- Research article
- Open Access
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
- Manganese oxide
- Electron transfer
- Mössbauer spectroscopy
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/) .
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.
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