Cation Effects on the Layer Structure of Biogenic Mn-Oxides

May 14, 2010 - This study investigates the impacts of cations (H+, Ni(II), Na+, and Ca2+) during biotic Mn(II) oxidation on the structure of Mn octahe...
0 downloads 11 Views 2MB Size
Download PDF
Environ. Sci. Technol. 2010, 44, 4465–4471

Cation Effects on the Layer Structure of Biogenic Mn-Oxides MENGQIANG ZHU,* MATTHEW GINDER-VOGEL,† SANJAI J. PARIKH,‡ XIONG-HAN FENG,§ AND DONALD L. SPARKS Department of Plant and Soil Sciences, Delaware Environmental Institute, University of Delaware, 152 Townsend Hall, Newark, Delaware 19716

Received October 30, 2009. Accepted April 26, 2010.

Biologically catalyzed Mn(II) oxidation produces biogenic Mnoxides (BioMnOx) and may serve as one of the major formation pathways for layered Mn-oxides in soils and sediments. The structure of Mn octahedral layers in layered Mn-oxides controls its metal sequestration properties, photochemistry, oxidizing ability, and topotactic transformation to tunneled structures. This study investigates the impacts of cations (H+, Ni(II), Na+, and Ca2+) during biotic Mn(II) oxidation on the structure of Mn octahedral layers of BioMnOx using solution chemistry and synchrotron X-ray techniques. Results demonstrate that Mn octahedral layer symmetry and composition are sensitive to previous cations during BioMnOx formation. Specifically, H+ and Ni(II) enhance vacant site formation, whereas Na+ and Ca2+ favor formation of Mn(III) and its ordered distribution in Mn octahedral layers. This study emphasizes the importance of the abiotic reaction between Mn(II) and BioMnOx and dependence of the crystal structure of BioMnOx on solution chemistry.

Introduction Mn-oxidizing and -reducing microorganisms are crucial drivers for biogeochemical cycling of Mn and accordingly those elements associated with Mn-oxides (1). A diverse group of microorganisms, including bacteria and fungi, in various natural settings, participate in Mn(II) oxidation to Mn(III, IV)-oxides and the reverse process (1). The formation of most naturally occurring Mn-oxides is likely, at least initially, microbially mediated because of rapid biological Mn(II) oxidation (2). Mn-oxides found in soils and marine nodules, hydrothermal deposits, desert vanishes, and sediments of fresh lakes and streams commonly exhibit association with Mn-oxidizing microorganisms (1). Biological Mn(II) oxidation proceeds in two steps: Mn(II) is first oxidized to Mn(III) and then to Mn(IV) by O2, and it is catalyzed by multicopper oxidase enzymes (2, 3). Mn(III) cations are believed to be attached on cell surfaces since intermediate Mn(III)-oxides are not detected (2, 4, 5). Synchrotron X-ray diffraction (SR-XRD) and extended X-ray absorption fine structure (EXAFS) spectroscopy reveal that * Corresponding author phone: (302) 831-1230; fax: (302) 8310605; e-mail: [email protected]. † Current address: Calera Corporation, 14600 Winchester Blvd., Los Gatos, CA 95030. ‡ Current address: Department of Land, Air and Water Resources, University of California, One Shields Ave, Davis, CA 9561. § Current address: College of Resources and Environment, Huazhong Agricultural University, Wuhan 430070, P. R. China. 10.1021/es1009955

 2010 American Chemical Society

Published on Web 05/14/2010

neo-formed BioMnOx are disordered nanoparticulate MnO2 (6). These nanoparticles abiotically transform to hexagonal and pseudo-orthogonal layered Mn-oxides, with structure similar to hexagonal and triclinic birnessite, or form Mn(II,III)-oxides at higher concentrations of Mn(II) (6, 7). This abiotic process occurs concurrently with the biotic process (8). Identities of thermodynamically stable Mn-oxide phases are independent of microorganism type and largely determined by solution chemistry instead, such as dissolved Mn(II) concentration, pH, and coexisting cations (6, 8-10). For instance, β-MnOOH is the thermodynamically favorable phase formed in a constant 1 mM Mn(II) solution at pH 7.8; in contrast, birnessite is the stable phase at lower concentrations of Mn(II) solution or at lower pH (5, 6, 8, 11-14). Coexisting cations also impact layer structure of BioMnOx. Na+ tends to cause layered Mn-oxides with hexagonal layer symmetry, while Ca2+ favors one with orthogonal layer symmetry (9); A todorokite-like, tunneled structure of Mnoxide is formed with [UO2]2+ as a coexisting cation (10). Mn-oxidizing microbes exist in a wide range of geochemical conditions, such as high temperature, high salinity (i.e., seawater), acidic or alkaline conditions, and in the presence of heavy metals (1). It is necessary to investigate how these geochemical conditions affect the crystal structure of BioMnOx. In the current study, the effects of H+ (pH), alkaline and alkaline earth metals (Na+ and Ca2+), heavy metals (Ni(II)), and combinations of these cations were systematically investigated with XAFS and SR-XRD. The goal of this study is to decipher how these common naturally occurring cations affect Mn-layer structure of BioMnOx.

Material and Methods Preparation of Mn-Oxidizing Biomaterials. Pseudomonas putida strain GB-1 was grown in 500 mL of Leptothrix growth media in 1800-mL Erlenmeyer flasks at 30 °C and orbital mixing at 200 rpm (11). High-speed mixing provides sufficient oxygen supply during the growth and is required for the cells to have Mn-oxidizing ability (15). Biomaterials, including cells and extracellular polymeric substances (EPS), were harvested from cultures after 19 h of growth, when cells possessed maximum oxidizing activity. The harvested biomaterials were washed with a pH 7, 10 mM HEPES solution containing 50 mM NaCl before use. Detailed bacterial preparation procedures are given in the Supporting Information (S1). Since the Mn-oxidizing activity of cells differed slightly from batch to batch, effects of variation in oxidizing activity were eliminated by using the same cell suspension to investigate a single factor. For example, investigation of pH effects was performed with the same biomaterials for pH 6, 7, and 8 in NaCl solutions, which is different from those used in CaCl2 solutions. The biomaterials used to study Ni(II) effects at pH 7 and 8 were from another culture. Biological Mn(II) Oxidation. The washed biomaterials collected from three 500-mL cultures were mixed then resuspended in three, 1000-mL solutions of pH 6, 7, and 8. The solutions contained 100 µM MnSO4 and 50 mM NaCl or 50 µM MnSO4 and 16.67 mM CaCl2 (ensuring the same ionic strength). The lower MnSO4 concentration was used for CaCl2 experiments to decrease the possibility of forming Mn(III)oxides, based on preliminary experiments showing Mn(II) oxidation was much slower in a CaCl2 solution. All solutions were buffered with 20 mM MES at pH 6 and 20 mM HEPES at pH 7 and 8, and the pH was adjusted by addition of 6 M NaOH. The solutions were autoclaved prior to addition of a filter-sterilized MnSO4 solution. The suspensions were placed on an orbital shaker at 200 rpm and at 30 °C. The Mn(II) VOL. 44, NO. 12, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

4465

FIGURE 1. Synchrotron X-ray diffraction patterns of (a) BioMnOx formed at varying pH and the cells only. From the top to the bottom: pH8Ca, pH7Ca, pH6Ca, pH8Na, pH7Na, pH6Na, and cells. (b) Acidified BioMnOx preformed at pH 8 for one week and acidification pH 7, 6, and 5. From the top to the bottom: pH81weekCa, pH8to7Ca, pH8to6Ca, pH8to5Ca, pH81weekNa, pH8to7Na, pH8to6Na, and pH8to5Na. (c) BioMnOx formed at varying Ni(II) to Mn(II) ratio at pH7 and 8. From the top to the bottom: pH8Ni02, pH8Ni005, pH8Ni0, pH7Ni1, pH7Ni02, pH7Ni005, and pH7Ni0. Each figure also contains its magnified part from 38 to 42° for close inspection of the characteristics indicating Mn-layer symmetry. concentration was measured by the formaldoxime colorimetric method (16) with a UV-visible spectrometer (Bausch & Lomb, Spectronic 21) by filtering 3-mL aliquots of suspension with 0.22 µm cellulose syringe filters. After 48 h, solids were collected after centrifuging about 10 times at 3,000 × g. At this centrifugation speed, a fraction of the EPS remained in the supernatant, and it was decanted until the final supernatant was visibly clear. Weak-Acid Treatment of BioMnOx. BioMnOx, formed in solutions containing 50 mM NaCl or 16.67 mM CaCl2 at pH 8 for 1 week, were prepared using the same procedure as described above. For acidification, these two preformed BioMnOx were then reacted with solutions at pH 5, 6, and 7 for 36 h at 25 °C. The solutions contained either 50 mM NaCl or 16.67 mM CaCl2, the same background electrolyte as used for the preformed BioMnOx. Sodium azide was added to prevent the cells from reoxidizing released Mn(II) from this weak-acid treatment. Biological Ni(II) and Mn(II) Coprecipitation. Another batch of harvested biomaterials was transferred to 500-mL polycarbonate bottles containing 300 mL of Ni(NO3)2 and MnSO4 solutions (50 µM) at Ni(II) to Mn(II) molar ratios of 0, 0.05, 0.20, and 1.0 at both pH 7 and 8 (without a Ni/Mn ratio of 1.0 at pH 8, where Ni(II) can precipitate in the bulk solution). The cell densities were approximately the same as those used to study the effect of pH. Under these conditions, neither Mn(II) nor Ni(II) precipitate in the bulk solutions, abiotically. The background electrolyte was 50 mM NaCl and the pH was buffered at pH 7 or 8 with 20 mM HEPES. The bottles were mixed on an orbital shaker at 150 rpm and 24 °C. Three-mL aliquots were taken and filtered at each time interval to monitor dissolved Mn(II) concentrations measured by ICP-MS (Agilent Technologies 7500) since the above colorimetric method for Mn(II) measurement can be affected by Ni(II) precipitation. X-ray Characterization. X-ray Diffraction. The solids prepared from the above procedures were stored at -20 °C prior to analyses. SR-XRD patterns were collected at a X-ray 4466

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 44, NO. 12, 2010

wavelength of 0.97584 Å with an image plate in transmission mode on beamline 11-3 at the Stanford Synchrotron Radiation Lightsource (SSRL). Wet sample slurries were placed in an aluminum sample cell with Lexan windows (polycarbonate). LaB6 was used to calibrate parameters for data processing with Fit2D (17). The background contributions from the Lexan windows and water in the sample were removed by subtraction using XRD-bs (Samuel Webb, SSRL). Two Mn-oxide references, triclinic birnessite (TcBir) and δ-MnO2 were synthesized using standard methods (S2). EXAFS Spectroscopy. Mn and Ni K-edge EXAFS spectra were collected from the BioMnOx samples in fluorescence mode using a Lytle detector and from reference Mn-oxides in transmission mode on beamline X11A equipped with a Si (111) crystal monochromator at the National Synchrotron Light Source (NSLS). The beam was detuned by 30% to suppress higher order harmonic X-rays. An EXAFS spectrum of Mn metal foil (E0 ) 6539.0 eV) was collected concurrently for energy calibration. The lack of self-absorption in the fluorescence spectra was confirmed by comparing XANES transmission and fluorescence spectra. SIXPack was used to perform EXAFS data reductions and analyses (18). Shell-byshell fitting of EXAFS spectra and determinations of Mn average oxidation states (AOS) and Mn(II) contents in the solids are described in detail in the Supporting Information.

Results X-ray Diffraction. The (31/02) and (20/11) reflections in XRD patterns are indicative of layered Mn-oxides (Figure 1) (6, 9, 11). Absent or weak (001) and (002) reflections suggest BioMnOx lacking 3-D periodic structure (19). Splitting of the (31/02) reflection can be used to differentiate orthogonal layer symmetry from hexagonal symmetry containing a single (31/02) reflection (6, 20). BioMnOx formed in the CaCl2 solution at pH 8 for 48 h (pH8Ca) has an orthogonal layer structure, while the rest of BioMnOx in Figure 1a exhibits hexagonal symmetry. A close inspection of (31/02) reflections

FIGURE 2. Fractions of Mn(II) determined from XANES linear combination fittings (a). Fractions of layer Mn(III) (b) and vacant sites (c) determined from EXAFS linear combination fittings, and RSF peak intensity ratio of edge-sharing Mn shell to the first O shell (d). The fraction values in LCF were normalized to the component sum. shows peak broadening for both the NaCl and CaCl2 series as formation pH increases from 6 to 8 (Figure 1a). OrthogonalstructuresareformedinbothCaCl2 (pH81weekCa) and NaCl (pH81weekNa) solutions (Figure 1b) for one-week reaction time. Upon acidification at pH 7, 6, and 5, one of the (31/02) splitting peaks of pH81weekNa gradually diminishes in amplitude (Figure 1b), suggesting a transformation from orthogonal to hexagonal symmetry and consistent with acidified transformation of synthetic triclinic birnessite (21-23). The two splitting peaks of pH81weekCa for orthogonal symmetry do not change much until solution pH is lowered to 5 (Figure 1b), suggesting Ca2+ retards the transformation. An orthogonal structure occurs at pH 8 in the absence of Ni(II) (pH8Ni0) based on splitting (31/02) reflections (Figure 1c). The presence of Ni(II) simply results in hexagonal structures (pH8Ni02). Additionally, (31/02) reflection peaks sharpen as the Ni(II) to Mn(II) ratio increases at pH 7 (Figure 1c). XANES. XANES line shapes and white line positions of BioMnOx samples vary under different formation conditions (Figure S3A, Supporting Information). XANES linear combination fittings (LCF) were best when the fitting references MnSO4(aq), δ-MnO2 and sample pH8Ca (not having Mn(II), see below) were included to estimate Mn(II) proportions (fMn(II)) (Figure S3B) (8). Replacement of pH8Ca with TcBir gave slightly worse fitting accuracies, but almost the same results. fMn(II) decreases as formation and acidification pH and Ni/Mn ratio increases except for pH8Na and pH6Na which have similar amounts of Mn(II) (Figure 2a), consistent with a XANES visual inspection (Figure S3). EXAFS Spectroscopy. δ-MnO2 has sharp peaks at ∼8.10 and 9.25 Å-1, while in TcBir, in addition to peak broadenings, the peak at ∼8.10 Å-1 splits and the peak at ∼9.25 Å-1 shifts to the left (Figure 3). The Mn layers in δ-MnO2 are hexagonal, having varying amounts of vacant sites but no Mn(III). The Mn layers of ideal TcBir are orthogonal and vacant site free,

containing orderly distributed Mn(III) with a molar fraction of 0.33 (23). The longer Mn(III)-O and Mn(III)-Mn bond distances contribute to the peak splitting and the left shift (24, 25). The possible presence of vacant sites in our TcBir sample, as shown by Ni(II) sorption in the companion paper (26), may cause the peak splitting to not be as pronounced as in previous studies (24, 25). Major variations in BioMnOx EXAFS spectra lie in the region of 7.5-9.5 Å-1 as formation conditions change (Figure 3). As the formation pH increases, the 8.10 and 9.25 Å-1 peaks undergo splitting or broadening and shift to lower k as well (Figure 3a and b). The region is more similar to δ-MnO2 than TcBir for samples formed at pH 6 and 7. The single 8.10 Å-1 peaks in pH8Na and pH81weekNa may indicate hexagonal symmetry, but remarkable broadening and slight shift to lower k of the 8.10 and 9.25 Å-1 peaks suggest similarities to TcBir. The EXAFS spectra of pH8Ca and pH81weekCa closely resemble TcBir, indicating orthogonal layer symmetry. Upon acidification, 8.10 and 9.25 Å-1 peaks of the pH81weekNa sample gradually sharpen and 9.25 Å-1 peaks slightly shift to the right (Figure 3a), corresponding to a transformation toward hexagonal layer symmetry (21-23). The splitting features in pH81weekCa remain until acidification pH decreases to pH 5 when a structure similar to a hexagonal birnessite is formed (Figure 3b) (21). The EXAFS spectrum of pH8Ni0 shows some features like TcBir. When Ni(II) is present, the 8.10 and 9.25 Å-1 peaks substantially sharpen and slightly shift to the right for both pH 7 and 8, suggesting changes toward hexagonal layer symmetry (Figure 3c). It is noteworthy that the EXAFS spectrum of the acidified form (pH8to5Na(Ca)) at pH 5 is very similar to the BioMnOx formed at pH 6 (pH6Na(Ca)). The 7.5-9.5 Å-1 region is sensitive to the amount (fMnIII), spatial distribution of Mn(III) in Mn-layers (24, 25) and probably the quantity of vacant sites (fvac), as described in the comparison between δ-MnO2 and TcBir. The observed changes in the 7.5-9.5 Å-1 region indicate varying fMnIII and VOL. 44, NO. 12, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

4467

FIGURE 3. EXAFS spectra of (a) BioMnOx formed in NaCl solutions at different pH and the acid-treated forms of BioMnOx preformed at pH 8 for one week, (b), BioMnOx formed in CaCl2 solutions and their acid-treated forms, and (c) BioMnOx formed in Ni(II) solutions with Ni to Mn ratio of 0 and 0.2 at pH 8 and 0 and 1 at pH 7. fvac, which are also underlying reasons for their different layer symmetry. Mn(III) cations are probably ordered instead of randomly distributed in Mn layers because a random Mn(III) distribution shifts the 8.10 and 9.25 Å-1 peaks to 7.95 and 9.05 Å-1, respectively (25). To quantitatively compare fMnIII and fvac in the BioMnOx samples, the EXAFS LCF used follows the method of Marcus et al. (24). LCF were best with two end members, the samples pH8Ca and pH8to5Ca (S4). The LCF of the TcBir and δ-MnO2 were not as good, probably because of the TcBir sample containing some vacant sites and δ-MnO2 not containing corner-sharing Mn, whereas it is prevalent in the samples (shown later). Assuming pH8Ca and pH8to5Ca have the same composition as triclinic birnessite (fMnIII ) 0.33 and fvac ) 0) and hexagonal birnessite (fMnIII ) 0.111 and fvac ) 0.167) (22), fMnIII and fvac in samples can be estimated. The obtained fMnIII and fvac demonstrate opposite trends, generally increasing and decreasing, respectively, as forma4468

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 44, NO. 12, 2010

tion and acidification pH increases and Ni/Mn ratio decreases (Figure 2b and c). Radial Structure Function (RSF) Analysis. In our acidified series, an increase in peak intensity ratio of the edge-sharing Mn shell to the first O shell (Mn/O ratio) (Figure 2d) and an increase in 5.2 Å (R + ∆R) peak intensities are observed as acidification pH decreases (Figure 3). The 5.2 Å (R + ∆R) peaks are derived mainly from Mn-Mn-Mn multiplescattering (MS), and their intensities are seriously attenuated by an ordered Mn(III) distribution (6, 19) and are also subject to impacts of vacant sites (6). Acidification results in more vacant sites and less layer Mn(III) (21-23). Therefore, the increasing Mn/O ratio and the MS peak intensities in acidified series are corresponding to an increase of fvac and a decrease of fMnIII. Thus, the finding that the decreases of Mn/O ratio and the MS peak intensities as formation pH increases and Ni/Mn ratio decreases may indicate presence of more layer

Mn(III) and less vacant sites. The changes in the Mn/O ratio as solution chemistry are basically consistent with fvac and fMnIII determined by EXAFS LCF (Figure 2c and d). The atomic shells at 3 Å (R + ∆R) correspond to cornersharing Mn(II,III) adsorbed on vacant sites in the interlayer regions (Figure 3) (27). The interlayer Mn contents decrease with increasing formation and acidification pH and Ni/Mn ratio based on the peak intensities (Figure 3). This conforms to the coordination numbers (CN) determined by shell-byshell fitting (S6). Mn(II,III) species bound to bacterial materials or in interlayers of BioMnOx (nonlayer Mn) can reduce RSF peak intensities (27). Intensities of the first O peaks and edgesharing Mn peaks are lower in samples containing larger amounts of nonlayer Mn species, such as pH6Na and pH6Ca (Figure 2a and Figure 3); however, the MS peak intensities are larger in these samples. Therefore, nonlayer Mn is not the reason for the MS peak intensity changes. Intercalation of Ca2+, Na+ and Ni(II) in interlayer regions are unlikely to affect Mn EXAFS spectra (19, 27).

Discussion Mn-Layer Compositions. Although XRD does not directly measure Mn oxidation state, the orthogonal layer symmetry suggests the presence of large amounts of ordered distributed Mn(III) and low vacant site contents, while hexagonal layer symmetry suggests less layer Mn(III) and more vacant sites (6, 9, 21, 22). Intercalation of cations in the interlayers may affect Mn layers, but intercalation of Ca2+, Na+, and Ni(II) has not been reported to change layer symmetry from hexagonal to orthogonal. Generally, the EXAFS analyses agree with the XRD results. Varying layer symmetry observed from XRD and EXAFS analyses suggests more Mn(III) and less vacant site contents as BioMnOx formation pH increases and Ni/Mn ratio decreases. This is further validated by the similarities in EXAFS spectra between the samples pH6Na(Ca) and pH8to5Na(Ca), indicating that BioMnOx formed at pH 8 have higher fMnIII and lower fvac than those formed at pH 6. Moreover, the change in fvac for the CaCl2 series is directly proven in a Ni(II) sorption study described in the companion paper (26). Neither chemical titration (11) nor Mn LII,III edge spectroscopy (28) can be used to quantify fMnIII because of the three mixed Mn valences in our samples, as well as interference from bacterial materials during titration (11). Mn/O ratio has been used to estimate fvac in a BioMnOx sample (13); however, the Mn and O peak intensities are also affected by Mn(III) (6). EXAFS multiple-scattering (MS) fitting can estimate fvac and bending angles, suggesting layer Mn(III) contents. Nevertheless, estimated bending angles determined by MS fittings have large uncertainties, and fvac is subject to interference from particle size (6, 19). In this study, simple LCF of Mn K-edge EXAFS spectra with two end members are used to estimate fMn(III) and fvac. Although their absolute values may not be accurate, the method indeed provides information about the relative fMn(III) and fvac, which also agrees well with EXAFS qualitative analyses and XRD results. However, one should be cautious to use this method without evidence from other techniques (e.g., XRD). Genesis of Mn(III). This study and previous studies (3, 6, 11, 13) demonstrate that layer Mn(III) cations are prevalent in BioMnOx. Intermediate Mn(III) cations during biotic Mn(II) oxidation are believed to be attached on cell surfaces (2, 4, 5). Therefore, Mn(III) in BioMnOx is more likely formed by abiotic reactions, such as reaction between residual dissolved Mn(II) and neoformed BioMnOx (2, 8, 29). BioMnOx formed at very early stages are MnO2 nanoparticles with hexagonal layer symmetry, containing significant amounts of vacant sites (6, 8). Therefore, the residual dissolved Mn(II) is immediately adsorbed on vacant sites

because of the rapid sorption kinetics (30). High pH favors metal sorption on mineral surfaces. The observed slower Mn(II) depletion kinetics at high pH (e.g., pH 8) (S8) further allows Mn(II) sorption to proceed more completely. Nevertheless, our results show that amounts of interlayer Mn decrease as formation pH increases. One possible explanation is that adsorbed Mn(II) on vacant sites at high pH is oxidized to Mn(III), which then enters the vacant sites and consequently, less Mn is present in the interlayers. Under the experimental conditions of this study, oxidation of the adsorbed Mn(II) by O2 in air is less likely because of its slow kinetics (31). Mn(II) is more likely oxidized by Mn(IV) surrounding the vacant sites on which Mn(II) is adsorbed. This explains the observations because high pH favors this conproportion reaction forming Mn(III) in solids (7) and results in less interlayer Mn. Enough Mn(III) accumulation in layers at high pH makes BioMnOx have a tendency to have orthogonal layer symmetry, as observed in this study and previous studies (6, 8, 9). Bargar and his co-workers (6, 8, 9) found in their time-series studies that orthogonal layer symmetry emerged from initial hexagonal phases in both biotic systems and an abiotic control, that is, a reaction between δ-MnO2 and 100 µmol dissolved Mn(II) at pH 8. Their findings support the reaction pathway proposed here. This pathway is analogous to Co(II) oxidation by hexagonal birnessite. Co(II) adsorbed at vacant sites is oxidized by the layer Mn(III), surrounding the vacant site, and then Co(III) is either incorporated into Mn-layers or may be still adsorbed on vacant sites (32). High pH may favor Mn(III) incorporation, provided that Mn(III) migrates out of Mn layers and is adsorbed on vacant sites at acidic pH (21). The Mn-Mncorner distances derived from EXAFS fittings decrease from ∼3.51 to ∼3.41 Å as the formation pH increases (Table S5, Supporting Information). The shorter Mn-Mncorner distances may suggest the presence of more Mn(III) than Mn(II) in the interlayer regions (21). This implies that part of the formed Mn(III) is still adsorbed at vacant sites. Effects of Metal Cations. Na+ and Ca2+ can compensate for the charge deficiency caused by the presence of Mn(III) in Mn-layers (20), and subsequently may act to stabilize Mn(III) in the layers. More Mn(III) is present in BioMnOx formed in CaCl2 than in NaCl solutions at the same pH (Figure 3b), which is consistent with previous studies (6). The difference may be related to the stronger adsorption behavior of Ca2+ than Na+ in interlayer regions (33). This may contribute to more Mn(III) in BioMnOx formed in the CaCl2 solution (Figure 3b), and a higher resistance to structural modifications induced by the weak-acid treatment (Figure 1b). The presence of Ni results in less Mn(III) in the Mn-layers. EXAFS results reveal that Ni(II) is adsorbed on vacant sites or incorporated into Mn-layers (S7). This can interfere with Mn(II) adsorption and subsequent oxidation and incorporation into vacant sites (34), resulting in less layer Mn(III). Adsorption of Ca2+ and Na+ on vacant sites is much weaker than heavy metals and may not be able to effectively compete with Mn(II). Furthermore, Ni(II) does not adsorb around Mn(III) as Ca2+ and Na+ do in interlayers (34) and is not able to stabilize Mn(III) in Mn-layers. Other heavy metals may have similar effects as Ni(II) because of their higher sorption affinities to birnessite than Mn(II) (35). Boofueng et al. (36) investigated Zn(II)/Mn(II) coprecipitation during bacterial Mn(II) oxidation and concluded that Zn(II) does not affect the BioMnOx structure. In fact, characteristic changes in their Mn EXAFS spectra around 8.10 and 9.25 Å-1 clearly suggest that the presence of Zn(II) decreases layer Mn(III) contents. Effects of Mn(II) Depletion Kinetics. In this study, the Mn(II) concentrations were not maintained constant and may contribute to the structural changes. Slower Mn(II) VOL. 44, NO. 12, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

4469

depletion allows longer contact time between dissolved Mn(II) and BioMnOx and may enhance Mn(III) formation (8, 9). Nevertheless, the observed structural changes are unlikely caused by the oxidation kinetics because in each experiment kinetics resemble each other (S8) but result in distinguishable crystal structures (e.g., pH6Na and pH8Na). The slow kinetics in the CaCl2 solution at pH 8 (S8) may contribute to orthogonal BioMnOx formation. Although no dissolved Mn(II) was present, aging may enhance the probability of adsorbed Mn(III) incorporation to form orthogonalBioMnOx,suchasthatobservedinthepH81weekNa sample. Aging effects should be investigated under other pH conditions in the future. Formation of an orthogonal structure in 50 µM Mn(II) (pH8Na) but not in 100 µM Mn(II) in NaCl solutions at pH 8 (pH8Ni0) (Figure 1) is probably because of more residual Ca2+ in the former system during the biomaterial washing procedure. Environmental Relevance. Vacant sites, large surface areas and strong oxidizing abilities make layered Mn-oxides one of the most reactive environmentally related minerals. The reactivity of layered Mn-oxides depends on Mn-layer structure and composition. Heavy metal sorption mainly occurs on vacant sites whose contents largely determine maximum metal sorption capacities. Vacant sites also affect photochemical dissolution of layered Mn-oxides (37). Mn(III) has been shown to affect oxidation mechanisms of Co(II) and Cr(III) by layered Mn-oxides (32, 38). Mn(III) is also a critical component for topotactic transformation of birnessite to todorokite, the most abundant tunneled Mn-oxide (39). The observations in this study suggest solution chemistry strongly controls layer structure and composition of BioMnOx and accordingly their environmentally related reactivity. Additionally, this study provides some clues on relations between layered Mn-oxide structure and the formation conditions. For instance, heavy metals are often present during Mn(II) oxidation in the environment, which may explain the dominance of hexagonal birnessite in nature (20). Field studies are required to construct a clear correlation between natural habitats of layered Mn-oxides and their crystal structure.

Acknowledgments M.Z. is grateful for a University of Delaware (UD) Institute of Soil and Environmental and Quality (ISEQ) Graduate Fellowship. M.Z. also thanks Dr. Samuel Webb at SSRL for his help with SR-XRD data collection and Dr. Bruce Ravel at NSLS for his help in EXAFS data analyses. This research was supported by Delaware EPSCoR with funds from National Science Foundation Grant EPS-0447610 and the State of Delaware and USDA Grant 2005-35107-16105. Use of the NSLS was supported by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences, under Contract No. DEAC02-98CH10886. The beamline X11 is supported by the Office of Naval Research and contributions from Participating Research Team (PRT) members. Portions of this research were carried out at the SSRL, a national user facility operated by Stanford University on behalf of the U.S. Department of Energy, Office of Basic Energy Sciences. The SSRL Structural Molecular Biology Program is supported by the Department of Energy, Office of Biological and Environmental Research, and by the National Institutes of Health, National Center for Research Resources, Biomedical Technology Program.

Supporting Information Available Bacterial preparation procedure, syntheses of the standard Mn-oxides, Mn XANES and EXAFS fitting results, Ni EXAFS and Mn(II) oxidation kinetics. This material is available free of charge via the Internet at http://pubs.acs.org. 4470

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 44, NO. 12, 2010

Literature Cited (1) Tebo, B. M.; Bargar, J. R.; Clement, B. G.; Dick, G. J.; Murray, K. J.; Parker, D.; Verity, R.; Webb, S. M. Biogenic manganese oxides: Properties and mechanisms of formation. Annu. Rev. Earth. Planet. Sci. 2004, 32, 287–328. (2) Webb, S. M.; Dick, G. J.; Bargar, J. R.; Tebo, B. M. Evidence for the presence of Mn(III) intermediates in the bacterial oxidation of Mn(II). Proc. Natl. Acad. Sci. U. S. A. 2005, 102 (15), 5558– 5563. (3) Pecher, K.; Mccubbery, D.; Kneedler, E.; Rothe, J.; Bargar, J.; Meigs, G.; Cox, L.; Nealson, K.; Tonner, B. Quantitative charge state analysis of manganese biominerals in aqueous suspension using scanning transmission X-ray microscopy (STXM). Geochim. Cosmochim. Acta 2003, 67 (6), 1089–1098. (4) Spiro, T. G.; Bargar, J. R.; Sposito, G.; Tebo, B. M. Bacteriogenic manganese oxides. Acc. Chem. Res. 2009, 43 (1), 2–9. (5) Bargar, J. R.; Tebo, B. M.; Villinski, J. E. In situ characterization of Mn(II) oxidation by spores of the marine Bacillus sp. strain SG-1. Geochim. Cosmochim. Acta 2000, 64 (16), 2775–2779. (6) Webb, S. M.; Tebo, B. M.; Bargar, J. R. Structural characterization of biogenic Mn oxides produced in seawater by the marine Bacillus sp. strain SG-1. Am. Mineral. 2005, 90, 1342–1357. (7) Mandernack, K. W.; Post, J.; Tebo, B. M. Manganese mineral formation by bacterial spores of the marine Bacillus, strain SG1: Evidence for the direct oxidation of Mn(II) to Mn(IV). Geochim. Cosmochim. Acta 1995, 59 (21), 4393–4408. (8) Bargar, J. R.; Tebo, B. M.; Bergmann, U.; Webb, S. M.; Glatzel, P. Biotic and abiotic products of Mn(II) oxidation by spores of the marine Bacillus sp. strain SG-1. Am. Mineral. 2005, 90, 143– 154. (9) Webb, S. M.; Tebo, B. M.; Bargar, J. R. Structural influences of sodium and calcium ions on the biogenic manganese oxides produced by the marine Bacillus sp. strain SG-1. Geomicrobiol. J. 2005, 22, 181–193. (10) Webb, S. M.; Fuller, C. C.; Tebo, B. M.; Bargar, J. R. Determination of uranyl incorporation into biogenic manganese oxides using x-ray absorption spectroscopy and scattering. Environ. Sci. Technol. 2006, 40 (3), 771–777. (11) Villalobos, M.; Toner, B.; Bargar, J.; Sposito, G. Characterization of the manganese oxide produced by Pseudomonas putida strain MnB1. Geochim. Cosmochim. Acta 2003, 67 (14), 2649–2662. (12) Miyata, N.; Tani, Y.; Maruo, K.; Tsuno, H.; Sakata, M.; Iwahori, K. Manganese(IV) oxide production by Acremonium sp. strain KR21-2 and extracellular Mn(II) oxidase activity. Appl. Environ. Microbiol. 2006, 72 (10), 6467–73. (13) Saratovsky, I.; Wightman, P. G.; Pasten, P. A.; Gaillard, J. F.; Poeppelmeier, K. R. Manganese oxides: Parallels between abiotic and biotic Structures. J. Am. Chem. Soc. 2006, 128 (34), 11188– 11898. (14) Villalobos, M.; Lanson, B.; Manceau, A.; Toner, B.; Sposito, G. Structural model for the biogenic Mn oxide produced by Pseudomonas putida. Am. Mineral. 2006, 91, 489–502. (15) Okazaki, M.; Sugita, T.; Shimizu, M.; Ohode, Y.; Iwamoto, K.; de Vrind de Jong, E. W.; de Vrind, J. P.; Corstjens, P. L. Partial purification and characterization of manganese-oxidizing factors of Pseudomonas fluorescens GB-1. Appl. Environ. Microbiol. 1997, 63 (12), 4793–9. (16) Burle, E.; Kirby-Smith, W. W. Application of formaldoxime colorimetric method for the determination of manganese in the pore water of anoxic estuarine sediments. Estuaries Coasts 1979, 2 (3), 198–201. (17) Hammersley, A. P. ESRF Internal Report, ESRF98HA01T, FIT2D V9.129 Reference Manual V3.1, 1998. (18) Webb, S. SIXPack: a graphical user interface for XAS analysis using IFEFFIT. Phys. Scr. 2005, T115, 1011–1014. (19) Bargar, J. R.; Fuller, C. C.; Marcus, M. A.; Brearley, A. J.; Rosa, M. P.D.l.; Webb, S. M.; Caldwell, W. A. Structural characterization of terrestrial microbial Mn oxides from Pinal Creek, AZ. Geochim. Cosmochim. Acta 2009, 73, 889–910. (20) Drits, V. A.; Lanson, B.; Gaillot, A.-C. Birnessite polytype systematics and identification by powder X-ray diffraction. Am. Mineral. 2007, 92, 771–788. (21) Silvester, E.; Manceau, A.; Drits, V. A. Structure of synthetic monoclinic Na-rich birnessite and hexagonal birnessite: II. Results from chemical studies and EXAFS spectroscopy. Am. Mineral. 1997, 82, 962–978. (22) Lanson, B.; Drits; V, A.; Silvester, E.; Manceau, A. Structure of H-exchanged hexagonal birnessite and its mechanism of formation from Na-rich monoclinic buserite at low pH. Am. Mineral. 2000, 85, 826–838.

(23) Drits, V. A.; Silvester, E.; Gorshkov, A. I.; Manceau, A. Structure of synthetic monoclinic Na-rich birnessite and hexagonal birnessite: I. Results from X-ray diffraction and selected-area electron diffraction. Am. Mineral. 1997, 82, 946–961. (24) Marcus, M. A.; Manceau, A.; Kersten, M. Mn, Fe, Zn and As speciation in a fast-growing ferromanganese marine nodule. Geochim. Cosmochim. Ac. 2004, 68 (14), 3125–3136. (25) Manceau, A.; Tommaseo, C.; Rihs, S.; Geoffroy, N.; Chateigner, D.; Schlegel, M.; Tisserand, D.; Marcus, M. A.; Tamura, N.; Chen, Z.-S. Natural speciation of Mn, Ni, and Zn at the micrometer scale in a clayey paddy soil using X-ray fluorescence, absorption, and diffraction. Geochim. Cosmochim. Acta 2005, 69 (16), 4007–4034. (26) Zhu, M.; Ginder-Vogel, M.; Sparks, D. L., Ni(II) sorption on biogenic Mn-oxides with varying Mn-octahedral layer structure Environ. Sci. Technol. 2010 (DOI 10.1021/es9037066). (27) Toner, B.; Manceau, A.; Webb, S. M.; Sposito, G. Zinc sorption to biogenic hexagonal-birnessite particles within a hydrated bacterial biofilm. Geochim. Cosmochim. Acta 2006, 70, 27–43. (28) Loomer, D. B.; Al, T. A.; Weaver, L.; Cogswell, S. Manganese valence imaging in Mn minerals at the nanoscale using STEMEELS. Am. Mineral. 2007, 92, 72–79. (29) Perez-Benito, J. F. Reduction of colloidal manganese dioxide by manganese(II). J. Colloid Interface Sci. 2002, 248 (1), 130– 135. (30) Fendorf, S.; Sparks, D. L.; Franz, J. A.; Camaioni, D. M. Electron paramagnetic resonance stopped-flow kinetic study of manganese(II) sorption-desorption on birnessite. Soil Sci. Soc. Am. J. 1993, 57, 57–62.

(31) Davies, S. H. R.; Morgan, J. J. Manganese (II) oxidation kinetics on metal oxide surfaces. J. Colloid Interface Sci. 1989, 129 (1), 63–76. (32) Manceau, A.; Drits, V. A.; Silvester, E.; Bartoli, C.; Lanson, B. Structural mechanism of Co2+ oxidation by the phyllomanganate buserite. Am. Mineral. 1997, 82, 1150–1175. (33) Murray, J. W. The interaction of metal ions at the manganese dioxide-solution interface. Geochim. Cosmochim. Acta 1975, 39 (4), 505–519. (34) Peacock, C. L.; Sherman, D. M. Sorption of Ni by birnessite: Equilibrium controls on Ni in seawater. Chem. Geol. 2007, 238, 94–106. (35) McKenzie, R. M. The adsorption of lead and other heavy metals on oxides of manganese and iron. Aust. J. Soil Res. 1980, 18, 61–73. (36) Boonfueng, T.; Axe, L.; Yee, N.; Hahn, D.; Ndiba, P. K. Zn sorption mechanisms onto sheathed Leptothrix discophora and the impact of the nanoparticulate biogenic Mn oxide coating. J. Colloid Interface Sci. 2009, 333 (2), 439–447. (37) Kwon, K. D.; Refson, K.; Sposito, G. Defect-induced photoconductivity in layered manganese oxides: A density functional theory study. Phys. Rev. Lett. 2008, 100, 146601. (38) Nico, P.; Zasoski, A. J. Importance of Mn(III) availability on the rate of Cr(III) oxidation on δ-MnO2. Environ. Sci. Technol. 2000, 34, 3363–3367. (39) Cui, H.; Liu, X.; Tan, W.; Feng, X.; Liu, F.; Ruan, H. D. Influence of Mn(III) availability on the phase transformation from layered buserite to tunnel-structured todorokite. Clay Clay Miner. 2008, 56 (4), 397–403.

ES1009955

VOL. 44, NO. 12, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

4471