Redox

Illustration of a redox reaction

Redox (shorthand for reduction-oxidation reaction) describes all chemical reactions in which atoms have their oxidation number (oxidation state) changed. This can be either a simple redox process such as the oxidation of carbon to yield carbon dioxide or the reduction of carbon by hydrogen to yield methane (CH₄), or it can be a complex process such as the oxidation of sugar in the human body through a series of very complex electron transfer processes.

The term redox comes from the two concepts of reduction and oxidation. It can be explained in simple terms:

• Oxidation describes the loss of electrons / hydrogen or gain of oxygen / increase in oxidation state by a molecule, atom or ion
• Reduction describes the gain of electrons / hydrogen or a loss of oxygen / decrease in oxidation state by a molecule, atom or ion

Though sufficient for many purposes, these descriptions are not precisely correct. Oxidation and reduction properly refer to a change in oxidation number — the actual transfer of electrons may never occur. Thus, oxidation is better defined as an increase in oxidation number, and reduction as a decrease in oxidation number. In practice, the transfer of electrons will always cause a change in oxidation number, but there are many reactions that are classed as "redox" even though no electron transfer occurs (such as those involving covalent bonds).

Non-redox reactions, which do not involve changes in formal charge, are known as metathesis reactions.

File:Redox Halves.png

The two parts of a redox reaction

Rusting iron

A bonfire. Combustion consists of redox reactions involving free radicals.

Oxidizing and reducing agents

Substances that have the ability to oxidize other substances are said to be oxidative and are known as oxidizing agents, oxidants, or oxidizers. Put another way, the oxidant removes electrons from another substance, and is, thus, reduced itself. And, because it "accepts" electrons, it is also called an electron acceptor. Reagents

Oxidants are usually chemical substances with elements in high oxidation numbers (e.g., H₂O₂, MnO₄⁻, CrO₃, Cr₂O₇²⁻, OsO₄) or highly electronegative substances that can gain one or two extra electrons by oxidizing a substance (O, F, Cl, Br).

Substances that have the ability to reduce other substances are said to be reductive and are known as reducing agents, reductants, or reducers. Put in another way, the reductant transfers electrons to another substance, and is, thus, oxidized itself. And, because it "donates" electrons it is also called an electron donor. Reductants in chemistry are very diverse. Metal reduction—electropositive elemental metals can be used (Li, Na, Mg, Fe, Zn, Al). These metals donate or give away electrons readily. Reactions-Reducing Agents Other kinds of reductants are hydride transfer reagents (NaBH₄, LiAlH₄), these reagents are widely used in organic chemistry, primarily in the reduction of carbonyl compounds to alcohols. Another useful method is reductions involving hydrogen gas (H₂) with a palladium, platinum, or nickel catalyst. These catalytic reductions are primarily used in the reduction of carbon-carbon double or triple bonds.

The chemical way to look at redox processes is that the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or reducing agent loses electrons and is oxidized, and the oxidant or oxidizing agent gains electrons and is reduced. The pair of an oxidizing and reducing agent that are involved in a particular reaction is called a redox pair.

Examples of redox reactions

A good example is the reaction between hydrogen and fluorine:

$\mathrm {H} _{2}+\mathrm {F} _{2}\longrightarrow 2\mathrm {HF}$

We can write this overall reaction as two half-reactions: the oxidation reaction

$\mathrm {H} _{2}\longrightarrow 2\mathrm {H} ^{+}+2e^{-}$

and the reduction reaction:

$\mathrm {F} _{2}+2e^{-}\longrightarrow 2\mathrm {F} ^{-}$

Analysing each half-reaction in isolation can often make the overall chemical process clearer. Because there is no net change in charge during a redox reaction, the number of electrons in excess in the oxidation reaction must equal the number consumed by the reduction reaction (as shown above).

Elements, even in molecular form, always have an oxidation number of zero. In the first half-reaction, hydrogen is oxidized from an oxidation number of zero to an oxidation number of +1. In the second half-reaction, fluorine is reduced from an oxidation number of zero to an oxidation number of −1.

When adding the reactions together the electrons cancel:

${\frac {\begin{array}{rcl}\mathrm {H} _{2}&\longrightarrow &2\mathrm {H} ^{+}+2e^{-}\\\mathrm {F} _{2}+2e^{-}&\longrightarrow &2\mathrm {F} ^{-}\end{array}}{\begin{array}{rcl}\mathrm {H} _{2}+\mathrm {F} _{2}&\longrightarrow &2\mathrm {H} ^{+}+2\mathrm {F} ^{-}\end{array}}}$

And the ions combine to form hydrogen fluoride:

$\mathrm {H} _{2}+\mathrm {F} _{2}\,\ \longrightarrow \ 2\mathrm {H} ^{+}+2\mathrm {F} ^{-}\ \longrightarrow \ 2\mathrm {HF}$

Displacement reactions

Redox occurs in single displacement reactions or substitution reactions. The redox component of this type of reaction is the change of oxidation state (charge) on certain atoms, not the actual exchange of atoms in the compounds.

For example, in the reaction between iron and copper(II) sulphate solution:

$\mathrm {Fe} +\mathrm {CuSO} _{4}\longrightarrow \mathrm {FeSO} _{4}+\mathrm {Cu}$

The ionic equation for this reaction is:

$\mathrm {Fe} +\mathrm {Cu} ^{2+}\longrightarrow \mathrm {Fe} ^{2+}+\mathrm {Cu}$

As two half-equations, it is seen that the iron is oxidized:

$\mathrm {Fe} \longrightarrow \mathrm {Fe} ^{2+}+2{e}^{-}$

And the copper is reduced:

$\mathrm {Cu} ^{2+}+2{e}^{-}\longrightarrow \mathrm {Cu}$

Other examples

• iron(II) oxidizes to iron(III):

Fe²⁺ → Fe³⁺ + e⁻

• hydrogen peroxide reduces to hydroxide in the presence of an acid:

H₂O₂ + 2 e⁻ → 2 OH⁻

overall equation for the above:

2Fe²⁺ + H₂O₂ + 2H⁺ → 2Fe³⁺ + 2H₂O

• denitrification, nitrate reduces to nitrogen in the presence of an acid:

2NO₃⁻ + 10e⁻ + 12 H⁺ → N₂ + 6H₂O

• iron oxidizes to iron(III) oxide and oxygen is reduced forming iron(III) oxide (commonly known as rusting, which is similar to tarnishing):

4Fe + 3O₂ → 2 Fe₂O₃

• Combustion of hydrocarbons, e.g., in an internal combustion engine, produces water, carbon dioxide, some partially oxidized forms such as carbon monoxide and heat energy. Complete oxidation of materials containing carbon produces carbon dioxide.
• In organic chemistry, stepwise oxidation of a hydrocarbon produces water and, successively, an alcohol, an aldehyde or a ketone, carboxylic acid, and then a peroxide.

Redox reactions in industry

The primary process of reducing ore to produce metals is discussed in the article on Smelting.

Oxidation is used in a wide variety of industries such as in the production of cleaning products and oxidising ammonia to produce nitric acid, which is used in most fertilizers.

Redox reactions are the foundation of electrochemical cells.

The production of compact discs depends on a redox reaction, which coats the disc with a thin layer of metal film.

Redox reactions in biology

Top: ascorbic acid (reduced form of Vitamin C)
Bottom: dehydroascorbic acid (oxidized form of Vitamin C)

Many important biological processes involve redox reactions.

Cellular respiration, for instance, is the oxidation of glucose (C₆H₁₂O₆) to CO₂ and the reduction of oxygen to water. The summary equation for cell respiration is:

C₆H₁₂O₆ + 6 O₂ → 6 CO₂ + 6 H₂O

The process of cell respiration also depends heavily on the reduction of NAD⁺ to NADH and the reverse reaction (the oxidation of NADH to NAD⁺). Photosynthesis is essentially the reverse of the redox reaction in cell respiration:

6 CO₂ + 6 H₂O + light energy → C₆H₁₂O₆ + 6 O₂

Biological energy is frequently stored and released by means of redox reactions. Photosynthesis involves the reduction of carbon dioxide into sugars and the oxidation of water into molecular oxygen. The reverse reaction, respiration, oxidizes sugars to produce carbon dioxide and water. As intermediate steps, the reduced carbon compounds are used to reduce nicotinamide adenine dinucleotide (NAD⁺), which then contributes to the creation of a proton gradient, which drives the synthesis of adenosine triphosphate (ATP) and is maintained by the reduction of oxygen. In animal cells, mitochondria perform similar functions. See Membrane potential article.

The term redox state is often used to describe the balance of NAD⁺/NADH and NADP⁺/NADPH in a biological system such as a cell or organ. The redox state is reflected in the balance of several sets of metabolites (e.g., lactate and pyruvate, beta-hydroxybutyrate and acetoacetate), whose interconversion is dependent on these ratios. An abnormal redox state can develop in a variety of deleterious situations, such as hypoxia, shock, and sepsis. Redox signaling involves the control of cellular processes by redox processes.

Redox proteins and their genes must be Co-located for Redox Regulation according to the CoRR Hypothesis for the function of DNA in mitochondria and chloroplasts.

Redox cycling

A wide variety of aromatic compounds are enzymatically reduced to form free radicals that contain one more electron than their parent compounds. In general, the electron donor is any of a wide variety of flavoenzymes and their coenzymes. Once formed, these anion free radicals reduce molecular oxygen to superoxide, and regenerate the unchanged parent compound. The net reaction is the oxidation of the flavoenzyme's coenzymes and the reduction of molecular oxygen to form superoxide. This catalytic behavior has been described as futile cycle or redox cycling.

Examples of redox cycling-inducing molecules are the herbicide paraquat and other viologens and quinones such as menadione.

Balancing redox reactions

Describing the overall electrochemical reaction for a redox process requires a balancing of the component half-reactions for oxidation and reduction. For reactions in aqueous solution, this generally involves adding H⁺ , OH^- ion, H₂O and electrons to compensate the oxidation changes.

Acid medium

In acid medium H⁺ ions and water are added to half reactions to balance the overall reaction. For example, when manganese (II) reacts with sodium bismuthate.

${\mbox{Reaction unbalanced: }}{\mbox{Mn}}^{2+}(aq)+{\mbox{NaBiO}}_{3}(s)\rightarrow {\mbox{Bi}}^{3+}(aq)+{\mbox{MnO}}_{4}^{-}(aq)\,$

${\mbox{Oxidation: }}{\mbox{4H}}_{2}{\mbox{O}}(l)+{\mbox{Mn}}^{2+}(aq)\rightarrow {\mbox{MnO}}_{4}^{-}(aq)+{\mbox{8H}}^{+}(aq)+{\mbox{5e}}^{-}\,$

${\mbox{Reduction: }}{\mbox{2e}}^{-}+{\mbox{6H}}^{+}(aq)+{\mbox{BiO}}_{3}^{-}(s)\rightarrow {\mbox{Bi}}^{3+}(aq)+{\mbox{3H}}_{2}{\mbox{O}}(l)\,$

The reaction is balanced by scaling the two half-cell reactions to involve the same number of electrons (i.e. multiplying the oxidation reaction by the number of electrons in the reduction step and vice versa). Addition gives:

${\mbox{8H}}_{2}{\mbox{O}}(l)+{\mbox{2Mn}}^{2+}(aq)\rightarrow {\mbox{2MnO}}_{4}^{-}(aq)+{\mbox{16H}}^{+}(aq)+{\mbox{10e}}^{-}\,$

${\mbox{10e}}^{-}+{\mbox{30H}}^{+}(aq)+{\mbox{5BiO}}_{3}^{-}(s)\rightarrow {\mbox{5Bi}}^{3+}(aq)+{\mbox{15H}}_{2}{\mbox{O}}(l)\,$

Reaction balanced:

${\mbox{14H}}^{+}(aq)+{\mbox{2Mn}}^{2+}(aq)+{\mbox{5NaBiO}}_{3}(s)\rightarrow {\mbox{7H}}_{2}{\mbox{O}}(l)+{\mbox{2MnO}}_{4}^{-}(aq)+{\mbox{5Bi}}^{3+}(aq)+{\mbox{5Na}}^{+}(aq)\,$

Similarly for a propane fuel cell under acidic conditions:

${\mbox{Reaction unbalanced: }}{\mbox{C}}_{3}{\mbox{H}}_{8}+{\mbox{O}}_{2}\rightarrow {\mbox{CO}}_{2}+{\mbox{H}}_{2}{\mbox{O}}\,$

${\mbox{Reduction: }}{\mbox{4H}}^{+}+{\mbox{O}}_{2}+{\mbox{4e}}^{-}\rightarrow {\mbox{2H}}_{2}{\mbox{O}}\,$

${\mbox{Oxidation: }}{\mbox{6H}}_{2}{\mbox{O}}+{\mbox{C}}_{3}{\mbox{H}}_{8}\rightarrow {\mbox{3CO}}_{2}+{\mbox{20e}}^{-}+{\mbox{20H}}^{+}\,$

Balancing the number of electrons involved gives:

${\mbox{20H}}^{+}+{\mbox{5O}}_{2}+{\mbox{20e}}^{-}\rightarrow {\mbox{10H}}_{2}{\mbox{O}}\,$

${\mbox{6H}}_{2}{\mbox{O}}+{\mbox{C}}_{3}{\mbox{H}}_{8}\rightarrow {\mbox{3CO}}_{2}+{\mbox{20e}}^{-}+{\mbox{20H}}^{+}\,$

Equation balanced:

${\mbox{C}}_{3}{\mbox{H}}_{8}+{\mbox{5O}}_{2}\rightarrow {\mbox{3CO}}_{2}+{\mbox{4H}}_{2}{\mbox{O}}\,$

Basic medium

In basic medium OH^- ions and water are added to half reactions to balance the overall reaction. For example, in the reaction between potassium permanganate and sodium sulfite:

${\mbox{Reaction unbalanced: }}{\mbox{KMnO}}_{4}+{\mbox{Na}}_{2}{\mbox{SO}}_{3}+{\mbox{H}}_{2}{\mbox{O}}\rightarrow {\mbox{MnO}}_{2}+{\mbox{Na}}_{2}{\mbox{SO}}_{4}+{\mbox{KOH}}\,$

${\mbox{Reduction: }}{\mbox{3e}}^{-}+{\mbox{2H}}_{2}{\mbox{O}}+{\mbox{MnO}}_{4}^{-}\rightarrow {\mbox{MnO}}_{2}+{\mbox{4OH}}^{-}\,$

${\mbox{Oxidation: }}{\mbox{2OH}}^{-}+{\mbox{SO}}_{3}^{2-}\rightarrow {\mbox{SO}}_{4}^{2-}+{\mbox{H}}_{2}{\mbox{O}}+{\mbox{2e}}^{-}\,$

Balancing the number of electrons in the two half-cell reactions gives:

${\mbox{6e}}^{-}+{\mbox{4H}}_{2}{\mbox{O}}+{\mbox{2MnO}}_{4}^{-}\rightarrow {\mbox{2MnO}}_{2}+{\mbox{8OH}}^{-}\,$

${\mbox{6OH}}^{-}+{\mbox{3SO}}_{3}^{2-}\rightarrow {\mbox{3SO}}_{4}^{2-}+{\mbox{3H}}_{2}{\mbox{O}}+{\mbox{6e}}^{-}\,$

Equation balanced:

${\mbox{2KMnO}}_{4}+{\mbox{3Na}}_{2}{\mbox{SO}}_{3}+{\mbox{H}}_{2}{\mbox{O}}\rightarrow {\mbox{2MnO}}_{2}+{\mbox{3Na}}_{2}{\mbox{SO}}_{4}+{\mbox{2KOH}}\,$

See also

• Organic reduction
• Hydrogenation
• Bessemer process
• Bioremediation
• Calvin cycle
• Citric acid cycle
• Electrochemical cell
• Electrochemistry

• Galvanic cell
• Membrane potential
• Oxidative addition and reductive elimination
• Reducing agent
• Thermic reaction
• Partial oxidation
• Reduction potential

External links

• Chemical Equation Balancer, balances chemical equations, including Reduction-Oxidation (redox) reactions and reactions with several distinct solutions. Also teaches how to balance an equation using matrices and linear algebra.
• Redox reactions calculator
• Redox reactions at Chemguide
• Online redox reaction equation balancer, balances equations of any half-cell and full reactions

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