Gibb's Energy and Nernst Equation

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. 4.28V

2. 6.42V

3. 8.56V

4. 2.14V

Q.

How does concentration affect EMF?

Q. For the given cell reaction :
$C{u}^{2+}\left(aq\right)+2C{l}^{-}\left(aq\right)\to Cu\left(s\right)+C{l}_2\left(g\right)$
$E_{cell}^{\circ }$ is given as $-1.02V$.
Which of the following is correct?

1. The given reaction is spontaneous
2. Reaction will occur whenever $C{u}^{2+}$and $C{l}^{-}$ are brought together in an aqueous solution.
3. The given reaction is non-spontaneous
4. The reaction is in equilibrium

Q. The standard electrode potentials of $C{u}^{2+}|C{u}^{+}$ and $C{u}^{+}|Cu$ electrodes are +0.18V and +0.50V. The potential of $C{u}^{2+}|Cu$ electrode is _________.

1. 0.44V
   
2. -0.34V
   
3. -0.44V
4. 0.34V
   

Q.

For an electrochemical cell, $\mathrm{Sn}\left(\mathrm{s}\right)/{\mathrm{Sn}}^{2+}\left(\mathrm{aq}\right),1\mathrm{M})//{\mathrm{Pb}}^2+(\mathrm{aq},1\mathrm{M})/\mathrm{Pb}(\mathrm{s})$. The ratio $\left[{\mathrm{Sn}}^{2+}\right]/\left[{\mathrm{Pb}}^{2+}\right]$ when this cell attains equilibrium is ______.

$\mathrm{Given}:{\mathrm{E}}^{°}\mathrm{Sn}\left(\mathrm{s}\right)/{\mathrm{Sn}}^{2+}=0.14\mathrm{V};{\mathrm{E}}^{°}{\mathrm{Pb}}^{2+}/\mathrm{Pb})/{\mathrm{Sn}}^{2+}=0.13\mathrm{V},2.303\mathrm{RT}/\mathrm{F}=0.06\mathrm{V}$

Q. The standard reduction potential of the half cells at $298K$ are,
$N{i}^{2+}\left(aq\right)+2e^{-}\to Ni\left(s\right);E^o=-0.28V$
$M{g}^{2+}\left(aq\right)+2e^{-}\to Mg\left(s\right);E^o=-2.37V$

The standard cell potential for a voltaic cell constructed using the two half reaction is

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. - 49215 J
4. 96500 J

Q. For the cell reaction
$2{\mathrm{Fe}}^{3+}\left(\mathrm{aq}\right)+2I^{-}\left(\mathrm{aq}\right)\to 2{\mathrm{Fe}}^{2+}\left(\mathrm{aq}\right)+I_2\left(\mathrm{aq}\right)E{°}_{\mathrm{cell}}=0.24V\mathrm{at}298K$ The standard Gibbs energy $△G^o$ of the cell reaction is:
[Given that Faraday constant F = 96500 C mol–1]

1. – 46.32 kJ $mo{l}^{-1}$
2. – 23.16 kJ $mo{l}^{-1}$
   
3. 46.32 kJ$mo{l}^{-1}$
4. 23.16 kJ$mo{l}^{-1}$

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. For the cell reaction
$2{\mathrm{Fe}}^{3+}\left(\mathrm{aq}\right)+2I^{-}\left(\mathrm{aq}\right)\to 2{\mathrm{Fe}}^{2+}\left(\mathrm{aq}\right)+I_2\left(\mathrm{aq}\right)E{°}_{\mathrm{cell}}=0.24V\mathrm{at}298K$ The standard Gibbs energy $△G^o$ of the cell reaction is:
[Given that Faraday constant F = 96500 C mol–1]

1. – 46.32 kJ $mo{l}^{-1}$
2. – 23.16 kJ $mo{l}^{-1}$
   
3. 46.32 kJ$mo{l}^{-1}$
4. 23.16 kJ$mo{l}^{-1}$

Q.

Zn | Zn2+ (C1) || Zn2+ (C2) | Zn. For this cell G is negative if

1. C1 = C2

2. C1> C2

3. C2> C1

4. None

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

2. $110.37kJmo{l}^{-1}$

3. $210.37kJmo{l}^{-1}$

4. $105.185kJmo{l}^{-1}$

Q. The Nernst equation represents the sum of the change in Gibb's free energies for a cell at a given temperature.

1. True
2. False

Q. The standard electrode potentials of $C{u}^{2+}|C{u}^{+}$ and $C{u}^{+}|Cu$ electrodes are +0.18V and +0.50V. The potential of $C{u}^{2+}|Cu$ electrode is _________.

1. 0.34V
   
2. 0.44V
   
3. -0.34V
   
4. -0.44V

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. 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. Given that:
$E_{F{e}^{2+}/Fe}^{\circ }=-0.44\text{V}\text{and}{E}_{F{e}^{3+}/F{e}^{2+}}^{\circ }=0.77\text{V}$
$E_{F{e}^{3+}/Fe}^{\circ }$ is:

1. $1.21\text{V}$
2. $0.33\text{V}$
3. $-0.036\text{V}$
4. $0.036\text{V}$

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.

Electra is almost bored now. Stupid Nernst making her solve so many problems.

But wait, this is her last one. If she completes this, she is through as an apprentice!
Quick. Tell her what is the Gibb's energy of a Daniel Cell.
Use $F=96487Cmo{l}^{-1}$

1. $21.227Jmo{l}^{-1}$
2. $21.227mJmo{l}^{-1}$
3. $21.227kJmo{l}^{-1}$
4. $21.227mo{l}^{-1}$

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