 O = 2 Na 2 O.2 SiO 2 . A1 2 O 3 + 4 HC1 The carbon burns to carbon monoxide and serves to render the material porous and easily accessible to the steam and air. Iron, which is present, volatilizes as ferric chloride, FeCl 3 . If the mixture of sodium silicate and aluminate is mixed with lime and heated, an insoluble calcium silicate and soluble sodium aluminate are formed : 2 Na 2 O.2 Si0 2 .Al 2 O 3 + 4 CaO = 2 Ca 2 SiO 4 + NaAlO 2 + Na 3 AlO 3 Calcium & , . . , Silicate Sodmm Alummate From the solution of sodium aluminate, aluminium hydroxide may be precipitated by carbon dioxide : 2 NaA10 2 + C0 2 + 3 H 2 O = 2 A1(OH) 3 + Na 2 CO 3 The process seems promising because it gives three valuable compounds, hydrochloric acid, sodium carbonate and aluminium ALUMINIUM: THERMITE 497 hydroxide, with the use of cheap raw materials (Journal of Industrial and Engineering Chemistry, 5, 331). Properties of Aluminium. Aluminium melts at 658.7 and boils at 1800. It has a specific gravity of only 2.6, almost the same as that of glass and scarcely more than one third that of iron. This makes it useful for the construction of apparatus which should be light. It does not tarnish readily and is used to some extent for cooking utensils. It is not attacked by water, even at the boiling point, but dissolves readily in alkalies or in acids, forming aluminates with the alkalies, such as NaAlO 2 , and salts with acids, such as A^SOJs. Aluminium is rather easily corroded by salt solutions. Aluminium which has been amalga- mated by bringing it into contact with a dilute solution of mer- curic chloride becomes active and will decompose water rapidly at ordinary temperatures. In this condition it is in very sharp contrast with amalgamated zinc, which is not attacked by hy- drochloric or sulfuric acid because of its homogeneous surface. The surface of amalgamated aluminium is gray and evidently nonhomogeneous. Aluminium is used to some extent for elec- tric conductors in place of copper. It is often used as an addi- tion to cast iron, greatly improving its quality. Alloys. The best known alloy is aluminium bronze, composed of copper with 5-12 per cent of aluminium. It resembles gold very closely in appearance and does not tarnish readily. Magna- lium, an alloy with a small amount of magnesium, is very light and is much more easily worked on a lathe than aluminium itself. Alloys containing from 2 to 10 per cent of copper are used for castings for automobiles and for other purposes where lightness is desirable. Goldschmidt's Thermite Process. Aluminium has a very strong affinity for oxygen, as shown by the difficulty with which it is reduced. The heat of combustion of aluminium is : 2 Al + 3 O = A1 2 O 3 + 380,000 calories That of iron, if it could be burned to ferric oxide, is : 2 Fe + 3 O = Fe 2 O 3 + 195,000 calories 498 A TEXTBOOK OF CHEMISTRY From these values it is evident that the reaction Fe 2 3 + 2 Al = A1 2 O 3 + 2 Fe may occur with the evolution of a large amount of heat. Gold- schmidt has made use of this principle for the production of very high temperatures and also for the reduction of chromic oxide and other refractory oxides. Because the aluminium oxide is not volatile, the heat of the reaction is not dissipated by the for- mation of a vapor, and a temperature high enough to melt iron or steel may be easily obtained for the welding of steel rails, per- foration of iron plates and similar purposes. The thermite process is also very useful for the production of chromium and other metals which it is difficult to obtain in other ways. A special advantage of the process, in some cases, is that the metals obtained in this way are free from carbon. Aluminium Chloride, A1C1 3 , is easily prepared by heating aluminium turnings in chlorine or in hydrochloric acid. It sublimes at 183 under atmospheric pressure and melts at 193 under increased pressure. Aluminium chloride dissolves in water with the evolution of considerable heat. If the solution is not too dilute, the addition of concentrated hydrochloric acid will cause a crystalline hydrate, A1C1 3 .H 2 O, to separate. On heating this hydrate it loses hydrochloric acid and aluminium oxide is left : 2 Aids + 3 H 2 O = A1 2 O 3 + 6 HC1 Aluminium chloride forms with many organic substances extremely reactive addition compounds which are used in a variety of syntheses. For these reactions water must be care- fully excluded, and the hydrate of the chloride cannot be used at all for such purposes. All of these facts indicate very clearly that the water of the hydrate is in a state of intimate chemical combination such as to greatly modify the relation which the chlorine and aluminium have in the anhydrous chloride. Solu- tions of aluminium chloride are partially hydrolyzed by water, and react strongly acid. Nearly all salts of trivalent or quadri- valent metals act in the same way. ALUMINIUM COMPOUNDS 499 Aluminium Fluoride, A1F 3 . Either aluminium or aluminium hydroxide dissolves easily in an aqueous solution of hydrofluoric acid, forming a supersaturated solution from which anhydrous aluminium fluoride slowly separates in small crystals. The double fluoride, Na 3 AlFe, is found in nature as the mineral cryolite, and was formerly used to dissolve the oxide for the electrolytic preparation of metallic aluminium. Aluminium Hydroxide, A1(OH) 3 , is precipitated from solutions of aluminium salts on the addition of an alkaline hydroxide. It forms a voluminous, gelatinous precipitate which dissolves either in solutions of strong acids or strong bases. For this reason it is called amphoteric, meaning that it has both basic and acid properties. It seems probable that in contact with an acid its hydroxyl combines with the hydrogen of the acid, while in contact with a base its hydrogen combines with the hydroxyl of the base. The compound may be considered, therefore, either as a triacid base or as a tribasic acid. It is, of course, very weak both as an acid and as a base. Salts in which it is the cation are hydrolyzed in solution and have an acid reaction : A1 2 (SO 4 ) 3 + 6 HOH ^T 2 A1(OH) 8 + 3 H 2 SO 4 Those salts in which the aluminium forms part of the anion are also hydrolyzed and have an alkaline reaction : Na 3 AlO 3 + 3 HOH ^ A1(OH) 8 + 3 NaOH When heated, aluminium hydroxide loses water and is con- verted into the oxide, A1 2 O 3 . It conducts itself very much as the silicic acids, however, losing a part of the water very easily at ordinary temperatures but requiring ignition at bright redness to expel the last portions. Aluminium Oxide, A1 2 O 3 , is found in nature as the mineral corundum. In its massive, not very pure forms, it is called emery. Next to the diamond it is the hardest mineral known and has been long used as an abrasive for grinding and polish- ing glass and metals. It has now been partly displaced for these uses by carborundum, SiC, which is much harder. Crystalline 500 A TEXTBOOK OF CHEMISTRY forms of corundum colored blue by some foreign substance are called sapphires, or other forms colored red by chromium are called rubies. The latter are now made artificially. A fused oxide, prepared in the electric furnace, is used, under the name of alundum, as an abrasive and also as a refractory material. Ignited aluminium oxide is insoluble in acids but may be brought into solution slowly by fusion with sodium pyrosulfate, Na 2 S 2 O7. Aluminium oxide is reduced by carbon at the tem- perature of the electric furnace, but cannot be reduced at lower temperatures. Before metallic aluminium was prepared by electrical processes, the anhydrous aluminium chloride used for the preparation of the metal was obtained by heating a mixture of the oxide with carbon in a current of chlorine : A1 2 O 3 + 3 C + 3 C1 2 = 2 A1C1 8 + 3 CO Aluminium Sulfate, A1 2 (SO 4 ) 3 .18 H 2 O, is prepared by the decomposition of clay with sulfuric acid. A more or less pure sulfate containing some ferric sulfate is extensively used under the name of " alum " for the clarification and purification of water. If a solution of the sulfate is mixed with a water con- taining calcium bicarbonate, an insoluble, gelatinous precipitate of aluminium hydroxide is formed : A1 2 (SO 4 ) 3 + 3 CaH 2 (C0 3 ) 2 = 2 A1(OH) 8 + 3 CaSO 4 + 3 CO 2 The precipitate collects fine particles of clay and also bacteria which are suspended in the water, and by suitable filtration clear water, nearly or quite free from disease germs, is obtained. The amount of aluminium sulfate required is so small that the calcium sulfate formed does not seriously increase the permanent hard- ness of the water. All of the aluminium added is removed in the filtration. Alums, M'M'"(SO 4 ) 2 .12H 2 O. By adding potassium sul- fate, K 2 SO 4 , to a solution of aluminium sulfate a compound having the composition KA1(SO 4 ) 2 .12H 2 O, and known since early times under the name of alum, is formed. It crystallizes in octahedra, which, with care, may be obtained in very perfect forms of large size. ALUMS, EARTHENWARE 501 Alum was formerly much used as a mordant in dyeing. The aluminium hydroxide formed by its hydrolysis combines with many coloring matters to form insoluble compounds called lakes. These compounds attach themselves strongly to the fibers of the cloth and cannot be removed by washing. Alu- minium sulf ate and the aluminates have largely displaced alum for such uses because the potassium sulfate is expensive and unnecessary. The potassium of alum may be replaced by ammonium or other univalent metals, the aluminium may be replaced by ferric iron or other trivalent metals and even the sulfate radical, SO4, may be replaced by the selenate radical, SeO 4 . This gives a great variety of alums, all of which crystallize in octahedra and are isomorphous. A crystal of any alum will grow in a supersaturated solution of any other. The following may be given as illustrations of the alums : Ammonium Alum NH 4 A1(SO 4 ) 2 . 12 H 2 O Ammonium Ferric Alum NH 4 Fe(SO 4 ) 2 . 12 H 2 O Chrome Alum KCr(SO 4 ) 2 . 12 H 2 O Rubidium Alum RbAl(SO 4 ) 2 . 12 H 2 O Brick, Earthenware, Porcelain. Aluminium silicate melts only at very high temperatures, but the presence of other com- mon metals, such as iron, calcium, magnesium, sodium or potas- sium, lowers the melting point. Ordinary clays contain com- pounds of these metals distributed as very fine particles through- out their mass, and when such clays are heated to a high temperature, the particles melt and cause the material to sinter together to a strong but very porous mass. In addition to this property of sintering without fully melting, the original clays become plastic when mixed with water, and in this condition may be molded into bricks or into the " biscuit " forms which furnish the basis of earthenware or porcelain. The plasticity of the clay seems to be closely connected with its colloidal character. In order to give them a surface which is smoother and im- 502 A TEXTBOOK OF CHEMISTRY pervious to water, articles of earthenware and porcelain must be covered with a glaze. Several methods of glazing are in use. One method is to throw salt into the furnace after the biscuit has been well burned. The sodium chloride reacts with the silica and alumina of the clay and the moisture of the air to form hydrochloric acid, which escapes, and a fusible silicate is formed, which melts and covers the surface with a glass. Other glazes are made from mixtures containing lead oxide. Some glazes of this type are not wholly insoluble in water, and England has enacted stringent laws requiring lead glazes to be highly insoluble when the articles are to be used to contain or cook food. Porcelains are usually glazed by the application of finely powdered feldspar and subjecting them to a temperature which causes it to melt and run into the surface. Ultramarine. Small quantities of a beautiful blue stone called lapis lazuli are found in nature. When powdered this stone gives a beautiful blue pigment which is not affected even by long exposure to the light. The mineral is so rare, however, that during the first years of the nineteenth century the pig- ment was sold to artists at $60 an ounce. In 1828 Gmelin dis- covered that the material can be made artificially by heating mixtures of clay, sodium sulfate, charcoal and sulfur. The artificial product is fully equal to the natural mineral for the uses to which it is applied and is now sold at a few cents per pound. By changing the method of manufacture, other com- pounds, ranging in color from reddish violet to bluish green, are also made. In spite of a very large amount of work devoted to the preparation and analysis of these compounds it has not been possible to assign definite formulas to them. The Rare Earths. These include the oxides of a number of elements all of which are characterized by being trivalent and by forming oxalates which are insoluble in dilute mineral acids. They resemble each other very closely in all their properties and in the types of compounds which they form. Their salts are isomorphous and do not as a rule differ greatly in solubility. Because of these slight differences they cannot, with the possible ALUMINIUM FAMILY. RARE EARTH METALS 503 exception of cerium, be separated from each other quantita- tively, and their preparation in a pure state requires a long-con- tinued series of fractional crystallizations or precipitations. In fact, some of them have not been obtained in a high state of purity. These elements may be divided into two groups : the Cerium group, including cerium, lanthanum, neodymium, pra- seodymium, samarium, europium and gadolinium ; and the Yttrium group, including terbium, dysprosium, holmium, yttrium, erbium, thulium, ytterbium, scandium and lutecium. The members of the cerium group may be separated roughly from those of the yttrium group by making use of the fact that the elements of the former form double sulfates with sodium, sulfates which are insoluble in a saturated solution of sodium sulfate, while the corresponding compounds of the yttrium group are soluble. * Scandium, Sc, 44.1. When Mendeleeff proposed the Periodic System of the elements in 1869, he predicted that several ele- ments not then known would probably be discovered in the future. Among these was an element which he called " eka- boron " which should have an atomic weight of about 44 and form compounds similar to those of aluminium. Ten years later Nilson found scandium among the elements found in gadolinite and euxenite, and shortly after Mendeleeff pointed out that this is in reality the " ekaboron " which he had pre- dicted. The hydroxide, Sc(OH) 3 , oxide, Sc 2 O 3 , sulfate, Sc 2 (SO 4 )3.6H 2 O, oxalate, Sc 2 (C 2 O 4 )3.6 H 2 O and other salts are known. * Yttrium, Y, 89, is found in gadolinite, xenotime and monazite. Its compounds resemble those of scandium. The oxide, Y 2 O 3 , chloride, YC1 3 , sulfide, Y 2 S 3 , phosphate, YPO 4 , and bromate, Y(BrO 3 ) 3 .9H 2 O, may be mentioned. * Lanthanum, La, 139, is the most positive of the rare earth metals. Its oxide, La^Os, combines with water, much as lime does, forming the hydroxide, La(OH) 3 , which turns litmus paper blue. The hydroxide also absorbs carbon dioxide from the air, forming the carbonate, La 2 (CO 3 )3. The oxalate, 504 A TEXTBOOK OF CHEMISTRY La 2 (C 2 O4) 3 .9 H 2 O, is difficultly soluble, as are the oxalates of all of the rare earth metals. * Ytterbium, Yb, 172, gives the sulfate, Yb 2 (SO 4 ) 3 .8 H 2 O, carbonate, Yb 2 (CO 3 ) 3 .4 H 2 O, acetate, Yb(C 2 H 3 O 2 ) 3 .4 H 2 O, oxalate, Yb 2 (C 2 O4) 3 .10 H 2 O, and many other salts. * Praseodymium, Pr, 140.6, and Neodymium, Nd, 144.3. In 1842 Mosander obtained from cerite (a silicate containing cerium, lanthanum, praseodymium and neodymium) an oxide of a metal to which he gave the name didymium. In 1885 Auer. v. Welsbach discovered that by a long series of crystalli- zations the double nitrate of the metal previously called didym- ium could be separated into two compounds, praseodymium ammonium nitrate, Pr 2 (NO 3 ) 3 .2 NH 4 NO 3 .4 H 2 O, and neodym- ium ammonium nitrate, Nd(NO 3 ) 3 .2 NH 4 NO 3 .4H 2 O. These two salts are isomorphous and do not differ very greatly in solubility. It was necessary, therefore, to repeat the crystalli- zation many hundreds of times in such a manner that the more soluble portions were transferred in one direction through the crystallizing dishes and the less soluble portions in the other direction, while portions of the same degree of separation were systematically united. This process of fractional crystalliza- tion has been much used in the separation of the metals of the rare earths. (See Auer. v. Welsbach, Monatshefte fur Chemie, 6, 477 ; Baxter and Chapin, J. Amer. Chem. Soc. 33, 5 ; James, ibid. 30, 182; 31, 913.) Typical compounds of praseodymium and neodymium are, Pr 2 O 3 , PrO 2 , Nd 2 O 3 , Pr 2 (SO 4 ) 3 .8 H 2 O, Nd 2 (SO 4 ) 3 .8 H 2 O, Pr 2 (C 2 O 4 ) 3 . 10 H 2 O, Nd 2 (C 2 O 4 ) 3 . 10 H 2 O, Pr(BrO 3 ) 3 . 9 H 2 O, Nd(Br0 3 ) 3 .9H 2 0. The oxalates of nearly all of the rare earth metals are very difficultly soluble and are often used as a means of separating these metals from those of other groups. The salts of praseodymium are green in color, those of neodymium are rose-colored, the two colors being comple- mentary very much as those of the salts of cobalt and nickel are. ALUMINIUM FAMILY. RARE EARTH METALS 505 * Samarium, Sm, 150.4, is found in samarskite, a columbate of metals of the rare earths, and was first partially separated from the " didymium " of that mineral by Lecoq de Boisbaudran. It is less basic than praseodymium and neodymium, and its double nitrate with magnesium is more easily decomposed by heat, a method sometimes used in separations. The oxide, Sm 2 O3, and solutions of its salts are yellow. It forms a chloride, SmCl 2 , in which the metal is bivalent, but in nearly all of its salts it is trivalent. A considerable number of salts have been prepared, such as samarium sulfate, Sn^SO^s-S H 2 O, the nitrate, Sm(NO3)3.6H 2 O and the carbonate, Sm 2 (CO 3 )3.3H 2 O. * Europium, Eu, 152, Gadolinium, Gd, 157.3, and Terbium, Tb, 159.2, form a group of weakly basic earths intermediate between the cerium earths on the one side and the yttrium earths on the other. They have been separated by tedious fractional crystallizations and finally identified by means of their spectra and determinations of their atomic weights. Europium oxide, Eu 2 O 3 , has a light rose color, and the sulfate, Eu 2 (SO4)3.8H 2 O, is also rose-colored. Gadolinium oxide, Gd 2 O 3 , and sulfate, Gd 2 (SO 4 ) 3 .8H 2 O, are white. Terbium oxide, Tb 2 Os, is also white ; but a higher oxide, possibly TbO 2 , but not yet obtained pure, is dark brown or black according to the method of preparation. * Holmium, Ho, 163.5, is obtained from euxenite, a columbate and titanate of the yttrium and cerium earths. The solubility of the double sulfate with ammonium seems to lie between those of yttrium and erbium (Holmberg). Apparently pure compounds have not yet been prepared. The metal is named from Stockholm. * Dysprosium, Dy, 162.5, has been obtained in its purest form by fractional crystallization of salts of ethyl sulfuric acid, H(C 2 H 5 )S0 4 . The dysprosium salt is Dy(C 2 H 5 SO 4 )3. The oxide, Dy 2 O 3 , is white. The bromate, Dy(BrO 3 )3.9 H 2 O, is yellow. * Erbium, Er, 167.7, is found among the yttrium earths from 506 A TEXTBOOK OF CHEMISTRY euxenite and other sources. The oxide, Er 2 3 , is rose-colored, as are also the sulfate, Er 2 (SO 4 ) 3 .8 H 2 O, and other salts. * Thulium, Tu, 168.5, is found in euxenite and other rare earth minerals. It seems to have been separated in the purest condition by the fractional crystallization of the bromates (James, J. Amer. Chem. Soc., 33, 1332). The oxide, Tu 2 O 3 , is white, with a faint green tint. The bromate, Tu(BrO 3 ) 3 . 9 H 2 O, separates in pale, bluish green prisms, isomorphous with the bromates of other rare earth metals. * Lutecium, Lu, 174, is one of the more recently discovered elements of this group and has been separated from the gado- linite earths. * Gallium, Ga, 69.9, was also predicted by Mendeleeff under the name of " eka-aluminium." Unlike the other metals of the group thus far described, it forms two classes of compounds, those in which it is bivalent and others in which it is trivalent. The chlorides are GaCl 2 and GaCl 3 ; the sulfates, GaSO 4 and Ga 2 (SO4) 3 .18 H 2 O. The former is oxidixed by potassium perman- ganate as ferrous sulfate is. The alum, NH 4 Ga(SO 4 ) 2 .12H 2 O, is isomorphous with ordinary alum. Indium, In, 114.8, was discovered by means of the blue line of its spectrum by Reid and Richter shortly after the methods of spectrum analysis had been developed by Bunsen and Kirchoff. The analysis of its oxide gave about 76 parts of indium for 16 parts of oxygen, and an atomic weight of 76 was at first assigned to the element. But Mendeleeff pointed out that this would place it between arsenic and selenium, where there is no vacant place in the Periodic System, and also that its properties did not agree with such a position in the table. He suggested, there- fore, that the formula of the oxide is In 2 O 3 and the atomic weight 114. The determination of the specific heat gave the value 0.056. This points to an atomic weight of 6.2/0.056 = 110, which agrees fairly well with an atomic weight of 114 but would not agree at all with the value of 76. The preparation of an ammonium alum, NH 4 In(SO 4 ) 2 .12H 2 O, soon after this, gave further support for the accepted formula. ALUMINIUM FAMILY. RARE EARTH METALS 507 Indium is a soft, white metal, which melts at 155. It gives three chlorides, InCl, InCl 2 , InCl 3 , but the compounds in which it is trivalent are most stable and best known. Thallium, Tl, 204, was also discovered by means of the spec- troscope. Crookes found it in 1861 in the slimes from sulfuric acid made at Tilkerode in the Harz. He named it thallium from the Latin word thallus, meaning a young twig, because of a brilliant green line in its spectrum. Metallic thallium is a bluish white, soft metal, somewhat resembling lead. Thallium forms thallous compounds in which it is univalent, and thallic compounds in which it is trivalent. Of the former, thallous oxide, T^O, thallous chloride, T1C1, thallous hydroxide, T1OH.H2O, and thallous sulfide, T1 2 S, may be mentioned. The last is a black precipitate nearly insoluble in acetic acid but soluble in mineral acids. Thallous iodide, Til, is also very difficultly soluble. Among the thallic compounds are thallic chloride, T1C1 3 .H 2 O, thallic nitrate, T1(NO 3 ) 3 .8 H 2 O, and thallic sulfide, T1 2 S 3 . EXERCISE Assuming the specific heat of aluminium oxide as 0.217, that of iron as 0.15, and the heat of fusion of iron as 23 calories per kilogram, what is the maximum temperature which could be reached by the reaction of a thermite consisting of ferric oxide and metallic aluminium ? CHAPTER XXIX TIN AND LEAD IT has been pointed out (p. 361) that tin and lead belong to the carbon group of the Periodic System and that each gives an oxide (SnC>2 and PbCy resembling carbon dioxide, CO2, and silicon dioxide, SiO2, in formula and in some other properties, especially in their acidic character. Both of these elements are clearly metals rather than non-metals in most of their properties. Tin (Sn, 119). Occurrence, Metallurgy. Tin is rather re- markable in that, although it is an element which is found in sufficient quantity so that it is a common metal for household and commercial use, there are only a very few localities in the world where its ores can be profitably mined. One of the oldest of these is Cornwall in England, where tin has been obtained for nearly or quite twenty centuries and which furnished a large part of the tin used in the world until comparatively modern times. The world's supply of tin for one year is approximately 80,000 tons, and this comes almost entirely from Banca and the East India islands, Tasmania, Bolivia and Cornwall. The only important ore is cassiterite, stannic oxide, SnO 2 . This is found sometimes in veins, sometimes as a heavy gravel, called " stream tin." The metallurgy is comparatively simple, consisting in the reduction of the oxide by means of charcoal or coal. On ac- count of its value, the recovery of tin from tin scrap has also assumed considerable commercial importance. Several methods are in use, one of the best being the treatment of the scrap with dry chlorine gas, which converts the tin into stannic chloride, SnCl 4 , but leaves the iron comparatively unattacked. Stannic chloride is volatile and can be easily separated. 508 TIN 509 Uses of Tin. Alloys. Tin Plate. Tin is not affected by dry or moist air or by water, even at the boiling point or higher. It is, for this reason, the most suitable of the cheaper metals for the tubes of condensers to be used in the ^reparation of distilled water. The principal use of the metal is in the manu- facture of tin plate sheet iron which has been covered with a thin coating by dipping the carefully cleaned metal in a vat of melted tin. Tin is more electropositive than iron, hence when the two metals are in contact with water or a dilute acid the tendency is for the iron to corrode while the tin is protected. For this reason tin vessels rust through rapidly as soon as the iron is exposed at any point exactly contrary to the conduct of iron coated with zinc (p. 482). " Terne plate," which is used for roofing purposes, is covered with an alloy of lead and tin, the lead being used because it is very much cheaper than tin. Solder is an alloy of lead and tin, used to join pieces of tin plate in making culinary vessels of all kinds. Common solder (contains equal parts of the metals, but fine solder, containing more tin, and coarse solder, containing more lead, are often used. Tin was formerly much used in gun metal and bell metal and is still used in statuary and ornamental bronzes, which are alloys of tin and copper, usually with a little lead and zinc. Tin, antimony and lead are the principal constituents of Britannia metal, pewter and Babbitt metal. Copper and other metals are sometimes added. Tin melts at 231.9. It is convenient to remember the melting points of tin (232), lead (327) and zinc (419) as about 100 apart, with tin the lowest and zinc the highest. Tin boils at a high temperature, but the boiling point has not been determined. The specific gravity is 7.30. The metal oxidizes slowly, when heated above its melting point in the air, to stannic oxide, SnO 2 . It dissolves as stannous chloride, SnCl 2 , in concentrated hydrochloric acid and is con- verted by nitric acid into a mixture of stannic and metastannic acids, SnO 2 .H 2 O, which is insoluble in an excess of the acid or in water. 510 A TEXTBOOK OF CHEMISTRY Compounds of Tin. Tin forms s tan nous compounds, such as SnCl 2 , in which it is bivalent, and stannic compounds, as SnCU, in which it is quadrivalent. In the former it is rather strongly basic and metallic, in the latter much weaker as a base and in some of the compounds distinctly acidic. * Stannous Oxide, SnO. If a solution of potassium carbonate, K 2 CO3, is added to a solution of stannous chloride, SnCl 2 , a white precipitate having the composition 2 SnO.H 2 O is formed ; but this loses water on heating the solution, especially if a little alkali is present, much as cupric hydroxide, Cu(OH)2, does (p. 431), and is changed to black insoluble stannous oxide, SnO. The original precipitate is amphoteric, like aluminium hydroxide, Al(OH)s, and dissolves either in acids or in alkalies. Stannous Chloride, SnCl2.2 H 2 O, is easily obtained by dis- solving tin in concentrated hydrochloric acid. The anhydrous salt can be prepared by heating this hydrate in a stream of hydrochloric acid. It boils at 606. Stannous chloride is frequently used in the laboratory as a reducing agent because of its strong tendency to take up chlorine or other elements and pass over to the stannic form. A solution ' of stannous chloride to which an excess of sodium hydroxide has been added, forming sodium stannite, NaHSnO 2 or Na 2 SnO 2 , is also a powerful reducing agent. Stannous chloride and several other compounds of tin are ex- tensively used as mordants in dyeing. Stannous Sulphide, SnS, is precipitated from acid solutions of stannous salts as a dark brown, almost black compound. It does not dissolve in colorless ammonium sulfide, (NH 4 ) 2 S, but the yellow ammonium or sodium sulfides, which contain poly- sulfides, (NH4) 2 S 2 , etc., dissolve it as ammonium or sodium sulfostannate, (NH 4 ) 2 SnS 3 or Na 2 SnS 3 . From such a solution acids precipitate yellow stannic sulfide, SnS 2 . Stannic Oxide, SnO 2 , is prepared by the oxidation of tin in air at a high temperature or by treating tin with nitric acid and igniting the mixture of stannic acids which is formed. It is also found as the crystalline mineral, cassiterite, in nature. STANNIC ACIDS 511 Stannic oxide does not dissolve in the melted silicates which form glass, and it has been sometimes used for the manufacture of white, opaque glass, but less expensive materials are usually employed. Stannic Acids. Tin resembles silicon in that several acids are derived from the same anhydride, SnO 2 , just as there is a long list of silicic acids derived .from silicon dioxide, SiO2. Berzelius discovered in 1817 that the stannic acid, obtained by precipitation from a solution of stannic chloride, SnCU, is very different in its properties from metastannic acid, which is formed by treating tin with nitric acid, and his study of these compounds led him to propose the word isomer to designate a compound having the same composition as some other compound which has different properties. At that time the compounds which we now call anhydrides were called acids, and the compounds which he considered isomeric were the anhydrides of the two acids rather than the compounds which we should now call stannic and metastannic acids. Later investigations have shown that neither the free acids, when dried in the air, nor their salts are isomeric as the term is used to-day. This will be clear from the following table. A third acid, parastannic acid, which was also discovered by Berzelius, is included. Stannic and metastannic acid are isomeric when dried in a vacuum. NAME FORMULA DRIED IN THE AlR FORMULA DRIED IN A VACUUM FORMULA OF POTASSIUM SALT CHLORIDE FORMED WITH HC1 Stannic acid H 2 SnOs.H 2 O EhSnOs K^nOa-EWD SnCU Metastannic acid . . . HzSnsOn.Q H 2 O H 2 Sn 5 On~4 H 2 O K 2 Sn 6 On.4 H 2 O SmOoCh .4H 2 O Parastannic acid . . . HzSnsOn.? HjO H 2 SnsOii.2 H 2 O K 2 Sn6On.2or3H 2 O SniiOCI 2 .2H 2 O Stannic Acid, H 2 SnO 3 .H 2 O, or H 4 SnO 4 , is obtained by pre- cipitating a solution of stannic chloride, SnCU, with ammonia or with calcium carbonate. It dissolves easily in strong acids or in alkalies. From its solution in alkalies it is reprecipitated 512 A TEXTBOOK OF CHEMISTRY i by acids. On drying it is partly changed to metastannic acid, and a failure to understand this has led to much confusion in the literature. Metastannic Acid, H 2 Sn 5 On.9 H 2 O or (H 4 SnO 4 )5, is the prin- cipal product formed by the action of warm nitric acid on tin. If the mixture obtained in this way is dissolved in a little sodium hydroxide, the addition of an excess of the alkali will cause the precipitation of sodium metastannate, while the sodium stannate will remain in solution. Metastannic acid is insoluble in nitric acid or sulfuric acid. When treated with concentrated hydrochloric acid it forms a chloride, Sn 5 O 9 Cl 2 .4 H 2 O, which dissolves in water but is re- precipitated by concentrated hydrochloric acid. Solutions of stannic chloride, SnCU, which have stood for some time, con- tain this compound, formed by hydrolysis and rearrangement or condensation. It is properly named metastannyl chloride. * Parastannic Acid, H 2 Sn 5 Ou.7 H 2 O, was obtained by Ber- zelius by heating metastannic acid with water at 100. It is quite similar to metastannic acid. Stannic Chloride, SnCl 4 , is a volatile liquid which boils at 114 and fumes strongly in the air, owing to its hydrolysis by the moisture of the air and the escape of hydrochloric acid. It dissolves in water and forms several hydrates, but it seems pretty certain -that these contain compounds of the same general char- acter as metastannyl chloride (see above) rather than hydrates of stannic chloride, as that would ordinarily be understood. But our knowledge of the structure of hydrates in general is still very imperfect. Stannic Sulfide, SnS2, separates as a yellow, amorphous pre- cipitate when hydrogen sulfide is passed into an acid solution of a stannic salt. It dissolves in concentrated hydrochloric acid, resembling antimony and differing from arsenic in this respect. It also is not precipitated from and dissolves in solu- tions of oxalic acid from which antimony can be precipitated as the sulfide, Sb 2 Ss. Stannic sulfide dissolves in ammonium sulfide as ammonium sulfostannate, (NH 4 ) 2 SnS 3 . FIREPROOFING COTTON GOODS 513 Fireproofing of Cotton Goods. Many fatal accidents occur every year from the burning of clothing, and serious accidents have occurred from the burning of curtains and fabrics in theaters. A variety of substances have been used to render fabrics less inflammable. One of the best of these is stannic oxide. Pro- fessor Perkin describes its application as follows : " The flan- nelette (or other material) is run through a solution of sodium stannate of approximately 45 Tw (sp. gr. 1.225) in such a man- ner that it becomes thoroughly impregnated. It is then squeezed to remove the excess of stannate solution, passed over heated copper drums to thoroughly dry it, after which it is run through a solution of ammonium sulfate of about 15 Tw, and again squeezed and dried. Apart from the precipitated stannic oxide, the material now contains sodium sulfate and this is removed by passage through water; the material is then dried and subjected to the ordinary processes of finishing. A long series of trials, carried out under the most stringent condi- tions, have conclusively proved that material, subjected to this process, is permanently fireproof. No amount of washing with hot soap and water will remove the fireproofing agent, or in other words, the property of resisting flame lasts so long as the material itself lasts." (Address before the Inter- national Congress of Applied Chemistry, New York City, Sep- tember, 1912.) Lead, Pb, 207.1. Occurrence, Metallurgy. Lead is most often found as the sulfide, PbS, in the form of galena, a heavy black mineral which crystallizes in cubes having a bright, metallic luster. Galena is usually associated with other minerals, es- pecially with sphalerite, ZnS, and pyrites, FeS 2 , and with a gangue of quartz, SiC>2, fluorite, CaF 2 , barite, BaSCh, or calcite, CaCO 3 . Lead is obtained from the ore by roasting it to convert a part of the sulfide to the oxide or sulfate : PbS + 3 O = PbO + SO 2 PbS + 4 O = PbS0 4 514 A TEXTBOOK OF CHEMISTRY If the mixture of sulfide with oxide or sulfate is heated with exclusion of air or in a reducing atmosphere, the compounds mutually reduce each other and sulfur dioxide escapes : 2 PbO + PbS = 3 Pb + SO 2 PbS0 4 + PbS = 2 Pb + 2 SO 2 This process is most suitable for very pure ores. Less pure ores, which are often reduced for the silver and other metals which they contain rather than for the lead, are usually reduced in a blast furnace (p. 541) by the combined action of coal or coke and iron. The iron combines with the sulfur of the galena, reducing it to metallic lead : PbS + Fe = FeS + Pb The recovery of silver, which is usually present in crude lead, has been discussed in a previous chapter. Properties and Uses of Lead. Alloys. Lead is the heaviest of the cheaper metals, having a specific gravity of 1 1 .34, which is even higher than that of silver. It is because of this property, and also because of the ease with which lead can be melted and cast, that it is used for bullets and shot. Lead melts at 327.4. It has a bright white luster when freshly cut, but tarnishes quickly in the air. Lead dissolves easily in nitric acid but only slowly and to a slight extent in hydrochloric acid. Even hot sulfuric acid scarcely attacks the metal till it has the specific gravity of 1.72. A more concentrated acid dissolves lead sulfate and attacks metallic lead strongly. Lead is so soft that it can be pressed through a die into the form of tubing, by means of hydraulic pressure. Such tubing is used for waste pipes from sinks and for similar purposes and is liked by plumbers because of the ease with which it can be worked. It is not suitable for pipes to convey drinking water, because a little of the lead is liable to dissolve, and all soluble lead compounds are very poisonous. Lead has the further, very dangerous, property, that it acts as a cumulative poison so that minute quantities taken daily for some weeks or months may finally produce fatal results. LEAD 515 Lead has such slight tenacity that it cannot be drawn into wire, in spite of its softness. It can be rolled into sheets and beaten into thin foil. Foil made from the alloy with tin has sometimes been substituted for pure tin foil to wrap around articles of food, but such use is strongly condemned because of the poisonous character of the lead. The more important alloys of lead, solder, Britannia metal, pewter, Babbitt and other antifriction metals, type metal, stereotype metal and fusible alloys used for safety fuses, have been mentioned in previous chapters. Oxides of Lead. There are three definite, well-characterized oxides of lead : lead monoxide, or litharge, PbO, lead plumbate, usually called red lead or minium, Pb 3 O4, and lead dioxide, 1 or plumbic anhydride, PbO2. Three other oxides, Pb2O, PbsOy and Pb2Oa, have been described by various authors, but there is considerable doubt whether these represent definite compounds or not. The evidence in favor of the existence of the oxide, Pb 2 O 3 (or PbPbO 3 , lead metaplumbate), is better than that for the other two. Lead Monoxide or Litharge, PbO, is readily formed by exposing lead at a red heat to the action of the air. At that tempera- ture the melted film of litharge constantly flows to the side, ex- posing a fresh surface to oxidation. The melted litharge is collected, allowed to solidify, and is then ground to a fine powder for the market. It forms a buff-colored powder used in the manu- facture of " boiled " linseed oil and flint glass, and for other purposes. For the manufacture of red lead, metallic lead is first oxidized at a lower temperature, such that the oxide formed does not melt. This oxide is then heated to dull redness, but below the melting point of litharge, with free access of air. The chemical character of red lead is most clearly shown by the action of dilute nitric acid upon it. This dissolves two 1 Often called lead peroxide. It is better to restrict the designa- tion peroxide to compounds having a structure similar to that of hydrogen peroxide, H O O H. 516 A TEXTBOOK OF CHEMISTRY thirds of the lead and leaves a dark brown residue of lead di- oxide, or plumbic anhydride, PbO 2 : Pb 3 O 4 + 4 HNO 3 = 2 Pb(NO 3 ) 2 + PbO 2 + 2 H 2 O This reaction shows that two thirds of the lead is basic and one third acidic in character, or, in other words, that red lead is a lead salt of plumbic acid, H4PbO4. The relation is clearer if we write the formula Pb 2 PbO 4 . When this is treated with nitric acid, the two lead atoms which form the positive ions, Pb ++ , of the salt, are easily exchanged for the hydrogen of the acid, but the plumbic acid, H 4 PbO 4 , which results, is unstable and decomposes at once to water and plumbic anhydride or lead dioxide, PbO 2 , just as carbonic acid, H 2 CO 3 , decomposes to carbon dioxide and water. Red lead is used as a pigment, as an oxidizing agent in glass manufacture and with linseed oil as a lute in plumbing. As litharge is oxidized to lead plumbate, Pb 2 PbO 4 , a mixture of lime, CaO, and litharge may be readily oxidized at low redness to calcium orthoplumbate, Ca 2 PbO 4 . Dilute acids decompose this with the separation of lead peroxide. Storage Batteries. A storage battery which has been charged contains two kinds of lead, one of which has been more or less completely changed to lead dioxide, while the other consists of spongy, metallic lead. During the discharge the electrons escape to the connecting wire from the metallic lead, Pb, chang- ing it to lead ions, Pb ++ , which combine with sulfate ions of the solution, forming insoluble lead sulfate, PbSO 4 : Pb + SO 4 " = PbSO 4 + 2- The electrons pass through the plate and connections to the other plate. At the other plate the two electrons combine with the lead of the lead dioxide, reducing it from the quadrivalent to the bivalent form, resulting in the formation of lead sulfate and the liberation of a sulfate ion, SO 4 ~~, from the sulfuric acid: 2- + PbO 2 + 2 H 2 SO 4 = PbSO 4 + 2 H 2 O + SO 4 ~ STORAGE BATTERIES 517 or Pb+ +++ O--O--- + 2- = Pb++O-- -f- O" O + H 2 SO 4 = H 2 + SO 4 = and Pb ++ O-" + H 2 S0 4 = Pb ++ SO 4 = + H 2 O During the discharge there is a difference of potential of about two volts between the two plates, and the passage of the current develops a very considerable amount of electrical energy. In charging the battery the reverse operations take place. At the cathode, which is connected with the negative pole of the dynamo, electrons combine with lead ions, Pb +4 ", reducing them to metallic lead, Pb, and leaving tWe sulfate ions, SO 4 ~~, of the lead sulfate free to pass into solution. Pb ++ SO 4 + 2- = Pb + SO 4 ~- At the anode the bivalent lead ions, Pb ++ , lose two electrons, giving tetravalent lead ions, Pb ++++ , which combine momentarily with another sulfate ion, SO 4 ~~~, to form lead tetrasulfate, Pb(SO 4 ) 2 . The lead tetrasulfate is at once hydrolyzed to lead peroxide, PbO 2 , and sulfuric acid. Pb + +SO 4 -" + SO 4 ~- = Pb ++++ (SO 4 --) 2 + 2- Pb(SO 4 ) 2 + 2 H 2 O = Pb0 2 + 2 H 2 SO 4 It will be noticed that in both processes sulfate ions, SO 4 = , are discharged at one plate and enter into combination with the lead, while at the other plate sulfate ions pass into solution. In charging the battery the discharged sulfate ions combine with the lead of lead sulfate, PbSO 4 , forming the tetrasulfate, Pb(SO 4 ) 2 . In discharging the discharged sulfate ions combine with the metallic lead, forming lead sulfate. At the same time sulfate ions must, of course, migrate through the solution in one direction in charging, in the opposite direction in discharging. The charging is, of course, accompanied by an absorp- tion of energy. Practically, a very high efficiency can be se- cured, the energy obtained during the discharge approaching closely to that absorbed in charging. Since the sulfate radicals are mostly in the form of sulfuric acid in the charged battery 518 A TEXTBOOK OF CHEMISTRY and in the form of lead sulfate in the battery after discharge, the amount of sulfuric acid in the electrolyte, which can be easily determined by a hydrometer, furnishes a pretty close indication of the condition of the cell. Lead dioxide when warmed with hydrochloric acid gives at first lead tetrachloride, PbCU, but this is unstable and decom- poses into lead chloride, PbCl 2 , and chlorine, C1 2 , exactly as manganese tetrachloride does (p. 101). It is noteworthy that this is entirely different from the conduct of barium peroxide, BaO 2 , or sodium peroxide, Na 2 O 2 , either of which gives hydrogen peroxide, H 2 t) 2 , with hydrochloric acid. (What does this dif- ference in conduct indicate as to the structure of these three oxides?) Lead Sulfide, PbS, is found in nature as the mineral galena and is formed as a black precipitate by the action of hydrogen sulfide on a solution of a lead salt in a dilute acid. It is not precipitated in the presence of much hydrochloric acid, or of much of any other strong acid. It dissolves easily in nitric acid. Lead Chloride, PbCl 2 , forms as a white, crystalline precipi- tate on adding hydrochloric acid to a solution of almost any soluble salt of lead. It dissolves in 125 parts of water at 18 and in 30 parts of boiling water. It is less soluble in dilute hydrochloric acid than in pure water, but is more soluble in concentrated acid, doubtless because of the formation of a complex compound with the acid, such as chloroplumbous acid, H 2 PbCl 4 . Lead Tetrachloride, PbCU. By dissolving lead peroxide, PbO 2 , in cold, concentrated hydrochloric acid a solution con- taining lead tetrachloride, or more likely chloroplumbic acid, H 2 PbCle, is obtained. On the addition of ammonium chloride, ammonium chloroplumbate, (NH 4 ) 2 PbCl 6 , separates. If this salt is dissolved in cold concentrated sulfuric acid, hydrochloric acid escapes and lead tetrachloride separates below the acid as a heavy yellow liquid. It is quite unstable, decomposing readily into lead chloride, PbCl 2 , and chlorine. It is hydrolyzed by LEAD SALTS 519 water to lead dioxide and hydrochloric acid. With a little hydrochloric acid it gives a yellow compound, chloroplumbic acid, H 2 PbCl 6 . A considerable number of salts of this acid are known. The ammonium salt is mentioned above. These salts are usually called double salts of lead tetrachloride with other chlorides, and are frequently written in the form PbCU.2 KC1 instead of K 2 PbCl 6 . Lead Sulfate, PbSO 4 , may be prepared by the precipitation of any soluble salt of lead with dilute sulfuric acid. It may also be obtained by roasting lead sulfide at a moderate tem- perature. The compound prepared in the latter manner has been used in America to a limited extent as a pigment in place of the ordinary white lead described below. It is a difficultly soluble salt, but is distinctly more soluble than barium sulfate. It dissolves easily in a solution of ammonium acetate. Lead Nitrate, Pb(NO 3 ) 2 , is a very easily soluble salt which may be prepared by dissolving either lead or litharge in dilute nitric acid. Either lead nitrate or lead acetate may be used for the preparation of the insoluble salts of lead, especially of the chromate, PbCrO 4 . Lead Acetate or Sugar of Lead, Pb(C 2 H 3 O 2 )2.3 H 2 O, is an easily soluble salt prepared by dissolving litharge in acetic acid. Sugar of lead has sometimes been used in hair dyes, but its use in this way is considered dangerous and liable to cause paralysis. Basic Lead Acetates, Pb(C 2 H 3 O 2 )OH and Pb(C 2 H 3 O 2 ) 2 . 2 Pb(OH) 2 , are formed by dissolving litharge, PbO, in a solu- tion of lead acetate. Such a solution is used to clarify dark- colored sugar solutions to prepare them for determinations with the polarimeter. Lead Carbonate, PbCO 3 . The normal salt may be precipi- tated by adding sodium carbonate to a solution of lead acetate. It is so difficultly soluble that it can also be precipitated by passing carbon dioxide through a solution of lead acetate, and lead acetate gives a turbid solution with ordinary distilled water because of its formation. 520 A TEXTBOOK OF CHEMISTRY Basic Lead Carbonate, or White Lead, Pb 3 (CO 3 )2(OH) 2 or 2 PbCO 3 .Pb(OH) 2 , has been manufactured for many centuries for use as a pigment. The principal process used, known as the " Dutch process," has been scarcely changed in principle for a very long time. Plates, or " buckles," of lead about one eighth of an inch thick and five and one half inches in diameter are cast in the form shown in Fig. 104. These plates are packed in pots, Fig. 105, having about 250 cc. of dilute acetic acid or vinegar in the Fig. 104 Fig. 105 t bottom. These pots are then packed in layers with alternate layers of spent tanbark until a large room is filled. The room is left to itself for about three months. The combined action of the vapors of acetic acid and air on the lead plates causes them to corrode superficially, and the basic acetate formed is changed to carbonate by the carbon dioxide which comes from the fermentation of the tanbark. As the carbonate is formed some acetic acid is continuously liberated, and this, with the air, carries on the corrosion till the " buckles " are almost completely changed to basic carbonate. At the end of three months the pots are emptied, the white lead is finely ground and separated from particles of lead an4 coarse particles of material by bolting and lixiviating with water. The fine powder is then dried and intimately incorporated with linseed oil for WHITE LEAD 521 the market. Another process is to mix the moist powder directly with linseed oil. Owing to the relation between the surface tension of white lead toward water and that toward linseed oil, the latter is able to displace the water and joins with the white lead to form a paint which is practically identical in composition and properties with the pigment prepared from the dry powder. Owing to the poisonous character of lead com- pounds, workmen in white lead factories and painters often suffer from a painful and sometimes fatal disease, called lead colic. Stringent laws have been passed by some states for the protection of the workmen from inhaling the dust and from poisoning in other ways. Some manufacturers spend large sums of money to protect their workmen from the danger of poisoning. White lead depends for its value on the fact that it is an amorphous, very fine, opaque white powder, the opacity being much greater than that of barium sulfate, which is sometimes used as an adulterant or substitute. The fact that water will not wet it when it is in contact with linseed oil is also a factor of prime importance. White lead is blackened by hydrogen sulfide and for that reason is less suitable than zinc white and lithopone for use in chemical laboratories or in localities where it is subjected to the action of sewer gas. CHAPTER XXX VANADIUM AND CHROMIUM GROUPS Group V. Vanadium, columbium (or niobium) and tantalum are the elements of Group V, which alternate with phosphorus, arsenic, antimony and bismuth. They are usually classed as rare elements, but vanadium and tantalum have acquired some industrial importance. Each forms a pentoxide, corresponding to P2O5, and salts of acids corresponding to themetaphosphates, MPO 3 , pyrophosphates, M 4 P 2 O7, and orthophosphates, M 3 PO4. Salts of the form MgV^C^ are also known for each. Vanadium, V, 51.0, is very widely diffused in nature, being found in small amounts in almost all clays and massive rocks. The most important mineral is vanadinite, Pb 5 (VO 4 ) 3 Cl, which corresponds in composition to apatite, Ca 5 (PO 4 )3F. Vanadium is a silvery white metal having a specific gravity of 5.5 and melting at about 1720. It is used as an addition to steel, increasing its hardness, malleability and tensile strength. Ferrovanadium is an alloy with iron, which is much more easily prepared than the pure metal. It is the commercial form used for addition to steel. * Vanadium forms compounds in which it is bivalent, tri- valent, quadrivalent and quinquivalent. The following may be mentioned : vanadous chloride, VC12, vanadous sulfate, VSO 4 .7 H2O, vanadous sulfide, VS, vanadic chloride, VC1 3 , vanadic sulfide, V2S 3 , vanadic sulfate, V2(SO 4 ) 3 , vanadium alum, KV(SO 4 ) 2 .12 H 2 O, vanadium tetrachloride, VC1 4 , vanadium pentachloride, VCls, vanadium oxychloride, VOC1 3 , sodium orthovanadate, Na 3 VO 4 .12H 2 O, sodium pyrovanadate, Na 2 V 2 O 7 .18H 2 O, sodium metavanadate, NaVO 3 .2H 2 O. More complex vanadates and other complex compounds of a great variety of forms have been prepared. 522 COLUMBIUM. TANTALUM 523 * Columbium, Cb,93.5 (or Niobium, Nb). In 1801 Hatchett discovered a new element in a mineral from Haddam, Con- necticut. He called the mineral columbite (from Columbia, the poetical name for America), and the element columbium. It is probable that the compounds which he prepared contained both columbium and tantalum, and the two elements were first clearly separated and characterized by H. Rose in 1844. Rose either overlooked or ignored the discovery of Hatchett and called the two metals tantalum and niobium. Columbite, the mineral in which columbium was discovered, is a ferrous metacolumbate, Fe(CbO 3 ) 2 , containing ferrous metatantalate, Fe(TaO 3 ) 2 , in isomorphous mixture. The formula is more properly written, Fe((Cb,Ta)O 3 ) or, in the oxide form, which is most frequently used by mineralogists, FeO((Cb 2 ,Ta 2 )0 5 ). Elementary columbium is still more metallic in its properties than vanadium. It has a specific gravity of 7.4 and melts at about 2200. The following are typical compounds : colum- bium trichloride, CbCl 3 , columbium pentafluoride, CbFs, colum- bium oxyfluoride, CbOF 3 , columbium pentoxide, Cb 2 O5, mag- nesium orthocolumbate, Mg 3 (CbC>4) 2 , calcium pyrocolumbate, Ca 2 Cb 2 O 7 , potassium hexacolumbate, 4 K 2 O.3 Cb 2 O 5 .16 H 2 O or K 8 Cb 2 Og.l6 H 2 O. Many complex columbates ana other com- plex compounds are known. * Tantalum, Ta, 181.5. In 1902 Ekeberg examined a mineral from Finland which closely resembles columbite and gave to it the name tantalite. He called the element which it contains tantalum, but the compounds which he prepared were doubtless mixtures containing both columbium and tantalum. The mineral which he studied always contains both elements, and it is most properly called columbite when the columbium is in excess and tantalite when there is more of the tantalum. As stated above, H. Rose first distinguished sharply between the two elements and cleared up the confusion of the earlier workers. Tantalum is a bright metal somewhat resembling platinum in appearance. It is ductile and can be drawn into fine wire. 524 A TEXTBOOK OF CHEMISTRY As the melting point is 2850 it is suitable for the filaments of electric lights, and was used for a short time in that way, but was quickly displaced by tungsten. The specific gravity of the metal is 16.5. Among the compounds of tantalum are the potassium fluo- tantalate, 2KF.TaF 5 or K 2 TaF 7 , by means of which tantalum can be best separated from columbium (E. F. Smith), tantalum pentachloride, Tads, tantalum pentoxide, Ta 2 Os, and a series of tantalates. Group VI. Chromium, molybdenum, tungsten and uranium alternate with sulfur, selenium and tellurium of Group VI, exactly as the three elements last considered alternate with the elements of the phosphorus family. The elements form compounds in which they appear bivalent, trivalent and sexivalent, especially, but also some in which they are quadrivalent. They resemble the sulfur family in such compounds as potassium chromate, K 2 CrO 4 , potassium dichromate, K 2 Cr 2 O 7 , and ammonium molybdate, (NH 4 ) 2 MoO 4 . Chromium, Cr, 52.0, is found chiefly as ferrous chromite, FeO.C^Os or Fe(CrO 2 ) 2 , which is known as chromite or chrome iron ore, a black mineral isomorphous with magnetite, Fe 3 O4 or (FeO.Fe 2 O 3 ). Chromium is also found as lead chromate, PbCrO 4 , and was first discovered in that mineral by Vauquelin in 1797. The name was given because of the colored compounds which it forms. Metallurgy, Uses. Chromium is now prepared by Gold- schmidt's thermite process (p. 497) by igniting a mixture of chromic oxide, Cr 2 O 3 , and aluminium. Metallic chromium is a white, crystalline, extremely hard metal as hard as corundum. It has a specific gravity of about 7.0. It melts at 1520 and boils at 2200. It is used as an addition to steel, making it extremely hard and resistant to the penetration of projectiles when used for armor plate. Such steel requires very careful heat treatment to bring out its best properties. Alloys with nickel or cobalt resist the action of acids and are proving useful for special purposes. CHROMIUM 525 Chromous Chloride, CrCl2.4 H 2 O. A light blue solution containing chromous chloride is easily prepared by the action of zinc and hydrochloric acid on a solution of chromic chloride, CrCl 3 , or of chromic anhydride, CrO 3 . Air must be carefully excluded, as air or oxygen, in the presence of hydrochloric acid, changes the compound back to green chromic chloride. The solid chromous chloride is white. Chromic Oxide, Cr 2 O 3 , is a green powder formed by igniting the hydroxide. A more or less pure chromic oxide or hydroxide is prepared in a variety of ways and used as a pigment under the name of chrome green or Guignet's green. The latter con- tains a little boric acid. A green pigment which is made by mixing chrome yellow, PbCrO 4 , and Prussian blue, Fe 4 (FeC6N 6 ) 3 , is also often sold and used under the name of " chrome green." Chromic Hydroxide, Cr 2 O 3 .4 H 2 O. From analogy with the formula of the chloride, the formula Cr(OH)3 is often given for the hydroxide. The compound formed by precipitating a chromic salt with ammonium hydroxide has the composition Cr 2 O 3 .4 H 2 O, however, when it is dried in a vacuum. A com- pound, or mixture, having the composition Cr(OH) 3 has been obtained in some cases, but it is not the usual composition of the hydroxide. Chromic hydroxide dissolves to some extent in alkalies, but does not show as marked acidic properties of aluminium hydroxide. A number of chromites have been pre- pared, however, and ferrous chromite, Fe(CrO 2 ) 2 , has been men- tioned above as the most important ore of chromium. Chromic Chloride, CrCl 3 , was formerly prepared by heating a mixture of chromic oxide, Cr 2 O 3 , and carbon in a current of chlorine. It sublimes in beautiful reddish pink or violet leaflets. These are almost insoluble in water, but dissolve easily to a green solution in water containing a trace of chromous chloride, CrCl 2 . Hydrates of Chromic Chloride. A solution containing chromic chloride is easily prepared by the reduction of chromic anhydride or a solution of a chromate : 2 CrO 3 + 6 HC1 + 3 C 2 H 5 OH = 2 CrCl 3 + 3 C 2 H 4 O + 6 H 2 O Alcohol Aldehyde 526 A TEXTBOOK OF CHEMISTRY Other reducing agents, such as sulfurous acid, may be used, but alcohol has the advantage that it leaves no other compound in the solution. By various methods it is possible to prepare two isomeric hydrates of chromic chloride, both of which have, in the crystal- line form, a composition corresponding to the formula CrCla.G H2O. One of these hydrates dissolves in water to a green solution, while the other gives a violet solution. Each can be changed more or less completely into the other, and solutions of chromic chloride usually contain the two compounds in equilibrium with each other. The green solution has a lower electrical conductivity than the violet. If silver nitrate is added to an ice-cold solution of the green hexahydrate, only one third of the chlorine is precipitated at first. The most satisfactory explanation of these facts is given by the theory of Werner. He supposes that in each compound six atoms or groups are directly combined with or arranged about the chromium atom to form a characteristic group. In the violet chloride he supposes that these six groups are all mole- cules of water and he calls this hydrate hexaaquochromic chloride and writes the formula (Cr(OH2)e)Cl3. In the, green chloride, on the other hand, he considers that four molecules of water and two atoms of chlorine form the six groups directly combined with the chromium, while one atom of chlorine and two mole- cules of water are less directly connected. He writes the formula of the green chloride, accordingly, cr* Cl + 2 H 2 O. [_ (OH2J4J Three chlorine atoms of the violet chloride are readily ionized, giving electrical conductivity and easy precipita- tion as silver chloride. But only one of the chlorine atoms of the green chloride ionizes to give electrical conductivity, and only this one atom is directly precipitated by silver nitrate. Werner and others have shown that a large number of com- plex inorganic compounds, containing water, ammonia and other groups, exhibit isomerism similar to that of these hydrates. CHROMATES 527 (See Werner, Neuere Anschauung auf dem Gebiet der anorgan- ischen Chemie.) Potassium Chromium Sulfate or " Chrome Alum," KCr(SO4)2.12 H 2 O. If potassium dichromate, K 2 Cr 2 O 7 , is warmed with alcohol and dilute sulfuric acid, the chromium is reduced to the form of chromic sulfate, 02(804)3, and on evapo- ration of the solution and cooling, the chromic sulfate and potas- sium sulfate will combine to form chrome alum, a dark violet salt, which is isomorphous with ordinary alum. Potassium Chromate, K 2 CrO 4 . When chrome iron ore, FeCr 2 O4, is roasted with potassium carbonate, the chromium is oxidized and forms potassium chromate. This is in accordance with the practically universal rule that metals assume a higher state of oxidation in the presence of a base than in the presence of an acid : 4 FeCr 2 O 4 + 8 K 2 CO 3 + 7 O 2 = 2 Fe 2 O 3 + 8 K 2 CrO 4 + 8 CO 2 The potassium chromate is easily soluble and can be easily separated from the insoluble ferric oxide. It is an easily soluble yellow salt, which gives a lemon-yellow solution. Potassium Dichromate, or Pyrochromate, K 2 Cr 2 O 7 , is formed on adding an acid, even a weak acid, to a solution of potassium chromate. It crystallizes in orange-red crystals, which are much less soluble than the crystals of the chromate. Potassium dichromate is the practical source of almost all other chromium compounds and exceeds all of the others in commercial impor- tance. It is used as an oxidizing agent, as a mordant in dyeing, and for the preparation of leather in chrome tanning, a process rapidly increasing in importance. Potassium dichromate corresponds to potassium pyrosulfate, K^Oy. The chromium in it is in the same state of oxidation as in the chromate. Lead Chromate, or Chrome Yellow, PbCrO 4 , is an insoluble salt obtained by precipitating a solution of lead acetate with potassium dichromate. It is a brilliant yellow compound which is used as a pigment and as a constituent of "chrome green " (p. 525). 528 A TEXTBOOK OF CHEMISTRY * Barium Chromate, BaCrO4, is easily prepared by precipita- tion. It is insoluble in water or acetic acid but dissolves in hy- drochloric or nitric acid. Chromium Trioxide or Chromic Anhydride, CrOa, separates in the form of dark red needles on adding an excess of concen- trated sulfuric acid to a saturated solution of potassium dichromate. When mixed with sulfuric or other acids, it is a powerful oxidizing agent and is much used for that purpose in the laboratory. To illustrate ; such a mixture with sulfuric acid will oxidize the carbon or graphite of cast iron or steel to carbon dioxide, and this method is often used for the determi- nation of these elements. Chromyl Chloride, CrO 2 Cl 2 , is prepared by distilling a mixture of potassium dichromate, salt and sulfuric acid. It is a volatile, fuming, dark red liquid, which is a very powerful oxidizing agent. Many organic compounds react with it with explosive violence. It may be considered as chromic acid, H- in which the two hydroxyl groups have been replaced by chlorine, Ck ,& y>Cr^ . It is readily hydrolyzed by water to hydrochloric Cl y \ and dichromic acids. The later may be written /O . \ <V 0> Molybdenum, Mo, 96, occurs chiefly as molybdenite, MoS 2 , a black mineral closely resembling graphite in appearance but having more than twice the specific gravity of that mineral. The metal is silver-white, after melting, and has a crystalline fracture. It melts at about 2500. * Molybdium Trioxide, or Molybdic Anhydride, MoOs, is a yellowish white or white powder, which dissolves easily in alka- lies, forming molybdates, such as sodium molybdate, Na 2 MoO4. MOLYBDENUM 529 It is nearly insoluble in acids, but if a solution of ammonium molybdate, (NH 4 ) 2 MoO 4 , is poured into dilute nitric acid (not the reverse), the molybdic acid or anhydride which is formed remains in solution either as a colloid or in a supersaturated solu- tion. Such a solution, known as " molybdic solution/' is used to precipitate solutions of orthophosphoric acid, H 3 PO4. On adding an excess of the molybdic solution to a solution contain- ing a soluble orthophosphate, or orthophosphoric acid, a very difficultly soluble, yellow precipitate separates. This has the composition (NH 4 ) 3 PO 4 .12 MoO 3 .nH 2 O. As the molybdic an- hydride will neutralize alkalies, the amount of the molybde- num and so, indirectly, that of the phosphorus may be de- termined by titration with a standard solution of potassium hydroxide : MoO 3 + 2 KOH = K 2 MoO 4 + H 2 O If the precipitate is dissolved in ammonia and an excess of dilute sulfuric acid added, the molybdenum maybe reduced to molybdic sulfate, Mo 2 (804)3, by means of zinc. The solution obtained in this manner may be oxidized quantitatively to a solution con- taining molybdic anhydride by potassium permanganate. This, again, gives an indirect determination of the amount of phos- phorus. The solution of molybdic sulfate is very easily oxi- dized by the air and the success of the operation depends on rapid work. * Compounds of Molybdenum. Molybdenum forms a series of oxides, the most important or best characterized being, Mo 2 Os, Mo 5 Oi 2 , Mo 2 O 5 , Mo 3 O 8 , MoO 2 and MoO 3 . The sulfides are Mo 2 S 3 , MoS 2 , MoS 3 and MoS 4 . * Molybdic anhydride, MoO 3 , forms a bewildering variety of complex compounds of which the ammonium phosphomolyb- date used in analysis and referred to above is an isolated exam- ple. Among the related compounds are many ammonium molybdates, such as triammonium dodekamolybdate 12 MoOs.3 NHs.12 H 2 O, and phosphomolybdic acids, as phos- phoduodecimolybdic acid, 24 MoO 3 .P 2 O5.4 H 2 O. 530 A TEXTBOOK OF CHEMISTRY A number of different ammonium phosphomolybdates have also been prepared; and it is only by careful attention to the proper conditions of temperature, acidity and concentration of the reacting solutions that a compound having the composition given above can be obtained. It is much easier to be sure that all of the phosphoric acid is precipitated from the solution than to secure the correct ratio between molybdic anhydride and phos- phoric acid. For this reason it is a common analytical practice to dissolve the precipitate in ammonia and precipitate the phos- phoric acid as magnesium ammonium phosphate, MgNH 4 PO4. Tungsten, W, 184.0. Although tungsten is usually classed among the less common elements, its compounds have been known since the middle of the eighteenth century. As a constit- uent of the old " Damascus blade " its valuable effect on steel was used long before its presence was recognized, and some of its compounds have long been used for fireproofing fabrics. The recent use for the filaments of electric lights has now made the name of the metal a household word. Tungsten is found chiefly in the form of wolframite, a ferrous manganese tungstate, (FeMn)WO 4 , the iron and manganese replacing each other as isomorphous constituents. Metallic tungsten can be obtained by the reduction of tungstic anhydride, WO 3 , with carbon or hydrogen at a high temperature, or by Goldschmidt's thermite process. Tungsten is a heavy, steel-gray, very hard metal. A good deal of difficulty was experienced in learning how to draw the metal into wire suitable for incandescent electric lights. The specific gravity is 18.64. Tungsten melts at 3000, the highest melting point of any element except carbon. Its use for electric lights depends, of course, on this property. The temperature for rapid volatilization is probably higher than that of carbon. A very important application of tungsten is in the manufacture of " high-speed tool steel." Ordinary steel cannot be used for any rapid lathe work because it would become so hot as to lose its temper. Some kinds of tungsten steel are very hard and will also retain their hardness even when almost red-hot. The TUNGSTEN. URANIUM 531 introduction of such tools has almost revolutionized shop practice in America. * Compounds of Tungsten. The element forms four chlo- rides : tungsten dichloride, WC12, tungsten tetrachloride, WCU, tungsten pentachloride, WCls, and tungsten hexachloride, WC1 6 . The last two are volatile, while the dichloride and tetrachloride are not. The two most important oxides of tungsten are the di- oxide, WO2, and the trioxide, or tungstic anhydride, WOa. Very many complex tungstates have been prepared. In addition to the normal sodium tungstate, Na 2 WO 4 , no less than thirteen complex sodium tungstates containing a smaller proportion of tungsten have been described. The compound 5 Na2O.12 WO 3 may be mentioned as an illustration of these compounds. Sodium tungstate has been used as a mordant in dyeing and for the fireproofing of fabrics. By the reduction of acid sodium tungstates by heating them with tin or hydrogen a series of yellow, blue, violet and purple compounds called " bronzes " has been prepared. These con- tain less oxygen than true tungstates should. Thus the composi- tion Na2W4Oi2 is given for the violet bronze, while the corre- sponding tungstate would be Na 2 W 4 Oi3 = Na 2 O.4 WOs. * Phosphotungstic Acid, H 3 PO 4 .12 WO 3 .18 H 2 O. This com- pound can be prepared by treating silver tungstate, Ag 2 WO 4 , mixed with the calculated amount of phosphoric acid, with hydrochloric acid. It crystallizes in rhombic crystals which are soluble in water. The solution gives characteristic precipitates with alkaloids and with proteins and is used as a reagent for these purposes. Several other complex phospho tungstic acids have been prepared. Uranium, U, 238.5. Uranium is found chiefly in the form of uraninite, or pitchblende, or UaOg. Several other minerals con- tain uranium and all of these are now of interest because of the connection with radium (p. 471). Cleveite, a mineral from Norway, resembles pitchblende, but contains also yttrium and other rare elements. It is the mineral in which Ramsay first discovered helium on the earth (p. 237). Carnotite is a uranate 532 A TEXTBOOK OF CHEMISTRY and vanadate of potassium K 2 O.2 UO 3 .V 2 O 5 .3 H 2 O. Samars- kite is a complex tantalo-columbate of uranium, yttrium, iron and other metals. Uranium is a white metal. When free from carbon it is not so hard as steel. Its specific gravity is 18.68. It melts at a higher temperature than platinum. The oxides of uranium are uranium dioxide, UO 2 , a green oxide, U 3 O 8 , of the same composition as uraninite and uranium trioxide or uranic anhydride, UO 3 . The chlorides are uranium trichloride, UC1 3 , uranium tetrachloride, UCU, uranium penta- chloride, UCls, and uranium hexachloride, UCle. Comparatively few salts of uranium are known in which metal- lic uranium replaces the hydrogen of an acid directly. In these few it is quadrivalent. One of the simplest is uranium sulfate, U(SO4) 2- 2 H 2 O. The more common salts of uranium contain the bivalent group uranyl, UO 2 . Thus uranyl nitrate is UO 2 (NO 3 ) 2 .3 H 2 Oand uranyl acetate is UO2(C 2 H 3 O 2 ) 2 .2 H 2 O. The compounds in which uranium acts as an acid forming ele- ment are mostly diuranates, corresponding to the dichromates. Potassium diuranate, K 2 U 2 O 7 , is an orange-yellow, almost in- soluble powder. As a radioactive element uranium has a " half-life " of about 6,000,000,000 years. CHAPTER XXXI MANGANESE Group VII. While there are three elements (V, Cb, Ta) al- ternating with phosphorus, arsenic, antimony and bismuth in Group V, and four elements (Cr, Mo, W, U) alternating with sulfur, selenium and tellurium in Group VI, manganese is the only element alternating with chlorine, bromine and iodine in Group VII. Not only is the halogen of atomic weight about 214 missing, but elements resembling manganese with atomic weights approximately 98, 187 and 242 have never been found. The radioactivity of uranium and the ephemeral life of niton (p. 474) which are found in this region of the table have recently given us a hint as to a possible reason for these gaps in the sys- tem. It seems likely that the structure of the atom which would give elements of these atomic weights is unstable, and that either these elements cannot exist at all or they are to be looked for among the radioactive elements of brief life. Manganese has properties such as its position in the table leads us to expect. As a metal it closely resembles iron, cobalt and nickel, all four elements forming a transition from the hard, difficultly fusible chromium to copper with its much lower melt- ing point (1083) and its great malleability and ductility. In its nonmetallic properties, on the other hand, it forms acidic oxides and acids, and the highest of these, permanganic acid, HMnC>4, corresponds to perchloric acid. Manganese, Mn, 54.93. Occurrence, Properties. Manganese is found chiefly as the dioxide in the mineral pyrolusite, MnC>2. It is also found in small amounts in most minerals and rocks, in practically all iron ores and in some natural waters. Except for scientific purposes the element is not prepared in the pure state. Pure manganese is a very hard, reddish gray metal with 533 534 A TEXTBOOK OF CHEMISTRY a specific gravity of 7.2. It melts at 1260. It dissolves easily in acids, even more easily than iron, forming manganous salts, in which the element is bivalent. Alloys of manganese with iron are easily manufactured, com- mercially, by reducing a mixture of the ores of the two metals in a blast furnace (p. 541). Those containing 10-15 per cent of manganese are white, with a brilliant metallic luster, and re- tain the carbon in the combined form. Because of the appear- ance of the surface, this alloy is called spiegeleisen (mirror-iron). It is used in the manufacture of steel (p. 548). An alloy con- taining 70-90 per cent of manganese with iron is called ferro- manganese and is used as an addition to cast iron. Manganese bronze is an alloy of manganese and copper containing 30 per cent of manganese. It is hard and has a high tensile strength. Compounds of Manganese. In compounds in which it acts as a metal, manganese is almost exclusively bivalent, as in manga- nous chloride, MnCl 2 , and manganous sulfate, MnSO 4 . There are a few unstable compounds, such as manganic chloride, MnCU, in which it is trivalent ; and compounds in which it is a quadri- valent, basic element have been prepared only in the form of double salts such as 2 KCLMnCl 4 , or K 2 MnCl 6 . The basic oxides are manganous oxide, MnO, and manganese sesquioxide, Mn 2 O 3 . As an acid-forming element manganese is quadrivalent in manganese dioxide, MnO 2 , and the manganites, such as calcium /\ manganite, Ca^ y>Mn=O. It seems to be sexivalent in the XX K (X .0 manganates, such as potassium manganate, /Mnz' , K<y ^o or K 2 MnO 4 , and septivalent in the permanganates, such as potassium permanganate, K O Mn=O. \) Manganous Manganic Oxide, Mn 3 O 4 , is of a mixed type, some- what similar to that of red lead, but the similarity in formulas MANGANESE 535 is probably superficial. Red lead is lead plumbate, or Pk, while the oxide, Mn 3 O 4 , is most likely derived from a hypothetical acid HMnO 2 . If this is true, the structure xO Mn=O is Mn(MnO 2 ) 2 or Mn<^ \)Mn=O * Manganous Hydroxide, Mn(OH) 2 , is a white precipitate which quickly turns brown from oxidation on exposure to the air, especially if an alkali is present. * Manganous Chloride, MnCl 2 .4 H 2 O, is a light pink, easily soluble salt. * Manganous Sulfate, MnSO 4 , forms hydrates with 1, 2, 3, 4, 5 or 7 molecules of water. The last, which corresponds in com- position to white vitriol, ZnSO 4 .7 H 2 O, and green vitriol or copperas, FeSO 4 .7 H 2 O, can only be obtained by crystalliza- tion at temperatures below 6. * Manganous Sulfide, MnS. Ammonium sulfide precipitates from alkaline solutions of manganous salts a flesh-colored precip- itate having the composition MnS.H 2 O or Mn . It is easily soluble in acids, even in acetic acid. The sulfide, MnS, is an olive green powder which is formed by the action of hydrogen sulfide on any of the oxides. Manganese Dioxide, MnO 2 , or Black Oxide of Manganese is found as the mineral pyrolusite in sufficient quantities to make it a very valuable commercial product. The compound has played a commanding role in the development of chemical knowledge and also in a variety of industrial processes. By means of it Scheele discovered chlorine, and it was early used in the prepara- tion of oxygen. It is still constantly used in laboratories to catalyze the decomposition of potassium chlorate (p. 21). During the latter half of the nineteenth century manganese 536 A TEXTBOOK OF CHEMISTRY dioxide was used in large quantities for the preparation of chlo- rine as an adjunct to the Leblanc soda process (p. 411). When the supply of manganese ore declined and the material grew ex- pensive, the Weldon process (p. 103) was introduced to conserve the manganese. At the present time, when electrolytic sodium hydroxide and chlorine are promising to complete the extinction of Leblanc soda and compete strongly with ammonia-soda, man- ganese is finding increasing use in the metallurgy of iron and steel, and iron manufacturers are searching eagerly for new sup- plies of the ore. It is worthy of notice that the Weldon process (p. 103) depends on the tendency of the metallic elements to assume a higher valence toward oxygen, especially in an alkaline solution, than their usual valence toward chlorine and other acid radicals in an acid solution. The sign of the valence does not, however, change. Manganese seems to be positive in all of its com- pounds. Manganates. When manganese dioxide, or indeed almost any compound of manganese, is fused with potassium or sodium carbonate and some oxidizing agent, as potassium nitrate or potassium chlorate, potassium manganate, K^MnO^ or sodium manganate, Na2MnO4, is formed. This has a green color and dissolves in water to a dark green solution. The solution is sometimes called " chameleon solution " because of the color changes which it undergoes so easily as the manganate changes to red permanganate or turns brown from the separation of man- ganese dioxide. The manganates are stable only in an alkaline solution, and free manganic acid, H 2 MnO4, does not, apparently, exist, even in solution. The addition of an acid, even of carbonic acid, to the solution causes an autoxidation and reduction some- what similar to that of hypochlorous acid (p. 127) or to that of potassium chlorate to potassium perchlorate and potassium chlo- ride (p. 128). H 2 MnO 4 -f 2 H 2 MnO 4 = 2 HMnO 4 + MnO 2 + 2 H 2 O Manganic Permanganic Acid Acid PERMANGANATES 537 (In developing this equation, notice, (1) that MnO 2 + H 2 O is equivalent to H 2 MnO 3 , hence one molecule of manganic acid gives one atom of available oxygen, and (2) that one atom of oxygen will oxidize two molecules of manganic acid to perman- ganic acid by the removal of two atoms of hydrogen.) Permanganates. Manganic acid is a very weak acid and its salts are strongly hydrolyzed in the absence of an alkali : K 2 MnO 4 + 2 HOH = H 2 MnO 4 + 2 KOH As has just been stated, the manganic acid formed is unstable and decomposes to manganese dioxide and permanganic acid, HMnO 4 . Permanganic acid is a comparatively strong acid and will at once react with the alkali present, forming a perman- ganate. Potassium Permanganate, KMnO 4 , is a dark red, moderately soluble salt, which dissolves in water to an intensely colored red solution. The color is so deep that the presence of the perman- ganate ion is evident even in exceedingly dilute solutions. By reduction in acid solutions the manganese is easily reduced to a manganous salt, such as MnSO 4 , which gives a practically color- less solution. In alkaline solutions the manganese is reduced to manganese dioxide, which is brown, leaving the solution above colorless. As the loss of color gives a very sharp indication of the end of the reaction, these properties are extensively used in quantitative analysis. As typical reactions of this character may be mentioned the oxidation of ferrous sulfate, FeSO 4 , to ferric sulfate, Fe 2 (SO 4 )3, in the presence of sulf uric acid ; the oxidation of sulf urous acid, H 2 SO 3 , to sulfuric acid ; of nitrous acid, HNO 2 , to nitric acid, HNOs ; of oxalic acid, H 2 C 2 O 4 , to carbon dioxide, CO 2 ; and of hydrogen peroxide, H 2 O 2 , to water, H 2 O, and oxygen, O 2 . As reactions in alkaline solutions may be given the oxidation of sodium sulfite, Na 2 SOs, to sodium sulfate, Na 2 SO 4 , and that of manganous sulfate, MnSO 4 , to manganese dioxide, MnO 2 . The student is advised to write the equations for all of these reactions. 538 A TEXTBOOK OF CHEMISTRY Manganese Heptoxide, or Permanganic Anhydride, Mn 2 07, is a dark greenish black, oily, volatile liquid formed when potas- sium permanganate is treated with cold, concentrated sulfuric acid. It is extremely unstable, exploding violently on slight provocation. Many organic substances take fire when brought into contact with it. It will be seen on looking back through the chapter that there are five well defined oxides of manganese, MnO, Mn 3 O4, Mn 2 O 3 , MnO 2 and Mn 2 O 7 . There is some evidence, also, but not entirely satisfactory, of the existence of a trioxide, MnO 3 . CHAPTER XXXII IRON, COBALT, NICKEL Group VIII. Between manganese and copper, molybdenum and silver, and between tungsten and gold there is, in each case, a group of three metals, closely resembling each other in proper- ties. These three groups are : the iron group iron, cobalt and nickel ; the ruthenium group ruthenium, rhodium and pal- ladium ; and the platinum group osmium, iridium and plati- num. The valence of these elements is variable, as with the other elements of this part of the table. Valences of two, three and four are most common, but nickel carbonyl, Ni(CO)4, ruthe- nium oxide, RuO 4 , and osmium oxide, OsO 4 , indicate a maximum valence of eight in some cases. Iron, Fe, 55.84. Iron is by far the most important of the metals more important than all of the other metals taken together. Several factors contribute to this importance. With the exception of aluminium, iron is more abundant than any other metal. As has been pointed out, oxygen forms about one half of the crust of the earth and silicon one fourth. Aluminium forms about one fourteenth and iron one twentieth, the four elements comprising about seven eighths of that part of the earth which we can examine. In addition to this, ores of iron con- taining 50 to 70 per cent of the metal are abundant and can be reduced on a large scale at a very low cost. ' Although alumin- ium is more abundant than iron, the latter can be produced at a very much lower cost. Finally, the forms of commercial iron contain carbon and other elements which greatly change its character, making it possible to prepare many forms of iron differing in hardness, malleability, tensile strength, permeability to magnetism and other properties which adapt the various forms to special uses. 539 540 A TEXTBOOK OF CHEMISTRY Occurrence of Iron. The ores of iron which are of primary value for manufacturing purposes are all of them oxides or com- pounds which can be readily converted into oxides by heat. The most common and important are hematite or ferric oxide, Fe 2 O3, magnetite, a ferrous-ferric oxide, FeaO^ limonite, ferric hydroxide, Fe 2 O 3 .Fe 2 (OH) 6 or 2 Fe 2 O 3 .3 H 2 O, and siderite or ferrous carbonate, FeCOa. Iron pyrites, FeS 2 , is of value primarily for the sulfur which it contains, sulfur being a much more expensive element than iron. The oxide of iron from the pyrite burners of sulfuric acid plants is occasionally used for the manufacture of a low grade of pig iron. The presence of small amounts of manganese, phosphorus and sulfur affect the quality of the iron and are often of great importance in determining the value of an iron ore. Meteorites often consist largely of metallic iron usually containing nickel, and dredgings from the bottom of the ocean show the presence of meteoric dust containing iron. The composition of meteorites, the presence of iron in the sun and density of the earth as a whole (about 5.53 as compared with 2.7 for the portion we can examine, omitting the ocean) all suggest the possibility that the central portions of the earth may contain large amounts of metallic iron. Metallurgy of Iron. In prehistoric times men learned how to reduce iron from its ores in a forge or small open hearth, with the use of a bellows or other device to secure a blast. The process was very wasteful of fuel and ore, but gave an iron of fair quality, sometimes approaching the properties of steel. An apparatus somewhat resembling a blast furnace seems to have been invented about the close of the fifteenth centirry, but for two and a half centuries the ore was reduced by means of charcoal. At one time this use of charcoal threatened to cause the destruction of the forests of England. In 1735 the reduction by means of coal was discovered, and the use of coal and coke gradually displaced the use of charcoal, though some charcoal iron is still made for special uses, because of its purity and freedom from sulfur. The more important features of a modern blast furnace for the manufacture of iron are shown in Fig. 106. A mixture of IRON: BLAST FURNACE 541 ore, fuel (usually coke) and limestone is introduced at the top of the furnace in such proportions that the oxides of iron are completely reduced, giving metallic iron. The iron combines with carbon, silicon and small amounts of phosphorus and sulfur to form the crude product known as pig iron. The lime of the limestone combines with the silica and other impurities of the ore, forming a fusible silicate, called a slag, which melts and collects in the bottom of the furnace on top of the melted iron. The air for the combustion of the fuel is forced in by means of a powerful blower through the openings near the bottom of the furnace, which are called tuyeres. The fuel remains in excess down to the very bottom of the furnace, with the result that carbon diox- ide and water, formed as the ore is reduced, are continu- ally reduced back to carbon monoxide and hydrogen. The reduction is brought about chiefly by the carbon monoxide and hydrogen in accordance with the versible reactions : re- Fig. 106 Fe 2 O 3 + 3 CO : 2 Fe + 3 CO 2 and Fe 2 O 3 + 3 H 2 :z2 Fe + 3 H 2 O 542 A TEXTBOOK OF CHEMISTRY The equilibrium of these reactions is very far toward the left, and it is only because the solid carbon of the fuel continually reduces the carbon dioxide and water back to carbon monoxide and hydrogen that the process can succeed. These conditions make it necessary to use such a proportion of fuel that the gases escaping from the top of the furnace still contain considerable amounts of carbon monoxide and hydrogen enough so that these blast furnace gases furnish a valuable fuel, retaining ap- proximately one half of the original energy of the coke. The gases are used to furnish the power for the blowing engines, either by burning them under boilers used to furnish steam for engines or by utilizing them directly in gas engines. The gases are also used in " stoves " in which the blast of air for the furnace is heated to about 800 before it enters the tuyeres. The use of the hot blast concentrates the reactions of the furnace in the lower part and greatly lessens the amount of fuel required in the charge. It has been discovered rather recently that a further saving of about 10 per cent in the amount of coke required per ton of iron is effected by first cooling the air for the blast to a low tem- perature so as to condense most of the moisture which it contains (Gayley ; see Journal of Industrial and Engineering Chemistry, 5, 241 (1913)). The materials used in a furnace working with an ore containing 60 per cent of iron are, approximately, in the proportion, one ton of ore, 0.6 ton of coke and 0.3 ton of limestone, but the amounts vary with the character of the gangue in the ore and the ash of the coke. Approximately five tons of air must be blown into the furnace for each ton of ore reduced. The process is carried on continuously for many months and sometimes for several years. At intervals of a few hours the melted iron which col- lects in the hearth of the furnace is drawn off through the tap- hole at the bottom into a large ladle, by means of which it is transferred to a mixer, where the iron from several furnaces is brought together and mixed before treating it further in Besse- mer converters or open hearth furnaces. The larger part of the iron is not allowed to cool until it is converted into steel rails, CAST IRON 543 steel or iron plates or structural materials of various forms. In the older practice, in the production of pig iron for foundry use or for the market, the iron was drawn out into a series of channels in the sand floor of the furnace room and allowed to solidify. It is then called pig iron. In modern practice pig iron is usually cast for the market in continuous casting machines. The slag either runs continuously from an opening above the surface of the iron in the hearth or more often is drawn off each time after the iron. The furnace slag is now extensively used for the manufacture of Portland cement and has sometimes been used to make cheap glass. Pig Iron. Cast Iron. The iron from the blast furnace is always a crude product containing manganese, carbon, partly combined with the iron, partly as graphite, silicon, sulfur and phosphorus. The larger portion of this crude iron is subjected to various methods of treatment which convert it into steel or refined irons which are more suitable for most purposes than the crude iron. A large amount of pig iron, however, is melted in cupola furnaces with or without the addition of aluminium, ferromanganese or other substances to improve its quality, and cast in sand moulds. Iron prepared is this way is gray in struc- ture. In the melted iron the carbon is probably all combined with the iron as iron carbide, FesC, which dissolves in the molten mass, forming a homogeneous solution. When the iron is cast as described, however, the larger part of the carbon sepa- rates from the iron as graphite during the slow cooling. Such an iron has a gray color and is known as gray cast iron. The presence of the graphite can be easily shown by dissolving the iron in a dilute acid. If the iron is cast in contact with a cold metallic surface, called a " chill," the carbon does not have time to separate as graphite, but remains combined with the iron, giving a very hard white iron, suitable for the rims of car wheels and other similar pur- poses. The addition of ferromanganese aids in holding the carbon in the combined form. On the other hand, silicon tends to cause the carbon to separate as graphite. 544 A TEXTBOOK OF CHEMISTRY The following analyses indicate the usual composition of the material : 1 ANALYSES OF PIG IRON l 2 Iron (by difference) 9429 92 72 0.55 061 Graphite 2.22 1 85 Silicon 1 84 2 57 0035 0044 Phosphorus 0.19 0.54 Titanium 0.074 0081 Manganese . 074 1 54 Copper 0.06 0.043 100.00 100.00 Gray cast iron melts at 1120 to 1230. Wrought Iron. While castings of gray iron can be finished by filing or turning in a lathe, the metal cannot be welded or rolled into bars or sheets. The puddling process for producing a nearly pure iron from the cast iron was invented by Henry Cort in England in 1784. The iron is melted and subjected to an oxidizing flame on a hearth lined with iron ore. It oxidizes to the magnetic oxide, Fe 3 O4, on the surface, and by stirring this through the mass of the iron with a rabble the oxygen of the oxide burns the carbon to carbon monoxide, CO, which escapes. The silicon is burned to silicon dioxide, which com- bines with ferrous oxide, FeO, forming a fusible ferrous silicate, Fe 2 SiO 4 . The phosphorus is burned to phosphorus pentoxide, P2O 5 , which combines with more ferrous oxide to ferrous phos- phate, Fes(PO4)2, and the sulfur burns to sulfur dioxide, which escapes. The process gives a malleable, ductile iron which 1 Samples of pig iron furnished by the U. S. Bureau of Standards. The analyses are published with the permission of the Director of the Bureau. STEEL 545 may be more than 99 per cent pure iron. This process, which was a very important one till near the close of the nineteenth century, has now been largely replaced by the Bessemer and Open Hearth processes, which are used to make mild irons as well as steel. Pure iron melts at 1530 and has a specific gravity of 7.55. It boils at 2950. Cementation Steel. Cast Steel. Wrought iron contains only a very small per cent of carbon. It is soft and malleable and can be welded,, but will not harden and is not suitable for the manufacture of knives and edge tools. By packing bars of wrought iron in charcoal, in long, earthenware boxes and heating them to 1000-! 100 for 8 to 10 days the iron absorbs approxi- mately one per cent of carbon and is changed to steel. The process is called cementation. The product is melted to render it homogeneous and is then known as cast steel. This process, which was practically the only method known for making a good grade of steel before 1855, has been more and more dis- placed by the Bessemer and open hearth processes. The cemen- tation process is tedious and very expensive, and is now used only for the manufacture of a very high grade of steel for mak- ing cutlery and for other uses where the articles manufactured are small and the labor expended in giving them their final form is the chief item in the cost of production. The name " cast steel " is now often used for steels made by other methods. When steel is heated to a temperature of bright redness, 700-800, and then cooled suddenly by quenching in water, it is rendered very hard. If such a steel is heated again to a temperature of 450-600 the steel becomes less hard and less brittle. This process is called tempering. By polishing the steel before the second heating, the color of the film of oxide which forms on the surface furnishes an indication as to when the proper temper has been obtained. Skilled workmen deter- mine the temper by watching the color of the surface of the steel. * The tempering of steel can be best understood in the light of the following facts : 546 A TEXTBOOK OF CHEMISTRY 1. Pure iron exists in three allotropic modifications called a-, /?- and y- ferrite. The word ferrite is simply a name used by metallurgists to designate pure iron. These forms of iron differ in properties, especially in being magnetic or nonmagnetic and in the amount of iron carbide, FesC, which they can hold in solid solution. Each is stable through a definite range of temperature. The ranges of stability and properties will be seen from the following table : ALLOTROPIC FORM TEMPERATURE OP STABILITY MAGNETIC PROPERTIES HARDNESS DUCTILITY SOLUBILITY OF FesC a-Ferrite . . /3-Ferrite . . y-Ferrite . . Below 750 750-860 Above 860 Magnetic Nonmagnetic Nonmagnetic Soft Hard Hard Ductile Brittle Ductile Little Little Considerable The existence of these three forms has been demonstrated by a study of the rate of cooling of iron. If a piece of iron which has been heated to a temperature above 860 is connected with a thermocouple to record the temperature, it will be found that, instead of a regular fall in temperature, as would be expected, when a transition point is reached the temperature falls more slowly, ceases to fall, or may even rise for a short time, because of the heat evolved when y-ferrite changes to /3-ferrite or when ^-ferrite changes to a-ferrite. In the case of steel a visible brightening of the mass, which is called recalescence, can be observed when one of the transition points is reached. 2. Iron combines with carbon to form a definite compound, iron carbide, FeaC, which is called by the metallurgists cement- ite. This carbide is soluble in y-ferrite but much less soluble in a-ferrite or /?-ferrite. 3. The presence of the iron carbide, FesC, lowers the tran- sition points so that in a steel containing 0.89 per cent of carbon this may all be held in a homogeneous solution in the y-ferrite at a temperature of 690, or above. Such a saturated solution of cementite or iron carbide in y-ferrite is called austenite. It is very hard, and a hardened steel of this composition consists BESSEMER STEEL 547 entirely of austenite. The quenching of the steel in water carries it so quickly by the transition points that the transforma- tion does not occur and the metal is left in the hardened form. 4. If such a steel is cooled slowly below the transition point to a-ferrite, the latter can no longer hold the cementite or carbide in solution, and the two separate into a mixture, which would consist, when the separation is complete, of about 13 per cent of cementite and 87 per cent of a-ferrite. As the latter is soft and forms the larger part of the mass, it gives character to the whole. On the other hand, if the hardened steel is heated to 450-600, it changes slowly to the mixture of a-ferrite and cementite, the change being more rapid and complete at the higher temperature. The more completely the change occurs, the softer will be the steel. Bessemer Steel. In 1852 an American by the name of Kelly patented a process for purifying iron by blowing air through it. Three years later, Besse- mer in England discovered inde- pendently and pat- ented a similar process and suc- ceeded in develop- ing it to practical success. After some litigation Mr. Kelly sold out his interest to Bes- semer, and the process is called by the latter's name. The apparatus used is shown in Fig. 107. The large, cylindrical Fig. 107 548 A TEXTBOOK OF CHEMISTRY vessel is at first turned on its side and a charge of several tons of pig iron introduced. A strong blast, which enters through one axis of the converter and through the tuyere holes at the bottom, is turned on and the converter brought to an upright position. The silicon of the iron is burned to silicon dioxide and the carbon to carbon monoxide. The silica combines with ferrous oxide to form a highly silicious slag. The heat from the combustion of the silicon raises the temperature of the iron so that it remains fluid even after the carbon and silicon are removed, although the melting point of the iron is considerably raised. When the carbon is gone, the flame suddenly drops and at this point the converter is again turned on its side and enough spiegeleisen added to give a steel containing the desired amount of carbon 0.40 to 0.45 per cent for steel rails. After mixing, the steel is poured into ingot moulds and from these it is placed in " soaking pits " to come to a uniform temperature throughout and then taken directly to the rolls and rolled into rails or structural iron or other forms of iron, without being allowed to cool. In the " acid " Bessemer process as described, in which the lining of the converter is of silicious materials, the slag contains more than 70 per cent of silica, SiO2, and very little of the phos- phorus is removed. A good quality of steel can be obtained only when the material used is quite pure. To make use of less pure ores and iron, the basic or Thomas-Gilchrist process was designed. For this the converter is lined with calcined dolomite and lime is added to the charge. In the presence of the basic lining and lime the phosphorus is removed as calcium or magnesium phosphate. The basic slag from the converter is valuable as a fertilizer because of the phosphorus which it contains. In America this process has been completely dis- placed by the basic open hearth process. Open Hearth or Siemens-Martin Process. During the last twenty years another process, illustrated in Figs. 108, 109 and 110, has grown rapidly in favor and at the present time more steel is manufactured in the United States by the open hearth OPEN HEARTH STEEL 549 550 A TEXTBOOK OF CHEMISTRY than by the Bessemer process. In Fig. 108 two chambers are represented filled with a checkerwork of brick. The producer gas (p. 297) and air used in the furnace pass up through one pair of these chambers, and after passing through the furnace room pass downward through a second pair of chambers, parting with their heat to the bricks. After an interval of twenty minutes or half an hour, the current is reversed, and now the gas and air take up heat from the bricks before entering the furnace proper. Such a furnace is called a regenerative furnace and works economically with a low grade of gas. In the furnace there is melted a mixture of cast iron, ore, steel scraps and sometimes lime or other fluxes. A charge of 50 to 75 tons may be used, and as the materials can be kept melted indefinitely, it is possible to take out a sample for analysis and secure a more accurate control than with the Bessemer converter. The process can also utilize an almost unlimited variety of material, and, especially, it is suitable for iron containing too much phos- phorus for the acid Bessemer and too little for the basic Besse- mer process. It is possible to stop the decarburization of the iron at any point desired instead of carrying it to completion and recarburizing, as is done in the Bessemer process. By all three processes grades of iron and steel containing from 0.1 to 1.0 per cent of carbon are made. Nails and sheet iron made from materials produced by these modern processes corrode, when subjected to moisture and air, very much more rapidly than when made from wrought iron from the puddling process. This seems to be because the latter is more homogeneous and does not give differences of potential between different parts of the iron. It has recently been found that the addition of a very small amount of copper (0.2 per cent, or less; Chamberlain J. Ind. and Eng. Chem. 5, 360 (1913)) renders the iron very much more resistant to corrosion. The effect is similar to that produced by amalgamating the surface of the zinc used in an electrical cell (p. 481). The following analyses illustrate the composition of iron and steel made by modern processes. IRON AND STEEL 551 . !> i I ss CD CD O5 CO iO O CO O LO O O LO O 1 I T-H O iO r-H O CO O O O CO ^ ^ O O _M CO OS O t^ T i <M r^ o CD o o CO CO t ^t 1 CO i *O CD iO Ol CO 000^ iO O O <M r^ CO CO O GO O O O o o o o CS| GO O rH O CO O O O O O O T-i co CO GO o o LO CO Oq CO O CO r-t CO T-H CO O O 0000 c^ o o o 0000 O 00 O O iO CO O GO i-H O GO O O O o SLO I rH !>. o o O CO GO CO GO O O O <M CO ss o o 8 11 J*.M| 111 I 111 I 11 -a $ 552 A TEXTBOOK OF CHEMISTRY Alloy Steels. The addition of chromium, manganese, tungsten and molybdenum to carbon steels lowers the point of decom- position of austenite (p. 546). By the use of suitable mixtures the point of decomposition may be brought below ordinary tem- perature, and such a steel will be hard at any temperature. Steels of this type have proved very useful for high-speed lathe tools, which can be run at such a rate as to become almost red- hot without losing their temper. Such steels are called " self- hardening." Compounds of Iron. When iron is dissolved in dilute acids, it forms ferrous salts, such as ferrous chloride, FeCU, or ferrous sulfate, FeSC>4. In these compounds the iron is apparently bivalent, but the vapor density of ferrous chloride at tempera- tures slightly above its boiling point points to a formula, Fe2Cl4, rather than FeCl2. If we assume this as the true formula, it still remains uncertain whether the iron or the chlorine is tri- valent or the iron, possibly, quadrivalent. The structure might be Ck /Cl /Cl=Ck >Fe=F< or Fe< >Fe CK X3 \C1=CK At present no means has been discovered of deciding posi- tively between these formulas. On exposure to the air, especially in neutral or alkaline solu- tions, ferrous compounds are oxidized to the ferric condition. If an acid is present, a ferric salt is formed. If the solution is alkaline, neutral, or acid with a weak acid, such as carbonic acid, ferric hydroxide or a basic salt is formed. Iron is ap- parently trivalent in the ferric salts, but the vapor density indi- cates that ferric chloride has, in part, the formula Fe2Cle at temperatures slightly above the point of sublimation. As with the ferrous salts, it is not known whether the molecule is held together by the iron or the chlorine atoms. The basic properties of iron are much weaker in the ferric than in the ferrous salts and while there is some hydrolysis in solutions of ferrous salts and the ferrous salts of strong acids IRON COMPOUNDS 553 react acid toward litmus, the hydrolysis of ferric salts is much more marked. This is well illustrated by the complete precipi- tation of the iron as ferric hydroxide when powdered barium carbonate, suspended in water, is added to a solution of ferric chloride : FeCl 3 + 3 HOH ^ Fe(OH) 3 + 3 HC1 2 HC1 + BaCO 3 = BaCl 2 + H 2 O + CO 2 The ferric hydroxide or a basic ferric salt, such as FeCl 2 OH or FeCl(OH) 2 , remains in collodial solution until the barium carbonate is added. When the barium carbonate reacts with the hydrochloric acid, the hydrolysis becomes complete and the barium chloride also assists in coagulating the colloidal ferric hydroxide (p. 362). Potassium Ferrate, K 2 FeO 4 , can be prepared by passing chlorine through a solution of potassium hydroxide in which ferric hydroxide is suspended. This gives a red solution from which the salt can be crystallized. Some other ferrates have been prepared, but all of these hydrolyze in water and decompose easily. Ferrous chloride, FeCl 2 .4 H 2 O, separates from concentrated solutions from which the air has been carefully excluded in clear blue crystals, which become green on exposure to the air. The hydrate dissolves in two thirds of its weight of water. In the presence of hydrochloric acid it is easily oxidized to ferric chloride, FeCl 3 , by the action of nitric acid, potassium perman- ganate, potassium dichromate, chlorine or almost any vigorous oxidizing agent. Ferrous Hydroxide, Fe(OH) 2 , forms as a pure white precipi- tate on precipitation of a ferrous salt with sodium hydroxide in a solution entirely free from dissolved oxygen or an oxidizing agent. The slightest exposure to the air causes the precipitate to turn green, and on longer exposure it is changed to reddish brown, ferric hydroxide, Fe(OH) 3 . * Ferrous Oxide, FeO, is formed as a black powder when ferrous oxalate, FeC 2 O4, is heated out of contact with the air. 554 A TEXTBOOK OF CHEMISTRY In some conditions it takes fire spontaneously on exposure to the air. If reduced by hydrogen at a low temperature it forms pyrophoric iron, which takes fire in the air. Ferrous Sulfate, Green Vitriol or Copperas, FeSO 4 .7 H 2 O. Metallic iron dissolves readily in dilute sulfuric acid and the hydrate known as green vitriol crystallizes from the solution on cooling. One hundred parts of water dissolves 38 parts of the hydrate at 10 or 48 parts at 20. Because the salt is much less soluble than the chloride, sulfuric acid is not so suitable as hydrochloric acid for the preparation of hydrogen sulfide from ferrous sulfide, FeS. Ferrous Carbonate, FeCOs, is found in nature in the mineral siderite. It crystallizes in rhombohedra and is isomorphous with calcite, CaCOs. An impure siderite, called clay iron stone, has been one of the most important iron ores used in England. Ferrous carbonate dissolves as ferrous bicarbonate, FeH2 (003)2, in waters which contain carbonic acid, exactly as calcium car- bonate does. Such waters are known as chalybeate waters, and some of these have been considered valuable for their medicinal properties. On exposure to the air the iron of such waters is oxidized to the ferric state and separates as ferric hydroxide. Beds of iron ore were doubtless formed, in many cases by a similar process. Ferrous Chloride and Nitric Oxide. Solutions of ferrous chloride or ferrous sulfate absorb nitric oxide readily, giving dark brown or black solutions. As one molecule of the ferrous salt will absorb one molecule of nitric oxide, it is supposed that the solution contains a compound of the formula FeCl2.NO.nH2O, 1 but it has never been possible to isolate the compound, and it is completely decomposed by boiling the solution. Ferric Chloride, FeCl 3 .6 H 2 O, can be obtained by passing chlorine into a solution of ferrous chloride and crystallizing the solution. It is a yellow, deliquescent solid, very easily soluble in water. The anhydrous chloride, FeCls, sublimes in dark 1 Manchot and Zechentmayer, Liebig's Annalen, 350, 368 (1906). IRON COMPOUNDS 555 green, iridescent scales which are red by transmitted light. It boils, or sublimes, at 280-285. At 448 the weight of the gram molecular volume is about 303 grams and at 750 it is 167 grams. As the molecular weight of Feds is 162.2, it seems that the mole- cules at the lower temperature are chiefly Fe 2 Cle and at the higher temperature, FeCls. Ferric chloride is easily reduced to ferrous chloride by hydro- gen sulfide, by nascent hydrogen, or by stannous chloride, SnCl 2 . Ferric Hydroxide, Fe(OH)3, is readily formed as a reddish brown, flocculent, amorphous precipitate when sodium hydroxide or ammonia is added to a solution of a ferric salt. It does not dissolve in an excess of the alkali, differing very markedly from aluminium hydroxide, A1(OH)3, in this regard. Ferric hydroxide loses water very easily, and all of the hydroxides in nature con- tain less water than would correspond to the formula Fe(OH)3. The most common natural hydroxide is limonite, Fe 2 O3.Fe 2 (OH) 6 , or as it is often written, 2 Fe 2 O 3 .3 H 2 O. Com- pounds of this type illustrate the somewhat artificial character of the distinction between hydroxides and hydrates. Dialyzed Iron. A solution of ferric chloride will dissolve a considerable quantity of ferric hydroxide. If the solution is dialyzed with a parchment membrane (p. 353), the hydrochloric acid formed by the hydrolysis of the chloride passes through the membrane, while the ferric hydroxide remains behind as a colloidal solution. Such a solution is called " dialyzed iron " and is sometimes used in medicine. It is particularly suitable as an antidote for arsenic poisoning. The arsenious oxide com- bines with the ferric hydroxide to form an insoluble arsenite. Ferric Oxide, Fe 2 O3, is prepared, artificially, by igniting the hydroxide. It is also found in nature as the mineral hematite. Ferric oxide prepared in various ways is sold as a pigment under the name Venetian red or as rouge. A mixture of calcium sulfate and ferric oxide prepared by calcination of a mixture of sulfate of iron and lime is sometimes sold under the same name. The best quality of rouge is obtained by calcining ferrous oxalate. The shade obtained depends on the temperature of calcination. 556 A TEXTBOOK OF CHEMISTRY * Ferric Sulfate, Fe2 (804)3, is formed by the oxidation of ferrous sulfate in the presence of sulfuric acid. It forms alums, of which ferric ammonium alum, NH4Fe(SC>4)2.12 H^O, is the best known. Ferric sulfate was formerly obtained by the oxi- dation of a pyrite-bearing shale at Nordhausen in Bohemia, and was decomposed by dry distillation for the preparation of fuming sulfuric acid. Magnetic Oxide of Iron, FesC^, is the product formed when iron burns in air or oxygen or when steam is passed over heated iron. It occurs also as the natural mineral magnetite, which is one of the purest of the iron ores in Norway and Sweden. It may be considered as ferrous ferrite, Fe(FeO2)2, or FeO.Fe2Os, formulas which bring out its relation to chromite, with which it is isomorphous. Ferrous Sulfide, FeS, is readily formed by heating a mixture of iron and sulfur. It may also be prepared by heating iron pyrites, FeS 2 , which loses half of its sulfur at a high temperature. It is formed as a black precipitate by the addition of ammonium or sodium sulfide to a solution of a ferrous salt. Ferrous sulfide dissolves easily in dilute sulfuric or hydrochloric acid, the reaction which is commonly used for the preparation of hydrogen sulfide. * Ferric Sulfide, Fe 2 S 3 , forms as a black precipitate on adding ammonium sulfide or sodium sulfide to a solution of a ferric salt : 2 FeCl 3 + 3(NH 4 ) 2 S = Fe 2 S 3 + 6 NH 4 C1 Ferric sulfide reacts with an ammoniacal solution of zinc chloride to give zinc sulfide and ferric hydroxide, while ferrous sulfide, FeS, reacts scarcely at all with the same solution. (Stokes, J. Am. Chem. Soc. 29, 304.) Iron Bisulfide, or Iron Pyrites, FeS 2 , is a bright yellow mineral sometimes called " fools' gold " because of its color and appear- ance. It is brittle and is easily recognized by its burning to sulfur dioxide and ferric oxide when heated in the air. It is the chief source of sulfur for the manufacture of sulfuric acid. Ferric Thiocyanate, Fe(CNS)a. When ammonium thio- cyanate, NH 4 CNS, is added to a solution of a ferric salt, ferric ^ COBALT 557 thiocyanate is formed and imparts a deep red color to the solu- tion. The reaction is used for the qualitative detection and sometimes for the colorimetric determination of iron. The ferrocyanides and ferricyanides have been described in a previous chapter. Cobalt, Co, 58.97. Cobalt is usually found associated with nickel and combined with arsenic or sulfur or both. The best known mineral is smaltite (CoNi)As2. Cobalt is a hard white metal, closely resembling iron. It is malleable and ductile and slightly magnetic. As it is not very abundant and has few properties which would make it distinctly more valuable than iron, it has not attained any considerable commercial use as a metal. An alloy with chromium is even more resistant to the attack of acids than nichrome (p. 560) and offers some promise of use for fruit knives and for spatulas for chemical laboratories (Haynes, J. Ind. and Eng. Chem. 2, 397). Cobalt melts at 1478 and has a specific gravity of 8.5. Compounds of Cobalt. Oxides. Cobalt forms cobaltous and cobaltic compounds, corresponding to the ferrous and ferric salts, but while iron tends, on the whole, to pass into the ferric state, the ordinary compounds of cobalt are the cobaltous salts. The four oxides are cobaltous oxide, CoO, cobaltic oxide, Co 2 O 3 , cobaltous-cobaltic oxide, CosO^ and cobalt dioxide, CoO2. * Cobaltous Hydroxide, Co(OH) 2 . On adding potassium hy- droxide to a solution of a cobaltous salt a blue, basic precipitate is formed. This changes on boiling to pink cobaltous hydroxide, Co (OH) 2. If cobaltous hydroxide is heated with exclusion of the air, it changes to light green cobaltous oxide, CoO. * Cobaltous Chloride, CoCl 2 .6 H 2 O, dissolves in water to the reddish or pink solutions which are characteristic of all cobaltous salts. The hydrated salt is also pink, but in dry air or on warm- ing it loses its water of hydration and changes to a deep green color. This is made the basis of a " sympathetic ink." Draw- ings made with a solution of cobalt chloride remain nearly in- visible in moist air, but come out to a clear green in dry air or on warming gently. 558 A TEXTBOOK OF CHEMISTRY * Cobalt Sulfide, CoS, is a black sulfide which dissolves only slowly in dilute hydrochloric acid. Nickel sulfide, NiS, conducts itself in the same way, and the property is often made the basis of a partial separation of these metals from the sul fides of iron, zinc and manganese, which dissolve easily and quickly in dilute acids. The difference seems to be due rather to the speed of solution than to a marked difference in the solubilities of the sulfides (A. A. Noyes, Bray and Spear, J. Am. Chem. Soc. 30 483). Cobalt sulfide dissolves easily in nitric acid. * Cobalt Nitrate, Co(NO 3 ) 2 .6 H 2 O, is also a pink, easily soluble salt. It is used in blowpipe analysis. Alumina gives a deep blue color when moistened with it and ignited, while zinc oxide gives a green color. Cobalt Glass. Cobalt compounds impart an intense blue color to the borax bead, and this is used as a delicate test for the element. They give a similar color to glass and are used in the manufacture of blue glass and in decorating porcelain. Smalt is a deep blue silicate of potassium and cobalt which has been disintegrated by pouring the melted glass into water and after- ward ground to a fine powder. It has a variety of technical uses. * Potassium Cobaltocyanide, K 4 CoC 6 N 6 , is a salt closely analogous to potassium ferrocyanide, and is formed in a similar manner by dissolving cobaltous cyanide, CoC2N 2 , in a solution of potassium cyanide. * Potassium Cobalticyanide, K 3 CoC 6 N 6 , is formed by evaporat- ing a solution of the cobaltocyanide exposed to the air. These compounds correspond to the potassium ferro- and ferricyanides (p. 320). Nickel forms no similar compounds Potassium Cobaltinitrite, K 3 Co(NO 2 )6.H 2 O, forms as a bright yellow, difficultly soluble precipitate on adding a concentrated solution of cobaltous acetate and sodium nitrite, or an adding a solution of sodium cobaltinitrite, Na 3 Co(NO 2 ) 6 , to a solution of a potassium salt. The reaction is used for the detection and quantitative estimation of either cobalt or potassium. For the determination of potassium the addition of silver nitrate, AgNO 3 , causes the formation of a still less soluble salt in which COBALT. NICKEL. 559 the potassium is partly replaced by silver, ] (Burgess and Kamm, J. Am. Chem. Soc. 34, 652). Nickel forms no similar compounds. Cobalt Ammines. Cobalt forms a very great number of com- plex compounds with ammonia. The compounds Luteocobalt chloride, [Co(NH 3 ) 6 ] C1 3 Purpureocobalt chloride, rCo(NH 3 ) 5 ~| C\z and L ci Roseocobalt chloride, rCo(NH 3 ) 5 l C1 3 L H 2 J may be given as illustrations. In these compounds the cobalt combines with ammonia, or ammonia and water, or sometimes with ammonia and chlorine or other halogens or acid radicals, to form complex groups which, in turn, combine with acid radicals to form salts. Some of these compounds show forms of isomerism recalling that of the hydrates of the chromic chlorides (p. 526). A study of these compounds by Werner, especially, has led to the development of new views of valence which are important, and which differ considerably from the points of view developed from the study of organic compounds. See Werner, Z. anorg. Chem. 3, 267-330 (1893) ; and Neuere Auschauungen auf dem Gebiete der anorganischen Chemie, 1905. Nickel, Ni, 58.68, is found always associated with some cobalt and combined with arsenic or sulfur, or both, or in the form of a silicate. The principal sources for the nickel of the world have been Sudbury in Ontario and New Caledonia, an island east of Australia. At Sudbury the ore is a complex sulfide containing copper, arsenic and other metals. The metallurgical process is complicated and the details are not well known out- side of the works where they are carried out. The first step is the preparation of a nickel-copper matte by smelting the ore, somewhat as is .done with copper ores (p. 429) . Nickel is a white metal resembling steel. It is magnetic, melts at 1452 and has a specific gravity of 8.8. It takes a high 560 A TEXTBOOK OF CHEMISTRY polish and does not rust or tarnish so readily as steel. It is extensively used for plating the ornamental parts of stoves, handle bars of bicycles and many other articles. For this pur- pose it is usually deposited from a solution of nickel ammonium sulfate. Nickel is used in several alloys, the most important being German silver, an alloy of nickel, copper and zinc, which is light colored and used as a basis for silver plated ware. Ni- chrome, an alloy of nickel and chromium, is resistant to acids and is also suitable for triangles for laboratory use, for thermo- couples, and for electric heating devices. Other alloys highly resistant to acids have been prepared and there is a prospect of the development of important new alloys of this type. The five-cent coin of the United States is 75 per cent copper and 25 per cent nickel. Compounds of Nickel. Nickel forms three oxides, nickelous oxide, NiO, nickelic oxide, Ni 2 O 3 , and nickelous nickelic oxide, Ni 3 O4, corresponding to the similar oxides of iron. The salts of nickel form green solutions which are complementary in color to the salts of cobalt. Mixtures of the two can be made which are nearly colorless. The chloride is NiCl2-6 H^O and the sulfate NiSO 4 .7 H 2 O. Nickel Dimethylglyoxime, CH 3 C =N O\ [~CH 3 C =NOH CH 3 C =N (X CH 3 C =NOH Nickel forms no precipitate with potassium nitrite, and cobalt can be separated from solutions containing it by this means. CH 3 C=NOH It is precipitated by dimethylglyoxime, , from CH 3 C=NOH an ammoniacal solution or from a solution containing a weak acid, in the form of a scarlet-red, highly characteristic compound, nickel dimethylglyoximine. Cobalt forms no similar precipitate, especially if the ammoniacal solution is first shaken in the air to oxidize the cobalt and convert it into a complex cobaltic NICKEL 561 ammine. The formation of this precipitate furnishes one of the best means for the detection and quantitative estimation of nickel. The precipitate may be sublimed without decomposition by careful heating (Chugaev, 1 Ber. 38, 2520; Z. anorg. Chem. 46, 144; Kraut, Z. angew. Chemie, 19, 1793). Nickel Carbonyl, Ni(CO) 4 . By passing carbon monoxide over finely divided nickel at a temperature below 80, nickel carbonyl is formed. It is a volatile liquid which boils at 43 and whose vapor decomposes explosively at 60. Cobalt forms no similar compound, and attempts have been made to use it industrially for the separation of nickel from other elements, but these attempts have not met with much success. Iron forms a similar compound, iron tetracarbonyl, Fe(CO) 4 , and also a pentacarbonyl, Fe(CO)5, but these are much less stable than the nickel carbonyl. However, it seems probable that traces of these iron compounds are formed from the carbon monoxide of illuminating gas and are the cause of the deposit of ferric oxide sometimes obtained from gas burners. 1 Also written Tschugaeff. CHAPTER XXXIII THE PLATINUM METALS THE following table gives the atomic weights, specific gravity, melting points and formulas of the oxides of the metals of Group VIII of the Periodic System : Fe Co Ni Atomic weight 55.84 58.97 58.68 Specific gravity 7.88 8.7 8.8 Melting point 1530 1478 1452 FeO CoO NiO n . , Fe 3 4 Co 3 O 4 Ni 3 O 4 ' ' Fe 2 3 Co 2 O 3 Ni 2 3 CoO 2 Ru Rh Pd Atomic weight 101.7 102.9 106.7 Specific gravity ..... 12.2 12.6 11.9 Melting point 2300 1940 1549 RuO RhO PdO , Ru 2 3 Rh 2 3 ^Ru0 2 Rh0 2 Pd0 2 RuO 4 Os Ir Pt Atomic weight 190.9 193.1 195.2 Specific gravity 22.48 22.4 21.4 Melting point ..... 2700 2300 1755 OsO Oxid .Os 2 3 Ir 2 3 PtO 1 ^Os0 2 Ir0 2 Pt0 2 OsO 4 The platinum metals are found almost exclusively in the free state, alloyed together in small grains or nuggets. Platinum forms two thirds to five sixths of the alloy. 562 THE PLATINUM METALS 563 * Ruthenium, Ru, 101.7, is found in the mineral laurite, (RuOs) 2 S 3 , as well as in the natural platinum alloys and in osmium-iridium. Ruthenium monoxide, RuO, is obtained by heating a mixture of ruthenium dichloride, RuCl 2 , and sodium carbonate. The sesquioxide, Ru 2 O 3 , is formed when the metal is heated in the air. It will be seen from this that ruthenium resembles iron rather than platinum. Ruthenium dioxide, RuO2, is prepared by heating ruthenium in a current of oxygen. Ruthenium tetroxide, RuO4, is a volatile compound formed by fusing ruthenium with potassium hydroxide and potassium nitrate. Its odor resembles that of ozone. Ruthenium forms three chlorides, RuCl2, RuCl 3 and RuCU, and such double chlorides as K 4 RuCl 6 (or 4 KCLRuCl 2 ) and K 2 RuCl 6 . The potassium ruthenate, K 2 RuO4.H 2 O, which gives orange-red solu- tions and perruthenate, KRuC>4.H 2 O, giving a green solution, correspond to the manganates and permanganates. * Rhodium, Rh, 102.9, is much less easily attacked by acids and other reagents than ruthenium. It is hard and has been sometimes used for the tips of gold pens. Rhodium monoxide, RhO, is obtained by heating the hydroxide, Rh (OH) 3 . The sesquioxide, Rh 2 O 3 , is prepared by heating the nitrate, and the dioxide, RhO 2 , by fusing rhodium with potassium hydroxide and nitrate. The chlorides are RhCl 2 and RhCl 3 . The double chloride with potassium is K 2 RhCl 5 .H 2 O or 2 KCl.RhCl 3 .H 2 O. There are many complex salts, such as roseorhodium chloride, (~Rh ( ^ T H ' )5 lci 3 , and luteorhodium chloride, Rh(NH 3 ) 6 Cl 3 . L H 2 J Palladium, Pd, 106.7, is always present in the natural platinum alloys. It is also frequently found in metallic silver. It re- sembles platinum or silver in appearance and occupies a posi- tion somewhat between them in its properties. It can be rolled into foil and drawn into wire. It is soluble in nitric acid, and finely divided palladium will dissolve in hydrochloric acid. One of the most remarkable properties of palladium is its absorption of hydrogen. If palladium foil is heated in an at- mosphere of hydrogen and then allowed to cool in a current of 564 A TEXTBOOK OF CHEMISTRY the gas, 100 grams of the metal will absorb about 0.64 gram of hydrogen or approximately 7 liters of the gas. As the volume of 100 grams of the metal is only 8.4 cc., it follows that the metal absorbs more than 800 times its volume of the gas. This property has been used as a convenient method of weighing hydrogen for the determination of its atomic weight. The hydrogen absorbed is in an active form. It is oxidized to water at once in contact with oxygen or the air. A solution of colloidal palladium also has a similar catalytic effect and has been used with hydrogen for the reduction of organic compounds. It has even been proposed to use this method commercially for the reduction of liquid or semiliquid fats containing glycerides of unsaturated acids to convert these into solid fats which are commercially much more valuable. * Palladium forms only two well-defined oxides, the monoxide, PdO, and the dioxide, PdO 2 . Palladium dichloride, PdCl 2 .2 H 2 O, is obtained by dissolving spongy palladium in hydrochloric acid. The solution is reduced by hydrogen and is sometimes used to absorb that gas in gas analysis. The double salt, K 2 PdCl4, dissolves in water to a dark red solution. Palladium tetrachloride, PdCl 4 , is known only in solution and is not very stable. The double salt, K 2 PdCl 6 or 2 KCl.PdCl 4 , is difficultly soluble in cold water and crystallizes in scarlet-red octahedra. Palladium forms a series of ammines similar to those of cobalt and rhodium. * Osmium, Os, 190.9, is the heaviest substance known. The name is given because of the strong odor of its volatile tetroxide, OsO 4 . It is found with platinum and iridium in the alloy called osmium-iridium, which is insoluble in aqua regia. From this alloy the osmium is obtained by heating in a current of oxygen, which converts the osmium into volatile osmium tetroxide, OsO 4 . The reguline metal has a bluish color somewhat resem- bling zinc. Osmium gives four oxides, OsO, Os 2 O 3 , OsO 2 and OsO 4 . The tetroxide is a white solid which dissolves slowly in water, but volatilizes from the solution. The vapor has a disagree- THE PLATINUM METALS 565 able, chlorine-like odor. It attacks the eyes strongly and is extremely poisonous. It melts at 40 and boils at about 100. It is often called an acid, but has no acid properties. The aqueous solution is sometimes used in histology to stain or harden tissues. The chlorides of osmium are OsCl 2 , OsCl 3 and OsCU. The double salts with potassium are OsCl 3 .3 KC1.6 H 2 O, or K 3 OsCl 6 .6 H 2 O and K 2 OsCl 6 . Potassium osmate, K 2 OsO 4 .2 H 2 O, crystallizes in rose-red or violet octahedra. * Indium, Ir, 193.1, melts about 550 higher than platinum, and as it does not oxidize in the air at high temperatures, it has proved especially useful for some forms of chemical apparatus. It is also used for the tips of gold pens, because of its hardness. Its color is between those of silver and tin. The oxides are the sesquioxide, Ir 2 O 3 , and the dioxide, IrO 2 . Both are decomposed at a high temperature into iridium and oxygen. The chlorides are IrCl 2 , IrCl 3 and .IrCl 4 . The double salts are K 3 IrCle.6 H 2 O and K 2 IrCle. The latter forms dark red octahedra, difficultly soluble in water. Platinum, Pt, 195.2, is very much the most important of the platinum metals. Its very high melting point (1745) and the fact that it does not dissolve in nitric, hydrochloric or sulfuric acid make it an almost indispensable metal in the laboratory. The fact that its coefficient of expansion is almost the same as that of some kinds of glass has led to its usje for the leading-in wires of electric light bulbs. Its use as a catalytic agent to cause the combination of hydrogen and oxygen, and also the combination of sulfur dioxide and oxygen in the contact process for sulfuric acid (p. 175), have been given. Platinum sponge is used as the filtering material in the Munroe-Neubauer crucibles. * Platinous Chloride, PtCl 2 , is a greenish compound formed by passing chlorine over platinum at 240-250. It dissolves in hydrochloric acid, giving chloroplatinous acid, H 2 PtCl4. The potassium salt, K 2 PtCl 4 , is used in photography. Chloroplatinic Acid, H 2 PtCle, is usually prepared by dissolving platinum in aqua regia and evaporating the solution repeatedly 566 A TEXTBOOK OF CHEMISTRY to expel the excess of nitric acid, but it is extremely difficult to obtain a pure product in this manner. The pure compound is best prepared by dissolving platinum black electrolytically in hydrochloric acid (Weber, J. Am. Chem. Soc. 30, 29). The acid is easily soluble in water, giving a yellow or reddish yellow solution, according to 'the concentration. With potassium or ammonium salts the solution gives yellow precipitates of potas- sium chloroplatinate, K 2 PtCle, and ammonium chloroplatinate, (NH^PtCle- These compounds are much used in analytical chemistry. Similar compounds, many of which, however, are more easily soluble, are formed with organic bases. Silver nitrate, AgNO 3 , precipitates silver chloroplatinate, Ag 2 PtCle, and not silver chloride, AgCl, from a solution of chloroplatinic acid. Platinic Chloride, PtCU, is obtained by heating chloroplatinic acid in a current of chlorine at 360. When dissolved in water, it- gives the compound H 2 PtCl 4 O.4 H 2 O or PtCl 4 .5 H 2 O. The first formula is justified by the fact that four molecules of water can be readily expelled, but the fifth cannot be removed without loss of chlorine. * Platinum Bisulfide, PtS 2 , is precipitated on passing hydrogen sulfide into a solution of chloroplatinic acid. It is a black pre- cipitate which dissolves in ammonium sulfide as ammonium sulfoplatinate. Platinum forms a long series of complex ammines. INDEX INDEX Abscissas, axis of, 43. Absolute, alcohol, 325 ; potential of elements, 436 ; temperature, 39 ; units, 33; zero, 40. Abstract sciences, 4. Acetamide, acid in liquid ammonia, 208. Acetanilide, 340. Acetic acid, acetyl chloride from, 245 ; manufacture, properties, con- stituent of " liquid smoke," 329 ; strength illustrated, 386 ; titration of, 389 ; solubility of calcium phos- phate in, 462 ; insolubility of calcium oxalate in, 465. Acetone, formation, preparation, uses, 328 ; solvent for acetylene, 293. Acetyl chloride, formed from acetic acid, 245. Acetylene, endothermic, 292, 294; formation, 292, preparation from calcium carbide, 293 ; in illu- minating gas, 295 ; light from, 293 ; liquid, explosive, 294 ; poly- merization, 294 ; solution in ace- tone not explosive, uses, 294 ; tetrabromide, 293. Acheson, graphite in colloidal solu- tion, 277. Acid, definition, 45, 168 ; properties from oxygen, 23 ; strong, defined, 168. Acids and bases, writing equations for reactions between, 156 ; degree of ionization, table, 383 ; dibasic, defined, 183 ; nomenclature of, 123; organic, structure, 328; strength of, definition, 167; strength of, illustration, 386; strength of in relation to solu- bility of sulfides, 168 ; tribasic, defined, 183 ; weak and strong, 386. Acid chlorides, defined, 189. Acidimetry, 185. Acidity or alkalinity of indicators at change of color, table, 388. Actinium, 475 ; series of elements, 475. Adiabatic cooling of air at higher levels, 232. Adjective dyes, 342. Adsorption, 278. Affinity, chemical, 29. Agate, 348. Air, absorption as mixed gas by water, 228;] a mixture, 228; amount of fresh, required per hour, 231 ; analysis of by nitric oxide, 230 ; calculated weight of 1 liter, 229 ; coefficient of ex- pansion, 38. Air, composition of by volume and weight, 228 ; demonstrated by Lavoisier, 19 ; determined by hydrogen, 227 ; determined by mercury, 227 ; determined by phosphorus, 227. Air, determination of moisture in by weighing, by dew point, and by moist bulb of thermometer, 232; liquefaction, 232; puri- fication before liquefaction, 234 ; sources of carbon dioxide in, 229 ; weight of gram molecular volume, 228 ; weight of 1 liter, 229. Air-slaked lime, 453. Alabaster, 457. Alberene, 349. Albite, trisilicate, 356. Albumen, 343. Albumoses, formed in digestion, 344. Alchemists, name for silver, 444. Alcohol, defined, 324 ; manufacture, properties, uses, absolute, dena- tured, 325. Alcoholic beverages, composition, 325. Aldehydes, 327. Alfalfa, fixation of nitrogen by, 199. Alizarin, manufacture, 341. Alkali industry, history of, 400. Alkali metals, general properties, 395. Alkali-earth metals, general proper- ties, 451. Alkalimetry, 185. Alkalinity or acidity of indicators at change of color, table, 388. Alkaloids, 342. Allotropic forms, definition, 98 ; of phosphorus, 241 ; of sulfur, 162. 570 INDEX Alloy steels, 552. Alloys, fusible, 269. Alum, ammonium gallium, 506 ; caesium, rubidium, 424 ; chrome, 527. Alums, 500, potassium, 500, ammo- nium, ammonium ferric, chrome, rubidium, 501. Alumina, blue color with cobalt nitrate, 558. Aluminium acetate, mordant, 342 ; amalgam, activity of, 497 ; bronze, from electric furnace, 495, com- position, 497 ; chloride, prepara- tion, anhydrous, hydrate, use, 498; exercises, 507. Aluminium, history of metallurgy of, 391 ; occurrence, formation of shales, clays, soils, 494 ; metal- lurgy, history, 495 ; manufacture, 496; properties, 497; alloys, thermite process, 497 ; compounds, 498 ; use in cast iron, 543 ; use in metallurgy, 391. Aluminium hydroxide, precipitation, 499 ; preparation from clay, 496 ; base and acid, 499 ; fluoride, prep- aration, 499 ; metachlproanti- monate, 268 ; oxide, dissolved by sodium pyrosulfate, 408 ; oxide, preparation for manufacture of aluminium, 495, 496 ; oxide, oc- currence, 494 ; artificial, 499 ; uses, 500 ; solution in sodium pyrosulfate, 500 ; sulfate, hydroly- sis, preparation, use, 499, 500. Amalgam, ammonium, 420 ; sodium, 487. Amalgamated zinc, conduct toward acids, 481. Amalgamation process for silver, 441. Amalgams, 486. Amethyst, 348. Amide, definition of, 206 ; ions in liquid ammonia, 208. Amine, definition of, properties, 205. Amino acids, formed in digestion, 344. Ammines, cobalt, 559. Ammonia, " associated " liquid, 204 ; combination with acids, 202 ; deriv- atives of, 205 ; detection with Nessler's reagent, 492. Ammonia, determination of com- position by volume by action of chlorine on, 208 ; by decomposition and recombination with electric discharge, 209. Ammonia, deviation from Boyle's law, 35 ; formation, 201 ; formed by action of zinc on nitric acid, 213 ; liquid, solutions in, 207. Ammonia, preparation by hydrolysis of a nitride, 202 ; from ammonium sulfate, 202 ; from aqua ammonia, 202. Ammonia, properties, solubility, 202 ; reaction between chlorine and, 209 ; synthesis of, 201. Ammonia soda process, discovery, 400. Ammoniacal gas liquors, 202. Ammonio-cadmium sulfate, 492; -cu- pric sulfate, 434 ; -cuprous chloride, 433 ; -zinc sulfate, 492. Ammonium, in ammonium amal- gam, 420 ; bicarbonate, 423 ; bi- carbonate, use in ammonia soda process, 412 ; carbonate, com- mercial, preparation, composition, use, 423 ; carbonate, formation, hydrolysis, use, 423 ; chloroaurate, 450. Ammonium chloride, formation from ammonia and hydrochloric acid, 203 ; manufacture, dissociation, volatilization of dry without dis- sociation, relation to Avogadro's hypothesis, 421 ; recovery of am- monia from, 413. Ammonium chloroplatinate, 566, 423. Ammonium chloroplumbate, 518 ; citrate, use in analysis of fertilizers, 331 ; citrate, use in determining citrate-soluble phosphoric acid, 461 ; ferric citrate, use with potas- sium ferricyanide in blue prints, 331. Ammonium hydroxide, dissociation to ammonia and water, 204 ; equili- bria in solutions of, 420 ; ioniza- tion, 203 ; structure according to election theory, 207 ; titration of, 389. Ammonium hydrosulfide, prepara- tion, formation of polysulfides from, 421 ; magnesium phosphate, decomposition of, 252 ; molybdate, use to determine phosphorus, 529, 530. Ammonium nitrate, decomposition of exothermic, use in explosives, 215, 423 ; preparation of nitrous oxide from, 214 ; preparation, proper- ties, use, 422. Ammonium nitrite, properties, 423 ; oxalate, use to precipitate cal- cium, 330 ; phosphomolybdate, use to determine phosphoric acid, 529, 530 ; polysulfides, formation, use, 422. INDEX 571 Ammonium salts, theory of, 203; sodium hydrogen phosphate, use, 423 ; sulfarsenite, 261. Ammonium sulfate from ammoniacal gas liquors, 202 ; preparation, use, 422 ; use as fertilizer, 199. Ammonium sulfide, preparation, 421, sulfide, formation of polysulfides from, uses, 422 ; sulfostannate, 512; trinitride, structure, 221. Ammono-dimercuric iodide, use in testing for ammonia, 493. Ammono-mercuric chloride, 492 ; compounds, 492 ; nitrate, 492. Amorphous, definition, 162. Ampere defined, 33. Amphibole, metasilicate, 355. Amphoteric compounds, definition of, 206, 483. Amyl acetate, use in lacquers, 338. Analysis, definition, 66 ; hydrogen sulfide basis of groups in, 166. Analytical chemistry, groups of, 166. Andrews, critical temperature, 232. Anesthesia, produced by nitrous oxide, 215. Anhydrite, 457, 458, soluble, 458 ; conditions for formation of, 458, 460 ; vapor pressure of systems con- taining, 459. Aniline, preparation, uses, 340. Animal charcoal, 278. Animal foods, 347. Anion, definition, 48. Anode, definition, 47, 113; deposit of silver peroxide on, 443. Anthracene, alizarin from, 341 ; from coal tar, use, 295. Anthracite coal, composition, 280. Antifebrine, 340. Antifriction metals, 264. Antimonic acids, 267. Antimonious acid, 265. Antimony, chlorides of, 267 ; hy- droxide, 265 ; oxides of, 265. Antimony, occurrence, preparation, 263; properties, uses, alloys, ex- plosive, 264. Antimony oxychloride, 267; penta- chloride, 267; pentoxide, 265; pentasulfide, 268. Antimony tetrachloride, 267; tetra- chloride, endothermic, 267 ; tet- roxide, 265 ; trichloride, explosive antimony from, 264 ; trichloride, preparation, properties, hydrolysis, 267 ; trioxide, 265 ; trioxide, tartar emetic from, 266 ; trisulfide, 268. Antimonyl, 266; chloride, 267; po- tassium tartrate, 266 ; sulfate, 266. Antipyrine, 340. Antitoxins, 345. Apatite, 153, 452. Apollinaris water, 309. Aqua ammonia, 202, 203. Aqua regia, 213 ; use to oxidize sulfur of sulfides, 213. Aragonite, 452. Argentum, 11. Argol, 330. Argon, atomic weight, 236 ; coeffi- cient of expansion, 38 ; discovery, 235 ; molecular weight, 236 ; per cent in ah*, 228 ; properties, 236. Argyrodite, discovery of germanium in, 361. Arsenic acid, oxidizing agent, 259 ; preparation, properties, salts, 259 ; transformation in steps to arsenic pentasulfide, 261. Arsenic disulfide, 260. Arsenic from smelting copper ores, 256; occurrence, 256. Arsenic pentasulfide, 260 ; pentoxide, 259 ; " poison " to platinum catal- ysis for preparation of sulfur trioxide, 175 ; preparation, prop- erties, uses, 257 ; sulfides of, 260. Arsenic trioxide, conversion to col- loidal arsenic trisulfide, 261 ; tri- oxide, formation, properties, 258 ; trisulfide, 260; trisulfide, col- loidal, 261. Arsenious acid, salts, 259 ; oxide, oxidation by nitrogen trichloride, 224 ; oxide, standard in iodimetry, 260. Arsenites, 259. Arsenopyrite, 256 ; arsenic from, 257. Arsine, Marsh's test, 257 ; com- pared with ammonia, 243. Asbestos, diaphragm of for alkali manufacture, 401 ; metasilicate, 355 ; platinized, preparation of, 62. Assaying, 440. Association, water and ammonia, 204. -ate, suffix, use, 47 ; use for salts, 124. Atmosphere, exercises, 239. Atomic theory, 14; volumes, curve for, 137 ; volumes, relation to periodic system, 136 ; weight of chlorine, determination of, 130. Atomic weights, selected by law of Dulong and Petit, 397 ; selection of, 16, 92 ; table of, 10 ; unit for, 68. Atoms, probably complex aggre- gates, 138 ; structure of, 473. Atropa belladonna, atropine from, 343. Atropine, 343. At water, respiration calorimeter, 313. 572 INDEX -Auric acid, 449. Aurum, 11. Austenite, relation to tempering of steel, 546. Avogadro's law, 89, 91 ; exercises, 99 ; related to law of Dulong and Petit, 397 ; relation to laws of Boyle and Charles, 94. Az-, prefix derived from azote, 220. Azo-, prefix derived from azote, 220. Azoimide, see hydronitric acid, 223. Azote, name for nitrogen, 220. Babbitt metal, 269. Bacteria, killed by radiations from radioactive elements, 476 ; re- moval from water, 83 ; nitrifying, 199. Badische Anilin Soda Fabrik, syn- thetic ammonia, 201 ; manufac- ture of indigo, 341. Baeyer, synthesis of indigo, 341. Baker, atomic weightof tellurium, 190. Baking powder, acid potassium tar- trate in, 330. Baking soda, 412. Banca tin, 508. Barite, 162, 468 ; use, 470. Barium carbonate, dissociation pres- sure, manufacture of barium oxide from, 468; chloride, 470; chro- mate, 528; exercises, 476; flame color, 471. Barium, hydroxide, properties, uses, 470 ; nitrate, preparation, use, 470 ; nitrate, barium oxide from, 469 ; occurrence, compounds, 468 ; oxide, manufacture, 468 ; uses, 469. Barium peroxide, contrast with lead dioxide, 518 ; dissociation pres- sure, use to prepare oxygen, for hydrogen peroxide, 469 ; peroxide hydrate, 469 ; hydrogen peroxide from, 84. Barium silicofluoride, insoluble, 350 ; sulfate, use, constituent of litho- pone, 470 ; sulfate, solubility, 471 ; sulfide, preparation, use, 470. Barometer, correction of readings for altitude, 37 ; correction of readings for glass and brass scales, 36; correction of readings for latitude, 37. Bases, definition, 121 ; derived from ammonia, 339 ; how formed from ammonia and amines, 206 ; weak and strong, 386 ; and acids, writing equations for reactions between, 156. Basicity, defined, 183. Batteries electric, use of zinc in, 481. Battery galvanic, reverse of electro- lytic cell, 439. Bausfield, value of the calorie at different temperatures, 33. Bauxite, 494. Baxter, separation of praseodymium, and neodymium, 504. Beads, borax and sodium metaphos- phate with blowpipe, 304. Bearings in machinery, phosphor bronze for, 431. Becquerel, discovery of rays, 471. " Bee hive," coke ovens, 278. Beef, use of potassium nitrate in salt, 418. Beet sugar, 333. Bell metal, 509. Benedict, respiration calorimeter, 313. Benzaldehyde, 328. Benzene in illuminating gas, 295 ; from coal tar, 294 ; properties, 295 ; structure, 285. Benzine, 289. Benzoic acid, cocaine a derivative of, 343 ; occurrence, manufacture, use as a food preservative, 331. Beryl, 451. Beryl, metasilicate, 355. Beryllium carbonate, 452 ; chloride, 451; hydroxide, 451 ; nitrate-, 452 ; occurrence, properties, compounds, 451 ; sulfate, 451. Berzelius, determination of composi- tion of water, 69 ; experience with hydrogen selenide, 190 ; use of isomer, 511. Bessemer converter, 547 ; use in metallurgy of copper, 429. Bessemer steel, history, 547 ; acid and basic or Thomas-Gilchrist, 548; soaking pits, 548. Bi-bivalent salts, law of solubility product not general for, 378. Bicarbonate ion, ionization of, 310. Bicarbonates, formation from car- bonic acid, 310. Bimolecular reactions, 150. Binary compounds, nomenclature of, 29. Biological sciences, 4. Biscuit, forms for earthenware, 501. Bisdiazoacetic acid, use in preparing hydrazine, 222. Bismuth, alloys of, 269 ; basic nitrates of, 270 ; in crude copper, 430 ; melt- ing point lowered by pressure, 269. Bismuth, occurrence, properties, uses, 268; nitrate, preparation, hy- drolysis, 270 ; oxides of, 269 ; oxy- chloride, 270; " subnitrate," 270; trichloride, 269 ; trisulfide, 270. INDEX 573 Bismuthyl chloride, 270. Bituminous coal, calculation of heat of combustion of, 44 ; composition, 280. Bivalent, definition, 64. Black ash in Leblanc soda process, 411. Black oxide of manganese, history of uses, 535. Blast, heating for blast furnace, 542 ; dry, 542. Blast furnace, 541 ; gas, 298 ; per- centage composition, 299 ; slag, use for cement, 454. Bleaching powder, 124 ; manufac- ture, 455 ; properties, uses, 456 ; use in purifying water, 83. Bleaching by sulfur dioxide, 173 ; with chlorine, 106. Blindness caused by methyl alcohol, 325. Blowpipe, construction and use, 304 ; oxyhydrogen, 61. Blueing, 321. Blue-print paper, 331. Blue vitriol, 433. Bodenstein, decomposition of hydrio- dic acid, 148 ; heat of formation of hydriodic acid, 153. Body, definition, 7. Bohemian glass, 467. Boiling point, criterion of pure sub- stance, 12 ; of solutions and os- motic pressure, 360. Boisbaudran, discovery of samarium, 505. Bolivia, tin from, 508. Bomb calorimeter, 25. Bone ash, 241. Bone black, 278. Borax beads in blowpipe flame, 304 ; use in blowpiping, 366. Borax, occurrence in California, 365 ; glass, 366 ; uses, 367. Boric acid, in Tuscany, 365 ; prepa- ration, properties, uses, 366 ; use as food preservative, 366 ; tests for, 367, 368. Bornite, 428. Boron, exercises, 368 ; fluoride, hydrolysis, 367; nitride, 367; oc- currence, preparation, properties, 365 ; sulfide, 367 ; trichloride, 367 ; trioxide, 365. Borosilicate glasses, 467. Boyle, law of, 34. Boyle's law, deviation of gases from, 35 ; illustration of, 35. Brandt, discovery of phosphorus, 242. Brandy, 325. Brass, 431. Brazil, diamonds from, 275. Breath, source of carbon dioxide, 229. Brick, 501. Brimstone, roll, 161. Brines, evaporation of, 405. Britannia metal, 264, 509. British thermal unit (B. T. U.), 26. Bromate, sodium, 144. Bromine, occurrence, 140 ; origin of name, 141 ; preparation, 140 ; properties, 141 ; uses, 142. Bronze age, relation to metallurgy, 390; phosphor, 431. Bronzes, 431, 509 ; sodium tungstate, 531. Brownian movement, 96. Bullion, lead, separation of silver from, 440. Bunsen, discovery of rubidium and caesium, 424 ; discovery of spec- trum analysis, 424. Bunsen burner, nature of flame, 300 ; burner, temperature of flame, 303 ; flame, separated, 301. Burgess, table of melting points of elements, 373. Burgess, use of potassium silver cobaltinitrite in analysis, 559. Burns produced by radioactive ele- ments, 476. Butadiene, structure, 285. Butane, normal, structure, 284. Butene, 283. Butine, 283. By-products coke ovens, 279. By-products, importance in Leblanc soda process, 401. Cadmium, abnormal ionization of, 491 ; compounds, effect of ammonium hydroxide .on solutions of, 491 ; hydroxide, 484. Cadmium, occurrence, properties, 483, use in fusible alloys, 484 ; separation from copper with po- tassium cyanide, 435 ; sulfate, 484 ; sulfate, use in Weston cell, 438; sulfide, conduct toward acids and potassium cyanide, 484 ; sulfide, solubility, 491. Caesium alum, 424 ; chloroiodide, use in purification of caesium, 424. Caesium chloroplatinate, 393, 424. Caesium, discovery, properties, 424 : tetrachloroantimonate, 267. Cailletet, liquefaction of air, 233. Calamine, 481. Calamine, orthosilicate, 355. Calcite, 452. Calcium acetate, manufacture, use, 574 INDEX 465; bicarbonate, hard waters, 310. Calcium carbide, hydrolysis to cal- cium hydroxide and acetylene, 293 ; manufacture, hydrolysis, use, 462 ; calcium cyanamide from, 462. Calcium carbonate, dissociation, phase rule, 453 ; dissociation pressures of, 453 ; in mortar, 454 ; solubility in pure water and in water containing carbonic acid, hard waters, 463. Calcium chlorate, formation, use in making potassium chlorate, 456 ; chlorate, use to prepare potassium chlorate, 127 ; chloroaurate, 450. Calcium chloride, by-product in am- monia soda process, 413 ; effect on the formation of gypsum, 460 ; moisture left in gas by, 54 ; prep- aration, hydrates, by-product in ammonia soda process, as drying agent, uses, 455. Calcium cyanamide, manufacture, hydrolysis, 462; use, 463; exer- cises, 476 ; flame color, 471 ; fluoride, properties, uses, 456 ; fluoride, occurrence, 153 ; hydride, 452 ; hydrosulfide, action of car- bonic acid on, 457 ; hydroxide, in setting of cement, 455 ; hy- pochlorite, manufacture, 455, prop- erties, uses, 456 ; manganite, formation in Weldon process, 103 ; metachloroantimonate, 268. Calcium nitrate, formation in soil, 199 ; occurrence, potassium nitrate from, 418 ; preparation, manu- facture, uses, 460. Calcium nitride, formation, hydroly- sis, 452 ; occurrence, preparation, properties, 452 ; orthoplumbate, 516. Calcium oxalate, precipitation, solu- bility in strong acids, insolubility in weak acids, use as test for calcium and oxalic acid, 465. Calcium oxide, manufacture, 452; equilibrium with carbon dioxide and calcium carbonate, 453. Calcium phosphates, occurrence, 460 ; superphosphates, 461 ; solubility, relation to fertilizing value, 461, 462. Calcium phosphate, solubility in weak acids, 462; selection of atomic weight, 397 ; silicate, prep- aration, occurrence, 466; stearate, 332. Calcium sulfate, forms, uses, 457; hydrates of and the phase rule, 458; in cement, 454; in hard waters, 463 ; permanent hardness from, 311. Calcium, sulfite, acid, preparation, use in paper manufacture, 457 ; sulfite, acid, use in paper making, 175; sulfida in Leblanc soda pro- cess, 411 ; sulfide, preparation, hydrolysis, 456, recovery of sulfur from by Chance process, 457 ; super- phosphate, manufacture, use, anal- ysis, 461. Calculation of formula of mineral, 356 ; of relative speed of reactions at equilibrium, 151. Calomel, preparation, vapor density, molecular weight, uses, 489. Calorie, definition, 26 ; value at different temperatures, 33 ; value in joules, 33. Calorimeter, 25. Calorimeter, respiration, 313, muscu- lar energy in, 315; mental work in, 316; cuts of, 314, 315. Camphor, in celluloid, 338. Candle power, standard, 293, of illuminating gas, 296, of acetylene, 293, of Welsbach light, 296. Cane sugar, 333. Cannel coals, 281. Canon Diablo, diamonds in meteo- rite from, 274. Caramel, 333. Carat, defined, 448. Carbohydrates, defined, 332. Carbolic acid, see phenol. Carbon, amorphous, heat of com- bustion, 276 ; preparation, prop- erties, 277. Carbon bisulfide and nitric oxide, flame of, 217 ; preparation, prop- erties, uses, formation of sulfo- carbonates from, 317. Carbon, chemical properties of, 281 ; cycle of in nature, 312. Carbon dioxide, amount formed in world by burning coal, 229 ; amount in the ocean, 230 ; amount in air kept constant by ocean, 230 ; coefficient of expansion, 38 ; conditions for escape from solu- tions, 376. Carbon dioxide, critical temperature of, 233; density, 308; accumula- tion in wells and caves, 309, diffusion in rooms, 309 ; deviation from Boyle's law, 35 ; exercises, 322 ; formerly considered poison- ous, 231 ; from burning charcoal, 23. Carbon dioxide, limit for in ventilated rooms, 231 ; per cent in air, 228 ; preparation from calcium car- INDEX 575 bonate, sodium bicarbonate, mag- nesium carbonate, 306 ; proper- ties, isothermals, 307. Carbon dioxide, ratio of specific heats of, 237 ; reduction in plants, 312 ; removed from air by plants, 229 ; solubility, 309 ; sources, 313 ; sources of in air, 229. Carbon, effect on decomposition of barium carbonate, 468 ; electrodes, 279 ; gas, 279 ; heat of combus- tion, 27 ; in steel, use of potassium cupric chloride in determining, 432. Carbon monoxide, absorbed by cu- prous chloride, 433 ; coefficient of expansion, 38 ; deviation from Boyle's law, 35 ; effect of forma- tion on decomposition of barium carbonate, 468 ; formation from calcium oxalate, 466. Carbon monoxide, formation in burning coal, preparation from oxalic acid, 311 ; properties, failure to burn when dry, 311 ; in water gas, danger, 297; poisonous, 312; formation of sodium formate from, 312. Carbon, occurrence, number of com- pounds, importance, 273 ; oxy- chloride, 316 ; oxysulfide, forma- tion from thiocyanates, properties, hydrolysis, 318 ; slowness of reac- tion, 280; suboxide, 316; tetra- chloride from methane, 287 ; va- lence, 282. Carbonate, ammonium, 423. Carbonate ion, r61e in decomposition of carbonates, 375. Carbonates, decomposition by acids, theory, 375 ; formation from car- bonic acid, 310 ; oxides prepared from, 392. Carbonic acid, carbonates and bi- carbonates from, 310 ; deter- mination of free and combined, 464 ; formation, properties, 309. Carbonyl chloride, preparation, prop- erties, 316, hydrolysis, urea from, 317. Carborundum, 349. Carboxyl, characteristic group of organic acids, 328. Carnallite, 414, 478; rubidium in, 424. Carnotite, uranium in, 531. Caro's acid (permonosulfuric acid), 188. Casein, a protein, 343. Cassiterite, 508. Cast iron, composition, gray, white, chilled, 543 ; analyses, 544. Castner-Kellner apparatus for alkali manufacture, 402. Cast steel, 545. Catalysis, 28; definition, 62; in sulfuric acid manufacture by oxides of nitrogen, 178 ; of preparation of sulfur trioxide, 175 ; of synthesis of ammonia, 201. Catalyzer, copper chloride for Deacon process, 103. Cathode, definition, 47, 113. Cation, definition, 48. Cavendish, analysis of air by nitric oxide, 230 ; nearly discovered argon, 235. Celestite, 468. Celluloid, 338. Cellulose, use as fuel, as food, in paper, 337. Cement, dental, composition, 482 ; manufacture, composition, 454; setting, 455. Cementation steel, 545. Cementite, iron carbide, 546. Centimeter-gram-second system, 33. Cerargyrite, 439. Cereal, composition of, 8. Cerium, alloy with iron, 364 ; group of rare earths, 503 ; occurrence, properties, 362 ; oxides, sulfate, double sulfate with sodium, 364; phosphate, 363. Chalcedony, 348. Chalcocite, 428. Chalcopyrite, 428. Chamberlain, use of copper to pre- vent corrosion of iron, 550. " Chamber process " for sulfuric acid, 177. Chameleon solution, 536. Chance process for recovering sulfur, 457. Chapin, separation of praseodymium and neodymium, 504. Charcoal, animal, 278 ; burning in oxygen, 23 ; composition, 280 ; manufacture, 277 ; properties, uses, adsorption by, 278. Charles, law of, 38. Chemical action of radioactive rays, 475. Chemical activity in solutions, 81. Chemical affinity, 29 ; relation to speed of chemical reactions, 149. Chemical energy, nature of, 27; defined, 34. Chemical reactions, equilibrium in, 108 ; effect of concentration on, 24 ; speed of, 149 ; unimolecular and bimolecular, 150. Chemistry, definition, 5 ; study of, 18. 576 INDEX Chicago, typhoid fever in, from water, 83. Chili saltpeter, 210 ; potassium ni- trate from, 418. Chlorates, 127. Chloric acid, 127 ; structure, 130. Chloride, acid, hydrolysis of, 317 ; of lime, manufacture, 455 ; proper- ties, uses, 456. Chlorides, hydrolysis of, 115; of acids, 189 ; preparation by use of sulfur monochloride, 188. Chlorination process for gold, 446. Chlorine and oxygen, comparison of heats of combination, 108 ; action on ammonia, 209 ; bleaching by, 106; burning in hydrogen, 118. Chlorine, combination with other ele- ments, 104 ; determination of atomic weight of, 130; effect of light on reaction with hydrogen, 105 ; effect of moisture on combination with other elements, 105 ; exercises, 116 ; from electrolysis in alkali manufac- ture, 402. Chlorine hydrate, phases, 107; list of oxides and oxygen acids of, 123 ; occurrence, 100 ; dioxide, 127. Chlorine, preparation by electrolysis of sodium chloride, 100 ; by the Deacon process, 102 ; from hydro- chloric acid and manganese dioxide, 101 ; from hydrochloric acid and potassium permanganate, 102 ; by oxidation of hydrochloric acid, 100 ; by Weldon process, 102. Chlorine, properties, 104; reaction with water, 106. Chlorites, 127. Chloroaurates, 450. Chloroauric acid, 450. Chlorocuprous acid, 432. Chloroform from methane, 287. Chloroplatinates from amines, 423. Chloroplatinic acid, preparation, 565 ; from platinum black, 566 ; use in photography, 445. Chloroplatinous acid, 565. Chloroplumbic acid, 518. Chloroplumbous acid, 518. Chlorosulfonic acid, preparation, properties, 189. Chlorous acid, 127 ; structure, 130. Cholera from impure water supply, 83. Chrome alum, preparation, 527. Chrome green, 525. Chrome iron ore, 524 ; decomposition of, 527. Chrome tanning, 527. Chrome yellow, 527 ; constituent of chrome green, 525. Chromic anhydride, preparation, use to oxidize carbon, 528. Chromic chloride, hydrates, 525; isomeric, 526 ; preparation, proper- ties, 525 ; theory of, 526. Chromic hydroxide, formation, com- position, properties, 525 ; chro- mites, 525. Chromic oxide, preparation, use as pigment, 525. Chromite, 524. Chromites, 525. Chromium, alloy with cobalt, 557; occurrence, metallurgy, properties, uses, 524 ; preparation by thermite process, 524, 498. Chromium trioxide, preparation, use to oxidize carbon, 528. Chromous chloride, preparation, properties, 525. Chromyl chloride, preparation, prop- erties, structure, hydrolysis to di- chromic acid, 528. Chugaev, separation of nickel and cobalt, 561. Cinnabar, 485, 489. Citral, use in making ionpne, 328. Citrate-soluble phosphoric acid, 461. Citric acid, structure, source, 330 ; salts, 331. Clark cell, electromotive force, 437. Clarke, F. W. Composition of the crust of the earth, 11. Classification, of metals, 370, 371 ; of the elements, 132. Clays, formation of, 494; manufac- ture of aluminium oxide, hydro- chloric acid and sodium carbonate from, with salt, 496. Cleveite, discovery of helium in, 237 ; uranium in, 531. Clinker, cement, 454. Clouds, conditions of formation, 232. Clover, fixation of nitrogen by, 199. Coal, energy of from sunlight, 230 ; formation, varieties, composition, 280. Coal gas, percentage composition, 299 ; tar, character, 294 ; tar dips, phenol in, 326. Coals, coking, noncoking and cannel, 281. Cobalt ammines, 559 ; glass, smalt, 558 ; nitrate, use in blowpipe analysis, 558. Cobalt, occurrence, properties, alloy with chromium oxides, 557 ; sepa- ration from nickel by dimethyl- glyoxime, 560 ; sulfide, formation, slow solubility, 558. Cobaltous chloride, properties, sym- INDEX 577 pathetic ink, 557 ; hydroxide, prep- aration, properties, 557. Cocaine, 343 ; similar synthetic alka- loids, 343. Coefficients of expansion of air, O 2 , N 2 , NO, H 2 , A, He, CO, CO 2 , SO 2 , 38. Coining value of gold, 447. Coins, gold, 448 ; nickel five cent, 560; silver, 442. Coke, composition, 280 ; manufac- ture, uses, 278. Coking coals, 281. Colemanite, boric acid from, 365. Collodion, 338. Collection and storage of gases, 22. Colloidal arsenic trisulfide, 261 ; silicic acid, 353 ; solutions, 262. Colloids, contrasted with crystalloids, 357 ; precipitated by bivalent ions, 263 ; properties, 262. Columbite, 523. Columbium (niobium), discovery, oc- currence, properties, compounds, 523. Combining volumes, law of, 89 ; weights, law of, 13. Combustion, 24 ; heat of, 25. Complex cyanides, 319. Complex ions, formation of, 378 ; ' evidence for existence of, 379. Composition of air demonstrated by Lavoisier, 19 ; of the crust of the earth, 11 ; of pure substances expressed in multiples of atomic weights, 17. Compounds, definition, 9 ; general methods of preparing, 372-379. Compressed gases, cooling on ex- pansion, 233. Comstock, dependence of mass on velocity, 5. Concentration and speed of reaction, 149 ; effect of on chemical reac- tions, 24. Congress water, 309. Coniine, 342. Conservation of energy, 6 ; of matter, 6. Constant proportion, law of, 12. Contact mass " for catalyzing formation of sulfur trioxide, 176. Converter, Bessemer, 547. Cooking of starchy foods, 336. Copper, addition to steel, 431 ; action of nitric acid on, 213 ; acetylide, see copper carbide. Copper, alloys of, 431 ; annual pro- duction, value, 430 ; properties, effect of impurities on conduc- tance of, 430; uses, 431 ; arsenite and acetate (Paris green), 259. Copper, basic carbonate of, 428; formation, 431. Copper carbide, preparation, 292 ; chloride, catalyzer for Deacon process, 103 ; detected by sodium metaphosphate, 253 ; electro- lytic refining, 429. Copper, exercises, 450 ; ferrocyanide, use in semipermeable membranes, 358 ; hydroxide, precipitation, de- composition, 431 ; in five cent piece, 560 ; metallurgy of, 428. Copper, occurrence, metallurgy, 428 ; electrolytic refining, 429. Copper oxide, from copper hydroxide, 431 ; from wire, nitrate, use, 432 ; in " oxone," 21 ; use in deter- mination of the composition of water, 69. Copper, precipitation by sodium thiosulfate, 409 ; precipitation by iron, 435 ; prevention of corro- sion of iron by, 550; pyrites, 161, 428 ; separation from cadium as cuprocyanide, 435. Copper, sulfate, hydrates, uses, 433 ; reaction with sulfuric acid, 173 ; titration of with potassium iodide and sodium thiosulfate, 433. Copperas, 554. Corn sirup, 334. Cornwall, tin from, 508. Corpuscle, same as electron, 181. Correction of volume of gas for temperature, 39 ; zero and stem, for thermometers, 486. Corrections for readings of barom- eter for altitude, 37 ; for read- ings of barometer for glass and brass scales, 36 ; for readings of barometer for latitude, 37. Corrosion of iron, prevention by copper, 550. Corrosive sublimate, preparation, 489 ; properties, uses, antidote for, 490. Cort, invention of puddling process, 544. Cotton goods, fire-proofing of, 513. Cowles Brothers, electric furnace, 495 ; use of electric furnace for aluminium, 391. Cowles, process for decomposition of clay, 496. Cowper-Cowles, Sherard, sherard- ized iron, 482. Cream of tartar, tartar emetic from, 266 ; use in jellies, in baking pow- ders, 330. Cretinism, connected with defi- ciency of iodine, 144. 578 INDEX Critical temperature, defined, 233; discovery of, 232 ; of carbon dioxide, 233 ; relation to lique- faction of air, 233. Crookes, discharge of electricity through rarefied gases, 471 ; dis- covery of thallium, 507. Crown glass, 467. Cryolite, 153, 495 ; use for manu- facture of aluminium, 495. Crystallization, 8 ; water of, 82. Crystallographic systems, 193. Crystalloids, 357. Crystals, definition of, 192. Cubic centimeter, true volume of, 31. Cupellation, 440. Cupric chloride, preparation, ioniza- tion, 432; nitrate, hydrates, de- composition of, 434 ; oxide, prep- aration, use, 432 ; sulfide, forma- tion, properties, 433. Cuprous chloride, preparation, prop- erties, 432; use to absorb car- bon monoxide, 433 ; cyanide, for- mation, complex salt with potas- sium cyanide, 434; iodide, forma- tion in titrating copper, 433; oxide, formation, in testing for glucose, hydrazine, etc., proper- ties, 432; sulfide in matte, 429; sulfide, occurrence, 428 ; forma- tion, 433. Cyanides, complex, 319; formation, preparation, 319; from ammoni- acal gas liquors, 319. Cyanide process, for gold, 446; for silver, 441. Cyclopentene, structure, 285. Cyclopropane, structure, 284. Dalton, atomic theory, 14 ; formula for water, 91 ; law of multiple proportion, 88; law of partial pressures, 41, 77 ; view of mole- cules of the elements, 93. Damascus blade, tungsten in, 530. Davy, discovery of metallic potassium, 415 ; discovery of metallic sodium and potassium, 399; injured by nitrogen trichloride, 224; safety lamp, 287 ; study of explosion of fire damp, 287. Deacon process for chlorine, 102; equilibrium of, 109. Debye, quantum theory, 398. Degree of ionization, measurement of , 380 ; table, acids, 383 ; bases, 383 salts, 384. Degrees of freedom, 77. Dehydration by sulfuric acid, 182. Dekahydronaphthalene, 285. Deliquescence, 82. Density, criterion of pure substance, 12 ; of gases, table, 95. Dental cement, composition, 482. Derivatives of ammonia, 205. Determination of weight of a liter of gas, 40. Detonating caps, 491. Developing in photography, 445. Dewar flasks for liquid air, 235. Dextrin, manufacture from starch, uses, 336. Dextrose, see Glucose. Dialysis, 357. Dialyzed iron, 555. Diamond, heat of combustion, 276 ; artificial, 275, natural, 275; uses, properties, 275. Diastase, 344. Dibasic acids, defined, 183. Dibromoethane, 291. Dicalcium phosphate, 249. Dichloroethane, 291. Dichromic acid, from hydrolysis of chromyl chloride, 528. Diet, salt essential in, 406. Dietary, average American, 347. Diffusion of gases, 56 ; law of, 59. Digestion, colloidal solutions, 263. Dimethylglyoxime, precipitant for. nickel, 560. Diphosphorus pentasulfide, 254. Disilicates, 356. Disilicic acid, 355. Disinfectant, formaldehyde, 327 ; sulfur dioxide, 174. Disintegration of atoms, possible by radium emanation, 475. Disodium phosphate, 249 ; hydroly- sis, alkaline reaction of, 251 ; phenolphthalein as indicator for, 251 ; uses, arsenic as impurity in, 410. Displacement of equilibrium, 152. Dissociation, definition, 59 ; pressures of calcium carbonate, 453 ; pres- sure of silver oxide, 443 ; of am- monium hydroxide to ammonia and water, 204 ; of calcium car- bonate, phase rule, 453 ; of sulfuric acid, 180 ; of water, 59. Distillation, 8. Dithionic acid, 188. Divariant, definition, 77. Dixon, explosion waves, 301. Dolomite, 478. Double decomposition, defined, 81. Double refraction of crystals, 196. Dry batteries, use of zinc in, 482. Dry plates, photographic, 444. Drying of gases, 54. INDEX 579 Dulong and Petit, law of, 396; relation to Ayogadro's law, 397. Dumas, determination of the compo- sition of water, 69. Durax glass, 467. Dust, explosion of with air, 289. Dutch process for white lead, 520. Dydimium, separation into praseo- dymium and neodymium, 504. Dyeing, lakes for, 501. Dyes, 340 ; use of sodium nitrite in manufacture of, 410 ; substantive, adjective, 342. Dysprosium, compounds, 505. Earth, composition of crust of, 11 ; mean density of, 540. Earthenware, 501, glazing, 502. Effervescent waters, 309. Efflorescence, 82. Egypt, early manufacture of iron in, 390 ; sodium carbonate from, 450. Einstein, quantum theory, 398. Eka-aluminium, same as gallium, 506. Ekaboron, same as scandium, 136, 503. Ekeberg, discovery of tantalite, 523. Electric furnace, carborundum in, 349 ; use for aluminium, 391. Electrical batteries, use of zinc in, 481, theory of, 435 ; horse power, 34 ; unit charge, 438 ; units, 33. Electrochemical theory, influence on formulas of minerals, 356. Electrodes, carbon, 279. Electrolysis, migration of ions in, 113 ; of dilute sulfuric acid, 47, 9 ; sodium hydroxide by, 402. Electrolyte, definition, 48. Electrolytic methods in metallurgy, 391. Electromotive force, relation to solu- tion pressure, 435. Electromotive series, 435, table, 436. Electron, relation to Faraday's law, 438. Electron theory, 181 ; as explanation of ionization, 182 ; relation to ionization, 206; relation to prop- erties of metals and non-metals, 370. Electronegative elements, denned, 437. Electroplating, copper, 433. Electropositive elements, defined, 437. Electrotyping, 433. Elements, absolute potential of, 436 ; atoms of probably complex aggre- gates, 138 ; classification of, 132 ; definition, 9 ; life of, 474 ; melting points of, absolute, 135; melting points of, 372, table, 373 ; metallic in periodic system, 136 ; missing in Group VII, possible reason, 533 ; molecules of, 93. Elements, non-metallic in periodic system, 136 ; radioactive, series of, 475 ; specific heat of, 397 ; symbols of, 11; table, of familiar, 371; table of groups of, 371 ; table of non-metallic, 348. Emery, 494; artificial, 500. Endothermic compounds, defined, 225 ; explosive decomposition, 225. Endothermic reactions, 215. Energy, conservation of, 6 ; defini- tion, 6; muscular, in respiration calorimeter, 315; of coal from sunlight, 230. Energy required to decompose a gram equivalent, 438 ; units of, 32 ; varieties of, 6. England, destruction of forests for iron manufacture, 540. English laws for glazes, 502. Enzymes, 344. Epsom salts, 478, 480. Equations, writing of, 49 ; writing of, for reactions between acids and bases, 156. Equilibrium between gaseous and solid phases, 443, silver, silver oxide and oxygen, 443 ; between water and water vapor, 76; dis- placement of, 152 ; effect on, of removing one of the reacting sub- stances, 152 ; for combination of nitrogen and hydrogen, 201 ; for formation of nitric oxide, 216; hydrogen, iodine and hydriodic acid, 146. Equilibrium in chemical reactions, 108 ; in gas flame, 300 ; illustra- tion of that between water and water vapor, 76 ; in ionization of orthophosphoric acid, 250 ; in neutralization, 385. Equilibrium of carbon monoxide, carbon dioxide, hydrogen and water vapor, 300 ; of Deacon process, 109 ; of reaction between hydro- chloric acid and oxygen, 108. Equivalents, relation to Faraday's law, 438. Erbium, compounds, 505. Erg, defined, 32. Ethane, structure, 284; substitu- tion products of, 291. Ethene, 290. Ether, ethyl, 290. Ethyl alcohol, ether and ethylene from, 290; ethyl chloride from, 580 INDEX 245 ; manufacture, properties, uses, 325; absolute, denatured, 325; structure, 323. Ethyl borate, 367; chloride, formed from ethyl alcohol, 245; ether, formation, 290; iodide, relation to structure of ethyl alcohol, 323. Ethylene, addition compounds, 291 ; bromide, 291 ; chloride, formation from ethylene, 291 ; in illuminat- ing gas, 295 ; preparation, prop- erties, uses, formation and de- composition, 290 ; ratio of specific heats of, 237 ; structure, 292. Eudiometer, description, 67. Europium, compounds, 505. Eutectic point, defined, 488. Evaporators, triple and multiple- effect, 405 ; Yaryan, 405. Exercises, aluminium, 507 ; atmos- phere, 239; Avogadro's law, 99; calcium, barium, 476 ; carbon dioxide, cyanides, 322 ; chlorine, 116; copper, silver, gold, 450. Exercises, Group V, 272 ; hydrocar- bons, 305 ; laws of gases, 43 ; mag- nesium and mercury, 493 ; nitro- gen, 225 ; phosphorus, 255 ; silicon, boron, 368; sulfur, 196; writing equations, 156. Exothermic reaction, definition of, 215. Explosions, definition, 62 ; of en- dothermic compounds, 225 ; of methane and air, 287; of dust and air, 289 ; waves, 301. Factories, humidity in, 232. Families of elements, table, 371. Faraday, injured by nitrogen tri- chloride, 224; law, 338; liquefac- tion of chlorine, 108. Fast colors, 340. Fats, composition, 331 ; use in making soaps, 332 ; reduction with the aid of colloidal palladium, 564 ; soft soap from, 414. Fatty acids, calcium salts of, 332. Fehling's solution, formula for, use, to detect glucose, 335. Feldspars, 348 ; use as glaze, 502. Ferric acetate, mordant, 342. Ferric chloride, hydrate, anhydrous, 554 ; molecular weight, 555 ; hydrolysis, precipitation of ferric hydroxide from, 553; theory of hydrolysis of, 386. Ferric ferrocyanide, 320 ; decomposi- tion with sodium hydroxide, 321 ; hydroxide, precipitation, 555; by barium carbonate, 553 ; hydroxide, to absorb hydrogen sulfide, 295; oxide, preparation, manufacture, use, 555. Ferric sulfate, fuming sulfuric acid from, 556 ; preparation, alums from, 556 ; reduction by hydrogen sul- fide, 171 ; sulfide, formation, prop- erties, 556 ; thiocyanate, test for iron, 322 ; thiocyanate, use as test for iron, 556. Ferrite, defined, a-, /3- and 7-, rela- tion to tempering, 546 ; table of properties, 546. Ferromanganese, use in chilled cast iron, 543. Ferrous bicarbonate, formation, in mineral waters, ores from, 554 ; carbonate, ore, 540 ; carbonate, properties, 554. Ferrous chloride, absorption of nitric oxide by solutions of, 554 ; prepara- tion, properties, 553 ; preparation, structure, 552. Ferrous chromite, 524 ; preparation of potassium chromate from, 527 ; ferricyanide, 321 ; hydroxide, for- mation, properties, 553 ; hydroxide, use to reduce indigo, 341 ; manga- nese tungstate, 530 ; metatantalate, 523 ; metacolumbate, 523 ; oxide, 553 ; silicate, formed in metallurgy of copper, 429. Ferrous sulfate, preparation, proper- ties, 554 ; sulfide, oxidation to ferric hydroxide in gas purifiers, 296 ; sulfide, preparation, prop- erties, use, 556. Ferro vanadium, 522. Ferrum, 11. Fertilizer, calcium cyanamide as, 463 ; slag from basic steel process as, 548 ; various phosphates in, 461. Filters, charcoal, inefficient, 83. Fire damp, 286, 287. Fireproofing fabrics, with sodium tungstate, 530 ; of cotton goods, 513 ; with water glass, 353. Fixing in photography, 445. Flame colors of calcium, strontium and barium, 471 ; gas and candle, structure, 299. Flame, oxidizing, blowpipe and Bun- sen burner, 304; reducing, blow- pipe and Bunsen burner, 304. Flames, cause of luminosity, 299 ; re- versed, 304 ; temperature of, 302 ; source of carbon in, 299. Flint, 348. Flint, atomic weight of tellurium, 190. Flint glass, 467. INDEX 581 Flour, explosive when mixed with air, 289. Flowers of sulfur, 160. Fluoboric acid, 367. Fluorine, occurrence, 153 ; prepara- tion, 153 ; properties, 154. Fluorite, 153, 452. Fluosilicic acid, preparation, proper- ties, 350. Foods, animal, vegetable, inorganic constituents, 347 ; average Ameri- can dietary, 347. " Fool's gold," 556. Forge, reduction of iron in, 540. Formaldehyde, preparation, use as disinfectant, 327 ; use in food for- bidden, 328 ; probable formation by reduction of carbon dioxide by plants, 312. Formalin, see formaldehyde. Formic acid, from mercuric fulminate, 491 ; preparation from carbon monoxide, uses, 329. Formula of mineral, calculation, 357. Formular solutions, definition of, 183. Formulas, meaning of, 16; struc- tural, basis for, 323. Franklinite, 481. Frasch, process for getting sulfur, 161. Fraunhofer lines, discovery, 424. Free metals, occurrence, 439. Freedom, degrees of, 77. Freezing points of solutions and os- motic pressure, 360 ; law for de- pression of, 112. French Revolution, relation to alkali industry, 400 ; sources of saltpeter during, 418. Friedel, discovery of carbon in meteo- rite from Canon Diablo, 274. Fructose, formation from cane sugar, properties, 335; from cane sugar, 334. Fruit, sulfuring of, 164. Furnace, electric for aluminium, 495 ; electric for calcium carbide, 462 ; electric for carbon disulfide, 317; electric for carborundum, 349 ; electric for phosphorus, 241 ; re- generative, for open hearth steel, 549, 550. Fusible alloys, 269. Gadolinium, compounds, 505. Galena, 161 ; lead ore, 513 ; silver in, 439. Gallium, compounds, same as eka- aluminium, 506. Galvanic cells, calculation of electro- motive force, 437 ; theory, 437. Galvanized iron, manufacture, theory of conduct, 482. Garnet, 349 ; calculation of formula of, 356 ; orthosilicate, 355. Gas, blast furnace ; use for heating blast, for gas engines, 542 ; carbon, manufacture, uses, 279. Gas, correction of volume for changes of pressure, 38 ; for temperature, 39. Gas, determination of weight of a liter of, 40 ; effect of pressure on, 34 ; effect of temperature on, 38 ; effect of water vapor on volume and pressure of, 76. Gas, illuminating, 295 ; " ideal," 94 ; iron pentacarbonyl in, 561 ; laws, graphical representation, 42, 43 ; lighters, cerium-iron, 364 ; liquors, ammoniacal, 202. Gas, oil, Pintsch, water, 296 ; pro- ducer, 297 ; producer, use in re- generative furnace for steel, 550. Gases, collection and storage of, 22 ; diffusion of, 56 ; drying of, 54 ; Henry's law for solution of, 165 ; spectra of, 427 ; kinetic theory of, 58 ; laws of, 34, 38 ; laws of, exer- cises, 43 ; monatomic, ratio of specific heats for, 236; table of densities, 95. Gasoline, 289. Gasometer, description of, 22. Gastric juice, hydrochloric acid in, 406. Gayley, dry blast for iron manufac- ture, 542. Gay Lussac's law of combining volumes, 89 ; tower, use in manu- facture of sulfuric acid, 179. Geology, relation of radiochemistry to, 475. Germanium, discovery, compounds, 361. German silver, 560. Germicides, sulfites as, 174. Gibbs, Willard, phase rule, 107. Glass, etching of, 154 ; manufacture, properties, 466 ; crown, flint, strass, paste, Bohemian, Jena, Re- sistanz, Non-sol, hard, borosilicate, Durax, 467 ; opaque, stannic oxide in, 511 ; soluble or water, 353. Glazes, for earthenware and porce- lain, 502. Glover tower, use in sulfuric acid manufacture, 179. Glucinum, name for beryllium, 451. Glucose, from hydrolysis of cane sugar and of starch, properties, 334 ; formation in diabetes, defec- tion, 335. 582 INDEX Gluten from cereals, 8, 335. Glycerol, source, use, 326; by-prod- uct in manufacture of soap, 322. Glyceryl nitrate, 326. Gmelin, manufacture of ultramarine, 502. Goiter, connected with deficiency of iodine, 144. Gold, chlorides of, 450 ; dioxide, 448 ; exercises, 450 ; monochloride, 450 ; dichloride, 450 ; trichloride, 45 ; monoxide, 448. Gold, occurrence, 445 ; in sea water, washing for, hydraulic mining, cyanide process, 446 ; annual pro- duction, coining value, 447 ; oxides of, hydroxide, 448; properties, alloys, use of term carat, 448. Gold, recovery from copper, 430 ; trichloride, use in photography, 445; trioxide, 448. Goldschmidt's thermite process, 497. Goldthwaite, jelly making, 337. Goodwin, preparation of metallic calcium, 452. Graebe, discovery of structure of alizarin, 341. Graham, crystalloids, colloids, 357. Graham, study of diffusion of gases, 57 ; of liquids, 357. Gram, definition, 31. Gram atom, definition, 25. Gram equivalent, definition of, 184. Gram molecular volume, definition, 94. Gram molecule, definition, 25. Granites, decomposition to shales, clays and soils, 494. Graphical representation of gas laws, 42, 43. Graphite, abnormal specific heat, 397 ; anodes for alkali manufac- ture, 402 ; heat of combustion, 276 ; in cast iron, 543, 544 ; in colloidal solution as lubricant, 277 ; sources, formation, manufacture, proper- ties, 276 ; uses, 276. Gravitation, law of, 1. Gravity cell, description of, electro- motive force of, 437. Greenland, cryolite from, 495. Green vitriol, 554. Grotto del Cano, carbon dioxide in, 308. Groups of elements, table of, 371. Group zero, 236. Group I, alternate metals of, general properties, 428. Group II, alternate metals of, general properties, 478 ; general properties, 451 ; solubility of sulfides of, 491. Group III, metals of, 494. Group IV, metallic elements of, 361, 362. Group V, alternate elements, 522; exercises, 272 ; general properties, 256 ; table of oxides, chlorides, acids, sulfur acids and hydrides, 271. Group VI, alternate elements, 524; metallic elements of, 192 ; proper- ties of elements of, 191. Group VII, metallic elements of, 156, 533 ; missing elements, possible reason, 533. Group VIII, general statement, 539 ; metals of, general properties, table, 562. Guignet's green, 525. Gun cotton in smokeless powder, 419 ; manufacture, properties, uses, 338. Gun metal, 509. Gunpowder, composition, 418 ; de- composition, manufacture, theory of burning and explosion, 419. Gypsum, 161, 452; plaster of Paris from, 457 ; phase rule, 458 ; con- ditions for formation of, 458, 460 ; vapor pressure of, table, 459. Haber, dissociation of hydrobromic acid, 146 ; synthesis of ammonia, 201. Haddam, columbite from, 523. Hall, manufacture of aluminium, 495 ; process, 496. Halogen acids, table of, 139 ; family, 139 ; meaning of name, 139. Halogens, compounds with hydrogen and oxygen, 139 ; general prop- erties of, 138; table of, 139. Hamburg, cholera in, from water, 83. Hampson, liquid air machine, 234. Harcourt, atomic weight of tellu- rium, 190. Hard glass, 467. Hard waters, 310 ; permanent hard- ness, temporary hardness, soften- ing by boiling, 463 ; by milk of lime, 464 ; by sodium salts, 464. Hatchett, discovery of columbium, 523. Haynes, cobalt chromium alloy, 557. Heat of combination of chlorine with H 2 , Na, Zn, Cu, P, 108; of oxygen with Ek, Na, Zn, Cu, P, 108. Heat of combustion, 25 ; of bitumin- ous coal, calculation of, 44 ; of C, S, P, Fe, Hg, 27 ; of hydrogen, 65. INDEX 583 Heat, mechanical equivalent of, 6 ; units of, 33. Helium, coefficient of expansion, 38 ; discovery in sun and in cleveite, 237; distribution, 238; from radium, 9. Hematite, 540. Hemihedral forms of crystals, 193. Henry's law, applied to carbon di- oxide, 309 ; for solution of gases, 165. Heptane, 283. Heptene, 283. Heptine, 283. Hexaaquochromic chloride, 526. Hexagonal System (crystallography), 194. Hexane, 283. Hexanitrocellulose, 338. Hexathionic acid, 188. Hexene, 283. Hexine, 283. Hillebrand, gases from uraninite, 237 ; use of sodium pyrosulfate for solution of alumina, 408. Holmberg, holmium, 505. Holmium source, 505. Homologue, definition, 289. Honey, invert sugar in, 334. Hornblende, metasilicate, 355. Horn silver, 439. Hulett, purification of mercury, 485. Humidity of rooms, 232. Hunyadi water, magnesium sulfate in, 480. Hydrargyrum, 11. Hydrates, defined, 81 ; of nitric acid, 211 ; of sulfuric acid, 181 ; vapor pressure of, 82. Hydration, water of, 82. Hydrazine hydrochloride, 222 ; prep- aration, 222; structure, 220; trinitride, structure, 221. Hydriodic acid, constant boiling solution, 146 ; equilibrium with hydrogen and iodine, 146 ; forma- tion from potassium iodide, 145; heat of formation of, 153 ; prep- aration from iodine, phosphorus and water, 145 ; rate of formation and decomposition, 150 ; reduc- tion of sulfuric acid by, 145. Hydrobromic acid, constant boiling solution of, 143 ; dissociation of, 146 ; preparation from hydrogen and bromine, 143 ; preparation from potassium bromide, 142 ; reduction of sulfuric acid by, 142 ; preparation from bromine, phos- phorus and water, 143. Hydrocarbonate ion, r61e in decom- position of carbonates, 375. Hydrocarbons, exercises, 305 ; table of, 283. Hydrochloric acid, and oxygen, equilibrium of reaction between, 108 ; by-product in making alumin- ium oxide from clay, 496 ; diffi- culty with in Leblanc soda process, 411. Hydrochloric acid, constant boiling solution of, 120. Hydrochloric acid, determination of composition by volume, 120 ; de- termination of composition of, 130 ; deviation from Boyle's law, 35 ; electrolysis to show composi- tion by volume, 120 ; formation from chlorine and hydrogen, 118; in gastric juice, 406 ; oxidation of, 100. Hydrochloric acid, preparation from salt and sulfuric acid, 118; prop- erties, 119; reactions with hy- droxides and oxides, 121 ; re- actions with metals, 120 ; reac- tions with oxidizing agents, 122; reaction with potassium perman- ganate, 159. Hydrocyanic acid, preparation, prop- erties, uses, 319. Hydrofluoric acid, constant boiling solution of, 155 ; etching of glass by, 154 ; molecular weight, 155 ; preparation, 154 ; properties, 154. Hydrogen, 45 ; apparatus for prep- aration of, 53 ; burning in chlo- rine, 118; chemical properties of, 59 ; coefficient of expansion, 38. Hydrogen, combination of iodine with reversible, 146 ; combination with oxygen reversible, 372 ; deter- mination of atomic weight, 72 ; deviation from Boyle's law, 35 ; formed in preparing phosphine, 243. Hydrogen, formerly unit for atomic weights, 68 ; heat of combustion of, 65 ; ions in liquid ammonia, 208 ; in the atmosphere, 45. Hydrogen, occurrence, 45 ; nascent, 213 ; palladium semipermeable membrane for, 358, 360. Hydrogen, preparation, 47 ; prepara- tion by " hy drone," 52; prepara- tion by sodium and potassium, 50; preparation by zinc and acids, 52 ; preparation from iron and steam, 49 ; properties of, 55 ; purification of, 54. Hydrogen peroxide as oxidizing agent, 84 ; as reducing agent, 85 ; 584 INDEX preparation, 83; properties, 85; uses, 85 ; structure, 86 ; tests for, 86. Hydrogen phosphide, liquid, 243; selenide, preparation, properties, 190; silicide, 349. Hydrogen sulfide, decomposition by heat, 164 ; decomposition of solu- tion in air, 166; formation from elements, 164 ; ionization of, 168 ; occurrence, 164. Hydrogen sulfide, Parsons apparatus for generating, 165 ; reaction with iodine, 171 ; reducing agent, 171 ; removal from hydrogen, 55 ; re- moval from illuminating gas, 295 ; solubility in water, 166. Hydrogen telluride, 190. Hydrolysis of chlorides, 115; of chlorides of phosphorus, 245 ; of salts, theory, 385; of sulfides, 171. " Hy drone," preparation of hydrogen by, 52. Hydronitric acid, preparation, 223 ; properties, 223; structure, 221. Hydrosulfate ion, role in decomposi- tion of salt, 375. Hydrosulfites, see hyposulfites. Hydrosulfuric acid, 167. Hydrosulfurous acid, 186. Hydrotetrachloroantimonic acid, 267. Hydroxide, meaning of name, 21. Hydroxides, oxides prepared from, 392; preparation from metals, 392; from salts, 393. Hydroxylamine, from mercuric ful- minate, 491 ; preparation, 221 ; properties, uses, 222. Hydroxylammonium sulfate, 222. Hygrometer, moist bulb, 232. Hypo-, prefix, 123. Hypobromite, sodium, 143. Hypochlorites, autoxidation to chlo- rates, 125. Hypochlorites, preparation, proper- ties, uses, 124. Hypochlorous acid, preparation, properties, 124; formation from chlorine and water, 106 ; structure, 130. Hypochlorous anhydride, 126. Hyponitrous acid, preparation, prop- erties, 221. Hypophosphites, use, preparation, 248; structure, 247. Hypophosphoric acid, formation, salts, 254. Hyposulfite, old name for thiosul- fate, 187. Hyposulfites, preparation, 186. Hypotheses, 2. lanke, George, value of the calorie at different temperatures, 33. -ic, suffix, meaning, 30 ; use for acids, 123. Ice machines, 204. Ice, vapor pressure of, 75. -ide, suffix, use, 47. Illuminating gas, ammonia from, 201 ; manufacture, composition, 295 ; removal of hydrogen sulfide from, 295. Imide, definition of, 206. Indestructibility of matter, 6. India, diamonds from, 275 ; potas- sium nitrate from, 418. India rubber as semipermeable mem- brane, 358. Indicators, acidity or alkalinity of at change of color, table, 388 ; chemical nature, 389 ; choice of an, 389 ; definition and list, 122 ; for determining free and combined carbonic acid, 464 ; use of, 387 ; for weak acids and bases, 389. Indigo, source, synthesis, use as dye, 341 ; use of sodamide in manu- facture of, 410. Indigo white, 342. Indium, discovery, atomic weight, compounds, 506. Inductive reasoning, 13. Infusorial earth, use for dynamite, packing and scouring, 351. Ingle, separated Bunsen flame, 301. Ink, sympathetic, 557. Insolubility, effect on a reaction, 376. International Bureau of Weights and Measures, 31 ; scale of tempera- Invariant, definition, 78. [tures, 32. Invert sugar, 334. lodic acid, 139. Iodine and starch, 145 ; com- bination with iodine reversible, 146; in thyroid gland, 144; monatomic at high temperatures, 144 ; occurrence, 144 ; liberated by nitrous acid, 145 ; liberated in titrating copper, 433 ; positive in iodine trinitride, 224. Iodine, properties, 144 ; reaction with hydrogen sulfide, 171 ; removal from hydriodic acid, 145 ; sodium tetrathionate formed by action of on thiosulfate, 409; solutions standardized by arsenious oxide, 260; tincture of, 144; trinitride, iodine positive in, 223 ; trinitride, preparation, 223. INDEX 585 Ion, definition, 48. Ionium, 475. lonization, calculation of degree of from freezing points of solutions, 381. lonization, degree of, table acids, 383 ; bases, 383 ; salts, 384 ; effect of degree of, neutralization, 384 ; effect on freezing points of solu- tions, 112, 381. lonization, evidence of, 112; ex- plained by the electron theory, 182 ; measurement of degree of, 380 ; of acids, 167 ; of acids, relation to hydrolysis of sugar, 381 ; of ammonium hydroxide, 203 ; of compounds of cadmium and mer- cury, 491 ; of first and second hydrogen atom of acids, 168 ; of hydrates of chromic chlorides, 526. lonization of oxalic acid, relation to solubility of calcium oxalate, 465, 466 ; of sulfuric acid, 181 ; of sulfurous acid, 174 ; of trimethyl ammonium hydroxide and tetra- methyl ammonium hydroxide, 204 ; of water, 171, 383 ; relation to elec- tron theory, 206. Indium, properties, uses, oxides, chlorides, double salts, 565 ; tetra- chloride, 565. Iron, burning in oxygen, 23. Iron carbide, 543, relation to tem- pering of steel, 546. Iron, dialyzed, 555 ; discovery of metallurgy of, 390 ; disulfide, 556 ; heat of combustion, 27 ; in pro- teins, 343 ; magnetic oxide of from burning iron, 23. Iron, precipitation of copper by, 435 ; pentacarbonyl, 561 ; pyrites, 556. Iron, reasons for importance of, 539 ; ores, history of use, 540, blast furnace, 541 ; cast iron, 543 ; wrought, 544 ; cementation and cast steel, tempering of steel, 545 ; Bessemer steel, 547 ; open hearth steel, 548 ; analyses, 544, 551 ; 'alloy steels, 552 ; compounds, 552 ; tetracarbonyl, 561. Isobutane, structure, 284. Isomer, definition, 323 ; history of name, 511. Isometric System (crystallography), 193. Isothermals of carbon dioxide, 307. -ite, use for salts, 124. Jakowin, reaction of chlorine with water, 106. James, separation of thulium, 506. Jasper, 348. Jelly, conditions for making, 337. Jena glass, 467. Johnson, sherardized iron, 482. Jqule, defined, 26 ; relation to calorie. 33. Joule-Thomson, effect on expansion of compressed gases, 233. Kalium, 11. Kalk-Stickstoff, 463. Kamm, use of potassium silver co- baltinitrite in analysis, 559. Kaolin, 349; formation, 494; ortho- silicate, 355. Kekule, formula for benzene, 285. Kelly, inventor of process for steel, 547. Kelvin, size of molecules, 16 ; theory of the source of oxygen in the air, 230. Kerosene, manufacture, flashing point, 290. Ketones, 327. Kieselguhr, 359. Kilogram-meter, defined, 32; value in ergs, 33. Kilowatt, defined, 34. Kimberly, diamonds from, 275. Kindling temperature, 24. Kinetic theory of gases, 58 ; rela- tion to osmotic pressure, 361. Kipp generator, 54. Kirchoff, discovery of rubidium and caesium, 424 ; discovery of spec- trum analysis, 424. Knietsch, history of catalytic sulfur trioxide, 175. Konigsberger, critical temperature of mercury, 486. Kremann, hydrates of nitric acid, 212. Krogh, estimate of coal burned annually, 229. Krypton, discovery, 238. Kurnakow, sodium amalgams, 488. Kiister and Kremann, hydrates of nitric acid, 212. Lacquers, use of nitrocellulose in, 338. Lactic acid, formation, structure, 330. Lactose, from milk, use, 334. Ladenburg, density of ozone, 97. Lake Superior region, copper from, 428. Lakes, for dyeing, 501. Lampblack, 277. Landolt, demonstration of conserva- tion of matter, 6. Langworthy, respiration calorimeter, 313. Lanthanum, compounds, 503. 586 INDEX Latent heat, former use of term, 74. Laudanum, 343. Laughing gas, 215. Lavoisier, demonstration of the com- position of air, 19 ; determination of the composition of air, 227; system of nomenclature applied to minerals, 356. Law, Avogadro's, 89. Law, Boyle's, 34 ; Boyle's, relation to Avogadro's law, 94; Charles, 38; Charles, relation to Avogadro's law, 94 ; Dulong and Petit, 396. Law, Faraday's, 438. Law for depression of freezing points of solutions, 112. Law, natural, 1 ; of combining vol- umes, 89 ; of combining weights, 13 ; of constant proportions, 12 ; of diffusion of gases, 58; of gravitation, 1; of "mass action," 149 ; of multiple proportions, 87 ; of partial pressures, 41, 77. Laws, graphical representation of gas, 42, 43 ; of gases, 34, 38. Lead acetate, preparation, uses, 519 ; basic, 519 ; carbonate, prepara- tion, 519 ; basic, manufacture, 520 ; comparison with lithopone, 470 ; chloride, formation, solubility, 518 ; chromate, 527 ; chromate, dis- covery of chromium in, 524. Lead dioxide, formation, use in storage batteries, 516 ; contrast with barium peroxide, 518 ; glazes, danger from, 502 ; in brass and bronze, 431. Lead monoxide, . manufacture, use, 515 ; nitrate, preparation, use, 519. Lead, occurrence, metallurgy, 513 ; properties, uses, 514 ; alloys, 515 ; oxides, 515 ; compounds, 518. Lead oxide, reduction by potassium cyanide, 321. Lead, treatment by Parke's process, alloys with zinc, 441. " Lead " pencils, 276. Lead peroxide, 515 ; plumbate, red lead, preparation, 515 ; structure, 516 ; plumbate, structure, 535 ; red oxide of, 515; sugar of, 519; sulfate in storage batteries, 516 ; sulfate, preparation, use as pig- ment, 519 ; sulfide, formation, solubility, 518; sulfide, theory of precipitation, 169 ; tetrachloride, preparation, hydrolysis, 518 ; tetra- sulfate in storage batteries, 518. Leather, chrome tanning of, 527. Leblanc, discovery of soda process, 450. Leblanc soda process, 411 ; dis- covery, 400 ; recovery of sulfur in, 456. Le Chatelier, explosion waves, 301 ; principle of van't Hoff-, 111. Legislation controling Leblanc soda process, 411, 456; manufacture of matches, 243 ; control of lead glazes, 502 ; to prevent poisoning by white lead, 521. Leguminous plants, fixation of nitro- gen by, 199. Length, unit of, 31. Levulose, see fructose. Lewis, equilibrium between silver oxide, silver and oxygen, 443 ; relation of law of Dulong and Petit to Avogadro's law, 397. Liebermann, discovery of structure of alizarin, 341. Life of an element, 474. Lignites, brown and black, composi- tion, 280. Ligroin, 289. Lime kilns, continuous and inter- mittent, 452; slaking of, 453. Lime, manufacture, 452 ; theory of formation from calcium carbonate, 453 ; in minerals, 357. Lime-nitrogen, 463. Lime-sulfur wash, 164. Lime, used to prepare absolute alcohol, 325. Linde, liquid air machine, 234. Liquid air, 232; preparation of oxygen from, 20. " Liquid smoke, " contains acetic acid, 329. Liter, defined, 31. Litharge, reduction with blowpipe, 304; separation of lead from silver by formation of, 439. Lithium, atomic weight of, 396 ; carbonate, 396 ; chloride, deter- mination of molecular weight of, 131 ; comparison with magnesium, 396; hydride, 395; nitride, 396; nitride, formation, 201. Lithium, occurrence, properties, 395 ; flame reaction, 396; perchlorate, use in determining the atomic weight of chlorine, 131 ; phos- phate, 396; urate, 396. Lithopone, 470; compared with white lead, 521. Liversidge, gold in sea water, 446. Lockyer, discovery of helium in sun, 237. Louisiana, sulfur in, 161. Luminosity of flames, cause of, 299. Luna, alchemical name for silver, 444. INDEX 587 Lunar caustic, 444. Lunge, theory of manufacture of sul- furic acid by chamber process, 178. Lutecium, rare earth metal, 506. Luteocobalt chloride, 559. Luteorhodium chloride, 563. H, definition, 262. iu/x, defined, 262. Mabery, announcement of Cowles furnace, 495. McCay, formation of arsenic penta- sulfide, 261. McCoy, atoms of metallic elements, 94. Maclnnes, table for degree of ioni- zation, 383, 384. Magnalium, 497. Magnesia usta, 479. Magnesite, 478. Magnesium ammonium arsenate, 259 ; ammonium chloride, an- hydrous magnesium chloride from, 480; ammonium phosphate, use, decomposition to magnesium pyro- phosphate, 481 ; carbonate in hard water, 463. Magnesium chloride, by-product in ammonia soda process, 413 ; ef- fect on the formation of gypsum, 460 ; hydrate, conduct on heating, 480 ; in salt, 405. Magnesium compounds, effect of ammonium hydroxide on solutions of, 491 ; diammonium phosphate, decomposition, 252 ; exercises, 493 ; hydroxide, formation, theory of solubility in solutions of ammo- nium salts, 479, 491 ; metaphos- phate, formation, 252. Magnesium nitride, formation, 201 ; occurrence, preparation, proper- ties, 478 ; uses, 479 ; oxide, prepa- ration, uses, 479. Magnesium pyrophosphate, forma- tion, 252 ; pyrophosphate, from magnesium ammonium phosphate, 481; silicide, 349; sulfide, hy- drolysis, 491 ; sulfide, preparation, hydrolysis, 480. Magnetic oxide of iron, formation from iron and steam, 49 ; from burning iron, 23 ; ore, structure, 556. Magnetite, 540. Malachite, 428; formation, 431. Malonic acid, carbon suboxide from, 316. Maltodextrin, from starch, 334. Maltose, alcohol from, 325 ; forma- tion from starch, hydrolysis, 334. Mammoth cave, saltpeter from, 418. Manchot, combination of ferrous sul- fate with nitric oxide, 217 ; ferrous chloride and nitric oxide, 554. Manganates, preparation, 536. Manganese dioxide, history of uses, 535 ; use in preparing oxygen, 21. Manganese heptoxide, properties, 538 ; list of oxides, 538 ; occur- rence, properties, 533 ; alloys, uses, compounds, 534 ; valence, structure of compounds, 534 ; tetrachloride, probable formation, 101. Manganic acid, change to perman- ganic acid, 536. Manganous chloride, 530. Manganous hydroxide, 535. Manganous manganic oxide, 534, structure, 535. Manganous sulfides, 535. Manometer, 41. Maple sugar, 333. Marsh's test for arsenic, 257. " Mass action," law of, 149. Mass and weight, relation, 32 ; de- pendent on velocity, 5. Matches, 242 ; from tetraphosphorus trisulfide, 243 ; law forbidding ordinary phosphorus in, 243. Matte, copper, 429 ; nickel, 559. Matter, conservation of, 6 ; defini- tion, 5 ; indestructibility of, 6. Mauve, discovery of, 340. Mechanical energy, units of, 33 ; equivalent of heat, 6. Medicine, relation of radiochemistry to, 475. Meerschaum, trisilicate, 356. Meker burner, temperature of, 303. Melting point, criterion of pure sub- stance, 12. Melting points of elements, 372, table, 373 ; in absolute tempera- ture, 135. Mendeleef, identification of scan- dium as ekaboron, 503 ; gallium as eka-aluminium, 506 ; periodic sys- tem, 136. Mental work in respiration calorim- eter, 316. Menzies, critical temperature of mer- cury, 485. Mercuric chloride, ionization of anomalous, 382; chloride, prep- aration, 489; properties, uses, antidote for, 490; cyanide, prep- aration, decomposition, 490 ; de- rivatives of ammonia, 492. Mercuric fulminate, use, hydrolysis to hydroxylamine, 491 ; iodide, for- mation, complex compound with 588 INDEX potassium iodide, use, 490 ; iodide, Nessler's reagent from, 492 ; ni- trate, 490. Mercuric oxide, decomposition, 9 ; formation and decomposition, 19 ; preparation, 488 ; yellow, 489. Mercuric sulfide, red and black, use, 489 ; solubility, 491. Mercurous chloride, formation in amalgamation process, 441 ; chlo- ride, preparation, vapor density, formula, uses, 489 ; nitrate, prep- aration, oxidation and reduction of, 490; basic, 490; oxide, 488; sulfate, use in Weston and Clark cells, 437. Mercury, abnormal ionization of compounds of, 491 ; exercises, 493 ; heat of combustion, 27 ; occur- rence, 484; metallurgy, purifica- tion, properties, critical tempera- ture, 485 ; uses, amalgams, 486 ; use in recovering gold, 446 ; use in Castner-Kellner process, 402. Metaboric acid, 366. Metachloroantimonates, 268. Metallic elements in periodic system, 136. Metallurgy, development of, 390 ; electrolytic methods in, 391 ; of aluminium, history, 391; roasting of sulfides in, 391 ; use of fuels in, 390. Metals, characteristics, 369. Metals, classification, 370, 371; preparation by thermite process, 495 ; spectra of, 427 ; systematic study of, 390. Metantimonic acid, 267. Metaphosphoric acid, formation properties, 253; hydrolysis, 249; properties, polymeric forms, 254. Metasilicates, 355. Metasilicic acid, 355. Metastannic acid, preparation, prop- erties, 512. Metastannyl chloride, 512. Metathesis, defined, 81. Meteorites, iron and nickel in, 540. Meter, definition, 31. Methane, kindling temperature, 288. Methane, limits for explosive mix- ture with air, 288 ; occurrence, preparation from sodium acetate, properties, 286; structure, 284; substitution products of, 287. Methyl alcohol, manufacture, 324, properties, use, 325; relation to structure of methyl ether, 324. Methyl amine, 339. Methyl ammonium iodide, 339. Methyl chloride, from methane, 287. Methylene chloride, from methane, 287. Methyl ether, determination of struc- ture, 324. Methyl iodide, relation to structure of methyl ether, 324. Methyl red, use in determining free carbonic acid, 465. Methyl silicate, colloidal silicic acid from, 353. Meyer, Lothar, preparation of hy- driodic acid, 145. Meyer, V., kindling temperature of methane, 288. Mho, defined, 380. Mica, 348 ; orthosilicate, 355. Michael, explosion waves, 301. Michelson, explosion waves, 301. Michigan, bromine from brines in, 140. Microcosmic salt, use, 423 ; formula, decomposition, use, 253. Micron, defined, 262. Migration of ions in electrolysis, 113. Milk sugar, source, use, 334. Millikan, number of molecules in Ice., 16, 96. Milliliter, used instead of cc., 31. Mineral, calculation of formula of, 356. Mirrors, tin amalgam for, 487. Mispickel, 256. Mixer for iron, 542. Mixtures and pure substances, 7. Moissan, artificial diamonds, 274 ; preparation of fluorine, 153. Moisture, determination of in air by weighing, and dew point, 232 ; effect on chemical reactions, 312 ; precipitation from air at an alti- tude, 232 ; presence in air, 231 Mol, definition, 183. Molar solutions, definition of, 183. Molecular weights, determined by measuring osmotic pressure, 360. Molecules of the elements, 93; number of in 1 cc., 95, 16. Molybdenite, 528. Molybdenum, occurrence, properties, 528 ; compounds, 529 ; trioxide, properties, 528; use in molybdic solution, 529. Molybdic anhydride, complex com- pounds from, 529. " Molybdic solution " preparation, use to determine phosphoric acid, 529, 530. Molybdic sulfate, reduction of molyb- dic acid to for phosphorus de- terminations, 529. INDEX 589 Monatomic gases, ratio of specific heats, 236. Monazite sand, cerium from, 363, thorium from, 364. Mo no calcium phosphate, 249 ; solu- bility, hydrolysis, 461. Monoclinic system (crystallography), 195. Monopotassium diarsenite, 259. Monosodium phosphate, 249 ; methyl orange as indicator for, 251. Montana, arsenic from smelting fur- nace in, 256. Moore, dissociation and ionization of ammonium hydroxide, 204. Mordants, 342 ; potassium pyrochro- mate, 527 ; titanium compounds as, 363. Morley, determination of the compo- sition of water by volume, 68 ; by weighing oxygen and hydrogen, Morphine, 343. Mortar, composition, hardening of, 454. Mother of vinegar, 329. Muffle furnace, 440. Multiple-effect evaporators, 405. Multiple proportion, law of, 87. Munroe-Neubauer crucibles, 565. Muscular energy in respiration calo- rimeter, 315 ; 346. Musgrave, development of Leblanc soda process, 400. Naphtha, 289. Naphthalene, from coal tar, use, 295. " Nascent," definition, 213. Natrium, 11. Natural gas, 286. Natural law, definition, 1. Nature of chemical energy, 27 ; of scientific knowledge, 1. Negatives in photography, 445. Neodymium, discovery, separation from praseodymium, compounds, 504. Neon, discovery, 238. Nernst, equilibrium for formation of nitric oxide, 216 ; quantum theory, 398 ; table for periodic system, 135 ; lamp, 363. Nessler's reagent, 490, 492. Neubauer-Munroe crucibles, 565. Neutrality, definition of, 385. Neutralization, definition, 121 ; in liquid ammonia solutions by union of hydrogen and amide ions, 208 ; theory of imperfect, 386, 384. New Caledonia, nickel from, 559. Newton's law of inertia, probable basis for principle of van't Hoff-Le Chatelier, 111. Nickel carbonyl, preparation, proper- ties, 561 ; chloride, 560 ; dimethyl- glyoxime, separation from cobalt, 560 ; occurrence, properties, 559 ; uses, alloys, compounds, 560 ; sul- fate, 560. Nicotine, 342. Niobium, see Columbium. Niton, discovery, 238; properties, 474 ; half-life of, 474. Nitrates, formation in soils, 199 : oxides prepared from, 392. Nitric acid, action on copper, 213; action on zinc, 213 ; and nitric oxide from nitrogen peroxide, 220 ; as oxidizing agent, 212 ; decom- position of, 212 ; detection with ferrous sulfate, 217 ; formation of nitrates from, 212. Nitric acid, hydrates of, 211 ; oxida- tion of feathers or wool by, 212 ; preparation from sodium nitrate, 210 ; properties, 212 ; structure according to electron theory, 207. Nitric oxide and carbon bisulfide, flame of, 217 ; and ferrous chloride, 554 ; and nitric acid from nitro- gen dioxide, 220 ; coefficient of ex- pansion, 38 ; combination with ferrous sulfate, 217 ; deviation from Boyle's law, 35 ; equilibrium of formation, 35; formation of nitrites from, 217 ; formed by re- duction of nitric acid with ferrous sulfate, 217 ; formed by union of nitrogen and oxygen, 215 ; prepa- ration by action of nitric acid on copper, 215 ; structure, 217 ; use in manufacture of sulfuric acid, 178. Nitrobenzene, use in preparing ani- line, 340. Nitrogen, coefficient of expansion, 38 ; combination with oxygen, 200 ; with hydrogen, 201 ; deviation from Boyle's law, 35 ; exercises, 225; iodide, preparation, proper- ties, 225; liquid, 234; list of oxides of, 214 ; pentoxide, prepa- ration, properties, 220 ; dioxide, formation from nitric oxide, 219 ; dioxide, from copper nitrate, 434 ; dioxide, from nitric oxide, 217; ratio of specific heats of, 237 ; tri- chloride, endothermic, 225 ; tri- chloride, equivalent to 3 Ch, 224 ; trichloride, formation by action of chlorine on ammonia, 209 ; occur- rence and natural history, 198; 590 INDEX trichloride, oxidation of arsenious oxide by, 224 ; trichloride, prepara- tion, properties, 224 ; peroxide, structure, 219 ; peroxide, use in manufacture of sulfuric acid, 178 ; preparation from air, 200 ; prepara- tion from ammonium nitrite, 200 ; from sodium nitrite, 200 ; proper- ties, 200 ; sources of for vegetable growth, 198 ; tetroxide, dissocia- tion, 219 ; tetrqxide, formation of nitrous and nitric acids from, 219 ; tetroxide, relation to nitrogen per- oxide, 219; tetroxide, structure, 219. Nitroglycerin, 326 ; explosion of, 327. Nitro Nitrogen Trichloride, probable existence, 225 ; hydrolysis of, 226. Nitrosyl chloride, properties and hydrolysis, 214. Nitrosylsulfuric acid, formation in manufacture of sulfuric acid, 178. Nitrous acid, formation, 218 ; anhy- dride, dissociation, 218 ; anhydride, preparation, 218 ; anhydride, use in manufacture of sulfuric acid, 178 ; oxide, deviation from Boyle's law, 35 ; oxide, preparation, properties, structure, uses, 214 ; oxide, use as anesthetic, 215. Nomenclature, Lavoisier's system applied to minerals, 356 ; of acids and salts, 123 ; of binary com- pounds, 29 ; of oxides, 29. Noncoking coals, 281. Nonmetallic elements in periodic system, 136. Non-metals, characteristics, 369. Non-sol glass, 467. Nordhausen sulfuric acid, 556. Normal solutions, definition of, 184. Norton, composition of silicic acids, 354. Norway, manufacture of nitrates in, 460. Noyes, A. A., cause of slight solu- bility of cobalt and nickel sulfides, 558. Noyes, W. A., determination of the composition of water, 71. Number of molecules in 1 cc., 16. Nutrition, diet, 345, 347; studied with respiration calorimeter, 345. Octahedron (crystal), 193. Octane, 283. Octene, 283. Octine, 283. Octovalent, definition, 64. Ohm defined, 33. Oil gas, 296 ; percentage composi- tion, 299. Oil of vitriol, 46, Oleic acid, source, 331. Olein, 332. Opal, composition, 354. Open hearth steel, 548 ; regenerative furnaces for, 549, 550. Opium, 343. Ordinates, axis of, 43. Ores, assay of for gold and silver, 440. Origin, mathematical definition, 43. Orpiment, 256 ; preparation, proper- ties, uses, 260. Orthoantimonic acid, 267. Orthoclase, trisilicate, 356. Orthophosphoric acid, classes of salts, 249 ; decomposition of salts of, 252 ; formation, 248 ; ionizatipn, 250 ; preparation, 249 ; properties, 249; solubility of salts of, 252; structure, 247. Orthosilicates, 355. Orthosilicic acid, 355. Oscillators, in quantum theory, 398. " Osmic acid," see osmium tetroxide, 565. Osmium, catalyzer for synthesis of ammonia, 201. Osmium-iridium, composition, prop- erties, 564 ; ruthenium in, 563. Osmium, occurrence, oxides, 564 ; chlorides, osmates, 565; tetroxide, use in histology, 565. Osmosis, 358. Osmotic pressure, connection with freezing points and boiling points of solutions, 360; defined, 360; measurement of, 359. -ous, suffix, meaning, 30 ; use for acids, 123. Oxalic acid and hydrazine from bis- diazoacetic acid, 222 ; carbon mo- noxide from, 311; decomposition, 466 ; ionization, relation to solu- bility of calcium oxalate, 465, 466 ; occurrence, manufacture from so- dium formate, 329 ; strength illus- trated, 386 ; use, decomposition, 330. Oxidation, definition, 63 ; reactions for potassium permanganate, 537 ; writing equations for reactions of, 171. Oxide of mercury, decomposition, 9. Oxides, nomenclature of, 29 ; of nitrogen, summary of methods of preparation, 214 ; preparation from metals, nitrates, carbonates and hydroxides, 392 ; by precipita- tion, 393 ; valence of elements in, 157. INDEX 591 Oxidizing flame, blowpipe and Bun- sen burner, 304. Oximes, 222. " Oxone," preparation of oxygen from, 21. Oxyacids of chlorine, structure, 130. Oxygen, 19 ; absorbed by molten silver, 443 ; and acid properties, 23 ; and chlorine, comparison of heats of combination, 108 ; basis of unit for atomic weights, 68 ; coefficient of expansion, 38 ; com- bination with hydrogen reversible, 372. Oxygen, determination of in air, 227 ; deviation from Boyle's law, 35 ; dissociation pressure of from barium peroxide and manufacture, 469 ; for medicinal use from liquid air, 235 ; liquid, 234 ; occurrence, 19 ; of air, may have come from carbon dioxide, 230 ; origin of name, 23. Oxygen, preparation from liquid air, 20 ; mercuric oxide, 19 ; " oxone," 21 ; potassium chlorate, 20 ; potas- sium chlorate with manganese dioxide, 21 ; sodium peroxide, 21. Oxygen, properties of, 22 ; weight of one liter in different latitudes, 22. Oxyhydrogen blowpipe, 61. Ozone, action on silver, 443 ; prepa- ration, 97 ; properties, 97 ; struc- ture, 98 ; from action of fluorine on water, 154. Palladium, semipermeable mem- brane for hydrogen, 358, 360; catalytic effect, use as catalyzer for reduction of fats, 564 ; di- chloride, use in gas analysis, 564 ; occurrence, properties, absorption of hydrogen, 563 ; oxides, chlo- rides, ammines, 564. Palmaer, absolute potential of ele- ments, 436. Palmitic acid, source, 331. Palmitin, 331. Paper, manufacture, 337; sizing, 338. Paraffin, 290. Parastannic acid, 512. Paregoric, 343. Paris green, 259. Parke's process for silver, 440. Parsons' apparatus for hydrogen sulfide, 165. Partial pressures, law of, 41, 77. Paste, glass, 467. Pattison's process for silver, 439. Peat, composition, 280. Pectin, relation to jelly, 337. Pectose, relation to jelly, 337. Pens, iridium for tips of, 565. Pentane, 283. Pentathionic acid, 188. Pentine, 283. Pepsin, 344. Per-, prefix, meaning, 30, 123. Perchlorates, 128. Perchloric acid, 128 ; structure, 130; structure of hydrated, 129. Perchloric anhydride, 129. Periodic law, exceptions to, 138. Periodic system, 132; tables, 134, 135. Perkin, discovery of mauve, 340. Perkin, fireproofing of cotton goods, 513. Permanent hardness, 311. Permanganates, preparation, 537. Permanganic acid, formed by use of sodium bismuthate, 269 ; from manganic acid, 537. Permanganic anhydride, 538. Permonosulfuric acid, 188. Peroxides, structure, 518. Perrin, estimate of number of mole- cules in 1 cc., 96. Persulfuric acid, preparation, uses, 187. Pertitanic acid, use in detecting titanium, 362. Petrolatum, 290. Petroleum ether, 289. Petroleum, occurrence, localities, varieties, refining, 289. Pewter, 509. Pharmaceutical extracts, use of al- cohol in, 325. Phase, effect of escape of a gaseous, 376; solid, effect of on reaction, 377. Phases, definition, 77; of chlorine hydrate, 107. Phase rule, 107 ; dissociation of calcium carbonate, 453 ; plaster of paris, 458 ; transition or quadruple point for sodium sulfate, 406. Phenacetine, 340. Phenol, source, manufacture, prop- erties, use as antiseptic, 326. Phenolphthalein, use as indicator in liquid ammonia, 208 ; use in determining free carbonic acid, 464. Phenylhydrazine, derivative of hydra- zine, 223. Phosgene, 316. Phosphate rock, 241. 592 INDEX Phosphine, compared with ammonia, arsine and stibine, 244; prepara- tion, properties, 243. Phosphomolybdic acids, 529. Phosphonium group, unstable, 244 ; iodide, interference in preparing hydriodic acid, 145 ; iodide, prep- aration, 243. Phosphor bronze, 431. Phosphoric acid, determination with molybdic solution, 529, 530; formed by burning phosphine, 243 ; water soluble, citrate-solu- ble, insoluble, in fertilizers, 461. Phosphorous acid, preparation, prop- erties, 248 ; structure, 247. Phosphorus, acids of, list, 247 ; acids of, basicity, 247 ; allotropic forms, 241 ; burning in oxygen, 23; chlorides of, hydrolysis, 245; exercises, 255. Phosphorus, heat of combustion, 27 ; in proteins, 343 ; occurrence, 240 ; oxides of, 246; oxychloride, formed by hydrolysis of the penta- chloride, 245; oxychloride, prep- aration from the trichloride, 246. Phosphorus pentachloride, action on hydroxyl compounds, 245 ; dis- sociation, 245; hydrolysis, 116, 245 ; preparation, properties, 244. Phosphorus pentasulfide, properties, uses, 254. Phosphorus pentoxide, efficiency as drying agent, 246 ; from burning phosphorus, 23 ; moisture left in gas by, 54 ; preparation, prop- erties, 246. Phosphorus, positive and negative valences, 248 ; preparation, 241 ; red, 241 ; sulfides of, 254 ; sulfide, use for matches, 243. Phosphorus tetroxide, 246 ; tet- roxide, hydrolysis, 254 ; trichlo- ride, hydrolysis, 1 15 ; trichloride, hydrolysis, 245 ; trichloride, prep- aration, properties, 244 ; triox- ide, preparation, properties, 246; valence in acids, 248 ; yellow, 241. Phosphotungstic acid, 531. Photographic plate, effect of Rontgen and Becquerel rays on, 471. Photography, 444, dry plates, posi- tives, negatives, developing, fixing, toning, 445. Photometry, stellar, use of selenium in, 190. Photosphere of sun, spectrum, 426. Physical sciences, 4. Physics, definition, 5. Pictet, liquefaction of air, 233. Pig iron, continuous casting ma- chines for, composition, gray, white, chilled, 543; analyses, 544. Pintsch gas, 296. Pitchblende, uranium in, 531. Plank, quantum theory, 398. Plaster of Paris in cement, 454 ; manufacture, use, 457 ; phase rule, 458. Plating, silver, 442. Platinic chloride, preparation, prop- erties, 566. Platinized asbestos, catalysis of union of O and H by, 62 ; prep- aration, 62. Platinous chloride, 565. Platinum ammines, 566 ; catalyzer for sulfur dioxide, 175 ; disulfide, 566; iridium electrodes in alkali manufacture, 401 ; metals, gen- eral properties, table, 562 ; prints in photography, 445; properties, uses, catalytic action, sponge, 565. Plticker tubes for spectra of gases, 427. Plumbic acid, 516. Plumbum, 11. Poisoning by white lead, 521 ; by lead water pipes, 514. Polarimeter, used to determine sugar, 333 Polarization of light by crystals, 196. Polarized light, effect of sugar on, 333. Poly sulfides, ammonium, 422. Polythionic acids, 188. Porcelain, 501; glazing, 502. Positives in photography, 445. Potassium aluminium sulfate, 500 ; argenticyanide, 320 ; use in silver plating, 321 ; aurate, 450 ; aurous cyanide, formation in cyanide process, 446 ; bicarbonate, prep- aration, use, 420. Potassium carbonate from wood ashes, 414 ; from beet sugar manu- facture, from wool, 419 ; prop- erties, 419 ; uses, 420. Potassium chlorate, composition, 21 ; manufacture, uses, 416 ; prepa- ration, 127 ; preparation of oxygen from, 20. Potassium chloroaurate, 450. Potassium chloride, measurement of degree of ionization of, 380 ; conductance of solutions of, 380 ; properties, use in fertilizers, 416 ; for manufacture of saltpeter, 416, 418 ; chloroplatinate, use in de- termining atomic weight of chlo- INDEX 593 rine, 130; chloroplatinate, 566; chloroplatinite, use in photogra- phy, 565 ; chloroplumbate, 519 ; chromate, preparation from chrome iron ore, 527 ; chromium sulfate, 527 ; cobaltinitrite, formation, prop- erties, use in analysis, 558, 419 ; colbalticyanide, 558 ; cobaltocyan- ide, 558 ; cupric chloride, use in iron analysis, 432 ; cuprocyanide, 435 ; cyanate, 321 ; cyanide, for- mation, preparation, use, 319 ; cyan- ide, preparation, use, 420. Potassium diuranate, 532; dichro- mate, mordant, 342 ; dichromate, reduction by hydrogen sulfide, 171.; ferrate, 553; ferricyanide, formation from potassium cya- nide, 320 ; ferricyanide, use in blue-print paper, 331 ; ferric ferro- cyanide, 321 ; ferrocyanide, forma- tion from potassium cyanide, 320 ; Prussian blue from, 320 ; ferro- cyanide, preparation, 319 ; flu- oride, acid, 155 ; fluotantalate, use in purifying tantalum, 524. Potassiumhydroxide.'preparationfrom the carbonate, 415 ; by electroly- sis, properties, 415 ; slight effect on glass, use in analysis, 416. Potassium iodate, reduction in pre- paring the iodide, 417 ; iodide, manufacture, uses, 417. Potassium manganate, preparation, conversion to permanganate, 536 ; mercuric iodide, 490 ; metachlo- roantimonate, 268 ; metallic, dis- covery, preparation, properties, 415. Potassium nitrate, formation in soil, 199 ; nitrate, relation to gunpowder, 417 ; sources, manu- facture, properties, uses, 418 ; nitrite, preparation, properties, use, 419. Potassium, occurrence, relation to minerals, clays, soils, plant growth, 414 ; osmate, 565 ; oxide, 415 ; perchlorate, preparation, 128 ; per- chlorate, preparation, properties, 416. Potassium permanganate, properties, uses, 537 ; typical oxidations with 537 ; reaction of hydrochloric acid with, 159 ; purification of hydrogen by, 55. Potassium perruthenate, 563 ; per- sulfate, preparation, 187 ; polyi- odides, formation, use in iodimetry, 417 ; pyrochromate, preparation, properties, uses, relation to pyro- sulfate, 527 ; use in chrome tanning, 527 ; pyrosulfate, decomposition, use in analysis, 417. Potassium, retention in soils by colloi- dal silicic acids, 354 ; ruthenate, 563; silver cobaltinitrite, 559, 419 ; silver cyanide, use in silver plating, 442. Potassium, source for shales, clays and soils, 494 ; sulfate, 417 ; sulfate, acid, preparation, pyro- sulfate from, 417 ; sulfocarbonate, preparation, 317, use, 318 ; tar- trate, acid, 330 ; tetroxalate, use as standard in alkalimetry, 330. Potassium thiocyanate, carbon oxy- sulfide from, 318 ; thiocyanate, preparation, use in testing for iron, 321 ; triiodide, 145 ; zincate, 483. Potential, absolute, of elements, 436 ; differences of in relation to corro- sion of iron, 550. Power, unit of, 33. Praseodymium, discovery, separa- tion from neodymium, compounds, 504. Precipitation, theory of, 376. Premier and Schupp, molecular weight of sulfur vapor, 163. Preparation of compounds, general methods, 372-379; of pure sub- stances, 8. Pressure, effect of on a gas, 34. Priestly, analysis of air by nitric oxide, 230. Primary salts, 249. Printing in photography, 445. Producer gas, 297 ; heat relations in manufacture, 298; percentage composition, 299 ; use in regenera- tive furnace for steel, 550. Propane, structure, 284. Propene, 283; structure, 284. Propine, 283. Propylene, structure, 284. Proteins, occurrence, 343 ; diges- tion, 344. Prussian blue, 320; use in chrome green, 525. Prussic acid, 319. Ptomaines, 343. Ptyalin, 344. Pure substances and mixtures, 7 ; composition of expressed in mul- tiples of atomic weights, 17 ; distinguished from mixtures, 12; preparation of, 8. Purpureocobalt chloride, 559. Pyridine solutions with semiper- meable membrane, 358. Pyrite, 161. 594 INDEX Pyrite burners, 179; use of oxide of iron from, 540. Pyroantimonic acid, 267. Pyroarsenic acid, 259. Pyroboric acid, 367. Pyrolusite, 533. Pyrophosphoric acid, hydrolysis, 248 ; preparation, properties, 253. Pyrosulfates, preparation, 186. Pyrosulfuric acid, 186. Quadrivalent, definition, 64. Quadruple point for sodium sulfate, 406. Qualitative analysis, basis for groups of, 166 ; definition, 66 ; groups of, 166 ; of water, 66. Qualitative synthesis of water, 66. Quantitative analysis, definition, 66. Quantitative synthesis of water by volume, 66. Quantum theory, 398. Quartz, 348 ; properties, fused, uses, 352. Quinine, 343. Quinquivalent, definition, 64. Radiations, penetrating, 471 ; kinds of, 472, 473. Radical, definition, 47 ; definition, relation to structure, 323. Radioactive elements, series of, 475. Radiochemistry, relation to geology and medicine, 475. Radiothorium, 364. Radium, an element, 9, 473 ; chemical action, 475 ; discovery, 471 ; dis- integration, 473 ; emanation, 474 ; evolution of heat by, 472 ; half- life, 474 ; nature of rays, 473 ; properties, 472. Rails, steel, manufacture, 548. Ramsay, discovery of argon, 235 ; discovery of helium, 237 ; discovery of helium from radium, 473 ; disso- ciation of nitrous anhydride, 218 ; possible disintegration of atoms by radium emanation, properties of niton, 474, 475 ; use of periodic sys- tem in discovery of noble gases, 136. Rare earth elements, position in periodic system, 134, 138. Rare earths, general, 502 ; groups of, 503 ; methods of separation, 503, 504. Raschig, theory of sulfuric acid manufacture, 179. Rayleigh, discovery of argon, 235. Rays, a, 0, 473 ; 7, 5, 474. Rays, chemical action of radioactive, 475. Reacting substances, effect of remov- ing one on equilibrium, 152. Reactions, bimolecular and unimolec- ular, 150 ; calculation of relative speed of at equilibrium, 151 ; effect of insolubility on, 376 ; effect of volatility on, 374 ; re- versible, 50 ; speed of chemical, 151. Realgar, 256; preparation, proper- ties, uses, 260. Reasoning, inductive, 13. Reciprocal ohms, definition, 380." Red lead, oxide, manufacture, 515; structure, 516. Reducing agent, sodium amalgam, zinc amalgam, 487 ; flame, blowpipe and Bunsen burner, 304. Reduction, definition, 63 ; writing equations for reactions of, 171. Refining, electrolytic, of copper, 429. Refrigeration by ammonia in ma- chines, 204. Regenerative furnace for open hearth steel, 549, 550. Regular system (crystallography) , 193. Reid, discovery of indium, 506. Resistanz glass, 467. Resonators in quantum theory, 398. Respiration calorimeter, 313 ; study of nutrition with, 345. Reversed flames, 304. Reversible reactions, 50; hydro- chloric acid and oxygen, 109 ; hydrogen and iodine, 146 ; ioniza- tion, 115; salt and sulfuric acid, 119; theoretically all reactions, 372. Rhodium, properties, oxides, chlo- rides, complex salts, 563. Rhombic dodecahedron (crystal), 193; hexahedron (crystal), 195; system (crystallography), 194. Richter, discovery of indium, 506. Roasting, defined, 263; sulfides in metallurgy, 391. Rochelle salt, use in Fehling's solu- tion, 335. Rock crystal, use for lenses, 352. Rock salt, mining and obtaining, 404. Rontgen, discovery of Rontgen rays, 471. Rosa, respiration calorimeter, 313. Roscoe and Schorlemmer, Treatise, standard for ventilation, 231. Rose, distinguished columbium and tantalum, 523. Roseocobalt chloride, 559. Roseorhodium chloride, 563. Rouge, 555. INDEX 595 Rubidium alum, 424 ; chloroplatinate 424, 393 ; discovery, occurrence in carnallite, properties, 424. Ruby, 494 ; artificial, 500. Rum, 325. Ruthenium, occurrence, oxides, chlo~ rides, double chlorides, ruthenates, perruthenates, 563. Rutherford, number of molecules in 1 cc., 16, 96 ; disintegration theory, 472. Saccharimeters, 333. Sackur, quantum theory, 398. Safety lamp, Davy. Safety matches, 242. Sainte-Claire-Deville, preparation of aluminium, 495. Saleratus, 420. Sal soda, 411. Salt and sulfuric acid, reversible re- action, 374 ; definition, 47, 121 ; mining and obtaining, 404. Saltpeter, Chili, 210; formation in soil, 199. Salts, nomenclature of, 124 ; hy- drolysis, general statement, 394 ; list of insoluble, 393 ; solubility of, general statement, 393 ; valence of elements in, 158. Samarium, discovery, compounds, 505. Sapphire, 494, 500. Scandium, discovery, same as eka- boron, compounds, 503 ; same as ekaboron, 136. Scheele, discovery of chlorine, 535. Science, subdivisions of, 3. Sciences, abstract, physical, bio- logical, psychological, 4. Scientific knowledge, Nature of, 1. Scott, determination of the composi- tion of water by volume, 68. Sea water, gold in, 446. Secondary salts, 249. Selection of atomic weights, 16, 92. Selenic acid, preparation, properties, 190. Selenite, 457. Selenium dioxide, decomposition by sulfur dioxide, 190 ; preparation, 190. Selenium, occurrence, allotropic forms, uses, 189. Self-hardening steels, 552. Semibituminous coals, composition, 280. Semipermeable membranes, 357 ; preparation, 358 ; mechanism of action, 358, 360 ; for pyridine, water, hydrogen, 358. Septivalent, definition, 64. Series of hydrocarbons, 283 ; elec- tromotive, 435, table, 436; of radioactive elements, 475. Serpentine, disilicate, 356. Sexivalent, definition, 64. Shales, formation of, 494. Sherardized iron, 482. Sicily, sulfur in, 160. Siderite, 540. Siemens-Martin process for steel, 548 ; regenerative furnaces for, 549, 550. Silicates, artificial, 352 ; list of com- mon, 348, 355; natural, 355. Silicic acid from silicon fluoride, 350 ; from silicon chloride, 351 ; insol- uble after drying, 354. Silicic acids, composition, 354 ; im- portance of colloidal in soils, 354 ; preparation, colloidal, dialysis of, 353 ; structure of, 354 ; various, 355. Silicon carbide, 349. Silicon dioxide, forms, 348, 351 ; formula, 354 ; properties, 352. Silicon, exercises, 368; fluoride, preparation, hydrolysis, 350 ; hexa- iodide, preparation, silicooxalic acid from, 351 ; occurrence, 348 ; prepa- ration, properties, 349 ; tetrachlo- ride, preparation, properties, hydrolysis, 351 ; tetrafluoride, formed in etching glass, 151 ; tetraiodide, silicon hexaiodide from, 351. Silicooxalic acid, preparation, proper- ties, structure, 351. Silicozirconates, 363. Silver, annual production and value, properties, alloys, plating, 442 ; argenticyanide, evidence of exist- ence of complex ions in, transfer- rence of ions of in electrolysis, 379 ; bromide, properties, conduct toward light, 444 ; use in photog- raphy, 445. Silver chloride, properties, conduct toward light, 444 ; use in pho- tography, 445 ; reduction by mer- cury, 441;* solubility product of, 373. Silver chloroplatinate, 566; dep- osition from potassium argenti- cyanide, 321; exercises, 450; for mirrors, 487. " Silver from Clay," applied to aluminium, 495. Silver hydroxide, 442 ; ionization of, 443 ; hydroxide, temporary forma- tion, 393 ; iodide, properties, con- 596 INDEX duct toward light, 444 ; iodide, use in photography, 445 ; molten, ab- sorption of oxygen by, 443. Silver nitrate, preparation, proper- ties, uses, 444 ; theory of reaction with salt, 376 ; use in determining the atomic weight of chlorine, 131. Silver nitrite, preparation, use, 444 ; oxide, dissociation pressure, 443 ; formation, 442 ; orthoarsenite, 259 ; peroxide, formation, 443 ; pyro- phosphate, 253. Silver, occurrence, metallurgy, Patti- son's process, 439 ; cupellation, Parke's process, 440 ; amalgama- tion process, electrolytic process, cyanide process, 441. Silver recovered from copper, 430, 439 ; separation from lead, 439, 440 ; sulfide, conversion to chloride in ores, 441 ; sulfide on coin, test for sulfur, 409 ; sulfate, prepara- tion, use, 444 ; trinitride, failure to react with iodine trinitride, 224. Sizes for paper, 338. Slag, blast furnace, use for cement, 454, 543. Slimes, from electrolytic refining of copper, 430. Smalt, 558. Smaltite, 557. Smith, Alex., formula of calomel, 489. Smith, E. F., potassium fluotantalate, 524 ; preparation of metallic cal- cium, 452. Smithells, separated Bunsen flame, 301. Smithsonite, 481. Smokeless powder, 419 ; gun cotton in, 338. Soaking pits, in steel manufacture, 548 Soap, manufacture, use, 332 ; water glass in, 353 ; soft, from lye of wood ashes, 414. Soapstone, metasilicate, 355. Soda, baking, 412; washing, 411. Soda lime, use in preparing methane, 286 ; to absorb carbon dioxide and water, 6. Sodamide, base in liquid ammonia, 208 ; hydronitric acid from with nitrous oxide, 223 ; preparation, properties, ionization, use in mak- ing indigo, 410. Soda water, 309. Soddy, discovery of helium from radium, 473. Sodium aluminate, hydrolysis, 499 ; decomposition by carbonic acid, 496 ; aluminate, manufacture from clay, 496; amalgams, composi- tion, formula for, 487 ; ammonium phosphate (microcosmic salt) , 253 ; antimonite, 265 ; atoms, rate of vibration, 426. Sodium bicarbonate, formation from carbonic acid, 360 ; loss of carbon dioxide on boiling solution, 376 ; Solvay or ammonia soda process, 412. Sodium bismuthate, oxidation of manganese by, 269 ; bisulfate, 408 ; bisulfite, 408; bromate, 144; bromate, use with potassium iodide to illustrate strength of acids, 386. Sodium carbonate, by-product in making aluminium oxide from clay, 496 ; decomposition by acids, theory, 375 ; formation from car- bonic acid, 310 ; hydrolysis of, 385 ; Leblanc soda process, 411 ; mono- hydrate, 411 ; transition point, dekahydrate, 412 ; uses, 412. Sodium chloride, crystallization, con- centration of solution of, 405 ; elec- trolysis of, 401 ; localities for, 398 ; properties, essential in diet, 406; sodium bicarbonate from, 412 ; sources, 404 ; theory of reaction with silver nitrate, 376. Sodium cobaltinitrite, 558 ; copper orthophosphate, 253 ; formate, oxalic acid from, 329 ; hydrosul- fide, formation, properties, 409. Sodium hydroxide, Castner-Kellner process for, 402 ; from sodium carbonate, 401 ; from sodium chloride, 401 ; properties, density of solutions, table, 403. Sodium hypobromite, 143 ; hypophos- phite, formed in preparing phos- phine, 243 ; hyposulfite, prepara- tion, uses, 408 ; hyposulfite, prep- aration, 186 ; hyposulfite, use for reduction of indigo, 341. Sodium " hyposulfite," old name for thiosulfate, 408. Sodium hypophosphate, acid, 254 ; iodide, separation of ions by cen- trifugal force, 114; manganate, preparation conversion to perman- ganate, 536 ; manufacture for aluminium, 495 ; manufacture of sodium peroxide from, 404 ; meta- phosphate, formation, use, 253 ; metaphosphate, from microcosmic salt, use, 423. Sodium nitrate, acid sodium sulfate from decomposition of, 408 ; for- mation in soil, 199 ; reduction to nitrite, 218; source, properties, use, 410. INDEX 597 Sodium nitrite, preparation, 218; preparation, use, 410. Sodium, occurrence, 398 ; metallurgy, properties, 399 ; uses, 400 ; oxide, preparation, properties, 404 ; per- borate, bleaching by, 367 ; perchlo- rate, use in preparing perchloric acid, 128. Sodium peroxide, hydrogen peroxide from, 84; preparation of oxygen from, 21 ; preparation, properties, hydrate, uses, 404. Sodium phosphites, 248 ; pyro- borate, 367 ; pyrosulfate, prepara- tion, 186 ; pyrosulfate, prepara- tion, use in analysis, 408 ; ses- quicarbonate, occurrence, 398 ; silicates, manufacture, use, 413 ; silicate, preparation, properties, uses, 353 ; sulfarsenate, 261. Sodium sulfate, acid, preparation, uses, pyrosulfate from, 408 ; manu- facture, hydrate, transition point for hydrate, 406 ; solubility curve, 407 ; constituent of Hunyadi water, 408. Sodium sulfide, preparation from sodium hydroxide, from sodium carbonate and sulfur, as test for sulfur, 409 ; sulfite, acid, use in cider, 175 ; sulfite, acid, prepara- tion, uses, 408. Sodium stannate, use in fireproofing cotton goods, 513 ; stannite, prep- aration, reducing agent, 510 ; tetrathionate, from sodium thio- sulfate, 187 ; tetrathionate, for- mation in iodine titrations, 409. Sodium thiosulfate, conversion to tetrathionate, 187 ; preparation, use in photography, 408 ; in lixiviation processes, 409 ; anhy- drous, 409 ; precipitant for copper, 409; structure, 409; use in photography, 445. Sodium trinitride, formation, 410 ; trinitride, preparation, 223 ; tung- state, 531 ; zincate, 483. Softening water, by boiling, 463 ; by use of sodium salts, 464 ; Clark's Erocess for, 464 ; " permanently " ard waters with sodium carbonate, phosphate, fluoride or borate, 311. Soils, formation of, 494 ; formation of nitrates in, 199 ; importance of colloidal silicic acids in, 353 ; po- tassium in, 414. Solder, 509. Soluble glass, 353; manufacture, uses, 413. Solubility of salts, general statement, 393 ; graphical representation, 80. Solubility product, 377 ; rule for, 378 ; relation to solubility of magnesium hydroxide, 479 ; relations for uni-bivalent and bi-bivalent salts, 378. Solute, defined, 79. Solution, defined, 79; formular, def- inition of, 183 ; molar, definition of, 183 ; normal, definition of, 184 ; standard, definition of, 183. Solution pressure, defined, 435. Solutions, chemical activity in, 81 ; differences in potential between, 437 ; in liquid ammonia, 207 ; supersaturated, 80. Solvay, discovery of ammonia soda process, 400. Solvent, defined, 79. Sommerfield, quantum theory, 398. Specific heat, fixes atomic weight of indium, 506; of elements, 397; of gases at constant volume and constant pressure, 236. Spectra, comparison of, 427 ; of gases, Pliicker tubes, 427 ; of metals, how obtained, 427. Spectroscope, direct vision, 427 ; de- scription of, 425. Spectrum analysis, 424. Spectrum, continuous, bright and dark lines, 425 ; theory of dark line, diffraction, solar, 426. Speed of chemical reactions, 149 ; relation to chemical affinity, 149 ; and concentration, 149. Speed relation of two reactions calculated at equilibrium, 151. Sphalerite, 161, 481. Stafford, molecular weight of sulfur vapor, 163. Standard solutions, definition of, 184. Stannate, us in fireproofing cotton goods, 513. Stannic acids, table of, 511. Stannic acid, preparation, properties, 511. Stannic chloride, preparation, con- duct in solution, 512 ; oxide, prep- aration, 510; use in glass, 511; sulfide, 512. _ Stannous chloride, preparation, uses, reducing agent, 510 ; oxide, 510 ; sulfide, 510 ; sulfide dissolved by ammonium polysulfide, 422. Stannum, 11. Starch, glucose from, 334 ; source, manufacture, varieties, 335 ; ap- pearance of granules, cooking of, 336. Stassfurt, potassium chloride from, 415. 598 INDEX Stearic acid, source, 331. Stearin, 331. Steel, cementation, analyses, 551 ; alloy, 552 ; cast, tempering, 545 ; theory of tempering, 546 ; Besse- mer, 547 ; acid and basic Bessemer, open hearth, Siemens-Martin, 548 ; resistance to corrosion increased by copper, 431. Stellar photometry, use of selenium in, 190. Stereotype metal, 264, 269, 515. Stewart, standard of ventilation, 231. Stibine, 264. Stibium, 11. Stibnite, 263. Stockholm, holmium named for, 505. Stokes, ferric sulfide, 556. Storage batteries, theory of, 516. Storage of gases, 22. Strass, 467. Strength of acids, definition, 168; illustration, 386. Strength of organic bases, 339. Strong acids and bases defined, 386. Strontianite, 468. Strontium carbonate, 468 ; com- pounds, 468 ; flame color, 471 ; hydroxide, 468 ; nitrate, use in fireworks, 468 ; occurrence, 467 ; sulfate, 468. Structural formulas, basis for, 323. Structure of compounds, relation to valence, 65 ; of the oxyacids of chlorine, 130. Strychnine, 343 ; salt of chloroauric acid, 450. Strychnos mix vomica, strychnine from, 343. Study of chemistry, 18. Subdivisions of science, 3. Substance, definition, 7. Substantive dyes, 342. Substitution, 287 ; use in determining structure, 324. Sudbury, nickel from, 559. Sugar, cane, occurrence, manufac- ture, properties, effect on polarized light, 333; hydrolysis, 334; puri- fication by bone black, 278 ; solu- tions clarified by basic lead acetate, 519 ; use of strontium hydroxide in manufacture of, 468. Sugar of lead, 519. Sulfantimonates, 268. Sulfantimonites, 268. Sulfarsenates, formation, 261. Sulfarsenites, formation, 261. Sulfates, list of insoluble, 183. Sulfides, basis of groups of qualita- tive analysis, 166; hydrolysis of, 171 ; of Group II, solubility, 491 ; roasting in metallurgy, 391 ; theory of precipitation by alkaline sul- fides, 170. Sulfites, as germicides, 174; prep- aration and uses, 175. Sulfocarbonates, formed from carbon bisulfide, 317. Sulfocarbonic acid, formation, de- composition, 318. Sulfocyanates, see Thiocyanates. Sulfur, allotropic forms of, 162 ; amorphous, 162 ; boiling point, 163 ; burning in oxygen, 23 ; by Chance process, 160 ; com- pounds containing halogens, 188. Sulfur dioxide, as germicide and disinfectant, 173 ; bleaching by, 173 ; catalysis of conversion to sulfuric acid by oxides of nitrogen, 178 ; coefficient of expansion, 38 ; deviation from Boyle's law, 35 ; from burning sulfur, 23 ; from iron pyrites, 177 ; preparation by burn- ing sulfur, 172 ; preparation by reduction of sulfuric acid, 173 ; preparation from acid sodium sulfite, 173 ; properties, 173 ; solu- tion in water, 174 ; uses, 173. Sulfur, exercises, 196 ; family, table of compounds, 192 ; flowers of, 160; gaseous, Ss, 82 and S, 163; group, 160 ; heat of combustion, 27. Sulfur hexafluoride, 188. Sulfur, in Louisiana, getting of, 161 ; in petroleum removed with copper oxide, 289 ; in proteins, 343 ; in Sicily, getting of, 160. Sulfur, lime-, wash, 164 ; mobile liquid (Sx), 162 ; monochloride, preparation, uses, 188 ; monoclinic, 162 ; native, source of, 160 ; occurrence, 160. Sulfur, production in U. S. and in world, 161 ; properties, 163 ; rhom- bic, 162 ; roll brimstone, 161 ; test for with sodium carbonate on charcoal, 409. Sulfur trioxide, absorption by con- centrated sulfuric acid, 176; for- mation from sulfur and oxygen reversible reaction, 175 ; poly- meric, 176 ; preparation, 175 ; properties, 176 ; sulfuric acid from, 45. Sulfur, uses, 164 ; viscous liquid (Su), 163. Sulfuric acid as dehydrating agent, 182^ "chamber acid," 180; chamber process for, 177 ; con- centration of, 180 ; directions for INDEX 599 dilution of, 182 ; dissociation of, 180 ; electrolysis of, 9, 47 ; " fum- ing," 176; fuming from ferric sul- fate, 556 ; Gay-Lussac tower, 179 ; hydrates of, 181 ; ionization of, 181 ; moisture left in gas by, 54 ; prep- aration by " chamber process," 177 ; properties, 180 ; reaction with copper, 173 ; reduction by hydroidic acid, 145 ; reduction by hydro- bromic acid, 142. Sulfuring " fruit, 164, 174. ulfurous acid as reducing agent, 174 ; formation, 174 ; ionization of, 174; "strength" of, 174; struc- ture, 174. ulfuryl chloride, hydrolysis, 188 ; preparation, properties, 188. un's corona, helium in, 237; photo- sphere, spectrum, 426. uperphosphate, calcium, manufac- ture, use in fertilizers, 461. upersaturated solutions, 80. plvite, 414. Symbols of elements, 11. ympathetic ink, 557. Synthesis, definition, 66. ynthesis, quantitative, of water by volume, 66. /stems of crystallography, 193. able of atomic weights, 10 ; abso- lute potentials of elements, elec- tromotive series, 436 ; acidity or alkalinity of indicators, 388 ; compounds of elements of Group V, 271 ; compounds of sulfur family, 192 ; correction of apparent weight of water to volume, 73 ; correction of readings of barometer for glass and brass scale, 36 ; correction of readings of barometer for latitude and altitude, 37 ; degree of ionization, acids, 383 ; bases, 383 ; salts, 384 ; density and volume of water, 73 ; density of gases, 95 ; density of solutions of sodium hy- droxide, 403 ; deviation of gases from Boyle's law, 35 ; deviation of gases from law of Charles, 38 ; dis- sociation of water, 61 ; elements in earth's crust, 11 ; elements of sul- fur group, 160 ; equilibrium be- tween nitrogen and oxygen, 216; equilibrium of hydrogen, iodine and hydriodic acid, 147; groups and fam- ilies of elements, 371 ; heat of com- bustion of C, S, P, Fe and Hg, 27 ; halogen acids, 139 ; halogen family, 139 ; nonmetallic elements, 348 ; Periodic System, 134, 135 ; proper- ties and compounds of elements of Group VIII, 562 ; rate of decom- position of hydriodic acid, 148; rate of formation and decomposi- tion of hydriodic acid, 150 ; stannic acids, 511 ; values of calorie, 10- 30 , 33 ; vapor pressure of ice and water, 75 ; vapor pressure of sys- tems containing calcium sulfate, 459 ; varieties of ferrite, 546. Tafel, preparation of hydroxylamine, 221. Talc, metasilicate, 355. Tannic acid, mordant, 342. Tantalum, discovery, properties, 523, use, compounds, 524, electric light, 524. Tartar emetic, 266. Tartaric acid, structure, source, salts, 330. Tasmania, tin from, 508. Teeth, amalgam for, 487. Telluric acid, 190. Tellurium, anomalous position in periodic table, 138 ; atomic weight of, 190 ; dioxide, 190 ; occurrence, preparation, compounds, 190. Tellurous acid, 190. Temperature, absolute, 39, critical, 233 ; effect of on a gas, 38 ; inter- national scale of, 32 ; kindling, 24 ; of flames, 302 ; of interior of earth, relation of radioactivity to, 476; units of, 32. Temperatures, thermometers for high, 486. Temporary hardness, 310. Terbium compounds, 505. Terne plate, 509. Tertiary salts, 249. Tetragonal pyramid (crystal), 194. Tetragonal system (crystallography), 194. Tetrahedron (crystal), 193. Tetrahexahedron (crystal), 193. Tetramethyl ammonium hydroxide, analogy with ammonium hydrox- ide, 206. Tetraphosphorus heptasulfide, 254. Tetraphosphorus trisulfide, proper- ties, 254 ; use for matches, 243. Tetrathionic acid, 188. Thallium, discovery, compounds, 507. Theorem of Le Chatelier, 111. Theory, atomic, 14 ; electrochemical, influence on formulas of minerals, 356; the electron, 181; electron, in relation to metals and non- metals, 370; electron, relation to ionization, 206 ; colloidal solu- tions, 262 ; decomposition of car- 600 INDEX bonates by acids, 375; chamber process for manufacture of sulfuric acid, 178 ; hydrolysis of salts, 385 ; neutralization, 384 ; the quantum, 398 ; storage batteries, 576 ; Thermite process, Goldschmidt's, 497 ; chromium by, 524 ; tungsten by, 530. Thermodynamic scale, melting points on, 373. Thermometers, glass for, 467 ; of fused quartz, 352 ; mercury and international scales for, 486 ; zero point correction, special for high temperatures, 486 ; point on, fixed by transition point of sodium sulfate, 406. Thio-, prefix, 187. Thiocyanate, potassium, 321. Thiosulfates, preparation, use, 186. Thiosulfuric acid, formation and de- composition, 187 ; formation from sulfur monochloride, 188. Thomson, heat of combustion of hydrogen, 65. Thompson, J. J., the electron theory, 181 ; atoms of metallic elements, 94. Thompson, Sir William, size of mole- cules, 16. Thorianite, 364. Thorium, in monazite sand, proper- ties^ dioxide, sulfate, nitrate, use in Welsbach mantles, 364; series of elements, 475. Thulium, compounds, 506. Thyroid gland, iodine in, 144. Tilkerode, thallium from, 507. Time, units of, 32. Tincture, defined, 144. Tin, occurrence, sources, metallurgy, 508 ; properties, alloys, uses, tin plate, 509 ; compounds, 510 ; re- covery from tin scrap, 508. Titanium, compounds of as mordants, 363 ; occurrence, properties, com- pounds, 362 ; oxide, solution in sodium pyrosulfate, 408; separa- tion from silica, detection, 362 ; tetrafluoride, 362 ; test for hydro- gen peroxide, 86. Tolman, separation of ions by cen- trifugal force, 114. Toluene, 283. Toning in photography, 445. Tool steel, high-speed, 530. Topaz, 349. Torpedoes, gun cotton in, 338. Tourmaline, 349. Toxins, 344. Transition or quadruple point, for sodium sulfate, 406. Transition points, for steel and ferrite, 546. Trautz, theory of sulfuric acid manu- facture, 179. Triammonium dodekamolybdate, 529. Tribasic acids, defined, 183. Tricalcium phosphate, 249. Triclinic system (crystallography), 196. Tridymite, 352. Trimethyl amine, 204. Triphosphorus hexasulnde, 254. Triple-effect evaporation, 405. Triple point, definition, 78. Triple point water-ice-water-vapor above , 407. Trisilver phosphate, 253. Trisodium phosphate, 249 ; aliza- rine green as indicator for, 251. Trisilicates, 356. Trisilicic acids, 355. Trisodium phosphate, hydrolysis, al- kaline reaction of, 252. Trithionic acid, 188. Trivalent, definition, 64. Trypsin, 344. Tuberculosis, relation to ventilation, 231. Tungsten, history, preparation, prop- erties, use in lamps, in tool-steel, 530. Tungsten bronze colors, 531. Turmeric paper, test for boric acid, 368. Turnbull's blue, 321. Tuyeres of blast furnace, 541. Type metal, 264, 175. Typhoid fever from impure water supply, 83. Ultramarine, natural and artificial, 502. Ultra-violet light, use in purifying water, 83. Uni-bivalent salts, law of solubility product not general for, 378. Unimolecular reactions, 150. Unit for atomic weights, 68 ; electri- cal charge, 438; of length, 31; of power, 33 ; of temperature, 32 ; of volume, 31 ; of weight, 31. Units, absolute, 33; electrical, 33; of energy, 32; of heat, 33; of mechanical energy, 33 ; of time, 32. Uni-univalent salts, law of solubil- ity product for, 378. Univalent, definition, 64. Univariant, definition, 78. Unsaturated compounds, definition, 291. Uraninite, 531 ; helium in, 237. Uranium chlorides, 532 ; catalyzer INDEX 601 for -synthesis of ammonia, 201; occurrence, 531, properties, com- pounds, 532 ; radium from, 473, half-life, 474 ; series of radioac- tive elements, 475 ; sulfate, 532. Uranyl acetate, 532 ; compounds, 532 ; nitrate, 532. Urea, formed in body and from am- monium cyanate, 345. Use of Indicators, 387 ; for weak acids and bases, 389. Vacuum desiccator, 84; high, by means of charcoal, 278. Valence, definition, 63; illustra- tion, 63 ; of elements in oxides and salts, 157 ; of elements in periodic system, 133 ; relation to equivalents and Faraday's law, 438 ; use in determining structure, 323 ; use in writing equations, 156. Vanadinite, 522. Vanadium, occurrence, properties, uses, compounds, 522. Vanadous compounds, 522. van't Hoflf, definition of osmotic pressure, 360. van't Hoff-Le Chatelier, principle of, 111 ; applied to reversible reac- tion, 148; applied in the prepara- tion of sulphur trioxide, 175 ; applied in synthesis of ammonia, 201. Vapor pressure, definition, 74; of hydrates, 82 ; of ice and water, 75. Vaselin, 290. Vauquelin, discovery of chromium, 524. Vegetable foods, 347. Venetian red, 555. Ventilation, 230 ; standard of , 231 ; lack of causes disease, 231. Vermilion, 489. Vinegar, 329. Vitriol, blue, 433 ; oil of, 46, 434 ; defi- nition, 434 ; green, 554 ; white, 483. Volatility, effect on reactions, 374. Volcanoes, source of carbon dioxide, 229. Volt, defined, 33. Volume, unit of, 31. Washing soda, 411. Water, as a solvent, 79 ; calculation of the composition by weight, 68 ; of crystallization, 82 ; degree of ionization, 383; density of at different temperatures, 73. Water, determination of composition by copper oxide, 69 ; by weighing oxygen and hydrogen, 71. Water, determination of, in air by weighing and dew point, 232. Water, dissociation of, 59 ; effect of on chlorides, 112; niters, charcoal not efficient in, 278 ; heat of fusion, 74 ; heat of vaporization, 74. Water gas, heat relations in manu- facture, 298 ; percentage composi- tion, 299; enriched, 297; per- centage composition, 299, 296 ; carbon monoxide in, 297. Water glass, 353 ; of hydration, 82 ; hydrolysis of chlorides by, 115; ionization of, 171, 383 ; maximum density, effect, 72 ; phases and triple point, 78 ; properties of, 72 ; purification of, 83. Water, qualitative analysis and syn- thesis of, 66 ; quantitative ioniza- tion of in relation to indicators, 387 ; sea, amount of carbon dioxide in, 230. Water, use of aluminium sulfate in purifying, 500. Water vapor, effect of, on the volume of a gas, 76. Water, vapor pressure of, 75 ; table, 75 ; weight of 1 liter at different temperatures, 73. Water-soluble phosphoric acid, 461. Waters, effervescent, 309 ; hard, 310 ; natural, 82 ; radioactive, mineral, 476. Watt defined, 34. Waves, explosion, 301. Weak acids and bases, defined, 386. Weber, preparation of chloroplatinic acid, 566. Weight and mass, relation, 32 ; unit of, 31. Weights and Measures, International Bureau of, 31. Welding by thermite process, 498; use of borax in, 367. Weldon process, 103. Welsbach, resolution of dydimium, 504. Welsbach mantles, manufacture, 364 ; theory of, 365. Wentzki, theory of sulfuric acid manufacture, 179. Werner, formula for ammoniocupric sulfate, 434; periodic table, 138; theory of isomeric hydrates of chromic chloride, 526 ; theory of valence, 559. Weston cell, electromotive force, 437, 438. Whisky, 325. White lead, manufacture, properties, 520 ; poisoning by, 521. 602 INDEX White vitriol, 483. Willemite, 481. Wilson, discovery of scandium, 503. Winkler, discovery of germanium, 361. Witherite, 468. Wohler, discovery of aluminium, 495 ; synthetic urea, 345. Wolfram, 11. Wolframite, 530. Wood alcohol, 325; charcoal, 277; oak, composition, 280. Wood spirit, 324. Wood's metal, 269. Woolrich, Sherardized iron, 482. Wrought iron, development of puddling process, 544 ; replace- ment by other irons, 545. Xenon, discovery, 238. X-rays, 474. Xylene, 283. Yaryan evaporator, 405. Yeast, fermentation of sugar by, 325. Ytterbium, compounds, 504. Yttrium group of rare earths, 503 ; occurrence, compounds, 503. Zechentmayer, ferrous chloride and nitric oxide, 554. Zero, absolute, 40. Zero group, 236. Zero point, corrections for ther- mometers, 486 ; depression of in thermometers, 467. Zinc chloride, preparation, properties, use as wood preservative, 483 ; compounds, effect of ammonium hydroxide on solutions of, 491 ; hydroxide, formation, amphoteric, 483. Zinc, in brass and bronze, 431 ; oxide, green with cobalt nitrate, 558. Zinc, occurrence, metallurgy, proper- ties, uses, galvanized iron, 481 ; preparation, properties, uses, 482 ; Sherardized iron, 482. Zinc sulfate, 483 ; use in Clark cell, 437. Zinc sulfide, fluorescence of, with radioactive elements, 472 ; forma- tion, conduct toward acids, 483 ; ^ solubility, 491. Zinc, use in Parke's process, alloys with lead, 441. . Zincic acid, salts of, 483. S Zirconates, 363. Zirconium dioxide, use in Nernst lamp, oxyhydrogen light and for crucibles, 363. Zirconium, occurrence, properties, compounds, uses, 363. Zymase, 344. > j o g o P O o 85 4^ to ^ * in -^ >/ So' If /r S >> CO t-t ^fc'SfrS* 8 s / 89 ? N'i-clSHo I lEi'S'-g III gJ*asJlfe -p SlvUt v r y 171927 FEB 9 1932 THIS BOOK IS DUE ON THE LAST DATE STAMPED BELOW AN INITIAL FINE OF 25 CENTS WILL BE ASSESSED FOR FAILURE TO RETURN THIS BOOK ON THE DATE DUE. 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