Permanganometry: Redox Titrations Using Potassium Permanganate

Permanganometry is redox titrimetry using standardized potassium permanganate, $\text{KMnO}_4$, as the oxidizing titrant. It is one of the most common oxidimetric methods taught at this level because $\text{KMnO}_4$ is a strong, reliable oxidant that reacts with many reducing agents by a single, clean, well-characterized half-reaction in acidic medium:

so its n-factor in acidic medium is $5$, giving an equivalent weight of $\tfrac{158}{5}=31.6\ \text{g equiv}^{-1}$.

Self-indicating behaviour. $\text{KMnO}_4$ solutions are an intense purple-violet, while its reduction product $\text{Mn}^{2+}$ is almost colourless at the low concentrations typically present near the end point. As permanganate is added from the burette, it is reduced (and decolourized) as fast as it is added, so the solution being titrated stays essentially colourless until every last trace of the reductant has been consumed. The very next drop of $\text{KMnO}_4$ then has nothing left to react with, and persists — giving a permanent, faint pink colour that marks the end point without needing any separate indicator added.

A worked calculation — titrating $\text{Fe}^{2+}$. Iron(II) is oxidized cleanly to iron(III) by permanganate in acidic medium, following $\text{MnO}_4^{-} + 5\text{Fe}^{2+} + 8\text{H}^{+} \to \text{Mn}^{2+} + 5\text{Fe}^{3+} + 4\text{H}_2\text{O}$. If $V_1$ mL of $M_1$ molar $\text{KMnO}_4$ exactly oxidizes $V_2$ mL of an $\text{Fe}^{2+}$ solution of unknown molarity $M_2$, the mole ratio from the equation (1 mol $\text{MnO}_4^-$ per 5 mol $\text{Fe}^{2+}$) gives $M_1 V_1 \times 5 = M_2 V_2 \times 1$, or equivalently, in equivalents, $N_1 V_1 = N_2 V_2$ with $N_1 = 5M_1$ and $N_2 = M_2$ (since $n=1$ for the simple one-electron $\text{Fe}^{2+}\to\text{Fe}^{3+}$ oxidation). …