Gibb's Free Energy

Q. If $E_{A{u}^{+}/Au}^{\circ }$ is 1.69 V & $E_{A{u}^{3+}/Au}^{\circ }$ is 1.40 V, then $E_{A{u}^{3+}/A{u}^{+}}^{\circ }$ will be:

1. 0.19 V
2. 2.945 V
3. 1.255 V
4. None

Q. The free energy change when 1 mole of NaCl is dissolved in water at ${25}^{o}C$ is $-xkJmo{l}^{-1}$ Lattice energy of NaCl = $777.8kJmo{l}^{-1}$; $△S$ for dissolution $=0.043kJmo{l}^{-1}$ and hydration energy of $NaCl=-774.1kJmo{l}^{-1}.$ Value of x (rounded off to nearest integer) is

Q. If $E_{A{u}^{+}/Au}^{\circ }$ is 1.69 V & $E_{A{u}^{3+}/Au}^{\circ }$ is 1.40 V, then $E_{A{u}^{3+}/A{u}^{+}}^{\circ }$ will be:

1. 0.19 V
2. 2.945 V
3. 1.255 V
4. None

Q. $E^{\circ }\left(SRP\right)$ of different half cells are given:-
$E_{C{u}^{2+}/Cu}^{\circ }=0.34volt;E_{Z{n}^{2+}/Zn}^{\circ }=-0.76volt$
$E_{A{g}^{+}/Ag}^{\circ }=0.8volt;E_{M{g}^{2+}/Mg}^{\circ }=-2.37volt$
Which of the following has the most negative $\Delta G^{\circ }$ value?

1. $Zn\left(s\right)|Z{n}^{2+}\left(1M\right)||M{g}^{2+}\left(1M\right)|Mg\left(s\right)$
2. $Zn\left(s\right)|Z{n}^{2+}\left(1M\right)||A{g}^{+}\left(1M\right)\left)|Ag\right(s)$
3. $Cu\left(s\right)|C{u}^{2+}\left(1M\right)||A{g}^{+}\left(1M\right)|Ag\left(s\right)$
4. $Ag\left(s\right)|A{g}^{+}\left(1M\right)||M{g}^{2+}\left(1M\right)|Mg\left(s\right)$

Q. The Gibbs energy for the decomposition of $A{l}_2{O}_3$ at ${500}^{0}C$ is as follows:
$\frac{2}{3}A{l}_2{O}_3\to \frac{4}{3}Al+O_2$
${\Delta }_{rxn}G=+960kJmo{l}^{-1}$.The potential difference needed for the electrolytic reduction of aluminum oxide$\left(A{l}_2{O}_3\right)$ at ${500}^{o}C$ is:

1. 4.5 V
2. 3.0 V
3. 2.5 V
4. 5.0 V

Q. $d\left(\Delta G\right)=-\left(\Delta S_{\text{reaction}}\right).dT$ $\Delta G=\Delta H-T\Delta S$ $\Delta G=\Delta H+T{\left(\frac{d\left(\Delta G\right)}{dT}\right)}_P$
At $300\text{K},\Delta H$ for the reaction,
$Z{n}_{\left(S\right)}+2AgC{l}_{\left(S\right)}\to ZnC{l}_{2_{\left(aq\right)}}+2A{g}_{\left(S\right)}$
is $-218\text{kJ/mol}$ while the e.m.f of the cell is $1.015\text{V}$. ${\left(\frac{dE}{dT}\right)}_P$ of the cell is:

1. $-4.2\times {10}^{-4}{\text{VK}}^{-1}$
2. $-3.81\times {10}^{-4}{\text{VK}}^{-1}$
3. $7.62\times {10}^{-4}{\text{VK}}^{-1}$
4. $0.11{\text{VK}}^{-1}$

Q. If the $E_{cell}^{\circ }$ for a given reaction has a negative value, which of the following gives the correct relationships for the values of $\Delta G^{\circ }$ and $K_{eq}$ ?

1. $\Delta G^{\circ }<0;K_{eq}>1$
2. $\Delta G^{\circ }<0;K_{eq}<1$
3. $\Delta G^{\circ }>0;K_{eq}<1$
4. $\Delta G^{\circ }>0;K_{eq}>1$

Q. $d\left(\Delta G\right)=-\left(\Delta S_{\text{reaction}}\right).dT$ $\Delta G=\Delta H-T\Delta S$ $\Delta G=\Delta H+T{\left(\frac{d\left(\Delta G\right)}{dT}\right)}_P$
Calculate $\Delta S$ for the given cell reaction in the previous question.

1. $-73.53{\text{JK}}^{-1}{\text{mol}}^{-1}$
2. $83.53{\text{JK}}^{-1}{\text{mol}}^{-1}$
3. $100{\text{JK}}^{-1}{\text{mol}}^{-1}$
4. None of these

Q. $d\left(\Delta G\right)=-\left(\Delta S_{\text{reaction}}\right).dT$ $\Delta G=\Delta H-T\Delta S$ $\Delta G=\Delta H+T{\left(\frac{d\left(\Delta G\right)}{dT}\right)}_P$
Calculate $\Delta S$ for the given cell reaction in the previous question.

1. $-73.53{\text{JK}}^{-1}{\text{mol}}^{-1}$
2. $83.53{\text{JK}}^{-1}{\text{mol}}^{-1}$
3. $100{\text{JK}}^{-1}{\text{mol}}^{-1}$
4. None of these

Q.

For a thermodynamically spontaneous process the change in free energy is __.

Q.

Calculate the free energy change when 1 mol of NaCl is dissolved in water at 298 K. given:
(i) Lattice energy of NaCl = -778 kJ $mo{l}^{-1}$
(ii) Hydration energy of NaCl = -778 kJ $mo{l}^{-1}$
(iii) Entropy change at 298 K = 43 J $mo{l}^{-1}$

1. -9.11 kJ

2. -11.9 kJ

3. -10.9 kJ

4. None of these

Q. $d\left(\Delta G\right)=-\left(\Delta S_{\text{reaction}}\right).dT$
$\Delta G=\Delta H-T\Delta S$
$\Delta G=\Delta H+T{\left(\frac{d\left(\Delta G\right)}{dT}\right)}_P$
The temperature coefficient of the e.m.f of a cell, ${\left(\frac{dE}{dT}\right)}_P$ is given by:

1. $\frac{nF}{\Delta S}$
2. $\frac{\Delta S}{nF}$
3. $\frac{\Delta S}{nFT}$
4. $-nFE$

Q. What is the $\Delta G^{\circ }$ for the following reaction?
$C{u}_{\left(aq\right)}^{2+}+2A{g}_{\left(s\right)}\to C{u}_{\left(s\right)}+2A{g}_{\left(aq\right)}^{+}$
Given:$E_{C{u}^{2+}/Cu}^0=0.34V,E_{Ag/A{g}^{+}}^0=-0.8V$

1. -44.5 kJ
2. 44.5 kJ
3. -89 kJ
4. 89 kJ

Q. The temperature coefficient of emf of a cell designed as $Zn|\underset{\left(1M\right)}{Z{n}^{2+}}||\underset{\left(0.1M\right)}{C{u}^{2+}}|Cu$ is $-1.4\times {10}^{-4}V$ per degree. Change in entropy for the reaction $Zn+\underset{0.1M}{C{u}^{2+}}\rightleftharpoons \underset{1M}{Z{n}^{2+}}+Cu$ at ${27}^{\circ }C$ is:$(E_{cell}^{\circ }=1.1V)$

1. 14 J
2. 27 J
3. -14 J
4. -27 J

Q. The carbon reduction method is mainly applicable for the extraction of how many of the following metals?
$Sn,Al,Cr,Mn,Pb,Ca,Na,Zn$

Q. For the reaction, $A{g}_{\left(\text{aq}\right)}^{+}+C{l}_{\left(\text{aq}\right)}^{-}\rightleftharpoons AgC{l}_{\left(\text{s}\right)}$
Given:
$\begin{array}{cc}\text{Species}& \Delta G_f^{\circ }\left(\text{kJ/mol}\right)\\ A{g}_{\left(\text{aq}\right)}^{+}& +77\\ C{l}_{\left(\text{aq}\right)}^{-}& -129\\ AgC{l}_{\left(\text{s}\right)}& -109\end{array}$
Calculate $E_{\text{cell}}^{\circ }$ at $298\text{K}$.
Also find the solubility product of $AgCl$.

1. $0.59\text{V}$ and ${\text{K}}_{\text{sp}}={10}^{-10}$
2. $0.79\text{V}$ and ${\text{K}}_{\text{sp}}=3\times {10}^{-5}$
3. $0.36\text{V}$ and ${\text{K}}_{\text{sp}}={10}^{-14}$
4. $-0.59\text{V}$ and ${\text{K}}_{\text{sp}}={10}^{-10}$