Editing Chlorite

{{short description|Oxyanion with as chemical formula ClO<sub>2</sub><sup>−</sup>}}
{{for multi|the clay mineral|Chlorite group|the neutral chemical compound|Chlorine dioxide}}

{{Chembox
| ImageFile = Chlorition.png
| ImageClass = skin-invert-image
| ImageSize = 100
| ImageAlt = 
| ImageName = The chlorite ion
| ImageFile1 = Chlorite-3D-vdW.png
| ImageClass1 = bg-transparent
| ImageSize1 = 100
| ImageAlt1 = 
| ImageName1 = The chlorite ion
| IUPACName = Chlorite
| OtherNames = 
| Section1 = {{Chembox Identifiers
| CASNo = 14998-27-7
| CASNo_Ref = {{cascite|correct|CAS}}
| UNII_Ref = {{fdacite|correct|FDA}}
| UNII = Z63H374SB6
|  EINECS = 215-285-9
|  PubChem = 197148
|  ChemSpiderID = 170734
| SMILES = [O-][Cl+][O-]
|  StdInChI = 1S/ClHO2/c2-1-3/h(H,2,3)/p-1
|  StdInChIKey = QBWCMBCROVPCKQ-UHFFFAOYSA-M
  }}
| Section2 = {{Chembox Properties
| Formula = {{chem|ClO|2|−}}
| MolarMass = 67.452
| Appearance = 
| Density = 
| MeltingPt = 
| BoilingPt = 
| Solubility = 
| ConjugateAcid = [[Chlorous acid]]
  }}
| Section3 = {{Chembox Hazards
| MainHazards =
| FlashPt =
| AutoignitionPt =
  }}
}}

The '''chlorite''' [[ion]], or chlorine dioxide [[anion]], is the [[halite (oxyanion)|halite]] with the [[chemical formula]] of {{chem|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 [[salt (chemistry)|salt]]s of [[chlorous acid]].

==Compounds==
{{See also|Category:Chlorites}}
The free acid, [[chlorous acid]] HClO<sub>2</sub>, 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]] (NaClO<sub>2</sub>) being the only commercially important chlorite. Heavy metal chlorites (Ag<sup>+</sup>, Hg<sup>+</sup>, Tl<sup>+</sup>, Pb<sup>2+</sup>, and also Cu<sup>2+</sup> and {{chem|NH|4|+}}) are unstable and decompose explosively with heat or shock.<ref name=Greenwood>{{cite book|last1=Greenwood|first1=N.N.|last2=Earnshaw|first2=A.|title=Chemistry of the elements|date=2006|publisher=Butterworth-Heinemann|location=Oxford|isbn=0750633654|page=861|edition=2nd}}</ref>

Sodium chlorite is derived indirectly from [[sodium chlorate]], NaClO<sub>3</sub>. First, the explosively unstable gas [[chlorine dioxide]], ClO<sub>2</sub> 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 pair]]s on the chlorine atom, with an O–Cl–O bond angle of 111° and Cl–O bond lengths of 156&nbsp;pm.<ref name=Greenwood />
Chlorite is the strongest oxidiser of the chlorine [[oxyanions]] on the basis of standard [[half cell]] potentials.<ref>{{Cotton&Wilkinson5th|page=564}}</ref>

{| class="wikitable"
|-
! Ion !! Acidic reaction !! ''E''° (V) !! Neutral/basic reaction !! ''E''° (V)
|-
| align="center" | [[Hypochlorite]] || H<sup>+</sup> + HOCl + e<sup>−</sup> → {{1/2}}&nbsp;Cl<sub>2</sub>(''g'') + H<sub>2</sub>O || align="center" |1.63 || ClO<sup>−</sup> + H<sub>2</sub>O + 2&nbsp;e<sup>−</sup> → Cl<sup>−</sup> + 2&nbsp;OH<sup>−</sup> || align="center" |0.89
|-
| align="center" | '''Chlorite''' || 3&nbsp;H<sup>+</sup> + HOClO + 3&nbsp;e<sup>−</sup> → {{1/2}}&nbsp;Cl<sub>2</sub>(''g'') + 2&nbsp;H<sub>2</sub>O || align="center" |1.64||{{chem|ClO|2|−}} + 2&nbsp;H<sub>2</sub>O + 4&nbsp;e<sup>−</sup> → Cl<sup>−</sup> + 4&nbsp;OH<sup>−</sup> || align="center" | 0.78
|-
| align="center" | [[Chlorate]] || 6&nbsp;H<sup>+</sup> + {{chem|ClO|3|−}} + 5&nbsp;e<sup>−</sup> → {{1/2}}&nbsp;Cl<sub>2</sub>(''g'') + 3&nbsp;H<sub>2</sub>O ||align="center" |1.47||{{chem|ClO|3|−}} + 3&nbsp;H<sub>2</sub>O + 6&nbsp;e<sup>−</sup> → Cl<sup>−</sup> + 6&nbsp;OH<sup>−</sup> || align="center" |0.63
|-
| align="center" | [[Perchlorate]] ||8&nbsp;H<sup>+</sup> + {{chem|ClO|4|−}} + 7&nbsp;e<sup>−</sup> → {{1/2}}&nbsp;Cl<sub>2</sub>(''g'') + 4&nbsp;H<sub>2</sub>O ||align="center" |1.42||{{chem|ClO|4|−}} + 4&nbsp;H<sub>2</sub>O + 8&nbsp;e<sup>−</sup> → Cl<sup>−</sup> + 8&nbsp;OH<sup>−</sup> || align="center" |0.56
|}

==Uses==
The most important chlorite is [[sodium chlorite]] (NaClO<sub>2</sub>), 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]] (ClO<sub>2</sub>), normally via a reaction with HCl:

:5 NaClO<sub>2</sub> + 4 HCl → 5 NaCl + 4 ClO<sub>2</sub> + 2 H<sub>2</sub>O

== 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 notation|parts per billion]] for chlorite in drinking water<ref>{{Cite web |title=Final Public Health Goal for Chlorite 
|url=https://oehha.ca.gov/water/public-health-goal/final-public-health-goal-chlorite |access-date=2023-08-08 |website=oehha.ca.gov}}</ref> after scientists in the state reported that exposure to higher levels of chlorite affect sperm and thyroid function, cause stomach ulcers, and caused red blood cell damage in laboratory animals.<ref>{{Cite web |last=Group |first=Environmental Working |title=EWG's Tap Water Database: Contaminants in Your Water |url=https://www.ewg.org/tapwater/reviewed-disinfection-byproducts.php |access-date=2023-08-08 |website=www.ewg.org |language=en}}</ref> Some studies have indicated that at certain levels chlorite may also be carcinogenic.<ref>{{Cite web |title=Public Health Goal for Chlorite in Drinking Water |url=https://oehha.ca.gov/media/downloads/water/chemicals/phg/chloritephgfinal052209.pdf |access-date=2023-08-08 |website=oehha.ca.gov}}</ref>  

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.<ref>{{Cite web |last=US EPA |first=OW |date=2015-10-13 |title=Stage 1 and Stage 2 Disinfectants and Disinfection Byproducts Rules |url=https://www.epa.gov/dwreginfo/stage-1-and-stage-2-disinfectants-and-disinfection-byproducts-rules |access-date=2023-08-08 |website=www.epa.gov |language=en}}</ref>

==Other oxyanions==
Several [[oxyanion]]s of chlorine exist, in which it can assume [[oxidation state]]s of −1, +1, +3, +5, or +7 within the corresponding anions Cl<sup>−</sup>, ClO<sup>−</sup>, {{chem|ClO|2|−}}, {{chem|ClO|3|−}}, or {{chem|ClO|4|−}}, known commonly and respectively as chloride, hypochlorite, chlorite, chlorate, and perchlorate. These are part of a greater family of other [[chlorine oxide]]s.

{| class="wikitable"
|-
! oxidation state
| −1
| +1
| +3
| +5
| +7
|-
! anion named
| [[chloride]]
| [[hypochlorite]]
| '''chlorite'''
| [[chlorate]]
| [[perchlorate]]
|-
! formula
| Cl<sup>−</sup>
| ClO<sup>−</sup>
| {{chem|ClO|2|−}}
| {{chem|ClO|3|−}}
| {{chem|ClO|4|−}}
|-
! structure
| [[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]]
| [[File:Perchlorate-3D-vdW.png|50px|The perchlorate ion]]
|}

==See also==
* [[Tetrachlorodecaoxide]], a chlorite-based drug
* [[Chloryl]], {{chem|ClO|2|+}}

==References==
{{reflist}}
*''Kirk-Othmer Concise Encyclopedia of Chemistry'', Martin Grayson, Editor, John Wiley & Sons, Inc., 1985

{{Chlorites}}

{{Authority control}}

[[Category:Chlorites| ]]
[[Category:Chlorine oxides]]