Calcium hydride, CaH₂

The hydride is prepared by passing hydrogen over calcium heated to 400˚C:

Ca + H₂ CaH₂

It reacts vigorously with water liberating hydrogen:

CaH₂ + 2H₂O Ca(OH)₂ + 2H₂

The compound is a useful source of hydrogen and 1 g with water yields 1070 cm³ at s.t.p. (lithium hydride, LiH, gives an even greater yield of 2800 cm³/g and has been used as a portable source of hydrogen for filling balloons.)

Calcium oxide, quicklime, CaO

The oxide is formed when calcium is burned in oxygen:

2Ca + O₂ 2CaO

It may also be prepared by heating calcium nitrate or carbonate. On the large scale, limestone (calcium carbonate) is heated in a lime-kiln, a tall tower into the top of which lumps of limestone are fed. Producer gas is introduced at the bottom and burns to raise the temperature of the descending lumps of limestone to about 1000˚C. The reversible reaction:

CaCO₃ ⇌ CaO + CO₂, it proceeds in the forward direction, as the carbon dioxide is carried away by the upward draught through the kiln.

Pure calcium oxide may be prepared by igniting calcite in a platinum crucible. It is a white substance of very high melting point (about 2570˚C). When heated in the oxy-coal gas flame, it becomes incandescent and emits a vivid white light (‘limelight’). On exposure to air, it quickly absorbs moisture and carbon dioxide:

CaO + CO₂ CaCO₃

CaO + H₂O Ca(OH)₂; ΔH = -63 kJ

The last equation shows that the reaction of quicklime and water is strongly exothermic. The addition of a limited amount of water is called ‘slaking’ and the product (solid calcium hydroxide) slaked lime.

Calcium oxide is a basic oxide and, in the presence of moisture, reacts with gaseous acidic oxides forming salts, e.g.

CaO + SO₂ CaSO₃

At high temperatures, it can react with solid acidic oxides, for example with silica and phosphorus pentaoxide:

CaO + SiO₂ CaSiO₃

(This reaction occurs in the manufacture of sulphuric acid)

3CaO + P₂O₅ Ca₃(PO₄)₂

(This reaction occurs during the process of steel manufacture). Calcium oxide is used in the laboratory as a drying agent, in desiccators and for drying organic liquids, e.g. ethyl alcohol; or to produce ammonia by heating with an ammonium salt. A combination of these uses is to be found when quicklime and ammonium chloride are placed together in desiccators for the drying of moist crystals of ‘ammine’ salts, e.g. hexammine-nickel(II) bromide, [Ni(NH₃)₆]Br₂. The quicklime absorbs the moisture and at the same time ammonia is liberated creating an atmosphere which prevents loss of ammonia from the ammine salt.

Calcium hydroxide, slaked lime, Ca(OH)₂

The hydroxide is made by the addition of a limited amount of water to quicklime, and hence is alternatively called hydrated lime. It is sparingly soluble is known as lime-water, which with a pH value greater than 7, is alkaline to litmus. A suspension of calcium hydroxide in water is called milk of lime.

When carbon dioxide is passed through lime-water, a precipitate of calcium carbonate is first formed:

Ca(OH)₂ + CO₂ CaCO₃ + H₂O

More carbon dioxide converts the normal carbonate to the soluble bicarbonate and so the precipitate dissolves:

CaCO₃ + H₂O + CO₂ Ca(HCO₃)₂

Calcium carbonate, CaCO₃

The carbonate exists in two crystalline forms (i.e. it is dimorphous) which are calcite and aragonite. Calcite is the common form, being found as limestone, chalk, Iceland spar (a very pure form), marble and in eff-shells and sea-shells. Calcite is the stable, and aragonite the metastable form at ordinary temperature, but aragonite occurs in shells and coral and separates from hot lime-water when carbon dioxide is passed through.

Calcium carbonate is only slightly soluble in pure water but dissolves slowly in water containing carbon dioxide, due to the formation of calcium bicarbonate:

CaCO₃ + CO₂ + H₂O Ca₂⁺ + 2HCO₃^-

If calcium carbonate is heated in air, decomposition to the oxide is complete, but if heated in a closed vessel, it proceeds to an equilibrium position only:

CaCO₃ CaO + CO₂

Calcium carbonate, like all carbonates, reacts with acids evolving carbon dioxide, e.g.

CaCO₃ + 2HCl CaCl₂ + H₂O + CO₂

Calcium chloride, CaCl₂.xH₂O

The hydrated chloride can be prepared in a pure state by the reaction of calcium carbonate with hydrochloric acid. The solution is evaporated to syrup and then cooled, when crystals of the deliquescent hexahydrate, CaCl₂.6H₂O, are formed. Calcium chloride is available commercially as a by-product of the Solvay process and it is not easy to find a use for all that is produced. It has been used instead of salt to melt ice on roads (a 40 percent solution in water freezes at -55˚C) but the concentrated solution formed is slightly acidic and therefore corrosive to vehicles. A solution of the salt helps to bind the surface of a shale tennis court.

The hexahydrate melts on heating and loses water to form the dehydrate, CaCl₂.2H₂O; further strong heating gives the anhydrous salt. The latter contains a small amount of calcium oxide due to the reaction which occurs at the temperature of dehydration:

Ca₂⁺ + 2Cl^- + H₂O CaO + 2HCl

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