Chemguide: Core Chemistry 14 - 16

The Rusting of Iron

This page looks at the rusting of iron, and the various ways it can be prevented or slowed down.

You will need to understand the page about oxidation and reduction in the Reactivity Series section in terms of electron transfer.

Rusting as a form of corrosion

Corrosion happens on the surface of a metal when the metal atoms react to form compounds such as oxides or carbonates or more complicated substances. Rusting is an example of this.

Rusting is the special name given to the corrosion of iron. Many metals react with oxygen to form a layer of metal oxide on the surface. The very strong, very thin, layer of aluminium oxide on the surface of aluminium helps to protect the aluminium.

Rust can be thought of as a hydrated iron(III) oxide, Fe₂O₃.xH₂O. "x" shows a variable amount of water attached to the iron(III) oxide.

The problem with iron is that, unlike aluminium, the rust layer is very porous and flaky. So rather than protect the iron underneath, it still allows air and water to reach the iron.

The conditions for the rusting of iron

The video below gives you all the information you want. It is a bit slow to start with, but is clearer than anything else I have been able to find.

And this short bit of time-lapse photography shows a steel plate rusting under more everyday conditions.

When iron rusts:

• You need the presence of both air and water.

• Rusting is accelerated by the presence of dissolved ionic substances like salt.

The photo shows the effect of long exposure to sea water on a steel boat.

The chemistry of the rusting of iron

This is surprisingly complicated, and to understand it properly, you will need to do chemistry for a few more years. All I can do here is to give a hint of the main processes, but without detail.

The main thing is that rusting involves the initial formation of Fe²⁺ ions as the iron atoms lose electrons:

Fe     Fe²⁺ + 2e^-

The iron is being oxidised by loss of electrons (OIL RIG).

The electrons are picked up by oxygen and water molecules to form hydroxide ions. These react with the iron(II) ions to give iron(II) hydroxide, which is further oxidised and rearranged to give rust. You do not need to know this at this level.

But you should know that rusting starts by the loss of electrons from iron atoms to give Fe²⁺ ions.

Preventing rusting

Stopping air and water getting at the iron

You can do this by

• painting the metal,

• covering the metal with oil or grease,

• coating the iron with another metal like the tin traditionally used in "tin" cans - steel cans with a thin coating of tin.

In all of these cases, if the coating gets damaged, the iron underneath will rust.

Alloying the iron

Stainless steel is iron mixed with chromium and usually other elements such as nickel. There are actually a whole family of stainless steels.

The photo shows two of my spades - one is stainless steel and the other has an ordinary steel blade, probably containing a small amount of carbon to make it harder.

Neither spade has had any special treatment. This is how they came out of my shed.

The iron in stainless steel doesn't rust because of the very strongly attached layer of chromium oxide which protects the whole surface. This is similar to the protection given to aluminium by the strong layer of aluminium oxide on its surface.

Galvanising - sacrificial protection

Galvanised iron is iron which is covered with a thin layer of zinc. When galvanised iron is new, the zinc just serves as a barrier, but it still keeps on protecting the iron even after it has been scratched or otherwise knocked about.

The watering can in the next photo has been outdoors in the rain for years, but is showing no signs of rust.

In fact, zinc doesn't have to cover the whole surface to be protective. Blocks of zinc are attached to steel hulls of boats to stop them rusting.

Zinc is more reactive than iron, and so oxidises to zinc ions more readily than iron does.

Zn(s)    Zn²⁺(aq) + 2e^-

Those electrons flow into the iron and prevent the similar reaction happening to the iron.

Fe(s)    Fe²⁺(aq) + 2e^-

The extra electrons react with any iron(II) ions formed and drive them back to iron metal.

This is called sacrificial protection. The zinc blocks are sacrificed to protect the steel hull. It is much easier to replace the zinc blocks than repair the hull.