Aqueous solution chemistry

   In this section, we are mainly concerned with redox processes in aqueous solution for a list of relevant topics already covered in the book. Values of E^o for halfreactions 1.1 can be measured directly for X = Cl, Br and I (Table 1.1) and their magnitudes are determined by the X^_X bond energies the electron affinities of the halogen atoms (Table 1.1) and the standard Gibbs energies of hydration of the halide ions (Table 1.1).

Table 1.1 Some physical properties of fluorine, chlorine, bromine and iodine.

This can be seen from scheme 1.2; for X = Br and I, an additional vaporization stage is needed for the element.

                                 (1.1)        

       (1.2)

   Dichlorine is a more powerful oxidizing agent in aqueous media than Br₂ or I₂, partly because of a more negative enthalpy of formation of the anion but, more importantly,because the Cl^- ion (which is smaller than Br^- or I^-)  interacts more strongly with solvent molecules. (In solid salt formation, the lattice energy factor similarly explains why chloride salts are more exothermic than corresponding bromides or iodides.)   Since F₂ liberates ozonized O₂ from water, the value of E^o for half-reaction 1.1 has no physical reality, but a value of 2.87V can be estimated by comparing the energy changes for each step in scheme 1.2 for X = F and Cl, and hence deriving the difference in E^o for half-equation 1.1 for X = F and Cl. Most of the difference between these E^o values arises from the much more negative value of Δ_hydG^o of the smaller F^- ion (Table 1.1).   Diiodine is much more soluble in aqueous iodide solutions than in water. At low concentrations of I₂, equation 1.3 describes the system; K can be found be partitioning I₂ between the aqueous layer and a solvent immiscible with water (e.g. CCl₄).

                      (1.3)

   Potential diagrams (partly calculated from thermochemical data) for Cl, Br and I are given in Figure 1.1. Because several of the oxoacids are weak, the effects of [H‏⁺] on values of some of the reduction potentials are quite complicated. For example, the disproportionation of hypochlorite to chlorate and chloride could be written as equilibrium 1.4 without involving protons.

                                      (1.4)

However, the fact that HOCl is a weak acid, while HClO3 and HCl are strong ones means that, in the presence of hydrogen ions, [OCl]^- is protonated and this affects the position of equilibrium 1.4: HOCl is more stable with respect to disproportionation than [OCl]^- .

    On the other hand, the disproportionation of chlorate into perchlorate and chloride is realistically represented by equilibrium 1.5. From the data in Figure 1.1, this reaction is easily shown to be thermodynamically favourable. Nevertheless, the reaction does not occur in aqueous solution owing to some undetermined kinetic factor.

                                 (1.5)

    Another example of the limitations of the data in Figure 1.1 is the inference that O₂ should oxidize I^- and Br^- at pH 0. Further, the fact that Cl₂ rather than O₂ is evolved when hydrochloric acid is electrolysed is a consequence of the high overpotential for O₂ evolution at most surfaces. Despite some limitations, Figure 1.1 does provide some useful information: for example, the more powerful oxidizing properties of periodate and perbromate than of perchlorate when these species are being reduced to halate ions, and the more weakly oxidizing powers of iodate and iodine than of the other halates or halogens respectively.  The fact that Figure 1.1 refers only to specific conditions is well illustrated by considering the stability of I(I).

 Hypoiodous acid is unstable with respect to disproportionation into [IO₃]^-  and I₂, and is therefore not formed when [IO₃]^- acts as an oxidant in aqueous solution.

Fig. 1.1 Potential diagrams for chlorine, bromine and iodine at pH = 0.

However, in hydrochloric acid, HOI undergoes reaction 1.6.

                             (1.6)

Under these conditions, the potential diagram becomes:

and I(I) is now stable with respect to disproportionation.