Redox Reactions as the Basis of Titrations: Equivalence and Normality

A redox titration uses a balanced redox equation the same way an acid-base titration uses a neutralization equation: as an exact stoichiometric bridge between a known volume and concentration of one reagent (delivered from a burette) and an unknown amount of the substance being determined, with the reaction's own electron transfer providing the fixed mole (or equivalent) ratio between them.

Two related quantities make these calculations fast: equivalent weight and normality. The equivalent weight of an oxidizing or reducing agent is its molar mass divided by the number of electrons it gains or loses per formula unit in a specified reaction — this number-of-electrons is called the n-factor (or valence factor) of that species in that reaction. For example, $\text{MnO}_4^{-}$ has $n=5$ in acidic medium (it gains 5 electrons, becoming $\text{Mn}^{2+}$), so its equivalent weight there is $\tfrac{158}{5} = 31.6\ \text{g equiv}^{-1}$ (using the molar mass of $\text{KMnO}_4$).

Normality ($N$) is concentration expressed in equivalents per litre: $N = n \times M$, where $M$ is molarity and $n$ is the n-factor. The single most useful feature of normality is that, at the equivalence point of any titration, the number of equivalents of the oxidant delivered exactly equals the number of equivalents of the reductant consumed — regardless of how different their individual mole-ratio coefficients are in the balanced equation:

This relationship is exactly equivalent to using the balanced equation's mole ratio directly (moles of oxidant $\times$ its n-factor = moles of reductant $\times$ its n-factor, since every electron released must be captured), but it avoids having to re-derive the mole ratio from a full balanced equation every time — which is especially convenient in a laboratory setting where different students may titrate against different oxidants. …