View source for Hypochlorite

{{Short description|Anion}}
{{For|the neutral chemical compound|chlorine monoxide}}
{{Chembox
| verifiedrevid = 442901462
| ImageFile1 = Hypochlorit-Ion.svg
| ImageSize1 = 100
| ImageClass1 = skin-invert 
| ImageAlt1 = The hypochlorite ion 2D
| ImageFile2 = Hypochlorite Lewis Structures V1.svg
| ImageClass2 = skin-invert 
| ImageAlt2 = The Lewis structure for the hypochlorite anion
| ImageSize2 = 125
| ImageFile3 = Hypochlorite-3D-vdW.png
| ImageClass3 = bg-transparent 
| ImageAlt3 = The hypochlorite ion 3D
| ImageSize3 = 100px
| OtherNames = Chloroxide
| SystematicName = Chlorate(I)
| IUPACName = Hypochlorite
|Section1={{Chembox Identifiers
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| StdInChI = 1S/ClO/c1-2/q-1
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| StdInChIKey = WQYVRQLZKVEZGA-UHFFFAOYSA-N
| CASNo = 14380-61-1
| CASNo_Ref = {{cascite|correct|CAS}}
| UNII_Ref = {{fdacite|correct|FDA}}
| UNII = T5UM7HB19N
| PubChem = 61739
| ChemSpiderID = 55632
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| UNNumber = 3212
| ChEBI_Ref = {{ebicite|correct|EBI}}
| ChEBI = 29222
| SMILES = [O-]Cl
| InChI = 1/ClO/c1-2/q-1
| InChIKey = WQYVRQLZKVEZGA-UHFFFAOYAZ
| Gmelin = 682}}
|Section2={{Chembox Properties
| ConjugateAcid = [[Hypochlorous acid]]}}
}}

In [[chemistry]], '''hypochlorite''', or '''chloroxide,''' is an [[oxyanion]] with the [[chemical formula]] ClO<sup>−</sup>. It combines with a number of [[cation]]s to form hypochlorite salts. Common examples include [[sodium hypochlorite]] (household [[bleach]]) and [[calcium hypochlorite]] (a component of bleaching powder, swimming pool "chlorine").<ref name=G&E>{{Greenwood&Earnshaw2nd}}</ref>  The Cl–O distance in ClO<sup>−</sup> is 1.69 Å.<ref>{{ cite journal | title = After 200 Years: The Structure of Bleach and Characterization of Hypohalite Ions by Single-Crystal X-Ray Diffraction | first1 = Filip | last1 = Topić | first2 = Joseph M. |last2 = Marrett | first3 = Tristan H. | last3 = Borchers | first4 = Hatem M. | last4 = Titi | first5 = Christopher J. | last5 = Barrett | first6 = Tomislav | last6 = Friščić | journal = [[Angewandte Chemie International Edition|Angew. Chem. Int. Ed.]] | volume = 60 | issue = 46 | year = 2021 | pages = 24400–24405 | doi = 10.1002/anie.202108843 | pmid = 34293249 | s2cid = 236199263 }}</ref>

The name can also refer to [[esters]] of [[hypochlorous acid]], namely [[organic compound]]s with a ClO– [[moiety (chemistry)|group]] [[covalently]] bound to the rest of the molecule.  The principal example is [[Tert-Butyl hypochlorite|''tert''-butyl hypochlorite]], which is a useful chlorinating agent.<ref name=mintz>{{cite journal|last=Mintz|first=M. J.|author2=C. Walling|title=t-Butyl hypochlorite|journal=Organic Syntheses|year=1969|volume=49|page=9|url=http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv5p0184|doi=10.15227/orgsyn.049.0009|url-access=subscription}}</ref>

Most hypochlorite salts are handled as [[aqueous solution]]s. Their primary applications are as bleaching, [[disinfection]], and [[water treatment]] agents.  They are also used in chemistry for [[Halogenation|chlorination]] and [[oxidation]] reactions.

==Reactions==

===Acid reaction===
Acidification of hypochlorites generates [[hypochlorous acid]], which exists in an equilibrium with chlorine. A lowered [[pH]] (i.e. towards acid) drives the following reaction to the right, liberating chlorine gas, which can be dangerous:

:2&nbsp;{{chem|H|+}} + {{chem|ClO|-}} + {{chem|Cl|-}} {{eqm}} {{chem|Cl|2}} + {{chem|H|2|O}}

===Stability===
Hypochlorites are generally unstable and many compounds exist only in solution. [[Lithium hypochlorite]] LiOCl, [[calcium hypochlorite]] Ca(OCl)<sub>2</sub> and [[barium hypochlorite]] Ba(ClO)<sub>2</sub> have been isolated as pure [[anhydrous]] compounds. All are solids. A few more can be produced as [[aqueous solution]]s.  In general the greater the dilution the greater their stability. It is not possible to determine trends for the [[alkaline earth metal]] salts, as many of them cannot be formed. Beryllium hypochlorite is unheard of. Pure magnesium hypochlorite cannot be prepared; however, solid Mg(OH)OCl is known.<ref name="inorgchem" />  Calcium hypochlorite is produced on an industrial scale and has good stability. Strontium hypochlorite, Sr(OCl)<sub>2</sub>, is not well characterised and its stability has not yet been determined.{{cn|date=September 2023}}

Upon heating, hypochlorite degrades to a mixture of [[chloride]], [[oxygen]], and [[chlorate]]s:

:2&nbsp;{{chem|ClO|-}} → 2&nbsp;{{chem|Cl|-}} + {{chem|O|2}}
:3&nbsp;{{chem|ClO|-}} → 2&nbsp;{{chem|Cl|-}} + {{chem|ClO|3|-}}

This reaction is exothermic and in the case of concentrated hypochlorites, such as LiOCl and Ca(OCl)<sub>2</sub>, can lead to dangerous [[thermal runaway]] and is potentially explosive.<ref>{{cite journal|last=Clancey|first=V.J.|title=Fire hazards of calcium hypochlorite|journal=Journal of Hazardous Materials|year=1975|volume=1|issue=1|pages=83–94|doi=10.1016/0304-3894(75)85015-1}}</ref>

The [[alkali metal]] hypochlorites decrease in stability down the [[Group (periodic table)|group]]. Anhydrous lithium hypochlorite is stable at room temperature; however, [[sodium hypochlorite]] is explosive as an anhydrous solid.<ref name=bret>{{cite book | first = Peter | last = Urben | name-list-style = vanc | date = 2006 | title = [[Bretherick's Handbook of Reactive Chemical Hazards]] | edition = 7th | volume = 1 | page = 1433 | isbn = 978-0-08-052340-8 }}</ref> The pentahydrate NaOCl·5H<sub>2</sub>O is unstable above 0&nbsp;°C,<ref>{{cite book|last=Brauer|first=G.|title=Handbook of Preparative Inorganic Chemistry; Vol. 1|year=1963|publisher=Academic Press|page=309|edition=2nd}}</ref> although the more dilute solutions encountered as household bleach are more stable. Potassium hypochlorite (KOCl) is known only in solution.<ref name="inorgchem">{{cite book|last=Aylett|first=founded by A.F. Holleman; continued by Egon Wiberg; translated by Mary Eagleson, William Brewer; revised by Bernhard J.|title=Inorganic chemistry|year=2001|publisher=Academic Press, W. de Gruyter.|location=San Diego, Calif. : Berlin|isbn=978-0123526519|page=444|edition=1st English ed., [edited] by Nils Wiberg.}}</ref>

[[Lanthanide]] hypochlorites are also unstable; however, they have been reported as being more stable in their anhydrous forms than in the presence of water.<ref>{{cite journal|last=Vickery|first=R. C.|title=Some reactions of cerium and other rare earths with chlorine and hypochlorite|journal=Journal of the Society of Chemical Industry|date=1 April 1950|volume=69|issue=4|pages=122–125|doi=10.1002/jctb.5000690411}}</ref> Hypochlorite has been used to oxidise [[cerium]] from its +3 to +4 [[oxidation state]].<ref name=cerium>{{cite book|last=V. R. Sastri |display-authors=etal |title=Modern Aspects of Rare Earths and their Complexes.|year=2003|publisher=Elsevier|location=Burlington|isbn=978-0080536682|page=38|edition=1st}}</ref>

[[Hypochlorous acid]] itself is not stable in isolation as it decomposes to form [[chlorine]].  Its decomposition also results in some form of oxygen.

===Reactions with ammonia===
Hypochlorites react with ammonia first giving [[monochloramine]] ({{chem|NH|2|Cl}}), then [[dichloramine]] ({{chem|NHCl|2}}), and finally [[nitrogen trichloride]] ({{chem|NCl|3}}).<ref name=G&E/>

: {{chem|NH|3}} + {{chem|ClO|-}} → {{chem|HO|-}} + {{chem|NH|2}}Cl

: {{chem|NH|2}}Cl + {{chem|ClO|-}} → {{chem|HO|-}} + {{chem|NHCl|2}}

: {{chem|NHCl|2}} + {{chem|ClO|-}} → {{chem|HO|-}} + {{chem|NCl|3}}

==Preparation==

===Hypochlorite salts===
Hypochlorite salts are formed by the reaction between [[chlorine]] and alkali and alkaline earth metal [[hydroxide]]s. The reaction is performed at close to room temperature to suppress the formation of [[chlorate]]s. This process is widely used for the industrial production of [[sodium hypochlorite]] (NaClO) and [[calcium hypochlorite]] (Ca(ClO)<sub>2</sub>).
:Cl<sub>2</sub> + 2&nbsp;NaOH  → NaCl + NaClO + H<sub>2</sub>O

:2&nbsp;Cl<sub>2</sub> + 2&nbsp;Ca(OH)<sub>2</sub> →  CaCl<sub>2</sub> + Ca(ClO)<sub>2</sub> + 2&nbsp;H<sub>2</sub>O

Large amounts of  sodium hypochlorite are also produced [[electrochemical]]ly via an un-separated [[chloralkali process]]. In this process brine is electrolyzed to form {{chem|Cl|2}} which dissociates in water to form hypochlorite. This reaction must be conducted in non-acidic conditions to prevent release of chlorine:
:2&nbsp;{{chem|Cl|-}} → {{chem|Cl|2}} + 2&nbsp;e<sup>&minus;</sup>

: {{chem|Cl|2}} + {{chem|H|2|O}} {{eqm}} {{chem|H|Cl|O}} + {{chem|Cl|-}} + {{chem|H|+}}

Some hypochlorites may also be obtained by a [[salt metathesis reaction]] between calcium hypochlorite and various metal [[sulfate]]s. This reaction is performed in water and relies on the formation of insoluble [[calcium sulfate]], which will [[precipitate]] out of solution, driving the reaction to completion.
: Ca(ClO)<sub>2</sub> + MSO<sub>4</sub> → M(ClO)<sub>2</sub> + CaSO<sub>4</sub>

===Organic hypochlorites===
[[image:TBuOCl.svg|thumb|162px|left|[[Tert-Butyl hypochlorite|''tert''-Butyl hypochlorite]] is a rare example of a stable organic hypochlorite.<ref>{{cite book|doi=10.1002/047084289X.rb388.pub2|chapter=t-Butyl Hypochlorite|title=Encyclopedia of Reagents for Organic Synthesis|year=2006|last1=Simpkins|first1=Nigel S.|last2=Cha|first2=Jin K.|isbn=0471936235}}</ref>]]
Hypochlorite esters are in general formed from the corresponding [[Alcohol (chemistry)|alcohol]]s, by treatment with any of a number of reagents (e.g. [[chlorine]], [[hypochlorous acid]], [[dichlorine monoxide]] and various acidified hypochlorite salts).<ref name=mintz/>

==Biochemistry==
===Biosynthesis of organochlorine compounds===
[[Chloroperoxidase]]s are [[enzyme]]s that catalyze the [[Chlorination reaction|chlorination]] of organic compounds.  This enzyme combines the inorganic substrates [[chloride]] and [[hydrogen peroxide]] to produce the equivalent of Cl<sup>+</sup>, which replaces a proton in hydrocarbon substrate:
:R-H  +  Cl<sup>−</sup>  +  H<sub>2</sub>O<sub>2</sub>  +  H<sup>+</sup>  →   R-Cl  +  2 H<sub>2</sub>O
The source of Cl<sup>+</sup> is hypochlorous acid (HOCl).<ref>{{cite journal|last1=Hofrichter|first1=M.|last2=Ullrich|first2=R.|last3=Pecyna|first3=Marek J.|first4=Christiane |last4= Liers|first5=Taina |last5=Lundell|journal=Appl Microbiol Biotechnol|year=2010|volume=87|issue=3|doi=10.1007/s00253-010-2633-0|pmid=20495915| pages= 871–897 | title=New and classic families of secreted fungal heme peroxidases|s2cid=24417282}}</ref> Many organochlorine compounds are [[biosynthesis|biosynthesized]] in this way.

===Immune response===
In response to infection, the human immune system generates minute quantities of hypochlorite within special [[white blood cell]]s, called [[neutrophil granulocyte]]s.<ref>{{cite journal|doi=10.1007/s00726-012-1361-4|pmid=22810731|pmc=3894431|title=Taurine and inflammatory diseases|journal=Amino Acids|volume=46|issue=1|pages=7–20|year=2014|last1=Marcinkiewicz|first1=Janusz|last2=Kontny|first2=Ewa}}</ref>   These granulocytes engulf viruses and bacteria in an intracellular vacuole called the [[phagosome]], where they are digested.

Part of the digestion mechanism involves an enzyme-mediated [[respiratory burst]], which produces reactive oxygen-derived compounds, including [[superoxide]] (which is produced by [[NADPH oxidase]]). Superoxide decays to oxygen and [[hydrogen peroxide]], which is used in a [[myeloperoxidase]]-catalysed reaction to convert [[chloride]] to hypochlorite.<ref>{{cite journal|author1=Harrison, J. E. |author2=J. Schultz|year = 1976|title = Studies on the chlorinating activity of myeloperoxidase|journal = Journal of Biological Chemistry|volume = 251|issue = 5|pages = 1371–1374|doi=10.1016/S0021-9258(17)33749-3|pmid = 176150|doi-access = free}}</ref><ref name=ref93>{{cite journal|author = Thomas, E. L.|year = 1979|title = Myeloperoxidase, hydrogen peroxide, chloride antimicrobial system: Nitrogen-chlorine derivatives of bacterial components in bactericidal action against ''Escherichia coli''|journal = Infect. Immun.|volume = 23|issue = 2|pages = 522–531|pmid = 217834|pmc = 414195|doi = 10.1128/IAI.23.2.522-531.1979}}</ref><ref>{{cite journal|last1=Albrich|first1=JM|last2=McCarthy|first2=CA|last3=Hurst|first3=JK|title=Biological reactivity of hypochlorous acid: implications for microbicidal mechanisms of leukocyte myeloperoxidase.|journal=Proceedings of the National Academy of Sciences of the United States of America|date=January 1981|volume=78|issue=1|pages=210–4|pmid=6264434|doi=10.1073/pnas.78.1.210|pmc=319021|bibcode=1981PNAS...78..210A|doi-access=free}}</ref>

Low concentrations of hypochlorite were also found to interact with a microbe's [[heat shock protein]]s, stimulating their role as [[Heat shock protein#Role as chaperone|intra-cellular chaperone]] and causing the bacteria to form into clumps that will eventually die off.<ref name=Winter>{{cite journal
 | last = Jakob
 | first = U.
 |author2=J. Winter |author3=M. Ilbert |author4=P.C.F. Graf |author5=D. Özcelik
 | title = Bleach Activates A Redox-Regulated Chaperone by Oxidative Protein Unfolding
 | journal = [[Cell (journal)|Cell]]
 | volume = 135
 | issue = 4
 | pages = 691–701
 | publisher = Elsevier
 | date = 14 November 2008
 | url= | doi =10.1016/j.cell.2008.09.024
 | pmid = 19013278
 | pmc = 2606091 }}</ref>  The same study found that low (micromolar) hypochlorite levels induce ''[[Escherichia coli|E. coli]]'' and ''[[Vibrio cholerae]]'' to activate a protective mechanism, although its implications were not clear.<ref name=Winter/>

In some cases, the base acidity of hypochlorite compromises a bacterium's [[lipid membrane]].{{citation needed|date=June 2018}}

==Industrial and domestic uses==
Hypochlorites, especially of [[sodium hypochlorite|sodium]] ("liquid bleach", "Javel water") and [[calcium hypochlorite|calcium]] ("bleaching powder") are widely used, [[industrial chemicals|industrially]] and [[household chemicals|domestically]], to whiten clothes, lighten hair color and remove [[stain]]s.  They were the first commercial bleaching products, developed soon after that property was discovered in 1785 by French chemist [[Claude Berthollet]].

Hypochlorites are widely used as broad spectrum [[disinfectant]]s and [[deodorizer]]s.  That application started soon after [[French people|French]] chemist [[Antoine Germain Labarraque|Labarraque]] discovered those properties, around 1820 (still before [[Louis Pasteur|Pasteur]] formulated his [[germ theory]] of disease).

==Laboratory uses==

===As oxidizing agents===
Hypochlorite is the strongest oxidizing agent of the chlorine oxyanions. This can be seen by comparing the standard [[half cell]] potentials across the series; the data also shows that the chlorine oxyanions are stronger oxidizers in acidic conditions.<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> + 2OH<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
|}

Hypochlorite is a sufficiently strong oxidiser to convert Mn(III) to Mn(V) during the [[Jacobsen epoxidation]] reaction and to convert {{chem|Ce|3+}} to {{chem|Ce|4+}}.<ref name=cerium />
This oxidising power is what makes them effective bleaching agents and disinfectants.

In [[organic chemistry]], hypochlorites can be used to oxidise [[primary alcohol]]s to [[carboxylic acid]]s.<ref>{{cite book|last3=Warren|first1=Jonathan|last1=Clayden|authorlink1=Jonathan Clayden|first2=Nick|last2=Greeves|first3=Stuart|title=Organic Chemistry|publisher=Oxford University Press|location=Oxford|isbn=978-0-19-927029-3|page=195|edition=2nd|year=2012}}</ref>

===As chlorinating agents===
Hypochlorite salts can also serve as [[Chlorination reaction|chlorinating agents]]. For example, they convert [[phenol]]s to [[chlorophenol]]s. [[Calcium hypochlorite]] converts [[piperidine]] to [[N-Chloropiperidine|''N''-chloropiperidine]].

==Related oxyanions==
Chlorine can be the nucleus of [[oxyanion]]s with [[oxidation state]]s of −1, +1, +3, +5, or +7. Chlorine can also assume oxidation state +4 as seen in the neutral compound [[chlorine dioxide]] ClO<sub>2</sub>.

{| class="wikitable"
|-
! Chlorine oxidation state
| −1
| +1
| +3
| +5
| +7
|-
! Name
| [[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|alt=A green sphere|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==
* [[Chlorine oxide]]

==References==
{{reflist}}

{{Hypochlorites}}

[[Category:Hypochlorites| ]]