Electrochemistry
'Electricity can drive non-spontaneous reactions — and spontaneous reactions can produce electricity. Electrochemistry is the bridge between CHEMICAL and ELECTRICAL energy.'
1. Chapter Overview
Electrochemistry explores the relationship between CHEMICAL REACTIONS and ELECTRICITY. Topics include: ELECTROCHEMICAL CELLS (voltaic and electrolytic — devices that convert chemical energy ⇌ electrical energy), the NERNST EQUATION (relating cell potential to concentration), the STANDARD ELECTRODE POTENTIAL (E°), the ELECTROCHEMICAL SERIES, CONDUCTANCE (molar and equivalent conductivity, Kohlrausch's law), BATTERIES (primary and secondary), FUEL CELLS, and CORROSION.
2. Electrochemical Cells
Galvanic (Voltaic) Cell
- A device that converts CHEMICAL energy into ELECTRICAL energy through spontaneous redox reactions.
- Daniel cell: Zn (anode, oxidation) | ZnSO₄ || CuSO₄ | Cu (cathode, reduction).
- Anode (−): Zn → Zn²⁺ + 2e⁻ (oxidation). Cathode (+): Cu²⁺ + 2e⁻ → Cu (reduction).
- Salt bridge: Completes the circuit, maintains electrical neutrality, prevents mixing of solutions.
Cell Representation
- Anode on LEFT, Cathode on RIGHT. Single line = phase boundary. Double line = salt bridge.
- Zn(s) | Zn²⁺(aq, 1M) || Cu²⁺(aq, 1M) | Cu(s).
3. Standard Electrode Potential
- Standard Hydrogen Electrode (SHE) : Reference electrode with E° = 0 V. H⁺(1M) | H₂(1 atm) | Pt.
- Standard cell potential: E°_cell = E°_cathode − E°_anode.
- Positive E°_cell: Spontaneous reaction (ΔG° = −nFE°_cell < 0). Negative: Non-spontaneous.
Electrochemical Series
- Metals with MORE NEGATIVE E° are STRONGER REDUCING AGENTS (more reactive).
- 'Li has the most negative E° (−3.04 V) — it is the STRONGEST reducing agent. F₂ has the most positive E° (+2.87 V) — it is the STRONGEST oxidising agent.'
| High Reactivity (Negative E°) | Med | Low Reactivity (Positive E°) |
|---|---|---|
| Li, K, Ca, Na, Mg, Al, Zn, Fe | Ni, Sn, Pb | Cu, Ag, Hg, Pt, Au |
4. Nernst Equation
For a General Reaction: aA + bB → cC + dD
- E_cell = E°_cell − (0.0591/n) log Q (at 25°C, in volts).
- E_cell = E°_cell − (RT/nF) ln Q (general form).
For a Single Electrode
- Mⁿ⁺ + ne⁻ → M: E = E° − (0.0591/n) log(1/[Mⁿ⁺]) = E° + (0.0591/n) log[Mⁿ⁺].
Worked Example 1
Problem: Calculate E_cell for the Daniel cell when [Cu²⁺] = 0.01 M and [Zn²⁺] = 1 M. E°_cell = 1.1 V. Solution: Zn + Cu²⁺ → Zn²⁺ + Cu, n = 2. E_cell = 1.1 − (0.0591/2) log(1/0.01) = 1.1 − (0.02955) log(100) = 1.1 − 0.02955×2 = 1.1 − 0.0591 = 1.041 V.
5. Conductance
| Property | Symbol | Formula | Unit |
|---|---|---|---|
| Resistance | R | R = ρ·l/A | Ω |
| Conductivity | κ | κ = 1/ρ | S/m (or Ω⁻¹m⁻¹) |
| Cell constant | G* | G* = l/A | m⁻¹ |
| Molar conductivity | Λ_m | Λ_m = κ/c | S·m²/mol |
Variation of Molar Conductivity
- Strong electrolytes: Λ_m DECREASES SLOWLY as concentration increases (due to ion-ion interactions). Λ° = limiting molar conductivity (at infinite dilution).
- Weak electrolytes: Λ_m INCREASES SHARPLY as concentration decreases (due to increased dissociation). Λ° is found using Kohlrausch's law.
Kohlrausch's Law
- Λ° = λ°(+) + λ°(−). 'The molar conductivity at infinite dilution is the SUM of the INDEPENDENT contributions of the cation and anion.'
- Application: Finding Λ° for weak electrolytes. Example: Λ°(CH₃COOH) = λ°(H⁺) + λ°(CH₃COO⁻).
6. Batteries
| Type | Characteristics | Examples |
|---|---|---|
| Primary | NON-RECHARGEABLE — irreversible reaction | Dry cell (Leclanche), Mercury cell |
| Secondary | RECHARGEABLE — reversible reaction | Lead-acid battery (car battery), Ni-Cd, Li-ion |
| Fuel cell | Continuous supply of reactants | H₂-O₂ fuel cell (produces electricity + water) |
Lead-Acid Battery
- Anode: Pb + SO₄²⁻ → PbSO₄ + 2e⁻. Cathode: PbO₂ + 4H⁺ + SO₄²⁻ + 2e⁻ → PbSO₄ + 2H₂O.
- Overall: Pb + PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O.
- Voltage: ~2 V per cell. Six cells connected in series = 12 V car battery.
H₂-O₂ Fuel Cell
- Anode: 2H₂ + 4OH⁻ → 4H₂O + 4e⁻. Cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻.
- Overall: 2H₂ + O₂ → 2H₂O + Electrical energy. 'The only product is WATER — clean and efficient.'
7. Corrosion
- 'The GRADUAL DESTRUCTION of metals by chemical or electrochemical reaction with the environment.'
- Rusting of iron: Fe → Fe²⁺ (oxidation at ANODIC region). O₂ + 2H₂O + 4e⁻ → 4OH⁻ (reduction at CATHODIC region).
- Prevention: Painting, galvanisation (coating with Zn), tinning, cathodic protection (sacrificial anode — Mg or Zn connected to iron).
8. Comparison Table: Electrolytic Cell vs Galvanic Cell
| Feature | Electrolytic Cell | Galvanic Cell |
|---|---|---|
| Energy conversion | Electrical → Chemical | Chemical → Electrical |
| Reaction | NON-spontaneous (driven by electricity) | SPONTANEOUS |
| Electrode charges | Anode (+) | Cathode (−) |
| Electron flow | From external supply | From anode to cathode (external circuit) |
| Applications | Electroplating, electrolysis | Batteries |
9. Common Mistakes
- Anode vs Cathode in different cells: In a GALVANIC cell, anode is NEGATIVE. In an ELECTROLYTIC cell, anode is POSITIVE. 'In galvanic, oxidation happens at the negative electrode. In electrolytic, oxidation happens at the positive electrode.'
- Nernst equation sign: E = E° − (0.0591/n) log Q. It is MINUS, not plus. A large Q (products > reactants) REDUCES cell potential.
- Units of conductivity: κ is in S/m (or Ω⁻¹m⁻¹). Λ_m is in S·m²/mol. They differ by concentration: Λ_m = κ/c.
- Fuel cells are NOT batteries: Batteries store energy chemically and have a limited capacity. Fuel cells CONVERT fuel continuously as long as reactants are supplied.
10. CBSE Exam Focus
- Electrochemical cells — Daniel cell, cell notation, salt bridge
- Standard electrode potential — EMF of cell, electrochemical series
- Nernst equation — numerical problems (E_cell = E°_cell − 0.0591/n log Q)
- Conductance — molar conductivity, Kohlrausch's law
- Batteries — primary vs secondary, lead-acid, fuel cells
- Corrosion — mechanism, prevention methods
11. Self-Test
Q1: Calculate E°_cell for Zn|Zn²⁺||Ag⁺|Ag. Given: E°(Zn²⁺/Zn) = −0.76 V, E°(Ag⁺/Ag) = +0.80 V. A1: E°_cell = E°_cathode − E°_anode = 0.80 − (−0.76) = 1.56 V.
Q2: Calculate Λ° for CH₃COOH given λ°(H⁺) = 349.8 and λ°(CH₃COO⁻) = 40.9 S·cm²/mol. A2: Λ°(CH₃COOH) = λ°(H⁺) + λ°(CH₃COO⁻) = 349.8 + 40.9 = 390.7 S·cm²/mol.
Q3: For the reaction 2Al + 3Cu²⁺ → 2Al³⁺ + 3Cu, E° = 2.0 V. Calculate ΔG°. A3: n = 6 (Al → Al³⁺ loses 3e⁻, 2Al atoms = 6e⁻). ΔG° = −nFE° = −6×96500×2 = −1,158,000 J = −1158 kJ.
Q4: A copper electrode is placed in 0.001 M CuSO₄ solution at 25°C. Calculate its reduction potential. (E° = +0.34 V) A4: Cu²⁺ + 2e⁻ → Cu. E = E° + (0.0591/2) log[Cu²⁺] = 0.34 + 0.02955×log(0.001) = 0.34 + 0.02955×(−3) = 0.34 − 0.0887 = 0.251 V.
Q5: A solution of 0.05 M KCl has conductivity 0.007 S/m at 25°C. Find the molar conductivity. A5: Λ_m = κ/c = 0.007/(0.05×1000) = 0.007/50 = 1.4×10⁻⁴ S·m²/mol (Note: c is converted to mol/m³: 0.05 M = 50 mol/m³).
12. Conclusion
Electrochemistry is the SCIENCE of ENERGY CONVERSION:
- CELLS: 'Galvanic cells produce electricity from spontaneous reactions. Electrolytic cells use electricity to drive non-spontaneous reactions.'
- NERNST: 'Cell potential depends on concentration — dilute the reactants and the voltage changes.'
- CONDUCTANCE: 'How well a solution conducts electricity depends on the number and mobility of ions.'
- BATTERIES and CORROSION: 'The practical faces of electrochemistry — one we USE, the other we FIGHT.'
'Electrochemistry powers our devices, protects our structures, and drives essential industrial processes — it is chemistry in ACTION.'
