Upper secondary · 15 min

Permanganometric titration of iron(II)

Titrate an acidified Fe²⁺ solution with KMnO₄: the purple oxidant is its own indicator and the titration curve shows a steep jump at equivalence.

Goal

Read the equivalence volume from the curve and compute the unknown Fe²⁺ concentration using the 1:5 stoichiometry of MnO₄⁻/Fe²⁺.

Apparatus and reagents

Burette of KMnO₄ 0.02 mol/L, flask with Fe²⁺ solution acidified by dilute H₂SO₄, white tile, magnetic stirrer.

Procedure

Set the analyte type to the strong-acid option (acid = 0) and the Fe²⁺ concentration C.
Sweep the added volume Vb from 0 to 50 mL and locate the steepest part of the curve.
At the first permanent pale pink, record Vb: this is the equivalence volume Ve.
Compute c(Fe²⁺) = 5 · c(KMnO₄) · Ve / V(sample) and compare with the value set in the widget.

What to observe

Before equivalence each drop of MnO₄⁻ is instantly decolourised by Fe²⁺; just after equivalence the excess MnO₄⁻ persists as a pale pink tint.
The curve stays almost flat then jumps steeply near the equivalence volume, mirroring the abrupt potential jump at the endpoint.

Explanation

In acid solution $\ce{MnO4^- + 5Fe^{2+} + 8H+ -> Mn^{2+} + 5Fe^{3+} + 4H2O}$ : one mole of permanganate consumes five moles of Fe²⁺. Because $\ce{MnO4^-}$ is intensely violet while all products are nearly colourless, no external indicator is needed — the first persistent pink drop signals equivalence (self-indicating titrant).

History of the experiment

Permanganometry was introduced by Fernand Margueritte in 1846 to assay iron ores, exploiting the fact that

\ce{MnO4^-}

announces its own endpoint. The same potentiometric jump underlies modern electrochemical sensors, where the electrode replaces the eye.

Chemists behind it

Joseph Louis Gay-Lussac Walther Nernst