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

If a current of 0.5 ampere flows through a metallic wire for 2 hours, then how many electrons would flow through the wire?

$1. 2.35 \times {10}^{-22}$
$2. 2.25 \times {10}^{22}$
$3. 1.34 \times {10}^{21}$
$4. 3.16 \times {10}^{24}$

What is the quantity of electricity in coulombs needed to reduce 1 mol of ${\mathrm{Cr}}_2{\mathrm{O}}_7^{2-}$? Consider the reaction: ${\mathrm{Cr}}_2{\mathrm{O}}_7^{2-} + 14{\mathrm{H}}^{+} +\hspace{0.17em}6{\mathrm{e}}^{-} \to 2{\mathrm{Cr}}^{+3 }+ 7{\mathrm{H}}_2\mathrm{O}$

1. 578992 C

2. 289461 C

3. 192974 C

4. 96487 C

What is the potential of hydrogen electrode in contact with a solution whose pH is 10?

1. 0.591 V

2. -0.591 V

3. 0.295 V

4. -0.295 V

What will be the emf of the cell in which the following reaction takes place? Given ${\mathrm{E}}_{\mathrm{cell}}^{°} = 1.05 \mathrm{V}$

${\mathrm{Ni}}_{\left(\mathrm{s}\right)}+2{\mathrm{Ag}}^{+}(0.002 \mathrm{M}) \to {\mathrm{Ni}}^{2+}(0.160 \mathrm{M}) + 2{\mathrm{Ag}}_{\left(\mathrm{s}\right)}$

1. 0.80 V

2. 0.91 V

3. 0.45 V

4. 0.36 V

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 molar conductivity of 0.025 mol L⁻¹ methanoic acid is 46.1 S cm² mol⁻¹.

What is the dissociation constant of methanoic acid? Given λ°(H⁺)= 349.6 S cm² mol⁻¹ and λ°(HCOO⁻) = 54.6 S cm² mol^-1

^{$1. 2.43 \times {10}^4 \mathrm{mol} {\mathrm{L}}^{-1}$
$2. 1.15 \times {10}^{-5} \mathrm{mol} {\mathrm{L}}^{-1}$
$3. 3.67 \times {10}^{-4} \mathrm{mol} {\mathrm{L}}^{-1}$
$4. 2.44 \times {10}^{-6 }\mathrm{mol} {\mathrm{L}}^{-1}$}

Mg_(s) | Mg²⁺_(0.001M) || Cu²⁺_{(0.0001 M)} | Cu_(s)

${\mathrm{E}}_{{\mathrm{Mg}}^{2+}/\mathrm{Mg}}^{\mathrm{o}}=-2.36 \mathrm{V}; {\mathrm{E}}_{{\mathrm{Cu}}^{2+}/\mathrm{Cu}}^{\mathrm{o}}=0.34\mathrm{V}$

The value of ${\mathrm{E}}_{\mathrm{cell} }$ for the above reaction is -

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.