Reaction of oxydoreduction

A reaction of oxydoreduction is a process of transfer of electron S of a species to another. One calls reducing the species which yields the electrons and oxidant the species which collects them during the reaction.

Examples of reactions of oxydoreduction

Oxydoreduction constitutes an big family of reactions, since it intervene in the Combustion S, certain proportionings, in the Métallurgie, the Corrosion of metals, the electrochemistry or the cellular Respiration. This variety is explained by mobility, lightness and omnipresence in all shapes of the matter of the electron. These reactions also play a fundamental role in Biologie, in the transformation of oxygen within the human body for example.

Definitions

First definition

Following experiments with the mercury, Lavoisier highlights in 1772 the role of the Dioxygène in certain reactions of oxydoreduction. It poses the first definitions:

oxidation means “combination with oxygen”. For example:

2 Hg + O₂ → 2 HgO

a reduction is “the extraction of a Métal of its oxide”, definition already used out of metallurgy. For example:

SO₂ → S + O₂

In the language running, oxidation is the Chemical reaction in which a compound combines with one or more Atome S of Oxygène. Such as for example the oxidation of the Iron which produces rust:

4Fe + 3O₂ → 2 Fe₂O₃.

It is only at the 20th century, after the discovery of the electron (J.J. Thomson, 1897) and the introduction of the atomic model of Bohr (1913) that the chemical reactions were re-examined in the light of these new models and that similarities observed made it possible to gradually release the current concept of oxydoreduction which is expressed in terms of transfers of electrons.

More modern definitions

To facilitate the study of the reactions, one uses a tool which associates (sometimes abstractedly) with each atom of a compound a number of oxidation (n.o.) which symbolizes the value of the load carried. (Fe²⁺ has a number of oxidation of 2.)

an oxidation is a loss of electrons (thus an increase in the n.o., electrons being negatively charged). For example:

Cu → Cu²⁺ + 2e^-

This gift of electrons does not occur that if there exists a body likely to accept them.
the opposite phenomenon (acceptance of the electrons) is called the reduction.

a reduction is a profit of electrons (thus a reduction in the n.o., electrons being negatively charged). For example:

I₂ + 2e^- → 2 I^-

Thus, the “combinations with oxygen” are only one particular case of the reactions of oxydoreduction. Here two reactions with copper:

Cu + ½ O₂ → CuO

Cu²⁺ + 2 Cl^- → CuCl₂

The first combines the Cuivre and dioxygene while the second combines copper and the ion chloride. The ion chloride and the dioxygene have a common point: they are electronegative elements more which copper.

In fact, the oxidation of a body is always accompanied by the reduction by another (the electrons cannot be trotted all alone and are necessarily collected), one speaks about a reaction of oxydoreduction . Oxidation is a half-reaction of oxydoreduction, and the reduction is the other half-reaction.

Vocabulary

In an oxydoreduction,

the element which loses an electron (S) is called reducing ,
the element which collects an electron (S) is called Oxydant .

The reducer oxidizes (reaction of Oxydation), oxidizing it is reduced (reaction of reduction). Oxydoreduction is thus composed of two half-reactions: an oxidation and a reduction.

reducing Oxidation
(1) → oxidant (1) + N e^-
Reduction
oxidant (2) + reducing N e^- → (2)
Oxydoreduction (“nap” of oxidation and the reduction)
oxidant (2) + reducing (1) → oxidant (1) + reducing (2)

Example:

Ce⁴⁺ + e^- → Ce³⁺

Fe²⁺ → Fe³⁺ +e^-

from where the reaction assessment:

Ce⁴⁺ + Fe²⁺ → Ce³⁺ + Fe³⁺

An oxidized reducer is oxidizing, and a reduced oxidant is a reducer. One thus defines the couple oxidant-reducer (in the past called “redox cell”) which is composed of oxidant and the combined reducer (the reduced oxidant). One notes it in the form: reducing oxidant / .

Note: some chemical compounds can behave as well while oxidizing as out of reducer. It is in particular the case of the hydrogen peroxide, which one says that it dismute, and who consequently cannot be preserved a long time:

H₂O₂ → 2:00 ⁺ + O₂ + 2e^- (oxidation)

H₂O₂ + 2:00 ⁺ + 2e^- → 2:00 ₂O (reduction)

Maybe with the final one:

2:00 ₂O₂ → 2:00 ₂O + O₂ (oxydoreduction)

There are for example the couples oxidant-reducer Cu²⁺/Cu and Zn²⁺/Zn, which gives the reaction in aqueous Solution:

Zn_(S) + Cu²⁺_(aq) → Zn²⁺_(aq) + Cu_(S) (oxydoreduction)

This reaction can break up into a reduction and an oxidation:

Zn_(S) → Zn²⁺_(aq) + 2e^- (oxidation)

Cu²⁺_(aq) + 2e^- → Cu_(S) (reduction)

The two half-reactions of oxidation and reduction can really be separate in certain cases (i.e. they do not occur at the same place), which makes it possible to generate a Electric current (it is what occurs in the Batteries). In the other cases, for example in the example given, they have only one formal interest (the free electrons do not exist in water).

Balance equations of reaction

When one writes the equations of reaction, it should be made sure that one always has the same number of atoms of each type of each side of the arrow; one calls this “the balancing of the equation”.

There are sometimes complex reactions which require to balance the stoechiometric coefficients of the half-equations. One can also have to add molecules or ions in solution (according to the medium) to balance.

For example for the reaction enters the Permanganate of potassium (couple MnO₄^-/Mn²⁺) and an iron solution (couple Fe³⁺/Fe²⁺):

Fe²⁺ → Fe³⁺ + e^-

MnO₄^- + 8:00 ⁺ + 5e^- → Mn²⁺ + 4:00 ₂O

MnO₄^- + 8:00 ⁺ + 5Fe²⁺ → Mn²⁺ + 4:00 ₂O + 5Fe³⁺

To balance the reaction redox, it is also to linearly combine the half-reactions (oxidation and reduction) so that the number of electrons given is exactly the number of accepted electrons: it is said that the réacton redox is a strict exchange of electrons (thermodynamically favorable).

For example:

Fe → Fe³⁺ + 3e^-;

O₂ + 4e^- → 2 O^2-2

In this case, it is thus a question of finding the lowest common multiple of 3 and of 4 is 12, so as to have an assessment of strict exchange: it is thus necessary to combine 4 times the first half-reaction (iron will provide 12 electrons) with 3 times the second half-reaction (the dioxygene will accept 12 electrons), that is to say:

4 Fe + 3O₂ → 4 Fe³⁺ + 6 O^2-

This constitutes the exchange of electrons which constitutes the phenomenon redox.

Then, it occurs an electrostatic attraction: the positive loads and the negative charges attract each other and are prepared so as to form a neutral ionic crystal:

4 Fe³⁺ + 6 O^2- → 2 Fe₂O₃

Attention, this is not a chemical reaction strictly speaking, but a simple rewriting corresponding to static attraction in the ionic crystal, called hematite!

Potential of oxydoreduction

The character “oxidizing” or “reducer” is not absolute, but relative, within the framework of a chemical reaction. A reducing element in a reaction can be oxidizing in another. But it is possible to build a scale of oxidizing force (or, in the other direction, of reducing force): it is the Potentiel of oxydoreduction, which is measured in Volt. Moreover, this potential can depend on the chemical context and in particular on pH, and even on the physical context: the effects of the light are made profitable as well by nature in the Photosynthèse, as by the man in the Photographie.

Some couples of oxidant-reducer

All the couples of oxidant-reducer are written in the Ox/Red form.

Ag⁺/Ag

Al³⁺/Al

Au³⁺/With the

Br₂/Br^-

CH₃CHO/C₂H₅OH

Cl₂/Cl^-

Cr₂O₇^2-/Cr³⁺

Cu²⁺/Cu

Fe³⁺/Fe²⁺

Fe²⁺/Fe

H⁺/H₂

Hg^2+/Hg

H₂O₂/H₂

I₂/I^-

Mg²⁺/Mg

MnO₄^-/Mn²⁺

Ni²⁺/Nor

NO₂^-/NO

NO₃^-/NO₂^-

O₂/H₂O

Pb²⁺/Pb

S₄O₆^2-/S₂O₃^2-

Sn²⁺/Sn

SO₄^2-/SO₂

Zn²⁺/Zn

See too

Oxidation and reduction in organic chemistry
Combustible Electrochemistry
Combustion
Combustive
Corrosion
Fire
Cupellation

Random links: