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{{short description|Chemical species that donates an electron to another species in a redox reaction}}
In [[chemistry]], a '''reducing agent''' (also known as a '''reductant''', '''reducer''', or '''[[electron donor]]''') is a [[chemical species]] that "donates" an [[electron]] to an {{em|electron recipient}} (called the {{em|[[oxidizing agent]]}}, {{em|oxidant}}, {{em|oxidizer}}, or {{em|[[electron acceptor]]}}).

Examples of substances that are common reducing agents include [[hydrogen]], [[carbon monoxide]], the [[alkali metal]]s, [[formic acid]],<ref>{{cite journal | doi = 10.1016/j.watres.2005.05.012 | title = Use of formic acid as reducing agent for application in catalytic reduction of nitrate in water | date = 2005 | last1 = Garron | first1 = Anthony | last2 = Epron | first2 = Florence | journal = Water Research | volume = 39 | issue = 13 | pages = 3073–3081 | pmid = 15982701 | bibcode = 2005WatRe..39.3073G }}</ref> [[oxalic acid]],<ref>{{cite web | url = https://chemed.chem.purdue.edu/genchem/topicreview/bp/ch19/oxred_3.php | title = Oxidizing and Reducing Agents | publisher = Purdue University }}</ref> and [[sulfite]] compounds.

In their pre-reaction states, reducers have extra electrons (that is, they are by themselves reduced) and oxidizers lack electrons (that is, they are by themselves oxidized). This is commonly expressed in terms of their oxidation states. An agent's [[oxidation state]] describes its degree of loss of electrons, where the higher the oxidation state then the fewer electrons it has. So initially, prior to the reaction, a reducing agent is typically in one of its lower possible oxidation states; its oxidation state increases during the reaction while that of the oxidizer decreases.
Thus in a [[redox]] reaction, the agent whose oxidation state increases, that "loses/[[Electron donor|donates]] electrons", that "is oxidized", and that "reduces" is called the {{em|reducer}} or {{em|reducing agent}}, while the agent whose oxidation state decreases, that "gains/[[Electron acceptor|accepts]]/receives electrons", that "is reduced", and that "oxidizes" is called the {{em|oxidizer}} or {{em|oxidizing agent}}.
<!-- A reducing agent is thus oxidized by an oxidizer when it loses electrons that are gained by this oxidizing agent; this oxidizing agent is itself simultaneously reduced by the reducer. -->

For example, consider the overall reaction for aerobic [[cellular respiration]]:

:{{chem2|C6H12O6(s) + 6O2(g) -> 6CO2(g) + 6H2O(l)}}

The [[oxygen]] ({{chem2|O2}}) is being reduced, so it is the oxidizing agent. The [[glucose]] ({{chem2|C6H12O6}}) is being oxidized, so it is the reducing agent. 

==Characteristics==

Consider the following reaction:

:2 {{chem|[Fe([[Cyanide|CN]]){{sub|6}}]{{sup|4−}}}} + {{chem|link=Chlorine|Cl|2}} → 2 {{chem|[Fe(CN){{sub|6}}]{{sup|3−}}}} + 2 {{chem|link=Chloride|Cl|-}}

The reducing agent in this reaction is [[ferrocyanide]] ({{chem|[Fe(CN){{sub|6}}]{{sup|4−}}}}). It donates an electron, becoming oxidized to [[ferricyanide]] ({{chem|[Fe(CN){{sub|6}}]{{sup|3−}}}}). Simultaneously, that electron is received by the oxidizer [[chlorine]] ({{chem|Cl|2}}), which is reduced to [[chloride]] ({{chem|Cl|-}}).

Strong reducing agents easily lose (or donate) electrons. An atom with a relatively large atomic radius tends to be a better reductant. In such species, the distance from the nucleus to the [[valence electron]]s is so long that these electrons are not strongly attracted. These elements tend to be strong reducing agents. Good reducing agents tend to consist of atoms with a low [[electronegativity]], which is the ability of an atom or molecule to attract bonding electrons, and species with relatively small [[Ionization energy|ionization energies]] serve as good reducing agents too.{{cn|date=February 2024}}

The measure of a material's ability to reduce is known as its [[reduction potential]].<ref name=sfe>{{cite web|url=https://www.eesemi.com/ox_potential.htm|title=Electrode Reduction and Oxidation Potential Values|website= www.EESemi.com|access-date=12 July 2021}}</ref> The table below shows a few reduction potentials, which can be changed to oxidation potentials by reversing the sign. Reducing agents can be ranked by increasing strength by ranking their reduction potentials. Reducers donate electrons to (that is, "reduce") [[oxidizing agent]]s, which are said to "be reduced by" the reducer. The reducing agent is stronger when it has a more negative reduction potential and weaker when it has a more positive reduction potential. The more positive the reduction potential the greater the species' affinity for electrons and tendency to be reduced (that is, to receive electrons). The following table provides the reduction potentials of the indicated reducing agent at 25&nbsp;°C. For example, among [[sodium]] (Na), [[chromium]] (Cr), [[cuprous]] (Cu<sup>+</sup>) and [[chloride]] (Cl<sup>−</sup>), it is Na that is the strongest reducing agent while Cl<sup>−</sup> is the weakest; said differently, Na<sup>+</sup> is the weakest oxidizing agent in this list while Cl is the strongest.{{cn|date=February 2024}}

{{center|{{ReductionPotentialTable}} }}

Common reducing agents include metals potassium, calcium, barium, sodium and magnesium, and also compounds that contain the [[hydride]] H<sup>−</sup> ion, those being [[NaH]], [[LiH]],<ref>{{cite journal|vauthors = Aufray M, Menuel S, Fort Y, Eschbach J, Rouxel D, Vincent B|title = New Synthesis of Nanosized Niobium Oxides and Lithium Niobate Particles and Their Characterization by XPS Analysis|journal = Journal of Nanoscience and Nanotechnology|volume = 9|issue = 8|pages = 4780–4789|year = 2009|doi = 10.1166/jnn.2009.1087|pmid = 19928149|url = http://maelenn.aufray.free.fr/pdf/nano.pdf|access-date = 2019-09-24|archive-date = 2020-07-29|archive-url = https://web.archive.org/web/20200729225854/http://maelenn.aufray.free.fr/pdf/nano.pdf|url-status = dead}}</ref> [[Lithium aluminium hydride|LiAlH<sub>4</sub>]] and [[Calcium hydride|CaH<sub>2</sub>]].

Some elements and compounds can be both reducing or [[oxidizing agent]]s. [[Hydrogen gas]] is a reducing agent when it reacts with non-metals and an oxidizing agent when it reacts with metals.

:2 Li<sub>(s)</sub> + H<sub>2(g)</sub> → 2 LiH<sub>(s)</sub>{{efn|[[Half reaction]]s: 2 Li<sup>0</sup><sub>(s)</sub> → 2 Li<sup>+</sup><sub>(s)</sub> + 2 e<sup>−</sup> ::::: H<sub>2</sub><sup>0</sup><sub>(g)</sub> + 2 e<sup>−</sup> → 2 H<sup>−</sup><sub>(g)</sub>}}
Hydrogen (whose reduction potential is 0.0) acts as an oxidizing agent because it accepts an electron donation from the reducing agent [[lithium]] (whose reduction potential is -3.04), which causes Li to be oxidized and hydrogen to be reduced.

:H<sub>2(g)</sub> + F<sub>2(g)</sub> → 2 HF<sub>(g)</sub>{{efn|[[Half reaction]]s: H<sub>2</sub><sup>0</sup><sub>(g)</sub> → 2 H<sup>+</sup><sub>(g)</sub> + 2 e<sup>−</sup> :::::
F<sub>2</sub><sup>0</sup><sub>(g)</sub> + 2 e<sup>−</sup> → 2 F<sup>−</sup><sub>(g)</sub>}}
Hydrogen acts as a reducing agent because it donates its electrons to [[fluorine]], which allows fluorine to be reduced.

==Importance==

Reducing agents and oxidizing agents are the ones responsible for [[corrosion]], which is the "degradation of metals as a result of electrochemical activity".<ref name=sfe/> Corrosion requires an [[anode]] and [[cathode]] to take place. The anode is an element that loses electrons (reducing agent), thus oxidation always occurs in the anode, and the cathode is an element that gains electrons (oxidizing agent), thus reduction always occurs in the cathode. Corrosion occurs whenever there's a difference in oxidation potential. When this is present, the anode metal begins deteriorating, given there is an electrical connection and the presence of an [[electrolyte]].{{cn|date=February 2024}}

==Examples of redox reaction==
{{redox_example.svg}}

Historically, reduction referred to the removal of oxygen from a compound, hence the name 'reduction'.<ref name=Britannica>{{cite encyclopedia|title=oxidation-reduction reaction|last=Olson|first=Maynard V.|author-link=Maynard V. Olson|encyclopedia=Britannica|access-date=3 May 2022|url=https://www.britannica.com/science/oxidation-reduction-reaction/Historical-origins-of-the-redox-concept|quote=In his [[Traité Élémentaire de Chimie|Traité élémentaire de chimie]], [[Antoine Lavoisier|he]] clearly established that combustion consists of a chemical combination between oxygen from the atmosphere and combustible matter [...]. By the end of the century, his ideas were widely accepted and had been successfully applied to the more complex processes of respiration and photosynthesis. Reactions in which oxygen was consumed were classified as oxidations, while those in which oxygen was lost were termed reductions.}}</ref> An example of this phenomenon occurred during the [[Great Oxidation Event]], in which biologically−produced molecular oxygen ([[Oxygen#Allotropes|dioxygen]] ({{chem2|O2}}), an oxidizer and electron recipient) was added to [[Great Oxidation Event#The early atmosphere|the early Earth's atmosphere]], which was originally a weakly [[reducing atmosphere]] containing reducing gases like [[methane]] ({{chem2|CH4}}) and [[carbon monoxide]] ({{chem2|CO}}) (along with other electron donors)<ref name="Kasting 2014 pp. 157–175">{{cite book|last=Kasting|first=J.F.|title=Treatise on Geochemistry|chapter=Modeling the Archean Atmosphere and Climate|publisher=Elsevier|year=2014|doi=10.1016/b978-0-08-095975-7.01306-1|pages=157–175|isbn=9780080983004 }}</ref> and practically no oxygen because any that was produced would [[Redox reaction|react]] with these or other reducers (particularly with [[iron]] [[Solubility|dissolved]] in [[sea water]]), resulting in their {{em|removal}}.
By using water as a reducing agent, aquatic [[Photosynthesis|photosynthesizing]] [[cyanobacteria]] produced this molecular oxygen as a waste product.<ref name="Buick, R. 2008">{{cite journal|last=Buick|first=Roger|date=August 27, 2008|title=When did oxygenic photosynthesis evolve?|journal=[[Philosophical Transactions of the Royal Society B]]|volume=363|issue=1504|pages=2731–2743|doi=10.1098/rstb.2008.0041|issn=0962-8436|pmc=2606769|pmid=18468984|bibcode=2008RSPTB.363.2731B }}</ref> This {{chem2|O2}} initially oxidized the ocean's dissolved [[ferrous]] [[iron]] (Fe(II) − meaning iron in its +2 oxidation state) to form [[insoluble]] [[ferric]] [[iron oxide]]s such as [[Iron(III) oxide]] (Fe(II) lost an electron to the oxidizer and became Fe(III) − meaning iron in its +3 oxidation state) that precipitated down to the ocean floor to form [[banded iron formation]]s, thereby removing the oxygen (and the iron). The rate of production of oxygen eventually exceeded the availability of reducing materials that removed oxygen, which ultimately led [[Earth]] to gain a strongly oxidizing atmosphere containing abundant oxygen (like the [[Atmosphere of Earth|modern atmosphere]]).<ref>{{cite book|title=Sustaining Life on Planet Earth: Metalloenzymes Mastering Dioxygen and Other Chewy Gases|last1=Sosa Torres|first1=Martha E.|last2=Saucedo-Vázquez|first2=Juan P.|last3=Kroneck|first3=Peter M.H.|publisher=Springer|year=2015|isbn=978-3-319-12414-8|editor1=Kroneck, Peter M.H.|editor2=Sosa Torres, Martha E.|series=Metal Ions in Life Sciences volume 15|pages=1–12|chapter=Chapter 1, Section 2: The rise of dioxygen in the atmosphere|volume=15|doi=10.1007/978-3-319-12415-5_1|pmid=25707464}}</ref> The modern sense of donating electrons is a generalization of this idea, acknowledging that other components can play a similar chemical role to oxygen.

The formation of [[iron(III) oxide]];
:4Fe + 3O<sub>2</sub> → 4Fe<sup>3+</sup> + 6O<sup>2−</sup> → 2Fe<sub>2</sub>O<sub>3</sub>

In the above equation, the [[Iron]] (Fe) has an oxidation number of 0 before and 3+ after the reaction. For [[oxygen]] (O) the oxidation number began as 0 and decreased to 2−. These changes can be viewed as two "[[half-reaction]]s" that occur concurrently:

# Oxidation half reaction: Fe<sup>0</sup> → Fe<sup>3+</sup> + 3e<sup>−</sup>
# Reduction half reaction: O<sub>2</sub> + 4e<sup>−</sup> → 2 O<sup>2−</sup>

Iron (Fe) has been oxidized because the oxidation number increased. Iron is the reducing agent because it gave electrons to the oxygen (O<sub>2</sub>).
Oxygen (O<sub>2</sub>) has been reduced because the oxidation number has decreased and is the oxidizing agent because it took electrons from iron (Fe).

==Common reducing agents==
{{Refimprove section|date=October 2016}}
* [[Lithium aluminium hydride]] ([[Lithium|Li]] [[Aluminium|Al]] [[Hydrogen|H]]<sub>4</sub>), a very strong reducing agent
* [[Red-Al]] (NaAlH<sub>2</sub>(OCH<sub>2</sub>CH<sub>2</sub>OCH<sub>3</sub>)<sub>2</sub>), a safer and more stable alternative to lithium aluminum hydride
* [[Hydrogen]] without or with a suitable [[catalyst]]; e.g. a [[Lindlar catalyst]]
* [[Sodium amalgam]] ([[Sodium|Na]]([[Mercury (element)|Hg]]))
* Sodium-lead alloy ([[Sodium|Na]] + [[Lead|Pb]])
* [[Amalgam (chemistry)#Zinc amalgam|Zinc amalgam]] ([[Zinc|Zn]]([[Mercury (element)|Hg]])) (reagent for [[Clemmensen reduction]])
* [[Diborane]]
* [[Sodium borohydride]] ([[Sodium|Na]] [[Boron|B]]H<sub>4</sub>)
* [[Ferrous]] compounds that contain the [[Iron|Fe]]<sup>2+</sup> ion, such as [[iron(II) sulfate]]
* [[Stannous]] compounds that contain the [[Tin|Sn]]<sup>2+</sup> ion, such as [[tin(II) chloride]]
* [[Sulfur dioxide]] (sometimes also used as an [[oxidizing agent]]), [[Sulfite]] compounds
* [[Dithionate]]s, e.g. [[Sodium|Na]]<sub>2</sub>[[Sulfur|S]]<sub>2</sub>[[Oxygen|O]]<sub>6</sub>
* [[Thiosulfate]]s, e.g. Na<sub>2</sub>S<sub>2</sub>O<sub>3</sub> (mainly in analytical chemistry)<ref>{{Cite web |title=Cathodic Stripping Voltammetric Procedure for Determination of Some Inorganic Arsenic Species in Water, Soil and Ores Samples |url=https://www.researchgate.net/publication/286167941}}</ref>
* [[Iodide]]s, such as [[potassium iodide]] ([[Potassium|K]] [[Iodide|I]]) (mainly in [[analytical chemistry]])
* [[Hydrogen peroxide]] ({{chem|H|2|O|2}}){{snd}}mostly an oxidant but can occasionally act as a reducing agent, typically in analytical chemistry{{citation needed|date=July 2024}}
* [[Hydrazine]] ([[Wolff–Kishner reduction]])
* [[Diisobutylaluminium hydride]] (DIBAL-H)
* [[Oxalic acid]] ({{chem|C|2|H|2|O|4}})
* [[Formic acid]] (HCOOH)
* [[Ascorbic acid]]  (C<sub>6</sub>H<sub>8</sub>O<sub>6</sub>)
* [[Reducing sugars]], such as [[erythrose]], see [[Aldose]]
* [[Phosphite]]s, [[hypophosphite]]s, and [[phosphorous acid]]
* [[Dithiothreitol]] (DTT) – used in biochemistry labs to avoid [[SS-bond]]s
* [[Carbon monoxide]] (CO)
* [[Cyanide]]s in hydrochemical metallurgical processes
* [[Carbon]] (C)
* [[Tris-2-carboxyethylphosphine hydrochloride]] (TCEP)

==See also==
* {{annotated link|Corrosion}}
* {{annotated link|Electrochemistry}}
* {{annotated link|Electrolyte}}
* {{annotated link|Electron acceptor}}
* {{annotated link|Electron donor}}
* {{annotated link|Electrosynthesis}}
* {{annotated link|Organic reduction}}
* {{annotated link|Oxidizing agent}}
* {{annotated link|Redox}}
* {{annotated link|Reducing equivalent}}
* {{annotated link|Salt-free reduction}}

==Notes==

{{notelist}}

==References==

{{reflist}}

==Further reading==
* "Chemical Principles: The Quest for Insight", Third Edition. Peter Atkins and Loretta Jones p. F76

==External links==
* {{webarchive |url=https://web.archive.org/web/20110611104855/http://www.usm.maine.edu/~newton/Chy251_253/Lectures/OxidtionReduction/OxiationReduction.html |date=June 11, 2011 |title=Table summarizing strength of reducing agents }}

{{DEFAULTSORT:Reducing Agent}}

[[Category:Chemical reactions]]
[[Category:Redox]]
[[Category:Reducing agents| ]]