Chapter 8, Problem 7P

Consider the half-reaction:

$\begin{array}{l}{\left[{\text{MnO}}_4\right]}^{-}(\text{aq})+8{\text{H}}^{+}(\text{aq})+5{\text{e}}^{-}\rightleftharpoons 2{\text{Mn}}^{2+}(\text{aq})\text{+}4{\text{H}}_{\text{2}}\text{O(l)}\\ {\text{ E}}^{\circ }=+1.51\text{V}\end{array}$

If the ratio of concentrations of ${\left[{\text{MnO}}_4\right]}^{-}$: ${\text{Mn}}^{2+}$ is 100:1, determine E at pH values of (a) 0.5; (b) 2.0; and (c) 3.5 (T = 298 K). Over this pH range, how does the ability of permanganate (VII) (when being reduced to ${\text{Mn}}^{\text{2+}}$) to oxidize aqueous chloride, bromide or iodide ions change?