Learning Objectives
- Understand electrochemical cells and electrode potentials
- Apply the Nernst equation and predict spontaneity
- Study electrolysis and Faraday's laws
- Learn about conductance, molar conductivity, and Kohlrausch's law
- Understand batteries, fuel cells, and corrosion
Key Concepts
Electrochemical Cells
Galvanic (Voltaic) cell: Converts chemical energy to electrical energy. Spontaneous reaction (ΔG < 0).
Oxidation at anode (negative terminal); Reduction at cathode (positive terminal).
Cell notation: Anode | Anode solution || Cathode solution | Cathode. Example: Zn(s)|Zn²⁺(aq)||Cu²⁺(aq)|Cu(s).
EMF (Cell potential): E°_cell = E°_cathode - E°_anode
If E°_cell > 0: reaction is spontaneous.
Standard Electrode Potential
Measured against Standard Hydrogen Electrode (SHE), E° = 0.00 V.
Higher E° → stronger oxidising agent (greater tendency to be reduced).
Electrochemical series: arranged in increasing order of E°.
Nernst Equation
E_cell = E°_cell - (RT/nF) ln Q = E°_cell - (0.0591/n) log Q (at 25°C)
At equilibrium: E_cell = 0, so E°_cell = (0.0591/n) log K.
Relation to Gibbs energy: ΔG° = -nFE°_cell = -2.303 RT log K.
F = 96485 C/mol (Faraday constant).
Conductance and Conductivity
Conductance (G): G = 1/R = κ × A/l. SI unit: siemens (S).
Conductivity (κ): κ = 1/ρ = G × l/A = G × cell constant. SI unit: S/m.
Molar conductivity: Λ_m = κ/c (S·cm²/mol, where c in mol/cm³) or Λ_m = 1000κ/M (M in mol/L).
On dilution: κ decreases (fewer ions per volume), but Λ_m increases (less interionic interaction).
Kohlrausch's Law
Λ°_m = ν₊λ°₊ + ν₋λ°₋ (limiting molar conductivity is sum of individual ion conductivities).
Applications: Calculate Λ°_m of weak electrolytes, determine degree of dissociation α = Λ_m/Λ°_m, find solubility of sparingly soluble salts.
Electrolysis and Faraday's Laws
Electrolytic cell: Uses electrical energy to drive non-spontaneous reactions. Anode (+), Cathode (-).
Faraday's First Law: m = ZIt = (MIt)/(nF), where Z = electrochemical equivalent, M = molar mass, n = number of electrons.
Faraday's Second Law: When same quantity of electricity passes through different electrolytes, masses deposited are proportional to their equivalent weights.
1 Faraday = 96485 C = charge to deposit 1 mole of monovalent ion.
Batteries and Fuel Cells
Primary (non-rechargeable): Dry cell (Zn-MnO₂, 1.5V), Mercury cell (Zn-HgO, 1.35V).
Secondary (rechargeable): Lead-acid (Pb-PbO₂, 2V per cell, 6 cells = 12V), Ni-Cd battery.
Fuel cell: H₂-O₂ fuel cell: 2H₂ + O₂ → 2H₂O. High efficiency (~70%), pollution-free.
Corrosion
Electrochemical degradation of metals. Iron rusting: Fe → Fe²⁺ + 2e⁻ (anode); O₂ + 2H₂O + 4e⁻ → 4OH⁻ (cathode).
Prevention: galvanising, painting, cathodic protection (sacrificial anode), alloying.
Summary
Electrochemistry connects chemical reactions with electrical energy. Galvanic cells produce electricity from spontaneous reactions. The Nernst equation gives cell potential under non-standard conditions. Conductivity and molar conductivity describe ionic conduction. Faraday's laws quantify electrolysis. Batteries store chemical energy; fuel cells convert it efficiently. Corrosion is an electrochemical process that can be prevented.
Important Terms
- EMF: Maximum potential difference between electrodes
- Nernst Equation: E = E° - (0.0591/n) log Q at 25°C
- Faraday Constant: F = 96485 C/mol
- Molar Conductivity: Conductivity per unit concentration
- Kohlrausch's Law: Λ°_m = sum of individual ionic conductivities
- Corrosion: Electrochemical destruction of metal
Quick Revision
- E°_cell = E°_cathode - E°_anode; ΔG° = -nFE°
- Nernst: E = E° - (0.0591/n) log Q
- At equilibrium: E = 0; E° = (0.0591/n) log K
- Faraday: m = MIt/(nF); 1F = 96485 C
- Kohlrausch: Λ°_m = ν₊λ°₊ + ν₋λ°₋
- Galvanic: anode(-), cathode(+); Electrolytic: anode(+), cathode(-)