Electrochemistry — chemistry Class 12 Notes (CBSE & HBSE)
Free NCERT chemistry notes for Electrochemistry (Class 12) on Siksha Sarovar, aligned to CBSE and Haryana Board (HBSE). This chapter is broken into 3 topics with clear explanations, formulas, solved examples and board-pattern practice — free to read, no sign-up required.
Board exam focus — Electrochemistry (CBSE & HBSE)
Electrochemical cells, Galvanic and electrolytic cells, standard electrode potential, Nernst equation, conductance, Kohlrausch's law, Faraday's laws, and batteries.
Electrochemical Cells and EMF
Electrochemistry
Electrochemical Cells
A device that converts chemical energy to electrical energy (galvanic/voltaic) or electrical energy to chemical energy (electrolytic).
Galvanic Cell (Voltaic Cell)
Spontaneous redox reaction produces electrical energy.
Daniel Cell (Zn-Cu cell): Zn electrode in ZnSO4 || CuSO4 Cu electrode Anode (oxidation): Zn → Zn2+ + 2e- (negative electrode) Cathode (reduction): Cu2+ + 2e- → Cu (positive electrode) Overall: Zn + Cu2+ → Zn2+ + Cu
Cell notation: Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s) (| = phase boundary, || = salt bridge)
Salt bridge: Provides electrical connectivity between two half-cells; prevents charge buildup.
Standard Electrode Potential (E0)
Measured vs Standard Hydrogen Electrode (SHE, E0 = 0.00 V by convention). SHE: Pt|H2(g, 1 atm)|H+(aq, 1M) at 25 C.
| Half-reaction | E0 (V) |
|---|---|
| F2 + 2e- → 2F- | +2.87 |
| MnO4- + 8H+ + 5e- → Mn2+ | +1.51 |
| Cu2+ + 2e- → Cu | +0.34 |
| 2H+ + 2e- → H2 | 0.00 |
| Zn2+ + 2e- → Zn | -0.76 |
| Na+ + e- → Na | -2.71 |
Higher (more positive) E0 = greater tendency to be reduced = better oxidizing agent. Lower (more negative) E0 = greater tendency to be oxidized = better reducing agent.
Cell EMF
E0cell = E0cathode - E0anode E0cell must be positive for spontaneous reaction.
Daniel cell: E0cell = +0.34 - (-0.76) = +1.10 V
Nernst Equation
Relates cell EMF to concentration: Ecell = E0cell - (RT/nF) ln Q At 25 C: Ecell = E0cell - (0.0592/n) log Q
where n = electrons transferred, Q = reaction quotient.
At equilibrium: Ecell = 0 and Q = Keq 0 = E0cell - (0.0592/n) log Keq log Keq = n x E0cell / 0.0592
Gibbs Energy and EMF
DeltaG = -nFEcell DeltaG0 = -nFE0cell Spontaneous: DeltaG < 0, Ecell > 0
Conductance and Electrolysis
Conductance
Types of Conductors
Metallic (electronic) conductors: Conduct by electron flow. Resistance increases with temperature. Electrolytic (ionic) conductors: Conduct by ion movement in solution/melt. Conductance increases with temperature.
Resistance and Conductance
Resistance R (ohm, Omega): Opposition to current flow. Conductance G = 1/R (siemens, S)
Specific conductance (conductivity): kappa = l/(R x A) = G x (l/A) = G x cell constant Units: S/cm or S/m
Molar conductance (Lambda_m): Lambda_m = (kappa x 1000) / M (in cm^2/mol) [if kappa in S/cm, M in mol/L] Or Lambda_m = kappa / C [if C in mol/m^3]
Equivalent conductance: Lambda_eq = kappa x 1000 / C_eq
Variation of Conductance with Concentration
Strong electrolytes: Lambda_m increases with dilution (more ions available, less inter-ionic forces). Debye-Huckel-Onsager equation: Lambda_m = Lambda_m^0 - (A + B Lambda_m^0) x sqrt(C)
Weak electrolytes: Lambda_m increases sharply with dilution (more dissociation).
Kohlrausch's Law
At infinite dilution (Lambda_m^0): Conductance of each ion is independent of other ions.
Lambda_m^0 = nu+ x lambda^0+ + nu- x lambda^0-
where nu+, nu- = stoichiometric numbers; lambda^0 = ionic molar conductance at infinite dilution.
Application: Calculate Lambda_m^0 for weak electrolytes (cannot be measured directly): Lambda_m^0(CH3COOH) = Lambda_m^0(CH3COONa) + Lambda_m^0(HCl) - Lambda_m^0(NaCl)
Degree of dissociation of weak electrolytes: alpha = Lambda_m / Lambda_m^0
Electrolysis
Application of external electrical energy to drive non-spontaneous redox reactions.
Electrolytic cell: Anode (+), Cathode (-) [opposite of galvanic cell for external connection]
At cathode: Reduction (metal deposition, H2 evolution) At anode: Oxidation (metal dissolution, O2 evolution, Cl2 evolution)
Preferential discharge: At cathode: Ion with higher reduction potential discharged first. At anode: Ion with lower oxidation potential discharged first.
Products of Electrolysis
Dilute H2SO4: Cathode: H2; Anode: O2 Dilute NaCl: Cathode: H2; Anode: O2 (H+ easier to reduce than Na+) Conc. NaCl (brine): Cathode: H2; Anode: Cl2 (Cl- outnumbers OH-, discharged first) Molten NaCl: Cathode: Na; Anode: Cl2
Faraday's Laws and Batteries
Faraday's Laws of Electrolysis
Faraday's First Law
The amount of substance deposited/dissolved at an electrode is proportional to the quantity of charge passed. m = Z x Q = Z x I x t where m = mass (g), Z = electrochemical equivalent, Q = charge (coulombs), I = current (A), t = time (s).
Faraday's Second Law
When the same quantity of electricity passes through different electrolytes, the masses of substances deposited are proportional to their equivalent masses. m1/m2 = E1/E2 (equivalent masses)
1 Faraday = 96,485 C ≈ 96,500 C = charge of 1 mole of electrons.
Moles of substance = Q / (n x F) = I x t / (n x F) where n = number of electrons per ion.
Example: Depositing Cu from CuSO4 (n=2): Moles of Cu = Q / (2 x 96500) Mass of Cu = moles x 63.5
Batteries (Commercial Cells)
1. Primary Cells (non-rechargeable):
Dry cell (Leclanche cell): Anode: Zn (outer case) Cathode: MnO2 + NH4Cl + C (rod) Electrolyte: Moist NH4Cl + ZnCl2 paste EMF: ~1.5 V Uses: Remotes, clocks, torches.
2. Secondary Cells (rechargeable):
Lead Storage Battery: Anode: Pb; Cathode: PbO2; Electrolyte: H2SO4 (38%) Discharge:
- Anode: Pb + SO4^2- → PbSO4 + 2e-
- Cathode: PbO2 + SO4^2- + 4H+ + 2e- → PbSO4 + 2H2O
EMF per cell: 2 V; 6 cells = 12 V battery (car battery)
Charging: Reverse reactions; PbSO4 → Pb and PbO2. H2SO4 concentration decreases on discharge (sp. gravity indicator).
Nickel-Cadmium (Ni-Cd): Longer life than lead acid; used in cameras, phones.
Lithium-ion battery: Lightest, highest energy density. Anode: graphite (Li intercalation); cathode: LiCoO2 or LiFePO4. Uses: Laptops, phones, electric vehicles.
3. Fuel Cells: Electrochemical cells that convert fuel (H2, CH4) directly to electricity. H2-O2 fuel cell: Anode: H2 + 2OH- → 2H2O + 2e- Cathode: O2 + 2H2O + 4e- → 4OH- Overall: 2H2 + O2 → 2H2O Efficiency: up to 70% (much higher than heat engines). Used in spacecraft, experimental vehicles.
Corrosion
Electrochemical oxidation of metals (iron rusting): Anode: Fe → Fe2+ + 2e- Cathode: O2 + 2H2O + 4e- → 4OH- Overall: 4Fe + 3O2 + xH2O → 2Fe2O3.xH2O (rust)
Prevention: Galvanizing (zinc coating), painting, alloying, cathodic protection (sacrificial anode).
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Concept explanations, key formulas and definitions, fully solved examples and board-pattern practice questions for Electrochemistry.