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Reducing agent

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A reducing agent (also called a reductant, reducer, or electron donor) is an element or compound that loses or "donates" an electron to an electron recipient (called the oxidizing agent, oxidant, or oxidizer) in a redox chemical reaction. The reducer's oxidation state increases while the oxidizer's decreases; this is expressed by saying that reducers "undergo oxidation" and "are oxidized" while oxidizers "undergo reduction" and "are reduced". Thus, reducing agents "reduce" oxidizers by decreasing their oxidation state while oxidizing agents "oxidize" reducers by increasing their oxidation state.

To clarify, in a redox reaction, the agent whose oxidation state increases, that "loses/donates electrons", that "is oxidized", and that "reduces" is called the reducer or reducing agent, while the agent whose oxidation state decreases, that "gains/receives electrons", that "is reduced", and that "oxidizes" is called the oxidizer or oxidizing agent. A reducing agent is thus oxidized by an oxidizer when it loses electrons that are gained by this oxidizing agent, which itself is simultaneously reduced by the reducer.

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). A reducing agent typically is in one of its lower possible oxidation states and is known as the electron donor. Examples of reducing agents include the earth metals, formic acid, oxalic acid, and sulfite compounds. For example, consider the overall reaction for aerobic cellular respiration:

C6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(l)

The oxygen (O2) is being reduced, so it is the oxidizing agent. The glucose (C6H12O6) is being oxidized, so it is the reducing agent.

In organic chemistry, reduction usually refers to the addition of hydrogen to a molecule, though the aforementioned definition still applies. For example, the oxidizing agent benzene is reduced to cyclohexane in the presence of a platinum catalyst:

C6H6 + 3 H2 → C6H12

Historically, reduction referred to the removal of oxygen from a compound, hence the name 'reduction'. An important example of this phenomenon was the Great Oxidation Event, in which biologically−produced molecular oxygen (dioxygen (O2), an oxidizer and electron recipient) was added to the early Earth's atmosphere, which was originally a reducing atmosphere containing reducing gases like methane (CH4) and carbon monoxide (CO) (along with other electron donors) and practically no oxygen because any that was produced would react with these or other reducers (particularly with iron dissolved in sea water), resulting in their removal. By using water as a reducing agent, aquatic photosynthesizing cyanobacteria produced this molecular oxygen as a waste product.[1] This O2 initially oxidized the ocean's dissolved ferrous iron (Fe(II) − meaning iron in its +2 oxidation state) to form insoluble ferric iron oxides (Fe(III) − meaning iron in its +3 oxidation state) that precipitated down to the ocean floor to form banded iron formations, 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 modern atmosphere).[2] The modern sense of donating electrons is a generalization of this idea, acknowledging that other components can play a similar chemical role to oxygen.

Characteristics

Consider the following reaction:

2 [Fe(CN)6]4− + Cl
2 → 2 [Fe(CN)6]3− + 2 Cl−

The reducing agent in this reaction is ferrocyanide ([Fe(CN)6]4−). It donates an electron, becoming oxidized to ferricyanide ([Fe(CN)6]3−). Simultaneously, that electron is received by the oxidizer chlorine (Cl
2), which is reduced to chloride (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 electrons 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 energies serve as good reducing agents too.

The measure of a material's ability to reduce is known as its reduction potential.[3] 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 agents, 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 °C. For example, among sodium (Na), chromium (Cr), cuprous (Cu+) and chloride (Cl−), it is Na that is the strongest reducing agent while Cl− is the weakest; said differently, Na+ is the weakest oxidizing agent in this list while Cl is the strongest.

Reduction potentials of various reactions[4] v
Oxidizing agent Reducing agent Reduction
Potential (V)
Li+ + e− ⇌ Li −3.04
Na+ + e− Na −2.71
Mg2+ + 2 e− Mg −2.38
Al3+ + 3 e− Al −1.66
2 H2O (l) + 2 e− H2 (g) + 2 OH− −0.83
Cr3+ + 3 e− Cr −0.74
Fe2+ + 2 e− Fe −0.44
2 H+ + 2 e− H2 0.00
Sn4+ + 2 e− Sn2+ +0.15
Cu2+ + e− Cu+ +0.16
Ag+ + e− Ag +0.80
Br2 + 2 e− 2 Br− +1.07
Cl2 + 2 e− 2 Cl− +1.36
MnO−4 + 8 H+ + 5 e− Mn2+ + 4 H2O +1.49
F2 + 2 e− 2 F− +2.87

Common reducing agents include metals potassium, calcium, barium, sodium and magnesium, and also compounds that contain the H− ion, those being NaH, LiH,[5] LiAlH4 and CaH2.

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

2 Li(s) + H2(g) → 2 LiH(s)[a]

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.

H2(g) + F2(g) → 2 HF(g)[b]

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".[3] 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.

Example of redox reaction

Example of a reduction–oxidation reaction between sodium and chlorine, with the OIL RIG mnemonic[6]

The formation of iron(III) oxide;

4Fe + 3O2 → 4Fe3+ + 6O2− → 2Fe2O3

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-reactions" that occur concurrently:

  1. Oxidation half reaction: Fe0 → Fe3+ + 3e−
  2. Reduction half reaction: O2 + 4e− → 2 O2−

Iron (Fe) has been oxidized because the oxidation number increased. Iron is the reducing agent because it gave electrons to the oxygen (O2). Oxygen (O2) has been reduced because the oxidation number has decreased and is the oxidizing agent because it took electrons from iron (Fe).

Common reducing agents

See also

Notes

  1. ^ Half reactions: 2 Li0(s) → 2 Li+(s) + 2 e− ::::: H20(g) + 2 e− → 2 H−(g)
  2. ^ Half reactions: H20(g) → 2 H+(g) + 2 e− ::::: F20(g) + 2 e− → 2 F−(g)

References

  1. ^ Buick, Roger (August 27, 2008). "When did oxygenic photosynthesis evolve?". Philosophical Transactions of the Royal Society B. 363 (1504): 2731–2743. doi:10.1098/rstb.2008.0041. ISSN 0962-8436. PMC 2606769. PMID 18468984.
  2. ^ Sosa Torres, Martha E.; Saucedo-Vázquez, Juan P.; Kroneck, Peter M.H. (2015). "Chapter 1, Section 2: The rise of dioxygen in the atmosphere". In Kroneck, Peter M.H.; Sosa Torres, Martha E. (eds.). Sustaining Life on Planet Earth: Metalloenzymes Mastering Dioxygen and Other Chewy Gases. Metal Ions in Life Sciences volume 15. Vol. 15. Springer. pp. 1–12. doi:10.1007/978-3-319-12415-5_1. ISBN 978-3-319-12414-8. PMID 25707464.
  3. ^ a b "Electrode Reduction and Oxidation Potential Values". www.EESemi.com. Retrieved 12 July 2021.
  4. ^ "Standard Electrode Potentials". hyperphysics.phy-astr.gsu.edu. Retrieved 29 March 2018.
  5. ^ Aufray M, Menuel S, Fort Y, Eschbach J, Rouxel D, Vincent B (2009). "New Synthesis of Nanosized Niobium Oxides and Lithium Niobate Particles and Their Characterization by XPS Analysis" (PDF). Journal of Nanoscience and Nanotechnology. 9 (8): 4780–4789. doi:10.1166/jnn.2009.1087. PMID 19928149.
  6. ^ "Metals". Bitesize. BBC. Archived from the original on 2022-11-03.

Further reading

  • "Chemical Principles: The Quest for Insight", Third Edition. Peter Atkins and Loretta Jones p. F76
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