Overview
This lecture provides a comprehensive overview of electrochemistry, including voltaic/galvanic and electrolytic cells, cell potentials, Gibbs free energy, equilibrium constants, cell notation, Nernst equation, electrochemical stoichiometry, identification of oxidizing/reducing agents, and balancing redox reactions under acidic and basic conditions.
Voltaic and Electrolytic Cells
- Voltaic (galvanic) cells produce electrical energy via spontaneous reactions; cell potential must be positive.
- Electrolytic cells use external energy to drive nonspontaneous reactions; cell potential can be negative or positive.
- Electrons flow from anode (oxidation, loses mass) to cathode (reduction, gains mass).
- Salt bridges maintain charge balance: cations go to the cathode, anions to the anode.
Calculating Cell Potentials and Gibbs Free Energy
- Standard cell potential (E°cell) is computed by adding half-cell potentials (after reversing the oxidation reaction).
- Under non-standard conditions, cell potential is calculated using the Nernst equation:
Ecell = E°cell – (0.0591/n) × log(Q), where Q = [products]/[reactants].
- ΔG° = –nFE°cell; n = electrons transferred, F = Faraday’s constant (96485 C/mol e⁻).
- ΔG° < 0 and E°cell > 0 indicate a spontaneous reaction.
Equilibrium Constant and Thermodynamic Relationships
- ΔG° = –RT ln K; K = e^(–ΔG°/RT).
- For spontaneous (product-favored) reactions: K >> 1, ΔG° < 0, E°cell > 0.
- Changing concentrations affects Ecell and ΔG (nonstandard), but K is only affected by temperature.
Identifying Oxidizing and Reducing Agents
- The substance oxidized is the reducing agent; the one reduced is the oxidizing agent.
- Metals are typically reducing agents; metal ions and nonmetals are often oxidizing agents.
Electrochemical Stoichiometry Problems
- Q = It (charge = current × time); 1 C = 1 A × 1 s.
- Faraday’s constant relates charge to moles of electrons.
- Use conversion factors from current and time to moles of metal deposited via the stoichiometry of the half-reaction.
Cell Notation
- Format: anode | anode ion (aq) || cathode ion (aq) | cathode.
- Double lines (||) represent the salt bridge; inert electrodes (e.g., Pt, C/graphite) are used when needed.
Balancing Redox Reactions
- Under acidic conditions: add H⁺ and H₂O to balance atoms and charges.
- Under basic conditions: after acidic balancing, add OH⁻ to neutralize H⁺, then combine to form H₂O and simplify.
Anode and Cathode Reaction Choices
- Reduction occurs at the cathode (species with the highest reduction potential); oxidation at the anode.
- Compare standard reduction potentials to determine likely reactions at each electrode.
Key Terms & Definitions
- Anode — Electrode where oxidation occurs (loses electrons).
- Cathode — Electrode where reduction occurs (gains electrons).
- Cell potential (Ecell) — Measured voltage of an electrochemical cell.
- Standard reduction potential (E°) — Voltage for a reduction reaction under standard conditions.
- Gibbs free energy (ΔG°) — Maximum useful work from a reaction; relates to spontaneity.
- Equilibrium constant (K) — Ratio of product to reactant concentrations at equilibrium.
- Nernst equation — Calculates cell potential under non-standard conditions.
- Oxidizing agent — Substance reduced; causes oxidation of another.
- Reducing agent — Substance oxidized; causes reduction of another.
Action Items / Next Steps
- Practice balancing redox reactions under both acidic and basic conditions.
- Solve sample problems calculating cell potentials, ΔG, K, and electrochemical stoichiometry.
- Review and memorize key definitions and standard electrode potentials.