Q. No.What is the standard cell potential when the Sn⁴⁺/Sn²⁺ (E° = +0.15 V) and Cr³⁺/Cr (E° = −0.74 V) half-cells are connected under standard conditions?
1. +0.89 V
2. +0.18 V
3. +1.83 V
4. +1.199 V
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| Assertion(A): | A standard hydrogen electrode is also called reversible electrode. |
| Reason(R) : | It can act on both as anode as well as cathode in an electrochemical 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. |
Match the Column-I with Column-II and mark the appropriate option:
| Column-I (Term) |
Column-II (Units) |
||
| A. | \(\Lambda_m\) | 1. | S cm-1 |
| B. | Ecell | 2. | m-1 |
| C. | κ | 3. | S cm2 mol-1 |
| D. | \(G^*\) | 4. | V |
(\(\Lambda_m\),\(E_{\text {cell }}\),\(\kappa\) and \(G^*\) represent the molar conductivity, the electromotive force of the cell, conductivity and Cell constant respectively)
Which of the following is correct?
1. A-2,B-1,C-4,D-3
2. A-2,B-4,C-1,D-3
3. A-3,B-1,C-4,D-2
4. A-3,B-4,C-1,D-2
| 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 1 and Statement 2 are true. |
| 4. | Both Statements 1 and Statement 2 are false. |
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\(\begin{array}{ll} \mathrm{Fe}^{2+} \rightarrow \mathrm{Fe}^{3+}+\mathrm{e}^{-} & \mathrm{E}_{\mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}}^0=0.77 \mathrm{~V} \\ 2 \mathrm{I}^{-} \rightarrow \mathrm{I}_2+2 \mathrm{e}^{-} & \mathrm{E}_{\mathrm{I}_2 / \text{I}^-}^{\circ}=0.54 \mathrm{~V} \end{array}\)
The standard electrode potential for the spontaneous reaction in the cell is \(x\times 10^{-2}V~\text{at}~ 298~K.\)
The value of \(x\) is: (Nearest Integer)
1. 32
2. 23
3. 45
4. 17
\(\mathrm{Ag}^{+}(\mathrm{aq})+\mathrm{e}^{-} \rightarrow \mathrm{Ag}(\mathrm{s}) \quad \mathrm{E}^{\circ}=+0.80 \mathrm{~V} \)
\(\mathrm{Fe}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightarrow \mathrm{Fe}(\mathrm{s}) \ \mathrm{E}^{\circ}=-0.44 \mathrm{~V}\)
The emf of the following cell is:
\(\mathrm{Fe}(\mathrm{s})+2 \mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+2 \mathrm{Ag}(\mathrm{s}) \)
| 1. | 1.16 V | 2. | 1.24 V |
| 3. | -1.24 V | 4. | -1.16 V |
| 1. | Y > X > Z | 2. | Z > X > Y |
| 3. | X > Y > Z | 4. | Y > Z > X |
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Match the items of Column I and Column II on the basis of data given below
| Column I | Column II | ||
| A | F2 | 1. | Metal is the strongest reducing agent |
| B | Li | 2. | Metal ion which is the weakest oxidising agent |
| C | Au3+ | 3. | Non-metal which is the best oxidising agent |
| D | Br– | 4. | Unreactive metal |
| E | Au | 5 | Anion that can be oxidised by Au3+ |
| F | Li+ | 6 | Anion which is the weakest reducing agent |
| G | F– | 7 | Metal ion which is an oxidising agent |
| Options: | A | B | C | D | E | F | G |
| 1. | 3 | 1 | 7 | 5 | 4 | 2 | 6 |
| 2. | 1 | 2 | 3 | 5 | 7 | 4 | 6 |
| 3. | 5 | 4 | 3 | 2 | 6 | 5 | 1 |
| 4. | 4 | 5 | 3 | 2 | 1 | 6 | 7 |
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The EMF of a Daniel cell at 298 K is E1 Zn|ZnSO4(0.01 M) || CuSO4(1.0 M)|Cu.
When the concentration of ZnSO4 is 1.0 M and that of CuSO4 is 0.01 M, the EMF is changed to E2. The correct relationship between E1 and E2 is:
| 1. | E1 > E2 | 2. | E1 < E2 |
| 3. | E1 = E2 | 4. | E2 = 0 ≠ E1 |
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Find the standard Gibbs energy change, ΔG°, for the following cell reaction at 298 K:
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Given:
E°cell = 2.0 V
F = 96000 C mol⁻¹
1. –384
2. 384
3. 192
4. –192
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