Corrosion

Corrosion is the name of the reaction that occurs between a metal and oxygen. Although, that would also classify it as combustion, because these reactions happen so slowly, over such a long time, it seems odd to lump them together with fires and explosions. But, whether odd or not, that’s what these reactions are. 

Corrosion Reactions

Here is the reaction for iron:

\(2~Fe + O_2 \rightarrow 2~FeO\)

Iron will then undergo another reaction that takes it to the +3 state, but for the purposes of this page, we’ll focus on the first reaction.

The reaction between iron and oxygen is very slow, but there is a catalyst for this reaction - water. Remember that one of the ways that a catalyst can speed up a reaction is to provide an alternate pathway for the reaction which has a lower activation energy.

When water is present, corrosion can happen as a mechanism whose first step is a pair of half reactions. Oxygen in the air, can react with the water and electrons to make hydroxide ions.

\(O_2 + 4~e^{-1} + 2~H_2O \rightarrow 4~OH^{-1}\)

The source for those electrons is the iron, which loses electrons to become \(Fe^{+2}\) ions:

\(Fe \rightarrow 2~e^{-1} + Fe^{+2}\)

The iron II ions and the hydroxide ions can then combine to give iron II hydroxide (\(Fe(OH)_2\)).

Doubling the iron reaction (to match the number of electrons) and then adding the two half-reactions give us:

\(2~Fe + O_2 + 2~H_2O \rightarrow 2~Fe(OH)_2\)

Iron II hydroxide can then decompose into water and iron II oxide:

\(Fe(OH)_2 \rightarrow FeO + H_2O\)

Putting It ALL Together

Let’s put all of those steps together as a mechanism:

\(Step~1~~~~2~Fe + O_2 + 2~H_2O \rightarrow 2~Fe(OH)_2\)
\(Step2~~~~2~Fe(OH)_2 \rightarrow 2~FeO + 2~H_2O\)
\(Overall~~2~Fe + O_2 \rightarrow 2~FeO\)

Which is, of course, the original reaction at the top of the page. In other words, we have accomplished the same reaction, but through a different mechanism using water as the catalyst.

Because the activation energy for these reactions is much lower than it is for the overall reaction occurring in a single step, the catalyzed reaction occurs much faster. This is why cars in the desert areas of the Southwestern U.S. rust much more slowly than do cars in the wet Southeastern U.S. Adding salt to the water makes it a better conductor for charge which reduces the activation energy even more, which is why cars in the Northeastern U.S. (where they salt the roads in the winter) rust even faster.

Protecting From Corrosion

Separation

Once we understand that corrosion is a reaction with oxygen, catalyzed by water, protection from corrosion becomes clear. Let’s think about our cars. We paint our cars and wax them because the pain forms an airtight and watertight seal over the metal. In addition, wax is non-polar, so water beads up and rolls off of the surface, rather than staying in contact with the metal. Both of these help prevent corrosion.

Aluminum

Another way we avoid having things corrode is to change the metal from which they are made. Aluminum is a more reactive metal than iron, but aluminum oxide is extremely insoluble and it forms an airtight and watertight seal over the metal. In other words, aluminum corrodes a TINY amount and then that corrosion prevents further corrosion.

Sacrificial Anode

A last way we have of preventing (or at least slowing) corrosion, is the use of a sacrificial anode. One of the most common examples of this is galvanized steel.

Galvanized Steel is steel (mostly iron) that has been plated with a thin layer of zinc. When oxygen reacts with the water on galvanized steel, there are now two possible reactions that could supply those electrons. Here they are, listed with their voyages:

\(Fe \rightarrow Fe^{+2} + 2~e^{-1}~~~~~~~~~~0.44~V\)

\(Zn \rightarrow Zn^{+2} + 2~e^{-1}~~~~~~~~~~0.76~V\)

Both of these are positive, which means both could happen readily, but nature will favor the easier (more positive) half-reaction. That means that as long as ANY zinc remains, the electrons will be taken from the zinc and the iron will remain untouched. The zinc “sacrifices itself” to save the iron.