1. Electrochemistry MAE-212
Dr. Marc Madou, UCI, Winter 2016
Class VII Pourbaix Diagrams

2. Table of content
Marcel Pourbaix
Pourbaix Diagrams

3. Marcel Pourbaix
Marcel Pourbaix provided the brilliant means to utilize thermodynamics more effectively in corrosion science and electrochemistry in general. This development resulted in four important books that interpret his work: Thermodynamics of Dilute Aqueous Solutions, Atlas of Electrochemical Equilibria in Aqueous Solutions (solid-aqueous equilibria); Lectures on Electrochemical Corrosion (a teaching text); and, in his last years, Atlas of Chemical and Electrochemical Equilibria in the Presence of a Gaseous Phase (solid-gaseous equilibria).
His outstanding work in thermodynamics provided one of the main underpinnings of electrochemistry, especially in corrosion science.
Marcel Pourbaix
b. 1904, Myshega, Russia
d. September 28, 1998, Uccle (Brussels), Belgium

4. Marcel Pourbaix Marcel Pourbaix (1904-1998)
Marcel Pourbaix was born in Russia, where his father, a Belgian engineer, was working at the time.
The significance of Marcel Pourbaix’s great achievement　was pointed out by Ulick R. Evans, widely recognized as the　“father of corrosion science,” in his foreword to Pourbaix’s　Thermodynamics of Dilute Aqueous Solutions: “During the last　decade (the 1940s) Dr. Marcel Pourbaix of Brussels has developed a graphical method, based on generalized thermodynamical equations, for the solution of many different kinds of scientific problems, involving numerous types of heterogeneous or homogeneous reactions and equilibria... Some of these problems have long been treated from the aspect of thermodynamics... The application of thermodynamics to typical corrosion reactions is a much newer development.”

5. For Mr. Zhang Zhongcheng
As a memory of his stay at Cebelcor in 1984/1985.
With all best wishes from
　　　　　　　　Marcel Pourbaix
　　　　　　　　　April 26, 1985

6. Prof Pourbaix, his wife, and Zhang Zhongcheng, in 1985

7. Prof Pourbaix with his colleagues at Cebelcor in 1984

8. The E-pH diagram of copper-water system

9. Pourbaix Diagrams
Through the use of thermodynamic theory (the Nernst equation), so-called Pourbaix diagrams can be constructed. These diagrams show the thermodynamic stability of species as a function of potential and pH. Although many basic assumptions must be considered in their derivation, such diagrams can provide valuable information in the study of corrosion phenomena. The diagram on the right represents a simplified version of the Pourbaix diagram for the iron- water system at ambient temperature. For the diagram shown, only anhydrous oxide species were considered and not all of the possible thermodynamic species are shown.   
How do we construct this diagram for water?

10. Pourbaix Diagrams
Use Nernst Equation:

11. Pourbaix Diagrams

12. Pourbaix Diagrams Potential H2O is stable H2 is stable 7 14
2H+ + 2e- = H2 Equilibrium potential falls as pH increases
2.0
1.6
0.8
1.2
-0.4
0.4
0.0
-1.6
-0.8
-1.2
2H2O = O2 + 4H+ + 4e-
Equilibrium potential
falls as pH increases
O2 is stable

13. Pourbaix Diagrams
We will consider Cu in an aqueous solution as the next exercise: five different reactions are involved. 7

14. Pourbaix Diagrams

15. Pourbaix Diagrams

16. Pourbaix Diagrams
The diagram shown here shows how the potentials for reduction and oxidation of water vary with pH for natural waters. These are the inner two lines that slope downward from low pH to high pH. Note that the pH scale only runs from 2-10 (we are talking here about natural waters). For both oxidation and reduction of water, an additional line is shown that lies 0.6V above (for oxidation of water) or below (for reduction) the theoretical E. This pair of lines represents the potentials including an approximation for the overvoltage. Lastly, there is a pair of vertical lines at pH=4 and 9. These are reflective of the fact that most natural waters have a pH somewhere between these limits. .
2008
MAE 217-Professor Marc J. Madou

17. Pourbaix Diagrams
A Pourbaix diagram is an attempt to overlay the redox and acid-base chemistry of an element onto the water stability diagram. The data that are required are redox potentials and equilibrium constants (e.g. solubility products). On the right is the Pourbaix diagram for iron. Below that is the same diagram showing only those species stable between the water limits.
2008
MAE 217-Professor Marc J. Madou

18. Pourbaix Diagrams Equilibrium Reactions of iron in Water
1. 2 e- + 2H+  = H2
2. 4 e- + O2  + 4H+  = 2H2O
3. 2 e- + Fe(OH)2  + 2H+  = Fe + 2H2O
4. 2 e- + Fe2+  = Fe
5. 2 e- + Fe(OH)3-  + 3H+  = Fe + 3H2O
6. e- + Fe(OH)3  + H+  = Fe(OH)2 + H2O
7. e- + Fe(OH)3  + 3H+  = Fe2+ + 3H2O
8. Fe(OH)3-  + H+  = Fe(OH)2 + H2O
9.  e- + Fe(OH)3  = Fe(OH)3-
10. Fe3+  + 3H2O  = Fe(OH)3 + 3H+
11. Fe2+  + 2H2O  = Fe(OH)2 + 2H+
12.  e- + Fe3+  = Fe2+
13. Fe2+  + H2O  = FeOH+ + H+
14. FeOH+  + H2O  = Fe(OH)2(sln) + H+
15. Fe(OH)2(sln)  + H2O  = Fe(OH)3- + H+
16. Fe3+  + H2O  = FeOH2+ + H+
17. FeOH2+  + H2O  = Fe(OH)2+ + H+
18. Fe(OH)2+  + H2O  = Fe(OH)3(sln) + H+
19. FeOH2+  + H+  = Fe2+ + H2O
20. e- + Fe(OH)2+  + 2H+  = Fe2+ + 2H2O
21.  e- + Fe(OH)3(sln)  + H+  = Fe(OH)2(sln) + H2O
22. e- + Fe(OH)3(sln)  + 2H+  = FeOH+ + 2H2O
23.  e- + Fe(OH)3(sln)  + 3H+  = Fe2+ + 3H2O
MAE 217-Professor Marc J. Madou
2008

19. Pourbaix Diagrams Some limitations of Pourbaix diagrams include:
No information on corrosion kinetics is provided by these thermodynamically derived diagrams.
The diagrams are derived for specific temperature and pressure conditions.
The diagrams are derived for selected concentrations of ionic species (10-6 M for the above diagram).
Most diagrams consider pure substances only - for example the above diagram applies to pure water and pure iron only. Additional computations must be made if other species are involved.
In areas where a Pourbaix diagram shows oxides to be thermodynamically stable, these oxides are not necessarily of a protective (passivating) nature.
2008
MAE 217-Professor Marc J. Madou

20. Pourbaix Diagrams MAE 217-Professor Marc J. Madou 2008
Use the definition of ∆G and we find that when Q=1, we have ∆G = ∆G˚. This is the way we designed the equation. What is more interesting is to see what happens when ∆G goes to 0. Here is when we reach equilibrium, the various concentrations in the reaction stop changing. The reaction seems to stop. In this case, we obtain a relation between ∆G˚ and the equilibrium constant.
The reaction doesn’t really stop; we simply have reached the situation where the forward and reverse reaction rates are equal. Chemical activity is still ongoing with just as much vigor as before, but it is happening equally in both directions — products are reacting to turn into reactants just as fast as reactants are reacting to turn into products. This is why we call it a dynamic equilibrium, rather than a static equilibrium.
MAE 217-Professor Marc J. Madou
2008