⚡ Redox Titration — Potentiometric Endpoint & the Nernst Equation

A redox titration determines the concentration of a reducible or oxidizable analyte by reacting it with a titrant of known concentration that undergoes a complementary electron-transfer reaction. Unlike acid-base titrations, the reaction transfers electrons rather than protons, and progress is followed by watching the solution's electrode potential rather than its pH.

pH tracks proton concentration, but a redox reaction doesn't necessarily change pH at all — it changes the ratio of oxidized to reduced forms of the species in solution. An inert electrode immersed in the solution develops a potential that reflects that ratio via the Nernst equation, so measuring E against a reference electrode is the direct way to follow a redox titration's progress.

The Nernst equation relates a half-cell's electrode potential to its standard potential E° and the ratio of oxidized to reduced species: E = E° + (0.0592/n)·log([Ox]/[Red]) at 25°C. Before the equivalence point, the solution potential is governed by the analyte's own redox couple; after the equivalence point, it's governed by the titrant couple, now present in excess.

As the titrated fraction approaches 1, the ratio of oxidized to reduced analyte swings from very small to very large over just a tiny addition of titrant, and the logarithm in the Nernst equation amplifies that swing into a steep, near-vertical rise in E. Exactly at equivalence, the potential passes through the average of the two couples' standard potentials.

The size of the jump is set mainly by ΔE°, the difference between the titrant's and analyte's standard potentials. A larger ΔE° produces a bigger, sharper, easier-to-detect jump, which is exactly why redox titrations are designed around couples with well-separated E° values.

A redox indicator is a dye whose own oxidized and reduced forms have different colors and whose standard potential lies close to the expected equivalence-point potential, so it flips color right as the sharp potential jump occurs, for example ferroin, used in cerium(IV) titrations of iron(II).

Potassium permanganate is intensely purple, while its reduction product Mn²⁺ is nearly colorless. During the titration the purple color is consumed as fast as it is added, so the solution stays clear; once past the equivalence point, the first slight excess of permanganate produces a faint, persistent pink that signals the endpoint — no separate indicator is needed.

The equivalence point is the exact stoichiometric point, calculated from the reaction's known mole ratio. The endpoint is the point an experimenter actually observes, usually a color change, and it should closely coincide with the equivalence point if the indicator or self-indicating couple is well chosen.

Permanganate titrations determine iron content in ores and other samples; iodometric titrations measure the concentration of oxidizing agents in industrial and water-quality analysis; and the Winkler method, a classic iodometric procedure, determines dissolved oxygen concentration in water samples.

A titration curve is really just a half-cell potential measurement, exactly like the electrode potentials that combine to give a galvanic cell's EMF. The difference is that here the concentrations are continuously changing as titrant is added, so the Nernst equation has to be re-evaluated point by point rather than for one fixed pair of concentrations.