Reducing agent

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In chemistry, a reducing agent (also known as a reductant, reducer, or electron donor ) is a chemical species that "donates" an electron to an electron recipient (called the oxidizing agent , oxidant, oxidizer, or electron acceptor ).

Contents

Examples of substances that are common reducing agents include hydrogen, the alkali metals, formic acid, [1] oxalic acid, [2] 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/donates electrons", that "is oxidized", and that "reduces" is called the reducer or reducing agent, while the agent whose oxidation state decreases, that "gains/accepts/receives electrons", that "is reduced", and that "oxidizes" is called the oxidizer or oxidizing agent.

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.

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.[ citation needed ]

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.[ citation needed ]

Reduction potentials of various reactions [4] v
Oxidizing agentReducing agentReduction
Potential (V)
Li+ + eLi−3.04
Na+ + eNa−2.71
Mg2+ + 2 eMg−2.38
Al3+ + 3 eAl−1.66
2 H2O (l) + 2 eH2 (g) + 2 OH−0.83
Cr3+ + 3 eCr−0.74
Fe2+ + 2 eFe−0.44
2 H+ + 2 eH20.00
Sn4+ + 2 eSn2++0.15
Cu2+ + eCu++0.16
Ag+ + eAg+0.80
Br2 + 2 e2 Br+1.07
Cl2 + 2 e2 Cl+1.36
MnO4 + 8 H+ + 5 eMn2+ + 4 H2O+1.49
F2 + 2 e2 F+2.87

Common reducing agents include metals potassium, calcium, barium, sodium and magnesium, and also compounds that contain the hydride 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) [lower-alpha 1]

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) [lower-alpha 2]

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.[ citation needed ]

Examples of redox reaction

Example of a reduction-oxidation reaction between sodium and chlorine, with the OIL RIG mnemonic Redox example.svg
Example of a reduction–oxidation reaction between sodium and chlorine, with the OIL RIG mnemonic

Historically, reduction referred to the removal of oxygen from a compound, hence the name 'reduction'. [7] An example of this phenomenon occurred during 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 weakly reducing atmosphere containing reducing gases like methane (CH4) and carbon monoxide (CO) (along with other electron donors) [8] 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. [9] 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 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 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). [10] 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 + 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)

Related Research Articles

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A chemical reaction is a process that leads to the chemical transformation of one set of chemical substances to another. When chemical reactions occur, the atoms are rearranged and the reaction is accompanied by an energy change as new products are generated. Classically, chemical reactions encompass changes that only involve the positions of electrons in the forming and breaking of chemical bonds between atoms, with no change to the nuclei, and can often be described by a chemical equation. Nuclear chemistry is a sub-discipline of chemistry that involves the chemical reactions of unstable and radioactive elements where both electronic and nuclear changes can occur.

<span class="mw-page-title-main">Electrochemistry</span> Branch of chemistry

Electrochemistry is the branch of physical chemistry concerned with the relationship between electrical potential difference and identifiable chemical change. These reactions involve electrons moving via an electronically-conducting phase between electrodes separated by an ionically conducting and electronically insulating electrolyte.

<span class="mw-page-title-main">Oxide</span> Chemical compound where oxygen atoms are combined with atoms of other elements

An oxide is a chemical compound containing at least one oxygen atom and one other element in its chemical formula. "Oxide" itself is the dianion of oxygen, an O2– ion with oxygen in the oxidation state of −2. Most of the Earth's crust consists of oxides. Even materials considered pure elements often develop an oxide coating. For example, aluminium foil develops a thin skin of Al2O3 that protects the foil from further oxidation.

<span class="mw-page-title-main">Electrolysis</span> Technique in chemistry and manufacturing

In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell. The voltage that is needed for electrolysis to occur is called the decomposition potential. The word "lysis" means to separate or break, so in terms, electrolysis would mean "breakdown via electricity."

In chemistry, a half reaction is either the oxidation or reduction reaction component of a redox reaction. A half reaction is obtained by considering the change in oxidation states of individual substances involved in the redox reaction. Often, the concept of half reactions is used to describe what occurs in an electrochemical cell, such as a Galvanic cell battery. Half reactions can be written to describe both the metal undergoing oxidation and the metal undergoing reduction.

<span class="mw-page-title-main">Redox</span> Chemical reaction in which oxidation states of atoms are changed

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<span class="mw-page-title-main">Hydride</span> Molecule with a hydrogen bound to a more electropositive element or group

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<span class="mw-page-title-main">Oxidizing agent</span> Chemical compound used to oxidize another substance in a chemical reaction

An oxidizing agent is a substance in a redox chemical reaction that gains or "accepts"/"receives" an electron from a reducing agent. In other words, an oxidizer is any substance that oxidizes another substance. The oxidation state, which describes the degree of loss of electrons, of the oxidizer decreases while that of the reductant increases; this is expressed by saying that oxidizers "undergo reduction" and "are reduced" while reducers "undergo oxidation" and "are oxidized". Common oxidizing agents are oxygen, hydrogen peroxide, and the halogens.

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<span class="mw-page-title-main">Chromate and dichromate</span> Chromium(VI) anions

Chromate salts contain the chromate anion, CrO2−
4
. Dichromate salts contain the dichromate anion, Cr
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O2−
7
. They are oxyanions of chromium in the +6 oxidation state and are moderately strong oxidizing agents. In an aqueous solution, chromate and dichromate ions can be interconvertible.

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<span class="mw-page-title-main">Single displacement reaction</span> Type of chemical reaction

A single-displacement reaction, also known as single replacement reaction or exchange reaction, is an archaic concept in chemistry. It describes the stoichiometry of some chemical reactions in which one element or ligand is replaced by atom or group.

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<span class="mw-page-title-main">Tin(II) chloride</span> Chemical compound

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<span class="mw-page-title-main">Aluminium compounds</span>

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References

  1. Garron, Anthony; Epron, Florence (2005). "Use of formic acid as reducing agent for application in catalytic reduction of nitrate in water". Water Research. 39 (13): 3073–3081. Bibcode:2005WatRe..39.3073G. doi:10.1016/j.watres.2005.05.012. PMID   15982701.
  2. "Oxidizing and Reducing Agents". Purdue University.
  3. 1 2 "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. Archived from the original (PDF) on 2020-07-29. Retrieved 2019-09-24.
  6. "Metals". Bitesize. BBC. Archived from the original on 2022-11-03.
  7. Olson, Maynard V. "oxidation-reduction reaction". Britannica. Retrieved 3 May 2022. In his Traité élémentaire de chimie, 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.
  8. Kasting, J.F. (2014). "Modeling the Archean Atmosphere and Climate". Treatise on Geochemistry. Elsevier. pp. 157–175. doi:10.1016/b978-0-08-095975-7.01306-1. ISBN   9780080983004.
  9. 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.
  10. 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.
  11. "Cathodic Stripping Voltammetric Procedure for Determination of Some Inorganic Arsenic Species in Water, Soil and Ores Samples".

Further reading