Page:EB1911 - Volume 04.djvu/998

 occurrence in chalk (Latin calx). It does not occur in nature in the free state, but in combination it is widely and abundantly diffused. Thus the sulphate constitutes the minerals anhydrite, alabaster, gypsum, and selenite; the carbonate occurs dissolved in most natural waters and as the minerals chalk, marble, calcite, aragonite; also in the double carbonates such as dolomite, bromlite, barytocalcite; the fluoride as fluorspar; the fluophosphate constitutes the mineral apatite; while all the more important mineral silicates contain a proportion of this element.

Extraction.—Calcium oxide or lime has been known from a very remote period, and was for a long time considered to be an elementary or undecomposable earth. This view was questioned in the 18th century, and in 1808 Sir Humphry Davy (Phil. Trans., 1808, p. 303) was able to show that lime was a combination of a metal and oxygen. His attempts at isolating this metal were not completely successful; in fact, metallic calcium remained a laboratory curiosity until the beginning of the 20th century. Davy, inspired by his successful isolation of the metals sodium and potassium by the electrolysis of their hydrates, attempted to decompose a mixture of lime and mercuric oxide by the electric current; an amalgam of calcium was obtained, but the separation of the mercury was so difficult that even Davy himself was not sure as to whether he had obtained pure metallic calcium. Electrolysis of lime or calcium chloride in contact with mercury gave similar results. Bunsen (Ann., 1854, 92, p. 248) was more successful when he electrolysed calcium chloride moistened with hydrochloric acid; and A. Matthiessen (Jour. Chem. Soc., 1856, p. 28) obtained the metal by electrolysing a mixture of fused calcium and sodium chlorides. Henri Moissan obtained the metal of 99% purity by electrolysing calcium iodide at a low red heat, using a nickel cathode and a graphite anode; he also showed that a more convenient process consisted in heating the iodide with an excess of sodium, forming an amalgam of the product, and removing the sodium by means of absolute alcohol (which has but little action on calcium), and the mercury by distillation.

The electrolytic isolation of calcium has been carefully investigated, and this is the method followed for the commercial production of the metal. In 1902 W. Borchers and L. Stockem (Zeit. für Electrochemie, 1902, p. 8757) obtained the metal of 90% purity by electrolysing calcium chloride at a temperature of about 780°, using an iron cathode, the anode being the graphite vessel in which the electrolysis was carried out. In the same year, O. Ruff and W. Plato (Ber. 1902, 35, p. 3612) employed a mixture of calcium chloride (100 parts) and fluorspar (16.5 parts), which was fused in a porcelain crucible and electrolysed with a carbon anode and an iron cathode. Neither of these processes admitted of commercial application, but by a modification of Ruff and Plato’s process, W. Ruthenau and C. Suter have made the metal commercially available. These chemists electrolyse either pure calcium chloride, or a mixture of this salt with fluorspar, in a graphite vessel which serves as the anode. The cathode consists of an iron rod which can be gradually raised. On electrolysis a layer of metallic calcium is formed at the lower end of this rod on the surface of the electrolyte; the rod is gradually raised, the thickness of the layer increases, and ultimately a rod of metallic calcium, forming, as it were, a continuation of the iron cathode, is obtained. This is the form in which calcium is put on the market.

An idea as to the advance made by this method is recorded in the variation in the price of calcium. At the beginning of 1904 it was quoted at 5s. per gram, £250 per kilogram or £110 per pound; about a year later the price was reduced to 21s. per kilogram, or 12s. per kilogram in quantities of 100 kilograms. These quotations apply to Germany; in the United Kingdom the price (1905) varied from 27s. to 30s. per kilogram (12s. to 13s. per ℔.).

Properties.—A freshly prepared surface of the metal closely resembles zinc in appearance, but on exposure to the air it rapidly tarnishes, becoming yellowish and ultimately grey or white in colour owing to the information of a surface layer of calcium hydrate. A faint smell of acetylene may be perceived during the oxidation in moist air; this is probably due to traces of calcium carbide. It is rapidly acted on by water, especially if means are taken to remove the layer of calcium hydrate formed on the metal; alcohol acts very slowly. In its chemical properties it closely resembles barium and strontium, and to some degree magnesium; these four elements comprise the so-called metals of the “alkaline earths.” It combines directly with most elements, including nitrogen; this can be taken advantage of in forming almost a perfect vacuum, the oxygen combining to form the oxide, CaO, and the nitrogen to form the nitride, Ca3N2. Several of its physical properties have been determined by K. Arndt (Ber., 1904, 37, p. 4733). The metal as prepared by electrolysis generally contains traces of aluminium and silica. Its specific gravity is 1.54, and after remelting 1.56; after distillation it is 1.52. It melts at about 800°, but sublimes at a lower temperature.

Compounds.—Calcium hydride, obtained by heating electrolytic calcium in a current of hydrogen, appears in commerce under the name hydrolite. Water decomposes it to give hydrogen free from ammonia and acetylene, 1 gram yielding about 100 ccs. of gas (Prats Aymerich, Abst. J.C.S., 1907, ii p. 460). Calcium forms two oxides—the monoxide, CaO, and the dioxide, CaO2. The monoxide and its hydrate are more familiarly known as (q.v.) and slaked-lime. The dioxide was obtained as the hydrate, CaO2·8H2O, by P. Thénard (Ann. Chim. Phys., 1818, 8, p. 213), who precipitated lime-water with hydrogen peroxide. It is permanent when dry; on heating to 130° C. it loses water and gives the anhydrous dioxide as an unstable, pale buff-coloured powder, very sparingly soluble in water. It is used as an antiseptic and oxidizing agent.

Whereas calcium chloride, bromide, and iodide are deliquescent solids, the fluoride is practically insoluble in water; this is a parallelism to the soluble silver fluoride, and the insoluble chloride, bromide and iodide. Calcium fluoride, CaF2, constitutes the mineral (q.v.), and is prepared artificially as an insoluble white powder by precipitating a solution of calcium chloride with a soluble fluoride. One part dissolves in 26,000 parts of water. Calcium chloride, CaCl2, occurs in many natural waters, and as a by-product in the manufacture of carbonic acid (carbon dioxide), and potassium chlorate. Aqueous solutions deposit crystals containing 2, 4 or 6 molecules of water. Anhydrous calcium chloride, prepared by heating the hydrate to 200° (preferably in a current of hydrochloric acid gas, which prevents the formation of any oxychloride), is very hygroscopic, and is used as a desiccating agent. It fuses at 723°. It combines with gaseous ammonia and forms crystalline compounds with certain alcohols. The crystallized salt dissolves very readily in water with a considerable absorption of heat; hence its use in forming “freezing mixtures.” A temperature of −55°C. is obtained by mixing 10 parts of the hexahydrate with 7 parts of snow. A saturated solution of calcium chloride contains 325 parts of CaCl2 to 100 of water at the boiling point (179.5°). Calcium iodide and bromide are white deliquescent solids and closely resemble the chloride.

Chloride of lime or “bleaching powder” is a calcium chlor-hypochlorite or an equimolecular mixture of the chloride and hypochlorite (see and ).

Calcium carbide, CaC2, a compound of great industrial importance as a source of acetylene, was first prepared by F. Wohler. It is now manufactured by heating lime and carbon in the electric furnace (see ). Heated in chlorine or with bromine, it yields carbon and calcium chloride or bromide; at a dull red heat it burns in oxygen, forming calcium carbonate, and it becomes incandescent in sulphur vapour at 500°, forming calcium sulphide and carbon disulphide. Heated in the electric furnace in a current of air, it yields calcium cyanamide (see ).

Calcium carbonate, CaCO3, is of exceptionally wide distribution in both the mineral and animal kingdoms. It constitutes the bulk of the chalk deposits and limestone rocks; it forms over one-half of the mineral dolomite and the rock magnesium limestone; it occurs also as the dimorphous minerals (q.v.) and  (q.v.). (q.v.) and travertine are calcareous deposits found in volcanic districts. Most natural waters contain it dissolved in carbonic acid; this confers “temporary hardness” on the water. The dissipation of the dissolved carbon dioxide results in the formation of “fur” in kettles or boilers, and if the solution is falling, as from the roof of a cave, in the formation of stalactites and stalagmites. In the animal kingdom it occurs as both calcite and aragonite in the tests of the foraminifera, echinoderms, brachiopoda, and mollusca; also in the skeletons of sponges and corals. Calcium carbonate is obtained as a white precipitate, almost insoluble in water (1 part requiring 10,000 of water for solution), by mixing solutions of a carbonate and a calcium salt. Hot or dilute cold solutions deposit minute orthorhombic crystals of aragonite, cold saturated or moderately strong solutions, hexagonal (rhombohedral) crystals of calcite. Aragonite is the least stable form; crystals have been found altered to calcite.

Calcium nitride, Ca3N2, is a greyish-yellow powder formed by heating calcium in air or nitrogen; water decomposes it with evolution of ammonia (see H. Moissan, Compt. Rend., 127, p. 497).

Calcium nitrate, Ca(NO3)2·4H2O, is a highly deliquescent salt,