Metal Adsorption Controls Stability of Layered Manganese Oxides

Article Cite This: Environ. Sci. Technol. 2019, 53, 7453−7462

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Metal Adsorption Controls Stability of Layered Manganese Oxides Peng Yang,† Jeffrey E. Post,‡ Qian Wang,† Wenqian Xu,§ Roy Geiss,∥ Patrick R. McCurdy,∥ and Mengqiang Zhu*,† †

Department of Ecosystem Science and Management, University of Wyoming, Laramie, Wyoming 82071, United States Department of Mineral Sciences, Smithsonian Institution, Washington, District of Columbia 20013, United States § X-ray Science Division, Advanced Photon Source, Argonne National Laboratory, Lemont, Illinois 60439, United States ∥ Department of Chemistry, Colorado State University, Fort Collins, Colorado 80523, United States Downloaded via KEAN UNIV on July 23, 2019 at 00:36:03 (UTC). See https://pubs.acs.org/sharingguidelines for options on how to legitimately share published articles.

‡

S Supporting Information *

ABSTRACT: Hexagonal birnessite, a typical layered Mn oxide (LMO), can adsorb and oxidize Mn(II) and thereby transform to Mn(III)-rich hexagonal birnessite, triclinic birnessite, or tunneled Mn oxides (TMOs), remarkably changing the environmental behavior of Mn oxides. We have determined the effects of coexisting cations on the transformation by incubating Mn(II)-bearing δ-MnO2 at pH 8 under anoxic conditions for 25 d (dissolved Mn < 11 μM). In the Li+, Na+, and K+ chloride solutions, the Mn(II)-bearing δ-MnO2 first transforms to Mn(III)-rich δ-MnO2 or triclinic birnessite (T-bir) due to the Mn(II)−Mn(IV) comproportionation, most of which eventually transform to a 4 × 4 TMO. In contrast, Mn(III)-rich δ-MnO2 and T-bir form and persist in the Mg2+ and Ca2+ chloride solutions. However, in the presence of surface adsorbed Cu(II), Mn(II)-bearing δ-MnO2 turns into Mn(III)-rich δ-MnO2 without forming T-bir or TMOs. The stabilizing power of the cations on the δ-MnO2 structure positively correlates with their binding strength to δ-MnO2 (Li+, Na+, and K+ < Mg2+ and Ca2+ < Cu(II)). Since metal adsorption decreases the surface energy of minerals, our finding suggests that the surface energy largely controls the thermodynamic stability of LMOs. Our study indicates that the adsorption of divalent metal cations, particularly transition metals, can be an important cause of the high abundance of LMOs, rather than the more stable TMO phases, in the environment.

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INTRODUCTION Manganese (Mn) oxides are ubiquitous in soils, sediments, and ocean nodules.1 They are metal scavengers and strong oxidants,2−5 thereby affecting the fate of metals, nutrients, and organic compounds in the environment.5−13 The most common naturally occurring Mn oxide is birnessite,1 a layered Mn oxide (LMO) consisting of stacked layers constructed by edge-sharing MnO6 octahedra. Mn(IV) is dominant in its MnO6 layers, but a portion of Mn sites are vacant or substituted by Mn(III).1,14,15 Mn(III) affects the transformation of birnessite to tunneled Mn oxides (TMOs),16−18 another family of common naturally occurring Mn oxides but more thermodynamically stable than birnessite. Both Mn(III) and vacancies can also strongly affect birnessite metal sorption properties, oxidizing activity, and bandgap energies pertinent to photochemical reduction of birnessite.19−24 Mn(II) often encounters birnessite, such as during oxidative precipitation of Mn(II),25,26 partial reductive dissolution of birnessite,6,13 and in many other environmental scenarios.27,28 In these situations, inevitable adsorption and oxidation of Mn(II) by birnessite can increase Mn(III) but decrease vacancy concentration of birnessite or even transform birnessite to other Mn oxide phases.29−33 The oxidation of Mn(II) by Mn(IV) of birnessite is also called comproportio© 2019 American Chemical Society

nation. The products of the Mn(II)-birnessite reactions depend on the Mn(II)/MnO2 ratios and pH.25,29−33 At relatively low Mn(II)/MnO2 ratios (0.05−0.24) and pH ≥ 8, the reactions increase Mn(III) concentration but decrease vacancy concentration in the layers of hexagonal birnessite, eventually leading to the formation of triclinic birnessite (Tbir) that contains 1/3 Mn as Mn(III) without vacancies.32 The incubation of Mn(II)-bearing δ-MnO2 at pH 6−8 and 21 °C leads to the surprisingly rapid formation of a 4 × 4 TMO with T-bir or Mn(III)-rich δ-MnO2 as an intermediate product.16 Upon aging at pH 4 with Mn(II)/MnO2 ratios of 0.04−0.54, birnessite rearranges into a superlattice structure by rotated stacking or displacing the layers.13,34 In contrast, at relatively high Mn(II)/MnO2 ratios (e.g., 0.30−3.04) and pH 7.0−8.5, birnessite transforms to feitknechtite (β-MnOOH), manganite (γ-MnOOH), and hausmannite (Mn3O4).25,29−31 The birnessite structure can host large numbers of metal cations due to its strong metal adsorption capability for alkali, alkaline earth, and transition metals.3−5,35−37 The strong metal Received: Revised: Accepted: Published: 7453

February 27, 2019 May 16, 2019 May 20, 2019 May 22, 2019 DOI: 10.1021/acs.est.9b01242 Environ. Sci. Technol. 2019, 53, 7453−7462

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Environmental Science & Technology

MnCl2 at pH 4 in a 100 mM NaCl solution for 2 h in air to reach the highest Mn(II) sorption loading achievable at pH 4 (Mn(II)/MnO2 = 0.14).16 With such low pH and the short reaction time, minimal oxidation occurred to adsorbed Mn(II) by either atmospheric O2 or Mn(IV) in δ-MnO2. A control system was prepared using the same procedures but without adding Mn(II). The solids were collected using vacuum filtration, rinsed with DI water acidified to pH 4 by HCl, and subsequently used for the incubation experiments. To examine the effects of Cu(II) on the Mn(II)-birnessite reactions, 7 mM CuCl2 and 5 mM MnCl2 were equilibrated with 4 g/L δ-MnO2 at pH 4 for 2 h, resulting in a Mn(II)/ MnO2 ratio of 0.08 and a Cu(II)/MnO2 ratio of 0.14. In another experiment, 4 g/L δ-MnO2 was equilibrated with 3 mM MnCl2 at pH 4 for 2 h, resulting in a similar Mn(II)/ MnO2 molar ratio (0.08) as that in the Cu(II) system. The preloading of Cu(II) on δ-MnO2 was to minimize otherwise precipitation of Cu(II) during incubation at pH 8 (see below) and to mimic naturally occurring birnessite that usually adsorbs transition metals in its structure.50 With the same Mn(II) loading, a comparison of the two systems can reveal the effects of Cu(II) on the transformation of Mn(II)-bearing δ-MnO2. Incubation of Mn(II)-Bearing δ-MnO2. The experimental conditions are summarized in Table 1. The incubation

sorption likely affects the rate and products of Mn(II)-induced birnessite transformation. At high Mn(II)/MnO2 ratios (0.43− 1.04) and in the presence of Ni(II) and Zn(II), birnessite transforms into Ni-substituted feitknechtite and Zn-substituted hausmannite at pH 7.5−8, respectively,36,38,39 and the Ni substitution inhibits further transformation of feitknechtite to manganite.38 With Mn(II)/MnO2 ratios of 0.04−0.54, Ni(II) and Zn(II) mainly compete with Mn(II) for adsorption without changing the mineral phase of birnessite at pH 4 and 7.40 However, the effects of metal cations on the Mn(II)birnessite reactions at relatively low Mn(II)/MnO2 ratios remain unknown, although the low ratios are common in the natural environment.41,42 In the present study, we determined the effects of selected alkali (Li+, Na+, and K+), alkaline earth (Mg2+ and Ca2+), and transition metal (Cu(II)) cations on the rate and products of Mn(II)-induced birnessite transformation at Mn(II)/MnO2 ratios of ≤0.14, which are considered as low Mn(II)/MnO2 ratios based on the Mn(II) adsorption loading.16 Mn(II)bearing δ-MnO2 was prepared first and then incubated in solutions of each alkali and alkaline earth metal cation at pH 8 in a N2/H2 atmosphere (95% N2 + 5% H2) for 25 d. In terms of Cu(II) effects, Mn(II)-bearing δ-MnO2 with Cu(II) adsorbed on the surface was incubated in NaCl solution. The three categories of cations were selected because they are common in the natural environment and their adsorption behaviors on birnessite differ markedly. Li+, Na+, and K+ occupy the interlayers of birnessite by forming outer-sphere complexes.43,44 Both Mg2+ and Ca2+ adsorb on birnessite probably as a mixture of inner- and outer-sphere complexes.45,46 Cu(II) adsorbs on vacancies or edges as innersphere complexes,12,37 enters the vacancies to be part of the layer,10,12 or forms polynuclear clusters on edges of birnessite.37,47 The binding strength of the three categories of cations to birnessite increases in the order of Li+, Na+, and K+ < Mg2+ and Ca2+ < Cu(II).48 The binding strength of the cations within the alkali or alkaline earth group also differs. As metal adsorption affects energetics of the adsorbent, we expect that the three categories of cations have distinct impacts on the transformation of Mn(II)-bearing δ-MnO2. We also examined the effects of Na+ concentration on the transformation of Mn(II)-bearing δ-MnO2. δ-MnO2 was chosen because it is a synthetic analogue to widely spread vernadite and biogenic Mn oxides in the environment.49 The weakly alkaline condition (pH 8) is prevalent in alkaline soils and marine environment. The N2/H2 atmosphere was used to mimic suboxic environments with a low O2 content and to avoid interference of atmospheric O2 with Mn(II)-birnessite reactions.16,32 This study provides important insights into how metal adsorption affects the stability of LMOs with respect to phase transformation to more crystalline birnessite and tunneled structures.

Table 1. Summary of the Experimental Conditionsa

Condition

Initial Mn(II) concentration (mM)

Mn(II)/MnO2 molar ratio in Mn(II)-bearing δMnO2

0Mn

0

0

3Mn

3

0.08

0Na 10Na

10 10

0.14 0.14

Li+

10

0.14

Na+

10

0.14

K+

10

0.14

Mg2+

10

0.14

Ca2+

10

0.14

Cu(II)

5 mM Mn(II) + 7 mM Cu(II)

0.08

Background electrolyte 100 mM NaCl 100 mM NaCl 0 mM NaCl 10 mM NaCl 100 mM LiCl 100 mM NaCl 100 mM KCl 33.3 mM MgCl2 33.3 mM CaCl2 100 mM NaCl

Ionic strength (mM) 100 100 0 10 100 100 100 100 100 100

a All experiments were carried out using δ-MnO2 (4 g/L) as starting material at pH 8 in a N2/H2 atmosphere.

procedures were the same as those in our previous study.16 All incubation plastic bottles were wrapped with aluminum foil to prevent potential photoreduction of δ-MnO2. The Mn(II)bearing δ-MnO2 prepared with 10 mM Mn(II) was incubated in LiCl, NaCl, KCl, MgCl2, and CaCl2 solutions in the anaerobic chamber to investigate the effects of these cations on the Mn(II)-birnessite reactions. The ionic strength was 100 mM, controlled by 100 mM alkali metal (Li+, Na+, and K+) chloride or 33.3 mM alkaline earth metal (Mg2+ and Ca2+) chloride. Cu/Mn(II)-bearing δ-MnO2 and Mn(II)-bearing δ-MnO2 prepared with 3 mM Mn(II) (3Mn) were incubated in a 100 mM NaCl solution. As a control, δ-MnO2 equilibrated with 0

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MATERIALS AND METHODS Materials. All chemicals were of A.C.S. reagent grade and used as received. Solutions used in incubation experiments were prepared in an anaerobic chamber (Coy Vinyl-A, 5% H2 + 95% N2) with degassed deionized (DI) water (18.2 MΩ· cm). δ-MnO2 was prepared via KMnO4 reduction by Mn(NO3)2 (Supporting Information, SI-1). Preparation of Mn(II)-Bearing δ-MnO2. Mn(II)-bearing δ-MnO2 was prepared according to our previous study.16 Asprepared δ-MnO2 (4 g/L) was equilibrated with 10 mM 7454

DOI: 10.1021/acs.est.9b01242 Environ. Sci. Technol. 2019, 53, 7453−7462

Article


Figure 1. XRD patterns of as-prepared δ-MnO2 (As-prep) and incubated samples collected from the Li+, Na+, K+ (a), 0Na, 10Na (b), Mg2+, Ca2+ (c), Cu(II), and 3Mn (d) systems. 3Mn and 10Mn indicate the Mn(II) concentration of 3 and 10 mM during the equilibration with δ-MnO2 at pH 4, respectively. Init stands for δ-MnO2 equilibrated with Mn(II) and Cu(II) at pH 4 without incubation at pH 8. 0Na and 10Na denote the systems with 0 and 10 mM NaCl as background electrolytes, respectively. In panels c and d, the vertical solid lines indicate the appearance of T-bir. The vertical dash lines at 4.6, 4.7, and 4.9 Å indicate the appearance of feitknechtite (β-MnOOH), pyrochroite (Mn(OH)2), and hausmannite (Mn3O4), respectively. The XRD pattern of the As-prep sample is overlapped with initial samples in all panels and with incubated samples in the Ca2+, 3Mn, and Cu(II) systems in panels c and d to illustrate the changes of XRD patterns.

were measured using inductively coupled plasma-optical emission spectrometry. The concentrations of both metals were very low (Table S1), indicating a negligible release of the metals into solution during the incubations. Solid Characterization. The obtained solids from the incubation experiments were subject to X-ray diffraction (XRD), atomic pair distribution function (PDF), transmission electron microscopy (TEM), and Raman and X-ray absorption spectroscopic analyses. TEM images were collected on a FEI Tecnai G2 F20 200 kV (in the Department of Geology and Geophysics, University of Wyoming) or a JEOL 2100F 200 kV (in the Department of Chemistry, Colorado State University) transmission electron microscope. The XRD patterns and the total X-ray scattering data for the PDF analysis were collected, respectively, using Xrays of 0.4521 Å at beamline 17-BM-B and 0.2114 Å at

mM Mn(II) (0Mn) was also incubated in a 100 mM NaCl solution. Mn(II)-bearing δ-MnO2 prepared with 10 mM Mn(II) was also incubated in 0 and 10 mM NaCl solutions to determine the effects of NaCl concentration on Mn(II)-birnessite reactions. The suspension pH was controlled at pH 8 by an automatic pH titrator (Metrohm 907) for the first 2 h and by manual adjustment thereafter. The incubation lasted for 25 d, during which suspension aliquots were collected at predetermined time intervals and filtered. The obtained solids were rinsed with 10 mL of DI water of pH 8, dried in the anaerobic chamber, and ground for the following characterization. The filtrates were acidified for measuring dissolved Mn concentration using the formaldoxime colorimetric method51 whereas the dissolved Mn and Cu concentrations in the Cu(II) system 7455


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Figure 2. TEM images of starting materials and 25 d samples from different reaction systems. As-prep and 10Mn Init stand for the as-prepared δMnO2 and that equilibrated with 10 mM Mn(II) at pH 4 but without being incubated at pH 8, respectively.

Figure 3. Changes of the PDF peak area at 5.3 Å with time in different reaction systems. The insets show the data points within 10 h. As-prep stands for the as-prepared δ-MnO2.

combination fitting (LCF) analysis of Mn K-edge X-ray absorption near edge structure (XANES) spectra. The XANES data were collected at beamline 10-BM-B at the APS. Cu Kedge extended X-ray absorption fine structure (EXAFS) spectra were collected at beamline 4-1 or 7-3 at the Stanford Synchrotron Radiation Lightsource (SSRL) to characterize the local coordination structure of Cu in the Cu(II) system. More

beamline 11-ID-B at the Advanced Photon Source (APS), Argonne National Laboratory. Raman spectra were collected using a HORIBA LabRAM HR Evolution Raman microscope in the Department of Mineral Sciences at the Smithsonian Institution. To estimate the fraction of each Mn oxidation state in the incubated solids, selected solids were analyzed by linear 7456


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Figure 4. XANES-LCF derived Mn(III) (a) and Mn(II) (b) molar percentages for selected reaction systems. As-prep stands for the as-prepared δMnO2.

Effects of Li+, Na+, and K+. Upon incubation at pH 8 for 25 d, Mn(II)-bearing δ-MnO2, which was prepared by equilibrating 4 g/L δ-MnO2 with 10 mM Mn(II) at pH 4 for 2 h, transforms into T-bir within 15 or 25 d and partially transforms into a 4 × 4 TMO after 15 d in all Li+, Na+, and K+ systems (Figure 1a); minor pyrochroite (Mn(OH)2) also forms in these systems (Figure 1a). MnOOH and Mn3O4 phases were not observed. The formation of T-bir, indicated by the splitting of the diffraction peaks around 2.4 and 1.4 Å and the arising of new peaks between (Figure 1a), is due to the comproportionation between adsorbed Mn(II) and Mn(IV) of δ-MnO2 at alkaline pH (i.e., pH 8),16 producing Mn(III) (Figure 4a) and decreasing Mn(II) and Mn(IV) concentrations (Figures 4b and S5). Particle growth via the oriented attachment between T-bir particles further increases particle sizes, resulting in well crystallized T-bir.32 These observations are consistent with our previous study conducted in a NaCl solution,16 suggesting that the three monovalent cations impose similar impacts on the transformation in spite of their different binding strength. However, the three systems differ in the crystallinity, amount, and formation rate of the transformation products. The crystallinity of T-bir decreases in the order of Li+ > Na+ > K+, indicated by the decreasing sharpness of the characteristic XRD peaks of T-bir (Figure 1a). The high crystallinity of Li+-exchanged T-bir was also observed previously.53 However, the amount of T-bir formed in the three systems does not follow that order. The Na+ system produces the largest amount of T-bir, followed by the Li+ and then K+ systems according to the XRD patterns (Figure 1a). The amounts of T-bir formed in those systems are positively proportional to the Mn(III) percentages (Figure 4a) estimated by the XANES-LCF analysis (Figure S6 and Table S2) and to the amounts of the decrease in the PDF peak area at 5.3 Å (Figures 3a and S3e,f, S4a). However, the solids in the Li+ system still contain a high number of Mn(III) (26% at 25 d), which is very close to that in the Na+ system (Figure 4a). Providing that the XRD peaks of T-bir are relatively weaker in the Li+ system than in the Na+ system, we may conclude that a small portion of Mn(III) adsorbs on vacancies without being incorporated into the layers, i.e., a formation of Mn(III)-rich δ-MnO2.16

details about the solid characterization are provided in the Supporting Information (SI-2).

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RESULTS AND DISCUSSION Equilibrating δ-MnO2 with Mn(II) and/or Cu(II) solutions at pH 4 slightly changes the structure and composition of δMnO2. The obtained Mn(II)-bearing or Mn(II)/Cu(II)bearing δ-MnO2 has a similar mineral phase and morphology to as-prepared δ-MnO2 according to the XRD (Figures 1 and S1a) and TEM analyses (Figures 2 and S2). Important changes pertinent to Mn(II)/Cu(II) adsorption occur to the peak at 5.3 Å in the PDFs (Figures 3 and S3, S4) and the dip at 2.0 Å in the XRD patterns (Figures 1 and S1a).52 The 5.3 Å peak corresponds to the atomic pairs between the Mn and/or Cu adsorbed on a vacancy and the second nearest layer Mn atoms surrounding the vacancy.6,16 The increased intensity of this PDF peak is consistent with the adsorption of Mn(II) and/or Cu(II) on vacancies after the equilibration. The adsorption of the metals on vacancies also makes the dip at 2.0 Å in the XRD pattern more pronounced (Figures 1 and S1a). In addition, the equilibration of δ-MnO2 with Mn(II) solutions slightly increases the Mn(III) percentages by 4−6% (Figure 4a), estimated by the XANES-LCF analysis (Table S2), and leads to the formation of minor feitknechtite (Figures 1 and S1a). These changes of δ-MnO2 upon reaction with Mn(II) and/or Cu(II) are consistent with our previous study.16 As a control system, the δ-MnO2 prepared by reacting 0 mM Mn(II) (i.e., 0Mn) at pH 4 was incubated at pH 8 for 25 d under anoxic conditions. Both the 2.0 Å XRD dip (Figure S1b) and the 5.3 Å PDF peak (Figures 3b and S3a) become slightly weaker with increasing incubation time, indicating decreasing numbers of Mn(II,III) adsorbed on vacancies,2,6,16,52 which is caused by Mn(III) incorporation into vacancies.6,16 As shown and discussed below, Mn(II)-bearing δ-MnO2, however, undergoes remarkable changes in mineral phase and Mn oxidation state composition during incubation in solutions of various background electrolytes. The transformation kinetics and products depend on the type of the metal cations, which can be ascribed to the different binding strength of the cations to δ-MnO2. 7457


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Environmental Science & Technology The formation of the TMO phase occurs earlier in both Li+ and K+ systems than in the Na+ system, which is clearly indicated in the earlier increase of the PDF peak area at 5.3 Å (Figure 3a) and earlier appearance of needle-like crystals (Figure S2). The PDF peak area at 5.3 Å increases because the TMO contains abundant Mn−Mn corner-sharing atomic pairs with an interatomic distance of 5.3 Å.16 The faster transformation to the TMO in the Li+ and K+ systems than that in the Na+ system can be ascribed to the formation of the Mn(III)-rich δ-MnO2 with Mn(III) adsorbed on vacancies, which acts as a viable precursor for the transformation.16 The transformation of well-crystallized T-bir to the TMO in the Na+ system is slower. The XRD peak intensities of T-bir decrease significantly in both Na+ and K+ systems at 25 d, indicating a more extended transformation of T-bir to the TMO. In contrast, T-bir persists in the Li+ system, probably due to its very high crystallinity. The greater increase of the PDF peak area at 5.3 Å (Figure 3a) and more needle-like particles at 25 d (Figure 2) in the K+ system suggest that more TMO forms in the K+ system than in the Na+ and Li+ systems. However, the strong XRD peak at 12 Å suggests that the TMO formed in the Na+ system has the highest crystallinity (Figure 1a), which could be due to the slower transformation with well-crystallized T-bir as the precursor, favoring the formation of the well-crystallized TMO. Raman spectroscopy was used to further identify the solid phases. In the Li+ system, Raman spectrum of the 10 h solid has two dominant bands at 574 and 651 cm−1 (Figure 5), indicating a layered structure.54 The spectrum of the 25 d solid has an absorption band around 730 cm −1 , probably corresponding to tunneled structures.54,55 Similarly, the solids at 75 or 10 h in both Na+ and K+ systems show more spectral features of δ-MnO2 or T-bir, while the 25 d samples have more spectral features of tunneled structures (Figure 5). These results agree with the XRD analyses (Figure 1a). Concentrations of monovalent cations also affect the Mn(II)-induced birnessite transformation. Higher NaCl concentration favors the formation of T-bir. Less T-bir forms in the 10 mM NaCl than in the 100 mM NaCl solution (Figure 1a,b); and well-crystallized T-bir is not detected in the 0 mM NaCl solution (Figure 1b). However, poorly crystalline T-bir or Mn(III)-rich δ-MnO2 might be abundant in the solutions of low NaCl concentrations, indicated by the sharp decrease of PDF peak area at 5.3 Å within 5 h (Figure 3b). The favorability of the high NaCl concentration for T-bir formation could be caused by a sufficient number of Na+ that is needed to compensate the negative layer charges of T-bir and by a higher ionic strength that favors particle aggregation and thus particle growth via the oriented attachment.32,56 The TMO (4 × 4) forms at all NaCl concentrations, starting earlier at the lower NaCl concentration (Figure S2), while at 25 d, the transformation to the TMO is more extended and the TMO is more crystalline at the higher NaCl concentration. The earlier formation of the TMO is consistent with the abundance of poorly crystalline T-bir and/or Mn(III)-rich δ-MnO2 as the precursors.16 Once primary TMO particles form, the growth of the TMO crystals via the oriented attachment9 can be favored at the higher ionic strength,56 leading to the more extended transformation to the TMO at 25 d in the 100 mM NaCl solution. The transformations are consistent with the changes of the PDF peak area at 5.3 Å (Figures 3 and S3c,d,f). The observed differences among Li+, Na+, and K+ systems may be caused by their different magnitudes of electrostatic

Figure 5. Raman spectra of T-bir, woodruffite (3 × 4), todorokite (3 × 3), and selected samples from the Li+, Na+, K+, Mg2+, Ca2+, and Cu(II) systems.

attraction with MnO6 layers. The ionic potential, φH = Z/rH, where rH is the radius of hydrated cations,57 measures the charge density of a cation,58 and the electrostatic forces between these hydrated cations and birnessite layers increase in the order of Li+ < Na+ < K+ (Table S3). However, except for the crystallinity of T-bir, other aspects of the transformation, such as the amounts of T-bir and the TMO and their formation rates, do not correlate with the order of the ionic potentials, suggesting existence of other controlling factors, which could be hydration energy and ionic radius of cations, in addition to electrostatic interactions. Effects of Mg2+ and Ca2+. Different from alkali metal cations, both Mg2+ and Ca2+ lead to the formation of T-bir but not TMO. T-bir forms earlier, faster and more extensively and is much more crystalline in the Mg2+ system than in the Ca2+ system (Figure 1c). Consistently, the peak area at 5.3 Å in the PDFs decreases less in the Ca2+ system (Figure 3a). The continuous decrease of the PDF peak area at 5.3 Å in both systems (Figures 3a and S4b,c) is consistent with the XRD and TEM results showing that T-bir forms and increases in quantity and no tunneled phase forms. The Raman spectra of the solids in the Mg2+ system show a gradual transformation 7458


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Environmental Science & Technology from δ-MnO2 to T-bir with incubation time (Figure 5). The 25 d solid in the Ca2+ system has weak Raman spectral features of T-bir. None of the solids in the Mg2+ or the Ca2+ system possess spectral features of tunneled structures (Figure 5). The disfavored formation of T-bir in the Ca2+ system might result from the exchange of adsorbed Mn(II) by Ca2+ so that less Mn(III) formed via the Mn(II)−Mn(IV) comproportionation reaction.59 The formation of minor hausmannite (4.9 Å XRD peak in Figure 1c) suggests that the released Mn(II) may react with the edge sites of δ-MnO2 because the edge sites are negatively charged at pH 8 (pKa of Mn−OH groups on the edges is lower than 6).35,60 Due to the weaker adsorption, Mg2+ may not exchange adsorbed Mn(II) as strongly as Ca2+ does; otherwise, the Mn(III) percentage in the Mg2+ system would not reach 34% (Figure 4a), the percentage value of an ideal T-bir.61 In addition, Ca2+ can induce a drastic Mn(III) rearrangement in the MnO6 layers and an interstratification of the MnO6 layers with different Mn(III) distribution in T-bir,62 which may further decrease the crystallinity of T-bir. T-bir formed in the Mg2+ system has a high order of layer stacking, indicated by the strong (00l) XRD peaks (Figure 1c), which can be caused by the strong hydrogen bonds between the MnO6 layers and hydrated Mg2+.46 The high hydration energy of Mg2+ prevents the dehydration of Mg2+ upon air drying, leading to the (003) diffraction peak at 3.2 Å.63,64 Compared to the high stacking order, the degree of the intralayer ordering of the a-b plane of T-bir in the Mg2+ system is surprisingly low, as indicated by the broad XRD peaks between 2.4−1.4 Å (Figure 1c). The strong interaction between Mg2+ and the layers may cause the poor Mn(III) arrangement in the layers.46 Effects of Cu(II). Mn(II)/Cu(II)-bearing δ-MnO2 does not undergo much phase transformation. The XRD pattern of the 25 d solid is very similar to that of δ-MnO2 (Figure 1d), and the Raman spectrum of the solid is dissimilar to both T-bir and TMOs (Figure 5), indicating the absence of both T-bir and TMO phases. The 3Mn system uses a similar Mn(II)/MnO2 ratio as the Cu(II) system, but a small amount of poorly crystalline T-bir forms and decent layer stacking is developed, as indicated by the (001) and (002) XRD peaks (Figure 1d) and the strong decrease of the PDF peak area at 5.3 Å (Figures 3b and S3b). In spite of no mineral phase changes, the weakened dip at 2.0 Å in the XRD patterns (Figure 1d) and the reduced PDF peak area at 5.3 Å within 10 h (Figures 3b and S4d) suggest that a portion of Mn(II) is oxidized and incorporated into vacancies.2,52,59 Thus, the Cu(II) system leads to a formation of Mn(III)-rich δ-MnO2 which, however, does not transform to T-bir and the 4 × 4 TMO, suggesting an inhibition of Cu(II) on the transformation to T-bir and the 4 × 4 TMO. The speciation of Cu(II) in the solids provides insights into the inhibition of Cu(II) on T-bir formation. Both Cu K-edge XANES and EXAFS spectra show significant differences between the incubated samples and spertiniite (Cu(OH)2) and tenorite (CuO) references (Figure S7a,b), indicating that Cu(II) does not precipitate although the pH is high (pH 8). The EXAFS spectra of the incubated solids differ only slightly from the initial sample (Figure S7b), suggesting that the Cu coordination environment remains largely unchanged during the incubation. EXAFS spectral fits show that each Cu atom is surrounded by about 4 O atoms (CN = 3.7) at 1.95−1.96 Å (Figure S7b,c and Table S4), corresponding to the four equatorial Cu−O bonds of the Jahn−Teller distorted CuO6

octahedron.65,66 The two axial O atoms were not included in the fitting because of the strong thermal motion of the long axial bonds at room temperature.67 The second shell consists of 1.5 Cu at 2.90−2.93 Å, likely due to the formation of polynuclear clusters;37,47 however, we cannot exclude the possibility that a small number of Cu incorporates into layer vacancies. The small CNs of the second edge-sharing Cu−Cu/ Mn shell at 2.90−2.93 Å and the third corner-sharing Cu−Cu/ Mn shell at 3.38−3.40 Å suggest that Cu mainly adsorbs at edge sites as double edge-sharing complexes and single/double corner-sharing complexes.37 The absence of T-bir may be ascribed to the adsorption of Cu(II) clusters on the edge sites, leading to Mn(III) disordered arrangement in the layers. The Cu(II) in the layers may impair intralayer electron transfer and thus Mn(III) rearrangement as well. Role of Metal Adsorption on Stability of Layered Mn Oxides. The degree of the transformation of Mn(II)-bearing δ-MnO2 decreases in the order of Li+, Na+, or K+ > Mg2+ or Ca2+ > Cu(II). The cation impacts on the transformation can be understood from thermodynamic perspectives by considering the contribution of surface energy to the stability of LMOs and how much surface energy metal adsorption decreases. LMOs have large total surface areas mainly from the basal planes for well-crystallized LMOs while from both edges and basal planes for small LMO particles. With decreasing basal plane sizes, the edge sites contribute increasingly more additional surface areas that can be substantial, such as for δMnO2. Because of the high surface area to volume ratios of LMOs, the surface energy contributes a large portion of the total free energy (including both bulk and surface), elevating the total free energy and probably driving their transformation to TMOs.68,69 Due to the large contribution from the edge sites, small LMO particles have high total free energy and readily transform to more crystalline birnessite phase or to TMOs. For example, δ-MnO2 grows bigger in particle sizes via the oriented attachment70 or transforms to cryptomelane (2 × 2) at room temperature,71,72 while acid birnessite with a larger particle size does not. However, ion adsorption on LMO surfaces (including both edges and basal planes) can decrease the surface energy and thus the total free energy, increasing their thermodynamic stability. The extent to which the surface energy can be decreased depends on the metal binding strength. In the present work, Cu(II) adsorbs very strongly on LMO surfaces and the surface and total free energy of LMOs can be decreased substantially, thereby stabilizing Mn(III)-rich δMnO2 without forming T-bir and the 4 × 4 TMO (Figure 6). Similarly, the adsorption of transition metals retards the transformation of LMOs to todorokite, a 3 × 3 TMO, under refluxing conditions.9,73−75 Mn(III)O6 is Jahn−Teller distorted, and its presence and ordered distribution in layers, such as in T-bir, can increase the surface and total free energy of LMOs. Cu(II) adsorption decreases the Gibbs free energy of Mn(III)-rich δ-MnO2,76 and the energy of Mn(III)-rich δMnO2 with Cu(II) adsorbed is likely lower than that of T-bir, disfavoring the transformation of Mn(III)-rich δ-MnO2 to Tbir. Thus, Mn(III)-rich δ-MnO2 still maintains its hexagonal layer symmetry (Figure 6). The stabilization of hexagonal birnessite by transition metals is also evidenced in the previous studies although unlike Cu(II), vacancies are the main adsorption sites of those metals.5,77 For example, during the biogenic formation of nanoparticulate birnessite, Na+ and Ca2+ favor the formation of T-bir while Ni(II) favors the formation 7459


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on the stability of LMOs need to be addressed in future studies.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.est.9b01242.

Figure 6. Transformation pathways of Mn(II)-bearing δ-MnO2 to other phases in different systems.

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Details of preparation of δ-MnO2; TEM characterization and XRD and Raman spectroscopic analyses; XRD patterns, TEM images, PDF data, XANES and EXAFS spectra, and corresponding fitting results; dissolved Mn and Cu concentrations; physicochemical atomic parameters of alkali and alkaline earth metal cations (PDF)

AUTHOR INFORMATION

Corresponding Author

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

of hexagonal birnessite.33 Zn(II) inhibits Mn(III) accumulation in the structure of δ-MnO2 during its formation78 or during its partial reduction by dissolved organic matter,39 and the Zn(II) adsorption on δ-MnO2 causes a migration of Mn(III) out of the layers.79 Note that metal adsorption may cause kinetic hindrance of the transformation of hexagonal birnessite, which could also contribute to the increased stability of LMOs. As to Mg2+ and Ca2+, their binding to LMOs is much weaker than those of transition metal cations. Thus, Mn(III)-bearing δ-MnO2 proceeds to form T-bir. The high crystallinity and abundance of T-bir in the Mg2+ system are consistent with its weaker binding strength. However, the adsorption of Ca2+ and Mg2+ in the interlayers stabilizes T-bir, retarding or preventing its transformation to TMOs at room temperature (Figure 6). High temperature would be required for the transformation of Mg2+-exchanged T-bir to TMOs.80 Monovalent cations Li+, Na+, and K+ weakly bind to LMOs, and their adsorption cannot much decrease the surface and total free energy of Mn(III)-bearing δ-MnO2. Thus, in their solutions, Mn(III)rich δ-MnO2 and T-bir readily form but are unstable and eventually transform to the TMO (Figure 6). However, other environmental factors may also contribute to the high abundance of LMOs in the natural environment, such as the adsorption of oxyanions and the rapid formation of biogenic birnessite. Environmental Implications. Layered Mn oxide minerals are common in the natural environment and play an important role in elemental cycles and pollutant dynamics. They have large surface areas on basal planes and edges, are not thermodynamically stable, and can transform to TMOs. Our findings suggest that the unusually long lifetimes of LMOs in the natural environment may be caused by the adsorption of divalent cations, probably also trivalent cations (e.g., Fe(III)), particularly transition metals, which increases their thermodynamic stability by decreasing their surface energies. Thus, the stabilization of LMOs by metal adsorption may be an important cause for the overall higher abundance of LMOs than TMOs in the natural environment. The increased stabilization by metal adsorption may also decrease their redox potentials because the surface energy shifts redox equilibria.81 Our results have important implications for understanding the environmental and geochemical impacts of Mn oxides because LMOs are more reactive than TMOs in terms of oxidation and metal adsorption reactivity.82 Effects of other environmental factors, such as adsorption of oxyanions,

ORCID

Mengqiang Zhu: 0000-0003-1739-1055 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This study was supported by the U.S. Department of Energy Experimental Program to Stimulate Competitive Research Office for financial support (DOE-EPSCoR DE-SC0016272). We thank Dr. Xiaoming Wang in the College of Resources and Environment, Huazhong Agricultural University, Wuhan, China for providing Cu K-edge X-ray absorption spectra of spertiniite (Cu(OH)2) and tenorite (CuO). This work utilized resources of the APS, a U.S. DOE Office of Science User Facility, operated for the DOE Office of Science by the Argonne National Laboratory under Contract No. DE-AC0206CH11357. Use of the SSRL, SLAC National Accelerator Laboratory, was supported by the U.S. DOE, Office of Science, Office of Basic Energy Sciences, under Contract No. DEAC02-76SF00515.

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Metal Adsorption Controls Stability of Layered Manganese Oxides

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