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For the mineral type, see chlorite group. For the neutral chemical compound, see chlorine dioxide.
The chlorite ion

The chlorite ion, or chlorine dioxide anion, is ClO2. A chlorite (compound) is a compound that contains this group, with chlorine in oxidation state +3. Chlorites are also known as salts of chlorous acid.

Oxidation states[edit]

Chlorine can assume oxidation states of −1, +1, +3, +5, or +7 within the corresponding anions Cl, ClO, ClO2, ClO3, or ClO4, known commonly and respectively as chloride, hypochlorite, chlorite, chlorate, and perchlorate. An additional oxidation state of +4 is seen in the neutral compound chlorine dioxide ClO2, which has a similar structure.

oxidation state −1 +1 +3 +5 +7
anion named chloride hypochlorite chlorite chlorate perchlorate
formula Cl ClO ClO2 ClO3 ClO4
structure The chloride ion The hypochlorite ion The chlorite ion The chlorate ion The perchlorate ion

Some chlorite compounds[edit]


The free acid, chlorous acid, HClO2, is only stable at low concentrations. Since it cannot be concentrated, it is not a commercial product. However, the corresponding sodium salt, sodium chlorite, NaClO2 is stable and inexpensive enough to be commercially available. The corresponding salts of heavy metals (Ag+, Hg+, Tl+, Pb2+, and also Cu2+ and NH4+) decompose explosively with heat or shock.

Sodium chlorite is derived indirectly from sodium chlorate, NaClO3. First, the explosively unstable gas chlorine dioxide, ClO2 is produced by reducing sodium chlorate in a strong acid solution with a suitable reducing agent (for example, sodium chloride, sulfur dioxide, or hydrochloric acid).


The main application of sodium chlorite is the generation of chlorine dioxide for the bleaching and stripping of textiles, pulp, and paper. It is also used for disinfection in a few municipal water treatment plants after conversion to chlorine dioxide. An advantage in this application, as compared to the more commonly used chlorine, is that trihalomethanes are not produced from organic contaminants. In the European Union, however, the use of Sodium chlorite has been phased out by 24-10-2009 for use as certain types of biocides: Product-type(PT) 2: Private area and public health area disinfectants and other biocidal products, PT 3: Veterinary hygiene biocidal products, PT 4: Food and feed area disinfectants and PT 5: Drinking water disinfectants.[1] Since September 2006, its application as molluscicide (PT 16) is forbidden as well. Sodium chlorite, NaClO2 also finds application as a component of contact lens cleaning solution under the trade name purite.


Health related information: see sodium chlorite. Sodium chlorite, like many oxidizers, should be protected from inadvertent contamination by organic materials to avoid the formation of an explosive mixture.

See also[edit]

an explanation of the Lewis dot structure of the chlorite ion


  1. ^ Commission Decision of 14 October 2008 concerning the non-inclusion of certain substances in Annex I, IA or IB to Directive 98/8/EC of the European Parliament and of the Council concerning the placing of biocidal products on the market (notified under document number C(2008) 5894)(Text with EEA relevance)(2008/809/EC) [1]
  • Chemistry of the Elements, N.N. Greenwood and A. Earnshaw, Pergamon Press, 1984.
  • Kirk-Othmer Concise Encyclopedia of Chemistry, Martin Grayson, Editor, John Wiley & Sons, Inc., 1985