{{Short description|Electro-chemical device}}
{{Use mdy dates|date=October 2020}}
[[File:ElectrochemCell.png|thumb|A demonstration electrochemical cell setup resembling the [[Daniell cell]]. The two half-cells are linked by a salt bridge carrying ions between them. Electrons flow in the external circuit.]]

An '''electrochemical cell''' is a device that either generates [[electrical energy]] from [[chemical reaction]]s in a so-called [[galvanic cell|galvanic]] or voltaic cell, or induces chemical reactions ([[electrolysis]]) by applying external electrical energy in an [[electrolytic cell]].<ref name=":23">{{Cite journal |last=Wenzel |first=Thomas J. |date=2013-07-30 |title=Douglas A. Skoog, Donald M. West, F. James Holler, and Stanley R. Crouch: Fundamentals of analytical chemistry, 9th ed., international ed. |url=http://dx.doi.org/10.1007/s00216-013-7242-1 |journal=Analytical and Bioanalytical Chemistry |volume=405 |issue=25 |pages=412–432 |doi=10.1007/s00216-013-7242-1 |s2cid=94566587 |issn=1618-2642|url-access=subscription }}</ref><ref>{{Citation |last1=Wendt |first1=Hartmut |title=Electrochemistry, 1. Fundamentals |date=2011-10-15 |url=https://onlinelibrary.wiley.com/doi/10.1002/14356007.a09_183.pub4 |encyclopedia=Ullmann's Encyclopedia of Industrial Chemistry |pages=a09_183.pub4 |editor-last=Wiley-VCH Verlag GmbH & Co. KGaA |access-date=2023-05-05 |place=Weinheim, Germany |publisher=Wiley-VCH Verlag GmbH & Co. KGaA |language=en |doi=10.1002/14356007.a09_183.pub4 |isbn=978-3-527-30673-2 |last2=Kolb |first2=Dieter M. |last3=Engelmann |first3=Gerald E. |last4=Ziegler |first4=Jörg C.|url-access=subscription }}</ref>

Both galvanic and electrolytic cells can be thought of as having two [[half-cell]]s: consisting of separate [[Redox|oxidation and reduction reactions]].

When one or more electrochemical cells are connected in parallel or series they make a [[Battery (electricity)|battery]]. A [[primary battery]] consists of single-use galvanic cells. [[Rechargeable batteries]] are built from [[#Secondary cells|secondary cells]] that use [[reversible reactions]] and can operate as galvanic cells (while providing energy) or electrolytic cells (while charging).

== Types of electrochemical cells ==

=== Galvanic cell ===
{{main|Galvanic cell}}
A galvanic cell (voltaic cell), named after [[Luigi Galvani]] ([[Alessandro Volta]]), is an electrochemical cell that generates electrical energy from spontaneous [[redox]] reactions.<ref name=":32">''Chemistry,'' Rice University, 2015. [Online]. Available: https://web.ung.edu/media/Chemistry2/Chemistry-LR.pdf</ref>
[[File:Galvanic cell with no cation flow.svg|thumb|Galvanic cell with no cation flow]]
A wire connects two different [[metal]]s (e.g. [[zinc]] and [[copper]]). Each metal is in a separate solution; often the [[Aqueous solution|aqueous]] [[Sulfate|sulphate]] or [[nitrate]] forms of the metal, however more generally metal salts and water which conduct [[Electric current|current]].<ref>{{Cite book |last=Ahmad |first=Dr. Zaki |title=Principles of corrosion engineering and corrosion control. |date=2013 |publisher=Butterworth-Heinemann |isbn=978-0-08-097134-6 |oclc=857524149}}</ref> A [[salt bridge]] or porous membrane connects the two solutions, keeping electric neutrality and the avoidance of charge accumulation. The metal's differences in oxidation/reduction potential drive the reaction until [[Chemical equilibrium|equilibrium]].<ref name=":23"/>

Key features:
* [[Spontaneous process|spontaneous reaction]]
* generates electric current
* current flows through a wire, and [[ion]]s flow through a salt bridge
* [[anode]] (negative), [[cathode]] (positive)

==== Half cells ====
Galvanic cells consists of two half-cells.  Each half-cell consists of an [[electrode]] and an [[electrolyte]]<ref>{{cite book |last=Chieh |first=Chung |date=2015-08-26 |time=11:04:13PM |title=Galvanic Cells |url=https://chem.libretexts.org/Bookshelves/Analytical_Chemistry/Supplemental_Modules_(Analytical_Chemistry)/Electrochemistry/Basics_of_Electrochemistry/Electrochemistry/Galvanic_Cells |location= |publisher=University of Waterloo |page=1 |isbn= |access-date=2026-01-09}} {{ISBN missing|date=January 2026}}</ref>

The chemical reactions in the cell involve the electrolyte, electrodes, and/or an external substance ([[fuel cell]]s may use [[Hydrogen|hydrogen gas]] as a [[Reagent|reactant]]). In a full electrochemical cell, species from one half-cell lose electrons ([[Redox|oxidation]]) to their electrode while species from the other half-cell gain electrons ([[Redox|reduction]]) from their electrode.<ref name=":32" />

A ''[[salt bridge]]'' (e.g., filter paper soaked in KNO3<sub>,</sub> NaCl, or some other electrolyte) is used to ionically connect two half-cells with different electrolytes, but it prevents the solutions from mixing and unwanted side reactions. An alternative to a salt bridge is to allow direct contact (and mixing) between the two half-cells, for example in simple [[electrolysis of water]].{{citation needed|date=December 2024}}

As electrons flow from one half-cell to the other through an external [[Electrical network|circuit]], a difference in charge is established. If no ionic contact were provided, this charge difference would quickly prevent the further flow of electrons. A salt bridge allows the flow of negative or positive ions to maintain a steady-state charge distribution between the oxidation and reduction vessels, while keeping the contents otherwise separate. Other devices for achieving separation of solutions are porous pots and gelled solutions. A porous pot is used in the [[Bunsen cell]].<ref>{{cite web |last1=Jensen |first1=William |title=The Grove and Bunsen Cells |url=https://homepages.uc.edu/~jensenwb/museum-notes/23.%20The%20Grove-Bunsen%20Cell.pdf |publisher=University of Cincinnati |access-date=27 July 2025 |date=December 2013}}</ref><ref>{{cite book |editor1-last=Rennie |editor1-first=Richard |editor2-last=Law |editor2-first=Jonathan |title=A Dictionary of Physics |date=2019 |publisher=Oxford University Press |isbn=9780198821472 |edition=8 |url=https://www.oxfordreference.com/display/10.1093/acref/9780198821472.001.0001/acref-9780198821472-e-355 |access-date=27 July 2025}}</ref>

==== Equilibrium reaction ====
Each half-cell has a characteristic voltage (depending on the metal and its characteristic reduction potential). Each reaction is undergoing an [[Chemical equilibrium|equilibrium]] reaction between different [[oxidation states]] of the ions: when equilibrium is reached, the cell cannot provide further [[voltage]]. In the half-cell performing oxidation, the closer the equilibrium lies to the ion/atom with the more positive oxidation state the more potential this reaction will provide.<ref name=":23" /> Likewise, in the reduction reaction, the closer the equilibrium lies to the ion/atom with the more ''negative'' oxidation state the higher the potential.{{citation needed|date=December 2024}}

==== Cell potential ====
The cell potential can be predicted through the use of [[electrode potential]]s (the voltages of each half-cell). These half-cell potentials are defined relative to the assignment of 0 [[volt]]s to the [[standard hydrogen electrode]] (SHE). (See [[table of standard electrode potentials]]). The difference in voltage between electrode potentials gives a prediction for the potential measured. When calculating the difference in voltage, one must first rewrite the half-cell reaction equations to obtain a balanced oxidation-reduction equation.<ref>{{Cite web |date=2016-07-12 |title=20.4: Cell Voltage |url=https://chem.libretexts.org/Courses/Heartland_Community_College/HCC%3A_Chem_162/20%3A_Electrochemistry/20.4%3A_Cell_Voltage |access-date=2025-08-31 |website=Chemistry LibreTexts |language=en}}</ref>

# Reverse the reduction reaction with the smallest potential (to create an oxidation reaction/overall positive cell potential)
# Half-reactions must be multiplied by integers to achieve electron balance.

Cell potentials have a possible range of roughly zero to 6 volts. Cells using water-based electrolytes are usually limited to cell potentials less than about 2.5 volts due to high reactivity of the powerful oxidizing and reducing agents with water which is needed to produce a higher voltage. Higher cell potentials are possible with cells using other [[solvent]]s instead of water. For instance, [[Lithium battery|lithium cells]] with a voltage of 3 volts are commonly available.<ref>{{cite journal |last1=Cervera |first1=J |last2=Levin |first2=M |last3=Mafe |first3=S |title=Top-down perspectives on cell membrane potential and protein transcription |journal=Sci Rep |date=2025-12-10 |volume=16 |issue=1 |pages=1996 |doi=10.1038/s41598-025-31696-6  |doi-access=free|pmid=41372487 |pmc=12808214}}</ref>

The cell potential depends on the [[concentration]] of the reactants, as well as their type. As the cell is discharged, the concentration of the reactants decreases and the cell potential also decreases.<ref>{{Cite web |date=2013-10-02 |title=Nernst Equation |url=https://chem.libretexts.org/Bookshelves/Analytical_Chemistry/Supplemental_Modules_(Analytical_Chemistry)/Electrochemistry/Nernst_Equation |access-date=2026-02-08 |website=Chemistry LibreTexts |language=en}}</ref>

=== Electrolytic cell ===
{{main|Electrolytic cell}}
An electrolytic cell is an electrochemical cell in which applied electrical energy drives a non-spontaneous [[redox]] reaction.<ref>{{Citation |last=Brett |first=C.M.A. |title=Standard Electrode Potentials and Application to Characterization of Corrosion Phenomena |date=2018 |url=http://dx.doi.org/10.1016/b978-0-12-409547-2.13389-x |encyclopedia=Encyclopedia of Interfacial Chemistry |pages=511–516 |access-date=2023-04-18 |publisher=Elsevier|doi=10.1016/b978-0-12-409547-2.13389-x |isbn=9780128098943 |url-access=subscription }}</ref>
[[File:Electrolytic Cell Diagram.jpg|thumb|A modern electrolytic cell consisting of two half reactions, two electrodes, a salt bridge, voltmeter, and a battery.]]
They are often used to decompose chemical compounds, in a process called [[electrolysis]]. (The Greek word "[[wikt:λύσις#Ancient Greek|lysis]]" (λύσις) means "loosing" or "setting free".)<ref>{{Cite web |title=Definition of LYSIS |url=https://www.merriam-webster.com/dictionary/lysis |access-date=2025-07-03 |website=www.merriam-webster.com |language=en}}</ref>

Important examples of electrolysis are the decomposition of water into [[hydrogen]] and [[oxygen]], and of [[bauxite]] into [[aluminium]] and other chemicals. [[Electroplating]] (e.g. of Copper, [[Silver]], [[Nickel]] or [[Chromium]]) is done using an electrolytic cell. Electrolysis is a technique that uses a [[Direct current|direct electric current]] (DC).{{citation needed|date=December 2024}}

The components of an electrolytic cell are:{{citation needed|date=December 2024}}
* an electrolyte: usually a solution of water or other solvents in which ions are dissolved. Molten salts such as [[sodium chloride]] are also electrolytes.
* two electrodes (a cathode and an anode) which are [[Electrical connector|electrical terminals]] consisting of a suitable substance at which oxidation or reduction can take place, and maintained at two different [[electric potential]]s.

When driven by an external [[voltage]] (potential difference) applied to the electrodes, the ions in the electrolyte are attracted to the electrode with the opposite potential, where charge-transferring (also called [[Faradaic current|faradaic]] or redox) reactions can take place. Only with a sufficient external voltage can an electrolytic cell decompose a normally stable, or [[Chemically inert|inert]] chemical compound in the solution. Thus the electrical energy provided produces a chemical reaction which would not occur spontaneously otherwise.<ref>{{Cite encyclopedia |title=Electrolytic cell — Encyclopædia Britannica |encyclopedia=Encyclopedia Britannica |url=https://www.britannica.com/technology/electrolytic-cell |access-date=2026-05-12}}</ref> Key features:
* non-spontaneous reaction
* generates current
* current flows through a wire, and ions flow through salt bridge
* anode (positive), cathode (negative)

=== Primary cell ===
[[File:Batteries comparison 4,5 D C AA AAA AAAA A23 9V CR2032 LR44 matchstick-1.jpeg|thumb|A variety of standard sizes of primary cells.  From left: 4.5V multicell battery, D, C, AA, AAA, AAAA, A23, 9V multicell battery, LR44 ''(top)'', CR2032 ''(bottom)''.]]
{{main|Primary cell}}

A primary cell produces current by irreversible chemical reactions (ex. small disposable batteries) and is not rechargeable.{{citation needed|date=December 2024}}

They are used for their portability, low cost, and short lifetime.{{citation needed|date=December 2024}}

Primary cells are made in a range of standard sizes to power small household appliances such as [[flashlight]]s and portable radios.

As chemical reactions proceed in a primary cell, the battery uses up the chemicals that generate the power; when they are gone, the battery stops producing electricity.<ref>{{Cite web |last=mischa |date=2016-02-25 |title=How a battery works |url=https://www.science.org.au/curious/technology-future/batteries |access-date=2025-12-30 |website=Curious |language=en}}</ref>
[[File:Diagram of a primary cell (battery).jpg|thumb|Circuit diagram of a primary cell showing difference in cell potential, and flow of electrons through a resistor.]]

Primary batteries make up about 90% of the $50 billion battery market, but secondary batteries have been gaining market share.  About 15 billion primary batteries are thrown away worldwide every year,<ref>{{Cite web |last=Communications |first=Cactus |title=What if we could recycle the energy remaining in discarded batteries? Scientists now know how |url=https://techxplore.com/news/2022-08-recycle-energy-discarded-batteries-scientists.html |access-date=2023-04-18 |website=techxplore.com |language=en}}</ref> virtually all ending up in landfills.  Due to the toxic [[Heavy metal (chemistry)|heavy metals]] and strong acids or alkalis they contain, batteries are [[hazardous waste]].  Most municipalities classify them as such and require separate disposal. The energy needed to manufacture a battery is about 50 times greater than the energy it contains.<ref name="Hill2">{{cite book |last=Hill |first=Marquita K. |url=https://archive.org/details/understandingenv0000hill |title=Understanding Environmental Pollution: A Primer |date=2004 |publisher=[[Cambridge University Press]] |isbn=978-0-521-82024-0 |page=274 |quote=Manufacturing a disposable battery takes about 50 times more energy than the battery provides when used. |url-access=registration}}</ref><ref name="Watts2">{{cite book |last=Watts |first=John |url=https://books.google.com/books?id=KFTvvwOOi64C&q=battery+energy+%2250+times%22&pg=PA63 |title=Gcse Edexcel Science |date=2006 |publisher=[[Letts and Lonsdale]] |isbn=978-1-905129-63-8 |pages=63}}</ref><ref name="Wastebusters2">{{cite book |last=Wastebusters Ltd. |url=https://books.google.com/books?id=LcX9AQAAQBAJ&q=battery+energy+%2250+times%22&pg=PA96 |title=The Green Office Manual: A Guide to Responsible Practice |date=2013 |publisher=[[Routledge]] |isbn=978-1-134-19798-9 |pages=96}}</ref><ref name="Danaher2">{{cite book |last1=Danaher |first1=Kevin |url=https://books.google.com/books?id=JQdZCwAAQBAJ&q=battery+energy+%2250+times%22&pg=PA199 |title=Building the Green Economy: Success Stories from the Grassroots |last2=Biggs |first2=Shannon |last3=Mark |first3=Jason |date=2016 |publisher=[[Routledge]] |isbn=978-1-317-26292-3 |pages=199 |author1-link=Kevin Danaher (activist)}}</ref> Due to their high pollutant content compared to their small energy content, the primary battery is considered a wasteful, environmentally unfriendly technology.  Mainly due to the increasing sales of [[wireless device]]s and [[Cordless|cordless tools]], which cannot be economically powered by primary batteries and come with integral rechargeable batteries, the secondary battery industry has high growth and has slowly been replacing the primary battery in high end products.

=== Secondary cell ===
[[File:Photo-CarBattery.jpg|thumb|Lead acid car battery (secondary cell)]]
[[File:Secondary Cell Diagram.svg|thumb|Circuit diagram of a secondary cell showing difference in cell potential, and flow of electrons through a resistor.]]
{{main|Rechargeable battery}}
A secondary cell produces current by reversible chemical reactions (ex. [[Lead–acid battery|lead-acid battery]] car battery) and is [[Rechargeable battery|rechargeable]].{{citation needed|date=December 2024}}

Lead-acid batteries are used in an automobile to start an engine and to operate the car's electrical accessories when the engine is not running. The alternator, once the car is running, recharges the battery.<ref>{{Cite web |date=2024-02-26 |title=The role of automotive lead acid battery |url=https://www.accord-power.com/news-media/blog/1761767556227534850.html |access-date=2025-12-01 |website=ACCORD |language=en}}</ref>

It can perform as a galvanic cell and an electrolytic cell. It is a convenient way to store electricity: when current flows one way, the levels of one or more chemicals build up (charging); while it is discharging, they reduce and the resulting electromotive force can do work.{{citation needed|date=December 2024}}

They are used for their high voltage, low costs, reliability, and long lifetime.{{citation needed|date=December 2024}}

=== Fuel cell ===
[[File:Solid oxide fuel cell protonic.svg|thumb|Scheme of a proton-conducting fuel cell]]
{{main|Fuel cell}}
A [[fuel cell]] is an electrochemical cell that reacts hydrogen fuel with oxygen or another oxidizing agent, to convert chemical energy to [[electricity]].<ref name=":0">{{Cite web |title=Fuel Cells |url=https://www.energy.gov/eere/fuelcells/fuel-cells |access-date=2025-02-02 |website=US Department of Energy |language=en}}</ref>

Fuel cells are different from [[Battery (electricity)|batteries]] in requiring a continuous source of fuel and oxygen (usually from air) to sustain the chemical reaction, whereas in a battery the chemical energy comes from chemicals already present in the battery.<ref name=":0" />

Fuel cells can produce electricity continuously for as long as fuel and [[oxygen]] are supplied.<ref name=":0" />

They are used for primary and backup power for commercial, industrial and residential buildings and in remote or inaccessible areas. They are also used to power [[fuel cell vehicle]]s, including [[forklift]]s, automobiles, buses, boats, motorcycles and submarines.{{citation needed|date=December 2024}}

Fuel cells are classified by the type of electrolyte they use and by the difference in startup time, which ranges from 1 second for [[proton-exchange membrane fuel cell]]s (PEM fuel cells, or PEMFC) to 10 minutes for [[solid oxide fuel cell]]s (SOFC).<ref>{{Cite web |title=Comparison of Fuel Cell Technologies |url=https://www.energy.gov/eere/fuelcells/comparison-fuel-cell-technologies |access-date=2025-05-06 |website=Energy.gov |language=en}}</ref><ref>{{Cite web |last=Vskills |first=Team |title=Principle and Working Of Fuel Cell |url=https://www.vskills.in/certification/tutorial/principle-and-working-of-fuel-cell/ |access-date=2025-05-06 |website=Tutorial |language=en-US}}</ref>

There are many types of fuel cells, but they all consist of:

;anode: At the anode a catalyst causes the fuel to undergo oxidation reactions that generate [[proton]]s (positively charged hydrogen ions) and electrons. The protons flow from the anode to the cathode through the electrolyte after the reaction. At the same time, electrons are drawn from the anode to the cathode through an external circuit, producing [[direct current]] electricity.{{citation needed|date=December 2024}}
;cathode: At the cathode, another catalyst causes hydrogen ions, electrons, and oxygen to react, forming water.{{citation needed|date=December 2024}}
;electrolyte: Allows positively charged hydrogen ions (protons) to move between the two sides of the fuel cell.{{citation needed|date=December 2024}}

A related technology are [[Flow battery|flow batteries]], in which the fuel can be regenerated by recharging. Individual fuel cells produce relatively small electrical potentials, about 0.7 volts, so cells are "stacked", or placed in series, to create sufficient voltage to meet an application's requirements.<ref>{{Cite journal |last1=Qi |first1=Zhaoxiang |last2=Koenig |first2=Gary M. |date=2017-07-01 |title=Review Article: Flow battery systems with solid electroactive materials |journal=Journal of Vacuum Science & Technology B |volume=35 |issue=4 |pages=040801 |doi=10.1116/1.4983210 |bibcode=2017JVSTB..35d0801Q |issn=2166-2746|doi-access=free }}</ref> In addition to electricity, fuel cells produce water, heat and, depending on the fuel source, very small amounts of [[nitrogen dioxide]] and other emissions. The [[Efficient energy use|energy efficiency]] of a fuel cell is generally between 40 and 60%; however, if waste heat is captured in a [[cogeneration]] scheme, efficiencies up to 85% can be obtained.{{citation needed|date=December 2024}}

In 2022, the global fuel cell market was estimated to be $6.3 billion, and is expected to increase by 19.9% by 2030.<ref>{{Cite web |title=Fuel Cell Market Size, Share & Trends Analysis Report, 2030 |url=https://www.grandviewresearch.com/industry-analysis/fuel-cell-market |access-date=2023-04-18 |website=www.grandviewresearch.com |language=en}}</ref> Many countries are attempting to enter the market by setting renewable energy [[Gigawatt|GW]] goals.<ref>{{Cite web |title=Renewable energy targets |url=https://energy.ec.europa.eu/topics/renewable-energy/renewable-energy-directive-targets-and-rules/renewable-energy-targets_en |access-date=2023-04-22 |website=energy.ec.europa.eu |language=en}}</ref>

== See also ==
{{div col begin|colwidth=15em}}
* [[Activity (chemistry)]]
* [[Cell notation]]
* [[Electrochemical potential]]
* [[Electrochemical engineering]]
* [[Battery (electricity)]]
* [[Rechargeable battery]]
* [[Fuel cell]]
* [[Flow battery]]
* [[Scanning flow cell]]
{{div col end}}
{{portal|energy}}

==References==
{{reflist|25em}}

== External links ==
* {{Britannica|183162|Electrolytic cell (device)}}
{{right| {{commons category|Electrochemical cells}} }}{{Galvanic cells}}

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