Animated Solution for Chemistry - Redox Reactions: Amount of oxalic acid present in a solution can be determined by its titration with ${\text{KMnO}}_4$ solution in the presence of ${\text{H}}_2{\text{SO}}_4$. The titration gives unsatisfactory result when carried out in the presence of HCl because HCl

Imagine you are standing in a chemistry lab, carefully performing a titration. In your burette, you have the vibrant purple potassium permanganate ($KMn{O}_4$), and in your conical flask, you have a colorless solution of oxalic acid ($H_2{C}_2{O}_4$).

This is a classic redox titration, where $KMn{O}_4$ acts as a powerful oxidizing agent, and oxalic acid is the reducing agent. But there is a crucial condition for this reaction to proceed smoothly: it requires an acidic medium.

Why do we need an acid? The permanganate ion ($Mn{O}_4^{-}$) needs a steady supply of $H^{+}$ ions to get reduced to the colorless manganese ion ($M{n}^{2+}$).

The half-reaction looks like this:

Normally, we use dilute sulfuric acid ($H_{2}S{O}_4$) to provide these $H^{+}$ ions. Sulfuric acid is stable and doesn't interfere with the redox process. But what if we run out of sulfuric acid and decide to use hydrochloric acid ($HCl$) instead?

This is where the trap lies! If we add $HCl$ to the flask, we are introducing chloride ions ($C{l}^{-}$) into the mix.

Here is the catch: the chloride ion is itself a very good reducing agent. In fact, it is a better reducing agent than oxalic acid under these conditions.

When the powerful oxidizing agent $KMn{O}_4$ drops into the flask, it faces a choice. Should it oxidize the oxalic acid, or should it oxidize the chloride ions? Because $HCl$ is a stronger reducing agent, the permanganate preferentially reacts with the chloride ions.

Instead of the intended reaction, a side reaction takes over. The permanganate ion oxidizes the chloride ions into chlorine gas ($C{l}_2$), while itself getting reduced to $M{n}^{2+}$.

The chemical equation for this side reaction is:

Because the $KMn{O}_4$ is being consumed by the $HCl$ rather than the oxalic acid, you will end up using a much larger volume of titrant than expected. This completely ruins your burette reading and gives an unsatisfactory, incorrect result.

Therefore, carrying out this titration in the presence of $HCl$ is a fundamental mistake. The $HCl$ reduces the permanganate to $M{n}^{2+}$, acting as a competing reducing agent.

As a side note, you might wonder if we could use nitric acid ($HN{O}_3$) instead. The answer is no, because nitric acid is a strong oxidizing agent and would oxidize the oxalic acid itself! This is why $H_{2}S{O}_4$ remains the undisputed champion for providing an acidic medium in permanganate titrations.