Hypochlorite

In chemistry, hypochlorite, or chloroxide, is an oxyanion with the chemical formula ClO⁻. It combines with a number of cations to form hypochlorite salts. Common examples include sodium hypochlorite (household bleach) and calcium hypochlorite (a component of bleaching powder, swimming pool "chlorine"). The Cl–O distance in ClO⁻ is 1.69 Å.

The name can also refer to esters of hypochlorous acid, namely organic compounds with a ClO– group covalently bound to the rest of the molecule. The principal example is tert-butyl hypochlorite, which is a useful chlorinating agent.

Most hypochlorite salts are handled as aqueous solutions. Their primary applications are as bleaching, disinfection, and water treatment agents. They are also used in chemistry for 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  + + +

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.

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. The pentahydrate NaOCl·5H₂O is unstable above 0 °C, although the more dilute solutions encountered as household bleach are more stable. Potassium hypochlorite (KOCl) is known only in solution.

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

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 (), then dichloramine (), and finally nitrogen trichloride ().]]

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). 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 white blood cells, called neutrophil granulocytes. 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.

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 that will eventually die off. 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.

{| class="wikitable"

|-

! Ion !! Acidic reaction !! E° (V) !! Neutral/basic reaction !! E° (V)

|-

| align="center" | Hypochlorite || H⁺ + HOCl + e⁻ →  Cl₂(g) + H₂O || align="center" |1.63 || ClO⁻ + H₂O + 2 e⁻ → Cl⁻ + 2OH⁻ || align="center" |0.89

|-

| align="center" | Chlorite || 3 H⁺ + HOClO + 3 e⁻ →  Cl₂(g) + 2 H₂O || align="center" |1.64|| + 2 H₂O + 4 e⁻ → Cl⁻ + 4 OH⁻ || align="center" | 0.78

|-

| align="center" | Chlorate || 6 H⁺ + + 5 e⁻ →  Cl₂(g) + 3 H₂O ||align="center" |1.47|| + 3 H₂O + 6 e⁻ → Cl⁻ + 6 OH⁻ || align="center" |0.63

|-

| align="center" | Perchlorate ||8 H⁺ + + 7 e⁻ →  Cl₂(g) + 4 H₂O ||align="center" |1.42|| + 4 H₂O + 8 e⁻ → Cl⁻ + 8 OH⁻ || 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 to .

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. Chlorine can also assume oxidation state +4 as seen in the neutral compound chlorine dioxide ClO₂.

{| class="wikitable"

|-

! Chlorine oxidation state

| −1

| +1

| +3

| +5

| +7

|-

! Name

| chloride

| hypochlorite

| chlorite

| chlorate

| perchlorate

|-

! Formula

| Cl⁻

| ClO⁻

|

|

|

|-

! Structure

| 50px|alt=A green sphere|The chloride ion

| 50px|The hypochlorite ion

| 50px|The chlorite ion

| 50px|The chlorate ion

| 50px|The perchlorate ion

|}

See also

• Chlorine oxide

References