Redox Reactions as the Basis for Titrations

Redox Titrations: Using Electron Transfer to Measure Concentration

Just as acid-base titrations use a pH indicator to signal the neutralisation point, redox titrations use the transfer of electrons to determine the concentration of an oxidising or reducing agent. The key is to find the equivalence point — the exact moment when the moles of oxidant added are stoichiometrically equal to the moles of reductant present (or vice versa). At this point, the reaction is complete, and the amount of unknown substance can be calculated from the known volume and concentration of the titrant.

The challenge is detecting this equivalence point. Redox titrations employ three main strategies to do so, each relying on a different kind of colour change.

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(i) Self-Indicators: The Reagent is Its Own Signal

Some reagents are so intensely coloured that they serve as their own indicator. The classic example is the permanganate ion, $\text{MnO}_4^-$, which is a deep purple colour.

In a typical titration, a standard solution of potassium permanganate ($\text{KMnO}_4$) is added from a burette into a solution of a reducing agent, such as iron(II) ions ($\text{Fe}^{2+}$) or oxalate ions ($\text{C}_2\text{O}_4^{2-}$). The reaction proceeds as follows (in acidic medium):

The $\text{Mn}^{2+}$ ion is nearly colourless (a very pale pink). As long as any reducing agent remains in the flask, every drop of $\text{MnO}_4^-$ that is added is immediately reduced to colourless $\text{Mn}^{2+}$. The solution in the flask remains clear.

The moment all the reducing agent has been consumed, the next drop of $\text{MnO}_4^-$ is not reduced. It remains as the purple $\text{MnO}_4^-$ ion, imparting a faint but permanent pink tinge to the entire solution. This is the end point.

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(ii) Redox Indicators: A Chemical Traffic Light

When the titrant itself is not intensely coloured (like dichromate, $\text{Cr}_2\text{O}_7^{2-}$, which is orange but not intensely so), a different approach is needed. Here, a redox indicator is added to the solution. This is a substance that has two different colours — one in its oxidised form and another in its reduced form.

The ideal redox indicator is chosen so that it is not oxidised (or reduced) by the titrant until after the equivalence point has been reached. At that precise moment, the very first excess of titrant oxidises the indicator, causing a dramatic colour change.

The best example is the use of diphenylamine as an indicator in titrations involving dichromate ($\text{Cr}_2\text{O}_7^{2-}$). The dichromate ion itself is not a self-indicator. However, just after the equivalence point, the first excess of $\text{Cr}_2\text{O}_7^{2-}$ oxidises the colourless diphenylamine to an intensely blue-coloured compound, clearly signalling the end point.

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(iii) Iodometric Titrations: An Indirect Method Using Starch

This is a powerful and common method, particularly useful for determining the concentration of oxidising agents like copper(II) ions ($\text{Cu}^{2+}$). It relies on two key facts:

1. Iodine ($\text{I}_2$) forms an intense blue-coloured complex with starch.
2. Iodine has a very specific and rapid redox reaction with thiosulphate ions ($\text{S}_2\text{O}_3^{2-}$).

The process is indirect and occurs in two distinct steps.

Step 1: Liberation of Iodine

The oxidising agent (e.g., $\text{Cu}^{2+}$) is reacted with an excess of iodide ions ($\text{I}^-$). The oxidising agent oxidises the iodide to iodine. For copper(II), the reaction is:

The iodine produced is not very soluble in water, but it dissolves readily in the excess potassium iodide solution to form the triiodide ion, $\text{I}_3^-$, which is stable in solution.

Step 2: Titration of Iodine with Thiosulphate

The liberated iodine is then titrated against a standard solution of sodium thiosulphate ($\text{Na}_2\text{S}_2\text{O}_3$). The reaction is a redox reaction where iodine is reduced back to iodide, and thiosulphate is oxidised to tetrathionate:

Detecting the End Point …