Lime

Calcium Oxide, CaO may be used as a drying agent. It is more rapidly exhausted than fused calcium chloride, but, owing to its greater porosity, it is more efficient.

Calcium oxide slaked with sodium hydroxide solution, and known as soda-lime, is used for various chemical operations in place of either constituent alone, being much more reactive than these. It is apparently an efficient absorbing agent for all manner of gases, for example, carbonyl chloride, sulphuretted hydrogen, antimony and arsenic hydrides, phosphorus vapour, cyanogen, cyanogen chloride, bromide, and iodide, and so on. Soda-lime becomes incandescent when a mixture of sulphuretted hydrogen and air is passed over it, owing to the vigorous reaction which it promotes between them. Guareschi suggests that it contains compounds such as Ca(ONa)₂, OHCaONa, or OH.CaO.CaONa.

Lime is very extensively employed in agriculture. Its value to the soil does not lie only in the neutralisation of the soil acidity. A considerable amount is adsorbed by the soil constituents either physically or chemically, and this adsorption appears to be followed by an increased adsorptive power of the soil for other bases - for example, potash and ammonia. Lime also improves the physical condition of the soil and throws out of action aluminium and iron salts which might be injurious to the plants. There is apparently an optimum value for the ratio of calcium oxide to magnesium oxide in the soil.

The most ancient use of lime is in the production of lime mortar. Ordinary building-mortar is composed of lime and sand. In modern practice one volume of lime is used to three of aggregate. In old mortars the average appears to have been one to one. Experiment seems to show that the ratio one to two gives the strongest mortar. The lime should be fairly rich and easily slaked, and the sand formed of sharp angular grains free from humus and clay. The lime is slaked before mixing the mortar, which is applied as a paste. The process of hardening consists essentially in the drying of the mortar. There is not sufficient water present to produce a crystalline calcium hydroxide. Where this is found, it results from accidental outside influences. There is, nevertheless, a tendency now to consider that the formation of crystalloids has a bearing on the hardening of mortar.

The inert material, sand, prevents the formation of cracks in the mortar during drying, both by reduction of the relative shrinkage and by the action of capillary forces at the surface of separation between the sand and the lime. At the exposed surface of the mortar a second reaction takes place, which is not, however, considered to be essential to the hardening process. Carbon dioxide is absorbed, resulting in a network of closely interlaced crystals of calcium carbonate, which do not extend beyond the outer layer, as can be seen by examining the mortar from old Roman buildings. If this has been undisturbed it still consists mainly of calcium hydroxide. It was formerly supposed that the hardening of mortar is in part due to the gradual formation of calcium silicate, but this view is now discredited. Silicates are found in old mortars, but there is no reason to suppose that they were not present in the original lime.

It is not advisable to use dolomitic limestone for the manufacture of lime for mortar. The hydration of the magnesia only takes place slowly, and may occur after the mortar is in position, and by expansion cause disintegration.

Lime is employed in the purification of sugar through the precipitation of impurities and the intermediate formation of mono-, di-, and tri-calcium saccharate, although the use of strontia has superseded that of lime in the treatment of molasses.

In the purification of coal-gas, carbon dioxide, sulphuretted hydrogen, and carbon bisulphide are removed by lime. There is some doubt as to the nature of the reaction with the carbon bisulphide. A certain amount of oxygen is necessary, but too much is as harmful as too little. Veley regarded calcium hydroxyhydrosulphide, Ca(OH)(SH), as the active compound, but Lewes suggested that the real agent is calcium disulphide, CaS₂, formed from the hydrosulphide by oxidation, because this would readily combine with carbon bisulphide to form a perthiocarbonate. This explains the necessity for oxygen. Excess oxygen carries the oxidation further to thiosulphate, and so prevents the reaction. An observation by Walker may have a bearing on the question of the absorption of carbon bisulphide. If the latter be shaken with milk of lime, orange crystals are obtained. Their composition corresponds to the formula CaCS₃.2Ca(OH)₂.6H₂O.

Lime is employed in the leather industry on account of its depilatory power. It is also used for the causticising of alkalies. Lime acts as a clarifying agent for sewage by coagulating the colloids, and in pharmacy lime-water mixed with linseed oil yields the so-called carron oil for the treatment of burns.