If the `DeltaG^(@)` of a cell reaction
`AgCl+e^(-) rarrAg+Cl^-` is `-21.20KJ`
standard emf of cell is

Options

A. 0.329V
B. 0.220V
C. `-0.220V`
D. `-0.110V`

Explanation

To find the standard EMF (E°) of the cell reaction given the standard Gibbs free energy change (ΔG°), we can use the following relationship:

\[ \Delta G° = -nFE° \]

Where:
- ΔG° = standard Gibbs free energy change (in joules)
- n = number of moles of electrons transferred in the reaction
- F = Faraday's constant (approximately 96500 C/mol)
- E° = standard EMF of the cell (in volts)

### Step-by-Step Solution:

1. **Identify the given values:**
- ΔG° = -21.20 kJ
- Convert ΔG° to joules:
\[ \Delta G° = -21.20 \, \text{kJ} \times 1000 \, \text{J/kJ} = -21200 \, \text{J} \]
- n = 1 (since one electron is involved in the reaction)

2. **Write the equation for ΔG°:**
\[ \Delta G° = -nFE° \]
Substituting the known values:
\[ -21200 = -1 \times 96500 \times E° \]

3. **Rearrange the equation to solve for E°:**
\[ E° = \frac{-21200}{-96500} \]

4. **Calculate E°:**
\[ E° = \frac{21200}{96500} \approx 0.220 \, \text{V} \]

5. **Final answer:**
The standard EMF of the cell is approximately 0.220 V.