Rusting of iron is envisaged as setting up of an electrochemical cell because:
1. Oxidation of iron takes place during rusting.
2. Reduction of iron takes place.
3. Iron shows disproportionation reaction.
4. Oxygen is getting oxidised.
1. Oxidation of iron takes place during rusting.
2. Reduction of iron takes place.
3. Iron shows disproportionation reaction.
4. Oxygen is getting oxidised.
| Statement 1: | The standard electrode potential of the system Mg2+ | Mg can be measured with respect to hydrogen electrode. |
| Statement 2: | The standard electrode potential for hydrogen electrode is zero. |
1. Statement 1 is true, statement 2 is false.
2. Statement 1 is false, statement 1 is true.
3. Both statements are true.
4. Both statements are false.
For a cell reaction involving a two-electron change, the standard Emf of the cell is found to be 0.295V at 25oC . The equilibrium constant of the reaction at 25oC will be :
1. 1 \(\times\) 10-10
2. 29.5 \(\times\) 10-2
3. 10
4. 1 \(\times\) 1010
\(\mathrm{MnO}_{4}^{-}+8 \mathrm{H}^{+}+5 \mathrm{e}^{-} \rightarrow \mathrm{Mn}^{2+}+4 \mathrm{H}_{2} \mathrm{O}, \)
\( \mathrm{E}_{\mathrm{Mn}^{2+}}^{\circ} / \mathrm{MnO}_{4}^{-}=-1.510 \mathrm{~V} \)
\( \frac{1}{2} \mathrm{O}_{2}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow \mathrm{H}_{2} \mathrm{O}, \)
\( \mathrm{E}_{\mathrm{O}_{2} / \mathrm{H}_{2} \mathrm{O}}^{\circ}=+1.223 \mathrm{~V}\)
Will the permanganate ion, \(\mathrm{MnO}_{4}^{-}\) , liberate \(\mathrm{O}_{2}\) from water in the presence of an acid?
1. No, because \(\mathrm{E}_{\text {cell }}^{\circ}=-2.733 \mathrm{~V}\)
2. Yes, because \(\mathrm{E}_{\text {cell }}^{\circ}=+0.287 \mathrm{~V}\)
3. No, because \(\mathrm{E}_{\text {cell }}^{\circ}=-0.287 \mathrm{~V}\)
4. Yes, because \(\mathrm{E}_{\text {cell }}^{\circ}=+2.733 \mathrm{~V}\)
Find the emf of the cell in which the following reaction takes place at 298 K:
\(\mathrm{Ni}(\mathrm{s})+2 \mathrm{Ag}^{+}(0.001 \mathrm{M}) \rightarrow \mathrm{Ni}^{2+}(0.001 \mathrm{M})+2 \mathrm{Ag}(\mathrm{s}) \)
\( \small{\text { (Given that } \mathrm{E}_{\text {cell }}^{\circ}=10.5 \mathrm{~V}, \frac{2.303 \mathrm{RT}}{\mathrm{F}}=0.059 \text { at } \ 298 \mathrm{~K})} \)
1. 1.05 V
2. 1.0385 V
3. 1.385 V
4. 0.9615 V
The three cells with their \(E^\circ_{\text{(cell)}}\) values are given below:
| Cells | \(E^\circ_{\text{(cell)}}/V\) | |
| (a) | Fe|Fe2+||Fe3+|Fe | 0.404 |
| (b) | Fe|Fe2+||Fe3+, Fe2+|Pt | 1.211 |
| (c) | Fe|Fe3+||Fe3+, Fe2+|Pt | 0.807 |
(F represents the charge on 1 mole of electrons.)
| 1. | -1.212 F, -1.211 F, -0.807 F |
| 2. | +2.424 F, +2.422 F, +2.421 F |
| 3. | -0.808 F, -2.422 F, -2.421 F |
| 4. | -2.424 F, -2.422 F, -2.421 F |
\(\land^o_m\) for NaCl, HCl and \(CH_3COONa \) are 126.4, 425.9, and 91.05 S cm2 mol-1 respectively. If the conductivity of 0.001028 mol L-1 acetic acid solution is \(4.95 \times 10^{-5} S ~cm^{-1} \), the degree of dissociation of the acetic acid solution is-
| 1. | 0.01233 | 2. | 1.00 |
| 3. | 0.1233 | 4. | 1.233 |
Two half cell reactions are given below:
\(\begin{aligned} &\mathrm{Co}^{3+}+e^{-} \rightarrow \mathrm{Co}^{2+}, \mathrm{E}_{\mathrm{Co}^{2+} / \mathrm{Co}^{3+}}^{\circ}=-1.81 \mathrm{~V} \\ &2 \mathrm{Al}^{3+}+6 e^{-} \rightarrow 2 \mathrm{Al}(\mathrm{s}), \mathrm{E}_{\mathrm{Al} / \mathrm{Al}^{3+}}^{\circ}=+1.66 \mathrm{~V} \end{aligned} \)
The standard EMF of a cell with feasible redox reaction will be:
| 1. | +7.09 V | 2. | +0.15 V |
| 3. | +3.47 V | 4. | –3.47 V |
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
is 1.1 V. Calculate the standard Gibbs energy change for the cell reaction.
(Given F = 96487 C mol-1)
| 1. | –200.27 kJ mol-1 | 2. | –212.27 kJ mol-1 |
| 3. | –212.27 J mol-1 | 4. | –200.27 J mol-1 |
The molar conductivity of 0.007 M acetic acid is 20 S cm2 mol-1. The dissociation constant of acetic acid is :
(\(\mathrm{\Lambda_{H^{+}}^{o} \ = \ 350 \ S \ cm^{2} \ mol^{-1} }\))
(\(\mathrm{\mathrm{\Lambda_{CH_{3}COO^{-}}^{o} \ = \ 50 \ S \ cm^{2} \ mol^{-1} }}\))
1. mol L-1
2. mol L-1
3. mol L-1
4. mol L-1
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