Gibb's Energy and Nernst Equation

Q. The rusting of iron takes place as follows $2H^{+}+2e^{-}+\frac{1}{2}{O}_2\to H_{2}O\left(l\right);E^{\circ }=+1.23V$ $F{e}^{2+}+2e^{-}\to Fe\left(s\right);E^{\circ }=-0.44V$ Calculate $\Delta G^{\circ }$ for the net process

1. $-322kJmo{l}^{-1}$
2. $-161kJmo{l}^{-1}$
3. $-152kJmo{l}^{-1}$
4. $76kJmo{l}^{-1}$

Q. The standard emf of the cell $Fe|F{e}^{2+}\left(aq\right)||C{d}^{2+}\left(aq\right)|Cd$ is 0.0372 V and temperature coefficient of the cell is $-1.25\times {10}^{-4}V{K}^{-1}$. Which is/are correct about the cell (at 300 K)?

1. $\Delta G°=7.18kJ,\Delta H°=-14.42kJ$
2. $\Delta G°=-7.18kJ,\Delta H°=57.50J$
3. $\Delta G°=-7.18kJ,\Delta H°=-14.42kJ$
4. $\Delta S°=-24.125{\mathrm{JK}}^{-1},\mathrm{reaction}\mathrm{is}\mathrm{spontaneous}$

Q. Consider the following cell reaction $C{d}_{\left(s\right)}+H{g}_{2}S{O}_{4\left(s\right)}+\frac{9}{5}{H}_2{O}_{\left(l\right)}\rightleftharpoons CdS{O}_4.\frac{9}{5}{H}_2{O}_{\left(s\right)}+2H{g}_{\left(l\right)}$ The value of $E_{cell}^0$ is $4.315V$ at ${25}^{0}C$ . If $\Delta H^0=-825.2$ kJ $mo{l}^{-1}$ , the standard entropy change $\Delta S^0$ in $J{K}^{-1}$ is _________ . ( Nearest integer ) $[Given:\text{Faraday constant}=96487Cmo{l}^{-1}]$
(JEE MAIN 2021

Q. The half-cell reactions for rusting of iron are,
$2H^{+}+\frac{1}{2}{O}_2+2e^{-}\to 2H_{2}O,E^{\circ }=+1.23V,$
$F{e}^{2+}+2e^{-}\to Fe\left(s\right);E^{\circ }=-0.44V.$
$\Delta G^{\circ }\left(inkJ\right)$ for the reaction is :

1. - 76
2. -322
3. -122
4. -176

Q.

The electrical work done during the reaction at 298 K:
$2H{g}_{\left(l\right)}+C{l}_{2\left(g\right)}\to H{g}_{2}C{l}_{2\left(s\right)},isGiventhat{E}_{\frac{C{l}_2}{C{l}^{-}}}^{\circ }=1.36V;E_{\frac{H{g}_{2}C{l}_2}{Hg,C{l}^{-}}}^{\circ }=0.27V;P_{C{l}_2}=1atm,$

1. $210.37kJmo{l}^{-1}$

2. $105.185kJmo{l}^{-1}$

3. $420.74kJmo{l}^{-1}$

4. $110.37kJmo{l}^{-1}$

Q.

On the basis of information available from the reaction:
$\frac{4}{3}Al+O_2\to \frac{2}{3}A{l}_2{O}_3\Delta G=-827KJ/mol$ of $O_2$. The minimum e.m.f. required to carry out
electrolysis of $A{l}_2{O}_3$ is

1. 2.14V

2. 4.28V

3. 6.42V

4. 8.56V

Q. $DeltaH$ and $DeltaS$ for the reaction $A{g}_2{O}_{\left(s\right)}$ $\to$ $2A{g}_{\left(s\right)}$+$\frac{1}{2}$ $O_{2\left(g\right)}$ are 30.56 kJ $mo{l}^{-1}$ and 66.00 $J{K}^{-1}mo{l}^{-1}$ respectively, the temperature at which the free energy change for the reaction will be zero is:

1. 463 K
2. 35440 K
3. 20 K
4. 483 K

Q.

The standard electrode potentials, K⁺/K= -2.93V, Ag⁺/Ag = 0.80V, the electrode which is negatively charged is

1. Ag⁺/Ag

2. K⁺/K

3. Any of the two

4. None of them

Q.

A voltaic cell has an E° value of (–1V). The cell reaction:

1. is spontaneous

2. has a positive ΔG°

3. has a negative ΔG°

4. has K = 1

Q. The electrochemical cell shown below is a concentration cell.$M|M^{2+}$ (Saturated solution of a sparingly soluble salt, $M{X}_2\left)||M^{2+}\right(0.001mold{m}^{-3})M.$ The emf of the cell depends on the difference in concentration of $M^{2+}$ ions at the two electrodes. The emf of the cell at 298K is 0.059

The value of $\Delta G\left(kJmo{l}^{-1}\right)$ for the given cell is: (take$1F=96500Cmo{l}^{-1})$

1. $-5.7$
2. $5.7$
3. $11.4$
4. $-11.4$

Q. For the following cell:
$Zn\left(s\right)|ZnS{O}_4\left(aq\right)||CuS{O}_4\left(aq\right)\left)|Cu\right(s)$
When the concentration of $Z{n}^{2+}$ is 10 times the concentration of $C{u}^{2+}$, the expression for $\Delta G\left(inJmo{l}^{-1}\right)$ is
[F is Faraday constant; R Is gas constant; T is temperature ; $E^{\circ }\left(cell\right)=1.1v]$

1. 2.303 RT + 1.1 F
2. 1.1 F
3. 2.303 RT - 2.2 F
4. -2.2 F

Q. For the following electrochemical cell at 298 K,
$Pt\left(s\right)|H_2(g,1bar)|H^{+}(aq,1M)||M^{4+}\left(aq\right),M^{2+}\left(aq\right)|Pt\left(s\right)E_{cell}=0.092Vwhen\frac{\left[M^{2+}\right(aq\left)\right]}{\left[M^{4+}\right(aq\left)\right]}={10}^{x}Given:E_{M^{4+}/M^{2+}}^0=0.151V;2.303\frac{RT}{F}=0.059V$
The value of x is:

1. -2
2. -1
3. 1
4. 2

Q.

The electrical work done during the reaction at 298 K:
$2H{g}_{\left(l\right)}+C{l}_{2\left(g\right)}\to H{g}_{2}C{l}_{2\left(s\right)},isGiventhat{E}_{\frac{C{l}_2}{C{l}^{-}}}^{\circ }=1.36V;E_{\frac{H{g}_{2}C{l}_2}{Hg,C{l}^{-}}}^{\circ }=0.27V;P_{C{l}_2}=1atm,$

1. $210.37kJmo{l}^{-1}$

2. $105.185kJmo{l}^{-1}$

3. $420.74kJmo{l}^{-1}$

4. $110.37kJmo{l}^{-1}$

Q. The standard emf of the cell $Fe|F{e}^{2+}\left(aq\right)||C{d}^{2+}\left(aq\right)|Cd$ is 0.0372 V and temperature coefficient of the cell is $-1.25\times {10}^{-4}V{K}^{-1}$. Which is/are correct about the cell (at 300 K)?

1. $\Delta G°=7.18kJ,\Delta H°=-14.42kJ$
2. $\Delta G°=-7.18kJ,\Delta H°=57.50J$
3. $\Delta G°=-7.18kJ,\Delta H°=-14.42kJ$
4. $\Delta S°=-24.125{\mathrm{JK}}^{-1},\mathrm{reaction}\mathrm{is}\mathrm{spontaneous}$

Q. For the following electrochemical cell at 298 K,
$Pt\left(s\right)|H_2(g,1bar)|H^{+}(aq,1M)||M^{4+}\left(aq\right),M^{2+}\left(aq\right)|Pt\left(s\right)E_{cell}=0.092Vwhen\frac{\left[M^{2+}\right(aq\left)\right]}{\left[M^{4+}\right(aq\left)\right]}={10}^{x}Given:E_{M^{4+}/M^{2+}}^0=0.151V;2.303\frac{RT}{F}=0.059V$
The value of x is:

1. -2
2. -1
3. 1
4. 2

Q. 5. If 50.2 then Which one is correct exactly A. 1 B. 1.0

Q. Give the standard electrode potentials:
K⁺/K = -2.93 V, Ag⁺/Ag = +0.80 V, Cr⁺³/Cr = -0.74 V
Out of these electrode which will be the strongest reducing agent?

Q. When two half cells X and Y combined then the standard cell potential of cell is.

1. $E_X+E_Y$
2. $E_X^0+E_Y^0$
3. -($\Delta G_X^0+\Delta G_Y^0$)/nF
4. $\Delta G_X^0+\Delta G_Y^0$

Q.

Given standard electrode potentials
$F{e}^{++}+2e^{-}\to Fe;E^{\circ }=-0.440V$
$F{e}^{+++}+3e^{-}\to Fe;E^{\circ }=-0.036V$
The standard electrode potential $\left(E^{\circ }\right)$ for $F{e}^{+++}+e^{-}\to F{e}^{++}$ is

1. - 0.476 V

2. - 0.404 V

3. + 0.404 V

4. + 0.772 V

Q. Which statement is true about a spontaneous cell reaction in galvanic cell?

1. $E_{cell}^{\circ }>0,\Delta G^{\circ }<0$
2. $E_{cell}^{\circ }<0,\Delta G^{\circ }>0$
3. $E_{cell}^{\circ }>0,\Delta G^{\circ }>0$
4. $E_{cell}^{\circ }=0,\Delta G^{\circ }=0$

Q. The electrochemical cell shown below is a concentration cell.$M|M^{2+}$ (Saturated solution of a sparingly soluble salt, $M{X}_2\left)||M^{2+}\right(0.001mold{m}^{-3})M.$ The emf of the cell depends on the difference in concentration of $M^{2+}$ ions at the two electrodes. The emf of the cell at 298K is 0.059

The value of $\Delta G\left(kJmo{l}^{-1}\right)$ for the given cell is: (take$1F=96500Cmo{l}^{-1})$

1. $-5.7$
2. $5.7$
3. $11.4$
4. $-11.4$

Q. Calculate standard free energy change for the reaction $\frac{1}{2}Cu\left(s\right)+\frac{1}{2}C{l}_2\left(g\right)\rightleftharpoons \frac{1}{2}C{u}^{2+}+C{l}^{-}$taking place at in a cell whose standard e.m.f. is $1.02volts$

1. - 98430 J
2. 98430 J
3. 96500 J
4. - 49215 J