Redox

Redox (shorthand for reduction/oxidation reaction) describes all chemical reactions in which atoms have their oxidation state) changed.

This can be either a simple redox process such as the oxidation of 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 ion
Reduction describes the gain of ion

However, these descriptions (though sufficient for many purposes) are not truly 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 which 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 metathesis reactions.

Oxidizing and reducing agents

Substances that have the ability to oxidize other substances are said to be oxidative and are known as electron acceptor.

Oxidants are usually chemical substances with elements in high oxidation numbers (e.g., Br).

Substances that have the ability to reduce other substances are said to be reductive and are known as 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 oxidising 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 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:

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

Other examples

iron(II) oxidizes to iron(III):

Fe²⁺ → Fe³⁺ + e⁻

hydroxide in the presence of an acid:

H₂O₂ + 2 e⁻ → 2 OH⁻

overall equation for the above:

2Fe²⁺ + H₂O₂ + 2H⁺ → 2Fe³⁺ + 2H₂O

denitrification, 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 carbon produces carbon dioxide.

In peroxide.

Redox reactions in industry

Oxidation is used in a wide variety of industries such as in the production of cleaning products.
Redox reactions are the foundation of electrochemical cells.

Redox reactions in biology

Top: Vitamin C)

Many important biological processes involve redox reactions.
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 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. adenosine triphosphate (ATP) and is maintained by the reduction of oxygen. In animal cells, Membrane potential article.
The term redox state is often used to describe the balance of hypoxia, shock, and sepsis. Redox signaling involves the control of cellular processes by redox processes.

Redox cycling

A wide variety of 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 menadione. [1]PDF (2.76 MiB)

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 general involves adding 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 sodium sulfite.

$\mbox{Reaction unbalanced: }\mbox{KMnO}_{4}+\mbox{Na}_{2}\mbox{SO}_3+\mbox{H}_2\mbox{O}\rightarrow\mbox{MnO}_{2}+\mbox{Na}_{2}\mbox{SO}_{4}+\mbox{KOH}\,$
$\mbox{Reduction: }\mbox{3e}^{-}+\mbox{2H}_{2}\mbox{O}+\mbox{MnO}_{4}^{-}\rightarrow\mbox{MnO}_{2}+\mbox{4OH}^{-}\,$
$\mbox{Oxidation: }\mbox{2OH}^{-}+\mbox{SO}^{2-}_{3}\rightarrow\mbox{SO}^{2-}_{4}+\mbox{H}_{2}\mbox{O}+\mbox{2e}^{-}\,$

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

$\mbox{6e}^{-}+\mbox{4H}_{2}\mbox{O}+\mbox{2MnO}_{4}^{-}\rightarrow\mbox{2MnO}_{2}+\mbox{8OH}^{-}\,$
$\mbox{6OH}^{-}+\mbox{3SO}^{2-}_{3}\rightarrow\mbox{3SO}^{2-}_{4}+\mbox{3H}_{2}\mbox{O}+\mbox{6e}^{-}\,$

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

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

References

^ Hudlický, Miloš (1996). Reductions in Organic Chemistry. Washington, D.C.: American Chemical Society, 429. ISBN 0-8412-3344-6.

^ Hudlický, Miloš (1990). Oxidations in Organic Chemistry. Washington, D.C.: American Chemical Society, 456. ISBN 0-8412-1780-7.

Categories: Soil chemistry

This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Redox". A list of authors is available in Wikipedia.