Objective
In this article, we will explore how to interpret OLI’s Pourbaix Diagrams, using a foundational case study under varying conditions.
Disclaimer: The user interface, calculations, and results displayed in this article are from OLI Studio: Corrosion Analyzer Version 12.0.0. Other software versions may appear different or present slightly distinct results due to continual developments to the software and thermodynamic databanks.
Schematic Representation
A Pourbaix Diagram, also known as a Potential vs pH (E vs pH) diagram or Stability Diagram, helps you understand which chemical species are stable in water under different pH levels and electrochemical potentials. However, it's important to note that this diagram does not provide information on reaction rates or kinetic effects.
The accompanying image shows a Pourbaix diagram for copper in water at 25°C and 1 atm, generated using OLI Studio: Corrosion Analyzer and annotated for instructional purposes.
Figure 1. Pourbaix Diagram for the Cu-H2O system at 25°C and 1 atm
- Gray Area: Represents the "immune to corrosion" region, which is where the metal (copper, in this case) is stable and will not corrode.
- Green Area: Indicates the region where passivation may occur. Passivation is a thin layer that forms on the metal’s surface, which might protect it from corrosion. It is crucial to assess whether this layer is protective, as this depends on the crystalline structure of the compound. In this case, CuO and Cu2O are the solid phases produced by the oxidation of Cu.
- Light Yellow Area: Denotes the region where corrosion is possible, as ionic metal species dissolve in solution. In this zone, neither the metal nor passivating solids are stable. Here, Cu2+ is identified as the most stable ionic species.
Understanding the Lines on the Diagram
Equilibrium Lines
- The lines dividing different species in the Pourbaix diagram indicate equilibrium conditions for chemical and electrochemical reactions.
Water Decomposition Lines
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Diagonal Dashed Lines (a and b): Represent water reduction and water oxidation lines, respectively.
- Below line (a): Water decomposes to form H2 gas.
- Above line (b): Water decomposes to form O2 gas.
- The region between lines (a) and (b) is known as the stability region of water.
Corrosion Without Oxygen
In the absence of oxygen, the most common reaction is the reduction of the proton (H+) to elemental hydrogen (as shown by line (a) in the plot). The pH level will determine the reaction:
- Acidic conditions: H+ + 2e- -> H2
- Basic conditions: 2H2O + 2e- -> H2 + 2OH-
For a corrosion process to proceed, a driving force must be present. Therefore, line (a) must lie above an equilibrium line that corresponds to an equilibrium between the metal and metal-containing ions.
Corrosion in Oxygen-Containing Solutions
In oxygen-containing solutions, O2 can be reduced to H2O, as shown by line (b). Again, the pH will determine the reaction:
- Acidic conditions: O2 + 4H+ + 4e- -> 2H2O
- Basic conditions: O2 + 2H2O + 4e- -> 4OH-
For a corrosion process to occur, a driving force must be present. Therefore, line (b) must lie above an equilibrium line that corresponds to an equilibrium between the metal and metal-containing ions. In Figure 1, this condition is met, so corrosion is predicted.
Passivation is likely if (b) lies above a line that corresponds to an equilibrium between the metal and a sparingly soluble compound.
Additional Information
- Natural pH and ORP: The diagram also displays the values of the Natural pH and the Oxidation Reduction Potential (ORP). The natural pH line represents the computed pH of the water sample before any adjustments with acid or base are made for the diagram. The ORP, indicated by a red circle, shows the initial electrochemical potential of the water phase before potential adjustment using tools like a potentiostat.
OLI Studio: Corrosion Analyzer enables users to create diagrams like these for any non-ideal solution, at various temperatures and pressures. You can also incorporate complexation effects of additional chemistries in the solution. For instance, ammonia (NH3) is known to form complexes with copper. In the image below, you can see how the equilibrium lines for various Cu-NH3 complexes are represented in the diagram.
(A) (B)
Figure 2. (A) Cu-H2O-NH3 system at 25°C and 1 atm in a 0.1 m NH3 solution. (B) Cu-H2O-NH3 system at 25°C and 1 atm in a 0.5 m NH3 solution.
Note the significant Light Yellow Area that was previously the stability field for copper oxides in Figure 1. Here, the corrosion-susceptible region is greater, as copper prefers to be in the complex form. This indicates that ammonia can disrupt the passivation layer formed by copper oxide in the presence of oxygen, especially as ammonia concentration increases, leading to more corrosion.