Figure 3-107. Iron corrosion (rusting).
In accord with the chemical equation in Figure 3-107 and, as follows from independent experiments, iron rusting occurs only in the presence of both
O. In pure oxygen but in the absence of water, iron does not corrode. In the presence of water but in the absence of O2
, iron does not corrode, either. Corrosion of iron and steel is a complex chemical and electrochemical process that is accelerated by certain impurities in the iron and in the water and air.
- Impurities in the water
can and do speed up corrosion. Those living by the sea know well how salty seawater affects the production of rust. The presence of ionic species in the water makes it a conductor of electricity, which favors the formation of iron ions and helps electrons move around.
- Air pollutants
such as SO2
as well as CO2
are the reason for acid rains
. If present in the air, these oxides dissolve in droplets of rainwater to form acids (H2
) that corrode iron, although in this case it is obviously iron salts, not iron hydroxides and/or oxides, that are formed.
- Impurities in iron
usually facilitate its corrosion. While high purity iron is known to rust slowly, corrosion is strongly promoted when the iron is in contact with another metal. If that metal is a weaker donor of electrons than iron (Figure 3-50), then the iron rusts faster, while the other metal stays intact. Conversely, if the metal touching the iron is a better electron donor than iron, it is that "sacrificial" metal that will corrode. These effects are electrochemical in nature.
To avoid iron rusting, the iron surface must be protected from air and moisture. One of the ways of doing that is to block access of water and air to iron by coating its surface with a special paint or by plating it with another metal. The two most widely used metals for plating onto iron are zinc (galvanization
) and tin (tinning
). While both metals prevent corrosive substances from reaching the underlying iron, there is a big difference between using Zn and Sn for rust protection. Galvanization
. Being a better donor of electrons than iron, zinc corrodes on the galvanized steel first. Even if the iron underneath the Zn film gets exposed due to a dent or a scratch, the remaining zinc keeps protecting it by serving as the sacrificial metal. The iron begins to rust only after all of the zinc has been oxidized. Tinning
is used to make tin-coated sheets of steel, known as tinplate. Tin cans are made of tinplate. The tinning protects the underlying iron very well as long as the tin layer is intact. However, if the tin coating gets damaged, the exposed iron rusts faster than in the absence of tin because a Fe atom is a better donor of electrons than a Sn atom (Figure 3-50).
A powerful method to avoid rusting is alloying iron/steel with another metal/substance that can form a protective film on the surface. Stainless steel
is a highly corrosion-resistant alloy of steel with chromium. The chromium on the surface of stainless steel is air-oxidized to chromium oxide (Cr2
) that forms a thin film. This film protects the underlying metal as efficiently as the surface layer of Al2
protects aluminum (see previous section). To ensure the sufficient strength, thickness, and uniformity of the chromium oxide film, the content of chromium in stainless steel must be 11% or higher.
A remarkable example of non-rusting iron is the famous Iron Pillar of Delhi
(Figure 3-108) that has successfully resisted corrosion since it was built approximately 1,600 years ago. The secret of such powerful corrosion resistance, as we now know, lies in the presence of approximately 0.25% phosphorus
in the iron the pillar is made of. The phosphorus and iron create a uniform film of ferric hydrogen phosphate hydrate that blocks access of air and moisture to the underlying iron atoms. Unfortunately, alloying steel with phosphorus in quantities required for good enough rust protection makes the steel too brittle and insufficiently ductile for many uses.