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Understanding the Roles of Dissolution and Diffusion in Cr(OH)3 Oxidation by -MnO2 Chao Pan, Huan Liu, Jeffrey G Catalano, Zimeng Wang, Ao Qian, and Daniel E. Giammar ACS Earth Space Chem., Just Accepted Manuscript • DOI: 10.1021/ acsearthspacechem.8b00129 • Publication Date (Web): 30 Jan 2019 Downloaded from http://pubs.acs.org on January 31, 2019

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ACS Earth and Space Chemistry

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Understanding the Roles of Dissolution and Diffusion in Cr(OH)3 Oxidation by 𝛿-MnO2

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Chao Pan†∥*, Huan Liu‡, Jeffrey G. Catalano§, Zimeng Wang⊥, Ao Qian∥, Daniel E. Giammar∥

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† Glenn

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Laboratory, Livermore, CA 94550, United States

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‡ State

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§ Department

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T. Seaborg Institute, Physical & Life Sciences, Lawrence Livermore National

Key Lab for Mineral Deposits Research, Nanjing University, Nanjing, P.R. China of Earth and Planetary Sciences, Washington University in St. Louis, St. Louis,

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Missouri, 63130 United States

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⊥

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China

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∥

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Geosciences, Wuhan, P. R. China

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∥ Department

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Louis, St. Louis, Missouri, 63130 United States

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*E-mail: [email protected]; Mail: Glenn T. Seaborg Institute, Physical & Life Sciences, Lawrence

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Livermore National Laboratory, 7000 East Avenue, Livermore, CA 94550, United States. Tel:

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(314) 608-8987

Department of Environmental Science and Engineering, Fudan University, Shanghai, 200433,

State Key Laboratory of Biogeology and Environmental Geology, China University of of Energy, Environmental and Chemical Engineering, Washington University in St.

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Submitted to

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September 2018

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ACS Paragon Plus Environment

ACS Earth and Space Chemistry 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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Abstract Manganese oxides are the major oxidants of Cr(III) to Cr(VI) in natural environments.

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This study evaluated the rate and extent of oxidation of Cr(III) released from Cr(OH)3 by 𝛿-

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MnO2 from pH 5 to 9 with a particular focus on quantifying the rate constant for Cr(III)

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oxidation on the MnO2 surface. The Cr(III) oxidation rate was initially fast, but it then slowed

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and ceased for pH 5 to pH 7, which agrees with previously reported inhibition of the redox

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reaction above pH 4 by precipitation of Cr(III) on the MnO2 surface. Above pH 7, overall Cr(VI)

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production was higher than at lower pH even though the dissolved Cr(III) concentration in

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equilibrium with Cr(OH)3 was lower. This is probably due to the reoxidation of aqueous Mn(II)

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by dissolved oxygen at higher pH, which made more manganese oxide available to oxidize

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Cr(III) to Cr(VI). Multichamber experiments were used to assess the role of solid-solid proximity

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in Cr(OH)3-MnO2 interactions at pH 5. The different rates and extents of Cr(VI) production in

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the multichamber reactor and completely mixed batch reactor indicate that mixing of Cr(OH)3

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and MnO2 solids plays a role in the rate and extent of Cr(VI) generation. Further Cr(VI)

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generation in this systems strongly depends on pH. Because of the importance of mass transfer of

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Cr(III) to MnO2 solids for the overall Cr(III) oxidation process, the net rate of Cr(VI) generation

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in natural environments with Cr(III)-containing solids and MnO2 can be orders of magnitude

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slower than observed in well-mixed laboratory-scale batch experiments.

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Keywords:

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production, oxidation inhibition

multichamber experiments, manganese oxides, kinetic model, Cr(VI)

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Introduction Chromium (Cr) is widespread in soils, sediments and water from natural and

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anthropogenic sources. The predominant oxidation states of Cr in aquatic and soil environments

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are Cr(III) and Cr(VI), with Cr(III) being less mobile and less toxic.1 The chromium content of

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natural solids varies widely with the type and nature of rocks or sediments, and the highest

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chromium contents are usually associated with the finest particles in soils and sediments2 where

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the Cr is generally present in the +III oxidation state.3, 4 In natural waters, the range of chromium

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concentrations is large and dissolved chromium concentrations have been observed as high as 4

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µM, around twice the U.S. drinking water standard of 100 µg/L (1.92 µM). These high

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chromium concentrations in natural waters are mostly Cr(VI) species, which are more soluble

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than Cr(III) species.2

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Manganese oxides, which are ubiquitous in soils, can rapidly oxidize Cr(III) to Cr(VI).5-7

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They provide the major geochemical pathway for Cr(VI) production from Cr(III) in

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groundwater, soils, and subseafloor environments.8, 9 Manganese oxides are believed to form

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primarily by aqueous Mn(II) oxidation via either direct or indirect microbial activity.10 The

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predominant type of biogenic manganese oxides formed at circumneutral pH are highly

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disordered and nanocrystalline phases, similar to hexagonal birnessite (its synthetic analogue is

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𝛿-MnO2).10-14 Several studies have investigated the kinetics of Cr(III) oxidation by various

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manganese oxides.15-19 X-ray absorption spectroscopy (XAS) of the products of this reaction

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indicated that dissolved Cr(III) species first diffuse towards Mn(IV) vacancies in the sheet of

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MnO6 octahedra;20 the coupled Cr(III) oxidation/ Mn(IV) reduction occurs through one-electron-

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transfer reactions according to X-ray photoelectron spectroscopy (XPS),21 and the Cr(VI)

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produced is then released into solution. However, a rapid rate decline and cessation of the

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reaction followed the initially fast Cr(VI) generation from Cr(OH)3 oxidation by manganese

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oxides, and this decline is thought to be due to the formation of a Cr(OH)3 precipitate on 𝛿-

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MnO2.22, 23 The initial rates of Cr(III) oxidation on the MnO2 surface depend on pH, temperature,

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and Cr(III) and manganese concentrations.22, 24-27 Most studies have focused on the Cr(III)

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oxidation kinetics at low pH, and in most of these cases only dissolved Cr(VI) was measured

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without quantification of the total Cr(VI) generated.7, 28

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The objectives of this study were to investigate the rates and mechanism of Cr(VI)

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production from Cr(OH)3 oxidation by 𝛿-MnO2 as a function of pH and to identify the final Mn-

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containing products. Total Cr(VI) concentrations, which include adsorbed Cr(VI) on solid

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surfaces and dissolved Cr(VI) in solution, were measured to provide assessments of the total

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Cr(III) oxidation. Our previous work has investigated the role of solid-solid interactions in

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CrxFe1-x(OH)3/MnO2 system. The multichamber experiments were operated here to test the

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effects of mixing of Cr(OH)3 and MnO2 solids on the rate and extent of the Cr(VI) production.

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Compared with the work of CrxFe1-x(OH)3/MnO2, Cr(OH)3 has much higher potential for Cr(III)

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release and this facilitated the development of a model of Cr(VI) generation that account for the

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rates of the relevant chemical reactions as well as the rates of solute transport. Multichamber

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experiments were operated to test the effect of mixing of Cr(OH)3 and MnO2 solids on the rate

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and extent of Cr(VI) production. The reaction products were characterized by transmission

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electron microscopy (TEM), X-ray diffraction (XRD) and X-ray photoelectron spectroscopy

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(XPS). The reaction rate constant, fate and transport of Cr(III), as well as the reaction products

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from Cr(OH)3, also give us insight into Cr mobility from other Cr(III)-containing phases relevant

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to natural environments.

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Materials and Methods

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Materials

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Ultrapure water (resistivity > 18.2 MΩ-cm) was used for the experiments, and the

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chemicals were analytical reagents of high purity. At pH 5 and 6, no buffer was added to the

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reactors and the pH was maintained by NaOH or HCl addition within the range of ± 0.2. At pH 7

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and 8, the pH of the suspensions was buffered by 5 mM 3-(N-morpholino) propanesulfonic acid

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(MOPS, pKa=7.2). At pH 9, 5 mM N-cyclohexyl-2-aminoethanesulfonic acid (CHES, pKa=9.3)

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was used. The pH buffers and their concentrations were chosen because of their minimal

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formation of complexes with Cr(III) and their stability against oxidation by MnO2.29-31 NaCl was

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added to provide 5 mM ionic strength as the background electrolyte because Na+ and Cl- do not

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interfere with the chemistry of Cr(III) oxidation.

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Mineral synthesis

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Synthetic δ-MnO2 was prepared by reacting KMnO4 with MnCl2 at a basic pH following

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the method described by Villalobos et al.14 XRD confirmed that the solid was δ-MnO2, and TEM

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provided evidence of δ-MnO2 morphology and sizes as we have presented previously.15

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Cr(OH)3(s) was synthesized by titrating CrCl3 solutions with NaOH solution to pH 7 and

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maintaining the pH for 24 hours.15, 32 The suspension was then washed five times with ultrapure

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water, and the supernatant was discarded after centrifugation. The final chromium concentration

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of the Cr(OH)3(s) suspension was measured by inductively coupled plasma-mass spectroscopy

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(ICP-MS, PerkinElmer ELAN DRC II) after nitric acid digestion. Cr(OH)3 exists partially as the

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crystalline solid (Cr(OH)3·3H2O) in suspension and it partially converted to amorphous Cr(OH)3

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upon drying.33 The broad humps in the XRD pattern indicate the formation of some amorphous

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Cr(OH)3 (Figure 4). The solubility of crystalline Cr(OH)3 is actually greater than that of the

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amorphous form (Table 1).34 The Cr(OH)3 stock suspension was sonicated for 5 minutes before

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use in experiments to disperse the particles before they were added to the reactors.

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Mixed batch experiments and multichamber reactor

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Completely mixed batch experiments were conducted in glass beakers filled with

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ultrapure water, 5 mM NaCl, a pH buffer, MnO2 suspension, and Cr(OH)3 suspensions with 40

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mg/L (770 µM) initial Cr(III) concentration to a total volume of 1 L. Multichamber experiments

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were used to assess the role of mixing and solid-solid contact on Cr(OH)3-MnO2 interactions.

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The multichamber reactor was the same as described in our group’s previous work (Figure S1).15,

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29

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particle with diameter of 2-3 nm) divided the reactor into two 110-mL chambers, eliminating the

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direct contact of the Cr(OH)3 and MnO2 solids but allowing dissolved species to diffuse across

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the membrane. The suspensions were completely mixed by magnetic stirring of each chamber.

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For both completely mixed batch experiments and multichamber experiments, samples were

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periodically collected and a portion of them were filtered with 0.05 μm polyethersulfone (PES)

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syringe filters (Tisch Environmental, OH) for dissolved chromium, dissolved Cr(VI), and

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dissolved manganese analysis. The remaining portions of the unfiltered samples were used for

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total Cr(VI) and total Mn(II) analysis. Both completely mixed batch experiments and

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multichamber experiments were performed under the ambient laboratory atmosphere.

Briefly, a dialysis membrane with a molecular weight cut off (MWCO) of 3500 (roughly a

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Aqueous and solid phase analysis

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Total and dissolved Cr and Mn concentrations were measured by ICP-MS (PerkinElmer

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ELAN DRC II). The instrument detection limits for Cr and Mn were 0.2 μg/L (0.0039 μM) and

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0.5 μg/L (0.009 μM), respectively. The samples for measuring dissolved Cr(VI) and dissolved

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Mn(II) were filtered and then measured by the diphenylcarbazide method and ICP-MS,

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respectively. Dissolved Mn(II) concentrations were assumed to equal the total dissolved Mn

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concentration because both Mn(IV) and Mn(III) are sparingly soluble even in the presence of

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buffers.35 Cr(VI) concentrations in the samples were determined spectrophotometrically

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(PerkinElmer-Lambda XLS) after reacting with diphenylcarbazide.36 The detection limit for

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Cr(VI) by this method was 5 μg/L (0.096 μM). Total Cr(VI) and total Mn(II), which included

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both adsorbed and dissolved species, were measured the same way as dissolved Cr(VI) and

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Mn(II) after extracting the surface-associated species into the solution. For total Cr(VI)

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concentration, adsorbed Cr(VI) was extracted by providing a 10 mM phosphate concentration in

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the suspension to displace adsorbed Cr(VI); the Cr(VI) was then measured by the

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diphenylcarbzide method. The efficiency of this extraction method was 90% ± 5% based on

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control experiments in which Cr(VI) adsorption onto Cr(III) hydroxide solids was followed by

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phosphate addition to induce Cr(VI) desorption. Another portion of the sample suspension was

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treated with 10 mM CuSO4 to extract adsorbed Mn(II). In this method Cu(II) preferentially

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adsorbs to MnO2 and induces Mn(II) desorption.15, 29 This method was shown to remove more

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than 90% of the adsorbed Mn(II) from a biologically reduced 𝛿-MnO2.37 The method only

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extracts adsorbed Mn(II) and would not mobilize any structurally incorporated Mn(II) or

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Mn(III).

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TEM samples were prepared by dropping approximately 30 μL of suspension onto 200

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mesh carbon-coated copper grids (Ted Pella, Inc.) followed by immediate evaporation of the

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remaining water at room temperature under vacuum. TEM micrographs were taken with a

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transmission electron microscope under 120 kV (FEI Spirit G2). The solid samples for XRD,

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XPS and TEM were prepared by centrifugation and freeze-drying. XRD patterns were collected

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using Cu Kα radiation (Bruker d8 Advance X-ray diffractometer with a LynxEye XE strip

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detector). XPS analyses were conducted using a Physical Electronics 5000 Versa Probe II

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Scanning ESCA Microprobe with an Al Kα X-ray source at 23.5 eV pass energy at a 100 μm X-

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ray spot size. The binding energy was calibrated using C 1s at 284.6 eV, and the XPS spectra

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were processed by using CasaXPS software (Version 2.3.15)38 with the Gaussian-Lorentzian

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function (50% G-50% L), and a Shirley background for peak fitting. The quantification of Mn

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valence state was made following a method in which the Mn 2p3/2 spectrum is divided into five

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multiplet peaks (total of 15 binding energies) of Mn(IV), Mn(III) and Mn(II).39 A value of 1.5

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for the full width of the peak at half the maximum peak height (fwhm) was assigned to fit the

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Mn2p3/2 spectrum for all of the multiplet binding energy spectra.40 In this study, the standard

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deviation of the compositional data calculated by CasaXPS is mostly around 5%, and no more

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than 10%. TEM observation was carried out using an FEI TF electron microscope operated at

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200 kV and energy dispersive X-ray (EDX) analysis was used to confirm the particle

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compositions. Selected area electron diffraction (SAED) patterns were collected to identify the

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Mn and Cr mineral phases present.

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Model for dynamics of Cr(VI) production

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A quantitative model for the dynamics of Cr(OH)3 oxidation by 𝛿-MnO2 in multichamber

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experiments was developed based on the dissolution rate of Cr(OH)3(s), the rates of aqueous

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Cr(III) and Cr(VI) transport across the membrane, the rate of dissolved Cr(III) oxidation by 𝛿-

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MnO2, and equilibrium Cr(VI) adsorption on Cr(OH)3 solids. The rate of aqueous Cr(III)

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oxidation on the surface of MnO2 was calculated by determining the parameters that provided the

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best fit of the model output to the experimental data.

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For the Cr(OH)3||MnO2 system, where “||” notes the separation of solids by a dialysis membrane,

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the governing equations are

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V

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V

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V

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V

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[MnO2] = [MnO2]0 ―1.5 ∙ [Cr(VI)]Mn ―1.5 ∙ [Cr(VI)tot]Cr

d[Cr(III)]Mn dt

d[Cr(III)]Cr dt

d[Cr(VI)]Mn dt

= ―V ∙ k ∙ [Cr(III)]Mn ∙ [MnO2] + νCr(III) ∙ A ∙ ([Cr(III)]Cr ― [Cr(III)]Mn) [Cr(III)]Cr

= V ∙ k′ ∙ (1 ― [Cr(III)]eq) ― νCr(III) ∙ A ∙ ([Cr(III)]Cr ― [Cr(III)]Mn)

(2)

= V ∙ k ∙ [Cr(III)]Mn ∙ [MnO2] ― νCr(VI) ∙ A ∙ ([Cr(VI)]Mn ― [Cr(VI)diss]Cr)

(3)

d[Cr(VI)tot]Cr dt

(1)

= νCr(VI) ∙ A ∙ ([Cr(VI)]Mn ― [Cr(VI)diss]Cr)

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(4) (5)

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V (100 mL or 10-4 m3) is the volume of solution in each chamber, A (20 cm2 or 0.002 m2) is the

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interfacial area of the dialysis membrane, and ν is the transmembrane mass transfer coefficient of

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dissolved chromium. [Cr(III)]Mn is the concentration of dissolved Cr(III) in the MnO2 chamber,

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while [Cr(III)]Cr is the concentration of dissolved Cr(III) in the Cr(OH)3 chamber. Similarly,

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Cr(VI)Mn is the concentration of Cr(VI) in the MnO2 chamber. We show that no Cr(VI) adsorbs

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onto MnO2 solids (Figure S9), which is expected since soluble Cr(VI) and the MnO2 surface are

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both negatively charged over the conditions studied. [Cr(VI)tot]Cr is the concentration of total

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Cr(VI) including both dissolved and adsorbed Cr(VI) in the Cr(OH)3 chamber and [Cr(VI)diss]Cr

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is the dissolved Cr(VI) in the Cr(OH)3 chamber.

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An overall strategy of determining the different parameters in the model used control

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experiments to isolate specific processes to first determine particular parameters, and it gradually

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added complexity as additional parameters were determined (Figure S5). The transmembrane

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mass transfer coefficient of Cr(VI) (νCr(VI)) was estimated from the Cr(VI) tracer experiment

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(Figure S6 in Supporting Information), and the aqueous Cr(III) transmembrane mass transfer

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coefficient (νCr(III)) was calculated based on a semi-empirical equation relating the molecular

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weight and mass transfer coefficients of Cr(VI) and Cr(III) species (eq S3 in Supporting

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Information). The flux (mol m-2 s-1) of an aqueous species is proportional to its concentration

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gradient across the membrane. The relationship between [Cr(VI)diss]Cr and [Cr(VI)tot]Cr was

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estimated by adsorption experiments of Cr(VI) onto Cr(OH)3 solids (eq S12, S13 and Figure S8

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in Supporting Information). [Cr(VI)]Mn is used to represent both dissolved and total Cr(VI)

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concentrations in the Mn chamber; for the Mn chamber the dissolved and total Cr(VI)

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concentrations are identical because no Cr(VI) adsorbed on MnO2 solids (Figure S9). The rate

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constant for Cr(OH)3 dissolution k’ was obtained from fitting the data from Cr(OH)3||water

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experiments (Figure S7 in Supporting Information).

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Once the other parameters in Equations 1-5 had been determined from the control

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experiments, the rate constant of dissolved Cr(III) oxidation by MnO2 (k) was determined. The

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value of k (3.9 L mol-1 s-1) was estimated by finding the value that provided the optimal fit of the

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experimentally measured [Cr(VI)]Mn, [Cr(VI)tot]Cr and [Cr(VI)diss]Cr concentrations to the

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modeling output from Equations 1-5. The equations in the model were solved by the ode45

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solver in Matlab 7.0. More detailed model derivation can be found in the Supporting

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Information. 9 / 26



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Results and discussion

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Cr(OH)3 oxidation by MnO2

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From pH 5 to 7, Cr(III) oxidation rates were initially rapid in completely mixed

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suspensions. However, the rapid oxidation was followed by a cessation of the reaction (Figure

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2). The Cr(VI) concentration was nearly constant after increasing to around 100 M within the

227

first ten hours. At pH 5 the ratio of Cr(VI) to Mn(II) produced was around 1.5. This ratio

228

corresponds to the stoichiometry of equation 6 with aqueous Mn(II) as the product of MnO2

229

reduction. The Mn(II) generated was lower than the total amount of Mn(II) that could have been

230

generated if the reaction had gone to completion. The Cr(VI) produced was also lower than the

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initial dissolved Cr(III) concentration of 240 M, which was measured before MnO2 was added

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to the Cr(OH)3 suspensions. These observations indicate that Cr(III) oxidation had stopped even

233

though there was still sufficient MnO2 and Cr(III) for further oxidation. This inhibitory effect of

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Cr(OH)3 oxidation by Mn-oxides has also been previously observed at pH values greater than 4

235

and with higher Cr(III) loadings than those in our study.22, 23, 25 Those previous studies attributed

236

the inhibition to the formation of a Cr(OH)3 precipitate on δ-MnO2.23 The presence of δ-MnO2

237

could have helped nucleated the amorphous Cr(OH)3 following the initial dissolution of the

238

crystalline phase (Figure 4). As the crystalline Cr(OH)3 solid has a higher solubility than

239

amorphous one, such a process would allow release of Cr(III) from crystalline Cr(OH)3 solid and

240

then preciptiation in a freshly formed amorphous Cr(OH)3 solid that only formed on the MnO2.

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As a result, the amount of Cr(III) oxidized may be governed by the surface area of MnO2 when

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Cr(III) is in excess.

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2Cr(OH)3(s) + 3MnO2 + 2H+ = 2CrO42- + 3Mn2+ + 4H2O

(6)

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Cr(OH)3(s) + 3MnO2 + H2O = CrO42- + 3MnOOH

(7)

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From pH 8 to 9 when Cr(OH)3 solubility is much lower than at pH 5-7 (Table 1), the

246

Cr(VI) concentrations increased more slowly, but they ultimately increased to higher levels than

247

at pH 5 to 7 without any observable inhibition. Overall, Mn(III) is the dominant reaction product

248

of MnO2 reduction with 120 μM Cr(VI) generated and almost negligible Mn(II) at pH 8 (Figure

249

2 and eq 7). The reaction proceeded until MnO2 was limiting according to the stoichiometry of

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250

the reaction. At pH 9, Cr(VI) can be as high as 200 M, which would require 600 M MnO2

251

with Mn(III) as the dominant product. As there was only 436 M MnO2 in the initial suspension,

252

aqueous Mn(II) might have been oxidized by dissolved oxygen to MnO2,41 which could continue

253

oxidizing Cr(III) at this high pH.32 Mn(II) oxidation happened within hours at high pH, consistent with

254

the Cr(III) oxidation rates in Figure 1. Considering the fast manganese oxide formation rates at high

255

pH, the Cr(VI) generation will be lower in anaerobic conditions than with the presence of oxygen

256

above pH 7 .

257

To better understand the inhibitory mechanism of Cr(OH)3 oxidation by manganese

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oxides at low pH, different amounts of MnO2 were dosed to Cr(OH)3 suspension at pH 5. With

259

different amounts of MnO2 initial concentration, Cr(VI) production rates were always high in the

260

initial stage followed by the cessation of the reaction (Figure 3a). When the reaction ceased, both

261

Cr(III) and MnO2 were still present at amounts that theoretically could react with one another.

262

Even though neither reactant was completely consumed, with higher MnO2 doses, more Cr(VI)

263

was generated. Total Cr(VI) concentrations achieved were proportional to the MnO2 dose

264

(Figure 3b). These observations indicate that the MnO2 concentration was the limiting factor for

265

Cr(III) oxidation. Manceau and Charlet proposed that mononuclear Cr(III) diffuses to the lattice

266

vacancies in the Mn-oxide structure and is complexed at these sites subsequent to the electron

267

transfer with Mn(IV) during Cr(III) oxidation by manganese oxides.20 As a result, the number of

268

vacancies in the MnO2 lattice structure determines the extent of Cr(III) oxidation and the amount

269

of Cr(III) oxidized is dependent on MnO2 concentration.

270 271

Kinetic modeling of Cr(III) oxidation in multichamber experiments

272

Cr(VI) occurs through chromium(III) dissolution from Cr(OH)3 and its oxidation on the

273

MnO2 surface. At pH 5 when the initial dissolved Cr(III) concentration was as high as 240 𝜇M,

274

the aqueous Cr(III) ions could quickly reach the MnO2 surface by diffusing through the

275

membrane in the multichamber reactor. The average Cr(VI) concentration in the multichamber

276

reactors (Figure 3a) is around 110 𝜇M, similar to the 95 𝜇M total Cr(VI) produced in completely

277

mixed batch experiments. At pH 8 when the Cr(III) solubility was much lower than at pH 5, very

278

little Cr(VI) was generated in the multichamber experiment while much more was produced

279

when the Cr(OH)3 and MnO2 were mixed in the same suspension. The concentrations of Cr(III) 11 / 26



280

in the Cr(OH)3–containing chamber of the multichamber reactor were too low at pH 8 to provide

281

enough flux of Cr(III) across the membrane into the other chamber where it could be oxidized by

282

the MnO2.

283

At pH 5 the rate constant for aqueous Cr(III) oxidation by MnO2 can be evaluated from

284

the multichamber experiments (Figure 3a). The chromium mass transfer coefficients (νCr(III) and

285

νCr(VI)) and the rate constant of Cr(OH)3(s) dissolution (k’) were determined independently in

286

control experiments and then included in the model for examination of Cr(VI) release in

287

multichamber experiments. For control experiments, chromium concentrations were measured in

288

Cr(VI)||water and Cr(OH)3||water multichamber experiments. The Cr(VI)||water experiment was

289

only affected by Cr(VI) diffusion across the membrane, so it was used to determine a Cr(VI)

290

mass transfer coefficient (νCr(VI)) of 2.3× 10-7 m/s, and νCr(III) was then calculated to be 3.0 × 10-7

291

m/s following a semi-empirical equation. The Cr(OH)3||water multichamber experiment had

292

Cr(III) concentrations that were only affected by dissolution and by mass transfer across the

293

membrane, so this experiment was used to estimate the dissolution rate constant of Cr(OH)3(s),

294

which was determined to be 7.5 × 10-10 mol L-1 s-1. Detailed discussions of parameter

295

determination are included in the Supporting Information. The rate constant for aqueous Cr(III)

296

oxidation by manganese oxides was then determined by finding the value that provided the best

297

fit of the model output to the measured values of the Cr(VI) concentrations in the MnO2 chamber

298

and the dissolved and total Cr(VI) concentrations in the Cr(OH)3 chamber at pH 5 (Figure 3a).

299

The rate constant of Cr(III) oxidation was then applied to the completely mixed batch

300

experiments without any terms for mass transfer between chambers since there was only one

301

chamber (equation S12-13). The model confirms that the experimentally measured observation

302

of Cr(VI) is orders of magnitude faster in the completely-mixed experiment than in the

303

multichamber experiment (Figure 3b). A close examination of the initial hours of the experiment

304

do indicate that the model underpredicts the exact Cr(VI) generation rate for the shortest reaction

305

times (Figure 3b).

306 307

Mn-Containing Products of the Reaction

308

Feitknechtite (MnOOH) was produced from reaction of Cr(OH)3 and MnO2 with more

309

generation at higher pH (Figure 4). Recent work has shown that hexagonal birnessite is subject to 12 / 26


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310

structural and mineralogical changes during reaction with aqueous Mn(II), and solution pH could

311

affect this interaction as Mn(II) adsorption depends on pH.42-45 Our work shows that Mn(II)

312

strongly adsorbed onto the solid phases at high pH (Figure S3), which could lead to bulk

313

transformation of the birnessite into feitknechtite and even a more stable manganite phase through

314

a reductive transformation process.33 The XRD patterns at pH 8 and pH 9 indicated that

315

feitknechtite was the dominant product and that no manganite formed within the reaction time in

316

this study. The feitknechtite formation is consistent with results in Figure 1 that Mn(III)-containing

317

solids were the dominant reduction products of 𝛿-MnO2 at high pH. The XRD pattern for the pH

318

9 sample (Figure 1) also has a small peak at ~12o 2θ, consistent with the (001) diffraction peak for

319

birnessites with increased layer stacking, including the triclinic form. Triclinic birnessite is a

320

known product of the reaction of δ-MnO2 with dissolved Mn(II) at alkaline pH conditions and

321

results from Mn(II) adsorption and comproportionation with structural Mn(IV), forming Mn(III)

322

in the manganese oxide sheet.43 At pH 5, MnO2 reduction by Cr(OH)3 did not introduce new

323

manganese oxides phases (Figure 4) despite generating substantial dissolved Mn(II) (Figure 2a).

324

TEM was also used to investigate the products of the Cr(OH)3-MnO2 reaction at pH 5 and pH

325

9 (Figure 5). At pH 5, the reaction products did not show any strong electron diffraction, which is

326

consistent with the results from XRD. In contrast a change of the morphology of the particles after

327

reaction at pH 9 indicated that a mineral phase transformation occurred and the formed minerals

328

were identified by their electron diffraction patterns. Somewhat surprisingly, feitknechtite was not

329

observed in TEM-SAED despite it being a dominant reaction product identified by XRD. Instead,

330

the SAED patterns for solids reacted at pH 9 are most consistent with the formation of triclinic

331

birnessite. This could be due to the limited area examined by TEM, or it could be because the

332

degree of crystallinity was higher for the feitnechtite than triclinic birnessite; because the degree

333

of crystallinity was higher for the feitknechtite than triclinic birnessite; consequently, a small

334

amount of feitknechtite could give stronger signal in the XRD patterns than a larger amount of

335

triclinic birnessite.

336

The valence of Mn on the solid surface was determined directly by analyzing the solid product

337

with Mn2p3/2 and Mn3s splitting energy intervals using XPS (Figure 6 and Figure S4). Table 2

338

summarizes the Mn oxidation state percentage in all the solids as determined by XPS 2p3/2, and

339

percentages determined by Mn3s splitting energy intervals gave similar results (Figure S4 and

340

Table S2). For the initial δ-MnO2, Mn(IV) and Mn(III) were present in the near-surface region as 13 / 26



341

94.9% and 5.1%, respectively (Table 2) and the average oxidation state of Mn is 3.95. For

342

manganese oxides after reacting with Cr(OH)3, the Mn(IV) had been reduced to Mn(II) and Mn(III)

343

on the solid surface at both pH 5 and pH 9. The solids for XPS measurements are from completely

344

mixed batch experiments of Cr(OH)3 and MnO2. At pH 9, Mn(II) produced from reduction of

345

Mn(IV) is more likely to adsorb on the surface of the solids, which can lead to phase transformation

346

of δ-MnO2. At pH 5 the XPS results show substantially more residual Mn(IV) than at pH 9. This

347

difference in solid-phase average Mn oxidation state along with the greater release of dissolved

348

Mn(II) likely reflects the pH dependence of Mn(II)-Mn(IV) comproportionation44 and is

349

responsible for the lack of mineralogical transformations. However, the XPS-derived distribution

350

of Mn species at pH 5 contains more Mn(III) and Mn(II) than is anticipated to be stable in a

351

birnessite-type solid and may either derive from a more reduced surface, given the high surface

352

sensitivity of XPS, or from a portion of this Mn(III) and Mn(II) being associated with solids, such

353

as via adsorption or coprecipitation.

354 355 356 357

Conclusions Our results demonstrate that oxidation of Cr(OH)3 by manganese oxides was highly

358

dependent on pH even in a relative narrow range of circumneutral pH. In a completely mixed

359

system, Cr(III) oxidation was inhibited after an initially rapid stage at pH 5-7. Cr(VI) production

360

was slower but ultimately reached higher concentrations at pH 8 and 9 than at pH 5-7. At pH 5

361

Cr(VI) generated in equilibrium was proportional to the amount of MnO2 added. Multichamber

362

reactors were used to test the role of mixing of Cr(OH)3 and MnO2 on Cr(III) oxidation, which

363

could simulate the presence of a contact barrier. At pH 5 when the dissolved Cr(III)

364

concentration was high, Cr(VI) concentrations in the multichamber reactor reached a similar

365

level to that in the completely mixed batch reactor within 50 hours. While at pH 8 when the

366

dissolved Cr(III) concentration is low, the Cr(VI) concentration in the multichamber reactor is

367

almost negligible (even after 400 hours) compared with that in the completely mixed batch

368

reactor. Our modeling work could successfully fit the Cr(VI) production rates at pH 5 in

369

completely mixed batch experiments based on parameters determined from multichamber

370

experiments. This work could help us predict the rate and extent of Cr(III) oxidation in natural

371

environments at different conditions. The Cr(VI) production rates were determined by the

14 / 26


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372

transport of dissolved Cr(III) in equilibrium with Cr(III)-containing solids upon reaction with δ-

373

MnO2 when there is no direct contact of the solids. Furthermore, δ-MnO2 was transformed to

374

feitknechtite and triclinic birnessite at high pH in the presence of Cr(OH)3, which would affect

375

the further reactivity of manganese oxides.

376 377

Supporting Information

378

Additional information regarding the multichamber experimental setup, TEM images of

379

Cr(III) and MnO2 reaction products, pH dependence of Mn(II) adsorption onto the solid phases

380

in Cr(OH)3-MnO2 completely mixed suspensions, XPS spectra of Mn 3s photoelectron lines for

381

the solid product and initial MnO2, binding energies of surface Mn species for fitting the Mn2p3/2

382

peak of the solid product, summary of Mn 3s splitting energy intervals of XPS for manganese

383

oxides, and derivation of multichamber model for Cr(OH)3 oxidation by MnO2 are provided in a

384

supplementary document.

385 386 387

Acknowledgements This research was supported by the U.S. National Science Foundation (CBET 1335613).

388

C.P. acknowledges financial support from school of Engineering Applied Science in Washington

389

University in St. Louis for a first year Ph.D. fellowship. We thank Dr. Wenlu Li for the help with

390

TEM analysis for solid samples. TEM and ICP-MS were provided by the Nano Research Facility

391

(NRF) at Washington University in St. Louis. XPS analysis were performed in facilities of the

392

Institute of Materials Science and Engineering at Washington University in St. Louis. Work at

393

LLNL was conducted under the auspices of the US Department of Energy at LLNL under

394

Contract DE-AC52-07NA27344.

15 / 26



395

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18. Kim, J. G.; Dixon, J. B.; Chusuei, C. C.; Deng, Y., Oxidation of chromium(III) to (VI) by manganese oxides. Soil Sci. Soc. Am. J. 2002, 66, 306-315. 19. Hausladen, D. M.; Fendorf, S., Hexavalent chromium generation within naturally structured soils and sediments. Environ. Sci. Technol. 2017, 51, 2058-2067. 20. Manceau, A.; Charlet, L., X-ray absorption spectroscopic study of the sorption of Cr(III) at the oxide-water interface: I. Molecular mechanism of Cr(III) oxidation on Mn oxides. J. Colloid Interface Sci. 1992, 148, 425-442. 21. Banerjee, D.; Nesbitt, H. W., Oxidation of aqueous Cr(III) at birnessite surfaces: Constraints on reaction mechanism. Geochim. Cosmochim. Acta 1999, 63, 1671-1687. 22. Landrot, G.; Ginder-Vogel, M.; Sparks, D. L., Kinetics of chromium(III) oxidation by manganese(IV) oxides using quick scanning X-ray absorption fine structure spectroscopy (Q-XAFS). Environ. Sci. Technol. 2009, 44, 143-149. 23. Fendorf, S. E.; Fendorf, M.; Sparks, D. L.; Gronsky, R., Inhibitory mechanisms of Cr(III) oxidation by δ-MnO2. J. Colloid Interface Sci. 1992, 153, 37-54. 24. Fendorf, S. E.; Zasoski, R. J., Chromium(III) oxidation by δ-manganese oxide (MnO2). 1. Characterization. Environ. Sci. Technol. 1992, 26, 79-85. 25. Kim, J. G.; Moon, H.-S., Oxidation of chromium(III) to chromium(VI) by a series of synthesized birnessites (σ-MnO2): Kinetics and oxidation capacity. Clay Science 1998, 10, 363-373. 26. Weaver, R. M.; Hochella, M. F., The reactivity of seven Mn-oxides with Cr3+aq: A comparative analysis of a complex, environmentally important redox reaction. Am. Mineral. 2003, 88, 2016-2027. 27. Johnson, C. A.; Xyla, A. G., The oxidation of chromium(III) to chromium(VI) on the surface of manganite (γ-MnOOH). Geochim. Cosmochim. Acta 1991, 55, 2861-2866. 28. Landrot, G.; Ginder-Vogel, M.; Livi, K.; Fitts, J. P.; Sparks, D. L., Chromium(III) oxidation by three poorly crystalline manganese(IV) oxides. 2. Solid phase analyses. Environ. Sci. Technol. 2012, 46, 11601-11609. 29. Wang, Z.; Lee, S.-W.; Kapoor, P.; Tebo, B. M.; Giammar, D. E., Uraninite oxidation and dissolution induced by manganese oxide: A redox reaction between two insoluble minerals. Geochim. Cosmochim. Acta 2013, 100, 24-40. 30. Pan, C.; Troyer, L. D.; Catalano, J. G.; Giammar, D. E., Dynamics of chromium(VI) removal from drinking water by iron electrocoagulation. Environ. Sci. Technol. 2016, 50, 13502-13510. 31. Wadhawan, A. R.; Livi, K. J.; Stone, A. T.; Bouwer, E. J., Influence of oxygenation on chromium redox reactions with manganese sulfide (MnS(s)). Environ. Sci. Technol. 2015, 49, 3523-3531. 32. Namgung, S.; Kwon, M. J.; Qafoku, N. P.; Lee, G., Cr(OH)3(s) oxidation induced by surface catalyzed Mn(II) oxidation. Environ. Sci. Technol. 2014, 48, 10760-10768. 33. Papassiopi, N.; Vaxevanidou, K.; Christou, C.; Karagianni, E.; Antipas, G. S. E., Synthesis, characterization and stability of Cr(III) and Fe(III) hydroxides. J. Hazard. Mater. 2014, 264, 490-497. 34. Rai, D.; Sass, B. M.; Moore, D. A., Chromium(III) hydrolysis constants and solubility of chromium(III) hydroxide. Inorg. Chem. 1987, 26, 345-349. 35. 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, 463470. 17 / 26



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36. U.S.EPA, SW-846, Method 7196A, Chromium, Hexavalent (Colorimetric). 37. Burdige, D. J.; Nealson, K. H., Microbial manganese reduction by enrichment cultures from coastal marine sediments. Appl. Environ. Microbiol. 1985, 50, 491-497. 38. Fairley, N.; Carrick, A., The casa cookbook. Acolyte Science Cheshire: 2005; Vol. 1. 39. Nesbitt, H.; Banerjee, D., Interpretation of XPS Mn(2p) spectra of Mn oxyhydroxides and constraints on the mechanism of MnO2 precipitation. Am. Mineral. 1998, 83, 305-315. 40. Lee, G.; Song, K.; Bae, J., Permanganate oxidation of arsenic(III): Reaction stoichiometry and the characterization of solid product. Geochim. Cosmochim. Acta 2011, 75, 4713-4727. 41. Morgan, J. J., Kinetics of reaction between O2 and Mn(II) species in aqueous solutions. Geochim. Cosmochim. Acta 2005, 69, 35-48. 42. 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. 43. 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. 44. 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, 1036410371. 45. 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. 46. Bouznik, V.; Kirik, S.; Solovyov, L.; Tsvetnikov, A., A crystal structure of ultra-dispersed form of polytetrafluoroethylene based on X-ray powder diffraction data. Powder Diffr. 2004, 19, 219-224. 47. Schecher, W.; MINEQL, D. M., A chemical equilibrium modeling system, Version 4.6. Environmental Research Software: Hallowell, ME 2007.

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513 514

Figure 1. Cr(OH)3 oxidation by manganese oxide from pH 5 to pH 9 with 770M initial Cr(III)

515

(40 mg/L) and 436 M initial MnO2 (40 mg/L MnO2) in completely mixed batch experiments.

516

19 / 26



517 518

Figure 2. Cr(OH)3 (770 M/ 40 mg/L) oxidation by MnO2 at pH 5 with different overall doses

519

of MnO2 as seen by (a) the Cr(VI) production versus time and (b) the relationship between the

520

total Cr(VI) produced after 8 hours reaction and the amount of MnO2 added. The dashed line is

521

the best fitting line from a least squares regression.

20 / 26


Page 20 of 26

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522

523 524

Figure 3. Concentration of Cr(VI) from Cr(OH)3-MnO2 reaction in (a) multichamber

525

experiments at pH 5, (b) multichamber and completely mixed experiments at pH 5, (c)

526

multichamber and completely mixed experiments at pH 8. Sufficient data for control

527

experiments were available to parameterize a model to simulate the reactions at pH 5 but not at

528

pH 8. For all multi-chamber experiments, Cr(III)0 = 1440 M (80 mg/L) in the chromium

529

chamber and MnO2 = 872 M (80 mg/L MnO2) in MnO2 chamber. The concentrations in an

530

individual chamber are twice as high as in the completely mixed experiments, which then

531

provides overall (amount of solid for the entire reactor volume) Mn and Cr concentrations that 21 / 26



532

are the same for the multichamber and completely mixed experiments. Cr(VI) in Mn chamber

533

represents both dissolved and total Cr(VI) as no Cr(VI) adsorption on MnO2 solids. The dashed

534

lines for Figure 3 a and b are simulated Cr(VI) concentration from modeling. The rate constant

535

of Cr(III) dissolution and Cr(III) oxidation by MnO2 and the transfer coefficient of Cr(III) and

536

Cr(VI) are included to fit the model.

22 / 26


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537 538

Figure 4. X-ray diffraction patterns of MnO2 and Cr(OH)3 reaction products after 200 hours at

539

pH 5, pH 8, and pH 9. The reference patterns for feitknechtite (044-1445 from the International

540

Crystal Diffraction Database) and Cr(OH)3 is included for comparison. The hashtag (#) indicates

541

the diffraction features of triclinic birnessite and asterisk (*) indicates the diffraction features

542

from PTFE abraded from the stir bar.46

23 / 26



543 544

Figure 5. TEM images of the reaction products of Cr(OH)3 and MnO2 at pH 5 (a) and pH 9 (b).

545

The inset figures are the SAED patterns obtained from the area of the red circle.

24 / 26


Page 24 of 26

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546 547

Figure 6. XPS spectra of Mn 2p3/2 photoelectron lines for the solid products of Cr(OH)3 and

548

MnO2 in mixed batch experiments (a) at pH 5 and (b) pH 9 as well as (c) the initial MnO2. The

549

Mn(IV) had been reduced to Mn(II) and Mn(III) on the solid surface at both pH 5 and pH 9. 25 / 26



550

551 552 553 554 555

Page 26 of 26

Table 1. Comparison of Cr(OH)3 solubility with oxidation extent in the presence of MnO2 pH value 5 6 7 8 9 1 Amorphous Cr(OH)3 solubility (𝜇M) 17 0.3 0.02 0.0025 0.1 2 Crystalline Cr(OH)3 solubility (𝜇M) 1905 19 Measured Cr solubility3 (𝜇M) 240 1.0 0.8 0.2 0.9 Cr(VI)total produced at 400 hours (𝜇M) 95 80 80 130 200 Mn(II)total produced at 400 hours (𝜇M) 156.1 40.2 16.1 8.3 8.2 Mn(II)total/Cr(VI)total 1.64 0.50 0.20 0.06 0.04 47 1. Cr(III) solubility was calculated with MINEQL+ . 2. Cr(III) solubility was calculated based on literature32. 3. Cr(III) concentration was measured after equilibrating with 40 mg/L Cr(III) added as Cr(OH)3 at different pH for 24 hours before adding 436 M MnO2 to the completely mixed reactor.

556 557

Table 2. Summary of Mn oxidation state percent in solids determined using XPS Mn 2p3/2 Sample

Mn(IV)

Mn(III)

MnO2 initial solids

94.9%

5.1%

Cr(OH)3 + MnO2 pH 5 Cr(OH)3 + MnO2 pH 9

40.9% 16.5%

36.0% 61.6%

558 559 560

Graphical abstract

561

26 / 26


Mn(II) 23.1% 21.9%

3 Oxidation by Î´-MnO2 - ACS Publications - American Chemical Society

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