Electrochemical Cell · aiwiki.page

An electrochemical cell is a system containing at least two electrodes, which conduct electrons, in contact with electrolytes, which conduct ions. It can generate electrical energy from chemical processes, use electrical energy to drive chemical change, or support measurements with little or no current flowing. Electrochemical cells are fundamental devices in electrochemistry and include power-producing cells, industrial reaction vessels, and analytical measuring systems. (goldbook.iupac.org)

Components and charge transport

In a cell carrying sustained current through electrode reactions, oxidation and reduction occur at different electrode interfaces. Together, these processes constitute a redox reaction: oxidation releases electrons, while reduction accepts them. Electrons travel through the external electronic circuit, whereas ions transport charge through the electrolyte. Sustained operation requires both pathways; ionic conduction prevents the bulk electrolyte compartments from accumulating charges that would rapidly oppose further reaction. (openstax.org)

An electrode may participate chemically, as when a metal dissolves, or primarily provide a conducting surface for reactions involving other substances. Platinum and graphite are commonly used for the latter purpose, although whether a material is effectively inert depends on the reaction conditions. In a divided cell, a salt bridge or ion-conducting separator connects the electrolyte compartments while limiting direct mixing. Separate containers and a salt bridge are not essential features of every electrochemical cell. (openstax.org)

Galvanic and electrolytic operation

A galvanic cell, also called a voltaic cell, operates through a spontaneous chemical reaction and converts chemical energy into electrical energy. An electrolytic cell uses an external electrical supply to drive a reaction that would not proceed spontaneously in the imposed direction. This electrically driven process is electrolysis. These terms describe operating modes: a rechargeable cell normally operates galvanically during discharge and electrolytically during charging. (goldbook.iupac.org)

The anode is the electrode where oxidation occurs, and the cathode is where reduction occurs. These definitions apply in both operating modes, but terminal polarity differs. In a galvanic cell, the anode is negative and the cathode positive; in an electrolytic cell, the anode is positive and the cathode negative. Consequently, “anode” and “cathode” designate reaction roles rather than permanently identifying particular physical electrodes in a reversible device. (openstax.org)

A zinc–copper example

A familiar galvanic arrangement places zinc metal in a solution containing zinc ions and copper metal in a solution containing copper ions. With an external conducting connection and an ionic connection, the electrode reactions are

\text{Anode:}\quad \mathrm{Zn(s)\rightarrow Zn^{2+}(aq)+2e^-}

\text{Cathode:}\quad \mathrm{Cu^{2+}(aq)+2e^-\rightarrow Cu(s)}.

The overall chemical reaction is therefore

\mathrm{Zn(s)+Cu^{2+}(aq)\rightarrow Zn^{2+}(aq)+Cu(s)}.

Zinc dissolves, copper deposits, and electrons flow through the external circuit from zinc toward copper. The electrode processes are called half-reactions because each represents one part of the complete electron-transfer reaction. (openstax.org)

A conventional cell diagram is

\mathrm{Zn(s)\,|\,Zn^{2+}(aq)\,||\,Cu^{2+}(aq)\,|\,Cu(s)}.

Single vertical lines indicate phase boundaries, and the double line indicates the junction between electrolyte compartments. The oxidation half-cell is conventionally written on the left. Under standard conditions, this cell has a potential of approximately $1.10\ \mathrm{V}$ . (openstax.org)

Cell potential and thermodynamics

Cell potential measures energy available per unit charge transferred. When both electrode potentials are expressed as reduction potentials relative to the same reference,

E^\circ_{\mathrm{cell}} =E^\circ_{\mathrm{cathode}}-E^\circ_{\mathrm{anode}}.

Standard electrode potentials are conventionally referenced to the standard hydrogen electrode, assigned zero potential. The standard states specify thermodynamic reference conditions; they do not, by themselves, prescribe one temperature. Electrode potentials are not multiplied by the coefficients used to balance half-reactions. (openstax.org)

The relationship to Gibbs free energy is

\Delta_r G=-nFE_{\mathrm{cell}},

where $n$ is the number of electrons transferred per balanced reaction and $F$ is the Faraday constant. Here, the cell potential is the reversible thermodynamic value, not necessarily the terminal voltage during current flow. A positive potential corresponds to a negative reaction Gibbs energy for the reaction direction specified. At chemical equilibrium, the reaction Gibbs energy and corresponding reversible cell potential are zero. (openstax.org)

The Nernst equation describes the effect of composition and temperature:

E_{\mathrm{cell}} =E^\circ_{\mathrm{cell}}-\frac{RT}{nF}\ln Q.

Here $R$ is the gas constant, $T$ is absolute temperature, and $Q$ is the reaction quotient, rigorously expressed using thermodynamic activities. Concentrations can approximate activities under suitable conditions. A concentration cell exploits this dependence, producing a potential from a composition difference between otherwise similar half-cells. (openstax.org)

Practical operation and applications

Working cells differ from ideal reversible systems. Overpotential is the displacement of an electrode potential from its equilibrium value needed to sustain a specified current. Electrolytic operation can therefore require more applied voltage than the thermodynamic reaction voltage alone suggests. For example, splitting water into hydrogen and oxygen has a standard reversible voltage requirement of approximately $1.23\ \mathrm{V}$ , but practical operation also involves kinetic and resistive losses. (goldbook.iupac.org)

A battery contains one or more electrochemical cells. Primary batteries are intended for non-rechargeable service; secondary batteries, including lithium-ion batteries, are designed for repeated charging and discharge. A fuel cell instead receives fuel and oxidant from external supplies during operation. Electrolytic cells are used for metal deposition, refining, and chemical production. (openstax.org)

Cells used in analytical chemistry need not produce useful power. Potentiometric measurements can operate at essentially zero current. Three-electrode arrangements separate the working electrode, where the process under study occurs, from a reference electrode used to measure potential and an auxiliary electrode used to carry current. This separates potential measurement from the current-carrying role of the second electrode. (goldbook.iupac.org)