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Effect of Solution and Solid-Phase Conditions on the Fe(II)Accelerated Transformation of Ferrihydrite to Lepidocrocite and Goethite Daniel D. Boland,† Richard N. Collins,† Christopher J. Miller,† Chris J. Glover,‡ and T. David Waite*,† †

School of Civil and Environmental Engineering, University of New South Wales, UNSW Sydney New South Wales 2052 Australia Australian Synchrotron, 800 Blackburn Rd, Clayton 3168 Australia

‡

S Supporting Information *

ABSTRACT: Aqueous ferrous iron (Fe(II)) accelerates the transformation of ferrihydrite into secondary, more crystalline minerals however the factors controlling the rate and, indeed, the underlying mechanism of this transformation process remain unclear. Here, we present the first detailed study of the kinetics of the Fe(II)-accelerated transformation of ferrihydrite to goethite, via lepidocrocite, for a range of pH and Fe(II) concentrations and, from the results obtained, provide insight into the factors controlling the transformation rate and the processes responsible for transformation. A reaction scheme for the Fe(II)-accelerated secondary mineralization of ferrihydrite is developed in which an Fe(II) atom attaches to the ferrihydrite surface where it is immediately oxidized to Fe(III) with the resultant electron transferred, sequentially, to other iron oxyhydroxide Fe(III) atoms before release to solution as Fe(II). This freshly precipitated Fe(III) forms the nuclei for the formation of secondary minerals and also facilitates the ongoing uptake of Fe(II) from solution by creation of fresh surface sites. The concentration of solid-associated Fe(II) and the rate of transport of Fe(II) to the oxyhydroxide surface appear to determine which particular secondary minerals form and their rates of formation. Lepidocrocite growth is enhanced at lower solid-associated Fe(II) concentrations while conditions leading to more rapid uptake of Fe(II) from solution lead to higher goethite growth rates.

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INTRODUCTION The nanoparticulate Fe(III) oxyhydroxide ferrihydrite is most often the first solid to precipitate during the hydrolysis of Fe(III) but is thermodynamically unstable with respect to more crystalline oxyhydroxides.1−4 In the absence of dissolved oxygen, Fe(II) may accelerate the crystallization of ferrihydrite to lepidocrocite, goethite and magnetite.5−11 These transformations have wider implications than just the cycling of iron alone as the varying properties of the Fe(III) phases, such as surface area and solubility, and their interaction with Fe(II) can affect the mobility of a wide range of contaminants that are either adsorbed to or coprecipitated with the iron oxyhydroxides.12−17 In general, the Fe(III) oxyhydroxide polymorphs lepidocrocite (γ-FeOOH) and/or goethite (αFeOOH) are observed to form at relatively low Fe(II) concentrations with the formation of these minerals considered to occur by dissolution and reprecipitation mechanisms.5,11,18,19 However, the exact mechanisms via which the presence of Fe(II) accelerates the transformation of ferrihydrite to the more crystalline forms and the factors contributing to the transformation rates are still a matter of debate. Previous work by Burton et al.20 on the transformation of schwertmannite to goethite across a range of pH values showed that the transformation rate was proportional to the © XXXX American Chemical Society

concentration of sorbed Fe(II). Other studies, including our own on the transformation of ferrihydrite to goethite21 and that of Yang et al.19 on the transformation of six-line ferrihydrite to goethite or magnetite, have shown that the degree of Fe(II) sorption cannot fully explain differences in transformation rates. Both these studies suggest that the rate of Fe(II) uptake by the solid is also important in determining transformation rates (and may even affect which secondary minerals form) and that the uptake of Fe(II) by a semiconducting iron oxyhydroxide results in the injection of electrons to the conduction band of the iron oxyhydroxide with the rate of electron injection determined, at least partially, by the concentration of reduced species in solution22 and the difference between the absolute energy of the solution electrons and the conduction band of the (oxyhydr)oxide. Particularly lacking in previous experiments is quantification of ferrihydrite transformation rates over a range of pH and Fe(II) concentrations. In this investigation, we extend our previous study,21 where particularly high Fe(II) (50 mM and Received: September 27, 2013 Revised: April 6, 2014 Accepted: April 11, 2014

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Table 1. Range of Experimental Conditions, Modeled Solid-Associated Fe(II) Concentrations [Fe(II)]solid, the Ratio of [Fe(II)]solid to the Concentration of Ferrihydrite [Fhy], Calculated Absolute Energy of Electrons in Solution and Fitting Parameters for Goethite Growth; Fhy = Ferrihydrite treatment no.

no. samples

1 2 3 4 5 6 7 8 9

7 8 10 9 8 8 9 5 9

pH ± std dev.

[Fe(II)]T (mM)

[Fhy]0 (g L−1)

[Fe(II)]solid (mM)

[Fe(II)]solid/ [Fhy]

electron energy (eV)

k1 (x10−5 s−1)

k2 (x10−5 s‑1)

k3 (x10−5 s−1)

k4 (x10−5 s−1)

R2

± ± ± ± ± ± ± ± ±

1 1 1 1 1 1 2 3 1

0.5 0.5 0.5 0.5 0.5 0.5 0.5 0.5 1

0.063 0.12 0.14 0.34 0.44 0.5 0.12 0.28 0.14

0.011 0.022 0.025 0.062 0.080 0.091 0.022 0.051 0.013

−4.49 −4.44 −4.43 −4.36 −4.33 −4.31 −4.43 −4.40 −4.46

3.1 5.2 4.1 4.7 7.4 4.1 3.5 8.2 5.0

1.7 2.8 2.1 2.1 1.4 0.3 6.7 7.0 2.1

0.2 0.6 1 1.8 10.7 13.8 2.4 5.3 1.1

1.4 2.4 2 7.5 6.0 10.8 24.8 1.2 2.9

0.97 0.98 0.97 0.99 0.99 0.98 0.99 0.96 0.99

6.17 6.42 6.5 6.92 7.13 7.26 6.36 6.5 6.35

0.08 0.05 0.05 0.08 0.04 0.03 0.12 0.05 0.15

1). The suspension was then centrifuged at 3000 rpm and washed with milli-Q water three times, with the resulting solid immediately frozen and freeze-dried. Equilibrium modeling of Fe(II) speciation in the system (including extent of sorbed Fe(II) by the Fe(III) oxyhydroxide) was undertaken using the geochemical code PHREEQC24 employing a two site surface complexation model with stability constants and ferrihydrite surface site density from Appelo et al.25 EXAFS Data Treatment and Analysis. The relative quantities of various Fe(III) oxyhydroxide phases present in the reaction products were determined with Fe K-edge extended X-ray absorption fine structure (EXAFS) spectroscopy. EXAFS measurements were collected at either bending magnet beamline BL-20B (Australian National Beamline Facility) at the Photon Factory (KEK Tsukuba, Japan) or on the XAS beamline (ID 12) at the Australian Synchrotron (Clayton, Australia). In both cases the energy was selected with a channel-cut Si(111) crystal and shifts in beam energy were corrected with an Fe reference foil. The powdered samples were diluted with boron nitride and packed into aluminum slides. Reference spectra of goethite and lepidocrocite, prepared by standard methods2 with purity confirmed by XRD (data not shown), and of pure ferrihydrite were also collected. Spectra were energy calibrated with the software Average and normalized and background corrected with standard features of the ATHENA software package.26 Linear combination fits, using the ferrihydrite, lepidocrocite and goethite references as standards, were also carried out in ATHENA over k = 3−12 Å−1 except for experimental treatments 3 and 8 where, due to poorer data quality, only a fit over k = 3−8 Å−1 was possible.

250 mM) and Fe(III) oxyhydroxide concentrations at pH 6.5 were used, to a wider range of conditions that are more relevant to natural aquatic systems (pH 6.17 to 7.26 and Fe(II) concentrations of 1 to 3 mM). In particular, the effects of pH and Fe(II) concentrations on the transformation kinetics of ferrihydrite to lepidocrocite and goethite are investigated here and, from the results obtained, insights presented regarding the factors controlling both the transformation rate and the nature of the crystalline iron oxyhydroxides produced .

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MATERIALS AND METHODS Sample Preparation. Ferrihydrite was prepared by the addition of 1 M KOH to 0.1 M Fe(NO3)3 and titration to pH 7 with the resulting suspension then centrifuged at 3000 rpm and washed three times with Milli-Q water.2 This stock suspension of ferrihydrite was refrigerated and used within two days. Transformation reactions were initiated within an anoxic chamber (Plas-Laboratories, N2:H2 = 95%:5%) with a palladium catalyst used to remove trace levels of oxygen. All solutions were sparged with high purity argon for two hours to remove oxygen before being transferred to the chamber. The range of experimental conditions examined is shown in Table 1. Transformation studies were undertaken over a range of pH values between 6.17 and 7.26 at a particular constant total ferrous iron concentration ([Fe(II)]T = 1 mM) and initial ferrihydrite concentration ([Fhy]0 = 0.5 g L−1) (treatments 1− 6). Studies were also undertaken at two higher [Fe(II)]T (treatments 7−8) and one higher [Fhy]0 (treatment 9). 50 mM solutions of the organic noncomplexing buffer MES (2-(Nmorpholino)ethanesulfonic acid) were prepared and then adjusted to the necessary pH for each treatment using 1 M KOH. A stock 100 mM Fe(II) solution was prepared inside the anoxic chamber by dissolving FeSO4·7H2O in deoxygenated water. One hundred mL aliquots of MES solution were spiked with the preprepared ferrihydrite and stock Fe(II) to give the requisite concentrations. The bottles of the suspensions were fitted with butyl rubber stoppers and crimp-sealed, removed from the chamber and then shaken on a reciprocating shaker table at 200 rpm. Each experimental treatment involved ∼8−10 sacrificial samples with each sample shaken for different lengths of time ranging from 2 to 70 h. Following reaction, the pH of each suspension was measured and aliquots removed from selected containers, immediately filtered (0.22 μm, Millipore) then analyzed for [Fe(II)] using the ferrozine method.23 Average pH values for each set of samples were considered to be representative of the pH of that set (and are given in Table

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RESULTS AND DISCUSSION Uptake of Aqueous Fe(II). Despite the expected reduction in available surface sites with increasing proportions of crystalline phases present, there were no appreciable consistent changes in pH or measurable change in aqueous Fe(II) concentrations over the length of each transformation experiment. Accordingly, Fe2+ “sorption” data collected for [Fe(II)]T = 1 mM and [Fhy]0 = 0.5 g L−1 were pooled and compared to model output from PHREEQC (Figure 1). In a review by Gorski and Scherer27 and in modeling studies by Hiemstra and van Riemsdijk,28 it has been suggested that the uptake of Fe(II) species by Fe(III) (oxyhydr)oxides in the classical framework employed by surface complexation models is actually more complicated, involving the oxidation of surface Fe(II) to Fe(III) B

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modeling were deemed to be applicable to predicting the loss of Fe(II) from solution, with application of the model extended to other conditions used in the present studies enabling the prediction of solid-associated Fe(II) concentrations under these conditions. The values used are given in Table 1. The maintenance of stable solid-associated Fe(II) concentrations during transformation may result from the higher acidity of goethite compared to ferrihydrite surface sites29 or the noted higher affinity of Fe(II) toward goethite rather than ferrihydrite, despite a reduction in surface area.28 These changing surface properties during transformation may also have implications for the sorption of cations other than Fe2+. Transformation of Fe(III) Phases. Ferrihydrite transformed to goethite via an intermediate lepidocrocite phase (Figure 2 and Supporting Information (SI) Figure S1) in accord with the results of studies undertaken elsewhere under similar conditions.5,9,30 The intermediate lepidocrocite phase was generally more prevalent at lower pH and lower [Fe(II)]T conditions while higher pH and [Fe(II)]T generally resulted in the faster production of goethite. All collected Fe K-edge EXAFS spectra are shown in the SI. Transmission electron micrographs of samples resembled those obtained previously5 with “star-shaped” crystalline assemblages growing out of a ferrihydrite core (Figure 3). Extrapolating backward in time from the first data measurement in each treatment (Figures 2 and S1), it appears that transformation from ferrihydrite does not begin immediately or, at least, is not detected by EXAFS, a technique which probes short-range order. Delays in commencement of transformation to more crystalline forms have however been observed previously11,19 using X-ray diffraction (a technique which probes structure over a larger

Figure 1. Measured solid-associated Fe(II) species (dots) and modeled (line) concentrations, for a total Fe(II) concentration of 1 mM and ferrihydrite loading of 0.5 g L−1, which corresponds to data treatments 1−6. Modeling was undertaken using the geochemical code PHREEQC24 employing a two site surface complexation model with stability constants from Appelo et al.25

followed by electron conduction through the solid. Bearing this in mind, we will refer to “solid-associated Fe(II)” rather than “adsorbed Fe(II)” as the measured loss of Fe(II) from solution indicates it is associated with the solid phase. However, given the good agreement between the measured and modeled “adsorbed” Fe(II) concentrations under these specific conditions, the stability constants used in the surface complexation

Figure 2. Proportion of ferrihydrite (blue square), lepidocrocite (red circle), and goethite (green triangle) over time for treatments 1−6. Each has [Fe(II)]T = 1 mM and [Fhy]0 = 0.5 g L−1, the pH for each treatment is given in the figure. The lines show the nonlinear least-squares fits to the derived expressions for the concentration of each mineral. The gray bar on the x-axis represents an approximate backward extrapolation of the goethite and lepidocrocite data, indicating that transformation does not begin immediately. C

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In this scheme, ferrihydrite (Fhy) initially undergoes transformation to form a reactive ferrihydrite entity (Fhy*) which is then able to transform to either lepidocrocite (Lpd) or goethite (Gth). Lepidocrocite is also able to transform to goethite. No assumption was made as to how these processes occur; simply that they can be considered as pseudo-first order processes with rate constants remaining constant throughout a reaction for each particular treatment. For such a scheme, the decrease in ferrihydrite concentration and change in the proposed “reactive” ferrihydrite over time can be expressed as:

Figure 3. Transmission electron microscope images of samples showing the “star-shaped” crystalline structures of goethite growing from a ferrihydrite core.

distance than EXAFS) and, as such, it may be more than simply an artifact of the analysis. One possibility is that this “induction period” is associated with the time required for dissolutioninduced reorganization of highly disordered surface coatings of highly hydrated ferrihydrite before the formation of more ordered phases. This process may be akin to that described by Banfield et al.31 where oriented iron oxide aggregates appeared on the surface of iron oxidizing bacteria prior to the formation of more crystalline iron oxides. Additionally, Dideriksen et al.32 observed a gradual increase in the size of 2-line ferrihydrite particles on exposure to aqueous Fe(II) before subsequent transformation to a secondary mineral. The core of ferrihydrite is recognized to be more defect-free than the surface layers33 with this more ordered environment representing a favorable substrate from which crystalline iron oxide minerals might develop once the highly disordered surface layers are reconfigured. Based on these general observations, the following scheme was developed in order to quantify the transformation kinetics of the various mineral phases.

[Fhy] = [Fhy]0 e−k1t [Fhy*] =

k1[Fhy]0 [e −k1t −e−(k 2 + k4)t ] k 2 + k4 − k1

where [Fhy]0 is the initial ferrihydrite concentration. As the reactive ferrihydrite, Fhy*, is not distinguishable from “regular” ferrihydrite, Fhy, in the LCF analysis, “total” ferrihydrite, FhyT, is defined as the sum of these two species: [Fhy T] = [Fhy] + [Fhy*] = [Fhy]0 e−k1t +

k1[Fhy]0 [e−k1t − e−(k 2 + k4)t ] k 2 + k4 − k1

⎡ k 2 + k4 ⎤ k1 = [Fhy]0 ⎢ e−(k 2 + k4)t ⎥ e−k1t − k 2 + k4 − k1 ⎣ k 2 + k4 − k1 ⎦

The time-varying concentration of lepidocrocite is given by

Figure 4. Dependence of k1−k4 on solid-associated Fe(II) concentrations for experimental conditions 1−6 ([Fhy]0 = 0.5 g L−1 = 5.5 mM, [Fe(II)]T = 1 mM with varying pH, data points from treatments 1−6 left to right in each figure). D

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driving electron flow from solution to the solid. As can be seen from Figure 5, plots of the deduced k3 and k4 values against the

k1k 2[Fhy]0 (k 2 + k4 − k1)(k 3 − k1)(k 3 − k 2 − k4) [(k 3 − k 2 − k4)e−k1t − (k 3 − k1)e−(k 2 + k4)t + (k 2 + k4 − k1)e−k3t ]

and the concentration of goethite can be found by difference: [Gth] = [Fhy]0 − [Fhy T] − [Lpd]

The values for the rate constants k1, k2, k3, and k4 were determined for each condition using a nonlinear least-squares regression with the Jacobian matrix approximated using centered finite-differences and the parameter estimates refined using a Levenberg−Marquardt algorithm. Assessing how these four rate constants vary with regard to the different experimental conditions should provide some insight into the factors influencing iron oxyhydroxide transformation. The derived rate constants are shown alongside the experimental conditions in Table 1. First, the results for experimental conditions 1−6 will be considered, with constant ferrihydrite (0.5 g L−1) and total Fe(II) concentrations (1 mM) but varying pH (from 6.17 to 7.26). As can be seen from Figure 4, the transformation rate constants of lepidocrocite and reactive ferrihydrite to goethite (k3 and k4, respectively) exhibit a strong positive correlation to solid-associated Fe(II) concentrations, while the rate constant for the transformation of reactive ferrihydrite to lepidocrocite (k2) exhibits a somewhat negative correlation with solidassociated Fe(II) concentrations. The rate constant for transformation of ferrihydrite to reactive ferrihydrite (k1) does not appear to show a strong correlation to solid-associated Fe(II) concentrations for the various conditions examined. The apparent proportionality between the concentration of solid-associated Fe(II) and the rate constants for the transformation of lepidocrocite or reactive ferrihydrite to goethite indicates the importance of this variable in determining transformation rates. The transformation rate constants can also be used to assess the importance of other factors in the Fe(II)-mediated transformation of ferrihydrite to more crystalline iron oxyhydroxides. Of particular interest is the importance of the rate of direct injection of electrons from solution into the conduction band of ferrihydrite via the socalled “semiconductor mechanism” mentioned earlier. The absolute energy of electrons in solution can be readily deduced using the Nernst equation: E(eV) = E 0 + 0.059log

Figure 5. Dependence of k3 and k4 on the calculated energy of electrons in solution for experimental treatments 1−6. ([Fhy]0 = 0.5 g L−1, [Fe(II)]T = 1 mM with varying pH, data points from treatments 1−6 left to right in each figure).

energy of electrons in solution exhibit similar relationships to that observed between these rate constants and the concentrations of solid-associated Fe(II): the magnitude of the transformation rate constants increases as the electron energy (and thus driving force for injection of electrons from solution) increases. It should be noted that the declining solubility of ferrihydrite and thus lower ferric ion activity with increasing pH affects the value of the absolute energy of electrons more than declining ferrous ion activity associated with increasing solid-associated Fe(II) with increasing pH. As such, higher pH values result in both higher solid-associated Fe(II) concentrations and more reducing solution conditions. It is entirely possible then that the solution electron energy, which plays a part in determining rates of electron injection, in addition to or instead of solid-associated Fe(II) concentrations, is important in determining transformation rates. The data from treatments 7 and 8, which have different total Fe(II) concentrations, are inconsistent with the relationships shown so far. Treatment 7, which is differentiated by its higher total Fe(II) concentration ([Fe(II)]T = 2 mM compared to [Fe(II)]T = 1 mM in treatments 1−6) has a low solidassociated Fe(II) concentration (0.12 mM) compared to other treatments but, as can be seen from the results presented in Table 1, exhibits a k4 value that is more than double that of the next highest in the set (that from treatment 6). This result suggests that the solid-associated concentration of Fe(II) alone cannot be used to predict transformation rates, especially the transformation from the reactive ferrihydrite species to goethite. This conclusion is in accord with the findings in our

{Fe 2 +} where E 0 = −5.21 eV 3+ {Fe }

where {Fe2+} is the activity of dissolved ferrous iron,{Fe3+} the activity of dissolved ferric iron and E0 = −5.21 eV is equivalent to 0.77 V on the standard hydrogen electrode scale. {Fe2+} is available from measurement while {Fe3+} can be readily deduced from a knowledge of the solubility of ferrihydrite (given below as Fe(OH)3); i.e. Fe(OH)3 + 3H+ ⇌ Fe3 + + 3H 2O,Fhy K sp = 103.2

If the electron energy is above the energy of the conduction band of the solid phase (−5.00 eV for ferrihydrite19 and −5.08 eV for lepidocrocite34), it is thermodynamically possible for electrons to flow from solution to solid. The difference between the solution electron energy and the energy of the conduction band can then be considered to be the magnitude of the force E

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(or uptake of Fe(II)) will then be proportional to the aqueous Fe2+ concentration and the site concentration:

earlier studies at higher Fe(II) and ferrihydrite concentrations where two treatments with the same solid-associated Fe(II) concentration but different total Fe(II) concentrations showed markedly different transformation rates.21 Similarly, the calculated absolute electron energy (−4.43 eV) for treatment 7 is at the lower end of the electron energies deduced for the range of treatments used in this study despite its high value for k4. Despite this low electron energy, the high rate of transformation can be accounted for by the direct injection mechanism described above since the electron injection rate from solution to a semiconductor is also partially a function of the concentration of the reduced species in solution22 (in this case Fe 2+ ), and treatment 7 had an aqueous Fe(II) concentration (1.88 mM) which was more than double that of treatments 1−6. While the direct injection mechanism appears capable of accounting for the results obtained, uncertainty remains as to the exact mechanism by which the electron enters the soliddoes it occur via formation of innersphere complexes of Fe(II) at surface sites or is it via looser (possibly outer-sphere) associations? Either way, it appears that the uptake of electrons from solution (via dissolved Fe(II)) is important in determining secondary mineral transformation rates. Treatment 8, on the other hand, has a low value for k4, especially given its high solid-associated Fe(II) (0.28 mM) and aqueous Fe(II) concentrations (2.72 mM). This may however be an artifact of the manner in which the transformation rate constants are determined. For example, given that the reaction steps are not completely independent of each other, the fitted constants may be correlated to each other, with multiple “solutions” in some cases giving fits of similar “goodness”. This is particularly likely for data sets with smaller numbers of measurements, as is the case with treatment 8. Treatment 9, in which 1 g L−1 of ferrihydrite was used (compared to 0.5 g L−1 in the other studies undertaken here) exhibits a k4 value that is slightly higher than the k4 values for treatments 2 and 3 which show similar solid-associated Fe(II) concentrations as observed for treatment 9. Values for k1 for treatments 7−9 do not deviate much from those observed in treatment 1−6 (Table 1), making it difficult to ascertain the factors influencing the duration of this initial “priming” period. The demonstrated importance of electron injection rates in determining the iron oxyhydroxide transformation kinetics suggests that a focus on solid-associated Fe(II) concentrations alone may well be misguided with the need to give additional attention to the processes which contribute to maintaining these solid-associated Fe(II) concentrations. Initially, adsorption of Fe(II) to a surface site of an Fe(III) oxyhydroxide (>FeIIIOH) will occur; that is

d ([Fe IIIOFe+]) = kads[>Fe IIIOH][Fe 2 +] dt

Consider a comparison of treatment 2 with treatment 7. Both of these treatments have the same steady-state solid-associated Fe(II) concentrations (0.12 mM) and the same ferrihydrite concentration (and thus the same concentration of surface sites). Treatment 7 has higher aqueous Fe2+ concentration (1.88 mM compared to 0.88 mM). So, despite the equivalence of the solid-associated Fe(II) concentrations, it is clear from the above analysis that treatment 7 will exhibit faster rates of uptake of Fe(II) by the solid. Consider further, treatments 3 and 9. Both have the same aqueous Fe2+ and solid-associated Fe(II) concentrations, but treatment 9 has a higher ferrihydrite concentration (and thus site concentration). These conditions result in a faster uptake of Fe(II) by the solid in treatment 9 than treatment 3. In both of these comparisons, the higher rate of uptake of Fe2+ from solution is associated with a higher value for k4. It appears that either from a “classical” perspective of Fe2+ sorption, based on the affinity of Fe2+ for Fe(III) oxhydroxide surface sites, or from a “semiconductor” perspective in which the uptake of Fe2+ is envisaged to result in injection of an electron into the conduction band of ferrihydrite, both the rate of uptake of Fe(II) as well as the concentration of solid-associated Fe(II), are important determinants of Fe(II)-mediated transformation rates, especially that of ferrihydrite to goethite. Competing Ferrihydrite Transformations to Lepidocrocite and Goethite. The extent of ferrihydrite transformation to either lepidocrocite or goethite can be examined by considering the ratio of k2 to k4 (Figure 6). For treatments

Figure 6. Ratio of rate constants k2 to k4. Experimental treatments 1−6 are shown as filled circles. Treatments 7 and 9 are marked and are represented by unfilled circles. The ratio for treatment 8 (5.9) is off the scale of the graph.

kads

>Fe IIIOH + Fe 2 + ⎯→ ⎯ >Fe IIIOFe+ + H+

Immediately following Fe2+ adsorption to produce the adsorbed species >FeIIIOFe+, Fe(II) will likely be oxidized with resultant release of an electron to an adjacent Fe(III) atom and, from this point, will begin inducing transformation of ferrihydrite to more crystalline phases. To maintain the stable observed aqueous Fe(II) concentrations and the underlying catalytic effect of Fe(II), Fe(II) must also be released to solution with the location of the particular Fe(II) atoms released being dependent upon the electron path within (or at the surface of) the solid matrix. The rate of creation of solid-associated Fe(II)

1−6, these two rate constants are of similar magnitude at low solid-associated Fe(II) concentrations (as a result of the low pH) but production of goethite dominates at higher concentrations. The ratios deviate somewhat from this relationship for treatments 7 and 9, with more dominant goethite growth than may be “expected” in view of their particular solid-associated Fe(II) concentrations. Note that the previously reported effect of carbonate on lepidocrocite versus goethite formation5,35 is not applicable in this study as the experiments were undertaken in an anaerobic chamber with F

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Figure 7. Overview of proposed mechanistic model of ferrihydrite transformation. [1] Ferrihydrite exists with aqueous Fe(II). [2] Fe(II) adsorbs to the surface of ferrihydrite. [3] The adsorbed Fe(II) is oxidized to Fe(III), losing an electron to the solid. [4] The electron moves through the solid and is eventually ejected as Fe(II). The Fe(III) will precipitate as “reactive” ferrihydrite. [5] Aqueous Fe(II) will adsorb to this new phase and [6] be oxidized to Fe(III), losing an electron to the solid. [7] The electron will be conducted through the solid and be ejected as Fe(II). The local instability caused by the electron conduction causes Fe(III) to precipitate as lepidocrocite or goethite (or as reactive ferrihydrite), which provide fresh surfaces for Fe(II) adsorption. [8] Steps [1]−[7] continue, in addition to interaction of lepidocrocite and Fe(II) to form goethite, until only goethite remains.

followed by electron conduction through the solid could also be described in terms of an electron injection from aqueous Fe(II) into semiconducting ferrihydrite. As described previously, the attached Fe(II) will immediately oxidize and release an electron to the solid phase. The local structural instability that is created by this electron injection results in reformation of an Fe(III) oxide with the particular mineral that forms (at the surface) apparently dependent upon the extent to which the matrix is destabilized with possibilities including reformation of reactive ferrihydrite (Fhy*OH), lepidocrocite (LpdOH) or goethite (GthOH). As descibed above, the electron will hop through the solid (or on the surface of the solid) and eventually be released.

sparged solutions. These results corroborate those obtained in previous studies where enhanced lepidocrocite formation was observed at relatively lower values of pH or [Fe2+]5,9,36 with these conditions favoring lower surface coverage by sorbed Fe(II). Similarly, Cornell et al.7 reported a shift in transformation of ferrihydrite to goethite rather than lepidocrocite with increasing amounts of an added reductant (cysteine) which would increase the concentration of ferrous iron present and, in turn, the concentration of solid-associated Fe(II). Toward a Mechanistic Model for Ferrihydrite Transformation. The results of this study suggest the following reaction pathways for the Fe(II)-accelerated transformation of ferrihydrite to secondary minerals (shown schematically in Figure 7): Fe2+ adsorbs to a ferrihydrite surface site (FhyOH):

Fhy*OFe+ + H+ → Fhy*OH + Fe 2 +

FhyOH + Fe 2 + → FhyOFe+ + H+

Fhy*OFe+ + H+ → LpdOH + Fe 2 +

Once on the surface, the Fe(II) will be immediately oxidized, presumably to the “reactive” ferrihydrite species (Fhy*OH) that appears to be a prerequisite for further secondary transformation.

Fhy*OFe+ + H+ → GthOH + Fe 2 +

It is important to note that these freshly formed minerals create new sites to which Fe2+ can attach. For example: Fhy*OH + Fe 2 + → Fhy*OFe+ + H+

FhyOFe+ + H+ → Fhy*OH + Fe2 +(reduced structural Fe(III))

The constant generation of fresh surface sites allows these reactions to continue indefinitely. In a similar fashion, the lepidocrocite formed may transform to goethite. From these reactions it can be recognized that both faster uptake of Fe(II) from solution and higher solid-associated Fe(II) concentrations will increase Fe cycling through this system and the associated iron oxyhydroxide transformations which accompany this cycling.

2+

→ Fhy*OH + Fe (aq)

The electron generated on oxidation of the surface Fe(II) will “hop” from one Fe(III) atom to adjacent Fe(III) atom(s) and will eventually be rereleased into the solution via desorption of an Fe(II) atom, thus completing a cycle whereby a constant concentration of aqueous Fe(II) is maintained (assuming the system is at sorptive steady state). The purported electron hopping has been shown to occur in ferrihydrite at rates similar to those in more crystalline (oxyhydr)oxides.37 Note that this Fe(II) uptake from solution

LpdOH + Fe 2 + → LpdOFe+ + H+ → GthOH + Fe 2 +

Exchange with solution Fe(II) would be expected to still occur with goethite but with no further mineral transG

dx.doi.org/10.1021/es4043275 | Environ. Sci. Technol. XXXX, XXX, XXX−XXX

Environmental Science & Technology

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formations.38 Which secondary minerals form, and at what rate they form, will depend on the partitioning of Fe(II) between aqueous and solid-associated forms, and the rate at which Fe(II) is taken up by the solid (and recycled through the solid back into solution). Lepidocrocite production will be highest at relatively low surface associated Fe(II) concentrations, with these conditions more likely at lower pH (