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Halogen Oxyacids and Their Salts

Halogen Oxyacids and Their Salts

The compounds HXO, HXO2, HXO3, and HXO4, where X represents Cl, Br, or I, are the hypohalous, halous, halic, and perhalic acids, respectively. The strengths of these acids increase from the hypohalous acids, which are very weak acids, to the perhalic acids, which are very strong. The table below lists the known acids, and, where known, their pKa values are given in parentheses.

Oxyacids of the Halogens
NameFluorineChlorineBromineIodine
hypohalousHOFHOCl (7.5)HOBr (8.7)HOI (11)
halous HClO2 (2.0)  
halic HClO3HBrO3HIO3 (0.8)
perhalic HClO4HBrO4HIO4 (1.6)
paraperhalic   H5IO6 (1.6)

The only known oxyacid of fluorine is the very unstable hypofluorous acid, HOF, which is prepared by the reaction of gaseous fluorine with ice:

\({\text{F}}_{2}(g)+{\text{H}}_{2}\text{O}(s)\;⟶\;\text{HOF}(g)+\text{HF}(g)\)

The compound is very unstable and decomposes above −40 °C. This compound does not ionize in water, and there are no known salts. It is uncertain whether the name hypofluorous acid is even appropriate for HOF; a more appropriate name might be hydrogen hypofluorite.

The reactions of chlorine and bromine with water are analogous to that of fluorine with ice, but these reactions do not go to completion, and mixtures of the halogen and the respective hypohalous and hydrohalic acids result. Other than HOF, the hypohalous acids only exist in solution. The hypohalous acids are all very weak acids; however, HOCl is a stronger acid than HOBr, which, in turn, is stronger than HOI.

The addition of base to solutions of the hypohalous acids produces solutions of salts containing the basic hypohalite ions, OX. It is possible to isolate these salts as solids. All of the hypohalites are unstable with respect to disproportionation in solution, but the reaction is slow for hypochlorite. Hypobromite and hypoiodite disproportionate rapidly, even in the cold:

\(3{\text{XO}}^{\text{−}}(aq)\;⟶\;2{\text{X}}^{\text{−}}(aq)+{\text{XO}}_{3}{}^{\text{−}}(aq)\)

Sodium hypochlorite is an inexpensive bleach (Clorox) and germicide. The commercial preparation involves the electrolysis of cold, dilute, aqueous sodium chloride solutions under conditions where the resulting chlorine and hydroxide ion can react. The net reaction is:

\({\text{Cl}}^{\text{−}}(aq)+{\text{H}}_{2}\text{O}(l)\;\stackrel{\phantom{\rule{0.5em}{0ex}}\text{electrical energy}\phantom{\rule{0.5em}{0ex}}}{\to }\;{\text{ClO}}^{\text{−}}(aq)+{\text{H}}_{2}(g)\)

The only definitely known halous acid is chlorous acid, HClO2, obtained by the reaction of barium chlorite with dilute sulfuric acid:

\(\text{Ba}{({\text{ClO}}_{2})}_{2}(aq)+{\text{H}}_{2}{\text{SO}}_{4}(aq)\;⟶\;{\text{BaSO}}_{4}(s)+2{\text{HClO}}_{2}(aq)\)

Filtering the insoluble barium sulfate leaves a solution of HClO2. Chlorous acid is not stable; it slowly decomposes in solution to yield chlorine dioxide, hydrochloric acid, and water. Chlorous acid reacts with bases to give salts containing the chlorite ion (shown in the figure below). Sodium chlorite finds an extensive application in the bleaching of paper because it is a strong oxidizing agent and does not damage the paper.

Three models of molecules are shown, each surrounded by brackets and each with a superscript negative sign outside the brackets. The left molecule shows a chlorine atom with two orbitals occupied by lone pairs of electrons. The chlorine atom is single bonded to two oxygen atoms, all of which are located at 109.5 degree angles from one another. The center molecule shows a space-filling model with a green atom labeled, “C l,” bonded to two red atoms labeled, “O.” The right molecule is a Lewis structure of a chlorine atom with two lone pairs of electrons surrounded by two oxygen atoms on either side, each with four lone pairs of electrons.

Chlorite ions, \({\text{ClO}}_{2}{}^{\text{−}},\) are produced when chlorous acid reacts with bases.

Chloric acid, HClO3, and bromic acid, HBrO3, are stable only in solution. The reaction of iodine with concentrated nitric acid produces stable white iodic acid, HIO3:

\({\text{I}}_{2}(s)+10{\text{HNO}}_{3}(aq)\;⟶\;2{\text{HIO}}_{3}(s)+10{\text{NO}}_{2}(g)+4{\text{H}}_{2}\text{O}(l)\)

It is possible to obtain the lighter halic acids from their barium salts by reaction with dilute sulfuric acid. The reaction is analogous to that used to prepare chlorous acid. All of the halic acids are strong acids and very active oxidizing agents. The acids react with bases to form salts containing chlorate ions (shown in the figure below). Another preparative method is the electrochemical oxidation of a hot solution of a metal halide to form the appropriate metal chlorates. Sodium chlorate is a weed killer; potassium chlorate is used as an oxidizing agent.

Three models of molecules are shown, each surrounded by brackets and each with a superscript negative sign outside the brackets. The left molecule shows a chlorine atom with one orbital occupied by a lone pair of electrons. The chlorine atom is single bonded to three oxygen atoms, all of which are located at 109.5 degree angles from one another. The center molecule shows a space-filling model with a green atom labeled, “C l,” bonded to three red atoms labeled, “O.” The right molecule is a Lewis structure of a chlorine atom with a lone pair of electrons surrounded by three oxygen atoms, each with four lone pairs of electrons.

Chlorate ions, \({\text{ClO}}_{3}{}^{\text{−}},\) are produced when halic acids react with bases.

Perchloric acid, HClO4, forms when treating a perchlorate, such as potassium perchlorate, with sulfuric acid under reduced pressure. The HClO4 can be distilled from the mixture:

\({\text{KClO}}_{4}(s)+{\text{H}}_{2}{\text{SO}}_{4}(aq)\;⟶\;{\text{HClO}}_{4}(g)+{\text{KHSO}}_{4}(s)\)

Dilute aqueous solutions of perchloric acid are quite stable thermally, but concentrations above 60% are unstable and dangerous. Perchloric acid and its salts are powerful oxidizing agents, as the very electronegative chlorine is more stable in a lower oxidation state than 7+. Serious explosions have occurred when heating concentrated solutions with easily oxidized substances.

However, its reactions as an oxidizing agent are slow when perchloric acid is cold and dilute. The acid is among the strongest of all acids. Most salts containing the perchlorate ion (shown in the figure below) are soluble. It is possible to prepare them from reactions of bases with perchloric acid and, commercially, by the electrolysis of hot solutions of their chlorides.

Two models of molecules are shown, both with a superscript negative sign. The left molecule shows a space-filling model with a green atom labeled, “C l,” bonded to four red atoms labeled, “O.” The right molecule is a Lewis structure of a chlorine atom surrounded by four oxygen atoms, each with four lone pairs of electrons. The Lewis structure is surrounded by brackets, and the superscript negative sign appears outside the brackets.

Perchlorate ions, \({\text{ClO}}_{4}{}^{\text{−}},\) can be produced when perchloric acid reacts with a base or by electrolysis of hot solutions of their chlorides.

Perbromate salts are difficult to prepare, and the best syntheses currently involve the oxidation of bromates in basic solution with fluorine gas followed by acidification. There are few, if any, commercial uses of this acid or its salts.

There are several different acids containing iodine in the 7+-oxidation state; they include metaperiodic acid, HIO4, and paraperiodic acid, H5IO6. These acids are strong oxidizing agents and react with bases to form the appropriate salts.

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