Editing Electrolytic cell

{{Short description|Cell that uses electrical energy to drive a non-spontaneous redox reaction}}
{{original research|date=November 2017}}
{{More citations needed|date=November 2023}}[[Image:Electrolyser 1884.png|thumb|right|190px| Nineteenth-century electrolytic cell for producing [[oxyhydrogen]]]]
An '''electrolytic cell''' is an [[electrochemical cell]] that utilizes an external source of [[electrical energy]] to force a [[chemical reaction]] that would otherwise not occur.<ref name=":0">{{Cite book |url=https://archive.org/details/workingmethodapp0000unse/ |title=A Working Method Approach for Introductory Physical Chemistry Calculations |publisher=[[Royal Society of Chemistry]] |year=1997 |isbn=0-85404-553-8 |veditors=Murphy B, Murphy C, Hathaway B |location=Cambridge, United Kingdom |language=en |chapter=Electrochemistry I: Galvanic Cells |url-access=registration |via=Internet Archive}}</ref>{{Rp|page=|pages=64, 89}}<ref>{{cite book |last=Harris |first=Daniel C. |url=https://openlibrary.org/books/OL25176031M/Quantitative_chemical_analysis |title=Quantitative Chemical Analysis |publisher=[[W. H. Freeman and Company]] |year=2010 |isbn=978-1-4292-1815-3 |edition=8th |publication-place=New York |language=en |oclc=540161465 |url-access=registration |via=Open Library}}</ref>{{Rp|page=GL7}} The external energy source is a [[voltage]] applied between the cell's two [[electrode]]s; an [[anode]] (positively charged electrode) and a [[cathode]] (negatively charged electrode), which are immersed in an [[electrolyte]] solution.<ref name=":0" />{{Rp|page=89}}<ref name=Skoog/>{{Page needed|date=November 2023}} This is in contrast to a [[galvanic cell]], which itself is a source of electrical energy and the foundation of a [[Electric battery|battery]].<ref name=":0" />{{Rp|page=64}} The net reaction taking place in an electrolytic cell is a non-spontaneous reaction (reverse of a [[Spontaneous process|spontaneous reaction]]), i.e., the [[Gibbs free energy]] is +ve, while the net reaction taking place in a galvanic cell is a [[Spontaneous process|spontaneous reaction]], i.e., the Gibbs free energy is - ve.<ref name="Skoog">{{cite book |last1=Skoog |first1=Douglas A. |title=Fundamentals of analytical chemistry |last2=West |first2=Donald M. |last3=Holler |first3=F. James |last4=Crouch |first4=Stanley R. |publisher=Brooks/Cole, Cengage Learning <!-- from WorldCat --> |year=2014 |isbn=978-0-495-55828-6 |publication-place=Belmont, CA |oclc=824171785}}</ref>{{Page needed|date=November 2023}}

==Principles==
In an electrolytic cell, a [[Electric current|current]] passes through the cell by an external [[voltage]], causing a non-spontaneous chemical reaction to proceed. In a galvanic cell, the progress of a spontaneous chemical reaction causes an electric current to flow. An [[Chemical equilibrium|equilibrium]] electrochemical cell exists in the state between an electrolytic cell and a galvanic cell. The tendency of a spontaneous reaction to push a current through the external circuit is exactly balanced by a [[counter-electromotive force]] so that no current flows. If this counter-electromotive force is increased, the cell becomes an electrolytic cell, and if it is decreased, the cell becomes a galvanic cell.<ref>{{cite book |last=Mortimer |first=Robert G. |url=https://archive.org/details/physical-chemistry-3rd-edition-2008/ |title=Physical chemistry |publisher=Academic Press/Elsevier |year=2008 |isbn=978-0-12-370617-1 |edition=3rd |publication-place=Amsterdam |language=en |oclc=196313033 |via=Internet Archive}}</ref>{{Rp|page=354}}

An electrolytic cell has three components: an [[electrolyte]] and two electrodes (a [[cathode]] and an [[anode]]). The [[electrolyte]] is usually a [[Solution (chemistry)|solution]] of [[water]] or other [[solvent]]s in which [[ion]]s are dissolved. [[Molten salt]]s such as [[sodium chloride]] can also function as electrolytes. When driven by an external voltage applied to the electrodes, the ions in the electrolyte are attracted to an electrode with the opposite [[Electric charge|charge]], where charge-transferring (also called faradaic or [[redox]]) reactions can take place. Only with an external [[electrical potential]] (i.e., voltage) of correct polarity and sufficient magnitude can an electrolytic cell decompose a normally stable, or [[Chemically inert|inert]] chemical compound in the solution. The electrical energy provided can produce a chemical reaction that would otherwise not occur spontaneously.

[[Michael Faraday]] defined the cathode of a cell as the electrode to which cations (positively charged ions, such as silver ions Ag{{Su|p=+}}) flow within the cell, to be reduced by reacting with electrons (negatively charged) from that electrode. Likewise, he defined the anode as the electrode to which anions (negatively charged ions, like chloride ions Cl{{Su|p=−}}) flow within the cell, to be [[Redox|oxidized]] by depositing electrons on the electrode. To an external wire connected to the electrodes of a galvanic cell (or battery), forming an electric circuit, the cathode is positive and the anode is negative. Thus positive electric current flows from the cathode to the anode through the external circuit in the case of a galvanic cell.

==Applications==
[[File:ElectrolyticReduction.ogg|thumb|right|A video describing the process of electrolytic reduction as used on [[Captain Kidd's Cannon]] at [[The Children's Museum of Indianapolis]]]]
Electrolytic cells are often used to decompose chemical compounds, in a process called [[electrolysis]]—with ''electro'' meaning ''electricity''<ref>{{Cite dictionary |title=electro&mdash; |encyclopedia=[[Collins English Dictionary]] (online) |url=https://www.collinsdictionary.com/dictionary/english/electro |access-date=November 4, 2023}}</ref> and the Greek word ''[[lysis]]'' means ''to break up''. Important examples of electrolysis are the decomposition of [[water]] into [[hydrogen]] and [[oxygen]], and [[bauxite]] into [[aluminium|aluminum]] 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 electric current (DC).

Commercially, electrolytic cells are used in the electrorefining and [[electrowinning]] of several non-ferrous metals. Most high-purity [[aluminium|aluminum]], [[copper]], [[zinc]], and [[lead]] are produced industrially in electrolytic cells.

As already noted, water, particularly when ions are added (saltwater or acidic water), can be ''electrolyzed'' (subjected to electrolysis). When driven by an external source of voltage, hydrogen (H{{Su|p=+}}) ions flow to the cathode to combine with electrons to produce hydrogen gas in a reduction reaction. Likewise, hydroxide (OH{{Su|p=−}}) ions flow to the anode to release electrons and a hydrogen (H{{Su|p=+}}) ion to produce oxygen gas in an oxidation reaction.

In molten sodium chloride (NaCl), when a current is passed through the salt the anode oxidizes chloride ions (Cl{{Su|p=−}}) to chlorine gas, it releases electrons to the anode. Likewise, the cathode reduces sodium ions (Na{{Su|p=+}}), which accepts electrons from the cathode and deposits them on the cathode as sodium metal.

Sodium chloride that has been dissolved in water can also be electrolyzed. The anode oxidizes the chloride ions (Cl{{Su|p=−}}), and produces chlorine (Cl<sub>2</sub>) gas, and, depending on the pH of the solution, can produce [[Hypochlorous acid]]. However, at the cathode, instead of sodium ions being reduced to sodium metal, water molecules are reduced to hydroxide ions (OH{{Su|p=−}}) and hydrogen gas (H<sub>2</sub>). The overall result of the electrolysis is the production of [[chlorine]] gas at the anode, aqueous hypochlorous acid as the [[anolyte]], hydrogen gas at the cathode, and aqueous [[sodium hydroxide]] (NaOH) as the catholyte. Industrially, this is known as the [[chloralkali process]].

==See also==
* [[Electrolysis]]
* [[Electrochemical cell]]
* [[Galvanic cell]]

== References ==
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