Metals

Corrosion of metals

The overall driving forces of nature work to return metals to their stable oxidised states, that is, combined with oxygen, sulphates, carbonates, sulphides and chlorides. Unoxidised or native metallic element is produced when metals are unbound from their compounds with oxygen, sulphate, carbonate, sulphide and chloride. For this to happen there must be a sufficient driving force available through a high energy intervention. This intervention can be a carbon reduction or smelting. When metal ores are processed to produce metals, they start to corrode.

The primary property of electrical conductivity of metals is due to the dispersed nature of the electrons in the structure of the metals. When an external voltage is applied to a metal, the electrons flow. This very same useful property is the underlying cause for the corrosion of metals, because the voltage applied to the metal by the presence of oxygen in a moist environment will cause the electrons to flow irreversibly from the metal into the oxygen—to form an oxide coating.

A simple overview of corrosion

The corrosion of metals consists of two separate reactions:

• an oxidation reaction; and

• a reduction reaction.

To explain these reactions, it is necessary to give a simple overview of the structure of atoms. Atoms are made up of a nucleus which contains neutral particles called neutrons and positively charged particles called protons. Electrons, which are negatively charged particles, orbit around the nucleus of the atom. The number and activity of the electrons will determine how readily the atoms will react with other atoms. Many metals, because of the way their molecules are structured, can readily lose electrons. When they do this, they are no longer atoms. They are positively charged and are called ions. Because of the charge, ions are not stable and combine readily to achieve a stable, electrically neutral state.

An oxidation reaction is one in which an atom loses electrons. This can be represented very simply by the equation:

         MMⁿ⁺ +ne^-

where 'n' represents the number of electrons lost

For example, copper—Cu—can be put into this equation. In an oxidation reaction:

• Cu Cu⁺+ e^-

It can be oxidised further:

• Cu⁺ Cu²⁺ + 2e^-

Iron and zinc are other examples of polyvalent metals.

In reduction reactions, atoms gain electrons. A simple representation of this would be:

M ⁿ⁺+ne ^- M

where 'n' represents the number of electrons needed

These electrolytic reactions are used to produce solid metals from their ionic solutions. The negative ions can be supplied by a range of materials. For example, if the metal object is in a seaside location, chloride ions—Cl^-—will combine readily with the metal ions.

They will also combine with:

• sulphides—SO₃^-—sulphates—SO₄^2—nitrates—NO₃^2-—from atmospheric pollutants; and

• oxygen.

If the metal combines with oxygen, it forms a metal oxide on the surface of the metal. If this metal oxide is continuous, then the overall corrosion rate of the underlying metal will slow down and it will become passivated or protected.

Corrosion cells

Corrosion cells are small areas on metal objects where electrical differences are set up. Electrons flow between the charged areas, just as an electrical current flows between the positively and negatively charged electrodes of a battery.

A corrosion cell is an electrochemical cell which acts very much like a battery. The corrosion of metals consists of two separate reactions:

• oxidation. The oxidation reactions are called anodic reactions; and

• reduction. The reduction reactions are called cathodic reactions.

In an electrochemical cell the anodic, oxidation, half of the cell produces electrons as the metal is oxidised, while at the cathodic half of the cell, reduction occurs. The electrons are taken and held by the oxidising agent, which in aerated environments is oxygen.

In a corrosion cell, these reactions can continue in a cycle. The localised corrosion activity causes pitting in the metal.

The rate at which the electrons move out of the metal and across into the oxygen molecules is the principal factor controlling the overall corrosion rate.

Organic acids—formed by the oxidation of oils and fats—are capable of attacking metals which rely on a protective oxide coating to produce a good corrosion resistance. To prevent this type of damage, avoid direct contact between the object and the source of the organic material. Some examples of this type of damage are leather objects with copper fittings. The gradual deterioration of old candle wax in leather-lubricating oils leads to organic acids penetrating the protective copper oxide film, and reacting with the underlying metal—to form outgrowths of bright green organic copper compounds.

Human sweat on metal objects causes corrosion. Bacterial reactions with sweat can produce sulphides as metabolic by-products, and convert inherently inert sulphate ions into reactive sulphide ions.

Uneven coatings of oil—from sweaty hands for instance—can alter the ease of access of oxygen to metal surfaces. This has two major effects. It hinders the formation of passivating layers of corrosion. It also alters the relative reactivities of areas of the metals; and so it causes one part of the metal to corrode at the expense of another.

Acids

Inorganic acids such as hydrochloric acid—derived from the decay of plastics like polyvinyl chloride—and nitric and sulphuric acids—derived from air pollution—will attack metals which are either in the same storage environment as the plastic or in the open air.

Anything that prevents direct contact between the metal surface and acidic solutions helps to prolong the life of the object. Therefore, vapour phase inhibitors, lacquers, waxes and other coatings minimise the damage from air pollution. The filtering of external air also greatly helps to minimise corrosion damage.

Sulphide pollutants

Normally unreactive metals such as copper and silver can suffer significant corrosion in the presence of sulphide ions. Common sources of sulphide ions are:

• hydrogen sulphide—H₂S—from the anaerobic decay of plant material; and

• carbonyl sulphide—COS—from the degradation of sulphur-containing proteins, such as those found in wool.

Base metals such as zinc and tin are also significantly affected by sulphide pollution and/or contamination. Small concentrations of sulphur compounds in damp, oxygenated conditions cause corrosion. The resulting metal sulphides can often form a protective patina, as in the case of tin sulphides which protect pewter objects.

Adsorption of the sulphur-containing species is an essential step in the overall corrosion process and any factor which inhibits adsorption helps minimise attack on the metal. Therefore, adsorption of organic materials, such as vapour phase corrosion inhibitors, greatly decreases the corrosion rate.