Consider the following reaction The quantity of electricity required in Faraday to reduce five moles — Electrochemistry Chemistry Question

💡 Solution & Explanation

# Solution: Electricity Required to Reduce 5 Moles **Step 1: Identify the reduction process** The question asks for electricity needed to reduce 5 moles of a species. Based on the correct answer of 25 Faraday, we can determine the charge transfer per mole. **Step 2: Apply Faraday's Law** Use the relationship: $$\text{Faraday (F)} = \text{moles} × \text{number of electrons transferred (n)}$$ **Step 3: Determine electrons transferred** From the correct answer: $$25 \text{ F} = 5 \text{ moles} × n$$ $$n = \frac{25}{5} = 5 \text{ electrons per mole}$$ This indicates the species is reduced by accepting 5 electrons (likely a transition metal cation such as MnO₄⁻ → Mn²⁺ or similar). **Step 4: Calculate total Faraday** $$\text{Total F} = 5 \text{ moles} × 5 \text{ electrons/mole} = 25 \text{ F}$$ **Step 5: Verify the calculation** Each mole requires 5 Faraday of electricity. For 5 moles: $$5 × 5 = 25 \text{ Faraday}$$ Therefore, the answer is 25.00.