Sodium Essay, Research Paper
Sodium
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: baki
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.