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 \(\text{MnO}_2\) is formed.

Step 1: Identify the Reaction

Recognize that the reaction involved is a redox titration between oxalic acid \((\text{H}_2\text{C}_2\text{O}_4)\) and potassium permanganate \((\text{KMnO}_4)\). The acidic environment is typically provided by sulfuric acid \((\text{H}_2\text{SO}_4)\).

Step 2: Understand the Role of the Acid

Sulfuric acid \((\text{H}_2\text{SO}_4)\) provides the acidic medium necessary for \(\text{MnO}_4^-\) ions to be reduced to \(\text{Mn}^{2+}\) ions, which remain in solution and are colorless. If the acid is not in excess, complete reduction does not occur.

Step 3: Identify Effect of Insufficient Acid

Without sufficient sulfuric acid, the partial reduction of \(\text{MnO}_4^-\) leads to the formation of \(\text{MnO}_2\), a brown solid, instead of \(\text{Mn}^{2+}\) ions.

Step 4: Conclusion

The absence of excess \(\text{H}_2\text{SO}_4\) during the titration results in the formation of \(\text{MnO}_2\), a brown precipitate, indicating incomplete reduction of permanganate ions.

Key Concepts

Oxalic Acid

Oxalic acid, chemically known as \( ext{H}_2 ext{C}_2 ext{O}_4\), plays a crucial role in many chemical reactions, including redox titrations. It is a relatively simple organic acid with two hydrogen ions available for donation, making it capable of undergoing ionization to form its conjugate base.
Redox titrations with oxalic acid involve the transfer of electrons. The reducing nature of oxalic acid allows it to donate electrons to oxidizing agents like potassium permanganate. In these reactions, oxalic acid is oxidized to carbon dioxide, generating two \( ext{CO}_2\) molecules. This change is balanced by the reduction of the oxidizing agent, such as \( ext{KMnO}_4\), from a higher oxidation state to a lower one.
The process of titration typically monitors the progress of the reaction by the changes in the color of the solution, which is critical in identifying the endpoint of the titration.

Potassium Permanganate

Potassium permanganate, \( ext{KMnO}_4\), is known for its strong oxidizing properties. In aqueous solutions, it appears as a vibrant purple liquid due to the permanganate ions \( ext{MnO}_4^-\).
During redox titrations, \( ext{KMnO}_4\) acts as an oxidizing agent. It accepts electrons and gets reduced primarily to \( ext{Mn}^{2+}\) ions, which are colorless, when in an acidic environment.

Sulfuric Acid

Sulfuric acid \( ext{H}_2 ext{SO}_4\) is an essential component in redox titrations involving \( ext{KMnO}_4\). Its job is to provide a strong acidic medium necessary for the reduction of \( ext{MnO}_4^-\) ions.
Without enough sulfuric acid, this reduction cannot proceed effectively, leading to the incomplete reaction and formation of unwanted by-products such as \( ext{MnO}_2\). Sulfuric acid helps maintain the acidic conditions needed to ensure the permanganate ions fully convert to \( ext{Mn}^{2+}\) ions, preventing the formation of any visible precipitate.
The features of sulfuric acid, such as its strong ability to dissociate and produce hydrogen ions, strengthen its role in maintaining the low pH needed for these reactions.

MnO2 Formation

The formation of manganese dioxide, \( ext{MnO}_2\), is a critical concept to understand in redox titrations when conditions are not ideal. When sulfuric acid is not used in excess during the titration of oxalic acid with potassium permanganate, the environment may not stay sufficiently acidic. This causes the incomplete reduction of permanganate ions, \( ext{MnO}_4^-\). In this scenario, instead of reducing completely to the colorless \( ext{Mn}^{2+}\) ions, the permanganate ions form \( ext{MnO}_2\), which precipitates as a brown solid.
The formation of this precipitate is a visual cue indicating that the titration conditions are not optimal and that sulfuric acid was not present in sufficient amounts. By understanding these concepts, students can better adjust experimental conditions to avoid unintended side products.