The e.m.f. of the cell $Ag | Ag^{+}(0.1 \ M) || Ag^{+}(1 \ M) | Ag$ at $298 \ K$ is ........... $V$.

For the reaction $A_{(s)} + 2B^{+}_{(aq)} \rightarrow A^{2+}_{(aq)} + 2B_{(s)}$,the value of $K_c$ is $10^{12}$. What is the value of $E^o_{cell}$ (in $, V$)?

For a cell reaction involving a two-electron change,the standard emf of the cell is found to be $0.295 \ V$ at $25 \ ^oC$. The equilibrium constant of the reaction at $25 \ ^oC$ will be

At $298 \ K$,the standard reduction potentials are $1.51 \ V$ for $MnO_4^- \ | \ Mn^{2+}$,$1.36 \ V$ for $Cl_2 \ | \ Cl^{-}$,$1.07 \ V$ for $Br_2 \ | \ Br^{-}$ and $0.54 \ V$ for $I_2 \ | \ I^{-}$. At $pH = 3$,permanganate is expected to oxidize $\left( \frac{RT}{F} = 0.059 \ V \right)$

What must be the concentration of $Ag^{+}$ in an aqueous solution containing $Cu^{2+} = 1.0 \ M$ so that both the metals can be deposited on the cathode simultaneously? Given that $E^0_{Cu^{2+}/Cu} = 0.34 \ V$ and $E^0_{Ag^{+}/Ag} = 0.812 \ V$ at $T = 298 \ K$.

At $298 \ K$,the following reaction takes place for a hydrogen electrode:
$H^{+}_{(aq)} + e^{-} \longrightarrow \frac{1}{2} H_2(1 \ bar)$
The solution $pH$ is $10.0$. What is the hydrogen electrode potential in volts?
$\left(\frac{2.303 RT}{F} = 0.06 \ V\right)$