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1. Determine the relationship between the standard cell potentials and equilibrium constants for the following reactions:

(a)	\(A+B \rightleftharpoons C+D \text { }\)  with  \(E^0=x\) volts and \(K_{e q}=K_1\)
(b)	\(2 A+2 B \rightleftharpoons 2 C+2 D \text { }\) with \(E^0=y\) volts and \(K_{e q}=K_2\)

1. \(x=y, K_1=K_2\)
2. \(x=2 y, K_1=2 K_2\)
3. \(x=y, K_1^2=K_2\)
4. \(x^2=y, K_1^2=K_2\)

2.

Assertion (A):	EMF of the cell is the potential difference between the electrode potentials of the cathode and anode when no current is drawn through the cell.
Reason (R):	The anode is kept on the right side and the cathode on the left side while representing the galvanic cell.

 

1.	Both (A) and (R) are True and (R) is the correct explanation of (A).
2.	Both (A) and (R) are True but (R) is not the correct explanation of (A).
3.	(A) is True but (R) is False.
4.	Both (A) and (R) are false

3. Consider the following electrochemical cell at \(298~ K:\)
\(\small {\begin{aligned} & \operatorname{Pt}(\mathrm{s})\left|\mathrm{H}_2(\mathrm{g}), (1 \mathrm{bar})\right| \mathrm{H}^{+}(\mathrm{aq}), (1 \mathrm{M}) \| \mathrm{M}^{4+}(\mathrm{aq}), |\mathrm{M}^{2+}(\mathrm{aq}) \mid \mathrm{Pt}(\mathrm{~s}) \end{aligned}}\)
When \(\frac{\left[\mathrm{M}^{2+}(\mathrm{aq})\right]}{\left[\mathrm{M}^{4+}(\mathrm{aq})\right]}=10^{\mathrm{x}},\) find the value of \(\mathrm {x}\).
[Given: \(\mathrm{E}_{\text {cell }}=0.092 \mathrm{~V}\); 
\(\mathrm{E}_{\text {cell}}^0=0.151 \mathrm{~V} ; 2.303 \frac{\mathrm{RT}}{\mathrm{~F}}=0.059 \mathrm{~V}\)]

1.	\(-2\),	2.	\(-1\)
3.	\(1\)	4.	\(2\)

4. Consider the given reaction:
\(\mathrm{Cu}(\mathrm{s})+\mathrm{Sn}^{2+}(0.001 \mathrm{M}) \rightarrow \mathrm{Cu}^{2+}(0.01 \mathrm{M})+\mathrm{Sn}(\mathrm{s})\)
What is the Gibbs free energy change for the above reaction at \(298 ~\text{K}\)?
[Given: E^⊖_Cell= \(-0.48 \mathrm{~V}\) and \(F=96500 \mathrm{C} \mathrm{mol}^{-1}\)]
1. \(383 \times 10^{-1} \mathrm{~kJ} / \mathrm{mol}\)
2. \(783 \times 10^{-1} \mathrm{~kJ} / \mathrm{mol}\)
3. \(983 \times 10^{-3} \mathrm{~kJ} / \mathrm{mol}\)
4. \(983 \times 10^{-1} \mathrm{~kJ} / \mathrm{mol}\)

5.

6. If the molar conductivity \(\left(\Lambda_{\mathrm{m}}\right)\) of a \(0.050 ~\text{mol}~ \text{L}^{-1}\) solution of a monobasic weak acid is \(90 ~\text{S} ~\text{cm}^2 ~\text{mol}^{-1}\), then its degree of dissociation will be: [Assume \(\Lambda_{+}^0=349.6~ \mathrm{S ~cm}^2 \mathrm{~mol}^{-1}\) and \(\mathrm{\Lambda}_{-}^{\circ}=50.4 \mathrm{~S} \mathrm{~cm}^2 \mathrm{~mol}^{-1} \)]

1.	\( 0.225 \)	2.	\( 0.215 \)
3.	\(0.115 \)	4.	\(0.125\)

7.

The half-cell reaction at the anode during the electrolysis of aqueous sodium chloride solution  is represented by : 
1. Na⁺(aq) + e^- ⟶ Na(s) ; \(E_{cell}^{o} \ = \ -2.71 \ V \)
2. 2H₂O(l) ⟶ O₂(g) + 4H⁺(aq) + 4e^- ; \(E_{cell}^{o} \) = 1.23 V
3. H⁺(aq) + e^- ⟶ \(\frac{1}{2}\)H₂(g) ; \(E_{cell}^{o} \) = 0.00 V
4. Cl^-(aq) ⟶ \(\frac{1}{2}\)Cl₂(g) + e^- ; \(E_{cell}^{o}\) \(= 1 . 36  V\)
$\mathrm{}$

8. The correct statement regarding the discharge of lead storage batteries is:
1. SO₂ is evolved
2. Lead sulphate is consumed
3. Lead is formed
4. Sulphuric acid is consumed

9. A solution of 0.02 M \(\text{CH}_3\text{COOH}\) has a specific conductance (\(\kappa\)) of \(5 \times 10^{-5} \, \text{S cm}^{-1}\). Given that the limiting molar conductance of \(\text{CH}_3\text{COOH}\) is \(400 \, \text{S cm}^2 \text{ mol}^{-1}\), what is the dissociation constant (\(\mathrm{K_a}\)) of \(\text{CH}_3\text{COOH}\)?

1.	\(8 \times 10^{-7}\)	2.	\(6 \times 10^{-7}\)
3.	\(10 \times 10^{-7}\)	4.	\(4 \times 10^{-7}\)

10. Which metal is used as a protective coating during the process of galvanization?

1.	Cu	2.	Zn
3.	Pb	4.	Cr