Salt with the highest electrolytic conductivity in solution is : 

1. K₂[PtCl₆] 

2. [Co(NH₃)₃(NO₂)₃]

3. K₄[Fe(CN)₆] 

4. [Co(NH₃)₄]SO₄

If hydrogen electrodes dipped in two solutions of pH = 4 and pH = 6 are connected by a salt bridge, the emf of the resulting cell is:

1. 0.177 V               

2. 0.3 V               

3. 0.118 V             

4. 0.104 V

Calculate the emf of the given cell: 

Zn(s) | Zn⁺² (0.1M) || Sn⁺² (0.001M) | Sn(s)

(Given ${\mathrm{E}}_{{\mathrm{Zn}}^{+2}/\mathrm{Zn}}^{\mathrm{o}}=-0.76$ $\mathrm{V},$ ${\mathrm{E}}_{{\mathrm{Sn}}^{2+}/\mathrm{Sn}}^{\mathrm{o}}=-0.14$ $\mathrm{V)}$

1. 0.62 V

2. 0.56 V

3. 1.12 V

4. 0.31 V

What is the standard cell potential (E°cell) for the reaction A(s) + 2B⁺ → A²⁺ + 2B(s), given that the equilibrium constant (Kc = equilibrium constant) is 10¹²?

Limiting molar conductivities, for the given solutions, are:

\(\lambda_{m}^{0} \left(\right. H_{2} S O_{4} \left.\right) = x\) \(S  c m^{2}\) \(m o l^{- 1}\)

\(\lambda_{m}^{0} \left(\right. K_{2} S O_{4} \left.\right) = y\) \(S  c m^{2}\) \(m o l^{- 1}\)

\(\lambda_{m}^{0} \left(\right. C H_{3} C O O K \left.\right) = z\) \(S  c m^{2}\) \(m o l^{- 1}\)

From the data given above, it can be concluded that \(\lambda_m^0 \) in (\(S\ cm^2\ mol^{-1}\)) for CH₃COOH will be:

Calculate the standard cell potential, E°cell, in volts and the standard Gibbs free-energy change, ΔG°, in kJ mol⁻¹, respectively, for a two-electron redox reaction at 298 K having an equilibrium constant of 3.8 × 10⁻³.

The specific conductance of 0.01 M solution of a weak monobasic acid is 0.20 x 10^-3 S cm^-1. The dissociation constant of the acid is:

[Given  ${\Lambda }_{\mathrm{HA}}^{\infty }$ = 400 S ${\mathrm{cm}}^2$ ${\mathrm{mol}}^{-1}$]

For a cell involving one electron \(E_{cell}^{\ominus} = 0 . 59  V\) at 298 K. The equilibrium constant for the cell reaction is :
\(\mathrm{[Given~ that~ \frac {2.303 ~RT}{F} = 0.059 ~V~ at~ T = 298 K]}\)

Given the following cell reaction:
 \(\mathrm{2Fe^{3+}(aq) \ + \ 2I^{-}(aq)\rightarrow 2Fe^{2+}(aq) \ + \ I_{2}(aq)}\)

[Given: \(F  = 96500\) \(C\) \(mol^{- 1}\)]

1. \(23 . 16\) \(kJ\) \(mol^{- 1}\)

2. \(- 46 . 32\) \(kJ\) \(mol^{- 1}\)

3. \(- 23 . 16\) \(kJ\) \(mol^{- 1}\)

4. \(46 . 32\) \(kJ\) \(mol^{- 1}\)

For the hypothetical electrochemical cell given below:

A | A⁺(x M) || B⁺(y M) | B .

The measured cell emf is +0.20 V. Find the overall cell reaction.

1. A⁺ + B → A + B⁺

2.  A⁺ + e^- → A ; B⁺ + e^- → B

3. The cell reaction cannot be predicted.

4. A + B⁺ → A⁺ + B