Hypochlorite

Source: Wikipedia, the free encyclopedia.
For the neutral chemical compound, see chlorine monoxide.
Hypochlorite
The hypochlorite ion 2D
The Lewis structure for the hypochlorite anion
The hypochlorite ion 3D
Names
IUPAC name
Hypochlorite
Systematic IUPAC name
Chlorate(I)
Other names
Chloroxide
Identifiers
3D model (
JSmol
)
ChEBI
ChemSpider
ECHA InfoCard
 100.235.795 Edit this at Wikidata
682
UNII
UN number 3212
  • InChI=1S/ClO/c1-2/q-1 checkY
    Key: WQYVRQLZKVEZGA-UHFFFAOYSA-N checkY
  • InChI=1/ClO/c1-2/q-1
    Key: WQYVRQLZKVEZGA-UHFFFAOYAZ
  • [O-]Cl
Properties
Conjugate acid
 Hypochlorous acid
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

In

cations to form hypochlorite salts. Common examples include sodium hypochlorite (household bleach) and calcium hypochlorite (a component of bleaching powder, swimming pool "chlorine").[1] The Cl-O distance in ClO− is 1.69 Å.[2]

The name can also refer to

tert-butyl hypochlorite, which is a useful chlorinating agent.[3]

Most hypochlorite salts are handled as

oxidation
reactions.

Reactions

Acid reaction

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

2 H+
 + ClO−
 + Cl−
 ⇌ Cl
2 + H
2O

Stability

Hypochlorites are generally unstable and many compounds exist only in solution. Lithium hypochlorite LiOCl, calcium hypochlorite Ca(OCl)2 and barium hypochlorite Ba(ClO)2 have been isolated as pure anhydrous compounds. All are solids. A few more can be produced as aqueous solutions. 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.[4] Calcium hypochlorite is produced on an industrial scale and has good stability. Strontium hypochlorite, Sr(OCl)2, is not well characterised and its stability has not yet been determined.[citation needed]

Upon heating, hypochlorite degrades to a mixture of chloride, oxygen, and chlorates:

2 ClO−
 → 2 Cl−
 + O
2
3 ClO−
 → 2 Cl−
 + ClO−
3

This reaction is exothermic and in the case of concentrated hypochlorites, such as LiOCl and Ca(OCl)2, can lead to a dangerous thermal runaway and potentially explosions.[5]

The alkali metal hypochlorites decrease in stability down the group. Anhydrous lithium hypochlorite is stable at room temperature; however, sodium hypochlorite is explosive as an anhydrous solid.[6] The pentahydrate (NaOCl·(H2O)5) is unstable above 0 °C;[7] although the more dilute solutions encountered as household bleach are more stable. Potassium hypochlorite (KOCl) is known only in solution.[4]

Lanthanide hypochlorites are also unstable; however, they have been reported as being more stable in their anhydrous forms than in the presence of water.[8] Hypochlorite has been used to oxidise cerium from its +3 to +4 oxidation state.[9]

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 (NH
2Cl), then dichloramine (NHCl
2), and finally nitrogen trichloride (NCl
3).[1]

NH
3 + ClO−
 → HO−
 + NH
2Cl
NH
2Cl + ClO−
 → HO−
 + NHCl
2
NHCl
2 + ClO−
 → HO−
 + NCl
3

Preparation

Hypochlorite salts

Hypochlorite salts formed by the reaction between chlorine and alkali and alkaline earth metal hydroxides. The reaction is performed at close to room temperature to suppress the formation of chlorates. This process is widely used for the industrial production of sodium hypochlorite (NaClO) and calcium hypochlorite (Ca(ClO)2).

Cl2 + 2 NaOH → NaCl + NaClO + H2O
2 Cl2 + 2 Ca(OH)2 → CaCl2 + Ca(ClO)2 + 2 H2O

Large amounts of sodium hypochlorite are also produced

electrochemically via an un-separated chloralkali process
. In this process brine is electrolyzed to form Cl
2 which dissociates in water to form hypochlorite. This reaction must be conducted in non-acidic conditions to prevent release of chlorine:

2 Cl−
 → Cl
2 + 2 e−
Cl
2 + H
2O ⇌ HClO + Cl−
 + H+

Some hypochlorites may also be obtained by a

precipitate
out of solution, driving the reaction to completion.

Ca(ClO)2 + MSO4 → M(ClO)2 + CaSO4

Organic hypochlorites

tert-butyl hypochlorite is a rare example of a stable organic hypochlorite.[10]

Hypochlorite esters are in general formed from the corresponding alcohols, by treatment with any of a number of reagents (e.g. chlorine, hypochlorous acid, dichlorine monoxide and various acidified hypochlorite salts).[3]

Biochemistry

Biosynthesis of organochlorine compounds

chlorination of organic compounds. This enzyme combines the inorganic substrates chloride and hydrogen peroxide
to produce the equivalent of Cl+, which replaces a proton in hydrocarbon substrate:

R-H + Cl− + H2O2 + H+ → R-Cl + 2 H2O

The source of "Cl+" is hypochlorous acid (HOCl).[11] Many organochlorine compounds are biosynthesized in this way.

Immune response

In response to infection, the human immune system generates minute quantities of hypochlorite within special

neutrophil granulocytes.[12] 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.[13][14][15]

Low concentrations of hypochlorite were also found to interact with a microbe's heat shock proteins, stimulating their role as intra-cellular chaperone and causing the bacteria to form into clumps (much like an egg that has been boiled) that will eventually die off.[16] The same study found that low (micromolar) hypochlorite levels induce E. coli and Vibrio cholerae to activate a protective mechanism, although its implications were not clear.[16]

In some cases, the base acidity of hypochlorite compromises a bacterium's

lipid membrane, a reaction similar to popping a balloon.[citation needed
]

Industrial and domestic uses

Hypochlorites, especially of

Claude Berthollet
.

Hypochlorites are also widely used as broad spectrum

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.[17]

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− + 2OH− 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

Hypochlorite is a sufficiently strong oxidiser to convert Mn(III) to Mn(V) during the Jacobsen epoxidation reaction and to convert Ce3+
 to Ce4+
.[9] This oxidising power is what makes them effective bleaching agents and disinfectants.

In organic chemistry, hypochlorites can be used to oxidise primary alcohols to carboxylic acids.[18]

As chlorinating agents

Hypochlorite salts can also serve as

chlorinating agents. For example, they convert phenols to chlorophenols. Calcium hypochlorite converts piperidine to N-chloropiperidine
.

Related oxyanions

Chlorine can be the nucleus of oxyanions with oxidation states of −1, +1, +3, +5, or +7. (The element can also assume oxidation state of +4 is seen in the neutral compound chlorine dioxide ClO2).

Chlorine oxidation state −1 +1 +3 +5 +7
Name chloride hypochlorite chlorite chlorate perchlorate
Formula Cl− ClO− ClO−
2 		ClO−
3 		ClO−
4
Structure A green sphere The hypochlorite ion The chlorite ion The chlorate ion The perchlorate ion

See also

References

  1. ^
    ISBN 978-0-08-037941-8
    .
  2. S2CID 236199263
    .
  3. ^
    doi:10.15227/orgsyn.049.0009
    .
  4. ^
    ISBN 978-0123526519.{{cite book}}: CS1 maint: multiple names: authors list (link
    )
  5. doi:10.1016/0304-3894(75)85015-1
    .
  6. ISBN 978-0-08-052340-8
    .
  7. ^ Brauer, G. (1963). Handbook of Preparative Inorganic Chemistry; Vol. 1 (2nd ed.). Academic Press. p. 309.
  8. doi:10.1002/jctb.5000690411
    .
  9. ^
    ISBN 978-0080536682
    .
  10. ISBN 0471936235
    .
  11. S2CID 24417282
    .
  12. PMID 22810731
    .
  13. PMID 176150
    .
  14. PMID 217834
    .
  15. PMID 6264434
    .
  16. ^
    PMID 19013278
    .
  17. ISBN 0-471-84997-9
  18. ISBN 978-0-19-927029-3
    .
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