CALCIUM, a metallic element first obtained in the free state by Sir Humphry Davy in 1808. Its compounds are exceedingly abundant and are widely distributed. Calcium carbonate, CaCOs, is familiar in its various forms of mar ble, chalk, limestone and calcite. The sulphate, CaSO., is also very common, and is perhaps best known in the form of gypsum, which contains two molecules of water, and therefore has the formula CaSO4+ 2H10. Calcium phosphate also occurs in nature in considerable quantities, both in the form of fossilized bones and as a constituent of apatite (q.v.) and its various modifications.
Metallic calcium may be obtained by the electrolysis of the fused chloride (which melts at a red heat), or by decomposing the iodide with metallic sodium. It is a white metal with a light yellow hue, has a hardness about equal to that of gold and is very malleable and duc tile. Its density is 1.548 and its melting point about 1455° F. Its chemical symbol is. Ca, its specific gravity is about 1.58 and its atomic weight is 40.0 (0=16). Perfectly dry air does not affect it at ordinary temperatures, but in moist air it becomes rapidly coated with the hydrate Ca(OH)3. When strongly heated in air it burns with a yellow flame, taking up oxygen to form the oxide CaO. It decomposes water rapidly, passing into the form of the hydrate, with evolution of hydrogen. It melts at a red heat, has a specific heat of about 0.169 and has an electrical resistance only about one-twelfth of that of mercury.
In its chemical relations calcium is a dyad. It combines with almost every known acid, and yields a vast number of compounds, many of which are of great industrial value. Of these the best known are the carbonate, oxide, hy drate, chloride, sulphate, phosphate, fluoride, carbide and bisulphide, and the indefinite mix ture of the chloride and hypochlorite known as bleaching-powder (q.v.).
The carbonate occurs native in large quan tities, as already noted. It is also commonly present in ground water as obtained from wells and springs. It is almost insoluble in pure water, but dissolves to a considerable extent when the water contains free carbon dioxide in solution. It is this compound that gives to water what is known as °temporary hardness.°
Upon boiling, the free carbon dioxide held in solution is expelled, and the lime carbonate is therefore precipitated also, so that the water loses that part of its hardness which is due to the presence of the carbonate. This effect is well illustrated, in regions where the soil is rich in by the crust of lime carbonate that is deposited upon the interior of household kettles that are used for heating water. Cal cium carbonate also gives rise, in steam boilers, to troublesome deposits that keep the water out of contact with the metal plates, which, in such cases, become overheated and seriously impaired in consequence. To prevent this action chem ists often recommend the addition to the water in the boiler of a certain amount of ammonium chloride (sal ammoniac). This compound com bines with the lime carbonate to form calcium chloride, which is exceedingly soluble, and ammonium carbonate, which is volatile, and therefore passes away with the steam. Beautiful as this process is in theory, it cannot be recom mended for adoption in practice, because if the sal ammoniac is present in any excess it induces rapid corrosion of the boiler plates.
When calcium carbonate (more familiarly known as carbonate of lime) is strongly heated in a current of air, it loses its carbon dioxide and becomes converted into a substance known to the chemist as calcium oxide CaO, and in the arts as quicklime, burnt lime or simply lime. Pure calcium oxide (or lime) is a white, amor phous substance, infusible, glowing with a dazzling white light when strongly heated, possessing caustic properties and acting as a powerful chemical base. When treated with about one-third of its own weight of water, lime passes into the form of the hydrate or hydrox ide, Ca(OH)2, with the evolution of much heat. The process of converting it into the hydrate by the addition of water is called slaking, and the resulting hydrate is known in the arts as slaked lime. Mortar is composed of a mixture of slaked lime and sand, the silica (or sand) slowly com bining with the lime to form a silicate after the mortar has been applied. Slaked lime, or cal cium hydrate, is somewhat soluble in water, its solution being known as lime water.