Gibb's Energy and Nernst Equation

Q. 22. The electrode potentials for Cu2+ +e- ---> Cu+ and Cu+ +e- --->Cu are +0.15V and +0.50V respectively.The value of E(Cu2+/Cu ) will be 1.0.150V 2.0.500V 3.0.325V 4.0.650V

Q.

Calculate the standard cell potentials of galvanic cell in which the following reactions take place :

(i) $2Cr\left(s\right)+3C{d}^{2+}\left(aq\right)\to 2C{r}^{3+}\left(aq\right)+3Cd$

(ii) $F{e}^{2+}\left(aq\right)+A{g}^{+}\left(aq\right)\to F{e}^{3+}\left(aq\right)+Ag\left(s\right)$

Calculate the ${\Delta }_r{G}^{\circ }$ and equilibrium constant for the reactions

Q. 47. If Zn2+/An electrode is diluted 1000 times then the change in electrode potential is 1. Increase of 29.5 mV 2. Increase of 59 mV 3. Decrease of 88.5 mV 4. Decrease of 91.1 mV.

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. – 322
2. – 122
3. – 176
4. − 76

Q. Standard electrode potentials are Fe2+/Fe; E° = -0.44 volts Fe3+/Fe2+; E° = 0.77 volts Fe2+, Fe3+ and Fe blocks are kept together, then (1) Fe2+ increases (2) Fe3+ decreases (3) Fe2+/Fe3+ remains unchanged (4) Fe2+ decreases​

Q. The rate of decomposition for methyl nitrite and ethylnitrite can be given in terms of rate constant $k_1$ and $k_2$ respectively. The energy of activation for the two reactions are $152.30{\text{kJ mol}}^{–1}$ and $157.7{\text{kJ mol}}^{–1}$ as well as frequency factors are ${10}^{13}$ and ${10}^{14}$ respectively for the decomposition of methyl and ethylnitrite. Calculate the temperature at which rate constant will be same for the two reactions.

1. $\text{354 K}$
2. $\text{674 K}$
3. $\text{256 K}$
4. $\text{282 K}$

Q. For the cell $Cu\left(s\right)|C{u}^{2+}\left(aq\right)\left(0.1M\right)||A{g}^{+}\left(aq\right)\left(0.01M\right)|Ag\left(s\right)$ the cell potential $E_1=0.3095V$
For the cell $Cu\left(s\right)|C{u}^{2+}\left(aq\right)\left(0.01M\right)||A{g}^{+}\left(aq\right)\left(0.001M\right)Ag\left(s\right)$ the cell potential = ____$\times {10}^{-2}V.$
(Round off to the Nearest Integer).
$[\text{Use}:\frac{2.303RT}{F}=0.059]$

Q.

Using the standard electrode potential valuesl predict if the reaction between the following is feasible :

(i) $F{e}^{3+}\left(aq\right)and{I}^{-}\left(aq\right)$

(ii) $A{g}^{+}\left(aq\right)andCu\left(s\right)$

(iii) $F{e}^{3+}\left(aq\right)andB{r}^{-}\left(aq\right)$

(iv) $Ag\left(s\right)$ and $F{e}^{3+}\left(aq\right)$

(v) $B{r}_2\left(aq\right)andF{e}^{2+}\left(aq\right)$

Given standard electrode potentials

$E_{\frac{1}{2}\frac{t_2}{1^{-}}}^{\circ }=+0.541V{E}_{\left(\frac{C{u}^{2+}}{Cu}\right)}^{\circ }=+0.34V{E}_{\left(\frac{1}{2}\frac{B{r}_2}{B{r}^{-}}\right)}^{\circ }=+1.09V{E}_{\left(\frac{A{g}^{+}}{Ag}\right)}^{\circ }=+0.80V{E}_{\left(\frac{F{e}^{3+}}{F{e}^{2+}}\right)}^{\circ }=+0.77V$

Q.

How does temperature affect Nernst equation?

Q. 16. If the standard electrode potential of Cu2+/Cu is .34V, what is the electrode potential of .1M conc of Cu2+?

Q.

Standard electrode potential of NHE at 298 K is

1. 0.05 V

2. 0.1 V

3. 0.00 V

4. 0.11 V

Q. The potential of standard hydrogen electrode is zero. This implies that

1. $\Delta G_f^0(H^{+},aq.)=0$
2. $\Delta H_f^0(H^{+},aq.)=0$
3. $\Delta G_f^0(H^{+},aq.)<0$
4. $\Delta G_f^0(H^{+},aq.)>0$

Q. The standard electrode potential of $C{u}^{2+}/Cu$ is $+0.34\text{V}$. At what concentration of $C{u}^{2+}$ ions will this electrode potential be zero?

Take:
$Antilog(-11.52)=3\times {10}^{-12}$

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.

For the decomposition of ${\mathrm{N}}_2{\mathrm{O}}_5$ it is given that:

$\underset{\mathrm{Dinitrogen}\mathrm{pentaoxide}}{\underbrace{2{\mathrm{N}}_2{\mathrm{O}}_5\left(\mathrm{g}\right)}}\to \underset{\mathrm{Nitrogen}\mathrm{dioxide}}{\underbrace{4{\mathrm{NO}}_2\left(\mathrm{g}\right)}}+\underset{\mathrm{Oxygen}}{\underbrace{{\mathrm{O}}_2\left(\mathrm{g}\right)}}$, Activation energy $\mathrm{Ea}$

$\underset{\mathrm{Dinitrogen}\mathrm{pentaoxide}}{\underbrace{{\mathrm{N}}_2{\mathrm{O}}_5\left(\mathrm{g}\right)}}\to \underset{\mathrm{Nitrogen}\mathrm{dioxide}}{\underbrace{2{\mathrm{NO}}_2\left(\mathrm{g}\right)}}+\underset{\mathrm{Oxygen}}{\underbrace{\frac{1}{2}{\mathrm{O}}_2\left(\mathrm{g}\right)}}$, Activation energy ${\mathrm{Ea}}^{}$then,

1. $\mathrm{Ea}={\mathrm{Ea}}^{}$

2. $\mathrm{Ea}>{\mathrm{Ea}}^{}$

3. $\mathrm{Ea}<{\mathrm{Ea}}^{}$

4. $\mathrm{Ea}=2{\mathrm{Ea}}^{}$

Q. Standard electrode potential of three metals A, B and C are -1.5 V, +0.7 V and -3.7 V respectively. The reducing power of these metals will be

1. B > A > C
2. B > C > A
3. C > A > B
4. A > B > C

Q.

If $\overrightarrow{a}=2\hat{i}+\hat{j}+2\hat{k}$ , then the value of $\hat{i}\times {(\overrightarrow{a}\times \hat{i})}^2+\hat{j}\times {(\overrightarrow{a}\times \hat{j})}^2+\hat{k}\times {(\overrightarrow{a}\times \hat{k})}^2$is equal to

Q. A standard hydrogen electrode has zero electrode potential because:

1. the electrode potential is assumed to be zero
2. hydrogen is less reactive
3. hydrogen atom has only one electron
4. hydrogen atom can gain or lose an electron

Q. Consider an electrochemical cell: $A\left(s\right)|A^{n+}(aq,2M)||B^{2n+}$ (aq, 1 M)|B(s). The value of $\Delta H^{\circ }$ for the cell reaction is twice that of $\Delta G^{\circ }$ at 300 K. If th emf of the cell is zero, the $\Delta S^{\circ }\left(inJ{K}^{-1}mo{l}^{-1}\right)$ of the cell reaction per mole of B formed at 300 K is$Jmo{l}^{-1}{K}^{-1}$
(Given: ln(2) =0.7, R (universal gas constant) 8.3 $J{K}^{-1}mo{l}^{-1}$. H, S and G are enthalpy, entropy and Gibbs energy, respectively.)

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

Q. Calculate the reduction potential for the following half cell at ${25}^{\circ }C$:
$Ag\left(s\right)|A{g}^{+}(aq,{10}^{-5}M);E_{A{g}^{+}/Ag}^0=-0.80V$

1. $E_{A{g}^{+}/Ag}=0.5V$
2. $E_{A{g}^{+}/Ag}=0.25V$
3. $E_{A{g}^{+}/Ag}=0.8V$
4. $E_{A{g}^{+}/Ag}=1.5V$

Q. the s†an dard electrode potential of Zn , Ag and Cu are -0.76 , 0.80 and 0.34 then 1.Ag can oixidise Zn andCu 2.Agcan reduce Zn+2 and Cu+2 3.Zn can reduce Ag+ and Cu+2 4. Cu can oxidise Zn andA

Q.

Cell reaction is spontaneous when

1. ΔG^θ is positive

2. E^θ_Red is negative.

3. ΔG^θ is negative

4. E^θ_Red is positive

Q. For the following cell,

$Zn\left(s\right)|ZnS{O}_4\left(aq\right)||CuS{O}_4\left(aq\right)|Cu\left(s\right)$

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; $E^{\circ }$(cell) = 1.1V]

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

Q. All the energy released from the reaction $X\to Y;\Delta G^{\circ }=-193{\text{kJ mol}}^{-1}$ is used for oxidizing $M^{+}$ as $M^{+}\to M^{3+}+2e^{-};E^{\circ }=-0.25V$. Under standard conditions the number of moles of $M^{+}$ is
(Given : $[F=96500Cmo{l}^{-1}]$)

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. -322 kJ mol^-1
2. -161 kJ mol^-1
3. -152 kJ mol^-1
4. -76 kJ mol^-1

Q. At $298\text{K}$ the standard free energy of formation of $H_{2}O\left(l\right)$ is $-257.20\text{kJ/mole}$ while that of its ionization into $H^{+}$ ion and hydroxyl ions is $80.35\text{kJ/mole}$. Then the emf of the following cell at $298\text{K}$ will be $(\text{take F}=96500\text{C})$:
$H_2(g,1\text{bar})|H^{+}\left(1M\right)||O{H}^{-}\left(1M\right)|O_2(g,1\text{bar})$

1. $0.40\text{V}$
2. $0.50\text{V}$
3. $1.23\text{V}$
4. $-0.40\text{V}$

Q. 53. What is the electrode potential of Mg^{+2} \vert Mg electrode at 25 degree Celsius in which the concentration of Mg^{+2} is 0.01M. E _{{Mg^{+2} \vert Mg } }= -2.36V

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

1. $-2.65V$
2. $-2.09V$
3. $+2.65$
4. $2.09V$