Oxygen Isotope Evidence for Mn(II)-Catalyzed ... - ACS Publications

Subscriber access provided by University of Otago Library

Article

#xygen Isotope Evidence for Mn(II)-Catalyzed Recrystallization of Manganite (#-MnOOH) Andrew J. Frierdich, Michael J Spicuzza, and Michelle M Scherer Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.6b01463 • Publication Date (Web): 01 Jun 2016 Downloaded from http://pubs.acs.org on June 2, 2016

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a free service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are accessible to all readers and citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

Environmental Science & Technology is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 20

Environmental Science & Technology

Oxygen Isotope Evidence for Mn(II)-Catalyzed Recrystallization of Manganite (γγ-MnOOH)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

Andrew J. Frierdich1,*, Michael J. Spicuzza2, and Michelle M. Scherer3

1

School of Earth, Atmosphere & Environment, Monash University, Clayton, VIC 3800, Australia 2

3

Department of Geoscience, University of Wisconsin, Madison, WI, 53706, United States

Department of Civil and Environmental Engineering, University of Iowa, Iowa City, IA, 52242, United States.

*Corresponding author: Tel.: +61 (0)3 9905 4899; Fax: +61 (0)3 9905 4903; E-mail: [email protected]

21 22 23 24 25 26 27 28 29

Submitted to ES&T March 2016 Revision submitted to ES&T May 2016

1 ACS Paragon Plus Environment


30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50

Page 2 of 20

ABSTRACT Manganese is biogeochemically cycled between aqueous Mn(II) and Mn(IV) oxides. Aqueous Mn(II) often coexists with Mn(IV) oxides and redox reactions between the two (e.g., comproportionation) are well known to result in the formation of Mn(III) minerals. It is unknown, however, whether aqueous Mn(II) exchanges with structural Mn(III) in manganese oxides in the absence of any mineral transformation (similar to what has been reported for aqueous Fe(II) and some Fe(III) minerals). To probe whether atoms exchange between a Mn(III) oxide and water, we use a 17O tracer to measure oxygen isotope exchange between structural oxygen in manganite (γ-MnOOH) and water. In the absence of aqueous Mn(II), about 18% of the oxygen atoms in manganite exchange with the aqueous phase, which is close to the estimated surface oxygen atoms (≈11%). In the presence of aqueous Mn(II), an additional 10% (for a total of 28%) of the oxygen atoms exchange with water, suggesting that some of the bulk manganite mineral (i.e., beyond surface) is exchanging with the fluid. Exchange of manganite oxygen with water occurs without any observable change in mineral phase and appears independent of the rapid Mn(II) sorption kinetics. These experiments suggest that Mn(II) catalyzes manganese oxide recrystallization and illustrate a new pathway by which these ubiquitous minerals interact with their surrounding fluid.

TOC ART

51


Page 3 of 20

52 53 54 55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 87 88 89 90


INTRODUCTION Manganese is the second most abundant redox-active metal in the Earth’s crust. Its oxide minerals are ubiquitous and important trace element scavengers in soils, sediments, and seawater.1-5 Biogenic and abiogenic processes in these environments cycle manganese between reduced and oxidized forms (e.g., aqueous Mn(II) (Mn(II)aq) and Mn(IV) minerals, respectively).6-9 Initial biogenic manganese-oxides formed via Mn(II)aq oxidation typically consist of disordered δ-MnO2-like phyllomanganates.10 When Mn(II)aq concentrations are high or the oxidation rate is slow, mineral products have increasing crystallinity and lower manganese average oxidation states,11-16 which likely occurs as the result of abiotic interactions between newly-formed Mn(IV) precipitates that back-react with residual Mn(II) in solution.13, 14 Exposure of phyllomanganates to low concentrations of Mn(II)aq, for example, increases Mn(III) contents and sheet stacking order in the solid,13, 14, 17, 18 and consequently, reduces the number of layer vacancies and metal uptake capacity.19, 20 Similar observations have been reported for naturally occurring phyllomanganates.5, 21 Abiotic reactions between Mn(II)aq and structural Mn(IV) thus appear to affect the stability of solid-phase Mn(III) and the ultimate crystallinity, composition, and trace element uptake capacity of phyllomanganate mineral assemblages. At high Mn(II)aq concentrations (e.g., Mn(II):Mn(IV)>1), δ-MnO2 may be completely reduced and transformed to Mn(II,III)- (hausmannite, Mn3O4) and Mn(III)-bearing minerals (manganite, γ-MnOOH),14, 22-25 resulting in extensive exchange of manganese atoms between the aqueous phase and solid-phase.23 This exchange, however, is accompanied by net reduction and phase transformation of the Mn(IV) mineral. It is unclear whether Mn(II)aq exchanges with structural Mn(III) in manganese oxide minerals in the absence of net reduction and mineral transformation, similar to what has been reported for aqueous Fe(II) and some Fe(III) minerals.26-35 In the case of iron, we used stable 57Fe-enriched tracers to directly track atom exchange. However, manganese is mono-isotopic which prevents a similar approach, and radioisotopes of manganese (e.g., 54Mn,23 half-life ~312 days) may not be practical if manganese exchange is much slower than natural decay. Our recent work has shown that iron atom exchange between aqueous Fe(II) and goethite catalyzes oxygen isotope exchange between the mineral and water. Here we use oxygen as a proxy for atom exchange between Mn(II)aq and structural Mn(III) in manganese oxides. Specifically, we react manganite, the most stable Mn(III) oxide,36, 37 with Mn(II)aq under anoxic conditions in 17O-enriched water to track oxygen isotope exchange between the mineral and fluid as a proxy for manganese exchange. Our objectives are to 1) determine if Mn(II)aq facilitates oxygen isotope exchange and fractionation between water and manganite, 2) identify how exchange relates to Mn(II) solution dynamics, and 3) monitor the mineralogy before and after reaction to identify possible phase transformations. METHODS



91 92 93 94 95 96 97 98 99 100 101 102 103 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 119 120 121 122 123 124 125 126 127 128 129 130

Page 4 of 20

Manganite Synthesis. Manganite was synthesized by titration of a manganous sulfate solution with ammonium hydroxide. The manganous hydroxide precipitate was then oxidized by hydrogen peroxide, forming a black-brown suspension that was heated at 70°C for 24 hr before being washed by ≥18.2 MΩ cm water (hereon water). Once dry, the solid was de-aggregated to a fine powder. Full details of the synthesis are provided in the Supporting Information (SI). X-ray diffraction (XRD) patterns of the resulting powder reveal a product consistent with manganite. Scanning electron microscopy (SEM) images show particles having an acicular “needle-like” morphology (SI Figure S1) with a mean length by width of 565 (±218 nm) by 31 (±6 nm) (1σ, n=48). Measured BET specific surface area was found to be 61 m2 g-1, which equates to roughly 11% of the atoms located at the surface (See SI and Table S1 for calculations). Isotopic Exchange Experiments. A brief description of experimental protocols is given here. Experiments were conducted in a similar fashion as our prior work measuring oxygen isotope exchange between water and goethite38 and thus readers are referred elsewhere for exhaustive details. All experiments were conducted inside an anoxic chamber (~3% H2-97% N2) with O2 maintained at less than 1 ppm by continual circulation of the atmosphere over palladium catalysts. All labware and reagents were equilibrated with the chamber atmosphere for ≥48 hrs prior to use. Water was also pre-sparged with N2 inside the chamber prior to use. Experiments were performed with two isotopically distinct waters (differing in δ18O by ~16‰) to confirm unique isotopic exchange trajectories (i.e., approach equilibrium from opposite directions). These waters were obtained from Fairbanks (FB), Alaska and Houston (H), Texas, USA. Each water was identically enriched (within error) by mixing it with 17O-enriched water (Isotec-90.7%). Together, they were used to conduct duplicate reactions as they are chemically identical and have the same 17O enrichment; they only differ in their δ18O composition, which has no impact on measured isotopic exchange.38 Reactors were setup in duplicate by adding 8.8 mL of 17O-enriched water (FB and H) to 15 mL tubes containing 20±0.2 mg of pre-weighed manganite. Solution pH was buffered by adding 1 mL of a 250 mM HEPES (4-(2-hydroxyethyl)-1-piperazineethanesulfonic acid)/KBr solution. Reactions were then initiated by addition of 0.1 mL of 100 mM MnCl2. Both the HEPES/KBr buffer and Mn(II) solutions were prepared from 17O-enriched water; HEPES had a negligible effect on the water oxygen isotope composition as oxygen derived from HEPES comprises less than 0.2% of all the oxygen atoms in the final reactors. All tubes were wrapped in Al foil and mixed end-over-end until sampling. Solution pH remained constant (7.5±0.1) during the experiments. The experimental conditions were chosen to obtain a sufficient amount of Mn(II) sorption. Since previous work has shown that HEPES reduces Mn(IV) to Mn(III),23, 39 we conducted experiments to determine if the reduction of Mn(III) by HEPES was also possible. These reactions were initiated as described in the preceding paragraph except that the HEPES concentration was varied between 0 and 100 mM. Similar amounts of Mn(II) release to solution were observed for experiments containing 0 and 25 mM HEPES, suggesting that Mn(III) reduction by HEPES is insignificant. However, slightly higher Mn(II) concentrations were 4 ACS Paragon Plus Environment

Page 5 of 20

131 132 133 134 135 136 137 138 139 140 141 142 143 144 145 146 147 148 149 150 151 152 153 154 155 156 157 158 159 160 161 162 163 164 165 166 167


observed in the presence of 100 mM HEPES, and given the high sorption capacity of manganite, minor reduction at the lower concentrations used in our oxygen exchange experiments cannot be entirely ruled out (see SI for details). Reactors were sacrificed at specified time intervals by pouring the entire 10 mL suspension into a syringe capped with a Swinnex removable filter assembly. The suspension was then filtered (0.2 µm), with the filtrate collected into a clean 15 mL tube. The filtrate was acidified to 0.4 M HCl (trace metal grade) and saved for chemical analysis by inductively coupled plasma mass spectrometry (ICP-MS). The manganite solid was washed by passing 10 mL of 17O-enriched water through the filter assembly to remove residual salts and loosely bound Mn(II) that may be hydrated and hence artificially increase the measured oxygen isotope value of manganite. Analysis of this wash solution revealed that approximately 16% of the “sorbed” Mn(II) was removed. Calculations indicate that the remaining Mn(II) could, at most, change the measured oxygen isotope exchange values of manganite by 3.6% (see SI for details). The filter containing the washed solid was then removed from the filter assembly, placed into a glass vial, air-dried, then stored in a vacuum desiccator until oxygen isotope analysis. Isotopic Analysis, Data Presentation, and Exchange Calculations. Oxygen isotope ratios were measured by gas-source mass spectrometry following laser fluorination of the solid and conversion of liberated O2 to CO2 by reaction with a graphite rod. Measured isotopic values are presented in delta notation (δ), in units of per mil (‰). To assess the amount of exchange, we define “δ45[CO2]calc” which is proportional to the relative amount of 45[CO2] derived from our 17 O tracer (i.e., 12C16O17O). This value is equivalent to “δ13Ccalc” defined in Frierdich et al. (2015) and is described in detail therein.38 The amount of oxygen isotope exchange that a mineral undergoes with a fluid can be derived from a mass-balance equation,38 which simplifies to: % ℎ =

# # !!" , $% ,

× 100 ≅

× 100

(1)

& & !!" and $% are the initial (i) isotopic values of the mineral (MnOOH) and fluid (aq) and the values at time, t, respectively. The temporal oxygen isotope composition & of water ($% ) is effectively constant because oxygen in the water represents a >1000 molar

& # is approximately equal to $% (i.e., excess relative to oxygen in the solid manganite. Hence, $% –23.53‰). Oxygen isotope fractionation was determined by measuring 18O/16O ratios and oxygen isotope compositions are presented in standard δ notation ( ,- ! / ,/ !)01234

() = *(,-!/ ,/!)

0565

− 18 × 1000

(2)

where (18O/16O)standard is the isotope ratio for VSMOW (Vienna Standard Mean Ocean Water). The average measured δ18O value of the UWG-2 (Gore Mountain garnet) oxygen isotope standard during the course of this study was δ18O = 5.70 ± 0.14‰ (2σ; n=10), consistent with the accepted value of 5.8‰.40 The external precision, based on the average standard deviations of



168 169 170 171 172 173 174 175 176 177 178 179 180 181 182 183 184 185 186 187 188 189 190 191 192 193 194 195 196 197 198 199 200 201 202 203 204 205 206 207

Page 6 of 20

replicate analyses for select samples, for manganite was 0.03‰ for δ45[CO2]calc and 0.27‰ (2σ; n=3) for δ18O. The oxygen yield for manganite samples was 95 ± 8% (2σ; n=24). Dissolved manganese in the aqueous phase was measured by a Thermo X-series 2 quadrupole ICP-MS operated in collision cell mode. Blanks and calibration standards (SigmaAldrich TraceCERT) were run with each sample set. Quantification of manganese was done by analysis of counts at mass-55. RESULTS AND DISCUSSION Oxygen Isotope Exchange between Water and Manganite. We have previously shown that structural oxygen in goethite exchanges with water when aqueous Fe(II) exchanges with goethite Fe(III).38 Here we use a similar 17O-tracer approach to determine if Mn(II)aq promotes oxygen isotope exchange between manganite and water. Experiments were initiated by introduction of Mn(II)aq into pH 7.5 manganite suspensions. The manganese concentration in solution rapidly decreases, with roughly half taken up by the solid, and attains a steady state (within error) after 1 day (Figure 1). Control experiments show that Mn(II) alone is stable in solution and that manganite alone releases minimal Mn(II) to solution in the absence of added Mn(II)aq (Figure 1 and Table 1). These control experiments suggest that Mn(II)aq loss from the solution in the presence of manganite is due to sorption rather than homogeneous oxidation and/or precipitation. To probe whether Mn(II)aq catalyzed oxygen exchange, we measured the change in oxygen isotope values in manganite exposed to 17O-enriched water with and without Mn(II)aq. In the absence of Mn(II)aq, the 17O content of the manganite increased significantly over the course of several months (up to 0.8‰ positive shift in δ45[CO2]calc) (Table 1). In the presence of Mn(II)aq, the 17O content of the manganite increased even more (up to 1.2‰ positive shift in δ45[CO2]calc). In both cases, the increase in the 17O content in manganite is significantly greater than the variability observed for the natural abundance baseline (±0.06‰, 2σ, n=18), which represents the variability in δ45[CO2]calc values for the analysis of natural abundance oxides and silicates. Therefore, the substantial changes in δ45[CO2]calc indicate manganite uptake of 17O from the fluid occurs in both the absence and presence of Mn(II)aq. The 17O data can be used to calculate the percentage of oxygen atoms in manganite that have exchanged with the fluid by inserting the δ45[CO2]calc values (Table 1) into eq 1. Because the water contains a 1000 molar excess of oxygen relative to the manganite, the shifts in the oxygen isotope values of manganite are directly proportional to the amount of exchange, and after complete exchange, the manganite oxygen isotope composition (as δ45[CO2]calc) will be equal to the water within the precision of our measurements. In the absence of Mn(II)aq, about 18% of the oxygen atoms in manganite exchange with the aqueous phase after about 5 months, which is a few percent more than the estimated surface oxygen atoms (≈ 11%) (Figure 2). In the presence of Mn(II)aq an additional 10% (for a total of 28%) of the oxygen atoms exchange with water suggesting that some of the bulk manganite mineral (i.e., beyond surface) is exchanging with the fluid. The total extent of oxygen exchange, and the difference in exchange between our 6 ACS Paragon Plus Environment

Page 7 of 20

208 209 210 211 212 213 214 215 216 217 218 219 220 221 222 223 224 225 226 227 228 229 230 231 232 233 234 235 236 237 238 239 240 241 242 243 244 245 246


Mn(II)-free control and Mn(II) reaction experiment, is significantly larger than what is possible due to Mn(II) sorption alone. Interestingly, the oxygen isotope exchange dynamics appear to be decoupled from those of the Mn(II)aq solution concentration (i.e., exchange continues despite an apparent Mn(II)aq steady-state), illustrating that such mineral-fluid exchange cannot be captured by concentration measurements alone. We characterized the manganite solid with SEM and XRD following reaction with and without Mn(II)aq to identify if the particle size, morphology, crystallinity, or mineral phase changed as a result of mineral-water exchange. Despite ≈ 28% oxygen isotope exchange, no change in the phase identity of the solid is observed even after 150 days of reaction with Mn(II)aq (Figure 3). There is, however, a narrowing of the XRD peaks following reaction with Mn(II)aq for 150 days (SI Figure S2) indicating a larger coherent scattering domain size (i.e., larger crystallites). The larger domain size observed in the XRD patterns appears to be accompanied by a physical coarsening (decreased length, increased width) of the particles based on SEM observations (SI Figure S1). Lesser, yet measureable, changes in particle size and crystallinity are also observed for manganite suspended in water without Mn(II)aq. These results are thus evidence for Mn(II)-catalyzed (i.e., no phase transformation or net redox change) recrystallization of a stable Mn(III) oxide. Oxygen Exchange as a Proxy for Manganese Exchange. Our recent work has shown that Fe(II)aq-goethite iron exchange catalyzes goethite-water oxygen exchange.38 It is thus conceivable that the Mn(II)-promoted manganite-water oxygen exchange observed here is the result of Mn(II)aq-manganite manganese exchange. Oxygen exchange with goethite and manganite, however, are quite different. For example, the extent of oxygen exchange for goethite appears to reach a steady-state within a few weeks38 whereas oxygen exchange with manganite appears to continue to increase after several months (Figure 2). It is also difficult to quantify manganese exchange because it is unclear whether oxygen exchange is stoichiometric with respect to manganese. Previously, we observed non-stoichiometric exchange between iron and oxygen with goethite.38 Such non-stoichiometry may occur in the manganese system as well since the exchange rates of oxygen functional groups on the manganite surface are likely much slower than water oxygens ligating Mn(II)aq (see Frierdich et al. 2015 for further discussion).38 Additionally, Mn(III) in manganite is Jahn-Teller distorted41 and potentially has more labile MnO bonds relative to non-distorted octahedra, as noted for Cu(II) hexaaquo species.42 While the stoichiometry of manganese and oxygen exchange remains unclear, we can use the extent of oxygen exchange to estimate a minimum threshold of manganese exchange to be ~25-30% (Figure 2). Direct evidence for atom exchange between Mn(II)aq and δ-MnO2 has been shown via a radiogenic 54Mn tracer,23 with the solid-phase transforming to feitknechtite (β-MnOOH) and ultimately manganite. Whether Mn(II)aq continued to exchange with manganite is unclear as isotopic mixing of manganese was near-complete upon the initial phase transformation of δMnO2.23 The significant amounts of oxygen isotope exchange between manganite and water



247 248 249 250 251 252 253 254 255 256 257 258 259 260 261 262 263 264 265 266 267 268 269 270 271 272 273 274 275 276 277 278 279 280 281 282 283 284 285 286

Page 8 of 20

during reaction with Mn(II)aq observed in this study indicate that apparently stable Mn(III) oxides also significantly interact with the fluid phase despite an apparent chemical equilibrium. Mechanistic Considerations. The incorporation of water oxygen into manganite during aqueous suspension of this mineral suggests that recrystallization (dissolution-reprecipitation) is occurring. The control experiments in this study show that manganite alone in solution does not appreciably dissolve (only ~0.01% of the manganese was detected in solution, see Table 1 and SI) and that Mn(II)aq alone does not precipitate. Hence, the reactions between Mn(II)aq and manganite appear to induce a balanced recrystallization process that we propose is similar to electron transfer and atom exchange phenomena previously noted for interactions between aqueous Fe(II) and Fe(III) oxides.26-28 Conceptually, this process likely consists of manganese oxidative growth (eq 3), and after electron migration within manganite to a different surface site, coupled to manganese reductive dissolution (eq 4). >OH− + Mn2+ + H2O → MnOOH + 2H+ + e− (3) + − − 2+ (4) MnOOH + 2H + e → >OH + Mn + H2O − III Note that >OH is a highly Mn -coordinated (slow-exchanging) oxygen species associated with the solid. As written, these reactions illustrate a process for the incorporation of oxygen into manganite during recrystallization and imply non-stoichiometric behavior of manganese and oxygen exchange during the growth and dissolution of a MnOOH fundamental unit. The stoichiometric Mn:O isotope exchange expectation of 1:2 is reduced to 1:1 because the >OH− species never resides in the fluid phase. Although the presence of Mn(II)aq promotes the extent of oxygen isotope exchange between water and manganite, there is considerable exchange even in Mn(II)-free solutions. This contrasts with our prior work with goethite which exhibited no oxygen exchange (within error) in the absence of aqueous Fe(II).38 The δ18O values of manganite in Mn(II)-free solutions vary as a function of exchange identically (within error) as those reacted in solutions containing Mn(II)aq (Figure 4), indicating that the exchanged oxygens in each case are of the same component rather than from two unique oxygen pools (e.g., interlayer water from a non-crystalline phyllomanganate impurity versus structural manganite oxygen). The manganite surface does contain more singly-coordinated oxygen functional groups than goethite due to its contrasting structure (i.e., single versus double chains of edge-sharing octahedra). Although singlycoordinated groups are expected to readily exchange with water,38 they only comprise a few percent of the total oxygen atoms (SI). Background dissolution-reprecipitation via coupled disproportionation-comproportionation reactions (eq 5)36, 37, 43 is another possible mechanism for oxygen exchange in the absence of added Mn(II). 2ΜnΟΟΗ + 2Η+ ↔ MnΟ2 + Μn2+ + 2Η2O (5) From this equilibrium it is possible for mineral-water oxygen exchange to occur, as previously noted during the disproportionation-comproportionation of hausmannite.44 Although disproportionation is enhanced under acidic conditions, the onset has been proposed to occur at pH 7.5 for manganite, resulting in Mn(II)aq concentrations in the µM range.45 Indeed, manganese in solution in the Mn(II)-free controls (i.e., MnOOH alone) is observed at levels expected for 8 ACS Paragon Plus Environment

Page 9 of 20

287 288 289 290 291 292 293 294 295 296 297 298 299 300 301 302 303 304 305 306 307 308 309 310 311 312 313 314 315 316 317 318 319 320 321 322 323 324 325 326


manganite in thermodynamic equilibrium with a pH 7.5 fluid according to eq 5 (see SI for calculations). Although this dissolved manganese amounts to little net dissolution (i.e., ~0.01%), the forward and reverse rate of disproportionation-comproportionation in eq 5 could be sufficiently large enough to turnover substantial amounts of the mineral to produce the observed levels of manganite-water oxygen exchange. Low levels of dissolved manganese are also observed in the absence of HEPES buffer (SI Figure S3), which suggests that the dissolved Mn(II) in our control experiment (i.e., MnOOH alone) originates from Mn(III) disproportionation rather than Mn(III) reduction by HEPES. However, the possibility that the observed low levels of dissolved Mn(II) in our control experiment originate from minor reduction of Mn(III) by HEPES or from residual Mn(II) that was not oxidized during the original synthesis of manganite, cannot be entirely ruled out. This endogenous Mn(II), regardless of its origin, may act as a positive feedback mechanism in our control experiment to further oxygen exchange through the Mn(II)-catalyzed mechanism noted above (eqs 3 and 4). Consequently, coupled disproportionation-comproportionation and Mn(II)-Mn(III) electron transfer and atom exchange reactions may operate simultaneously with the overall contribution of each determined by pH, the concentration of Mn(II)aq, and mineralogical controls (e.g., grain size, crystal defects). In any case, particle coarsening (SI Figure S1) and XRD line narrowing (SI Figure S2) suggests that recrystallization, without a mineral phase change, may be energetically favored by reducing the manganite surface free energy. Environmental Implications. The oxygen isotope values of naturally occurring manganese oxides may reflect conditions of their formation environment, source water, and depositional mechanisms.44 However, oxygen isotope measurements of natural, and synthetic samples are limited.44, 46-48 Furthermore, proper interpretation of measured isotopic values of natural minerals requires that mineral-water equilibrium oxygen isotope fractionation factors are known and the diagenetic history is constrained since post-depositional alteration may “reset” the isotopic signature of the mineral. The equilibrium fractionation factor refers to the difference in the isotopic composition of two components that are at isotopic equilibrium. The measurable amounts of oxygen isotope exchange between manganite and water observed here suggests that the experimental calibration of the manganite-water equilibrium fractionation factor (∆18OMnOOH-H2O = δ18OMnOOH – δ18OH2O) is possible. Indeed, changes in the δ18O values of manganite (Table 1) occur with time and these variations are directly proportional to the percent of oxygen exchange (Figure 4). Manganite δ18O values shift by several per mil during isotopic exchange and reactions in waters having unique isotopic compositions (in δ18O, see Methods) causes manganite values to shift in discrete trajectories (Figure 4). The linear δ18OMnOOH vs exchange pathways may be used to assess the manganite-water equilibrium fractionation factor by extrapolation to complete exchange. The similarity in ∆18OMnOOH-H2O, which was approached from above and below the isotopic values of the two waters, suggests that the final extrapolated values are close to equilibrium values. Combination of the fractionation factors obtained from both waters yields a weighted average for ∆18OMnOOH-H2O of +3.97‰ (±1.45‰, 2σ) at 22°C (See SI Table S2 for details). This value is remarkably close to the 9 ACS Paragon Plus Environment


327 328 329 330 331 332 333 334 335 336 337 338 339 340 341 342 343 344 345 346 347 348 349 350 351 352 353 354 355 356 357 358 359 360 361 362 363 364 365 366

Page 10 of 20

measured fractionation between water and manganite from a marine sediment (+3.6‰) observed previously.46 Nevertheless, the extensive amounts of oxygen isotope exchange and fractionation between manganite and water observed here suggests that the oxygen isotope values of naturally occurring manganese oxides could be ‘reset’ under reducing conditions or during repeated aging in aqueous fluids of varying isotopic composition. Hence, the interpretation of oxygen isotope signatures of manganese oxides in the rock record require accurate constraints on their diagenetic history. The reactivity of manganite with water may complicate the use of its oxygen isotope composition as a suitable environmental proxy. However, the high reactivity of manganese oxide minerals in general makes them among the strongest trace elements scavengers in the environment. The manganite recrystallization, as observed from oxygen isotope exchange, illustrates that bulk atoms within the structure are accessible to the fluid even in the absence of a phase change. Therefore, trace elements and contaminants associated with manganese oxides may undergo mineral-fluid repartitioning during reactions with Mn(II)-bearing fluids, as has been previously described for iron oxides in the presence of Fe(II)aq.49, 50 As a result, these reactions could alter the capacity of manganese oxides to act as sources or sinks for important metals. ASSOCIATED CONTENT Supporting Information. Included are supplementary calculations, experimental details, data tables, and figures. This material is available free of charge via the internet at http://pubs.acs.org. AUTHOR INFORMATION Corresponding Author *E-mail: [email protected] Notes The authors declare no competing financial interest. ACKNOWLEDGEMENTS This material is based upon work supported by the United States National Science Foundation (NSF) under Award No. 1347848 to A.J.F. and NSF Grant No. EAR-1123978 to M.M.S. Any opinions, findings, and conclusions or recommendations expressed in this material are those of the authors and do not necessarily reflect the views of the NSF. Additional funding was provided by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences Award Number DE-FG02-93ER14389 to J.W. Valley, who assisted in stable isotope analysis and reviewed an early draft of this paper. The authors acknowledge use of facilities within the Monash Centre for Electron Microscopy and the Monash X-ray Analytical Platform. We thank Thomas Lapen and Tom Trainor for providing water from Houston, Texas and Fairbanks, Alaska, respectively.


Page 11 of 20

367 368 369 370 371 372 373 374 375 376 377 378 379 380 381 382 383 384 385 386 387 388 389 390 391 392 393 394 395 396 397 398 399 400 401 402 403 404 405 406 407 408 409 410 411


REFERENCES 1. Manceau, A.; Drits, V. A.; Silvester, E.; Bartoli, C. l.; Lanson, B., Structural mechanism of Co(II) oxidation by the phyllomanganate buserite. Am. Mineral. 1997, 82, 1150–1175. 2. Hlawatsch, S.; Garbe-Schonberg, C. D.; Lechtenberg, F.; Manceau, A.; Tamura, N.; Kulik, D. A.; Kersten, M., Trace metal fluxes to ferromanganese nodules from the western Baltic Sea as a record for long-term environmental changes. Chem. Geol. 2002, 182, 697–709. 3. Burns, R. G., The uptake of cobalt into ferromanganese nodules, soils, and synthetic manganese (IV) oxides. Geochim. Cosmochim. Acta 1976, 40, 95-102. 4. Peacock, C. L.; Sherman, D. M., Sorption of Ni by birnessite: Equilibrium controls on Ni in seawater. Chem. Geol. 2007, 238, 94-106. 5. Frierdich, A. J.; Hasenmueller, E. A.; Catalano, J. G., Composition and structure of nanocrystalline Fe and Mn oxide cave deposits: Implications for trace element mobility in karst systems. Chem. Geol. 2011, 284, 82-96. 6. Skinner, H. C. W.; Fitzpatrick, R. W., Biomineralization processes of iron and manganese: modern and ancient environments. Catena-Verlag: Cremlingen-Destedt, 1992; Vol. Catena Supplement 21, p 432. 7. Tebo, B. M.; Johnson, H. A.; McCarthy, J. K.; Templeton, A. S., Geomicrobiology of manganese(II) oxidation. Trends Microbiol. 2005, 13, 421-428. 8. Myers, C. R.; Nealson, K. H., Bacterial manganese reduction and growth with manganese oxide as the sole electron acceptor. Science 1988, 240, 1319-1321. 9. Blöthe, M.; Wegorzewski, A.; Müller, C.; Simon, F.; Kuhn, T.; Schippers, A., Manganese-cycling microbial communities inside deep-sea manganese nodules. Environ. Sci. Technol. 2015, 49, 7692-7700. 10. Tebo, B. M.; Bargar, J. R.; Clement, B. G.; Dick, G. J.; Murray, K. J.; Parker, D.; Verity, R.; M.Webb, S., Biogenic manganese oxides: Properties and mechanisms of formation. Annu. Rev. Earth Planet. Sci. 2004, 32, 287-328. 11. Mandernack, K. W.; Post, J.; Tebo, B. M., Manganese mineral formation by bacterial spores of the marine Bacillus, strain SG-1: Evidence for the direct oxidation of Mn(II) to Mn(IV). Geochim. Cosmochim. Acta 1995, 59, 4393-4408. 12. Toner, B.; Fakra, S.; Villalobos, M.; Warwick, T.; Sposito, G., Spatially Resolved Characterization of Biogenic Manganese Oxide Production within a Bacterial Biofilm. Appl. Env. Microbiol. 2005, 71, 1300-1310. 13. Learman, D. R.; Wankel, S. D.; Webb, S. M.; Martinez, N.; Madden, A. S.; Hansel, C. M., Coupled biotic-abiotic Mn(II) oxidation pathway mediates the formation and structural evolution of biogenic Mn oxides. Geochim. Cosmochim. Acta 2011, 75, 6048-6063. 14. Bargar, J. R.; Tebo, B. M.; Bergmann, U.; Webb, S. M.; Glatzel, P.; Chiu, V. Q.; Villalobos, M., Biotic and abiotic products of Mn(II) oxidation by spores of the marine Bacillus sp. strain SG-1. Am. Mineral. 2005, 90, 143-154. 15. 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, 1089-1098. 16. Mann, S.; Sparks, N. H. C.; Scott, G. H. E.; Jong, E. W. d. V.-d., Oxidation of manganese and formation of Mn3O4 (hausmannite) by spore coats of a marine Bacillus sp. Appl. Env. Microbiol. 1988, 54, 2140-2143.



412 413 414 415 416 417 418 419 420 421 422 423 424 425 426 427 428 429 430 431 432 433 434 435 436 437 438 439 440 441 442 443 444 445 446 447 448 449 450 451 452 453 454 455 456 457

Page 12 of 20

17. Zhu, M.; Ginger-Vogel, M.; Parikh, S.; Feng, X.-H.; Sparks, D. L., Cation effects on the layer structure of biogenic Mn-oxides. Environ. Sci. Technol. 2010, 44, 4465-4471. 18. Zhao, H.; Zhu, M.; Li, W.; Elzinga, E. J.; Villalobos, M.; Liu, F.; Zhang, J.; Feng, X.; Sparks, D. L., Redox Reactions between Mn(II) and Hexagonal Birnessite Change Its Layer Symmetry. Environmental Science & Technology 2016, 50, 1750-1758. 19. 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, 44, 4472-4478. 20. Droz, B.; Dumas, N.; Duckworth, O. W.; Peña, J., A Comparison of the Sorption Reactivity of Bacteriogenic and Mycogenic Mn Oxide Nanoparticles. Environ. Sci. Technol. 2015, 49, 4200-4208. 21. Bargar, J. R.; Fuller, C. C.; Marcus, M. A.; Brearley, A. J.; De la Rosa, M. P.; Webb, S. M.; Caldwell, W. A., Structural characterization of terrestrial microbial Mn oxides from Pinal Creek, AZ. Geochim. Cosmochim. Acta 2009, 73, 889-910. 22. Elzinga, E. J., Reductive Transformation of Birnessite by Aqueous Mn(II). Environ. Sci. Technol. 2011, 45, 6366–6372. 23. Elzinga, E. J.; Kustka, A. B., A Mn-54 Radiotracer Study of Mn Isotope Solid-Liquid Exchange during Reductive Transformation of Vernadite (δ-MnO2) by Aqueous Mn(II). Environ. Sci. Technol. 2015, 49, 4310–4316. 24. Lefkowitz, J. P.; Rouff, A. A.; Elzinga, E. J., Influence of pH on the reductive transformation of birnessite by aqueous Mn(II). Environ. Sci. Technol. 2013, 47, 10364-10371. 25. Perez-Benito, J. F., Reduction of Colloidal Manganese Dioxide by Manganese(II). J. Colloid Interf. Sci. 2002, 248, 130-135. 26. Frierdich, A. J.; Helgeson, M.; Chengshuai, L.; Wang, C.; Rosso, K. M.; Scherer, M. M., Iron atom exchange between hematite and aqueous Fe(II). Environ. Sci. Technol. 2015, 49, 8479−8486. 27. Handler, R. M.; Frierdich, A. J.; Johnson, C. M.; Rosso, K. M.; Beard, B. L.; Wang, C.; Latta, D. E.; Neumann, A.; Pasakarnis, T.; Premaratne, W. A. P. J.; Scherer, M. M., Fe(II)catalyzed recrystallization of goethite revisited. Environ. Sci. Technol. 2014, 48, 11302−11311. 28. Gorski, C. A.; Handler, R. M.; Beard, B. L.; Pasakarnis, T.; Johnson, C. M.; Scherer, M. M., Fe Atom Exchange between Aqueous Fe2+ and Magnetite. Environ. Sci. Technol. 2012, 46, 12399−12407. 29. Wu, L.; Beard, B. L.; Roden, E. E.; Johnson, C. M., Stable iron isotope fractionation between aqueous Fe(II) and hydrous ferric oxide. Environ. Sci. Technol. 2011, 45, 1847–1852. 30. Williams, A. G. B.; Scherer, M. M., Spectroscopic evidence for Fe(II)-Fe(III) electron transfer at the iron oxide-water interface. Environ. Sci. Technol. 2004, 38, 4782–4790. 31. Larese-Casanova, P.; Scherer, M. M., Fe(II) sorption on hematite: New insights based on spectroscopic measurements. Environ. Sci. Technol. 2007, 41, 471–477. 32. Jones, A. M.; Collins, R. N.; Rose, J.; Waite, T. D., The effect of silica and natural organic matter on the Fe(II)-catalysed transformation and reactivity of Fe(III) minerals. Geochim. Cosmochim. Acta 2009, 73, 4409-4422. 33. Pedersen, H. D.; Postma, D.; Jakobsen, R.; Larsen, O., Fast transformation of iron oxyhydroxides by the catalytic action of aqueous Fe(II). Geochim. Cosmochim. Acta 2005, 69, 3967–3977. 34. Handler, R. M.; Beard, B. L.; Johnson, C. M.; Scherer, M. M., Atom exchange between aqueous Fe(II) and goethite: An Fe isotope tracer study. Environ. Sci. Technol. 2009, 43, 1102– 1107. 12 ACS Paragon Plus Environment

Page 13 of 20

458 459 460 461 462 463 464 465 466 467 468 469 470 471 472 473 474 475 476 477 478 479 480 481 482 483 484 485 486 487 488 489 490 491 492 493 494 495 496 497


35. Neumann, A.; Wu, L.; Li, W.; Beard, B. L.; Johnson, C. M.; Rosso, K. M.; Frierdich, A. J.; Scherer, M. M., Atom exchange between aqueous Fe(II) and structural Fe in clay minerals. Environ. Sci. Technol. 2015, 49, 2786−2795. 36. Murray, J. W.; Dillard, J. G.; Giovanoli, R.; Moers, H.; Stumm, W., Oxidation of Mn(II): Initial mineralogy, oxidation state and ageing. Geochim. Cosmochim. Acta 1985, 49, 463-470. 37. Hem, J. D.; Lind, C. J., Nonequilibrium models for predicting forms of precipitated manganese oxides. Geochim. Cosmochim. Acta 1983, 47, 2037-2046. 38. Frierdich, A. J.; Beard, B. L.; Rosso, K. M.; Scherer, M. M.; Spicuzza, M. J.; Valley, J. W.; Johnson, C. M., Low temperature, non-stoichiometric oxygen-isotope exchange coupled to Fe(II)-goethite interactions. Geochim. Cosmochim. Acta 2015, 160, 38-54. 39. Simanova, A. A.; Peña, J., Time-resolved investigation of cobalt oxidation by Mn(III)rich δ-MnO2 using quick X-ray absorption spectroscopy. Environ. Sci. Technol. 2015, 49, 1086710876. 40. Valley, J. W.; Kitchen, N.; Kohn, M. J.; Niendorf, C. R.; Spicuzza, M. J., UWG-2, a garnet standard for oxygen isotope ratios: Strategies for high precision and accuracy with laser heating. Geochim. Cosmochim. Acta 1995, 59, 5223-5231. 41. Kohler, T.; Armbruster, T.; Libowitzky, E., Hydrogen Bonding and Jahn–Teller Distortion in Groutite,α-MnOOH, and Manganite,γ-MnOOH, and Their Relations to the Manganese Dioxides Ramsdellite and Pyrolusite. J. Solid State Chem. 1997, 133, 486-500. 42. Powell, D. H.; Helm, L.; Merbach, A. E., 17O nuclear magnetic resonance in aqueous solutions of Cu2+ : The combined effect of Jahn–Teller inversion and solvent exchange on relaxation rates. J. Chem. Phys. 1991, 95, 9258-9265. 43. Hem, J. D., Redox processes at surfaces of manganese oxide and their effects on aqueous metal ions. Chem. Geol. 1978, 21, 199-218. 44. Mandernack, K. W.; Fogel, M. L.; Tebo, B. M.; Usui, A., Oxygen isotope analyses of chemically and microbially produced manganese oxides and manganates. Geochim. Cosmochim. Acta 1995, 59, 4409-4425. 45. Duckworth, O. W.; Sposito, G., Siderophore−manganese(III) interactions II. Manganite dissolution promoted by desferrioxamine B. Environ. Sci. Technol. 2005, 39, 6045-6051. 46. Yapp, C. J., Oxygen and hydrogen isotope variations among goethites (α-FeOOH) and the determination of paleotemperatures. Geochim. Cosmochim. Acta 1987, 51, 355-364. 47. Savin, S. M.; Epstein, S., The oyxgen and hydrogen isotope geochemistry of ocean sediments and shales. Geochim. Cosmochim. Acta 1970, 34, 43-63. 48. Bar-Matthews, M.; Matthews, A., Chemical and stable isotope fractionation in manganese oxide—phosphorite mineralization, Timna Valley, Israel. Geological Magazine 1990, 127, 1-12. 49. Frierdich, A. J.; Catalano, J. G., Controls on Fe(II)-activated trace element release from goethite and hematite. Environ. Sci. Technol. 2012, 46, 1519−1526. 50. Frierdich, A. J.; Luo, Y.; Catalano, J. G., Trace element cycling through iron oxide minerals during redox-driven dynamic recrystallization. Geology 2011, 39, 1083–1086.

498 499



Page 14 of 20

500

TABLES

501 502

Table 1. Summary of aqueous Mn(II) concentrations, oxygen isotope values for manganite, and calculated extents of oxygen isotope exchange between manganite and water. δ18O (‰) b Time (d) [Mn(II)]aq (µ µM)a δ45[CO2]calc (‰)a % O Exchange FB H Mn(II) Control: No MnOOH 0 931 (19) NAc 1 921 (22) NA 7 930 (3) NA 100 945 (48) NA MnOOH Rxn with 1 mM Mn(II) at pH 7.5 0 931 (19) -27.68 (0.01) 1 511 (19) -27.46 (0.04) 7 444 (8) -27.35 (0.00) 15 428 (3) -27.27 (0.08) 30 360 (46) -27.10 (0.02) 60 458 (44) -26.92 (0.02) 100 404 (12) -26.70 (0.03) 150 363 (36) -26.48 (0.07) MnOOH Control: No added Mn(II) 0 0.00 -27.68 (0.01) 15 0.42 (0.06) -27.54 (0.01) 60 1.24 (0.04) -27.33 (0.01) 150 2.05 (0.07) -26.90 (0.01)

503 504 505 506 507 508 509

NA NA NA NA

NA NA NA NA

NA NA NA NA

0 5.1 (0.9) 7.6 (0.0) 9.4 (2.0) 13.4 (0.5) 17.6 (0.5) 22.5 (0.7) 27.5 (1.6)

-5.32 -5.09 -5.24 -5.81 -6.49 -6.83 -7.56 -7.74

-5.32 -4.76 -4.71 -4.46 -4.41 -4.30 -4.07 -3.40

0.00 3.3 (0.2) 8.0 (0.3) 17.9 (0.3)

-5.32 -5.71 -5.82 -6.58

-5.32 -5.16 -4.68 -4.08

a

Values derived from duplicate reactions in 17O-enriched waters from Fairbanks (FB), Alaska and Houston (H), Texas. The 17O enrichment, illustrated by δ45[CO2]calc, is identical within error for both waters. The δ45[CO2]calc values equal -23.53‰ and -23.52‰ for FB and H water, respectively. Reported values are the mean of the duplicate analyses and the numbers in parentheses represent the standard deviation. b Calculated from eq 1 using the measured δ45[CO2]calc values of manganite. c Not applicable.


Page 15 of 20


510

FIGURE CAPTIONS

511

Figure 1. Temporal manganese concentrations in solution during suspension of MnOOH (2 g L-

512

1

513

1 mM Mn(II) solution with no added solid MnOOH. Solution pH for all reactions is 7.5. All data

514

points represent the mean value from duplicate experiments; error bars are one standard

515

deviation and smaller than symbol if not visible.

) in the presence and absence of 1mM Mn(II)aq. Also shown is the manganese concentration of a

516 517

Figure 2. Temporal oxygen isotope exchange between water and manganite (γ-MnOOH) during

518

reaction with 1 mM Mn(II) at pH 7.5. Percent oxygen exchange is calculated according to eq 1

519

based on changes in the measured δ45[CO2]calc of the solid which results from incorporation of

520

17

521

suspended in 17O-enriched water. All data points represent the mean value from duplicate

522

experiments; error bars are one standard deviation and smaller than symbol if not visible.

O from enriched water. Also shown are Mn(II)-free controls where the solid minerals were

523 524

Figure 3. X-ray diffraction (XRD) patterns for the initial starting material (b), the starting

525

material after suspension in pH 7.5 buffered 17O-enriched water for 150 days in the absence (c)

526

and presence (d) of 1 mM Mn(II). Also shown are calculated XRD peak positions and relative

527

intensities for manganite (γ-MnOOH) (a) and hausmannite (Mn3O4) (e).

528 529

Figure 4. Oxygen isotope values for manganite during suspension in 17O-enriched water from

530

Houston, Texas (circles) and Fairbanks, Alaska (triangles) in the presence (filled symbols) and

531

absence (open symbols) of aqueous Mn(II). Black line represents the linear fit to the Mn(II)

532

reacted samples and the black dashed lines compose the error envelopes (2σ). Blue dotted lines

533

represent the δ18O values for the respective waters. Oxygen exchange fraction of the mineral is

534

calculated from eq 1 based on changes in δ45[CO2]calc for the mineral.



Figure 1. Temporal manganese concentrations in solution during suspension of MnOOH (2 g L-1) in the presence and absence of 1mM Mn(II)aq. Also shown is the manganese concentration of a 1 mM Mn(II) solution with no added solid MnOOH. Solution pH for all reactions is 7.5. All data points represent the mean value from duplicate experiments; error bars are one standard deviation and smaller than symbol if not visible. 91x90mm (300 x 300 DPI)

ACS Paragon Plus Environment

Page 16 of 20

Page 17 of 20


Figure 2. Temporal oxygen isotope exchange between water and manganite (γ-MnOOH) during reaction with 1 mM Mn(II) at pH 7.5. Percent oxygen exchange is calculated according to eq 1 based on changes in the measured δ45[CO2]calc of the solid which results from incorporation of 17O from enriched water. Also shown are Mn(II)-free controls where the solid minerals were suspended in 17O-enriched water. All data points represent the mean value from duplicate experiments; error bars are one standard deviation and smaller than symbol if not visible. 91x89mm (300 x 300 DPI)



Figure 3. X-ray diffraction (XRD) patterns for the initial starting material (b), the starting material after suspension in pH 7.5 buffered 17O-enriched water for 150 days in the absence (c) and presence (d) of 1 mM Mn(II). Also shown are calculated XRD peak positions and relative intensities for manganite (γ-MnOOH) (a) and hausmannite (Mn3O4) (e). 89x90mm (300 x 300 DPI)


Page 18 of 20

Page 19 of 20


Figure 4. Oxygen isotope values for manganite during suspension in 17O-enriched water from Houston, Texas (circles) and Fairbanks, Alaska (triangles) in the presence (filled symbols) and absence (open symbols) of aqueous Mn(II). Black line represents the linear fit to the Mn(II) reacted samples and the black dashed lines compose the error envelopes (2σ). Blue dotted lines represent the δ18O values for the respective waters. Oxygen exchange fraction of the mineral is calculated from eq 1 based on changes in δ45[CO2]calc for the mineral. 94x90mm (300 x 300 DPI)



Table of Contents Art 113x54mm (300 x 300 DPI)


Page 20 of 20

Oxygen Isotope Evidence for Mn(II)-Catalyzed ... - ACS Publications

Recommend Documents