Galvanic and Electrolytic Cells: The Electrochemistry of Batteries and Plating

Volta's Discovery Changed Civilization

In 1800, Alessandro Volta announced the voltaic pile — a stack of alternating zinc and silver discs separated by brine-soaked cloth that produced a continuous electric current. Within weeks of his announcement, William Nicholson and Anthony Carlisle used a voltaic pile to electrolyze water, splitting it into hydrogen and oxygen. Within a decade, Humphry Davy had used electrolysis to isolate sodium, potassium, calcium, barium, magnesium, and strontium — elements previously unknown. The voltaic pile did not merely power lamps; it reorganized the elemental table. Electrochemistry is inseparable from the history of element discovery and the history of energy storage.

Oxidation and Reduction: Fundamental Definitions

Electrochemistry rests on the complementary processes of oxidation (loss of electrons) and reduction (gain of electrons). Two mnemonics dominate chemistry education:

• OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons).
• AN OX RED CAT: Oxidation occurs at the ANode; REDuction at the CAThode.

These apply universally to both galvanic (spontaneous, electricity-generating) and electrolytic (driven, electricity-consuming) cells — with the polarity of the electrodes reversed:

Standard Reduction Potentials and the Standard Hydrogen Electrode

The electromotive force (EMF) of a galvanic cell — the voltage it produces — is the difference between the reduction potentials of its two half-cells. Absolute half-cell potentials cannot be measured; only differences are measurable. The standard hydrogen electrode (SHE) is assigned E° = 0.000 V by convention and serves as the universal reference:

2H+(aq, 1M) + 2e− → H2(g, 1 atm)    E° = 0.000 V

Cell EMF: E°cell = E°cathode − E°anode. For a Daniel cell (Zn/Zn2+ || Cu2+/Cu): E° = +0.34 − (−0.76) = +1.10 V. A positive E°cell corresponds to a spontaneous reaction (ΔG° = −nFE°cell, where n = moles of electrons transferred and F = Faraday constant = 96,485 C/mol).

The Daniel Cell: Electrochemistry in Action

The Daniel cell, invented by John Frederic Daniell in 1836, consists of a zinc electrode in zinc sulfate solution and a copper electrode in copper sulfate solution, connected by a salt bridge (typically KNO3 in agar gel). At the anode: Zn(s) → Zn2+(aq) + 2e−. The electrons travel through the external circuit to the copper cathode: Cu2+(aq) + 2e− → Cu(s). The salt bridge maintains electrical neutrality by allowing ion migration between half-cells without allowing the solutions to mix directly.

As the cell operates, zinc dissolves and copper deposits — the electrode masses change in measurable ways consistent with Faraday's laws. The driving force is the 1.10 V potential difference, which reflects the greater tendency of zinc to lose electrons (negative reduction potential) relative to copper (positive reduction potential). Nature moves electrons from the stronger reductant (zinc) to the weaker reductant (copper). Thermodynamics drives the machine.

The Nernst Equation: Non-Standard Conditions

Standard reduction potentials apply only at specific conditions: 298 K, 1 atm for gases, 1 M for solutes. Real cells operate under non-standard conditions. The Nernst equation gives cell potential at any concentration:

E = E° − (RT/nF) × ln Q

At 298 K, this simplifies to: E = E° − (0.0592/n) × log Q

where Q is the reaction quotient. As a cell discharges, reactant concentrations fall and product concentrations rise — Q increases, making the term (RT/nF)lnQ larger, and cell voltage decreases. When Q = K (equilibrium), E = 0 — the cell is dead. This relationship explains why battery voltage drops as it discharges and why a battery read in open circuit (no current draw) reads higher than under load. Concentration is voltage in electrochemical terms.

Faraday's Laws of Electrolysis

Michael Faraday formulated two empirical laws of electrolysis in 1833–1834 that remain exactly accurate today:

• First law: The mass of substance deposited or dissolved at an electrode is directly proportional to the quantity of charge passed (Q = It, where I is current in amperes and t is time in seconds).
• Second law: For the same quantity of electricity, masses of different substances deposited are proportional to their molar masses divided by the number of electrons involved in their electrode reaction (equivalents).

The mass deposited: m = (Q × M) / (n × F) = (I × t × M) / (n × F). This formula enables precise calculation of plating thickness, electrorefining purity, and electrolytic production rates. To deposit 1 gram of gold (M = 197, n = 3): charge required = (1 × 3 × 96485) / 197 = 1,469 C. At 1 ampere, this takes about 24.5 minutes. Faraday's laws are exact in practice — the only electrochemical relationships with virtually no exceptions.

Lithium-Ion Battery Electrochemistry

The lithium-ion battery, commercialized by Sony in 1991 (following foundational work by Whittingham, Goodenough, and Yoshino — Nobel 2019), stores and releases energy through the reversible intercalation of lithium ions between electrode materials.

• Cathode (positive electrode, discharge): LiCoO2 (or LiFePO4, NMC, NCA variants). During discharge: LiCoO2 → Li1−xCoO2 + xLi+ + xe−. Lithium ions deintercalate from cathode.
• Anode (negative electrode, discharge): Graphite (layered structure). During discharge: LixC6 → C6 + xLi+ + xe−. Lithium ions release from graphite layers.
• Electrolyte: Lithium salt (LiPF6) in organic carbonate solvents — ion-conducting but electronically insulating.

Cell voltage: approximately 3.6–3.7 V (LiCoO2/graphite), far higher than lead-acid (2.0 V) or NiMH (1.2 V). Energy density: 150–250 Wh/kg, enabling portable electronics and electric vehicles. The solid electrolyte interface (SEI) — a passivating film formed on graphite during first charge from electrolyte reduction — is essential for long cycle life; its chemistry determines aging rate. Solid-state electrolytes (replacing liquid with ceramic or polymer) are under intensive development to eliminate flammability risk and enable lithium metal anodes.

Electroplating and Corrosion as Galvanic Cells

Electroplating deposits a thin metal layer on a conductive substrate using electrolytic current. The object to be plated is the cathode; the plating metal (or inert electrode in some baths) is the anode; the electrolyte contains the plating metal ion.

Industrial chromium plating uses chromic acid baths (H2CrO4): Cr3+ reduced at cathode deposits hard, corrosion-resistant chrome. Thickness of deposit is controlled precisely by Faraday's first law: to achieve 10 μm chrome on 1 cm2 (Cr, M=52, n=3, density 7.19 g/cm3): mass = 10×10^−4 cm × 1 cm2 × 7.19 g/cm3 = 7.19 × 10^−3 g; charge required = (7.19×10^−3 × 3 × 96485) / 52 = 40.1 C.

Corrosion is an uncontrolled galvanic process. Iron-copper contact in seawater creates a galvanic cell: iron (more negative reduction potential) acts as the anode and dissolves; copper is protected as the cathode. Galvanic corrosion between dissimilar metals in marine environments causes billions in infrastructure damage annually. Cathodic protection — attaching a more reactive metal (zinc or magnesium "sacrificial anode") to the structure to be protected — reverses the polarity: the sacrificial anode dissolves, protecting the steel cathode. Ship hulls, pipelines, and bridge foundations rely on this electrochemical principle. Controlled sacrifice prevents uncontrolled loss.