(a) Suppose that tests of a municipal water system reveal the presence of bromate ion, \(\mathrm{BrO}_{3}^{-} .\) What are the likely origins of this ion? (b) Is bromate ion an oxidizing or reducing agent?

The bromate ion (\(\mathrm{BrO}_{3}^-\)) can be formed in a municipal water system through various sources. Some of the possible origins include: 1. Disinfection byproducts: Bromate ions can be produced as a byproduct during the disinfection process when ozone is used to treat bromide-containing water. The reaction between ozone (\(\mathrm{O}_{3}\)) and bromide ions (\(\mathrm{Br}^-\)) can generate bromate ions as follows: \[2\mathrm{O}_{3} + \mathrm{Br}^- \rightarrow \mathrm{BrO}_{3}^- + 2\mathrm{O}_{2}\] 2. Contamination from industrial waste: Industries that use bromine or produce bromine-containing compounds can release wastewater containing bromate ions, which may enter the municipal water system. 3. Natural occurrence: Bromate ions can also form naturally in small amounts, when bromide minerals in the environment react with oxygen.

To determine whether the bromate ion (\(\mathrm{BrO}_{3}^-\)) is an oxidizing or reducing agent, we need to look at its redox potential: - The half-reaction for the bromate ion acting as an oxidizing agent is: \[\mathrm{BrO}_{3}^{-} + 6\mathrm{H}^{+} + 6\mathrm{e}^{-} \rightarrow \mathrm{Br}^{-} + 3\mathrm{H}_{2}\mathrm{O}\] - The half-reaction for the bromate ion acting as a reducing agent is: \[\mathrm{BrO}_{3}^{-} \rightarrow \mathrm{Br}^{-} + 3\mathrm{O}_{2} + 6\mathrm{H}^{+} + 6\mathrm{e}^{-}\] Comparing their standard reduction potentials, we see that the reaction with the bromate ion acting as an oxidizing agent (first half-reaction) is more favorable: \[E^{\circ}_{\mathrm{BrO}_{3}^{-}/\mathrm{Br}^{-}} = 1.50\text{ V}\] Since the reaction with the bromate ion acting as an oxidizing agent has a higher standard reduction potential, we can conclude that the bromate ion (\(\mathrm{BrO}_{3}^-\)) predominantly acts as an oxidizing agent in redox reactions.