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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 |
| ΔfH0gas |
43.93 kJ/mol |
| ΔfH0liquid |
−557.33 kJ/mol |
| ΔfH0solid |
−635.09 kJ/mol |
| S0gas, 1 bar |
219.71 J/mol·K |
| S0liquid, 1 bar |
62.31 J/mol·K |
| S0solid |
38.19 J/mol·K |
| Hazards |
| MSDS |
External MSDS |
| NFPA 704 |
|
| Supplementary data page |
Structure and
properties |
n, εr, 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 lime, quicklime or burnt lime, 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). This is accomplished by heating the material to above 825°C,[1] a process called calcination or lime-burning, so as to remove the carbon dioxide. 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: see limekiln.
As hydrated or slaked lime, Ca(OH)2 (mineral name: portlandite), 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, aluminium and other non-ferrous metals) industries to remove impurities such 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. 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, concrete, paints and the food industry, where it is sometimes used (in conjunction with water) to heat items like MREs (Meals Ready To Eat) and coffee.
[edit] See also
- Calcium hydroxide
- Common chemicals
- Limelight
- Magnesium oxide
[edit] External links
[edit] References
- ^ Merck Index of chemicals and Drugs , 9th ed. monograph 1650
- Link page to external chemical sources.
