Lepidocrocite reaction

Similar photo-induced reductive dissolution to that reported for lepidocrocite in the presence of citric acid has been observed for hematite (a-Fe203) in the presence of S(IV) oxyanions (42) (see Figure 3). As shown in the conceptual model of Faust and Hoffmann (42) in Figure 4, two major pathways may lead to the production of Fe(II)ag i) surface redox reactions, both photochemical and thermal (dark), involving Fe(III)-S(IV) surface complexes (reactions 3 and 4 in Figure 4), and ii) aqueous phase photochemical and thermal redox reactions (reactions 11 and 12 in Figure 4). However, the rate of hematite dissolution (reaction 5) limits the rate at which Fe(II)aq may be produced by aqueous phase pathways (reactions 11 and 12) by limiting the availability of Fe(III)aq for such reactions. The rate of total aqueous iron production (d[Fe(aq)]T/dt = d [Fe(III)aq] +... 

Mn(II) oxidation is enhanced in the presence of lepidocrocite (y-FeOOH). The oxidation of Mn(II) on y-FeOOH can be understood in terms of the coupling of surface coordination processes and redox reactions on the surface. Ca2+, Mg2+, Cl, S042-, phosphate, silicate, salicylate, and phthalate affect Mn(II) oxidation in the presence of y-FeOOH. These effects can be explained in terms of the influence these ions have on the binding of Mn(II) species to the surface. Extrapolation of the laboratory results to the conditions prevailing in natural waters predicts that the factors which most influence Mn(II) oxidation rates are pH, temperature, the amount of surface, ionic strength, and Mg2+ and Cl" concentrations. 

This paper discusses the oxidation of Mn(II) in the presence of lepidocrocite, y-FeOOH. This solid was chosen because earlier work (18, 26) had shown that it significantly enhanced the rate of Mn(II) oxidation. The influence of Ca2+, Mg2+, Cl", SO,2-, phosphate, silicate, salicylate, and phthalate on the kinetics of this reaction is also considered. These ions are either important constituents in natural waters or simple models for naturally occurring organics. To try to identify the factors that influence the rate of Mn(II) oxidation in natural waters the surface equilibrium and kinetic models developed using the laboratory results have been used to predict the... 

Fig. 12.11 Reaction schemes suggested for the photochemical reduction of lepidocrocite in the presence of citrate (left) (Waite Morel, 1984, with permission) and ofgoethite in the presence of oxalate (right) (Cornell Schindler, 1987, with permission).

Dos Santos Alfonso and Stumm (1992) suggested that the rate of reductive dissolution by H2S of the common oxides is a function of the formation rate of the two surface complexes =FeS and =FeSH. The rate (10 mol m min ) followed the order lepidocrocite (20) > magnetite (14) > goethite (5.2) > hematite (1.1), and except for magnetite, it was linearly related to free energy, AG, of the reduction reactions of these oxides (see eq. 9.24). A factor of 75 was found for the reductive dissolution by H2S and Fe sulphide formation between ferrihydrite and goethite which could only be explained to a small extent by the difference in specific surface area (Pyzik Sommer, 1981). 

Under otherwise similar conditions, low oxidation rates appear to promote magnetite and goethite, whereas high rates favor lepidocrocite. Magnetite formation probably requires slow oxidation because complete dehydroxylation of the precursor (green rust) prior to complete oxidation is only possible if sufficient time is available if, on the other hand, complete oxidation is fast and precedes dehydroxylation, lepidocrocite forms in preference to magnetite (Schwertmann Taylor, 1977). Dehydroxylation and oxidation appear to be competing reaction steps. 

Ferrihydrite has indeed been found in association with magnetite in Magnetospirillum magnetotacticum (Frankel et ah, 1983). It seems essential that the cell solution is sufficiently buffered to maintain a neutral pH and thus ensure that the solubility product of magnetite is always exceeded. This is a reaction which easily takes place in a purely inorganic system at ambient temperature (see chap. 14). Lepidocrocite has also been suggested as a magnetite precursor (Abe et al. 1983). 

Processings of iron oxides at room temperature. II. Mechanochemical reaction effects on the structure and surface of pure, synthetic lepidocrocite. Mat. Res. Bull. 17 1017-1023... 

Iron frequently has been postulated to be an important electron acceptor for oxidation of sulfide (58, 84,119, 142, 152). Experimental and theoretical studies have demonstrated that Fe(III) will oxidize pyrite (153-157). Reductive dissolution of iron oxides by sulfide also is well documented. Progressive depletion of iron oxides often is coincident with increases in iron sulfides in marine sediments (94, 158, 159). Low concentrations of sulfide even in zones of rapid sulfide formation were attributed to reactions with iron oxides (94). Pyzik and Sommer (160) and Rickard (161) studied the kinetics of goethite reduction by sulfide thiosulfate and elemental S were the oxidized S species identified. Recent investigations of reductive dissolution of hematite and lepidocrocite found polysulfides, thiosulfate, sulfite, and sulfate as end products (162, 163). 

Recently we presented (23) the results of an experimental study on the kinetics and mechanisms of the reaction of lepidocrocite (y-FeOOH) with H2S. With respect to the interaction between iron and sulfur, lepidocrocite merits special attention. It forms by reoxidation of ferrous iron under cir-cumneutral pH conditions (24), and it can therefore be classified as a reactive iron oxide (19). The concept of reactive iron was established by Canfield (19), who differentiated between a residual iron fraction and a reactive iron fraction (operationally defined as soluble in ammonium oxalate). The reactive iron fraction is rapidly reduced by sulfide or by microorganisms. 

In this study we performed initial rate experiments, reacting H2S with lepidocrocite (23). The consumption of H2S was measured continuously by using a pH2S electrode cell (25). To avoid interferences of pH buffer solutions with the iron oxide surface, the pH was stabilized by using a pH-stat that added appropriate amounts of HC1 to the solution. The added volume, which was also continuously monitored, provided information about the amount of protons consumed during the reaction. Dissolved iron was measured only in some runs. 

Figure 1. The oxidation rate of H2S by lepidocrocite is pseudo-first-order with respect to H2S. The experimental pseudo-first-order rate constant k<,/ is plotted as a function of the surface area concentration of y-FeOOH. The reaction rate depends on the surface area (A).

Figure 2. Experimental pseudo-first-order rate constant k0z,v (normalized to the surface area concentration A) for the reaction of H2S with lepidocrocite plotted as a function of pH. Straight lines a and b correspond to eqs 8 and 9, respectively. kw and kj, are the empirical rate constants.

In summary, the reaction of H2S with y-FeOOH is a fast surface-controlled process. Equations 8 and 9 can be used to estimate an upper limit of sulfide oxidation rates in sediments with reactive iron (assuming reactive iron to be represented by lepidocrocite). The surface-area concentration A of reactive iron can be calculated according to... 

In addition to a better understanding of the reaction of sulfide with ferric oxides and its role in pyrite formation, a more exact definition of the term reactive iron is critical. Does reactive iron mean a different iron oxide fraction for bacterial dissolution (e.g., weathering products such as goethite or hematite) than for reaction with sulfide (e.g., reoxidized lepidocrocite) In other words, is there a predigestion of ferric oxides by bacteria that allows a subsequent rapid interaction of sulfide with ferric oxides ... 

Lepidocrocite (-y-FeOOH) has also been used as a catalyst for Fentonlike reactions [54]. First-order decomposition of hydrogen peroxide was observed in the presence of this catalyst. Peroxide decay at 20 g/L catalyst was found to be pseudo-first-order and pH-dependent, with rate constant values reported from 0.102 hr-1 at pH 3.3 to 0.326 hr-1 at pH 8.9. In this system benzoic acid degradation was fastest at the low pH value. Under these conditions, acid dissolution of the lepidocrocite was observed to produce... 

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