Geochemical Stability of Dissolved Mn(III) in the Presence of

Subscriber access provided by UNIV OF LOUISIANA

Environmental Processes

Geochemical Stability of Dissolved Mn(III) in the Presence of Pyrophosphate as a Model Ligand: Complexation and Disproportionation Ao Qian, Wen Zhang, Cheng Shi, Chao Pan, Daniel E. Giammar, Songhu Yuan, Hongliang Zhang, and Zimeng Wang Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.9b00498 • Publication Date (Web): 11 Apr 2019 Downloaded from http://pubs.acs.org on April 13, 2019

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 31

Environmental Science & Technology

Geochemical Stability of Dissolved Mn(III) in the Presence of Pyrophosphate as a Model Ligand: Complexation and Disproportionation

Ao Qian1, Wen Zhang2, Cheng Shi3, Chao Pan4, Daniel E. Giammar4, Songhu Yuan1, Hongliang Zhang3, Zimeng Wang2, 5*

1 State

Key Laboratory of Biogeology and Environmental Geology, China University of Geosciences, Wuhan, Hubei, China

2 Department

of Environmental Science and Engineering, Fudan University, Shanghai, China

3 Department

of Civil and Environmental Engineering, Louisiana State University, Baton Rouge, LA, United State

4 Department

of Energy, Environmental and Chemical Engineering, Washington University in St. Louis, St. Louis, MO, United

State 5 Shanghai

Institute of Pollution Control and Ecological Security, Shanghai, China.

* corresponding

author: [email protected]

Phone: +86-21-31248978 Revised Manuscript submitted to Environmental Science & Technology

0

ACS Paragon Plus Environment


Page 2 of 31

Abstract 1

Dissolved Mn(III) species have recently been recognized as a significant form of Mn in redox transition zones,

2

but their speciation, stability and reactivity are poorly understood. Besides acting as the intermediate for Mn redox

3

chemistry, Mn(III) can undergo disproportionation producing insoluble Mn oxides and aqueous Mn(II). Using

4

pyrophosphate(PP) as a model ligand, we evaluated the thermodynamic and kinetic stability of Mn(III) complexes.

5

They were stable at circumneutral pH and were prone to (partial) disproportionation at acidic or basic pH. With an

6

initial lag phase, the kinetics of Mn(III)-PP disproportionation was auto-catalytic with the produced Mn oxides

7

promoting the disproportionation. X-ray diffraction and the average Mn oxidation state indicated that the solid

8

products were not pure Mn(IV) oxides but a mixture of triclinic birnessite and -MnO2. Addition of synthetic

9

analogs of the precipitates eliminated the lag phase, confirming their catalytic roles. Thermodynamic calculations

10

adequately predicted the stability regime of Mn(III)-PP. The present results refined the constant for Mn(PP)25-

11

formation, which allows a consistent and quantitative prediction of equilibrium speciation of Mn(III)-Mn(II)-

12

birnessite with PP. A simple and robust model, which incorporated the thermodynamic constraints, autocatalytic

13

rate law, and verified reaction stoichiometry, successfully simulated all kinetic data.

1


Page 3 of 31

14


INTRODUCTION

15

Manganese (Mn) is the third most abundant transition metal in the Earth’s crust and is important to many

16

biogeochemical redox processes.1, 2 In natural aquatic systems, Mn exists in three oxidation states, Mn(II), Mn(III),

17

and Mn(IV), which are present over a wide range of redox conditions in the environment.1 Mn(IV) is insoluble and

18

exists as solid phases that can be potent oxidants.3, 4 Mn(II) is thought to be the primary dissolved form of Mn and

19

can be abiotically and biotically oxidized to Mn(III)5, 6 and Mn(IV).7, 8 The intermediate oxidation state, Mn(III), is

20

unstable in water because it can immediately disproportionate to Mn(II) and Mn(IV).9 Mn(III) forms from both one-

21

electron oxidation of Mn(II)5, 6 and from reduction of Mn(IV)10-12, and the reaction between Mn(II) and Mn(IV)

22

could also produce solid phases of Mn(III) through comproportionation.13-15 Dissolved Mn(III) species have been

23

underappreciated for a long time in the conventional paradigm of Mn environmental geochemistry.16

24

Recent reports have highlighted the ubiquitous presence of dissolved Mn(III) in a variety of aquatic systems

25

including anoxic/suboxic seawaters17, sediment porewaters18-21, oxygenated surface waters17, 22, 23 and engineered

26

treatment systems24. The stabilization of dissolved Mn(III) is enabled by complexation with naturally occurring

27

high-affinity ligands that include natural organic matter (and its degradation products)22, 25, biogenic siderophores26,

28

and cell-lysis products.27 The composition of these ligands in natural waters is not well known except for some

29

indirect information about the range of their overall binding strengths falling between those of pyrophosphate and

30

desferrioxamine-B.28 Although complexation will lower the redox potential of Mn(III) below its standard value

31

(EH°(Mn3+/Mn2+) = 1.56V), Mn(III) is presumably still an important geochemical oxidant from the thermodynamic

32

point of view29-33. For example, we previously demonstrated that UO2(s) was rapidly oxidized to U(VI) and

33

subsequently dissolved.34 It was also reported that dissolved Mn(III) did not oxidize Cr(III) unless or until

34

disproportionation happened, even though the energetics of the reaction was favored.35, 2


36

Therefore, the


35

Page 4 of 31

environmental geochemistry of dissolved Mn(III) remains a significant knowledge gap.

36

Even in simplified model ligand systems (e.g., pyrophosphate, citrate and desferrioxamine B.37-40), the speciation,

37

stability and reactivity of dissolved Mn(III) species are far from quantitatively understood.27, 34 The stability of

38

soluble Mn(III) is affected by the strength of complexation with the ligands, the solution pH and the total ligand:Mn

39

ratio (which determines the structural configuration of the complexes).41-43 Depending on the binding affinity of the

40

ligands, an excess of ligand is generally needed to maintain the stability of Mn(III) complexes.41, 42, 44 With natural

41

and biogenic occurrences, pyrophosphate (PP) is an inorganic, redox-inert ligand that can form stable complexes

42

with Mn(III) under neutral pH with representative binding affinity in reference to Mn(III) ligands in natural waters.28

43

Mn(III)-pyrophosphate complexes disproportionate at alkaline and acidic pH, and at high Mn(III):PP ratios.34, 43

44

Citrate forms complexes with Mn(III) that are metastable with respect to possible disproportionation and to

45

reduction of the Mn(III) center by the organic carbon in the citrate.42, 43 Different from citrate, desferrioxamine B

46

stabilizes Mn(III) at alkaline pH (7.0 -11.3), but at lower pH Mn(III)-desferrioxamine B complexes decompose by

47

intermolecular electron transfer to yield Mn(II) and oxidized siderophores, and at pH > 11 the complex degrades

48

by disproportionation to yield Mn(II) and Mn oxide.41 Nonetheless, those processes involving destabilization of

49

Mn(III) complexes lack quantitative kinetic data and models based on well controlled lab experiments.

50

With a prolonged lifetime when complexed by ligands, dissolved Mn(III) may be a critical means of transporting

51

oxidized Mn across redox interfaces21 and ultimately producing Mn(IV) in redox transition environments45

52

Although recent research has advanced our understanding of the formation and stability of dissolved Mn(III)

53

complexes,24, 47 less consideration has been specifically given to the process of disproportionation. Traditionally

54

Mn(II) and Mn(IV) are considered as the products from the disproportionation of Mn(III).9 Mn(II) exists mainly as

55

soluble species and insoluble Mn(IV) quickly hydrolyzes and polymerizes to precipitate as Mn(IV) oxides.46 Both 3


46.

Page 5 of 31


56

Mn(III) complexes and dissolved Mn(II) produced by disproportionation have the ability to adsorb to Mn oxides.42,

57

48

58

complexes are not well understood.

In addition, the mineral phases and properties of Mn oxides produced from disproportionation of Mn(III)

59

The objectives of this study were to (i) assess the stabilization and disproportionation of dissolved Mn(III)-PP at

60

environmentally relevant pH; (ii) characterize the Mn solids produced from Mn(III)-PP disproportionation; (iii)

61

quantify the kinetics of Mn(III)-PP disproportionation and (iv) develop and refine self-consistent thermodynamic

62

and kinetic models for the Mn(III)-PP system. In this study, we used pyrophosphate, a dimer of orthophosphate as

63

a model Mn(III) ligand because of its representative binding affinity28, redox-inertness and occurrence in natural

64

environments from biogenic and synthetic sources.27

65 66

MATERIALS AND METHODS

67

Materials. Manganese(III) acetate dihydrate (> 97%), manganese(II) chloride tetrahydrate (> 99%), potassium

68

permanganate (> 99%) and sodium pyrophosphate decahydrate (> 99%) were purchased from Sigma-Aldrich.

69

Mn(III)-PP stock solution was prepared following the procedures of Kostka.21 In this procedure sodium

70

pyrophosphate decahydrate was first dissolved in ultrapure water (resistivity > 18.2 1OP- ? with a concentration

71

of 50 mM and the pH was adjusted to 8.0. Mn(III) acetate dehydrate solid was then added into the sodium

72

pyrophosphate solution under vigorous stirring to reach a final concentration of 10 mM Mn(III). Note that previous

73

experience showed that at least 4 times excess of PP relative to Mn(III) was required to maintain the complexes

74

against disproportionation.29 The final pH after dissolving the aliquot of Mn(III) acetate was around 7.0. This

75

solution was filtered (0.22 Q polyethersulfone, Tisch Environmental, OH) to remove any possible Mn precipitates.

76

It served as the stock solution of Mn(III) for preparation of the experimental solutions summarized in Table 1. The 4



Page 6 of 31

77

average oxidation state of the Mn(III)-PP stock solution determined by comparing iodometric titration results and

78

total Mn concentrations confirmed the purity of Mn(III) above 98%.49 H2-sparged anoxic ultrapure water was used

79

to prepare all solutions. Diluted HNO3 and NaOH were used for pH adjustment.

80

The synthesis of R;1 22 and triclinic birnessite followed the procedures described in Villalobos et al.50. Briefly,

81

a volume of 320 mL of MnCl2 (0.095 mol) was added to 680 mL of solution containing 0.175 mol NaOH and 0.063

82

mol KMnO4. The resulting suspension was stirred for 2 h and then left to settle for 6 h. Following the precipitation

83

of the solid, R;1 22 was washed with 1 M NaCl, shaken for 1 h and centrifuged. The supernatant solution was then

84

discarded. The NaCl washing procedure was repeated five times. The centrifuged paste was washed with similar

85

procedures with DI water for another five times. The suspension was then dialyzed (3500 Da MWCO) in DI water

86

until the conductivity of water outside the dialysis bag remained below the detection limit for 12 hours. Triclinic

87

birnessite was synthesized as follows: 160 mL of MnCl2 (0.08 mol) was first added to 180 mL of NaOH (1.375

88

mol) under stirring to form a pink gel precipitate of Mn(OH)2, and then 160 mL of KMnO4 (0.032 mol) was added

89

slowly to the above mixture while stirring vigorously to form a dark gray precipitate. After another hour of stirring,

90

the container was covered and placed in an oven at 55 oC for one day. The suspension was centrifuged at room

91

temperature and the supernatant solution was discarded. The centrifuged paste was washed with 1 M NaCl for five

92

times in the same way as the R;1 22. The centrifuged paste was then resuspended and washed with DI water for

93

multiple times until the pH dropped to 9.8. Both R;1 22 and triclinic birnessite were transferred to polypropylene

94

bottles as concentrated slurries and stored at 4 °C prior to use in experiments. X-ray powder diffraction (XRD)

95

confirmed the success of the syntheses.

96

Batch Experiments. Batch experiments on dissolved Mn(III) disproportionation were performed in an anaerobic

97

chamber (Coy Lab Products Inc., MI) where the O2 concentration was controlled below 1 ppmv. Two groups of 5


Page 7 of 31


98

batch experiments were conducted: (1) spontaneous dissolved Mn(III) disproportionation and (2) dissolved Mn(III)

99

disproportionation with initial addition of different Mn oxide minerals ( -MnO2 or triclinic birnessite). The ratio of

100

total PP:Mn(III) was 5 and the excess PP (pKa = 0.9, 2.0, 6.6, 9.4)51 served as a pH buffer that could be effective

101

over the pH range of interest.34 Solution pH was first adjusted to the value known to be able to reach the desired

102

final pH value (upon trial and error) and then aliquots of Mn(III)-PP stock solution were added to provide an initial

103

concentration of 400 QM Mn(III) and 2 mM PP. Those concentrations might appear to be higher than in natural

104

waters, but the results and models using PP as a model ligand could be extrapolated to geochemical contexts.

105

Addition of concentrated acid or base into the Mn(III) solutions was avoided based on previous experience that

106

undesired chemical transformation might happen before the solution became well mixed.34 For the experiments

107

with initial seeding of Mn minerals, different amounts of the Mn minerals were added as concentrated slurries into

108

the prepared volume (100 mL) of 400 Q1 Mn(III)-PP. No pH shift was observed after the addition of solids. Careful

109

and prompt handling of the experimental solutions ensured accurate recording of the time-resolved data. Samples

110

were periodically taken from the well-mixed reactors using syringes. Complementary experiments to examine

111

comproportionation between Mn(II) and Mn minerals in the presence of pyrophosphate were also conducted. The

112

detailed experimental conditions and key results are compiled in Table 1.

113

Analysis Methods. Samples from the batch experiments were filtered through a polyethersulfone (PES) filter

114

with a pore size of 0.05 Q

Half of the filtrate was preserved in 1% HNO3 for total dissolved Mn analysis by

115

inductively coupled plasma mass spectrometry (ICP-MS, PerkinElmer ELAN DRC II). The other half was

116

immediately analyzed for the concentration of dissolved Mn(III)-PP complexes using a widely adopted method29,

117

34, 52

118

Lambda XLS, S258nm > 6000 M-1cm-1, detection limit < 5 Q1? Selected solid samples were collected on a filter and

based on the solution’s absorbance at 258 nm measured with a UV-vis spectrophotometer (PerkinElmer-

6



Page 8 of 31

119

then freeze dried. They were characterized by X-ray diffraction (XRD, Cu KU radiation Bruker d8 Advance X-ray

120

diffractometer) to identify the mineralogical composition. Iodometric titration was used to determine the average

121

Mn oxidation state (AOS) of the solid products.49, 50 Hydrodynamic diameter and zeta potential of the precipitating

122

solids were measured on a Zetasizer Nano (Malvern). For dynamic light scattering analysis, disproportionation was

123

occurred with in the cuvette that was continuously measured in the instrument to generate time resolved results.

124

Thermodynamic and Kinetic modeling. The thermodynamic calculations of the Mn(III)-PP system were based

125

on critically selected equilibrium constants for relevant reactions in the literature and new experimental data

126

observed in the present study. The pe-pH diagrams were generated with the Act2 subprogram of Geochemist’s

127

Workbench (GWB, V.11.0.8, Aqueous Solution LLC, IL). The general thermodynamic database used in the

128

calculations was the “thermo_ladder” file published by Bethke et al.53, a reliable source where the Mn redox

129

equilibrium was critically crafted. PP4- was added as a basis component species with multiple protonated species

130

with different pKa values. Redox and hydrolysis reactions of PP species were ignored in the calculations, but the

131

acid-base speciation of PP at various pH was considered. Choices of solid phases to include or suppress in the

132

calculations were made based on low-temperature environmental relevance and experimental observation (details

133

in discussions). With the predominant species identified, calculations of the energetics and speciation of the

134

controlling equilibrium reactions were enabled by the Rxn subprogram of GWB, and a single equilibrium constant

135

for the dominant Mn-PP complex were optimized to simulate equilibrium Mn(III) concentrations in the present

136

experiments. Based on rational assumptions and experimental evidence, a set of differential equations were solved

137

to model the kinetic behaviors of Mn(III) disproportionation where rate constants were optimized to simulate the

138

time series datasets. Models were calibrated by minimizing the residual sum of squares (RSS) between the

139

experimental and predicted values using a spreadsheet solver.54 7


Page 9 of 31

140


RESULTS AND DISCUSSIONS

141

Effect of pH on the Stability of Mn(III)-PP Complexes. The stability of Mn(III)-PP complexes with respect to

142

disproportionation was sensitive to the solution pH (Figure 1a). At pH 6, 7, 8 and 8.3, Mn(III)-PP was very stable

143

and no loss of Mn(III)-PP from the solution was observed throughout the 60-hour experiment. It was consistent

144

with the literature that Mn(III)-PP complexes were stable at neutral pH for over two months under anaerobic

145

conditions.29, 44 This observed range of pH for Mn(III) stability overlaps with the pH range of common groundwaters

146

and surface waters. However, pH values outside of this range (i.e., 5.0, 8.6, 8.8 and 9.0) induced decreases in the

147

Mn-PP concentrations. The solution changed from clear pink to turbid brown, and the final solution after the brown

148

precipitates settled out still exhibited a light pink color, suggesting remaining Mn(III)-PP complexes, consistent

149

with the measured plateau of dissolved Mn(III) concentration.34 The total dissolved Mn concentration measured by

150

ICP-MS, which included dissolved Mn(III) and Mn(II), followed a similar trend. (Figure S1).

151

The time series of dissolved Mn(III) concentrations for the experiments with disproportionation suggested an

152

autocatalytic reaction in which the formation of an initial amount of Mn(IV)-containing solid accelerated

153

subsequent Mn(III) disproportionation. Mn(III) concentration declined slowly during the initial period followed by

154

a much more dramatic decrease (Figure 1b). The duration of the lag phase was shorter for those experiments with

155

a higher extent of Mn(III) disproportionation. The final dissolved Mn(III) concentration plateaued at various levels

156

at different pH. The production of Mn(II) could not be responsible for the autocatalytic effect given its much lower

157

affinity to PP,9 and this was confirmed by a control experiment with initially added Mn(II) (Figure S2).

158

Autocatalysis of aerobic Mn oxidation has been well known55, and Mn(III) disproportionation under anaerobic

159

conditions showed similar kinetics, suggesting a catalytic role of the produced Mn oxides.

8



Page 10 of 31

160

Evolution of the Particle Size and Zeta Potential. The time resolved monitoring of the average hydrodynamic

161

diameter of the precipitates upon disproportionation at pH 5 and pH 9 suggested different production and growth

162

behaviors of the Mn oxide particles (Figure 2). Despite the substantial noise of the DLS data with dynamic particle

163

generation and growth, moving average results identified significantly different features. The DLS data could not

164

provide concentration values of the colloidal system (the number concentration may be estimated using the Mn(III)

165

and total dissolved Mn concentration data), but it could assess the average particle size as the reaction proceeded.

166

Both colloidal systems had very negative zeta potential, as expected for phyllomanganate minerals56, 57, at which

167

particle aggregation should be negligible and the particles would presumably nucleate and grow individually. Steric

168

effects caused by ligand complexation on the Mn oxide surface may further stabilize the colloidal system. At pH 5,

169

the DLS hydrodynamic diameter showed an initial growth period that was consistent with the lag phase of the loss

170

of dissolved Mn(III) concentration, suggesting that relatively slower disproportionation was taking place with the

171

growing Mn oxides precipitates as catalytic surfaces to destabilize dissolved Mn(III). At low pH with slower rates

172

(i.e., higher activation energy), dissolved Mn(III) may prefer to catalytically disproportionate on existing Mn oxide

173

particles than in the bulk solution, resulting in the growing feature of the DLS particle size. At pH 9, no distinct

174

growing phase was observed, consistent with the faster disproportionation observed in the solution chemistry data.

175

The particle size at pH 9 was generally smaller than at pH 5, although the total number concentration at pH 9 was

176

much higher.

177

Characterization of the precipitates. X-ray diffraction results suggested that the precipitates produced during

178

Mn(III) disproportionation were a mixture of triclinic birnessite and -MnO2 (Figure 3). No peak corresponding to

179

pure Mn(III) (manganite or bixbyite) was present, excluding non-redox precipitation of Mn(III). The XRD pattern

180

of -MnO2 shows two predominant diffraction peaks at 37.3° (2.4 Å) and 66.4° (1.4 Å) Z and a broad feature 9


Page 11 of 31


181

between them, which are characteristic of poorly crystalline phyllomanganates.58, 59 The precipitates produced by

182

spontaneous Mn(III) disproportionation exhibited higher crystallinity and the solid products harvested at 13 hours

183

and 56 hours showed similar XRD patterns. Four peaks at 12.4°, 24.9°, 42.0°, and 50.0° matched triclinic birnessite,

184

and the two at 12.4° (7.2 Å) and 24.9° (3.6 Å) Z can be attributed to the (001) and (002) reflections.60 The two

185

peaks around 37° and 66° indicated the presence of -MnO2. When 100 Q1 of -MnO2 was initially added, the final

186

products of disproportionation still showed evidence of triclinic birnessite, while the relative abundance of the two

187

Mn oxides were different. Iodometric titration of the spontaneous disproportionation products (Exp. 8) gave an

188

average Mn oxidation state (AOS) of 3.4, whereas the values for synthetic -MnO2 and triclinic birnessite were 3.8

189

and 3.4, respectively. Even considering the possible amount of sorbed Mn(II), the AOS number suggested that the

190

products may more closely resemble triclinic birnessite than -MnO2. While the ideal stoichiometry of Mn(III)

191

disproportionation presumes a solid product of pure Mn(VI),9 our results suggest that the products (including those

192

produced after both 13 hours and 55 hours) still contained a substantial amount of Mn(III).

193

Based on the limited XRD results it was a challenge to unravel the in-situ and initial mineralogical transformation

194

pathways when disproportionation, mineral growth, adsorption of Mn(II) and structural migration of Mn(III) were

195

happening simultaneously. It is reported that the presence of Mn(II) can promote transformation of hexagonal

196

birnessite to triclinic birnessite in alkaline pH through comproportionation between Mn(II) and Mn(IV).13,

197

Conversely, triclinic birnessite can transform to hexagonal birnessite via either Mn(III) migration into the interlayers

198

or Mn(III) disproportionation to Mn(IV) and Mn(II) at acidic pH.62 Our solid characterization results shed light on

199

the possibility that the final products of Mn(III) disproportionation are controlled by chemical equilibrium as will

200

be discussed in a later section on modeling.

10


61


Page 12 of 31

201

Disproportionation with Pre-added Mn Oxides. Knowing that the precipitates produced upon spontaneous

202

disproportionation were a mixture of -MnO2 and triclinic birnessite as suggested by XRD results, experiments with

203

addition of pre-synthesized precipitates indicated that both -MnO2 and triclinic birnessite can indeed act as the

204

catalysts for disproportionation (Figures 4, S3). We also note that for the most stable Mn(III)-PP solution (e.g, pH

205

7, Exp. 3) Mn oxides still did not induce disproportionation when -MnO2 or triclinic birnessite was added (Figure

206

S4). At pH 9, the 2-hour lag phase was eliminated when -MnO2 or triclinic birnessite was added at the beginning

207

of the experiments. Nevertheless, the dissolved Mn(III) concentration for all experiments at a given pH with

208

different amounts of solid present plateaued at the same level, suggesting an equivalent thermodynamic equilibrium

209

that was only governed by pH (comparing Figures 1, 4, S3 and S4, considering the activity of solids as a constant).

210

For the same concentration (Mn atom normalized) of the two Mn oxides, -MnO2 appeared to be a stronger catalyst

211

than triclinic birnessite with respect to promoting Mn(III) disproportionation. Partial disproportionation of dissolved

212

Mn(III) complexes when the produced Mn oxides remained in the system was reported previously and could be

213

interpreted as the concurrent comproportionation between Mn oxides and Mn(II) and subsequent stabilization of

214

Mn(III) by PP.34, 43 Such comproportionation reaction has been previously investigated in ligand-free systems, and

215

the produced Mn(III) resided in solid phases such as feitknechtite and manganite.14,

216

experiments reacting aqueous Mn(II) and -MnO2 or triclinic birnessite in the presence of PP (Figure S5) confirmed

217

that comproportionation could occur in a time scale of hours as a mechanism to maintain residual dissolved Mn(III)

218

in the experiments just discussed. The experimental data above allowed us to develop thermodynamic and kinetic

219

models to describe the dynamic behaviors of dissolved Mn(III).

220 221

Thermodynamic Modeling.

15, 63-65

Complementary

Despite the lack of understanding of Mn(III) environmental speciation, there are

a few studies in the chemistry literature that reported the equilibrium constants of Mn(III)-PP complexes.66-70 We 11


Page 13 of 31


222

compiled each set of those constants (Table S1) and calculated the corresponding pe-pH diagrams for each set of

223

constants and reactions. The selection of Mn oxide mineral phases was based on the likelihood of natural presence

224

in low temperature soil and groundwater systems; consequently, pyrolusite, bixbyite and todorokite were suppressed

225

in the calculations.71, 72 The equilibrium constants related to Mn mineralogy were from Bethke et al.53 It should be

226

noted that the thermodynamic calculations here utilized birnessite (Mn8O19H10(s)) in the database, which was a good

227

analog of the disproportionation product for its mineralogy and average valance state of Mn (+3.5) noted previously.

228

Three sets of Mn(III)-PP equilibrium constants retrieved from the literature turned out to predict predominance of

229

dissolved Mn(III) only at highly acidic pH and above the stability limit of water (Figures S6, S7). This was

230

reasonable as previous studies of Mn(III) complexation used to determine these constants were performed at very

231

low pH. However, they could not predict the observed stability of Mn(III) at circumneutral pH.

232

The only exception was the set of equilibrium constants from Gordienko et al.66 in which a Mn(PP)25- species

233

was predicted to be stable over an environmentally relevant regime (Figure 5). Multiple calculations were performed

234

by varying the total PP to total Mn ratio, and the magnitude of the equilibrium constant for Mn(PP)25-. While the

235

prevalence regime shrinks with lower PP to Mn ratio, as long as the equilibrium constant is large enough, Mn-PP

236

can maintain its thermodynamic stability at circumneutral pH. At higher or lower pH beyond its predominance

237

regime, Mn(II) or Mn(IV) oxides are more favorable and thermodynamics drives disproportionation of Mn(III)

238

complexes. Mn(III)-PP speciation calculations with redox decoupled showed that non-redox decomplexation of

239

Mn-PP complex occurred only below pH 3, confirming that the observed Mn(III) loss should be attributed to

240

disproportionation (Figure S8).

241

While they could determine the predominant species, pe-pH diagrams could not quantify the equilibrium

242

concentration of each species that may co-exist. Given that our experimental results suggested that equilibrium was 12



Page 14 of 31

243

approached (plateau of the residual Mn(III)), they can be used to verify or refine the equilibrium constants. We

244

consider a controlling equilibrium reaction with the predominant species at specific pH as below.

245

2Mn(PP)25- + 19/6 H2O + 4/3 H+ = 1/6 Mn8O19H10(s) + 2/3 Mn2+ + 4 H2PP2- , 2.0 < pH < 6.6

(1)

246

2Mn(PP)25- + 19/6 H2O = 1/6 Mn8O19H10(s) + 2/3 MnPP2- + 10/3 HPP3- + 4/3 H+, 6.6 < pH < 9.4

(2)

247

2Mn(PP)25- + 19/6 H2O = 1/6 Mn8O19H10(s) + 2/3 MnPP2- +10/3 PP4- + 14/3 H+, pH > 9.4

(3)

248

The stoichiometric relationships and a stability constant of Mn(PP)25- could predict the equilibrium Mn speciation

249

for a given water chemistry condition, including the cases with partial disproportionation (Figure 6). It was

250

remarkable that a single stability constant of Mn(PP)25- with a logK of 28.9, which is lower than the value of 30.9

251

from the literature66, was able to quantitatively predict the final equilibrium concentration of Mn(III) with partial

252

disproportionation at alkaline pH. The experiment duration (60 h) at pH 5 did not capture an unambiguous plateau

253

so that the pH 5 results were not used to constrain the stability constant of Mn(PP)25-. Nevertheless, the logK of

254

28.9 predicted the favorability of disproportionation at pH 5 but not at pH 6 (Figures 1 and 6), which was consistent

255

with our observations. The fact that initial addition of pre-synthesized Mn oxides did not change the final

256

equilibrium Mn(III) concentration (Figures 1, 4, S3 and S4) was in line with the thermodynamic calculations above.

257

These calculations and their consistency with the experimental data suggest that the commonly observed co-

258

existence of Mn(II)-Mn(III)-Mn(IV) species14, 15, 34, 43 over a wide range of pH might be predominantly governed

259

by thermodynamics.

260

Kinetic Modeling. With the final equilibrium Mn(III) concentrations at each pH constrained by thermodynamics,

261

we developed a mathematical model to describe the kinetics of Mn(III) disproportionation using formulations for

262

autocatalysis analogous to Mn(II) oxidation by dissolved oxygen as by Morgan.73, 74 Several assumptions were made

263

to simplify the model: (1) the homogenous disproportionation rate of Mn(III) was first order; (2) the surface 13


Page 15 of 31


264

catalyzed disproportionation rate was first order with respect to dissolved Mn(III) and first order with respect to Mn

265

oxide concentration; (3) the reaction followed the stoichiometry defined in Reactions 1-3 for the respective pH,

266

ignoring surface associated Mn(II) and Mn(III); (4) both homogenous and heterogeneous disproportionation were

267

constrained by the thermodynamic equilibrium. We acknowledge that the current model could be further refined by

268

considering surface and aqueous speciation, but with only two fitting parameters, the relative importance of various

269

processes could be quantified as below

270

Mn(III)]

=

(1

[Mn(III)] [Mn(III)]eq

)[Mn(III)] (1


)[Mn(III)][MnO ]

(4)

271

where k is the homogenous disproportionation rate constant (M-1·s-1), k’ is the surface-catalyzed disproportionation

272

rate constant (M-2·s-1), and [Mn(III)]eq is the equilibrium Mn(III) concentration based on the preceding

273

thermodynamic analysis. The third assumption allows the growing value of the concentration of the Mn-containing

274

solid product birnessite (quantified on the basis of Mn atoms) to be expressed by [Mn(III)] based on the

275

stoichiometric mass balance, such that

276

Mn(III)]

=

(1


)[Mn(III)] (1


)[Mn(III)]

2 3([Mn(III)]0

[Mn(III)]) (5)

277

where [Mn(III)]0 denotes the initial concentration. Numerical solution of Eq. 5 should reproduce the experimental

278

data of Mn(III) lost during disproportionation at various pH as shown in Figure 1. As the [Mn(III)]eq as a function

279

of pH was determined by the thermodynamic model, the kinetic model only had k and k’ as parameters to be

280

determined by finding the best fit to the data. The model calibrated using dissolved Mn(III) time series data provided

281

an excellent fit to the total dissolved Mn concentration (Figures S1, S3), suggesting that the model assumption of

282

the reaction stoichiometry, i.e., Reaction 2, (the average valence state of Mn and the quantity of dissolved Mn(II)

283

production) was acceptable. Slight underestimation of total dissolved Mn loss might be due to unaccounted

284

adsorption of Mn(II). Despite the success of the model with the assumption of a single Mn(PP)25- species in 14



Page 16 of 31

285

simulating the present disproportionation data, the real stoichiometry and protonation states of Mn(III)-PP

286

complexes may vary with water chemistry conditions and thus exhibit different kinetic behaviors.

287

The kinetic model framework, upon a change of the initial Mn oxide concentration, could simulate the

288

elimination of the initial lag phase when synthetic pure Mn oxide (triclinic birnessite or R;1 22) was added at the

289

beginning of the experiments (Figures 4 and S3). The homogenous disproportionation rate constant (k) was kept

290

intact at the value obtained from the spontaneous disproportionation experiments and the surface catalyzed

291

disproportionation rate constants (k’) were adjusted to provide an optimal fit of the kinetic data of Mn(III) loss.

292

Model calibration selected the time window of the initial first hour, before the accumulation of the

293

disproportionation products would make the solids a mixture. Comparing the extrapolation of the model to longer

294

durations with the experimental data over the full duration provided information about how different the in-situ

295

formed precipitate was from the initially seeded solids with respect to their catalytic properties. The k’ obtained

296

from the spontaneous disproportionation was higher than that determined from the experiments with addition of

297

pre-synthesized triclinic birnessite, and it was lower than that with R;1 22 addition. These observations are

298

consistent with the fact that the precipitates from spontaneous disproportionation were a mixture of the two Mn

299

oxides. As k’ was expressed by molar concentration on the basis of Mn, the larger value for R;1 22 might be

300

attributed to its larger specific surface area that would provide more catalytic sites per unit mass of solid.50 The

301

model calibrated with the first hour of Mn(III) data with R;1 22 overestimated later Mn(III) loss (Figures 4a, S3a),

302

suggesting that the precipitating products (likely the mixture mentioned above) were overall less reactive than pure

303

R;1 22 for catalyzing Mn(III) disproportionation . In contrast, the quality of the model fits for triclinic birnessite

304

was better over the entire experimental duration (Figures 4a, S3c). Combining these observations with the

15


Page 17 of 31


305

comparison of the k’ values for those experiments (Exp. 8-14), it appeared that the products might behave more

306

analogously to triclinic birnessite.

307

As discussed in previous sections, the autocatalytic mechanism looked analogous to that of Mn(II) oxidation in

308

which fast adsorption of Mn precursor on Mn oxides lowers the activation energy. Our kinetic model indicated that

309

a majority of the Mn(III) loss was driven by the surface catalyzed process (Figure S9). The identity of the final solid

310

products (mixture of MnIII and MnIV) and Mn(III) loss not accounted for by disproportionation (note the gap between

311

model and data) are in accord with this mechanism, although autocatalytic processes involving Mn often do not

312

yield unequivocal identification of the reaction mechanisms.55

313

Environmental Implications. With the emerging evidence for the presence of dissolved forms of Mn(III) in

314

aquatic environments and their roles in the fundamental mechanisms of Mn redox processes, understanding Mn(III)

315

stability and reactivity become critical to advancing knowledge of Mn environmental biogeochemistry. The aqueous

316

stability of dissolved Mn(III) would significantly enhance the mobility of oxidized Mn species in subsurface porous

317

media in contrast to particulate or colloidal Mn oxides.34, 75, 76 Disproportionation can occur when the solution pH

318

shifts out of its stability regime and when complexing ligands undergo chemical or microbial degradation. The

319

presence of ligands that can form stable aqueous complexes with Mn(III) allows the coexistence of three oxidation

320

states of Mn over a relatively wide pH range. Such coexistence of multiple oxidation states allows dynamic atomic

321

exchange and electron transfer at mineral-water interfaces that can affect the mobility of associated trace elements.63,

322

64, 77

323

The experimental approach using PP as a model ligand may be extended to other Mn(III) complexes. The pH

324

range of stability as well as the equilibrium residual Mn(III) can provide information about the binding affinity of

325

the ligands for Mn(III). Other ligands may be studied using the same approach to overcome the difficulty in 16



Page 18 of 31

326

measuring trace concentrations of free Mn(III). Considering the heterogeneity and time scales relevant to subsurface

327

environments, we anticipate that the autocatalytic disproportionation of dissolved Mn(III) could readily reach

328

equilibrium, which could be incorporated in future geochemical models that consider Mn(III) dynamics. Natural

329

organic matter (NOM) may have specific functional groups such as hydroxylamine and catechol groups that can

330

complex with Mn(III).21, 75 It is possible to follow a bottom up approach by using model ligands to individually

331

establish the speciation of Mn(III) complexes and then extrapolate to a continuum mixture of ligands present in

332

NOM. Caution should be taken when extrapolating the results of Mn-PP system to organic ligands because of the

333

potential intramolecular electron transfer from the organic molecules to Mn(III), which is more favorable at acidic

334

pH; in contrast PP cannot reduce Mn(III) itself, but it can undergo hydrolysis and destabilize Mn(III) towards

335

disproportionation. Therefore, consumption of Mn(III) may occur through different pathways for different ligands.

336

Disproportionation of Mn(III) has been recently recognized as a critical step in microbial production of Mn

337

oxides.45,

46

338

“controlled” disproportionation of Mn(III) in the presence of excess complexing ligands. The products of Mn(III)

339

disproportionation may have properties that depend on the disproportionation rates that are affected by the

340

concentration and Mn(III)-complexing affinity of ligands present 47. The reactivity of soluble Mn(III) species as

341

geochemical oxidants warrants detailed investigation to fill the knowledge gaps of environmental occurrence,

342

chemical speciation, redox activity, and biological roles of aqueous Mn(III) complexes. Those knowledge gaps may

343

challenge our current assessments of the relativities of actual environmental Mn oxides based on laboratory studies

344

using synthetic analogs.

The present study highlighted the mixture of triclinic birnessite and R;1 22 as the products of

345 346

ASSOCIATED CONTENT 17


Page 19 of 31


347

Supporting Information. The Supporting Information is available free of charge on the ACS Publications website

348

at DOI ---”.

349

Data and models for total dissolved Mn variation with time; additional control experiments; additional data and

350

model simulations for the experiments with initially seeded Mn oxides; complementary experiments for

351

comproportionation, aqueous speciation equilibrium of Mn(III)-PP, additional pe-pH diagram calculations using

352

other sets of equilibrium reactions and constants; model calculated homogenous and surface catalyzed

353

disproportionation rates, and compilation of the model reactions and constants. (PDF)

354 355

ACKNOWLEDGMENTS

356

We respectfully acknowledge insightful suggestions and comments of Profs. James Morgan and Alan Stone.

357

This study received financial support from National Natural Science Foundation of China (21806021, 41807187)

358

and Louisiana Board of Regents (LEQSF(2017-20)-RD-A-07). A.Q.’s participation was sponsored by a visiting

359

student fellowship from China Scholarship Council. Comments and suggestions of Associate Editor Prof. Timm

360

Strathmann and three other anonymous reviewers improved the quality of the manuscript.

361 362

REFERENCES

363 364 365 366 367 368 369 370 371 372 373 374

1. Morgan, J. J., Manganese in natural waters and earth’s crust: Its availability to organisms. In Met. Ions Biol. Syst., Sigel, A.; Sigel, H., Eds. Marcel Dekker New York: 2000; Vol. 37, pp 1-34. 2. Cox, P. A., Inorganic Chemistry in the Environment. Oxford University Press: New York 1995. 3. Lewis, B.; Landing, W., The biogeochemistry of manganese and iron in the Black Sea. Deep Sea Res. Pt. 1 1991, 38, S773-S803. 4. Luther, G. W., The role of one-and two-electron transfer reactions in forming thermodynamically unstable intermediates as barriers in multi-electron redox reactions. Aqua. Geochem. 2010, 16, (3), 395-420. 5. Murray, J. W.; Dillard, J. G.; Giovanoli, R.; Moers, H.; Stumm, W., Oxidation of Mn (II): Initial mineralogy, oxidation state and ageing. Geochim. Cosmochim. Acta 1985, 49, (2), 463-470. 6. Junta, J. L.; Hochella, M. F., Manganese (II) oxidation at mineral surfaces: A microscopic and spectroscopic study. Geochim. Cosmochim. Acta 1994, 58, (22), 4985-4999. 7. Mandernack, K. W.; Post, J.; Tebo, B. M., Manganese mineral formation by bacterial spores of the marine 18



375 376 377 378 379 380 381 382 383 384 385 386 387 388 389 390 391 392 393 394 395 396 397 398 399 400 401 402 403 404 405 406 407 408 409 410 411 412 413 414 415 416

Page 20 of 31

Bacillus, strain SG-1: Evidence for the direct oxidation of Mn (II) to Mn (IV). Geochim. Cosmochim. Acta 1995, 59, (21), 4393-4408. 8. Bargar, J. R.; Tebo, B. M.; Bergmann, U.; Webb, S. M.; Glatzel, P.; Chiu, V. Q.; Villalobos, M., Biotic and abiotic products of Mn (II) oxidation by spores of the marine Bacillus sp. strain SG-1. Am. Mineral. 2005, 90, (1), 143-154. 9. Davies, G., Some aspects of the chemistry of manganese (III) in aqueous solution. Coord. Chem. Rev. 1969, 4, (2), 199-224. 10. Nesbitt, H.; Canning, G.; Bancroft, G., XPS study of reductive dissolution of 7Å-birnessite by H 3 AsO 3, with constraints on reaction mechanism. Geochim. Cosmochim. Acta 1998, 62, (12), 2097-2110. 11. Banerjee, D.; Nesbitt, H., Oxidation of aqueous Cr (III) at birnessite surfaces: Constraints on reaction mechanism. Geochim. Cosmochim. Acta 1999, 63, (11), 1671-1687. 12. Lin, H.; Szeinbaum, N. H.; DiChristina, T. J.; Taillefert, M., Microbial Mn (IV) reduction requires an initial one-electron reductive solubilization step. Geochim. Cosmochim. Acta 2012, 99, 179-192. 13. Zhao, H.; Zhu, M.; Li, W.; Elzinga, E. J.; Villalobos, M.; Liu, F.; Zhang, J.; Feng, X.; Sparks, D. L., Redox reactions between Mn(II) and hexagonal birnessite change its layer symmetry. Environ. Sci. Technol. 2016, 50, (4), 1750-1758. 14. Elzinga, E. J., Reductive transformation of birnessite by aqueous Mn(II). Environ. Sci. Technol. 2011, 45, (15), 6366-6372. 15. Lefkowitz, J. P.; Rouff, A. A.; Elzinga, E. J., Influence of pH on the reductive transformation of birnessite by aqueous Mn(II). Environ. Sci. Technol. 2013, 47, (18), 10364-10371. 16. Johnson, K. S., Manganese redox chemistry revisited. Science 2006, 313, (5795), 1896-1897. 17. Oldham, V. E.; Owings, S. M.; Jones, M. R.; Tebo, B. M.; Luther, G. W., Evidence for the presence of strong Mn (III)-binding ligands in the water column of the Chesapeake Bay. Mar. Chem. 2015, 171, 58-66. 18. Trouwborst, R. E.; Clement, B. G.; Tebo, B. M.; Glazer, B. T.; Luther, G. W., Soluble Mn(III) in suboxic zones. Science 2006, 313, (5795), 1955-1957. 19. Yakushev, E. V.; Pollehne, F.; Jost, G.; Kuznetso, I.; Schneider, B.; Urnlauf, L., Analysis of the water column oxic/anoxic interface in the Black and Baltic seas with a numerical model. Mar. Chem. 2007, 107, (3), 388-410. 20. Dellwig, O.; Schnetger, B.; Brumsack, H.-J.; Grossart, H.-P.; Umlauf, L., Dissolved reactive manganese at pelagic redoxclines (part II): Hydrodynamic conditions for accumulation. Journal of Marine Systems 2012, 90, (1), 31-41. 21. Madison, A. S.; Tebo, B. M.; Mucci, A.; Sundby, B.; Luther, G. W., Abundant porewater Mn (III) is a major component of the sedimentary redox system. Science 2013, 341, (6148), 875-878. 22. Oldham, V. E.; Miller, M. T.; Jensen, L. T.; Luther, G. W., Revisiting Mn and Fe removal in humic rich estuaries. Geochim. Cosmochim. Acta 2017, 209, 267-283. 23. Jones, M. R.; Oldham, V. E.; Luther, G. W.; Mucci, A.; Tebo, B. M., Distribution of desferrioxamine-Bextractable soluble manganese(III) and particulate MnO2 in the St. Lawrence Estuary, Canada. Mar. Chem. 2019, 208, (20), 70-82. 24. Johnson, K. L.; McCann, C. M.; Wilkinson, J.-L.; Jones, M.; Tebo, B. M.; West, M.; Elgy, C.; Clarke, C. E.; Gowdy, C.; Hudson-Edwards, K. A., Dissolved Mn (III) in water treatment works: Prevalence and significance. Water Res. 2018, 140, 181-190. 25. Oldham, V. E.; Mucci, A.; Tebo, B. M.; Luther, G. W., Soluble Mn (III)–L complexes are abundant in oxygenated waters and stabilized by humic ligands. Geochim. Cosmochim. Acta 2017, 199, 238-246. 19


Page 21 of 31

417 418 419 420 421 422 423 424 425 426 427 428 429 430 431 432 433 434 435 436 437 438 439 440 441 442 443 444 445 446 447 448 449 450 451 452 453 454 455 456 457 458


26. Harrington, J. M.; Parker, D. L.; Bargar, J. R.; Jarzecki, A. A.; Tebo, B. M.; Sposito, G.; Duckworth, O. W., Structural dependence of Mn complexation by siderophores: Donor group dependence on complex stability and reactivity. Geochim. Cosmochim. Acta 2012, 88, (0), 106-119. 27. Wang, Y.; Stone, A. T., Phosphonate- and carboxylate-based chelating agents that solubilize (hydr)oxidebound MnIII. Environ. Sci. Technol. 2008, 42, (12), 4397-4403. 28. Luther III, G. W.; Madison, A. S.; Mucci, A.; Sundby, B.; Oldham, V. E., A kinetic approach to assess the strengths of ligands bound to soluble Mn(III). Mar. Chem. 2015, 173, (20), 93-99. 29. Kostka, J. E.; Luther, G. W.; Nealson, K. H., Chemical and biological reduction of Mn (III)-pyrophosphate complexes: potential importance of dissolved Mn (III) as an environmental oxidant. Geochim. Cosmochim. Acta 1995, 59, (5), 885-894. 30. Hu, E.; Zhang, Y.; Wu, S.; Wu, J.; Liang, L.; He, F., Role of dissolved Mn (III) in transformation of organic contaminants: Non-oxidative versus oxidative mechanisms. Water Res. 2017, 111, 234-243. 31. Sherigara, B.; Bhat, K. I.; Pinto, I.; Gowda, N. M., Oxidation of L aspartic acid and L glutamic acid by manganese (III) ions in aqueous sulphuric acid, acetic acid, and pyrophosphate media: A kinetic study. Int. J. Chem. Kinet. 1995, 27, (7), 675-690. 32. Szeinbaum, N.; Lin, H.; Brandes, J. A.; Taillefert, M.; Glass, J. B.; DiChristina, T. J., Microbial manganese (III) reduction fuelled by anaerobic acetate oxidation. Environ. Microbiol. 2017, 19, (9), 3475-3486. 33. Jones, M. E.; Nico, P. S.; Ying, S.; Regier, T.; Thieme, J. r.; Keiluweit, M., Manganese-driven carbon oxidation at oxic–anoxic interfaces. Environ. Sci. Technol. 2018, 52, (21), 12349-12357. 34. Wang, Z.; Xiong, W.; Tebo, B. M.; Giammar, D. E., Oxidative UO2 dissolution induced by soluble Mn(III). Environ. Sci. Technol. 2014, 48, (1), 289-98. 35. Perez-Benito, J. F.; Arias, C., The pyrophosphate-assisted reduction of chromium (VI) by manganese (II) and its reverse reaction. New J. Chem. 2001, 25, (11), 1438-1446. 36. Murray, K. J.; Tebo, B. M., Cr (III) is indirectly oxidized by the Mn (II)-oxidizing bacterium Bacillus sp. strain SG-1. Environ. Sci. Technol. 2007, 41, (2), 528-533. 37. Gunary, D., Pyrophosphate in soil; some physico-chemical aspects. Nature 1966, 210, 1297-1298. 38. Turner, B. L.; Newman, S., Phosphorus cycling in wetland soils. J. Environ. Qual. 2005, 34, (5), 1921-1929. 39. Renshaw, J. C.; Robson, G. D.; Trinci, A. P.; Wiebe, M. G.; Livens, F. R.; Collison, D.; Taylor, R. J., Fungal siderophores: structures, functions and applications. Mycol. Res. 2002, 106, (10), 1123-1142. 40. Kiss, T.; Farkas, E., Metal-binding ability of desferrioxamine B. J. Inclusion Phenom. Mol. Recognit. Chem. 1998, 32, (2-3), 385-403. 41. Duckworth, O. W.; Sposito, G., ) . ` manganese (III) interactions II. Manganite dissolution promoted by desferrioxamine B. Environ. Sci. Technol. 2005, 39, (16), 6045-6051. 42. Wang, Y.; Stone, A. T., The citric acid–Mn III, IV O 2 (birnessite) reaction. Electron transfer, complex formation, and autocatalytic feedback. Geochim. Cosmochim. Acta 2006, 70, (17), 4463-4476. 43. Klewicki, J. K.; Morgan, J. J., Kinetic behavior of Mn(III) complexes of pyrophosphate, EDTA, and citrate. Environ. Sci. Technol. 1998, 32, (19), 2916-2922. 44. Klewicki, J.; Morgan, J., Dissolution of a;1 22 particles by ligands: pyrophosphate, ethylenediaminetetraacetate, and citrate. Geochim. Cosmochim. Acta 1999, 63, (19), 3017-3024. 45. Soldatova, A. V.; Tao, L.; Romano, C. A.; Stich, T. A.; Casey, W. H.; Britt, R. D.; Tebo, B. M.; Spiro, T. G., Mn (II) oxidation by the multicopper oxidase complex Mnx: a binuclear activation mechanism. J. Am. Chem. Soc. 2017, 139, (33), 11369-11380. 20



459 460 461 462 463 464 465 466 467 468 469 470 471 472 473 474 475 476 477 478 479 480 481 482 483 484 485 486 487 488 489 490 491 492 493 494 495 496 497 498 499 500

Page 22 of 31

46. Soldatova, A. V.; Romano, C. A.; Tao, L.; Stich, T. A.; Casey, W. H.; Britt, R. D.; Tebo, B. M.; Spiro, T. G., Mn (II) oxidation by the multicopper oxidase complex Mnx: a coordinated two-stage Mn (II)/(III) and Mn (III)/(IV) Mechanism. J. Am. Chem. Soc. 2017, 139, (33), 11381-11391. 47. Jung, H.; Chadha, T. S.; Min, Y.; Biswas, P.; Jun, Y.-S., Photochemically-Assisted Synthesis of Birnessite Nanosheets and Their Structural Alteration in the Presence of Pyrophosphate. ACS Sustainable Chemistry & Engineering 2017, 5, (11), 10624-10632. 48. Banerjee, D.; Nesbitt, H., XPS study of reductive dissolution of birnessite by oxalate: Rates and mechanistic aspects of dissolution and redox processes. Geochim. Cosmochim. Acta 1999, 63, (19), 3025-3038. 49. Murray, J. W.; Balistrieri, L. S.; Paul, B., The oxidation state of manganese in marine sediments and ferromanganese nodules. Geochim. Cosmochim. Acta 1984, 48, (6), 1237-1247. 50. Villalobos, M.; Toner, B.; Bargar, J.; Sposito, G., Characterization of the manganese oxide produced by pseudomonas putida strain MnB1. Geochim. Cosmochim. Acta 2003, 67, (14), 2649-2662. 51. McElroy, W. D.; Glass, B., Phosphorus Metabolism. Johns Hopkins University Press: Baltimore, 1951; Vol. I. 52. Webb, S. M.; Dick, G. J.; Bargar, J. R.; Tebo, B. M., Evidence for the presence of Mn(III) intermediates in the bacterial oxidation of Mn(II). Proc. Natl. Acad. Sci. U. S. A. 2005, 102, (15), 5558-5563. 53. Bethke, C. M.; Sanford, R. A.; Kirk, M. F.; Jin, Q.; Flynn, T. M., The thermodynamic ladder in geomicrobiology. Am. J. Sci. 2011, 311, (3), 183-210. 54. Xie, X.; Giammar, D. E.; Wang, Z., MINFIT: A spreadsheet-based tool for parameter estimation in an equilibrium speciation software program. Environ. Sci. Technol. 2016, 50, (20), 11112-11120. 55. Stumm, W.; Morgan, J. J., Aquatic Chemistry. 3rd ed.; Wiley Interscience: New York, 1996. 56. Wang, Z.; Lee, S.-W.; Catalano, J. G.; Lezama-Pacheco, J. S.; Bargar, J. R.; Tebo, B. M.; Giammar, D. E., Adsorption of uranium(VI) to manganese oxides: X-ray absorption spectroscopy and surface complexation modeling. Environ. Sci. Technol. 2013, 47, (2), 850-858. 57. Zhao, W.; Cui, H.; Liu, F.; Tan, W.; Feng, X., Relationship between Pb2+ and average Mn oxidation state in synthetic birnessites. Clays Clay Miner. 2009, 57, (5), 513-520. 58. Villalobos, M.; Lanson, B.; Manceau, A.; Toner, B.; Sposito, G., Structural model for the biogenic Mn oxide produced by Pseudomonas putida. Am. Mineral. 2006, 91, (4), 489-502. 59. Webb, S. M.; Tebo, B.; Bargar, J., Structural characterization of biogenic Mn oxides produced in seawater by the marine Bacillus sp. strain SG-1. Am. Mineral. 2005, 90, (8-9), 1342-1357. 60. Lopano, C. L.; Heaney, P. J.; Post, J. E.; Hanson, J.; Komarneni, S., Time-resolved structural analysis of Kand Ba-exchange reactions with synthetic Na-birnessite using synchrotron X-ray diffraction. Am. Mineral. 2007, 92, (2-3), 380-387. 61. Zhu, M.; Ginder-Vogel, M.; Parikh, S. J.; Feng, X.-H.; Sparks, D. L., Cation effects on the layer structure of biogenic Mn-oxides. Environ. Sci. Technol. 2010, 44, (12), 4465-4471. 62. Lanson, B.; Drits, V. A.; Silvester, E.; Manceau, A., Structure of H-exchanged hexagonal birnessite and its mechanism of formation from Na-rich monoclinic buserite at low pH. Am. Mineral. 2000, 85, (5-6), 826-838. 63. Elzinga, E. J., 54Mn radiotracers demonstrate continuous dissolution and reprecipitation of vernadite >R;1 2 ? during interaction with aqueous Mn (II). Environ. Sci. Technol. 2016, 50, (16), 8670-8677. 64. Hinkle, M. A.; Dye, K. G.; Catalano, J. G., Impact of Mn (II)-manganese oxide reactions on Ni and Zn speciation. Environ. Sci. Technol. 2017, 51, (6), 3187-3196. 65. Wang, Q.; Yang, P.; Zhu, M., Structural transformation of birnessite by fulvic acid under anoxic conditions. 21


Page 23 of 31

501 502 503 504 505 506 507 508 509 510 511 512 513 514 515 516 517 518 519 520 521 522 523 524 525 526


Environ. Sci. Technol. 2018, 52, (4), 1844-1853. 66. Gordienko, V. I.; Sidorenko, V. I.; Mikhailyuk, Y. I., Amperometric investigation of Mn(III) pyrophosphate complexes. Russ. J. Inorg. Chem 1970, 15, 1241-1244. 67. Martell, A. E.; Smith, R. M.; Motekaitis, R. J., NIST Critically Selected Stability Constants of Metal Complexes. In NIST Standard Reference Database 46, Version 6.0. NIST, Gaitherburg: 2001. 68. Ciavatta, L.; Palombari, R., On the equilibria of complex formation between manganese(III) and pyrophosphate ions. Gazz. Chim. Ital. 1983, 113, 557-552. 69. Kolthoff, I.; Watters, J., Polarographic determination of manganese as tridihydrogen pyrophosphatomanganiate. Ind. Eng. Chem. Anal. Edit. 1943, 15, (1), 8-13. 70. Bogdanovich, N.; Pechurova, N.; Martynenko, L.; Pinnova, V., Reaction of manganese (III) with complexones. Russ. J. Inorg. Chem. 1971, 16, 1337-1339. 71. Wang, Z.; Giammar, D. E., Metal contaminant oxidation mediated by manganese redox cycling in subsurface environment. In Advances in the Environmental Biogeochemistry of Manganese Oxides, Feng, X. H.; Li, W.; Zhu, M.; Sparks, D. L., Eds. American Chemical Society: Washington DC, 2015; Vol. 1197, pp 29-50. 72. Tebo, B. M.; Bargar, J. R.; Clement, B. G.; Dick, G. J.; Murray, K. J.; Parker, D.; Verity, R.; Webb, S. M., Biogenic manganese oxides: Properties and mechanisms of formation. Annu. Rev. Earth Planet. Sci. 2004, 32, (1), 287-328. 73. Morgan, J. J., Kinetics of reaction between O2 and Mn(II) species in aqueous solutions. Geochim. Cosmochim. Acta 2005, 69, (1), 35-48. 74. Morgan, J. J. Chemistry of Aqueous Manganese II and IV: A Thesis. Harvard University, 1964. 75. Li, Q.; Xie, L.; Jiang, Y.; Fortner, J. D.; Yu, K.; Liao, P.; Liu, C., Formation and stability of NOM-Mn (III) colloids in aquatic environments. Water Res. 2019, 149, 190-201. 76. Wang, Z.; Tebo, B. M.; Giammar, D. E., Effects of Mn(II) on UO2 dissolution under anoxic and oxic conditions. Environ. Sci. Technol. 2014, 48, (10), 5546-5554. 77. Hinkle, M. A.; Flynn, E. D.; Catalano, J. G., Structural response of phyllomanganates to wet aging and aqueous Mn (II). Geochim. Cosmochim. Acta 2016, 192, 220-234.

527 528 529

22



530

Table of Content Art

531

23


Page 24 of 31

Page 25 of 31


532

533 534

Figure 1. Kinetics of Mn(III)-PP disproportionation at various pH. Panel (b) provides a zoom-in view of the

535

initial dynamics of the autocatalytic behavior. Symbols and lines were experimental data and model simulations,

536

respectively. Error bars indicated standard deviation of at least duplicate experiments, some of which might be

537

shown smaller than the size of the symbols. Data points for pH 6.0, 7.0, 8.0 and 8.3 had substantial overlapping.

538

Detailed experimental conditions are summarized in Table 1 (Exp. 1-8).

539

24




Page 26 of 31

Page 27 of 31


546 547 548

Figure 3. Mineralogical identification of the disproportionation products by X-ray diffraction (Blue: Exp. 8;

549

Green: Exp. 11)

550

26



Page 28 of 31

551

552 553

Figure 4. Kinetics of Mn(III) disproportionation with initially added R;1 22 (Exp. 9) and triclinic birnessite

554

(Exp. 12) at different concentrations. Data and models for Exp. 8 (no Mn mineral) were included for a

555

comparison. Panels a and b presents dissolved Mn(III) and total dissolved Mn, respectively. Symbols and lines

556

were experimental data and model simulation, respectively. Error bars indicated standard deviation of at least

557

duplicate experiments, some of which might be shown smaller than the size of the symbols. Note that the models

558

were calibrated only using the kinetic data in the initial first hour before the solids developed into a mixture of the

559

two oxides. Additional data using different dosages of R;1 22 and triclinic birnessite are presented in Figure S3.

560

27


Page 29 of 31




Page 30 of 31

Residual MnIII ( M)

500 400 300 200 100 0 4

6

8 pH

10

12

571 572

Figure 6. Experimentally observed and model predicted final residual Mn(III) concentration at different pH.

573

Initial total Mn(III) concentration was 400 Q

574

separate the pH ranges with different equilibrium reactions considered in the model calculations. The

575

experimentally observed Mn(III) concentration profile at pH 5 after 60 hours could not rule out even lower

576

equilibrium concentration, as denoted by the downward arrow. The residual concentration value for pH 5 as of 60

577

hours might merely be the upper bound and was not included in the optimization of the thermodynamic model.

and the total PP concentration was 2 mM. The vertical dashed lines

578

29


Page 31 of 31

579 580 581


Table 1. Summary of the Experimental Conditions, Key Results and Thermodynamic and Kinetic Model Parameters Exp

582 583 584 585 586 587 588

pH

MnIII

PP

MnII

R;1 22

Tric. Birn.

k (×103)

>Q1? a

(1/s) b

1/(M s)

>Q1?

>Q1?

>Q1?

>Q1?a

k' (×103) b

Measured MnIII

eq

>Q1?

Predicted MnIIIeq >Q1?c

1

5

400

2000

0

0

0

0.09

0.03

N.A. d

57.5

2

6

400

2000

0

0

0

N.A.

N.A.

400

399

3

7

400

2000

0

0

0

N.A.

N.A.

400

400

4

8

400

2000

0

0

0

N.A.

N.A.

400

400

5

8.3

400

2000

0

0

0

N.A.

N.A.

400

381

6

8.6

400

2000

0

0

0

0.92

2.18

297

276

7

8.8

400

2000

0

0

0

0.99

2.75

214

227

8

9

400

2000

0

0

0

2.20

3.30

180

181

9

9

400

2000

0

20

0

2.20

6.37

181

181

10

9

400

2000

0

50

0

2.20

7.29

184

181

11

9

400

2000

0

100

0

2.20

8.51

180

181

12

9

400

2000

0

0

20

2.20

2.04

217

181

13

9

400

2000

0

0

50

2.20

2.10

197

181

14

9

400

2000

0

0

100

2.20

2.12

187

181

15

9

0

2000

200

200

0

N.A.

N.A.

121

N.A.

16

9

0

2000

200

0

200

N.A.

N.A.

76.5

N.A.

a. The concentration of initially seeded R;1 22 and triclinic birnessite was quantified on the basis of Mn atom. b. k and k’ denote the rate constants for homogeneous and surface catalyzed disproportionation pathways. c. Prediction of the final Mn(III) concentrations was enabled by the thermodynamic calculation using the optimized logK for Mn(PP)25-. d. For pH 5, our experimental results up to 60 hours didn’t capture the plateau of the residual Mn(III) concentration approaching the predicted value. The observed residual Mn(III) after 60 hours was not applicable to compare with the calculated value.

30


Geochemical Stability of Dissolved Mn(III) in the Presence of

Recommend Documents