Electrochemistry is one of the most important chapters in Class 12 Chemistry because it connects electricity + chemical reactions. This chapter is very scoring in Boards and also very useful for NEET/JEE.
In these notes, we will cover every topic step-by-step like electrochemical cell, galvanic cell, electrolytic cell, Nernst equation, EMF, conductance, Kohlrausch law, batteries and corrosion with all important formulas and concepts.
Class 12 Electrochemistry Notes PDF | Nernst Equation, EMF, Numericals & Important Questions
✅1. What is Electrochemistry?
Electrochemistry is the branch of chemistry which deals with:
- Conversion of chemical energy into electrical energy
- Conversion of electrical energy into chemical energy
👉 It includes:
- Electrochemical cells (Galvanic cells) → produce electricity
- Electrolytic cells → use electricity to perform reactions
✅ 2. Redox Reaction and Electron Transfer
Electrochemical reactions are always Redox reactions.
Oxidation
- Loss of electrons
- Increase in oxidation number
Reduction
- Gain of electrons
- Decrease in oxidation number
📌 Example:
Zn \rightarrow Zn^{2+} + 2e^- \quad (Oxidation)
Cu^{2+} + 2e^- \rightarrow Cu \quad (Reduction)
✅ 3. Electrochemical Cell
An electrochemical cell is a device that converts chemical energy into electrical energy.
Main Parts
- Two electrodes: Anode and Cathode
- Electrolyte solutions
- Salt bridge / porous partition
- External wire for electron flow
✅ 4. Galvanic Cell (Voltaic Cell)
A Galvanic cell is a cell in which spontaneous redox reaction produces electricity.
🔥 Example: Daniell Cell
It consists of:
- Zn electrode in ZnSO₄ solution
- Cu electrode in CuSO₄ solution
Cell notation:
Zn(s) | Zn^{2+}(aq) || Cu^{2+}(aq) | Cu(s)
Working of Daniell Cell
- Zn is oxidized (Anode)
- Cu²⁺ is reduced (Cathode)
Anode Reaction:
Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-
Cathode Reaction:
Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)
Overall Reaction:
Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)
✅ 5. Anode and Cathode (Most Important)
📌 In Galvanic Cell
- Anode = Oxidation = Negative electrode
- Cathode = Reduction = Positive electrode
📌 In Electrolytic Cell
- Anode = Oxidation = Positive electrode
- Cathode = Reduction = Negative electrode
👉 Rule Always Same: ✅ Oxidation at Anode
✅ Reduction at Cathode
✅ 6. Electron Flow and Current Flow
- Electrons always flow from Anode → Cathode
- Conventional current flows from Cathode → Anode
✅ 7. Salt Bridge
A salt bridge is a U-shaped tube containing a gel of electrolyte like:
- KCl
- KNO₃
- NH₄NO₃
Functions of Salt Bridge
✅ Maintains electrical neutrality
✅ Completes the circuit
✅ Prevents mixing of two solutions directly
✅ Allows ion migration
👉 In Daniell cell:
- Anions move towards anode
- Cations move towards cathode
✅ 8. Electrode Potential
Each electrode has a tendency to:
- lose electrons (oxidation)
- gain electrons (reduction)
This tendency is measured by electrode potential.
✅ 9. Standard Electrode Potential (E°)
Electrode potential measured under standard conditions:
Standard Conditions
- Concentration = 1 M
- Pressure (for gases) = 1 atm
- Temperature = 298 K
📌 Reference electrode: Standard Hydrogen Electrode (SHE)
Its electrode potential is:
E^\circ_{SHE} = 0.00V
✅ 10. Standard Cell Potential / EMF (E°cell)
E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}
📌 For Daniell Cell:
E^\circ_{cell} = 0.34 - (-0.76) = 1.10V
✅ 11. Relationship Between Cell Potential and Spontaneity
A reaction is spontaneous if:
E^\circ_{cell} > 0
If:
E^\circ_{cell} < 0✅ 12. Gibbs Free Energy and Cell Potential
Very important formula:
\Delta G^\circ = -nFE^\circ_{cell}
Where:
- n = number of electrons transferred
- F = Faraday constant = 96500 C mol⁻¹
- E°cell = standard cell potential
📌 If → → Spontaneous
✅ 13. Nernst Equation (Most Important Topic)
Nernst equation tells cell potential at non-standard conditions.
General Nernst Equation
E_{cell} = E^\circ_{cell} - \frac{RT}{nF}\ln Q
At 298 K:
E_{cell} = E^\circ_{cell} - \frac{0.0591}{n}\log Q
Where:
- Q = reaction quotient
- n = electrons transferred
🔥 Nernst for Daniell Cell
Reaction:
Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu
Q = \frac{[Zn^{2+}]}{[Cu^{2+}]}
So,
E_{cell} = E^\circ_{cell} - \frac{0.0591}{2}\log\left(\frac{[Zn^{2+}]}{[Cu^{2+}]}\right)✅ 14. Equilibrium Constant (K) and EMF
At equilibrium:
E_{cell} = 0
So,
E^\circ_{cell} = \frac{0.0591}{n}\log K✅ 15. Faraday’s Laws of Electrolysis
Electrolysis = chemical change caused by electricity.
First Law
Mass deposited is directly proportional to charge passed:
m \propto Q
m = ZQ
Second Law
For same charge, masses deposited are proportional to their equivalent weights.
✅ 16. Charge and Faraday Constant
Charge passed:
Q = It
- I = current (ampere)
- t = time (seconds)
Faraday constant:
F = 96500\ C/mol
✅ 17. Electrochemical Equivalent (Z)
Z = \frac{E}{F}
✅ 18. Conductance (G)
Conductance is the ability of a solution to conduct electricity.
G = \frac{1}{R}
Unit: Siemens (S) or ohm⁻¹
✅ 19. Specific Conductance / Conductivity (κ)
\kappa = \frac{1}{\rho}
\kappa = \frac{l}{RA}
Unit: S m⁻¹
✅ 20. Molar Conductivity (Λm)
\Lambda_m = \frac{\kappa \times 1000}{C}
Unit: S cm² mol⁻¹
Where C = concentration in mol/L
Trend:
- Λm increases on dilution
- For strong electrolyte: increases slowly
- For weak electrolyte: increases sharply
✅ 21. Kohlrausch’s Law
At infinite dilution:
\Lambda_m^\circ = \lambda_+^\circ + \lambda_-^\circ
Meaning: limiting molar conductivity is sum of ionic conductivities.
✅ 22. Degree of Dissociation (α)
For weak electrolyte:
\alpha = \frac{\Lambda_m}{\Lambda_m^\circ}
✅ 23. Ostwald Dilution Law
For weak electrolyte:
K_a = \frac{C\alpha^2}{1-\alpha}
✅ 24. Batteries (Important in Boards)
(A) Primary Batteries
Not rechargeable.
Example: Dry cell
- Anode: Zn
- Cathode: carbon rod
- Electrolyte: NH₄Cl paste
(B) Secondary Batteries
Rechargeable.
Example: Lead storage battery Used in cars.
✅ 25. Fuel Cell
Example: Hydrogen-Oxygen Fuel Cell
- Produces electricity by combining H₂ and O₂.
- Eco-friendly and efficient.
✅ 26. Corrosion
Corrosion is slow destruction of metals due to chemical reactions.
Example: Rusting of Iron
4Fe + 3O_2 + 6H_2O \rightarrow 4Fe(OH)_3
Rust = hydrated iron(III) oxide.
Prevention
- Painting
- Greasing
- Galvanization
- Alloying
- Cathodic protection
✅ Most Important Formulas-Class 12 Electrochemistry Notes PDF(Quick Revision)
📌 Cell potential:
E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}
📌 Gibbs free energy:
\Delta G^\circ = -nFE^\circ_{cell}
📌 Nernst equation:
E = E^\circ - \frac{0.0591}{n}\log Q
📌 Charge:
Q = It
📌 Conductance:
G = \frac{1}{R}
📌 Molar conductivity:
\Lambda_m = \frac{\kappa \times 1000}{C}
📌 Kohlrausch law:
\Lambda_m^\circ = \lambda_+^\circ + \lambda_-^\circ
✅ Important Repeated Questions (Board + NEET)
1 Mark Questions
- Define electrochemistry.
- What is EMF of a cell?
- Write formula of Nernst equation.
- What is SHE?
- Unit of conductivity?
2-3 Marks Questions
- Differentiate between galvanic and electrolytic cell.
- Explain role of salt bridge.
- Define standard electrode potential.
- Write Daniell cell notation and reaction.
- Why molar conductivity increases on dilution?
5 Marks Questions
- Derive Nernst equation for a cell.
- Explain lead storage battery with reactions.
- Explain corrosion and its prevention methods.
- Explain Kohlrausch’s law and its applications.
✅ Conclusion-Class 12 Electrochemistry Notes PDF
Electrochemistry is a high-scoring chapter if you understand the basic concepts like oxidation-reduction, EMF, Nernst equation, conductance and batteries. These notes are designed to help Class 12 students prepare for Boards + competitive exams in a simple and complete way.