Electrochemistry – Core Vocabulary

Flashcards summarise essential vocabulary and definitions from the Electrochemistry lecture, covering cell types, thermodynamic relations, conductivity concepts, batteries, fuel cells, and corrosion.

Electrochemistry

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

Galvanic (Voltaic) Cell

Electrochemical cell in which a spontaneous redox reaction produces electrical energy.

Electrolytic Cell

Electrochemical cell that uses an external source of electricity to drive a non-spontaneous redox reaction.

Daniell Cell

Classic galvanic cell with a Zn/Zn²⁺ anode and Cu/Cu²⁺ cathode separated by a salt bridge; E° ≈ 1.10 V.

Half-Cell (Redox Couple)

One electrode of an electrochemical cell together with its surrounding electrolyte where either oxidation or reduction occurs.

Anode

Electrode where oxidation takes place; negative in a galvanic cell and positive in an electrolytic cell.

Cathode

Electrode where reduction takes place; positive in a galvanic cell and negative in an electrolytic cell.

Electrode Potential

Potential difference between an electrode and its electrolyte arising from 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.

Cell Potential (EMF)

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

Standard Hydrogen Electrode (SHE)

Reference half-cell assigned potential 0 V; Pt black electrode in 1 M H⁺ with H₂ gas at 1 bar.

Nernst Equation

Relation that gives electrode or cell potential at any concentration: E = E° – (RT/nF) ln Q.

Gibbs Energy–EMF Relation

ΔG = –nFEcell and ΔG° = –nFE°cell, linking electrical work to thermodynamics.

Equilibrium Constant–EMF Relation

E°cell = (0.059 V / n) log K at 298 K, connecting cell potential to reaction equilibrium constant.

Resistivity (ρ)

Intrinsic resistance of a material; resistance of a cube with 1 m sides, units Ω m.

Conductivity (κ)

Reciprocal of resistivity; conductance of a 1 m³ block, units S m⁻¹.

Cell Constant (G*)

Geometric factor l/A for a conductivity cell; determined using a KCl standard solution.

Conductivity Cell

Special vessel with platinised Pt electrodes used to measure solution resistance with AC.

Molar Conductivity (Λm)

Conductivity of an electrolyte divided by concentration; Λm = κ / c, units S m² mol⁻¹.

Limiting Molar Conductivity (Λm°)

Molar conductivity extrapolated to infinite dilution where ion interactions are negligible.

Kohlrausch’s Law

At infinite dilution, Λm° equals the sum of independent ionic contributions: Λm° = ν+λ+° + ν–λ–°.

Degree of Dissociation (α)

For weak electrolytes, fraction ionised; α ≈ Λm / Λm°.

Faraday Constant (F)

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

Faraday’s First Law

Mass of substance deposited ∝ quantity of electricity passed.

Faraday’s Second Law

For equal charge, masses deposited are proportional to equivalent weights of substances.

Overpotential

Extra potential required beyond thermodynamic value to drive a kinetically slow electrode reaction.

Primary Battery

Disposable galvanic cell that cannot be recharged; e.g., dry (Leclanché) cell.

Dry Cell (Leclanché)

Zn anode, MnO₂/C cathode, NH₄Cl–ZnCl₂ paste electrolyte; E ≈ 1.5 V.

Mercury Cell

Zn-Hg amalgam anode, HgO cathode in KOH/ZnO paste; constant 1.35 V.

Secondary Battery

Rechargeable electrochemical cell; e.g., lead storage or Ni–Cd battery.

Lead Storage Battery

Pb anode, PbO₂ cathode in 38 % H₂SO₄; reversible reaction yields ~2 V per cell.

Nickel–Cadmium Cell

Cd anode and NiO(OH) cathode in KOH; rechargeable with long cycle life.

Fuel Cell

Device that continuously converts chemical energy of fuel (e.g., H₂) and oxidant (O₂) directly into electricity.

Hydrogen–Oxygen Fuel Cell

H₂ anode and O₂ cathode in alkaline electrolyte; only product is water, efficiency ~70 %.

Electronic Conductivity

Charge transport via electrons in metals or semiconductors.

Ionic Conductivity

Charge transport through movement of ions in electrolyte solutions.

Superconductor

Material exhibiting zero resistivity below a critical temperature.

Corrosion

Electrochemical deterioration of metals, e.g., rusting of iron to Fe₂O₃·xH₂O.

Sacrificial Protection

Preventing corrosion by attaching a more easily oxidised metal (e.g., Mg) that corrodes instead.

Salt Bridge

Electrolyte connection maintaining charge neutrality between half-cells and completing the circuit.

Cell Notation

Conventional shorthand for galvanic cells: anode | electrolyte || electrolyte | cathode.

Activity

Effective concentration used in thermodynamic calculations; approximates molarity in dilute solutions.