Alessandro Volta's voltaic pile (1799) — alternating discs of zinc and copper separated by brine-soaked cardboard — was the first battery, a chemical device that produced steady electric current. Within a year, William Nicholson and Anthony Carlisle had used a voltaic pile to electrolyze water into hydrogen and oxygen; Humphry Davy used larger piles to isolate six new elements (sodium, potassium, calcium, magnesium, strontium, barium) by electrolysis, Michael Faraday (Davy's protégé) formalized the quantitative laws relating charge to mass deposited (1834), and Ostwald, Nernst, and Arrhenius completed the theoretical framework around 1900. Redox chemistry — the chemistry of electron transfer — turned out to underlie everything from rusting iron to cellular respiration to the lithium-ion battery.
Oxidation is the loss of electrons and reduction is the gain — the two always happen together — so an oxidizing agent gets reduced and a reducing agent gets oxidized, with oxidation states tracking electron transfers in complex molecules. The standard reduction potential E° of a half-reaction, measured against the standard hydrogen electrode (2H⁺ + 2e⁻ → H₂, E° = 0 by definition), tabulates the tendency to be reduced; F₂ + 2e⁻ → 2F⁻ at E° = +2.87 V is the strongest common oxidizer, Li⁺ + e⁻ → Li at E° = −3.04 V the strongest common reducer, and E°_cell = E°_cathode − E°_anode, with positive E°_cell meaning spontaneous reaction. The connection to thermodynamics is ΔG° = −nFE°, so positive E°_cell ↔ negative ΔG° ↔ K > 1, while the Nernst equation E = E° − (RT/nF)·ln(Q) gives the actual cell potential under non-standard conditions and explains why battery voltages depend on state of charge. Galvanic cells (batteries) produce electricity from a spontaneous redox reaction; electrolytic cells drive a non-spontaneous reaction with applied voltage. The electrochemical series organizes metals by reduction potential — more easily oxidized metals sacrifice themselves to protect less active ones, the principle behind galvanic corrosion and cathodic protection. The same redox machinery runs biology and industry: cellular respiration is a cascade of redox steps where electrons drop through NADH, FADH₂, coenzyme Q, cytochrome c, oxygen, pumping protons across the mitochondrial membrane to drive ATP synthesis; industrial electrochemistry includes chlor-alkali production, aluminum smelting (the Hall-Héroult process), copper refining, and electroplating.
Lithium-ion batteries (Goodenough, Whittingham, Yoshino, Nobel 2019) are the most consequential battery chemistry of the twenty-first century, powering phones, laptops, electric vehicles, and grid storage — lithium ions shuttle between a graphite anode and a metal-oxide cathode — and solid-state batteries are the next major target. Fuel cells (galvanic cells that consume H₂, methanol, or ammonia continuously) are the long-promised clean alternative for vehicles and stationary power; green hydrogen (from renewable-electricity-driven electrolysis) and electrochemical CO₂ reduction are the central technologies of the electrified-decarbonized energy future. Corrosion costs the world ~3.4% of GDP annually, with redox science underwriting all corrosion-protection strategies, and biosensors (including the home glucose meter) read current from enzyme-catalyzed redox reactions.