Electrochemistry Class 12 Chemistry Notes

An electrochemical cell converts chemical energy into electrical energy (galvanic cell, spontaneous reaction) or uses electrical energy to drive a non-spontaneous reaction (electrolytic cell). In every half-cell, oxidation occurs at the anode and reduction at the cathode.

• Reactivity series: a metal with a more negative E° is a stronger reducing agent and is oxidised more easily.
• More positive E° ⇒ the ion is a stronger oxidising agent and is reduced more readily.
• E°cell = E°cathode − E°anode; E°cell > 0 means the cell reaction is spontaneous (ΔG° < 0).

MCQ Questions

The Daniel cell is the classic galvanic cell: zinc in ZnSO₄ and copper in CuSO₄ joined by a salt bridge (KCl or KNO₃ in a gel). The zinc electrode dissolves (anode) and copper plates out (cathode).

• Anode: Zn(s) → Zn²⁺(aq) + 2e⁻ (oxidation, electrode loses mass).
• Cathode: Cu²⁺(aq) + 2e⁻ → Cu(s) (reduction, electrode gains mass).
• Overall: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s), E°cell = +1.10 V.

MCQ Questions

The Nernst equation relates electrode (or cell) potential to ion concentrations. It shows why a cell's EMF falls as the product concentrations build up, until Q = K and E = 0.

For a Daniel cell Zn(s) + Cu²⁺(aq) ⇌ Zn²⁺(aq) + Cu(s): E = 1.10 − (0.0591/2) log₁₀([Zn²⁺]/[Cu²⁺]).

• ΔG° = −nFE°cell links EMF to Gibbs energy; ΔG = ΔG° + RT ln Q.
• At equilibrium, E = 0 and Q = K: E°cell = (0.0591/n) log₁₀ K.
• Concentration cells generate a small EMF from a concentration difference across two half-cells.

MCQ Questions

Electrolysis forces a non-spontaneous redox reaction by passing direct current through an electrolyte. The molten or aqueous electrolyte dissociates into ions, which migrate to the electrodes and discharge in order of discharge potential.

• F = 96500 C per mole of electrons = 1 faraday (1 F deposits 1 gram-equivalent).
• Products of aqueous NaCl: H₂ at cathode, Cl₂ at anode (water is discharged before Na⁺).
• Products of molten NaCl: Na at cathode, Cl₂ at anode — electrolysis of the melt is used commercially.

MCQ Questions

Electrolytes conduct electricity by ion migration. The power of an electrolyte solution to conduct is measured by conductivity (κ) and molar conductivity (Λm = κ/c). Conductance rises with dilution as ions spread out, while κ itself falls.

• Strong electrolytes: Λm rises slowly with dilution; Λm vs √c is a straight line extrapolated to Λm°.
• Weak electrolytes: Λm rises steeply at high dilution; Λm° is found by Kohlrausch's law.
• Applications: dissociation constants of weak acids, solubility of sparingly soluble salts.

MCQ Questions

Batteries are galvanic cells designed for storage and portable power. Primary cells are not rechargeable; secondary cells are rechargeable by reversing the current; fuel cells continuously consume a fuel.

• Dry cell (Leclanché): Zn anode, graphite cathode in MnO₂ + NH₄Cl paste; E ≈ 1.5 V. Primary cell.
• Lead storage cell: Pb and PbO₂ in H₂SO₄; 2 V per cell; rechargeable secondary cell used in cars.
• Ni–Cd cell: NiO(OH) cathode, Cd anode in KOH; rechargeable, used in electronics.
• Fuel cell: H₂–O₂ cell with KOH electrolyte produces water and ~1.23 V, used in space missions.

MCQ Questions

Exercises (NCERT Q1–Q4)

1. 1
   Define electrochemical cell and differentiate between galvanic and electrolytic cells.
2. 2
   Write the Nernst equation and calculate the emf of the following cell at 298 K: Fe(s) | Fe²⁺ (0.001 M) || H⁺ (0.01 M) | H₂ (1 bar) | Pt(s). Given: E°Fe²⁺/Fe = -0.44 V.
3. 3
   State Kohlrausch's law of independent migration of ions. How is it used to determine the limiting molar conductivity of a weak electrolyte?
4. 4
   Calculate the quantity of electricity (in coulombs) required to reduce 1 mol of Cr₂O₇²⁻ to Cr³⁺. How much electricity is needed to reduce 0.05 mol of Cr₂O₇²⁻?

Answers at a glance

1. 1
   An electrochemical cell converts chemical energy to electrical energy or vice versa. Galvanic cells produce electricity spontaneously; electrolytic cells require external electricity to drive non-spontaneous reactions.
2. 2
   EMF of the cell = 0.41 V
3. 3
   Kohlrausch's law: Λ°m = Σνiλ°i. Used for weak electrolytes by summing ionic contributions from strong electrolytes.
4. 4
   For 1 mol Cr₂O₇²⁻: 5.79 × 10⁵ C; for 0.05 mol: 2.89 × 10⁴ C

MCQ Questions

1 Mark Questions

1. 1
   Define electrochemical cell.
2. 2
   What is the Nernst equation for a cell reaction?
3. 3
   State Faraday’s first law of electrolysis.

2 Mark Questions

1. 1
   Calculate the standard cell potential of the Daniel cell given E°(Cu2+/Cu) = +0.34 V and E°(Zn2+/Zn) = –0.76 V.
2. 2
   How is the molar conductivity of an electrolyte related to its conductivity?

3 Mark Questions

1. 1
   Derive the relationship between ΔG° and E°cell for an electrochemical cell.
2. 2
   Explain how the limiting molar conductivity of a weak electrolyte can be determined using Kohlrausch’s law of independent migration of ions.

5 Mark Questions

1. 1
   The conductivity of 0.001028 mol L−1 acetic acid is 4.95 × 10−5 S cm−1. Calculate its dissociation constant if Λ°(acetic acid) = 390.5 S cm2 mol−1.
2. 2
   For the reaction 2Ag+(aq) + Cu(s) → 2Ag(s) + Cu2+(aq), the standard cell potential is found to be 0.46 V at 298 K. Calculate the equilibrium constant of the reaction. (R = 8.314 J K−1 mol−1, F = 96487 C mol−1)