Rusting of iron refers to the formation of rust, a mixture of iron oxides, on the surface

of iron objects or structures. This rust is formed from a redox reaction between oxygen
and iron in an environment containing water (such as air containing high levels of
moisture). The rusting of iron is characterized by the formation of a layer of a red, flaky
substance that easily crumbles into a powder.

Rusting of Iron
This phenomenon is a great example of the corrosion of metals, where the surfaces of
metals are degraded into more chemically stable oxides. However, the term ‘rusting’ is
generally used to refer to the corrosion of objects made of iron or iron-alloys.

What is the Chemistry Behind the Rusting of Iron?

The exposure of iron (or an alloy of iron) to oxygen in the presence of moisture leads to
the formation of rust. This reaction is not instantaneous, it generally proceeds over a
considerably large time frame. The oxygen atoms bond with iron atoms, resulting in the
formation of iron oxides. This weakens the bonds between the iron atoms in the
object/structure.

The reaction of the rusting of iron involves an increase in the oxidation state of iron,
accompanied by a loss of electrons. Rust is mostly made up of two different oxides of
iron that vary in the oxidation state of the iron atom. These oxides are:
 1. Iron(II) oxide or ferrous oxide. The oxidation state of iron in this compound is +2
and its chemical formula is FeO.
2. Iron(III) oxide or ferric oxide, where the iron atom exhibits an oxidation state of
+3. The chemical formula of this compound is Fe 2O3.

Oxygen is a very good oxidizing agent whereas iron is a reducing agent. Therefore, the
iron atom readily gives up electrons when exposed to oxygen. The chemical reaction is
given by:

Fe → Fe2+ + 2e–

The oxidation state of iron is further increased by the oxygen atom when water is
present.

4Fe2+ + O2 → 4Fe3+ + 2O2-

Now, the following acid-base reactions occur between the iron cations and the water
molecules.

Fe2+ + 2H2O ⇌ Fe(OH)2 + 2H+

Fe3+ + 3H2O ⇌ Fe(OH)3 + 3H+

The hydroxides of iron are also formed from the direct reaction between the iron cations
and hydroxide ions.

O2 + H2O + 4e– → 4OH–

Fe2+ + 2OH– → Fe(OH)2

Fe3+ + 3OH– → Fe(OH)3

The resulting hydroxides of iron now undergo dehydration to yield the iron oxides that
constitute rust. This process involves many chemical reactions, some of which are listed
below.

1. Fe(OH)2 ⇌ FeO + H2O

2. 4Fe(OH)2 + O2 + xH2O → 2Fe2O3.(x+4)H2O
3. Fe(OH)3 ⇌ FeO(OH) + H2O
4. 2FeO(OH) ⇌ Fe2O3 + H2O
One similarity between all the chemical reactions listed above is that all of them are
dependent on the presence of water and oxygen. Therefore, the rusting of iron can be
controlled by limiting the amount of oxygen and water surrounding the metal.

Why is Rusting an Undesirable Phenomenon?

Rusting causes iron to become flaky and weak, degrading its strength, appearance and
permeability. Rusted iron does not hold the desirable properties of iron. The rusting of
iron can lead to damage to automobiles, railings, grills, and many other iron structures.

The collapse of the Silver Bridge in 1967 and the Mianus River bridge in 1983 is
attributed to the corrosion of the steel/iron components of the bridge. Many buildings
made up of reinforced concrete also undergo structural failures over long periods of
time due to rusting.

Rusted iron can be a breeding ground for bacteria that cause tetanus. Cuts from these
objects that pierce the skin can be dangerous.

Since rusting occurs at an accelerated rate in humid conditions, the insides of water
pipes and tanks are susceptible to it. This causes the pipes to carry brown or black
water containing an unsafe amount of iron oxides.

Factors that Affect the Rusting of Iron

Many factors speed up the rusting of iron, such as the moisture content in the
environment and the pH of the surrounding area. Some of these factors are listed
below.

 Moisture: The corrosion of iron is limited to the availability of water in the

environment. Exposure to rains is the most common reason for rusting.
 Acid: if the pH of the environment surrounding the metal is low, the rusting
process is quickened. The rusting of iron speeds up when it is exposed to acid
rains. Higher pH inhibits the corrosion of iron.
 Salt: Iron tends to rust faster in the sea, due to the presence of various salts.
Saltwater contains many ions that speed up the rusting process via
electrochemical reactions.
 Impurity: Pure iron tends to rust more slowly when compared to iron containing a
mixture of metals.
The size of the iron object can also affect the speed of the rusting process. For example,
a large iron object is likely to have small deficiencies as a result of the smelting process.
These deficiencies are a platform for attacks on the metal from the environment.

How can Rusting be Prevented?

Iron and its alloys are widely used in the construction of many structures and in many
machines and objects. Therefore, the prevention of the corrosion of iron is very
important. Some preventive methods are listed below.

Alloys that are Resistant to Rusting

Some alloys of iron are rust-resistant. Examples include stainless steel (which features a
layer of chromium(III) oxide) and weathering steel.

COR-TEN steel rusts at a relatively slower rate when compared to normal steel. In this
alloy, the rust forms a protective layer on the surface of the alloy, preventing further
corrosion.

Galvanization
 Galvanization is the process of applying a protective layer of zinc on a metal. It is
a very common method of preventing the rusting of iron.
 This can be done by dipping the metal to be protected in hot, molten zinc or by
the process of electroplating.
 Zinc is a relatively cheap metal that sticks to steel easily. It also offers cathodic
protection to the iron surface by acting as an anode. The zinc layer is corroded
instead of the iron due to this.
 The disadvantages of galvanization are that it only provides protection from
corrosion for a limited amount of time since the zinc layer is eaten up in the
process. It is not very effective in highly corrosive areas (where cadmium coating
can be used instead).

Cathodic Protection
 Providing the metals with an electric charge can help inhibit the electrochemical
reactions that lead to rusting.
 This can be done by making the iron/steel a cathode by attaching a sacrificial
anode to it.
 This sacrificial anode must have an electrode potential that is more negative than
that of iron.
 Metals that are commonly used as sacrificial anodes are magnesium, zinc, and
aluminium. Once they are corroded away, they must be replaced in order to
protect the iron/steel.
Coatings
Many types of coatings can be applied to the surface of the exposed metal in order to
prevent corrosion. Common examples of coatings that prevent corrosion include paints,
wax tapes, and varnish.

Smaller objects are coated with water-displacing oils that prevent the rusting of the
object. Many industrial machines and tools made of iron are coated with a layer of
grease, which lubricates the metal to reduce friction and prevents rusting at the same
time