The number of faradays required to reduce one mole of `Cu^(2+)` to metallic copper is

Options

A. one
B. two
C. three
D. four

Answer

To solve the problem of determining the number of Faradays required to reduce one mole of \( \text{Cu}^{2+} \) to metallic copper, we can follow these steps:

### Step 1: Understand the Reduction Reaction
The reduction reaction for copper ions can be represented as:
\[ \text{Cu}^{2+} + 2e^- \rightarrow \text{Cu} \]
This equation shows that one mole of \( \text{Cu}^{2+} \) ions requires two moles of electrons to be reduced to one mole of copper metal.

### Step 2: Determine the Number of Electrons Required
From the balanced equation, we see that:
- 1 mole of \( \text{Cu}^{2+} \) requires 2 moles of electrons (2 \( e^- \)) for reduction.

### Step 3: Relate Electrons to Faradays
One Faraday (F) is defined as the charge of one mole of electrons, which is approximately \( 96485 \) coulombs. Therefore, to reduce 2 moles of electrons, we need:
\[ \text{Number of Faradays} = \text{Number of moles of electrons} \]
Since we need 2 moles of electrons, we require 2 Faradays.

### Step 4: Conclusion
Thus, the number of Faradays required to reduce one mole of \( \text{Cu}^{2+} \) to metallic copper is:
\[ \text{Answer} = 2 \text{ Faradays} \]

### Final Answer
The number of Faradays required to reduce one mole of \( \text{Cu}^{2+} \) to metallic copper is **2**.

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