U.ofiLPvw THE UNIVERSITY OF ILLINOIS LIBRARY 546 08 '. pi- Return this book on or befoie the Latest Date stamped below. A charge is made on all overdue books. U. of I. Library 1 MAY -4 '36 : ; "■ ' ' 10 1903 ijlflR 30 136 S - 9324-S / ~vfoK C .CjoUIaJL- Inorganic Chemistry Syllabus BY HUBERT C. CAREL, B. S., Instructor in Medical Chemistry UNIVERSITY OF MINNESOTA. PUBLISHED BY University Book Store MINNEAPOLIS 1897 rVH * • \ I 8 8 3 X C\?A INTRODUCTORY DEFINITIONS. Science — Classified knowledge and deduced relation; natural science objective not subjective. Divided — Into two classes which shade into each other: Chemistry Inanimate Zoology Animate j Biology Botany Physics Mineralogy, etc. Physics and Chemistry — Treat of changes of matter; i. e., that which occupies space. Physical — Change in place, condition or properties, substance un- changed. Chemical — Change of substance and properties. Chemistry — Treats of what substances are composed, what occurs when they change and the dependence of properties on composition- chemical knowledge depends on analysis, and synthesis. Analysis — Chemical decomposition of complex, to simple substances. Synthesis — Chemical combination of simple to complex substances. Element — A substance as yet undecomposed. All matter is formed from one element or a combination of elements. Allotropic — Modification of an element is a substance of same elementary composition but differing in chemical and physical properties. Isomerism and Isomers — The above phenomenon in compounds is termed isomerism and such bodies are isomers. Physical — Form of elements— gases, liquids and solids. Chemical — Division — metals and non-metals. Metal (in general ) — Chemical element which unites with the elements of water to form a base (p. 15). Except hydrogen, metals in general are opaque solids, more or less malleable, ductile and tenacious, conductors of heat and electricity and possessed of a peculiar lus- tre termed metallic. Non=Metals (in general) — Are, as the name suggests, the direct opposites of metals, i. e., they unite with elements of water to form an acid. Three of the non-metals are gases : Oxygen, Nitrogen, and Chlorine — one probably a gas, Fluorine — oneisa liquid, Bromine — the rest are solids. Compound- Union of two or more elements in simple proportions, to form a new substance. Mixture — Mechanical intermingling of matter in any or all propor- tions, without change of substance. The original materials may be recovered by mechanical means. Chemical Affinity — Tendency of elements to unite — C. A. of a given element may vary in uniting with different elements, but is always constant with any particular element. Some elements unite directly to form a compound A+B=AB. More often an element unites with a compound A+CD=AD+C; or two compounds react to form two new compounds AB-j-CD=AC-l-BD. Metathesis— The last reaction is known as metathesis (Gr., to set over), and occurs especially in solution, where one of the produces is insoluble, or by heat, where one of the products is volatile. Heat and Chemical Energy — Save nitrogen, all union of elements yields heat; that is, chemical energy transforms to heat. Decomposition of com- pounds absorbs heat which is transformed into chemical energy. Acid (in general ) — Compound of hydrogen which yields salts by replacing its hydrogen with a metal. Acid reddens blue litmus and possesses characteristic “acid” taste. That part of acid which unites with the metal is called “acid radical.” Base (in general ) — Oxide or hydroxide (OH) of a metal which exchanges its metal for the hydrogen of an acid forming a salt. Base blues litmus and possesses characteristic “alkali” taste. Salt— The generally neutral result of the combination of metal of base with radical of acid. (base) (acid) (salt) 2NaOH + H 2 S0 4 = Na 2 S0 4 + 2H 2 0 (met- (acid (met- (acid al) radical) al) radical) Neutralization — This reaction of an acid body with alkali forming a sub- stance neither acid nor alkaline is termed neutralization. LAWS OF CHEMICAL PROPORTIONS. Composition of a given substance is invariable. Definite Proportions — Chemical combination always takes place between definite masses. Any excess is unacted on. Multiple Proportions — When elements unite in more than one proportion, the ratio of succeeding compounds are simple multiples of the first ratio. Example : N 2 0 — N 2 0 2 — N 2 C> 3 — N 2 0 4 — N 2 0s. Atomic Theory. Above laws depend on the Atomic Theory that every body is an aggregate of atoms or ultimate indivisible particles. The absolute weight of an atom is of course unobtainable, but we have relative atomic weights with hydrogen as the unit. Atom vs. Molecule. Atom = smallest indivisible particle. Molecule = smallest divisible particle. Law of Charles — Volume of gases varies directly as the absolute temperature. Law of Mariotte — Volume of gases varies inversely as pressure. Law of Avagadro — At same temperature and pressure equal volumes of all gases contain the same number of molecules. Certain Physical Properties — Of bodies are always the same for the same substance, i. e.: (1) The boiling point of liquids. (2) The melting point of solids. (3) The crystalization of solids. SOLUTION. Some substances require particular solvents, as ether for fats, carbon disulphide for yellow phosphorus, etc., but the great solvent is water. In stud 3 ung the phenomena of solution we notice, (1) the uniform distribution of dissolved substance, (2) the almost infinite subdivision of material, (3) that sub- stances vary as to solubility, and that the presence of foreign matter often affects solution. (4) Evaporate a solution and (a) solidsremain, (b) gases escape, (c) liquids pass off at their boiling points. (5) Solution is of great advantage in chemical reactions where often the new substance is insoluble and can be separ- ated. GROUPING OF ELEMENTS WITH REGARD TO PROPERTIES. Non=Meta!s — Group I— Cl, Br, I, F. Group II — 0, S, Se, Te. Group III — N, P, As, Sb, Bi. Group IV — B. Group V— Si, C. Metals — Group I — H. Group II — K, Rb, Cs, Na, Li. Group III— Ca, Sr, Ba. Group IV— Be, Mg, Zn, Cd, Hg. Group V — Pb, Tl, Cu, Ag, Au. Group VI— Y, La, Cr, Di, Sr. Group VII — Al, Ga, In. Group VIII— Mn, Fe, Co, Ni, Cr, Mo, W, U. Group IX — Sn, Ti, Zr, Th. Group X— V, Ta, Nb. Group XI — Pt, Ru, Rh, Pd, Ir, Os. OXYGEN (to generate acid). At. Wt. 16. Val. II. Symbol 0. Occurrence — % of water, Ys of air, Y2 earth’s crust, % organized matter. Preparation — By heating ( 1 ) mercuric oxide HgO, ( 2 ) manganese dioxide Mn02, ( 3 ) potassium chlorate KCIO3. ( 4 ) Decomposition of water (H2O) by electricity. ( 5 ) In the laboratory, by heating a mixture of KCIO3 + Mn02- (6) Commercially obtained by heating barium oxide BaO to a red heat in the air. This forms barium peroxide Ba02, which heated to a white heat loses half its oxygen and is reduced again to the oxide. BaC >2 — BaO + 0. Properties — Colorless, inodorous, tasteless gas — the great supporter of combustion and life — combines with all elements save fluorine to form oxides. At ordinary temperatures oxygen is mod- erately active, but its chemical affinity increases with heat. 0 is slightly soluble in water — with low temperature and great pressure condenses to a liquid. Detection — Distinguished from all gases but nitrous oxide N2O by its spark. From N2O ( 1 ) by the solubility' of 0 in potassium pyro- gallate, ( 2 ) Sulphur burns in 0 but not in N2O, ( 3 ) 0 , exploded with two volumes of hydrogen (H), loses all its volume, while N2O leaves a residue of N. Combustion — Rapid oxidation accompanied by light and heat. Decay — Slow oxidation of organic matter assisted by bacteria. Kindling Point — Temperature at which bodies combine with light and heat — Interposition of such bodies as wire gauze serves to reduce the temperature below the kindling point. In Davy Safety Lamp, a wire gauze jacket surrounds the flame. The heated gases passing through such a jacket are cooled below the kind- ling point, and there can be no combustion. Small quantities of “fire damp” passing through the gauze from without, are heated within the jacket to their K. P. and cause a small ex- plosion which warns of a dangerous proximity. Flame — A burning gas whose luminosity is usually due to red hot un- oxidized carbon (C). Intense light is produced by the glowing of some non-volatile bodies in a flame as calcium oxide (CaO) in calcium light. Other things being equal, combustion produces less light and more heat in proportion as oxygen is increased. In illuminating tips, gas burns by oxygen obtained at tip of tube, while in the colorless “Bunsen” the same gas is well mixed with air before it ignites. In the oxv hydrogen lamp or blow- pipe Oxygen alone is mixed with Hydrogen giving an intense non-luminous heat. The comparative heat of burning bodies is known and Carbon (C) yields the best fuel by volume. Oxidation — Adding of 0 to any substance. Reduction— Subtracting of 0 from any substance. Reactions of Oxides — Oxides may have one of three reactions to litmus. Acid — Basic — Neutral — so 2 , k 2 o, h 2 o, no 2 , Na 2 0, MgO, etc. etc. etc. Red. Blue. Purple. 9 OZONE, 0 3 . Oxygen occurs in two alio tropic forms, ordinary oxygen O 2 and ozone O 3 . Occurrence — Whenever oxygen is prepared, small amounts of ozone form. It is supposed to exist in the air, but we have no proof. Preparation — (1) By the silent discharge of an electric machine. (2) Slow oxidation of phosphorus. (3) In the laboratory by treating barium peroxide with sulphuric acid. BaC>2 + H2SO4 — 0(3) + H2O + BaS04. Properties — “Condensed oxygen” is a heavy gas — poisonous — taste and odor like weak chlorine — at ordinary temperature is a powerful oxidizing, bleaching and disinfecting agent. Ozone acts with its extra atom of 0 — O 3 = 02 + 0. Here 0 does the work, and 0 2 (ordinary oxygen) is set free. Ozone condenses to a blue liquid — at 300° breaks to O 2 . Detection — (1) O 3 blues potassium iodide (KI) starch paper (common to Cl, NO 2 and H 2 O 2 ). ( 2 ) Blues red KI litmus paper. This is distinctive in the absence of ammonia (NH 3 ). Chief use of 03 is to bleach old pictures. HYDROGEN-H. v. At. Wt. 1. Yal. I. History — Discovered by Cavendish in 1766. Occurrence — Free in volcanoes, oil-wells, and from decomposing organic matter; chiefly combined with oxygen in H 2 O and organic bodies. Preparation — (1) From H 2 O by (a) electricity, (b) metallic sodium or 10 potassium, (c) hot iron. (2) Laboratory method, by replacing H of an acid by a metal. Zn + 2HC1 = H 2 + ZnCl 2 . Properties — Tasteless, inodorous, colorless gas — does not support com- bustion, but burns with a blue flame to H 2 O — mixed with oxy- gen explodes at K. P. — lightest of gases, H is the unit of specific gravity for gases — highly diffusible and a strong reducing agent — H is a gaseous metal playing the same part in acids as the other metals play in salts— conducts heat and electricity, forms alloys, and in electric decompositions, goes to the negative (metal) pole. Nascent gas — Condition at immediate generation when chemical affinity is much stronger than usual — the atoms have not yet united to form the molecule. Nascent atoms. — Molecule. H + H = H 2 . WATER— H 2 0. At ordinary temperature a transparent fluid devoid of taste or smell — thin layers colorless, large masses blue — cooled to a certain temperature water c^stallizes, forming ice — this is the 0° C.— heated to a certain temperature water boils — this is the 100° C. On cooling water contracts till 4° C., which is its maximum density — below 4° it expands, hence ice floats, and water pipes break. H 2 0 evaporates at all temperatures, ice vaporizes slowly and without changing to liquid form. Heated water changes rapidly to an invisible gas called Steam, which partially recondenses to the liquid state when in contact with air. Composition of water — The electric current resolves water into two volumes of H and one of O ; conversely, by the electric current — or by heat — two volumes of H and one volume of O always combine to H 2 0 without residue, and the H 2 0 thus formed may be heated to two volumes of steam. Thus water by volume = H 2 + 0, or by weight = 2 H to 16 O, or 1 H to 8 0. Distilled water — Natural waters contain many substances in solution and to obtain water free of solids it must be boiled and the steam condensed, i. e., distilled. Water of Crystallization — Molecules of water which enter into definite combination with many salts, when these crystallize from their water solu- tions. It often influences the color, and crystalline form, but may be driven off by heat without alteration of substance. Efflorescence — Some salts yield their water of crystallization to the air as Na 2 C0 3 , 10H 2 6, etc. Deliquescence — Some salts take up water from the air and ultimately melt, as CaCl 2 , NaNC> 3 , etc. Unit of density — Water at 15° C. is unit of density for solids and liquids. Hydrometer — An instrument for determining specific gravity of liquids with reference to water (1000) HYDROGEN PER0XIDE-H 2 0 2 . History — Discovered by Thenard in 1818. Occurrence — In small quantities in the air and formed wherever ozone oxidizes in presence of H 2 0. Preparation — When barium peroxide Ba0 2 is treated with H 2 S 04 a double reaction takes place. Ozone 0( 3 ) is generated, which then oxidizes the water formed to H 2 0 2 . Ba0 2 + H 2 S0 4 = 0(3) + H 2 0 + BaSO*. 0 3 + H 2 0 = H 2 0 2 + 0 2 . Properties — “Ozonic ether,” “golden fluid,” “a oxygenated water,” H 2 0 2 is a thick colorless liquid. Soluble in H 2 0 — astringent taste — dilute Cl odor — strong oxidizing and disinfecting agent, contains more O than any other substance (94%), concentrated solution easily breaks to H 2 O + O — dilute slightly acid solution is more stable. Detection— (1) Blues KI starch paper. (2) When shaken with potassium bichromate (K^C^CL) and ether (CaHs^ 0 it gives the latter a blue color. Use- Antiseptic spray, and to bleach old pictures and hair. CHLORINE— Cl. At. Wt. 35.5. Val. I. History — Discovered by Scheele in 1774. Occurrence — Chiefly as common salt NaCl in sea waters, springs, and especially rock salt deposits. Never free, because of strong affinities. Preparation — From NaCl, H 2 SO 4 and Mn 02 a triple reaction. t (hydro- chloric) 2NaCl + H 2 SO 4 = 2HC1 + Na 2 S0 4 . (manganese (manganese dioxide) tetrachloride) Mn0 2 + 4HC1 — MnCL + 2H 2 0. MnCU (heated) = Cl 2 + MnCl 2 breaking up of MnCU shown by change of color — gas collected by displacement, or over hot water or salt solution. Properties — Green \^ellow gas — violent odor, poisonous; strong chemical agent at ordinary temperature. For example, Cl extracts H from such hydrocarbons as turpentine — with some metals as antimony Sb it combines with light and heat; with H it unites in direct sunlight, forming HC1, or Cl will burn in H gas to form HC1 and vice versa. Cl is a strong disinfecting agent and indi- rectly bleaches through the action of nascent 0 which it frees from H2O. Cl is easily soluble in H2O, forming a solution which retains the properties of Cl gas. Cl used in commerce chiefly as bleaching powder. Bleaching Powder (see Ca) — Is formed by running Cl gas over slaked lime Ca( 0 H) 2 — its approximate composition is CaCl(OCl). It disinfects and bleaches chiefly through nascent chlorine which is freed by dilute acids. 2CaOCl 2 + 2HC1 m 2HC10 + CaCl 2 . HCIO + HC1 — Cl 2 + H 2 0. Detection — By its odor, bleaching power, and by turning KI starch paper blue. CHLORINE OXIDES. Chlorine forms a series of oxides similar to that of nitrogen, 2. e., CI2O, CIO2, CI2O3, CI2O5. All are formed indirectly and unite with water to form acids. CHLORIC ACID— HCIO3. Preparation — The anhydride CI2O5 is not yet known in the free state. The K salt of the acid is formed by passing Cl into KOH, 6KOH + 6C1 — KCIO 3 + 5KC1 + 3H 2 0. The K salt is then treated with fluosilicic acid H2SiF6 2KC10s + H 2 SiF 6 = 2 HCIO 3 + K 2 SiF 6 . Properties — A liquid of faint odor; strongly acid reaction; easily decom- posed ; a strong oxadizing and bleaching agent. The K salt is a source of oxygen supplying enough to burn some com- bustible substances. White gunpowder iscomposed of KCIO3 and sugar. HYPOCHLOROUS ACID— HClO. The anhydride CI2O is a red yellow gas with a Cl odor; explosive; condenses to a liquid; is prepared by the action of Cl on HgO. * mercuric oxide HgO + 4C1 = C1 2 0 + HgCl 2 . 14 - The gas dissolves in water forming HC10 C1 2 0 + H 2 0 = 2HC10. The Ca salt Ca(C10)2 is highly important being the active principle of “bleaching powder.” Hypochlorites in general are very unstable and act as strong oxidizing agents. The other oxides of Cl are unimportant. HYDROCHLORIC ACID— H Cl. Occurrence — HC1 found free, in the gastric juice and in volcanoes — com- bined, chiefly as NaCl. Preparation — Generally prepared by action of sulphuric acid on common salt. H 2 SO 4 + 2NaCl = 2HC1 + Na 2 S0 4 . May be formed by direct union of H and Cl in the sunlight or by heat or electricity. Properties — Absolute HC1 is a colorless, suffocating gas — non-combust- ible, and non-supporter — fumes in damp air — is eagerly absorbed by water, forming a 33 per cent solution which is the concen- trated HC1 of the laboratory — with silver nitrate (AgNOs) hy- drochloric acid forms a curdy, white, insoluble precipitate of silver chloride, AgCl, which is the test for HC1. BROMINE— Br. At. Wt. 80. Yal. I. History — Discovered by Balard in mother liquor of salt works, 1827, Occurrence — Never free — usually as K, Na or Mg bromide — in salt water (1 gr. per gal.), rock salt deposits, sea-weeds and springs — ac- companies chlorine and is obtained from mother liquor of salt works — Br is one of the less common elements. Properties — Dark red liquid — red brown vapor — irritant poison — char- acter similar to Cl but with weaker affinities being freed from its compounds by Cl — Chemical unions of Br often accompanied by light — Br dissolves in water, alcohol, ether and chloroform. Preparation — Freed from its salts by (1) chlorine. 2NaBr + Cl 2 = Br 2 + 2NaCl. (2) Manganese dioxide and sulphuric acid. 2NaBr + Mn0 2 + 2H 2 S0 4 = Br 2 + K 2 S0 4 + MnS0 4 + 2H 2 0. ( Same stages in this reaction as in preparation of Cl. ) Detection — Add Cl water and CS 2 = red brown color to CS 2 . Br is set free by Cl — CS 2 dissolves free Br with a red brown color. Br gives starch an orange color. Use — Br is used in dye factories, photography, medicine. HYDROBROMIC ACID— HBr. Properties — Colorless gas — fumes in moist air, eagerly absorbed by water, forming HBr 2 H 2 O. Analogous to HC1, but less stable, is de- composed by H2SO4. Bromides freed by chlorine. Preparation — (1) Direct union of H and Br at red heat. (2) Generally prepared by action of phosphorus tri bromide and water. (Phosphorous acid) PBr 3 + 3H 2 0 = H 3 PO 3 + 3HBr ( 3 ) Cannot be made pure from salt by H2SO4, as resulting HBr is decomposed bjr H2SO4. Solubility — Most Bromides are soluble in H 2 O OXIDE AND CHLORIDE OF Br. Bromine chloride, BrCl — Br absorbs Cl forming a yellow unstable liquid with bleach- ing properties. Above 10° C. changes to Br and Cl. Bromine oxides — Bromine and oxygen unite with difficulty to form a series entirely analogous to Cl and made in same way — Br 2 0, — , Br0 2 , Br 2 03 , Br 2 0s. IODINE— I. At. Wt. 127— Val. I. History — Discovered by Courtois in 1811 . Occurrence — Not free — compounds accompanying Cl and Br, in sea-water, springs, etc. — Chief source is the ash of sea-weeds called “kelp.” Prope rties — Grey black solid with violet vapor — weak Cl odor — difficultly soluble in water — dissolves easily in alcohol or a solution of KI — chemically similar to Cl and Br but with weaker affinities, is freed by either from its compounds. Preparation- Kelp solution is partly evaporated — I remains in the mother liquor— solution is then distilled with MnC>2 and H2SO4. 2KI + Mn0 2 + 2H 2 S0 4 = I 3 + M 11 SO 4 + K 2 S0 4 + H 2 0 ' (Game reaction as for chlorine). Detection — (1) With Cl water and CS2 gives violet-color. ( 2 ) With cold starch paste, a blue color. Used- In dye factories and photography. HYDRIODIC ACID — HI. Properties— Colorless gas with suffocating odors — fumes in moist air — absorbed by water— easily decomposed; hence, cannot be made with H2SO4— Strong reducing agent. Preparation — ( 1 ) From Phos-tri-iodide and water. PI 3 + 3H 2 0 = H 3 PO 3 + 3HI (strong acid). (2) I with H2S gives weak acid. I 2 + h 2 S = S + 2HI. 17 OTHER COMPOUNDS OF IODINE. Iodine Chlorides (IC1, IC1 8 ) — Two compounds IC1, ICI 3 formed by direct action of the elements; used to add I to organic bodies. Iodine Bromide (IBr) — IBr is a solid similar to I and formed as IC1. Iodine Oxides — Iodine forms the same oxides as Br or Cl; I 2 O, IO 2 , I 2 O 3 , l 20 £. Iodic Acid (HI0 3 )— A white solid formed by boiling I in HNO3; permanent in air, but with heat changes to I2O5 +H2O; forms salts similar to chlorates. FLUORINE— FI. At. Wt.19. Val. 1. History — Discovered by Davy in 1812, but not isolated till 1886 by Moissan. Occurrence — Chiefly in fluor spar or fluorite, CaF 2 , and cryolite AlFla, 3NaFl. Properties — Colorless gas — deadly poisdn — attacks most substances, but does not combine with O, C, Pb or Pt, nor attacks guttapercha — properties analogous to, chlorine but much stronger, its atomic weight being much less. F frees Cl, Br or I from their com- pounds. ’• HYDROFLlfbRIC ACID, HF. Preparation in Et retort •% *»i From fluor sparjfov j daft + Ullft ^aS0 4 + 2HP. Properties— , . L N Colorless, mobile liquid — fumfes in air — corrosive poison — dissolves all metals ex^pi J^b;: Pt and Au, and all oxides, in- • *' 1 * "' -*'**■ IS eluding SiC> 2 . Must be kept in gutta percha bottles. Used to etch glass and detected by its etching quality. SULPHUR— S. At. Wt. 32. Val.II. History — Used by the ancients for fumigation and medicine — Consid- ered by alchemists to be the principle of combustion. Occurrence — (1) Native in volcanic countries, especially Sicily. (2) Sulphates and Sulphides. (3) Organic compounds. The amount of S produced is about 375,000 tons, of which nine-tenths comes from Sicily. Extraction from Ores — (1) Melting ores in a pot, S floats and is dipped out. (2) Burning ores in small supply of air so that most S melts. Purified — By distillation — That which passes over first cools more rapidly forming flowers and is less pure than the roll which forms afterwards — Flowers contain H2SO4; H 2 S and SO2 as impurities. Properties— Strong affinities — Union with metals often produces light and heat. Moist Fe -f- S = FeS + heat — unites with haloids at ordinary temperature — Slightly heated S + P explodes — H run into melted S = H 2 S; Compounds of S alone are called sul- phides. S is wholly analagous to O and when oxides of an element are soluble, usually sulphides are. With O, S burns with blue flame to SO 2 . Uses — Disinfectant; making H2SO4, SO2, matches and gunpowder. HYDROGEN AND SULPHUR. Two compounds well known— hydrogen sulphide, H 2 S, and persulphide, H 2 S 2 . The latter is unimportant. Hydrogen Sulphide (H 2 S)— Occurrence- Free in mineral springs, and when organic bodies decay. Preparation — (1) From a sulphide and an acid — FeS + 2HC1 = H 2 S + FeCl 2 . (2) Chemically pure — CaS + 2HC1 = H 2 S + CaCl 2 Collect over warm water or salt solution. (3) Formed by heating elements to 400° C. Properties— Colorless gas — bad odor — poisonous — combustible, burns to S0 2 and H 2 O— mixture with air explodes— unstable, oxidizes in air at ordinary temperature to S + H 2 0. Easily unites with metals to form sulphides. Dissolves in H 2 0 (3 volumes to 1), and water solution has properties of the gas, decomposing in light to H 2 0 + S. HoS is a dibasic acid, hence forms two series of Salts, viz., KHS =. potassium sulphydrate, K 2 S = potassium sulphide. SULPHUR HALOIDS. S unites with Cl in three proportions— S 2 C1 2 (SCI), SC1 2 , SCI 4 . The monochloride S 2 C1 2 is a red yellow liquid formed by passing Cl over molten S — used in vulcanizing rubber; SCl 2 and SCU are highly unstable. With Br and I, S forms similar compounds. SULPHUR DIOXIDE, S0 2 . Occurrence — Native in volcanic gases. Preparation — (1) Oxidizing sulphur — S + 0 2 =r S0 2 . (2) Reducing sulphuric acid (H 2 S 04 )— Cu + 2H 2 S0 4 = CuS0 4 + S0 2 + H 2 0. S0 2 is called “Sulphurous acid” because the acid proper, H 2 S0 3 , is unstable. Properties — Colorless* suffocating gas — inhibits combustion — poisonous to plants and animals — disinfecting agent — bleaches, but color is restored by alkalies — easily condensed without pressure — highly soluble in H 2 0 (50 volumes to 1)— salts easily oxidize to sulphates, hence good reducing agents. H 2 SO 3 is a weak, dibasic acid, little stronger than carbonic— S 0 2 does not unite directly with O but with ozone forms SO 3 . SULPHUR TRIOXIDE— S 0 3 . Preparation — Sulphuric anhydride formed from S0 2 by ( 1 ) ozone, ( 2 ) O over divided Pt, (3) electric spark, (4) distilling disulphuric acid, H 2 S 2 O 7 . Properties — White, silky, crystalline needles — fumes in damp air— acid only with H 2 0. Decomposes at high temperature to S0 2 + O. Unites with H 2 0 with great heat. SULPHURIC ACID-H2SO4. History — Known for centuries — alchemists prepared it from S and nitric acid (HNO 3 ) or by heating FeS 04 , calling it “oil of vitriol” — first lead chambers built in 1746. Occurrence — Free in volcanic rivers and the fluids of some mollusca — combined as sulphates in large quantity as CaS 04 , etc. Preparation — ( 1 ) Oxidation of S by HNO 3 directly. (2) Oxidation of S0 2 by HNO 3 in presence of H 2 0. — S0 2 from burning FeS 2 (iron sulphide), with O from the air, is passed into lead chambers, through which HNO 3 and H 2 0 are streaming — HNO 3 is decomposed to the oxides of N — these yield up O to S0 2 forming SO 3 , which unites with H 2 0 to form H 2 SO 4 . The lower oxides of N reunite with O of the air, forming the higher, N0 2 and N 2 O 3 , which again oxidize a new portion of S0 2 . The dilute chamber acid (sp. gr. 1.5(c) is concentrated in Pt or glass vessels to commercial strength ( 1 . 830 ). Crude H0SO4 contains as impurities PbSCU, HNO3, often As from S or FeS2 used. Impurities are generally removed by distillation, but an arsenic-free acid is obtained only by using arsenic-free materials. * Properties— Heavy, svrup-like liquid, colorless when pure, usually brown from charred organic matter — odorless — corrosive — in- tensely acid. Heated above boiling, gradually breaks up to PI2O + SO3 — mixed with water great heat is evolved, due to formation of higher hydrates — H2SO4 + H2O or 2H2O. Sul- phates of strong bases are not easily broken up by heat. FUMING OR NORDHAUSEN ACID— H 2 S 2 0 7 . By running SO 3 into concentrated H 2 SO 4 we get H 2 S 2 O 7 , disulphuric acid, a white solid stronger than H 2 SO 4 and used in indigo and alizarin manufacture. SELENIUM (Se) AND TELLURIUM (Te). At. Wt. 79. At. Wt. 128. Two rare elements closely analogous to Sulphur — Tellu- rium was discovered in 1782 in ores of gold — physical pro- perties are metallic, chemically it acts like S. — Selenium discov- ered in 1817 by Berzelius from the mud on floor of lead cham- bers. HNO3 with Te or Se forms respectively tellurious acid, H 2 Te03 or selenious acid, H2Se03 — with KNO3 the elements form K2Te04 or K2SeC>4. Selenium colors the flame blue and H2Se is remarkable for its odor. Selenium changes its resis- tance to an electric current when exposed to light. SULPHUR GROUP. 0 16— S 32— Se 79.4— Te 128. O a gas — others solids — O and S are non-metals — Se partly metallic — Te a metal. All unite with H2 — the three solids form with H 2 volatile strong smelling gases of acid nature — Stabil- ity of H compounds, varies inversely as atomic weights, of O compounds directly — O is always bivalent, the others vary in valence 22 NITROGEN— N. At. Wt. 14. Val. III. or V. History — Discovered in 1772 by Rutherford — Scheele and Lavoisier showed that air = O + N, called azote by the French and Ital- ians. « Occurrence — % of the air and in many organic substances. Preparation — (1) Removing O from air by (a) Phosphorus — (b) red-hot Cu. (2) Heating Ammonium nitrite (H4N)N02 (H 4 N)N0 2 = N 2 + 2H 2 0. Properties — Characterized by its inertness at ordinary temperature — colorless — inodorous, tasteless gas — non-combustible — non-sup- porter — not poisonous — at white-heat unites with metals— many of its compounds are explosives, while all explosives contain N. AIR. A comparatively constant mixture of % N, % 0, with appreciable quantities of H 2 O and CO 2 and trace of NH3 com- pounds, NaCl, and dust. N and 0 are quite constant; the oth- ers vary. Impurity of air is usually determined by amount of CO 2 present. CO 2 of itself is not harmful in small quantity, but ex- perience has shown that it bears a constant ratio to the amount of ammonia compounds present. This “organic waste,” so- called, is intensely poisonous, although it exists in quantity too small for easy determination. Air has the same physical properties as the gases of which it is composed, being a mixture, not a compound, because (a) it varies somewhat in composition, (b) proportion of O and N is not governed by atomic weight, (c) absorbed by water as free gases 35 vol- O : 65 vol. N, not 1 vol. 0 : 4 vol. N. AMMONIA— NH 3 . History — Called “spirits of hartshorn” because formerly obtained 23 from horns of deer. Gas discovered by Priestley, who called it alkaline air. Occurrence — Not free in nature as NH3. Found as salt in animal secre- tions, especially urine. Found also when nitrogenous organic bodies decay or are distilled, hence found in liquor of gasworks, as coal contains N compounds. Preparation — Commercially obtained from gas works. In laboratory from ammonium chloride (H4N) Cl and calcium hydrate Ca ( 0 H) 2 . x 2 H 4 NCI -I- Ca(OH ) 2 = 2NH 3 + CaCl 2 + 2H 2 0. Properties — Colorless gas — pungent odor — strong alkali — non-supporter of combustion — burns feebly in the air. Decomposed by electric- ity or by chlorine in sunlight. 4NH 3 + 3C1 — 3H 4 NC1 + N. NH3 is highly soluble in H 2 0 , forming H4NOH, ammonium hydrate, a strong base in which the radicle, H4N, deports itself like the metals Na or K, and will be treated with them. NITROGEN HALOIDS Do not form directly from elements — all are highly explosive — most important is NCI3 — a dark red liquid, formed by action of Cl on H4NCI. H 4 NCI + 6C1 = NCls + 4HC1. NBr3 and NI3 are similar. NITROGEN OXIDES. Five known compounds — Nitrous oxide. Nitrous anhydride. Nitric oxide. Nitrogen peroxide. Nitric anhydride N 2 0 . N 2 0 3 . N 0 (N 2 0 2 ) N 0 2 (N 2 0 4 ). n 2 o 5 . The first, second and fifth act as anhydrides, forming hypo- nitrous, nitrous, nitric acids. The third and fourth form a mix- ture of HN 0 2 and HNO3. All the oxides are prepared from HN 0 3 . Nitric Anhydride N 2 0s — Discovered by Deville. Made by passing Cl over dry AgN03— apparatus entirely of glass, melted together. 24 2AgN0 3 + Cl 2 = 2AgCl + N 2 0 5 + O. White, unstable solid — with H2O forms HNO3. N 2 0 5 + H 2 0 = 2HN0 3 . NITRIC ACID— HN 0 3 . History — Known early as 9 th century when it was made by distilla- tion of ZnSC>4 + KNO3 and called ‘‘Aqua fortis,” “Parting water,” or “Spiritus nitri fumans Glauberi.” Occurrence — Not free — in salts as nitrates, is widely distributed in the earth especially in Chili as NaN(>3 or Chili Saltpeter — in small quantities in air and water. Formation — From elements by electric spark in presence of H2O — also in process of decay, occasioned by bacteria, the so-called “nitrifi- cation,” which takes place best in the dark and is stopped by destruction of the bacteria through chloroform or boiling. Preparation — Action of H2SO4 on K salt. 2KN0 3 + H 2 S0 4 — 2HN0s + K 2 S0 4 . Impurities — Ordinary KNO3 contains KC 1 ; hence, HC 1 is formed with crude HNO3 — another common impurity is H2SO4. Properties — Pure HNO3 is a colorless volatile liquid leaving no residue- no stronger acid has been yet obtained — so powerful an oxidiz- ing agent that it deflagrates with easily combustible substances — exposed to sunlight it gradually decomposes to its oxides and H2O; hence, strong acid is often colored by NO2 — a powerful mono-basic acid; it dissolves most of the metals forming nitrates all of which are soluble in water, some insoluble in HNO3, all decompose at high heat. Used— Commercially to etch copper. 25 Aqua Regia — Formed by mixing HNO3 + HC 1 ; acts as nascent chlorine, forming a chloride. 2HN0 3 + 6HC1 = 4H 2 0 + 2N0C1 + Cl 4 . A long known mixture of dark yellow color and suffocating fumes and odor — called “royal water” because it dissolves “no- ble” metals, gold and platinum. Nitrogen Peroxide NO2 — Whenever NO is generated it combines with 0 of air to NO2 — red brown gas — poisonous, suffocating odor — very corrosive — colors organic bodies yellow — strong oxidizing agent, C and P burn in its vapor — easily liquified and solidified. NO2 does not form a hydrate, but with H2O gives 2N0 S + H 2 0 = HN0 2 + HNOs. Nitrous acid HNO2 — Hydrate of N2O3 and formed by heating HNO3 with starch, sugar, or other easily oxidizable substance, which reduces HNO3. In same way nitrites are prepared by reduction of nitrates — KNO 3 + Pb = PbO + KN0 2 . Nitrites are mostly soluble in H2O. Nitrous anhydride N2O3 — Formed by decomposition of nitrites — a red brown gas easily condensed to indigo blue liquid. Passing the gas into alkali hydrates forms corresponding nitrites — 2K0H + N 2 0 3 == 2KN0 2 + H 2 0. Nitric Oxide NO — Never found free, as it combines with 0 of air to form NO2. Made by action of HNO3 on Cu. 3Cu + 8 HNO 3 = 3Cu(N0 3 ) 2 + 2N0 + 4H 2 0. At high temperature N2O is formed. Colorless gas, irrespirable, poisonous — oxidizes in air to NO2 — slightly soluble in H2O. NO supports combustion of P and S but not of ordinary substances, as wood. Myponitrous Acid HNO — Discovered in 1871 by reducing KNO3 with NaHg. Known only as salt. Nitrous oxide NoO — Discovered by Priestley 1776 . Colorless gas — sweet taste — will not support life, but can be breathed a short time, producing intoxication and anaes- thesia (laughing gas). N2O supports combustion like O but extinguishes burning S. Condenses easily and forms with ether (C2Hs)20 the coldest known mixture — distinguished from O by leaving a residue of N when exploded with H. Prepared by heating (H4N)NC>3 (ammonium nitrate) — (H 4 N)N0 3 = n 2 0 + 2H 2 0. At high temperature NO forms with explosion. Valence is that property of an element by which it combines with defi- nite masses of other elements. Unit of valence is hydrogen. Elements are classified according to their power of uniting with or replacing different proportions of the unit. Thus chlorine unites with one H to form HC 1 and is a univalent element. Oxygen forms H2O, requiring two parts of H and is a Divalent element. Nitrogen forms H3N, acting as a trivalent element. Carbon forms H4C, acting as a quadrivalent element. In a comparatively few instances valence requires five, six or seven InTlrogens. Under the same conditions, valence of a given element is always constant, but it may vary for the same substance in different compounds, depending on what element it unites with and under what conditions union takes place. Generally when a substance is present in smaller mass it unites with higher valence. Thus a small amount of carbon heated in presence of much oxygen forms CO2, where C is quadrivalent, but where a large amount of C is heated with small 0 the gas formed is CO, where C acts as a bivalent element. The greatest variation appears in the oxides and chlorides. The oxides of N for exam- ple— N20,N202,N203,N204,N205— give to nitrogen five distinct valences ; and other elements form similar series. In general, valence may be considered less as a property of the specific elements and more as a function incident to their combination. PHOSPHORUS— P. At. Wt. 31. Val. Ill or V. History — Discovered in 1667 by Brand when searching for philoso- pher’s stone in urine — which was the only source till 1750 — later Scheele obtained it from bones. Occurrence — In soil as apatite and pyromorphite — extracted from soil by plants — then to animals in bones, brain and urine. Properties — Occurs in two important allotropic forms, yellow crystal- line and red amorphous phosphorus. Yellow is converted by heat ( 250 °) to red — ordinary P exposed to light becomes red on surface — common or yellow P is a wax-like solid — ozone odor — luminous in the dark — oxidizes easily— burns at 44 ° — a mixture of P and KCIO3 detonates when struck — causes chronic poisoning — almost insoluble in water, but dissolves easily in CS2 — preserved in H2O. . Red P is a weaker chemical agent — odorless, non-phosphorescent — not poisonous — does not dissolve in CS2 — is stable at ordinary temperatures — at 260 ° C red changes back to yellow. Preparation — Bones are burned to remove organic matter ( 55 % whole). Ash treated with H2SO4 — this changes insoluble Ca3(P04)2 of the bones to a soluble phosphate which heated with charcoal gives free P — the essential reaction with Cis 2P2O5 + 5 C = P4 + 5CO2. P is purified by redistillation. Cannot be success- fully prepared on small scale. Use— Chiefly for making matches. Safety matches = mixture of potassium-chlorate (KCIO3) and antimony trisulphide (Sb2Ss) — the surface = red phosphorus + Mn(>2. Ordinal match = yellow P, some oxidizing agent as MnC>2 + KCIO3 and glue. PHOSPHINE— PH 3 . Properties — The only important compound of P and H — a gas of disa- greeable odor and highly poisonous — chemical action similar to NH3— forms phosphonium compounds analogous to H4N com- pounds. P and H form also P2H4, a liquid, and P4H2, a gas, but these are unimportant. PH 3 does not ignite spontaneous- ly, except in presence of P2H4. Preparation — From KOH, P and H2O forming PH3 and an- acid of P. 3 KOH + P 4 + 3 H 2 0 = PH 3 + 3 KH 2 PO 2 . Phosphorus trichloride PCI3 — The more important chloride of P — prepared by action of Cl on melted P — a colorless liquid which fumes in moist air and with water forms phosphorus acid. 2 PC1 3 + 6 H 2 0 = 6 HC1 + 2 H 3 PO 3 . There is also a penta-chloride prepared in same way but with excess of Cl — a white crystalline substance. — I and Br form similar compounds, all of which are much used in the arts. Phosphorus Oxides — P forms two oxides by burning in greater or less quantity of dry air, P2O3 and PO5 — both white feathery solids — P2O5 is more important ancr generally called phosphoric ant^dride — unites eagerly with water in different proportions to form acids. ( 1 ) P2O6 + H2O — 2HPO3 (Meta- or glacial phosphoric acid). ( 2 ) P2O5 + 2H2O = H4P2O7 (Pyrophosphoric acid). ( 3 ) P2O5 + 3H2O = 2H3PO4 (Orthophosphoric acid). Orthophosphoric acid H3PO4 — P is oxidized with HNO3 and residue evaporated — difficult operation as oxidation is slow and does not work with weak while it explodes with strong HNO3 — can be made also ‘from the Ca salt but is difficult to purify. Salts — Form three classes — M2PO4, M2HPO4, M3PO4. Most sta- ble are the alkaline M2HPO4 or of heavy metals M3PO4. By ignition, M3PO4 unchanged. B3’ ignition, M2HPO4 changed to pyrophosphate M4P2O7. By ignition, MH2PO4 changed to metaphosphate MPO3. Pyrophosphoric Acid H4P2O7— Prepared by heating H3PO4 or adding H2O to P2O5 as above. Gradually absorbs H2O forming ortho- acid. H4P2O7 is a tetrabasic acid but acts as dibasic, forming M4P2O7 and M2H2P2O7. All normal salts are soluble save those of alkalies. All ortho- salts change with heat to pyro- salts. Metaphosphoric Acid HPO3 — Ortho- acid slightly ignited gives pyro- and strongly gives meta- — the melted mass solidifies to “glacial” meta- acid — this dissolves in H^Oand acts like P2O5 in H2O— solution coagulates albumen and gives white precipitate with AgN03 or BaCl2 — it is gradually changed to ortho-. Phosphorous Acid H3PO3 — 1) . PC1 3 + 3H 2 0 = H 3 PO 3 + 3HC1. 2 ) Slow oxidation of P in moist air. P 2 O 3 + 3H 2 0 = 2 H 3 PO 3 . Strong reducing agent, oxidizing to H3PO4. Forms two salts, MH2PO3 and M2HPO3. Hypophosphorous Acid H3PO2— Anhydrite P2O has not been obtained— free acid formed by treating Ba salt with H2SO4 — Ba salt formed in same reaction as for PH 3 . 3Ba(0H ) 2 + 8 P + 6H 2 0 = 3 Ba(H 2 P0 2 ) f , + 2PH 3 . White crystalline mass — a monobasic acid — Ca salt used in medicine. ARSENIC— As. At. Wt. 75. Yal. I, III or V. History — Element long known. Occurrence — Widely distributed as native As and compounds, especially sulphides AS2S2 realgar, AS2S3 orpiment and mispickel FeAsS— As known in commerce as “cobalt” or “fly stone.” 30 Properties — Two allotropic modification — amorphous black powder, and the common crystalline form— a brittle, steel grey solid N with metallic lustre— vapor is yellow with garlic odor— burns with blue flame to AS2O3 — metallic As is not poisonous. Preparation — (1) Heating mispickel in clay cylinders. \ ( 2 ) Heating AS2O3 with charcoal and resubliming the pro- duct. Use- Pigments— flypaper — also forms alloys as with Pb to make shot. HYDROGEN ARSENIDE, ARSINE— AsH 3 . history — Discovered by Scheele. Properties — Colorless gas — disagreeable odor — violent poison — burns with blue flame to AS2O3 + H2O — with limited supply of air to As + H2O — Decomposed at red heat to H + As which is deposited as a mirror giving a delicate test for As. As in reagents — FeS2 usually contains some arsenic, whenc e it passes into H2SO4, HC 1 , and any reagents made from these acids. Preparation — ( 1 ) By heating the alloy As2Zn3 with an acid. AS2Ztl3 + 3H2SO4 = 2Astt3 + Z11SO4. ( 2 ) Formed when As is treated with nascent H. As + H 3 = AsH 3 . ( 3 ) When compounds of As are treated with organic bodies especially decaying bodies, thus ASH3 is formed by As in wall paper. Arsenic trichloride (ASCI 3 ) — Formed by (a) burning As in Cl. (b) As 2 0 3 6HC1 = 2 AsC 1 3 + 3H 2 0. — Heavy, colorless, fuming liquid— important from its vola- tility as it may cause loss of As in analysis. 31 As — OXIDES — ACIDS AND SULPHIDES. Oxides — The oxides AS2O3 and AS2O5 correspond to those of P and N and are of acid nature — AS2O5 combines directly with water, forming H3ASO4, while AS2O3 does not easily dissolve; hence is known as arsenious acid . Arsenic trioxide (AS2O3) — The common compound of As — commercially known as “arsenic” or “white arsenic” — a white amorphous powder — formed when As burns — slightly soluble in H2O with faint acid reaction — odorless — sweet metallic taste — deadly poison — AS2O3 unites with bases, forming arsenites — H2S precipitates AS2S3, which is soluble in (H4N)2S. Arsenic pentoxide (AS2O5) — A colorless deliquescent mass — formed when As2C>3is heated to a red heat— not when As burns— its water solution forms arsenic acid (H3ASO4.) Arsenic Acid (H3ASO4) — White crystalline solid, formed when AS2O3 is digested in HN 0 3 . (ortho-arsenic acid) As 2 0 3 + 2HN0 3 + 2H 2 0 = 2H 3 As0 4 + N2O3. If heat is applied the pyro- and metarsenic acids are formed corresponding in formula to the phosphoric series — Cu3(As(>4)2 is blue — Ag3As04 brown — with H2S arsenic acid is generally first reduced to AS2O3 then precipitated as AS2S3. D 2 H 3 ASO 4 + 2H 2 S = AS2O3 + 5H 2 0 + 2S. As 2 0 3 + 3H 2 S = As 2 S 3 + 3H 2 0. ARSENIC SULPHIDE. Chief is Orpiment AS2S3 — volatile without decomposition — heated with reducing agent to As — soluble in alkaline sulphides, forming salts — As 2 S 3 + 3(H 4 N) 2 S = 2(H 4 N) 3 AsS 8 . A less important sulphide is Realgar (AS2S2). ANTIMONY— Sb. At. Wt. 120. Val. I. III. or V. History — One of old metals — called Stibium by Pliny. Occurrence — Occurs native and in combination — common ore is stibnite (Sb 2 S 3 ). Preparation — (1) Heat stibnite with Fe. Sb 2 S 3 + 3Fe = Sb 2 + 3FeS. (2) Heat stibnite with Na 2 C0 3 + C. 2Sb 2 Ss + 6Na 2 C0 3 + 3C = 4Sb + 6Na 2 S + 9C0 2 . Properties — Silver white metal, highly crystalline, hard and brittle — volatizes at red heat — expands on cooling — oxidizes slowly in moist air, rapidly in HNO3 to Sb 2 0 3 — Precipitates from its (tartaric acid) solution black by Zn — soluble instrong HC1 and H 2 (C4H40e) — insoluble in HNO3 — Sb compounds are poisons. Used— reason of its expansion on cooling Sb is used in tj r pe metal. Stibine (H 3 Sb)^ Analogous to H3N and H 3 As — discovered by Marsh process — formed under same conditions as AsH 3 , and has same proper- ties — chiefly important in detection of Sb. ANTIMONY HALOIDS. Sb forms two compounds with Cl and F — SbCls, SbCls— SI3F3, SbF 5 — with Br and I it forms SbBr 3 and SM3. Antimony trichloride (SbCls)— A soft colorless body— commercially called “Butter of Anti- mony”— prepared: (1) By heating in Cl gas (note PCI 3 ). 2Sb + 3C1 2 = 2SbCl 3 . •(2) By dissolving Sb in HC1. 2Sb + 6HC1 = 2SbCl 3 + 3H 2 . 33 SbCls is decomposed by H2O with precipitation of the oxy- chloride or basic chloride SbOCl. SbCl 3 + H 2 0 = SbOCl + 2HC1. This white precipitate is a characteristic reaction for anti- mony— SbOCl is called “Powder of Algaroth.” Antimony pentachloride (SbCls) — Prepared by action of chlorine in excess on SbCl3, and puri- fied by distilling in current of Cl — SbCls is a fuming colorless liquid, decomposed by water to Sb02Cl Other haloid com- pounds of Sb are analogous to those of P and As. OXIDES OF Sb. Three compounds— Sb203, Sb204, Sb 2 05. Sb 2 0 3 acts as base with strong acids. Sb 2 0 4 base with strong acid — acid with strong base. Sb 2 0 5 acts as acid. Antimony trioxide (Sb 203 ) — Found free in prismatic crystals. Prepared: ( 1 ) Heating Sb in the air — Sb203 sublimes as white cr} r stals. ( 2 ) Oxidizing Sb with HNO3. ( 3 ) “ “ “ KN 0 3 . Sb 2 0 3 is isomorphous with AS2O3 — acts as a base — soluble in strong HC 1 , forming SbCl3, which with H2O gives SbOCl. — Antimonic anhydride Sb 2 05 is a pale yellow powder which re- duces to Sb 2 03 with heat. Antimonic Acid H3Sb0 4 — -f Analogous to H3ASO4. and H3PO4, — derived from Sb 2 05. The Na salt is insoluble in H 2 0 and of importance in separating As from Sb. Antimony Sulphides — Two known sulphides, Sb 2 S3 and Sb 2 Ss, analogous to sul- phides of As. The trisulphide,, stibnite, Sb 2 S3 is most important ore of antimony — separated from impurities by melting — crude Sb 2 S3 called “antimony” oxidizes with heattoSb 2 03and Sb 2 0 4 . Antimony pentasulphite Sb 2 Ss is an orange colored powder formed by passing H 2 S into antimonic acid. 34 BISMUTH— Bi. At. Wt. 210. Val. I. III. or V. Occurrence — Mostly as free Bi in Saxony also as Bi203 and B^S^ — com- paratively rare substance. Properties — Brittle, hard, faint red- white color, with metallic lustre — not easily oxidized but at red heat con verted to BLO3 — not attacked by dilute H2SO4 or HC 1 , but easily dissolved by HNO3— melting point lowered by alloying with other substances, forming fusible metals, as Woods metal which melts in hot water — Bi does not combine with H — unites with Cl and Bi, in two proportions — with I in one. BISMUTH CHLORIDES. Bi forms two chlorides, BLCI4 and BiCL. Bismuthous Chloride (BLCli) is a black, water-absorbing powder, formed by heating Bi and Hg2Cl2- 2Hg 2 Cl 2 + 2Bi = 2Hg 2 + Bi 2 Cl 4 . Bismuth Chloride (Bids) is volatile like SbCls and formed in same way, also by dissolving the oxide in HC 1 — a soft unsta- ble mass, decomposed by H2O to BiOCl. Bismuth oxide (B^Os) — Formed by ( 1 ) Heating Bi in the air. ( 2 ) Heating Bi0N03. A yellow powder — forms salts with acids which are color- less when the acid is colorless — All salts decomposed by much Hb20 — chief compound is the basic nitrate BiONOs , bismuth sub-nitrate, or bismuthyl nitrate, which is much used in medi- cine, and as cosmetic. Often contains a dangerous amount of As. BORON— B. At. Wt. 11. Val. III. History — Sodium salt long known as “Tinkal.” Occurrence — Not free— combined as boracic acid H3BO3, and borax Na2 B4O7. Preparation — Treat Boric anhydride with Na. B2O3 + 6Na = 3Na20 + B2 (amorphous). Properties — Two varieties, amorphous and crystalline. Amorphous is a brown powder — odorless, tasteless, insoluble in H2O, non-conductor of electricity — on heating combines di- rectly with many substances asS, Cl, Br — Boron burns in air to B2O3 and N. Crystalline variety obtained by igniting B with Cl — thus formed is hard as the diamond. BORACIC ACID— H3BO3. Occurrence — In volcanic districts, of Tuscany certain hot springs rise, called “fumaroles” — these carry H3BO3 mechanically. Steam from the fumaroles is passed into water and 1% solution of H3BO3 obtained— this is evaporated by heat of the springs and purified through its sodium salt. Properties — Acid and its salts form, on heating, a colorless glass bead, which dissolves oxides with characteristic colors. H3BO3 itself forms no salts but on heating it loses water, condensing to other acids which form salts. At 100° H3BO3 = H 2 0 + HB 0 3 . Higher 4H3BO3 == 5H2O + H2B4O7. (Borax = Na2B4.07.) (See Na) Still higher H2B4O7 = H 2 0 + B 2 0 3 . Borates are identified by a green color to the flame. Alkali borates only are soluble in water. Boron trioxide (B2O3) — The only oxide of Boron — called Boracic anhydride. 3H2O + B2O3 =2H3BC>3. B2O3 is the source of all B compounds, and is prepared by heating B in the air or heating H3BO3. 2H3BO3 = B2O3 + 3H 2 0. Boron Haloids — B unites with all but I — forms the tri-compounds BiCL.etc. — Liquids of low b.p. and easily decomposed by H2O— formed by passing the haloid over heated B or B2O3+C. SILICON— Si. At. Wt. 28. Val. II or IV. History— Isolated by Berzelius in 1823 . Amorphous variety found first. Occurrence — Always combined with oxygen in Silicon dioxide (Si 02 ) as quartz, sand, etc.— All geological formations except chalk con- tain Si. Preparation — Three varieties of Silicon. ( 1 ) Amorphous, made by (a) fusing potassium-fluo-silicate K2SiF6 with K or NA. K 2 SiF 6 + 4K = Si + 6 KF. (b) use Silicon tetra-chloride and K. SiCl 4 + 4K — Si + 4KC1 (Berzelius). ( 2 ) Graphitic, by fusing the amorphous with Al. ( 3 ) Crystalline, by fusing the graphitic. Properties — Si has strong affinity for O — the amorphous variety burns to SiC>2 — Si is soluble in Aluminum (Al) just asC is in iron. Also soluble in HF forming SiF4 + H4. Silicic Hydride (SiFG) — A spontaneously lighting gas — burns to SiCL and HoO— analogous to methane (CH4). Silicon Fluoride (SiF4). Prepared from CaF2, SiC>2 and H2SO4. 2CaF 2 + S 1 O 2 + 2 H 2 SO 4 = 2CaS0 4 + SiF 4 + 2 H s O. Etching of glass depends on this reaction as the Si02 may be in form of glass. Si0 2 + 4HF = SiF 4 + 2H 2 0. Properties — Colorless gas — bad odor— deadly poison — fumes with H2O, forming silicic acid and hydro-fluo-silicic acid. 3SiF 4 + 3H 2 0 = H 2 Si0 3 + 2H 2 SiF 6 . Hydro=Fluo=Silicic Acid (HoSiFo) — Prepared by passing Sip 4 into H 2 O. K salt is insoluble in HoO. Chemical deportment similar to hydrogen-halogen acids (HC1, etc). Silicon Dioxide Si02 — “Silica” — At first supposed to be an earth — called vitreous earth from use in glassmaking — finally shown to be a weak acid. Found both free and combined in nature. Free as quartz (anhydrous), opal (hydrous), amethyst quartz (colored with Mn), agate, flint, etc. In combination as silicates forming most of the “rocks”. By evaporating to dryness with acids, silicates decompose with a residue oFSiOo insoluble in acids ; this constitutes a test * for silica. Silicates are insoluble save the Na or K salt. Silicic Acid H 2 Si0 3 — There are many silicic acids but this is the most common form — Heated it loses water to form SiC> 2 . CARBON— C. At. Wt. 12.' Val. II. or IV. History — First referred to as an element by Lavoisier 1780. He showed that “mephitic air” or “carbonic acid” contained C and O, also that C was the important part of charcoal. Occurrence — Free in three modifications — crystalline as diamond and graphite — amorphous as various kinds of coal — as CO 2 in air and water — as carbonates in all soil, sometimes forming moun- tains. C crystallizes easily as graphite, but not easily as dia- mond. Obtained as amorphous C from organic bodies. Diamond — ♦ The diamond is found mostly in old rock strata chiefly in India, Borneo, Brazil and South Africa. Usually colorless — when crude covered with dull crust — when polished exhioits great lustre— high refraction — extreme hardness — brittle— poor conductor of heat and electricity — not easily attacked by oxi- 38 dizing agents — unchanged by heat alone — burns easily in oxy- hydrogen blowpipe. Carbonado or black diamonds occur in large masses — have hardness but not fire or water. Graphite — Widely distributed as a mineral— called also plumbago and “black lead” — Till 1798 thought to contain Pb — usually con- tains about five % ash (Si0 2 + Fe 2 0 3 ). A black substance — metallic lustre — very friable but exceed- ingly hard — Good conductor of heat and electricity — Crystal- lizes in hexagonal crystals from melted Fe — not affected by O at ordinary temperatures — at high temperature oxidizes to C0 2 — insoluble in all ordinary liquids. Used for lead pencils — lubricants — electroU'ping — as preser- vative from rust — mixed with clay is used for crucibles. CHARCOAL. Preparation — Not easily obtained pure — made by distilling wood in re- torts or by burning wood in small amount of air, getting rid of volatile substances and saving most of the carbon. Properties — Charcoal is brittle because of cell structure of wood which is retained. Physical properties are important — absorbs and condenses great varietj^ of gases and vapors, especially such as dissolve in H 2 0 — action is not mere absorption, but oxidation of offen- sive gases, hence charcoal is a valuable disinfectant, by hasten- ing destruction of putrescible organic matter — is not an anti- septic or preservative agent — C likewise decolorizes, being es- pecially used in purification of sugar, and in chemical and phar- maceutical preparations. Bone Black — Prepared by destructive distillation of bones — particularly good for all uses of charcoal because of large surface. Coke— By destructive distillation of soft coal — chemically classed between graphite and charcoal. — 39 - Lam p=black — The kind of charcoal, from burning oils in small amount of air. Gas Carbon- Carbon of gas retorts is formed by long heating of the crust which condenses on the interior of gas retorts — very hard, com- pact and dense— conducts heat and electricity — employed in galvanic batteries and as pencils for electric lamps. Chemically classed between charcoal and graphite. Coal- Chief varieties are Anthracite, Bituminous and Peat. The amount of carbon decreases while water and S increase in the given order. CARBON AND HYDROGEN. Form an immense number of compounds known as hydro- carbons. One is formed by direct union of elements — C2H2, acetylene — formed with C electrodes in atmosphere of H. METHANE— CH 4 (MARSH GAS). Occurrence — Found where vegetable matter decays under water — hence called Marsh gas — Found also in coal mines where it often causes explosions, hence named “fire damp” — chief constituent of gas wells. Preparation — Distilling NaC 2 Hs 02 with NaOH — NaC 2 H 3 0 2 + NaOH - Na 2 C0 3 + CH 4 . Synthetically by CS2 + H2S + red hot Cu — CS 2 + 2 H 2 S + 8 Cu = CH 4 + 4 Cu 2 S. Properties — Colorless, inodorous, tasteless gas (condensed by 188 at.) — slightly soluble in water — burns with slight luminous flame — mixture with air explodes (“fire damp”), forming H2O + CO2, “choke damp.” CARBON MONOXIDE— CO. Properties — A colorless gas, inodorous, insoluble in H2O — burns with — 40 - pale blue flame to CO2 — highly poisonous— a powerful reducing agent, forming CO2. Preparation— Always formed by incomplete combustion of C. At first prepared by distilling ZnO with C and supposed to contain H. a) Pass steam over red-hot coal. c + H 2 0 = CO + h 2 . This is “water gas,” one form of illuminating gas — the other or coal gas, formed by destructive distillation of coal, is less poi- sonous. b) CO is formed by decomposition of many organic sub- stances, especially oxalic acid. H2C2O4 = C 0 2 + CO + H 2 0 . CARBON DIOXIDE— C 0 2 . History — Known in earliest times through action of vinegar on chalk —first distinguished as a gas by V. Helmont, who called it “gas sylvestre” — Lavoiser termed it carbonic acid and showed its composition. Occurrence — Widely distributed in atmosphere and soil — occurs in liquid form in some crystalline substances — generally found united with bases, especially in CaC03. Preparation — By action of acids on carbonates — CaC0 3 + 2HC1 = CaCl 2 + C0 2 + H 2 0. Properties — Colorless gas — acid odor — heavy — incombustible — extin- guishes flame — will not support respiration — liquid form is col- orless — becomes solid on rapid evaporation, and forms with ether a freezing mixture — CO2 is absorbed most by water at low temperature, making carbonic acid waters — its presence in wine and beer is due to fermentation — true carbonic acid is H2CO3, a very unstable substance forming two classes of salts — carbonates of alkalies only are stable, but acid salts of alkalies are decomposed by heat. 2NaHC0 3 = Na 2 C0 3 + H 2 0 + C0 2 . CARBON DISULPHIDE— CS 2 . History — Discovered 1796 by Lampadius and called “sulphur alco- hol.” Thought to contain S, H, C, N, but in 1811 shown to be only C and S. Preparation — Pass vapor of S over red-hot C. Properties — Heavy, colorless, strongly refracting liquid — etherial odor when fresh — volatile, with poisonous vapor, having cumulative effect — easily ignited and burns to C0 2 + S0 2 — vapor bums with NO — mixed with air its vapor explodes — remarkable anti- septic power — Decomposed by heat to its elements — CS 2 forms salts of thiocarbonic acid H 2 CS 3 . CS 2 used chiefly as solvent especially of S and P, also of Br and I. CYANOGEN— C 2 N 2 . H istory — Discovered in 1815 by Gay-Lussac while investigating “prussic acid” — called it cyanogen from prussian blue. Preparation — C and N unite only with difficulty — the compounds are ob- tained indirectly and all contain the group CN, which was the first “radicle” discovered. 1) C 2 N 2 formed when organic bodies are heated with K or Na — detection of N in organic bodies depends on this. 2) Heat Hg(CN) 2 = Hg + C 2 N 2 . Properties — Colorless gas — strong odor — intensely poisonous — burns with red flame to C0 2 + N ; CN is recognized by its formation of Prussian blue — when in free state the radicle (CN) exists as C 2 N 2 Dicyanogen, which closely resembles the Haloids. HYDROCYANIC ACID-HCN Occurs in nature in kernels of almonds, peaches, plums, etc.; blossoms and leaves of the peach tree, and several other trees • 4.2 and shrubs— a liquid similar in properties to HC1, but intensely poisonous — generally prepared by treating K salt with HC1, KCN + HC1 = HCN + KC1. ACETIC ACID— H(C 2 H 3 0 2 ) Pyroligneous acid a bi-product in making charcoal — clear, colorless liquid — characteristic taste and odor. Pure acid is called “glacial” — dilute, impure form is known as vinegar — H(C 2 H 3 0 2 ) is a monobasic acid and forms acetates M(C 2 H 3 0 2 ). OXALIC ACID— H 2 C 2 0 4 . Crystalline white solid — poisonous — good reducing agent. H 2 C 2 0 4 is a dibasic acid formed by action of HN0 3 on many organic substances, especially the sugars, starch, etc. — whole- sale by action of KOH on sawdust — occurs widely distributed, especially in plants of oxalis variety and in urinary calculi. CHROMIUM— Cr. At. Wt. 52.4. Yal. II or III. History — In 1797 was discovered by Vanquelin in mineral “crocoisite” (PbCr0 4 ). Occurrence — Not widely distributed — chief ore is “chromite” or chrome iron ore (FeO * Cr 2 0 3 ). Extraction — Ore is pulverized and heated with K 2 C0 3 and CaO — K 2 Cr0 4 potassium chromate is formed and extracted with water — H 2 S0 4 is added to form K 2 Cr 2 0 7 potassium bichromate, which is the commercial salt. The metallic Cr is obtained by heating Cr 2 0 3 (chromic oxide) with C. Properties — A hard, difficultly fusible metal — similar to iron — with heat it slowly oxidizes to Cr 2 0 3 — burns in oxygen with bright light — soluble in HC1 and H 2 S0 4 — not altered by HN0 3 — forms chromous and chromic compounds — Cr forms an immense num- ber of compounds, in some it is acidic, in some basic, or its acid and basic oxides may unite, forming a salt of itself. 43 CHROMIUM HALOIDS— CrCl 2 , Cr 2 Cl 6 , CrF 6 . Chromous chloride — A white powder soluble in water — formed by reducing Cr 2 Cl 6 with H — chromous compounds are hard to obtain and easily oxidize to chromic. Chromic Chloride Cr 2 Cle — A violet sublimate obtained by ignition of Cr 2 03 with C •in a current of Cl — insoluble when pure but easily dissolves if a trace of CrCl 2 is present — the hydrated salt Cr 2 Cl 2 12 H 2 0 forms green deliquescent crystals. Chromic Fluoride CrFe — Chiefly interesting as showing the sexivalent character of Cr. CHROMIUM OXIDES. CrO — Cr 2 03 — Cr03 — first two basic — last acidic. Chromous Oxide CrO — Known only in its Hydrate Cr(OH) 2 , which is made from CrCl 2 + KOH and rapidly changes to chromic oxide. Chromic Oxide Cr 2 03 — Found in nature as “chromite” — formed as a green powder by ignition of ammonium di-chromate (H 4 N) 2 Cr 2 07 = Cr 2 0 + 4H 2 0 + N. The ignited oxide is insoluble in acids — fused with silicates it gives them an emerald green color — used to color glass and porcelain — when freshly precipitated the hydrate is soluble in acids. It is used as a paint, “Gingnet’s green” — like all sesqui- oxides, Cr 2 03 is normally basic, but will not give salts with weak acids. On the contrarv, with strong bases it exhibits an acid character, forming chromites. Chromic anhydride Cr03— Red deliquescent crystals — active oxidizing agent — highly corrosive — made from the K chromate or bichromate by treat- ing with H9SO4 — K 2 Cr0 4 + H 2 S0 4 = Cr0 3 + H 2 0 + K 2 S0 4 The water solution contains the unstable chromic acid, H 2 Cr04, which is analogous to H 2 S04 but known only through its salts the chromates and bichromates — the neutral chromates are gen- 44 erally 3^ellow, the bichromates red — basic Cr 2 03 unites with acidic CrOa to form chromium chromate, Cr 2 Cr04 — The Ba, Pb, A g, and Hg salts of chromic acid are insoluble. Chrome alum Cr 2 'K 2 (S 0 4 )4 + 24H 2 0— The chief salt where Cr acts as a base — formed as violet octohedra, when a solution of K 2 Cr 2 07 -f H 2 S04 is treated with S0 2 , and evaporated below 80 ° — above 80 ° the solution turns green and refuses to crystallize. Chromium sulphides CrS and Cr 2 S3 — Cr forms two sulphides, CrS and Cr 2 S3 corresponding to the oxides but cannot be formed in the wet way. MOLYBDENUM-MO. At. Wt. 96. Val. II, IV, or VI. H istory — Name comes from the Greek for graphite, with which the mineral Molybdenite was confused — M0O3 was separated in 1778 by Scheele and Mo in 1798 by Hjelm. Occurrence — Chief ore is molybdenite MoS 2 . Properties — Hard, silver white metal — very stable — when pure, infusible — after long heating converted to Mo 2 03 — dissolves in HNO3 and aqua regia. COMPOUNDS OF Mo. Chief compounds are the acid H 2 Mo04 and its (NH4) salt (NH 4 ) 2 Mo 0 4 . Jlolybdic Oxide M0O3 — The anhydride of mofybdic acid is M0O3, made by roasting molybdenite — it combines readily with bases, forming molyb- dates. COMPOUNDS OF Mo. Molybdic acid H 2 Mo 04 — Is formed wnen molybdates are treated with dilute HNO3 — Dissolved in strong ammonia it forms (H4N) 2 M04, ammonium molybdate, which in HNO3 solution is used to precipitate H3PO4 in the form of ammonio-phospho-molybdate, a delicate TEST for PHOSPHATES. TUNGSTEN— W. At. Wt. 184. Val. II, IV, or VI. Analogous to Mo — forms similar compounds — Mo com- pounds crystallize easier — Tungstic acid similar to molvbdic acid — common salts are Na and H 4 N tungstates — the former used to make fabrics fire-proof — chief ore is wolframite, FeWC> 4 . Calcium tungstate, CaWC> 4 , is used as a fluorescent screen. TIN-Sn. At. Wt. 118. Val. II or IV. H istory — Used by the Phoenicians, who got it from England. Occurrence — Native in small amounts — chief ore is tin stone SnOo. Extraction — By reducing the ore with carbon and remelting. Properties — Lustrous white metal — crystalline — when bent emits a pe- culiar sound (tin cry) — ductile — malleable, but more so at 100° — brittle at 200° — not oxidized by the air till its melting point — at white heat burns brilliantly to SnC> 2 — forms salts with HC 1 and H2SO4 — forms SnCU with HNO3 or with the alkalies, when it acts as an acid. sodium stannate 2 Sn +-6NaOH = 2Na 3 Sn0 3 + 3H 2 0. Sn is not affected by H 2 S. Use — Chiefly in “tinware” which is iron covered with a layer of tin — also in bronze (Cu, Sn, Zn), solder (Sn, Pb), Brittania metal (Sn, Sb), tin amalgam (Sn, Hg) — the last is used in sil- vering mirrors. TIN CHLORIDES. Stannous chloride SnCL — The anhydrous form is made by HC1 gas on metallic tin — hydrated by treating excess of Sn with HC1 — SnCU comes on the 46 market as tin salt — is a white crystalline substance — soluble in small amount of H2O, in large amount forms the basic chlor- ide — has a strong tendency to unite with Cl, hence used as a re- ducing agent to form SnCU- Stannic Chloride SnCU — A colorless, fuming liquid, hence name of “Spiritus fumans Libavii” — made by action of Cl on Sn or SnCU — combines with metallic chlorides to form double salts, of which “pink salt” SnCU : 2H4NCI is most important — both chlorides of tin are used as mordants — the bromides and iodides are analogous. TIN OXIDES. Stannous Oxide SnO — The basis of stannous salts — formed by heating the hydrate — SnO is a white powder which easily oxides to Sn02. Stannic Oxide Sn02 — “Tinstone” is a crystalline solid formed by heating Sn or SnO — insoluble in acids fused with NaOH or KOH it forms soluble stannates — Sn02 acts also as a weak base forming with H2SO4 stannic sulphate Sn (SO4U ACIDS OF TIN. Stannic Acid H 2 Sn03— Separates as a white precipitate when HC 1 is added to po- tassium stannate (K^SnOs) — dissolves easily in HN03,HC1 and the alkalies — left under water or in vacuo, stannic gradually changes to metastannic acid, an insoluble modification of the same composition. This isomer is also formed when tin is treated with HNO3, concentrated — metastannates are entirely different from stannates. Stannic Phosphate Sn 3 (P 04 ) 2 — Tin forms with phosphoric acid a compound, Sn3(PCU)2> which is insoluble in HNO3 — used in analysis to remove H3PO4. TIN SULPHIDES. Stannous Sulphide SnS — Formed by direct union at high temperature — also precipi- tated from stannous salts by H2S in brown amorphous form. 47 Stannic Sulphide S11S2 — Precipitated as a yellow powder, by H2S, from stannic salts — when sublimed it forms a bright yellow crystalline mass known as mosaic gold which is used in bronzing. STANNOUS SALTS. If acid is colorless the salt is either colorless or yellow — stannous salts have metallic taste — absorb O from the atmos- phere — change easily to stannic — KOH gives white Sn(OH ) 2 soluble in excess. H4N(0H) gives white Sn(OH ) 2 insoluble in excess. AuCL gives purple of Cassius, characteristic. HgCL gives Hg2Cl2, characteristic. TITANIUM Ti— ZIRCONIUM Zr— THORIUM Th. At. Wt. 48. Yal. IV. At. Wt. 91. Val. IV. At. Wt. 232. Val. IV. Three rare metals closely resembling tin but more basic in character and forming NCkous salts — The most important is Ti which generally accompanies iron, especially in Titanic Iron, FeTi03 — metal is obtained by decomposing potassium flno- titanate, K^TiFe, with K (see silicon) — a magnetic dark gray powder — burns in air and chlorine. Ziconium closely follows Ti but has been obtained in crys- talline form, which looks like antimony but is harder — the oxide Zr02 is used for lime light. POTASSIUM— K. At. Wt. 39.— Val. I History— “Potash” from ashes known to ancients — not distinguished from “Soda” till middle of 18 century— metal K isolated by Davy in 1807 . Occurrence — The salts are found in rocks and all cultivated ground — In plants K occurs as oxalate and tartrate — In animals as chlor- ide and phosphate — Sweat on sheeps wool is one source of K compounds — Largely obtained from “Argol” an acid, potas- sium tartrate (HKC4H4O6) which mixed with coloring matter 48 deposits on side of wine casks — Another source is plant ashes which yield a crude K carbonate called potash — At Stassfort, large amounts are obtained from mineral carnallite (KC 1 ). Preparation — When Argol is heated in a closed retort it breaks up to cal- cium and potassium carbonate and carbon — on farther heating metallic K distills as a green vapor — essential reaction is K 2 C0 3 + 2C = 2K-f 3C0 Properties — A soft bluish white metal — green vapor — phosphoresces in dark — Of all the elements K has the greatest affinity for oxygen and chlorine at ordinary temperatures — it burns with a violet blue flame, but oxidizes in the air without taking fire — oxidizes on water, and decomposes it. H 2 0 + K = K 2 0 + Ho. K 2 0 + H 2 0 = 2K0H. K and Sodium are used to reduce metals from their oxides. Potassium Oxide K2O — The only important oxide of Potassium is K^O,^ gray solid formed by action of K on KOH — difficult to secure as it eagerly unites with HoO forming KOH. Potassium Hydrate KOH — Formed when K acts on H2O. Generally prepared by treating a salt of K with hydrate of some metal which will form an insoluble salt with the radicle of the K salt. K 2 C0 3 + Ca(OH) 2 = 2K0H + CaC0 3 — insoluble. Crude KOH obtained from wood ashes. KOH is best example of an alkali and is strongest of all bases — it neutralizes acids to form salts — decomposes fats and oils to soaps hence destroys animal tissue— is soluble in V2 its weight of water, forming a crystalline hydrate KOH * 2H2O. KN 0 3 (Saltpeter). Occurrence — Spoken of under HNO3. 4-9 Preparation — Usually by treating Chili saltpeter NaNOs with KCL NaN0 3 + KC1 = KNO s + NaC.l. NaCl being less soluble is removed by evaporation. Properties — Crystallizes in white, anhydrous, striated prisms, with cool- ing taste, and produces cold by its solution — large amounts act as poison — heated yields KNO2 and oxygen, heated higher de- composes to KOH. 2KN0 3 = K 2 0 + 2N0 2 + 0. I< 2 + H 2 0 = 2K0H. Melted in drops is “sal prunelle.” KNO3 oxidizes most substances save a few metals — with C it deflagrates. 4 KNO 3 + 5C = 2K 2 C0 3 + 3C0 2 + 2N 2 . Use — Used in medicine and preparation of HNO3. Chief use is in making gunpowder. Used instead of NaNOa because KNO3 is not deliquescent. The theoretical equation in discharge of gunpowder is — 2KN0 3 + S + C 3 te K 2 S + N 2 + 3C0 2 . Theoretical proportions are — , KNOs s c 74.8 11.8 13.4 parts. POTASSIUM HALOIDS. 2 Potassium chloride KC 1 — Called digestive salt, also sal febrifugium sylvii — found in sea- water and springs — occurs as sylvite and earn alii te at Stassfurt — prepared from K2CO3 and HC 1 — K 2 C0 3 + 2 HC1 = 2 KC1 + C0 2 + H 2 0. Crystallizes in colorless cubes — easily soluble in H2O — solution absorbs heat — used largety in making alum, and K2CO3 — crude salt used as a fertilizer. Potassium bromide KBr — Colorless cubes — soluble in H2O — made ( 1 ) directly from elements. ( 2 ) By action of Br on KOH — 6 KOH + 6 Br — 5 KBr + KBr 03 + 3 H 2 0. KB 1 O 3 = KBr + 0. (Ferrous Bromide) (Ferrous Carbonate) Usually (3) FcBr 2 + K 2 C0 3 = 2 KBr + FeC0 3 . Used in medicine. 0 Potassium Iodide KI — White crystalline solid, used in medicine and photography —usually made by action of K2CO3 on HI or Fel2~ (1) K 2 CO 3 + 2HI = 2KI + H 2 0 + C0 2 . ( 2 ) K 2 CO 3 + Fel 2 - 2 KI + FeO + C0 2 . Potassium Chlorate KCIO3 — White crystalline solid made by treating KOH with Cl — 6 KOH + 6 Cl - 5 KC1 + KC10 3 + 3 H 2 0. Substance treated under chlorine. Potassium Perchlorate KCIO4 — White solid formed by decomposition of KCIO3 — 2 KCIO 3 = KCIO 4 + KCI + 0 2 . Notable for its slight solubility. Potassium Sulphide K2S — Potassium and sulphur form many compounds all soluble in water — chief is K2S — prepared by fusing K2SO4 with C. K 2 SO 4 + 2 C = K 2 S + 2C0 2 . When fused is a red mass, but crystallizes from water solu- tions in colorless deliquescent prisms — “Hepar Sulphuris” or “liver of sulphur” is a mixture of the Polysulphides of K (K2S, K 2 S 4 ,K 2 S 5 ) and K2SO4, obtained by fusing K2CO3 + S. Potassium Sulphydrate KHS — White, crystalline, solid, with' alkaline reaction, prepared by action of H2S on KOH. KOH + H 2 S = KHS + H 2 0. Potassium Sulphate K2SO4 — White, crystalline, solid — native as Kainite — prepared from K 2 CO 3 + H 2 SO 4 = K 2 SO 4 + h 2 o = co 2 . Soluble in H 2 0 — insoluble in absolute alcohol. Kainite is used as fertilizer. Potassium Nitrite KNO2 — Formed b\ r reduction of KNO3 by heat. Used as source of HNO2 and in analysis. Potassium Arsenite K3ASO3 — Formed by action of K2CO3 011 AS2O3. A dilute solution used in medicine under name of “Fowler’s solution.” •51 Potassium Pyro=Antimoniate K 2 HoSb0 7 — Made from KN0 3 and Sb. Used in testing for Na. Potassium Carbonate K 2 C0 3 — Deliquescent salt— strongly alkaline — insoluble in alcohol. Prepared by ignition of K salts of organic acids. Large quantities obtained from wood ashes— also from KC 1 as in the LeBlanc process for sodium carbonate, which see. Potassium Hydrogen Carbonate KHCO3— Called “Bicarbonate of potash” and prepared bypassing CO2 through a solution of K2CO3 — K 2 CO 3 + C0 2 + H 2 0 = 2 KHCO 3 . KHCO3 is less soluble than K2CO3 and its solution o-jves a neutral reaction. Potassium Cyanide KCN— A poisonous salt— treated under Cyanogen. Potassium Thiocyanate KCNS— Formed by treating molten KCN with S— KCN + S - KCNS. Used as a reagent for Iron (Fe). SODIUM— Na-I. History — Metal discovered in 1807 by Davy. Occurrence — Not free found in many minerals— occurs especiallv NaCl m sea water, springs and as rock salt, and as NaN0 3 Chih saltpeter. Preparation — Metal prepared as K, but easier, hence cheaper. Properties — Similar to potassium— silver-white metal with purple vapor affinity for Cl, O, etc., only little less than that of K— decom poses H s O with explosion*, and if water is warm burns with a yellow flame-forms with K the only liquid alloy not contain- ing Hg— alloy has same appearance as Ho- -M j > U<3L “>• of melted NaOH finally touching the h 2 o, SODIUM CARBONATE— Na 2 C0 3 - 10 H 2 0. Occurrence — Found in nature in sea-plants just as K 2 C0 3 is found in land plants. Preparation — Formerly obtained from ashes of sea plants — may be gotten by heating a solution of the primary carbonate, HNaC0 3 — 2 HNaCOs = Na 2 C0 3 + 2 C0 2 + 2 H 2 0. Commercially in two ways — ( 1 ) LeBlanc’s process from NaCl. (a) NaCl to Na 2 S 04 by heating with H 2 S 04 — 2 NaCl + H 2 S0 4 = Na 2 S0 4 + 2 HC1. (b) Na 2 S 04 to Na 2 S by heating with C— Na 2 S0 4 + 2 C — Na 2 S -f- 2 C0 2 . (c) Na 2 S to Na 2 C0 3 by heating with CaC0 3 — Na 2 S + CaC0 3 = Na 2 C0 3 + CaS. CaS is a waste product. 2) Ammonia or Solvay process, which is mor£ modern— Mono-am. carbonate. (a) NaCl to HNaC0 3 by treating with HNH 4 C0 3 . NaCl + HNjHf 0 3 = HNaCOs + NH 4 C1. (b) HNaC0 3 to Na 2 C0 3 by heating. 2HNaC0 3 = Na 2 C0 3 + C0 2 + H 2 0. The C0 2 is passed into ammonia forming again the acid ammonium carbonate HNH 4 C 0 3 . The NH 4 CI in (a) is heated with lime or magnesia and NH 3 set free as H^NOH — into ammonia water thus formed, C0 2 from (b) is run forming HNH 4 C 0 3 again. The Solvay process has no troublesome residue like theCaS of Le Blanc. Properties — Na 2 C0 3 is an alkaline salt, efflorescent and very soluble in H 2 0 . Use- Essential in glass and soap making and a reagent in the labor at ory. 53 Sodium Sulphate Na2S04. . IOHoO — ‘‘Glauber’s salt” occurs in nature especially in certain nat- ural w aters, as Carlsbad and Freidrichsbad springs— is a by- product in making HNO3. 2NaNO s + H 2 SO 4 = Na 2 S0 4 + 2HN0 3 . Crystallized in large colorless monoclinic crystals easily soluble in water and easil}^ forming supersaturated solutions. Used as a purgative in medicine and for producingartificial cold in the laboratory. Sodium Thiosulphate Na2S2C>3 — “Hypo” is a colorless crystalline solid easily soluble in H2O — may be made by adding S to boiling solution of Na2SC>3. Na 2 S0 3 + S = Na 2 S 2 0 3 . Chiefly used in photography and in bleaching as “anti-chlor.” Sodium Nitrate NaNC>3 — “Chili saltpeter” is a deliquescent salt (see KNO3) of value as the source of HNO3 and KNO3, also used in coarser grades of gunpow der. Di=sodium Phosphate Na2HP(>4/ I2H2O — White rhombic prisms — soluble in H2O with slight alkaline reaction — formed from Na2C03 + H3PO4 — Na 2 C0 3 + H 3 P0 4 = Na 2 HP0 4 + C0 2 + H 2 0. The crude H3PO4 containing CaSC>4 and H2SO4 can be used because Na2SC>4 will not crystallize together with Na2HPC>4. Acid Sodium Phosphate NaH2P04‘H 2 0— White rhombic prisms easily soluble in H2O, with strong acid reaction — chief cause of acid reaction in urine — made from Na3PC>4 and H3PO4 — Na 3 P0 4 + 2H 3 P0 4 = 3NaH 2 P0 4 . Sodium Arsenite Na3As03 — An uncrystallized substance formed from AS2O3 and NaOH. As 2 0 3 + 6NaOH = 2Na 3 As0 3 + 3H 2 0. Sodium Arseniate Na3As04— Arseniate of soda forms in large crystals — prepared by igni- tion of AS2O3 and NaOH with NaN03. Used in dissecting room as an antiseptic. Sodium Tetraborate NaL>B407 . 10H 2 O— Borax is formed as “tinkal” in India and California, but is mostly made artificially in Tuscany — heated, it loses water — melted, dissolves most of the metallic oxides, hence its use in soldering, bead and blowpipe work. Sodium Oxides — Na forms oxides like K, chief being Nai> 0 , Sodium oxide, which in preparation and properties resembles K2O. Sodium Hydrate NaOH — Closely resembles KOH. Pure is fjrepared from Na + PUO, commercially from Na2COs + Ca(OH)2. Na 2 C0 3 + Ca(OH) 2 = 2NaOH + CaC0 3 . Purified by dissolving in alcohol. Sodium Chloride NaCl — “Common salt” is similar to KC 1 — Crystallizes in cubes without water — reaction of its solution is neutral — temperature makes but slight difference in its solubilitj^ — slightly soluble in alcohol — insoluble in HC 1 con. Sodium Hydrogen Carbonate NaHCOs — Saleratus or bicarbonate of soda is less soluble than Na2C03 and is prepared by running CO2 into a solution of carbonate — Na 2 C0 3 + H 2 0 + C0 2 = 2 HNaC0 3 . Used in medicine and preparation of effervescing drinks. Sodium Sulphide Na 2 S — Most important sulphide is Na2§, sodium sulphide — analo- gous to K2S — has alkaline reaction — formed by reduction of Na2SC>4 with C — Na 2 S0 4 + 4 C = Na 2 S + 4 CO. Occasionally used in analysis. Detection of Na salts — Na salts detected by the yellow sodium flame or by precipi- tating as sodium antimoniate. All Na salts except NaNOs are efflorescent. AMMONIUM— H 4 N. History — The group H4N has not been isolated, as it breaks up to H3N + H — after the discovery ofNa and K attempts were made at H4N, but were unsuccessful — by decomposition of H4N salt with Hg electrode, ammonium forms a voluminous unstable amalgam with Hg. AMMONIUM CHLORIDE H 4 NCe. History — “Sal ammoniac” was first made in Egypt by sublimation of camel’s dung. Occurrence — Small amounts found native in volcanic regions Preparation — Formed from NH 3 + HC 1 = (H 4 N)C 1 . Obtained from the liquors of gas works (see NH3) — the gas> water is treated with H2SO4. 2 H 3 N + H 2 S0 4 = (H 4 N)2S0 4 . ( H4N ) 2SO4 is then sublimed with NaCl. (H 4 N) 2 S0 4 + 2NaCl = 2H 4 NC1 + Na 2 S0 4 . Properties — Colorless, octohedral, crystals, highly soluble in water — caus- ing cold by solution — volatile when heated by its dissolution into NH3 + HC 1 . H4NCI united easily with other salts to form double salts The other haloid salts of H4N are analogous, to H4NCI. Used in medicine and the arts. H 4 N SALTS. Sulphate (H 4 N ) 2 SO4— Is amorphous with K2SO4 and highly soluble in water — its preparation from gas water given above. Nitrate (H 4 N) N 0 3 — Is amorphous with KNO3 and deliquesces in air — decom- posed by heat to N 2 0 and H 2 0 . H 4 N NO 3 = N 2 0 “I - 2H 2 0. Highly soluble in water, forming a cooling mixture. Nitrite (H 4 N)N0 2 — White crystalline solid prepared by action of AgN0 2 on (H 4 N)C1— AgN0 2 + (H 4 N)C1 — AgCl + (H 4 N)N0 2 . Its decomposition by heat yields pure N — (H 4 N)N0 2 = N 2 + 2 H 2 0. Ammonium=Sodium=Phosphate HNaNH 4 P0 4 ' 4H 2 0 — “Microcosmic salt” is formed in stale urine — may be pre- pared from HNa 2 P0 4 + (H 4 N)C1— HNa 2 P0 4 + H 4 NC1 = HNaNH 4 P0 4 + NaCl. Carbonate — White or transparent, hard mass giving off H 3 N to the air — formed by action of CaCC >3 on (H 4 N) 2 S0 4 — (H 4 N) 2 S0 4 + CaC0 3 = CaS0 4 + (H 4 N) 2 C0 3 . Commercial salt contains also the bicarbonate, HNH 4 CO 3 , and the carbamate, NH 2 (NH 4 )C0 2 , which are converted to (H 4 N) 2 C 03 by solution in (H 4 N)0H. Thiocyanate (H 4 N)CNS— Prepared by action of CS 2 on H 3 N. CS 2 + 4H 3 N — H 4 NCNS + (H 4 N) 2 S. Decomposed by heat to thio-urea. po-NHa Lb "NH 2 Sulphydrate (H 4 N)SH — Prepared by running H 2 S into ammonia water. (H 4 N)0H + H 2 S = (H 4 N)HS + h 2 o. H 4 N0H will now change it to (H 4 N) 2 S. (H 4 N)HS + H 4 N0H = (H 4 N) 2 S + h 2 o. The hydrosulphide dissolves S, forming the yellow ammonium sulphide, which contains the polysulphides and is valuable as a reagent — (H 4 N) 2 S decomposes gradually. (H 4 N) 2 S + H 2 0 + 0 = 2(H 4 N)0H + s. While any (H 4 N) 2 S remains, the S is dissolved, but finally pre- cipitates. LITHIUM. History — Discovered in 1017 by Arfoedsen in “Petalite.” Occurrence — In small amounts quite widely distributed — chiefly as a compound silicate in lepidolite or “lithia mica,” found also in some mineral springs, and in ashes of many plants, notably to- bacco and the beet. Preparation — Metal obtained by electroUses of the chloride. Properties — Lightest of all metals (sp. gr. .59) — Oxidizes without melt- ing in H 2 O — Most Li salts are soluble in water and give a pur- ple red color to the flame. Li characterized by the small solu- biIit\ T of Li 3 P 04 also of LiCOs (.75 in 100). Otherwise analo- gous to Na and K. CESIUM— Cs— (sky-blue) . RUBIDIUM— Rb— (dark red) . At. Wt. 132.9. Val. I. At. Wt. 85.4. Val. I. Rare metals named from color of lines they give in spectro- scope, by means of which they were in 1860 discovered by Bun- son and Kirchhoff. Both are perfect analogues of K, and though rare are widely distributed. They frequently accom- pany K in mineral springs and plant ashes, and occur in larger quantities in lepidolite (Rb. 5%) and pollucite (Cs. 30%). METALS OF THE ALKALIES. Symbol: Li. h 4 n. Na. K. At. Wt.: 7 18 23 39.1 MsSbO*: soluble insoluble insoluble soluble MHC 4 .H 4 .Oe: soluble soluble soluble insoluble MoPtCh: soluble insoluble soluble insoluble MNO s : deliquescent deliquescent deliquescent permanent M 2 CO 3 : difficultly sol. soluble efflorescent deliquescent I/3PO4: difficultly sol. insoluble soluble soluble Salts in general: permanent volatile efflorescent deliquescent CALCIUM— Ca. At. Wt. 40. Val. II. History — In 1800 Davy obtained the metal from CaO. Occurrence — Widely distributed in large quantities — asCaO inmost min- erals — as CaCOa in limestone, chalk, marble -as CaSC>4 in gyp- sum, alabaster, etc. — as CaSC>4 and CaCoa in most natural waters — likewise contained as phosphate and carbonate in shells and bones of animals, and all ashes of plants. Preparation — Metal best obtained by electrolyses of CaCU. Properties — Yellow ductile metal — decomposes water — heated in the air Ca burns to CaO. CALCIUM OXIDES. Calcium forms two oxides, CaO and Ca 02 . Ca02 is formed by heating CaC03 in streatn of 0 . The more important is CaO, commonly called “lime” or “quicklime.” Pure CaO is a white, infusible solid, best prepared b}^ heating marble or Iceland spar (CaCOs). heated CaCOs = CaO + C0 2 . When exposed to moist air CaO again takes up water, forming Ca(OH)2, which then unites with CO2, forming the carbonate. By reason of its infusibilitv CaO gives with the compound blow-pipe the intense calcium light. CALCIUM HYDRATE— Ca(OH) 2 . Preparation — “Slacked lime” is formed b\^ * ( 1 ) Adding H2O to CaO. CaO + H 2 0 ==' Ca(OH) 2 . Reaction is accompanied by considerable heat, sufficient to explode powder or char wood. ( 2 ) Treating CaCL with KOH. CaCl 2 + 2KOH = Ca(OH) 2 + 2KC1. Properties — Soft white powder — attracts CO2 of the air to form CaCOa — at red heat the hydrate breaks up to CaO + H2O. Ca(OH)2 is slightly soluble in cold, and less in hot water, whence a sat- tira ted cold solution becomes cloudy when heated. When CaQ is added to water in excess, a portion of the hydrate goes into solution, the remainder is precipitated. The clear supernatant liquid is used as lime water, the bulky precipitace as “milk of lime.” Lime water has an alkaline reaction — unites easity with CO2 — this is used as a test for CO2 in the air — crystallized Ca(0H)2 may be obtained from lime water — the solubility of Ca(0H)2 is increased by presence of sugar and diminished by alkalies. Slacked lime is used largely in ordinary mortar, a mixture of lime, water and sand, in which Ca(0H)2 forms CaC03 from the CO2 of air, and likewise forms with the sand a calcium sili- cate. Hydraulic mortar is produced by igniting limestone with aluminium silicate, the composition being chiefty Ca and A 1 silicates. Ca(0H)2 is cheapest of all bases and hence used wherever possible in preparing other hydrates. CALCIUM CHLORIDE CaCl 2 . Made by treating CaC03 with HC 1 . CaC0 3 + 2HC1 = CaCl 2 + C0 2 + H 2 0. Some Ca(0H)2 must be added to precipitate Fe or Mn present as impurities. — Chloride obtained by evaporation has formula CaCL ‘ 6H2O — is very soluble inH20and is used in freezing mix- tures. — Anhydrous CaCL is a white crystalline solid used to dry gases. Bleaching Powder CaOCL — Treat slaked lime with Cl gas. Ca(OH) 2 + 2C1 2 = CaOCl 2 + CaCl 2 + 2H 2 0. Formula of bleaching powder is undecided, but is approximate- ly CaClOCl (CaOCL) — commercially called “chloride of lime,’ r “chlorinated lime,” “chemick.” Properties — White porous powder — odor of chlorine — CO2 of the air decomposes it, freeing HC 10 — by action of strong acids Cl is set free. -60 Ca(C10) 2 + 2HC1 — CaCl 2 + 2HC10. HC10 + HC1 = Cl 2 + H 2 0. A weak acid as CO 2 sets free HC10 — its strong bleaching and disinfection powers are due to these reactions — in bleaching, an alkaline bath, also an anti-chlor bath should follow. anti-chlor 2Na 2 S 2 0 3 + Cl 4 = 4NaCl + 3S0 2 + S. CALCIUM SULPHATE CaS0 4 2H 2 0. Occurrence — Anhydrous forms occur as “gypsum” of which alabaster is one variet} r — natural waters contain CaS0 4 in solution. Properties — Not easily soluble in hot or cold water — presence of H 4 N or Na salts or acids, markedly increases its solubility. Insoluble in C 2 H 6 O (alcohol). Heated to 100° it loses most of its crystal water, forming plaster of Paris, an amorphous powder — this will again unite with water to form the crystalline variety — the “setting” is due to crystallization, which can be slowed by Na2B 4 0 7 , and hastened by K 2 S0 4 . If heated above 200° it becomes too dense to set. CALCIUM PHOSPHATES. Tertiary or Neutral Ca3(P0 4 )2 — Most important natural occurrence is in bones (see P). May be obtained from CaCl 2 and NasP0 4 . 3CaCl 2 + 2Na 3 P0 4 = Ca 3 (P0 4 ) 2 + 6NaCl. Or from the di-sodic phosphate (Na 2 HP0 4 ) in alkaline solution. 2HNa 2 P0 4 + 3CaCI 2 + 2NH 3 = Ca 3 (P0 4 ) 2 + 4NaCl + 2(H 4 N)C1. When fresh is a gelatinous mass which dries to a white powder highly insoluble in water but easily converted to the soluble forms. Secondary CaHP0 4 — A white crystalline salt formed by treating Na 2 HP0 4 with CaClo. Na 2 HP0 4 + CaCl 2 = CaHP0 4 + 2NaCl. 61 Primary or Acid H 4 Ca(P 04 ) 2 — White crystalline salt — soluble in H2O with acid reaction — prepared by dissolving HCaPC>4 in HNO3 and evaporating. 2HCaP0 4 + 2HN0 3 = CaH 4 (P0 4 ) 2 + Ca(N0 3 ) 2 . Largely used as a fertilizer. Hypophosphite — White crystalline salt — obtained by action of P on Ca(OH)2 — (see preparation of H3P). Calcium Nitrate Ca(N 03 ) 2 ~ A white deliquesent salt of little importance. Calcium Carbonate CaC 03 — Widely distributed as limestone, chalk, marble, etc., purest in marble and Iceland spar — occurs likewise in forms of plant and animal life — insoluble in pure water, it dissolves somewhat in H2CO3, hence found in all natural waters — Heated it readily yields CO2 and CaO and is thechief source of Ca salts and CO2. Calcium Sulphide-^ Ca forms polysulphides of which CaS is the most important — Calcium sulphide is a yellowish mass, formed by heating CaSC>4 with C — chiefly notable from it phosphoresence, hence used in 'preparation of luminous paints. REACTIONS OF CALCIUM SALTS. NaOH and KOH: HiNOH: HzCrCU: H2SO4.: (if 2 iV) 2 p 2 0 4 : Flame: Blowpipe: precipitate Ca salts as Ca(0H)2. gives no precipitate. gives no precipitate (compare Ba). gives no precipitate (compare Ba and Sr). precipitates in alkaline solutions. P insoluble in H(C2Hs02) — soluble in mineral acids, reaction is a red color when salt is moistened with HC1. Ca salts give a luminous residue. STRONTIUM— Sr. At. Wt. 87.5. Val. II. The earth Strontia discovered in 1793 — the metal by Davy in 1808 b}^ electrolysis of the hydrate. A rare metal found as the carbonate and sulphate. Properties — In general analogous to Ca. Strontium Oxide SrO — Prepared like CaO from the carbonate. Strontium hydrate Sr(OH )2 — A strong base resembling Ca(OH) 2 but more soluble in water. Forms salts like Ca(OH) 2 . REACTIONS OF Sr SALTS. NaOH and KOH: gives precipititate of Sr(OH) 2 only in concentrated solution. H 2 SO 4 : or soluble sulphates, precipitate SrSCb ATa 2 C0 3 : precipitates SrC 03 . Na 2 HPO±: precipitates SrHP 04 . The only important salt is the nitrate, much used in fire- works. BARIUM. At. Wt. 137. Val. II. H istory — Metal obtained by Davy 1808. Occurrence — Chiefly as the carbonate (witherite) and the sulphate (bar- ite). Preparation — Best obtained by electrolysis of BaCl 2 . C will not reduce the metal from BaCOs, as it does with the alkali carbonates. Properties — Analogous to Sr and Ca, but a stronger base. Barium Oxide BaO — Prepared (1) by igniting Ba(NOs) 2 . Ba(N0 3 ) 2 = BaO + 2N0 2 + O'. (2) Heating BaCOs with C. BaCOs + c = BaO + 2C0. Properties are similar to those of CaO. Barium Peroxide Ba0 2 — Made by heating BaO in stream of O. BaO + O — Ba0 2 . If strongly heated, Ba0 2 yields its O. Ba0 2 = BaO H - 0. 63 This is one source of commercial 0. Ba02 is used also in preparing Ozone and Hydrogen peroxide. Barium Hydrate Ba(0H)2~r BaO + H 2 0 = Ba(OH ) 2 + heat. Ba(OH )2 crystallizes with 8 mols. of water, Ba(0H)2'8H20 — with water forms an aquous solution called “Baryta water” which is a strong base and like Ca(0H)2 rapidly absorbs CO 2 from air — is considered a somewhat better reagent for carbonic acid. REACTIONS OF Ba SALTS. Barium salts similar to Ca salts but more poisonous. NaOH : H^NOH : H 2 SO 4 , : K 2 CrO 4 : Na 2 C0 3 : precipitates Ba(OH ) 2 only in strong solutions, gives no precipitate. precipitates the most dilute solutions, j Differ from Ca and Sr precipitates insoluble BaC 03 . The sulphides of Ba, Sr, and Ca are all phosphorescent. Ba(N 0 s )2 gives a green fire in pyrotechnics. METALS OF THE ALKALINE EARTHS. Symbol: Ca. Sr. Ba. At. Wt: 40 87.5 137 M(OH) 2 : insoluble medium soluble M 2 C0 3 : insoluble insoluble insoluble MCh: soluble soluble soluble soluble in C 2 HrO soluble in C 2 HeO insoluble in C 2 HeO deliquescent permanent permanent M(N0 3 ) 2 : deliquescent permanent permanent M 2 SO±: 1 : 400 1 : 6895 1 : 685,000 M 2 CrO±: soluble soluble insoluble MAGNESIUM— Mg. At. Wt. 24. Val. II. History — In 1750 Black distinguished magnesia, MgO, from lime, CaO. I 11 1808 Davy showed that it was the oxide of a metal. Later Wohler obtained the pure metal. Occurrence — Compounds less abundant than those of Ca but as widely distributed — contained in most soils and all plant ashes— in / 64 large quantity in Hornblende and meerschaum (silicates) and dolomite (carbonate of Mg and Ca). Preparation — Metallic Mg prepared by electrolysis of the chloride heated to fusion, or heating the chloride with Na. Properties — Lustrous silver white metal — medium hard, ductile, and mal- leable. At ordinary temperature Mg gradually oxidizes on the surface— ignited burns with an intense light. Mg is easily at- tacked by dilute acids— slightly by the haloids — does not de- compose H2O, as Mg(0H)2 is insoluble. Uses — Metallic Mg is made chiefly to burn , and comes next to lime in strength of light — used largely for “flashlights” in photogra- phy — forms many alloys, the one with Zn is often used in place of pure Mg as it burns with an equally bright light. Magnesium Oxide MgO— “Magnesia”— A white, very light powder, prepared by heating magnesia alba (MgC03* Mg(0H)2). Commercial magnesia contains traces of Si02, FeO and CO2. — MgO forms with H2O an insol- uble hydrate Mg(0H)2. Magnesium Hydrate Mg(0H)2-^ Formed as above, or by treating any Mg salt with NaOH. Mg(0H)2 is a white amorphous powder, attracting CO2 from the air — a strong base, almost insoluble in water, but sol- ' uble in (H4N)C1 — is used withFe2(S04)3as antidote for arsenic. riagnesium Chloride MgCl 2 — A deliquescent salt, occurring in sea water and mineral springs — prepared from MgC03 and HC 1 . MgCOs + 2HC1 =* MgCl 2 + C0 2 + H 2 0. Heated, the deliquescent crystals yield water and break as below. MgCl 2 + H 2 0 = MgO + 2HC1. If H4NCI should be present we would obtain anhydrous MgCl 2 — The commercial salt occurs as a bi-product in making KC 1 and is much used in the arts. 65 Magnesium Sulphate MgSCLThLO — “Epsom salts” occur in sea-water and mineral springs, especially at Stassfurt. Pure MgS04 forms in rhombic efflores- cent crystals, but usually contains some MgCl2. MgSC>4 is high- ly soluble in water with bitter saline taste — insoluble in alcohol. Used- In medicine — also to weight cotton cloth— a saturated solu- tion with dextrine, will give a crystalized surface on glass. Magnesium Carbonate MgCC>3 — Occurs in compact masses as magnesite, also with CaC03 as dolomite. The “magnesia alba” of medicine is a basic carbonate formed by treating an Mg salt with NaOH. Some CO2 escapes and a white precipitate falls, which when dried at a low tem- perature has the formula Mg(OH)2' 4MgC03 + 4H2O. REACTIONS OF Mg SALTS. precipitates Mg(OH) 2 . “ “ incompletely; with (H 4 N)C 1 no precip. Mg(0H) 2 -MgC0 3 . “ “ “ incompletely; never with (H 4 N) Cl no precipitate ; compare Ca. precipitate MgHP(> 4 , which with H 4 NCI changes to MgH 4 NP0 4 BERYLLIUM. At. Wt. 9. Val II. A rare metal found native only as the oxide — common form is Beryl (3 Be02Al203 * 6 Si 02 ) — when green it is the emerald — Salts resemble those of Mg. ZINC— Zn. At. Wt. 64.9. Yal. II. History — Brass or bronze was known to the ancients as evidenced in old coins — Zn was first recognized as a peculiar metal by Para- celsus who introduced the name, Zinc. Occurrence — Is comparatively rare — found native only in Australia — chief ores are “Smithsonite” (ZnCC>3) and zinc blende (ZnS)— Qres usually accompanied by Cadmium. Na(OH): (H±N)OH : Na 2 COa\ (IL N) 2 CO s : (H4N) 2 C 2 0±: Na 2 HPO±: G6- Extraction from ores — By roasting in the air and then igniting the resulting oxide with C — ZnO + C = Zn + CO. Zinc is quite volatile and hence the crucibles must be connected with a condenser — As in the case of S the Zn vapor condenses at first to a fine dust called Zinc dust, which contains Zinc, generally cadmium, and all the volatile impurities. — The commercial Zinc which forms after the dust, contains also many impurities especially Pb, Fe, C, sometimes S and Cd, and small amounts of As and Sb — it is purified by repeated distilla- tion. Detection — Most characteristic compound is ZnS— before the blow-pipe it gives with Cu(N0s)2 a green color known as Rinman’s green — the white oxide becoming yellow on ignition, is also charac- teristic. Properties — Blue white, crystalline metal — physical properties vary with temperature — brittle at ordinary temperature — at 100-150° can be welded or drawn — 205° breaks under the hammer — melted and poured into cold water forms granulated Zn — not affected by air till strongly heated when it burns with strong blue flame to ZnO — pure Zn is slightly, commercial Zn is easily, attacked by acids and alkalis. Used- In galvanizing sheet-iron — forming alloys, especially brass — in galvanic batteries — with acids and alkaline hydrates as reducing agent. Zinc Oxide ZnO— White powder, which becomes yellow with heat. Prepared as “Zinc white,’ ’ a pigment made by distilling metallic Zn— pre- pared for medical use by igniting the carbonate — for dental use by igniting the nitrate, the last two methods avoid As — ZnO oc- curs in nature as Zincite. 67 Zinc Hydrate Zn(0H)2 — White powder — formed by treating Zn salt with KOH, and soluble in excess of alkali. ZnCl 2 + 2KOH = Zn(OH) 2 + 2KC1. Zn(OH) 2 + 2K0H = Zn0 2 K 2 = 2H 2 0. The same K zincate formed when metallic Zn is treated with KOH. Zn + 2K0H v= Zn0 2 K 2 + 2H 2 0. This reaction often used in alkaline reduction. Zinc Chloride ZnCl2 — 1) Heating Zn in current of Cl gas. 2) Dissolving Zn in HC1. Zn + 2HC1 =? ZnCl 2 + H 2 . White deliquescent mass — fuses with heat— vaporizes at red heat — at high heat forms ZnO Hr HC1 (compare MgCU) — the ZnO unites with unchanged ZnCl 2 to form basic chloride — a concentrated solution of ZnCU + ZnO hardens, with evolution of heat to ZnOHCl, which is used as a dental filling. ZnCU is used as disinfectant, antiseptic, and wood preservative. ZnBr 2 and Znl 2 are analogous to ZnCl 2 . Zinc Sulphate ZnS 04 * 7 H 2 0 — “White vitriol” is found in Zn mines — prepared (a) by oxi- dation of blende (ZnS). ZnS + 04 = ZnS 04 . (b) by solution of Zn in H2SO4. Zn + H 2 S0 4 = ZnS0 4 + H 2 . By ignition it forms the oxide ZnO. Used in the arts and medicine, in battery fluids, etc. — used also to form colors by igniting with metallic salts— emetic. Zinc Carbonate ZnCOs — Formed in company with the hydrate, as basic Zn carbonate, when a Zn salt is treated with a soluble carbonate. ZnCl 2 + Na 2 C0 3 = ZnC0 3 + 2NaCl. ZnCl 2 + Na 2 C0 3 + H 2 0 = Zn(OH) 2 + 2NaCl + C0 2 . Notice tendency to form hydrates also in Mg. 68 Phosphates — Most important are the basic phosphates — when metaphos- phoric acid is mixed with ZnO it forms a cement, which is a basic metaphosphate. ZnO + HP0 3 = Zn(0H)P0 3 . Sulphide ZnS— Zinc blende in nature is yellow — artificial is white — is readily formed from Zn dust and powdered sulphur. ZnS is soluble in mineral acids and alkalies — a little is formed by action of H 2 S on ZnS(> 4 , but the acid freed soon dissolves it. Zinc Salts— Are colorless if the acid is colorless — normal and soluble salts redden litmus — Zn salts are poisonous, have metallic and a stringent taste and used as emetics. CADMIUM-Cd. At. Wt. 112. Val. II. History — Discovered in 1817 by Hermann and Stromeyer simultane- ously. Occurrence — One of rarer metals — generally accompanies Zn ores — Z 11 blendes often contain from .2 to 3 percent Cd— the rare mineral “Greenockite” is CdS. Preparation — Cd is obtained in the first part of Zn distillate, i. e., zinc dust, and is purified by fractional distillation, as it is more vol- atile than Zn. Properties— A white, brilliant metal — malleable and tenacious, but the smallest trace of Zn makes it brittle. — Cd is not oxidized till at high temperature, when it burns to CdO. — Cd is soluble in H2SO4, HC1 and best in HNO3. — A pure Cd salt is precipitated completely by H 2 S. — An amalgam of Zn + Cd is used as a dental filling. Oxide CdO — The only oxide of Cd is formed by burning the metal or ig- niting the nitrate or carbonate — the hydrates are formed by KOH acting on a Cd salt. G9 Sulphate CdSCU — The sulphate Cd(S04)2 is the most common salt of Cd and is anhydrous. Nitrate Cd(N0 3 )2-4H 2 0— The nitrate Cd(N0s)2'4H20 is ver}' soluble and gives the oxide on ignition. Chloride CdCl 2 — The chloride CdCU is similar to ZnCl 2 but efflorescent. Cadmium Sulphide CdS — The most important Cd salt occurs native a Greenockite — commonly made from H 2 S and a Cd salt — not easily from the elements. — CdS is scarcely attacked by dilute acids, but decom- posed by strong — forms a good yellow pigment — differs from AS2S3 in volatility and solubility in alkalies. CADMIUM SALTS. In general resemble Zn salts. NaOH and KOH: precipitate Cd(0H)2 insol. in excess (note difference from Zn). {H±N)OH: precipitates Cd(0H)2 soluble in excess. Na 2 COs: “ insoluble basic carbonate. Na 2 HP0 4! : “ “ CdHP0 4 . iVa 2 C 2 0 4 : “ “ CdC 2 0 4 . K 2 CrO±: does not precipitate dilute solutions (difference from Zn). H 2 S: precipitates yellow CdS insol. in acids (difference from Zn). BaCOz ’■ precipitates CdC 03 completely (difference from Zn). COPPER— Cu. At. Wt. 67.5. Val. II and I (-ic and -ous salts). History— Known to the ancients before iron and used for arms — name from Cyprus, where the Greeks and Romans got it. Occurrence — Native in large quantities — ores are generally the oxide and sulphide, especially chalcopyrite (CuFeS2). Extraction — Two processes, the “English’ * and Mansfield, the latter used when ores are abundant but poor.— In English process the divided CuFeS2 is first roasted in the air — this converts it par- tially to the oxide— then the ore is ignited with silica fluxes and carbon — the iron reduces to the oxide and is dissolved in the slag 70 — bv repeating the process a blistered mass is obtained which contains much S — this is repeatedly roasted, melted and “poled” and Cu run into iron moulds. The Mansfield process is the same, save that a blast furnace is used and more silicious flux. Properties — Red crystalline metal — malleable — tenacious — weldable — conducts heat and electricity — not attacked by dilute acids save in presence of air — dissolves slightly in HC 1 , more in hot H2SO4. Cu + 2 H 2 SO 4 = CuS0 4 + S0 2 2 H 2 0. best in HNO3. 3Cu + 8 HNO 3 - 3Cu(N0 3 ) 2 + 4H 2 0 + 2N0. In dry air Cu is unaffected — in moist air, it forms CuO — with CO2 is coated with CUCO3 — H 2 S blackens it, forming CuS.— Cu is blue green when melted, and expands on cooling — melted Cu always contains gas which escapes when it solidifies. — This is called “spitting” and prevents casting. Pure Cu is used as wire, roofing, etc. — All casting must be made from alloys, of which brass, german silver and the various bronzes are examples. Copper forms four oxides, CU4O, CU2O, CuO, Cu02, of which CU4O is of least importance. Copper Suboxide CU4O — An olive green powder formed by reducing cupric hydrate. Cuprous Oxide CU2O — In nature as cuprite, and formed when cupric hydrate is re- duced by grape sugar. 2 Cu( 0H) 2 — 0 = Cu 2 0 + 2 H 2 0. Cuprous hydrate Cu2(OH)2 is a yellow precipitate which loses water on boiling to form the red CU2O. CU2O is used to give red color to glass. CUPROUS SALTS are not important and are easily converted to cupric form — they are generally colorless, red or yellow. KOH: precipitates the yellow hydrate Cu 2 (OH) 2 . Na 2 COs:-.. precipitates the yellow carbonate CU 2 CO 3 . (HityOH: dissolves salts colorless, but turns blue by oxidation. Cupric Oxide CuO — Found as the mineral cuprite — may be formed by oxida- tion of Cuor by heating the hydrate or nitrate — remarkable for its ease of reduction. — The hydrate Cu(OH)2 is a strong base which is changed to the oxide even when in water. Copper Peroxide CuC>2 — The peroxide is a yellow brown powder formed by treating Cu(OH) 2 with H 2 0 2 .~ CUPRIC SALTS. Generally blue or green when crystallized — colorless when anhydrous — color the flame blue or green — tartaric acid, sugar and many organic compounds prevent precipitation of Cu(0H)2. NaOH: precipitates blue Cu( 0 H )2 (save as above). H&NOH: .... precipitates blue Cu( 0 H) 2 , this is soluble in excess of reagent and the solution thus formed dissolves cellulose. NazCO^: .... precipitates basic CUCO 3 . H 2 S: precipitates black CuS. K^Fe(CN)e precipitates red Cu 2 Fe(CN) 6 - Zn, Fe } Pb: precipitates metallic Cu from solutions. Cuprous Chloride CU2CI2 — The only cuprous salt of importance— formed from CU2O and HC 1 . Cu 2 0 + 2HC1 = Cu 2 Cl 2 + H 2 0. A white powder, insoluble in water, — absorbs carbon monox- ide CO, hence is used in analysis. Cupric Chloride CuCL - Formed by dissolving the oxide or carbonate in HC 1 . CuC0 3 + 2HC1 = CuCl 3 + H 2 Q + C0 2 . Cupric Nitrate Cu(N0s)2— Formed by action of HNO3 on CUCO3. CuC0 3 + 2HN0 3 = Cu(N0 3 ) 2 + H 2 0 + C0 2 . Cupric Sulphate CuS04‘5H 2 0— Found native in mines, where CuS is oxidized — may be made by dissolving Cu in H2SO4 — CUSO4 is soluble in three parts cold cr one-half part hot water — “Blue Vitriol” is much used in the arts in pigments, copper platirg, galvanic batteries, etc. Copper Arsenite CuHAs0 3 — A greenish yellow precipitate, formed when cupric sulphate is added to potassium arsenite— commonly known as Scheele’s Grefn. “Paris green” is also a copper arsenite, mixed with copper acetate. These colors are very poisonous and give off AsH 3 in the presence of organic matter— colors may .be removed from fabrics by (H 4 N)0H. Cupric Carbonate CuC0 3 — The neutral salt CuC0 3 is not known — when Na2C0 3 is added to Cu salt the basic carbonate is precipitated as either a blue precipitate Cu(0H)2 + 2CuC0 3 called “azurite” or a green Cu(OH) 2 * CuC0 3 called “malachite.” Cupric Ferrocyanide Ct^FeCeNe — When acetic acid and potassium ferrocyanide are added to a Cu salt, a red brown precipitate of Cu 2 FeC 6 N 6 is formed, which constitutes the test for Cu. Cupric Sulphide CuS— A black compound found in nature, or precipitated by H 2 S or (H 4 N) 2 S from solution of Cu salt. Insoluble in dilute acids —slowly changes in moist air to CuS0 4 . ALLOYS OF COPPER. The addition of Sn or Zn to Cu prevents “spitting” and forms an alloy harder than Cu, more durable, and of better color — likewise more sonorous. Ancient bronze = Sn + Cu. — Modern = Cu + Sn + Zn. — German silver = Cu + Zn + Ni. — Brass = Cu + Zn, in different proportions and sometimes contains not more than .5 per cent of Pb, which facilitates working — a varity of brass alloys is produced by varying the amounts of Cu and Zn — the later in- creases hardness and fusibility, but decreases specific gravity and malleability. Aluminum bronze = A1 -f- Cu in varying proportions. In electro-plating, the article is immersed in a CuS0 4 bath —the negative pole of an electric battery is connected with the object— CuS0 4 decomposes and precipitates copper on the object — in electro-typing, a mould of plaster of Paris is covered with graphite, and electro-plated as above. SILYER-Ag. At. Wt. 108. Val. I. Occurrence — Widely distributed — often native — the chief ore is the sulphide Ag2S which occurs sometimes native, but generally with some other sulphides, especially galenite PbS. From argentiferous (silver bearing) lead ores, most commercial silver is obtained. Extraction — Three processes are in use: ( 1 ) (Pattinson Process.) Gal- ena, a mixture of Pb and Ag sulphides, is strongly heated to melting — much of the lead separates out and some is united as PbAg alloy — the pure lead crystallizes first and is removed, leaving a readily fusible alloy of PbAg — by repetitions an alloy rich in Ag is obtained which is then cupelled. 2 ) (Parkes or Zinc process). The molten alloy is treated withZn — the alloy of zinc and silver thus obtained is treated with superheated steam, which oxidizes the Zn, leaving Ag unchanged. 3 ) (Amalgama- tion process ) Galena is roasted with NaCl, forming AgCl. H2O and Fe reduce it to Ag, which is separated by amalga- mating with Hg, and then distilling. Properties — The whitest metal — harder than gold and softer than Cu — very malleable and tenacious — at high temperature forms a light blue vapor — melted Ag absorbs O from the air and con- tracts on cooling — its allo3 T s with Cu do not contract and polish better than Ag — silver is not oxydized at ordinary tem- peratures but O3 changes it to the oxide Ag 20 — Ag is not affected by alkalies or KNO3 — unites with Cl, Br and I at ordi- nary temperature and decomposes H2S forming Ag2S, which is soluble in KCN, — Ag is soluble in HNO3, and H2SO4 cone., but AgNC>3 is insoluble in strong HNO3, hence dilute HNO3 should be used in testing for chlorides. Silver has three oxides, Ag40, Ag 20 and AgO, of which the first is of least importance. — —74 ' Silver Suboxide Ag 40 — A brown powder formed like (JU2O and easily converted to Ag'jO and Ag2. Silver Oxide Ag20— The most important oxide of silver is a dark brown precip- itate formed when KOH is added to AgNC>3— any Ag hydrate formed is easily decomposed, leaving Ag20, which acts as if it were Ag(OH) — a strong base, slightly soluble in water and re- duced by light, heat orH at 100 °. Its soluble salts are poison- ous and have metallic taste. Peroxide AgO- Black crystals with metallic lustre formed by passing ozone over Ag or Ag20. Silver Chloride AgCl — Found in nature as “Horn silver,” and formed when HC 1 is added to a soluble Ag salt — AgCl forms as a white, curdy pre- cipitate insoluble in dilute acids, readily soluble in (H4N)0H,. whence it c^stallizes in large octohedrals — Fused AgCl sodifies- to a horn-like mass, hence the name. Silver Bromide AgBr — Separates from silver salts on addition of HBr or a soluble bromide — a bright yellow precipitate similar to the chloride but less soluble in (H4N)0H. Silver Iodide Agl — Yellow, curdy precipitate formed from the action of HI ora soluble iodide, on silver salts in solution— insoluble in H4NOH. Photography — Certain chemical compounds are sensitive to sunlight or other chemically active rays, (as Mg or P light) — Prominent among these are the Argentic haloids (AgX), which form vio- let changing to black argentous compounds (Ag2X) — This is the basis of photography. AgCl is the most sensitive to light, but does not develop well; hence the “wet” plate is sensitized by a mixture of AgBr and Agl — The “dry” plate with AgBr. The exposed plate is treated with some “developer” like pyro- gallic acid, which reduces the argentous haloid (Ag2Br) formed to metallic silver and dissolves free bromine — Metallic silver unites with Argentic Bromide, forming argentous, which is again reduced by developer — Process continues till sufficient ‘ ‘molecular” silver is formed to give a distinct image on the plate — After sufficient development the plate is transferred to a “fixer” of sodium hyposulphite (“hypo”), which reacts with unchanged silver haloid, freeing Br and forming the soluble Na2S203 . Ag2S203. Silver Sulphate Ag 2 S 04 — Obtained by dissolving Ag in hot H2SO4 — is used in refining of Ag — slightly soluble in H2O. SILVER NITRATE— AgN 0 3 . Preparation — “Lunar caustic” is prepared by dissolving Ag in HNO3. Properties — A white crystalline solid, soluble in water, ether and alcohol — a brittle substance, but less so if accompanied by AgCl, which is generally present in the commercial article — when pure is unaffected by light but in presence of organic matter is reduced to metallic Ag. Silver Fulminate Ag2(CN)202 — Made by dissolving AgNOs in C2H5OH and HNO3 — more explosive than the Hg salt. Cyanide AgCN— AgCN is formed when HCN is added to solution of AgNOs — used in electro-plating. Carbonate Ag 2 C 03 — Grayish white precipitate, formed when Na2C03 is added to a silver salt. Chromate Ag 2 Cr 04 — Ag2Cr04 is precipitated by K2Cr04 from soluble Ag salt — a yellow precipitate, which is an indicator in analysis, where Ag salt is used. Silver Sulphide Ag 2 S — Occurs native as. Argentite, or precipitated by H2S or (H4N)2S from Ag salts — a black precipitate used in silvering mirrors. 76 Alloys — Silver is generally used as an alloy with copper — In coin 90% is Ag, in silverware about 75%, “sterling” silver 92%. SILVER SALTS. NaOH gives precipitate of Ag20. (H^N) OH gives precipitate of AgoO, soluble in excess; (in presence of free acid no precipitate). Na2C03 gives precipitate of 3'ellowish white Ag2C03. NazHPO± gives precipitate of yellowish Ag2HP04- H2S and (//4AO2S give precipitate of black Ag2S. HC 1 gives precipitate of white curdy AgCl [sol. in (H4N)0H, KCN, and Na2S203 — insoluble in acids.] KI gives precipitate Agl [insoluble in (H4N)0H, but soluble in Na2S203-] FeSO± gives precipitate of metallic Ag. Tannic Acid......... gives precipitate of metallic Ag. Zn and Fe gives precipitate of metallic Ag. All Ag salts are anhydrous. MERCURY-Hg. At. Wt. 200. Val. I and II. History — Mercury has been known since the earliest times. Occurrence — Not widely distributed nor in large quantities— comes most- ly from Spain and California — native in small quantity — chief ore in Cinnabar HgS— Rare ores are “Amalgam” AgHg and and “Horn mercury” 2Hg2CL. Extraction — Native Hg is filtered through leather — cinnabar is heated with O. HgS + 2 0 = Hg + S0 2 . Commercial is sometimes pure, but often contains Pb, Sn, Bi and Cu dissolved — Pure Hg can be made from HgCls + Fe = Hg + FeCl 2 . Properties — The only liquid metal— bluish white color, metallic lustre — at 40° crystallizes in octohedra — evaporates at ordinary tem- peratures — Hg resembles AuandPtin its small affinity for 0 and large affinity for Cl— requires high temperature to form the •77 oxide, HgO, and at still greater heat yields it up again (com- pare Ba02) — Ozone and chlorine attack it at normal tempera- ture, but boiling HC 1 is harmless — Hg is easil}’ divided bv fats and chalk. Used in manufacture of thermometers and barometers — tin amalgam for mirrors — extracting gold and silver from ores — useful also in electric connections. Mercurous Oxide Hg20 — Mercury forms two oxides, Hg20 and HgO — Hg20, cor- responding to Ag20, is a heavy black powder formed from Hg 2 (N0 3 )2 and KOH— easily decomposed by light to HgO + O also by heat first to HgO + O, then to Hg +O2. — HC 1 converts it to Itg2Cl2 ; Hg20 is the basis of mercurous salts. MERCUROUS SALTS. KOH : precipitates black Hg 2 0. : precipitates black Hg 2 0 containing NH 3 . Na 2 CO%\ precipitates yellow unstable Hg 2 C 03 (which changes to. the black oxide when heated). H 2 S and (H 4 ,N) 2 S: precipitates black Hg 2 S. KI : precipitates green Hg 2 I 2 . HC1 : precipitates white Hg 2 Cl 2 . SnCl 2 : precipitates white Hg 2 Cl 2 (afterwards black Hg). Cu : precipitates metallic Hg. Mercurous salts are mostly insoluble — have a metallic taste and act less violently than mercuric salts (Hg). riercuric Oxide HgO— Is obtained as the red or yellow allotropic modification ac- cording to preparation — Red oxide by heating Hg(N0 3 )2. Hg(N0 3 ) 2 — HgO + 2NO. Yellow oxide in the wet way from HgCL and NaOH. HgCl 2 + 2 NaOH = HgO + 2NaCl + H 2 0. With heat HgO changes to the black oxide Hg20, but recovers its oxygen on cooling — The yellow and red differ slightly in chemical action ; for example, the yellow is at once converted to the oxalate by H2C2O4, but the red must be long heated — both are strong oxidizing agents and form the basis of mercuric salts. REACTIONS OF MERCURIC SALTS. KOH precipitates the yellow oxide HgO. H^NOH... gives a “white precipitate” in very dilute solutions. KI gives a yellow precipitate of Hgl, which turns red. H 2 S in large amounts gives black HgS. H 2 S in small amounts gives yellow precipitates HgX 2 . HgS.* Cu is covered with a deposit of metallic Hg. Mercurous Chloride Hg 2 Cl2— “Calomel” is prepared in ( 1 ) Dry way Hg + HgCl 2 = Hg 2 Cl 2 (Crystalline). (2) Wet way White powder. Hg 2 (N0 3 ) 2 + 2HC1 = Hg 2 Cl 2 + 2HN0 2 . Hg2Cl2 sublimes without melting — somewhat less volatile than HgCL — insoluble in H2O, ether and alcohol — not easily attacked by acids, but boiled with HC 1 is converted to HgCL. Mercuric Chloride HgCL — “Corrosive sublimate” is formed (1) from HgO and HC 1 — HgO + 2HC1 = HgCl 2 + H 2 0. (2) HgS0 4 + 2NaCl = HgCl 2 + Na 2 S0 4 . HgCL is a highlv volatile substance — soluble in cold water, better in hot — soluble also in alcohol or ether — aqueous solu- tion reacts acid and in light decomposes to Hg2CL — when boiled some of HgCL is volatilized — mercuric chloride is precip- itated by albumenoid substances, hence these are the best anti- dotes. Mercurous Nitrate Hg2(N0s)2 — Made by action of dilute HNO3 on excess of Hg — unless metallic Hg is present it gradually changes to Hg(N0s)2. Mercuric Nitrate Hg(N0s)2 — Solution may be obtained by dissolving Hg, or HgO in ex- cess of strong HNO3 ; acid must be in excess or basic salt will separate — the basic salt is reconverted by boiling in H2O. Mercuric Cyanide Hg(CN)2 — Chief source of (CN)2 gas and is prepared from HgO and HCN. HgO + 2HCN Hg(CN) 2 + H 2 0. *X stands for any non-metal. 79 Mercuric Fulminate HgC2N2(>2 — Analogous to the Ag salt but less explosive — used in percus- sion caps. * Tlercuric Sulphide HgS — Occurs in Cinnabar and Vermillion— Is produced as black amorphous mass, by action of H2S on a mercuric salt — If this black sulphide is heated it sublimes as a red crystalline mass. SUBSTITUTED H 4 N COMPOUNDS. When H4NOH is poured over Hg2Cl2 it blackens with the following reaction. amid o-m ercurous-chloride Hg 2 Cl 2 + 2 NH 3 = NH 2 Hg 2 Cl + H 4 NCI. The new compound is regarded as ammonium chloride in which two hydrogens are replaced by Hg2. This one of several substitu- tion products formed by the entrance of Hg into ammonium compounds. METALS OF THE Cu GROUP. Cu Ag Hg Specific Gravity 8.95 10.57 13.5*9 Melting Point 1054° 954°......... 39° At. Wt 63.4 107.6 199.8 Ag resembles somewhat the alkali metals — group chlorides with symbols MCI2 are soluble in H2O. Chlorides M2CI2 are insoluble. Sulphides are black, and insoluble in acids or alkalies but CuS is slightly soluble in (H 4 N)2S. LEAD-Pb. At. Wt. 208. Val. II. History — One of earliest known metals. Occurrence — Occasionally native — chief ore is Galena PbS, a very com- mon and widely distributed mineral. Extraction — ( 1 ) By heating the sulphide with Fe. PbS + Fe = FeS + Pb. 80 ( 2 ) By roasting until the sulphide is partially converted to PbO and PbS 0 4 . PbS + 30- J‘b0 + S0 2 . And PbS + 0 4 - PbS0 4 . By ignition, the lead now separates. PbS + 2PbO = 3Pb + S0 2 . And PbS + PbS0 4 = 2Pb + 2S0 2 . Pure lead is obtained by heating PbC20 4 with C. Properties — Lustrous, blue white metal — at 1 , 600 ° gives off poisonous vapors — all its salts are poisonous — at red heat is easily oxi- dized — at normal temperature, a thin coat of the oxide forms, which protects the body of metal. — Pb is dissolved by HNO3 but not by HC 1 , nor dilute H2SCL because of insoluble salts formed — From its nitrate, Pb(N0s)2, lead is precipitated in metallic form, by Zn, forming a lead tree. WhenPb is in con- tact with air and water the somewhat soluble hydrate Pb20(0H)2 is formed. This is especially dangerous where pure H2O or rain water stands in lead cisterns and pipes. With hard water the action is very slight. Salts and organic matter modify, CO2 and CaH2( 003)2 and sulphates diminish, chlorides and nitrates and especially nitrogenous organic matter increase, corrosion. Lead in con- tact with wood is rapidly corroded. Lead Suboxide Pb 203 — Lead forms four oxides, Pb20, PbO, Pb203 and Pb02. Pb20 is the gray coating on metallic Pb — it decomposes to Pb + PbO. Lead Oxide PbO — Found occasionally in nature — artificial is of various colors accordingto preparation, i.e., red, yellow and white, due proba- bly to crystallization. By ignition of Pb we get the yellow which becomes red when rubbed — when strongly heated becomes yellowish red, forming “litharge”. — PbO is a strong base resembling BaO and SrO — it absorbs CO2 from the air, forming PbC03. Lead oxides are much used in lead plaster, glazing earthen ware, making glass, etc. 81 Lead Trioxide and Dioxide Pb 2 C> 3 , Pb0 2 — Pb 2 0 3 isof no importance. Pb0 2 is formed on positive pole in the electrolysis of a lead salt — also by action of Cl on PbC0 3 — a dark brown powder, which conducts itself somewhat like Mn0 2 , as for example, when heated with HC1, chlorine gas is formed. iTinium Pb 3 C> 4 — “Red lead” is a mixture of PbO and Pb 2 0 3 — important as a pigment, and made b} T oxidizing Pb. Lead Hydrate Pb 2 0(0H) 2 — When a Pb Salt is treated with KOH it does not form Pb(OH) 2 , but a basic hydrate of approximately Pb 2 0(0H) 2 — In presence of strong bases, Pb 2 0(0H) 2 acts as an acid, hence is soluble in KOH. Lead Chloride PbCl 2 — A white precipitate formed when HC1, or a soluble chlo- ride, is added to a cold solution of Pb salt. Soluble in hot water, but is largely deposited in needle shaped crystals when cool — melted PbCl 2 solidifies to a horn like mass. Bromide PbBr 2 — Iodide Pbl 2 — PbBr 2 is very similar to the chloride — chief characteristic of Pbl 2 is that it crystallizes from its hot water solution, in yellow lustrous laminte. Lead Sulphate Pb(S 04 ) — Found in nature, or prepared by adding a soluble sulphate to a lead salt — almost insoluble in water and dilute H 2 SC> 4 — soluble in HN0 3 and (H 4 N) C 2 H 3 0 2 . Nitrate Pb(N0 3 ) 2 — The most common soluble salt of PbO is formed by dissolv- ing lead or its oxide in HN0 3 . Lead Chromate PbCr 04 — A yellow precipitate, formed when a soluble chromate reacts on a soluble Pb salt — it is the test for either — it is used as a pig- ment and is highly poisonous from both of its ingredients. Sulphide PbS — “Galena” is the chief ore of lead — occurs in cubical crystals with metallic lustre — is precipitated by H 2 S from Pb salt. White Lead — Is one of the oldest pigments— A basic lead carbonate of variable composition, made by passing CO2 through solution of Pb(C2H302)2 (French method), or exposing rolls of Pb to atmosphere of H(C2H 3 02) and C 0 2 — often adulterated with BaSC>4 and CaSO*. LEAD SALTS. NaOH : precipitates Pb20(0H)2, soluble in excess. H&NOH: “ Pb20(0H)2, insoluble in excess. Na 2 C0 2 : “ white PbC03, soluble in excess. H 2 S : “ black PbS. KI: “ yellow PbL, soluble in hot H2O. K 2 CrO 4: “ yellow PbCr 0 4 . HC1: “ PbCb, in strong solutions. H 2 SO±: “ PbS 0 4 , insoluble in H2O or dilute H2SCU— soluble in H2S0 4 concentrated. H^NC 2 H 2 0 2 \ decomposes PbS 0 4 . Fe and Zn : ... separate crystalline Pb from solutions. Blowpipe : .... gives a yellow coating and malleable lead. THALLIUM— Te. At. Wt. 204 . Val. I or III. A spectrum metal found by Crookes in the mud from Swed- ish H2SO4 chambers. Occurs in minute quantities in combina- tion with other metals — chemical character is partly like Na, partly like Pb Oxidizes rapidly in the air, forming Te20 and and Te20 3 — most salts are soluble. ALUMINIUM — Al. At. Wt. 274 . Yal. III. or IV. History — In 1817 the metal was first obtained by Wohler from Al 3 Cl 6 and K. Occurrence — Widely distributed in large quantities as the oxide, in ruby, sapphire, corundum and emery — as silicate, in clay, mica and most crytalline rocks. Preparation — ( 1 ) By fusing the chloride with metallic Na. ( 2 ) Electrolysis of the chloride, AI2CI6. 83 Properties — Silver white, highly malleable, ductile and sonorous — a light, tenacious metal— in bulk, A 1 is stable in ,the air — easily soluble in HC 1 , or boiling H2SO4, not in HNO3. NaOH dissolves it forming an aluminiate. A1 + 3NaOH = Na 3 A10 3 + 3H. Unites directly with S group, P, As, Si, and Haloids. In foil, A 1 burns in the air, and decomposes H2O. At 100 ° A 1 alloys easily with some metals, especially with Cu in aluminium bronze. Metallic A 1 is used in scientific instruments, watch-springs, ornaments, etc. — Aluminium bronze, in watches, spoons, etc. — Aluminium silicate or clay, in pottery, bricks, etc. — Aluminium acetate, as a mordant in dyeing cotton cloth. Al COMPOUNDS. Aluminium Oxide AI2O3 “Alumina” — Found in nature as ruby, sapphire, etc. — prepared by heating the hydrate Al2(OH)e or igniting an alum — The oxide is infus- ible, save in oxy hydrogen flame — if ignited, is insoluble in H2O and acids — before ignition, is soluble in acids, and KOH. Aluminium Hydrate A12 (OH)g — A voluminous, white precipitate, formed when (H4N)0H is added to solution of an Al salt — If any vegetable coloring mat- ter should be present, it would be precipitated with the hydrate forming an insolublecompound, technically known as a “Lake” — Al2(OH)e freshly formed, dissolves in excess of KOH or in acids. After long standing its solubility diminishes — heated, it forms AI2O3. Aluminium Chloride ALCle — Formed when Cl gas is passed over hot Al — the hydrated salt is formed by dissolving Al2(OH)6 in HC 1 — ALCle is a color- less deliquescent salt — forms double salts with other metals — aqueous solution is a disinfectant. Aluminium Sulphate Al2(S04)3 * I8H2O — A white crystalline solid, obtained from solution of Al2(OH) 6 in H2SO4, or industrially, from clay and H2SO4 — easily soluble 84 in water— becomes anhydrous with heat — at high temperature decomposes to AI2O3 + SO4. Alums — Aluminium sulphate combines with the alkali sulphates to form “alums.” The chief alum is potassium aluminium sulphate K2 AI3 (SCLH + 24 Hl> 0 , “common alum.” Alum is a general term, applied to a series of double salts of similar composition and crystalline form. Iron, chromium and man- ganese form similar derivatives by replacing the AI2 — a further series is formed, by replacing K2 with Na2 or (H4N)2, thus we have iron alum. Fe2K2(S04)4 + 24H2O, chrome alum Cr2K2- (S04)4 + 24H2O, and others of similar formula. ALUMINIUM SALTS. The haloids, sulphates, nitrates and acetates are soluble. NaOH or KOH : gives Al2(0H)6 soluble in excess. H±NOH\ gives Al2(0H)6 insoluble in excess. NazCOs : “ “ (H 4 W) 2 S “ “ “ “ “ Al silicates. Occur ranee — Granite contains both Al silicate and alkaline silicates. When it disintegrates, the alkaline silicates are washed away, leaving insoluble aluminum silicate, clay — purest form of clay in Kao- lin — Topaz, Beryl and Lapis Lazuli, are important silicates, es- pecially the latter, of which the powder was the old source of Ultramarine. Ultramarine — Is now made by heating a mixture of clay, Na2C03, S and C. — this gives a colorless compound, turning green — by gently heating with sulphur, it forms the blue variety — variation in heat and ingredients gives a variet\ r of colors. Porcelain — Is essentially an Al glass made from kaolin — true porcelain is translucent — slightly attacked by reagents. Pottery— Is an impure porcelain made from clay — its red color due to FeO. — 85 - GLASS. Properties — A mixture of silicates of the alkalies and other metals — a thin liquid at high temperatures, becoming viscous as it cools. Amorphous, transparent solid, little affected by acids or water. No single silicate has all these properties. The alkali silicates are amorphous and transparent but soluble. Others are in- soluble, but crystalline. A combination has the required prop- erties. There are four classes of glass. I — Bohemian or Hard Glass — K and Ca silicate, which melts at high temperature and re- sists reagents — Used in laboratories. II— French Glass— “Crown Glass”— Na and Ca silicate of blue-green color, harder than Bohe- mian but easier melted. Another name is “crown glass.” III— Bottle Glass — Impure glass, colored by Fe203 — a silicate of the alkalies, Ca, Mg and Al. IV — Lead or Crystal Glass— K and Pb silicates — Pb increases refraction, hence used as “strass” glass to imitate jewels. Used in optics as “flint” glass. Colored Glass — Contains Cu20(red), or Au(ruby), CuO(green), Co(blue), Mn( violet), etc. GALLIUM Ga. INDIUM In. At. Wt. 70 . Val. Ill or IV. At. Wt. 113 . 4 . Val. Ill or IV. Two metals discovered through the spectroscope, and close- ly analogous to aluminium — the aluminium group, Al, Ga, In, are true metals — their oxide M2O3, is a weak base, acting as an acid in presence of strong alkalies — the metals do not decom- pose water — are easily dissolved by haloid acids with evolution of H. -Their valence is III. or IV. AI2O3 Ga2C>3 I112O3 AI2CI6 Ga2Cl6 Iti2 Cle Alums are easily formed from their sulphates. -86 MANGANESE-Mn. At. Wt. 55. Val. II. and IV. Occurrence — Never free — chief one is Pyrolusite, (M11O2) Extraction — By igniting the oxide with C. Properties— Hard, gray, metal — fuses with difficulty — not stable in the air — metal has no technical use — forms an alloy with Fe which is used in Bessemer process. OXIDES OF MANGANESE. MnO, Mn203, Mn304, Mn 02 , Mn207. Manganous Oxide MnO — Formed by reduction of higher oxides or by ignition of the oxalate — reduced by C to the metal — MnO is the basis of com- mon manganese salts— its hydrate Mn(OH)2 is a white precip- itate easily oxidized— like Mg(OH)2 it is not precipitated in H4NCI solution. Manganic Oxide Mn 2 03~ A black powder — made by igniting the other oxides in oxy- gen — its hydroxide, Mn(OH)3 is prepared from Mn(OH)2. Treated with dilute acids, the oxide or hydroxide gives mangan- ous salts. Mn 2 0 3 + H2SO4 — MnS0 4 + Mn0 2 + H 2 0 . M/32O3 occurs as mineral “braunite.” Mangano=Hanganic Oxide Mn 3 04 or MnO * Mn 2 Os— Found as a red-brown powder, by ignition of the carbon- ate — acts like a mixture of MnO + M^Os — is isomorphous with magnetite, Fe304. M3O4 occurs as mineral “hausmanite.” rianganese Dioxide Mn02 — Occurs as “pyrolusite’ ’ — made by igniting manganous nitrate Mn(N03)2~ Mn02 is a black powder, which, heated, yields O — treated with HC 1 , gives off chlorine — its hydroxide is Mn(OH)2, from which the unstable manganites are derived. OTHER Mn COMPOUNDS. Hanganous Ammonium Phosphate M11H4NPO4 — Like the Mg salt is formed when H4NOH and alkaline phos- phate are added to a Mn salt. Manganous Carbonate MnCOs — Precipitated by alkaline carbonates from manganous salts — a white powder which easily oxidizes turning brown — like CaC03, is soluble in natural waters which contain CO2. Manganous Sulphide MnS — Precipitated from manganous solution by the alkaline sul- phides — a flesh colored precipitate turning brown, then green. MnS2 corresponding to Mn02 is likewise found in nature as “hauerite.” Mn ACIDS. Manganic Acid H 2 Mn 04 — The oxide Mn(>3 is not known — The hydrated oxide, man- ganic acid H2MnC>4 breaks up in solution to permanganic acid and Mn02. 3H 2 Mn0 4 = 2HM11O4 + M«0 2 + 2H 2 0 Manganic acid is dibasic, corresponding to H2SO4 or H2C2O4 and the salts of these three acids are isomorphous. Potassium Manganate K^MnCL — The K salt is formed from Mn02+K0H. 3Mn0 2 + 2K0H = K 2 Mn0 4 + Mn 2 0 3 + H 2 0. Manganates are of green color and easily decomposed by acids to permanganates. Permanganic Acid HMnCL — Known both in salts and free state — the free acid is made from its Ba salt. HMnC>4 is a powerful oxidizing agent — de- composes above 40 ° — salts are purple and generally made by decomposition of manganates with acids. Potassium Permanganate KMnCL (K 2 Mn 2 08) — Chief salt of permanganic acid, is formed when CO2 acts on K manganate, till the green color changes to red. 3K 2 Mn0 4 + 2C0 2 = 2KMn0 4 + 2K 2 C0 3 + M«0 2 . KMnC>4 crystallizes in dark red prisms — isomorphous with KCIO4. 88 A powerful oxidizing agent — in oxidizing it is reduced to the colorless manganate, hence the persistence of the permanganate color tells when an oxidation is complete — Na-permanganate is used commercially as a disinfectant under name of Condy’s fluid. OTHER Mn COMPOUNDS. flanganous Chloride MnCL — Red deliquescent crystals — formed when Cl is made from Mn02 and HC 1 . Mn0 2 + 4HC1 = MnCl 2 + Cl 2 + 2H 2 0. Manganous Sulphate MnS 04 — Bright red or pink crystals, formed from H2SO4 and MnO. MnO + H 2 S0 4 = M 11 SO 4 + H 2 0. If higher oxides are used, oxygen will separate. MANGANOUS SALTS. Color — reddish when crystalline — otherwise colorless. NaOH and KOH give white Mn(OH) 2 becoming brown. iYa 2 COs gives “ “ “ “ BaCOs precipitates “ only from MnS0 4 . (H 4 AT) 2 S “ MnS, flesh-colored, turning green. H 2 S “ MnS but slightly, in neutral solution. KNOs and iVa 2 C 03 heated with -ous salts give K 2 Mn0 4 (characteristic). Na< 2 ,B±Oi “ “ “ amvthyst color. IRON— Fe. At. Wt. 56. Val II. and III. History — The use of iron dates back to the early legends, when the source was probably meteoric. Later it was obtained from ores, as the great slag fields of India give witness. Occurrence — The most widely distributed metal — seldom found native save in meteors, where it generally contains some nickel — occurs in the blood and mineral waters — chief ores are magnetite Fe3 O4, hematite Fe203, limonite Fe203 +H20,and siderite FeC03. 89 IRON ORES. Hagnetite Fes04— (Fe203, FeO)— Is tlie richest and best of iron ores — occurs in crystalline and amorphous forms — is reduced with difficulty, but yields excellent iron — magnetite may have Fe 20 s replaced by Mn203, and FeO by ZnO — it is then an ore of zinc called franklinite — or may have its Fe203 replaced by Cr203, and is then chromite. Hematite Fe 203 — One of the most common ores of iron — occurs also in a crys- talline modification known as specular iron ore. Limonite — Or bog iron is a mixture of Fe203 and Fe2(OH)6 — yellow ochre is the clayey variety. Siderite — Known also as spathic iron ore, and when mixed with clay as clay iron stone, is the chief English ore. Extraction — Ores are pulverized and roasted, to drive out water, car- bonic acid, or sulphur, and to convert oxides to ferric oxide Fe 20 s, which easily reduces. Ores are then heated with C and fluxes in the blast furnace. The furnace is charged with alter- nate layers of fuel (charcoal, coke or anthracite), ore and flux — a powerful blast of hot air oxidizes the C to CO2 — this is re- duced by hot C farther up the pile to CO — CO passing through the hot ore reduces Fe203 to a spongy mass of metallic Fe, which, mixed with flux and earthy impurities, settles to hotter part of furnace. Here the iron forms a fusible compound with C, and drops to bottom of furnace to be drawn off— the flux and impurities melt to a liquid “slag” which floats above the molten iron. VARIETIES OF IRON. Pig=Iron or Cast=Iron is the crude iron as obtained from the furnace and may contain carbon, phosphorus, sulphur, silicon, manganese, etc. If cooled rapidly when taken from the furnace most of its carbon remains in combination and it is then known as white cast iron — this leaves no residue when dissolved in acid, as all the C unites 90 with the H set free. If theiron is cooled slowly theC separates as graphite and is so left when iron is dissolved — this variety is gray cast iron. If the ore contains much manganese, this is renuced at same timeand such manganese-containing iron, may combine with a greater auantity of carbon, and is known as spiegel iron. The presence of C, Si and P renders cast iron brittle, hence they must be removed before it can be welded — these substances are oxidized in the puddling furnance, which yields wrought iron containing less than .6 per cent of C. Steel — Is intermediate between cast and wrought iron — may be made by adding C directly to wrought iron as in the Cementa- tion process, which produces very good steel, or in Bessemer process, by decarbonizing cast iron to form wrought iron, then adding enough cast iron to produce the desired percentage of C for steel. Properties — Commercial iron is never pure. — The several varieties differ chiefly by reason of impurities, especially C as given above — purest form is wrought iron, especially in piano wire which con- tains but three per cent, of impurities. Pure Fe may be obtained by igniting the oxide in a current of H — it is then a silver white metal — softer and more malleable than wrought iron, but less tenacious. Iron oxidizes easily in moist air or under water to Fe3C>4 — finely divided it decomposes H2O at 100° — dilute acids dissolve Fe with evolution of H — concentrated HNO3 does not affect it, due probably to thin coat of the oxide — CO2 in water forms FeCOs — Fe unites directly with haloids with S, C, Si, and the metals. OXIDES OF Fe. Iron forms four oxides FeO , Fe^O^, FesO±, and the hypo- thetical FeOz. FeO and Fe203 are basic and form salts — FeO, a strong base, corresponds to CaO — Fe203, a weak base, corre- sponds to AI2O3 — Fe03, like Mn03, exists only in the K salt. Ferrous Oxide FeO — Does not occur free but always as a salt, in nature — difficult to make, as it absorbs 0 — may be formed by reducing Fe203 with H — a black powder. 91 Ferrous Hydrate Fe(OH)2 — The base of ferrous (Fe) salts — a white precipitate thrown out of ferrous salts by action of the alkalies in a current of H. On exposure to air it oxidizes to green, then red-brown color — insoluble in water^-a strong base. FERROUS SALTS— Fe. White when anhydrous — blue-green when hydrated. KOH and NaOH form white Fe(0H)2, easily oxidized. K2CO3 precipitates white FeCOs, easily oxidized. H 2 S “ black FeS, only in dilute neutral solutions. (H4AO2S “ “ FeS, soluble in acids. K^FcCqNq “ white Fe^FeCeNe, turns blue by oxidation. KqFc 2 C\ 2 N \ 2 ' “ blue Fe3Fe2Ci2A f i2 “Turnbulls blue.” KCNS gives no color when no Fe2 salts are present. Na 2 FLFO± precipitates white Fe3(P04)2, turns blue by oxidation. Tannic acid gives no color when no Fe2 salts are present. Na 2 B^07 gives a colorless or yellow bead. Ferrous salts absorb NO in the cold, turning black, and are used as a reagent for HNO3. Ferric Oxide Fe 203 — Occurs in nature as hematite — prepared by ignition of fer- rous sulphate — 2 FeS0 4 = Fe 2 0 3 + S0 3 + S0 2 . Thus prepared Fe203 is a dark red powder, used as a paint under name “Colcothar” — Fe203 is a very hydroscopic sub- stance and like Mn02 aids in preparation of 0 from KCIO3. Ferric Hydrate Fe 2 (OH) 6 — The basis of ferric salts — red brown voluminous precipitate, made by treating a ferric salt with an alkali hydrate — red heat renders it insoluble in acids.. FERRIC SALTS. Color — If anhydrous the neutral (Fe2) salts are without color — when crystallized they are yellow or reddish yellow, save the nitrate and fluoride which are colorless, and the acetate and thiocyanate which are deep red — All ferric salts are reduced to ferrous by H2S. KOII, K 2 COs\ precipitates red brown Fe 2 (OH )6 soluble in acids— this pre- and BnCO$: j cipitation is prevented by sugar. HzS: reduces ferric to ferrous salts, setting free S. Fe 2 Cl 6 + H 2 S — 2 FeCl 2 + 2 HC1 + S. (H 4 N) 2 S: gives FeS. Fe 2 Cl 6 + (H 4 N) 2 S = 2 FeS + 6 H 4 NC1 + S. K^FeCeNe : gives blue Fe 2 FeCeN6, “Prussian blue.” KCNS: gives red Fe 2 (CNS)6, organic acids hinder. Tannic acid:.... gives black iron tannate (ink). Ferric Acid FUFeCU — Corresponds to H2SO4 and H2Mn04 — free acid is not known as it is unstable. Potassium ferrate K 2 Fe 04 is formed when iron filings are fused with KNO3 — (see K^MnCU) — crys- tals of K salt form in dark red prisms. Ferrous Chloride FeCU — The white anhydrous salt is formed when HC1 gas passes over excess of heated iron— green deliquescent FeC^ is formed by the solution of Fe in HC1 — oxidizes easily to Fe 2 Cl 6 — all ferrous salts absorb NO. Ferrous Bromide FeBr 2 — Made by action of Br on excess of Fe in the presence of H 2 O — used in preparation of KBr. Ferrous Sulphate FeS 04 * 7 H 20 — “Green vitriol” is a bi-product of Cu works — obtained pure from FeS and H2SO4. FeS + H 2 S0 4 = FeS0 4 + H 2 S. Crystals effloresce in dry air, and if heated lose six molecules of water to form the colorless salt FeSC^'H^O — at still greater heat FeS 04 decomposes. 2FeS0 4 = Fe 2 0 3 + S0 3 + S0 2 . From this reaction the “fuming” or “Nordhausen” sulphuric acid is obtained. Ammonium Ferrous Sulphate FeSC> 4 ‘ (H 4 N) 2 S 04 . 6 H 20 — A stable salt, prepared bv evaporating together equal parts of FeS0 4 -7H 2 0 and (H 4 N) 2 S 04 . Ferrous Nitrate Fe(N0s)2 — Unstable, green salt, made by action of HNO3 on excess of iron, or from Ba(N 03)2 and FeSC> 4 . Ba(N0 3 ) 2 + FeS0 4 = Fe(N0 3 ) 2 + BaS0 4 . 93 Ferric Chloride Fe2Clc> — Anhydrous, by heating iron in current of Cl — hydrated Fe2- Cle ' 6H2O bypassing Cl gas into solution of FeCb—anhydrous is brown — hydrated, red — both are deliquescent — soluble — vola- tile at high heat — forms basic salts easily — used as a reagent for organic substances. Other (Fe)2 salts are formed by oxidation of the (Fe) fer- rous. CYANOGEN COMPOUNDS OF IRON. Ferrous Cyanide Fe(CN)2 — Is found as a white unstable precipitate when KCN is added to the solution of a ferrous salt — with an excess of KCN the double salt Fe(CN)2, 4 KCN or K^FeCeNe is formed. Potassium Ferrocyanide ^FeCeNe — This is known commercially as yellow prussiate of potash, and made by heating K2CO3 + Fe with nitrogenous bodies to a red heat — a sweetish, non-poisonous salt — used in dyeing and as a test for Fe2 and CN salts — with concentrated HC 1 it forms the acid, hydrogen ferrocyanide, or ferrohydrocyanic acid H4 Fe(CN)6, which is unstable — with ferric salts, ferrocyanide forms prussian blue. Potassium Ferricyanide KeFe2Ci2Ni2 — Is formed when the ferrocyanide is treated with Cl. 2K 4 Fe(CN) 6 + Cl 3 = K 6 Fe 2 (CN)i 2 + 2KC1. Separates in red prisms, known as red prussiate of potash — its water, solution reduces to ferrocyanide by light — its alkali solution is a strong oxidizing agent — with ferrous salts, ferricy- anide forms “Turnbulls blue.” IRON SULPHIDES. Ferrous Sulphide FeS — Gray, metallic mass, made by fusing Fe and S — or merely moistening a mixture of Fe and S at ordinary temperature. — Is obtained pure through precipitation from ferrous salts by (H4N)2S — FeS is soluble in acid and largely used for making H 2 S. FeS + 2HC1 = H 2 S + FeCl 2 . 94 Ferric Sulphide Fe2S3 — Black sulphide analogous to Fe2C>3, made by heating Fe and S in proper proportions. “Pyrite” FeS 2 — A natural sulphide not analogous to any known oxide. NICKEL— Ni. At. Wt. 58.6— Val. II. Occurrence — Generally accompanies Co — chief ore is NiAs“Kupfernickel” — with Fe only in meteorites. Properties — Hard white metal — malleable— weldable — ductile — more te- nacious than iron — magnetic — but slightly affected by the air, hence used to plate other metals — soluble in mcst acids but like Fe, is passive in HNO3 concentrated. Oxides and Salts — Of nickel are closely analogous to those of Co. — Salts are green, complementary to the red of Co salts — the hydrate Ni 02 H2 is somewhat soluble in H4NOH, with a blue color similar to Cu. COBALT-Co. At. Wt. 58.6. Val. II or III. History — The name “kobold” was given first to minerals which col- ored glass blue, afterwards to the metal which was discov- ered i i them. Occurrence — Never in the free state, usually in combination with As and S— chief ores are “smaltite,” C0AS2, and “cobaltite,” CoAs2* C0S2. Extraction — By roasting the ore a crude arseniate is obtained, “zaffre this treated with KHSO4 gives the crude oxide which is reduced with carbon to metallic Co. Properties — A reddish metal — lustrous — malleable— ductile — difficultly fusible — harder and more tenacious than iron — can be magnet- ized— not attacked by air or water — slightly by HClandH2S04 — easily by HNO3. Chief compounds are cobaltous, i. e., where Co is bivalent. COBALT OXIDES. CobaltQus Oxide CoO — Cobalt forms three oxides, CoO, C02O3, C03O4. CoO is a brown powder formed by heating C02O3 with H — is the basis of cobaltus salts — with water it forms Co(OH) 2, Cobaltous hydrate, a rose colored precipitate turning brown — Co (OH) 2 like Fe02H2 must be made in an atmosphere of H. Cobaltic Oxide C02O3 — A black powder formed by gentle ignition of Co(N03)2 — it is basis of cobaltic salts, which are ver} r unstable. Cobaltous=Cobaltic Oxide C03O4 — A black pow der formed by strong ignition of the nitrate — it is not magnetic and forms no salts. OTHER Co COMPOUNDS. Cobaltous Chloride C0CI2 — Obtained by dissolving the oxide or carbonate in HC 1 — it crystallizes with 6H2O to form red crystals, which, when anhy- drous are blue — the red crystals are also turned blue by H2SO4, which takes up its water of crystallization — characters written with the red solution are almost invisible but turn to distinct blue when heated (sympathetic ink). The Br and I compounds are similar. Cobaltous Sulphate C0SO4 7H2O — Crystallizes in dark red prisms, and resembles ferrous sulphate. Cobaltous Nitrate Co(NOs)2 ' 6 H 2 0 — Forms in red deliquescent prisms — is much used in blowpipe work, and gives characteristic blowpipe reaction swith Zn (Rin- maus green) and A1 (Thenard’s blue) also a blue with phos- phates. Cobalt Silicate — No silicates are found in nature, but when glass is fused with a Co salt it forms a dark blue silicate, called “smalt” — when reduced to powder is used as a paint. COBALT CYANIDES. Cobalt Cyanides — Are formed by dissolving the hydrate in KCn — in excess of KCN a double salt is formed, Co(CN)2.2KCn — when this is boiled with an oxidizing agent a cobalticyanide, KeCo 2 (CN )i 2 , is formed and from this KOH does not precipitate Co2(0H)e. COBALT SULPHIDES. Cobalt Sulphides CoS— The most common is CoS, a black precipitate, formed when (H 4 N) 2 Sis added to a Co salt — insoluble in alkalies or cold dilute HC1 — soluble in strong acids. GOLD— Au. At. Wt. 196-7 Val. 1. and III. History — Long known to the ancients and called the king of metals — chief object of alchemists, was the transmutation of base metals to gold. Occurrence — Chiefly as native gold, alloyed with more or less Ag,Cu,Pb and Bi — is widely distributed, but generally occurs in small quantities — its ores are unimportant. Properties — Most malleable and ductile of metals — highlj^ tenacious — at high temperature is volatile — melted has a red yellow color — on cooling, Au contracts more than other metals — no single acid dissolves Au, but with aqua regia (HC 1 + HNO3) it forms the chloride — the other haloids also attack it, and melted KNO3 forms the oxide. 97 Auric Chloride A11CI3 — Formed by solution of Au in Aqua regia — a reddish brown, deliquescent mass of crystals, which dissolve readily in alcohol and ether — All reducing agents are oxidized by AuCU, with precipitates of gold — thus FeSC>4 precipitates gold lrom the chloride as a lustreless brown precipitate — and stannous chlor- ide SnCU, precipitates the oxide AU2O as “purple of Cassius.” Aurous Chloride AuCl. When auric chloride is heated to 180 ° it forms aurous chloride, AuCl, a white powder in H2O— ignited AuCl decom- poses to Au + CL Aurous Oxide AU2O — Formed in “purple of cassius” also when KOH acts on aurous chloride— dark violet powder — decomposed by heat — changed to Xu and AUCI3 by HC 1 . Auric Oxide AU2O3 — When A-UCI3 is heated with MgO a brown precipitate is formed. If excess of MgO in is removed by concentrated HNO3 auric oxide is left as a brown powder. Auric Hydroxide — If MgO is removed with dilute HNO3, Au(OH)3 remains as a yellow-red powder — oxide and hydroxide are insoluble in water and acids, but have acid properties and are soluble in alkalies. The hydroxide is called auric acid and forms aurates, which are derived from meta-auric acid, HAUO2. Auric Sulphide AU2S3 — H2S or soluble sulphides precipitate auric sulphide, as a brown precipitate, soluble in alkaline sulphides. Gold Cyanide AuCN — When gold or its oxide is dissolved in KCN the colorless double salt, AuCN . KCN, may be crystalized out — from this metallic Au is easily precipitated by electrolysis, hence used in gilding. 98 PLATINUM— Pt. At. Wt. 194.4. Val. II. and IV. Occurrence — Found only in metallic state, generally alloyed with pal- ladium, iridium and other rare metals. Extraction — The ore is dissolved in aqua regia, and precipitated by (H4 N)C 1 — this forms by ignition a double salt of iridium bearing platinum in a spongy mass ( platinum sponge.) This is used directly in making platinum vessels— if desired the iridium, which diminishes malleability but increases hardness and re- sistance to reagents, may be farther removed. Properties — Heavy, lustrous, gray-white metal — softer than Ag — not affected by air nor common acids — dissolved by aqua regia and chlorine water — attacked by fusing caustic alkalies — finely divided as in Pt sponge, and Pt black, it remarkably influences the chemical combination of gases. PLATIUM COMPOUNDS. Platinic Chloride PtCL — Is the most important salt, made by dissolving Pt in aqua- regia. It dissolves in water to a red yellow solution, important because it forms with H±N or K an insoluble salt — the oxides and salts of Pt in general are formed, and act in the same way as those of Au. — Their formula is different because Pt acts with valence of two or four, thus: PtCl2, PtCU, PtS2,etc — all Pt salts leave, on ignition, a residue of metallic Pt.