The standard $emf$ of a galvanic cell involving $2$ moles of electrons in its redox reaction is $0.59 \ V$. The equilibrium constant for the redox reaction of the cell is

The cell potential for the given cell at $298 \, K$ is $Pt \mid H_2 (g, 1 \, bar) \mid H^{+}_{(aq)} \parallel Cu^{2+}_{(aq)} \mid Cu_{(s)}$. The cell potential is $0.31 \, V$. The $pH$ of the acidic solution is $3$,and the concentration of $Cu^{2+}$ is $10^{-x} \, M$. The value of $x$ is $.....$ (Given: $E^{\ominus}_{Cu^{2+}/Cu} = 0.34 \, V$ and $\frac{2.303 RT}{F} = 0.06 \, V$)

Which of the following is the correct representation of the Nernst equation for the reduction reaction $M^{n+} + ne^- \rightarrow M$?

Calculate the $emf$ of the cell in which the following reaction takes place:
$Ni_{(s)} + 2Ag^{+}(0.002 \, M) \rightarrow Ni^{2+}(0.160 \, M) + 2Ag_{(s)}$
Given that $E^{\Theta}_{(cell)} = 1.05 \, V$ (in $, V$)

For the cell reaction,$3 Sn^{4+} + 2 Cr \longrightarrow 3 Sn^{2+} + 2 Cr^{3+}$,$E^{\circ}_{cell}$ is $0.89 \ V$. Then $\Delta G^{\circ}$ for the reaction is

The $EMF$ of a hydrogen electrode in terms of $pH$ is (at $1 \ atm$ pressure).