NEET PYQ Electrochemistry | Chapter-Wise Solutions

Q 1 :
The standard cell potential of the following cell $Zn | {Zn}_{(a q)}^{2 +} {||Fe}_{(a q)}^{2 +} | Fe$ is 0.32 V. Calculate the standard Gibbs energy change for the reaction ${Zn}_{(s)} + {Fe}_{(a q)}^{2 +} ⟶ {Zn}_{(a q)}^{2 +} + {Fe}_{(s)}$ (Given : 1 F = 96487 C) [2024]

$- 61.75$ $kJ$ ${mol}^{- 1}$
$+ 5.006$ $kJ$ ${mol}^{- 1}$
$- 5.006$ $kJ$ ${mol}^{- 1}$
$+ 61.75$ $kJ$ ${mol}^{- 1}$

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(1)

$Δ G ° = - n F E_{cell}^{\circ}$

Given, $E_{cell}^{\circ} = 0.32$ $V$

and 1F = 96487 C

For the given reaction, $n = 2$

$Δ G ° = - 2 \times 96487 \times 0.32 = - 61751.68$ $J$ ${mol}^{- 1}$ $= - 61.75$ $kJ$ ${mol}^{- 1}$

Q 2 :
The $E °$ value for

${Al}^{+} / Al = + 0.55$ $V and {Tl}^{+} / Tl = - 0.34$ $V$

${Al}^{3 +} / Al = - 1.66$ $V and {Tl}^{3 +} / Tl = + 1.26$ $V$

Identify the incorrect statement. [2023]

Al is more electropositive than Tl.
${Tl}^{3 +}$ is a good reducing agent than ${Tl}^{+}$ .
${Al}^{+}$ is unstable in solution.
Tl can be easily oxidised to ${Tl}^{+}$ than ${Tl}^{3 +}$ .

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(2)

${Tl}^{3 +}$ acts as an oxidising agent and not reducing agent.

Q 3 :
Given below are two statements : One is labelled as Assertion A and the other is labelled as Reason R:

Assertion A : In equation $Δ_{r} G = - n F E_{cell}$ , value of $Δ_{r} G$ depends on $n$ .

Reason R : $E_{cell}$ is an intensive property and $Δ_{r} G$ is an extensive property.

In the light of the above statements, choose the correct answer from the options given below. [2023]

A is true but R is false.
A is false but R is true.
Both A and R are true and R is the correct explanation of A.
Both A and R are true and R is not the correct explanation of A.

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(4)

$Δ_{r} G = - n F E_{cell}$

$E_{cell}$ is an intensive parameter but $Δ_{r} G$ is an extensive thermodynamic property and the value of $Δ_{r} G$ depends on $n$ .

Q 4 :
Given below are half-cell reactions:

${MnO}_{4}^{-} + 8 H^{+} + 5 e^{-} ⟶ {Mn}^{2 +} + 4 H_{2} O;$ $E_{{Mn}^{2 +} / {MnO}_{4}^{-}}^{\circ} = - 1.510$ $V$

$\frac{1}{2} O_{2} + 2 H^{+} + 2 e^{-} ⟶ H_{2} O;$ $E_{O_{2} / H_{2} O}^{\circ} = + 1.223$ $V$

Will the permanganate ion, ${MnO}_{4}^{-}$ liberate $O_{2}$ from water in the presence of an acid? [2022]

Yes, because $E_{cell}^{\circ} = + 0.287$ $V$
No, because $E_{cell}^{\circ} = - 0.287$ $V$
Yes, because $E_{cell}^{\circ} = + 2.733$ $V$
No, because $E_{cell}^{\circ} = - 2.733$ $V$

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(1)

$E_{cell}^{\circ} = E_{cathode}^{\circ} - E_{anode}^{\circ} = + 1.510 - 1.223 = + 0.287$ $V$

As $E_{cell}^{\circ}$ is positive, hence the reaction is feasible.

Q 5 :
At 298 K the standard electrode potentials of ${Cu}^{2 +} / Cu, {Zn}^{2 +} / Zn, {Fe}^{2 +} / Fe$ and ${Ag}^{+} / Ag$ are 0.34 V, - 0.76 V, - 0.44 V and 0.80 V respectively.

On the basis of standard electrode potential, predict which of the following reaction cannot occur? [2022]

$C u S O_{4 (a q)} + Z n_{(s)} ⟶ Z n S O_{4 (a q)} + C u_{(s)}$
$C u S O_{4 (a q)} + F e_{(s)} ⟶ F e S O_{4 (a q)} + C u_{(s)}$
$F e S O_{4 (a q)} + Z n_{(s)} ⟶ Z n S O_{4 (a q)} + F e_{(s)}$
$2 C u S O_{4 (a q)} + 2 A g_{(s)} ⟶ 2 C u_{(s)} + A g_{2} S O_{4 (a q)}$

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(4)

The values of standard reduction potential of Cu and Ag suggest that Cu would undergo oxidation (lower reduction potential) and Ag would undergo reduction (higher reduction potential). Hence, the cell reaction will be $Cu + 2 {Ag}^{+} ⟶ {Cu}^{2 +} + 2 Ag$

Q 6 :
Find the emf of the cell in which the following reaction takes place at 298 K

${Ni}_{(s)} + 2 {Ag}^{+} (0.001 M) ⟶ {Ni}^{2 +} (0.001 M) + 2 {Ag}_{(s)}$

Given that $E_{cell}^{\circ} = 10.5$ $V, \frac{2.303 R T}{F} = 0.059 at 298 K$ [2022]

1.0385 V
1.385 V
0.9615 V
1.05 V

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(3)

According to Nernst equation,

$E = E_{cell}^{\circ} - \frac{0.059}{n} \log \frac{[{Ni}^{2 +}]}{{[{Ag}^{+}]}^{2}}$

$E_{cell}^{\circ} = 1.05 (Given)$

Note : Please read 10.5 as 1.05 in question.

$E = 1.05 - \frac{0.059}{2} \log \frac{0.001}{{(0.001)}^{2}}$

$= 1.05 - \frac{0.059}{2} \log 10^{3} = 1.05 - \frac{0.059 \times 3}{2}$

$= 1.05 - 0.0885 = 0.9615$ $V$

Q 7 :
For the cell reaction:

$2 {Fe}_{(a q)}^{3 +} + 2 I_{(a q)}^{-} ⟶ 2 {Fe}_{(a q)}^{2 +} + {I_{2}}_{(a q)}$

$E_{cell}^{\circ} = 0.24 V at 298 K.$ The standard Gibbs' energy $(Δ_{r} G °)$ of the cell reaction is

[Given that Faraday constant, F = 96500 C ${mol}^{- 1}$ ] [2019]

23.16 kJ ${mol}^{- 1}$
- 46.32 kJ ${mol}^{- 1}$
- 23.16 kJ ${mol}^{- 1}$
46.32 kJ ${mol}^{- 1}$

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(2)

The standard Gibbs' energy, $(Δ G °) = - n F E_{cell}^{\circ}$

Value of $n = 2$

$∴$ $Δ G ° = - 2 \times 96500 \times 0.24 = - 46320$ $J$

$= - 46.32$ $kJ/mol$

Q 8 :
For a cell involving one electron, $E_{cell}^{\circ} = 0.59$ $V$ at 298 K, the equilibrium constant for the cell reaction is [Given that $\frac{2.303 R T}{F} = 0.059 V at T = 298 K$ ] [2019]

$1.0 \times 10^{30}$
$1.0 \times 10^{2}$
$1.0 \times 10^{5}$
$1.0 \times 10^{10}$

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(4)

According to Nernst equation,

$E_{cell} = E_{cell}^{\circ} - \frac{0.059}{n} \log Q_{c}$

At equilibrium $E_{cell} = 0,$ $∴ Q_{c} = K_{c}$

$E_{cell}^{\circ} = \frac{0.059}{n} \log K_{c} \Rightarrow 0.59 = \frac{0.059}{1} \log K_{c}$

$K_{c} = antilog 10 \Rightarrow K_{c} = 1 \times 10^{10}$

Q 9 :
In the electrochemical cell:

$Z n | Z n S O_{4} (0.01 M) | | C u S O_{4} (1.0 M) | C u,$

the emf of this Daniell cell is $E_{1}$ . When the concentration of $Z n S O_{4}$ is changed to 1.0 M and that of $C u S O_{4}$ changed to 0.01 M, the emf changes to $E_{2}$ .

From the followings, which one is the relationship between $E_{1}$ and $E_{2}$ ? (Given, $\frac{R T}{F} = 0.059$ ) [2017, 2003]

$E_{1} < E_{2}$
$E_{1} > E_{2}$
$E_{2} = 0 \neq E_{1}$
$E_{1} = E_{2}$

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(2)

$E_{cell} = E_{cell}^{\circ} - \frac{0.059}{n} \log \frac{[{Zn}^{2 +}]}{[{Cu}^{2 +}]}$

$E_{1} = E ° - \frac{0.059}{2} \log \frac{0.01}{1}$

$E_{1} = E ° - \frac{0.059}{2} (- 2) = E ° + 0.059$

$E_{2} = E ° - \frac{0.059}{2} \log \frac{1}{0.01} = E ° - 0.059$

Hence, $E_{1} > E_{2}$

Q 10 :
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 $Δ G °$ and $K_{e q}$ ? [2016, 2011]

$Δ G ° > 0; K_{e q} < 1$
$Δ G ° > 0; K_{e q} > 1$
$Δ G ° < 0; K_{e q} > 1$
$Δ G ° < 0; K_{e q} < 1$

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(1)

$Δ G ° = - n F E_{cell}^{\circ}$

If $E_{cell}^{\circ} = - ve$ then $Δ G ° = + ve i . e .; Δ G ° > 0$ .

$Δ G ° = - n R T \log K_{e q}$

For $Δ G ° = + ve, K_{e q} = - ve i . e ., K_{e q} < 1$ .

Topic Question Set

Q 1 :
The standard cell potential of the following cell $Zn | {Zn}_{(a q)}^{2 +} {||Fe}_{(a q)}^{2 +} | Fe$ is 0.32 V. Calculate the standard Gibbs energy change for the reaction ${Zn}_{(s)} + {Fe}_{(a q)}^{2 +} ⟶ {Zn}_{(a q)}^{2 +} + {Fe}_{(s)}$ (Given : 1 F = 96487 C) [2024]

Q 2 :
The $E °$ value for

${Al}^{+} / Al = + 0.55$ $V and {Tl}^{+} / Tl = - 0.34$ $V$

${Al}^{3 +} / Al = - 1.66$ $V and {Tl}^{3 +} / Tl = + 1.26$ $V$

Identify the incorrect statement. [2023]

Q 5 :
At 298 K the standard electrode potentials of ${Cu}^{2 +} / Cu, {Zn}^{2 +} / Zn, {Fe}^{2 +} / Fe$ and ${Ag}^{+} / Ag$ are 0.34 V, - 0.76 V, - 0.44 V and 0.80 V respectively.

On the basis of standard electrode potential, predict which of the following reaction cannot occur? [2022]

Q 6 :
Find the emf of the cell in which the following reaction takes place at 298 K

${Ni}_{(s)} + 2 {Ag}^{+} (0.001 M) ⟶ {Ni}^{2 +} (0.001 M) + 2 {Ag}_{(s)}$

Given that $E_{cell}^{\circ} = 10.5$ $V, \frac{2.303 R T}{F} = 0.059 at 298 K$ [2022]

Q 7 :
For the cell reaction:

$2 {Fe}_{(a q)}^{3 +} + 2 I_{(a q)}^{-} ⟶ 2 {Fe}_{(a q)}^{2 +} + {I_{2}}_{(a q)}$

$E_{cell}^{\circ} = 0.24 V at 298 K.$ The standard Gibbs' energy $(Δ_{r} G °)$ of the cell reaction is

[Given that Faraday constant, F = 96500 C ${mol}^{- 1}$ ] [2019]

Q 8 :
For a cell involving one electron, $E_{cell}^{\circ} = 0.59$ $V$ at 298 K, the equilibrium constant for the cell reaction is [Given that $\frac{2.303 R T}{F} = 0.059 V at T = 298 K$ ] [2019]

Q 10 :
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 $Δ G °$ and $K_{e q}$ ? [2016, 2011]

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Topic Question Set

Q 1 : The standard cell potential of the following cell Zn|Zn(aq)2+||Fe(aq)2+|Fe is 0.32 V. Calculate the standard Gibbs energy change for the reaction Zn(s)+Fe(aq)2+⟶Zn(aq)2++Fe(s) (Given : 1 F = 96487 C) [2024]

Q 2 : The E° value for Al+/Al=+0.55 V and Tl+/Tl=-0.34 V Al3+/Al=-1.66 V and Tl3+/Tl=+1.26 V Identify the incorrect statement. [2023]

Q 4 : Given below are half-cell reactions: MnO4-+8H++5e-⟶Mn2++4H2O ; EMn2+/MnO4-∘=-1.510 V 12O2+2H++2e-⟶H2O ; EO2/H2O∘=+1.223 V Will the permanganate ion, MnO4- liberate O2 from water in the presence of an acid? [2022]

Q 5 : At 298 K the standard electrode potentials of Cu2+/Cu,Zn2+/Zn,Fe2+/Fe and Ag+/Ag are 0.34 V, - 0.76 V, - 0.44 V and 0.80 V respectively. On the basis of standard electrode potential, predict which of the following reaction cannot occur? [2022]

Q 6 : Find the emf of the cell in which the following reaction takes place at 298 K Ni(s)+2Ag+(0.001M)⟶Ni2+(0.001M)+2Ag(s) Given that Ecell∘=10.5 V, 2.303RTF=0.059 at 298 K [2022]

Q 7 : For the cell reaction: 2Fe(aq)3++2I(aq)-⟶2Fe(aq)2++I2(aq) Ecell∘=0.24 V at 298 K. The standard Gibbs' energy (ΔrG°) of the cell reaction is [Given that Faraday constant, F = 96500 C mol-1] [2019]

Q 8 : For a cell involving one electron, Ecell∘=0.59 V at 298 K, the equilibrium constant for the cell reaction is [Given that 2.303RTF=0.059 V at T=298 K] [2019]

Q 10 : If the Ecell∘ for a given reaction has a negative value, which of the following gives the correct relationships for the values of ΔG° and Keq? [2016, 2011]

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Q 1 :
The standard cell potential of the following cell $Zn | {Zn}_{(a q)}^{2 +} {||Fe}_{(a q)}^{2 +} | Fe$ is 0.32 V. Calculate the standard Gibbs energy change for the reaction ${Zn}_{(s)} + {Fe}_{(a q)}^{2 +} ⟶ {Zn}_{(a q)}^{2 +} + {Fe}_{(s)}$ (Given : 1 F = 96487 C) [2024]

Q 2 :
The $E °$ value for

${Al}^{+} / Al = + 0.55$ $V and {Tl}^{+} / Tl = - 0.34$ $V$

${Al}^{3 +} / Al = - 1.66$ $V and {Tl}^{3 +} / Tl = + 1.26$ $V$

Identify the incorrect statement. [2023]

Q 5 :
At 298 K the standard electrode potentials of ${Cu}^{2 +} / Cu, {Zn}^{2 +} / Zn, {Fe}^{2 +} / Fe$ and ${Ag}^{+} / Ag$ are 0.34 V, - 0.76 V, - 0.44 V and 0.80 V respectively.

On the basis of standard electrode potential, predict which of the following reaction cannot occur? [2022]

Q 6 :
Find the emf of the cell in which the following reaction takes place at 298 K

${Ni}_{(s)} + 2 {Ag}^{+} (0.001 M) ⟶ {Ni}^{2 +} (0.001 M) + 2 {Ag}_{(s)}$

Given that $E_{cell}^{\circ} = 10.5$ $V, \frac{2.303 R T}{F} = 0.059 at 298 K$ [2022]

Q 7 :
For the cell reaction:

$2 {Fe}_{(a q)}^{3 +} + 2 I_{(a q)}^{-} ⟶ 2 {Fe}_{(a q)}^{2 +} + {I_{2}}_{(a q)}$

$E_{cell}^{\circ} = 0.24 V at 298 K.$ The standard Gibbs' energy $(Δ_{r} G °)$ of the cell reaction is

[Given that Faraday constant, F = 96500 C ${mol}^{- 1}$ ] [2019]

Q 8 :
For a cell involving one electron, $E_{cell}^{\circ} = 0.59$ $V$ at 298 K, the equilibrium constant for the cell reaction is [Given that $\frac{2.303 R T}{F} = 0.059 V at T = 298 K$ ] [2019]

Q 10 :
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 $Δ G °$ and $K_{e q}$ ? [2016, 2011]