What Is Electrochemistry?
Electrochemistry studies the interconversion between chemical energy and electrical energy. It encompasses two main types of cells:
- Galvanic (voltaic) cells: Convert chemical energy → electrical energy (spontaneous reactions)
- Electrolytic cells: Convert electrical energy → chemical energy (non-spontaneous reactions, driven by external power)
Both rely on oxidation-reduction (redox) reactions — reactions involving the transfer of electrons from one species to another.
Oxidation and Reduction
In every redox reaction: - Oxidation = loss of electrons (OIL — Oxidation Is Loss) - Reduction = gain of electrons (RIG — Reduction Is Gain)
The species that is oxidized is the reducing agent (electron donor); the species that is reduced is the oxidizing agent (electron acceptor). Oxidation and reduction always occur together.
Example: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) - Zn is oxidized: Zn → Zn²⁺ + 2e⁻ - Cu²⁺ is reduced: Cu²⁺ + 2e⁻ → Cu - Zn is the reducing agent; Cu²⁺ is the oxidizing agent
Galvanic Cells: Harnessing Spontaneous Reactions
A galvanic cell separates the two half-reactions so that electron transfer occurs through an external circuit, generating electrical current.
Anatomy of a galvanic cell: - Anode: Electrode where oxidation occurs (negative terminal in galvanic cell) - Cathode: Electrode where reduction occurs (positive terminal in galvanic cell) - Salt bridge (or porous membrane): Maintains electrical neutrality by allowing ion flow between half-cells - External circuit: Path for electrons to flow from anode to cathode
Memory aid: AN OX, RED CAT (ANodee = OXidation; REDuction = CAThodee)
For the Zn/Cu cell: - Anode half-cell: Zn(s) | Zn²⁺(aq) — zinc dissolves, releasing electrons - Cathode half-cell: Cu²⁺(aq) | Cu(s) — copper ions deposit as metal
Cell Potential (EMF)
The cell potential (E_cell) measures the driving force of the electrochemical reaction in volts (V). It is the difference in reduction potentials between cathode and anode:
E°_cell = E°_cathode − E°_anode
Standard reduction potentials (E°) are measured relative to the standard hydrogen electrode (SHE), which is assigned E° = 0.000 V by convention. Tables of standard reduction potentials allow calculation of E°_cell for any galvanic cell.
- E°_cell > 0: Spontaneous reaction (ΔG < 0)
- E°_cell < 0: Non-spontaneous (requires external energy)
For Zn/Cu cell: E°(Cu²⁺/Cu) = +0.34 V, E°(Zn²⁺/Zn) = −0.76 V E°_cell = 0.34 − (−0.76) = +1.10 V (spontaneous)
Connection to Gibbs Free Energy
The relationship between cell potential and Gibbs free energy:
ΔG° = −nFE°_cell
Where n is moles of electrons transferred and F is Faraday's constant (96,485 C/mol ≈ 96,500 C/mol).
Combined with ΔG° = −RT ln K:
E°_cell = (RT/nF) ln K = (0.0257/n) ln K at 298 K
The Nernst Equation
Under non-standard conditions (concentrations ≠ 1 M), cell potential is calculated using the Nernst equation:
E = E° − (RT/nF) ln Q = E° − (0.0592/n) log Q at 298 K
Where Q is the reaction quotient. As a cell discharges, reactant concentrations decrease and product concentrations increase, Q increases, and E decreases — until E = 0 at equilibrium (a dead battery).
Application: Concentration cells exploit the Nernst equation — even identical electrodes can generate a voltage if the ion concentrations in the two half-cells differ.
Electrolytic Cells: Driving Non-Spontaneous Reactions
In electrolytic cells, an external voltage source drives a non-spontaneous reaction. The electrode designations reverse from galvanic cells: - Cathode (connected to negative terminal): Reduction occurs - Anode (connected to positive terminal): Oxidation occurs
Applications of electrolysis: - Electroplating: Depositing a thin layer of metal (chromium, gold, nickel) onto a surface for corrosion resistance or aesthetics - Hall-Héroult process: Industrial production of aluminum by electrolysis of molten Al₂O₃ - Chlor-alkali process: Electrolysis of brine (NaCl + H₂O) to produce Cl₂, NaOH, and H₂ — essential industrial chemicals - Water electrolysis: 2H₂O → 2H₂ + O₂ — hydrogen fuel production
Faraday's Laws of Electrolysis
First Law: The mass of substance deposited or dissolved at an electrode is proportional to the total charge passed.
Second Law: For the same charge, the masses of different substances deposited are proportional to their equivalent weights.
Quantitatively: m = (M × Q) / (n × F)
Where m = mass (g), M = molar mass (g/mol), Q = charge (coulombs), n = electrons per atom.
Example: How much copper is deposited when 2 A flows for 10 minutes? Q = 2 × 600 = 1200 C. m = (63.55 × 1200) / (2 × 96485) = 0.395 g Cu.
Real-World Applications
- Batteries: From simple Daniell cells to modern lithium-ion batteries — all are galvanic cells. Li-ion batteries achieve ~3.7 V per cell because of large ΔE° between lithium (strong reducing agent) and cathode materials like LiCoO₂
- Fuel cells: H₂ + ½O₂ → H₂O generates electricity at ~1.23 V per cell with only water as a byproduct
- Corrosion: Galvanic corrosion occurs when two dissimilar metals are in electrical contact in an electrolyte — the more active metal (lower reduction potential) corrodes preferentially
- pH meters: Use electrochemical cells where E varies with H⁺ concentration (Nernst equation)
- Electrochemical sensors: Breathalyzers, blood glucose meters, and dissolved oxygen sensors all use electrochemical principles
Summary
Electrochemistry bridges chemistry and electricity through the controlled transfer of electrons. Galvanic cells convert spontaneous chemical reactions into useful electrical energy, while electrolytic cells use electrical energy to drive otherwise impossible reactions. The quantitative framework — standard potentials, the Nernst equation, Faraday's laws — provides powerful tools for designing batteries, plating metals, and understanding corrosion.