ELECTROCHEMISTRY

MULTIPLE CHOICE QUESTIONS

1. A cathode and an anode are the most common components of an electrochemical cell. Which of the
following claims about the cathode is correct?
a) Oxidation occurs at the cathode b) Electrons move into the cathode
c) Usually denoted by a negative sign d) Is usually made up of insulating material
2. Which of the following claims about electrochemical cells is true?
a) Cell potential is an extensive property b) Cell potential is an intensive property
c) The Gibbs free energy of an electrochemical cell is an intensive property
d) Gibbs free energy is undefined for an electrochemical cell
3. Which of the following does not belong in the category of electrochemical cells?
a) Voltaic cell b) Photovoltaic cell c) Electrolytic cell d) Fuel Cell
4. Which of the following assertions about the main cell is correct?
a) An example of a primary cell is a mercury cell
b) An example of a primary cell is a nickel-cadmium storage cell
c) The electrode reactions can be reversed d) It can be recharged
5. In a dry cell, which of the following is the electrolyte?
a) Potassium hydroxide b) Sulphuric acid c) Ammonium chloride d) Manganese dioxide
6. Which of the following statements about a lead storage cell (or a lead-acid battery) is false?
a) It is a primary cell b) The cathode is made up of lead(IV) oxide
c) The anode is made up of lead d) The electrolyte used is an aqueous solution of sulphuric acid
7. The conductivity of electrolytic conductors is due to __________
a) Flow of free mobile electrons b) Movement of ions
c) Either movement of electrons or ions d) Cannot be said
8. The process of transmitting electric current through an electrolyte’s solution to decompose it is known as
__________
a) Electrolyte b) Electrode c) Electrolysis d) Electrochemical cell
9. In a fuel cell, which of the following can be utilized as a fuel?
a) Nitrogen b) Argon c) Hydrogen d) Helium
10. Which of the following is given to a fuel cell’s cathode?
a) Hydrogen b) Nitrogen c) Oxygen d) Chlorine
11. The quantity of charge required to obtain one mole of aluminium from Al2O3 is ___________.
(i) 1F (ii) 6F (iii) 3F (iv) 2F
12. The cell constant of a conductivity cell _____________.
(i) changes with change of electrolyte. (ii) changes with change of concentration of electrolyte.
(iii) changes with temperature of electrolyte. (iv) remains constant for a cell.
13. While charging the lead storage battery ______________.
(i) PbSO4 anode is reduced to Pb. (ii) PbSO4 cathode is reduced to Pb.
(iii) PbSO4 cathode is oxidised to Pb. (iv) PbSO4 anode is oxidised to PbO2.

14.
15. In the electrolysis of aqueous sodium chloride solution which of the half cell reaction will occur at anode?

ASSERTION AND REASON TYPE QUESTIONS

(i) Both assertion and reason are true and the reason is the correct explanation of assertion.
(ii) Both assertion and reason are true and the reason is not the correct explanation of assertion.
(iii) Assertion is true but the reason is false.
(iv) Both assertion and reason are false.
(v) Assertion is false but reason is true.
16. Assertion: In an electrochemical cell anode and cathode are respectively negative and positive electrodes.
Reason: At anode oxidation takes place and at cathode reduction takes place.
17. Assertion: Galvanised iron does not rust.
Reason: Zn has more (–) ve electrode potential than Fe.
18. Assertion: Conductivity decreases with dilution.
Reason: The number of ions per unit volume decreases on dilution.
19. Assertion: Ni/Ni2+(1.0M)||Au3+(1.0M)|Au , for this cell emf is 1.75 V if Eo Au3+|Au = 1.50, EoNi/Ni2+ = – 0.25 .
Reason: Emf of the cell = Eo cathode – Eoanode.
20. Assertion: The resistivity for a substance is its resistance when it is one meter long and its area of cross-
section is one square meter.
Reason: The SI units of resistivity is ohm meter.
21. Assertion: Salts like KCl, KNO3 i.e., inert electrolytes are used in salt bridge.
Reason: An inert electrolyte can easily be filled in the U[1]tube.
22. Assertion: On increasing dilution, the specific conductance keeps on increasing.
Reason: On increasing dilution, the degree of ionisation of weak electrolyte increases, and molality of ions
also increases.
23. Assertion: Emf and potential difference are the same for cells.
Reason: Both give the difference in electrode potential under any condition.
24. Assertion: Galvanised iron does not rust.
Reason: Zinc has a more negative electrode potential than iron.
25. Assertion: Current stops flowing when Ecell = 0.
Reason: Equilibrium of the cell reaction is attained.
26. Assertion : Cu is less reactive than hydrogen.

Reason : is negative.
27. Assertion : ECell should have a positive value for the cell to function.
Reason : Ecathode < Eanode
28. Assertion : Conductivity of all electrolytes decreases on dilution.
Reason : On dilution number of ions per unit volume decreases.
29. Assertion : Λm for weak electrolytes shows a sharp increase when the electrolytic solution is diluted.
Reason : For weak electrolytes degree of dissociation increases with dilution of solution.
30. Assertion : Mercury cell does not give steady potential.
Reason : In the cell reaction, ions are not involved in solution.
31. Assertion : Electrolysis of NaCl solution gives chlorine at anode instead of O2.
Reason : Formation of oxygen at anode requires overvoltage.
32. Assertion : For measuring resistance of an ionic solution an AC source is used.
Reason : Concentration of ionic solution will change if DC source is used.
33. Assertion : Current stops flowing when ECell = 0.
Reason : Equilibrium of the cell reaction is attained.
34. Assertion : EAg+/Ag increases with increase in concentration of Ag+ ions.
Reason : EAg+/Ag has a positive value.
35. Assertion : Copper sulphate can be stored in zinc vessel.
Reason : Zinc is less reactive than copper.
NERNST EQUATION AND GIBBS ENERGY
36. Two half cell reactions of an electrochemical cell are given below :
MnO–4(aq) + 8H+ (aq) + 5e– → Mn2+ (aq) + 4H2O (I), E° = + 1.51 V
Sn2+ (aq) → 4 Sn4+ (aq) + 2e–, E° = + 0.15 V
Construct the redox equation from the two half cell reactions and predict if this reaction favours formation of
reactants or product shown in the equation. (ANSWER = 1.36 V)
37. Given that the standard electrode potentials (E°) of metals are :
K+/K = -2.93 V, Ag+/Ag = 0.80 V, Cu2+/Cu = 0.34 V,
Mg2+/Mg = -2.37 V, Cr3+/Cr = -0.74 V, Fe2+/Fe = -0.44 V.
Arrange these metals in increasing order of their reducing power. More negative the value of standard
electrode potentials of metals is, more will be the reducing power.
38. Determine the values of equilibrium constant (Kc) and ΔG° for the following reaction :
Ni(s) + 2Ag+ (aq) → Ni2+ (aq) + 2Ag(s), E° = 1.05 V (1F = 96500 C mol-1) (ANSWER = -202650 J/mol and
K = 0.35 *107)
39. The standard electrode potential (E°) for Daniel cell is +1.1 V. Calculate the ΔG° for the reaction
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
(1 F = 96500 C mol-1). (ANSWER = -212.3 KJ/mol)
40. A copper-silver cell is set up. The copper ion concentration in it is 0.10 M. The concentration of silver ion is
not known. The cell potential is measured 0,422 V. Determine the concentration of silver ion in the cell.
Given : E°Ag+/Ag = + 0.80 V, E° Cu2+/Cu = + 0.34 V. (ANSWER = 13.93 M)
41. A voltaic cell is set up at 25°C with the following half cells :
Al/Al3+ (0.001 M) and Ni/Ni2+ (0.50 M)
Write an equation for the reaction that occurs when the cell generates an electric current and determine the
cell potential. E0Ni2+/Ni=−0.25V and E0Al3+/Al=−1.66V (Log 8 × 10-6 = -0.54) (ANSWER = 1.46 V)
42. The cell in which the following reaction occurs :
2Fe3+ (aq) + 2I– (aq) → 2Fe2+ (aq) + I2 (s) has E0cell = 0.236V at 298K. Calculate the standard Gibbs energy
and the equilibrium constant of the cell reaction.
(Antilog of 6.5 = 3.162 × 106; of 8.0 = 10 × 108; of 8.5 = 3.162 × 108) (ANSWER = 0.529 V)
43. Calculate the emf of the following cell at 25°C : Ag(s) | Ag+ (10-3 M) || Cu2+ (10-1 M) | Cu(s) Given E0cell =
+0.46 V and log 10n = n. (ANSWER = 0.31 V)

CONDUCTIVITY
44. Express the relation among cell constant, resistance of the solution in the cell and conductivity of the
solution. How is molar conductivity of a solution related to its conductivity?
45. Question 13.
The molar conductivity of a 1.5 M solution of an electrolyte is found to be 138.9 S cm2 mol-1. Calculate the
conductivity of this solution. (ANSWER = 0.20835 S cm-1)
46. The conductivity of 0.20 M solution of KCl at 298 K is 0.025 S cm-1. Calculate its molar conductivity.
(ANSWER Λm = 125 S cm2 mol-1)
47. The conductivity of 0.001 M acetic acid is 4 × 10-5 S/cm. Calculate the dissociation constant of acetic acid, if
molar conductivity at infinite dilution for acetic acid is 390 S cm2/mol. (1.46 *10-6)
48. State Kohlrausch law of independent migration of ions. Why does the conductivity of a solution decrease
with dilution?
49. Define conductivity and molar conductivity for the solution of an electrolyte. Discuss their Variation with
change in temperature.
50. Calculate the degree of dissociation (a) of acetic acid if its molar conductivity (Λm) is 39.05 S cm2 mol-1.
Given: λ°(H+) = 349.6 S cm2 mol-1 and λ°(CH3COO–) = 40.9 S cm2 mol-1 (ANSWER = 0.1)
51. The electrical resistance of a column of 0.05 M NaOH solution of diameter 1 cm and length 50 cm is 5.55 ×
103 ohm. Calculate its resistivity, conductivity and molar conductivity. (ANSWER = 229.6 S cm2/mol)
52. When a certain conductance cell was filled with 0.1 M KCl, it has a resistance of 85 ohms at 25°C. When the
same cell was filled with an aqueous solution of 0.052 M unknown electrolyte, the resistance was 96 ohms.
 Calculate the molar conductance of the electrolyte at this concentration.
[Specific conductance of 0.1 M KCl = 1.29 × 10-2 ohm-1 cm-1] (ANSWER = 219.23 S cm2/mol)
53. Conductivity of 2.5 × 10-4 M methanoic acid is 5.25 × 10-5 S cm-1. Calculate its molar conductivity and
degree of dissociation. Given : λ0(H+) = 349.5 Scm2 mol-1 and λ0(HCOO–) = 50.5 Scm2 mol-1. (ANSWER =
0.525)

ELECTROLYSIS
54. (a) Following reactions occur at cathode during the electrolysis of aqueous silver chloride solution :
Ag+(aq) + e– → Ag(s) E° = +0.80 V
H+(aq) + e– → 12H2(g) E° = 0.00 V
On the basis of their standard reduction electrode potential (E°) values, which reaction is feasible at the
cathode and why?
55. Calculate the time to deposit 1.27 g of copper at cathode when a current of 2A was passed through the
solution of CuSO4. (ANSWER = 1930 s) (Molar mass of Cu = 63.5 g mol-1,1 F = 96500 C mol-1)
56. Following reactions may occur at cathode during the electrolysis of aqueous CuCl2 solution using Pt
electrodes:

On the basis of their standard electrode potential values, which reaction is feasible at cathode and why?
57. How many coulombs are required to reduce 1 mole Cr2O72- to Cr3+?
58. How much electricity in terms of Faradays is required to produce 20 g of calcium from molten CaCl2?
59. A current was passed for 5 hours through two electrolytic cells connected in series. The first cell contains
AuCl3 and second cell CuSO4 solution. If 9.85 g of gold was deposited in the first cell, what amount of
copper gets deposited in the second cell? Also calculate magnitude of current in ampere.
Given: Atomic mass of Au = 197 amu and Cu = 63.5 amu. (AMOUNT = 4.76 g of Cu and I = 0.8 A)
60. (i) Calculate the mass of Ag deposited at cathode when a current of 2 amperes was passed through a solution
of AgNO3 for 15 minutes. [Given: Molar mass of Ag = 108 g mol-1 1F = 96,500 C mol-1)
(ii) Define fuel cell. (ANSWER = 2.0145 g)
61. (a) How many moles of mercury will be produced by electrolysing 1.0 M Hg(NO3)2 solution with a current
of 2.00 A for 3 hours? [Hg(NO3)2 = 200.6 g mol-1] (ANSWER = 0.1119 mole)
(b) A voltaic cell is set up at 25 °C with the following half-cells Al2+ (0.001 M) and Ni2+ (0.50 M). Write an
equation for the reaction that occurs when the cell generates an electric current and determine the cell
potential. (Given : E0Ni2+/Ni = – 0.25 V, E0Al3+/Al = – 1.66 V) (ANSWER = 1.46 V)
62. Predict the products of electrolysis in each of the following:
(i) An aqueous solution of AgNO3 with platinum electrodes.
(ii) An aqueous solution of H2SO4 with platinum electrodes.
(b) Estimate the minimum potential difference needed to reduce Al2O3 at 500°C. The Gibbs energy change
for the decomposition reaction 23 Al2O3 → 43 Al + O2 is 960 kJ
(F = 96500 C mol-1) (ANSWER = -2.487 V)

63. CELL AND BATTERIES

64. The chemistry of corrosion of iron is essentially an electrochemical phenomenon. Explain the reactions
occurring during the corrosion of iron in the atmosphere.
65. Write the reactions taking place at cathode and anode in lead storage battery when the battery is in use. What
happens on charging the battery?
66. Set up Nernst equation for the standard dry cell. Using this equation show that the voltage of a dry cell has to
decrease with use.
67. From the given cells: Lead storage cell, Mercury cell, Fuel cell and Dry cell
Answer the following: (i) Which cell is used in hearing aids? (ii) Which cell was used in Apollo Space
Programme?
(iii) Which cell is used in automobiles and inverters? (iv) Which cell does not have long life?
68. What are fuel cells? Explain the electrode reactions involved in the working of H2 – O2 fuel cell.
69. Why an electrochemical cell stop working after some time? The reduction potential of an electrode depends
upon the concentration of solution with which it is in contact.