Question

How many electrons are involved in the following redox reaction? \[ \text{Cr}_2\text{O}_7^{2-} + \text{Fe}^{2+} + \text{C}_2\text{O}_4^{2-} \rightarrow \text{Cr}^{3+} + \text{Fe}^{3+} + \text{CO}_2 \quad (\text{Unbalanced}) \] (A) \( 3 \)

• A. \( 3 \)
• B. \( 4 \)
• C. \( 6 \)
• D. \( 5 \)

Hint:

When balancing redox reactions, determine oxidation states and use the half-reaction method to balance the number of electrons transferred.

Answer 1

The Correct Option is A

Step 1: Identify the Oxidation and Reduction Half-Reactions
The given reaction consists of multiple species undergoing oxidation and reduction: 1. Reduction Half-Reaction (Chromium): \[ \text{Cr}_2\text{O}_7^{2-} \rightarrow \text{Cr}^{3+} \] Chromium changes from \( +6 \) (in \( \text{Cr}_2\text{O}_7^{2-} \)) to \( +3 \) (in \( \text{Cr}^{3+} \)). Since each chromium atom gains 3 electrons, and there are 2 Cr atoms, the total electrons gained: \[ 2 \times 3 = 6 \text{ electrons}. \] 2. Oxidation Half-Reaction (Iron): \[ \text{Fe}^{2+} \rightarrow \text{Fe}^{3+} \] Iron changes from +2 to +3, meaning it loses 1 electron per Fe atom. 3. Oxidation Half-Reaction (Oxalate Ion): \[ \text{C}_2\text{O}_4^{2-} \rightarrow 2 \text{CO}_2 \] Each carbon in \( \text{C}_2\text{O}_4^{2-} \) changes from +3 to +4, losing 1 electron per carbon atom. Since there are 2 carbon atoms, the total electrons lost: \[ 2 \times 1 = 2 \text{ electrons}. \]
Step 2: Balancing the Electrons
The total electrons gained in the reduction step = 6.
The total electrons lost = 1 (Fe) + 2 (C) = 3. To balance the loss and gain, we need 3 electrons. \[ \text{Total electrons involved} = 3. \] Final Answer: The total number of electrons involved in the redox reaction is 3.