Unit 2 – Electrochemistry: Key Vocabulary

Vocabulary flashcards summarising essential terms, concepts and devices discussed in the Electrochemistry unit.

Electrochemistry

Branch of chemistry that studies conversion between chemical energy and electrical energy through redox reactions.

Electrochemical cell

Any device that converts chemical energy to electrical energy or vice-versa by means of redox reactions occurring at two electrodes.

Galvanic (Voltaic) cell

An electrochemical cell in which a spontaneous redox reaction produces electrical energy.

Electrolytic cell

An electrochemical cell that uses external electrical energy to drive a non-spontaneous chemical reaction.

Daniell cell

Classic galvanic cell consisting of a Zn/Zn²⁺ half-cell and a Cu/Cu²⁺ half-cell connected by a salt bridge, emf ≈ 1.1 V under standard conditions.

Half-cell (Redox couple)

One electrode of an electrochemical cell where either oxidation or reduction occurs, described by an electrode and its ionic solution.

Anode

Electrode where oxidation occurs; negative in a galvanic cell, positive in an electrolytic cell.

Cathode

Electrode where reduction occurs; positive in a galvanic cell, negative in an electrolytic cell.

Electrode potential

Potential difference developed between an electrode and its electrolyte due to charge separation at equilibrium.

Standard electrode potential (E°)

Electrode potential measured under standard conditions (1 M, 1 bar, 298 K) relative to the standard hydrogen electrode.

Standard hydrogen electrode (SHE)

Reference electrode assigned E° = 0 V, consisting of a platinum black electrode in 1 M H⁺ with H₂ gas at 1 bar.

Cell potential (emf)

Difference between cathode and anode potentials when no current flows; Ecell = Eright – Eleft.

Cell notation

Conventional way to write a galvanic cell: anode | electrolyte || electrolyte | cathode, left-to-right representing electron flow.

Salt bridge

Ionic conductor (often a U-tube with agar–salt gel) that completes the circuit and maintains electroneutrality between half-cells.

Nernst equation

Expression giving electrode or cell potential as a function of temperature and reactant activities: E = E° – (0.059/n) log Q at 298 K.

Gibbs free-energy–emf relation

ΔG° = –nFE°cell, linking cell potential with standard Gibbs energy of the cell reaction.

Equilibrium constant–emf relation

E°cell = (0.059/n) log K at 298 K, connecting standard cell potential with the reaction’s equilibrium constant.

Conductance (G)

Reciprocal of resistance (R); measured in siemens (S).

Resistance (R)

Opposition offered by a conductor to electric current; R = ρ l/A, unit ohm (Ω).

Resistivity (ρ)

Intrinsic resistance of a material per unit length and cross-section; unit Ω m.

Conductivity (κ)

Reciprocal of resistivity; conductance of a 1 m cube of material; unit S m⁻¹.

Cell constant (G*)

Geometric factor l/A for a conductivity cell; κ = G*/R.

Molar conductivity (Λm)

Conductivity of an electrolyte solution divided by its molar concentration, Λm = κ/c; units S m² mol⁻¹ (or S cm² mol⁻¹).

Limiting molar conductivity (Λm°)

Molar conductivity extrapolated to zero concentration (infinite dilution) where the electrolyte is fully dissociated.

Kohlrausch’s law

Λm° of an electrolyte equals the sum of the limiting ionic conductivities of its cation and anion: Λm° = ν⁺λ⁺° + ν⁻λ⁻°.

Ionic conductivity

Electrical conduction in electrolytes via migration of ions.

Electronic conductivity

Conduction through movement of electrons as in metals and semiconductors.

Faraday constant (F)

Charge of one mole of electrons; F ≈ 96 487 C mol⁻¹.

Faraday’s first law

Mass of substance deposited or liberated at an electrode is proportional to the quantity of electricity passed.

Faraday’s second law

For the same charge, masses of substances liberated are proportional to their equivalent weights.

Overpotential

Extra potential beyond the thermodynamic requirement needed to drive an electrode reaction at a finite rate.

Primary battery

Non-rechargeable galvanic cell; reaction proceeds only once, e.g., dry cell.

Dry cell (Leclanché cell)

Common 1.5 V primary battery with Zn anode, MnO₂/C cathode, and NH₄Cl-ZnCl₂ paste electrolyte.

Mercury cell

Compact primary battery using Zn-Hg amalgam anode and HgO cathode; emf ≈ 1.35 V with nearly constant voltage.

Secondary battery

Rechargeable cell that can undergo multiple discharge–charge cycles, e.g., lead storage battery.

Lead storage battery

Secondary battery with Pb anode, PbO₂ cathode, and H₂SO₄ electrolyte; overall discharge forms PbSO₄ and water.

Nickel–cadmium cell

Rechargeable cell in which Cd and NiO(OH) act as active materials; noted for long cycle life.

Fuel cell

Galvanic device continuously supplied with reactants (fuel and oxidant) and delivering electricity directly from their redox reaction.

Hydrogen–oxygen fuel cell

Fuel cell where H₂ is oxidised and O₂ reduced in alkaline electrolyte, producing water and electricity (η ≈ 70 %).

Corrosion

Undesired electrochemical oxidation of metals by environmental agents, leading to deterioration (e.g., rusting of iron).

Sacrificial anode

More reactive metal attached to a structure to corrode preferentially, protecting the main metal from corrosion.

Hydrogen economy

Energy system concept where H₂ produced via electrolysis serves as a clean, renewable fuel replacing fossil fuels.

Superconductor

Material that exhibits zero resistivity (infinite conductivity) below a critical temperature.

Semiconductor

Material with conductivity between that of conductors and insulators; conductivity can be modified by doping.