Fe(II) reduction of pyrolusite (β-MnO2) and secondary mineral evolution
© The Author(s) 2017
Received: 30 August 2017
Accepted: 27 November 2017
Published: 5 December 2017
Iron (Fe) and manganese (Mn) are the two most common redox-active elements in the Earth’s crust . Reactions between Fe and Mn species, as well as with other common groundwater constituents, have significant impacts on mineral formation and dissolution , trace metal sequestration , and contaminant transformations [4, 5]. Our understanding of the health effects of Mn exposure to humans is also evolving, and recent research indicates that elevated Mn concentrations in drinking water may lead to developmental disorders in children, among other adverse health effects [6–9]. The present study focuses on redox reactions of ferrous iron (Fe(II)) with oxidized Mn(IV) solids at circumneutral pH. Thermodynamics predict that in the presence of Fe(II), all manganese species would exist as reduced Mn(II) as opposed to oxidized Mn(IV). Complex environmental systems, however, do not always adhere to the compositions implied by thermodynamic constraints, especially in complex media such as soil aggregates . For example, microorganisms can significantly impact the speciation of Fe and Mn between reduced and oxidized forms [11, 12] and lead to high local concentrations of dissolved Fe(II) or to Mn(IV) solids that form and persist in the presence of Fe(II) on transient but relevant time scales.
Summary of experimental results of previous studies of Fe(II) reacted with Mn-oxides
Fe(III) stays in solution for pH < 4 and XRD inconclusive
Lepidocrocite and trace goethite
Higher pH produces more goethite, noncrystalline Fe-oxide
Lepidocrocite and goethite
Less goethite under anoxic conditions compared to oxic
Fe(OH)3; 6-line ferrihydrite
Column pH 2.5–6
Natural Mn-oxide coated sand
Pyrolusite coated silica sand
2-line ferrihydrite and jacobsite (MnFe2O4)
Fe(III) precipitates inhibit reductive dissolution of pyrolusite by Fe(II)
Pyrolusite coated quartz
Schwertmannite or sulfate-substituted ferrihydrite
Poorly crystalline MnO2 (similar to birnessite) and freshwater sediment
Amorphic Fe(III) oxide
Fe phase proposed in equation but not characterized
Formation of an Fe(III) surface coating on Mn oxide solids may impact the rate or overall ability of Mn oxides to remain redox-active phases in environmental systems. In simulated acid-mine drainage systems, Mn(II) production from Mn oxides reacted with Fe(II) decreases with time, suggesting that evolution of a new Fe oxide surface interferes with the ability of underlying Mn oxides to accept electrons from aqueous Fe(II) by creating a passivating Fe-oxide layer . Further studies in this experimental system attributed changing rates of Fe(II) loss and Mn(II) production from batch reactors to Langmuir-type blocking of Mn(IV) surface sites by Fe(III) oxide precipitates using model simulations . In these studies, it was also difficult to concretely ascertain the composition of resulting Fe(III) reaction products. Fe(II)/Mn(IV) redox activity may decrease the oxidation capacity of Mn oxides, which have been demonstrated to be important oxidants for a variety of environmental processes including abiotic release of organic nitrogen in soil  and contaminant remediation processes [22, 23]; formation of an amorphous Fe(III) precipitate has previously been shown to inhibit Cr(III) oxidation by birnessite at pH 5.5 .
Our guiding hypothesis was that precipitation of Fe(III) minerals at the Mn oxide surface would lead to partial passivation of the Mn oxide reactivity. Thus, we evaluated the effect of aqueous Fe(II) on electron transfer reactions at Mn oxide surfaces by subjecting pyrolusite to successive exposures of Fe(II) at pH 7.5. Pyrolusite was chosen as a model Mn-oxide for this study because it is the most thermodynamically stable Mn mineral phase, and therefore represents the end-member case for Mn(IV) reduction by Fe(II). Many investigations involving Fe(II) and Mn oxides have occurred at lower solution pH values between 3–6 in order to simulate acid mine drainage conditions. Evaluation of Fe/Mn redox chemistry at circum-neutral pH values is also important, as anoxic Fe(II) plumes may persist in neutral pH environments in the presence of Mn oxides .
Alongside traditional methods of analysis (XRD, scanning electron microscopy, chemical Fe and Mn analyses), we utilized 57Fe Mössbauer spectroscopy in conjunction with isotopically enriched 57Fe(II) in order to increase the Fe signal (natural Fe contains ~ 2.2 mol% 57Fe). To further examine Fe(III) surface precipitate morphology and what effect this phase has on subsequent redox reactions with Fe(II), we exposed Mn oxide particles to a series of solutions buffered at pH 7.5 which contained either 57Fe(II) (Mössbauer-visible) or 56Fe(II) (Mössbauer-transparent). In this manner, we could subject Mn oxide solids to a series of Fe(II) exposures, but only a particular “pulse” of Fe(II) would be visible with Mössbauer spectroscopy throughout the experiment. Isotope labeling allowed us to track the chemical changes that occurred to a specific set of Fe atoms, even as more Fe(II) was introduced to the reactor.
Mn oxide solids characterization
Sequential batch experiments with isotopically-enriched aqueous Fe(II)
All reagents were used as received. Experiments were performed in an anoxic chamber with a 95% N2, 5% H2 atmosphere. The chamber contained multiple palladium catalysts to scavenge trace O2 and maintained an O2 level below 1 ppmv. All solutions were made with deionized water (> 18.2 MΩ-cm) that had been deoxygenated by N2 sparging and degassing in the anaerobic chamber. Aqueous Fe(II) stock solutions were prepared by dissolving enriched 56Fe or 57Fe metal (Chemgas, 99 and 96%, respectively) in 0.5 M HCl . To initiate Fe(II) redox experiments, 18 mL of a pH 7.5 buffer solution [25 mM 4-(2-hydroxyethyl)-1-piperazineethanesulfonic acid (HEPES) + 25 mM KBr] was spiked with either a 57Fe or 56Fe stock solution to yield an initial aqueous Fe concentration of approximately 3 mM. Prior to Fe addition, reactors were counter-spiked with an equivalent volume of 0.5 M NaOH to maintain initial pH. Reactors were equilibrated for 1 h before filtering through a 0.2-µm syringe filter to remove any potential Fe precipitates resulting from trace oxidants. Initial Fe(II) concentration was then measured, and 18 mg of pyrolusite was added to initiate the experiment (solids loading 1 g L−1, Fe/Mn molar ratio ~ 0.26). Reactors were placed on an end-over-end rotator and mixed in the dark. Periodically, small aliquots (~ 150 µL) of suspension were withdrawn, filtered with 0.2-µm nylon syringe filters, and used for chemical Fe and Mn analyses. Experiments were typically allowed to run for ~ 90 min. If solids for a particular experiment were scheduled to receive more than 1 treatment in an aqueous solution, experimental reactors were allowed to stand for a short amount of time to allow Mn solids to settle, where they could be easily removed with a pipette. Solids were placed in a microcentrifuge tube and centrifuged inside the anoxic chamber to pellet solids and facilitate removal of the residual aqueous supernatant. Mn solids were then resuspended in a new buffer solution containing 3 mM 57Fe(II), 56Fe(II), or no Fe, depending on the particular experiment, and an additional experiment was performed to investigate the movement of aqueous Fe and Mn into or out of solution. Solids were resuspended in new buffer solutions with or without additional aqueous Fe(II) from 1 to 9 times.
To reconcile the amount of Fe(II) lost from solution with the production of Mn(II) into solution, acid extractions were performed on recovered solids to measure total Fe and Mn species. Control reactors with an identical buffer system, Mn solids loading, and Fe/Mn ratio were mixed for 90 min before solids were collected and resuspended in deionized water. 5 M HCl was added to different reactors to obtain a distribution of pH values between ~ 1–2. Extraction reactors were allowed to mix for ~ 150–300 h, periodically removing samples for Fe and Mn analyses. Additional controls of unreacted pyrolusite in HCl resulted in no measurable Mn in solution.
Aqueous Fe(II) was measured photometrically using 1,10-phenanthroline at 510 nm . Fluoride was used to remove interferences from aqueous Fe(III) . The amount of Fe(III) in solution was determined by the difference of measured Fe(II) content and the total Fe concentration measured by reducing Fe(III) to Fe(II) with hydroxylamine HCl. Aqueous Mn was determined by modifying the formaldoxime method outlined in Morgan and Stumm  and Abel  using phenanthroline to complex interfering aqueous Fe.
Solids characterization with SEM and Mössbauer spectroscopy
At the end of each experiment solids were captured by filtration through a syringe filter with a removable 0.45-µm filter disc. A small portion of recovered solids (~ 1 mg) were removed from the filter disc and rinsed with deionized water to remove residual aqueous Fe, Mn, and buffer salts. Rinsed solids were placed on an aluminum microscopy stub and fixed with carbon tape. Imaging of resulting particles and surface precipitates was performed with a Hitachi S-4800 scanning electron microscope (SEM). Remaining Mn/Fe solids recovered after sequential reaction experiments were wrapped in Kapton oxygen-impermeable tape prior to analysis with 57Fe Mössbauer spectroscopy. Mössbauer spectra were collected in transmission mode using a 57Co source and a Janis cryostat with temperature control to 13 K. Mössbauer spectra were collected at room temperature, 140, 77, and 13 K, and data was calibrated to a spectrum of α-Fe foil collected at room temperature. Spectral fitting was performed with the Recoil Software package (http://www.isapps.ca/recoil/) .
Results and discussion
Fe(II) oxidation by Mn(IV) oxide and transformation of the secondary Fe oxide
Previous work has found that both the identity and morphology of Fe(III) precipitates formed from oxidation of Fe(II) varies depending on the Mn substrate (MnO2, Mn2O3, MnOOH) and solution conditions [pH, ionic strength, Fe(II) concentration] (Table 1). Most studies of Fe(II) reacting with Mn-oxides have been performed at acidic pH, while studies at circumneutral pH focused on biological mechanisms involved with the reaction. Most of these studies identified the formation of discrete Fe phases (only one study reported mixed Mn/Fe jacobsite (MnFe2O4) formation ) that were predominantly hydroxylated Fe phases (Fe(OH)3 or FeOOH) with only one study reporting magnetite (Fe3O4) formation . Our observations of lepidocrocite formation from Fe(II) oxidation by pyrolusite are consistent with previous results at pH 6 where lepidocrocite was produced across a range of Mn/Fe molar ratios .
To probe the evolution and continued transformation of the Fe-coated pyrolusite particles, we reacted the particles with additional aqueous Fe(II). During the second and third exposure of these particles to Fe(II), Fe(II) loss from and Mn release to solution still occurred, but decreased with each exposure (Fig. 2). Images of particles taken after each Fe(II) exposure reveal changes in particle morphology from rod-like structure (1 exposure) to a mixture of rod-like and spherical structures (Fig. 2). The proportion of spherical particles increases between exposures 2 and 3, and spherical morphology is consistent with magnetite particles . Magnetite formation after reaction of Fe(II) with lepidocrocite has been observed previously , although high sulfate concentrations (not present in our study) may inhibit this process [2, 38].
Evolution of secondary Fe oxide
To confirm the secondary formation of magnetite or maghemite (hereafter referred to as magnetite) analysis of pyrolusite particles reacted with different sequences of isotopically-labeled Fe(II) was performed using 57Fe Mössbauer spectroscopy. Mössbauer spectroscopy is specific to the 57-isotope of Fe and we employed both enriched 56Fe(II) and 57Fe(II) in different reaction sequences. Iron isotope labeling allowed for the use Mössbauer spectra to track a specific set of Fe atoms (the 57Fe atoms) through an experiment without spectral contribution from 56Fe atoms. Mössbauer spectra were collected at 77 K to differentiate magnetite from lepidocrocite by minimizing errors due to superparamagnetic behavior of magnetite . At 77 K, lepidocrocite [and most Fe(III) oxides] display doublet spectral features, whereas magnetite exhibits multiple sextets.
Relative abundances of lepidocrocite and magnetite/maghemite appearing in 57Fe Mössbauer spectra at 77 K
Appearance of 57Fe(II) in Seriesa
Mn-oxides are powerful natural oxidants, and Mn redox cycling plays a major role in contaminant fate and transport . Here we show that Fe-oxide coatings that form through the abiotic reaction of Fe(II) with Mn-oxide alter the surface properties of the Mn-oxide mineral, but do not shut down the particles’ redox activity. Our findings suggest that surface passivation through the formation of Fe-oxides may not be as extensive or complete as previously thought. In our experiments, we show that the conversion of the initially precipitated Fe-oxide (lepidocrocite) to magnetite is coincident with excess Mn release either from the underlying Mn-oxide or Mn incorporated in the lepidocrocite. These experiments were performed with pyrolusite, the most thermodynamically stable Mn(IV) oxide; thus we expect Mn-oxide reduction by Fe(II) to be a process applicable to a variety of Mn(III/IV)-oxides under environmental conditions. Our findings raise the interesting question of whether sustained redox reactivity in the presence of surface coatings is restricted to Fe(II)/Fe(III) interactions or extends to other environmentally important constituents such as reduced groundwater contaminants.
MMS lead project conceptualization and all authors contributed to the research design. MVS and RMH collected data, all authors analyzed data and contributed to writing the manuscript. All authors read and approved the final manuscript.
The authors would like to thank David Cwiertny, Chris Gorski, and Drew Latta for experimental help and Samantha Ying for helpful discussions. Portions of this work were funded by the US National Science Foundation through a Graduate Research Fellowship Program Grant (DGE-114747) and NIRT Grant (EAR-0506679).
The authors declare that they have no competing interests.
Ethics approval and consent to participate
Springer Nature remains neutral with regard to jurisdictional claims in published maps and institutional affiliations.
Open AccessThis article is distributed under the terms of the Creative Commons Attribution 4.0 International License (http://creativecommons.org/licenses/by/4.0/), which permits unrestricted use, distribution, and reproduction in any medium, provided you give appropriate credit to the original author(s) and the source, provide a link to the Creative Commons license, and indicate if changes were made. The Creative Commons Public Domain Dedication waiver (http://creativecommons.org/publicdomain/zero/1.0/) applies to the data made available in this article, unless otherwise stated.
- Klein C, Hurlbut CSJ (1999) Manual of mineralogy. Wiley, New YorkGoogle Scholar
- Hansel CM, Benner SG, Fendorf S (2005) Competing Fe(II)-induced mineralization pathways of ferrihydrite. Environ Sci Technol 39:7147–7153. https://doi.org/10.1021/es050666z View ArticleGoogle Scholar
- Bowell RJ (1994) Sorption of arsenic by iron oxides and hydroxides in soils. Appl Geochem 9:279–286View ArticleGoogle Scholar
- Pecher K, Haderlein SB, Schwarzenbach RP (2002) Reduction of polyhalogenated methanes by surface-bound Fe(II) in aqueous suspensions of iron oxides. Environ Sci Technol 36:1734–1741. https://doi.org/10.1021/es011191o View ArticleGoogle Scholar
- Ulrich HJ, Stone AT (1989) Oxidation of chlorophenols adsorbed to manganese oxide surfaces. Environ Sci Technol 23:421–428View ArticleGoogle Scholar
- Wasserman GA, Liu X, Parvez F et al (2006) Water manganese exposure and children’s intellectual function in Araihazar, Bangladesh. Environ Health Perspect 114:124–129Google Scholar
- Menezes-Filho JA, de Novaes CO, Moreira JC et al (2011) Elevated manganese and cognitive performance in school-aged children and their mothers. Environ Res 111:156–163. https://doi.org/10.1016/j.envres.2010.09.006 View ArticleGoogle Scholar
- Khan K, Wasserman GA, Liu X et al (2012) Manganese exposure from drinking water and children’s academic achievement. Neurotoxicology 33:91–97. https://doi.org/10.1016/j.neuro.2011.12.002 View ArticleGoogle Scholar
- Ljung K, Vahter M (2007) Time to re-evaluate the guideline value for manganese in drinking water? Environ Health Perspect 115:1533–1538. https://doi.org/10.1289/ehp.10316 View ArticleGoogle Scholar
- Ying SC, Masue-Slowey Y, Kocar BD et al (2013) Distributed microbially- and chemically-mediated redox processes controlling arsenic dynamics within Mn-/Fe-oxide constructed aggregates. Geochim Cosmochim Acta 104:29–41. https://doi.org/10.1016/j.gca.2012.08.020 View ArticleGoogle Scholar
- Di-Ruggiero J, Gounot AM (1990) Microbial manganese reduction mediated by bacterial strains isolated from aquifer sediments. Microb Ecol 20:53–63. https://doi.org/10.1007/BF02543866 View ArticleGoogle Scholar
- Lovley DR, Phillips EJP (1988) Novel mode of microbial energy metabolism: organic carbon oxidation coupled to dissimilatory reduction of iron or manganese. Appl Environ Microbiol 54:1472–1480Google Scholar
- Widerlund A, Ingri J (1996) Redox cycling of iron and manganese in sediments of the Kalix River estuary, Northern Sweden. Aquat Geochem 2:185–201. https://doi.org/10.1007/BF00121631 View ArticleGoogle Scholar
- van der Zee C, Slomp CP, Rancourt DG et al (2005) A Mossbauer spectroscopic study of the iron redox transition in eastern Mediterranean sediments. Geochim Cosmochim Acta 69:441–453View ArticleGoogle Scholar
- Stollenwerk KG (1994) Geochemical interactions between constituents in acidic groundwater and alluvium in an aquifer near Globe, Arizona. Appl Geochem 9:353–369. https://doi.org/10.1016/0883-2927(94)90058-2 View ArticleGoogle Scholar
- Villinski JE, O’Day PA, Corley TL, Conklin MH (2001) In situ spectroscopic and solution analyses of the reductive dissolution of MnO2 by Fe(II). Environ Sci Technol 35:1157–1163. https://doi.org/10.1021/es001356d View ArticleGoogle Scholar
- Villinski JE, Saiers JE, Conklin MH (2003) The effects of reaction-product formation on the reductive dissolution of MnO2 by Fe(II). Environ Sci Technol 37:5589–5596. https://doi.org/10.1021/es034060r View ArticleGoogle Scholar
- Postma D (1985) Concentration of Mn and separation from Fe in sediments—I. Kinetics and stoichiometry of the reaction between birnessite and dissolved Fe(II) at 10 °C. Geochim Cosmochim Acta 49:1023–1033. https://doi.org/10.1016/0016-7037(85)90316-3 View ArticleGoogle Scholar
- Postma D, Appelo CAJ (2000) Reduction of Mn-oxides by ferrous iron in a flow system: column experiment and reactive transport modeling. Geochim Cosmochim Acta 64:1237–1247. https://doi.org/10.1016/S0016-7037(99)00356-7 View ArticleGoogle Scholar
- Krishnamurti GSR, Huang PM (1988) Influence of manganese oxide minerals on the formation of iron-oxides. Clays Clay Miner 36:467–475View ArticleGoogle Scholar
- Reardon PN, Chacon SS, Walter ED et al (2016) Abiotic protein fragmentation by manganese oxide: implications for a mechanism to supply soil biota with oligopeptides. Environ Sci Technol. https://doi.org/10.1021/acs.est.5b04622 Google Scholar
- Laha S, Luthy RG (1990) Oxidation of aniline and other primary aromatic amines by manganese dioxide. Environ Sci Technol 24:363–373. https://doi.org/10.1021/es00073a012 View ArticleGoogle Scholar
- Eary LE, Rai D (1987) Kinetics of chromium(III) oxidation to chromium(VI) by reaction with manganese-dioxide. Environ Sci Technol 21:1187–1193View ArticleGoogle Scholar
- Amacher MC, Baker DA (1982) Redox reactions involving chromium, plutonium, and manganese in soils. PhD thesis, The Pennsylvania State UniversityGoogle Scholar
- McCobb TD, LeBlanc DR, Walter DA, et al. (1999) Phosphorus in a ground-water contaminant plume discharging to Ashumet Pond, Cape Cod, Massachusetts. US Geological Survey Water Resources Investigations Report 02-4306 70pGoogle Scholar
- Williams AGB, Scherer MM (2004) Spectroscopic evidence for Fe(II)–Fe(III) electron transfer at the iron oxide–water interface. Environ Sci Technol 38:4782–4790. https://doi.org/10.1021/es049373g View ArticleGoogle Scholar
- Fortune WB, Mellon MG (1938) Determination of iron with o-phenanthroline—a spectrophotometric study. Ind Eng Chem Anal Ed 10:0060–0064View ArticleGoogle Scholar
- Tamura H, Goto K, Yotsuyanagi T, Nagayama M (1974) Spectrophotometric determination of iron(II) with 1, 10-phenanthroline in the presence of large amounts of iron(III). Talanta 21:314–318View ArticleGoogle Scholar
- Morgan JJ, Stumm W (1965) Analytical chemistry of aqueous manganese. J (Am Water Works Assoc) 57:107–119Google Scholar
- Abel R (1998) Scavenging of particulate and dissolved lead compounds by coprecipitation with manganese oxyhydroxides. Electronic M.S. thesis, Virginia TechGoogle Scholar
- Rancourt DG, Ping JY (1991) Voight-based methods for arbitrary-shape static hyperfine parameter distributions in Mössbauer spectroscopy. Nucl Instrum Methods Phys Res B58:85–97View ArticleGoogle Scholar
- Krishnamurti GSR, Huang PM (1987) The catalytic role of birnessite in the transformation of iron. Can J Soil Sci 67:533–543View ArticleGoogle Scholar
- Morgan JJ, Stumm W (1964) Colloid-chemical properties of manganese dioxide. J Colloid Sci 19:347–359View ArticleGoogle Scholar
- Elzinga EJ (2011) Reductive transformation of birnessite by aqueous Mn(II). Environ Sci Technol 45:6366–6372. https://doi.org/10.1021/es2013038 View ArticleGoogle Scholar
- Murad E, Cashion J (2004) Mossbauer spectroscopy of environmental materials and their industrial utilization. Kluwer Academic Publishers, DordrechtView ArticleGoogle Scholar
- Cornell RM, Schwertmann U (2003) The iron oxides: structure, properties, reactions, occurrences and uses, 2nd edn. Wiley-VCH, WeinheimView ArticleGoogle Scholar
- Tamaura Y, Ito K, Katsura T (1983) Transformation of γ-FeO(OH) to Fe3O4 by adsorption of iron(II) ion on γ-FeO(OH). J Chem Soc Dalton Trans 2:189–194. https://doi.org/10.1039/DT9830000189 View ArticleGoogle Scholar
- Tamaura Y, Buduan PV, Katsura T (1981) Studies on the oxidation of iron(II) ion during the formation of Fe3O4 and α-FeO(OH) by air oxidation of Fe[OH]2 suspensions. J Chem Soc Dalton Trans 9:1807–1811. https://doi.org/10.1039/DT9810001807 View ArticleGoogle Scholar
- Gorski CA, Scherer MM (2009) Influence of magnetite stoichiometry on Fe(II) uptake and nitrobenzene reduction. Environ Sci Technol 43:3675–3680View ArticleGoogle Scholar
- Handler RM, Beard BL, Johnson CM, Scherer MM (2009) Atom exchange between aqueous Fe(II) and goethite: an Fe isotope tracer study. Environ Sci Technol 43:1102–1107View ArticleGoogle Scholar
- Gorski CA, Handler RM, Beard BL et al (2012) Fe atom exchange between aqueous Fe2+ and magnetite. Environ Sci Technol 46:12399–12407. https://doi.org/10.1021/es204649a View ArticleGoogle Scholar
- Frierdich AJ, Helgeson M, Liu C et al (2015) Iron atom exchange between hematite and aqueous Fe(II). Environ Sci Technol. https://doi.org/10.1021/acs.est.5b01276 Google Scholar
- Neumann A, Wu L, Li W et al (2015) Atom exchange between aqueous Fe(II) and structural Fe in clay minerals. Environ Sci Technol 49:2786–2795. https://doi.org/10.1021/es504984q View ArticleGoogle Scholar
- Wang Y, Morin G, Ona-Nguema G, Brown GE (2014) Arsenic(III) and arsenic(V) speciation during transformation of lepidocrocite to magnetite. Environ Sci Technol 48:14282–14290. https://doi.org/10.1021/es5033629 View ArticleGoogle Scholar
- De Grave E, Da Costa GM, I’w LH et al (1996) 57Fe Mossbauer effect study of A1-substituted lepidocrocitest. Clays Clay Miner 44:214–219View ArticleGoogle Scholar
- Frierdich AJ, Catalano JG (2012) Fe(II)-mediated reduction and repartitioning of structurally incorporated Cu Co, and Mn in iron oxides. Environ Sci Technol 46:11070–11077. https://doi.org/10.1021/es302236v View ArticleGoogle Scholar
- Kato S, Hashimoto K, Watanabe K (2012) Microbial interspecies electron transfer via electric currents through conductive minerals. PNAS 109:10042–10046. https://doi.org/10.1073/pnas.1117592109 View ArticleGoogle Scholar
- Ying SC, Kocar BD, Griffis SD, Fendorf S (2011) Competitive microbially and Mn oxide mediated redox processes controlling arsenic speciation and partitioning. Environ Sci Technol 45:5572–5579. https://doi.org/10.1021/es200351m View ArticleGoogle Scholar
- Oze C, Bird DK, Fendorf S (2007) Genesis of hexavalent chromium from natural sources in soil and groundwater. Proc Natl Acad Sci USA 104:6544–6549View ArticleGoogle Scholar
- Myers CR, Nealson KH (1988) Microbial reduction of manganese oxides: interactions with iron and sulfur. Geochim Cosmochim Acta 52:2727–2732View ArticleGoogle Scholar