An Introduction to Environmental Chemistry

78 Chapter Four

Box 4.3 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:

eqn. 1

can be considered to consist of two half-

reactions:

eqn. 2

eqn. 3

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), Al^3 +=Al(+3), Cl-=

Cl(–1).

3 Oxygen has an oxidation number of –2 in

all compounds except O 2 , 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 CCl 4 are +4 and –1

respectively and the oxidation number of

chlorine in Cl 2 is 0.

312 6Oe O 2 +Æ--^2

412 4Fe-Æe-+Fe^3

432 Fe()metal+ÆO 22 ()g Fe O 3 ()s

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 CrO 42 - , 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:

eqn. 4

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 Sn^2 +

solution with Fe^3 +solution involves two half-

reactions:

eqn. 5

eqn. 6

These combine to give a positive E°, which

shows that the forward reaction (eqn. 7) is

favoured:

eqn. 7

E° for this reaction=0.77 – 0.15=0.62 V.

The ability of any natural environment to

bring about oxidation or reduction processes

is measured by a quantity called its redox

potentialor Eh (see Box 5.4).

22 Fe^32 +++Æ +Sn Fe^24 ++Sn

Sn^42 +-+Æ 20 e Sn+E∞=. 15 V tabulated value( )

Fe^32 +-+Æe Fe+E∞= 077. V tabulated value( )

(^12) He H 2 -Æ-+