The standard electrode potential a `Ag^(+)//Ag` is +0.80 V and of `Cu^(2+)//Cu` is +0.34 V. These electrodes are connected through a salt bridge and if :

Options

A. copper electrode acts as cathode, then `E^(∘)cell` is +0.46 volt
B. Silver electrode acts as anode, then `E^(∘)cell` is -0.34 volt
C. Copper electrode acts as anode, then `E^(∘)cell` is +0.46 volt
D. Silver electrode acts as cathode, then `E^(∘)cell` is -0.34 volt

Explanation

To solve the problem, we need to analyze the given standard electrode potentials and determine the reactions that will occur when the two electrodes are connected through a salt bridge.

### Step-by-Step Solution:

1. **Identify the Standard Electrode Potentials**:
- The standard electrode potential for the silver half-reaction \( \text{Ag}^+ + e^- \rightarrow \text{Ag} \) is \( +0.80 \, \text{V} \).
- The standard electrode potential for the copper half-reaction \( \text{Cu}^{2+} + 2e^- \rightarrow \text{Cu} \) is \( +0.34 \, \text{V} \).

2. **Determine Which Electrode Will Be Reduced and Which Will Be Oxidized**:
- The half-reaction with the higher standard electrode potential will be reduced, while the one with the lower potential will be oxidized.
- Since \( +0.80 \, \text{V} \) (Ag) is greater than \( +0.34 \, \text{V} \) (Cu), silver ions will be reduced to silver metal, and copper metal will be oxidized to copper ions.

3. **Write the Half-Reactions**:
- Reduction at the cathode (where reduction occurs):
\[ \text{Ag}^+ + e^- \rightarrow \text{Ag} \quad (E^\circ = +0.80 \, \text{V}) \]
- Oxidation at the anode (where oxidation occurs):
\[ \text{Cu} \rightarrow \text{Cu}^{2+} + 2e^- \quad (E^\circ = +0.34 \, \text{V}) \]

4. **Determine the Overall Cell Reaction**:
- To balance the number of electrons, we can multiply the silver half-reaction by 2:
\[ 2\text{Ag}^+ + 2e^- \rightarrow 2\text{Ag} \]
- The overall cell reaction becomes:
\[ 2\text{Ag}^+ + \text{Cu} \rightarrow 2\text{Ag} + \text{Cu}^{2+} \]

5. **Calculate the Standard Cell Potential (\(E^\circ_{\text{cell}}\))**:
- The standard cell potential is calculated using the formula:
\[ E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} \]
- Substituting the values:
\[ E^\circ_{\text{cell}} = 0.80 \, \text{V} - 0.34 \, \text{V} = 0.46 \, \text{V} \]

6. **Identify the Anode and Cathode**:
- The copper electrode acts as the anode (oxidation occurs here).
- The silver electrode acts as the cathode (reduction occurs here).

### Conclusion:
- The copper electrode acts as the anode, and the overall cell potential is \( +0.46 \, \text{V} \).