|Formula weight||56.1 amu|
|Melting point||3200 K (2927 °C)|
|Boiling point||3773 K (3500 °C)|
|Density||3.3 ×103 kg/m3|
|S0gas, 1 bar||219.71 J/mol·K|
|S0liquid, 1 bar||62.31 J/mol·K|
|Ingestion||Dangerous—causes GI irritation, larger doses could be fatal.|
|Inhalation||Dangerous - irritation; chemical bronchitis or even death for larger exposures.|
|Skin||Irritation and possible burns.|
|Eyes||May cause permanent damage.|
|More info||Hazardous Chemical Database|
|SI units were used where possible. Unless otherwise stated, standard conditions were used.|
Calcium oxide is usually made by the thermal decomposition of CaCO3, heating materials such as limestone to around 500°C and so removing the carbon dioxide in a reversible reaction. It is one of the first chemical reactions discovered by man and was known in prehistory: see limekiln.
As hydrated or slaked lime, Ca(OH)2, it was used in mortar and plaster to increase the rate of hardening. Hydrated lime is very simple to make as lime is a basic anhydride and reacts vigorously with water. Lime was also used in glass production and its ability to work with silicates is also used in modern metal production (steel, magnesium, aluminum and other non-ferrous metals) industries to remove impurities as slag.
It is also used in water and sewage treatment to reduce acidity, to soften, as a flocculant and to remove phosphates and other impurities; in paper making to dissolve lignin, as a coagulant and in bleaching; in agriculture to improve acid soils; and in pollution control - in gas scrubbers to desulphurize waste gases and to treat many liquid effluents. It is a refactory and a dehydrating agent and is used to purify citric acid, glucose, dyes and as a CO2 absorber. It is also used in pottery, concrete, paints and the food industry.