Monday, February 14, 2011

REDOX


Illustration of a redox reaction

Redox (shorthand for oxidation-reduction) reactions describe 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 (CO2) or the reduction of carbon by hydrogen to yield methane (CH4), or a complex process such as the oxidation of sugar (C6H12O6) in the human body through a series of complex electron transfer processes.

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

  • Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion.
  • Reduction is the gain of electrons or a 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.

The two parts of a redox reaction
Rusting iron
A bonfire. Combustion consists of redox reactions involving free radicals.


Oxidizing and reducing agents

In redox processes 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.

Oxidizers

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

Oxidants are usually chemical elements or substances with elements in high oxidation numbers (e.g., H2O2, MnO
4
, CrO3, Cr2O2−
7
, OsO4) or highly electronegative substances/elements that can gain one or two extra electrons by oxidizing an element or substance (O, F, Cl, Br).

Reducers

Substances that have the ability to reduce other substances are said to be reductive or reducing and are known as reducing agents, reductants, or reducers. That is, the reductant transfers electrons to another substance, and is thus itself oxidized. And, because it "donates" electrons, it is also called an electron donor. Electron donors can also form charge transfer complexes with electron acceptors.

Reductants in chemistry are very diverse. Electropositive elemental metals, such as lithium, sodium, magnesium, iron, zinc, aluminium, carbon, are good reducing agents. These metals donate or give away electrons readily. Hydride transfer reagents, such as NaBH4 and LiAlH4, are widely used in organic chemistry,[1][2] primarily in the reduction of carbonyl compounds to alcohols. Another method of reduction involves the use of hydrogen gas (H2) with a palladium, platinum, or nickel catalyst. These catalytic reductions are used primarily in the reduction of carbon-carbon double or triple bonds.

Examples of redox reactions

A good example is the reaction between hydrogen and fluorine in which hydrogen is being oxidized and fluorine is being reduced:

H2 + F2 → 2 HF

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

the oxidation reaction:

H2 → 2 H+ + 2 e

and the reduction reaction:

F2 + 2 e
→ 2 F

Analyzing 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:

H2 2 H+ + 2 e
F2 + 2 e
2 F

H2 + F2 2 H+ + 2 F

And the ions combine to form hydrogen fluoride:

H2 + F2 → 2 H+ + 2 F → 2 HF

Displacement reactions

Redox occurs in single displacement reactions or substitution reactions. The redox component of these types of reactions 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) sulfate solution:

Fe + CuSO4FeSO4 + Cu

The ionic equation for this reaction is:

Fe + Cu2+ → Fe2+ + Cu

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

Fe → Fe2+ + 2 e

And the copper is reduced:

Cu2+ + 2 e
→ Cu

Other examples

  • The oxidation of iron(II) to iron(III) by hydrogen peroxide in the presence of an acid:
Fe2+ → Fe3+ + e
H2O2 + 2 e → 2 OH
Overall equation:
2 Fe2+ + H2O2 + 2 H+ → 2 Fe3+ + 2 H2O
2 NO3 + 10 e + 12 H+ → N2 + 6 H2O
Iron rusting in pyrite cubes
  • Oxidation of elemental iron to iron(III) oxide by oxygen (commonly known as rusting, which is similar to tarnishing):
4 Fe + 3 O2 → 2 Fe2O3

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, H2O, and electrons to compensate for the oxidation changes.

Acidic media

In acidic media, H+ ions and water are added to half reactions to balance the overall reaction.

For example, when manganese(II) reacts with sodium bismuthate:

Unbalanced reaction: Mn2+(aq) + NaBiO3(s) → Bi3+(aq) + MnO4 (aq)
Oxidation: 4 H2O(l) + Mn2+(aq) → MnO
4
(aq) + 8 H+(aq) + 5 e
Reduction: 2 e
+ 6 H+ + BiO
3
(s) → Bi3+(aq) + 3 H2O(l)

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

8 H2O(l) + 2 Mn2+(aq) → 2 MnO
4
(aq) + 16 H+(aq) + 10 e
10 e
+ 30 H+ + 5 BiO
3
(s) → 5 Bi3+(aq) + 15 H2O(l)

Adding these two reactions eliminates the electrons terms and yields the balanced reaction:

14 H+(aq) + 2 Mn2+(aq) + 5 NaBiO3(s) → 7 H2O(l) + 2 MnO
4
(aq) + 5 Bi3+(aq) + 5 Na+(aq)

Basic media

In basic media, 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:

Unbalanced reaction: KMnO4 + Na2SO3 + H2OMnO2 + Na2SO4 + KOH
Reduction: 3 e
+ 2 H2O + MnO4MnO2 + 4 OH
Oxidation: 2 OH + SO32−SO42− + H2O + 2 e

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

6 e
+ 4 H2O + 2 MnO4 → 2 MnO2 + 8 OH
6 OH + 3 SO32− → 3 SO42− + 3 H2O + 6 e

Adding these two half-cell reactions together gives the balanced equation:

2 KMnO4 + 3 Na2SO3 + H2O → 2 MnO2 + 3 Na2SO4 + 2 KOH

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