hypochlorite

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The hypochlorite ion

The hypochlorite ion is ClO⁻. A hypochlorite compound is a chemical compound containing this group, with chlorine in oxidation state +1. Because of this oxidation state, the hypochlorite ion can be alternately named the chlorate(I) ion.

Hypochlorites are the salts of hypochlorous acid. Common examples include sodium hypochlorite (chlorine bleach or bleaching agent) and calcium hypochlorite (bleaching powder). 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 sodium chloride and sodium chlorate. Heating of NaClO solution also causes this reaction. Hypochlorites decompose in sunlight, giving chlorides and oxygen.

Due to their low stability, hypochlorites are very strong oxidizing agents. They react with many organic and inorganic compounds. Reaction with organic compounds is very exothermic and may cause ignition, so hypochlorites should be handled with care. They can oxidize manganese compounds, converting them to permanganates.

Preparation

The sodium salt of the hypochlorite ion, NaClO, is formed by the disproportionation of chlorine gas bubbled through dilute aqueous sodium hydroxide at room temperature:

Cl₂ (g) + 2 NaOH (aq) → NaCl (aq) + NaClO (aq) + H₂O (l)

The reaction of chlorine with hot, concentrated sodium hydroxide forms chlorates of a higher oxidation state:

Cl₂ (g) + 6 NaOH (aq) → 5 NaCl(aq) + NaClO₃ (aq) + 3 H₂O (l)

Chemistry

Acid reaction

Hypochlorites generate chlorine gas when mixed with dilute acids. Hypochlorite and chloride are in equillibrium with chlorine gas:

2 H⁺ (aq) + OCl⁻ (aq) + Cl⁻ (aq) Cl₂ (g) + H₂O (l)

Therefore, by Le Chatelier's principle, a high pH drives the reaction to the left by consuming 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

Hypochlorites are used as bleaches to remove dyes.^{[citation needed]}

As an oxidising agent

Hypochlorite is the strongest oxidising agent of the generalized chlorates;^{[citation needed]} able to oxidize almost any reducing agent. For example, it oxidises Mn²⁺ to permanganate:

2 Mn²⁺ + 5 ClO⁻ + 6 OH⁻ → 2 MnO−4 + 3 H₂O + 5 Cl⁻

Stability

Hypochlorite is the least stable of the chlorine oxyanions. Many hypochlorite compounds exist only in solution, and are nonexistent in a pure form, as is also the case with hypochlorous acid (HClO) itself.

Hypochlorite is unstable with respect to disproportionation. Upon heating, it degrades to a mixture of chloride, oxygen and other chlorates:

2 OCl⁻ (aq) → 2 Cl⁻ (aq) + O₂ (g)

3 OCl⁻ (aq) → 2 Cl⁻ (aq) + ClO−3 (aq)

Other oxyanions

If a Roman numeral in brackets follows the word "chlorate", this indicates the oxyanion contains chlorine in the respective oxidation state, namely:

Using this convention, "chlorate" means any chlorine oxyanion. An unqualified "chlorate" typically refers only to the +5 oxidation state.^{[citation needed]}

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