How Electrochemistry Works: Batteries, Corrosion, and Electrolysis

The Relationship Between Chemistry and Electricity

Electrochemistry is the branch of chemistry that studies the interconversion of chemical energy and electrical energy. At its core lies a simple but profound idea: certain chemical reactions involve the transfer of electrons from one substance to another, and if that electron transfer is made to flow through an external circuit rather than happening directly, the moving electrons constitute an electric current that can do useful work. Conversely, an externally supplied electric current can drive chemical reactions that would not occur spontaneously. These two directions—chemical energy to electricity, and electricity to chemical energy—define the two pillars of electrochemistry: galvanic cells and electrolytic cells.

The reactions that underpin electrochemistry are oxidation-reduction reactions, or redox reactions. Oxidation is the loss of electrons by a species; reduction is the gain of electrons. The two half-reactions always occur together: if one species loses electrons, another must gain them. A useful mnemonic is OIL RIG: Oxidation Is Loss, Reduction Is Gain. The species that is oxidized acts as the reducing agent (it donates electrons), and the species that is reduced acts as the oxidizing agent (it accepts electrons).

Galvanic Cells: Converting Chemical Energy to Electricity

A galvanic cell (also called a voltaic cell, after Alessandro Volta) is an electrochemical cell in which a spontaneous redox reaction generates electrical energy. The classic demonstration is the Daniell cell, invented in 1836. A strip of zinc is immersed in a zinc sulfate solution in one compartment, and a strip of copper is immersed in a copper sulfate solution in another. The two compartments are connected by a salt bridge (a tube filled with a concentrated electrolyte solution that allows ion flow without mixing), and the two metal strips are connected by an external wire.

At the zinc electrode (the anode), zinc is oxidized: Zn → Zn²⁺ + 2e⁻. Zinc atoms dissolve into solution, leaving two electrons behind on the electrode. Those electrons flow through the external wire toward the copper electrode. At the copper electrode (the cathode), copper ions from solution are reduced: Cu²⁺ + 2e⁻ → Cu. Copper metal deposits on the electrode. The net reaction—Zn + Cu²⁺ → Zn²⁺ + Cu—is spontaneous because zinc is a stronger reducing agent than copper. The Daniell cell generates about 1.1 volts.

The voltage (electromotive force, or EMF) of a galvanic cell reflects the difference in the tendency of the two half-reactions to proceed. Chemists tabulate these tendencies as standard reduction potentials (E°), measured in volts relative to the standard hydrogen electrode. The standard cell EMF is calculated as E°cell = E°cathode − E°anode. Half-reactions with more positive standard reduction potentials are more likely to occur as reductions; those with more negative potentials preferentially occur as oxidations.

Standard Reduction Potentials of Common Half-Reactions

Half-Reaction	E° (V)
F₂ + 2e⁻ → 2F⁻	+2.87
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

Batteries: Portable Electrochemical Power

A battery is a collection of galvanic cells arranged to deliver useful electrical energy. The first true battery was Volta's pile of 1800—alternating discs of zinc and silver separated by brine-soaked cloth. Modern batteries are far more sophisticated but obey the same electrochemical principles.

The lead-acid battery, invented in 1859 and still used as the automotive starter battery, consists of lead (Pb) anodes and lead dioxide (PbO₂) cathodes immersed in sulfuric acid electrolyte. During discharge, both electrodes react with sulfate ions to form lead sulfate. The reaction is highly reversible, enabling the battery to be recharged hundreds of times by applying an external voltage. A 12-volt car battery contains six lead-acid cells in series, each delivering about 2 volts.

The lithium-ion battery, commercialized by Sony in 1991, has revolutionized portable electronics and is now transforming transportation and grid energy storage. In a lithium-ion cell, lithium ions migrate between a graphite anode and a lithium metal oxide cathode (commonly lithium cobalt oxide, LiCoO₂, or lithium iron phosphate, LiFePO₄) through a lithium salt electrolyte. The cell chemistry is more complex than older battery types and requires careful engineering to prevent thermal runaway—a dangerous self-reinforcing overheating that can cause fires. Lithium-ion cells deliver 3.6–3.7 volts per cell, have high energy density, and can withstand hundreds to thousands of charge-discharge cycles.

Comparison of Common Battery Types

Battery Type	Nominal Voltage	Energy Density (Wh/kg)	Rechargeable	Key Application
Zinc-carbon (Leclanché)	1.5 V	65–85	No	Remote controls, clocks
Alkaline	1.5 V	100–150	Limited	Flashlights, toys
Lead-acid	2.0 V/cell	30–50	Yes	Automotive starter
Nickel-metal hydride (NiMH)	1.2 V/cell	60–120	Yes	Hybrid vehicle packs
Lithium-ion	3.6–3.7 V/cell	150–300	Yes	Smartphones, EVs
Solid-state lithium	3.7 V/cell	300–500 (projected)	Yes	Next-gen EVs (emerging)

Corrosion: Electrochemistry in the Wild

Corrosion is the electrochemical degradation of metals by reaction with their environment. It is, in essence, an uncontrolled galvanic cell operating at the surface of a metal. Iron rusting is the most familiar example. When iron is exposed to water and oxygen, small electrochemical cells form on its surface: iron acts as the anode (Fe → Fe²⁺ + 2e⁻) at active sites, while electron-conducting regions of the surface act as the cathode where oxygen is reduced (O₂ + 4H⁺ + 4e⁻ → 2H₂O). The Fe²⁺ ions eventually react with oxygen and water to form hydrated iron oxide, Fe₂O₃·xH₂O—the reddish-brown substance we call rust.

Corrosion costs the global economy an estimated 2.5 trillion dollars per year in infrastructure damage, equipment replacement, and maintenance. Strategies to prevent or retard corrosion are therefore of enormous practical importance. These include:

• Protective coatings: Paints, lacquers, and polymer coatings physically exclude water and oxygen from the metal surface.
• Galvanization: Coating steel with a layer of zinc. Even if the zinc coating is scratched, zinc preferentially oxidizes (acts as a sacrificial anode) because it has a more negative reduction potential than iron, protecting the underlying steel.
• Cathodic protection: Connecting the metal to be protected to a more active metal (the sacrificial anode) or to a power supply that forces the protected metal to act as a cathode, preventing it from being oxidized. Ship hulls, underground pipelines, and offshore oil rigs use this approach.
• Alloying: Stainless steel contains at least 10.5% chromium, which spontaneously forms a thin, adherent chromium oxide (Cr₂O₃) layer on the surface that passivates the alloy, greatly slowing further oxidation.

Electrolysis: Driving Chemistry with Electricity

In an electrolytic cell, an external electric power source drives a non-spontaneous chemical reaction. The power source acts as an electron pump, forcing electrons from the anode (where oxidation occurs) through the external circuit to the cathode (where reduction occurs). The minimum voltage required to drive the electrolysis is at least equal to the reverse of the spontaneous cell potential, plus overpotential—additional voltage needed to overcome energy barriers at the electrode surfaces.

Electrolysis is used in a number of vital industrial processes. The chlor-alkali process electrolyzes brine (concentrated NaCl solution) to produce chlorine gas (at the anode), hydrogen gas (at the cathode), and sodium hydroxide solution. These three products are the feedstocks for an enormous range of chemicals including PVC plastic, paper pulp bleaching agents, and pharmaceuticals. The Hall-Héroult process electrolyzes aluminum oxide dissolved in molten cryolite to produce aluminum metal; this energy-intensive process consumes roughly 5% of all electricity generated in the United States. Electroplating—depositing a thin layer of one metal onto the surface of another—uses electrolysis to create decorative finishes (gold plating, chrome plating) and protective coatings (nickel plating of steel components).

Fuel Cells: A Different Kind of Electrochemistry

A fuel cell is a galvanic cell that generates electricity by continuously consuming fuel and oxidant supplied from an external source, rather than from stored chemical reactants like a conventional battery. The hydrogen fuel cell combines hydrogen gas (at the anode) with oxygen from air (at the cathode) to produce electricity, heat, and water—with no combustion products. The overall reaction is the same as burning hydrogen (H₂ + ½O₂ → H₂O), but because the reaction is electrochemical rather than thermal, the theoretical efficiency is far higher than that of a heat engine.

Fuel cells power spacecraft (NASA has used them since the Gemini program), are being deployed in buses, trucks, and trains, and are under development as stationary power sources for buildings and data centers. The key challenges are the cost of platinum catalysts at the electrodes, the logistics of producing and transporting hydrogen, and the durability of membranes and components over years of operation.

Conclusion

Electrochemistry reveals that the flow of electrons and the transformations of matter are deeply intertwined. Galvanic cells convert the chemical potential energy of spontaneous redox reactions into electricity—powering everything from wristwatches to electric vehicles. Corrosion reminds us that electrochemical forces are constantly at work in the natural world, silently degrading the metals on which modern infrastructure depends. Electrolysis allows us to reverse those natural tendencies using electric power, enabling the large-scale production of aluminum, chlorine, hydrogen, and a host of electroplated products. Fuel cells offer a glimpse of an energy future in which chemical oxidation is tamed into a quiet, efficient, emission-free flow of electrons. All of these phenomena share a common electrochemical foundation that makes this branch of chemistry indispensable to modern science and technology.