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

Option (a) is correct; a brown precipitate of MnO2 is formed.

Step 1: Understanding the Reaction

The titration involves a redox reaction between oxalic acid ( H_2C_2O_4) and potassium permanganate (KMnO_4) in an acidic medium, usually sulfuric acid ( H_2SO_4). The KMnO_4 acts as an oxidizing agent.

Step 2: Analyzing the Role of H2SO4

The presence of H_2SO_4 provides the acidic environment necessary for the correct progression of the redox reaction, ensuring complete reduction of MnO_4^- ions to Mn^{2+} without side reactions.

Step 3: Consequence of Insufficient Acid

If H_2SO_4 is not in excess, the KMnO_4 is not fully reduced to Mn^{2+}. Instead, MnO_2 precipitates as a brown solid due to intermediate oxidation states of manganese.

Step 4: Identifying the Correct Outcome

The lack of sufficient H_2SO_4 results in the formation of a brown precipitate of MnO_2 instead of the clear solution when Mn^{2+} ions are formed. Hence, option (a) is correct, indicating a brown precipitate.

Key Concepts

Oxalic Acid

Oxalic acid is a naturally occurring organic compound that can be found in several plants like rhubarb and spinach. In chemical terms, it is represented by the formula \(\mathrm{H}_2\mathrm{C}_2\mathrm{O}_4\), often existing as a dihydrate \(\mathrm{H}_2\mathrm{C}_2\mathrm{O}_4\cdot 2\mathrm{H}_2\mathrm{O}\).
In a laboratory setting, oxalic acid is commonly used as a reducing agent. That means it has the ability to donate electrons to other compounds, effectively reducing them.

Oxalic acid is especially useful in redox titrations due to its predictable reactivity. Typically, it reacts with strong oxidizing agents, such as potassium permanganate, which is exactly what occurs in the titration exercise we are discussing. This reaction with potassium permanganate is performed under acidic conditions, normally provided by sulfuric acid. The oxalic acid donates electrons to the permanganate ion, reducing it from a higher oxidation state to a lower one.
Some of the key considerations when using oxalic acid in such titrations include:

• The need for a controlled acidic environment to ensure complete reactions.
• The measurement of both oxalic acid and the oxidizing agent to ensure a proper stoichiometric balance.
• Understanding that oxalic acid can react at higher temperatures leading to better reaction rates.

Potassium Permanganate

Potassium permanganate, known by its chemical formula \(\mathrm{KMnO}_4\), is a powerful oxidizing agent used extensively in redox reactions.
This compound is visually striking with its deep purple color, which helps indicate the endpoint of titrations.

In redox titrations, potassium permanganate is often employed as the titrant, especially when analyzing reducing agents like oxalic acid. Under acidic conditions, \(\mathrm{KMnO}_4\) reacts by gaining electrons from the reducing agent, converting from the permanganate ion \(\mathrm{MnO}_4^-\) to the manganese ion \(\mathrm{Mn}^{2+}\). This reaction depends heavily on the presence of an acidic medium like sulfuric acid to proceed efficiently.
The importance of using \(\mathrm{KMnO}_4\) in titration includes:

• The need for the solution to remain acidic; otherwise, incomplete reduction leads to unwanted by-products like \(\mathrm{MnO}_2\), as seen in the exercise.
• The visually noticeable color change from purple to colorless signaling the titration endpoint.
• Its strong oxidizing nature, which ensures the rapid oxidation of the reducing agent.

Sulfuric Acid

Sulfuric acid, denoted as \(\mathrm{H}_2\mathrm{SO}_4\), is a strong mineral acid widely used in chemical reactions. It plays a crucial role in redox titrations involving compounds like oxalic acid and potassium permanganate.
In this context, sulfuric acid facilitates the setting of an acidic environment necessary for the redox reaction to progress smoothly and completely.

During a titration process, sulfuric acid serves several important functions, such as:

• Providing a medium that helps stabilize the permanganate ions, facilitating their reduction to manganese ions.
• Preventing the formation of precipitates like \(\mathrm{MnO}_2\), which occurs if the acidity is insufficient, as noted in the original exercise.
• Helping in the dissociation of the reactants, ensuring a greater number of effective collisions and faster reaction rates.

The absence or insufficient amount of sulfuric acid risks incomplete reduction of \(\mathrm{MnO}_4^-\) ions, leading to side reactions and useless precipitates. Ensuring sulfuric acid is in excess is essential for a successful reaction outcome, as emphasized in the original problem solution.