Chapter-3 Electrochemistry
Important Questions
Q1. Why does the conductivity of a solution decrease with dilution
A1. The conductivity of a solution is the conductance of ions present in a unit volume of the solution. On dilution, the number of ions per unit volume decreases. Hence, the conductivity decreases.
Q2) What does the negative sign the expression E0Zn2+ / Zn = -0.76 V mean?
A2) Negative sign shows that zinc is more reactive than hydrogen. This means that when the zinc electrode is connected to SHE, Zn will be oxidized to Zn2+ and H+ will be reduced to H2.
Q3) State Kohlrausch laws of independent migration of ions?
A3) It states that “ At infinite dilution when the dissociation of the electrolyte is complete, each ion makes a definite contribution to the total molar conductivity of the electrolyte irrespective of the nature of the other ion with which it is associated.
Q4) Write two advantages of H2O2 fuel cell over ordinary cell?
or
What advantage do the fuel cells have over primary and secondary batteries?
or
Name the type of cell which was used in the Apollo space program for providing electrical power?
A4) (a) Fuel cells can be run continuously so long as the reactants are supplied, primary batteries become dead and secondary batteries take a long time for recharging.
(b) Fuel cells do not cause any pollution H2–O2 fuel cell was used in the Apollo Space program for providing electrical power.
Q5) Express the relation between the conductivity of a solution in the cell, the cell constant, and the resistance of solution in the cell?
A5) K = 1 x l
R A
Where K = Conductivity
l/A = Cell constant
R = Resistance.
Q6) What is electrode potential?
A6) It is defined as the tendency of an electrode to lose or gain electrons when it is in contact with a solution of its own ions.
Q7) Is it safe to stir AgNO3 solution with a copper spoon? Why or Why not?
Given :- Eo Ag+/Ag = 0.8 volt and Eo cu2+ /Cu = 0.34 volt
A7) The given values of reduction potentials show that Cu is more reactive than Ag. i.e Cu reacts with AgNO3 solution. Hence it is not safe to stir AgNO3 solution with a copper spoon.
Q8) What is the role of ZnCl2 in a dry cell?
A8) ZnCl2 combines with the NH3 produced to form the complex salt [Zn (NH3) 2 Cl2] as otherwise, the pressure developed due to NH3 would crack the seal of the cell.
Q9) State faraday’s first law of electrolysis?
A9) The mass of any substance (w) deposited or liberated at an electrode is directly proportional to the quantity of electricity (Q) passed through the electrolyte or W α Q
W = ZQ
Z = electrochemical equivalent
Q10. Calculate the potential of the hydrogen electrode in contact with a solution whose pH is 10.
A10. H⁺ + e → ½ H2
Applying Nernst eqn
EH+/1/2 H2 = EoH+ ½ H2 -0.0591 log _1_
n [H⁺]
= 0 – 0.0591log _1_
1 10⁻¹⁰
because pH = 10 means [H+] = 10⁻¹⁰ M.
= - 0 .0591 x -10
= + 0.591V
Q11. Mention the reactions occurring at (i) Anode (ii) Cathode, during the working of a mercury cell. Why does the voltage of a mercury cell remain constant during its operation?
A 11.
At Anode :-
Zn (Hg) + 2OH- → ZnO (s) + H2O + 2e-
At Cathode
HgO + H2O + 2e- → Hg (l) + 2OH⁻¹
______________________________
Zn(Hg) + HgO (s) ZnO (s) + Hg(l)
The voltage of a mercury cell remains constant during its life as the overall reaction does not involve any ion in a solution whose concentration can change during its life.
Q12. Calculate G for the reaction, Mg (s) + Cu2+ (aq) Mg2+ (aq) + Cu (s)
Given Eo cell = + 2.71V 1F = 96500 (Cmol -1)
A12. For the reaction
Mg(s) + Cu2+ (aq) → Mg 2+ (g) + Cu (s)
n = 2
△G°= -nFE° cell
= -2 x 96500 x 2.71
= - 523030 J mol⁻¹
Q13. Calculate the degree of dissociation of acetic acid at 298K given that
λm° (CH3 COOH) = 11.7 Scm² mol⁻¹
λm° (CH3 COO-) = 40.9 Scm² mol⁻¹
λm° (H⁺) = 349.1 Scm² mol⁻¹
A13. Degree of dissociation
α = λm
λ°m
λM° = λM° (CH3 COO-) + λM (H⁺)
= 40.9 + 349.1
= 390 Scm² mol⁻¹
α = λm = 11.7
λm° 390
α = 3 x 10⁻²
Q14. A solution of Ni(NO3)2 is electrolyzed between platinum electrodes using a current of 5 A for 20 min. What mass of nickel will be deposited at the cathode?
Given : Atomic mass of Ni = 58.7 gmol⁻¹ 1F = 96,500 (mol⁻¹)
A14. Given :- I = 5A Time (t) = 20 x 60 = 1200s
Q = I x t = 5 x 1200 = 6000C.
M = Z x I x t
M = Eq. wt x I x t
96,500
M = 58.7 x 6000 = 1.82g
2 x 96500
Q15. (a) Define electrochemical series.
(b) Given that the standard electrode potentials (E°) metals are:-
K⁺/K = -2.93V, Ag⁺/Ag = 0.8V, Cu ²⁺/Cu = 0.34V, Mg²⁺/Mg = -2.37V,
Cr³⁺/Cr = - 0.74, Fe²⁺/Fe = -0.44V.
Arrange these metals in increasing order of their reducing power.
A15. a) The arrangement of the various electrodes in order of their increasing values of standard
reduction potentials are called electrochemical series.
(b) Greater the negative value of the standard electrode potential (Eo) greater is the reducing power of the electrode. Thus the increasing order of reducing power is
Ag⁺/Ag < Cu²⁺/Cu < Fe²⁺/Fe < Cr³⁺/Cr < Mg²⁺/Mg < K⁺/K
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