For the cell,$Mn_{(s)}|Mn_{(aq)}^{2+}(0.4\,M)||Sn_{(aq)}^{2+}(0.04\,M)|Sn_{(s)}$,calculate the free energy change $(\Delta G)$ at $298\,K$ in $kJ$.
Given: $E_{Mn^{2+}|Mn}^o = -1.18\,V$; $E_{Sn^{2+}|Sn}^o = -0.14\,V$; $\frac{2.303\,RT}{F} = 0.06$

What is the oxidation potential of $0.05 \, M \, H_2SO_4$ in volts?

Calculate the equilibrium constant $(K_c)$ of the reaction: $Ni_{(s)} + 2Ag_{(aq)}^{+} \rightarrow Ni_{(aq)}^{2+} + 2Ag_{(s)}$; $E_{cell}^{\circ} = 1.05 \ V$. (Given: $\frac{2.303 \ RT}{F} = 0.06$)

For the redox reaction occurring in a cell: $Zn_{(s)} + Cu^{2+}(0.1 \ M) \to Zn^{2+}(1 \ M) + Cu_{(s)}$,if $E^o_{cell} = 1.10 \ V$,calculate the value of $E_{cell}$ in $V$. (Given: $2.303 \frac{RT}{F} = 0.0591$)

$Cu^{2+} + 2e^- \to Cu$. On increasing $[Cu^{2+}]$ concentration,electrode potential

One half cell in a voltaic cell is constructed by dipping a silver rod in an $AgNO_3$ solution of unknown concentration,and the other half cell is a $Zn$ rod dipped in a $1 \text{ M}$ solution of $ZnSO_4$. $A$ voltage of $1.60 \text{ V}$ is measured at $298 \text{ K}$ for this cell. What is the concentration of $Ag^+$ ions in terms of $\log x$ (where $x = [Ag^+]$)? Given: $E^\ominus_{Zn^{2+}/Zn} = -0.76 \text{ V}$,$E^\ominus_{Ag^+/Ag} = +0.80 \text{ V}$,and $\frac{2.303RT}{F} = 0.059 \text{ V}$.