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Reductive Transformation of Birnessite by Aqueous Mn(II) Evert J. Elzinga* Department of Earth & Environmental Sciences, Rutgers University, Newark, New Jersey 07102, United States
bS Supporting Information ABSTRACT: Reaction of aqueous Mn(II) with hexagonal birnessite at pH 7.5 causes reductive transformation of birnessite into feitknechtite (β-MnIIIOOH) and manganite (γ-MnIIIOOH) through interfacial electron transfer from adsorbed Mn(II) to structural Mn(IV) atoms and arrangement of product Mn(III) into MnOOH, summarized by Mn(II) + Mn(IV)O2 + 2 H2O f 2 Mn(III)OOH + 2 H+. Feitknechtite is the initial transformation product, and subsequently converted into the more stable manganite polymorph during ongoing reaction with Mn(II). Feitknechtite production is observed at Mn(II) concentrations 2 orders of magnitude below thermodynamic thresholds, reflecting uncertainty in thermodynamic data of Mn-oxide minerals and/ or specific interactions between Mn(II) and birnessite surface sites facilitating electron exchange. Under oxic conditions, feitknechtite formation through surface-catalyzed oxidation of Mn(II) by O2 leads to additional Mn(II) removal from solution relative to anoxic systems. These results indicate that Mn(II) may be an important moderator of the reductive arm of Mn-oxide redox cycling, and suggest a controlling role of Mn(II) in regulating the solubility and speciation of phyllomanganate-reactive metal pollutants including Co, Ni, As, and Cr in geochemical environments.
’ INTRODUCTION Manganese oxides are among the most reactive mineral phases found in the environment and exert a strong control on the chemistry of soils and sediments through adsorption and redox reactions1. Hexagonal birnessites dominate Mn mineralogy in natural aquatic systems. These layered Mn(IV) oxides consist of stacked sheets of edge-sharing MnO6 octahedra and carry negative structural charge caused by Mn site vacancies and substitution of Mn(III) for Mn(IV), with charge neutralization provided by cations that coordinate along the basal planes.1,2 The interlayer region between sheets is accessible to water molecules and dissolved ions, allowing for interaction of aqueous species with the highly reactive vacancy sites, which are capable of multidentate inner-sphere adsorption of divalent metals such as Zn(II) and Pb(II) that coordinate above and below the sites.18 Complexation of divalent metal cations at birnessite edge sites, and structural incorporation into the octahedral phyllomanganate sheets have been observed as well.79 In addition to adsorption and incorporation reactions, Mn(III, IV)-oxides also engage in redox reactions. Birnessite oxidizes Se(IV) to Se(VI), Cr(III) to Cr(VI), and As(III) to As(V),1013 and thus strongly influences the solubility and speciation of these pollutants in geochemical environments. Abiotic formation of Mn(III, IV)-oxides through oxidation of Mn(II) is thermodynamically favorable but kinetically hindered by the high activation energy of the oxidation of Mn(II) by O2. Mineral surfaces catalyze the reaction,1416 but oxidation of Mn(II) mediated by bacteria and fungi is generally fast relative to abiotic Mn(II) oxidation reactions, and most naturally occurring environmental Mn oxides are believed to be either derived directly from biogenic Mn(II) oxidation processes or from r 2011 American Chemical Society
subsequent alteration of biogenic oxides.2 Recent studies have shown that bacterially and fungally mediated oxidation of Mn(II) results in the initial formation of hexagonal birnessite-like phyllomanganate minerals with predominantly tetravalent structural Mn, and substantial structural disorder due to site vacancies and turbostratic sheet stacking.1722 These poorly crystalline and highly reactive initial biooxides may readily undergo (abiotic) transformation into secondary Mn mineral products of higher crystallinity and stability,21,22 but the (bio)geochemical controls on these secondary transformation reactions remain poorly characterized. The current study focuses on the reaction of hexagonal birnessite with aqueous Mn(II), which has been shown to affect solid state Mn mineralogy and structure in previous studies. Work by Tu et al.23 and Mandernack et al.24 indicated that sorptive interactions of Mn(II) with hexagonal birnessites under oxic conditions resulted in the formation of various Mn(III)containing mineral phases, including feitknechtite and manganite at neutral pH and higher,23,24 and groutite, cryptomelane, nsutite, and ramsdellite at pH 6.0, 4.0, and 2.4, respectively.23 Bargar et al.21 showed that biogenic hexagonal birnessite reacts with aqueous Mn(II) to form triclinic birnessite (which contains substantial structural Mn(III)) at Mn(II) concentrations below 500 μM, and feitknechtite at higher Mn(II) concentrations in oxic solutions of pH 7.77.8. Similarly, Zhu et al.25 observed formation of triclinic birnessite during reaction of Mn(II) with Received: April 16, 2011 Accepted: June 15, 2011 Revised: June 12, 2011 Published: June 15, 2011 6366
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Environmental Science & Technology hexagonal birnessite. The mechanisms proposed to explain formation of Mn(III)-containing mineral phases in these systems have differed. Tu et al.23 and Mandernack et al.24 ascribed formation of Mn(III) phases to surface-catalyzed oxidation of Mn(II) by molecular O2, whereas more recent studies have proposed electron exchange between sorbed Mn(II) and structural Mn(IV) to explain formation of Mn(III).13,21,22,2527 Since oxidation of Mn(II) at phyllomanganate surfaces affects surface reactivity and possibly Mn-oxide bulk structure, a detailed understanding of the Mn(II)-birnessite surface interaction and resulting reaction products will lead to improved understanding of the environmental structure and reactivity of Mn(III,IV)-hydroxide minerals. The aims of the current study were to systematically characterize sorption products resulting from reaction of Mn(II) with hexagonal birnessite as a function of (1) Mn(II) concentration; (2) reaction time; and (3) the presence (or absence) of O2.
’ MATERIALS AND METHODS Birnessite Substrate. Synthesis and characteristics of the birnessite substrate used in this study are described in the Supporting Information (SI). Mn(II)-Birnessite Isotherm Experiments. The majority of Mn(II)-birnessite experiments were run in an anaerobic glovebox under strictly anoxic conditions to prevent Mn(II) oxidation by O2; the experimental procedures applied to exclude O2 are described in the SI. Anoxic aqueous suspensions of birnessite were prepared in 0.1 M NaCl and maintained at pH 7.0 or 7.5 using 20 mM HEPES buffer dissolved in the reaction electrolyte. The first set of experiments involved reaction of 174 mg L1 birnessite suspensions (500 mL volume, prepared in opaque polyethylene containers) with Mn(II) (as MnCl2) added at concentrations of 2 and 4 mM at pH 7.0. An identical suspension spiked with 4 mM MgCl2 instead of MnCl2 was prepared as control sample, and a blank sample consisting of a 4 mM MnCl2 solution prepared in 0.1 M NaCl and maintained at pH 7.0 using 20 mM HEPES was run in parallel to the sorption experiments. No solid formation was observed in the blank at any time during the experiment. The samples were placed on a rotary shaker inside the glovebox and equilibrated for 10 days. Following reaction, the suspensions were centrifuged, and the solids were washed five times with anoxic deionized (DI) water by centrifugation and decantation, and then freeze-dried. The dried solids were stored in Al-foil wrapped bottles inside the glovebox prior to further characterization. A second set of samples was prepared where birnessite suspensions (0.05 g L1 in 0.1 M NaCl maintained at pH 7.5 using 20 mM HEPES) were reacted for 9 days under anoxic conditions with Mn(II) added at (initial) concentrations in the range of 501800 μM. The volume of each sample was 30 mL and was contained in a 50 mL polyethylene centrifuge tube wrapped in Al foil. The samples were spiked with Mn(II), and placed on a rotary shaker inside the glovebox for 9 days. Following reaction, the samples were syringe-filtered through 0.22 μm nitrocellulose membranes, and the filtered solids (1.5 mg per sample) were syringe-washed with 50 mL of anoxic DI water pushed through the solids film remaining on the filter membrane; the washed solids were then dried inside the glovebox. Filtered reaction solutes were analyzed for dissolved Mn(II) using the formaldoxime method,28 and removal of Mn(II) from solution was calculated as the difference between the initial and final Mn(II) concentrations. A third set of samples was prepared identical to the second sample
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Figure 1. XRD spectrum of birnessite (a) compared to the spectra obtained following reaction of this material with Mg(II) (spectrum b), and Mn(II) (c, d) for 10 days at pH 7.0. The * symbols in spectra c and d indicate XRD peaks consistent with feitknechtite (β-MnOOH) and manganite (γ-MnOOH), respectively. The # symbol in spectrum a indicates an XRD contribution from the sample holder also seen in spectra b and d.
set, except that preparation was done outside the glovebox under atmospheric conditions to investigate effects of O2 on occurring reactions. Mn(II)-Birnessite Kinetic Experiment. A kinetic experiment was run to monitor Mn(II) removal from solution and concurrent changes in Mn oxide mineralogy during a 20 day time frame under anoxic conditions at pH 7.5, using procedures similar to the isotherm experiments described above. Experimental details are described in the SI. Mn(II)-Feitknechtite Adsorption Experiment. Feitknechtite was obtained from the Mn(II)-birnessite sorption experiment described above involving reaction of an anoxic 174 mg L1 birnessite suspension with 2 mM Mn(II) at pH 7.0 for 9 days. The washed and freeze-dried material (confirmed by XRD as feitknechtite; see below) was resuspended inside the glovebox in an anoxic 0.1 M NaCl solution buffered at pH 7.5 by 20 mM HEPES to yield a suspension density of 0.05 g L1. Three 30-mL suspension volumes were transferred to opaque 50 mL polyethylene centrifuge tubes. Two of the suspensions were spiked with Mn(II) (100 μM and 250 μM initial concentrations), whereas the third sample was run as a control and received no Mn(II). The tubes were equilibrated on a rotary shaker inside the glovebox for 10 days. Following reaction, the samples were filtered, and the solids were washed with anoxic DI water, and dried inside the glovebox. Substrate Analyses. The reacted birnessite solids from the experiments described above were characterized using X-ray diffraction (XRD) and attenuated total reflectance Fourier transform infrared (ATR-FTIR) spectroscopic analyses. Details of these analyses are provided in the SI.
’ RESULTS AND DISCUSSION XRD results. Figure 1 presents the XRD spectra of the original nonreacted birnessite substrate along with the spectra of the Mn(II)- and Mg(II)-reacted birnessites reacted under anoxic conditions at a birnessite particle loading of 174 mg L1 with 2 or 4 mM Mn(II) or Mg(II). The XRD spectrum of nonreacted birnessite is consistent with previous studies,19,29 showing broad 6367
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Figure 2. Mn(II) sorption in 0.05 g L1 birnessite suspensions at pH 7.5 under oxic and anoxic conditions following 9 days of reaction plotted as a function of the remaining Mn(II) solution concentration.
hk bands corresponding to 2.4 and 1.4 Å d-spacings, which confirm the hexagonal symmetry of the sample. The two peaks at 12.3 2θ (7.2 Å) and 24.8 2θ (3.6 Å) arise from 001 and 002 reflections, respectively, and indicate an interlayer separation of 7.18 Å.29 The spectrum of the Mg(II)-reacted birnessite is very similar to that of nonreacted birnessite (Figure 1), indicating that the birnessite structure did not notably change during reaction with Mg(II). In contrast, the XRD spectra of the Mn(II)-reacted birnessites are radically different from the spectrum of the starting material, and contain XRD peaks indicating the presence of feitknechtite (β-Mn(III)OOH; sample reacted at [Mn(II)] = 2 mM) and manganite (γ-Mn(III)OOH; sample with [Mn(II)] = 4 mM) whereas reflections associated with birnessite are no longer seen. The possibility of Mn(III)OOH forming in these samples as a result of surface-catalyzed oxidation of Mn(II) by molecular O2 can be excluded as the experiments were performed under anaerobic conditions. Instead, appearance of manganite and feitknechtite concurrent with disappearance of birnessite indicates structural transformation of birnessite during reaction with Mn(II)(aq). Two major reactions may account for the observed transformation of birnessite into β-MnOOH and γ-MnOOH in the Mn(II)-birnessite samples. The first involves reductive transformation of birnessite according to 2MnðIVÞO2 + H2 O f 2MnðIIIÞOOH + 0:5O2
ð1Þ
In the current experiments, this reaction would be driven to the right as a result of the anoxic experimental conditions employed. However, operation of this reaction as a major pathway of birnessite transformation can be excluded based on the absence of feitknechtite or manganite product in the blank sample involving reaction of birnessite with Mg(II) under identical conditions as the Mn(II)-birnessite samples; in addition, as will be shown below, the same transformations occur under oxic conditions. Transformation of birnessite into Mn(III)OOH appears instead to be driven by reaction with Mn(II), involving interfacial electron transfer from adsorbed Mn(II) to structural Mn(IV) atoms and arrangement of product Mn(III) into Mn(III)OOH, a reductive transformation process that can be summarized by the following comproportionation reaction: MnðIIÞ + MnðIVÞO2 + 2H2 O f 2MnðIIIÞOOH + 2H+ ð2Þ
Figure 3. Normalized ATR-FTIR spectra of Mn-oxide sample solids from anoxic (spectra dk) and oxic (spectra nq) isotherm experiments at pH 7.5, run at a birnessite suspension density of 0.05 g L1. Spectrum a is the spectrum of nonreacted birnessite. Spectra b and c were obtained for the Mn(II)- birnessite samples from Figure 1 (prepared anoxically at a birnessite suspension density of 0.174 g L1), and spectra l and m for anoxic samples reacted at a low birnessite suspension density of 0.0025 g L1 (see text). Initial and final (9 day reaction) solution Mn(II) concentrations are indicated along each spectrum; the final Mn(II) solution concentration of the sample represented by spectrum m was below the detection limit of the formaldoxime method.
Secondary observations consistent with this transformation pathway were notable increases in Mn-oxide solids mass and substantial production of acidity during the course of reaction. The operation of reaction 2 has been proposed under oxic conditions as well.21 The difference in Mn(III)OOH polymorphs seen between the samples suggests that the concentration of Mn(II) controls the type of transformation product formed. Mn(II)-Birnessite Isotherms. Systematic characterization of the impact of Mn(II) concentration on occurring reactions is provided by the isotherm experiments. The macroscopic results of these experiments are presented in Figure 2, where the concentrations of Mn(II) removed from solution are plotted as a function of the remaining Mn(II) solution concentrations for both the anoxic and oxic experiments. Removal of aqueous Mn(II) in these systems takes on the form of a high affinity isotherm, with strong preference for solid phase partitioning at low Mn(II) concentrations. Manganese(II) removal plateaus at higher concentrations in the anoxic experiments, whereas no sorption maximum is reached under oxic conditions, resulting in notably higher Mn(II) removal rates in the oxic than in the anoxic experiments at Mn(II) concentrations >200 μM (Figure 2). The results of ATR-FTIR analysis of the reacted solids of the isotherm samples are presented in Figure 3, which compares the 1250850 cm1 spectral regions of the sample spectra to those of nonreacted birnessite, and the Mn(II)-reacted birnessite 6368
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Environmental Science & Technology samples from Figure 1. Nonreacted birnessite (spectrum a) lacks structural IR bands in the spectral region shown in Figure 3, consistent with previous studies.30 The IR bands seen in the birnessite sample reacted with Mn(II) at 4 mM from Figure 1 (spectrum b) shows band at 1150, 1116, and 1087 cm1, which represent the in-plane and out-of-plane bending modes of structural OH in manganite.31 The birnessite sample from Figure 1 reacted with Mn(II) at a concentration of 2 mM (spectrum c in Figure 3) contains the OH bending modes of feitknechtite (at 1067 cm1 and 946 cm1) consistent with the XRD results of this sample (Figures 1 and 3). Inspection of the spectra obtained for the anoxic and oxic isotherm samples (spectra dk and nq, respectively) indicates that in both data sets, feitknechtite is the dominant Mn(III) phase in the samples reacted at low Mn(II) concentrations, whereas manganite grows in at higher concentrations; samples with intermediate Mn(II) concentrations contain a mixture of manganite and feitknechtite. The predominance of feitknechtite at relatively low Mn(II) concentrations and manganite at high concentrations is consistent with the results presented in Figure 1. Comparison of the IR spectra of oxic and anoxic samples spiked with the same initial Mn(II) concentrations (spectra nq versus dk in Figure 3) indicates a higher proportion of feitknechtite relative to manganite in the oxic samples. This suggests that the additional Mn(II) removal seen in the oxic versus anoxic experiments at Mn(II) concentrations >200 μM (Figure 2) is due to formation of feitknechtite produced by surface-catalyzed oxidation of Mn(II) by O2. Surface-catalyzed oxidation of Mn(II) by O2 at the surfaces of hematite and goethite has been shown to result in formation of feitknechtite as well,14 and MnOOH formation was noted during autocatalytic oxidation of Mn(II) in oxic solutions.15 The initial and final aqueous Mn(II) concentrations of the solutes reacted with birnessite in the isotherm experiments are indicated alongside the ATR-FTIR spectra in Figure 3. A potential concern is that aqueous Mn(II) concentrations in the isotherm samples dropped sharply during the course of the experiment, in particular (in relative terms) for the low [Mn(II)] samples, allowing the possibility that relatively high initial [Mn(II)] drives formation of initial products that may still be present but no longer thermodynamically stable after 9 days. To address this concern, two anoxic samples were prepared having a low birnessite suspension density of 0.0025 g L1, and reacted at (low) initial Mn(II) concentrations of 10 μM and 50 μM for 9 days. The IR spectra of these samples (spectra l and m in Figure 3) indicate formation of feitknechtite, consistent with the results of the isotherm samples. It is instructive to compare the Mn(II) concentrations indicated in Figure 3 with the Mn(II) activities predicted from comproportionation reaction 2 assuming reversible thermodynamic equilibrium, as outlined in ref 24. The equilibrium constant of reaction 2 can be calculated from the standard Gibbs free energy change of reaction, ΔGR0, which is calculated from the ΔGf0 values of the reactants and products partaking in reaction (ΔGR0 = ∑(ΔGf0)products ∑(ΔGf0)reactants). Values of ΔGf0 for Mn2+(aq) and H2O(l) are reported in ref 32, those of manganite and birnessite in ref 33, and that of feitknechtite in ref 34. These thermodynamic data can be used to calculate two values of ΔGR0 for the comproportionation reaction: one assuming formation of feitknechtite, the other assuming manganite. From these ΔGR0 values, the equilibrium constant Keq of reaction 2 can be calculated for systems with manganite or feitknechtite (ln Keq = ΔGR0 /RT), yielding Keq = 1012.14
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Figure 4. ATR-FTIR data of the Mn-oxide solids obtained from the kinetic Mn(II)-birnessite sorption experiment (anoxic conditions; pH 7.5; initial aqueous Mn(II) concentration of 800 μM; birnessite suspension density of 0.05 g L1). Reaction times are indicated along the spectra.
assuming formation of feitknechtite, and Keq = 107.01 assuming manganite. These values allow calculation of the equilibrium aqueous Mn(II) activities at given pH for systems with feitknechtite and manganite, according to a(Mn(II)) = a(H+)2/Keq (assuming pure solids with activity = 1). At pH 7.5, these calculations predict an aqueous Mn(II) equilibrium activity of 1.02 108 M for systems containing manganite, and a Mn(II) activity of 1.38 103 M for systems with feitknechtite; these activities translate into Mn(II) solution concentrations of 3 108 M and 4 mM, respectively, under the conditions of the current experiments after accounting for effects of ionic strength and MnCl+(aq) complex formation. The lower Mn(II) equilibrium concentrations calculated for systems containing manganite relative to those with feitknechtite reflect the higher stability of manganite compared to feitknechtite.34 Comparison of the thermodynamic calculations to the experimental results reveals two notable discrepancies between the theoretical and experimental data. First, the thermodynamic calculations suggest a Mn(II) solution concentration threshold for feitknechtite formation that is slightly above the maximum (initial) Mn(II) concentration used in the sorption experiments. Therefore, it would be reasonable to assume that feitknechtite, if it formed, would occur only in the high Mn(II) samples with solution concentrations near the limiting Mn(II) solution concentration, leaving manganite to dominate in the low Mn(II) samples where formation of feitknechtite is not thermodynamically favorable. The experimental results (Figure 3) exhibit in fact the exact opposite trend, with feitknechtite dominating in the low Mn(II) samples, and manganite at high [Mn(II)]. A second discrepancy lies in the actual Mn(II) solution concentrations where feitknechtite formation is expected to occur. Comparison of the thermodynamic Mn(II) equilibrium concentrations to the actual Mn(II) concentrations measured in the reaction solutes (Figure 3) shows that, while experimental Mn(II) concentrations are above the manganite equilibrium values (indicating a thermodynamic driving force for manganite formation in these samples), formation of feitknechtite occurs at Mn(II) solution concentrations that are well over 2 orders of magnitude below the 6369
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Figure 5. IR spectra of sample solids obtained following reaction of feitknechtite (0.05 g L1) with 0, 100, and 250 μM Mn(II) at pH 7.5 under anoxic conditions. Final dissolved Mn(II) concentrations (10 day reaction) are indicated in brackets; solution Mn(II) in the blank sample was below detection.
equilibrium Mn(II) concentration calculated from thermodynamics. Clearly, thermodynamic equilibrium arguments have limited merit in explaining the current results. Data from the kinetic Mn(II)-birnessite and the Mn(II)-feitknechtite sorption experiments address this issue further. Mn(II)-Birnessite Kinetic and Mn(II)-Feitknechtite Sorption Experiments. The kinetic Mn(II)-birnessite adsorption experiment monitored Mn(II) removal and concurrent mineralogical transformations over a 20 day reaction period. The results are shown in Figure 4 presenting the IR spectra of the reacted sample solids obtained during the course of reaction; time-dependent removal of Mn(II) from solution is shown in SI Figure S1. The data in Figure 4 indicate that the initial mineral transformation product formed from comproportionation of structural Mn(IV) and adsorbed Mn(II) is feitknechtite, with first appearance after 1 day of reaction, whereas manganite grows in later, with the first clear indication of its formation seen after 6 days (Figure 4). Of note is the ultimate disappearance of feitknechtite while manganite bands become increasingly defined as time progresses. This indicates that the initial feitknechtite phase formed during the early stages of reaction transforms into manganite over time, and implies that manganite formation does not proceed directly in these systems but requires formation of feitknechtite as a precursor phase. This finding is consistent with previous studies which have shown that feitknechtite is metastable and converts into manganite during aging in aqueous solution.33,34 Figure 4 shows ongoing transformation beyond the 9 day reaction time allowed for the isotherm samples, indicating that the results shown in Figure 3 do not represent the final state of the Mn(II)-birnessite reaction, which is expected to ultimately produce manganite as the dominant Mn(III) mineral transformation product. The results of the Mn(II)-feitknechtite adsorption experiments are presented in Figure 5, and show that the feitknechtite samples reacted with Mn(II) for 10 day partially converted into manganite, whereas the blank sample with no added Mn(II) remained unchanged. These results indicate that reaction with Mn(II) promotes the conversion of relatively disordered feitknechtite into crystalline manganite. Combined with the
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information from the kinetic Mn(II)-birnessite experiment, the dependence of Mn(III) mineralogy on Mn(II) concentration seen in the isotherm experiments (Figure 3) can now be explained. The initial transformation product in the Mn(II)birnessite samples is feitknechtite, which is subsequently converted into manganite. The predominance of manganite in isotherm samples spiked with higher Mn(II) concentrations is due to relatively high Mn(II) solution concentrations in these samples that promote transformation of the initial feitknechtite phase into manganite. While the current observations do not resolve the mechanistic details of the role of Mn(II) in accelerating feitknechtite transformation, they do resemble results of studies on Fe(II) reactivity with ferrihydrite (amorphous Fe(OH)3) showing that Fe(II) catalyzes conversion of ferrihydrite into crystalline Fe(III)-(hydr)oxides as a result of interfacial electron transfer reactions between (soluble) Fe(II) and (insoluble) Fe(III) accelerating Fe dissolution and reprecipitation (e.g., refs 35 and 36). The results presented here are consistent with similar electron exchange reactions between (insoluble) Mn(III) in feitknechtite and (soluble) Mn(II) catalyzing transformation of feitknechtite into manganite. Formation of feitknechtite is observed at Mn(II) solution concentrations well over 2 orders of magnitude below thermodynamic thresholds calculated assuming reversible comproportionation between Mn(II) and Mn(IV) at the birnessite surface according to reaction 2. Several factors may be in play. First, uncertainties in the thermodynamic data of the Mn-oxides involved may contribute. The difference in the ΔGf0 values of feitknechtite (ΔGf0 = 543.1 kJ mol1) and manganite (ΔGf0 = 557.7 kJ mol1) is relatively small (14.6 kJ mole1, or
