$A$ galvanic cell is set up from a zinc bar weighing $50 \ g$ and $1.0 \ L$,$1.0 \ M$ $CuSO_4$ solution. How long would the cell run,assuming it delivers a steady current of $1.0 \ A$? (Answer in $hrs$)

Consider the cell reaction at $300 \ K$: $A_{(s)} + B^{2+}_{(aq)} \rightleftharpoons A^{2+}_{(aq)} + B_{(s)}$. Its $E^{\circ}$ is $1.0 \ V$. The $\Delta_{r}H^{\circ}$ of the reaction is $-163 \ kJ \ mol^{-1}$. What is $\Delta_{r}S^{\circ}$ (in $J \ K^{-1} \ mol^{-1}$) of the reaction? $(F = 96500 \ C \ mol^{-1})$

$A$ negative $e.m.f.$ for a cell represents:

The standard reduction potentials of $2H^{+}/H_2$,$Cu^{2+}/Cu$,$Zn^{2+}/Zn$,and $NO_3^{-}, H^{+}/NO$ are $0.0 \ V$,$0.34 \ V$,$-0.76 \ V$,and $0.97 \ V$ respectively. Observe the following reactions:
$I$. $Zn + HCl \rightarrow$
$II$. $Cu + HCl \rightarrow$
$III$. $Cu + HNO_3 \rightarrow$
Which reactions do not liberate $H_{2(g)}$?

Calculate the equilibrium constant at $298 \ K$ for the following reaction and also calculate the maximum work that can be obtained from this cell:
$Mg(s) \ | \ Mg^{2+}(aq) \ || \ Ag^{+}(aq) \ | \ Ag(s)$
Given: $E_{Mg^{2+} \mid Mg}^{o} = -2.37 \ V$ and $E_{Ag^{+} \mid Ag}^{o} = 0.80 \ V$

The rusting of iron takes place as follows. Calculate $\Delta G^o$ for the net process in $kJ \ mol^{-1}$.
$2H^{+} + 2e^- + \frac{1}{2}O_2 \longrightarrow H_2O_{(l)} ; E^o = +1.23 \ V$
$Fe^{2+} + 2e^- \longrightarrow Fe_{(s)} ; E^o = -0.44 \ V$