The Pennsylvania State University

The Graduate School

College of Earth and Mineral Sciences

IN SITU pE-pH ANALYSIS ON THE OXIDATION KINETICS OF MANGANESE

BEARING SOLUTION

A Thesis in

Geosciences

by

William C. Ethier Colón

© 2014 William C. Ethier Colón

Submitted in Partial Fulfillment of the Requirements

for the Degree of

Master of Science

December 2014

The thesis of William C. Ethier Colón was reviewed and approved* by the following:

Hiroshi Ohmoto Professor Emeritus of Geochemistry Thesis Adviser

Peter J. Heaney Professor of Mineral Sciences Associate Head of Undergraduate Programs

James Kubicki Professor of Geosciences

Demian Safer Professor of Geosciences Interim Associate Head of Graduate Programs and Research

*Signatures are on file in the Graduate School.

ii

ABSTRACT

The study of manganese(II) oxidation in oxygenated solutions has been well documented over

long timescales (> 30 days) and over varying temperature ranges. Experimental setup for this

study involved the addition of MnCl2∙4H2O solid to a basic (pH 9 - 10.25) aqueous solution to

produce an initial MnII concentration ranging 1 – 100 mM. We apply a method of in situ pH and

pE analysis to study the major reactions occurring during Mn oxidation. The experimental

system is not run under equilibrium conditions. pH is neither titrated nor buffered, where doing

so would interfere with the pH and pE of the solution. This type of analysis has not been

previously performed. Calculation of an experimental rate constant of the system confirmed that

changes in pE/pH slope of the solution reflected changes in the rate of the reaction. Changes in

rate of reaction are due to alteration of the major reaction taking place.

From solution chemistry, two major reaction stages were noted. Stage 1 can be further broken

down into stages 1A and 1B, with 1A lasting one minute and consisted of the MnII hydration

reaction to Mn(OH)2. Stage 1B continued with Mn hydration as well as the slower 1 electron

oxidation Mn(OH)2 to Mn(OH)3. Both manganese hydroxides are amorphous phases, confirmed

by both XRD and elemental analysis. However with increasing extent of reaction, the

crystallinity of the solid increased. Since rate of reaction decreases with decreasing pH, and Mn

hydration is fast, the expected transformation pathway for the production of observed groutite,

feitknechtite and hausmannite is by oxidation of amorphous phases Mn(OH)2 to Mn(OH)3, and

transformation to crystalline phases. Specifically rate of production of feitknechtite is slightly

fast compared to groutite, but groutite is metastable and its rate of transformation to hausmannite

is faster than its rate formation. Comparatively, the transformation of feitknechtite to

iii

hausmannite is slower, therefore feitknechtite transformation primarily dictates the rate of

reaction.

[OH-] was found to be second order with respect to the rate of Mn oxidation, which is in

accordance with the rate law for Mn oxidation given by Stumm and Morgan, 1996: -d[MnII(aq)]/dt

= k1 [MnII(aq)][OH-

(aq)]2[O2 (aq)] + k2[MnII(aq)][MnOx (s)][OH-

(aq)]2[O2 (aq)]. However,

experimental MnII(aq) concentration data over time, gathered from this study revealed [MnII] to be

second order with respect to rate of reaction, which countered Stumm and Morgan, (1996). [OH-]

was noted to be the major contributor to the Mn oxidation. From this, an initial approximation of

the rate constant was calculated as Mn oxidation reaction to be pseudo-second order with respect

to [OH-], written as –d[MnII]/dt = kOH[O2][OH-]2 . kOH was then calculated and used to iteratively

solve for the apparent rate constant for the overall rate equation, kapp = [MnII]2[O2][OH-]2. kapp for

the oxidation of Mn was determined to be 6.2*1015 M-4 Hr-1. Complete comparison to the

Stumm and Morgan, (1996) equation was not possible as BET surface area studies were not

performed during this study. However, [MnII] over the course of the experiments was never

reduced more than by 33%, suggesting Mn underwent autocatalytic oxidation over the course of

the experiment.

iv

TABLE OF CONTENTS

LIST OF EQUATIONS . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . vii LIST OF TABLES . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .ix LIST OF FIGURES . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . x

ACKNOWLEDGEMENTS . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . xii 1. INTRODUCTION. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .1 1.1. Background. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .1 1.2. Objectives. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .2 2. METHODS. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5 2.1. Water Equilibration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5 2.2. Experimental set-up . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .6 2.3. Determination of MnII in solution and calculation of precipitated Mn . . . . . . . . . . . . 7 2.4. Mn solid collection . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8 2.5. XRD Solid sample preparation and analysis. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9 2.6. Determining wt. % Mn content in collected solids. . . . . . . . . . . . . . . . . . . . . . . . . . .10 3. RESULTS . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .11 3.0. Visual evidence for Mn oxidation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11 3.1. Variation of initial MnII concentration. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .14 3.2. Effect of dissolved oxygen equilibrium reaction on slope. . . . . . . . . . . . . . . . . . . . .16 3.3. XRD Spectra. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .. . . . . . . . . . . 17 3.3.1. Determination of relative precipitate degree of crystallinity . . . . . . . . . . . 18 3.3.2. Feitknechtite over time. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .18 3.3.3. Groutite over time. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19 3.3.4. Hausmannite over time. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .. . . . . . . . .19 3.4. Mn content analyses. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .. . . . . . . . . . . . . .23 3.4.1. Mn content analysis of precipitated solid. . . . . . . . . . . . . . . . . . . . . . . . . . 23 3.4.2. Precipitation rate vs. wt. % Mn in precipitate. . . . . . . . . . . . . . . . . . . . . . . 25 3.5. Changes in the pH and pE of experimental solutions: experimental series 9 - 13. . . 26 3.5.1. Experimental series 9 – 11, 75 hour overlay. . . . . . . . . . . . . . . . . . . . . . . . 26 3.5.2. Experimental series 12 and 13: investigation of micro slopes. . . . . . . . . . .27 4. DISCUSSION. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 29 4.1. Solution chemistry reaction stages. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 29 4.2. Determining reaction rate order . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .30 4.3. Mineral phase production over time . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 37 4.4. Future Studies. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .40

v

5. CONCLUSION. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 40 REFERENCES . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .41

APPENDIX A: Supplemental Tables. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 43 APPENDIX B: Supplemental Methods. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .46 APPENDIX C: Supplemental Photos. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 47 APPENDIX D: Supplemental Figures. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .52

vi

LIST OF EQUATIONS

Equation 1: Stumm and Morgan Mn rate equation. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .1

Equation 2: Oxidation of water. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4

Equation 3: Oxidation of MnII(aq) – precipitation of MnIVO2 (s) . . . . . . . . . . . . . . . . . . . . . . . . . . . .4

Equation 4: Solid phase oxidation of MnIIIOOH (s) to MnIVO2 (s) . . . . . . . . . . . . . . . . . . . . . . . . . . 4

Equation 5: Oxidation of MnII(aq) – precipitation of MnIIIOOH (s) . . . . . . . . . . . . . . . . . . . . . . . . . 4

Equation 6: Triple point of MnII(aq), MnIIOMnIII

2O3 (s), MnIIIOOH (s) . . . . . . . . . . . . . . . . . . . . . . . 4

Equation 7: Oxidation of MnII(aq) – precipitation of MnIIOMnIII

2O3 (s) . . . . . . . . . . . . . . . . . . . . . .4

Equation 8: Solid phase oxidation of MnII(OH)2 (s) to MnIIOMnIII2O3 (s) . . . . . . . . . . . . . . . . . . . . .4

Equation 9: Hydration: MnII(aq) – MnII(OH)2 (s) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .4

Equation 10: Oxidation of hydrogen . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .4

Equation 11: Solid phase oxidation of MnII(OH)2 - β-MnIIIOOH (s) & MnIIOMnIII2O3 (s) . . . . . . . .4

Equation 12: Solid phase oxidation of MnII(OH)2 - β-MnIIIOOH (s) . . . . . . . . . . . . . . . . . . . . . . . .4

Equation 13: Triple point of MnIVO2 (s), MnII(aq), β-MnIIIOOH (s) . . . . . . . . . . . . . . . . . . . . . . . . . . 4

Equation 14: ORP to Eh conversion. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5

Equation 15: Eh to pE conversion. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .5

Equation 16: Calculated mass Mn precipitate. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .8

Equation 17: Collected weight percent Mn content in precipitate. . . . . . . . . . . . . . . . . . . . . . . . 10

vii

Equation 18: Calculated weight percent Mn in precipitate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10

Equation 19: Calculated amount of solid precipitate. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10

Equation 20: Calculated mass of observed mineral phases. . . . . . . . . . . . . . . . . . . . . . . . . . . . . .11

Equation 21: Mn mass in observed mineral phases. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .11

Equation 22: log Keq. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .16

Equation 23: Nernst equation. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .16

Equation 24: Solid phase oxidation of Mn(OH)2 (s) to Mn(OH)3 (s). . . . . . . . . . . . . . . . . . . . . . . . 29

Equation 25: kapp. . . . . . . . . . . . . . . . . . . . . .. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .34

Equation 26: kOH. . . . . . . . .. . . . . . . . . . . . . . . . . . . . . . . . .. . . . . . . . . . . . . . . . . . . . . . . . .. . . . .34

Equation 27: Solid phase oxidation of Mn(OH)2 (s), Mn(OH)3 (s) to MnOOH (s). . . . . . . . . . . . . . .37

Equation 28: Solid phase oxidation of Mn(OH)2 (s), Mn(OH)3 (s) to Mn3O4 (s) . . . . . . . . . . . . . . . .37

Equation 29: Triple Point of MnII(aq), MnOOH (s), and Mn3O4 (s . . . . . . . . . . . . . . . . . . . . . . . . . 37

viii

LIST OF TABLES

Table 1: Pertinent reactions. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4

Table 2: Rate constant approximation. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .35

Supplemental table 1: Experimental solution data. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 43

Supplemental table 2: Experimental solid phase data. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .44

Supplemental table 3: Mineral phase XRD peak list. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .45

ix

LIST OF FIGURES

Figure 1: Time dependent solution color change. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .12

Figure 2: Visual opacity increase over time . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .12

Figure 3: Visual extent of reaction. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .13

Figure 4: pH-pE diagram. [MnII] variation pHi = 9. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15

Figure 5: XRD experiments 9 – 11, 4 hour overlay. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .21

Figure 6: XRD experiments 9 – 11, 40 hour overlay. . . . . . . . . . . . . . . . . . . . . . . . . . .. . . . . . . . .22

Figure 7: XRD experiments 9 – 11, 75 hour overlay . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 23

Figure 8: Mn in solution and calculated Mn in solid over time. . . . . . . . . . . . . . . . . . . . . . . . . . .24

Figure 9: Solid precipitation rate vs. wt. % Mn in solid. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 25

Figure 10: pH-pE diagram. 75 Hour Overlay. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .26

Figure 11: pH-pE diagram. Experiments 12 and 13, pH 10, 10.25. . . . . . . . . . . . . . . . . . . . . . . .28

Figure 12: Inverse concentration over time. . . . . . . . . . . . . . . . . . . . .. . . . . . . . . . . . . . . . . . . . . 31

Figure 13: Stage 1, 2 comparison of [OH-]-1 vs. time. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 31

Figure 14: Apparent rate/[MnII] vs. pH. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 36

Figure 15: Figure 15: Mineral phase transformation scheme. . . . . . . . . . . . . . . . . . . . . . . . . . . .37

Figure 16: Mineral precipitation rate and slope vs. pH. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39

Supplemental figure 1: Experimental method schematic. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .46

x

Supplemental figure 2: Additional experiment photos. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .47

Supplemental figure 3: Individual experiment pH-pE data. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 52

Supplemental figure 4: Individual experiment slope over time. . . . . . . . . . . . . . . . . . . . . . . . . . .61

Supplemental figure 5: Individual experiment XRD spectra. . . . . . . . . . . . . . . . . . . . . . . . . . . . .63

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ACKNOWLEDGEMENTS

I would like to give specific thanks: my advisor Dr. Hiroshi Ohmoto whose guidance and

feedback was invaluable, and Dr. Akane Miyazaki, without loaning me her pH, ORP and DO

meters this project could not have been accomplished. Further thanks go to my committee

members, Peter Heaney and James Kubicki, my research group Andrew Chorney and Jamie

Brainard. Finally, I would like to thank the funding sources NSF Sulfur Isotope 424-07 67NF0,

PSARC Signatures 424-07 58DF, Penn State Foundation, and Diodato Geosciences Fund that

made this research possible.

xii

1. INTRODUCTION

1.1. Background

Understanding the oxidation of manganese is important in studying the conditions that produce

unwanted manganese precipitates in water supplies as well as in studying the formation of

ferromanganese nodules in the deep ocean. Aqueous oxidation of manganese(II) has been

studied extensively. Past Mn oxidation studies determining the rate of reaction have been

performed at constant pH, either by titration with weak base or in the presence of a buffer

solution and bubbling O2 into solution. Coughlin and Matsui, (1976) modeled Mn oxidation as a

function of amount base added over the course of the experiment. Approximately 2.8 mole H+

were produced per mole Mn consumed. Near the end of the reaction (> 80% Mn consumed) the

ratio dropped to ~ 1.8 mole H+ per mole Mn. A possible net reaction for producing this observed

MnII/H+ ratio could be the oxidation/precipitation of hausmannite, Mn3O4 (see table 1, equation

7). Kinetic studies concluded MnII undergoes oxidation autocatalytically when pH is kept

constant (Morgan 1967; Hem 1981):

-d[MnII(aq)]/dt = k1 [MnII

(aq)][OH-(aq)]2[O2 (aq)] + k2[MnII

(aq)][MnOx (s)][OH-(aq)]2[O2 (aq)] (1)

where -d[MnII(aq)]/dt is the rate change in MnII

(aq) concentration over time in M*d-1, [MnII(aq)] is

the concentration of MnII(aq) in Molar, and [MnOx (s)] is the concentration of precipitated Mn

oxide in solution in Molar. Rate constants were later reported: k1 = 4 x 1012 M-3 d-1 and k2 = 1018

M-4 d-1 at pH 9 and p[O2] of 1 atm (Diem and Stumm, 1984). The k1 term describes

homogeneous oxidation of MnII(aq) and the k2 term describes heterogeneous oxidation occurring

on manganese oxide surfaces. When [MnOx (s)] > [MnII(aq)] the reaction is autocatalytic. When

[MnOx (s)] < [MnII(aq)], it can be treated as pseudo first-order, written as -d[MnII

(aq)]/dt = k1

1

[MnII(aq)][OH-

(aq)]2[O2 (aq)]. The rate of MnII oxidation can also be influenced positively by the

presence of other hydrous metal oxides. MnII oxidation half-life reported to reduce from

hundreds of hours to hours in the presence of lepidocrocite, γ-FeIIIOOH (Sung and Morgan 1981;

Matsui 1973; Coughlin and Matsui 1976).

Variation of reaction conditions may alter the relative ratios of produced Mn oxides. Mn oxides

produced from oxidation by O2 have been identified as hausmannite, sometimes accompanied by

an unspecified form of MnIIIOOH and birnessite, MnO2. Bubbling of a solution with O2 yielded

more prevalent manganite, γ-MnIIIOOH. Acidification of suspended Mn3O4 in the presence of O2

produced birnessite (Bricker, 1965; Hem, 1978). Feng et al., (2004) performed a synthesis of

birnessite by mixing anoxic 200 mL 0.5 M MnCl2 solution with oxygenated 250 mL 5.5 M

NaOH and bubbling with O2 and concluded two possible pathways of mineral formation:

(1) Mn(OH)2 (amorphous) feitknechtite, β-MnIIIOOH buserite, Na4Mn14O27 ·21H2O birnessite, (Na0.3Ca0.1K0.1)(MnIV,MnIII)2O4·1.5 H2O (2) Mn(OH)2 (amorphous) hausmannite, γ-Mn3O4 feitknechtite buserite birnessite.

Increasing O2 flow rate or stirring rate increased the Mn(OH)2 oxidation rate. At 3.0 L/min O2

flow rate birnessite was the only phase observed in the final product. When O2 flow rate or

stirring rate were reduced, feitknechtite and hausmannite were observed along with the major

phase birnessite. The importance of feitknechtite as an intermediate Mn oxide phase has been

proposed in the literature, where feitknechtite is considered a meta-stable intermediate (Elzinga

et al., 2013).

1.2. Objectives

Constraining the pertinent Mn oxidation and solid phase transformation reactions responsible for

the observed mineral phase alteration over time has not been performed. This study attempts to

2

achieve this by correlating changes in accumulated amounts of observed mineral phases to

changes in pH and pE, a proxy for monitoring solution chemistry. Such a study cannot be

completed at constant pH as Mn oxidation generates H+ and e-, thus altering pH and pE.

In this experiment, it is hypothesized that long-term variation of pE/pH slope, over the entire

course of the experiment, logs the extent of reaction of Mn oxidation. Micro-scale alterations to

pE/pH slope, on the order of ~30 minute time-scales reflect alteration of the major redox reaction

taking place. Alteration of the major reaction should also reflect as a change in the accumulated

mineral phase ratios over time. If the experimental system is simply a homogenous oxidation of a

single Mn oxide phase, without further solid phase transformation, over short-term intervals

should not exhibit changes in slope. However, if the alteration of the major reaction is taking

place, as is expected, then short-term slope variation will occur and fluctuate according to the

relative rate of that major reaction. Under these conditions, solutions are supersaturated with

respect to [MnII(aq)] and precipitate Mn oxides. This experimental system is not at equilibrium.

Reaction kinetics dominates the production of mineral phases and subsequent transformation

reactions that may over the course of the experiment. Comparison of reaction rates, non-

oxidative precipitation should occur fastest followed by aqueous phase oxidative precipitation,

and solid phase transformations and oxidations.

Table 1 provides Mn redox reactions that are pertinent to the experimental system and are used

to construct the Pourbaix diagram for Mn.

3

Table 1: Pertinent reactions. Reactions 2 – 5, 7 - 10 are the equilibrium reactions present in the Pourbaix diagram (figure 4).

[MnII] = 100 mM Reaction Nernst

Equation Reaction Slope Eqn

# 2H2O(l) ⇌ O2 (aq) + 2H+

(aq) + 2e- pE =21.03 – pH (1 atm ⇌ PO2) pE =20.83 – pH (.2 atm ⇌ PO2)

Oxidation of Water

-1 2a, b

MnII(aq) + 2H2O(l) ⇌ MnIVO2 (s) + 4H+

(aq) + 2e-

pE =21.21 – 2pH

Oxidative Precipitation

-2 3

MnIIIOOH (s) + 2H2O (l) ⇌ MnIVO2 (s) + H+

(aq) + e- pE =16.538 – pH

Solid Phase Oxidation

-1 4

MnII(aq) + 2H2O(l) ⇌ MnIIIOOH(s) +

3H+(aq) + e-

pE =25.82 – 3pH

Oxidative Precipitation

-3 5

2MnIIIOOH (s) + MnII(aq) ⇌

MnIIOMnIII2O3 (s) + 2H+

(aq) pE =12.97 – pH Triple Point -1 6

3MnII(aq) + 4H2O(l) ⇌ MnIIOMnIII

2O3 (s) + 8H+

(aq) + 2e- pE =32.324 - 4pH

Oxidative Precipitation

-4 7

3MnII(OH)2 (s) ⇌ MnIIOMnIII2O3 (s) +

2H2O(l) + 2H+(aq) + 2e-

pE =7.879 – pH Solid Phase Oxidation

-1 8

MnII(aq) + 2H2O(l) ⇌ MnII(OH)2 (s)

+2H+(aq)

pH = 8.22 Hydration/ Dehydration

0 9

H2 (g) ⇌ 2H+(aq) + 2e- pE =-pH Oxidation of

Hydrogen -1 10

4MnII(OH)2 (s) ⇌ β-MnIIIOOH (s) + MnIIOMnIII

2O3 (s) + 2H2O (l) + 3H+(aq) +

3e-

pE =10.42 – pH Solid Phase Oxidation

1 11

MnII(OH)2 (s) ⇌ β-MnIIIOOH(s) + H+(aq) +

e- pE =46.6 - pH Solid Phase

Oxidation -1 12

MnIVO2 (s) + MnII (aq) + H2O (l) ⇌ 2β-

MnIIIOOH (s) +2H+(aq)

(Elzinga et al., 2011)

Triple Point 0 13

4

2. METHODS

The experimental method has been summarized in Appendix B, supplemental figure 1.

2.1. Water equilibration

All experiments were performed in a glovebox with an atmosphere of 80/20 N2/O2, void of CO2

to remove its buffering effect on pH. 18 Ω water was left stirring to equilibrate in the glove box

for 100 hours, wherein CO2 degassed from the solution and raised pH from 5.8 to 7.8. O2

dissolved raising the pE of the solution from 7.0 to 9.1. The equilibrated water was then filtered

through a 0.22 μm Omnipore© PTFE filter to remove any solids present in the solution. The

filtered water was titrated using a pH 11.8 NaOH solution to a desired pH value (experiments 3,

6 and 10, titrated to pH 9; experiments 9, titrated to pH 9.5; experiments 11 and 12, titrated to pH

10l and experiments series 13, titrated to pH 10.25). A Laqua Benchtop model F-74 monitored

pH and ORP. A Laqua model F-55 DO meter monitored dissolved oxygen content. The pH

probe was calibrated using pH standards of 4, 7, 10 provided by Acros Organics. The ORP probe

was calibrated to +220 mV at 25 oC using an ORP solution standard provided by Thermo

Scientific. After titration, the water was left to equilibrate further the glove box atmosphere for

another 24 hours. After 24 hours, the water was capped and stored. ORP values were converted

to pE following equations 14 and 15, based on the half-cell potential for a platinum electrode

(Van Loon and Duffy, 2011):

Eh (V) = (ORP + 220 - .7*(T-25))/1000 (14) pE = Eh* 𝐹𝐹

2.303𝑅𝑅𝑅𝑅 (15)

5

ORP is the oxidation-reduction potential in mV. Eh is the reduction potential for the solution in

Volts, and T is temperature in oC. In equation 15, F is the Faraday constant in s∙Amol

, R is the gas

constant in Jmol∙K

, and T is the solution temperature in K.

2.2. Experimental set-up

pH, pE, and DO data was collected in situ for the solution prior to the addition of MnII. MnII was

added to the solution in the form of solid manganese (II) chloride tetrahydrate, MnCl2∙4H2O,

supplied by Acros Organics. Experimental series 3, 6, 9-13 had initial MnII concentrations of 1

mM (54.93 ppm), 10 mM (549.3 ppm), and 100 mM (5493 ppm), respectively. For experimental

series 3, 6, 9 – 11, 1 L of solution was tested while experiments 12 and 13 used 500 mL per

experiment. Time zero was defined immediately before MnCl2∙4H2O addition. pH, and pE was

recorded over the entirety of the experiment, while DO was recorded at the start and end of each

experiment. Experiments ranged 2.77 – 206 hours.

The purpose of these experiments to compare changes in solution conditions over time to

changes in gathered precipitated solid samples. Because each collection of solid material was a

destructive process, three separate experiments for each pH value were performed for

experimental series 9, 10 and 11. Experiments were run at initial pH values of 9.5 (experiment

9), 9 (experiment 10) and 10 (experiment 11) all with an initial concentration of 100 mM MnII.

Each experimental series was run for 4, 40 and 75 hours. Experiments 12 and 13 studied the

variations in micro-slope over the course of Mn oxidation for the first 25 hours at pHi = 10 and

10.25, respectively.

pE and pH values of the solutions altered with time due to the poisoning of the electrodes by Mn

oxides precipitating on to the glass surfaces This introduced a pH and pe error of ±1%.

6

At the end of each experiment, the beaker, stir bar and filtration apparatus were cleaned in a bath

of 10 wt. % hydroxylamine hydrochloride solution for 24 hours, followed by 2 acid baths of 10

wt. % nitric acid, 24 hours each. The pH, ORP and DO probes were cleaned of any visible Mn

oxide that had precipitated on the glass surfaces of the probes in a 1M HCl bath for 30 minutes.

2.3. Determination of MnII in solution and calculation of precipitated Mn

Quantification of [MnII(aq)]f determined the amount Mn precipitated over the course of the

reaction. First, the solution at the end of the experiment was massed to account for water loss due

to evaporation, which was < 3% for any experiment. The average evaporation rate for all

experiments was 0.34 ±.1 mL/Hr. A 1.5 - 2 mL aliquot of solution was taken and filtered through

a 0.22 μm nylon syringe filter, provided by VWR, to remove any solid particles in solution. It

was assumed Mn particles < 0.22 μm consisted solely of MnII(aq). 1 mL of the filtered solution

was then acidified with 1 mL of concentrated hydrochloric acid and was left to reaction in a 50

oC water bath and sonicated overnight.

For experiments 9 – 11, the digested solutions underwent a 100,000x serial dilutions with 2%

ultrapure nitric acid, provided by Sigma Aldrich, to an estimated concentration of 55 ppb.

Experiments 9 – 11 were analyzed for manganese concentration using a Thermo X-Series II

Quadruple ICP-MS. For experiments 12 – 13, the digested solutions underwent a 1000x serial

dilution with 2% ultrapure nitric acid to an estimated concentration of 5 ppm. Experiments 12 –

13 were analyzed with a Perkin-Elmer Optima 5300DV ICP-AES. Mn calibration curve

standards were prepared from either a 10 or 100 ppm Mn stock solution provided by High Purity

Standards. Serial dilutions were performed to produce an overall calibration curve with solution

7

concentrations of 0, .01, .1, 1, 5, 10 ppb Mn for experiments 9 – 11 or 0, 0.01, .1, 1, 5, 10 ppm

for experiments 12 and 13.

The mass of Mn precipitated over the course of the reaction was defined as the difference

between the final and initial Mn content in solution, defined in the following equation 16.

Mass of Mn(ppt) = [Mn(initial, ppm) – Mn(final ppm)] /1000 gm (16)

2.4. Mn solid collection

The probes were rinsed with water to clear them of precipitated Mn oxide. The beaker containing

the entire experimental solution was then removed from the glove box and was quickly

vacuumed filtered through a 0.22 μm Omnipore© PTFE filter. It was assumed that solids > 0.22

μm were defined as solid Mn precipitate. The experimental beaker, stir bean and vacuum filter

funnel were rinsed repeatedly with water to remove precipitated solid from the walls. Only 1-2%

(~ 1 - 2 mg) of the entire solid particles were recovered from the experimental system; the

remaining 99% of the solid particles (~gram) probably remained on the beaker walls even after

rinsing. The major of the recovered solids consisted of suspended solids in solution, while solid

precipitated on the walls most likely precipitated prior to the precipitation of suspended solids.

Collected solid was then allowed to dry overnight at room temperature on the filter disk. The

difference in mass between the filter disk before and after filtration was logged as the amount of

collected manganese oxide solid. The dry solid was scraped off the surface of the filter disk and

was stored in a nitrogen atmosphere for future analysis.

8

2.5. XRD Solid sample preparation and analysis

Solid mineral phases were identified by X-ray diffraction and performing a whole pattern fitting

of the spectra to calculate the percent mineral content of the each phase present. Solid sample

was prepared by grinding collected solid was ground using an agate mortar and pestle and

acetone as lubricant. Once ground, the sample was loaded into a 0.81 mm internal diameter

polyimide capillary tube provided by Cole-Parmer. Both sides of the tube were capped with

modelling clay and were then ready for analysis. Spectra were taken using a Rigaku DMAX-

Rapid II microdiffractometer fitted with a Mo source and set to transmission mode. The detector

was a curved image plate. Collimator size was .3 mm. Each sample was run for 10 minutes. Φ

was rotated ±30o at an angular velocity of 1 degree per second. The output source was set to 50

kV at 40 mA. Collected XRD spectra were analyzed using Jade software and the observed

mineral spectra was matched to mineral phases present in the ICSD-Mineral library.

Feitknechtite, hausmannite, and groutite were observed with varying mineral percentages.

Relative intensity ratio (RIR) data were available for both hausmannite and groutite, however an

RIR value does not exist for feitknechtite (Amann et al., 1950; Ben Yahia et al., 2013). A RIR

value of 3.6 was assigned to feitknechtite and is the calculated average of RIR values for similar

Mn oxy(hydr)oxides. A whole pattern fit for each spectra was produced in Jade between 2Θ

values of 5o and 27o. Below 5o 2Θ, no discernable peaks were observed. Above 27o 2Θ, no

feitknechtite XRD data exists presently in the literature. Each whole pattern fit was beneath a

maximum R-value < 10%. It was assumed that the collected solid contained only crystalline

species when performing the whole pattern fitting. Supplemental table 3 describes peak position

data for observed phases.

9

2.6. Determining wt. % Mn content in collected solids

Wt. % Mnppt was obtained by acidifying approximately 10 μg of powdered solid in 1 mL of

concentrated hydrochloric acid. The solution was capped and left to react overnight in a 50 oC

water bath and sonicated. After acid digestion, the solution was diluted to 10 mL with a 2%

ultrapure nitric acid solution to an estimated Mn concentration of 5 ppm. Mn content in the

collected solid was determined using a Perkin-Elmer Optima 5300DV ICP-AES.

The weight percent mineral content of Mn in the collected solid was calculated by the following

equation 17:

wt. % Mncollected solid = 100*Mndigested solid / Σdigested solid (17)

wt. % Mncollected solid is the weight percent Mn present in the collected solid, Mndigested solid is the

amount Mn present in the digested solid in μg, Σdigested solid is the amount digested solid in μg. Wt.

% Mn for entire precipitate is written as equation 18:

wt. % Mncollected solid = wt. % MnΣppt (18)

wt. % Mncollected solid is the weight percent of Mn present in the collected solid and wt. % MnΣppt is

the weight percent value of Mn present in the total precipitate. Calculated total precipitate was

calculated using equation 19:

Σsolid ppt = Mnppt * wt. % MnΣppt/100 (19)

Σsolid ppt is the total amount of solid precipitated over the course of the experiment in grams,

Mnppt is the amount of Mn precipitated over the course of the reaction in grams calculated from

equation 16, and wt. % MnΣppt is the wt. % Mn content present in the total precipitate calculated

from equation 18. Calculation of the mass of each observed mineral phase present in the

collected solid was performed using equation 20:

10

Min PhaseF, G or H = wt. % Fcollected solid * Σcollected solid (20)

Min PhaseF, G, or H represent the mass of the observed mineral phase in the sum total precipitate in

mg. F, G and H are the mineral phases feitknechtite, groutite and hausmannite, respectively. wt.

% Fcollected solid is the calculated weight percent value of the observed mineral phase from the

whole pattern fitting, and Σcollected solid is the sum total of collected solid in mg, obtained from

vacuum filtration of the experimental solution. Since it is assumed that the collected solid

mineral content is representative of the total precipitate mineral content, the calculated weight

percent value for each observed mineral phase would the same in the total precipitate. Therefore,

the amount of each mineral phase present in the total precipitate would be calculated in the same

fashion as equation 20, but in g as opposed to mg. Mn content present in each observed mineral

phase was calculated using equation 21:

MnF, G or H = Min PhaseF, G or H * WMn/F, G, or H (21)

where MnF, G or H is the Mn content present in either feitknechtite, F; groutite, G; or hausmannite,

H in mg; Min PhaseF, G or H is the amount of either feitknechtite, F; groutite, G; or hausmannite,

H present in the collected solid in mg, and WMn/F, G, or H is the mass fraction of Mn present in each

observed mineral phase, which ideally would be 62 wt. % Mn in Feitknechtite, 62 wt. % Mn in

Groutite and 72 wt. % in Hausmannite.

3. RESULTS

3.0. Visual evidence for Mn oxidation

Oxidation of MnII in solution was confirmed visually. Three important observations were noted

from visual inspection. 1. color change from pink to brown (figure 1). Pink indicates the

presence of MnII(aq)

in solution, while the transition to brown indicates the production of oxidized

Mn precipitate.

11

Figure 1: Time dependent solution color change. Left: Experiment 3, 1 mM Mn t = 1 Hr. Right: Experiment 3, 1 mM Mn, tend = 161 Hr. Over the course of the reaction, the solution color turns from pink to brown (left to right).

2. Increased opacity over time (figure 2). An increase in opacity means the concentration of

suspended brown Mn oxide precipitate increased over time, evidence that oxidation persists.

Figure 2: Visual opacity increase over time. Left: Experiment 9A, 100 mM Mn, pH 9.5, t = 40 Hr. Right: Experiment 9A, 100 mM Mn, pH 9.5, t = 75 Hr. Comparison of images of the same solution taken at 40 and 75 hours reveals that the opacity of the solution increases over time.

3. The rate increase in opacity positively correlated with higher initial concentrations of MnII ,

but more importantly, to higher initial pH values (figure 3). This follows the rate law for Mn

oxidation proposed by Stumm and Morgan (eq. 1), i.e., the rate increases with [Mn][O2][OH-]2.

A higher pHinitial and [MnII(aq)]initial increases the rate of reaction.

12

Experiment 6, 10 mM Mn, pH 9, t = 5 min. Experiment 6, 10 mM Mn, pH 9, t = 1 Hr.

Experiment 10A, 100 mM Mn, pH 9, t = 10 min. Experiment 10A, 100 mM Mn, pH 9, t = 4 Hr.

Experiment 11A, 10 mM Mn, pH 10, t = 10 min. Experiment 11A, 10 mM Mn, pH 10, t = 4 Hr.

Figure 3: Visual extent of reaction. Between experiments 6 at 5 minutes and 1 hour, the opacity did not noticeably changed. Comparing the top two images to experiment 10A at 10 min, the experiment

13

10A solution was already more opaque than the experiment 6 solution at 1 Hr. Comparing the middle two images to the bottom two images, it is noticeable to tell the increased rate in opacity due to the differences in initial pH conditions of the reactions.

3.1. Variation of initial MnII concentration

pH and pE data were collected over the course of every experiment. Figure 4 is a Pourbaix

diagram for [MnII] = 100, 10 and 1 mM. If a solution is at equilibrium, solutions should travel in

pH and pE space along the lines that delineate aqueous and mineral phase boundaries (dashed

lines). If a system is not at equilibrium then the position of the solution in pH-pE space will not

plot along such phase boundaries. LeChatelier’s principle indicates that a system in non-

equilibrium will tend to return to its equilibrium state. If a solution plots rightward of the

aqueous/mineral phase boundary, the solution is oversaturated with respect to MnII in solution

and will precipitate solid if the reaction is fast. If a solution plots leftward of the aqueous/mineral

phase boundary, the solution is undersaturated with respect MnII and will dissolve precipitated

solid.

Experiments 3, 6, and 10 began with an pHinitial =9 and [MnII]initial = 100, 10, 1 mM. The pH-pE

was divided into two stages. Stage 1 having a positive slope and stage 2 having a negative slope.

Stage 1 was further broken down into stages 1A and 1B. 1A occurs within the first minute, and

1B occurs within 0.3 Hr of the start of the experiment. This was observed in all experiments. A

global pE minimum marks the change in the sign of pE/pH slope and the start of stage 2.

Increasing the starting concentration of MnII in solution resulted in a less pronounced drop in pE

and a more pronounced drop in pH during stage 1. An increased [MnII]initial should increase the

rate of reaction (Stumm and Morgan, 1996). The observed pE drops in stage 1 were 2.53, 1 mM;

1.44, 10 mM; and .686, 100 mM. The observed pH drop for stage 1 for experiments 3, 6, and 10

were .369, 1 mM; .961, 10 mM; and 1.711, 100 mM. Slopes for stages 1A and 1B are available

14

in Appendix A, supplemental table 1. During stage 2, an overall increase in pE and decrease in

pH produced a more acidic and more oxidized solution. During stage 2, the rate of reaction

decreased over the course of the experiment as pH decreased due to proton production from Mn

oxidation. A higher initial [MnII] should produce a faster rate of reaction. This would be

evidenced by a lower [MnII]end and lower pH and higher pE values.

Figure 4: pH-pE diagram. [MnII] variation pHi = 9. Experiment 3, 1 mM Mn, red circles; Experiment 6, 10 mM MnII, pink circles, and experiment 10, 100 mM MnII

, orange circles. All pertinent reactions demarcating equilibrium reaction boundaries can be found in table 1. Aqueous phase/mineral phase boundaries are represented by long dashed lines for 100 mM MnII, short dashed lines for 10 mM MnII

, and dotted lines for 1 mM MnII (reactions 3, 5, 7, and 9 in red). Reactions 4, 6 and 8 describe mineral phase/mineral phase boundaries. Solid lines 2 and 10 represent the stability field of water.

Note that the phase diagram presents idealized [MnII(aq)] and neglects activity coefficients for

each pertinent species in solution. Accounting for activity coefficients should shift the phase

boundaries rightward. An estimated activity coefficient of MnII could be .2 (Whitfield, 1975).

Therefore, the experimental solutions during stage 2 may be nearing the aqueous phase/mineral

15

phase boundary. The closer an experiment plots to that line, the closer that experiment is to

reaching chemical equilibrium.

A nearly linear relationship between pE and pH was observed during stage 2 in each

experimental solution. This linear relationship, modeled as pE = m*pH + b, is the Nernst

equation for the forward Mn redox reaction only when equilibrium is established (see table 1). At

thermodynamic equilibrium, for the reaction MnII(aq) oxidizing to a generic MnOx (s) species Keq

equals:

Keq = [H+]m * [e-]/AMn log Keq = m*log [H+] + log [e-] – log AMn (22)

-log[H+] = pH and –log[e-] = pE, solving for pE yields the Nernst equation, equation 23:

pE = -m*pH - log Keq – log AMn (23)

Log Keq is the log equilibrium constant for the overall reaction, m is the pE/pH slope, and log

AMnII is the log activity of MnII in solution. The y-intercept, b, is the sum of log Keq and log AMn.

If equilibrium is not attained, the slope "m" cannot be related to the equilibrium constant".

In supersaturated solutions, Ksp, the solubility product is greater than the equilibrium constant.

The pE/pH slopes of experiments 3, 6, 10 varied from -1.7 to -2.4, compared to the slopes of -3

for reaction 3 and -4 for reaction 7. However, all experiments plot rightward of the

aqueous/mineral phase boundary. Therefore, this experimental system is supersaturated with

respect to Mn-bearing minerals and cannot be treated as if it were in equilibrium.

3.2. Effect of dissolved oxygen equilibrium reaction on slope

Comparison of experiments run at different initial pH values, initial solutions in the absence of

Mn plot on the Pourbaix diagram with a slope of -1, see figures 10 and 11. If the true equilibrium

16

was established, the initial solutions should lie on line 2a. However, all laboratory and natural

solutions line on lines (-1 slope) far below the 2a/b lines at apparent log pO2 values < -20.

Experimental DO values remained essentially constant at 5.5 ppm and varied < 10% throughout

stages 1 and 2. DO is treated as a constant over the course of all experiments.

The pE/pH slope of this reaction is -1. Slope changes sign from positive to negative at the pE

minimum, which denotes the end of stage 1B and the beginning of stage 2. Here, the rate of DO

dissolution is treated as fast compared to rate of Mn oxidation, the DO equilibration contribution

to pE/pH slope is -1.

Since DO = 5.5 ppm = 5/32 mM O2 << MMn, mixing of MnCl2 into O2-bearing solutions will

result in instantaneous consumption of all the dissolved O2 in solution by the reaction MnII(aq) +

.25O2 (aq) +1.5H2O(l) MnOOH(s) + 2H+(aq). This process leads to the consumption of O2 and the

generation of H+. If the supply rate of O2 was much slower than the supply rate of Mn, the

dissolved O2 would have been consumed, and the electrons decreased due to the H2O/O2/e-

reaction. This could have caused the decrease in pE and would explain stage 1. Stage 2 starts

with increasing additions of O2 into MnII-rich solutions. Increasing O2 causes decreasing [H+]*

[e-]: because O2 +2 H+ + 2e- = H2O. If Mn hydration occurs, equation 9, [H+] will increase; then

[e-] must decrease (i.e., pE increase). Thus, the slope pE/pH of the combined reactions in stage 2

will be negative.

3.3. XRD Spectra

Individual XRD spectra of collected solids were compared to one another to determine the

pertinent mineral phases produced over time. From Stumm and Morgan, (1996) rate of reaction

for Mn oxidation is said to increase with higher pHi. Therefore, comparison of experiments

17

holding reaction time constant and varying pH is proportional to comparing experiments holding

pH constant and varying reaction time.

Comparison of recovered solid to calculated solid, the percent recovery for this study was an

average of 1.32% (see supplemental table 2). The remaining 98-99% of the solids probably

adhered on the beaker walls and was not recovered. It is assumed that the recovered precipitate

mineral chemistry is representative of the entirety of the precipitated solid.

3.3.1. Determination of relative precipitate degree of crystallinity

Mineral phase crystallinity is inversely related to the peak full-width half-max value, FWHM, of

XRD spectra. FWHM values for the Feitknechtite 100% peak at d(Å) = 4.62, the most intense

signal, decreased with increasing reaction time as well as with increasing the initial pH of the

solution. Therefore, crystallinity of the collected solid was observed to increase over time. For

specific FWHM values, see Appendix A, supplemental table 2.

Appendix A, supplemental table 2 contains the wt. % calculated amounts of each mineral phase

observed. Appendix B, supplemental figure 3 contains all individual XRD spectra.

3.3.2. Feitknechtite over time

Prominent feitknechtite peaks are labeled in figures 5 – 7 as 2-F (main peak); 9-F; 11-H, F; and

18-H, F. From the pHi = 9, 4 Hr. experiment, 10C, the spectra is highly amorphous with humps

that have apexes at 2-F; 11-H, F; 18-H, F. Because 10C has a high degree of amorphous

character, it is difficult to distinguish the presence of other phases. In Figure 5, hump 2 with

increasing pH shifts left towards 9-F, indicating feitknechtite presence increases dramatically.

Feitknechtite/amorphous phase ratios are presented in supplemental table 2.

18

3.3.3. Groutite over time

Prominent groutite peaks are 3-G (main peak); 4-G; 6-G; and 8-G. Groutite is not observed in

experiments with pHi > 10. Groutite is observed only in the 4 and 40 hour experiments with pHi

< 9.5. Decreased intensity of 3-G in the 40 hour pHi = 9, experiment 10B, and its absence in the

40 hour pHi = 9.5, experiment 9B, suggest groutite production occurs early and is meta-stable

with respect to hausmannite and feitknechtite. Experiments 12 and 13, pH 10 and 10.25, did not

observe groutite. Since rate of reaction is dependent on pH, at these higher pHi it is expected that

groutite formation occurs and is consumed, but is not observed as the rate of reaction is

exponentially faster at these higher pHi experiments.

3.3.4. Hausmannite over time

Prominent hausmannite peaks are 1-H; 7-H; 10-H (main peak), 11-H, F; 15-H; and 17-H.

Hausmannite was not observed at pHi < 9 less than 40 Hr. for pHi = 9.5. In experiments with pHi

>10, hausmannite was not observed. It is believed along with the occurrence and dissapearence

of groutite in pHi < 9.5 experiments; the rate of reaction in pHi > 10 was much too fast to

observe the transition from groutite to hausmannite. Stage 2 pH values experiments pH9.5, 40 Hr =

6.251 and pH10, 4 Hr = 6.855. Overall change in pH from initial pH values, which can be used to

compare extent of reaction since Mn oxidation generates protons, was ∆pH9.5, 40 Hr = 3.249 and

∆pH10, 4 Hr = 3.145. Rate generation of hausmannite may increase with respect to rate of reaction

of groutite with higher pHi. In Figure 6, hausmannite peak intensities clearly increase with

increasing pHi, specifically, the intensity increases of 10-H and 1-H. In figure 6, hump 2 intensity

is further increased and 10-H is even more prominent. The leftward shift on hump 2 in figure 4,

19

then followed by a rightward shift towards 10-H suggests the order of production of hausmannite

follows the initial production of feitknechtite and after the decrease in presence of groutite.

Feng et al., (2004) proposed a mineral transformation pathway that began with the production of

amorphous Mn(OH)2 followed by the production of crystalline phases, by either pathway 1.

hausmannite feitknechtitebuseritebirnessite or pathway 2. Feitknechtite buserite

birnessite. The observed increased crystallinity with increasing extent of reaction in this data set

coincides with the conclusions of Feng et al., (2004).

Hausmannite may have also been formed via the dehydration of feitknechtite when samples were

left to dry in air overnight. However if that was the case, the presence of hausmannite should

have been detected in all samples and hausmannite was not detected in any pHi = 9 experiment

and was only detected for the 75 hour pHi = 9.5 experiment. Therefore, the contribution to the

production of hausmannite by this pathway was negligible.

20

Figure 5: XRD experiments 9 – 11, 4 hour overlay. Initial conditions: 1000 mL solution. [MnII(aq)]i =

100 mM. Red spectra, pHi = 9, experiment 10C; Brown spectra, pHi = 9.5, experiment 9C; Black Spectra, pHi = 10, experiment 11C. F indicates a peak identified as feitknechtite, H indicates a peak identified as a hausmannite phase and G indicates a peak identified as a groutite phase.

21

Figure 6: XRD experiments 9 – 11, 40 hour overlay. Initial conditions: 1000 mL solution. [MnII(aq)]i

= 100 mM. Green spectra, pHi = 9, experiment 10B; Red spectra, pHi = 9.5, experiment 9B; Blue Spectra, pHi = 10, experiment 11B. F indicates a peak identified as feitknechtite, H indicates a peak identified as a hausmannite phase and G indicates a peak identified as a groutite phase.

22

Figure 7: XRD experiments 9 – 11, 75 hour overlay. Initial conditions: 1000 mL solution. [MnII(aq)]i

= 100 mM. Green spectra, pHi = 9, experiment 10A; Red spectra, pHi = 9.5, experiment 9A; Blue Spectra, pHi = 10, experiment 11A. F indicates a peak identified as feitknechtite, H indicates a peak identified as a hausmannite phase and G indicates a peak identified as a groutite phase.

3.4. Mn content analyses

3.4.1. Mn content analysis of precipitated solid

Amount Mn remaining in solution decreased over time. Calculated amount Mn oxidized is the

difference between the experimental MnII remaining in solution and the initial amount after t=0.

When comparing like reaction times, the amount Mn remaining in solution decreased with

increasing pHi. Therefore the extent of reaction increases with higher pHi, which confirms that

23

rate of reaction increases with increasing pH of the solution and the rate of reaction is faster with

higher pHi (figure 8).

Figure 8: Mn in solution and calculated Mn in solid over time. The left-hand vertical axis pertains to experiments 9 – 11. The right-hand vertical axis pertains to experiments 12 and 13. The following line colors describe experiments 9 – 13: orange, experiments 10 (pHi = 9); green, experiments 9 (pHi = 9.5); blue, experiments 11 (pHi = 10); red, experiments 12 (pHi = 10); and black, experiments 13 (pHi = 10.25). Amount MnII is represented by circles. Squares represent Mn in solid ppt.

Mn concentration dropped about 33% over the course of the reaction for pHi = 10, (experiment

11A), over 75 hours. Experiment 12, pHi = 10, runs of 19.75 and 25 hours displayed unusually

low [MnII(aq)]. It is expected for experiment 12 runs to have Mn concentrations higher than

experiment 13, pHi = 10.25, runs when comparing similar unit times. This irregularity may have

been due to a source of contamination, where the MnII concentration had been overly diluted

during the serial dilution of ICP-MS analysis. This would result in a lower than expected final

MnII in solution and a higher than expected calculated amount precipitated Mn.

24

3.4.2. Precipitation rate vs. wt. % Mn in precipitate

Low pHi experiments produced low wt. % Mn content in the solid (e.g., 15-20 wt. %, experiment

10C). Since wt. % Mn content in observed minerals is 62 wt. % Mn for feitknechtite and

groutite and 72 wt. % for hausmannite, amorphous manganese oxy(hydr)oxides,

MnxO(OH)y·nH2O must constitute a significant percentage within the collected solids at the start

of the reaction (figure 9). From figure 8, increasing pHi of the solution resulted in higher degrees

of Mn accumulation per unit time. With further reaction time for a given experiment, rate Mn

accumulation in solid decreased over time, except for experiment 10, pHi 9, where rate generally

remained the same. This is expected considering a lower pH should produce a slower rate of

reaction, see equation 1. Increasing pHi correlated to increasing wt. % Mn content (figure 9).

From this, it can be inferred that rate of reaction may positively correlate with wt. % Mn in solid.

Figure 9: Solid precipitation rate vs. wt. % Mn in solid. Data labels give the reaction time of each data point. Orange squares, pHi = 9, experiment 10. Green squares, pHi = 9.5, experiment 9. Blue

25

squares, pHi = 10, experiment 11. Red squares, pHi = 10, experiment 12. Black squares, pHi = 10.25, experiment 13.

Assuming feitknechtite and/or groutite is the major crystalline phase; the ratio of crystalline to

amorphous phases present in the collected solids decrease with increasing pHi, e.i. pHi 9 and 10

at 4 hours have crystalline/amorphous ratios of .274 and .839, respectively. Additional data is

presented in Appendix A, supplemental table 2.

3.5. Changes in the pH and pE of experimental solutions: experimental series 9 - 13

3.5.1. Experimental series 9 – 11, 75 hour overlay

Figure 10: pH-pE diagram. 75 Hour Overlay. Orange, green and blue circles represent experiments 10 (pHi = 9), 9 (pHi = 9.5), and 11 (pHi = 10), respectively.

The experimental solutions plotted in Figure 10, as well as all other experiments <100 Hr.

plotted within the MnIIIOOH(s) stability field of the Pourbaix diagram. If under strict

26

thermodynamically dominated conditions, only MnOOH(s) would precipitate. If the system were

still thermodynamically dominated, but precipitating two phases simultaneously, it would require

the solution to plot at the thermodynamic triple point for all three species. This scenario is

unlikely because for the given conditions, these solutions plot leftward/undersaturated with

respect to [MnII(aq)] of the triple point. Instead, this experimental system is dominated by kinetics

at non-equilibrium conditions.

In stage 1, a higher pHi led to a more severe pH drop to the global pE minimum. The pE/pH

slope of stage 1 was less positive in higher pHi experiments. From figure 8, increasing pHi

produced a further extent of reaction. Since the Mn oxidation reaction generates e- and H+, higher

pHi experiments will in turn generate more e-. This is evidenced by higher pEend values (figure

10). Theoretically, if left to react longer, solutions would continue to react until they attained

equilibrium. This would be defined as the rate oxidation going to zero, by either the depletion of

[MnII(aq)] or the pH of the solution to be too acidic. During stage 2, slope becomes less negative

at the end of 75 hours in higher pHi experiments. During stage 2, slope indicates the rate of

reaction. Slope over time indicates the extent of the reaction. A faster reaction would produce a

steeper pE/pH slope because less electrons are generated with respect to proton generation per

unit time.

3.5.2. Experimental series 12 and 13: investigation of micro slopes

Experiments 12 and 13 were performed at initial pH values of 10 and 10.25, respectively, to

study the micro scale variations in slope observed in experiments 9 – 11 during stage 2

specifically to understand the importance of observed local minima and local maxima. A More

negative slope suggests that the rate of H+ production was less than the rate of oxidation by O2. It

27

is hypothesized, that differences in slope indicate a change in the overall apparent rate of

reaction. During stage 2, where rate of Mn oxidation is slower than rate of O2 dissolution,

changes in slope could be due to the changes in the emergence of a new dominant reaction.

Overall, for experiments 12 and 13, the slope during phase 2 is about -2, in detail, the slope

changes from -2.3 to -1.5 during phase 2, indicating the rate of Mn oxidation has decreased over

time (figure 11). Prior to reaching a local maxima, slope becomes more negative. After reaching

a global maxima, slope becomes less negative (see local maxima hump for experiment 13 at 7.5

hours and 10.97 hours). The opposite occurs when approaching passing a local minima.

Figure 11: pH-pE diagram. Experiments 12 and 13, pH 10, 10.25. Experiment 12, pHi = 10. Experimetn 13, pHi = 10.25.

pH-pE plots for every experiment can be found in Appendix D, supplemental figure 3.

Additional slopes can be found in Appendix A, supplemental figure 1.

28

4. DISCUSSION

4.1. Solution chemistry reaction stages

Stage 1 is evidence for precipitation. The fastest reaction expected in the system is the hydration

of MnII to form Mn(OH)2, by equation 9: MnII(aq) + 2H2O(l) ⇌ MnII(OH)2 (s) +2H+

(aq). This

reaction does not generate electrons, therefore in the forward direction, reaction slope would

equal = 0. Stage 1A reflects this slope trend with a slightly positive slope (~.5). Stage 1B has a

much more positive slope than stage 1A, which implies oxidation is occurring. Along with the

continued hydration of MnII, the second fastest reaction in the system is the oxidation of

Mn(OH)2 to Mn(OH)3 (equation 24).

MnII(OH)2 (am) + H2O (l) ⇌ MnIII(OH)3 (am) + H+(aq) + e- (24)

At higher initial pH, the slope of stage 1a is less positive and produces a higher pH drop.

Equation 9 is dependent on solution pH and the rate of reaction will increase with increasing pH.

Differentiation of stage 1 from stage 2 occurs where the slope changes from positive to negative

at the global pE minimum. At this point, rate of O2 dissolution is faster than the rate of Mn

oxidation. If the reaction system were closed off from the atmosphere, wherein the system had a

finite amount of O2, then the solution should continue to plot with a positive slope till the rate

Mn oxidation becomes zero, when dissolved O2 is fully consumed. In this experimental system,

DO values vary by < 10% and [O2] in solution is assumed to be constant. This means that the

rate of O2 dissolution was the same as the rate of O2 consumption. If the O2 consumption rate by

Mn oxidation was less than the DO supply, the pE value continues to rise.

29

4.2. Determining reaction rate order

Stumm and Morgan, (1996) state Mn oxidation with respect to [MnII] is first order At const. O2

and pH., see equation 1. However, plotting 1/[MnII(aq)] vs. time for this data set yielded slightly

higher R2 values than ln([MnII]) vs. time. This second order rate equation would then be written

as -d[MnII(aq)]/dt = k1 [MnII

(aq)]2[OH-(aq)]2[O2 (aq)] + k2[MnII

(aq)]2[MnOx (s)][OH-(aq)]2[O2 (aq)].

Therefore, in this experimental system, [MnII] is second order with respect to rate of oxidation. A

similar rate order has been derived by Park and Dempsey, (2005) modeling the rate order of

heterogeneous FeII oxidation at low DO and neutral pH, where they proposed rate of Fe

oxidation is second order with respect to [FeII]. They were able to distinguish dissolved FeII

concentrations from concentration of FeII sorbed onto an iron oxide surface. This may explain

the [MnII] rate order relationship from this study, as the presence of Mn oxides surfaces are

known to autocatalytically enhance the rate of reaction (Matsui 1973; Coughlin and Matsui

1976; Sung and Morgan 1981; Diem and Stumm 1984). In order to determine rate order for Mn

oxidation applying the Stumm and Morgan rate equation, the [MnOx] term should be treated as a

reaction surface. BET studies to determine surface area of the solids were not performed in this

study. Large variances in surface area of the precipitates could have a profound effect on the rate

of reaction. Therefore, only the homogenous term of equation 1 will be used to derive an

approximation for the apparent rate constant. Hem, (1981) deduced that when [MnII] < [MnOx]

the system will behave auto-catalytically and when [MnII] > [MnOx], the reaction will behave

pseudo-first order. Observation of [OH-]-2 vs. time for experiment 13 in figure 13 between 10

and 20 hours noticed an enhanced increase in [OH-]-2. This could be a possible explanation of

autocatalysis. The major contributing species to the rate of reaction of Mn oxidation was

30

determined by plotting 1/([OH-]2[MnII]2[O2]) vs. time and comparing the curves to [Mn]-1 vs.

time and [OH-]-2 vs. time.

31

Figure 12: Inverse concentration over time. Top: 1/([Mn]2[O2][OH-]2) vs. time. Middle: [MnII]-1 vs. time. Bottom: [OH-]-2 vs. time. [OH-]2 is the major contributor to rate of reaction for this proposed system. Orange, pHi = 9, experiment 10. Green, pHi = 9.5, experiment 9. Blue, pHi = 10, experiment 11. Red , pHi = 10, experiment 12. Black, pHi = 10.25, experiment 13.

[OH-] is the major contributor to the experimental rate equation because the shape of the [OH-]-2

vs. time plot mirrors the 1/([OH-]2[MnII]2[O2]) vs. time plot more than the [Mn-]-1 vs time plot.

32

Figure 13: Stage 1, 2 comparison of [OH-]-1 vs. time. Orange, pHi = 9, experiment 10. Green, pHi = 9.5, experiment 9. Blue, pHi = 10, experiment 11. Red , pHi = 10, experiment 12. Black, pHi = 10.25, experiment 13. Triangles in the middle figure represent stage 1A of the reaction while squares represent stage 1B of the reaction.

33

At stage 1, two slopes indicating different k values are observed, one for stage 1A and the other

for stage 1B. Plotting 1/[OH-] over time for stage 2 results in abruptly different slopes over the

course of each experiment, leading to the conclusion that different reactions are occurring over

time. Since alteration of [OH-] in solution is due to the oxidation and/or hydration of Mn, this is

therefore evidence that changes in the slope of pE/pH indicate the alteration of major reactions

occurring in solution.

To estimate the different rate constants observed in each experiment from plotting [OH-]-1 vs.

time, an initial approximation of the rate constant can be taken by assuming a pseudo second

order reaction with respect to [OH-] since [MnII] changes <30% over the course of any reaction.

-d[MnII]/dt = Rate = kapp [MnII]2[O2][OH-]2 (25)

the experimentally derived rate law above can be simplified to a first approximation of k as:

-d[MnII]/dt = Rate = kOH [O2][OH-]2 (26)

where kapp is the apparent rate constant for the reaction and kOH is the first approximation derived

rate constant. From an initial guess calculated from (equation 26) and an experimentally derived

[MnII] at time t, the rate constant for equation 25. Table 2 contains data of the derived reaction

rates and rate constants for every experiment.

34

Table 2: Rate constant approximation. First approximated kOH- and the iterated rate constant, kapp for every experiment.

First approximation of the rate constant for stage 1 produced an average kOH- for stage 1A was

4.67*107 M-1 Hr-1, and for stage 1B kOH = 1.08*106 M-1 Hr-1. kapp could not be solved as stage 1

[Mn] data does not exist. It is expected this stage should observed the fastest rates of reaction.

A larger rate constant implies a faster reaction. In addition, a more negative pE/pH slope implies

a faster rate of reaction. Therefore, calculated rate constant and pE/pH slope are inversely related

to each other. At points of local minima and maxima, slope is shallow, therefore rate of reaction

Exp Run time (Hr.) pHi pHf kOH- - d[Mn]dt exp kapp slope10C 4 9.0 7.222 2.40E+05 1.02E-04 2.89E+13 -2.99410B 40 9.0 6.818 3.00E+05 7.41E-05 1.65E+13 -1.29310A 75 9.0 6.532 4.00E+05 8.32E-05 4.18E+13 -2.2759C 4 9.5 7.126 1.00E+06 3.82E-04 1.03E+14 -2.4539B 40 9.5 6.187 2.00E+06 3.15E-04 7.75E+15 -1.9129A 75 9.5 6.019 8.50E+05 1.37E-04 2.22E+13 -2.265

11C 4 10.00 6.865 5.00E+06 1.29E-03 4.17E+14 -2.34811B 40 10.00 6.086 1.00E+06 3.80E-04 1.49E+16 -1.33911A 75 10.00 6.045 4.00E+05 1.63E-04 7.69E+15 -1.685

12_4.4 4.4 10.00 6.843 4.00E+06 1.18E-03 1.41E+15 -2.19612_5.6 5.6 10.00 6.763 2.00E+06 2.85E-03 4.94E+15 -1.97512_7.8 7.8 10.00 6.632 3.00E+06 6.67E-04 2.11E+15 -1.75712_13 13 10.00 6.476 2.00E+06 8.93E-05 5.81E+14 -1.350

12_16.5 16.5 10.00 6.415 1.00E+06 -1.32E-03 -1.43812_19.75 19.75 10.00 6.372 1.00E+06 1.01E-03 1.06E+16 -1.438

12_21 21 10.00 6.362 7.40E+05 -7.66E-03 -1.53612_25 25 10.00 6.340 4.40E+05 1.67E-03 2.02E+16 -2.204

13_2.75 2.75 10.25 6.734 5.00E+06 2.17E-03 4.30E+15 -2.25213_3.83 3.83 10.25 6.676 3.00E+06 4.30E-03 8.52E+15 -0.52313_4.5 4.71 10.25 6.632 3.00E+06 -2.51E-03 -0.01713_7.5 7.5 10.25 6.528 2.00E+06 1.67E-04 8.56E+14 -1.82413_9.3 9.3 10.25 6.464 3.00E+06 -1.29E-04 -1.646

13_10.97 10.97 10.25 6.412 3.00E+06 -1.49E-03 -1.48113_13 13 10.25 6.360 1.00E+06 3.52E-04 3.90E+15 -1.58413_42 42 10.25 6.176 4.00E+05 1.59E-04 4.11E+15 -1.443

35

is slow and the rate constant is so accordingly. At midpoints, the rate of reaction is fast ( see

between 3.5 and 10.97 Hr. for experiment 13, figure16).

This experimental system varies both in [MnII] and [OH-] over time. Calculation of the apparent

rate and normalizing for [MnII] for each experiment, can be used to compare the extent of

reaction over all experiments. Here, rates over a 1.25 pH log units is proportional (figure 14).

Figure 14: Apparent rate/[MnII] vs. pH. Experiment 10, pHi = 9 orange dot; experiment 9, pHi = 9.5, green dot; experiment 11, pHi = 10, blue dot; experiment 12, pHi = 10, red dot; experiment 13, pHi = 10.25, black dot.

0.00E+00

1.00E-18

2.00E-18

3.00E-18

4.00E-18

5.00E-18

66.26.46.66.877.2

rate

app/

[MnII ]

pH

rateapp/[MnII] vs. pH

36

4.3. Mineral phase production over time

Figure 15: Mineral phase transformation scheme. Arrows in black refer to reactions occurring during stage 1. Reactions in red occur during stage 2.

H2O(l) ⇌ O2 (aq) + 2H+(aq) + e- (1)

MnII(aq) + 2 H2O(l) ⇌ MnII(OH)2 (am) + 2H+

(aq) (9)

MnII(OH)2 (am) + H2O (l) ⇌ MnIII(OH)3 (am) + H+(aq) + e- (24)

MnII(OH)2 (am) + MnIII(OH)3 (am) ⇌ 2MnOOH(s) + H2O(l) + H+(aq) + e- (27)

MnII(OH)2 (am) + 2 MnIII(OH)3 (am) ⇌ MnIIOMnIII2O3 (s) + 4 H2O(l) (28)

2 MnOOH(s) + MnII(aq) ⇌ MnIIOMnIII

2O3 (s) + 2H+(aq) (29)

The production of amorphous solids must occur first, which then transform to crystalline phases

that increase the wt. % Mn content observed in collected solids with furthering extent of reaction

as evidence by XRD and increasing wt. % Mn content in solid over time. The production of

37

amorphous solids can be summarized in figure 15, and are produced by reactions 9 and 24 during

stage 1. Specifically Mn(OH)2 is the fastest reaction and happens during stage 1A and then

transforms at a slower rate to amorphous Mn(OH)3 during stage 1B. Soon after, these

amorphous species transform to either groutite or feitknechtite via equation 27, MnII(OH)2 (am) +

MnIII(OH)3 (am) ⇌ 2MnOOH(s) + H2O(l) + H+(aq) + e-. The rate of groutite formation is slightly

slower than the rate of feitknechtite formation, as can be seen in figure 16 where the rate of

feitknechtite accumulation is > those of any other observed phase. Groutite is also transformed in

pH 9.5 experiments similarly in time with the rise of hausmannite. Therefore, a possible outcome

is groutite being metastable and subject to transformation to hausmannite by equation 29.

Feitknechtite transformation to hausmannite must be slow with respect to the generation of

feitknechtite by reaction 27, while this transformation is fast in the groutite pathway. If

transformation from feitknechtite to hausmannite was fast, then hausmannite would be the

expected main phase.

It could be possible that hausmannite would form from the oxidation and dehydration of

amorphous phases (equation 28) when solids were dried in open air. However if this reaction

were to have occurred, then hausmannite should have been observed in all experiments.

Comparison of experiment 13 from figure 16 to the pE-pH diagram noticed that at midpoints,

slope was at its most negative, meaning reaction rate was at its highest. Comparison of slope to

rate mineral phase accumulation confirmed that the rate of reaction in phase two was dominated

by the formation of feitknechtite. An estimated rate constant for this reaction (27) would be

6.2*1015 M-5 Hr-1.

38

`

Figure 16: Mineral precipitation rate and slope vs. pH. Slope is the black line. Rate is in g/Hr. Rate precipitation of feitknechtite is the red line, precipitation rate of groutite is the orange line, precipitation of hausmannite is the blue line. Blue numbers denote in hours the reaction time. Left Axes define rate precipitate of a mineral, (g/hr). The right Axes define pE/pH slope, and the horizontal axis is pH

39

4.4. Future studies

Future research would involve developing an experimental RIR value for feitknechtite. This will

increase the accuracy of the whole pattern fit since no RIR value currently exists. Further,

application of a different XRD technique besides capillary would help reduce peak broadening,

which would again produce more accurate whole pattern fittings.

Thirdly, it would be necessary to look at the mineral structure of solids produced prior to 4 hours

for all starting pH values, however this is limited by low crystallinity of the produced solid

because MnII oxidation to an amorphous solid is first and fastest reaction.

5. CONCLUSION

Under unbuffered conditions, which have not been previously studied, the rate of Mn oxidation

can be treated as pseudo second order with respect to [OH-] to later constrain the rate constants

of pertinent reactions. Main conclusions drawn from this thesis: Precipitation of Mn ocurrs at pH

values as low as 6.5, the inverse correlation between wt. % content of Mn in solid and

accumulation rate of Mn in solid (figure 9), the second order relationship between [MnII] and

oxidation rate, and the ability to perform in situ measurements of pH-pE-DO, combined with

solution and mineral analyses to produce an effective method to study the details of redox

reactions among aqueous and solid phases. pH/pE slope does not apply to the equilibrium pH/pE

ratio of each pertinent reaction under these experimental conditions because the system is not at

equilibrium. Formation of Mn(OH)2 from MnII is the fastest reaction, therefore the major

pathway for production of crystalline phases is solid phase transformation of amorphous

Mn(OH)2 to MnII-III phases, and not direct precipitation in this kinetically dominated system.

40

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Bricker, O. P. (1965) Some stability relationships in the system Mn–O2–H2O at 25 oC and 1 atmosphere total pressure. American Mineralogist 50, 1296-1354.

Coughlin R. W., and Matsui I. (1976) Catalytic oxidation of aqueous Mn(II). Journal of Catalysis 41, 108-123.

Davies S. H. R., Morgan J. J. (1988) Manganese(II) oxidation kinetics on metal oxide surfaces. Journal of Colloid and Interface Science 129, 66-77.

Diem D., and Stumm W. (1984) Is dissolved Mn2+ being oxidized by O2 in absence of Mn- bacteria or surface catalysts? Geochimica et Cosmochimica Acta 48, 1571-1573.

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Hem J. D. (1981) Rates of manganese oxidation in aqueous systems. Geochimica et Cosmochimica Acta 45, 1369-1374.

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Lefkowitz J. P., Rouff A. A., and Elzinga E. J. (2013) Influence of pH on the reductive transformation of birnessite by aqueous Mn(II) Environmental Science & Technology 47, 10364-10371.

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Morgan J. J. (2004) Kinetics of reaction between O2 and Mn(II) species in aqueous solutions. Geochimica et Cosmochimica Acta 69, 35-48.

Park B. and Dempsey B. A. (2005) Hetergenous Oxidation of FeII on Ferric Oxide at Neutral pH and a Low Partial Pressure of O2. Environmental Science and Technology 39, 6494–6500.

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Sung W., and Morgan J. J. (1981) Oxidative removal of Mn(II) from solution catalyzed by the γ- FeOOH (lepidocrocite) surface. Geochimica et Cosmochimica Acta 45, 2377-2383.

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42

APPENDIX A: Supplemental Tables

Supplemental table 1: Experimental solution data.

Exp Run time (Hr.)

pHi pEi DOi ppm

pH(pE

min) pE(pE

min)

DO(pE

min)

ppm pHf pEf DOf

ppm Max/min

Stage 1A

Slope

Stage 1B

Slope Stage 2 Slope

3 161 9.006 7.109 5.31 8.643 4.751 5.32 5.372 12.206 5.34 3.182 6.971 -2.299 6 206.7 8.949 6.938 5.39 7.988 0.21 5.40 5.298 11.534 5.34 0.870 -7.049 -2.324

10C 4 8.991 7.000 5.85 7.308 6.514 7.224 6.790 5.85 0.626 -9.432 -1.845 10B 40 9.010 7.132 5.89 7.290 6.542 6.730 7.595 5.85 0.603 -15.31 -1.371 10A 75 8.995 7.315 5.94 7.291 6.443 6.532 8.077 5.79 0.401 15.19 -1.692 9C 4 9.504 6.446 5.96 7.566 5.748 7.028 6.928 5.85 0.047 -44.66 -2.341 9B 40 9.500 6.546 5.89 7.563 5.749 6.251 8.449 5.80 0.544 160.5 -1.840 9A 75 9.497 6.579 5.87 6.579 7.566 6.019 9.057 5.80 0.211 67.35 -1.960

11C 4 10.010 6.165 5.58 7.930 4.944 6.865 7.479 5.32 0.575 1.910 -2.464 11B 40 10.009 6.203 5.89 7.693 5.531 6.086 8.688 5.76 0.121 3.260 -1.743 11A 75 10.007 6.254 5.85 7.713 5.579 6.045 8.736 5.65 0.300 17.494 -1.453

12_4.4 4.4 9.998 6.962 5.61 7.792 5.376 6.669 7.852 5.35 min 0.435 -22.20 -2.078 12_5.6 5.6 9.988 6.262 5.45 7.830 5.123 6.600 7.834 5.22 mid 0.528 -2.025 -2.135 12_7.8 7.8 10.012 6.131 5.48 7.958 5.223 5.40 6.632 8.154 5.20 max 0.338 0.206 -2.299 12_13 13 10.002 6.179 5.50 7.969 5.243 6.539 8.167 5.22 mid 0.151 2.563 -0.567

12_16.5 16.5 9.999 5.762 5.41 7.826 5.317 5.39 6.450 7.887 5.21 min 0.169 -9.518 -1.197 12_19.75 19.75 10.003 6.541 5.47 7.808 5.220 6.366 8.450 5.22 max 1.080 0.359 -1.584

12_21 21 10.015 6.569 5.56 7.829 5.366 6.374 8.350 5.25 min 0.514 1.190 -2.328 12_25 25 10.065 6.770 5.45 7.851 5.056 6.340 7.934 5.23 mid 0.644 7.828 -2.204

13_2.75 2.75 10.236 6.020 5.71 7.828 5.332 6.887 7.444 5.65 mid 0.293 0.382 -2.373 13_3.83 3.83 10.260 6.510 5.61 7.926 4.761 6.829 6.907 5.40 min 0.423 -20.14 -1.359 13_4.5 4.71 10.285 6.775 5.68 7.839 5.452 6.807 7.578 5.42 mid 0.276 3.551 -2.017 13_7.5 7.5 10.231 7.317 5.67 7.835 5.053 6.670 7.730 5.38 max 0.305 3.502 -1.463 13_9.3 9.3 10.263 6.944 5.61 7.732 5.337 6.533 7.341 5.38 mid 0.253 7.681 -0.235

13_10.97 10.97 10.251 6.789 5.67 7.774 5.559 6.516 8.112 5.36 min 0.325 1.510 -0.442 13_13 13 10.241 6.531 5.71 7.826 5.149 6.474 8.044 5.41 mid 0.200 12.26 -1.261 13_42 42 10.281 6.419 5.56 7.742 5.357 6.176 8.357 5.19 mid 0.260 6.538 -1.443

43

Supplemental table 2: Experimental solid phase data.

Exp Run time (Hr.)

∑MnII(aq)

(mM) % Solid

Recovery ∑solid

(g) Wt. % Mnsolid

∑Mnsolid (g)

Feitknechtite (g)

Hausmannite (g)

Groutite (g)

Crystalline/Amorphous Phase Ratio

Feitknechtite FWHMo at d(Å) = 4.62

10C 4 5471 1.92 0.044 17 0.020 0.044 0 0.274 3.370

10B 40 5325 0.25 0.631 20 0.171 0.458 0 0.173 0.323 2.074

10A 75 5165 0.28 1.116 15 0.324 0.910 0 0.205 0.242 1.064

9C 4 5410 0.21 0.508 46 0.087 0.329 0 0.179 0.742 1.168

9B 40 4787 0.10 3.359 27 0.656 2.301 0 1.058 0.435 0.849

9A 75 4523 0.03 4.500 29 0.851 4.127 0.153 0.221 0.468 0.814

11C 4 5209 2.85 0.548 52 0.284 0.383 0.165 0.839 0.926

11B 40 4184 0.49 2.563 51 1.310 2.432 0.131 0.823 0.787

11A 75 3692 0.48 3.447 52 1.799 3.241 0.207 0.839 0.927

12_4.4 4.4 5210 1.77 0.342 42 0.144 0.310 0.031 0.677 0.978

12_5.6 5.6 5022 0.78 0.851 28 0.235 0.745 0.106 0.452 0.722

12_7.8 7.8 4941 1.35 0.516 54 0.278 0.428 0.089 0.871 0.510

12_13 13 4915 1.04 0.520 56 0.290 0.461 0.059 0.903 1.125

12_16.5 16.5 5169 1.25 0.388 41 0.159 0.355 0.033 0.661 0.797

12_19.75 19.75 4542 0.43 1.452 33 0.475 1.220 0.232 0.532 0.891

12_21 21 5068 1.41 0.439 49 0.214 0.409 0.030 0.790 1.046

12_25 25 4062 0.27 2.350 31 0.720 2.035 0.315 0.500 0.826

13_2.75 2.75 5166 3.46 0.323 51 0.165 0.301 0.022 0.823 1.053

13_3.83 3.83 4910 2.02 0.563 52 0.292 0.515 0.048 0.839 0.933

13_4.5 4.71 5032 2.21 0.506 46 0.232 0.457 0.049 0.742 0.711

13_7.5 7.5 4877 2.66 0.500 62 0.309 0.444 0.056 1 0.836

13_9.3 9.3 4889 1.58 0.706 43 0.303 0.617 0.088 0.694 0.924

13_10.97 10.97 5026 3.23 0.405 58 0.234 0.369 0.036 0.935 0.939

13_13 13 4770 1.68 0.743 49 0.362 0.646 0.097 0.790 0.818

13_42 42 4517 0.85 1.502 32 0.487 1.342 0.159 0.516 0.959

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Supplemental table 3: Mineral phase XRD peak list.

Feitknechtite XRD Peak ListPeak Angle d(Å) I%(f) ( h k l)

1 8.805 4.62 100 ( 0 0 2)2 15.47 2.635 50 ( 3 1 1)3 17.29 2.36 20 ( 3 1 2)4 20.85 1.96 10 ( 4 0 2)5 26.45 1.55 1 ( 0 0 6)6 27.35 1.5 1 ( 4 4 1)Hausmannite XRD Peak List

Peak Angle d(Å) I%(f) ( h k l) Peak Angle d(Å) I%(f)( h k l)1 8.265 4.9211 25.1 ( 1 0 1) 21 29.724 1.3827 2.6 ( 4 1 1)2 13.2 3.0868 38 ( 1 1 2) 22 29.824 1.3782 1.4 ( 4 0 2)3 14.14 2.8815 15.7 ( 2 0 0) 23 30.515 1.3477 4 ( 3 0 5)4 14.74 2.7654 75.7 ( 1 0 3) 24 31.287 1.3152 0.1 ( 1 0 7)5 16.4 2.4866 100 ( 2 1 1) 25 31.525 1.3055 1.2 ( 3 3 2)6 16.57 2.4605 14.1 ( 2 0 2) 26 31.949 1.2887 1.1 ( 4 2 0)7 17.26 2.364 18.1 ( 0 0 4) 27 32.229 1.2777 6.9 ( 4 1 3)8 20.05 2.0375 19.8 ( 2 2 0) 28 33.148 1.2433 2.9 ( 4 2 2)9 20.48 1.9952 0.3 ( 2 1 3) 29 33.509 1.2303 3.4 ( 4 0 4)

10 21.72 1.8826 0.1 ( 3 0 1) 30 33.777 1.2208 0.7 ( 3 2 5)11 22.38 1.8276 5.6 ( 2 0 4) 31 34.485 1.1965 3.9 ( 2 1 7)12 22.77 1.7969 20.3 ( 1 0 5) 32 34.616 1.1921 2.1 ( 3 1 6)13 24.08 1.7005 8.5 ( 3 1 2) 33 34.92 1.182 3.1 ( 0 0 8)14 24.97 1.6404 7.5 ( 3 0 3) 34 36.115 1.1441 0.3 ( 5 0 1)15 26.01 1.576 24 ( 3 2 1) 35 36.534 1.1315 1.5 ( 4 2 4)16 26.57 1.5434 45.2 ( 2 2 4) 36 36.783 1.1241 4.5 ( 4 1 5)17 26.9 1.5247 0.8 ( 2 1 5) 37 37.441 1.105 0.1 ( 3 0 7)18 27.92 1.4699 2.1 ( 1 1 6) 38 37.644 1.0993 1.2 ( 5 1 2)19 28.5 1.4408 16 ( 4 0 0) 39 37.847 1.0936 0.6 ( 2 0 8)20 28.81 1.4256 0.8 ( 3 2 3)

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APPENDIX B: Supplemental Methods

Supplemental figure 1: Experimental method schematic.

46

APPENDIX C: Supplemental Photos

Exp 3. 1 mM Mn. No Mn present

Exp 3. 1 mM Mn. t=0. Mn added.

Exp 3. 1 mM Mn. t=5 min.

Exp 3. 1 mM Mn. t=10 min.

Exp 3. 1 mM Mn. t=15 min.

Exp 3. 1 mM Mn. t=20 min.

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Exp 6. 10 mM Mn. No Mn present

Exp 6. 10 mM Mn. t=0. Mn added.

Exp 6. 10 mM Mn. t=15 min.

Exp 6. 10 mM Mn. t=10 min.

Exp 6. 10 mM Mn. t=206 Hr.

Exp 10A. 100 mM Mn. pH 9. No Mn

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Exp 10A. 100 mM Mn. pH 9. t = 0 min

Exp 10A. 100 mM Mn. pH 9. t = 20 min

Exp 10A. 100 mM Mn. pH 9. t = 15 min

Exp 10A. 100 mM Mn. pH 9. t = 75 Hr.

Exp 10A. 100 mM Mn. pH 9. t = 40 Hr.

Exp 9A. 100 mM Mn. pH 9.5. No Mn

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Exp 9A. 100 mM Mn. pH 9.5. t = 0 min

Exp 9A. 100 mM Mn. pH 9.5. t = 10 min

Exp 9A. 100 mM Mn. pH 9.5. t = 15 min

Exp 9A. 100 mM Mn. pH 9.5. t = 20 min

Exp 9A. 100 mM Mn. pH 9.5. t = 4 Hr.

Exp 10A. 100 mM Mn. pH 10. No Mn added.

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Exp 11A. 100 mM Mn. pH 10. T = 0 Hr.

Exp 11A. 100 mM Mn. pH 10. T = 5 min

Exp 11A. 100 mM Mn. pH 10. T = 15 min

Exp 11A. 100 mM Mn. pH 10. T = 40 Hr.

Exp 11A. 100 mM Mn. pH 10. T = 75 Hr.

Supplemental figure 2: Additional experiment photos.

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APPENDIX D: Supplemental Figures

Mn2+ (aq)

52

53

54

55

56

57

58

59

Supplemental figure 3: Individual experiment pH-pE data.

60

61

Supplemental figure 4: Individual experiment slope over time.

62

63

64

65

66

67

68

69

70

71

72

73

74

75

76

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Supplemental figure 5: Individual experiment XRD spectra.

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