Class 12 Chemistry Chapter 3 Notes

Electrochemistry is a fine combination of, or the relationship between electrical and chemical reactions. The class 12 Chemistry chapter 3 notes cover a brief outline of electrochemical cells, Nernst equation, Gibbs energy of cell reaction, conductivity, Kohlrausch law and its applications, electrolysis, etc. The notes are quite helpful to students who would want to get in-depth knowledge on the topic. Chemistry Class 12 Chapter 3 notes will be helpful for an easy revision of the derivation of the Nernst equation and its uses. 

Chemistry Class 12 chapter 3 notes are an efficient way to revise and evaluate yourself in the electrochemistry unit. The galvanic cells convert the chemical energy of the spontaneous reaction into electrical energy. But, a Daniel cell converts the chemical energy produced during redox reactions into electrical energy. Chapter 3 chemistry class 12 notes constitute all the essential formulas of the electrochemistry unit. 

Students can access the Class 12 Chemistry Chapter 3 notes on Extramarks.

Key Topics Covered In Class 12 Chemistry Chapter 3 Notes

The class 12 chemistry chapter 3 notes is a comprehensive material. We start with an introduction to electrochemistry, the electrical properties of conductors and their types and continue with their functions, applications, and relationship with temperature. Next, we will explain electrode potential, batteries, fuel cells, and corrosion. The main topics covered in chemistry chapter 3 class 12 notes are given in detail below.  

Introduction to ElectroChemistry

Electrochemistry is the phenomenon of generating electricity from the energy produced during a spontaneous chemical reaction and electrical energy to bring about nonspontaneous chemical transformations.

  • It is used for the purification of metals.
  • The process is used in batteries and fuel cells that convert chemical energy into electrical energy and several instruments and devices.
  • This process is also used in electroplating.
  • The reactions carried out using the process of electrochemistry are energy effective and less polluting.

Electrochemical Cells

A spontaneous chemical reaction can occur on its own, and in such a reaction, the system’s Gibbs energy decreases. The energy is then converted into electrical energy. It is also feasible to force nonspontaneous processes by providing external energy with electrical energy. These Cells are used to carry out these interconversions. A simple example of an electrochemical cell is a standard 1.5V cell used to power many electrical appliances such as tv remotes, toys, and watches. Such cells capable of producing an electric current from the chemical reactions are called Galvanic or Voltaic cells. Those that create chemical reactions via electrolysis are known as electrolytic cells. Students may refer to Class 12 Chemistry Chapter 3 notes for a quick understanding of electrochemical cells.

Electrochemical cells usually consist of a cathode and an anode.

  • Cathode: It represents a positive sign since electrons are absorbed here. A reduction reaction always occurs in the cathode of an electrochemical cell.
  • Anode:  It represents a negative sign since electrons are released here. Oxidation reaction occurs here, and electrons move out of the anode. 

Half cells and cell potential

Under this section of class 12 chemistry chapter 3 notes, students will learn about cell potential. Electrochemical cells connect two half cells, each consisting of an electrode dipped in the electrolyte. This same electrolyte can be applied to half cells. A Salt bridge is associated with these half cells, which come up with a platform to form ionic contact between them, not allowing them to fuse. An excellent example of a salt bridge is a filter paper dipped in potassium nitrate solution, KCl Or NaCl solutions.

One of the half cells of the electrochemical cell loses electrons due to oxidation, and the other gains one electron in a reduction process. As a result, an equilibrium reaction occurs in both the half cells, and once the equilibrium reaches net voltage becomes zero, the cell stops producing electricity.

The tendency of an electrode in contact with an electrolyte to gain or lose an electron is explained by its electrode potential. Therefore, the value of these potentials can predict the overall cell potential. Commonly, the electrode potentials are measured with the help of the standard hydrogen electrode as a reference electrode (known electrode potential), given in  Class 12 Chemistry Chapter 3 notes.

The two types of electrochemical cells are

1. Galvanic cells (also known as Voltaic cells)

2. Electrolytic cells

Galvanic cells convert chemical energy into electrical energy, and electrolytic cells convert electrical energy into chemical energy. The key differences between Galvanic cells and electrolytic cells are shown below.

Galvanic Cells: In a galvanic cell, a spontaneous reaction occurs, and the chemical energy is converted into electrical energy. It is also known as the Voltaic cell or Daniel cell. A spontaneous chemical process or reaction extracts cell energy and transforms it into an electric current. The redox reactions that take place are spontaneous. The anode charged is negative in these electrochemical cells, and the cathode is positively charged. The electrons originate from the species that undergo oxidation.

E.g. a Daniell Cell is a Galvanic Cell in which the redox reaction is carried out using Zn and Cu.

               In redox reaction:  Zn(s) + Cu2+ (aq)  → Zn2+ (aq )+ Cu(s)

               Oxidation Half reaction(Anode) :  Zn(s)  → Zn2 + (aq) + 2e

              Reduction Half reaction (Cathode): Cu2 + (aq) + 2e  →  Cu(s)

The reducing agent is Zn2 +, and the oxidising agent is Cu2+.

An anode is the oxidation half, as well a Cathode is a reduction half. A cathode is used to describe a type of electrode. In the external circuit, electrons transfer from anode to cathode. The negative polarity is assigned to the anode, and the positive polarity is given to the cathode. Daniell Cell is a fictional character developed by Daniell Cell. The anode is Zinc, while the cathode is Copper.

Electrolytic Cell: Generally, these electrodes are submerged in an electrolytic solution comprising cations and anions. When current is passed, the ions migrate towards electrodes of opposite polarity, where they undergo simultaneous reduction and oxidation. 

The nonspontaneous redox reaction is carried out by electrical energy in an electrolytic cell. Energy input is required for the redox reactions to proceed in these cells, i.e. the reactions are nonspontaneous. These cells feature an anode positively charged and a negatively charged cathode. Electrons originate from an external source  (such as a battery).

It will help students understand the brief knowledge of electrochemical cells. Students may refer to various study materials in addition to Class 12 Chemistry Chapter 3 notes.

Measurement of electrode potential

The electrode is connected to a (SHE) standard hydrogen electrode to constitute a cell. The electrode comprising the negative terminal of the cell is allotted a negative value of electrode potential, whereas the electrode comprising the positive terminal of the cell is allotted a positive value of electrode potential.

The potential difference developed between the two terminals is measured using a potentiometer. The direction in which the flow of electric current in the external circuit is identified using a galvanometer.

Ecell = Ecathode – Eanode

In the case of Daniel’s cell

At the anode: Oxidation shows loss of electron Zn –> Zn2+ + 2e

At cathode: Reduction shows gain of electrons Cu2+ + 2e  –> Cu

The overall reaction of this cell is the sum of above two half-cell reactions.and we obtain the equation is: Zn(solid) + Cu2+ (aq) –> Zn2+ (aq) + Cu(s)

The difference between the electrode potential of 2 electrodes constituting an electrochemical cell is called the electromotive force(EMF) of a cell.

Emf of the cell is

                    E0cell = E0cathode  –  E0anode

     Emf of the cell  = 0.34 V – (- 0.76) V = 1.10 Volt

(Measured Emf of Cu is 0.34 Volt and that of Zn is 0.76 Volt).

Sometimes metals like Pt or Au are used as inert electrodes. They do not involve the reaction but provide their surface for oxidation and reduction reactions. Students may refer to NCERT Solutions parallel to Class 12 Chemistry Chapter 3 notes for a more detailed explanation.

Standard Electrode potential: The potential difference created between a metal electrode and solutions of ions of unit molarity(1M) at one atmosphere and 250 C is called standard electrode potential, denoted as E0.

Reference Electrode: Known potential of the electrode is called the reference electrode. It may be a primary reference electrode, i.e. Hydrogen electrode or a secondary reference electrode, i.e. calomel electrode.

SHE, also known as normal hydrogen electrode (NHE), consists of platinum wire carrying Pt foil coated with finely divided Pt black. The wire is sealed into the glass tube and placed in a beaker containing 1M acidic solution(HCl). The pure hydrogen gas 1 atm is bubbled through the solution at 298 K.

Half cell is Pt H2(1 atm) H+(1M)

Ecell = Ecathode – Eanode

Ecell = Ecathode – 0 = Ecathode

The measured Emf of the cell:

Pt| H2 (g, 1 bar)| H+ (aq, 1M) || Cu2+ (aq, 1M)| Cu is 0.34 V.

The positive significance of the standard electrode potential signifies the easy reduction of Cu2+ ions than H+ ions.

The measured Emf of the cell

Pt| H2 (1 bar)| H+ (1M) || Zn2+ (1M)| Zn is -0.76 V.

The negative value of the standard electrode potential signifies that the H+ ions oxidise the Zn ions).

Nernst Equation

Under this section of Class 12 Chemistry Chapter 3 notes, students get a detailed explanation of the Nernst Equation. The Nernst equation is named by” German physicist Walther Nernst”. This equation guarantees cell potential under nonstandard conditions, relates the measured potential to the reaction quotient, and permits the exact measurement of equilibrium constants.

Let us consider an electrochemical reaction shown in the following type

                            aA + bB    ⇾  cC + dD

The equation can be written as follows.  

Ecell =  Ecell – RT/ nF lnQ

       = Ecell – RT/ nF ln [C]c [D]d


                                                              [A]a [B]b

E= emf of cell,  n = Number of electron exchange, F = Faraday constant

 E cell – RT/ 2F ln [Zn2+]



In the case of Daniel’s cell, the Nernst equation is shown as follows.

Ecell = Ecell – RT/ nF lnQ

This equation implies that the value increases with the increase in the concentration of Copper ion and decreases the concentration of Zn ion. Substituting the values of R, F at T= 298 K. the equation becomes 

 E cell – 0.059/ 2 log [Zn2+]



If the circuit in Daniel cell is closed:

Zn(solid) + Cu2+ (aq) →Zn2+ (aq) + Cu(s)

There is no change in the concentration of Cu2+ and Zn2+ ions after some time, and the voltmeter gives zero reading. At this point, equilibrium has been attained.

So the Nernst equation may be written as:

E =  E cell – 2.303 RT/ 2F log [Zn2+]



Or Ecell = 2.303 RT/ 2F log [Zn2+]



But at equilibrium,  [Zn2+]



and at T = 298 K. The above equation can be rewritten as


 Ecell = 0.059 V/2 log Kc                                                                              {Ecell =1.1V}

 Ecell = log Kc =1.12 x 2 /0.059 V =37.28

Kc =2 x 1037 at 298 K

Students may refer to various study materials such as NCERT Solutions, and CBSE revision notes, in addition to Class 12 Chemistry Chapter 3 notes for a more detailed explanation.

Electrochemical cell and Gibbs Energy reaction:

Electrical work done in 1 second equals electrical potential multiplied by the total charge passed. Passing the charges reversibly through the galvanic cell results in maximum work. Reversible work done using a galvanic cell is equivalent to a decrease in Gibbs energy; therefore, if the emf of the cell is E and nF is the amount of charge passed, and ΔrG = Gibbs energy of the reaction, then.

        ΔrG = -nFEcell

For the reaction, Ecell is an intensive parameter, but Gibbs energy is an extensive thermodynamic property, and the value depends on n,

Thus if we write the reaction

        Zn(s) + Cu2+ (aq) –> Zn2+ (aq) + Cu(s)

        rG = -2FEcell  ]

But when we write the equation 

2Zn(s) + 2Cu2+ (aq) –> 2Zn2+ (aq) + 2Cu(s)

  [ΔrG = -4FEcell ]

Electrolytic conductors (electrolysis)

The class 12 chemistry chapter 3 notes define electrolysis in dept to benefit the students. Conductance is because of the movement of ions towards oppositely charged electrodes. It causes the transfer of matter and hence decomposition.

  • Electrolytes that are ionised completely are strong electrolytes, for example, sodium chloride, Hydrochloric acid, potassium chloride, sodium hydroxide and ammonium nitrate etc.
  • Electrolytes that do not ionise completely in solutions are weak electrolytes. For example, Acetic acid, carbonic acid, Zinc chloride, Mercuric chloride, Hydrogen cyanide, Ammonium hydroxide etc.
  1. i) Ohm law: The current (I) carried by a conductor or electrolytic solution is directly proportional to the potential difference(V) between the two ends of the conductor.

           Vα I

           V = RI

  1. ii) Electrical Resistance: The electrical resistance is directly proportional to its length, Inversely proportional to its A, area of cross-section. It is denoted by the symbol ‘R’. It is measured in a SI unit called ohm (Ω) or volt/ ampere with the help of a Wheatstone bridge.

                   R α I/A

               So  R = ⲣ I/ A

In the equation, the constant of proportionality denoted by rho is resistivity. It is measured in units called ohm metres.When length is 1m and cross-sectional area = 1m2 . Then resistivity becomes the resistance.

1 Ω m = 100 Ω cm

1 Ω cm = 0.01 Ω m

iii) Conductance: It is the property of a conductor(metallic or electrolytic) that facilitates the flow of electricity. It is inversely proportional to resistance, and the si unit is ohm-1 or mho.

G = I /R = A / ⲣ I 

  1. iv) Specific conductance (conductivity): It is a reciprocal of specific resistance.

? = 1 / ⲣ  =  I /R .A = 1/R. I/A

It is denoted by the symbol (Greek, kappa). It is measured in a unit called Sm–1.

Then conductivity becomes the conductance.

The conductivity of an electrolytic(ionic) solution depends on

  • The nature of the electrolyte added
  • Size of the ions formed and their solvation
  • The nature of the solvent as well as its viscosity
  • The concentration of the electrolyte
  •  Temperature

There are various study materials available in addition to class 12 Chemistry Chapter 3 notes for a more detailed explanation of conductivity.

Measurement of resistance of an electrolyte solution 

The cell constant: For any cell, the ratio of the distance between electrodes (l) and the area of electrodes  (A) is constant, known as the cell constant. The symbol G* denotes it.

           G*= (I/ A)

              = I/ ?A

It is determined by measuring the resistance of the cell containing a solution of known conductivity.

Variation of Conductivity and Molar Conductivity along with concentration: They depend on the concentration of the electrolyte. The Conductivity and Molar Conductivity of both weak and strong electrolytes decrease with a decrease in concentration as the number of ions per unit volume carrying the current in a solution decrease on dilution.

The conductivity of a solution at a specific concentration = Conductance of solution placed in between the two platinum electrodes where

Volume of solution = 1 unit

Cross sectional area of electrodes = 1unit

Distance = 1unit

G = ?A / I = V =  ?

Molar conductivity of a solution at a specific concentration = Conductance of solution where 

Volume of solution = V

Concentration of electrolyte = 1 mole

Cross sectional area of electrodes = A

Distance = 1unit

       Λm =  ?A / I =  ?

When the concentration approaches zero, the molar conductivity is known as limiting molar conductivity and is represented by the symbol Λ0m

Kohlrausch’s law

Class 12 Chemistry Chapter 3 notes also offer a lot of information on Kohlrausch’s law of independent migration of ions. Kohlrausch’s law of independent migration of ions expresses that limiting the molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and cation of the electrolyte. Limiting the equivalent molar conductivity of an electrolyte is the algebraic sum of determining the molar equivalent conductivities of its constituent ions.


Λo electrolytes = λ0+ + λ0 

Where λo+ = limiting equivalent conductivities of cation and λo is limiting comparable conductivities of the anion.

Faraday’s laws of electrolysis

Michael Faraday explained the quantitative features of electrolysis and came forward with two laws:

First law: The amount of substance deposited or liberated at any electrode during electrolysis by a current is proportional to the quantity of electricity moved through the electrolyte (solution or melt).

Second law: The amounts of various substances liberated by the same quantity of electricity passing through the electrolytic solution are proportional to their chemical equivalent weights

Therefore, the Atomic Mass of Metal / Number of electrons is required to reduce the cation.

                      Q = It

here Q = coulomb

I = ampere and

t = second.

This quantity of electricity is called Faraday and is represented by the symbol F. In the above example, the application of more current results in the deposition of more copper from the anode to the cathode. Apart from class 12 Chemistry Chapter 3 notes, students may refer to various other study materials such as CBSE previous year question papers, important questions and CBSE revision notes while preparing for the examination.

Products of Electrolysis

The product of electrolysis produced depends on the nature of the material being electrolysed and the type of electrodes being used. Inert electrodes, e.g., platinum or gold, does not participate in a chemical reaction and act as a source or sink for electrons. If the electrode is reactive, it engages in the electrode reaction. It also depends on the different oxidising and reducing species present in the electrolytic cell and their standard electrode potentials. Electrolysis of molten sodium chloride results in sodium metal and chlorine gas production.

Electrolysis  reactions may be summarized as:

NaCl (aq) –> Na+ (aq) + Cl (aq)

  Cathode: H2O(l) + e  –> 1/2  H2(g) + OH (aq)

 Anode: Cl (aq) –>1/2 Cl2(g) + e

Net reaction may be summarised as:

NaCl(aq) + H2O(l) –> Na+(aq) + OH(aq) + 1/2H2(g) +1/2 Cl2(g)

While class 12 Chemistry chapter 3 notes provide a detailed explanation, students may also refer to various other study materials.


A battery is a collection of two or more electrochemical cells connected in series. These are the following types of cells(commercial):

a)Primary battery: That battery that cannot be recharged is a primary cell or battery. E.g. Dry cells,  Mercury cells. The dry cell consists of a zinc anode and the carbon (graphite) cathode surrounded by powdered MnO2 and carbon. The area between the two electrodes is filled with a soaking paste of NH4Cl and ZnCl2.

                 Zn(s) ⎯ –>Zn2+ + 2e              In   Anode     

                  MnO2+ + NH4++ e–     ⎯ →    MnO (OH) + NH3     In Cathode

In the reaction at the cathode, Mn  is reduced from + 4 oxidation state to +3 oxidation state.

 Ammonia is produced and forms a complex with Zn to give [Zn (NH3)4]2+. The cell potential is 1.5 V. Mercury cell consists of Zinc – mercury amalgam as anode and a paste of HgO and carbon as the cathode.

A paste of potassium hydroxide and zinc oxide electrolyte is used as an electrolyte.

As in class 12 Chemistry chapter 3 notes, the Electrode reactions for the cell are shown below:

In  Anode:  Zn(Hg) + 2OH →  ZnO(s) + H2O + 2e

In  Cathode: HgO + H2O + 2e →  Hg(l) + 2OH

The net  reaction is represented by

Zn (Hg) + HgO(s) –> ZnO(s) + Hg(l)

The cell potential is approx 1.35 V, which remains constant throughout its lifetime.

Secondary battery: Secondary battery is one of the main topics under class 12 Chemistry chapter 3 notes. The reaction in a secondary battery happens many times. Once it drains, it can be recharged and again used. It can be restored by passing an electric current through it in the opposite direction. For  Example, secondary cells are Pb and Ni-Cd cells. They are commonly used in automobiles and inverters. The most crucial secondary cell is the lead storage battery. The cell contains a lead anode and lead dioxide (PbO2) as cathode, and Sulphuric acid is used as an electrolyte (38 %solution).

The cell reactions if the battery is in use

In  Anode: Pb(s) + SO42–(aq) –> PbSO4(s) + 2e–  (E0=0.296 V)

In Cathode: PbO2(s) + SO42–(aq) + 4H+(aq) + 2e –> PbSO4 (s) + 2H2O (l)  (E0=1.628V)

i.e., net reaction is-

          Pb(s)+PbO2(s)+2H2SO4(aq) –> 2PbSO4(s) + 2H2O(l) (E0=1.924V)

When charging, the reaction  is reversed and PbSO4(s) on anode and cathode is converted into Pb and PbO2, respectively. Its life is longer than lead storage cell and it requires more expenses to

Another important secondary cell is Ni-Cd The overall reaction is:

Cd (s)+ 2Ni(OH)3 (s) –> CdO (s) + 2Ni(OH)2 (s) +H2O(l)

Students may refer to NCERT Solutions along with class 12 Chemistry chapter 3 notes for a more detailed explanation.

Fuel cells: Galvanic cells designed to convert the energy of combustion of fuel, i.e. hydrogen, methane, methanol, etc., directly into electrical energy are known as fuel cells.

Hydrogen oxygen fuel cell: Electrode is porous graphite impregnated with catalyst(Pt, Ag or metal oxide). An electrolyte is an aqueous solution of KOH or NaOH.

The  reaction of Hydrogen and oxygen is given below

At Cathode: O2 (g) + 2H2O (l) + 4e –> 4OH(aq) (Reduction half cell reaction)

At Anode: 2H2 (g) + 4OH(aq) ⎯ –> 4H2O(l) + 4e (Oxidation half cell reaction)

The overall reaction reaction is:

2H2(g) + O2(g) →2 H2O(l)

The EMF of the cell is 1 V .

Corrosion: Corrosion is one of the most vital topics of class 12 Chemistry chapter 3 notes. Metals react with atmospheric oxygen and produce basic metal oxides because they react with water to form bases.

In the case of rusting of iron, the iron reacts with the oxygen present in air and moisture and develops rust (hydrated iron (III) oxide). More details of corrosion properties are shown in NCERT class 12 chapter 3 Electrochemistry.

4Fe + 3O2 + 2 H2O  → 2Fe2O3

Factors that promote Corrosion:

  • Presence of impurity
  • Reactivity of metal ion
  • Presence of air and moisture
  • Strains in metals
  • Presence of electrolyte

Prevention of corrosion:

  • Barrier protection: Apply a thin film of oil, grease, paint, copper-tin etc. if the coating is broken, rust will appear iron.
  • Covering with more reactive metal(Sacrificial electrode): This is the electrochemical method to cover iron’s surface by layering other more active metals like Mg, Zn, etc. 
  • Galvanisation, electrical protection, and anti-rust solutions methods are also used to prevent corrosion.

Students may refer to CBSE revision notes, CBSE sample papers, important questions and more apart from Class 12 Chemistry Chapter 3 notes for more details on fuel cells and corrosion.

Class 12 Chemistry Chapter 3: Exercises & Answer Solutions

Students may refer to class 12 Chemistry chapter 3 notes on Electrochemistry on Extramarks. The exercise and answer solutions are explained in detail to help students understand the various concepts mentioned in the chapter. Every minute detail that a student may need to understand electrochemistry is clearly explained in the notes.

Furthermore, students can click on the links provided below to access the study material they may need to prepare for the examination. In addition, to exercise and answer solutions, at Extramarks, we provide sample question papers, past year question papers, essential questions, revision notes, and more short notes. 

CBSE Syllabus

CBSE Revision Notes

CBSE Sample Papers

CBSE Extra Questions

CBSE Previous Year Question Papers

Students may find various information under one channel, making studying easier, especially during board exams. 

NCERT Class 12 Chemistry Exemplar for Chapter 3 

The Class 12 Chemistry Exemplar for Chapter 3 Electrochemistry adds to students’ learning skills and tests their information recall, comprehension, analytical thinking, and problem-solving ability. It is the most comprehensive study material students can rely on to study, practice various questions, and prepare for board and various other entrance examinations.

NCERT Class 12 Chemistry exemplar for chapter 3 has questions and answers from the exemplar book together with extra questions, electrochemistry Class 12 problems, various chemistry problems and solutions for Class 12, MCQs, exercises, assignments, and important questions.

Key Features of NCERT Solutions Class 12 Chemistry Chapter 3 Notes

ElectroChemistry is an essential topic for exams like IIT, JEE or NEET. Hence while studying Chemistry, the best quality notes and solutions provided by Extramarks are helpful. Class 12 chemistry chapter 3 notes can help students get a deeper understanding of the various topics covered under the chapter. 

  • They are prepared by educated and experienced subject matter experts.
  • These important notes include a detailed explanation of every topic present in the chapter. 
  • They can be used to study right after school or used as revision notes during examinations. 
  • The class 12 chemistry chapter 3 notes are helpful for all types of school examinations or various competitive examinations.

FAQs (Frequently Asked Questions)

1. How important is Chemistry chapter 3 for the CBSE Board exam?

Chemistry may seem like an intimidating subject. In addition, chapter 3 is a vital topic that might seem very difficult. It is crucial to note that Electrochemistry has a lot of weightage, and students can expect questions from this chapter in the board examinations. 

2. Does class 12 chemistry chapter 3 notes cover all topics of Electrochemistry?

Electrochemistry is a vast topic, but every detail is covered in the Extramarks class 12 chemistry chapter 3 notes.