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Of all the classes of reaction which take place in the environment one the most important is that which involves reduction and oxidation. Since reduction and oxidation always go hand in hand the reactions are referred to as redox reactions.

Their importance lies in the fact that they can influence the movement of materials through the environment and affect the availability of nutrients and pollutants to the biota. For example, certain metals, such as iron and manganese as well as many of the heavy metals, can exist in more than one oxidation, or valency, state. Thus, iron can exist as iron(II), or ferrous, iron in which the valency is two (i.e. the iron forms two bonds with other atoms/ions). It can also exist as iron(III), or ferric, iron in which the valency is three (i.e. the iron forms three bonds with other atoms/ions). Now the compounds of iron(II) are generally soluble in water while the iron(III) compounds tend to be insoluble in water. What is more, iron(II) compounds are not stable in the presence of oxygen and they will be converted into iron(III) compounds. Thus, in the environment, where oxygen is present any iron will exist as iron(III) and will be insoluble in water. On the other hand, when iron(III) compounds enter anoxic conditions the iron will be converted into iron(II) and will form water-soluble compounds. Given that living organisms absorb their nutrient from aqueous solution then iron will only be available to the biota if it is in the iron(II) form. Therefore, the nutrient iron will normally only enter food chains and webs via anaerobic organisms.

Inorganic suspended solids in water bodies are made up principally of iron(III) and manganese(III) oxides. These solid particles will have other species, for example heavy metals, adsorbed on to their surfaces. Bound as they are to these solid particles of iron and manganese oxides the heavy metals are not available to living organisms. However, should the suspended solids enter an anoxic area the iron(III) and manganese(III) oxides will be reduced to iron(II) and manganese(II) compounds which are soluble in the water. The particles dissolve so releasing the adsorbed metal ions which are now available for uptake by living organisms. Sulphur is widely distributed throughout then Earth’s crust in pyritic ores, for example iron(II) pyrites, where it occurs as metal sulphides. During mining operations these ores are exposed to the air when there is a reaction between the ore and oxygen. The result is that the sulphides are oxidised to sulphur oxides which dissolve in any water passing over the rocks to form sulphuric acid. Consequently, tailings emerging from old mines are often acidic and carry with them iron(III) oxide formed as the iron(II) is oxidised as well. This is why the course of the tailings is often stained with rust.

Reduction, Oxidation and Redox Reactions

Redox reactions involve reduction and oxidation, where reduction is the gain of electrons and oxidation is the loss of electrons. Thus, in the following half reactions:

O2  +  4e- ® 2O2-   (1)

 Oxygen         oxide

Fe2+ ® Fe3+  +  e-   (2)

 Iron(II)  iron(III)

 (ferrous iron)    (ferric iron)

the oxygen in (1) has accepted electrons and so has been reduced to oxide. In (2) the iron(II) ions have lost electrons and so have been oxidised to iron(III) ions.

Since reduction and oxidation involve the transfer of electrons it follows that a reduction must always be accompanied by an oxidation, and vice versa. In other words, the electrons accepted by the species being reduced must have been given up by a second species which, consequently, has become oxidised. For example, in the environment reactions (1) and (2) are invariably linked. That is, iron(II) reacts with oxygen as follows:

4Fe2+  +  8OH-  +  O2 ® 2Fe2O3  +  H2O (3)

As is well known, iron metal itself readily reacts with oxygen in exactly the same way, particularly in damp conditions, to form hydrated iron(III) oxide which is better known as rust.

4Fe  +  3O2 ® 2Fe2O3 (4)


The two half reactions involved in this process are:

4Fe ® 4Fe3+  +  12e- (oxidation) (5)


3O2  +  12e- ® 6O2- (reduction) (6)

Note that oxygen is such an electronegative element that when it binds to any other element (other than fluorine) it tends to pull the bonding electrons away from that element; that is it oxidises the element. Conversely, when non-metallic elements bind to hydrogen then, since they are almost invariably more electronegative than the hydrogen, they pull the bonding electrons away from it causing it to become oxidised. The non-metallic elements themselves are, therefore, reduced. For these reasons reduction is sometimes defined as loss of oxygen or gain of hydrogen while oxidation is defined as gain of oxygen or loss of hydrogen. So, consider the following transformations of sulphur:

S  +  O2 ® SO2 (7)

2SO2  +  O2 ® 2SO3 (8)

SO3  +  2H+ ® SO2  +  H2O (9)

SO2  +  2H+ ® S  +  H2O (10)

S  +  2H+ ® H2S (11)

In transforms (7) and (8) sulphur is acquiring oxygens (SO2 = sulphur dioxide; SO3 = sulphur trioxide) and so it is becoming nprogressively more oxidised. Conversely, in transformations (9) and (10) the sulphur is losing oxygens while in (11) it is gaining hydrogen. Therefore, reactions (9), (10) and (11) show sulphur becoming progressively more reduced.

3. Oxygen and Life

Life is dominated by two processes, namely:

 (i) photosynthesis,

 (ii) respiration.

(a) Photosynthesis

This is the process by which light energy is trapped and stored in organic compounds such as the carbohydrate glucose.

6CO2 + 6H2X + hn ® C6H12O6 + 6XO (12)

where hn represents light energy and H2X is a source of hydrogen.

Organisms which can photosynthesise are called producers while those organisms which cannot photosynthesise are consumers. Consumers are dependent on the producers for their energy needs.

The major producers are the green plants and these use water, H2O, as the hydrogen source in the reduction of carbon dioxide:

6CO2 + 6H2O + hn ® C6H12O6 + 6O2 (13)

The evolution of green plants has led to the presence of oxygen in the atmosphere and, consequently, to the evolution of aerobic organisms, which have come to dominate most habitats.

(b) Respiration

This is the process by which energy is released from organic matter (primarily carbohydrates). There are two forms of respiration namely aerobic and anaerobic respiration, with the anaerobic form being known also as fermentation. Aerobic respiration is the exact reverse of photosynthesis since it involves the oxidation of carbon in carbohydrates:

C6H12O6 + 6O2 ® 6CO2 + 6H2O + ENERGY (14)

Note that all organisms respire, even the producers.

Examination of equation 14 shows that during aerobic respiration organisms not only utilise the energy trapped by the producers but they consume free oxygen as well.

4. Aquatic Environments