Calculate the standard potential of the cell ,If the standard electrode potentials of `Zn^(2+)//Zn and Ag^(+) //Ag` are -0.763 V and + 0.799 V respectively .

Options

A. 0.036 V
B. 1.56 V
C. `-1.562V`
D. 0.799V

Answer

To calculate the standard potential of the cell, we will use the standard electrode potentials provided for the half-reactions involving zinc and silver.

### Step-by-Step Solution:

1. **Identify the Standard Electrode Potentials:**
- For the reduction of silver ions:
\[ \text{Ag}^+ + e^- \rightarrow \text{Ag} \quad (E^\circ = +0.799 \, \text{V}) \]
- For the reduction of zinc ions:
\[ \text{Zn}^{2+} + 2e^- \rightarrow \text{Zn} \quad (E^\circ = -0.763 \, \text{V}) \]

2. **Determine the Anode and Cathode:**
- The half-reaction with the higher standard reduction potential acts as the cathode (where reduction occurs), and the half-reaction with the lower standard reduction potential acts as the anode (where oxidation occurs).
- Here, silver (Ag) has a higher potential (+0.799 V) than zinc (Zn) (-0.763 V), so:
- **Cathode:** Ag (Ag^+/Ag)
- **Anode:** Zn (Zn^2+/Zn)

3. **Use the Formula for Standard Cell Potential:**
The standard cell potential (E°cell) can be calculated using the formula:
\[ E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} \]

4. **Substitute the Values:**
\[ E^\circ_{\text{cell}} = E^\circ_{\text{Ag}} - E^\circ_{\text{Zn}} = (+0.799 \, \text{V}) - (-0.763 \, \text{V}) \]

5. **Calculate the Standard Cell Potential:**
\[ E^\circ_{\text{cell}} = 0.799 + 0.763 = 1.562 \, \text{V} \]

### Final Answer:
The standard potential of the cell is **1.562 V**.