Chemistry of corrosion
Chemistry of Corrosion
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Corrosion chemistry
Corrosion Chemistry
Module Two of CCE 281 Corrosion: Impact, Principles, and Practical Solutions
Lesson Objectives
Explain the instability of metals
Discuss the factors that can trigger corrosion
Explain the chemistry of corrosion
Review the definition of acidity
Compare corrosion reactions of some common metals
Required Reading
This Module consists of three Web pages of required reading. The pagination is visible at the bottom of each page with
direct links to adjacent pages.
Additional information can be found in sections 2.1, 2.2, 2.3, and 2.4 of the reference textbook (Corrosion
Engineering: Principles and Practice).
Introduction
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.
See Why Metals Corrode?
Example problem 2.1
Compare the energy required to produce one metric ton of magnesium from its oxide to the energy required
to convert enough copper oxide to produce one ton of metallic copper. http://www.corrosion-doctors.org/Chemistry-of-Corrosion/Introduction.htm[10/12/2010 9:01:38 AM]
Corrosion chemistry
Example problem 2.2
Discuss the energy values presented in the Table shown on the page describing why metals corrode in
relation to the order in which metals and associated alloys appeared in the history of mankind. When discussing the ionic content of an aqueous medium, the question often arises as to how acid (or alkaline) is the
solution. Quite simply, this refers to whether there is an excess of H + (hydrogen) or OH - (hydroxyl) ions present. The
H+ ion is acid while the hydroxyl ion is alkaline or basic. The other ionic portion of an acid or alkali added to water
can increases its conductivity or change other properties of the liquid, but does not increase or decrease its acidity. For
instance, whether a given amount of H+ ion is produced in water by introducing hydrochloric (HCl), sulfuric (H2SO 4),
or any other acid is immaterial. The pH of the solution will be the same for the same number of dissolved hydrogen
atoms. (reference)
The pH may be measured with a meter or calculated if certain parameters are established. Water itself dissociates to a
small extent to produce equal quantities of H+ and OH - ions displayed in the following equilibrium:
pH , originally defined by Danish biochemist Søren Peter Lauritz Sørensen in 1909, is a measure of the
concentration of hydrogen ions. The term pH was derived from the manner in which the hydrogen ion concentration is
calculated, it is the negative logarithm of the hydrogen ion (H+) concentration:
where log is a base-10 logarithm and a H+ is the activity (related to concentration) of hydrogen ions. The "p" in
Equation stands for the German word for "power", potenz, so pH is an abbreviation for "power of hydrogen".
Example problem 2.3
A solution is made up to contain 0.01 M HCl. What is its pH? Example problem 2.4
A solution is made up to contain 0.01 M NaOH. What is its pH? A higher pH means there are fewer free hydrogen ions, and that a change of one pH unit reflects a tenfold change in
the concentrations of the hydrogen ion. For example, there are 10 times as many hydrogen ions available at pH 7 than
at pH 8. The pH scale commonly quoted ranges from 0 to 14 with a pH of 7 considered to be neutral.
Substances with a pH less that 7 are considered to be acidic and substances with pH equal to or greater than 7 to be
basic or alkaline. Thus, a pH of 2 is very acidic and a pH of 12 very alkaline. However, it is technically possible to
have very acidic solutions with a pH lower than zero and concentrated caustic solutions with a pH greater than 14.
Such solutions are in fact typical of many ore extracting processes that require the digestive power of caustics and
acids.
Low pH acid waters accelerate corrosion by supplying hydrogen ions to the corrosion process. Although even
absolutely pure water contains some free hydrogen ions, dissolved carbon dioxide (CO2) in the water can increase the
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Corrosion chemistry
hydrogen ion concentration. Dissolved CO 2 may react with water to form carbonic acid (H2CO 3) as shown in equation.
where Keq is the reaction equilibrium expressed as a ratio.
Carbonic acid subsequently dissociates in bicarbonate and carbonate ions as expressed respectively in the following
equations:
Example problem 2.5
A solution contains a mixture of sodium bicarbonate (0.05 M) and sodium carbonate (0.2 M).What is its pH? Care must be taken when quoting and using the dissociation constant in equation. This equilibrium value is correct for
the H2CO 3 molecule, and shows that it is a stronger acid than acetic acid or formic acid as might be expected from the
influence of the electronegative oxygen substituent. However, carbonic acid only exists in solution in equilibrium with
carbon dioxide, and so the concentration of H2CO 3 is much lower than the concentration of CO 2, reducing the
measured acidity. The equation may be rewritten as follows:
Even more acidity is sometimes encountered in mine waters and in water contaminated by industrial wastes. Many
salts added to an aqueous system also have a direct effect on the pH of that mixture through the following process of
hydrolysis shown here for the addition of ferric ions to water:
In this particular example the equilibrium is established between ferric ions, water, ferric hydroxide or Fe(OH)3 and
the acidity of the water. This particular example is quite useful to explain the severity of a situation that can develop in
confined areas such as corrosion pitting and crevices.
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Page 1 of 3
Information Module
Introduction
Corrosion in acids
Corrosion in neutral or alkaline environments
See also CCE 513: Corrosion Engineering
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Corrosion in acids
Module Two of CCE 281 Corrosion: Impact, Principles, and Practical Solutions
Corrosion in Acids
One of the common ways of generating hydrogen in a laboratory is to place zinc into a dilute acid, such as
hydrochloric or sulfuric. When this is done, there is a rapid reaction in which the zinc is attacked or “dissolved” and
hydrogen is evolved as a gas.
Rapid evolution of hydrogen bubbles during the corrosion of a zinc strip in a 1 M HCl acid solution
These reactions are described in the following equations to:
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Corrosion in acids
These equations are the chemical shorthand for the statement: One zinc atom + two hydrochloric acid molecules
dissociated as ions H+ and Cl - and becomes one molecule of zinc chloride in the first equation and written as a soluble
salt in the form of Zn2+ and Cl - ions in the second equation + one molecule of hydrogen gas which is given off as
indicated by the vertical arrow. It should be noted that the chloride ions do not participate directly in this reaction,
although they could play an important role in real corrosion situations.
Similarly, zinc combines with sulfuric acid to form zinc sulfate (a salt) and hydrogen gas as shown in the following
equations:
Note that each atom of a substance that appears on the left-hand side of these equations must also appear on the righthand side. There are also some rules that denote in what proportion different atoms combine with each other. As in the
preceding reaction, the sulfate ions that are an integral part of sulfuric acid do not participate directly to the corrosion
attack and therefore one could write these equations in a simpler form:
Many other metals are also corroded by acids often yielding soluble salts and hydrogen gas as shown in Equations and
for respectively iron and aluminum:
Note that zinc and iron react with two H+ ions, whereas aluminum reacts with three. This is due to the fact that both
zinc and iron, when corroding, each lose two electrons and display two positive charges in their ionic form. They are
said to have a valence of +2 or II, whereas aluminum loses three electrons when leaving an anodic surface and hence
displays three positive charges and is said to have a valence of +3 or III. Some metals have several common valences,
others only one. The following Figure shows Some of the oxidation states found in compounds of the transition-metal
elements.
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Corrosion in acids
Oxidation states found in compounds of the metalic elements. A solid circle represents a common oxidation
state, and a ring represents a less common (less energetically favorable) oxidation state
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Corrosion in Neutral or Alkaline Environments
Module Two of CCE 281 Corrosion: Impact, Principles, and Practical Solutions
Corrosion in Neutral or Alkaline Environments
The corrosion of metals can also occur in fresh water, seawater, salt solutions, and alkaline or basic media. In almost
all of these environments, corrosion occurs importantly only if dissolved oxygen is also present. Water solutions
rapidly dissolve oxygen from the air, and this is the source of the oxygen required in the corrosion process. The most
familiar corrosion of this type is the rusting of iron when exposed to a moist atmosphere. (reference)
In this equation, iron combines with water and oxygen to produce an insoluble reddish-brown corrosion product that
falls out of the solution, as shown by the downward pointing arrow.
During rusting in the atmosphere, there is an opportunity for drying, and this ferric hydroxide dehydrates and forms the
familiar red-brown ferric oxide (rust) or Fe2O3, as shown below:
Similar reactions occur when zinc is exposed to water or moist air followed by natural drying.
The resulting zinc oxide is the whitish deposit seen on galvanized pails, rain gutters, and imperfectly chrome-plated
bathroom faucets. It also familiarly called 'white rust' a non-protective and even destructive form of corrosion that
attacks incompletely passivated galvanized steel material or galvanized components subjected to marine atmospheres.
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Corrosion in Neutral or Alkaline Environments
White rust on seaside road railing
As discussed previously, the iron that took part in the reaction with hydrochloric acid in had a valence of 2, whereas
the iron that takes part in the reaction shown in the previous equation has a valence of 3. The clue to this lies in the
examination of the equation for the corrosion product Fe(OH)3. Note that water ionized into H+ and OH - . It is further
known that hydrogen ion has a valence of 1 (it has only one electron to lose). It would require three hydrogen ions
with the corresponding three positive charges to combine with the three OH - ions held by the iron. It can thus be
concluded that the iron ion must have been Fe3+ or a ferric ion.
Also note that there is no oxidation or reduction (electron transfer) during either reaction. In both cases the valences of
the elements on the left of each reaction remain what it is on the right. The valences of iron, zinc, hydrogen, and
oxygen elements remain unchanged throughout the course of these reactions, and it is consequently not possible to
divide these reactions into individual oxidation and reduction reactions.
Answers to example problems
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