Page 712 - Corrosion Engineering Principles and Practice

666 C h a p t e r 1 5 H i g h - Te m p e r a t u r e C o r r o s i o n 667

15.2 Thermodynamic Principles

15.2.1 Standard Free Energy of Formation
In a corrosion context, the thermodynamic properties of a metallic
material describe the energy that is lost or gained in the form of heat
as the metal reacts to form oxides or other compounds. The stability
of the various oxides can also differ.
Two important points are involved. First, the more energy released
when forming the compound, such as an oxide, the more stable the
compound. Secondly, as the temperature is raised, more of the
oxidizing species is required to maintain the oxide or other film.
As an illustration, let us consider a bar of copper and a bar of iron
at 980°C in an atmosphere made up of 20 percent H and 80 percent
2
H O. This atmosphere is partly oxidizing and partly reducing. Under
2
these conditions, copper will not corrode and will remain bright. Iron,
on the other hand, will become coated with a layer of iron oxide.
In the case of copper, the reaction with water that produces copper
oxide and hydrogen is thermodynamically not favored.
2Cu H O (80%) ← Cu O H (20 %) (15.1)
+
+
2
2
2
In fact, the reverse reaction would occur if some copper oxide
were on the copper; that is, the copper oxide would be reduced by the
hydrogen to form H O and metallic copper.
2
In the case of iron, however, the reaction with water to form iron
oxide and hydrogen in the same conditions is thermodynamically
favored and would therefore proceed.
Fe H O (80%) → FeO H (20 %) (15.2)
+
+
2
2
The 80 percent H O-20 percent H mixture would not permit H to
2
2
2
reduce FeO to form Fe and H O, as in the case of copper. Thus, we
2
may say that copper is corrosion resistant under these circumstances,
or is more noble than iron. Iron, however, is thermodynamically
subject to corrosive attack.
It is possible to use plots of the free energy of formation of metal
oxides versus temperature to predict the temperatures at which a
metal is stable and the temperatures at which it will spontaneously
oxidize. For temperatures at which the free energy of formation of the
oxide is positive, the reverse reaction is favored and the oxide will
spontaneously decompose to the metal.
Plots of the standard free energy of reaction (∆G ) as a function of
0
temperature, commonly called Ellingham diagrams, can help to
visualize the relative stability of metals and their oxidized products.
Figure 15.3 shows an Ellingham diagram for many simple oxides.
The ∆G values on an Ellingham diagram are expressed as kilojoule
0
per mole of O to normalize the energy scale and compare the stability
2
of these oxides directly, that is, the lower the position of the line on