Corrosion Theory

Humans have most likely been trying to understand and controlcorrosion for as long as they have been using metal objects. The most important periods of prerecorded history are named for the metals that were used for tools and weapons (Iron Age, Bronze Age). With a few exceptions, metals are unstable in ordinary aqueous environments. Metals are usually extracted from ores through the application of a considerable amount of energy. Certain environments offer opportunities for these metals to combine chemically with elements to form compounds and return to their lower energy levels.

Corrosion is the primary means by which metals deteriorate. Most metals corrode on contact with water (and moisture in the air), acids, bases, salts, oils, aggressive metal polishes, and other solid and liquid chemicals. Metals will also corrode when exposed to gaseous materials like acid vapors, formaldehyde gas, ammonia gas, and sulfur containing gases.

Corrosion specifically refers to any process involving the deterioration or degradation of metal components. The best known case is that of the rusting of steel. Corrosion processes are usually electrochemical in nature, having the essential features of a battery. When metal atoms are exposed to an environment containing water molecules they can give up electrons, becoming themselves positively charged ions, provided an electrical circuit can be completed. This effect can be concentrated locally to form a pit or, sometimes, a crack, or it can extend across a wide area to produce general wastage. Localized corrosion that leads to pitting may provide sites for fatigue initiation and, additionally, corrosive agents like seawater may lead to greatly enhanced growth of the fatigue crack. Pitting corrosion also occurs much faster in areas where microstructural changes have occurred due to welding operations.

Corrosion is the disintegration of metal through an unintentional chemical or electrochemical action, starting at its surface. All metals exhibit a tendency to be oxidized , some more easily than others. A tabulation of the relative strength of this tendency is called the galvanic series. Knowledge of a metal's location in the series is an important piece of information to have in making decisions about its potential usefulness for structural and other applications.

The corrosion process (anodic reaction) of the metal dissolving as ions generates some electrons, as shown above, that are consumed by a secondary process (cathodic reaction). These two processes have to balance their charges. The sites hosting these two processes can be located close to each other on the metal's surface, or far apart depending on the circumstances. This simple observation has a major impact in many aspects of corrosion prevention and control, for designing new corrosion monitoring techniques to avoiding the most insidious or localized forms of corrosion. (more advanced reading)

The electrons (e^- in this figure) produced by the corrosion reaction will need to be consumed by a cathodic reaction in close proximity to the corrosion reaction itself. The electrons and the hydrogen ions react to first form atomic hydrogen, and then molecular hydrogen gas. If the acidity level is high (low pH), this molecular hydrogen will readily become a gas as it is demonstrated by exposing a strip of zinc to a sulfuric acid solution.

As hydrogen forms, it could inhibit further corrosion by forming a very thin gaseous film at the surface of the metal. This "polarizing" film can be effective in reducing water to metal contact and thus in reducing corrosion. Yet it is clear that anything which breaks down this barrier film tends to increase the rate of corrosion. Dissolved oxygen in the water will react with the hydrogen, converting it to water, and destroying the film.

High water velocities tend to sweep the film away, exposing fresh metal to the water. Similarly, solid particles in the water can brush the hydrogen film from the metal. Other corrosion accelerating forces include high concentrations of free hydrogen ions (low pH) which speed the release of the electrons, and high water temperatures, which increase virtually all chemical reaction, rates. Thus a variety of natural and environmental factors can have significant effects on the corrosion rate of metals, even when no other special conditions are involved.

Simple experiments to illustrate the theory:

Corrosion current, Corrosion noise, Galvanic coupling, Rust preventive testing