Box 43 Oxidation and reduction redox

Oxidation and reduction (redox) reactions are driven by electron transfers (see Section 2.2). Thus the oxidation of iron by oxygen:

can be considered to consist of two half-reactions:

where e- represents one electron.

Oxidation involves the loss of electrons and reduction involves the gain of electrons

In equations 1-3, oxygen—the oxidizing agent (also called an electron acceptor) — is reduced because it gains electrons.

Equation 1 shows that elements like oxygen can exist under set conditions of temperature and pressure in more than one state, i.e. as oxygen gas and as an oxide. The oxidation state of an element in a compound is assigned using the following rules:

1 The oxidation number of all elements is 0.

2 The oxidation number of a monatomic ion is equal to the charge on that ion, for example: Na+ = Na(+1), Al3+ = Al(+3), Cl- = Cl(-1).

3 Oxygen has an oxidation number of -2 in all compounds except O2, peroxides and superoxides.

4 Hydrogen has an oxidation number of +1 in all compounds.

5 The sum of the oxidation numbers of the elements in a compound or ion equals the charge on that species.

6 The oxidation number of the elements in a covalent compound can be deduced by considering the shared electrons to belong exclusively to the more electronegative atom (see Box 3.2). Where both atoms have the same electronegativity the electrons are considered to be shared equally. Thus the oxidation numbers of carbon and chloride in CCl4 are +4 and -1 respectively and the oxidation number of chlorine in Cl2 is 0.

Oxidation states are important when predicting the behaviour of elements or compounds. For example, chromium is quite insoluble and non-toxic as chromium (III), while as chromium (VI) it forms the soluble complex anion CrO42-, which is toxic. As with most simple rules, those for oxidation state assignment apply to most but not all compounds.

Since redox half-reactions involve electron transfer, they can be measured electrochemically as electrode potentials, which are a measure of energy transfer (Box 4.8). The reaction:

is assigned an electrode potential (E°) of zero (at standard temperature and pressure) by international agreement. All other electrode potentials are measured relative to this value and are readily available as tabulated values in geochemical texts and on the Internet.

A positive E° shows that the reaction proceeds spontaneously (e.g. the reduction of fluorine gas (oxidation state 0) to fluoride (F-, oxidation state -1). A negative E° shows that the reaction is spontaneous in the reverse direction (e.g. the oxidation of Li to Li+).

To calculate the overall E° for a reaction the relevant half-reactions are combined (regardless of the stoichiometry of the reactions). For example, the reaction of Sn2+ solution with Fe3+ solution involves two half-reactions:

These combine to give a positive E°, which shows that the forward reaction (eqn. 7) is favoured:

The ability of any natural environment to bring about oxidation or reduction processes is measured by a quantity called its redox potential or Eh (see Box 5.4).

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