Sodium and its compounds

Sodium is a typical alkali metal. It occurs naturally in its ores chiefly as sodium chloride (common salt), sodium trioxonitrate(V) (Chile saltpetre) and sodium trioxocarbonate(IV) (soda ash). Sodium is extracted industrially by the Downs Process. Here, common salt is electrolysed in the molten condition. Calcium chloride is added to common salt to serve as a flux (to bring the melting point of common salt from 800°C to about 600°C). In the molten sodium chloride are Na+ and Cl- ions.

Sodium is used:

i. In the manufacture of lead(IV)tetraethyl, an anti-knock addictive used in petrol.

ii. In sodium vapour lamps (having intensity yellow illumination).

iii. In the production of titanium needed for the manufacture of heat-resistant alloys in rockets.

iv. In an alloy for making coolants for nuclear reactors.

Sodium has some important compounds (e.g NaCl, NaOH, Na2SO4, Na2CO3) that are useful industrially.

1. Sodium chloride (NaCl): When pure, it is not deliquescent (dampness is due to impurities (like MgCl2) which are deliquescent). It is a convenient starting material for manufacturing compounds like caustic soda, washing soda, baking soda and salt-cake. It is also used for the preparation of chlorine gas and hydrogen chloride needed for the manufacture of bleaching agents and fine chemicals. In addition, sodium chloride is used in salting out process in soap making or in solvent extraction of organic substance from water.

2. Sodium hydroxide (NaOH) is prepared industrially by the electrolysis of brine (a saturated solution of sodium chloride). The pellet is deliquescent. Its dissolution in water is exothermic. The resulting aqueous solution is alkaline (turns red litmus blue).

Sodium hydroxide is used in the manufacture of soaps (saponification). It is also used in the textile industries for treatment of cotton. Also, an aqueous sodium hydroxide solution is used in qualitative analysis to form the hydroxides of metal ions in solution as precipitates. This facilitates the identification of some actions in unknown samples.

3. Sodium tetraoxosulphate(VI) (Na2SO4):

It is prepared in the laboratory by neutralizing sodium hydroxide solution by dilute H2SO4. The hydrated form is the Glauber's salt, Na2SO4.10H2O.

Sodium tetraoxosulphate(VI) is used:

i. In the manufacture of glass (a mixture of two silicates e.g. Na2SiO3 and CaSiO3)/

ii. In the manufacture of sodium sulphide (by heating it with coke).

Sodium sulphide is used for stripping the hair from hides and for making shaving powder.

iii. In drying the ethereal layer during solvent extraction of organic compounds from water.

4. Sodium trioxocarbonate(IV) (Na2CO3):

It is obtained industrially by the Solvay process. The raw materials are concentrated brine (NaCl), ammonia gas (NH3) and limestone (CaCO3) which yields quicklime (CaO) and carbon(IV)oxide (CO2) when heated. The process involves essentially the following steps:

i. formation of ammoniacal brine when brine is saturated with ammonia gas.

ii. Reaction between the ammoniacal brine and carbon (IV) oxide (produced by heating limestone) to form sodium hydrogen trioxocarbonate(IV) and ammonium chloride.

iii. Sodium hydrogen trioxocarbonate (IV) (not very soluble in water) is filtered from the reaction mixture and the washed.

iv. Heating of sodium hydrogen trioxocarbonate (IV) to sodium trioxocarbonate (IV) and carbon (IV) oxide (which is recycled). Sodium trioxocarbonate (IV) is stable to heat i.e. not decomposed on heating.

v. If the anhydrous sodium trioxocarbonate (IV) (soda ash) is dissolved in hot water and allowed to cool, it gives the crystalline form (washing soda).

The important uses of sodium trioxocarbonate (IV) are:

i. Softening of water for domestic purposes.

ii. Manufacture of glass (formation of silicates). The mixture of Na2SiO3, CaSiO3 and unreacted SiO2 constituent the glass.

iii. Manufacture of caustic soda.

Calcium and its compounds

Calcium is a typical alkali earth metal. It occurs as calcium trioxocarbonate (IV) in different forms: chalk, limestone, marble, calcite, Iceland spar and aragonite. It also occurs as calcium tetraoxosulphate (VI) e.g. gypsum, anhydrite. However, it is extracted industrially by the electrolysis of calcium chloride (a by-product from Solvay process), using graphite anode (also the container) and iron cathode. It is a silvery white metal. The density is 1.55g per cm3 and the melting point is about 850°C. It is not as reactive as sodium and so it is not necessary to keep it below the surface of petroleum oil (in contrast to sodium). However, a white film of oxide is formed on the surface on exposure to air. Calcium compounds generally burn with s brick-red flame.

Some of the important reaction of calcium are:

i. In air: it forms calcium oxide (quicklime).

ii. With water: calcium metal is used in the extraction of thorium and in steel casting (as deoxidizer). The important compounds of calcium are calcium oxide (CaO), calcium hydroxide (Ca(OH)2), calcium trioxocarbonate (IV) (CaCO3), calcium tetraoxosulphate (VI) (CaSO4) and calcium chloride (CaCl2).

1. Calcium oxide (lime, quicklime) CaO, is a white solid, made industrially by the action of strong heat on limestone (CaCO3).

It is a refractory substance (it will not melt even when heated to a very high temperature) hence its use in producing light (limelight). Its important reactions are:

i. With water, slaked lime Ca(OH)2 is formed.

This reaction is called slaking. The slaked lime is changed to lime-water if water is added and filtered..

Because quicklime is basic and hygroscopic, it is used as the drying agent for ammonia gas.

ii. With silica: CaO + SiO2 —> CaSiO3

2. Calcium hydroxide, Ca(OH)2 is a white powder, prepared by addition of water to calcium oxide. Its dissolution in water is exothermic. It is sparingly soluble in water but the solubility decreases with increasing temperature. It is basic hence gives reactions expected of bases (e.g. neutralization with aqueous mineral acids). The compound is useful in many ways:

i. In agriculture, to reduce acidity in soils (liming).

ii. In preparation of calcium trioxocarbonate (IV) e.g. chalk.

iii. In the making of mortar by builders.

3. Calcium trioxocarbonate (IV), CaCO3, occurs naturally as marble, chalk, limestone and in many other forms such as stalagmites and stalactites from the roof and floor of carvers. It is a white solid which is practically insoluble in water. Some of its reactions are:

i. With dilute acids (HCl and HNO3); it evolves carbon (IV) oxide (CO2).

ii. With water containing dissolved CO2, it dissolves to give soluble calcium hydrogentrioxocarbonate(IV):

Calcium trioxocarbonate (IV) is used:

i. In the extraction of iron using the blast furnace.

ii. In the manufacture of cement.

iii. In the Solvay process of manufacturing sodium trioxocarbonate (IV).

4. Calcium tetraoxosulphate (VI), CaSO4, is a white solid. It occurs naturally as anhydrite (CaSO4) and gypsum (CaSO4.2H2O), known as plaster of Paris (POP). This involves heating gypsum to a temperature between 100°C and 200°C.

Plaster of Paris is used medically in making casts to maintain joints in a fixed position. It is used industrially for cement and wall-plasters.

5. Calcium chloride, CaCl2, is a white solid. The solution is prepared by adding marble to dilute HCl until a little of marble remains.

The mixture is filtered. The evaporation of the filtrate to dryness gives the fused solid (anhydrous CaCl2). This is very deliquescent hence its use as a drying agent for most gases (except ammonia with which it reacts).

Calcium and its compounds

Calcium is a typical alkali earth metal. It occurs as calcium trioxocarbonate (IV) in different forms: chalk, limestone, marble, calcite, Iceland spar and aragonite. It also occurs as calcium tetraoxosulphate (VI) e.g. gypsum, anhydrite. However, it is extracted industrially by the electrolysis of calcium chloride (a by-product from Solvay process), using graphite anode (also the container) and iron cathode. It is a silvery white metal. The density is 1.55g per cm3 and the melting point is about 850°C. It is not as reactive as sodium and so it is not necessary to keep it below the surface of petroleum oil (in contrast to sodium). However, a white film of oxide is formed on the surface on exposure to air. Calcium compounds generally burn with s brick-red flame.

Some of the important reaction of calcium are:

i. In air: it forms calcium oxide (quicklime).

ii. With water: calcium metal is used in the extraction of thorium and in steel casting (as deoxidizer). The important compounds of calcium are calcium oxide (CaO), calcium hydroxide (Ca(OH)2), calcium trioxocarbonate (IV) (CaCO3), calcium tetraoxosulphate (VI) (CaSO4) and calcium chloride (CaCl2).

1. Calcium oxide (lime, quicklime) CaO, is a white solid, made industrially by the action of strong heat on limestone (CaCO3).

It is a refractory substance (it will not melt even when heated to a very high temperature) hence its use in producing light (limelight). Its important reactions are:

i. With water, slaked lime Ca(OH)2 is formed.

This reaction is called slaking. The slaked lime is changed to lime-water if water is added and filtered..

Because quicklime is basic and hygroscopic, it is used as the drying agent for ammonia gas.

ii. With silica: CaO + SiO2 —> CaSiO3

2. Calcium hydroxide, Ca(OH)2 is a white powder, prepared by addition of water to calcium oxide. Its dissolution in water is exothermic. It is sparingly soluble in water but the solubility decreases with increasing temperature. It is basic hence gives reactions expected of bases (e.g. neutralization with aqueous mineral acids). The compound is useful in many ways:

i. In agriculture, to reduce acidity in soils (liming).

ii. In preparation of calcium trioxocarbonate (IV) e.g. chalk.

iii. In the making of mortar by builders.

3. Calcium trioxocarbonate (IV), CaCO3, occurs naturally as marble, chalk, limestone and in many other forms such as stalagmites and stalactites from the roof and floor of carvers. It is a white solid which is practically insoluble in water. Some of its reactions are:

i. With dilute acids (HCl and HNO3); it evolves carbon (IV) oxide (CO2).

ii. With water containing dissolved CO2, it dissolves to give soluble calcium hydrogentrioxocarbonate(IV):

Calcium trioxocarbonate (IV) is used:

i. In the extraction of iron using the blast furnace.

ii. In the manufacture of cement.

iii. In the Solvay process of manufacturing sodium trioxocarbonate (IV).

4. Calcium tetraoxosulphate (VI), CaSO4, is a white solid. It occurs naturally as anhydrite (CaSO4) and gypsum (CaSO4.2H2O), known as plaster of Paris (POP). This involves heating gypsum to a temperature between 100°C and 200°C.

Plaster of Paris is used medically in making casts to maintain joints in a fixed position. It is used industrially for cement and wall-plasters.

5. Calcium chloride, CaCl2, is a white solid. The solution is prepared by adding marble to dilute HCl until a little of marble remains.

The mixture is filtered. The evaporation of the filtrate to dryness gives the fused solid (anhydrous CaCl2). This is very deliquescent hence its use as a drying agent for most gases (except ammonia with which it reacts).