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]
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.
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 salt metathesis reaction between calcium hypochlorite and various metal sulfates. 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.
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
Chloroperoxidases are enzymes that catalyzes the 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 white blood cells, called neutrophil granulocytes.[12] These granulocytes engulf viruses and bacteria in an intracellular vacuole called the phagosome, where they are digested.
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 sodium ("liquid bleach", "Javel water") and calcium ("bleaching powder") are widely used, industrially and domestically, to whiten clothes, lighten hair color and remove stains. They were the first commercial bleaching products, developed soon after that property was discovered in 1785 by French chemist Claude Berthollet.
Hypochlorites are also widely used as broad spectrum disinfectants and deodorizers. That application started soon after French chemist Labarraque discovered those properties, around 1820 (still before 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.[17]
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.
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).
^Topić, Filip; Marrett, Joseph M.; Borchers, Tristan H.; Titi, Hatem M.; Barrett, Christopher J.; Friščić, Tomislav (2021). "After 200 Years: The Structure of Bleach and Characterization of Hypohalite Ions by Single-Crystal X-Ray Diffraction". Angew. Chem. Int. Ed.60 (46): 24400–24405. doi:10.1002/anie.202108843. PMID34293249. S2CID236199263.
^ abAylett, founded by A.F. Holleman ; continued by Egon Wiberg ; translated by Mary Eagleson, William Brewer ; revised by Bernhard J. (2001). Inorganic chemistry (1st English ed., [edited] by Nils Wiberg. ed.). San Diego, Calif. : Berlin: Academic Press, W. de Gruyter. p. 444. ISBN978-0123526519.((cite book)): CS1 maint: multiple names: authors list (link)
^Clancey, V.J. (1975). "Fire hazards of calcium hypochlorite". Journal of Hazardous Materials. 1 (1): 83–94. doi:10.1016/0304-3894(75)85015-1.
^Brauer, G. (1963). Handbook of Preparative Inorganic Chemistry; Vol. 1 (2nd ed.). Academic Press. p. 309.
^Vickery, R. C. (1 April 1950). "Some reactions of cerium and other rare earths with chlorine and hypochlorite". Journal of the Society of Chemical Industry. 69 (4): 122–125. doi:10.1002/jctb.5000690411.
^ abV. R. Sastri; et al. (2003). Modern Aspects of Rare Earths and their Complexes (1st ed.). Burlington: Elsevier. p. 38. ISBN978-0080536682.
