The conductivity of 0.00241 M acetic acid is 7.896 × 10^–5 S cm^–1. If ${\mathrm{\Lambda }}_{\mathrm{m}}^0$ for acetic acid is 390.5 S cm² mol^–1, the dissociation constant will be 
1. \(2.45 \times 10^{-5} \mathrm{~mol} \ \mathrm{~L}^{-1} \)
2. \(1.86 \times 10^{-5} \mathrm{~mol} \ \mathrm{L^{-1}} \)
3. \(3.72 \times 10^{-5}\mathrm{~mol} \mathrm{~L^{-1}} \)
4. \(2.12 \times 10^{-5}\mathrm{~mol} \mathrm{~L^{-1}}\)

A solution of Ni(NO₃)₂ is electrolysed between platinum electrodes using a current of 5 amperes for 20 minutes. What mass of Ni is deposited at the cathode?

1. 1.82 g

2. 1.95 g

3. 2.01 g

4. 1.16 g

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 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

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 : 

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