Basics of Corrosion

Updated: Jan 6

Excerpted from Corrosion Basics

Published by the National Association of Corrosion Engineers,

1440 South Creek Drive

Houston, Texas

77084

Supplied by Norman Nock, British Car Specialists


Introduction


It is imperative to learn the basic mechanisms of the corrosion process in order to properly analyze corrosion problems and arrive at effective solutions.


This information is fundamental to an understanding of corrosion and rust. Thus, it is suggested that the reader proceed carefully, concentrating on learning the component parts of a corrosion cell and their interrelationships.


Why Metals Corrode


The driving force that causes metals to corrode is a natural consequence of their temporary existence in metallic form. To reach this metallic state from their occurrence in nature in the form of various chemical compounds (ores), it is necessary for them to absorb and store up for later return by corrosion, the energy required to release the metals from their original compounds. The amount of energy required and stored varies from metal to metal. It is relatively high for metals such as magnesium, aluminum, and iron and relatively low for metals such as copper and silver. Table 2.1 lists some commonly used metals in order of diminishing amounts of energy required to convert them from their ores to metal.


A typical cycle is illustrated by iron. The most common iron ore, hematite, is an oxide of iron (Fe2O3). The most common product of the corrosion of iron, rust, has the same chemical composition. The energy required to convert iron ore to metallic iron is returned when the iron corrodes to form the same compound. Only the rate of energy change is different.


The energy difference between metals and their ores can be expressed in electrical terms which are related to heats of formation of the compounds. The difficulty of extracting metals from their ores in terms of the energy required, and the consequent tendency to release this energy by corrosion, is reflected by the relative positions of pure metals in a list, which is discussed later as the electromotive series.


Forms of Corrosion


Destruction by corrosion takes many forms, depending on the nature of the metal or alloy: the presence of inclusions or other foreign matter at the surface; the homogeneity of its structure; the nature of the corrosion medium; the incidental environmental factors such as the presence of oxygen and its uniformity, temperature, and velocity of movement; and other factors such as stress (residual or applied, steady or cyclic); oxide scales (continuous or broken); porous or semiporous deposits on surfaces, built-in crevices; galvanic effects between dissimilar metals; and the occasional presence of stray electrical currents from external sources.


Except in rare cases of a grossly improper choice of material for a particular service, or an unanticipated drastic change in the corrosive nature of the environment or complete misunderstanding of its nature, failures of metals by rapid general attack (wasting away) are not often encountered. Corrosion failures are more often localized in the form of pits, intergranular corrosion, attack within crevices, etc. These and several other forms of attack are discussed in Chapter 5.


Corrosion Products


The term corrosion products refers to the substances produced during a corrosion reaction.

These can be soluble, such as zinc chloride or zinc sulfate in the examples cited earlier, or insoluble compounds such as iron oxide or hydroxide.


The presence of corrosion products is one way corrosion is detected (e.g. rust). However, it should be noted that insoluble corrosion products are not always visible. Upon exposure to air, aluminum forms an almost invisible oxide film which protects it from extensive atmospheric corrosion. It is invisible because it is so thin. This explainsthe widespread use of aluminum in storm windows, gutters, and automobile trim.


The products of the anodic and cathodic processes frequently migrate through the solution and meet to enter further reactions that yield many of our common visible corrosion products. For example, with iron in water, the hydroxyl ions from the cathodic reaction, in their migration through the electrolyte towards the anodic surfaces, encounter ferrous ions moving in the opposite direction. These ions combine to form ferrous hydroxide which subsequently reacts further with oxygen in solution to form ferric hydroxide. This is illustrated in Figure I and represents a form of iron rust with which we are all quite familiar.


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