Structural Transformation of Birnessite by Fulvic Acid under Anoxic

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Structural Transformation of Birnessite by Fulvic Acid under Anoxic Conditions Qian Wang, Peng Yang, and Mengqiang Zhu Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b04379 • Publication Date (Web): 22 Jan 2018 Downloaded from http://pubs.acs.org on January 27, 2018

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Environmental Science & Technology

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Structural Transformation of Birnessite by Fulvic Acid

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under Anoxic Conditions

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Qian Wang, Peng Yang, and Mengqiang Zhu*

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Department of Ecosystem Science and Management, University of Wyoming, Laramie, WY

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82071, United States

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Submitted to Environmental Science & Technology

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January 10, 2018

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*Corresponding author: Mengqiang Zhu

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E-mail: [email protected]

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Phone: 307-766-5523

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1 ACS Paragon Plus Environment


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ABSTRACT

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The structure and Mn(III) concentration of birnessite dictate its reactivity, and can be

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changed by birnessite partial reduction, but effects of pH and reductant/birnessite ratios on the

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changes by reduction remain unclear. We found that the two factors strongly affect structure of

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birnessite (δ-MnO2) and its Mn(III) content during its reduction by fulvic acid (FA) at pH 4 – 8

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and FA/solid mass ratios of 0.01 – 10 under anoxic conditions over 600 hours. During the

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reduction, the structure of δ-MnO2 is increasingly accumulated with both Mn(III) and Mn(II) but

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much more with Mn(III) at pH 8, whereas the accumulated Mn is mainly Mn(II) with little

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Mn(III) at pH 4 and 6. Mn(III) accumulation, either in layers or over vacancies, is stronger at

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higher FA/solid ratios. At FA/solid ratios ≥ 1 and pH 6 and 8, additional hausmannite and

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MnOOH phases form. The altered birnessite favorably adsorbs FA because of the structural

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accumulation of Mn(II, III). Like during oxidative precipitation of birnessite, the dynamic

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changes during its reduction are ascribed to the birnessite-Mn(II) redox reactions. Our work

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suggests low reactivity of birnessite co-existing with organic matter and severe decline of its

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reactivity by partial reduction in alkaline environment.

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INTRODUCTION

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Birnessite is a layered manganese (Mn) oxide and common in sediments and soils.1, 2 It is

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a metal scavenger and a strong oxidizer, and affects the environmental fate of toxic metals and

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organic pollutants,3-7 and biogeochemical cycles of carbon, nitrogen, sulfur and iron.5, 8-11 The

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high reactivity of birnessite stems from its open-framework (layered) structure,12 abundant

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defects (vacancies), and the high oxidizing potential of tetravalent Mn (Mn(IV)).1, 13-15

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Partial reduction of birnessite by organic matter can remarkably increase Mn(III) content


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of birnessite,4, 16-19 which can profoundly change birnessite metal sorption capacity and behavior,

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oxidizing reactivity and bandgap, and affect its transformation to tunneled structures.10, 12, 20-23

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Reduction of birnessite by humate or low molecular weight (LMW) organics under oxic

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conditions results in enrichment of both Mn(II) and Mn(III) in the reacted solid.16-18,

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Mn(III) proportion can be increased to surprisingly high levels (37.9% – 56%) based on analyses

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of X-ray photoelectron spectroscopy (XPS) and Mn K-edge X-ray absorption near-edge structure

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(XANES) spectroscopy.4, 16-19 However, the mechanism for the accumulation of such high levels

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of Mn(III), and the effects of pH and reductant/birnessite ratios on the Mn(III) accumulation and

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birnessite structural transformation remain unclear. It is poorly understood if Mn(III) resides on

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vacancies, in the layers, or both in the altered birnessite structure, which is important to know as

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the specific location of Mn(III) affects birnessite properties differently.20-22 It is unknown

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whether the reacted solid remains as birnessite with structural alteration, or converts to other Mn

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mineral phases of lower Mn oxidation states, such as hausmannite (Mn3O4) and MnOOH phases.

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The lack of the solid phase information also hinders understanding of the reaction mechanisms

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behind the changes in Mn oxidation state composition. Moreover, these studies were conducted

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in air where O2 could participate in the redox reaction, limiting the transferability of the results to

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suboxic environment.

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The

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Birnessite and aqueous divalent Mn (Mn(II)(aq)) coexist during reduction of birnessite and

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can react with each other. Adsorption and oxidation of Mn(II)(aq) by birnessite can dramatically

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change birnessite structure and Mn(III) concentration. Under circumneutral or alkaline

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conditions, Mn(III) is formed with a concomitant decrease in vacancy concentration at a low

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Mn(II)(aq)/MnO2 ratio due to the conproportionation or redox reaction between adsorbed Mn(II)

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and Mn(IV) in birnessite;6,

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at a high Mn(II)(aq)/MnO2 ratio, the birnessite-Mn(II)



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conproportionation converts birnessite into hausmannite and MnOOH phases.1, 26, 27, 29-31 The

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above processes are disfavored under acidic conditions.28, 29 Thus, we hypothesize that the above

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reductive transformation of birnessite by Mn(II) is responsible for Mn(III) accumulation in the

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solid phase during birnessite reduction, and the mineral phase of the reacted products depends on

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pH and the ratio of the reductant to birnessite.

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In the present study, we characterized temporal changes of the structure and Mn

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oxidation state composition of δ-MnO2 during its reaction (adsorption and oxidation) with fulvic

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acid (FA) and determined FA removal kinetics at various FA/solid mass ratios (R) and pH values

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under anoxic conditions. δ-MnO2, a nanocrystalline hexagonal birnessite, is analogous to

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natually occuring vernadite and biogenic Mn oxides.2 Fulvic acid, an important fraction of

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natural organic matter (NOM), was used as a model reductant to approximate NOM. The

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experiments were conducted at pH 4, 6 or 8, representing the common pH conditions of suboxic

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environment.28, 32, 33 Unlike in the previous studies,1, 4, 29 pH buffers were not used as they may

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reduce birnessite and complicate the interpretation of the results.27, 34 The mineral phases, local

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to intermediate range structures, and Mn oxidation state compositions of the reacted soilds

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collected at different time were thoroughly characterized using synchrotron-based X-ray

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diffraction (XRD), X-ray atomic pair distribution function (PDF) analysis, and Mn K-edge

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XANES spectroscopy. The observations are interpreted in light of the most recent progress on

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birnessite reductive transformaton by reacting with Mn(II)(aq).1, 26, 29 We find that accumulation

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and the location of Mn(III) in birnessite structure strongly depends on pH and R (FA/solid mass

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ratios), and the reacted solids include hausmannite and MnOOH phases in addition to Mn(III)-

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rich birnessite. These results advance our understanding of the dynamic changes in birnessite

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structure, composition and reactivity during birnessite reduction, an important process during Mn


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redox cycling, and shed light on the role of Mn oxides in mineralization and stabilization of

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organic carbon.

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MATERIALS AND METHODS

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Materials. δ-MnO2 was synthesized via reduction of KMnO4 by Mn(NO3)2 (Supporting

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Information, SI-1). Suwannee River FA (1S101F) was purchased from the International Humic

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Substance Society (SI-1). All other chemicals used were of analytical grade. The O2-free

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deionized (DI) water was used for preparation of chemical solutions (SI-1).

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Reaction of FA with δ-MnO2. Batch experiments were conducted in a COY vinyl

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anaerobic chamber (95% N2/5% H2) over a Pd catalyst at 22 ± 0.1 oC. A total 500 mL suspension

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in a capped polyethylene bottle containing δ-MnO2 (0.03, 0.3 or 3 g/L, equivalent to 0.343, 3.43

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or 34.3 mM Mn), FA (0, 0.03 or 0.3 g/L, equivalent to 0, 0.0142 or 0.142 mM FA based on the

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empirical molecular weight of 2114 Da35) and NaCl (50 mM) was maintained at pH 4, 6 or 8

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over 600 hours (h) while being vigorously stirred with a magnetic stirring bar. Thus, R was 0,

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0.01 (0.03: 3), 0.1 (0.3:3), 1 (0.3:0.3) or 10 (0.3: 0.03). Those conducted at R = 0 were the

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control experiments. The bottles were wrapped with aluminum foil to prevent exposure to light

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that can promote reductive dissolution of birnessite.8, 24, 36

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Prior to mixing the reactants, both the FA solution and the δ-MnO2 suspension were

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maintained at the target pH in the anaerobic chamber for 24 h by manually adding 0.1 M NaOH

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or HCl. An aliquot of the δ-MnO2 suspension was sampled (0-h sample) right before the mixing.

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The pH of the mixed suspension was quickly raised to the target value within ~ 0.02 h using an

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auto pH titrator (Metrohm 907 Titrando) when another aliquot of the suspension was sampled

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(0.02-h sample). Then the pH of the suspension was maintained for another 5 h using the titrator,

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during which the pH fluctuated slightly (± 0.01), before switching to manual pH adjustment for 5 ACS Paragon Plus Environment


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the remaining period of the experiments that lasted for 600 h. At pre-determined time intervals,

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aliquots (~ 20 mL) were collected and the solid and solution were separated using 0.2 µm

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membrane filters. The solid particles aggregated extensively and the amount of particles passing

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through the filter was negligible. The concentration of dissolved Mn ([Mn]dis) in the filtrates was

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measured by inductively coupled plasma optical emission spectrometry. The obtained solids

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were air dried in the anaerobic chamber, ground in an agate mortar and stored in a closed glass

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vial in the chamber prior to the following characterization. The FA remained in the solution was

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quantified as total dissolved organic carbon (DOC) (SI-2). Both specific UV absorbance at 254

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nm (SUVA254) and the ratio of the UV absorbance at 250 nm to that at 365 nm (E2/E3 index)

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based on UV-vis spectra of the filtrates were used to indicate compositional changes of FA (SI-2).

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A thorough characterization of reacted FA is beyond the scope of this work. LMW organic

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products in the filtrates were not measured.

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Characterization of Reacted Solids. Both XRD and total X-ray scattering data for the

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PDF analysis were collected from the air-dried solids using X-rays of 58.6491 keV (λ= 0.2114 Å)

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at beamline 11-ID-B at the Advanced Photon Source (APS), with the sample-to-detector

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distances of ~ 95 cm and ~ 15 cm, respectively (SI-3). The data collection and processing were

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as described in Wang et al.12

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The relative molar fractions of Mn(II), Mn(III) and Mn(IV) in the reacted solids were

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estimated using linear combination fitting (LCF) of Mn K-edge XANES spectra. The XANES

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spectra were collected from the solids in fluorescence mode using a monochromator equipped

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with Si (111) double crystal at beamline 10-BM-B at APS (SI-4). The LCF analysis was similar

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to the Combo method,37 and four Mn reference standards were finally chosen, i.e., MnSO4 solid

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for Mn(II), feitknechtite (β-MnOOH) for Mn(III), and a combination of potassium birnessite


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(KBi) and polymeric manganese oxides (PolyMnO2) for Mn(IV). The first three reference

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spectra were reported in Manceau et al.37 PolyMnO2 contains entirely Mn(IV)38, 39 and has a

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particle size smaller than that of δ-MnO2.40 As particle sizes can affect Mn K-edge XANES

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spectra,12 using PolyMnO2 in the fitting was to account for decreased particle size of δ-MnO2

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due to the reduction. A sensitivity analysis of the fitting strategy to the relative fraction of Mn(III)

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was conducted following the method in Grangeon et al.,41 showing the fitting results were

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reliable (SI-5). The obtained Mn oxidation state composition and the measured dissolved Mn(II)

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concentration were combined to quantify the net electron transfer from FA to δ-MnO2 (SI-6).

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A portion of the Mn(III) in the layers of birnessite migrates into interlayers upon

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acidification,6 leading to structural changes of birnessite. The changes can be used as another

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indicator for the presence of Mn(III) in the layers of birnessite. The end solid products reacted

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for 600 h at pH 6 and 8 were acidified and the structural changes were characterized (SI-7).

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RESULTS

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Dissolved Mn. In the control experiments (R = 0), < 1 µM dissolved Mn, likely Mn(II),

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is detected at pH 4 but at neither pH 6 nor pH 8 (Figure S1). The presence of dissolved Mn is due

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to both destabilization of structural Mn(III) in δ-MnO2 to form Mn(II) and disfavored Mn(II)

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adsorption under highly acidic conditions (pH 4).6

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At pH 6 and 8, R = 0.01, low [Mn]dis is detected because the low concentration of FA

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produces little Mn(II), most of which is resorbed onto δ-MnO2 (Figure 1B and 1C). The

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resorption also results in the initial low [Mn]dis under other conditions (Figure 1). Higher [Mn]dis

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is detected at R = 0.1 and 1, which increases with time but does not reach plateaus within 600 h

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(Figure 1). At R = 10 and pH 8, the solid is completely dissolved within 72 h (Figure 1C).

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Overall, at a given pH, a higher R results in higher [Mn]dis because more FA participates in the



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reduction (Figure 1); for a given R, a higher pH leads to much lower [Mn]dis and Mn release rates

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(Figure 1), ascribed to both disfavored FA adsorption due to increasing negativity of both FA

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and surface, producing less Mn(II), and enhanced Mn(II) sorption at the higher pH.

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Adsorption and Oxidation of FA. The DOC and solid C concentrations at different pH

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and R are given in Figure 2A, 2B, S2 and S3. Both solid C and DOC concentrations display slow

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removal kinetics of FA from solution and neither reach plateaus within 600 h (Figure 2A, 2B, S2

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and S3). Both lower pH and lower R lead to faster removal kinetics, a higher C removal from

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solution, and more C accumulated in the solids (Figure 2A, 2B, S2 and S3).

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The UV-vis absorbance of the FA in the filtrates decreases at all wavelengths

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simultaneously with increasing time (Figure S4), suggesting that FA adsorption on the solid

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plays a more important role for the FA removal from solution, rather than FA oxidative

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decomposition. Otherwise, the UV-vis absorbance would change only at the wavelengths

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belonging to the oxidized FA moieties.

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The band at 254 nm, corresponding to aromatic functional groups, becomes sharper with

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time (Figure S4), suggesting that remaining FA differs in composition from the initial FA. The

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SUVA254 and E2/E3 indices are used to characterize the compositional changes. Generally,

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SUVA254 decreases and E2/E3 increases significantly with increasing time, indicating that both

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aromaticity and molecular weight of remaining FA decrease as the reaction proceeds (Figure 2C,

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2D and S5), i.e., preferentially removing the fraction of FA high in both aromaticity and

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molecular weight, consistent with previous studies.42,

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remained in the solution are caused by both adsorption and oxidation of FA, probably more by

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the former which is well known to cause FA fractionation.42,

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fractionation are stronger at both lower R and lower pH, at which more FA is also removed

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The compositional changes of FA

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The FA oxidation and


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(Figure 2C, 2D and S5).

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Mineral Phases. The XRD patterns of the solids in the control experiments in the

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absence of FA (R = 0) consist of the characteristic peaks of δ-MnO2 (Figure S6A).40 After 600 h,

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these peaks change slightly, indicating subtle structural alteration of δ-MnO2. The changes of the

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XRD patterns at R = 0.01 are also subtle because of much lower concentration of FA used than

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that of δ-MnO2 (Figure S6B and S6C).

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Figure 3 shows the XRD patterns of the reacted δ-MnO2 at R > 0.01. The (002) peak

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broadens significantly with increasing reaction time under all conditions, indicating fewer

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stacked MnO6 layers and/or a decrease of the degree of parallelism between adjacent layers,12

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which can be caused by Mn(II) and/or FA adsorption on δ-MnO2.29 At R = 0.1, the solid remains

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as δ-MnO2 at all pH values during the reaction (Figure 3A-3C). A dip at ~ 1.97 Å becomes

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increasingly pronounced, indicating more Mn (II, III) capping the vacancies as the reaction

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proceeds.4, 23, 41, 44 At a given reaction time, the dip is more pronounced at lower pH, indicating

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more Mn (II, III) on the vacancies,6, 41, 44, 45 as higher [Mn]dis is produced.

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Similar but stronger changes are observed at R = 1 than at R = 0.1 (Figure 3D and 3E). In

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addition, a small portion of δ-MnO2 is transformed to other mineral phases at R = 1. At pH 6,

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minor hausmannite and groutite (α-MnOOH) form at 10 h (Figure 3D), but groutite disappears

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with time. At pH 8, feitknechtite (β-MnOOH) is detected at 24 h and then disappears over time;

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hausmannite forms at 120 h and becomes more abundant with time (Figure 3E). Hausmannite

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forms earlier, but its amount is less at pH 6 than at pH 8.

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At R = 10, both MnOOH phases form in addition to hausmannite. Hausmannite forms

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more rapidly and extensively at R = 10 than at R = 1, indicated by the earlier appearance of the

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stronger XRD peaks of hausmannite (Figure 3F). No solids remain after 72 h, indicating that 9 ACS Paragon Plus Environment


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both hausmannite and MnOOH phases are reductively dissolved completely by FA.

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Local Structural Changes. The PDF analysis is sensitive to structural changes of short

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and intermediate range, which may not be revealed unambiguously by XRD. The two peaks at ~

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2.86 Å ((Mn-Mn)L) and ~ 5.27 Å (MnL-MnIL), derived from the edge-sharing correlation in the

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layer (L) and the corner-sharing correlation between layers and interlayers (IL) (SI-3),

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respectively, are used to indicate the abundance of layer Mn(III) and interlayer Mn(II, III) of the

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reacted solids that remain as birnessite. The peak position corresponds to the interatomic distance

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of the two atoms involved in the correlation. A higher concentration of layer Mn(III) leads to a

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longer (Mn-Mn)L interatomic distance because Mn(III) has a larger ionic radius than Mn(IV).40

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Likewise, the MnL-MnIL peak is located at a higher r with a higher proportion of Mn(II) on

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vacancies because Mn(II) has a larger radius than Mn(III). In addition, when the vacancies are

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capped with more Mn(II, III), the MnL-MnIL peak becomes stronger. Thus, the PDF analysis

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allows for detailed characterization of the relative proportions of vacancy-adsorbed Mn(II) and

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Mn(III), and of the location (layers versus vacancies) of the accumulated Mn(III) with time

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during δ-MnO2 reduction.

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No changes in PDFs are observed for the control samples (R = 0) at pH 6 and 8 (Figure

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S7B). However, at pH 4, the (Mn-Mn)L peak slightly shifts to lower r, and the MnL-MnIL peak

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becomes stronger after 600 h of reaction. These slight changes suggest migration of layer Mn(III)

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onto vacancies (Figure S7B). Similarly at R = 0.1, with increasing time, the Mn(III) migration

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leads to the left shift (towards lower r) of the (Mn-Mn)L peak at pH 4 and 6 with less shift at pH

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6 (Figure 4A, 4B). In contrast, at R = 0.1 and pH 8, the (Mn-Mn)L peak shifts to the right

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(towards higher r) with time (Figure 4C), indicating that the layers contain more Mn(III).

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With increasing time, the MnL-MnIL peak shifts to the right and the peak intensity


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increases for all pH at R = 0.1 and 0.01 with less changes at R = 0.01 (Figure 4 and S8),

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indicating more Mn (II, III) capping the vacancies and a higher proportion of Mn(II) as the

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reaction proceeds. At a given reaction time, with increasing pH from 4 to 8 at R = 0.1, the MnL-

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MnIL peak area decreases and the peak position shifts to lower r (Figure 4 and 5). These changes

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indicate that at a higher pH, less Mn is adsorbed on vacancies, and the adsorbed Mn contains a

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higher proportion of Mn(III).

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To determine the kinetics of the structural alteration of δ-MnO2, the (Mn-Mn)L peak

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position and the MnL-MnIL peak area are plotted versus reaction time (Figure 5). Both change

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rapidly within 72 h and then slow down (Figure 5). At R = 1 and pH 6, the (Mn-Mn)L peak

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position shifts to the shorter distance up to 5 h, ascribed to the migration of layer Mn(III); the

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abrupt jump to a longer distance at 10 h is ascribed to the presence of groutite (Figure 3D, 5A

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and S8D) that contains longer Mn-Mn pairs. At R = 1 and pH 8, the (Mn-Mn)L interatomic

240

distance increases before the appearance of feitknechtite at 5 h, indicating more layer Mn(III) in

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δ-MnO2 (Figure 5A and S8E). After 5 h, the increase of the distance is due to both Mn(III)

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enrichment in δ-MnO2 and the presence of feitknechtite and hausmannite.

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Mn Oxidation States and Net Electron Transfer. The LCF analysis is used to

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determine the molar fraction of each Mn oxidation state in the reacted solids. A comparison of

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the fits to the spectra shows good agreement (Figure S10). At R = 0.1, the Mn(II) fraction

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reaches 9.3% at pH 6 and 10.9% at pH 4 after 600 h; however, the Mn(III) fractions (0 – 3.6%)

247

are very low and change negligibly (Figure 6 and Table S1). At pH 8 and R = 0.1 (Figure 6 and

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Table S1), the Mn(III) fraction increases from 0.4% to 14.2% with negligible changes in the

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Mn(II) fraction (0.4 to 4.0%). At pH 8 and R = 1, the fractions of Mn(III) and Mn(II) reach

250

27.9% and 11.2%, respectively, within 600 h (Figure 6 and Table S1). Our fitting may



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underestimate the Mn(III) fractions. For example, the LCF analysis shows that triclinic

252

birnessite, which ideally consists of one third of Mn as Mn(III), contains only 26.3% Mn(III);

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and a chemical titration analysis shows that as-synthesized δ-MnO2 contains 6% Mn(III) (SI-1),

254

but Mn(III) is not detected by the LCF analysis (Table S1).

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Figure 7 shows the estimated net electron transfer from FA to δ-MnO2 during the reaction

256

based on dissolved Mn(II) and the LCF-determined Mn oxidation state composition (Table S1).

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The net electron transfer versus reaction time better represents the progress of the redox reaction

258

than based on only dissolved Mn(II) because the solids contain high reduced Mn (i.e., Mn(II) and

259

Mn(III)). At R = 0.1, the net electron transfer at pH 4 is much higher than at pH 6 and pH 8; and

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interestingly, no significant difference occurs between pH 6 and pH 8 at R = 0.1 (Figure 7).

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DISSCUSSION

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Adsorption and Oxidation of FA on δ-MnO2. Adsorption of NOM on mineral surfaces

263

can involve electrostatic interactions, ligand exchange, cation bridging, and hydrophobic

264

interactions.42, 46 The electrostatic interactions make little contribution to adsorption of FA on

265

δ-MnO2 because both FA and δ-MnO2 are negatively charged at pH 4 - 8. The carboxylic and/or

266

phenolic groups of FA are responsible for ligand exchange, likely with >Mn-OH at the edge sites

267

of the δ-MnO2 layers because the hydroxyls associated with vacancies are inert for ligand

268

exchange as each is coordinated to two Mn atoms. However, adsorbed Mn(II, III) on vacancies

269

can enhance FA adsorption via cation bridging.47 Hydrophobic interactions could also contribute

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to the adsorption of FA on δ-MnO2.42

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Both the surface properties of the solid and the FA composition (because of FA oxidation)

272

are changing during reaction of FA with δ-MnO2, affecting FA removal from solution by

273

adsorption. The enrichment of Mn(II, III) increases PZC of δ-MnO2, and both hausmannite and 12 ACS Paragon Plus Environment

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the MnOOH phases have higher PZC values than birnessite.48 These lead to less negatively

275

charged solids, favoring FA adsorption. In addition, complexation of FA with dissolved Mn in

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solution decreases the negative charge of FA and favors FA adsorption as well. Thus, the system

277

is changing towards favoring FA adsorption with increasing time, which keeps increasing the FA

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adsorption capacity of the solid. Consequently, it is hard for FA adsorption to reach plateaus with

279

time, i.e., displaying the slow FA removal kinetics (Figure 2).

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In addition, the reduction rate of δ-MnO2 by FA decreases over time, which is also

281

observed during birnessite reduction by other reductants.18, 23, 36, 49 The decreased rate can be

282

caused by the depletion of oxidizable functional groups of FA, and the reduced oxidizing

283

potential of MnO2/Mn(II)(aq).49 Surface passivation of δ-MnO2 by adsorbed Mn(II, III) and

284

adsorption of oxidized FA could also contribute.18, 23 In fact, the accumulation of Mn(III) in δ-

285

MnO2 with time (e.g., R = 0.1 and pH 8 in Figure 5) indicates lower reactivity of Mn(III) than

286

Mn(IV) in oxidizing FA. The accumulation of MnOOH and hausmannite at R = 0.1 (Figure 3)

287

also indicates the lower oxidizing reactivity of these phases than birnessite, and thus their

288

formation slows down the overall reduction.

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Low pH favors both adsorption and oxidation of FA by δ-MnO2. Both FA and δ-MnO2

290

surfaces become more protonated at lower pH, decreasing their negative charge and favoring FA

291

adsorption and further oxidation (Figure 1A and 2A), consistent with Allard et al.42 Lower pH

292

also favors FA oxidation from a thermodynamic perspective because the redox reaction

293

consumes H+.50 However, similar amounts of electron transfer at pH 6 and 8 are observed

294

(Figure 7), probably due to uncertainties in the calculation of electron transfer that relies on Mn

295

oxidation state composition. In terms of dissolved Mn, low pH supresses Mn(II) back adsorption

296

to δ-MnO2,6 increasing dissolved Mn concentration and thus favorable to FA adsorption.



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Moreover, structural modification and phase transformation of δ-MnO2 during oxidizing FA

298

highly depend on pH, which further affects FA adsorption and oxidation as the reaction proceeds.

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Mn(III) in the Structure of δ-MnO2. Trivalent Mn remarkably affects birnessite

300

properties.10, 12, 20-23 Our XRD and PDF analyses indicate that at R = 0.1, δ-MnO2 is enriched in

301

Mn(III) at pH 8 but not at pH 4 and 6 under anoxic conditions (Figure 3, 4 and 5). The

302

accumulation of Mn(III) in the layers at pH 8 is further supported by the acidification experiment

303

showing that a considerable amount of the layer Mn(III) in the solids at pH 8 migrates out of

304

layers upon acidification (SI-7).

305

Two mechanisms can lead to the Mn(III) accumulation in δ-MnO2 at pH 8 during

306

δ-MnO2 reduction by FA. In the first mechanism, Mn(III) is formed as an intermediate product

307

during Mn(IV) reduction to Mn(II) by FA. A portion of the Mn(III) is not reactive enough to

308

oxidize FA while being stabilized on the surface by FA complexation.16, 17 Since FA most likely

309

attacks the edge sites, the intermediate Mn(III) initially formed must also be located at the edge

310

sites. Our PDF data indicate that a significant portion of the accumulated Mn(III) resides in the

311

layers of δ-MnO2. Therefore, if this mechanism applies, the Mn(III) formed at the edge sites

312

must be redistributed in the layers, likely through intra-layer electron transfer between Mn(IV)

313

and Mn(III). However, such mechanism cannot explain the presence of abundant Mn(III) on

314

vacancies. In the second mechanism, Mn(III) forms as a result of the conproportionation between

315

Mn(IV) and Mn(II).23 In this mechanism, Mn(II) produced by Mn(IV) reduction adsorbs on a

316

vacancy and then reacts with a Mn(IV) cation surrounding the vacancy to form two Mn(III)

317

cations, with one in the layer and the other adsorbed on the vacancy or incorporated into the

318

layer.6 This conproportionation mechanism does not require FA complexation to stabilize the

319

Mn(III). Numerous recent studies support the conproportionation mechanism.1,

26, 27

This 14

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320

mechanism can also better explain the presence of Mn(III) on vacancies than the intermediate

321

mechanism.

322

Interestingly, at pH 4 and 6, the content of layer Mn(III) decreases but the content of the

323

all structural Mn(III) (both layered and interlayered) remains almost constant or increases

324

slightly at low levels with increasing time. This suggests that layer Mn(III) migrates into the

325

interlayer without further disproportionation. Grangeon et al. shows that adsorption of Zn2+ and

326

Ni2+ on vacancies induces the migration of layer Mn(III) at pH < 7.51, 52 Mn(II) adsorbs on

327

vacancies at low pH is not supposed to reduce layer Mn(IV) and it probably behaves similarly as

328

Zn2+ and Ni2+ at pH 4 and 6, leading to decreased layer Mn(III) content. Alternatively, the

329

leftwards peak shifts at ~ 2.87 Å in the PDFs are not derived from the decrease of the layer

330

Mn(III) content. The shift could be caused by decreasing particle sizes that lead to structural

331

contraction.12 Regardless of the mechanism leading to the PDF peak shift, we do not observe

332

significant accumulation of Mn(III) in the structure of δ-MnO2 at either pH 4 or 6 during FA

333

interaction with δ-MnO2.

334

The low Mn(III) content in δ-MnO2 structure at pH 4 and 6 contrast to the results of the

335

previous studies under oxic conditions based on XPS analyses, showing 34% - 56% Mn(III) after

336

δ-MnO2 reacted with humate and oxalate at pH ~ 5.2.16,

337

exceptionally strong Mn(III)-humate or Mn(III)-oxalate complexes are proposed to form on the

338

solid surface, inhibiting further reduction of Mn(III) to Mn(II).16, 17 However, both humate and

339

oxalate are weak ligands, and their complexes with Mn(III) are not strong enough to be stable at

340

acidic pH.53 Indeed, even the highly stable Mn(III)-Desferrioxamine B complexes exist only at

341

alkaline pH.54, 55

342

17

In those previous studies,

The previously observed high Mn(III) fractions16-18 are more likely due to the presence of



Page 16 of 31

343

O2, formation of low valence of Mn phases, uncertainties in Mn(III) quantification, and/or

344

properties of starting materials. Oxidation of adsorbed Mn(II) on birnessite surfaces by

345

atmospheric O2, and formation of MnOOH and Mn3O4 can both increase Mn(III) concentration.

346

The XPS analysis based on Mn 2p3/2 spectra and its probing chemical composition of surface

347

structures, not the bulk structure, may lead to overestimation of Mn(III) fractions.56, 57 In contrast,

348

a more recent work on the oxidation of bisphenol A by Mn(III)-rich δ-MnO2 at pH 7 under oxic

349

conditions shows essentially no accumulation of either bulk or surface Mn(III) using XANES

350

spectroscopy and more accurate XPS analyses based on Mn 3p spectra.4 No Mn(III)

351

accumulation in the previous study4 could be caused by the already high Mn(III) content (31%)

352

in the starting δ-MnO2 material that cannot accommodate additional Mn(III), and/or by pH 7 that

353

is not high enough to cause significant Mn(III) accumulation.

354

Formation of MnOOH and Hausmannite. The transformation product of δ-MnO2

355

depends on the molar ratio of produced Mn(II) to δ-MnO2. High Mn(II)/MnO2 ratios result in

356

transformation of birnessite to MnOOH and hausmannite (Mn3O4),1, 29, 58 consistent with the

357

observations in the present work (Figure 3). The formation of MnOOH needs a lower

358

Mn(II)/MnO2 ratio than the formation of Mn3O4,29 accounting for the observed earlier formation

359

of MnOOH. However, groutite and feitknechtite are metastable and can convert into

360

hausmannite in the presence of Mn(II)(aq),29, 59 or they are dissolved by FA reduction (e.g., at R =

361

10), both leading to their disappearance with time.

362

Groutite and hausmannite also form at R = 1 and pH 6 (Figure 3) where the molar ratio of

363

[Mn]dis to the remaining Mn(IV) is 0.03 and 0.05 – 0.5, respectively. However, Lefkowitz et al.29

364

find that hausmannite forms at pH 8 but not at pH < 7 during the reaction of [Mn]dis (0.2 – 2.2

365

mM) with acid birnessite (0.05 g/L) ([Mn(II)]/[Mn(IV)] = ~ 0.35 – 3.8) under anoxic conditions.


Page 17 of 31


366

The difference could be caused by the different Mn(II)/Mn(IV) ratios in the two studies or the

367

lower reactivity of acid birnessite than δ-MnO2.26 The adsorbed FA may also affect the

368

transformation products.

369

Environmental Implications. Both birnessite and NOM are common compounds in

370

suboxic sediments.1,

2, 60

371

concentration, and accordingly decrease birnessite metal sorption and oxidizing reactivity and

372

change metal sorption behavior.20, 21, 23 Our work shows that pH and FA/birnessite mass ratios

373

strongly affect both the content and location of the accumulated Mn(III) in birnessite structure,

374

which was not revealed previously.4, 16-18 Results indicate that Mn(III) accumulates much more

375

both in layers and on vacancies with less vacancy-adsorbed Mn(II) under alkaline than acidic

376

conditions. Such detailed information is important because Mn(III) adsorbs more strongly than

377

Mn(II) and the presence of abundant Mn(III) on vacancies can strongly decrease birnessite

378

sorption capacities for metals and force the metals to adsorb on birnessite edge sites.20, 21 Thus,

379

the finding suggests that the decline of birnessite reactivity by partial reduction is more severe in

380

alkaline environment (e.g. ocean sediments) than in acidic environment (e.g., acid mine

381

drainage). On the other hand, more Mn(III) accumulated in birnessite in alkaline environment is

382

more favorable to birnessite transformation to todorokite as Mn(III) is required for the

383

transformation.22 Our work for the first time shows that high NOM/birnessite ratios lead to not

384

only Mn(III)-rich birnessite but also Mn3O4 and MnOOH phases around circumneutral pH.

385

These low-valence Mn oxides can contribute to the decline of birnessite oxidizing reactivity

386

during reduction, and the reduction-induced formation from birnessite is a new formation

387

pathway of these mineral phases in natural environment. Our study implies that Mn oxides co-

388

existing with NOM in suboxic environment, either as Mn(III,II)-rich birnessite or the low-

Partial reduction of birnessite by NOM can increase Mn(III)



Page 18 of 31

389

valence Mn oxides, have low metal adsorption and oxidizing reactivity. This work also shows

390

that the altered birnessite due to partial reduction has a high NOM adsorption capacity because

391

adsorbed Mn(II, III) makes the birnessite surface less negative and favors NOM adsorption via

392

cation bridging. Thus, birnessite not only can change quality of organic carbon but also is a

393

potent adsorbent for organic carbon.

394

Dynamic mineralogical and Mn oxidation state changes of Mn oxides during

395

microbially-mediated oxidative precipitation have been ascribed to the Mn(II)-birnessite redox

396

reactions.15, 61, 62 Birnessite partial reduction by FA leads to similar dynamic changes, which can

397

also be well explained by the redox reactions. Such reaction mechanisms depend less on the type

398

of the reductant, thus the findings in this study are likely applicable to reduction of birnessite by

399

other reducing substances, including both organics and inorganics, although the extent and rate

400

of the changes during reduction may depend on the type of reductants. For instance, NOM

401

occurring in sediments are more reducing than surface-environment extracted FA used in this

402

work,60, 63 and thus faster reduction and phase transformation of birnessite in sediments may

403

occur.

404

ASSOCIATED CONTENT

405

Supporting Information.

406

Synthesis of δ-MnO2, acidification experiments, Mn XANES spectra and LCF fits,

407

organic matter characterization, and some of the XRD and PDF data are available free of charge

408

at http://pubs.acs.org.

409

AUTHOR INFORMATION

410

Corresponding Author

411

*M. Zhu. E-mail: [email protected]; Phone: 307-766-5523


Page 19 of 31

412 413 414


Notes The authors declare no competing financial interest. ACKNOWLEDGMENTS

415

This work was supported by the U.S. National Science Foundation under Grant

416

EAR-1529937. We acknowledge beamline scientists Olaf J. Borkiewicz, Kevin A. Beyer and

417

Karena W. Chapman at beamline 11-ID-B and John Katsoudas and Carlo Segre at beamline 10-

418

BM-B at APS for their technical assistance in data collection. This research used resources of the

419

Advanced Photon Source, a U.S. Department of Energy (DOE) Office of Science User Facility

420

operated for the DOE Office of Science by Argonne National Laboratory under Contract No.

421

DE-AC02-06CH11357. We thank four anonymous reviewers and Associate Editor T. David

422

Waite for constructive comments and suggestions.



Page 20 of 31

423

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424 425 426 427 428 429 430 431 432 433 434 435 436 437 438 439 440 441 442 443 444 445 446 447 448 449 450 451 452 453 454 455 456 457 458 459 460 461 462 463 464 465 466 467

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Spectroscopic investigation of interfacial interaction of manganese oxide with triclosan, aniline, and phenol. Environ. Sci. Technol. 2016, 50 (20), 10978-10987. (19) Jokic, A.; Frenkel, A. I.; Vairavamurthy, M. A.; Huang, P. M. Birnessite catalysis of the Maillard reaction: Its significance in natural humification. Geophys. Res. Lett. 2001, 28 (20), 3899-3902. (20) Simanova, A. A.; Kwon, K. D.; Bone, S. E.; Bargar, J. R.; Refson, K.; Sposito, G.; Peña, J. Probing the sorption reactivity of the edge surfaces in birnessite nanoparticles using nickel(II). Geochim. Cosmochim. Acta 2015, 164, 191-204. (21) Hinkle, M. A. G.; Dye, K. G.; Catalano, J. G. Impact of Mn(II)-manganese oxide reactions on Ni and Zn speciation. Environ. Sci. Technol. 2017, 51 (6), 3187-3196. (22) Feng, X. H.; Zhu, M.; Ginder-Vogel, M.; Ni, C.; Parikh, S. J.; Sparks, D. L. Formation of nano-crystalline todorokite from biogenic Mn oxides. Geochim. Cosmochim. Acta 2010, 74 (11), 3232-3245. (23) Lafferty, B. J.; Ginder-Vogel, M.; Zhu, M.; Livi, K. J. T.; Sparks, D. L. Arsenite oxidation by a poorly crystalline manganese-oxide. 2. Results from X-ray absorption spectroscopy and X-ray diffraction. Environ. Sci. Technol. 2010, 44, 8467-8472. (24) Waite, T. D.; Wrigley, I. C.; Szymczak, R. Photoassisted dissolution of a colloidal manganese oxide in the presence of fulvic acid. Environ. Sci. Technol. 1988, 22 (7), 778-785. (25) Hinkle, M. A. G.; Flynn, E. D.; Catalano, J. G. Structural response of phyllomanganates to wet aging and aqueous Mn(II). Geochim. Cosmochim. Acta 2016, 192, 220-234. (26) 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. Environ. Sci. Technol. 2016, 50, 1750-1758. (27) 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 (7), 4310-4316. (28) Elzinga, E. J. 54Mn radiotracers demonstrate continuous dissolution and reprecipitation of vernadite (δ-MnO2) during interaction with aqueous Mn (II). Environ. Sci. Technol. 2016, 50, 8670-8677. (29) 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 (18), 1036410371. (30) 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 (21), 4393-4408. (31) Tu, S.; Racz, G. J.; Goh, T. B. Transformations of synthetic birnessite as affected by pH and manganese concentration. Clays Clay Miner. 1994, 42 (3), 321-330. (32) Nordstrom, D. K.; Alpers, C. N.; Ptacek, C. J.; Blowes, D. W. Negative pH and extremely acidic mine waters from Iron Mountain, California. Environ. Sci. Technol. 2000, 34 (2), 254-258. (33) Sánchez-Andrea, I.; Knittel, K.; Amann, R.; Amils, R.; Sanz, J. L. Quantification of Tinto River sediment microbial communities: importance of sulfate-reducing bacteria and their role in attenuating acid mine drainage. Appl. Environ. Microbiol. 2012, 78 (13), 4638-4645. (34) 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 (18), 10867-10876. 21 ACS Paragon Plus Environment


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(54) Klewicki, J. K.; Morgan, J. J. Kinetic behavior of Mn(III) complexes of pyrophosphate, EDTA, and citrate. Environ. Sci. Technol. 1998, 32 (19), 2916-2922. (55) Duckworth, O. W.; Sposito, G. Siderophore-manganese(III) interactions. I. Air-oxidation of manganese(ll) promoted by desferrioxamine B. Environ. Sci. Technol. 2005, 39 (16), 6037-44. (56) Ilton, E. S.; Post, J. E.; Heaney, P. J.; Ling, F. T.; Kerisit, S. N. XPS determination of Mn oxidation states in Mn (hydr)oxides. Appl. Surf. Sci. 2016, 366, 475-485. (57) Yin, H.; Tan, W.; Zheng, L.; Cui, H.; Qiu, G.; Liu, F.; Feng, X. Characterization of Nirich hexagonal birnessite and its geochemical effects on aqueous Pb2+/Zn2+ and As(III). Geochim. Cosmochim. Acta 2012, 93 (0), 47-62. (58) Lefkowitz, J. P.; Elzinga, E. J. Impacts of aqueous Mn(II) on the sorption of Zn(II) by hexagonal birnessite. Environ. Sci. Technol. 2015, 49 (8), 4886-4893. (59) Colón, W. C. E. In situ pE-pH analysis on the oxidation kinetics of manganese bearing solution. M. S. Thesis, The Pennsylvania State University, University Park, PA, 2014. (60) Canfield, D. E.; Thamdrup, B.; Hansen, J. W. The anaerobic degradation of organic matter in Danish coastal sediments: Iron reduction, manganese reduction, and sulfate reduction. Geochim. Cosmochim. Acta 1993, 57 (16), 3867-3883. (61) 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. Nat. Acad. Sci. U. S. A. 2005, 102, 5558-5563. (62) 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 (1), 143-154. (63) Segarra, K. E. A.; Comerford, C.; Slaughter, J.; Joye, S. B. Impact of electron acceptor availability on the anaerobic oxidation of methane in coastal freshwater and brackish wetland sediments. Geochim. Cosmochim. Acta 2013, 115, 15-30.

585



Page 24 of 31

586 6 (A)

pH 4

4

R = 0.1

2 0

(mM)

dis [Mn]

R=1

0.5

R = 0.1 R = 0.01

0.0 (C)

0.1 Log scale

pH 6

(B)

1.0

R = 10

pH 8 R=1

0.01 1E-3

R = 0.1 R = 0.01

1E-4 1E-5 0

587 588 589 590 591 592 593 594 595 596

100

200

300

600

Time (h)

Figure 1. Dissolved Mn concentration ([Mn]dis) during reduction of δ-MnO2 by fulvic acid (FA) at pH 4, 6 and 8 and FA/solid mass ratios (R) of 0.01 - 10. (A) pH 4, (B) pH 6, and (C) pH 8. Note that the logarithmic scale is used for [Mn]dis in panel C. The error bars result from repeated concentration measurements using ICP-OES.


Page 25 of 31


150

(A)

(C) R=1

140

R=1

3.6

pH 8

pH 8

3.5 pH 6

3.4 pH 6

SUVA254

DOC (mg/L)

130

120 (B)

150

R = 0.1

3.3 4.0

(D) R = 0.1

pH 8

pH 8

3.2 pH 6

100 pH 6

2.4

pH 4

pH 4

50

1.6 0

597 598 599 600 601 602

100

200

300

600

0

100

200

300

600

Time (h)

Time (h)

Figure 2. Dissolved organic carbon (DOC) (A, B) and SUVA254 (C, D) of fulvic acid (FA) remained in the solution during reduction of δ-MnO2 by FA at pH 4, 6 and 8 and FA/solid mass ratios (R) of 0.1 and 1. The error bars for DOC and SUVA254 result from repeated concentration measurements using a TOC analyzer and UV-vis spectroscopy.



(A)

(B)

pH 4, R = 0.1

Page 26 of 31

(C)

pH 6, R = 0.1

pH 6, R = 1

(11, 20) (31, 20)

H

(002) (22, 40) 600 h

600 h

600 h

120 h

240 h

72 h 24 h

24 h

1h

5h

120 h G

24 h 10 h

Intensity (a. u.)

1h 0.02 h 0h

(D)

(E)

pH 8, R = 0.1

1h

0.02 h

0.02 h

0h

0h

pH 8, R = 1

(F) H

600 h 600 h

5h

5h

0.02 h

0h

605 606 607 608 609 610 611

0.02 h

1h

0.02 h

d spacing (Å)

24 h

24 h

24 h

603 604

H G

120 h

120 h

2

FG

240 h F

360 h

86 4

pH 8, R = 10

H

H

0h

0h

86 4

2

d spacing (Å)

86 4

2

d spacing (Å)

Figure 3. Synchrotron-based X-ray diffraction (XRD) patterns of δ-MnO2 reacted with fulvic acid (FA) at pH 4, 6 and 8 and FA/solid mass ratios (R) of 0.1, 1 and 10. For each condition, XRD patterns are normalized by adjusting (11, 20) peak height and superimposed with the patterns of the 0-h sample. The arrow indicates the position of the dip at ~ 1.97 Å. The labels of H, G and F represent hausmannite, groutite and feitknechtite, respectively.


Page 27 of 31


(A)

(Mn-O)

2.857 Å

(Mn-Mn)L

pH 4, R = 0.1 (MnL-MnIL)

5.318 Å

600 h 24 h 1h 0.02 h 0h

(B)

(Mn-O)

2.858 Å

(Mn-Mn)L

pH 6, R = 0.1

G(r) (Å ) (a.u.)

(MnL-MnIL)

5.308 Å

600 h

-2

240 h 72 h 24 h 5h 0.5 h 0.02 h 0h

(C)

(Mn-O)

2.876 Å

(Mn-Mn)L

pH 8, R = 0.1 (MnL-MnIL) 5.272 Å 600 h 240 h 72 h 24 h 5h 1h 0.02 h 0h

1.6

612 613 614 615 616 617 618

2.4

3.2

4.0

4.8

5.6

r (Å)

Figure 4. X-ray atomic pair distribution functions within 6 Å for δ-MnO2 reacted with fulvic acid (FA) at FA/solid mass ratio (R) of 0.1 and pH 4 (A), 6 (B) or 8 (C). The PDF data are superimposed with the 0-h sample at each pH (black). The highlighted peaks show clear shifts with increasing reaction time under each condition, and the number is the peak position of the 600-h sample. The data for other conditions are provided in Figure S8.



2.885

0.30

(A)

(B) pH 8, R = 1

pH 8, R = 1

2.880

0.25 pH 6, R = 1

pH 8, R = 0.1

2.875

Area

Distance (Å)

Page 28 of 31

2.870

0.20 pH 4, R = 0.1

0.15

pH 6, R = 0.1

0.10

pH 8, R = 0.1

pH 6, R = 1

2.865 pH 6, R = 0.1

2.860

pH 4, R = 0.1

0

619 620 621 622 623 624 625

50

0.05 100

150

200

250 600

0

50

Time (h)

100

150

200

250 600

Time (h)

Figure 5. The nearest (Mn-Mn)L interatomic distances (A) and the area of the MnL-MnIL peak located at ~ 5.27 Å (B) derived from the atomic pair distribution functions for δ-MnO2 reacted with fulvic acid (FA) at 4, 6 or 8 and FA/solid mass ratios (R) of 0.1 or 1.


Page 29 of 31


(A)

80

Mn(II) Mn(III) Mn(IV)

60

Red: pH 4 Black: pH 6

40

80 60

Mn(II) Red: R = 0.1 Mn(III) Black: R = 1 Mn(IV)

40

0

0 0

627 628 629

pH 8, R = 0.1 and 1

20

20

626

(B)

100

pH 4 and 6, R = 0.1

Mn Fractions (%)

Mn Fractions (%)

100

100

200

300

400 600

0

100

200

300

400 600

Time (h)

Time (h)

Figure 6. Molar fractions of Mn(II), Mn(III) and Mn(IV) during δ-MnO2 reduction by fulvic acid at pH 4, 6 or 8 and FA/solid mass ratios (R) of 1 and 0.1, determined from Mn K-edge XANES linear combination fitting analysis. (A) R = 0.1, pH 4 and 6. (B) pH 8, R = 1 and 0.1.



Page 30 of 31

Electron transfer (mM)

18 16

pH 4, R = 0.1

14 12 10

pH 6, R = 0.1

8 6

pH 8, R = 0.1

4 pH 8, R = 1

2 0 0

630 631 632 633

100

200

300

400 600

Time (h)

Figure 7. The amount of net electron transfer from 0.3 g/L fulvic acid (FA) to 3 g/L (R = 0.1) or 0.3 g/L (R = 1) δ-MnO2 during reduction of δ-MnO2 by FA at pH 4, 6 or 8.


Page 31 of 31

634


TOC Art

635

636


Structural Transformation of Birnessite by Fulvic Acid under Anoxic

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