The resistance of a cell containing 0.001 M KCl solution at 298 K is 1500 Ω. The conductivity is 0.146 × 10^–3 S cm^–1. The cell constant would be-

$1.$ $0.12$ ${\mathrm{cm}}^{-1}$

$2.$ $0.56$ ${\mathrm{cm}}^{-1}$

$3.$ $0.22$ ${\mathrm{cm}}^{-1}$

$4.$ $1.36$ ${\mathrm{cm}}^{-1}$

The amount of charge required for the reduction of 1 mol of ${\mathrm{MnO}}_4^{-}$ to ${\mathrm{Mn}}^{2+}$ is -

$1.$ $4.8$ $\times$ ${10}^5$ $\mathrm{C}$
$2.$ $3.2$ $\times$ ${10}^6$ $\mathrm{C}$
$3.$ $1.8$ $\times$ ${10}^5$ $\mathrm{C}$
$4.$ $4.1$ $\times$ ${10}^4$ $\mathrm{C}$

The number of Faradays required to produce 20.0 g of Ca from molten CaCl₂ is-

1. 2F

2. 1F

3. 4F

4. 3F

Three electrolytic cells A, B, C containing solutions of ZnSO₄, AgNO₃, and CuSO₄, respectively are connected in series.

A steady current of 1.5 amperes was passed through them until 1.45 g of silver was deposited at the cathode of cell B. The current flow time is-

1. 14 minutes

2. 25 minutes

3. 20 minutes

4. 11 minutes

The cell in which the following reactions occurs:

$2{{\mathrm{Fe}}^{3+}}_{\left(\mathrm{aq}\right)}+2{{\mathrm{I}}^{-}}_{\left(\mathrm{aq}\right)}\to 2{{\mathrm{Fe}}^{2+}}_{\left(\mathrm{aq}\right)}+{\mathrm{I}}_{2\left(\mathrm{s}\right)}$ has  \(E_{cell}^{o}\) = 0.236 V at 298 K.

The equilibrium constant of the cell reaction is : 

$1.$ $9.57$ $\times$ ${10}^7$
$2.$ $8.43$ $\times$ ${10}^6$
$3.$ $3.68$ $\times$ ${10}^{-7}$
$4.$ $1.74$ $\times$ ${10}^5$

The cell that will measure the standard electrode potential of a copper electrode is:

The electrode potential for Mg electrode varies according to the equation

\(E_{Mg^{2+}/Mg}\ = \ E_{Mg^{2+}/Mg}^{o} \ - \ \frac{0.059}{2}log\frac{1}{[Mg^{2+}]}\)

The graph of E_{Mg²⁺ / Mg} vs log [Mg²⁺] among the following is:

1.		2.
3.		4.

The correct statement among the following options is: 

The difference between the electrode potentials of two electrodes when no current is drawn through the cell is called:

The incorrect statement about an inert electrode in a cell is: