Chlorite

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

Chlorite

Names
IUPAC name Chlorite
Identifiers
CAS Number	14998-27-7
Jmol 3D model	Interactive image
SMILES [O-]Cl=O
Properties
Chemical formula	ClO2−
Molar mass	67.452
Vapor pressure	{{{value}}}
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

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.

Compounds

See also: Category:Chlorites

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 ony commercially important chlorite. Heavy metals chlorites (Ag⁺, Hg⁺, Tl⁺, Pb²⁺, and also Cu²⁺ and NH₄⁺) are unstable and 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).

Structure and Properties

The chlorite ion adopts a bent molecular geometry, due to the effects of the lone pairs on the chlorine atom, with a O-Cl-O bond angle of 111° and Cl-O bond lengths of 156pm. When compared to other chlorine oxyanions (on the basis of standard half cell potentials) chlorite is one strongest oxidizing agents in the series.

Ion	Acidic reaction	E° (V)	Neutral/basic reaction	E° (V)
Hypochlorite	H+ + HOCl + e− → ½Cl2(g) + H2O	1.63	ClO− + H2O + 2e− → Cl− + 2OH−	0.89
Chlorite	3H+ + HOClO + 3e− → ½Cl2(g) + 2H2O	1.64	ClO2− + 2H2O + 4e− → Cl− + 4OH−	0.78
Chlorate	6H+ + ClO3− + 5e− → ½Cl2(g) + 3H2O	1.47	ClO3− + 3H2O + 6e− → Cl− + 6OH−	0.63
Perchlorate	8H+ + ClO4− + 7e− → ½Cl2(g) + 4H2O	1.42	ClO4− + 4H2O + 8e− → Cl− + 8OH−	0.56

Uses

The most important chlorite is sodium chlorite; 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⁻, ClO₂⁻, ClO₃⁻, or ClO₄⁻, 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−	ClO2−	ClO3−	ClO4−
structure

See also

• sodium chlorite, NaClO₂
• chlorous acid, HClO₂
• Tetrachlorodecaoxide, a chlorite based drug
• Chloryl, ClO₂⁺

References

• Kirk-Othmer Concise Encyclopedia of Chemistry, Martin Grayson, Editor, John Wiley & Sons, Inc., 1985