Electrochemistry

Electrochemical Processes

  • Oxidation-reduction reactions convert energy to electricity or vice versa.

Oxidation Numbers

  • Definition: Charge an atom would have if electrons were fully transferred.

  • Free elements = oxidation number of 0.

  • Monatomic ions: oxidation number equals ion charge.

  • Oxygen usually has an oxidation number of -2; exceptions in peroxides.

  • Hydrogen typically +1; -1 when bonded to metals in binary compounds.

  • Group IA = +1, IIA = +2; fluorine always -1.

Balancing Redox Equations

  • Write unbalanced ionic equations and split into half-reactions.

  • Balance non-O and non-H atoms first.

  • Balance O with H2O, and H with H+.

  • Add electrons to balance charges.

  • Equalize electron count in both half-reactions if needed.

  • Combine the half-reactions ensuring atoms and charges are balanced.

  • Adjust for basic solutions with OH-.

Standard Reduction Potentials (E°)

  • E°: voltage for reduction at 1 M solute and 1 atm gas.

  • Higher E° indicates greater tendency to be reduced.

Electrolysis

  • Uses electrical energy to drive nonspontaneous reactions.

  • Example: Electrolysis of Na2SO4 produces oxygen at the anode and hydrogen at the cathode.

Concentration Cells

  • Galvanic cells with the same materials but different ion concentrations.

Nernst Equation

  • Relates cell potential to concentrations: E=E∘−0.0257Vnln⁡QE = E^\circ - \frac{0.0257 V}{n} \ln Q at 298 K.

  • Used to calculate spontaneity and emf in non-standard conditions.

Examples and Calculations

  • Predict reactions based on standard reduction potentials and calculate cell emf.

  • Use the relationship between ΔG°, K, and Ecell to determine spontaneity and equilibrium constants.

  • Quantify gas production in electrolysis using current, time, and Faraday's laws.