Metal Adsorption Controls Stability of Layered Manganese Oxides

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Metal Adsorption Controls Stability of Layered Manganese Oxides Peng Yang, Jeffrey E. Post, Qian Wang, Wenqian Xu, Roy H. Geiss, Patrick R McCurdy, and Mengqiang Zhu Environ. Sci. Technol., Just Accepted Manuscript • Publication Date (Web): 22 May 2019 Downloaded from http://pubs.acs.org on May 28, 2019

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Environmental Science & Technology

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Metal Adsorption Controls Stability of Layered Manganese Oxides

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Peng Yang,† Jeffrey E. Post,‡ Qian Wang,† Wenqian Xu,§ Roy Geiss,∥ Patrick R. McCurdy,∥ and

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Mengqiang Zhu†,*

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†Department

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Wyoming 82071, United States

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‡Department

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20013, United States

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

of Mineral Sciences, Smithsonian Institution, Washington, District of Columbia

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§X-ray

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Illinois 60439, United States

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∥Department of Chemistry, Colorado State University, Fort Collins, Colorado 80523, United States

Science Division, Advanced Photon Source, Argonne National Laboratory, Lemont,

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*Corresponding author: E-mail: [email protected]

ACS Paragon Plus Environment

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Abstract

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Hexagonal birnessite, a typical layered Mn oxide (LMO), can adsorb and oxidize Mn(II)

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and thereby transform to Mn(III)-rich hexagonal birnessite, triclinic birnessite, and/or tunneled Mn

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oxides (TMOs), remarkably changing the environmental behavior of Mn oxides. We have

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determined the effects of co-existing cations on the transformation by incubating Mn(II)-bearing

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δ-MnO2 at pH 8 under anoxic conditions for 25 d (dissolved Mn < 11 μM). In the Li+, Na+ or K+

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chloride solution, the Mn(II)-bearing δ-MnO2 first transforms to Mn(III)-rich -MnO2 and/or

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triclinic birnessite (T-bir) due to the Mn(II)-Mn(IV) comproportionation, most of which eventually

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transform to a 4  4 TMO. In contrast, Mn(III)-rich -MnO2 and T-bir form and persist in the Mg2+

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or Ca2+ chloride solution. However, in the presence of surface adsorbed Cu(II), Mn(II)-bearing

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δ-MnO2 turns into Mn(III)-rich -MnO2 without forming T-bir or TMOs. The stabilizing power of

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the cations on the -MnO2 structure positively correlates with their binding strength to -MnO2

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(Li+, Na+ or K+ < Mg2+ or Ca2+ < Cu(II)). Since metal adsorption decreases the surface energy of

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minerals, our finding suggests that the surface energy largely controls the thermodynamic stability

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of LMOs. Our study indicates that the adsorption of divalent metal cations, particularly transition

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metals, can be an important cause of the high abundance of LMOs, rather than the more stable

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TMO phases, in the environment.


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INTRODUCTION

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Manganese (Mn) oxides are ubiquitous in soils, sediments and ocean nodules.1 They are

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metal scavengers and strong oxidants,2-5 thereby affecting the fate of metals, nutrients and organic

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compounds in the environment.5-13 The most common naturally-occurring Mn oxide is birnessite,1

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a layered Mn oxide (LMO) consisting of stacked layers constructed by edge-sharing MnO6

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octahedra. Mn(IV) is dominant in MnO6 layers, but a portion of Mn sites are vacant or substituted

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by Mn(III).1,14,15 Mn(III) affects the transformation of birnessite to tunneled Mn oxides (TMOs),16-

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18

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than birnessite. Both Mn(III) and vacancies can also strongly affect birnessite metal sorption

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properties, oxidizing activity, and bandgap energies pertinent to photochemical reduction of

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birnessite.19-24

another family of common naturally-occurring Mn oxides but more thermodynamically stable

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Mn(II) often encounters birnessite, such as during oxidative precipitation of Mn(II),25,26

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partial reductive dissolution of birnessite,6,13 and in many other environmental scenarios.27,28 In

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these situations, inevitable adsorption and oxidation of Mn(II) by birnessite can increase Mn(III)

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but decrease vacancy concentration of birnessite, or even transform birnessite to other Mn oxide

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phases.29-33 The oxidation of Mn(II) by Mn(IV) of birnessite is also called comproportionation.

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The products of the Mn(II)-birnessite reactions depend on the Mn(II)/MnO2 ratios and pH.25,29-33

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At relatively low Mn(II)/MnO2 ratios (0.05 – 0.24) and pH ≥ 8, the reactions increase Mn(III)

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concentration but decrease vacancy concentration in the layers of hexagonal birnessite, eventually

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leading to formation of triclinic birnessite (T-bir) that contains 1/3 Mn as Mn(III) without

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vacancies.32 The incubation of Mn(II)-bearing δ-MnO2 at pH 6 – 8 and 21 oC leads to surprisingly

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rapid formation of 4 × 4 TMO with T-bir or Mn(III)-rich δ-MnO2 as intermediate products.16 Upon

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aging at pH 4 with Mn(II)/MnO2 ratios of 0.04 – 0.54, birnessite rearranges into a superlattice


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structure by rotated stacking or displacing the layers.13,34 In contrast, at relatively high

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Mn(II)/MnO2 ratios (e.g., 0.30 – 3.04) and pH 7.0 – 8.5, birnessite transforms to feitknechtite

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(-MnOOH), manganite (-MnOOH) and hausmannite (Mn3O4).25,29-31

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The birnessite structure can host large numbers of metal cations due to its strong metal

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adsorption capability for alkali, alkaline earth and transition metals.3-5,35-37 The strong metal

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sorption likely affects the rate and products of Mn(II)-induced birnessite transformation. At high

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Mn(II)/MnO2 ratios (0.43 – 1.04) and in the presence of Ni(II) and Zn(II), birnessite transforms

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into Ni-substituted feitknechtite and Zn-substituted hausmannite at pH 7.5 – 8, respectively,36,38,39

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and Ni substitution inhibits further transformation of feitknechtite to manganite.38 With an

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Mn(II)/MnO2 ratio of 0.04 – 0.54, Ni(II) and Zn(II) mainly compete with Mn(II) for adsorption

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without changing the mineral phase of birnessite at pH 4 and 7.40 However, the effects of metal

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cations on the Mn(II)-birnessite reactions at relatively low Mn(II)/MnO2 ratios remain unknown,

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although the low ratios are common in the natural environment.41,42

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In the present study, we determined the effects of selected alkali (Li+, Na+, and K+), alkaline

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earth (Mg2+ and Ca2+) and transition metal (Cu(II)) cations on the rate and products of

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Mn(II)-induced birnessite transformations at Mn(II)/MnO2 ratios of ≤ 0.14, which are considered

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as low Mn(II)/MnO2 ratios based on the Mn(II) adsorption loading.16 Mn(II)-bearing δ-MnO2 was

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prepared first and then incubated in solutions of each alkali and alkaline earth metal cation at pH

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8 in a N2/H2 atmosphere (95% N2 + 5% H2) for 25 d. In terms of Cu(II) effects, Mn(II)-bearing

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δ-MnO2 with Cu(II) adsorbed on the surface was incubated in NaCl solution. The three categories

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of cations were selected because they are common in the natural environment and their adsorption

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behaviors on birnessite differ markedly. Li+, Na+ and K+ occupy the interlayers of birnessite by

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forming outer-sphere complexes.43,44 Both Mg2+ and Ca2+ adsorb on birnessite probably as a


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mixture of inner- and outer-sphere complexes.45,46 Cu(II) adsorbs on vacancies or edges as inner-

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sphere complexes,12,37 enters the vacancies to be part of the layer,10,12 or forms polynuclear clusters

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on edges of birnessite.37,47 The binding strength of the three categories of cations to birnessite

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increases in the order of Li+, Na+, or K+ < Mg2+ or Ca2+ < Cu(II).48 The binding strength of the

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cations within the alkali or alkaline earth group also differs. As metal adsorption affects energetics

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of the adsorbent, we expect that the three categories of cations have distinct impacts on the

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transformation of Mn(II)-bearing δ-MnO2. We also examined the effects of Na+ concentration on

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the transformation of Mn(II)-bearing δ-MnO2. δ-MnO2 was chosen because it is a synthetic

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analogue to widely spread vernadite and biogenic Mn oxides in the environment.49 The weakly

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alkaline condition (pH 8) is prevalent in alkaline soils and marine environment. The N2/H2

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atmosphere was used to mimic suboxic environments with a low O2 content and to avoid

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interference of atmospheric O2 with Mn(II)-birnessite reactions.16,32 This study provides important

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insights into how metal adsorption affects the stability of LMOs with respect to phase

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transformation to more crystalline birnessite and tunneled structures.

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MATERIALS AND METHODS

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Materials. All chemicals were of A.C.S. reagent grade and used as received. Solutions

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used in incubation experiments were prepared in an anaerobic chamber (Coy Vinyl-A, 5% H2 +

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95% N2) with degassed deionized (DI) water (18.2 MΩ·cm). δ-MnO2 was prepared via KMnO4

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reduction of by Mn(NO3)2 (Supporting Information, SI-1).

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Preparation of Mn(II)-Bearing δ-MnO2. Mn(II)-bearing δ-MnO2 was prepared

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according to our previous study.16 As-prepared -MnO2 (4 g/L) was equilibrated with 10 mM

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MnCl2 at pH 4 in 100 mM NaCl solution for 2 h in air to reach the highest Mn(II) sorption loading

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achievable at pH 4 (Mn(II)/MnO2 = 0.14).16 With such low pH and the short reaction time, minimal


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oxidation occurred to adsorbed Mn(II) by either atmospheric O2 or Mn(IV) in -MnO2. A control

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system was prepared using the same procedures but without adding Mn(II). The solids were

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collected using vacuum filtration, rinsed with DI water acidified to pH 4 by HCl, and subsequently

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used for the incubation experiments.

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To examine the effects of Cu(II) on the Mn(II)-birnessite reactions, 7 mM CuCl2 and 5 mM

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MnCl2 were equilibrated with 4 g/L -MnO2 at pH 4 for 2 h, resulting in a Mn(II)/MnO2 ratio of

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0.08 and a Cu(II)/MnO2 ratio of 0.14. In another experiment, 4 g/L -MnO2 was equilibrated with

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3 mM MnCl2 at pH 4 for 2 h, resulting in a similar Mn(II)/MnO2 molar ratio (0.08) as that in the

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Cu system. The preloading of Cu(II) on -MnO2 was to minimize otherwise precipitation of Cu(II)

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during incubation at pH 8 (see below) and to mimic naturally-occurring birnessite that usually

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adsorbs transition metals in its structure.50 With the same Mn(II) loading, a comparison of the two

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systems can reveal the effects of Cu(II) on the transformation of Mn(II)-bearing -MnO2.

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Incubation of Mn(II)-Bearing δ-MnO2. The experimental conditions are summarized in

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Table 1. The incubation procedures were the same as those in our previous study.16 All incubation

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plastic bottles were wrapped with aluminum foil to prevent potential photoreduction of -MnO2.

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The Mn(II)-bearing δ-MnO2 prepared with 10 mM Mn(II) was incubated in LiCl, NaCl, KCl,

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MgCl2, or CaCl2 solutions in the anaerobic chamber to investigate the effects of these cations on

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the Mn(II)-birnessite reactions. The ionic strength was 100 mM, controlled by 100 mM alkali

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metal (Li+, Na+ or K+) chloride or 33.3 mM alkaline earth metal (Mg2+ or Ca2+) chloride.

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Cu/Mn(II)-bearing -MnO2 and Mn(II)-bearing -MnO2 prepared with 3 mM Mn(II)

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(3Mn) were incubated in 100 mM NaCl solution. As a control, -MnO2 equilibrated with 0 mM

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Mn(II) (0Mn) was also incubated in 100 mM NaCl solution.

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Mn(II)-bearing -MnO2 prepared with 10 mM Mn(II) was also incubated in 0 or 10 mM


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NaCl solutions to determine the effects of NaCl concentration on Mn(II)-birnessite reactions.

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The suspension pH was controlled at pH 8 by an automatic pH titrator (Metrohm 907) for

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the first 2 h and by manual adjustment thereafter. The incubation lasted for 25 d, during which

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suspension aliquots were collected at pre-determined time intervals and filtered. The obtained

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solids were rinsed with 10 mL DI water of pH 8, dried in the anaerobic chamber, and ground for

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the following characterization. The filtrates were acidified for measuring dissolved Mn

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concentration using the formaldoxime colorimetric method51 whereas the dissolved Mn and Cu

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concentrations in the Cu system were measured using inductively coupled plasma-optical emission

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spectrometry. The concentrations of both metals were very low (Table S1), indicating negligible

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release of the metals into solution during the incubations.

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Solid Characterization. The obtained solids from the incubation experiments were subject

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to X-ray diffraction (XRD), atomic pair distribution function (PDF), transmission electron

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microscopy (TEM), Raman spectroscopic, and X-ray absorption spectroscopic analyses.

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TEM images were collected on a FEI Tecnai G2 F20 200 kV (at the Department of Geology

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and Geophysics, University of Wyoming) or a JEOL 2100F 200 kV (at the Department of

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Chemistry, Colorado State University) transmission electron microscope. The XRD patterns and

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the total X-ray scattering data for the PDF analysis were collected, respectively, using X-rays of

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0.4521 Å at beamline 17-BM-B and 0.2114 Å at beamline 11-ID-B at the Advanced Photon Source

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(APS), Argonne National Laboratory. Raman spectra were collected using a HORIBA LabRAM

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HR Evolution Raman microscope in the Department of Mineral Sciences at the Smithsonian

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Institution.

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To estimate the fraction of each Mn oxidation state in the incubated solids, selected solids

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were analyzed by linear combination fitting (LCF) analysis of Mn K-edge X-ray absorption near


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edge structure (XANES) spectra. The XANES data were collected at beamline 10-BM-B at APS.

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Cu K-edge extended X-ray absorption fine structure (EXAFS) spectra were collected at beamline

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4-1 or 7-3 at the Stanford Synchrotron Radiation Lightsource (SSRL) to characterize the local

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coordination structure of Cu in the Cu system. More details about the solid characterization are

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provided in the supporting information (SI-2).

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RESULTS AND DISCUSSION

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Equilibrating -MnO2 with Mn(II) and/or Cu(II) solutions at pH 4 slightly changes the

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structure and composition of -MnO2. The obtained Mn(II)-bearing or Mn(II)/Cu(II)-bearing

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δ-MnO2 has a similar mineral phase and morphology to as-prepared δ-MnO2 according to the XRD

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(Figures 1 and S1a) and HR-TEM analyses (Figures 2 and S2). Important changes pertinent to

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Mn(II)/Cu(II) adsorption occur to the peak at 5.3 Å in the PDFs (Figures 3, S3 and S4) and the dip

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at 2.0 Å in the XRD patterns (Figures 1 and S1a).52 The 5.3 Å peak corresponds to the atomic pairs

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between the Mn and/or Cu adsorbed on a vacancy and the second nearest layer Mn atoms

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surrounding the vacancy.6,16 The increased intensity of this PDF peak is consistent with adsorption

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of Mn(II) and/or Cu(II) on vacancies after the equilibration. The adsorption of the metals on

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vacancies also makes the dip at 2.0 Å in the XRD pattern more pronounced (Figures 1 and S1a).

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In addition, the equilibration of -MnO2 with Mn(II) solutions slightly increases Mn(III)

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percentages by 4 – 6% (Figure 4a) estimated by the XANES-LCF analysis (Table S2) and leads to

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the formation of minor feitknechtite (Figures 1 and S1a). These changes of -MnO2 upon reaction

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with Mn(II) and/or Cu(II) are consistent with our previous study.16

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As a control system, the δ-MnO2 prepared by reacting 0 mM Mn(II) (i.e., 0Mn) at pH 4

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was incubated at pH 8 for 25 d under anoxic conditions. Both the 2.0 Å XRD dip (Figure S1b) and

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the 5.3 Å PDF peak (Figures 3b and S3a) become slightly weaker with increasing incubation time,


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indicating decreasing numbers of Mn(II,III) adsorbed on vacancies,2,6,16,52 which is caused by

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Mn(III) incorporation into vacancies.6,16 As shown and discussed below, Mn(II)-bearing δ-MnO2,

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however, undergoes remarkable changes in mineral phases and Mn oxidation state composition

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during incubation in solutions of various background electrolytes. The transformation kinetics and

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products depend on the type of the metal cations, which can be ascribed to the different binding

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strength of the cations to -MnO2.

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Effects of Li+, Na+, and K+. Upon incubation at pH 8 for 25 d, Mn(II)-bearing δ-MnO2,

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which was prepared by equilibrating 4 g/L -MnO2 with 10 mM Mn(II) at pH 4 for 2 h, transforms

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into T-bir within 15 or 25 d and partially transforms into 4 × 4 TMO after 15 d in all Li+, Na+, and

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K+ systems (Figure 1a); minor pyrochroite (Mn(OH)2) also forms in these systems (Figure 1a).

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MnOOH and Mn3O4 phases were not observed. The formation of T-bir, indicated by the splitting

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of the diffraction peaks around 2.4 and 1.4 Å and the arising of new peaks between (Figure 1a), is

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due to the comproportionation between adsorbed Mn(II) and Mn(IV) of -MnO2 at alkaline pH

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(i.e., pH 8),16 producing Mn(III) (Figure 4a) and decreasing Mn(II) and Mn(IV) concentrations

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(Figures 4b and S5). Particle growth via oriented attachment between T-bir particles further

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increases particle sizes, resulting in well crystallized T-bir.32 These observations are consistent

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with our previous study conducted in NaCl solution,16 suggesting that the three monovalent cations

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impose similar impacts on the transformation in spite of their different binding strength. However,

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the three systems differ in the crystallinity, amount and formation rate of the transformation

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products.

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The crystallinity of T-bir decreases in the order of Li+ > Na+ > K+, indicated by the

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decreasing sharpness of the characteristic XRD peaks of T-bir (Figure 1a). The high crystallinity

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of Li+-exchanged T-bir was also observed previously.53 However, the amount of T-bir formed in


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the three systems does not follow that order. The Na+ system produces the largest amount of T-bir,

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followed by the Li+ and then K+ systems according to the XRD patterns (Figure 1a). The amounts

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of T-bir formed in those systems are positively proportional to the Mn(III) percentages (Figure 4a)

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estimated by the XANES-LCF analysis (Figure S6 and Table S2) and to the amounts of the

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decrease in the PDF peak area at 5.3 Å (Figures 3a, S3e,f, and S4a). However, the solids in the Li+

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system still contain high numbers of Mn(III) (26% at 25 d), which is very close to that in the Na+

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system (Figure 4a). Providing that the XRD peaks of T-bir are relatively weaker in the Li+ system

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than in the Na+ system, we may conclude that a small portion of Mn(III) adsorbs on vacancies

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without being incorporated into the layers, i.e., formation of Mn(III)-rich -MnO2.16

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The formation of the TMO phase occurs earlier in both Li+ and K+ systems than in the Na+

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system, which is clearly indicated in the earlier increase of the PDF peak area at 5.3 Å (Figure 3a)

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and earlier appearance of needle-like crystals (Figure S2). The PDF peak area at 5.3 Å increases

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because TMO contains abundant Mn-Mn corner-sharing atomic pairs with an interatomic distance

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of 5.3 Å.16 The faster transformation to TMO in the Li+ and K+ systems than that in the Na+ system

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can be ascribed to the formation of the Mn(III)-rich -MnO2 with Mn(III) adsorbed on vacancies,

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which acts as a viable precursor for the transformation.16 The transformation of well-crystallized

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T-bir to TMO in the Na+ system is slower. The XRD peak intensities of T-bir decrease significantly

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in both Na+ and K+ systems at 25 d, indicating more extended transformation of T-bir to TMO. In

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contrast, T-bir persists in the Li+ system, probably due to its very high crystallinity. The greater

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increase of the PDF peak area at 5.3 Å (Figure 3a) and more needle-like particles at 25 d (Figure

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2) in the K+ system suggest that more TMO forms in the K+ system than in the Na+ and Li+ systems.

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However, the strong XRD peak at 12 Å suggests that the TMO formed in the Na+ system has the

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highest crystallinity (Figure 1a), which could be due to the slower transformation with well-


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crystallized T-bir precursor, favoring formation of the well-crystallized TMO.

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Raman spectroscopy was used to further identify the solid phases. In the Li+ system, Raman

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spectrum of the 10-h solid has two dominant bands at 574 and 651 cm–1 (Figure 5), indicating a

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layered structure.54 The spectrum of the 25-d solid has an absorption band around 730 cm–1,

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probably corresponding to tunneled structures.54,55 Similarly, the solids at 75 h or 10 h in both Na+

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and K+ systems show more spectral features of δ-MnO2 or T-bir, while the 25-d samples have

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more spectral features of tunneled structures (Figure 5). These results agree with the XRD analyses

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(Figure 1a).

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Concentrations of monovalent cations also affect the Mn(II)-induced birnessite

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transformations. Higher NaCl concentration favors the formation of T-bir. Less T-bir forms in 10

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mM NaCl than in 100 mM NaCl solution (Figure 1a,b); and well-crystallized T-bir is not detected

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in 0 mM NaCl solution (Figure 1b). However, poorly crystalline T-bir or Mn(III)-rich -MnO2

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might be abundant in the solutions of low NaCl concentrations, indicated by the sharp decrease of

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PDF peak area at 5.3 Å within 5 h (Figure 3b). The favorability of high NaCl concentration for

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T-bir formation could be caused by a sufficient number of Na+ that is needed to compensate the

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negative layer charges of T-bir and by a higher ionic strength that favors particle aggregation and

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thus particle growth via oriented attachment.32,56 4 × 4 TMO forms at all NaCl concentrations,

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starting earlier at the lower NaCl concentration (Figure S2), while at 25 d, the transformation to

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TMO is more extended and the TMO is more crystalline at the higher NaCl concentration. The

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earlier formation of TMO is consistent with the abundance of poorly crystalline T-bir and/or

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Mn(III)-rich -MnO2 as the precursors.16 Once primary TMO particles form, the growth of TMO

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crystals via oriented attachment9 can be favored at higher ionic strength,56 leading to more

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extended transformation to the TMO at 25 d in the 100 mM NaCl solution. The transformations


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are consistent with the changes of the PDF peak area at 5.3 Å (Figures 3 and S3c,d,f).

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The observed differences among Li+, Na+ and K+ systems may be caused by their different

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magnitudes of electrostatic attraction with MnO6 layers. The ionic potential, φH = Z/rH, where rH

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is the radius of hydrated cations,57 measures the charge density of a cation,58 and the electrostatic

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forces between these hydrated cations and birnessite layers increase in the order of Li+ < Na+ < K+

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(Table S3). However, except for the crystallinity of T-bir, other aspects of the transformation, such

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as the amounts of T-bir and TMO and their formation rates, do not correlate with the order of the

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ionic potentials, suggesting existence of other controlling factors, which could be hydration energy

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and ionic radius of cations, in addition to electrostatic interactions.

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Effects of Mg2+ and Ca2+. Different from alkali metal cations, both Mg2+ and Ca2+ lead to

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formation of T-bir but not TMO. T-bir forms earlier, faster and more extensively and is much more

250

crystalline in the Mg2+ system than in the Ca2+ system (Figure 1c). Consistently, the peak area at

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5.3 Å in the PDFs decreases less in the Ca2+ system (Figure 3a). The continuous decrease of PDF

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peak area at 5.3 Å in both systems (Figures 3a and S4b,c) is consistent with the XRD and TEM

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results showing that T-bir forms and increases in quantity and no tunneled phase forms. The Raman

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spectra of the solids in the Mg2+ system show a gradual transformation from δ-MnO2 to T-bir with

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incubation time (Figure 5). The 25-d solid in the Ca2+ system has weak Raman spectral features of

256

T-bir. None of the solids in Mg2+ nor Ca2+ systems possess spectral features of tunneled structures

257

(Figure 5).

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The disfavored formation of T-bir in the Ca2+ system might result from the exchange of

259

adsorbed Mn(II) by Ca2+ so that less Mn(III) formed via the Mn(II)-Mn(IV) comproportionation

260

reaction.59 The formation of minor hausmannite (4.9 Å XRD peak in Figure 1c) suggests that the

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released Mn(II) may react with the edge sites of -MnO2 because the edge sites are negatively


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charged at pH 8 (pKa of Mn-OH groups on the edges is lower than 6).35,60 Due to weaker

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adsorption, Mg2+ may not exchange adsorbed Mn(II) as strongly as Ca2+ does; otherwise, the

264

Mn(III) percentage in the Mg2+ system would not reach 34% (Figure 4a), the percentage value of

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an ideal T-bir.61 In addition, Ca2+ can induce a drastic Mn(III) rearrangement in MnO6 layers and

266

interstratification of MnO6 layers with different Mn(III) distribution in T-bir,62 which may further

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decrease the crystallinity of T-bir.

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T-bir formed in the Mg2+ system has a high order of layer stacking, indicated by strong

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(00l) XRD peaks (Figure 1c), which can be caused by the strong hydrogen bonds between MnO6

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layers and hydrated Mg2+.46 The high hydration energy of Mg2+ prevents dehydration of Mg2+ upon

271

air drying, leading to the (003) diffraction peak at 3.2 Å.63,64 Compared to the high stacking order,

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the degree of the intra-layer ordering of the a-b plane of T-bir in the Mg2+ system is surprisingly

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low, as indicated by the broad XRD peaks between 2.4 – 1.4 Å (Figure 1c). The strong interaction

274

between Mg2+ and the layers may cause the poor Mn(III) arrangement in the layers.46

275

Effects of Cu(II). Mn(II)/Cu(II)-bearing δ-MnO2 does not undergo much phase

276

transformation. The XRD pattern of the 25-d solid is very similar to that of -MnO2 (Figure 1d),

277

and the Raman spectrum of the solid is dissimilar to both T-bir and TMOs (Figure 5), indicating

278

the absence of both T-bir and TMO phases. The 3Mn system uses a similar Mn(II)/MnO2 ratio as

279

the Cu system, but a small amount of poorly crystalline T-bir forms and decent layer stacking is

280

developed, as indicated by the (001) and (002) XRD peaks (Figure 1d) and the strong decrease of

281

the PDF peak area at 5.3 Å (Figures 3b and S3b). In spite of no mineral phase changes, the

282

weakened dip at 2.0 Å in the XRD patterns (Figure 1d) and the reduced PDF peak area at 5.3 Å

283

within 10 h (Figures 3b and S4d) suggest that a portion of Mn(II) is oxidized and incorporated into

284

vacancies.2,52,59 Thus, the Cu(II) system leads to formation of Mn(III)-rich -MnO2 which,


13


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285

however, does not transform to T-bir and 4 × 4 TMO, suggesting an inhibition of Cu(II) on the

286

transformation to T-bir and 4 × 4 TMO.

287

The speciation of Cu(II) in the solids provides insights into the inhibition of Cu(II) on T-

288

bir formation. Both Cu K-edge XANES and EXAFS spectra show significant differences between

289

the incubated samples and spertiniite (Cu(OH)2) and tenorite (CuO) references (Figure S7a,b),

290

indicating that Cu(II) does not precipitate although the pH is high (pH 8). The EXAFS spectra of

291

the incubated solids differ only slightly from the initial sample (Figure S7b), suggesting that the

292

Cu coordination environment remains largely unchanged during the incubation. EXAFS spectral

293

fits show that each Cu atom is surrounded by about 4 O atoms (CN = 3.7) at 1.95 – 1.96 Å (Figure

294

S7b,c and Table S4), corresponding to four equatorial Cu-O bonds of the Jahn-Teller distorted

295

CuO6 octahedron.65,66 The two axial O atoms were not included in the fitting because of the strong

296

thermal motion of the long axial bonds at room temperature.67 The second shell consists of 1.5 Cu

297

at 2.90 – 2.93 Å, likely due to the formation of polynuclear clusters;

298

exclude the possibility that a small number of Cu incorporates into layer vacancies. The small CNs

299

of the second edge-sharing Cu-Cu/Mn shell at 2.90 – 2.93 Å and the third corner-sharing Cu-

300

Cu/Mn shell at 3.38 – 3.40 Å suggest that Cu mainly adsorbs at edge sites as double edge-sharing

301

complexes and single/double corner-sharing complexes.37 The absence of T-bir can be ascribed to

302

the adsorption of Cu(II) clusters on the edge sites, leading to Mn(III) disordered arrangement in

303

the layers. The Cu(II) in layers can impair intra-layer electron transfer and thus Mn(III)

304

rearrangement as well.

37,47

however, we cannot

305

Role of Metal Adsorption on Stability of Layered Mn Oxides. The degree of the

306

transformation of Mn(II)-bearing -MnO2 decreases in the order of Li+, Na+ or K+ > Mg2+ or Ca2+

307

> Cu(II). The cation impacts on the transformation can be understood from thermodynamic


14

Page 15 of 35


308

perspectives by considering the contribution of surface energy to the stability of LMOs and how

309

much surface energy metal adsorption decreases.

310

LMOs have large total surface areas mainly from basal planes for well-crystalized LMOs

311

while from both edges and basal planes for small LMO particles. With decreasing basal plane size,

312

edge sites contribute increasingly more additional surface areas that can be substantial, such as for

313

-MnO2. Because of the high surface area to volume ratios of LMOs, surface energy contributes a

314

large portion of the total free energy (including both bulk and surface), elevating the total free

315

energy and probably driving their transformation to TMOs.68,69 Due to the large contribution from

316

edge sites, small LMO particles have high total free energy and readily transform to more

317

crystalline birnessite phase or to TMOs. For example, -MnO2 grows bigger in particle sizes via

318

oriented attachment70 or transforms to cryptomelane (2  2) at room temperature,71,72 while acid

319

birnessite with a larger particle size does not. However, ion adsorption on LMO surfaces (including

320

both edges and basal planes) can decrease the surface energy and thus the total free energy,

321

increasing their thermodynamic stability. The extent to which the surface energy can be decreased

322

depends on metal binding strength.

323

In the present work, Cu(II) adsorbs very strongly on LMO surfaces and the surface and

324

total free energy of LMOs can be decreased substantially, thereby stabilizing Mn(III)-rich -MnO2

325

without forming T-bir and 4 × 4 TMO (Figure 6). Similarly, the adsorption of transition metals

326

retards the transformation of LMOs to todorokite, a 3×3 TMO, under refluxing conditions.9,73-75

327

Mn(III)O6 is Jahn-Teller distorted, and its presence and ordered distribution in layers, such as in

328

T-bir, can increase the surface and total free energy of LMOs. Cu(II) adsorption decreases the

329

Gibbs free energy of Mn(III)-rich -MnO2,76 and the energy of Mn(III)-rich -MnO2 with Cu(II)

330

adsorbed is likely lower than that of T-bir, disfavoring the transformation of Mn(III)-rich -MnO2


15


Page 16 of 35

331

to T-bir. Thus, Mn(III)-rich -MnO2 still maintains its hexagonal layer symmetry (Figure 6). The

332

stabilization of hexagonal birnessite by transition metals is also evidenced in the previous studies

333

although unlike Cu(II), vacancies are the main adsorption sites of those metals.5,77 For example,

334

during the biogenic formation of nanoparticulate birnessite, Na+ and Ca2+ favor the formation of

335

T-bir while Ni(II) favors the formation of hexagonal birnessite.33 Zn(II) inhibits Mn(III)

336

accumulation in the structure of -MnO2 during its formation78 or during its partial reduction by

337

dissolved organic matter,39 and Zn(II) adsorption on -MnO2 causes migration of Mn(III) out of

338

layers.79 Note that metal adsorption may cause kinetic hindrance of the transformation of

339

hexagonal birnessite, which could also contribute to the increased stability of the LMOs.

340

As to Mg2+ and Ca2+, their binding to LMOs is much weaker than those of transition metal

341

cations. Thus, Mn(III)-bearing -MnO2 proceeds to form T-bir. The high crystallinity and

342

abundance of T-bir in the Mg2+ system are consistent with its weaker binding strength. However,

343

the adsorption of Ca2+ and Mg2+ in the interlayers stabilizes T-bir, retarding or preventing its

344

transformation to TMOs at room temperature (Figure 6). High temperature would be required for

345

transformation of Mg2+-exchanged T-bir to TMOs.80 Monovalent cations Li+, Na+ and K+ weakly

346

bind to LMOs and their adsorption cannot much decrease the surface and total free energy of

347

Mn(III)-bearing -MnO2. Thus, in their solutions, Mn(III)-rich -MnO2 and T-bir readily form but

348

are unstable and eventually transform to the TMO (Figure 6). However, other environmental

349

factors may also contribute to the high abundance of LMOs in the natural environment, such as

350

the adsorption of oxyanions and the rapid formation of biogenic birnessite.

351

Environmental Implications. Layered Mn oxide minerals are common in the natural

352

environment and play important role in elemental cycles and pollutant dynamics. They have large

353

surface areas on basal planes and edges and are not thermodynamically stable and can transform


16

Page 17 of 35


354

to TMOs. Our findings suggest that the unusually long lifetimes of LMOs in the natural

355

environment may be caused by the adsorption of divalent cations, probably also trivalent (e.g.,

356

Fe(III)), particularly transition metals, that increases their thermodynamic stability by decreasing

357

their surface energies. Thus, the stabilization of LMOs by metal adsorption may be an important

358

cause for the overall higher abundance of LMOs than TMOs in the natural environments. The

359

increased stabilization by metal adsorption may also decrease their redox potentials because

360

surface energy shifts redox equilibria.81 Our results have important implications for understanding

361

environmental and geochemical impacts of Mn oxides because LMOs are more reactive than

362

TMOs in terms of oxidation and metal adsorption reactivity.82 Effects of other environmental

363

factors, such as adsorption of oxyanions, on the stability of LMOs need to be addressed in future

364

studies.

365

ASSOCIATED CONTENT

366

Supporting Information.

367

The Supporting Information are available free of charge at http://pubs.acs.org, including

368

the details of preparation of δ-MnO2, TEM characterization, and XRD and Raman spectroscopic

369

analyses; XRD patterns, TEM images, PDF data, XANES and EXAFS spectra and corresponding

370

fitting results; dissolved Mn and Cu concentrations; and physicochemical atomic parameters of

371

alkali and alkaline earth metal cations.

372

Corresponding Author

373 374 375

*M. Zhu. E-mail: [email protected]; Phone: 307-766-5523 Notes The authors declare no competing financial interest.


17


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Page 18 of 35

ACKNOWLEDGEMENTS

377

This study was supported by the U.S. Department of Energy Experimental Program to

378

Stimulate Competitive Research Office for financial support (DOE-EPSCoR DE-SC0016272). We

379

thank Dr. Xiaoming Wang in the College of Resources and Environment, Huazhong Agricultural

380

University, Wuhan, China for providing Cu K-edge X-ray absorption spectra of spertiniite

381

(Cu(OH)2) and tenorite (CuO). This work utilized resources of APS, a U.S. DOE Office of Science

382

User Facility, operated for the DOE Office of Science by Argonne National Laboratory under

383

Contract No. DE-AC02-06CH11357. Use of SSRL, SLAC National Accelerator Laboratory, was

384

supported by the U.S. DOE, Office of Science, Office of Basic Energy Sciences, under Contract

385

No. DE-AC02-76SF00515.

386

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ray scattering and transmission electron microscopy perspective on reaction mechanisms.

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Geochem. Trans. 2015, 16 (1), 12.

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(73) Cui, H.; Feng, X.; Tan, W.; He, J.; Hu, R.; Liu, F. Synthesis of todorokite-type manganese

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oxide from Cu-buserite by controlling the pH at atmospheric pressure. Micropor. Mesopor. Mater.

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2009, 117 (1), 41-47.

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(74) Cui, H.-J.; Liu, F.; Tan, W.-F.; Feng, X.-H. Effect of cobalt-doped framework on formation

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of todorokite from layered manganese oxides with Mg2+/Co2+ ions as template. Pedosphere 2011,

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21 (6), 730-737.

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(75) Wu, Z.; Peacock, C. L.; Lanson, B.; Yin, H.; Zheng, L.; Chen, Z.; Tan, W.; Qiu, G.; Liu, F.;

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Feng, X. Transformation of Co-containing birnessite to todorokite: Effect of Co on the

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transformation and implications for Co mobility. Geochim. Cosmochim. Acta 2019, 246, 21-40.

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Results from powder and polarized extended X-ray absorption fine structure spectroscopy.

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Cosmochim. Acta 2012, 85, 302-313.

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Page 28 of 35

Figures (001)

2.4 1.4

(001)

(a)

2.4

(b) 1.4

4.7 (002)

4.7

10Na 25d

Li 25d

10Na 15d

Intensity

Li 15d Li 5d Li 50h Na 25d Na 15d

(002)

10Na 5d 10Na 50h 10Na 10h 0Na 25d

Na 5d Na 50h

0Na 15d 0Na 5d

K 25d K 15d K 5d 10Mn Init As-prep

2.0 4.6

12

0Na 50h

(001) (002)

6 (001)

(c)

2.4

As-prep

4.6

12

6

1.4

2.4

(d) 1.4

(002)

(003)

Mg 25d

3Mn 25d

Mg 15d

3Mn 15d

Intensity

Mg 5d 2.0

Mg 50h Ca 25d

4.6

Cu 25d

Ca 15d Ca 5d

2.0

10Mn Init

12

587 588 589 590 591 592 593 594 595 596 597

12

6

Cu 15d Cu Init As-prep

As-prep

4.6

3Mn 5d 3Mn Init As-prep

4.9

2.0

0Na 10h 10Mn Init

2.0

6

d-spacing (Å)

d-spacing (Å)

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 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 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 panels a and b and with incubated samples in the Ca2+, 3Mn, and Cu(II) systems in panels c and d to illustrate the changes of XRD patterns.


28

Page 29 of 35


598 599

Figure 2. The TEM images of starting materials and 25-d samples from different reaction systems.

600

As-prep and 10Mn Init stand for as-prepared δ-MnO2 and that equilibrated with 10 mM Mn(II) at

601

pH 4 but without being incubated at pH 8, respectively.


29


0.05

0.04

As-prep Li Na K Mg Ca

0.02

0.01

0.0

0.1

0.2

0.3

Page 30 of 35

(a) 0.05

0.04

As-prep 0Mn 3Mn 0Na 10Na Cu

0.02

0.01

0.4

0.0

0.03

0.02

0.02

0.01

0.01

0.2

0.3

0.4

Area

0.03

0.1

(b)

0.00

0.00 0

602

5

10

15

25

0

Time (d)

5

10

15

25

Time (d)

603

Figure 3. Changes of the PDF peak area at 5.3 Å with time in different reaction systems. The

604

insets show the data points within 10 h. As-prep stands for as-prepared δ-MnO2.


30

Page 31 of 35


(a)

8

Mn(II) percentage (%)

30

Mn(III) percentage (%)

(b)

10

35

25

20

As-prep 0Mn 3Mn

15

Li Na Mg

6 4 2 0

10 0

605

1

2

3

25

0

Time (d)

3

6

9

12

24

Time (d)

606

Figure 4. XANES-LCF derived Mn(III) (a) and Mn(II) (b) molar percentages for selected reaction

607

systems. As-prep stands for as-prepared δ-MnO2.


31


Page 32 of 35

730

Woodruffite Todorokite T-bir Cu 25d

Intensity (counts)

Ca 25d Mg 25d Mg 15d Mg 5d Mg 50h Mg 10h K 25d K 10h Na 25d Na 75h Li 25d Li 10h 574 651

10Mn Init As-prep

100 200 300 400 500 600 700 800 900 1000 –1

608

Wavenumber (cm )

609

Figure 5. Raman spectra of T-bir, woodruffite (3 × 4), todorokite (3 × 3), and selected samples

610

from the Li+, Na+, K+, Mg2+, Ca2+, and Cu(II) systems.


32

Page 33 of 35


Triclinic birnessite

4 × 4 TMO

δ-MnO2 Mn(III)-rich δ-MnO2 +

Mn(II)

+,K

a

+,N

Li

Mn(II)-bearing δ-MnO2

Mg2+, Ca2+

Cu

Triclinic birnessite

Mn(II) Mn(III)

(II )

Mn(III)-rich δ-MnO2

Mn(IV) Cu(II)

611 612

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

613

systems.


33


Page 34 of 35

614

Table(s)

615

Table 1. A summary of the experimental conditions. All experiments were carried out using

616

δ-MnO2 (4 g/L) as starting material at pH 8 in a N2/H2 atmosphere.

Condition

Initial Mn(II) concentration (mM)

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

Background electrolyte

Ionic strength (mM)

0Mn

0

0

100 mM NaCl

100

3Mn

3

0.08

100 mM NaCl

100

0Na

10

0.14

0 mM NaCl

0

10Na

10

0.14

10 mM NaCl

10

Li+

10

0.14

100 mM LiCl

100

Na+

10

0.14

100 mM NaCl

100

K+

10

0.14

100 mM KCl

100

Mg2+

10

0.14

33.3 mM MgCl2

100

Ca2+

10

0.14

33.3 mM CaCl2

100

Cu(II)

5 mM Mn(II) +

0.08

100 mM NaCl

100

7 mM Cu(II)

617


34

Page 35 of 35

618


TOC

619


35

Metal Adsorption Controls Stability of Layered Manganese Oxides

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