Early Stage Formation of Iron Oxyhydroxides during Neutralization of

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Early Stage Formation of Iron Oxyhydroxides during Neutralization of Simulated Acid Mine Drainage Solutions Mengqiang Zhu,†,* Benjamin Legg,‡ Hengzhong Zhang,‡ Benjamin Gilbert,† Yang Ren,§ Jillian F. Banfield,†,‡ and Glenn A. Waychunas† †

Earth Sciences Division, Lawrence Berkeley National Laboratory, Berkeley, California 94720, United States Department of Earth and Planetary Science, University of California-Berkeley, Berkeley, California, 94720, United States § Advanced Photon Source, Argonne National Laboratory, Argonne, Illinois 60439, United States ‡

S Supporting Information *

ABSTRACT: The phases and stability of ferric iron products formed early during neutralization of acid mine drainage waters remain largely unknown. In this work, we used in situ and time-resolved quick-scanning X-ray absorption spectroscopy and X-ray diffraction to study products formed between 4 min and 1 h after ferric iron sulfate solutions were partially neutralized by addition of NaHCO3 ([HCO3−]/[Fe3+] < 3). When [HCO3−]/[Fe3+] = 0.5 and 0.6 (initial pH ∼ 2.1 and 2.2, respectively), the only large species formed were sulfatecomplexed ferrihydrite-like molecular clusters that were stable throughout the duration of the experiment. When [HCO3−]/[Fe3+] = 1 (initial pH ∼ 2.5), ferrihydrite-like molecular clusters formed initially, but most later converted to schwertmannite. In contrast, when [HCO3−]/[Fe3+] = 2 (initial pH ∼ 2.7), schwertmannite and larger ferrihydrite particles formed immediately upon neutralization. However, the ferrihydrite particles subsequently converted to schwertmannite. The schwertmannite particles formed under both conditions aggregated extensively with increasing time. This work provides new insight into the formation, stability and reactivity of some early products that may form during the neutralization of natural acid mine drainage.

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INTRODUCTION

Fe oxyhydroxide phase stability and their behavior in the environment. It is generally believed that solvated Fe3+ undergoes hydrolysis with subsequent formation of low-molecular-weight Fe3+ species (hydrated and hydrolyzed monomers, dimers, trimers, etc.).6 Fe oxyhydroxide particles may nucleate from the polynuclear species6 and grow via atom-by-atom addition or oriented aggregation.7 In addition to hydrolysis, olation and oxolation processes are responsible for these condensation processes.8 The resultant Fe oxyhydroxide phases depend on anion type and pH. The impact of anions in phase organization originates from their Fe3+ binding ability, and this is strongly affected by pH. In general, weak Fe3+ binding anions, such as NO3− and ClO4−, favor ferrihydrite formation. Stronger Fe3+binding ligands, such as Cl− and SO42‑, lead to formation of akaganeite and schwertmannite, respectively, under acidic conditions. Under neutral and alkaline pH conditions, ferrihydrite tends to be the resultant phase because OH− outcompetes Cl− and SO42‑ to bind Fe3+. In the presence of very strong Fe3+-binding ligands, such as phosphate and silicate,

When exposed to air and water, metal sulfide deposits and coal mines produce iron-rich sulfuric acid solutions that can contain a broad range of toxic elements.1,2 These solutions, referred to as acid mine drainage (AMD), degrade environmental quality. Neutralization of AMD occurs when solutions mix with natural waters, such as natural rivers and streams, during wetland treatment,3 through reaction with rock surfaces, or following lime addition in remediation processes. As a result, ferric iron (Fe3+) hydrolyzes and precipitates as Fe oxyhydroxide nanoparticles such as ferrihydrite and schwertmannite. These Fe oxyhydroxide nanoparticles can transform to more thermodynamically stable phases goethite and hematite with time. During the formation of Fe oxyhydroxide nanoparticles, toxic elements in solution can be incorporated into the nanoparticle structure. During or after formation, contaminants can also be adsorbed onto nanoparticle surfaces or trapped in interfacial regions between aggregated nanoparticles. Such processes influence contaminant fate and transport.4 Nanoparticles also can interact with microorganisms and their extracellular polymers,5 and hence extracellular substances, and other solution species may promote or retard particle growth, transformation and aggregation. These processes can control © 2012 American Chemical Society

Received: Revised: Accepted: Published: 8140

March 30, 2012 June 25, 2012 July 5, 2012 July 5, 2012 dx.doi.org/10.1021/es301268g | Environ. Sci. Technol. 2012, 46, 8140−8147

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and two-line ferrihydrite reference materials were synthesized using the procedures described by Schwertmann and Cornell.14 UV-Vis Spectroscopy and pH Measurement. Each of the mixed solutions was immediately transferred to a quartz cuvette with 1 mm path length for time-resolved UV-vis spectra collection using an Agilent UV-vis spectrophotometer (model 8453). UV-vis spectra were continuously recorded with 5 s of integration time. Time-dependent pH values of the above solutions were automatically recorded every 20 s. The pH meter was calibrated using pH 1, 1.68, 2, and 4 buffer solutions. Synchrotron X-ray Diffraction (SXRD). Time-resolved SXRD 2-D images were collected at beamline 11-ID-B at the Advanced Photon Source (APS). The high flux of the X-rays of this beamline is the key for collecting high signal-to-noise ratio XRD data from particles in dilute particle suspensions. A detailed description for XRD collection and processing is provided in the Supporting Information (SI-1). QEXAFS Spectroscopy. Each of the mixed solutions was transferred to a 2 mm thick, 15 mm long, and 10 mm wide cell made from acrylic plastic plate. Both sides of the cell were sealed with Kapton film to hold the solutions. QEXAFS spectra were collected at one scan per 2 s in transmission mode. Fifteen or 14 continuous scans of QEXAFS spectra were averaged to improve counting statistics, resulting in a time-resolution of 30 s. The measurements were conducted at beamline X18A at the National Synchrotron Light Source (NSLS), using a specially equipped monochromator with a Si (111) crystal. The monochromator was detuned by 35% with respect to I0 to minimize higher order harmonic X-rays. The detailed beamline and monochromator setup for the QEXAFS measurement can be found elsewhere.15 Powder samples, including the Fe oxyhydroxide references and the solid precipitates in the mixed solutions at h = 1 and 2, were ground finely and uniformly spread on Scotch tape. The tapes were folded to four or eight layers for EXAFS data collection in transmission mode at beamline 4-1 at the Stanford Synchrotron Radiation Lightsource (SSRL), using a Si (220) monochromator crystal. After background removal and normalization the EXAFS spectra were converted into the k-weighted function, k3χ(k), and Fourier transforms |χ(R)| were calculated over the k range of 3−14 Å−1 using the Bessel-Kaiser window function. Linear combination fitting (LCF) analysis of QEXAFS spectra was performed over 3−12 Å−1 in k space. In the fits, the component weight sum was forced to equal 1.0. The program Athena16 was used for all of the above data reduction and processing. Further EXAFS shell-by-shell fitting was performed using the SixPack program suite17 to obtain the detailed local atomic environment of iron atoms for representative samples. Ferron Assay. Ferron (8-hydroxy-7-iodo-5-quinoline sulfonic acid) has been used for semiquantification of Fe3+ monomers, polymers and precipitates based on distinct ferron−Fe complex formation rates. Fe3+ monomers are complexed by ferron rapidly, but the breakdown of larger Fe3+ species to form ferron−Fe complexes occurs much more slowly.18 It has been assumed that reactions of monomeric Fe3+ species with ferron are complete within 1 min (denoted as Fe_monomer).18 Species that form ferron−Fe complexes over the next 12 h were interpreted as Fe polymers, and any Fe species that remained unreactive after 12 h were considered as stable precipitates.18 Following this method, we used the ferron assay to estimate Fe3+ precipitation reaction progress during neutralization of ferric salts by NaHCO3, based on Fe_mo-

Fe oxyhydroxide particle formation can be substantially impaired.9,10 Aqueous speciation in pure Fe3+ solutions is relatively well understood. In recent years, much has been learned about nanoparticle structure, growth, and phase transformation behavior.11 However, a major gap in knowledge concerns the nucleation pathway whereby aqueous monomeric Fe3+ species are converted into molecular clusters and nanoparticles. Prior work attempted to study early stage Fe oxyhydroxide formation kinetics and mechanisms in partially neutralized Fe3+ solutions using various methods, including titration, UV-vis spectroscopy, wide-angle X-ray scattering diffraction (WAXS), small-angle Xray scattering (SAXS), extended X-ray absorption fine structure (EXAFS) spectroscopy, electron microscopy, and analytical centrifugation.6,12,13 Effects of anion type, ratio of neutralization base to Fe3+, Fe3+ concentration, aging period, and temperature, were examined. However limited knowledge was obtained about the precipitate phase, average local atomic structure, particle size and morphology, and aggregate structure. In particular, due to the limited time-resolution of the experimental methods, these investigations were confined to late reaction stages (>30 min) when particle aggregation and phase transformation comprise the major events. Although the final precipitate phases could be identified in these studies, the forms and structures of the early formed products in the neutralization process remained largely unknown, and hence their formation kinetics and mechanism are yet to be explored. In the current study, we were able to investigate the formation kinetics and phase identities of early formed products in partially neutralized Fe3+-sulfate solutions using time-resolved UV-vis spectroscopy, quick-scanning EXAFS (QEXAFS) spectroscopy, and synchrotron XRD, by combining these rapid methods with a slowed neutralization process.

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EXPERIMENTAL SECTION Sample Preparation. All chemicals used were of analytical grade. A 0.8 M Fe3+ stock solution ([SO42‑]/[Fe3+] = 1.5) was prepared by adding 4.16 g of ferric iron sulfate powder (containing 21−23% Fe by weight) to 16 mL deionized (DI) water. The Fe solution was used within 48 h after preparation, although there were no visible changes in color and turbidity over 1 month. Reactions were initiated by slowly pumping 4 mL DI water, 0.2 M, 0.24 M, 0.4 M, or 0.8 M NaHCO3 solution at 5 mL·min−1 using a syringe pump, into a 30 mL plastic bottle containing 4 mL of 0.4 M Fe3+ sulfate solution (2 mL 0.8 M Fe stock solution plus 2 mL DI water) to achieve [HCO3−]/[Fe3+] molar ratios (represented by h) of 0, 0.5, 0.6, 1, and 2. An acidified Fe solution was prepared by adding 4 mL of concentrated HNO3 to 4 mL of 0.4 M Fe3+ sulfate solution, referred to as FeAcid. This solution is supposed to contain mainly Fe(H2O)63+. The final concentration of total Fe3+ in all solutions was 0.2 M. Vigorous stirring was enforced during the base addition to minimize local oversaturation with respect to Fe oxyhydroxide phases. Each base addition step took 48 s. Upon completion, one additional minute of stirring was used to remove CO2 gas resulting from NaHCO3 decomposition during the neutralization reaction. Then, the solutions were immediately transferred for the following time-resolved analyses. Measurements of reaction time started from the first drop of base addition. When the first data sets were recorded, about 2−5 minutes had elapsed, depending on time-resolved analytical approaches (see below). All experiments were conducted at room temperature of 24−25 °C. Schwertmannite 8141

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addition, ferric Fe species with an extended structure, such as Fe oxyhydroxide particles, have a characteristic absorption feature called electron pair transition (EPT) band.22 The EPT band is significant for particles but very weak for molecular clusters, and thus can be used to differentiate the two types of Fe species. Figure 2a−d shows UV-vis spectra of the solutions at different h. For clarity, the spectra were presented in two different wavelength ranges. The spectra of h = 1 and h = 2 solutions evolved with time whereas others remained the same. The acidified solution (FeAcid) only had a very weak band at ∼806 nm (Figure 2b). The h = 0 solution had two broad absorption bands centered at approximately 425 and 830 nm. For the h = 0.5 and 0.6 solutions, three bands at about 430, 448, and 850 nm can be discerned (Figure 2b and d), and the h = 0.6 solution bands are more intense. The barely changed spectra of the above solutions indicate that Fe species in these solutions were stable within the experimental time period. The first spectrum of the h = 1 solution also showed the same three bands as the h = 0.5 solution (Figure 2a and b). However, the first two bands gradually disappeared with reaction time, and the 850 nm band decreased its intensity and shifted its position to 862 nm. Another band at 485 nm, not obvious in the first spectrum of the h = 1 solution, emerged and became increasingly pronounced (Figure 2b). This band did not exist in the spectra of the h = 0, 0.5, and 0.6 solutions. The observed bands at 425, 430, and 448 nm can be assigned to 6A1 g → (4A1, 4Eg) transitions and the bands at 830, 850, and 862 nm to 6A1 g→ 4T1 g transitions.11,21 The substantial existence of all these bands indicates formation of Fe clusters or particles.21,23 The 485 nm band is assigned to EPT.22 The absence of this band in the h = 0, 0.5, 0.6 and the initial h = 1 solutions suggests that the Fe species were molecular clusters. The disappearance of the three bands of the h = 1 solution shows that the Fe cluster was decreasing in abundance, whereas the increasing EPT absorbance indicates that a Fe oxyhydroxide phase was forming and increasing in abundance. The UV-is spectrum of the filtered h = 1 solution after 2 h of reaction time still showed the three bands, although they were much less intense (Figure 2b and d), indicating that the Fe clusters had not completely disappeared. Regarding the h = 2 solution, the EPT band at 485 nm is pronounced in the first collected spectrum, indicating that Fe oxyhydroxide particles formed from the beginning. A band at 865 nm (assigned to 6A1 g→ 4T1 g) decreased in intensity with time without significant position change. In addition, the absorbance between 525 and 600 nm decreased with time, suggesting disappearance of a metastable phase. The filtered h = 2 solution after 2 h had a weak peak at 850 nm besides the 485 nm band (Figure 2b and d), indicating the solution also had a small amount of Fe clusters similar to those present in the h = 1 solutions. The lack of obvious 430 and 448 nm bands in this case could be due to the overlap with the intense absorbance tail (300−600 nm) from the larger particles. Phase Analysis by Synchrotron XRD (SXRD). To identify the phases of particles being formed, in situ SXRD data were collected from the neutralized solutions. As shown in Supporting Information Figure SI-1, the first XRD pattern of the h = 1 solution cannot be firmly attributed to known phases due to the low particle abundance and/or small particle size. The first pattern of the h = 2 solution presented broad peaks, indicating formation of very fine crystalline particles. With increasing reaction time, the initial broad peaks of both h = 1

nomer fractions measured at different times as the neutralization reaction proceeded. The specific experimental procedures were described in the Supporting Information (SI-2). The accuracy of ferron assay-based inferences was re-evaluated in light of other results of this study (see below).

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RESULTS Color and Turbidity. The visible appearance of the solutions immediately following bicarbonate addition was dependent on the base to ferric iron molar ratio (h = [HCO3−]/[Fe3+]). All h ≥ 0.5 solutions displayed red wine color, and the color deepened as h increased. Moreover, the cloudiness (or turbidity) of the h = 1 and h = 2 solutions increased over time, whereas other solutions remained clear during the experimental period. To monitor the turbidity changes with time, the measured absorbance at 600 nm was used as a turbidity indicator, as no samples had UV-vis absorption bands at this wavelength (see below). Assuming that Rayleigh scattering applies, increased absorbance was due to increase of size and/or numbers of scattering objects.19 As shown in Figure 1, absorbance at 600 nm increased with

Figure 1. UV-vis absorbance at 600 nm versus reaction time for Fe sulfate solutions at various h values (h = [NaHCO3]/[Fe3+], i.e., neutralization ratio). Since no significant UV-vis absorption bands exist at this wavelength, the measured absorbance was due to particle scattering, indicating turbidity of the solutions.

increasing h when reaction time < 27 min, indicating particle formation. For the h = 1 solution, the 600 nm absorbance increased in the first ∼10 min (see inset) and dramatically after ∼26 min. For the h = 2 solution, an initially subtle decrease was followed by a dramatic increase in scattering after ∼28 min (Figure 1). The dramatic increases correlated with the transitions from clear red wine-like solutions to cloudy yellow suspensions. The cloudiness was due to the massive formation of large particle aggregates.19 The aggregates then settled to form precipitates that are referred to as Ppt.h1 and Ppt.h2 for h = 1 and h = 2, respectively. Fe3+ Molecular Clusters and Particles Determined by UV-vis Spectroscopy. Fe3+ monomers, molecular clusters and particles have different optical excitation features that can be used for identification. Ligand-field or d → d transitions are spin forbidden and thus are extremely weak in Fe 3+ monomers.20,21 By contrast, the ligand-field bands of Fe3+ clusters and particles are intense mainly due to antiferromagnetic coupling between next neighboring Fe atoms.11,21 In 8142

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Figure 2. UV-vis spectra of Fe sulfate solutions at various h values. The spectra of the h = 1 and 2 solutions were time dependent and the arrows indicate the reaction directions with increased reaction time (2 min < t < 27 min). The “filtered” spectra were collected after filtration of the cloudy solutions to remove solids. For closer inspection, the UV-vis spectra were presented in two different wavelength regimes (note that the absorbance scales differ). The spectral amplitude of the h = 1 and 2 solutions increased over the whole wavelength range due to particle scattering, correlating with the transition from clear to cloudy solutions due to onset of precipitation. The spectra collected after the solutions became cloudy (top black curves in panels (a) and (c)) are not shown in (b) and (d).

Figure 3. Fe K-edge EXAFS spectra of HNO3 acidified Fe sulfate solution (FeAcid) (brown), the neutralized Fe sulfate solutions at different reaction times and different h values (pink, green, cyan, blue), the final precipitates at h = 1 and 2 (red), with 2-line ferrihydrite (Fhy2L) and schwertmannite (Schw) reference spectra. (a) k-weighted χ(k) and (b) the magnitude (|χ(R)|) of their Fourier transforms.

and h = 2 solutions sharpened (2.4 Å−1) and split (4 Å−1 and 5.5 Å−1), increasingly resembling those of the schwertmannite reference phase. The SXRD results showed that the precipitates (Ppt.h1 and Ppt.h2) formed in the two solutions were primarily schwertmannite (Supporting Information Figure SI-1). However, XRD phase characterization for particles in solutions based on poorly defined peaks has large uncertainties due to solution interference. In the following two sections, QEXAFS spectroscopic analyses, including EXAFS curve fitting and linear

combination fitting (LCF), were used to further identify the particle phases and quantify their amounts versus reaction time. Local Atomic Environment of Fe. Fe K-edge EXAFS spectroscopy provides information about the local atomic environment around central Fe (within 6 Å). Figure 3 displays the results of in situ QEXAFS spectroscopy acquired from samples under the same conditions as the XRD data of Supporting Information Figure SI-1. The peaks at 1.5−1.6 Å (R +ΔR) in |χ(R)| are assigned to the nearest O shell at a Fe−O distance of ∼2 Å, corresponding to Fe−O bond lengths. As the 8143

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Figure 4. Molar fractions of ferrihydrite (red circles), schwertmanite (blue triangles), and h = 0.5 solutions (green triangles) were obtained from linear combination fitting for (a) h = 1 and (b) h = 2. Values measured via the ferron assay are indicated by black boxes. Comparison of the inferred molar fraction of ferrihydrite plus the h = 0.5 solution (open boxes) to the ferron assay values indicates that the ferron assay actually measures both of these components at h = 1.

introduces further diversity.25,26 As it is impossible to identify and prepare LCF standards for all possible Fe species in the neutralized solutions, we assume that the h = 0.6, 1, and 2 solutions can be represented by a combination of solutions at lower neutralization ratios (representing monomeric Fe3+ and small Fe3+ clusters) and Fe oxyhydroxide solid phases approximating particles in aqueous solutions. FeAcid, the h = 0, 0.2, and 0.5 solutions, schwertmannite, Fhy2L, goethite, lepidocrocite, hematite, Ppt.h1, Ppt.h2, and their combinations, thus were tested with LCF analysis to determine if well-defined general fits could be obtained with a minimum set of components. It was found that a combination of the h = 0.5 solution, Ppt.h2 and Fhy2L gave the best fits with the least number of components (fits in Supporting Information SI-6 and results in Figure 4). Use of Ppt.h2 and h = 0.5 solution only in the fits decreased the goodness of fit (in terms of reduced χ2) up to 33% (the first spectrum of the h = 1 solution), showing necessity of the Fhy2L component. The LCF analysis indicates that Ppt.h1 contained Fhy2L at a phase fraction 17% and Ppt.h2 at 83%. The h = 0.6 solution contained Fhy2L at 4% and the h = 0.5 solution at 96%, but this was subject to large uncertainties due to the low abundance of Fhy2L. Although the EXAFS spectrum of Ppt.h2 was almost identical to that of schwertmannite reference phase (Supporting Information SI-5), replacing Ppt.h2 with the schwertmannite reference slightly decreased the fitting accuracies but gave similar component fraction values. Henceforth, for clarity, we use “schwertmannite” to represent “Ppt.h2” in further discussion of the LCF results that were summarized in Figure 4. The Fhy2L component accounted for ∼20% of Fe species initially in the h = 1 solution whereas the schwertmannite quantity was negligible (Figure 4a). With time, the Fhy2L component fraction gradually decreased to ∼5% whereas schwertmannite developed to 15−20% (Figure 4a). About 80% of the Fe could be represented by the h = 0.5 solution and this quantity stayed almost constant with time, indicating formation of schwertmannite from conversion of Fhy2L. For the h = 2 solution, both Fhy2L and schwertmannite accounted for ∼30%, respectively, initially (Figure 4b). With time, the Fhy2L component gradually disappeared and the schwertmannite fraction increased to ∼70%. The h = 0.5 solution component decreased from 40% to 30% with time.

neutralization degree increased from the acidified solution to the h = 2 solutions, the Fe−O bond length shortened, the O peaks broadened, and the amplitudes decreased (Figure 3). The bond length shortening can be due to the formation of particles and clusters that usually have shorter Fe−O bonds in Fe−OH and Fe−OFe moieties compared to those in Fe−OH2 of monomeric Fe species. These changes also widen the overall Fe−O bond length distribution and cause destructive interference in the χ(k) functions and accordingly the reduced O peak amplitudes in their Fourier transforms.24 The wider Fe−O bond distribution is consistent with the increasing Debye−Waller factors (σ2) obtained from fitting using a single O shell (Supporting Information SI-3). The peaks at ∼2.6 and ∼3.05 Å (R+ΔR)| are assigned to the edge- and corner-sharing Fe shells at the Fe−Fe distances of ∼3.0 and ∼3.35 Å, respectively. Sulfate can complex with Fe25,26 and sulfur scattering may slightly modify the peaks corresponding to approximately these distances. As the neutralization degree increases, the two Fe peaks become more and more pronounced, indicating Fe polymerization. FeAcid did not exhibit the edge-sharing Fe peak (Figure 3), showing that the high concentration of acid suppressed formation of Fe polymers. The pronounced peaks between 2.8 and 3.75 Å (R+ΔR) may result from both multiplescattering paths within the first O shell of Fe octahedra and single-scattering paths from the O atoms in the second hydration shell.27 The peak at 3.1 Å (R+ΔR) overlaps with the corner-sharing Fe peak, making it hard to discern the Fe shell when particle abundance is low, such as the spectra of the h < 1 solutions and the initial spectrum of the h = 1 solution. For the h = 1 and 2 solutions, with increasing reaction time, the O shells did not change significantly, but the two Fe shells underwent substantial increases in peak amplitude and shifts in peak positions (representative spectra in Figure 3 and the complete set in Supporting Information SI-4). As revealed by further analysis described below, these changes are mainly due to particle phase transformation. Time-Dependent Evolution of Fe3+ Species. EXAFS linear combination fitting (LCF) analysis can be used to identify and quantify each species in a system containing mixed Fe phases.28 We anticipate that the neutralized solutions contain many different Fe species, and complexation by sulfate 8144

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Figure 5. (a) The fraction of Fe species (Fe_monomer) that reacts with ferron within 1 min, and (b) the change in pH over time for h = 0 to h = 2 solutions. The empty symbols in (a) are replicates measured after filtration to remove precipitates. The solid symbols indicate the fraction of Fe species measured by ferron assay for unfiltered solutions.

schwertmannite over time. Combining the above two analyses indicates that the initially formed Fhy2L component, detected as molecular clusters in the h = 1 solution and as particles in the h = 2 solution, transformed to schwertmannite over time. UVvis and EXAFS results are also consistent with the SXRD data that schwertmannite particles were present initially in the h = 2 solution and not in the h = 1 solution, but the SXRD analysis failed to identify the ferrihydrite particles. The XRD peak sharpening and splitting seen in both solutions were due to the transformation of the Fhy2L particles and clusters, showing poorly defined XRD peaks, to schwertmannite particles showing better defined peaks. The clusters in the h = 1 solution were too small to contribute much to the turbidity measured at 600 nm (I ∼ d6/λ4). Therefore, the observed gradual turbidity increase of the h = 1 solution was mainly due to the formation of schwertmannite particles. Transformation of Ferrihydrite to Schwertmannite. The transformation of the ferrihydrite clusters and particles to schwertmannite indicates that schwertmannite was the thermodynamically more stable phase in the h = 1 and h = 2 solutions. This is anticipated, given the presence of sulfate anions. However, the transformation in the h = 1 solution was not complete (5% remained). In addition, ferrihydrite clusters existed in the h = 0.5 and 0.6 solutions without transforming to schwertmannite within the studied time period. These observations indicate that ferrihydrite clusters were more stable than schwertmannite in the presence of sulfate under more acidic conditions, although their abundances were low. Given the large structural difference of the two phases, the transformation likely occurred via dissolution and reprecipitation. The EXAFS LCF result shows that 70% of Fe existed as schwertmannite eventually in the h = 2 solution, greater than the sum (60%) of the initial ferrihydrite and schwertmannite fractions. A second nucleation event occurred at 30 min in the h = 2 solution, as suggested by the sudden drop of the dissolved Fe concentration in the ferron assay (Figure 5a). Therefore, the additional 10% schwertmannite resulted from the formation of new nuclei rather than growth of the existing particles. This indicates that schwertmannite can directly form from dissolved Fe species, and that ferrihydrite is not an indispensible precursor for its formation. According to a previous study using an electric field jump relaxation kinetic technique,29 hydrolysis reactions of monomeric Fe3+ are extremely fast and should be complete within

As the two Fe peaks of Fhy2L have lower amplitude and longer corner-sharing Fe−Fe distance than those of schwertmannite (Figure 3), the relative changes of the two Fe phases with time resulted in increasing amplitude of the two Fe peaks and shortening of the corner-sharing Fe distance (Supporting Information SI-4). For the same reason, the Fe peaks of the h = 1 solution were not significant initially, though Fhy2L accounted for 20%. Reaction Progress Measured by Ferron Assay and pH Changes. In addition to the above EXAFS LCF analysis, the ferron assay was also used to determine reaction progress of particle formation. Figure 5a shows the time-dependent change in the proportion of Fe3+ species that react with the ferron probe within one minute, assumed to represent monomeric ferric iron. Initially, almost all Fe was in the Fe_monomer fraction for the h = 1 solution, whereas less than half of the Fe was in this fraction for the h = 2 solution (Figure 5b). In addition, the ferron-Fe formation rates of h = 1 and 2 solutions gradually decreased with the increasing neutralization reaction time (Supporting Information Figure SI-2). In contrast, for the h = 0, 0.5, and 0.6 solutions, ferron−Fe complex formation reaction completed almost instantaneously (data not shown) and the rates did not change with the neutralization reaction time, indicating that the solution was dominated by the Fe_monomer fraction. It is noteworthy that for the h = 2 solution, a second abrupt decrease in Fe_monomer occurred at ∼30 min when the solution became cloudy, possibly due to a second nucleation event. Fe3+ oxyhydroxide particle formation is a hydrolysis and condensation process, releasing H+. Therefore, the pH change over reaction time reflects particle formation progress. As Figure 5b shows, the pH of the h = 1 and 2 solutions slowly decreased with the reaction time, suggesting particle formation.

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DISCUSSION According to the UV-vis results, Fe molecular clusters formed initially in the h = 1 solution but transformed to Fe oxyhydroxide phases over time; in contrast, the initially formed Fe species in the h = 2 solution were not molecular clusters but Fe oxyhydroxide particles, some of which decreased and others increased in abundance with time. The QEXAFS LCF analysis shows that the Fhy2L component was the only phase formed initially in the h = 1 solution, whereas for the h = 2 solution, both Fhy2L and schwertmannite phases were present initially; the Fhy2L components in both solutions transformed to 8145

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seconds. The slow pH decreases of the h = 1 and 2 solutions must be due to other slower chemical reactions. In this case, the pH decrease can be attributed to the phase transformation, that is schwertmannite formation from ferrihydrite clusters and particles leads to a net release of protons. Phase transformationinduced pH decreases have been observed previously in ferrihydrite transformation to goethite, hematite, and lepidocrocite.30 Invalidation of Ferron Assay for Fe Monomers. The LCF-determined soluble Fe fraction (represented by the h = 0.5 solution) in the h = 2 solution is basically consistent with the ferron assay result (Figure 4b). The large discrepancy between the two measurements for the h = 1 solution (Figure 4a) is removed if the ferron assay actually measures both Fhy2L (clusters) and the Fe species approximated by the h = 0.5 solution. The finding that the Fhy2L-like clusters reacted with ferron within one minute invalidates the ferron assay for quantification of Fe monomers. There is less evidence for reaction of ferron with the Fhy2L particles in the h = 2 solution within one minute, consistent with the conclusion that the main component was Fhy2L particles. The decreased ferron−Fe complex formation rates with the neutralization reaction time showed that the developing schwertmannite particles were less reactive than ferrihydrite clusters and particles regarding ferron−Fe formation. The Ferrihydrite Components. The Fhy2L-like cluster in the h = 1 solution might resemble the Fe13 motif in the ferrihydrite structure, that is, [FeO4Fe12(OH)24(H2O)12]7+.31 This molecular cluster was proposed by Bradley and Kydd32 decades ago based on infrared spectroscopy, although few studies followed this suggestion. These authors found this cluster was unstable, but could be stabilized by sulfate binding. The structure of the aluminum Keggin cluster (Al13), [AlO4Al12(OH)24(H2O)12]7+, was determined after crystallization of the cluster from a solution as a sulfate salt.33 It is not surprising that sulfate binding could stabilize similar Fe13 molecular clusters in solution, possibly by preventing their growth into larger ferrihydrite particles and their dissolution by acid as well. The sulfate binding becomes weaker at higher pH.34 Therefore, when adding more base for the h = 2 case, OH− could outcompete sulfate to bind with Fe3+ facilitating the growth of ferrihydrite clusters to large particles. This is consistent with the fact that ferrihydrite can be synthesized by rapidly increasing Fe3+ sulfate solution pH to pH 7. On the other hand, the weak sulfate binding at h = 2 was not able to protect ferrihydrite from reaction, and accordingly ferrihydrite converted completely as the pH was still very acidic (pH < 2.65). The Fhy2L-like cluster was also present in the h = 0.5 and 0.6 solutions, according to UV-vis spectroscopy. Since the h = 0.5 solution was one of the three components used in the LCF analyses of the h = 0.6, 1, and 2 solutions, the abundances of the Fhy2L-like clusters in these solutions were underestimated. This means that the clusters were also present in the h = 2 solution, consistent with the UV-vis spectrum of the filtered solution, although the majority of the ferrihydrite units were particles. Abundance of the Clusters. Assuming the cluster contains 13 Fe atoms, based on the LCF-determined fractions and UV-vis absorbance differences at 850 nm, we estimated its absorption coefficient (μ) as 44.5 mol−1·cm−1 using the h = 1 and h = 0.5 results and 52.0 mol−1·cm−1 using the h = 0.6 and h = 0.5 results, with an average value of 48.3 ± 3.7 mol−1·cm−1.

Using this value the cluster concentrations can be estimated as 2.4 ± 0.2, 3.1 ± 0.2, and 5.3 ± 0.4 mM in h = 0.5, 0.6 and the initial h = 1 solutions, respectively, according to their absolute absorbance at 850 nm. Environmental Implications. For the first time, the identities and abundances of iron molecular clusters and particles formed early in ferric iron solution neutralization have been determined. The result provides new insights into the precipitation process and demonstrates the necessity of using in situ fast diffraction and spectroscopic measurements for analysis of rapid ferric iron oxyhydroxide particle formation processes. The concentrations of ferric iron and sulfate used in this study are comparable to those in AMD solutions, and the particle phases identified are commonly found in AMD-impacted environments. In addition, the various neutralization ratios are applicable for natural and engineered AMD neutralization processes. Therefore, the formation and phase transformation processes revealed in this study are likely relevant in a variety of AMD environments, though are most directly related to those with minimal ferrous iron and low concentration of heavy metal(loid)s and dissolved organic matter. Given that AMD can contain a wide variety of toxic elements that are attenuated by adsorption onto particle surfaces or through coprecipitation during neutralization, knowledge of early stage particle formation and transformation processes contribute to an improved understanding of attenuation mechanisms.

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ASSOCIATED CONTENT

S Supporting Information *

SXRD, the ferron assay procedures, all time-series QEXAFS spectra, QEXAFS shell-by-shell fitting results and comparisons of QEXAFS spectra with their linear combination fits. This material is available free of charge via the Internet at http:// pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Phone: (510) 643-9120; fax: (510) 486-5686; e-mail: mzhu@ lbl.gov. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the Director, Office of Science, Office of Basic Energy Sciences, Chemical Sciences, Geosciences, and Biosciences Division, of the U.S. Department of Energy under Contract No. DE-AC02-05CH11231. The authors are grateful to beamline scientists Dr. Syed Khalid at beamline X18B at NSLS, Brookhaven National Laboratory (BNL), Dr. Karena Chapman, Dr. Olaf Borkiewicz and Dr. Peter Chupas at beamline 11-ID-B at APS, Argonne National Laboratory (ANL), and Dr. John Bargar at beamline 4-1 at Stanford Synchrotron Radiation Lightsource (SSRL) for providing technical help on data collection. Use of the National Synchrotron Light Source, Brookhaven National Laboratory was supported by the U.S. DOE Office of Science, Office of Basic Energy Sciences, under Contract No. DE-AC0298CH10886. Use of APS was supported by the U.S. DOE Office of Science under Contract No. DE-AC02-06CH11357. Portions of this research were carried out at the Stanford Synchrotron Radiation Laboratory, a national user facility 8146

dx.doi.org/10.1021/es301268g | Environ. Sci. Technol. 2012, 46, 8140−8147

Environmental Science & Technology

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dx.doi.org/10.1021/es301268g | Environ. Sci. Technol. 2012, 46, 8140−8147