NCERT Solutions for Class 12 Chemistry Chapter 3 Electrochemistry plays a pivotal role in the CBSE class 12 Chemistry examination. Chemistry Class 12 solutions chapter 3 is a comprehensive material that has answers to the textbook questions, important questions from previous papers. By studying chemistry class 12 NCERT solutions chapter 3, you will be able to solve different kinds of questions you can expect to appear in the main examination and entrance examinations.
Electrochemistry is the branch of chemistry that deals with the relationship between chemical energy and electrical energy produced in a redox reaction and how they can be converted into each other. The solutions are prepared by the best subject experts. In essence, these solutions can be useful for those students preparing for class 12 board exams and also for JEE advance and other medical entrance exams. Students can successfully answer for numerical problems based on electrochemistry by downloading the free pdf.
Class 12 NCERT Solutions for Electrochemistry
The NCERT solutions for chapter 3 – Electrochemistry has mainly been designed to help the students in preparing well and score good marks in CBSE class 12 Chemistry paper. Further, the solutions consist of well thought or structured questions along with detailed explanations to help students learn and remember concepts easily.
Subtopics for Class 12 Chemistry Chapter 3 – Electrochemistry
 Electrochemical Cells

Galvanic Cells
 Measurement of Electrode Potential

Nernst Equation
 Equilibrium Constant from Nernst Equation
 Electrochemical Cell and Gibbs Energy of Reaction

The conductance of Electrolytic Solutions
 Measurement of the Conductivity of Ionic Solutions
 Variation of Conductivity and Molar Conductivity with Concentration

Electrolytic Cells and Electrolysis
 Products of Electrolysis

Batteries
 Primary Batteries
 Secondary Batteries
 Fuel Cells
 Corrosion
After studying electrochemistry class 12 important questions and solutions, you will be able to describe an electrochemical cell and differentiate between galvanic and electrolytic cells. You will study the application of the Nernst equation for calculating the emf of a galvanic cell and define the standard potential of the cell.
This chapter has derivations of the relation between the standard potential of the cell, Gibbs energy of cell reaction and its equilibrium constant. This solution will give the definition of resistivity (p), conductivity (K) and molar conductivity ( Am) of ionic solutions; differentiate between ionic (electrolytic) and electronic conductivity; describe the method for measurement of conductivity of electrolytic solutions and calculation of their molar conductivity; justify the variation of conductivity and molar conductivity of solutions with change in their concentration and define Aom (molar conductivity at zero concentration or infinite dilution); enunciate Kohlrausch law and learn its applications; understand quantitative aspects of electrolysis; describe the construction of some primary and secondary batteries and fuel cells and explain corrosion as an electrochemical process.
Class 12 Chemistry NCERT Solutions (Electrochemistry) – Important Questions
Q 3.1:
In the order of their reactivity, i.e how they displace each other from their salt solutions, align the metals in decreasing order. Cu, Fe, Al, Zn and Mg.
Answer:
According to their reactivity, the given metals replace the others from their salt solutions in the said order: Mg, Al, Zn, Fe, Cu.
Mg: Al: Zn: Fe: Cu
Q 3.2:
Standard electrode potentials given as,
Mg^{2+}/Mg = −2.37 V, Hg^{2+}/Hg = 0.79V, Cr^{3+}/Cr = − 0.74V, Ag^{+}/Ag = 0.80V, K^{+}/K = −2.93V
In the order of increasing of reducing power arrange the given metals accordingly.
Ans:
The reducing power increases with the lowering of reduction potential. In order of given standard electrode potential (increasing order) : K^{+}/K < Mg^{2+}/Mg < Cr^{3+}/Cr < Hg^{2+}/Hg < Ag^{+}/Ag
Thus, in the order of reducing power, we can arrange the given metals as Ag< Hg < Cr < Mg < K
Q 3.3 :
Represent the galvanic cell in which the following reaction takes place.
Zn(s) + 2Ag^{+}(aq) → Zn^{2+}(aq) + 2Ag(s)
Also find :
(i) The negatively charged electrode ?
(ii) Current carriers in the cell.
(iii) At each electrode, the individual reaction.
Ans :
The galvanic cell in which the given reaction takes place is depicted as:
(i) The negatively charged electrode is the Zn electrode (anode)
(ii) The current carriers in the cell are ions. Current flows to zinc from silver in the external circuit.
(iii) Reaction at the anode is given by :
Reaction at the anode is given by :
Q 3.4:
With the following reactions given, find the standard cell potentials of galvanic cells with given reactions.
(i) 2Cr_{(s)} + 3Cd^{2+}_{(aq)} → 2Cr^{3+}_{(aq)} + 3Cd
(ii) Fe^{2+}_{(aq)} + Ag^{+}_{(aq)} → Fe^{3+}_{(aq)} + Ag_{(s)}
Calculate the ∆_{r}G^{θ} and equilibrium constant of the reactions.
Ans :
(i)
The galvanic cell of the given reaction is depicted as :
Now, the standard cell potential is
= – 0.40 – ( 0.74 )
= + 0.34 V
In the given equation, n = 6
F = 96487 C mol^{−1}
Then,
= −196833.48 CV mol−1
= −196833.48 J mol−1
= −196.83 kJ mol−1
Again,
= 34.496
K = antilog (34.496) = 3.13 × 10^{34}
The galvanic cell of the given reaction is depicted as:
Now, the standard cell potential is
Here, n = 1.
Then,
= −1 × 96487 C mol^{−1} × 0.03 V
= −2894.61 J mol^{−1}
= −2.89 kJ mol^{−1}
Again,
= 0.5073
K = antilog (0.5073)
= 3.2 (approximately)
Q 3.5:
Write the Nernst equation and emf of the following cells at 298 K:
(i) Mg_{(s)}  Mg^{2+}(0.001M)  Cu^{2+}(0.0001 M)  Cu_{(s)}
(ii) Fe_{(s)}  Fe^{2+}(0.001M)  H^{+}(1M)H_{2(g)}(1bar)  Pt_{(s)}
(iii) Sn_{(s)}  Sn^{2+}(0.050 M)  H^{+}(0.020 M)  H_{2(g) }(1 bar)  Pt_{(s)}
(iv) Pt_{(s)}  Br_{2(l)}  Br^{−}(0.010 M)  H^{+}(0.030 M)  H_{2(g)} (1 bar)  Pt_{(s)}.
Answer
(i) For the given reaction, the Nernst equation can be given as:
= 2.7 − 0.02955
= 2.67 V (approximately)
(ii) For the given reaction, the Nernst equation can be given as:
= 0 – ( – 0.14) –
= 0.52865 V
= 0.53 V (approximately)
(iii) For the given reaction, the Nernst equation can be given as:
= 0 – ( – 0.14) –
= 0.14 − 0.0295 × log125
= 0.14 − 0.062
= 0.078 V
= 0.08 V (approximately)
(iv) For the given reaction, the Nernst equation can be given as:
= 0 – 1.09 –
= 1.09 – 0.02955 x
= 1.09 – 0.02955 x
= 1.09 – 0.02955 x
= 1.09 – 0.02955 x (0.0453 + 7)
= 1.09 – 0.208
= 1.298 V
Q 3.6:
The following reaction takes place in the button cells widely used in watches and other devices:
For the given reaction calculate
Ans:
We know that,
= −2 × 96487 × 1.04
= −213043.296 J
= −213.04 kJ
Q 3.7:
For the solution of an electrolyte describe its conductivity and molar conductivity. Also put some light on how they vary with concentration.
Answer
The conductivity of a solution is defined as the conductance of a solution of 1 cm in length and area of crosssection 1 sq. cm. Specific conductance is the inverse of resistivity and it is represented by the symbol κ. If ρ is resistivity, then we can write:
At any given concentration, the conductivity of a solution is defined as the unit volume of solution kept between two platinum electrodes with the unit area of crosssection at a distance of unit length.
When concentration decreases there will a decrease in Conductivity. It is applicable for both weak and strong electrolyte. This is because the number of ions per unit volume that carry the current in a solution decreases with a decrease in concentration.
Molar conductivity –
Molar conductivity of a solution at a given concentration is the conductance of volume V of a solution containing 1 mole of the electrolyte kept between two electrodes with the area of crosssection A and distance of unit length.
Now, l = 1 and A = V (volume containing 1 mole of the electrolyte).
Molar conductivity increases with a decrease in concentration. This is because the total volume V of the solution containing one mole of the electrolyte increases on dilution. The variation of
Q 3.8:
The conductivity of 0.20 M solution of KCl at 298 K is 0.0248 Scm^{−1}. Find its molar conductivity.
Ans :
Given, κ = 0.0248 S cm^{−1} c
= 0.20 M
Molar conductivity,
= 124 Scm^{2}mol^{1}
Q 3.9:
Considering the case of a conductivity cell having 0.001 M KCl solution at 298 K is 1500 Ω. If given, the conductivity of 0.001M KCl solution at 298 K is 0.146 × 10^{−3} S, find the cell constant?
Answer
Given,
Conductivity, k = 0.146 × 10^{−3} S cm−1
Resistance, R = 1500 Ω
Cell constant = k × R
= 0.146 × 10^{−3} × 1500
= 0.219 cm^{−1}
Q 3.10:
The conductivity of NaCl at 298 K has been found at different concentrations and the results are given below:
Concentration/M 0.001 0.010 0.020 0.050 0.100
10^{2} × k/S m^{−1} 1.237 11.85 23.15 55.53 106.74
for all concentrations and draw a plot between
Molar conductivity of
Calculate
Ans:
Given,
κ = 1.237 × 10^{−2} S m−1, c = 0.001 M
Then, κ = 1.237 × 10^{−4} S cm^{−1}, c^{1⁄2} = 0.0316 M^{1/2}
= 123.7 S cm^{2} mol^{−1}
Given,
κ = 11.85 × 10^{−2} S m^{−1}, c = 0.010M
Then, κ = 11.85 × 10^{−4} S cm^{−1}, c^{1⁄2} = 0.1 M^{1/2}
= 118.5 S cm^{2} mol^{−1}
Given,
κ = 23.15 × 10^{−2} S m^{−1}, c = 0.020 M
Then, κ = 23.15 × 10^{−4} S cm^{−1}, c^{1/2} = 0.1414 M^{1/2}
= 115.8 S cm^{2} mol^{−1 }
Given,
κ = 55.53 × 10^{−2} S m^{−1}, c = 0.050 M
Then, κ = 55.53 × 10^{−4} S cm^{−1}, c^{1/2} = 0.2236 M^{1/2}
= 111.1 1 S cm^{2} mol^{−1}
Given,
κ = 106.74 × 10^{−2} S m^{−1}, c = 0.100 M
Then, κ = 106.74 × 10^{−4} S cm^{−1}, c^{1/2} = 0.3162 M^{1/2}
= 106.74 S cm^{2} mol^{−1}
Now, we have the following data :
Since the line interrupts
Q 3.11:
Find the molar conductivity of acetic acid if its conductivity is given to be 0.00241 M. Also, if the value of
Ans:
Given, κ = 7.896 × 10^{−5} S m^{−1} c
= 0.00241 mol L^{−1}
Then, molar conductivity,
=
= 32.76S cm^{2} mol^{−1}
Again,
=
Now,
= 0.084
Dissociation constant,
=
= 1.86 × 10^{−5} mol L^{−1}
Q 3.12:
How much charge is required for the following reductions of 1 mol of :
(i) Al^{3+} to Al.
(ii) Cu^{2+} to Cu.
(iii)
Ans :
(i)
Required charge = 3 F
= 3 × 96487 C
= 289461 C
(ii)
Required charge = 2 F
= 2 × 96487 C
= 192974 C
(iii)
i.e
Required charge = 5 F
= 5 × 96487 C
= 482435 C
Q 3.13:
In terms of Faraday, how much electricity is required to produce :
(i) From molten CaCl_{2}, 20.0 g of Ca.
(ii) From molten Al_{2}O_{3,} 40.0 g of Al.
Ans:
(i) From given data,
Electricity required to produce 40 g of calcium = 2 F
Therefore, electricity required to produce 20 g of calcium = (2 x 20 )/ 40 F
= 1 F
(ii) From given data,
Electricity required to produce 27 g of Al = 3 F
Therefore, electricity required to produce 40 g of Al = ( 3 x 40 )/27 F
= 4.44 F
Q 3.14:
Calculate the amount of electricity required for the oxidation of 1 mol of the following in coulombs :
(i) H_{2}O to O_{2}.
(ii)FeO to Fe_{2}O_{3}.
Ans :
(i) From given data,
We can say that :
Electricity required for the oxidation of 1 mol of H_{2}O to O_{2} = 2 F
= 2 × 96487 C
= 192974 C
(ii) From given data,
Electricity required for the oxidation of 1 mol of FeO to Fe_{2}O_{3} = 1 F
= 96487 C
Q 3.15:
For 20 minutes, a current of 5 A is applied to between platinum electrodes to electrolyze a solution of Ni(NO_{3})_{2}. Find the amount of Ni deposited at the cathode?
Ans :
Given,
Current = 5A
Time = 20 × 60 = 1200 s
Charge = current × time
= 5 × 1200
= 6000 C
According to the reaction,
Nickel deposited by 2 × 96487 C = 58.71 g
Therefore, nickel deposited by 6000 C =
= 1.825 g
Hence, 1.825 g of nickel will be deposited at the cathode.
Q 3.16:
Solutions of 3 electrolytic cells are ZnSO_{4}, AgNO_{3} and CuSO_{4}, cells are connected in series. Of the cells, A, B, C respectively, after passing a steady current of 1.5 amperes, 1.45 g of silver was found deposited at the cathode of cell B. How much time did the current flow? What amount of zinc and copper were deposited?
Ans :
According to the reaction:
i.e., 108 g of Ag is deposited by 96487 C.
Therefore, 1.45 g of Ag is deposited by =
= 1295.43 C
Given,
Current = 1.5 A
Time = 1295.43/ 1.5 s
= 863.6 s
= 864 s
= 14.40 min
Again,
i.e., 2 × 96487 C of charge deposit = 63.5 g of Cu
Therefore, 1295.43 C of charge will deposit
= 0.426 g of Cu
i.e., 2 × 96487 C of charge deposit = 65.4 g of Zn
Therefore, 1295.43 C of charge will deposit
= 0.439 g of Zn
Q 3.17:
Using the standard electrode potentials given in Table 3.1, predict if the reaction between
the following is feasible:
(i)
Fe^{3+}(aq) and I^{−}(aq)
(ii) Ag^{+} (aq) and Cu(s)
(iii) Fe^{3+} (aq) and Br^{−} (aq)
(iv) Ag(s) and Fe^{3+} (aq)
(v) Br_{2} (aq) and Fe^{2+} (aq).
Ans :
(i)
(ii)
E^{0 } is positive, hence reaction is feasible.
(iii)
E^{0 } is negative, hence reaction is not feasible.
(iv)
E^{0 } is negative, hence reaction is not feasible.
(v)
E^{0 } is positive, hence reaction is feasible.
Q 3.18:
Predict the products of electrolysis in each of the following :
(i) An aqueous solution of AgNO_{3} with silver electrodes.
(ii) An aqueous solution of AgNO_{3}with platinum electrodes.
(iii) A dilute solution of H_{2}SO_{4}with platinum electrodes.
(iv) An aqueous solution of CuCl_{2} with platinum electrodes.
Ans:
(i) At cathode:
The following reduction reactions compete to take place at the cathode.
The reaction with a higher value of E^{0} takes place at the cathode. Therefore, deposition of silver will take place at the cathode.
At anode:
The Ag anode is attacked by
(ii) At cathode:
The following reduction reactions compete to take place at the cathode.
The reaction with a higher value of E^{0} takes place at the cathode. Therefore, deposition of silver will take place at the cathode.
At anode:
Since Pt electrodes are inert, the anode is not attacked by
(iii) At the cathode, the following reduction reaction occurs to produce H_{2} gas.
At the anode, the following processes are possible.
For dilute sulphuric acid, reaction (i) is preferred to produce O_{2} gas. But for concentrated sulphuric acid, reaction (ii) occurs.
(iv) At cathode:
The following reduction reactions compete to take place at the cathode.
The reaction with a higher value of takes place at the cathode. Therefore, deposition of copper will take place at the cathode.
At anode:
The following oxidation reactions are possible at the anode.
At the anode, the reaction with a lower value of E^{0} is preferred. But due to the over potential of oxygen, Cl^{−} gets oxidized at the anode to produce Cl_{2} gas.
Also Access 
NCERT Exemplar for class 12 Chemistry Chapter 3 
CBSE Notes for class 12 Chemistry Chapter 3 
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