$Emf$ of the following cell at $298 \,K$ in $V$ is $x \times 10^{-2}$. $Zn | Zn^{2+}(0.1 \,M) || Ag^{+}(0.01 \,M) | Ag$. The value of $x$ is .... . (Rounded off to the nearest integer) [Given: $E^{0}_{Zn^{2+}/Zn} = -0.76 \,V$; $E^{0}_{Ag^{+}/Ag} = +0.80 \,V$; $\frac{2.303 RT}{F} = 0.059$]

The cell potential for the following cell notation is approximately
$M_{(s)} | M^{3+}(aq, 0.01 \ M) || N^{2+}(aq, 0.1 \ M) | N_{(s)}$
$E_{M^{3+} / M}^0 = 0.6 \ V$ and $E_{N^{2+} / N}^0 = 0.1 \ V$ (in $V$)

Consider the electrochemical cell shown in the figure where a metal electrode $(M)$ undergoes a redox reaction by forming $M^{+}$ $(M \rightarrow M^{+} + e^{-})$. The cation $M^{+}$ is present in two different concentrations $c_{1}$ and $c_{2}$. Which of the following statements is correct for generating a positive cell potential?

The $emf$ of a $Daniell$ cell at $298 \, K$ is $E_1$:
$Zn | ZnSO_4 \,(0.01 \, M) || CuSO_4 \,(1.0 \, M) | Cu$
When the concentration of $ZnSO_4$ is $1.0 \, M$ and that of $CuSO_4$ is $0.01 \, M$,the $emf$ changes to $E_2$. What is the relation between $E_1$ and $E_2$?

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 pressure of $H_2$ required to make the potential of $H_2$-electrode zero in pure water at $298 \ K$ is