Chlorite
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| Names | |
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| IUPAC name
Chlorite
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| Identifiers | |
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3D model (JSmol)
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PubChem CID
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| Properties | |
| ClO− 2 | |
| Molar mass | 67.452 |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Template:Chembox Footer/trackingTemplate:Short description
The chlorite ion, or chlorine dioxide anion, is the halite with the chemical formula of ClO−
2. A chlorite (compound) is a compound that contains this group, with chlorine in the oxidation state of +3. Chlorites are also known as salts of chlorous acid.
Compounds
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The free acid, chlorous acid HClO2, is the least stable oxoacid of chlorine and has only been observed as an aqueous solution at low concentrations. Since it cannot be concentrated, it is not a commercial product. The alkali metal and alkaline earth metal compounds are all colorless or pale yellow, with sodium chlorite (NaClO2) being the only commercially important chlorite. Heavy metal chlorites (Ag+, Hg+, Tl+, Pb2+, and also Cu2+ and NH+
4) are unstable and decompose explosively with heat or shock.[1]
Structure and properties
The chlorite ion adopts a bent molecular geometry, due to the effects of the lone pairs on the chlorine atom, with an O–Cl–O bond angle of 111° and Cl–O bond lengths of 156 pm.[1] Chlorite is the strongest oxidiser of the chlorine oxyanions on the basis of standard half cell potentials.[2]
| Ion | Acidic reaction | E° (V) | Neutral/basic reaction | E° (V) |
|---|---|---|---|---|
| Hypochlorite | H+ + HOCl + e− → Page Template:Fraction/styles.css has no content.1⁄2 Cl2(g) + H2O | 1.63 | ClO− + H2O + 2 e− → Cl− + 2 OH− | 0.89 |
| Chlorite | 3 H+ + HOClO + 3 e− → Page Template:Fraction/styles.css has no content.1⁄2 Cl2(g) + 2 H2O | 1.64 | ClO− 2 + 2 H2O + 4 e− → Cl− + 4 OH− |
0.78 |
| Chlorate | 6 H+ + ClO− 3 + 5 e− → Page Template:Fraction/styles.css has no content.1⁄2 Cl2(g) + 3 H2O |
1.47 | ClO− 3 + 3 H2O + 6 e− → Cl− + 6 OH− |
0.63 |
| Perchlorate | 8 H+ + ClO− 4 + 7 e− → Page Template:Fraction/styles.css has no content.1⁄2 Cl2(g) + 4 H2O |
1.42 | ClO− 4 + 4 H2O + 8 e− → Cl− + 8 OH− |
0.56 |
Uses
The most important chlorite is sodium chlorite (NaClO2), used in the bleaching of textiles, pulp, and paper. However, despite its strongly oxidizing nature, it is often not used directly, being instead used to generate the neutral species chlorine dioxide (ClO2), normally via a reaction with HCl:
- 5 NaClO2 + 4 HCl → 5 NaCl + 4 ClO2 + 2 H2O
Health risks
In 2009, the California Office of Environmental Health Hazard Assessment, or OEHHA, released a public health goal of maintaining amounts lower than 50 parts per billion for chlorite in drinking water.[3] Some studies have indicated that at certain levels chlorite may also be carcinogenic.[4]
The federal legal limit in the United States allows chlorite up to levels of 1,000 parts per billion in drinking water, 20 times as much chlorite as California’s public health goal.[5]
Other oxyanions
Several oxyanions of chlorine exist, in which it can assume oxidation states of −1, +1, +3, +5, or +7 within the corresponding anions Cl−, ClO−, ClO−
2, ClO−
3, or ClO−
4, known commonly and respectively as chloride, hypochlorite, chlorite, chlorate, and perchlorate. These are part of a greater family of other chlorine oxides.
| oxidation state | −1 | +1 | +3 | +5 | +7 |
|---|---|---|---|---|---|
| anion named | chloride | hypochlorite | chlorite | chlorate | perchlorate |
| formula | Cl− | ClO− | ClO− 2 |
ClO− 3 |
ClO− 4 |
| structure | The chloride ion | The hypochlorite ion | The chlorite ion | The chlorate ion | The perchlorate ion |
See also
- Tetrachlorodecaoxide, a chlorite-based drug
- Chloryl, ClO+
2
References
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- ^ a b Page Module:Citation/CS1/styles.css has no content.Greenwood, N.N.; Earnshaw, A. (2006). Chemistry of the elements (2nd ed.). Oxford: Butterworth-Heinemann. p. 861. ISBN 0750633654.
- ^ Page Module:Citation/CS1/styles.css has no content.Cotton, F. Albert; Wilkinson, Geoffrey (1988), Advanced Inorganic Chemistry (5th ed.), New York: Wiley-Interscience, p. 564, ISBN 0-471-84997-9
- ^ Page Module:Citation/CS1/styles.css has no content."Final Public Health Goal for Chlorite". oehha.ca.gov. Retrieved 2023-08-08.
- ^ Page Module:Citation/CS1/styles.css has no content."Public Health Goal for Chlorite in Drinking Water" (PDF). oehha.ca.gov. Retrieved 2023-08-08.
- ^ Page Module:Citation/CS1/styles.css has no content.US EPA, OW (2015-10-13). "Stage 1 and Stage 2 Disinfectants and Disinfection Byproducts Rules". United States Environmental Protection Agency. Retrieved 2023-08-08.
- Kirk-Othmer Concise Encyclopedia of Chemistry, Martin Grayson, Editor, John Wiley & Sons, Inc., 1985
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