$A$ constant current $(0.5 \, A)$ is passed for $1 \, hour$ through $(i)$ aqueous $AgNO_3$,$(ii)$ aqueous $CuSO_4$ and $(iii)$ molten $AlF_3$,separately. The ratio of the mass of the metals deposited on the cathode is $[M_{Ag}, M_{Cu}, M_{Al}$ are molar masses of the respective metals.]

In order to separate oxygen from one mole of $H_2O$,the required quantity of electricity in coulombs is:

$CuSO_4$ solution is electrolysed for $15 \ minutes$ to deposit $0.4725 \ g$ of copper at the cathode. The current in amperes required is (Faraday $= 96,500 \ C \ mol^{-1}$,atomic weight of copper $= 63$)

$A$ 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 $STP = 22.4 \ L$]

$0.50 \, L$ of a $0.60 \, M$ $CuSO_4$ solution is electrolyzed for a period of $30.0 \, min$ using a current of $4.60 \, A$. If inert electrodes are used,what is the final concentration of $Cu^{2+}$ remaining in the solution? $M$