Electrochemistry in AP Chemistry (Unit 9): From Redox Reactions to Batteries and Electrolysis

Electrochemistry

The study of redox (oxidation–reduction) reactions and how electron transfer can produce electrical energy (galvanic cells) or be driven by electricity to cause chemical change (electrolytic cells).

Redox reaction

A reaction involving simultaneous oxidation (loss of electrons) and reduction (gain of electrons).

Oxidation

Loss of electrons by a species (often accompanied by an increase in oxidation number).

Reduction

Gain of electrons by a species (often accompanied by a decrease in oxidation number).

OIL RIG

Mnemonic for redox: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).

Half-cell

One part of an electrochemical cell containing an electrode and an electrolyte where either oxidation or reduction occurs.

Electrode

A conductive solid (metal or inert material like Pt/graphite) where electron transfer occurs in an electrochemical cell.

Salt bridge

A connection (or porous barrier) that allows ions to migrate between half-cells to maintain electrical neutrality and sustain current flow.

Anode

The electrode where oxidation occurs (true for both galvanic and electrolytic cells).

Cathode

The electrode where reduction occurs (true for both galvanic and electrolytic cells).

Galvanic (voltaic) cell

An electrochemical cell that produces electrical energy from a spontaneous redox reaction.

Electrolytic cell

An electrochemical cell that uses an external power source to drive a nonspontaneous redox reaction.

Electron flow (in a cell)

Electrons travel through the external wire from anode to cathode.

Galvanic cell sign convention

In a galvanic cell, the anode is negative (produces electrons) and the cathode is positive (consumes electrons).

Electrolytic cell sign convention

In an electrolytic cell, the anode is positive (connected to + terminal of power supply) and the cathode is negative (connected to − terminal).

Cell notation (cell diagram)

Shorthand for a galvanic cell: single line | = phase boundary; double line || = salt bridge; anode written on the left, cathode on the right (e.g., Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s)).

Cell potential (E) / emf

The voltage that measures the driving force pushing electrons through an external circuit; measured in volts (V).

Standard reduction potential (E°red)

Tabulated potential for a reduction half-reaction under standard conditions (1.0 M, 1.0 atm, pure solids/liquids, typically 25°C).

Standard hydrogen electrode (SHE)

Reference half-cell with E° = 0.00 V: 2H+(aq) + 2e− → H2(g).

Standard cell potential (E°cell)

Cell potential under standard conditions, computed using reduction potentials: E°cell = E°cathode − E°anode.

Intensive property (of E°)

A property that does not depend on amount; E° values are not multiplied by stoichiometric coefficients when balancing redox equations.

Faraday constant (F)

Charge per mole of electrons: approximately 96485 C/mol e−.

Gibbs free energy relationship (ΔG = −nFE)

Connects electrical work and thermodynamics: ΔG = −nFE (and under standard conditions ΔG° = −nF E°cell).

Reaction quotient (Q)

An expression like K but using current concentrations/pressures; products over reactants with exponents from coefficients (omit pure solids/liquids).

Nernst equation

Relates nonstandard cell potential to standard potential and Q: E = E° − (RT/nF)lnQ (at 25°C: E = E° − (0.0592/n)logQ).