ElectrochemistryHard
Question
The EMF for the cell: Ag(s)|AgCl(s)|KCl(0.2 M)||KBr (0.001 M)|AgBr(s)|Ag(s) at 25°C is (Ksp(AgCl) = 2.0 × 10−10; Ksp(AgBr) = 4.0 × 10−13, 2.303 RT/F = 0.06, log 2 = 0.3)
Options
A.0.024 V
B.−0.024 V
C.−0.24 V
D.− 0.012 V
Solution
Assuming the cell as concentration cell, the cell reaction may be written as
$Ag^{+}\left( C_{1} = \frac{4 \times 10^{- 13}}{0.001} = 4 \times 10^{- 10}\text{ M} \right) \rightarrow Ag^{+}\left( C_{2} = \frac{2 \times 10^{- 10}}{0.2} = 10^{- 9}\text{ M} \right)$
Now, $E_{cell} = 0 - \frac{0.06}{1}.\log\frac{10^{- 9}}{4 \times 10^{- 10}} = - 0.024\text{ V}$
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