MP Board Class 12 Chemistry Chapter 2: Electrochemistry — Complete Notes for 2027 Exam

Chapter 2: Electrochemistry is a high-weightage chapter in the MP Board Class 12 Chemistry syllabus, typically contributing 12–16 marks in the board exam. This chapter bridges physical chemistry with real-world applications — from how batteries power your phone to how metals are purified. These comprehensive notes cover electrochemical cells, Nernst equation, conductance, electrolysis, and batteries with clear derivations and MP Board–focused exam tips.

⚡ Electrochemical Cells

An electrochemical cell is a device that converts chemical energy into electrical energy (or vice versa). There are two types:

Galvanic Cell (Voltaic Cell): Spontaneous chemical reaction produces electricity. Example: Daniel cell.
Electrolytic Cell: Electrical energy drives a non-spontaneous chemical reaction. Example: Electrolysis of water.

🧪 Daniel Cell — The Classic Example

The Daniel cell consists of a zinc electrode dipped in ZnSO₄ solution (anode) and a copper electrode dipped in CuSO₄ solution (cathode), connected by a salt bridge (KCl or KNO₃).

Component	Details
Anode (Oxidation)	Zn(s) → Zn²⁺(aq) + 2e⁻
Cathode (Reduction)	Cu²⁺(aq) + 2e⁻ → Cu(s)
Net Cell Reaction	Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Cell Notation	Zn(s) \| Zn²⁺(aq) \|\| Cu²⁺(aq) \| Cu(s)

🔑 Key Terms

Salt bridge: Maintains electrical neutrality and completes the circuit. Contains inert electrolyte like KCl, KNO₃, or NH₄NO₃.
Electrode potential: The potential difference between the electrode and its electrolyte. Measured in volts (V).
Standard electrode potential (E⁰): Measured under standard conditions — 1 M concentration, 1 atm pressure, 298 K.
Standard Hydrogen Electrode (SHE): Reference electrode with E⁰ = 0.00 V. Used to measure all other electrode potentials.

🎯 Exam Tip: In MP Board exams, you are often asked to draw a labelled diagram of the Daniel cell and write the cell notation. Practise the cell representation format — it carries 2–3 marks in long-answer questions.

📐 Nernst Equation & EMF of Cell

The Nernst equation relates the electrode potential to the concentration of ions (or pressure of gases) at a given temperature. For a general cell reaction:

E_cell = E⁰_cell − (RT / nF) × ln Q

At 298 K, using log₁₀:

E_cell = E⁰_cell − (0.0591 / n) × log Q

Symbol	Meaning	Unit
E_cell	Cell potential under non-standard conditions	Volts (V)
E⁰_cell	Standard cell potential	Volts (V)
R	Gas constant (8.314 J mol⁻¹ K⁻¹)	J mol⁻¹ K⁻¹
T	Temperature	Kelvin (K)
n	Number of electrons transferred	—
F	Faraday constant (96485 C mol⁻¹)	C mol⁻¹
Q	Reaction quotient	—

🔋 EMF of a Cell

The standard EMF (E⁰_cell) of a cell is the difference between the standard reduction potentials of the cathode and anode:

E⁰_cell = E⁰_cathode − E⁰_anode

📘 Gibbs Free Energy Relation

ΔG = −nFE_cell and ΔG⁰ = −nFE⁰_cell
When ΔG = 0, the cell reaches equilibrium: E⁰_cell = (0.0591/n) log K_eq

🎯 Exam Tip: MP Board 2027 — A numerical question on EMF calculation using the Nernst equation is almost certain (3–4 marks). Memorise the formula E = E⁰ − (0.0591/n) log Q and practise with different concentrations.

📊 Conductance of Electrolytic Solutions

Conductance (G) is the measure of how easily electricity flows through an electrolytic solution. It is the reciprocal of resistance (R). The key parameters are:

Term	Symbol	Unit	Formula
Resistance	R	Ω (ohm)	R = ρ × ℓ / A
Conductance	G	S (siemens)	G = 1/R
Resistivity	ρ	Ω m	ρ = R × A / ℓ
Conductivity	κ (kappa)	S m⁻¹	κ = 1/ρ
Molar Conductivity	Λ_m	S m² mol⁻¹	Λ_m = κ / C

📏 Cell Constant

Cell constant (G*) = ℓ / A, where ℓ is the distance between electrodes and A is the cross-sectional area. Experimentally: G* = κ × R. It has units of m⁻¹.

📘 Key Trend: As concentration decreases, molar conductivity (Λ_m) increases because interionic attractions weaken, allowing ions to move more freely. For strong electrolytes, Λ_m approaches Λ⁰_m (limiting molar conductivity) at infinite dilution.

📜 Kohlrausch’s Law

Kohlrausch’s Law of Independent Migration of Ions: At infinite dilution, the molar conductivity of an electrolyte is equal to the sum of the contributions of its individual ions.

Λ⁰_m = ν₊λ⁰₊ + ν₋λ⁰₋

Where ν₊ and ν₋ are the number of cations and anions per formula, and λ⁰₊, λ⁰₋ are their limiting molar conductivities.

✅ Applications of Kohlrausch’s Law

Calculate Λ⁰_m for weak electrolytes (e.g., CH₃COOH)
Determine degree of dissociation: α = Λ_m / Λ⁰_m
Calculate dissociation constant K_a = Cα² / (1 − α)
Calculate solubility of sparingly soluble salts: s = κ × 1000 / Λ⁰_m

🎯 Exam Tip: Kohlrausch’s Law is a favourite in MP Board 2027 — expect a 3-mark question on its application to calculate Λ⁰_m of a weak acid or the degree of dissociation.

💡 Electrolysis & Faraday’s Laws

Electrolysis is the process of decomposing an electrolyte by passing an electric current through its solution (or molten state). Products are obtained at the electrodes through oxidation and reduction.

🧪 Faraday’s First Law

The mass of a substance deposited (or liberated) at an electrode is directly proportional to the quantity of electricity passed through the electrolyte.

m = Z × Q = Z × I × t

Where: m = mass (g), Z = electrochemical equivalent (g C⁻¹), Q = charge (Coulomb), I = current (A), t = time (s).

🧪 Faraday’s Second Law

When the same quantity of electricity is passed through different electrolytes, the masses of substances deposited are proportional to their chemical equivalents.

m₁ / m₂ = E₁ / E₂

Where E is the equivalent weight of the substance. Also: w = (E × I × t) / 96500

📘 1 Faraday (F) = 96485 C = charge of 1 mole of electrons. This can deposit 1 gram-equivalent of any substance.

⚙️ Products of Electrolysis

The products depend on the nature of the electrolyte and the electrode material:

Electrolyte	At Cathode	At Anode
Molten NaCl	Na(s)	Cl₂(g)
Aqueous NaCl	H₂(g)	Cl₂(g)
Aqueous Na₂SO₄	H₂(g)	O₂(g)
Aqueous CuSO₄ (Pt electrodes)	Cu(s)	O₂(g)
Aqueous CuSO₄ (Cu electrodes)	Cu(s)	Cu²⁺(aq) (anode dissolves)

🔋 Batteries & Fuel Cells

🔋 Primary Batteries

Non-rechargeable. Example: Dry cell (Leclanché cell) — Zn anode, graphite cathode, MnO₂ + NH₄Cl paste as electrolyte. Produces 1.5 V.

🔋 Secondary Batteries

Rechargeable. Example: Lead-acid battery — Pb anode, PbO₂ cathode, H₂SO₄ electrolyte. Produces 2 V per cell (6 cells = 12 V car battery). Nickel-Cadmium (Ni-Cd) battery — NiO(OH) cathode, Cd anode, KOH electrolyte.

⛽ Fuel Cells

Fuel cells convert chemical energy of a fuel (like H₂) directly into electrical energy. The Hydrogen-Oxygen fuel cell uses H₂ at the anode and O₂ at the cathode with KOH as electrolyte. Produces electricity + water as the only by-product — clean and efficient (60–70% efficiency).

Feature	Primary Battery	Secondary Battery	Fuel Cell
Rechargeable?	No	Yes	Continuous fuel supply
Example	Dry cell, Mercury cell	Lead-acid, Ni-Cd, Li-ion	H₂-O₂ fuel cell
Lifespan	Single use	500–1000 cycles	As long as fuel is supplied

🛡️ Corrosion

Corrosion is the process where metals are oxidized by atmospheric oxygen, moisture, or other chemicals. The most common example is rusting of iron:

Anode: Fe → Fe²⁺ + 2e⁻
Cathode: O₂ + 4H⁺ + 4e⁻ → 2H₂O (in acidic medium)
Overall: 2Fe + O₂ + 2H₂O → 2Fe(OH)₂ → Fe₂O₃·xH₂O (rust)

🛡️ Methods to Prevent Corrosion

Galvanization: Coating iron with zinc. Zn acts as a sacrificial anode.
Tinning: Coating with tin (used in food cans).
Electroplating: Depositing a protective metal layer (Cr, Ni) on the surface.
Sacrificial anode method: Attaching a more reactive metal (Mg, Zn) to protect the main structure.
Painting or oiling: Creating a barrier against air and moisture.
Alloying: Mixing with other metals (e.g., stainless steel = Fe + Cr + Ni).

📝 Important Questions for MP Board 2027

✏️ Very Short Answer (1 Mark)

What is the function of a salt bridge in a galvanic cell?
Define molar conductivity.
What is the unit of cell constant?
Give one example of a primary battery.
State Faraday’s First Law of Electrolysis.

📝 Short Answer (2–3 Marks)

Write the cell reaction and calculate E⁰_cell for the Daniel cell. Given: E⁰_Cu²⁺/Cu = +0.34 V, E⁰_Zn²⁺/Zn = −0.76 V.
Explain why molar conductivity increases with dilution.
State and explain Kohlrausch’s Law with one application.
Distinguish between a primary and secondary battery.
Calculate the mass of copper deposited when 0.5 A current is passed through CuSO₄ solution for 30 minutes. (Atomic mass of Cu = 63.5 g, F = 96500 C)

📚 Long Answer (5–6 Marks)

Explain the working of a Daniel cell with a labelled diagram. Derive the expression for its EMF using the Nernst equation.
What is electrolysis? Explain Faraday’s Laws of Electrolysis with equations and one numerical example.
Describe the construction and working of a hydrogen-oxygen fuel cell. What are its advantages?
Explain the mechanism of corrosion of iron. Discuss any four methods to prevent corrosion.

📘 Quick Revision Tips

Memorise the standard reduction potential series — it helps predict cell reactions
Nernst equation derivation appears every year — practice with different concentrations
For electrolysis, remember: cathode attracts cations (reduction), anode attracts anions (oxidation)
Practice at least 3 numerical problems: EMF calculation, Faraday’s laws, Kohlrausch’s Law

📖 More Resources: Download MP Board Class 12 Chemistry previous year papers and practice sets from mpboard.ai. For video explanations and chapter-wise PYQ discussions, visit our Class 12 Chemistry Course.