Chlorite

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.

For the mineral type, see chlorite group. For the neutral chemical compound, see chlorine dioxide.
Chlorite
Names
IUPAC name
Chlorite
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.123.477
EC Number
  • 215-285-9
UNII
Properties
ClO−
2
Molar mass 67.452
Conjugate acid Chlorous acid
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

Compounds

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]

Sodium chlorite is derived indirectly from sodium chlorate, NaClO3. First, the explosively unstable gas chlorine dioxide, ClO2 is produced by reducing sodium chlorate with a suitable reducing agent such as methanol, hydrogen peroxide, hydrochloric acid or sulfur dioxide.

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]

IonAcidic reactionE° (V)Neutral/basic reactionE° (V)
HypochloriteH+ + HOCl + e− → ​1⁄2 Cl2(g) + H2O1.63ClO− + H2O + 2 e− → Cl− + 2 OH−0.89
Chlorite3 H+ + HOClO + 3 e− → ​1⁄2 Cl2(g) + 2 H2O1.64ClO−
2 + 2 H2O + 4 e− → Cl− + 4 OH−		0.78
Chlorate6 H+ + ClO−
3 + 5 e− → ​1⁄2 Cl2(g) + 3 H2O		1.47ClO−
3 + 3 H2O + 6 e− → Cl− + 6 OH−		0.63
Perchlorate8 H+ + ClO−
4 + 7 e− → ​1⁄2 Cl2(g) + 4 H2O		1.42ClO−
4 + 4 H2O + 8 e− → Cl− + 8 OH−		0.56

Uses

The most important chlorite is sodium chlorite (NaClO2); this is 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

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
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See also

References
  1. Greenwood, N.N.; Earnshaw, A. (2006). Chemistry of the elements (2nd ed.). Oxford: Butterworth-Heinemann. p. 861. ISBN 0750633654.
  2. Cotton, F. Albert; Wilkinson, Geoffrey (1988), Advanced Inorganic Chemistry (5th ed.), New York: Wiley-Interscience, p. 564, ISBN 0-471-84997-9
  • Kirk-Othmer Concise Encyclopedia of Chemistry, Martin Grayson, Editor, John Wiley & Sons, Inc., 1985
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