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Life could not exist without compounds of sodium. These compounds hold

water in body tissues, and a severe deficiency of sodium can cause death.

Blood contains sodium compounds in solution. Sodium compounds are used in

industry in the manufacture of chemicals and pharmaceuticals, in metallurgy,

in sodium vapor lamps, and in the production of hundreds of every day

products. One of the most common sodium compounds is table salt, or sodium

chloride. In its pure form sodium is a silver-white, soft and waxy metallic

element. It is the sixth most abundant element on Earth and occurs in more

that trace amounts in the stars and sun.

The secret that led to low-cost production was learned in 1789, when the

French chemist Nicolas Leblanc discovered how to make soda out of common

salt. The compound called soda is sodium carbonate. Crude sodium carbonate

is called soda ash. The carbonate also combines with water in crystals known

as washing soda, or sal soda. Soda is used in manufacturing soap, glass,

dyestuffs, and explosives and as the basis for making other sodium compounds.

Other sodium compounds, with some of their uses, are: baking soda(sodium

bicarbonate), an ingredient of baking powder; borax (sodium borate), a food

preservative; and caustic soda, or lye (sodium hydroxide), used in

soapmaking. Some properties of sodium are: Symbol Na, Atomic Number 11,

Atomic Weight 22.9898, Boiling Point 1,621.2 F, and Melting Point 208.06 F.

Sodium belongs to the group of elements known as alkali metals. It is never

found uncombined in nature and was first isolated by the English scientist

Sir Humphry Davy in 1807. Lighter than water, pure sodium can be cut with a

knife at room temperature and is brittle at low temperatures. It conducts

heat and electricity easily and exhibits a photoelectric effect, that is, it

emits electrons when exposed to light. In its pure form, sodium oxidizes

instantly when exposed to the air and reacts vigorously with water, seizing

the oxygen and a part of the hydrogen to form sodium hydroxide. The

remaining hydrogen is liberated and may ignite from the heat of the reaction.

Pure metallic sodium –usually obtained by the electrolysis of sodium

hydroxide– must be stored in kerosene to keep it from air and moisture. One

of the few uses of pure sodium is in vapor lamps along highways.


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