Page:EB1911 - Volume 07.djvu/126

Rh production had fallen off to 25,000 tons. Great Britain, though she had made half the world’s copper in 1830, held second place in 1860, making from native ores 15,968 tons; in 1900 her production was 777 tons, and in 1907, 711 tons. The United States made only 572 tons in 1850, and 12,600 tons in 1870; but she today makes more than 60% of the world’s total. In 1879, Spain was the largest producer, but now ranks third.

The estimated total production for each decade of the 19th century in metric tons is here shown:—

The following table gives the output of various countries and the world’s production for the years 1895, 1900, 1905, 1907:—

As the stock on hand rarely exceeds three months’ demand, and is often little more than a month’s supply, it is evident that consumption has kept close pace with production.

The large demand for copper to be used in sheathing ships ceased on the introduction of iron in shipbuilding because of the difficulty of coating iron with an impervious layer of copper; but the consumption in the manufacture of electric apparatus and for electric conductors has far more than compensated.

Alloys of Copper.—Copper unites with almost all other metals, and a large number of its alloys are of importance in the arts. The principal alloys in which it forms a leading ingredient are brass, bronze, and German or nickel silver; under these several heads their respective applications and qualities will be found.

Compounds of Copper.—Copper probably forms six oxides, viz. Cu4O, Cu3O, Cu2O, CuO, Cu2O3 and CuO2. The most important are cuprous oxide, Cu2O, and cupric oxide, CuO, both of

which give rise to well-defined series of salts. The other oxides do not possess this property, as is also the case of the hydrated oxides Cu3O22H2O and Cu4O35H2O, described by M. Siewert.

Cuprous oxide, Cu2O, occurs in nature as the mineral (q.v.). It may be prepared artificially by heating copper wire to a white heat, and afterwards at a red heat, by the atmospheric oxidation of copper reduced in hydrogen, or by the slow oxidation of the metal under water. It is obtained as a fine red crystalline precipitate by reducing an alkaline copper solution with sugar. When finely divided it is of a fine red colour. It fuses at red heat, and colours glass a ruby-red. The property was known to the ancients and during the middle ages; it was then lost for several centuries, to be rediscovered in about 1827. Cuprous oxide is reduced by hydrogen, carbon monoxide, charcoal, or iron, to the metal; it dissolves in hydrochloric acid forming cuprous chloride, and in other mineral acids to form cupric salts, with the separation of copper. It dissolves in ammonia, forming a colourless solution which rapidly oxidizes and turns blue. A hydrated cuprous oxide, (4Cu2O, H2O), is obtained as a bright yellow powder, when cuprous chloride is treated with potash or soda. It rapidly absorbs oxygen, assuming a blue colour. Cuprous oxide corresponds to the series of cuprous salts, which are mostly white in colour, insoluble in water, and readily oxidized to cupric salts.

Cupric oxide, CuO, occurs in nature as the mineral (q.v.), and can be obtained as a hygroscopic black powder by the gentle ignition of copper nitrate, carbonate or hydroxide; also by heating the hydroxide. It oxidizes carbon compounds to carbon dioxide and water, and therefore finds extensive application in analytical organic chemistry. It is also employed to colour glass, to which it imparts a light green colour. Cupric hydroxide, Cu(OH)2, is obtained as a greenish-blue flocculent precipitate by mixing cold solutions of potash and a cupric salt. This precipitate always contains more or less potash, which cannot be entirely removed by washing. A purer product is obtained by adding ammonium chloride, filtering, and washing with hot water. Several hydrated oxides, e.g. Cu(OH)2·3CuO, Cu(OH)2·6H2O, 6CuO·H2O, have been described. Both the oxide and hydroxide dissolve in ammonia to form a beautiful azure-blue solution (Schweizer’s reagent), which dissolves cellulose, or perhaps, holds it in suspension as water does starch; accordingly, the solution rapidly perforates paper or calico. The salts derived from cupric oxide are generally white when anhydrous, but blue or green when hydrated.

Copper quadrantoxide, Cu4O, is an olive-green powder formed by mixing well-cooled solutions of copper sulphate and alkaline stannous chloride. The trientoxide, Cu3O, is obtained when cupric oxide is heated to 1500°–2000° C. It forms yellowish-red crystals, which scratch glass, and are unaffected by all acids except hydrofluoric; it also dissolves in molten potash. Copper dioxide, CuO2H2O, is obtained as a yellowish-brown powder, by treating cupric hydrate with hydrogen peroxide. When moist, it decomposes at about 6° C., but the dry substance must be heated to about 180°, before decomposition sets in (see L. Moser, Abst. J.C.S., 1907, ii. p. 549).

Cuprous hydride, (CuH)n, was first obtained by Wurtz in 1844, who treated a solution of copper sulphate with hypophosphorous acid, at a temperature not exceeding 70° C. According to E. J. Bartlett and W. H. Merrill, it decomposes when heated, and gives cupric hydride, CuH2, as a reddish-brown spongy mass, which turns to a chocolate colour on exposure. It is a strong reducing agent.

Cuprous fluoride, CuF, is a ruby-red crystalline mass, formed by heating cuprous chloride in an atmosphere of hydrofluoric acid at 1100°–1200° C. It is soluble in boiling hydrochloric acid, but it is not reprecipitated by water, as is the case with cuprous chloride. Cupric fluoride, CuF2, is obtained by dissolving cupric oxide in hydrofluoric acid. The hydrated form, (CuF2, 2H2O, 5HF), is obtained as blue crystals, sparingly soluble in cold water; when heated to 100° C. it gives the compound CuF(OH), which, when heated with ammonium fluoride in a current of carbon dioxide, gives anhydrous copper fluoride as a white powder.

Cuprous chloride, CuCl or Cu2Cl2, was obtained by Robert Boyle by heating copper with mercuric chloride. It is also obtained by burning the metal in chlorine, by heating copper and cupric oxide with hydrochloric acid, or copper and cupric chloride with hydrochloric acid. It dissolves in the excess of acid, and is precipitated as a white crystalline powder on the addition of water. It melts at below red heat to a brown mass, and its vapour density at both red and white heat corresponds to the formula Cu2Cl2. It turns dirty violet on exposure to air and light; in moist air it absorbs oxygen and forms an oxychloride. Its solution in hydrochloric acid readily absorbs carbon monoxide and acetylene; hence it finds application in gas analysis. Its solution in ammonia is at first colourless, but rapidly turns blue, owing to oxidation. This solution absorbs acetylene with the precipitation of red cuprous acetylide, Cu2C2, a very explosive compound. Cupric chloride, CuCl2, is obtained by burning copper in an excess of chlorine, or by heating the hydrated chloride, obtained by dissolving the metal or cupric oxide in an excess of hydrochloric acid. It is a brown deliquescent powder, which rapidly forms the green hydrated salt CuCl2, 2H2O on exposure. The oxychloride Cu3O2Cl2·4H2O is obtained as a pale blue precipitate when potash is added to an excess of cupric chloride. The oxychloride Cu4O3Cl2, 4H2O occurs in nature as the mineral atacamite. It may be artificially prepared by heating salt with ammonium copper sulphate to 100°. Other naturally occurring oxychlorides are botallackite and tallingite. “Brunswick green,” a light green pigment, is obtained from copper sulphate and bleaching powder.

The bromides closely resemble the chlorides and fluorides.

Cuprous iodide, Cu2I2, is obtained as a white powder, which suffers little alteration on exposure, by the direct union of its components or by mixing solutions of cuprous chloride in hydrochloric acid and potassium iodide; or, with liberation of iodine, by adding potassium iodide to a cupric salt. It absorbs ammonia, forming the compound Cu2I2, 4NH3. Cupric iodide is only known in combination, as in CuI2, 4NH3, H2O, which is obtained by exposing Cu2I2, 4NH3 to moist air.

Cuprous sulphide, Cu2S, occurs in nature as the mineral chalcocite or (q.v.), and may be obtained as a black brittle mass by the direct combination of its constituents. (See above, Metallurgy.) Cupric sulphide, CuS, occurs in nature as the mineral covellite. It may be prepared by heating cuprous sulphide with sulphur, or triturating cuprous sulphide with cold strong nitric acid, or as a dark brown precipitate by treating a copper solution with sulphuretted hydrogen. Several polysulphides, e.g. Cu2S5, Cu2S6, Cu4S6, Cu2S3, have been described; they are all unstable, decomposing into cupric sulphide and sulphur. Cuprous sulphite, CuSO3·H2O, is obtained as a brownish-red crystalline powder by treating cuprous hydrate with sulphurous acid. A cuproso-cupric sulphite, Cu2SO3, CuSO3,2H2O, is obtained by mixing solutions of cupric sulphate and acid sodium sulphite.

Cupric sulphate or “Blue Vitriol,” CuSO4, is one of the most important salts of copper. It occurs in cupriferous mine waters and as the minerals chalcanthite or cyanosite, CuSO4·5H2O, and boothite, CuSO4·7H2O. Cupric sulphate is obtained commercially by the