Carbothermic reactions involve the reduction of substances, often metal oxides, using carbon as the reducing agent. These chemical reactions are usually conducted at temperatures of several hundred degrees Celsius. Such processes are applied for production of the elemental forms of many elements. Carbothermic reactions are not useful for some metal oxides, such as those of sodium and potassium. The ability of metals to participate in carbothermic reactions can be predicted from Ellingham diagrams.
Carbothermal reactions produce carbon monoxide and sometimes carbon dioxide. The facility of these conversions is attributable to the entropy of reaction: two solids, the metal oxide and carbon, are converted to a new solid (metal) and a gas (CO), the latter having high entropy. Heat is required for carbothermic reactions because diffusion of the reacting solids is otherwise slow.
The most prominent example is that of iron ore smelting. Many reactions are involved, but the simplified equation is usually shown as:
3 + 3C → 4Fe + 3CO2
On a more modest scale, about 1 million tons of elemental phosphorus is produced annually by carbothermic reactions. Calcium phosphate (phosphate rock) is heated to 1,200–1,500 °C with sand, which is mostly SiO
2, and coke (impure carbon) to produce P
4. The chemical equation for this process when starting with fluoroapatite, a common phosphate mineral, is:
3F + 18SiO
2 + 30C → 3P
4 + 30CO + 18CaSiO
3 + 2CaF
Sometimes carbothermic reactions are coupled to other conversions. One example is the chloride process for separating titanium from ilmenite, the main ore of titanium. In this process, a mixture of carbon and the crushed ore is heated at 1000 °C under flowing chlorine gas, giving titanium tetrachloride:
3 + 7Cl
2 + 6C → 2TiCl
4 + 2FeCl
3 + 6CO
For some metals, carbothermic reactions do not afford the metal, but instead give the metal carbide. This behavior is observed for titanium, hence the use of the chloride process. Carbides also form upon high temperature treatment of Cr
3 with carbon. For this reason, aluminium is employed as the reducing agent.
- Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9. "Figure 8.19 Ellingham diagram for the free energy of formation of metallic oxides" p. 308
- Diskowski, Herbert; Hofmann, Thomas (2005). "Phosphorus". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a19_505. ISBN 9783527306732.
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