Pourbaix diagram iron species

Pourbaix diagrams (Pourbaix, 1963) indicate graphically the conditions of redox potential (Eh) and pH under which different types of corrosion behaviour may be expected. These plots of potential vs. pH indicate the phase and species in equilibrium with iron under various conditions (see Chap. 8). The solid phases indicated are those that are thermodynamically the most stable owing to kinetic factors other phases may be present during the initial stages of the corrosion process. What the different regions show, however, are the predominant oxidation states to be expected. 

The Pourbaix diagram shown in Figure 13.1 illustrates the thermodynamic stability of iron species in aqueous solutions of a few organic substrates. The relative position of each substrate shows that the reaction between iron and the corresponding organic is thermodynamically favorable. Three elementary reactions involved in the reductive dechlorination of organic compounds are shown in Figure 13.2. 

Fig. 16.1. Simplified Pourbaix diagram for iron in pure water—ionic activities 1(T6 (a) The diagram in terms of most stable species (b) The diagram in terms of the type of reaction that can occur.

Reference to the Pourbaix diagrams for water and copper shows the stable species to be Cu2+, H+ and H2O. Corrosion is possible under these conditions. Let us consider the case of iron metal at —0.750 V vs SCE in a solution of pH 5.0 and a ferrous ion activity of 10 6 at 25°C. The electrode potential with respect to SHE is 0.509 V and reference to Pourbaix diagrams for water and iron shows the stable species to be Fe2+ and H2, and that corrosion is possible under these conditions. 

Next we consider the Pourbaix diagram for iron, which is, of course, of paramount importance for the understanding of corrosion of ferrous alloys such as the many types of steel and stainless steel. This is a rather complex diagram, since two oxidation states of iron exist both in the liquid and the solid state and the metal is amphoteric to some extent. Figure 16M is a simplified version of the diagrams shown in the original work of Pourbaix. The two soluble species in acid solutions are Fe and Fe. The relevant equilibria are... 

Any species capable of oxidizing either Fe " or HjS at pH 6 cannot survive in this environment. According to the Pourbaix diagram for iron (Figure 5.11), the potential for the Fe(OH)j/Fe couple at pH 6 is approximately 0.3 V. Using the Latimer diagrams for sulfur in acid and base (see Resource Section 3), the S/H2S potential at pH 6 can be calculated as follows ... 

The many corrosion equilibria involving iron can be conveniently summarized in an equilibrium or Pourbaix diagram (15). These diagrams map regions of stability for various ionic and solid species as functions of pH and on the basis of equilibrium potentials. A simplified equilibrium diagram is shown in Figure 2. At sufficiently negative potential, iron is thermodynamically stable, and corrosion cannot occur. The domain of stability of water is defined by the dashed lines a and b below a,... 

Figure 4 is the Pourbaix Diagram for iron and some of its ionic species and compounds in contact with water at 25 C. The equilibrium potential of the iron-ferrous ion reaction falls outside the region of stability of water (dashed lines). Therefore, any attempt to measure the equilibrium potential will fail since the solvent will undergo electrochemical reduc-... 

Figure S3.4 Simplified Pourbaix diagram showing the stable species in the iron-water-oxygen system (a) oxidation-reduction boundaries that are not pH-dependent (b) acid-base boundaries that are not oxidation-dependent (c) as part (b), including a sloping boundary that is dependent on pH and oxidation potential (d) complete diagram, showing further sloping boundaries that delineate regions that are dependent on pH and oxidation potential...

The Pourbaix diagram for iron is constructed by using nine reduction and oxidation reactions over a pH range of 0-14. All dissolved species are assumed to have activities of 1 M. The diagram is plotted at 25 °C. The following reactions were used to constmct the diagram Reaction (1)... 

Figure 5. Potential-pH equilibrium diagrams for the system iron-water at 298 K [considering as solid substances only Fe, Fe(OH)2, and Fe(OH)3] for (a) the bulk-to-bulk equilibria (Pourbaix diagrams) and (b) the surface-to-bulk equilibria (Drazic-Pourbaix diagrams), for a concentration of all soluble species in solution of 1 M.

FIGURE 3.13 Pourbaix diagram for iron species at various reduction potentials E and pH in aqueous environment at 25 °C. Only the species indicated in the fields were considered. The dashed lines represent the thermodynamic thresholds for the emission of (upper) and Hj (lower). 

A further important feature of the hydration reactions of cement with water is that the resulting pore solution in concrete is highly alkaline [refer to Eq. (2.28) above]. In addition to calcium hydroxide, sodium and potassium hydroxide species are also formed, resulting in a pH of the aqueous phase in concrete that is typically between 12.5 and 13.6. Under such alkaline conditions, reinforcing steel tends to display completely passive behavior, as fundamentally predicted by the Pourbaix diagram for iron. In the absence of corrosive species penetrating into the concrete, ordinary carbon steel reinforcing thus displays excellent corrosion resistance. 

Potential-pH (Pourbaix) diagrams for iron and ruckel in aqueous solutions above the critical temperature of water ate shown in Figs. 4A and 4B, respectively. " The lack of detail, compared to their ambient temperature coimterparts, reflects the lack of thermodynamic data for species at supercritical temperatures, particularly for hydrolyzed iottic species, and the lack of stability of dissolved iottic species in the low dielectric constant supercritical medium. 

In terms of the Pourbaix potential/pH diagrams, the theoretical scale compares the potentials of immunity of the different metals, while the practical scale compares the potentials of passivation. But this is not enough either. The real scale depends on the environment with which the structure will be in contact during service. Passivity, as we have seen, depends on pH. It also depends on the ionic composition of the electrolyte, particularly the concentration of chloride ions or other species that are detrimental to passivity. Finally, one must remember that construction materials are always alloys, never the pure metals. The tendency of a metal to be passivated spontaneously can depend dramatically on alloying elements. For example, an alloy of iron with 8% nickel and 18% chromium (known as 304 stainless steel) is commonly used for kitchen utensils. This alloy passivates spontaneously and should be ranked, on the practical scale of potentials, near copper. If... 

---