Surfaces ferrous iron oxidation

Ferrous iron oxidation by manganese oxide was found to be especially fast as long as no iron oxyhydroxide precipitates, which presumably blocks reactive sites on the manganese oxide surface (Postma 1985). The oxidation of ferrous iron by manganese oxide has been proven to be important for the interpretation of pore water profiles and the precipitation of authigenic phases (Canfield et al. 1993a Haese et al. 2000 van der Zee 2005). In Fig. 7.16 pore water profiles of iron and manganese reveal concurrent liberation of the... 

Hydrogenis prevented from forming a passivating layer on the surface by an oxidant additive which also oxidizes ferrous iron to ferric iron. Ferric phosphate then precipitates as sludge away from the metal surface. Depending on bath parameters, tertiary iron phosphate may also deposit and ferrous iron can be incorporated into the crystal lattice. When other metals are included in the bath, these are also incorporated at distinct levels to generate species that can be written as Zn2Me(P0 2> where Me can represent Ni, Mn, Ca, Mg, or Fe. 

Iron oxide yellows can also be produced by the direct hydrolysis of various ferric solutions with alkahes such as NaOH, Ca(OH)2, and NH. To make this process economical, ferric solutions are prepared by the oxidation of ferrous salts, eg, ferrous chloride and sulfate, that are available as waste from metallurgical operations. The produced precipitate is washed, separated by sedimentation, and dried at about 120°C. Pigments prepared by this method have lower coverage, and because of their high surface area have a high oil absorption. 

Sulfides are intermixed with iron oxides and hydroxides on carbon steels and cast irons. The oxides are also produced in the corrosion process (Reaction 6.6). Although theoretical stoichiometry of 1 to 3 is often suggested between sulfide and ferrous hydroxide, empirically the ratio of iron sulfide to ferrous hydroxide is highly variable. Sulfide decomposes spontaneously upon exposure to moist air. Additionally, corrosion-product stratification is marked, with sulfide concentration being highest near metal surfaces. 

Iron is the most abundant, useful, and important of all metals. For example, in the 70-kg human, there is approximately 4.2 g of iron. It can exist in the 0, I, II, III, and IV oxidation states, although the II and III ions are most common. Numerous complexes of the ferrous and ferric states are available. The Fe(II) and Fe(III) aquo complexes have vastly different pAa values of 9.5 and 2.2, respectively. Iron is found predominantly as Fe (92%) with smaller abundances of Fe (6%), Fe (2.2%), and Fe (0.3%). Fe is highly useful for spectroscopic studies because it has a nuclear spin of. There has been speculation that life originated at the surface of iron-sulfide precipitants such as pyrite or greigite that could have caused autocatalytic reactions leading to the first metabolic pathways (2, 3). 

Iron and the Ferrous Alloys. Iron exposed to a moist environment reacts with water and atmospheric oxygen to form rust, a brown, crumbly corrosion product (composed of hydrated iron oxide). Initially rust forms a surface layer that is usually held in much disfavor. If an ancient iron object... 

It is unfortunately the case that when we incubate apoferritin with a certain number of iron atoms (for example as ferrous ammonium sulfate), the product, after elimination of non-protein-bound iron, does not have a homogeneous distribution of iron molecules which were able to (i) take up iron rapidly through the three fold channels, (ii) quickly transfer it and form a diiron centre on a ferroxidase site, and (iii) to transfer the iron inward to a nucleation site, where (iv) it will begin to catalyse iron oxidation on the surface of the growing crystallite, will accumulate iron much more rapidly, and in much greater amounts than molecules in which steps (i), (ii) and (iii) are slower, for whatever reasons (perhaps most importantly subunit composition, and the disposition of subunits of the two types H and L, one with regard to the other). This polydispersity makes the analysis of the process of iron uptake extremely difficult. 

Table 28.1. Surface complexation reactions considered by Liger et al. (1999) in describing the kinetics of the catalytic oxidation ofuranyl by ferrous iron...

Fig. 28.1. Results (symbols) and simulations (lines) of an experiment at 25 °C by Liger et al. (1999 their Fig. 6) in which uranyl was oxidized by ferrous iron in the presence of nanoparticulate hematite, which served as a catalyst. Vertical axis is amount of NaHCCE-extractable uranyl, which includes uranyl present in solution as well as that sorbed to the nanoparticles in the experiment, nearly all the uranyl was sorbed. Broken line shows results of a simulation assuming uranyl forms a single surface complex, >Fe0U020H, which is catalytically active solid line shows simulation in which a non-catalytic site of this stoichiometry is also present. Inset is an expanded view of the first few hours of reaction.

Dissolved iron(III) is (i) an intermediate of the oxidative hydrolysis of Fe(II), and (ii) results from the thermal non-reductive dissolution of iron(III)(hydr)oxides, a reaction that is catalyzed by iron(II) as discussed in Chapter 9. Hence, iron(II) formation in the photic zone may occur as an autocatalytic process (see Chapter 10.4). This is also true for the oxidation of iron(II). As has been discussed in Chapter 9.4, the oxidation of iron(II) by oxygen is greatly enhanced if the ferrous iron is adsorbed at a mineral (or biological) surface. Since mineral surfaces are formed via the oxidative hydrolysis of Fe(II), this reaction proceeds as an autocatalytic process (Sung and Morgan, 1980). Both the rate of photochemical iron(II) formation and the rate of oxidation of iron(II) are strongly pH-dependent the latter increases with... 

When a spadeful of wet, anaerobic soil is brought to the surface and allowed to dry, air enters through drying cracks and the soil tends to become uniformly oxidized and turn a uniform brown. Whereas when oxidation occurs without drying—as, for example, near a root releasing O2 into wet soil—it is far less uniform and reddish-brown ferric oxide deposits form on and near the oxidizing source. The difference depends on the relative rates of movement of O2 into the soil and of ferrous iron and other reductants in the opposite direction, and the rates of reaction. 

Pecher et al. (2002) show how the uptake of ferrous iron from aqueous solution, by iron oxides, leads to the formation of a variety of reactive surface species that are capable of reducing polyhalogenated methanes (PHMs). The iron oxides used in the experiments and their characteristics are shown in Table 16.2. The PHMs studied include bromodichloromethane (CHBrCl ), chlorodibromomethane (CHBr Cl), bromoform (CHBr ), tetrachloromethane (CCl ), hexachloroethane (HCE), fluorotribromomethane (CFBrj), bromotrichloromethane (CBrCl ) and dibromodichloromethane (CBr Cy. 

Core and valence level photoemission studies of iron oxide surfaces and the oxidation of iron. Surface Sd. 68 459—468 Bruno, J. Sturam, J.A. Wersin, P. Brand-berg, E. (1992) On the influence of carbonate on mineral dissolutions I. The thermodynamics and kinetics of hematite dissolution in bicarbonate solutions at T = 25°C. Geo-chim. Cosmochim. Acta 56 1139—1147 Brusic.V. (1979) Ferrous passivation. In Corrosion Chemistry, 153—184 Bruun Hansen, H.C. Raben-Lange, R. Rau-lund-Rasmussen, K. Borggaard, O.K. 

The dissolution rate of goethite by sulfide was found to increase with surface area and proton concentration. Pyzik and Sommer (21) suggested that HS" is the reactive species that reduces surface ferric iron after exchanging versus OH . A subsequent protonation of surface ferrous hydroxide would lead to dissolution of a surface layer. Elemental sulfur was the prominent oxidation product polysulfides and thiosulfate were found to a lower extent. The dissolution rate R (in moles per square meter per second) of hematite by sulfide was demonstrated to be proportional to the surface concentration of the surface complexes >FeHS and >FeS (22). 

A black coloration that appeared during the experiments at pH values >6.5 indicated the formation of FeS. However, the black disappeared again toward the end of the experiments. Apparently most of the sulfide stored in FeS was also oxidized, and only a small portion of the sulfide may have remained as FeS. This development is not surprising because we worked with excess ferric oxide, in contrast to Pyzik and Sommers study (21) in which FeS could accumulate during the experiments. We concluded that the ferrous iron released after redissolution of FeS adsorbs to the ferric oxide surface and forms a surface complex >FeO-Fe+ (reaction 7), to which poly sulfides may bond. 

Preliminary work (10) on the transition from oxidized surface sediment to reduced subsurface sediment in Milltown Reservoir showed that the redox transition occurs in the upper few tens of centimeters. Strong chemical gradients occur across this boundary. Ferrous iron in sediment pore water (groundwater and vadose water) is commonly below detection in the oxidizing surface zone and increases with depth. Arsenic is also low in pore water of the oxidized zone, but increases across the redox boundary, with As(III) as the dominant oxidation state in the reduced zone. Copper and zinc show the opposite trend, with relatively high concentrations in pore water of the oxidized surface sediment decreasing across the redox boundary. 

Redox Reactions. Aquatic organisms may alter the particular oxidation state of some elements in natural waters during activity. One of the most significant reactions of this type is sulfate reduction to sulfide in anoxic waters. The sulfide formed from this reaction can initiate several chemical reactions that can radically change the types and amounts of elements in solution. The classical example of this reaction is the reduction of ferric iron by sulfide. The resultant ferrous iron and other transition metals may precipitate with additional sulfide formed from further biochemically reduced sulfate. Iron reduction is often accompanied by a release of precipitated or sorbed phosphate. Gardner and Lee (21) and Lee (36) have shown that Lake Mendota surface sediments contain up to 20,000 p.p.m. of ferrous iron and a few thousand p.p.m. of sulfide. The biochemical formation of sulfide is undoubtedly important in determining the oxidation state and amounts of several elements in natural waters. 

Thus, while fresh, unaltered granite matrix rock has little or no iron in the form of iron oxides, alteration zones around fractures do, and technetium sorption may occur there. It should also be noted that sorption of technetium is limited to specific mineral surfaces. Thus, it appears that the reduction of TcO, to a lower oxidation state occurs at or near the surface of the iron oxide and not in the bulk of the solution, by dissolved ferrous ions. 

The relationship of the stirring rate in these experiments to the rates of hydrolysis reactions of basalt phases is indicative of surface-reaction controlled dissolution (21). First order kinetics are not inconsistent with certain rate-determining surface processes (22). Approximate first order kinetics with respect to dissolved oxygen concentration have been reported for the oxidation of aqueous ferrous iron (23) and sulfide (24), and in oxygen consumption studies with roll-type uranium deposits(25). 

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