A redox titration is a titration of a reducing agent by an oxidizing agent or titration of an oxidizing agent by a reducing agent. Typically, this type of titration involves a redox indicator or a potentiometer.
- Analytical titrations using redox reactions were introduced shortly after the development of acid–base titrimetry.
- The earliest Redox titration took advantage of the oxidizing power of chlorine. In 1787, Claude Berthollet introduced a method for the quantitative analysis of chlorine water (a mixture of Cl2, HCl, and HOCl) based on its ability to oxidize indigo, a dye that is colorless in its oxidized state.
- In 1814, Joseph Gay-Lussac developed a similar method for determining chlorine in bleaching powder. In both methods the end point is a change in color.
- Before the equivalence point the solution is colorless due to the oxidation of indigo. After the equivalence point, however, unreacted indigo imparts a permanent color to the solution.
For example, a redox titration may be set up by treating an iodine solution with a reducing agent to form the iodide. A starch solution can then be used as a color-change indicator to detect the titration endpoint. In this case, the solution begins blue and disappears at the endpoint when the iodine is all reacted.
Types of Redox Titrations
Redox titrations are named according to the titrant that is used:
- Bromometry uses a bromine (Br2) titrant.
- Cerimetry employs cerium(IV) salts.
- Dichrometry uses potassium dichromate.
- Iodometry uses iodine (I2).
- Permanganometry uses potassium permanganate.
The number of redox titrimetric methods increased in the mid-1800s with the introduction of MnO4–, Cr2O72–, and I2 as oxidizing titrants, and of Fe2+ and S2O32– as reducing titrants.
Even with the availability of these new titrants, redox titrimetry was slow to develop due to the lack of suitable indicators.
A titrant can serve as its own indicator if its oxidized and reduced forms differ significantly in color. For example, the intensely purple MnO4– ion serves as its own indicator since its reduced form, Mn2+, is almost colorless.
Other titrants require a separate indicator. The first such indicator, diphenylamine, was introduced in the 1920s. Other redox indicators soon followed, increasing the applicability of redox titrimetry.
REDOX TITRATION CURVE
To evaluate a redox titration we need to know the shape of its titration curve. In an acid–base titration or a complexation titration, the titration curve shows how the concentration of H3O+ (as pH) or Mn+ (as pM) changes as we add titrant. For a redox titration it is convenient to monitor the titration reaction’s potential instead of the concentration of one species.
You may recall from that the Nernst equation relates a solution’s potential to the concentrations of reactants and products participating in the redox reaction. Consider, for example, a titration in which a titrand in a reduced state, Ared, reacts with a titrant in an oxidized state, Box.
After each addition of titrant the reaction between the titrand and the titrant reaches a state of equilibrium. Because the potential at equilibrium is zero, the titrand’s and the titrant’s reduction potentials are identical.
This is an important observation because we can use either half-reaction to monitor the titration’s progress.
Before the equivalence point the titration mixture consists of appreciable quantities of the titrand’s oxidized and reduced forms. The concentration of unreacted titrant, however, is very small. The potential, therefore, is easier to calculate if we use the Nernst equation for the titrand’s half-reaction