Corrosion Mechanisms
11
2.2.4 Oxygen Concentration Cells
The oxygen-reduction reaction that occurs in neutral or basic solutions, O 2 +
2Η 2 O + 4e − → 40H − , plays a significant role in many corrosion processes. It not
only contributes to sustaining a cathodic reaction, but also can induce one.
This occurs when substantial differences in dissolved oxygen content exist
at one area on the metal surface relative to another. The natural tendency
is to equal concentrations, and the means of achieving this by corrosion is
to lower the oxygen concentration at the region where it is the highest. The
oxygen-reduction reaction accomplishes this but the area where this occurs
becomes cathodic to the lower oxygen concentration region. Because of the
current flow created by this action, corrosion will occur at the anodic or low
oxygen concentration site.
2.2.5 Metal ion Concentration Cells
Metal ion concentration cells can also develop and fuel the corrosion process.
This situation arises when a significant difference in metal ion concentration
exists over a metal surface. The tendency is to reach equilibrium ion concentration, and in a corrosive environment this is managed by putting more
metal ions into solution at the low-concentration area. This area becomes the
anode, and the current flow generated by this process can result in plating
out metal ions at the cathodic or high metal ion concentration region.
2.3 Potential–pH Diagrams
Potential–pH diagrams, also known as Pourbaix diagrams, are graphical representations of the stability of a metal and its corrosion products as a function of the potential and pH (acidity or alkalinity) of the aqueous solution.
The potential is shown on the vertical axis and the pH on the horizontal axis.
Such diagrams are constructed from calculations based on the Nernst equation and the solubility data for various metal compounds. The potential–pH
diagram for an Fe–Η 2 Ο system is shown in Figure 2.3. In the diagram, the
horizontal lines represent pure electron transfer reactions dependent solely
on potential, but independent of pH:
1.
Fe Fe
e
=
+
+
−
2
2
(2.3)
2.
Fe
Fe
e
2+
=
+
+
−
3
(2.4)
11
2.2.4 Oxygen Concentration Cells
The oxygen-reduction reaction that occurs in neutral or basic solutions, O 2 +
2Η 2 O + 4e − → 40H − , plays a significant role in many corrosion processes. It not
only contributes to sustaining a cathodic reaction, but also can induce one.
This occurs when substantial differences in dissolved oxygen content exist
at one area on the metal surface relative to another. The natural tendency
is to equal concentrations, and the means of achieving this by corrosion is
to lower the oxygen concentration at the region where it is the highest. The
oxygen-reduction reaction accomplishes this but the area where this occurs
becomes cathodic to the lower oxygen concentration region. Because of the
current flow created by this action, corrosion will occur at the anodic or low
oxygen concentration site.
2.2.5 Metal ion Concentration Cells
Metal ion concentration cells can also develop and fuel the corrosion process.
This situation arises when a significant difference in metal ion concentration
exists over a metal surface. The tendency is to reach equilibrium ion concentration, and in a corrosive environment this is managed by putting more
metal ions into solution at the low-concentration area. This area becomes the
anode, and the current flow generated by this process can result in plating
out metal ions at the cathodic or high metal ion concentration region.
2.3 Potential–pH Diagrams
Potential–pH diagrams, also known as Pourbaix diagrams, are graphical representations of the stability of a metal and its corrosion products as a function of the potential and pH (acidity or alkalinity) of the aqueous solution.
The potential is shown on the vertical axis and the pH on the horizontal axis.
Such diagrams are constructed from calculations based on the Nernst equation and the solubility data for various metal compounds. The potential–pH
diagram for an Fe–Η 2 Ο system is shown in Figure 2.3. In the diagram, the
horizontal lines represent pure electron transfer reactions dependent solely
on potential, but independent of pH:
1.
Fe Fe
e
=
+
+
−
2
2
(2.3)
2.
Fe
Fe
e
2+
=
+
+
−
3
(2.4)