Before you read this, I suggest you read posts 16.33 and 16.39.

When iron rusts, it reacts (post 16.33) with oxygen molecules (post 16.30) to form a mixture of iron (II) ions (Fe²⁺), iron (III) ions (Fe³⁺) and oxide ions (O^2-). To find out more about ions, see post 16.39. The resulting mixture is sometimes called red iron oxide (because of its colour) and is often represented as Fe₃O₄, because for every Fe²⁺ ion, there are two Fe³⁺ ions (a total of 3 iron ions) and four O^2- ions. The overall electrical charge of the mixture is then 2 + (2 X 3) – (4 X 2) = 0; so the mixture has no overall electrical charge. Red iron oxide occurs naturally in a form of iron ore called magnetite.

So, when iron, rusts it reacts with oxygen – reaction with oxygen is called oxidation. The reaction can be represented (post 16.33) by

3Fe + 2O₂ → Fe₃O₄.

When iron is extracted from its ore it is usually heated, in a blast furnace, with a form of carbon called coke; coke is coal with most of its tar removed. The carbon reacts with the oxygen in magnetite to form carbon dioxide, leaving iron. The reaction can be represented by

Fe₃O₄ + 2C → 3Fe + 2CO₂.

In this reaction, oxygen has been removed from magnetite – a reaction in which oxygen is removed is called reduction.

Now let’s look at what’s happening in more detail. When iron rusts, the oxidation process makes Fe²⁺ and Fe³⁺ ions. When these ions are formed, an iron atom gains electrical charge (post 16.25) which is another way of saying that it loses electrons (post 16.39). When magnetite is reduced, Fe²⁺ and Fe³⁺ ions lose electrical charge, which is another way of saying they gain electrons.

So, we define oxidation more generally as atoms or ions gaining charge (losing electrons) and reduction more generally as atoms or ions losing charge (gaining electrons).

Now let’s look at the production of iron from its ore, in a blast furnace, using our original definitions of oxidation and reduction. As iron oxide is reduced, oxygen is reacting with carbon to form carbon dioxide. Carbon dioxide does not contain ions – it is a neutral molecule. How can we reconcile this observation with our second definition of oxidation – a reaction in which an atom gains charge? When carbon forms carbon dioxide, it shares some of its electrons with oxygen atoms (post 16.30). It has given up some of its share in these negatively charged electrons and so, in a sense, it has gained some positive charge.

In this reaction, carbon is reducing magnetite – it is acting as a reducing agent. But the magnetite is also oxidising the carbon – it is acting as an oxidising agent. Oxidation and reduction are occurring simultaneously – magnetite is being reduced and carbon is being oxidised.

Did you ever do an experiment at school in which you dropped small pieces of the metal zinc into dilute acid? The zinc dissolves and bubbles of hydrogen gas are produced. Acids contain hydrogen ions (post 17.49), H⁺ (post 16.39). In the reaction of acids with zinc, the positive charge of the hydrogen ions transfers to the zinc atoms to form zinc ions, Zn²⁺. The zinc ions dissolve in water because it has a high dielectric constant (post 16.39). The hydrogen atoms combine to form hydrogen molecules (post 16.30). The reaction can be represented by

Zn +2H⁺ → Zn²⁺ + H₂.

The zinc atoms have gained charge – so zinc is oxidised in this reaction. Hydrogen has lost charge – it is reduced.

The positive charges of H⁺ the and Zn²⁺ ions must be balanced by negatively charged ions – because solutions are not electrically charged. If you dropped the zinc into hydrochloric acid, for example, there would be chloride, Cl^–, ions present that were unchanged by the reaction. Showing all types of ions present, we get

Zn +2H⁺ + 2Cl^– → Zn²⁺ + 2Cl^– + H₂.

So, you may have been taught to represent this reaction as

Zn + 2HCl → ZnCl₂ + H₂

even though it is a reaction between ions and not a reaction between molecules.

In this reaction, electrical charge flows from hydrogen to zinc. So zinc must be at a lower electrical potential than hydrogen – otherwise the charge would not have flowed in this direction (post 17.44).

Oxidation and reduction reactions provide another example of charges moving when a potential difference exists – just as they do in electrical circuits.

 

Related posts

19.1 Hard water
18.1 pH
17.50 Alkalis
17.49 Acids
17.48 Moles
16.39 Ions
16.33 Chemical reactions