Electrochemistry · Chemistry · NEET
MCQ (Single Correct Answer)
From the following select the one which is not an example of corrosion.
The standard cell potential of the following cell $$\mathrm{Zn}\left|\mathrm{Zn}^{2+}(\mathrm{aq})\right| \mathrm{Fe}^{2+}(\mathrm{aq}) \mid \mathrm{Fe}$$ is $$0.32 \mathrm{~V}$$. Calculate the standard Gibbs energy change for the reaction:
$$\mathrm{Zn}(\mathrm{s})+\mathrm{Fe}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{Fe}(\mathrm{s})$$
(Given : $$1 \mathrm{~F}=96487 \mathrm{C}$$)
Match List I with List II.
| List I (Conversion) |
List II (Number of Faraday required) |
||
|---|---|---|---|
| A. | 1 mole of H$$_2$$O to O$$_2$$ | I. | 3F |
| B. | 1 mol of MnO$$_4^-$$ to Mn$$^{2+}$$ | II. | 2F |
| C. | 1.5 mol of Ca from molten CaCl$$_2$$ | III. | 1F |
| D. | 1 mol of FeO to Fe$$_2$$O$$_3$$ | IV. | 5F |
Choose the correct answer from the options given below :
Mass in grams of copper deposited by passing 9.6487 A current through a voltmeter containing copper sulphate solution for 100 seconds is (Given : Molar mass of $$\mathrm{Cu}: 63 \mathrm{~g} \mathrm{~mol}^{-1}, 1 \mathrm{~F}=96487 \mathrm{C}$$)
The $$\mathrm{E}^{\Theta}$$ values for
$$\begin{aligned} & \mathrm{Al}^{+} / \mathrm{Al}=+0.55 \mathrm{~V} \text { and } \mathrm{Tl}^{+} / \mathrm{Tl}=-0.34 \mathrm{~V} \\ & \mathrm{Al}^{3+} / \mathrm{Al}=-1.66 \mathrm{~V} \text { and } \mathrm{T}^{3+} / \mathrm{Tl}=+1.26 \mathrm{~V} \end{aligned}$$
Identify the incorrect statement
Molar conductance of an electrolyte increase with dilution according to the equation:
$$\Lambda_{\mathrm{m}}=\Lambda_{\mathrm{m}}^{\circ}-\mathrm{A} \sqrt{\mathrm{c}}$$
Which of the following statements are true?
(A) This equation applies to both strong and weak electrolytes.
(B) Value of the constant $$\mathrm{A}$$ depends upon the nature of the solvent.
(C) Value of constant $$\mathrm{A}$$ is same for both $$\mathrm{BaCl}_2$$ and $$\mathrm{MgSO}_4$$
(D) Value of constant $$\mathrm{A}$$ is same for both $$\mathrm{BaCl}_2$$ and $$\mathrm{Mg}(\mathrm{OH})_2$$
Choose the most appropriate answer from the options given below:
The correct value of cell potential in volt for the reaction that occurs when the following two half cells are connected, is
$$\begin{aligned} & \mathrm{Fe}_{(\mathrm{aq})}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Fe}(\mathrm{s}), \mathrm{E}^{\circ}=-0.44 \mathrm{~V} \\ & \mathrm{Cr}_2 \mathrm{O}_7^{2-} \text { (aq) }+14 \mathrm{H}^{+}+6 e^{-} \rightarrow 2 \mathrm{Cr}^{3+}+7 \mathrm{H}_2 \mathrm{O} \\ & \mathrm{E}^{\circ}=+1.33 \mathrm{~V} \end{aligned}$$
The conductivity of centimolar solution of $$\mathrm{KCl}$$ at $$25^{\circ} \mathrm{C}$$ is $$0.0210 ~\mathrm{ohm}^{-1} \mathrm{~cm}^{-1}$$ and the resistance of the cell containing the solution at $$25^{\circ} \mathrm{C}$$ is $$60 ~\mathrm{ohm}$$. The value of cell constant is -
Given below are two statements: one is labelled as Assertion A and the other is labelled as Reason R:
Assertion A : In equation $$\mathrm{\Delta_rG=-nFE_{cell}}$$, value of $$\mathrm{\Delta_rG}$$ depends on n.
Reason R : $$\mathrm{E_{cell}}$$ is an intensive property and $$\mathrm{\Delta_rG}$$ is an extensive property.
In the light of the above statements, choose the correct answer from the options given below:
Two half cell reactions are given below.
$$C{o^{3 + }} + {e^ - } \to C{o^{2 + }},\,\,\,\,\,\,\,\,\,E_{C{o^{2 + }}/C{o^{3 + }}}^0 = - 1.81\,V$$
$$2A{l^{3 + }} + 6{e^ - } \to 2Al(s),\,\,\,E_{Al/A{l^{3 + }}}^0 = + 1.66\,V$$
The standard EMF of a cell with feasible redox reaction will be :
Standard electrode potential for the cell with cell reaction
Zn(s) + Cu2+(aq) $$\to$$ Zn2+(aq) + Cu(s)
is 1.1 V. Calculate the standard Gibbs energy change for the cell reaction. (Given F = 96487 C mol$$-$$1)
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 V and 0.80 V, respectively.
On the basis of standard electrode potential, predict which of the following reaction cannot occur?
Given below are half cell reactions:
$$MnO_4^ - + 8{H^ + } + 5{e^ - } \to M{n^{2 + }} + 4{H_2}O$$,
$$E_{M{n^{2 + }}/MnO_4^ - }^o = - 1.510\,V$$
$${1 \over 2}{O_2} + 2{H^ + } + 2{e^ - } \to {H_2}O$$
$$E_{{O_2}/{H_2}O}^o = + 1.223\,V$$
Will the permanganate ion, $$MnO_4^ - $$ liberate O2 from water in the presence of an acid?
Find the emf of the cell in which the following reaction takes place at 298 K
Ni(s) + 2Ag+ (0.001 M) $$\to$$ Ni2+ (0.001 M) + 2Ag(s)
(Given that E$$_{cell}^o$$ = 10.5 V, $${{2.303\,RT} \over F} = 0.059$$ at 298 K)
[$$\Lambda _{{H^ + }}^o$$ = 350 S cm2 mol$$-$$1
$$\Lambda _{C{H_3}CO{O^ - }}^o$$ = 50 S cm2 mol$$-$$1]
[Given that $${{2.303RT} \over F}$$ = 0.059 V at T = 298 K ]
2Fe3+(aq) + 2I– (aq) $$ \to $$ 2Fe2+(aq) + I2(aq)
$${E_{cell}^\Theta }$$ = 0.24 V at 298 K. The standard Gibbs energy ($$\Delta $$rGo) of the cell reaction is :
[Given that Faraday constant F = 96500 C mol–1]
Then the species undergoing disproportionation is :
$$Zn\left| {ZnS{O_4}\left( {0.01\,M} \right)} \right|$$$$\left| {CuS{O_4}\left( {1.0M} \right)} \right|Cu,$$
the emf of this Daniell cell is E1. When the concentration of ZnSO4 is changed to 1.0 M and that of CuSO4 changed to 0.01 M, the emf changes to E2. From the followings, which one is the relationship between E1 and E2? (Given, RT/F = 0.059)
Mn2+ + 2e$$-$$ $$ \to $$ Mn, Eo = $$-$$1.18 V
Mn2+ $$ \to $$ Mn3+ + e$$-$$, Eo = $$-$$ 1.51 V
The $$E$$o for the reaction 3 Mn2+ $$ \to $$ Mno + 2Mn3+, and possibility of the forward reaction are respectively
Zn(s) + Ag2O(s) + H2O(l) $$\rightleftharpoons$$ 2Ag(s) + Zn2+(aq) + 2OH$$-$$(aq)
If half cell potentials are
Zn2+(aq) + 2e$$-$$ $$ \to $$ Zn(s); Eo = $$-$$0.76 V
Ag2O(s) + H2O(l) + 2e$$-$$ $$ \to $$ 2Ag(s) + 2OH$$-$$(aq), Eo = 0.34 V
The cell potential will be
$${2 \over 3}$$ Al2O3 $$ \to $$ $${4 \over 3}$$ Al + O2
$$\Delta $$rG = +960 kJ mol$$-$$1
The potential difference needed for the electrolytic reduction of aluminium oxide (Al2O3) at 500oC is at least
F2(g) + 2e$$-$$ $$ \to $$ 2F$$-$$(aq) ; Eo = + 2.85 V
Cl2(g) + 2e$$-$$ $$ \to $$ 2Cl$$-$$(aq) ; Eo = + 1.36 V
Br2(l) + 2e$$-$$ $$ \to $$ 2Br$$-$$(aq) ; Eo = + 1.06 V
I2(s) + 2e$$-$$ $$ \to $$ 2I$$-$$(aq) ; Eo = + 0.53 V
The strongest oxidising and reducing agents 23 respectively are
$$\left[ {} \right.$$i.e. $$\Lambda _{m\left( {N{H_4}OH} \right)}^0$$$$\left. {} \right]$$ is equal to
The favourable redox reaction is
and Cu+(aq) + e$$-$$ $$ \to $$ Cu(s) are + 0.15 V and + 0.50 V respectively.
The value of Eocu2+/cu will be
(i) EMF of cell = (Oxidation potential of anode) $$-$$ (Reduction potential of cathode)
(ii) EMF of cell = (Oxidation potential of anode) + (Reduction potential of cathode)
(iii) EMF of cell = (Reductional potential of anode) + (Reduction potential of cathode)
(iv) EMF of cell = (Oxidation potential of anode) $$-$$ (Oxidation potential of cathode)
Which of the above relations are correct?
(F = 96500 C mol$$-$$1)
(i) Cu2+ + 2e$$-$$ $$ \to $$ Cu, Eo = 0.337 V
(ii) Cu2+ + e$$-$$ $$ \to $$ Cu+, Eo = 0.153 V
Electrode potential, Eo for the reaction,
Cu+ + e$$-$$ $$ \to $$ Cu, will be
[Fe(CN)6]4$$-$$ $$ \to $$ [Fe(CN)6]3$$-$$ + e$$-$$; Eo = $$-$$0.35 V
Fe2+ $$ \to $$ Fe3+ + e$$-$$; Eo = $$-$$0.77 V
Cu(s) + 2Ag+(aq) $$ \to $$ Cu2+(aq) + 2Ag(s);
Eo = 0.46 V at 298 K is
$$A\left| {{A^ + }\left( {xM} \right)} \right|\left| {{B^ + }\left( {yM} \right)} \right|B$$
The emf measured is + 0.20 V. The cell reaction is
(Atomic mass : Al = 27)
4/3Al + O2 $$ \to $$ 2/3Al2O3, $$\Delta $$G = $$-$$ 827 kJ mol$$-$$1 of O2,
the minimum e.m.f. required to carry out an electrolysis of Al2O3 is
(F = 96500 C mol$$-$$1)
When the concentration of ZnSO4 is 1.0 M and that of CuSO4 is 0.01 M, the e.m.f. changed to E2. What is the relationship between E1 and E2?
Fe2+/Fe [Eo = $$-$$0.44] and
Fe3+/Fe2+[ Eo = 0.77];
If Fe2+, Fe3+ and Fe blocks are kept together, then
2Cu+ $$ \to $$ Cu2+ + Cu, Eo is
(Given Eo for Cu2+/Cu is 0.34 V and
Eo for Cu2+/Cu+ is 0.15 V.)