Problem 108 - A Complete Resource Book in Chemistry for JEE Main

Question

In the titration between oxalic acid and \(\mathrm{KMnO}_{4}\) if \(\mathrm{H}_{2} \mathrm{SO}_{4}\) is not used in excess (a) A brown precipitate due to \(\mathrm{MnO}_{2}\) is formed (b) A red precipitate is formed. (c) There is not effect (d) A yellow precipitate is formed

Step-by-Step Solution

Answer

(a) A brown precipitate due to \(\mathrm{MnO}_{2}\) is formed.

Step 1: Understanding the Reaction

When titrating oxalic acid with potassium permanganate (\(\mathrm{KMnO}_{4}\)), the reaction occurs in an acidic medium usually provided by sulfuric acid \(\mathrm{H}_{2} \mathrm{SO}_{4}\). The reaction relies on the acidic environment to ensure that the reduction of \(\mathrm{MnO}_{4}^{-}\) to \(\mathrm{Mn}^{2+}\) proceeds without interference.

Step 2: Role of Excess Sulfuric Acid

\(\mathrm{H}_{2} \mathrm{SO}_{4}\) is used in excess to maintain an acidic environment. If \(\mathrm{H}_{2} \mathrm{SO}_{4}\) is not in excess, then the medium might not be sufficiently acidic. The expected reaction would not proceed fully as intended in such an environment.

Step 3: Effect of Insufficient Sulfuric Acid

Without sufficient \(\mathrm{H}_{2} \mathrm{SO}_{4}\), \(\mathrm{KMnO}_{4}\) does not reduce properly and instead produces \(\mathrm{MnO}_{2}\). \(\mathrm{MnO}_{2}\) is a brown precipitate. This occurs because \(\mathrm{MnO}_{4}^{-}\) can convert to \(\mathrm{MnO}_{2}\) in a less acidic medium.

Key Concepts

Role of sulfuric acid in titration

Sulfuric acid ( H_2 S O_4 ) plays a critical role in the titration of oxalic acid with potassium permanganate ( KMnO_4 ). It acts to maintain an acidic environment, which is crucial for the reaction to occur properly. In the context of this redox reaction, sulfuric acid provides the necessary hydrogen ions ( H^+ ) to keep the solution at the right level of acidity.
This acidic medium is essential because it allows the permanganate ions ( MnO_4^- ) to effectively undergo reduction to manganese ( Mn^{2+} ). Without a sufficiently acidic medium, the oxidative power of potassium permanganate cannot be properly harnessed, which may affect the titration's accuracy.

• An excess of H_2 S O_4 ensures the solution remains acidic throughout the titration, preventing unwanted side reactions.

The presence of sulfuric acid thus facilitates the smooth progression of the intended reaction between oxalic acid and potassium permanganate.

Effect of acidic medium on redox reactions

The acidic medium, typically maintained by H_2 S O_4 , greatly impacts the redox reactions involved in the titration process. For the redox reaction between oxalic acid and KMnO_4 , an acidic environment is necessary to ensure the correct chemical species are formed.
In an acidic solution, MnO_4^- ions can be converted to Mn^{2+} ions, a process which is essential for this titration's success. An alkaline or neutral medium would alter the reaction pathway, potentially hindering the desired reaction.
Without the proper acidic conditions, complete reduction of permanganate may not occur, leading to inaccurate titration results. The acidic medium stabilizes the reaction, allowing oxalic acid to act as an effective reducing agent.

Formation of MnO2 precipitate

One of the visible signs that the acidic medium is insufficient during a titration is the formation of a brown precipitate. This precipitate is manganese dioxide ( MnO_2 ). When the medium is not acidic enough, MnO_4^- might not fully reduce to Mn^{2+} .
Instead, it partially reduces to MnO_2 , which is insoluble in the solution. The appearance of this brown precipitate can act as an indicator of improper titration conditions, implying that more sulfuric acid is needed to create the necessary acidic environment.
To avoid the formation of MnO_2 , ensure that the sulfuric acid is added in excess before starting the titration. This precaution helps to keep the reaction following the intended pathway and ensures reliable results.

Permanganate ion reduction

The reduction of permanganate ions (MnO_4^-) is a crucial aspect of the titration process. In a strong acidic medium, provided by H_2 S O_4, the MnO_4^- ions are reduced to Mn^{2+}. This reduction is what allows for the reaction with oxalic acid to proceed effectively.
The chemical equation for this reduction can be represented as:
\[MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O\]

• The presence of H^+ ions is critical here. They facilitate the transfer of electrons, turning MnO_4^- purple ions to Mn^{2+} nearly colorless ions.
• Completing this reaction correctly ensures that the oxidation of oxalic acid to carbon dioxide (CO_2) is also achieved.

This careful reduction in a properly maintained acidic environment ensures the titration's success and the reliability of the results.