Class 12 Chemistry Chapter 3 Electrochemistry Notes

The Electrochemistry chapter in CBSE Class 12 Chemistry is one of the most important physical chemistry topics. It explains how chemical energy is converted to electrical energy and vice versa through redox reactions.

This chapter is essential for understanding electrochemical cells, galvanic and electrolytic processes, standard electrode potential, the Nernst equation, and batteries.

These Class 12 Electrochemistry Notes are designed to simplify tough formulas, clarify reactions, and help you solve numericals confidently. Use them for regular revision and to strengthen your concepts for boards and competitive exams like NEET, JEE, etc.

Electrochemistry Class 12 Notes Material PDF Download

This study material for Class 12 explains the basics of Electrochemistry in an easy-to-understand way. Download the PDF to learn key concepts and prepare well for your exams.

Below we have provided the links to downloadable PDFs of class 12 science Ch 3 notes and get an in-depth explanation and understanding of the chapter.

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Important Notes from Electrochemistry Class 12

Let’s revise your chapter 3 Electrochemistry together, and be better at what we’re studying:

1. Redox Reactions and Electrochemical Cells

• Redox Reaction: A reaction where oxidation and reduction occur simultaneously.
• Oxidation: Loss of electrons
• Reduction: Gain of electrons

Electrochemical Cell: A device that converts chemical energy into electrical energy via redox reactions.

Types:

• Galvanic (Voltaic) Cell: Spontaneous redox reaction produces electric current.
• Electrolytic Cell: Non-spontaneous redox reaction driven by external electricity.

2. Representation of a Galvanic Cell

Cell Notation (for Zn-Cu cell):
Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

• Anode: Zn (oxidation)
• Cathode: Cu (reduction)
• Salt Bridge: Maintains charge neutrality by allowing ion flow without mixing solutions.

3. Electrode Potential

The potential difference between a metal electrode and its ion solution.

Types:

1. Standard Electrode Potential (E°): Measured under standard conditions (1M, 298K, 1 atm).

2. Electrochemical Series: A list of elements arranged by their E° values.

• Higher E° → better oxidizing agent
• Lower E° → better reducing agent

4. Cell Potential and EMF

Cell Potential (Ecell): The voltage generated by the redox reaction in the cell.

Ecell = E°(cathode) – E°(anode)

• If Ecell > 0 → Reaction is spontaneous
• If Ecell < 0 → Reaction is nonspontaneous

5. Nernst Equation

Used to calculate the electrode potential under non-standard conditions.

E = E° – (0.0591/n) × log([products]/[reactants]); where:

• E = electrode potential under given conditions
• E° = standard electrode potential
• n = number of electrons transferred
• T = temperature (K)
• [ ] = concentrations

Example:

For Zn²⁺ + 2e⁻ → Zn,

E = E° – (0.0591/2) log (1/[Zn²⁺])

6. Electrochemical Series & Applications

• Helps in predicting the direction of redox reactions.
• Determines feasibility of redox reactions.
• Predicts strengths of oxidizing and reducing agents.

7. Electrolytic Cells & Faraday’s Laws

1. Faraday’s First Law: The mass of substance deposited or liberated is directly proportional to the charge passed. m = Z × Q = Z × I × t

2. Faraday’s Second Law: When the same quantity of electricity is passed through different electrolytes, the masses of substances deposited are proportional to their equivalent weights.

8. Batteries

• Primary Cells: Non-rechargeable (e.g., dry cell, mercury cell)
• Secondary Cells: Rechargeable (e.g., lead-acid battery, lithium-ion battery)

Example – Lead Acid Battery:

• Anode: Pb
• Cathode: PbO₂
• Electrolyte: H₂SO₄
• Recharged by reversing the current.

9. Fuel Cells

Electrochemical cells that convert the energy of combustion directly into electricity.

Hydrogen-Oxygen Fuel Cell:

• Reactants: H₂ and O₂
• Product: Water
• Highly efficient and eco-friendly.

10. Corrosion

• Electrochemical process involving oxidation of metals in presence of air and moisture.
• Example: Rusting of iron.
• Prevention: Galvanization, painting, coating with oil, sacrificial anodes.

Important Questions from Electrochemistry Class 12

Given below are important questions from the chapter, which you should know by heart:

1. Very Short Answer Questions (1 Mark Each)

Q: Define cell potential.
Ans: Cell potential is the difference in electric potential between two electrodes of a galvanic cell.

Q: What is the role of a salt bridge?
Ans: A salt bridge maintains electrical neutrality by allowing the flow of ions between the half-cells.

Q: What does a positive E° value signify?
Ans: A positive E° value indicates that the electrode has a greater tendency to gain electrons (undergo reduction).

2. Short Answer Questions (2–3 Marks Each)

Q: Differentiate between galvanic and electrolytic cells.
Ans: 

‍

Q: Write the Nernst equation and define the terms.
Ans: Nernst Equation: E = E∘ − (00.0591/n) ​logQ; where:

E = electrode potential under non-standard conditions

E∘ = standard electrode potential

n = number of electrons transferred

Q = reaction quotient

Q: How is the electrochemical series useful?

Ans: Electrochemical Series Usefulness:

• Predicts feasibility of redox reactions
• Helps determine oxidizing/reducing strength
• Used to calculate cell EMF values

3. Long Answer Questions (4–5 Marks Each)

Q: Explain construction and working of the lead storage battery.

Ans: Lead Storage Battery:
Construction: Consists of lead (Pb) as anode and lead dioxide (PbO₂) as cathode, immersed in sulphuric acid. 

Working:

• Discharge:
  Pb + PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O
• Recharge: Reverse of discharge reaction using external current.

Q: What is corrosion? Explain its mechanism and prevention methods.

Ans: Corrosion is the slow deterioration of metals due to chemical reactions with the environment.
Mechanism (e.g., rusting of iron):

• At anode: Fe →  Fe²⁺  + 2e^-
• At cathode: O₂ + 4H⁺  +  4e^-   → 2H₂O
• Overall: 4Fe + 3O₂  + 6H₂O → 4Fe(OH)₃  → rust

 Prevention Methods:

• Painting/coating
• Galvanization
• Alloying
• Using sacrificial anodes (e.g., Zn with Fe)

Common Mistakes to Avoid

🚫 Reversing the electrode potentials when calculating EMF (anode and cathode confusion).

🚫 Ignoring ion concentrations in Nernst equation applications.

🚫 Assuming all batteries are rechargeable – know which are primary and secondary.

🚫 Miscalculating ‘n’ (number of electrons) in redox reactions.

🚫 Forgetting unit consistency (especially time in seconds, current in amperes).

Creative Ways to Make Notes for Electrochemistry Chapter

Just writing down pointers is not enough, notes should be catchy enough to remember things thoroughly, here are ways to make notes more useful for you guys:

1. Reaction Maps: Show electron flow in galvanic/electrolytic cells.
2. Quick-reference Tables: For standard electrode potentials (E° values).
3. Process Diagrams: Labelled illustrations for batteries, electrolysis setups.
4. Formula Sheets: Collect Nernst, Faraday’s Law, EMF equations.
5. Redox Flowcharts: Highlight oxidation and reduction steps visually.

How Can Notes Help?

• Boost conceptual clarity for one of the most numerical-heavy chapters.
• Provide clear formula applications and reaction-based summaries.
• Help recall standard values and proper signs of E° in EMF calculations.
• Sharpen accuracy for competitive exams and board numericals.
• Speed up last-minute revision with key point summaries.

Electrochemistry forms the backbone of understanding modern energy systems and electrochemical processes in real life. With well-structured notes, you can master redox reactions, batteries, and EMF calculations effectively. Whether you're preparing for CBSE board exams or JEE/NEET, these notes on Electrochemistry are your go-to resource for clarity and precision.

Frequently Asked Questions

1. What is Electrochemistry?

Answer: Electrochemistry is the branch of chemistry that deals with the relationship between chemical reactions and electrical energy. It involves the study of redox reactions, where oxidation and reduction occur, and the conversion of chemical energy into electrical energy (as in galvanic cells) or the use of electrical energy to drive nonspontaneous chemical reactions (as in electrolytic cells).

2. What is a Galvanic Cell?

Answer: A Galvanic cell (or Voltaic cell) is a device that converts chemical energy into electrical energy through spontaneous redox reactions. It consists of two half-cells, each containing a metal electrode and an electrolyte, connected by a wire and a salt bridge to allow the flow of electrons.

3. What is an Electrolytic Cell?

Answer: An Electrolytic cell is a device that uses electrical energy to drive non-spontaneous chemical reactions. In this cell, electrical energy is applied to drive reactions like the electrolysis of water, electroplating, or the extraction of metals from their ores. It consists of two electrodes and an electrolyte, with a power supply driving the reaction.

4. What is Electrode Potential?

Answer: Electrode potential is the potential difference between an electrode and its surrounding electrolyte solution when the electrode is at equilibrium with the ions in the solution. It indicates the tendency of an electrode to gain or lose electrons and is measured in volts. Standard electrode potentials are given relative to the standard hydrogen electrode (SHE).

5. What is the Concept of the Standard Hydrogen Electrode (SHE)?

Answer: The Standard Hydrogen Electrode (SHE) is a reference electrode used to measure the electrode potentials of other electrodes. It is defined as having a potential of 0 volts at all temperatures. The SHE consists of a platinum electrode in contact with 1 M H⁺ ions (usually in the form of H₂ gas at 1 atm pressure) at 25°C.