$A$ current of $15.0 \ A$ is passed through a solution of $CrCl_2$ for $45 \ min$. The volume of $Cl_2$ (in $L$) obtained at the anode at $1 \ atm$ and $273 \ K$ is around ($1 \ F = 96500 \ C \ mol^{-1}$,At. wt. of $Cl=35.5$,$R=0.082 \ L \ atm \ K^{-1} \ mol^{-1}$) (in $.7$)

Electricity is passed through an acidic solution of $Cu^{2+}$ until all the $Cu^{2+}$ is exhausted, leading to the deposition of $300 \ mg$ of $Cu$ metal. Subsequently, a current of $600 \ mA$ is passed through the same solution for another $28 \ minutes$ while keeping the total volume of the solution fixed at $200 \ mL$. The total volume of oxygen evolved at $STP$ during the entire process is . . . . . . $mL$. (Nearest integer) [Given: $Cu^{2+}_{(aq)} + 2e^{-} \rightarrow Cu_{(s)}$, $E_{red}^0 = +0.34 \ V$; $O_{2(g)} + 4H^{+} + 4e^{-} \rightarrow 2H_2O$, $E_{red}^0 = +1.23 \ V$; Molar mass of $Cu = 63.54 \ g \ mol^{-1}$; Faraday Constant $= 96500 \ C \ mol^{-1}$; Molar volume at $STP = 22.4 \ L$]

$5 \ A$ current is passed through a solution of zinc sulphate for $40 \ \text{minutes}$. Find the amount of zinc deposited at the cathode in $gm$.

The amount of copper metal deposited at the cathode on passing an electric current of $500 \, mA$ for $20 \, minutes$ in a cupric chloride solution is ............ $g$.

The time required (in hours) to reduce $3 \ mol$ of $Fe^{3+}$ ions to $Fe^{2+}$ ions with $2.0 \ A$ of current is $\left(1 \ F = 96500 \ C \ mol^{-1}\right)$ (in $.2$)

How much charge in coulombs is required for the reduction of one mole of $Al^{3+}$ to $Al$?