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.

Chlorite

Names
IUPAC name Chlorite
Identifiers
CAS Number	14998-27-7
3D model (JSmol)	Interactive image
ChemSpider	170734
ECHA InfoCard	100.123.477
EC Number	215-285-9
PubChem CID	197148
UNII	Z63H374SB6
CompTox Dashboard (EPA)	DTXSID7021522
InChI InChI=1S/ClHO2/c2-1-3/h(H,2,3)/p-1 Key: QBWCMBCROVPCKQ-UHFFFAOYSA-M
SMILES [O-][Cl+][O-]
Properties
Chemical formula	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

CompoundsEdit

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 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 propertiesEdit

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− → ​1⁄2 Cl2(g) + H2O	1.63	ClO− + H2O + 2 e− → Cl− + 2 OH−	0.89
Chlorite	3 H+ + HOClO + 3 e− → ​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− → ​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− → ​1⁄2 Cl2(g) + 4 H2O	1.42	ClO− 4 + 4 H2O + 8 e− → Cl− + 8 OH−	0.56

UsesEdit

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 oxyanionsEdit

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

This article uses material from the Wikipedia article  Metasyntactic variable, which is released under the  Creative Commons Attribution-ShareAlike 3.0 Unported License.