Iron oxide reduction thermodynamics

Although thermodynamically favorable, reductive dissolution of Fe(III)(hydr)oxides by some metastable ligands (even those, such as oxalate, that can form surface complexes) does not occur in the absence of light. The photochemical pathway is depicted in Fig. 9.3e. In the presence of light, surface complex formation is followed by electron transfer via an excited state (indicated by ) either of the iron oxide bulk phase or of the surface complex. (Light-induced reactions will be discussed in Chapter 10.)... 

Based upon thermodynamic data given in Table I, oxidant strength decreases in the order NijO > Mn02 > MnOOH > CoOOH > FeOOH. Rates of reductive dissolution in natural waters and sediments appear to follow a similar trend. When the reductant flux is increased and conditions turn anoxic, manganese oxides are reduced and dissolved earlier and more quickly than iron oxides (12, 13). No comparable information is available on release of dissolved cobalt and nickel. 

Eh is a measure of oxidation-reduction potential in the solution. The chemical reactions in the aqueous system depend on both the pH and the Eh. While pH measures the activity (or concentration) of hydrogen ions in the solution, Eh is a measure of the activity of all dissolved species. Aqueous solutions contain both oxidized and reduced species. For example, if iron is present in the solution, there is a thermodynamic equilibrium between its oxidized and reduced forms. Thus, at the redox equilibrium, the reaction is as follows ... 

Table 2.1 lists equilibrium ratios for the reduction of selected metal oxides [4], while Figure 2.2 provides a complete phase diagram for the reduction of iron oxide at different temperatures [3, 5], In order to reduce bulk iron oxide to metallic iron at 600 K, the water content of the hydrogen gas above the sample must be below a few percent, which is easily achieved. However, in order to reduce Cr2C>3, the water content should be as low as a few parts per billion, which is much more difficult to realize. The data in Table 2.1 also illustrate that, in many cases, only partial reduction to a lower oxide may be expected. Reduction of Mn2C>3 to MnO is thermodynamically allowed at relatively high water contents, but further reduction to manganese is unlikely. 

Metallic corrosion occurs because of the coupling of two different electrochemical reactions on the material surface. If, as assumed in the discussion of iron dissolution kinetics above, only iron oxidation and reduction were possible, the conservation of charge would require that in the absence of external polarization, the iron be in thermodynamic equilibrium. Under those conditions, no net dissolution would occur. In real systems, that assumption is invalid, and metallic dissolution occurs with regularity, keeping corrosionists employed and off the street. 

It has been generally assumed that iron is transported across biological membranes in the ferrous form and that ferric iron would have to be reduced before it can be used by the organism. Thus, based on nutritional studies it has long been recognized that Fe(II) is1 more effectively absorbed than Fe(III), and this has been attributed to differences in the thermodynamic and kinetic stability of the complexes and chelates formed by these cations (for review, see Ref. 2). The experimental proof of a transport in the ferrous form has, however, not been given until quite recently in studies of iron transport in isolated mitochondria (23) as well as in enterobacteria (33). In rat liver mitochondria we have found that Fe(III) donated from a metabolically inert water soluble complex of sucrose interacts with the respiratory chain at the level of cytochrome c (and possibly cytochrome a) (23, 32) (Figure 1 B), which has a oxidation-reduction potential of around +250 mV (34) and is localized to the outer phase of the mitochondrial inner membrane (35). 

The pH dependence of cytochrome c oxidation-reduction reactions and the studies of modified cytochrome c thus demonstrate that the coordination environment of the iron and the conformation of the protein are relatively labile and strongly influence the reactivity of the metallo-protein toward oxidation and reduction. The effects seen may originate chiefly from alterations in the thermodynamic barriers to electron transfer, but the conformation changes are expected to affect the intrinsic barriers also. One such conformation change is the opening of the heme crevice referred to above. The anation and Cr(II) reduction studies provide an estimate of 60 sec 1 for this process in Hh(III) at 25°C (59). To date, no evidence has been found for a rapid heme-crevice opening step in ferrocytochrome c. 

In the analysis of the deposition of iron sediments it has already been mentioned that quite likely both iron silicates and carbonates and amorphous iron hydroxide were formed, which could convert to other forms both during the formation of the sediment and in subsequent diagenesis. Reduction of hydroxide could have been controlled by external (atmospheric) or internal (organic matter, free carbon in the sediment) oxidation-reduction buffer systems. All these variants need additional consideration in the thermodynamic analysis of diagenetic processes. 

Figure 2 illustrates the thermodynamic limitations for the reduction of iron oxide by means of producer gas at 1100 K. It shows that the reduction of Fc304 to FeO with CO only proceeds as long as the CO2/CO ratio is smaller than 18. Correspondingly, reduction only occurs until the H2O/H2 ratio in the gas reaches a value of 19. This means that about 95% of the CO and H2, respectively, can be converted to CO2 and H2O, respectively. 

Figure 28. Hypothetical anaerobic nitrogen cycle based on the following thermodynamically permissible reactions (1) ammonium oxidation to dinitrogen by carbon dioxide,. sulfate or ferric iron (no evidence at present, possibly kinetically limited) (2) dinitrogen fixation by various organic and inorganic reductants (known) (3) ammonium oxidation by nitrite or nitrate producing dinitrogen (known) (4) denitrification (known) (5) nitrite or nitrate respiration (known) (6) ferric iron oxidation of ammonium to nitrite or nitrate (no evidence at present) (7) nitrate assimilation (known) (8) ammonium assimilation and di.s,similation (known) (Fenchel etai, 1998).

The sequence of reactions involved in the overall reduction of nitric acid is complex, but direct measurements confirm that the acid has a high oxidation/reduction potential, -940 mV (SHE), a high exchange current density, and a high limiting diffusion current density (Ref 38). The cathodic polarization curves for dilute and concentrated nitric acid in Fig. 5.42 show these thermodynamic and kinetic properties. Their position relative to the anodic curves indicate that all four metals should be passivated by concentrated nitric acid, and this is observed. In fact, iron appears almost inert in concentrated nitric acid with a corrosion rate of about 25 pm/year (1 mpy) (Ref 8). Slight dilution causes a violent iron reaction with corrosion rates >25 x 1()6 pm/year (106 mpy). Nickel also corrodes rapidly in the dilute acid. In contrast, both chromium and titanium are easily passivated in dilute nitric acid and corrode with low corrosion rates. 

With supported metal catalysts that have to be treated in a reducing gas flow at elevated temperatures to convert the catalytic precursor into the desired metal, it is important to assess the extent of reduction. Often the oxidic phase of the cata-lytically active precursor is stabilized by interaction with the support. It is even possible for a finely divided precursor to react with the support to a compound much more stable than the corresponding metal oxide. An example is cobalt oxide, which can react with alumina to form cobalt aluminate, which is very difficult to reduce to metallic cobalt and alumina. Another example is silica-supported iron oxide. Usually the reduction of iron(III) to iron(II) proceeds readily, because the reduction to iron(II) is hardly thermodynamically limited by the presence of water vapor. Iron(ll), however, reacts rapidly with silica to iron(II) silicate, which is almost impossible to reduce. 

All refined metals have a tendency to revert to a thermodynamically more stable form such as those in which they occur naturally on earth. Thus one of the corrosion products of iron is iron oxide (Fea03(s)) which is one form of iron ore. Almost all types of corrosion can be explained in terms of electrochemistry (oxidation-reduction reactions) for this reason we will consider corrosion as an example of the application of redox chemistry or electrochemistry to a practical situation. We will not present a detailed quantitative analysis of corrosion and the design of corrosion-control systems. Other texts should be consulted for this type of information. 

Oxidation-reduction reactions involving iron and manganese include changes in oxidation states, relative stabilities of iron and manganese compounds, and the energetics and kinetics of oxidized and reduced compounds. Using thermodynamic data (from free energies of formation, see Chapter 4),... 

The most numerous thermodynamic data on iron, ruthenium and osmium diimine complexes, are the oxidation-reduction (redox) potentials of the MeLi /MeLi (and similar) couples ). The relation... 

Under environmental conditions, the Fe -H20 interface has a surface layer of corrosion products that develops due to the thermodynamic instability of Fe in the presence of water. Long-term batch and colunm studies have shown that this layer evolves with time into a complex mixture of amorphous iron oxides, iron oxide salts, and other mineral precipitates (2-S). Because this material lies at the metal-water interface, it must, in some manner, mediate the reduction of contaminants by the underlying metal. Understanding the mechanism by which metals reduce contaminants in the presence of a substantial layer of oxides is one of the critical, remaining challenges for researchers in this field. The goal of the following analysis is... 

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