Question

In the cell $Ag|A{g}^{+}||C{u}^{2+}|Cu$ containing 1 M $C{u}^{2+}$ and 1 M $A{g}^{+}$ ions, 9.65 ampere current is passed for 1 hour. Select the correct statement.

• A. Ag will oxidise to $A{g}^{+}$ and new $\left[A{g}^{+}\right]=1.36M$
• B. $A{g}^{+}$ will reduce to Ag and new $\left[A{g}^{+}\right]=0.64M$
• C. $C{u}^{2+}$ will reduce to Cu and new $\left[C{u}^{2+}\right]=0.82M$
• D. Cu will oxidise to $C{u}^{2+}$ and new $\left[C{u}^{2+}\right]=0.82M$

Answer 1

Answer: (A), (C)

Ag will oxidise to $A{g}^{+}$ and new $\left[A{g}^{+}\right]=1.36M$
$C{u}^{2+}$ will reduce to Cu and new $\left[C{u}^{2+}\right]=0.82M$

On passing the current through the given cell Ag will oxidise to Ag${}^{+}$ and Cu${}^{2+}$ will be reduced to Cu.

As same charge is passing through both solutions.

According to Faraday's law:

$\frac{w}{E}=\frac{i.t}{96500}=\frac{9.65\times 3600}{96500}$

$=0.36$ eq. of $A{g}^{+}=0.36$ mol of $A{g}^{+}$

$=0.36$ eq. of $C{u}^{2+}$

$=\frac{0.36}{2}$ mole of $C{u}^{2+}$ (as mole $=$ equivalent/valency)

$\therefore \left[A{g}^{+}\right]=1+0.36=1.36M$

$\left[C{u}^{2+}\right]=1.0-0.18=0.82M$