Electrochemistry

Electrochemistry is a branch of physical chemistry concerned with the relationships between electrical phenomena and chemical change. It examines how reactions generate electricity, how electrical energy drives reactions, and how charge moves across interfaces between electronic and ionic conductors. Its central systems contain electrodes in contact with an electrolyte, linking electron flow in an external circuit to ionic conduction within the system. Electrochemistry provides the scientific basis for batteries, fuel cells, electrolysis, and numerous chemical measurement techniques. (goldbook.iupac.org)

Reactions and electrochemical cells

Many electrochemical processes involve a redox reaction: oxidation releases electrons, while reduction consumes them. These complementary processes can occur at separate electrodes. The electrode where oxidation predominates is the anode; the electrode where reduction predominates is the cathode. These names describe reaction direction, not a permanently assigned electrical polarity. (old.iupac.org)

An electrochemical cell contains at least two electrodes connected through an ionic conducting medium. Electrons travel through the electronic circuit, whereas ions carry charge through the electrolyte. In divided cells, a salt bridge or an ion-conducting separator provides ionic continuity between compartments. A galvanic cell, also called a voltaic cell, converts chemical energy into electrical energy through a spontaneous reaction. An electrolytic cell uses supplied electrical energy to drive chemical change. (goldbook.iupac.org)

A familiar galvanic example pairs zinc oxidation with copper-ion reduction:

The overall reaction is $\mathrm{Zn+Cu^{2+}\rightarrow Zn^{2+}+Cu}$. During galvanic operation, the zinc anode is negative and the copper cathode positive. In an electrolytic cell, the anode is positive and the cathode negative. (old.iupac.org)

Thermodynamics and electrode potential

Thermodynamics determines whether a proposed cell reaction is energetically favorable, but does not by itself establish how rapidly it proceeds. For a reversible cell reaction at constant temperature and pressure, the reaction Gibbs free-energy change is related to cell potential by

where $n$ is the number of electrons transferred in the balanced reaction and $F$ is the Faraday constant, approximately $96{,}485\ \mathrm{C\,mol^{-1}}$. A positive reversible cell potential corresponds to a negative reaction free-energy change in the specified direction. (old.goldbook.iupac.org)

Electrode potentials are reported relative to a reference electrode rather than measured as isolated absolute values. The standard hydrogen electrode provides the conventional reference for standard electrode potentials. When both half-reactions are written as reductions, the cell potential is the cathode reduction potential minus the anode reduction potential. (iupac.org)

The Nernst equation describes the dependence of reversible cell potential on composition:

Here $R$ is the gas constant, $T$ is absolute temperature, and $Q$ is the reaction quotient, constructed from dimensionless activities. Replacing activities with concentrations is an approximation; formal potentials can incorporate effects of specified solution conditions. At chemical equilibrium, the reaction has no net thermodynamic driving force and its reversible cell potential is zero. (goldbook.iupac.org)

Charge, kinetics, and electrode interfaces

Faraday’s laws connect chemical conversion to transferred electric charge. For a specified reaction transferring $n$ electrons per reacting entity, charge $q$ corresponds to a reacted amount $q/(nF)$. This relationship underlies quantitative electrolysis and coulometric analysis. Its application to a particular product requires accounting for competing electrode reactions. (old.goldbook.iupac.org)

Actual electrode operation also depends on reaction kinetics. An overpotential is the deviation of an electrode’s potential from its equilibrium value needed to sustain a specified current. Consequently, equilibrium potentials alone cannot predict the operating voltage of a device carrying appreciable current. (goldbook.iupac.org)

Electrocatalysis changes electrode-reaction rates through catalytic materials or surface species. It is important in hydrogen evolution, oxygen reduction, and water oxidation. Current associated with oxidation or reduction is called faradaic current; distinguishing it from other contributions is important when interpreting electrochemical measurements. (goldbook.iupac.org)

Experimental methods

Electrochemical experiments measure or control potential, current, charge, and their variation with time. A three-electrode arrangement separates functions: the working electrode hosts the process under investigation, the counter electrode carries current, and the reference electrode provides a potential reference. This arrangement supports controlled measurements of working-electrode behavior. (goldbook.iupac.org)

In cyclic voltammetry, electrode potential is swept between selected limits and reversed while current is recorded. The resulting current–potential curve helps investigate electrode reactions, their mechanisms, and electrochemical parameters. Electrochemical impedance spectroscopy examines frequency-dependent electrical response, providing complementary information about electrochemical reaction behavior. These techniques form part of analytical chemistry. (goldbook.iupac.org)

Applications and historical development

A battery generates electrical energy through electrochemical reactions. A fuel cell instead receives fuel and oxidant from external supplies and can continue producing electricity while they are supplied. In a hydrogen fuel cell, separated oxidation and reduction reactions yield water, electrical energy, and heat. Conversely, electrolysis can split water into hydrogen and oxygen using electricity. Its energy requirements and environmental consequences depend partly on the electricity source and electrolyzer performance. (energy.gov)

Alessandro Volta developed the voltaic pile around 1800, providing a continuous electrical source for experiments. In the early 1830s, Michael Faraday established the laws of electrolysis and helped introduce terminology including electrode, cathode, and ion. These developments connected electrical measurement with quantitative chemical transformation. (acs.org)