Comprehensive notes on Electrochemistry, Redox Reactions, Oxidation States, Balancing Redox Equations (Ion-Electron Method), Electrolytic and Galvanic Cells, Electrode Potential, Electrochemical Series, and Modern Batteries.
Oxidation State: The apparent charge on an atom in a molecule/ion.
Rules:
1. Free element (e.g., Na, O2) = 0.
2. Group IA = +1, IIA = +2, Al = +3.
3. Hydrogen: +1 (with non-metals), -1 (in metal hydrides like NaH).
4. Oxygen: -2 (usually), -1 (peroxides), +2 (OF2).
5. Sum of O.S. in neutral molecule = 0; in ion = charge on ion.
Steps:
1. Split equation into two half-reactions (Oxidation & Reduction).
2. Balance atoms other than O and H.
3. Balance O by adding H2O.
4. Balance H by adding H+ (Acidic medium).
5. Balance charge by adding electrons (e-).
6. Equalize electrons and add half-reactions.
For Basic Medium: Add OH- to both sides to neutralize H+.
Cells taking electrical energy to drive non-spontaneous chemical reactions.
Electrolysis of Molten NaCl:
Anode (Oxidation): 2Cl- → Cl2 + 2e-.
Cathode (Reduction): 2Na+ + 2e- → 2Na.
Nelson's Cell (Aq. NaCl): Produces NaOH, H2, Cl2. Steel cathode, Graphite anode.
Cells generating electrical energy from spontaneous redox reactions.
Daniel Cell: Zn anode, Cu cathode.
Anode: Zn → Zn2+ + 2e-.
Cathode: Cu2+ + 2e- → Cu.
Cell Notation: Zn | Zn2+ (1M) || Cu2+ (1M) | Cu.
Salt Bridge: Maintains electrical neutrality by allowing ion migration.
Potential of an electrode relative to Standard Hydrogen Electrode (SHE) at 25°C, 1 atm, 1M conc. SHE acts as reference (E0 = 0.00 V).
Electrochemical Series: Arrangement of elements by increasing reduction potential. Elements above H have -ve E0 (strong reducing agents); below H have +ve E0 (strong oxidizing agents).
Lead Accumulator (Car Battery): Secondary cell (rechargeable).
Anode: Pb. Cathode: PbO2. Electrolyte: 30% H2SO4.
Discharge: Pb + PbO2 + 2H2SO4 → 2PbSO4 + 2H2O.
Fuel Cell: Converts chemical energy of fuel (H2, CH4) directly to electricity. H2-O2 Fuel Cell produces water and electricity.