Oxoacids of chlorine, bromine and iodine

Table 1.1 lists the families of oxoacids known for Cl, Br and I.

Table 1.1 Oxoacids of chlorine, bromine and iodine.

   The hypohalous acids, HOX, are obtained in aqueous solution by reaction 1.1 (compare reactions 1.9 and 16.42).

          (1.1)

   All are unknown as isolated compounds, but act as weak acids in aqueous solutions (pKa values: HOCl, 4.53; HOBr, 8.69; HOI, 10.64). Hypochlorite salts such as NaOCl, KOCl and Ca(OCl)₂ (equation 1.2) can be isolated; NaOCl can be crystallized from a solution obtained by electrolysing aqueous NaCl in such a way that the Cl₂ liberated at the anode mixes with the NaOH produced at the cathode. Hypochlorites are powerful oxidizing agents and in the presence of alkali convert [IO₃]^- to [IO₄]^-, Cr³⁺‏ to [CrO₄]^2- , and even Fe^3‏+ to [FeO₄]^2-. Bleaching powder is a non-deliquescent mixture of CaCl₂, Ca(OH)₂ and Ca(OCl)₂ and is manufactured by the action of Cl₂ on Ca(OH)₂; NaOCl is a bleaching agent and disinfectant.

                 (1.2)

   All hypohalites are unstable with respect to disproportionation (equation 1.3); at 298 K, the reaction is slow for [OCl]^-, fast for [OBr]^- and very fast for [OI]^- . Sodium hypochlorite disproportionates in hot aqueous solution (equation 1.4), and the passage of Cl₂ through hot aqueous alkali yields chlorate and chloride salts rather than hypochlorites. Hypochlorite solutions decompose by reaction 1.5 in the presence of cobalt(II) compounds as catalysts.

                         (1.3)

                                          (1.4)

                              (1.5)

   Like HOCl, chlorous acid, HClO₂, is not isolable but is known in aqueous solution and is prepared by reaction 1.6; it is weak acid (pK_a = 2.0). Sodium chlorite (used as a bleach) is made by reaction 1.7; the chlorite ion has the bent structure 1.1.

                                                                                    (1.6)

                                    (1.7)

(1.1)

   Alkaline solutions of chlorites persist unchanged over long periods, but in the presence of acid, a complex decomposition occurs which is summarized in equation 1.8.

                                              (1.8)

   Chloric and bromic acids, HClO₃ and HBrO₃, are both strong acids but cannot be isolated as pure compounds. The aqueous acids can be made by reaction 1.9 (compare with reaction 1.6).

      (1.9)

   Iodic acid, HIO₃, is a stable, white solid at room temperature, and is produced by reacting I₂O₅ with water (equation in below) or by the oxidation of I₂ with nitric acid.

Crystalline iodic acid contains trigonal HIO₃ molecules connected by extensive hydrogen bonding. In aqueous solution it is af airly strong acid (pK_a = 0.77).

   Chlorates are strong oxidizing agents; commercially, NaClO₃ is used for the manufacture of ClO₂ and is used as a weedkiller, and KClO₃ has applications in fireworks and safety matches. Chlorates are produced by electrolysis of brine at 340 K, allowing the products to mix efficiently (scheme 1.10); chlorate salts are crystallized fromthe mixture.

                (1.10)

Fig. 1.1 Structures of (a) perchloric acid, in which one Cl^_O bond is unique, and (b) perchlorate ion, in which all Cl^_O bonds are equivalent. Colour code: Cl, green; O, red; H, white.

    Anodic oxidation of [OCl]^- produces further [ClO₃]^-.   Bromates are made by, for example, reaction 1.11 under alkaline conditions. Reaction 1.12 is a convenient synthesis of KIO₃.

                                       (1.11)

                                           (1.12)

    Potassium bromate and iodate are commonly used in volumetric analysis. Very pure KIO₃ is easily obtained, and reaction 1.13 is used as a source of I₂ for the standardization of thiosulfate solutions (reaction 15.113).

                                    (1.13)

(1.2)

    Halate ions are trigonal pyramidal (1.2) although, in the solid state, some metal iodates contain infinite structures in which two O atoms of each iodate ion bridge two metal centres.† The thermal decomposition of alkali metal chlorates follows reaction 1.14, but in the presence of a suitable catalyst, KClO₃ decomposes to give O₂. Some iodates (e.g. KIO₃) decompose when heated to iodide and O₂, but others (e.g. Ca(IO₃)₂) give oxide, I₂ and O₂. Bromates behave similarly and the interpretation of these observations is a difficult problem in energetics and kinetics.

                             (1.14)

   Perchloric acid is the only oxoacid of Cl that can be isolated, and its structure is shown in Figure 1.1a. It is a colourless liquid (bp 363K with some decomposition), made by heating KClO₄ with concentrated H₂SO₄ under   reduced pressure. Pure perchloric acid is liable to explode when heated or in the presence of organic material, but in dilute solution, [ClO₄]^- is very difficult to reduce despite the reduction potentials (which provide thermodynamic but not kinetic data) shown in equations 1.15 and 1.17. Zinc, for example, merely liberates H₂, and iodide ion has no action.   

   Reduction to Cl^- can be achieved by Ti(III) in acidic solution or by Fe(II) in the presence of alkali.

              (1.15)

               (1.17)

    Perchloric acid is an extremely strong acid in aqueous solution . Although [ClO₄]^- (Figure 1.1b) does form complexes with metal cations, the tendency to do so is less than for other common anions. Consequently, NaClO₄ solution is a standard medium for the investigation of ionic equilibria in aqueous systems. Alkali metal perchlorates can be obtained by disproportionation of chlorates (equation 1.14) under carefully controlled conditions; traces of impurities can catalyse decomposition to chloride and O₂. Perchlorate salts are potentially explosive and must be handled with particular care; mixtures of ammonium perchlorate and aluminium are standard missile propellants, e.g. in the space shuttle. When heated, KClO₄ gives KCl and O₂, apparently without intermediate formation of KClO₃.   Silver perchlorate, like silver salts of some other very strong acids (e.g. AgBF₄, AgSbF₆ and AgO₂CCF₃), is soluble in many organic solvents including C₆H₆ and Et₂O owing to complex formation between Ag‏ and the organic molecules.   The best method of preparation of perbromate ion is by reaction 1.18. Cation exchange can be used to give HBrO₄, but the anhydrous acid has not been isolated.

                                     (1.18)

    Potassium perbromate has been structurally characterized and contains tetrahedral [BrO₄]^- ions (Br^_O = 161 pm).  Thermochemical data show that [BrO₄]^- (half-reaction 1.19) is a slightly stronger oxidizing agent than [ClO₄]^- or [IO₄]^- under the same conditions. However, oxidations by [BrO₄]^- (as for [ClO₄]^-) are slow in dilute neutral solution, but more rapid at higher acidities.

                  (1.19)

    Several different periodic acids and periodates are known; Table 1.1 lists periodic acid, HIO₄ and orthoperiodic acid, H₅IO₆ (compare with H₆TeO₆). Oxidation of KIO₃ by hot alkaline hypochlorite yields K₂H₃IO₆ which is converted to KIO4 by nitric acid; treatment with concentrated alkali yields K₄H₂I₂O₁₀, and dehydration of this at 353K leads to K₄I₂O₉. Apart from [IO₄]^- (1.3) and [IO₅]^3- and [HIO₅]^2-  (which are square-based pyramidal), periodic acids and periodate ions feature octahedral I centres, e.g. H₅IO₆ (1.4), [H₂I₂O₁₀]^4-  (1.5) and [I₂O₉]^4- (1.6).

(1.3)                                (1.4)                             (1.5)                                         (1.6)

    The relationships between these ions may be expressed by equilibria 1.20, and aqueous solutions of periodates are therefore not simple systems.

                           (1.20)

     Orthoperiodic acid is obtained by electrolytic oxidation of iodic acid, or by adding concentrated nitric acid to Ba₅(IO₆)₂, prepared by reaction 1.21.

                                       (1.21)

     Heating H₆IO₆ dehydrates it, first to H₄I₂O₉, and then to HIO₄. In aqueous solution, both H₆IO₆ (pK_a = 3.3) and HIO₄ (pK_a=  1.64) behave as rather weak acids. Periodate oxidizes iodide (equation 1.22) rapidly even in neutral solution (compare the actions of chlorate and bromate); itliberates ozonized O₂ from hot acidic solution, and oxidizes Mn(II) to [MnO₄]^-.

                           (1.22)