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Chlorite

Ion or compound containing chlorine

The chlorite ion, or chlorine dioxide anion, is the halite with the chemical formula of . 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

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 ) are unstable and decompose explosively with heat or shock.

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. Chlorite is the strongest oxidiser of the chlorine oxyanions on the basis of standard half cell potentials.

Ion	Acidic reaction	E° (V)	Neutral/basic reaction	E° (V)
Hypochlorite	H+ + HOCl + e− → Cl2(g) + H2O	1.63	ClO− + H2O + 2 e− → Cl− + 2 OH−	0.89
Chlorite	3 H+ + HOClO + 3 e− → Cl2(g) + 2 H2O	1.64	+ 2 H2O + 4 e− → Cl− + 4 OH−	0.78
Chlorate	6 H+ + + 5 e− → Cl2(g) + 3 H2O	1.47	+ 3 H2O + 6 e− → Cl− + 6 OH−	0.63
Perchlorate	8 H+ + + 7 e− → Cl2(g) + 4 H2O	1.42	+ 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.{{Cite web |title=Final Public Health Goal for Chlorite

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.

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−, , , or , known commonly and respectively as chloride, hypochlorite, chlorite, chlorate, and perchlorate. These are part of a greater family of other chlorine oxides.

oxidation state	anion named	formula	structure
−1	+1	+3	+5
chloride	hypochlorite	chlorite	chlorate
Cl−	ClO−
[[File:Chloride-ion-3D-vdW.png	50px	The chloride ion]]	[[File:Hypochlorite-3D-vdW.png	50px	The hypochlorite ion]]	[[File:Chlorite-3D-vdW.png	50px	The chlorite ion]]	[[File:Chlorate-3D-vdW.png	50px	The chlorate ion]]

References

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

References

(2006). "Chemistry of the elements". Butterworth-Heinemann.
{{Cotton&Wilkinson5th
"Public Health Goal for Chlorite in Drinking Water".
US EPA, OW. (2015-10-13). "Stage 1 and Stage 2 Disinfectants and Disinfection Byproducts Rules".

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chlorites chlorine-oxides

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