Overview
This lecture provides an overview of electrochemistry, focusing on voltaic cells, balancing redox reactions, calculating cell potentials, Gibbs free energy, equilibrium constants, and performing stoichiometry and electrolysis calculations.
Voltaic (Galvanic) and Electrolytic Cells
- Voltaic (galvanic) cells generate spontaneous reactions with positive cell potentials.
- Electrolytic cells require external energy, allowing non-spontaneous reactions (cell potential can be negative).
- Electrons flow from anode (oxidation) to cathode (reduction).
- Anode loses mass (oxidation), cathode gains mass (reduction).
- The salt bridge maintains charge balance: cations move toward cathode, anions toward anode.
Cell Potentials and Redox Reactions
- Standard reduction potential (E°) is measured under 1 M ion concentrations.
- To get the net cell reaction, reverse the oxidation half-reaction and sum potentials.
- Cell potential (E°cell) = E°cathode − E°anode; net potential must be positive for voltaic cells.
- Metals are typically reducing agents, metal ions and nonmetals act as oxidizing agents.
- Reducing agent is oxidized; oxidizing agent is reduced.
Gibbs Free Energy, Equilibrium Constant, and the Nernst Equation
- ΔG° = −nFE°cell, where n = electrons transferred, F = 96,485 C/mol e⁻, E = cell potential.
- Positive E°cell → negative ΔG° (spontaneous); K (equilibrium constant) > 1 (product-favored).
- ΔG° = −RT ln K; solve for K: K = exp(−ΔG°/RT).
- Nernst equation for non-standard conditions:
Ecell = E°cell − (0.0591/n) × log(Q), where Q = [products]/[reactants] (excluding solids).
- If reactant concentration increases or product decreases, Ecell increases.
Electrolysis and Stoichiometry Calculations
- Q (charge) = current (I, in amps) × time (t, in seconds); 1 coulomb = 1 amp × 1 second.
- Faraday's constant links moles of e⁻ and total charge: 96,485 C/mol e ⁻.
- Use stoichiometry and balanced half-reactions to relate deposited or consumed material to charge transferred and current.
- Example: mass deposited = (It × molar mass)/(n × F).
Identifying Strongest Oxidizing/Reducing Agents
- Metals are strongest reducing agents; metal cations and nonmetals are strong oxidizing agents.
- The strongest reducing/oxidizing agent corresponds to the species with most positive/negative (reversed) E°.
Cell Notation and Half-Cell Reactions
- Cell notation: anode (oxidation) | anode ion || cathode ion | cathode (reduction).
- Use inert electrodes (Pt, C) if no solid is present at the electrode.
Balancing Redox Reactions (Acidic/Basic)
- Acidic: balance elements, then O with H2O, H with H⁺, balance charge with e⁻.
- Basic: balance as acid, then add OH⁻ to both sides to neutralize H⁺, simplify water.
Key Terms & Definitions
- Anode — Electrode where oxidation (electron loss) occurs.
- Cathode — Electrode where reduction (electron gain) occurs.
- Standard Reduction Potential (E°) — Voltage for reduction at 1 M concentration.
- Nernst Equation — Calculates Ecell under non-standard conditions.
- Faraday’s Constant (F) — 96,485 C/mol e⁻, charge per mole of electrons.
- Q (Reaction Quotient) — Ratio of product to reactant concentrations (excluding solids).
- Reducing Agent — Substance oxidized; donates electrons.
- Oxidizing Agent — Substance reduced; accepts electrons.
Action Items / Next Steps
- Practice balancing redox reactions under acidic and basic conditions.
- Solve practice problems on cell potential calculations using the Nernst equation.
- Review examples of identifying oxidizing and reducing agents.
- Complete homework on electrolysis and stoichiometry involving current, time, and mass calculations.