The reaction is performed at close to room temperature to suppress the formation of chlorates. Hypochlorite salts formed by the reaction between chlorine and alkali and alkaline earth metal hydroxides. Hypochlorites react with ammonia first giving monochloramine ( NHĢ), and finally nitrogen trichloride ( NClģ Preparation Hypochlorite salts Its decomposition also results in some form of oxygen. Hypochlorous acid itself is not stable in isolation as it decomposes to form chlorine. Hypochlorite has been used to oxidise cerium from its +3 to +4 oxidation state. Lanthanide hypochlorites are also unstable however, they have been reported as being more stable in their anhydrous forms than in the presence of water. Potassium hypochlorite (KOCl) is known only in solution. (H 2O) 5) is unstable above 0 ☌ although the more dilute solutions encountered as household bleach possess better stability.Anhydrous lithium hypochlorite is stable at room temperature however, sodium hypochlorite is explosive as an anhydrous solid. The alkali metal hypochlorites decrease in stability down the group. 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. Upon heating, hypochlorite degrades to a mixture of chloride, oxygen, and chlorates: Strontium hypochlorite, Sr(OCl) 2, is not well characterised and its stability has not yet been determined. Calcium hypochlorite is produced on an industrial scale and has good stability. Pure magnesium hypochlorite cannot be prepared however, solid Mg(OH)OCl is known. It is not possible to determine trends for the alkaline earth metal salts, as many of them cannot be formed. In general the greater the dilution the greater their stability. A few more can be produced as aqueous solutions. Lithium hypochlorite LiOCl, calcium hypochlorite Ca(OCl) 2 and barium hypochlorite Ba(ClO) 2 have been isolated as pure anhydrous compounds. Hypochlorites are generally unstable and many compounds exist only in solution. towards acid) drives the following reaction to the right, liberating chlorine gas, which can be dangerous: Reactions Acid reaction Īcidification of hypochlorites generates hypochlorous acid, which exists in an equilibrium with chlorine. They are also used in chemistry for chlorination and oxidation reactions. Their primary applications are as bleaching, disinfection, and water treatment agents. Most hypochlorite salts are handled as aqueous solutions. The principal example is tert-butyl hypochlorite, which is a useful chlorinating agent. 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. Common examples include sodium hypochlorite (household bleach) and calcium hypochlorite (a component of bleaching powder, swimming pool "chlorine"). It combines with a number of cations to form hypochlorite salts. In chemistry, hypochlorite, or chloroxide is an anion with the chemical formula ClO −.
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