Question

$\left(a\right)$ A current of $1.50A$ was passed through an electrolytic cell containing ${AgNO}_3$ solution with inertelectrodes. The weight of $Ag$ deposited was $1.50g$. How long did the current flow?
$\left(b\right)$ Write the reactions taking place at the anode and cathode in all the above cell.
$\left(c\right)$ Give reactions taking place at the two electrodes if these are made up of $Ag$.

Answer 1

$\left(a\right)$ Reaction involved in deposition of $Ag$ on electrode,

${Ag}^{+}+e\to {Ag}_{\left(s\right)}$

$1$ mol of $e^{-}$ is required for the reduction of $1$ mol of ${Ag}^{+}$, i.e., $1$ mol of ${Ag}^{+}$ require $1$ mol of $e^{-}\left(1F\right)$

$1F$ is required to deposit $108kg$ of $Ag$.

Deposited $Ag=1.5g$

So, if $108g$ $Ag\to 1F=96500C$ of charge

$1.5g$ $Ag\to \frac{96500}{108}\times 1.5C$

$=1340.277C$.

We know,

$Q=IK$

$t=\frac{Q}{I}=\frac{1340.277}{1.5}=893.518sec$

$\left(b\right)$ Reaction at Anode,

${2H}_{2}O\to O_2+{4H}^{+}+{4e}^{-}$

Reaction at Cathode,

${Ag}_{\left(ag\right)}^{+}+e^{-}\to {Ag}_{\left(s\right)}$

$\left(c\right)$ If both electrodes are of $Ag$, then at anode, $Ag$ is converted to ${Ag}^{+}$ ions

${Ag}_{\left(s\right)}\to {Ag}^{+}+e^{-}$

atcathode, $Ag$ is deposited

${Ag}^{+}+e^{-}\to {Ag}_{\left(s\right)}$

Thus $Ag$ is dissolved at anode and deposited at cathode.