Calcium oxide

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Calcium oxide

General
Systematic name	Calcium Oxide
Molecular formula	CaO
Molar mass	56.1 g/mol
Appearance	White solid
Properties
Density and phase	3350 kg/m³, solid
Solubility in water	reacts
Melting point	2572 °C (2845 K)
Boiling point	2850 °C (3123 K)
Structure
Crystal structure	Face-Centered Cubic
Thermochemistry
Δ_fH⁰_gas	43.93 kJ/mol
Δ_fH⁰_liquid	−557.33 kJ/mol
Δ_fH⁰_solid	−635.09 kJ/mol
S⁰_{gas, 1 bar}	219.71 J/mol·K
S⁰_{liquid, 1 bar}	62.31 J/mol·K
S⁰_solid	38.19 J/mol·K
Hazards
MSDS	External MSDS
NFPA 704
Supplementary data page
Structure and properties	n, ε_r, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references

Calcium oxide (CaO), commonly known as burnt lime, caustic lime, lime or quicklime, is a widely used chemical compound. It is a white, caustic and alkaline crystalline solid. As a commercial product lime often also contains magnesium oxide, silicon oxide and smaller amounts of aluminium oxide and iron oxide.

Calcium oxide is usually made by the thermal decomposition of materials such as limestone, that contain calcium carbonate (CaCO₃; mineral name: calcite) in a lime kiln. This is accomplished by heating the material to above 825°C,^[1] a process called calcination or lime-burning, to liberate a molecule of carbon dioxide(CO₂); leaving CaO. This process is reversible, since once the quicklime product has cooled, it immediately begins to absorb carbon dioxide from the air, until, after enough time, it is completely converted back to calcium carbonate. Calcination of limestone is one of the first chemical reactions discovered by man and was known in prehistory.

As hydrated or slaked lime, Ca(OH)₂ (mineral name: portlandite), it is 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 is also used in glass production and its ability to react with silicates is also used in modern metal production (steel, magnesium, aluminium 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 acidic soils; and in pollution control - in gas scrubbers to desulfurize waste gases and to treat many liquid effluents. It has traditionally been used in the burial of bodies in open graves, to hide the smell of decomposition, as well as in forensic science, to reveal fingerprints. It is a refractory and a dehydrating agent and is used to purify citric acid, glucose, dyes and as a CO₂ absorber. It is also used in pottery, paints and the food industry. It is also used in mass burials, as often occurred when epidemics or plagues struck, such as the Black Death in England.

A relatively inexpensive substance, CaO produces heat energy by the formation of the hydrate, as in the following equation:^[2]

CaO + H₂O ↔ Ca(OH)₂ + 488 BTU/lb of CaO

The hydrate can be reconverted to calcium oxide by removing the water in the reversible equation. If the hydrated lime is heated to redness, the CaO will be regenerated to reverse the reaction. One pound of water combines with approximately 3 1/9 pounds of calcium oxide to give calcium hydroxide plus 1618 BTU's of energy. This process can be used to provide a convenient portable source of heat, as for on-the-spot food warming.

World lime production is around 130 million tonnes, with the USA and China vying for first place at around 20 million tonnes each[1].