Corrosion

1. Introduction
Corrosion results in huge losses to industrial countries every year. An SSINA - Battelle study estimated that the combined economic loss to the US in 1995 due to metal and alloy corrosion (excluding non-metallics) was $300 billion. Obviously engineers and designers must have an understanding of corrosion and of methods to minimise the losses.

Corrosion is the deterioration and loss of material due to chemical attack. The conditions that promote corrosion involve both chemical and electronic changes and are always with us.

2. Corrosion by Solution
The simplest corrosion is chemical solution, eg dissolving sugar in water. Most engineering materials are not particularly soluble, but some polymers, such as some types of rubber, can be dissolved by strong solvents such as petrol.
There are some generalisations that can be made about solutions:

Small molecules and ions dissolve most readily - apparent exceptions to this involving polymers occur where the polymer is easily de-polymerised - and the material goes into solution as small molecules.
Solution occurs more readily when the solvent and solute are structurally similar. Within these categories, similarity of solvent and solute structures produces greater solubility. Copper is more soluble in liquid zinc than liquid lead. Copper and nickel are similar sized molecules and are totally mutually soluble.
The presence of more than one solute may produce greater solubility than one alone. As an example, calcium carbonate (limestone) is nearly insoluble in pure water, but the addition of CO₂ to form coarbonic acid in the water greatly increases the CaCO₃ solubility.
The rate of solution increases with temperature.

3. Electro-chemical Oxidation
The most common type of corrosion involves the electro-chemical process of metal oxidation. Strictly 'oxidation' is the removal of electrons from an artom.
The oxidation of iron to ferrous ions:

Fe gives Fe²⁺ + 2e^- .... (i)
The oxidation of ferrous to ferric ions:
Fe²⁺ gives Fe³⁺ + e^- .... (ii) The combination of chemical reaction and electron release causes other reactions, such as rust formation. Rust may be considered to be ferric hydroxide and is formed by the overall reaction: 4Fe + 3O₂ +6H₂O gives 4Fe(OH)₃ For iron to rust these reactions must occur and both oxygen and moisture must be present. Iron will not rust in oxygen free water nor in an atmosphere containing only oxygen. The amount of moisture required to produce rusting is quite small, but at a relative humidity of less than 60% corrosion of most metals is slow. However significant daily changes in temperature, leading to condensation, can cause rapid corrosion even at low humidity.

Different metals have different 'oxidation potentials' ie the energy required to remove electrons varies from metal to metal. For example electrons are removed from iron when both oxygen and water are present and they are removed from aluminium atoms when chloride ions are present.

4. Electrode Potential
Most corrosion occurs through the interaction of the two processes of 'solution' and 'oxidation'. With slight modifications the mechanism of corrosion of iron described by the two equations above, (i) and (ii) may be applied to all metals and even non-metals.

As iron goes into solution, excess electrons are produced and in static liquids, equilibrium is usually reached quickly as ions and electrons soon combine at the same rate at which they form.
The production of ions and electrons builds up an electrical potential called an 'electrode potential', which depends upon the nature of the metal and the nature of the solution. All metal does not oxidise to ions and electrons with equal facility. For example, atoms along a grain boundary are less stably located than those in the crystal lattice, so they ionise more readily. Also the reaction (i) will produce equilibrium with a higher electrode potential if the metal ions enter a solution in which they are relatively stable. Positive iron ions are more stable in a uniformly Cl^- concentrated solution than in a uniformly dilute Cl^- solution.

To measure the electrode potential of any material (and therefore its tendency to corrode) the voltage difference between the metal and a standard hydrogen electrode need to be found. With hydrogen the equation of equilibrium is given by:

H₂ gives (either to the left or to the right) 2H⁺ + 2e^-

The potential difference between the iron and hydrogen electrodes is + 0.44 volts.
The alkali and alkaline earth metals, which hold their outer shell electrons rather loosely, show a greater potential difference than iron. Conversely, the noble metals, such as silver, gold and platinum, produce fewer electrons than hydrogen and are lower on the potential scale.

Table - Electrode Potentials of Metals - 25^oC, Molar solutions of metal ions:

Metal ion	Potential
Li⁺ (base)	+2.96 (anodic)
K⁺	+2.92
Ca²⁺	+2.90
Na⁺	+2.71
Mg²⁺	+2.40

Al³⁺	+1.70
Zn²⁺	+0.76
Cr²⁺	+0.56
Fe²⁺	+0.44
Ni²⁺	+0.23

Sn²⁺	+0.14
Pb²⁺	+0.12
Fe³⁺	+0.045
H⁺	0.000 (reference)
Cu²⁺	-0.34

Cu⁺	-0.47
Ag⁺	-0.80
Pt⁴⁺	-0.86
Au⁺ (noble)	-1.50 (cathodic)

5. Galvanic Corrosion
When a single electrode is immersed in a static liquid, the dissolution of ions quickly stops due to equilibrium being reached. However if two electrodes with different electrode potentials are immersed in the conducting liquid (eg iron and hydrogen) and connected, then electrons flow into the external connecting circuit, leaving Fe²⁺ to go into solution at the anode and electrons from the external circuit enter hydrogen terminal (cathode) and combine with H⁺ ions from the water to produce gaseous hydrogen.

In practice a number of additional reactions and constraints are operating. Although corrosion occurs at the anode, rust is more often found at the cathode, because electrons and smaller Fe³⁺ ions move to the cathode more readily than the larger OH^- ions can move to the anode.

6. Electroplating
If a piece of copper and a piece of iron are immersed in an conducting liquid and the iron is used as a cathode and the copper as an anode, and electrons are forced to the cathode (by a battery) then the corrosion process is reversed and electro-plating occurs (electro-plating is the opposite of corrosion). Corrosion always occurs at the anode and electro-plating always occurs at the cathode.

7. Types of Galvanic Cell
There are three different groups of galvanic cell:

composition cells
stress cells
concentration cells

Each produces corrosion because one half of the couple acts as an anode and the other half, with a lower electrode potential, serves as the cathode.
i) Composition Cells
A composition cell may be established between any two dissimilar materials. The metal higher in the electromotive series acts as the anode. For example with a sheet of galvanised steel, the zinc coating acts as the anode and protects the underlying iron even if the surface is not completely covered, because the exposed iron is the cathode and does not corrode. So long as zinc remains it provides protection to adjacent exposed iron. This is made us of with ships and barges where zinc anodes are fixed at intervals to the steel hull.
Conversely a tin coating on iron or steel only provides protection so long as the surface is completely covered. Since tin is only slightly above hydrogen in the electromotive series the rate of corrosion of the tin is slow. However if the tin coating becomes punctured, the tin becomes the cathode. The exposed iron, which is above tin in the electromotive series, acts as the anode and the resulting galvanic couple corodes the iron. Since the small anodic area must supply electrons to a large cathode surface, very rapid local corrosion can result.

There is no size limit to galvanic cells so these can be set up in two phase alloys when they are exposed to an electrolyte. Often the potential difference between two similar phases is quite small. Ferrite and iron carbide have electrode potential values sufficiently close together for plain carbon steel to normally have a lower corrosion rate than a steel brass combination.

Heat treatment may affect the corrosion rate by altering the microstructure of the metal. After quenching a thin section carbon steel component the structure is a single phase, martensite. After tempering ferrite and carbide produced provide many galvanic cells and the corrosion rate is increased. Tempering at higher temperatures, gives agglomeration of the carbides and reduces the number of galvanic cells, which gives a significant reduction in corrosion rates.

David J Grieve, 28th October 2003.