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{{Short description|Anion}}
{{For|the neutral chemical compound|chlorine monoxide}} {{For|the neutral chemical compound|chlorine monoxide}}
{{Refimprove|date=February 2009}}

{{Chembox {{Chembox
| verifiedrevid = 402151117 | verifiedrevid = 442901462
| ImageFile1 = Hypochlorit-Ion.svg
| ImageFile = Hypochlorite-3D-vdW.png
| ImageSize1 = 100
| IamgeAlt = The hypochlorite ion
| ImageAlt1 = The hypochlorite ion 2D
| IUPACName = Hypochlorite
| ImageFile2 = Hypochlorite Lewis Structures V1.svg
| Section1 = {{Chembox Identifiers
| ImageAlt2 = The Lewis structure for the hypochlorite anion
| StdInChI = 1S/ClO/c1-2/q-1
| ImageSize2 = 125
| ImageFile3 = Hypochlorite-3D-vdW.png
| 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_Ref = {{stdinchicite|correct|chemspider}}
| StdInChIKey = WQYVRQLZKVEZGA-UHFFFAOYSA-N | StdInChIKey = WQYVRQLZKVEZGA-UHFFFAOYSA-N
| CASNo = 14380-61-1 | CASNo = 14380-61-1
| CASNo_Ref = {{cascite|correct|CAS}} | CASNo_Ref = {{cascite|correct|CAS}}
| UNII_Ref = {{fdacite|correct|FDA}}
| PubChem = 61739
| UNII = T5UM7HB19N
| PubChem_Ref = {{pubchemcite|correct|??}}
| ChemSpiderID = 55632 | PubChem = 61739
| ChemSpiderID = 55632
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| UNNumber = 3212
| ChEBI = 29222 | UNNumber = 3212
| ChEBI_Ref = {{ebicite|correct|EBI}}
| SMILES = Cl
| ChEBI = 29222
| InChI = 1/ClO/c1-2/q-1
| SMILES = Cl
| InChIKey = WQYVRQLZKVEZGA-UHFFFAOYAZ
| InChI = 1/ClO/c1-2/q-1
| Gmelin = 682}}
| InChIKey = WQYVRQLZKVEZGA-UHFFFAOYAZ
| Gmelin = 682}}
|Section2={{Chembox Properties
| ConjugateAcid = ]}}
}} }}


In ], '''hypochlorite''', or '''chloroxide''' is an ] with the ] ClO<sup>−</sup>. It combines with a number of ]s to form hypochlorite salts. Common examples include ] (household ]) and ] (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 = ] | volume = 60 | issue = 46 | year = 2021 | pages = 24400–24405 | doi = 10.1002/anie.202108843 | pmid = 34293249 | s2cid = 236199263 }}</ref>
The '''hypochlorite''' ], also known as '''chlorate(I)''' anion is ]]<sup>−</sup>. A hypochlorite compound is a ] containing this group, with chlorine in ] +1.


The name can also refer to ] of hypochlorous acid, namely ]s with a ClO– ] ] bound to the rest of the molecule. The principal example is ], 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}}</ref>
Hypochlorites are the ] of ]. Common examples include ] (chlorine ] or ]) and ] (bleaching powder or swimming pool chlorination compound). Hypochlorites are frequently quite unstable — for example, sodium hypochlorite is not available as a solid, since removal of the water from NaClO solution converts it to a mixture of ] and ]. Heating of NaClO solution also causes this reaction. ] in sunlight, giving chlorides and oxygen.


Most hypochlorite salts are handled as ]s. Their primary applications are as bleaching, ], and ] agents. They are also used in chemistry for ] and ] reactions.
Due to their low stability, hypochlorites are very strong ]s. They react with many organic and inorganic compounds. Reaction with organic compounds is very exothermic and may cause ], so hypochlorites should be handled with care. They can oxidize manganese compounds, converting them to ]s.

==Reactions==

===Acid reaction===
Acidification of hypochlorites generates ], which exists in an equilibrium with chlorine. A lowered ] (ie. 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. ] LiOCl, ] Ca(OCl)<sub>2</sub> and ] Ba(ClO)<sub>2</sub> have been isolated as pure ] compounds. All are solids. A few more can be produced as ]s. In general the greater the dilution the greater their stability. It is not possible to determine trends for the ] 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 ], ], and ]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 ] 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 ] hypochlorites decrease in stability down the ]. Anhydrous lithium hypochlorite is stable at room temperature; however, ] is explosive as an anhydrous solid.<ref name=bret>{{cite book | first = Peter | last = Urben | name-list-style = vanc | date = 2006 | title = ] | edition = 7th | volume = 1 | page = 1433 | isbn = 978-0-08-052340-8 }}</ref> The pentahydrate (NaOCl·(H<sub>2</sub>O)<sub>5</sub>) 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., by Nils Wiberg.}}</ref>

] 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 ] from its +3 to +4 ].<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>

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

===Reactions with ammonia===
Hypochlorites react with ammonia first giving ] ({{chem|NH|2|Cl}}), then ] ({{chem|NHCl|2}}), and finally ] ({{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}}


Covalent hypochlorites, such as ] are also known, and are typically unstable.
==Preparation== ==Preparation==
The sodium salt of the hypochlorite ion, NaClO, is formed by the ] of chlorine gas bubbled through dilute aqueous sodium hydroxide at room temperature:


===Hypochlorite salts===
:{{chem|Cl|2}} (g) + 2 NaOH (aq) → NaCl (aq) + NaClO (aq) + {{chem|H|2|O}} (l)
Hypochlorite salts formed by the reaction between ] and alkali and alkaline earth metal ]s. The reaction is performed at close to room temperature to suppress the formation of ]s. This process is widely used for the industrial production of ] (NaClO) and ] (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
The reaction of chlorine with hot, concentrated sodium hydroxide forms ] of a higher oxidation state:


Large amounts of sodium hypochlorite are also produced ]ly via an un-separated ]. 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:
:3 {{chem|Cl|2}} (g) + 6 NaOH (aq) → 5 NaCl(aq) + {{chem|NaClO|3}} (aq) + 3 {{chem|H|2|O}} (l)
: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|+}}
==Chemistry==

===Acid reaction===
Some hypochlorites may also be obtained by a ] between calcium hypochlorite and various metal ]s. This reaction is performed in water and relies on the formation of insoluble ], which will ] out of solution, driving the reaction to completion.
Hypochlorites generate chlorine gas when mixed with dilute acids. Hypochlorite and chloride are in equilibrium with chlorine gas:
: Ca(ClO)<sub>2</sub> + MSO<sub>4</sub> → M(ClO)<sub>2</sub> + CaSO<sub>4</sub>

===Organic hypochlorites===
] 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 ]s, by treatment with any of a number of reagents (e.g. ], ], ] and various acidified hypochlorite salts).<ref name=mintz/>

==Biochemistry==
===Biosynthesis of organochlorine compounds===
]s are ]s that catalyzes the ] of organic compounds. This enzyme combines the inorganic substrates ] and ] 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 ] in this way.

===Immune response===
In response to infection, the human immune system generates minute quantities of hypochlorite within special ]s, called ]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 ], where they are digested.

Part of the digestion mechanism involves an enzyme-mediated ], which produces reactive oxygen-derived compounds, including ] (which is produced by ]). Superoxide decays to oxygen and ], which is used in a ]-catalysed reaction to convert ] 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 ]s, stimulating their role as ] and causing the bacteria to form into clumps (much like an egg that has been boiled) 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 = ]
| 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 '']'' and '']'' 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 ], a reaction similar to popping a balloon.{{citation needed|date=June 2018}}

==Industrial and domestic uses==
Hypochlorites, especially of ] ("liquid bleach", "Javel water") and ] ("bleaching powder") are widely used, ] and ], to whiten clothes, lighten hair color and remove ]s. They were the first commercial bleaching products, developed soon after that property was discovered in 1785 by French chemist ].


Hypochlorites are also widely used as broad spectrum ]s and ]s. That application started soon after ] chemist ] discovered those properties, around 1820 (still before ] formulated his ] of disease).
:2 {{chem|H|+}} (aq) + {{chem|OCl|-}} (aq) + {{chem|Cl|-}} (aq) {{eqm}} {{chem|Cl|2}} (g) + {{chem|H|2|O}} (l)


==Laboratory uses==
Therefore, by ], a high ] drives the reaction to the left by consuming {{chem|H|+}} ions, promoting the disproportionation of chlorine into chloride and hypochlorite, whereas a low pH drives the reaction to the right, promoting the release of chlorine gas.


===Bleaching action=== ===As oxidizing agents===
Hypochlorite is the strongest oxidizing agent of the chlorine oxyanions. This can be seen by comparing the standard ] potentials across the series; the data also shows that the chlorine oxyanions are stronger oxidizers in acidic conditions.<ref>{{Cotton&Wilkinson5th|page=564}}</ref>
Hypochlorites are used as ] to remove dyes.


{| class="wikitable"
===As an oxidising agent===
|-
Hypochlorite is the strongest oxidising agent of the generalized ]s{{Citation needed|date=November 2009}}. For example, it oxidises {{chem|Mn|2+}} to ]:
! 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" | ] || 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" | ] || 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" | ] ||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 ] reaction and to convert {{chem|Ce|3+}} to {{chem|Ce|4+}}.<ref name=cerium />
:2 {{chem|Mn|2+}} + 5 {{chem|ClO|-}} + 6 {{chem|OH|-}} → 2 {{chem|MnO|4|-}} + 3 {{chem|H|2|O}} + 5 {{chem|Cl|-}}
This oxidising power is what makes them effective bleaching agents and disinfectants.


In ], hypochlorites can be used to oxidise ]s to ]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>
==Stability==
Hypochlorite is the least stable of the ]{{Citation needed|date=July 2010}}. Many hypochlorite compounds exist only in solution, as is also the case with ] (HClO) itself.


===As chlorinating agents===
Hypochlorite is unstable with respect to ]. Upon heating, it degrades to a mixture of ], ] and other ]:
Hypochlorite salts can also serve as ]. For example, they convert ]s to chlorophenols. ] converts ] to ].


==Related oxyanions==
:2 {{chem|OCl|-}} (aq) → 2 {{chem|Cl|-}} (aq) + {{chem|O|2}} (g)
Chlorine can be the nucleus of ]s with ]s of −1, +1, +3, +5, or +7. (The element can also assume oxidation state of +4 is seen in the neutral compound ] ClO<sub>2</sub>).


{| class="wikitable"
:3 {{chem|OCl|-}} (aq) → 2 {{chem|Cl|-}} (aq) + {{chem|ClO|3|-}} (aq)
|-
! Chlorine oxidation state
| −1
| +1
| +3
| +5
| +7
|-
! Name
| ]
| '''hypochlorite'''
| ]
| ]
| ]
|-
! Formula
| Cl<sup>−</sup>
| ClO<sup>−</sup>
| {{chem|ClO|2|−}}
| {{chem|ClO|3|−}}
| {{chem|ClO|4|−}}
|-
! Structure
| ]
| ]
| ]
| ]
| ]
|}


==See also== ==See also==
*] * ]
*]
*]
*]


==References==
{{reflist}}


] {{Hypochlorites}}
]


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