Electrochemistry: Oxidation, Reduction, and How Batteries Work

Alessandro Volta announced his pile — the first true electric battery — in a letter to the Royal Society in 1800. He had stacked alternating discs of zinc and silver separated by saltwater-soaked cloth and found that the stack produced a continuous electrical current. He understood little of the chemistry involved, but the discovery launched a century of electrochemical investigation that eventually gave scientists the tools to isolate potassium, sodium, aluminum, and chlorine, and gave the world the telegraph, electroplating, and ultimately every rechargeable battery in modern electronics.

Oxidation and Reduction: The Electron Transfer Foundation

Every electrochemical process is fundamentally a redox reaction — a simultaneous oxidation and reduction. Oxidation is the loss of electrons; reduction is the gain of electrons. The two always occur together because electrons released by one species must be accepted by another. The mnemonic OIL RIG captures this: Oxidation Is Loss, Reduction Is Gain.

In a zinc-copper galvanic cell, zinc atoms at the anode give up two electrons to become Zn²⁺ ions in solution — zinc is oxidized. The electrons travel through the external circuit to the copper cathode, where Cu²⁺ ions in solution accept them and deposit as copper metal — copper(II) is reduced:

• Anode (oxidation): Zn → Zn²⁺ + 2e⁻    E° = +0.76 V
• Cathode (reduction): Cu²⁺ + 2e⁻ → Cu    E° = +0.34 V
• Overall cell potential: E°_cell = E°_cathode − E°_anode = 0.34 − (−0.76) = 1.10 V

The standard reduction potential E° (measured in volts against the standard hydrogen electrode) quantifies the tendency of a species to gain electrons. A large positive E° means strong tendency to be reduced (a good oxidizing agent). A large negative E° means strong tendency to be oxidized (a good reducing agent). The table of standard reduction potentials — the electrochemical series — allows chemists to predict whether any given redox reaction will proceed spontaneously.

The Nernst Equation: Potential Under Real Conditions

Standard potentials assume unit activity (approximately 1 M concentrations, 1 atm gas pressures, 25°C). Real cells operate under different conditions. The Nernst equation corrects the cell potential for actual concentrations:

E = E° − (RT / nF) ln Q

where R is the gas constant (8.314 J/mol·K), T is temperature in Kelvin, n is the number of electrons transferred, F is Faraday's constant (96,485 C/mol), and Q is the reaction quotient. At 25°C this simplifies to:

E = E° − (0.0592 V / n) log Q

The Nernst equation explains why battery voltage drops as it discharges (reactant concentrations fall, Q increases, E decreases) and why temperature affects battery performance (cold weather reduces lithium-ion battery capacity).

Standard Reduction Potentials of Common Half-Reactions

Half-Reaction	E° (V vs. SHE)
F2 + 2e− → 2F−	+2.87 (strongest oxidizer)
MnO4− + 8H+ + 5e− → Mn2+ + 4H2O	+1.51
O2 + 4H+ + 4e− → 2H2O	+1.23
Cu2+ + 2e− → Cu	+0.34
2H+ + 2e− → H2	0.00 (reference)
Zn2+ + 2e− → Zn	−0.76
Al3+ + 3e− → Al	−1.66
Li+ + e− → Li	−3.04 (strongest reducer)

Electrolysis: Driving Non-Spontaneous Reactions

When an external voltage drives electrons through a cell in the direction opposite to spontaneous flow, non-spontaneous reactions occur. This is electrolysis. The external power supply forces oxidation at the anode and reduction at the cathode regardless of what the standard potentials prefer.

Faraday's laws of electrolysis (1833) quantify this process. The mass of substance produced at an electrode is proportional to the charge passed (Q = It, in coulombs) and to the molar mass divided by the number of electrons in the half-reaction:

m = (M × I × t) / (n × F)

Electrolysis is the basis of aluminum production (Hall-Héroult process: molten Al₂O₃ at ~960°C, yielding Al at cathode), electroplating (copper, nickel, chromium, gold deposition on substrate metals), chlorine and caustic soda production (chlor-alkali process from brine), and water splitting for hydrogen fuel production.

Lithium-Ion Batteries: Intercalation Chemistry

Unlike older battery chemistries where the electrode material dissolves and deposits during cycling, lithium-ion batteries work by intercalation — lithium ions move in and out of the crystal structure of electrode materials without dissolving them. This preserves electrode structure and allows thousands of charge-discharge cycles.

• Anode: Graphite, where Li⁺ ions insert between graphene layers during charging. LiC₆ at full charge. E° ≈ −3.0 V vs. Li/Li⁺.
• Cathode: Lithium metal oxides — LiCoO₂ (original Sony cell), LiFePO₄ (safer, longer life), LiNiMnCoO₂ (NMC, high energy density). Li⁺ deintercalates during charging.
• Electrolyte: Lithium salt (LiPF₆) dissolved in organic carbonates (ethylene carbonate, dimethyl carbonate). Must be non-aqueous — water would be electrolyzed at lithium potentials.
• Solid electrolyte interphase (SEI): A thin layer that forms on the anode during the first charge, passivating the surface and enabling stable cycling. Its composition determines cycle life and safety.

Battery Chemistry	Nominal Voltage	Energy Density (Wh/kg)	Cycle Life	Primary Application
Lead-acid	2.0 V/cell	30–50	200–500	Automotive starter batteries
Nickel-Metal Hydride	1.2 V/cell	60–120	500–1,000	Hybrid vehicles (Toyota Prius)
LiCoO2 / graphite	3.6 V	150–200	300–500	Consumer electronics
LiFePO4 / graphite	3.2 V	90–140	2,000–5,000	EVs, stationary storage
NMC / graphite	3.7 V	150–250	1,000–2,000	Electric vehicles

Corrosion: Electrochemistry's Destructive Side

Iron rusting is a galvanic process. Iron in contact with oxygen and moisture forms micro-galvanic cells: iron acts as the anode (Fe → Fe²⁺ + 2e⁻), and oxygen reduction occurs at cathodic sites (O₂ + 2H₂O + 4e⁻ → 4OH⁻). The Fe²⁺ and OH⁻ combine and oxidize further to produce Fe₂O₃·H₂O — rust. Corrosion costs the global economy roughly $2.5 trillion per year — about 3.4% of global GDP. Cathodic protection, galvanizing (zinc coating), and sacrificial anodes (attaching zinc or magnesium blocks to steel structures) all exploit electrochemical principles to prevent this destruction.