Dissertation-TZ Final

Mechanisms of Microbial Formation and Photodegradation of Methylmercury in the Aquatic Environment

Tong Zhang

Department of Civil and Environmental Engineering Duke University

Date:______Approved:

______Heileen Hsu-Kim, Supervisor

______Marc Deshusses

______Mark Wiesner

______Claudia Gunsch

______Heather Stapleton

Dissertation submitted in partial fulfillment of the requirements for the degree of Doctor of Philosophy in the Department of Civil and Environmental Engineering in the Graduate School of Duke University

2012 ABSTRACT

Mechanisms of Microbial Formation and Photodegradation of Methylmercury in the Aquatic Environment

Tong Zhang

Department of Civil and Environmental Engineering Duke University

Date:______Approved:

______Heileen Hsu-Kim, Supervisor

______Marc Deshusses

______Mark Wiesner

______Claudia Gunsch

______Heather Stapleton

An abstract of a dissertation submitted in partial fulfillment of the requirements for the degree of Doctor of Philosophy in the Department of Civil and Environmental Engineering in the Graduate School of Duke University

2012

Abstract

Methylmercury is a bioaccumulative neurotoxin that severely endangers human health. Humans are exposed to methylmercury through consumption of contaminated aquatic fish. To date, effective strategies for preventing and remediating methylmercury contamination have remained elusive, mainly due to the lack of knowledge in regard to how methylmercury is generated and degraded in the aquatic environment. The goal of this dissertation was to study the mechanisms of two transformation processes that govern the fate of methylmercury in natural settings: microbial mercury methylation and methylmercury photodegradation. The role of mercury speciation (influenced by environmental conditions) in determining the reactivity of mercury in these biological and photochemical reactions was the focus of this research.

Methylmercury production in the aquatic environment is primarily mediated by anaerobic bacteria in surface sediments, particularly sulfate reducing bacteria (SRB). The efficiency of this process is dependent on the activity of the methylating bacteria and the availability of inorganic divalent mercury (Hg(II)). In sediment pore waters, Hg(II) associates with sulfides and dissolved organic matter (DOM) to form a continuum of chemical species that include dissolved molecules, polynuclear clusters, amorphous nanoparticles and after long term aging, bulk-scale crystalline particles. The methylation potential of these mercury species were examined using both pure cultures of SRB and sediment slurry microcosms. The results of these experiments indicated that the activity of SRB was largely determined by the supply of sulfate and labile carbon, which significantly influenced the net methylmercury production in sediment slurries. The

availability of mercury for methylation decreased during aging. Dissolved Hg-sulfide

(added as Hg(NO 3)2 and Na 2S) resulted in the highest methylmercury production.

Although the methylation potential of humic-coated HgS nanoparticles decreased with an increase in the age of nanoparticle stock solutions, nano-HgS was substantially more available for microbial methylation relative to microparticulate HgS, possibly due to the smaller size, larger specific surface area and more disordered structure of the nanoparticles. Moreover, the methylation of mercury derived from nanoparticles cannot be explained by equilibrium speciation of mercury in the aqueous phase (<0.2 µm, the currently-accepted approach for assessing mercury bioavailability for methylation).

Instead, the methylation potential of mercury sulfides appeared to correlate with the extent of dissolution and their reactivity in thiol ligand exchange. Additionally, partitioning of mercury to a diverse group of bulk-scale mineral particles and colloids

(especially FeS) may be an important process controlling the mercury speciation and subsequent methylmercury production in natural sediments.

In surface waters, sunlight degradation is believed to be the predominant pathway for the decomposition of methylmercury. The mechanism of this process was investigated in a series of photodegradation experiments under natural sunlight and

UV-A radiation, and in the presence of DOM and selective quenchers for photo- generated reactive intermediates. The results suggested that singlet oxygen generated from photosensitization of DOM drove the photodecomposition of methylmercury. The rate of methylmercury degradation depended on the type of methylmercury (CH 3Hg +) binding ligand present in the water. CH 3Hg -thiol (e.g., glutathione, mercaptoacetate,

DOM) complexes were significantly more reactive in photodegradation compared to other methylmercury complexes (CH 3HgCl or CH 3HgOH), which may be because thiol- binding can effectively decrease the activation energy and thus enhance the reactivity of methylmercury molecules toward the Hg-C bond breaking process. These findings challenge the long-accepted view that water chemistry characteristics do not affect the kinetics of methylmercury sunlight degradation, and help explain recent field observation that methylmercury photodegradation occurred rapidly in freshwater lakes

(where CH 3Hg-DOM dominate methylmercury speciation) but relatively slowly in sea water (where CH 3Hg-Cl control methylmercury speciation).

Overall, this dissertation has demonstrated that chemical speciation of inorganic mercury and methylmercury determines their availability for microbial methylation and sunlight degradation, respectively. The abundance of these available mercury species is influenced by a variety of environmental parameters (e.g., DOM). This dissertation work contributes mechanistic knowledge toward understanding the occurrence of methylmercury in the aquatic environment. This information will ultimately help construct quantitative models for accurately predicting and assessing the risks of mercury contamination.

Dedication

To my parents and Yao for being a steady source of love and support in my life.

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Contents

Abstract...... iv

List of tables...... xii

List of figures...... xiii

Acknowledgements ...... xvi

Chapter 1. Introduction ...... 1

1.1 Motivation ...... 1

1.2 The problem of methylmercury in the environment...... 5

1.2.1 Methylmercury and its toxicological effects ...... 5

1.2.2 Sources and transformations of methylmercury ...... 7

1.3 Microbial formation of methylmercury in natural sediments...... 8

1.3.1 Mechanisms of microbial mercury methylation...... 8

1.3.2 Chemical speciation and reactivity of inorganic mercury in natural sediments ...... 15

1.4 Photodegradation of methylmercury in natural waters ...... 24

1.4.1 Mechanisms of sunlight induced degradation ...... 24

1.4.2 Photochemistry of dissolved organic matter ...... 29

1.4.3 Chemical speciation and reactivity of methylmecury in natural waters ...... 32

1.5 Research objectives ...... 34

Chapter 2. Methylation of mercury by bacteria exposed to dissolved, nanoparticulate, and microparticulate mercuric sulfides...... 38

2.1 Introduction...... 38

2.2 Materials and Methods...... 40

2.2.1 Microorganisms and Culture Conditions...... 40

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2.2.2 HgS Particle Preparation ...... 41

2.2.3 HgS Particle Characterization ...... 42

2.2.4 Mercury Methylation Bioassay ...... 43

2.2.5 Chemical Analysis...... 44

2.2.6 Mercury Fractionation by Filtration ...... 45

2.2.7 TEM Analysis of HgS-Amended Cultures...... 45

2.2.8 Mercury Ligand Exchange Reactivity ...... 46

2.3 Results and Discussion...... 47

2.3.1 Methylation of Mercury Sulfides ...... 47

2.3.2 Mercury Fractionation in Methylating Cultures...... 53

2.3.3 Mercury Speciation by Competitive Ligand Exchange...... 58

2.3.4 Environmental Implications...... 59

Chapter 3. Net methylation of mercury in estuarine sediment slurry microcosms amended with dissolved, nanoparticulate, and microparticulate mercuric sulfides ...... 63

3.1 Introduction...... 63

3.2 Materials and Methods...... 65

3.2.1 Sediment and water collection...... 65

3.2.2 Pore water characterization...... 66

3.2.3 Sediment slurry preparation ...... 67

3.2.4 HgS particle preparation ...... 67

3.2.5 Mercury methylation experiments ...... 68

3.2.6 Chemical analysis...... 70

3.2.7 Equilibrium speciation calculations...... 72

3.3 Results and Discussion...... 72

3.3.1 Surface water and sediment characteristics ...... 72

3.3.2 Net MeHg production in sediment slurry microcosms ...... 74

3.3.3 Mercury speciation in sediment slurry microcosms ...... 81

3.3.4 Environmental implications...... 87

Chapter 4. Photolytic degradation of methylmercury enhanced by binding to natural organic ligands...... 91

4.1 Introduction...... 91

4.2 Materials and Methods...... 93

4.2.1 Materials ...... 93

4.2.2 Sample matrix...... 94

4.2.3 Photodegradation experiments...... 94

4.2.4 MeHg decomposition by superoxide ...... 96

4.2.5 Chemical analyses...... 97

4.2.6 Quantification of rate constants ...... 99

4.2.7 Dark controls and direct photolysis...... 102

4.3 Results and Discussion...... 103

4.3.1 MeHg photodegradation by NOM...... 103

4.3.2 Speciation-dependent MeHg photodegradation rate...... 108

4.3.3 Environmental Implications...... 113

Chapter 5. Conclusions...... 115

5.1 Summary ...... 115

5.2 Implications and future research...... 119

5.2.1 Towards an understanding of mercury methylation potential ...... 119

5.2.2 Other metal sulfide nanoparticles...... 124

5.2.3 Implications for MeHg photodegradation ...... 126

5.2.4 Photochemistry in the aquatic environment...... 128

5.2.5 Predictive model for methylmercury ...... 130

Appendix A Supporting information for Chapter 2 ...... 131

Appendix B Supporting information for Chapter 3...... 143

Appendix C Supporting information for Chapter 4...... 147

Reference ...... 157

Biography ...... 185

List of tables

Table 2.1: Average size and surface area of humic-HgS nanoparticles and HgS microparticles utilized in methylation bioassays...... 52

Table 3.1: Surface water and sediment characteristics ...... 75

Table 4.1: Degradation rate constants for methylmercury complexes...... 111

Table B 1: Stability constants utilized to calculate saturation indices for HgS (s) (metacinnabar), FeS (s) (mackinawite) and FeS 2(s) (pyrite)...... 143

Table C 1: Effects of D 2O, sodium azide, and β-carotene on the steady state 1O2 concentration during UV-A irradiation of simulated water containing Suwannee River humic acid...... 147

Table C 2: Stability constants utilized for methylmercury speciation calculations in MINEQL+ (I= 0 M, 25 °C)...... 148

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List of figures

Figure 1.1: Phylogeny of bacterial strains capable of methylating and demethylating mercury for those strains with available 16S rRNA gene sequences...... 11

Figure 1.2: The transformations of mercuric sulfides in sediments...... 20

Figure 1.3 : 1O2 distribution around a 1.3-nm raidus DOM macromolecule...... 31

Figure 1.4: MeHg speciation calculated for a range of chloride concentrations...... 33

Figure 2.1: Net MeHg production in SRB cultures exposed to different forms of mercuric sulfides...... 50

Figure 2.2: Percentages of mercury in solution after filtration of D. propionicus 1pr3 cultures and TEM images of the cultures...... 56

Figure 2.3: Filtration of mercury-amended bacteria-free media (for culturing D. propionicus 1pr3)...... 57

Figure 2.4: Hydrophilic mercury in HgS-amended media after competitive ligand exchange with GSH or DEDC and C 18 solid-phase extraction...... 60

Figure 3.1: Mercury fractionation using centrifugation and ultracentrifugation...... 77

Figure 3.2: Net MeHg production in sediment slurries 2.2-2.8 days after amendment of

50 nmol mercury (either dissolved Hg(NO3)2 + Na 2S, HgS nanoparticles or HgS microparticles)...... 78

Figure 3.3: Net MeHg production in the slurry microcosms exposed to 2 nmol g -1 (dw) dissolved Hg(NO 3)2 and Na 2S, HgS nanoparticles or HgS microparticles in Experiment 2...... 80

Figure 3.5: Calculated saturation indices for the water samples withdrawn from the sediment (BG30) slurries...... 88

Figure 3.6: Calculated saturation indices for the water samples withdrawn from the sediment (LSB129S) slurries...... 89

Figure 4.1: Concentration dependence of MeHg photodegradation rates...... 105

Figure 4.2: MeHg degradation via generation of 1O2 from photosensitized humic acid.107

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Figure 4.3: Effect of ligand complexation on singlet oxygen-induced of MeHg...... 110

Figure 4.4: Effect of ligand complexation on direct UV-C degradation of MeHg...... 112

Figure A 1: Net MeHg production in D. propionicus 1pr3 cultures exposed to different forms of mercuric sulfides...... 134

Figure A 2: Bacterial growth in SRB cultures exposed to different forms of mercuric sulfides...... 135

Figure A 3: Intensity-weighted size distribution of HgS particles measured by dynamic light scattering...... 136

Figure A 4: TEM and SAED patterns of HgS nanoparticles...... 137

Figure A 5: X-ray photoelectron spectroscopy (XPS) and X-ray diffraction (XRD) results of HgS particles...... 138

Figure A 6: MeHg degradation in D. propionicus 1pr3 cultures...... 139

Figure A 7: TEM images and EDX spectra of D. propionicus 1pr3 cultures...... 140

Figure A 8: Net MeHg production in D. propionicus 1pr3 cultures exposed to different forms of mercuric sulfides...... 141

Figure A 9: Centrifugation and ultracentrifugation of mercury-amended D. propionicus 1pr3 cultures...... 142

Figure B 1: Total organic carbon (TOC) concentrations in water samples withdrawn from -1 the sediment slurries exposed to 2 nmol g (dw) dissolved Hg(NO 3)2 and Na 2S, HgS nanoparticles or HgS microparticles...... 144

Figure B 2: Sulfate reduction in slurry microcosms exposed to 2 nmol g -1 (dw) dissolved

Hg(NO 3)2 and Na 2S, HgS nanoparticles or HgS microparticles in Experiment 2...... 145

Figure C 1: Suwannee River humic acid (SRHA)-induced photodegradation of MeHg...... 149

Figure C 2: Reaction between MeHg and superoxide (O2-)...... 150

Figure C 3: Recovery of reactants in dark controls corresponding to competition kinetics experiments for 1O2 degradation of CH 3Hg-thiol complexes...... 151

xiv

Figure C 4: MeHg degradation by 1O2 generated by UV-A irradiation of rose bengal or Eosin-Y...... 152

Figure C 5: Effect of ligand complexation on 1O2-induced degradation of MeHg in seawater...... 153

Figure C 6: Simultaneous degradation of methylmercury and thiols during reactions • with 1O2 and OH...... 154

Figure C 7: Recovery of reactants for dark experiments with hydrogen peroxide...... 155

Figure C 8: MeHg speciation calculated for a range of chloride concentrations...... 156

Acknowledgements

I am sincerely grateful for all the people who have helped make this long journey less lonely and more joyful. My first thanks are to my advisor Heileen Hsu-Kim for being an amazing mentor and inspiring role model. During the past five years, she has taught and helped me more than I could ever ask for. Also, I would like to acknowledge my committee members, Drs Marc Deshusses, Mark Wiesner, Claudia Gunsch and

Heather Stapleton for their guidance and encouragement. I wish to thank all my dear friends, both near and far, for the selfless care that they give and for the much laughter that we share. Finally, special thanks go to my parents and Yao for always putting my best interests before theirs. Without them, I would have stopped long ago.

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Chapter 1. Introduction

Part of this chapter was reproduced from Hsu-Kim H., Kucharzyk K., Zhang T.,

Deshusses M.A. Mechanisms of mercury bioavailability and microbial methylation in the aquatic environment: a critical review. Environmental Science & Technology , Submitted.

1.1 Motivation

Methylmercury is a potent neurotoxin that accumulates in aquatic food webs and poses a significant risk to human health(1). Individuals with developing neurological systems (i.e., fetuses, infants and young children) are particularly susceptible to the adverse (and in extreme cases, fatal) effects caused by methylmercury poisoning. In the

U.S., 10,000’s to 100,000’s of children are born each year with in utero methylmercury exposure exceeding health guidelines(2) as methylmercury can pass through the placenta from the mother to the fetus. The major route of human exposure to methylmercury is consumption of fish(3) contaminated by methylmercury that originates from both natural waters(4) and sediments(5). Despite the critical importance of controlling the methylmercury levels in these aquatic settings, the mechanisms by which methylmercury is generated and degraded in the aquatic environment are largely unknown. As a result, accurate prediction of methylmercury accumulation and efficient remediation strategy of methylmercury contamination have remained elusive.

The biogeochemical cycling of mercury that results in bioaccumulation of methylmercury is highly complex and dynamic(6). Methylmercury concentrations in aquatic systems vary widely and are seldom a simple function of the total mercury input(7). Instead, methylmercury accumulation reflects the net result of many sources and sinks in the ecosystem. In particular, microbial formation from inorganic divalent mercury in anaerobic sediments accounts for the majority of the methylmercury that is potentially taken up by aquatic biota(8), while photodegradation by natural sunlight is a predominant removal pathway of methylmercury in surface water(9, 10). The significance of these two processes in methylmercury cycling in the aquatic environment has been demonstrated in previous work, yet the exact mechanisms that govern these processes are not well understood.

Microbial formation of methylmercury is primarily driven by sulfate-reducing bacteria in anaerobic sediments(11). Methylmercury production rates are generally related to two main factors: the productivity of the methylating bacteria and the availability of inorganic mercury species for microbial uptake and methylation. The uptake of mercury into methylating organisms is thought to occur through passive diffusion of neutrally-charged dissolved mercury-sulfide complexes (i.e., Hg(HS) 20 and

HgS 0) (12). With this basis, chemical equilibrium models are utilized to predict the concentrations of these mercury species and correlate with methylmercury production(13-15). However, a growing body of evidence has suggested that the

underlying assumptions of this approach are questionable. For example, instead of passive diffusion, methylating bacteria may take up mercury through an active transport mechanism(16, 17). Moreover, recent advances in nanogeochemistry helped demonstrate that HgS 0 is more likely to be mercury sulfide nanoparticles rather than a mononuclear aqueous complex(18). These discrepancies point to the need for re- examining the current mercury methylation paradigm and this was the first goal of my dissertation. The potential role of mercury sulfide nanoparticles in microbial methylation is of particular research interest and may provide new perspectives for understanding the mechanisms of this process.

Like the formation process of methylmercury, there still exist many gaps in our knowledge regarding the mechanism and pathway of methylmercury photodegradation.

In natural waters, the degradation of methylmercury occurs mainly through sunlight induced pathways, as indicated by previous in situ incubation experiments(9). This process is known to occur through an indirect mechanism(19) that may involve a variety of reactive intermediates generated from sunlight-activated chromophores in surface waters. However, the nature and source of these reactive intermediates are unclear.

Initial field studies have reported first-order degradation rates of methylmercury that were of similar magnitude in different types of freshwater systems and appeared to be independent of other water chemistry characteristics (e.g., TOC, pH) (9, 10, 20, 21). Thus, sunlight intensity has been proposed to be the only important factor in methylmercury

photodegradation. This statement has been seriously challenged by two recent observations. First, sunlight degradation of methylmercury occurred to a much lesser extent in a coastal marine environment compared to freshwater lakes(22). Second, theoretical calculations indicated that methylmercury-thiol complexes (dominant methylmercury species in freshwater) tend to be more reactive than methylmercury- chloride complexes (dominant methylmercury species in seawater) toward decomposition by electrophiles (e.g., photo-generated reactive species)(23). The results from this theoretical study, although they need to be verified in laboratory experiments, provide a plausible explanation for the first observation. Thus, the second focus of this dissertation was to address these questions concerning the mechanism of methylmercury photodegradation.

This research effort aims at developing a better mechanistic understanding of how methylmercury accumulation occurs in the aquatic environment. Our overall hypothesis is that the various mercury species present in natural settings are not equally available for microbial mercury methylation and methylmercury photodegradation. To what extent these two processes occur is largely determined by the chemical speciation of mercury that is subject to the environmental conditions. In order to examine this hypothesis, two key research questions are addressed in this dissertation: (1) what mercury species are readily available for microbial formation and photodegradation of methylmercury; and (2) how the environmental conditions influence the occurrence of

these reactive mercury species. The results from this research will allow for accurate prediction of the “hot spots” of net methylmercury accumulation (high methylation potential and low demethylation potential) based on measurable environmental parameters, and ultimately the elimination of the hazard of methylmercury contamination to human beings.

1.2 The problem of methylmercury in the environment

1.2.1 Methylmercury and its toxicological effects

Mercury can exist in three oxidation states: elemental/metallic (Hg 0), mercurous

(Hg(I)) and mercuric (Hg(II)). Elemental mercury (Hg 0) strongly partitions to the air, and thereby is capable of traveling across long distances (24, 25), and causes widespread contamination of mercury, even in remote areas where no point sources are identified.

Under ambient conditions, mercurous mercury (Hg(I)) is highly unstable and rarely considered in environmental studies(26). Mercuric mercury (Hg(II) has been found in both inorganic and organic forms. Organic mercury compounds are defined as mercury bound to carbon based structures, and the most commonly encountered organ mercuric species in the aquatic environment is monomethylmercury (MeHg).

MeHg is a developmental neurotoxin(1, 3) that can transport across the blood- brain barrier and placenta in mammalian animals. The bioaccumulative properties of

MeHg in mammals is thought to be based on specific uptake of the MeHg-L-cysteine complex that is mistaken for methionine and transported across membranes by a neutral

amino acid membrane transport protein (27-30). Compared to adults, fetuses and infants are particularly susceptible to MeHg poisoning due to their developing neurological systems(1, 31). Elevated level of MeHg causes neurological disorders and brain dysfunction, possibly by generating reactive oxygen species (e.g., superoxide, hydrogen peroxide and nitric oxide) or compromising intracellular calcium homeostasis and glutamate homeostasis(32, 33). The symptoms of MeHg poisoning range from mild numbness to blindness, loss of balance, and in extreme cases, death(34). While the brain is the primary organ affected by MeHg poisoning, adverse cardiovascular effects, such as acute myocardial infarction, have also been linked to MeHg exposure(35, 36).

MeHg is better retained by higher-level organisms than other mercury species and is the predominant form of mercury that bioaccumulates in the aquatic food chain(37). Previous evidence suggests that nearly all of the mercury (>85%) in the muscle tissue of fish occurs as MeHg(38-40). Because of the large biomagnification factors, fish body burdens for MeHg can be as high as 10 6 times the MeHg concentration in the surrounding water(41, 42). Dietary intake of contaminated fish is the major exposure route for MeHg in humans(3). As of 2010, more than 4500 fish consumption advisories have been issued in the U.S. These advisories span across 50 states, one territory, and three tribes. 81% of all advisories are issued, at least in part, because of unsafe levels of mercury(43). With the application of mercury isotopes, recent studies have demonstrated that the MeHg in fish originates from both aquatic sediments(5) and

surrounding water(4). Thus, the processes that determine the MeHg concentrations in these two media are critical for reducing the risk to human health posed by MeHg.

1.2.2 Sources and transformations of methylmercury

Mercury is a global pollutant that is released into the atmosphere from both natural and anthropogenic sources. Natural sources include volcanic eruptions, forest fires, biomass burning, and low-temperature volatilization(44). Anthropogenic sources include fossil fuel combustion, mining, waste disposal, and chemical production(44).

The mercury emission from these sources is mostly gaseous Hg 0. The major pathway by which Hg 0 enters the aquatic environment is first oxidation to inorganic Hg(II), followed by transportation to water bodies via wet or dry deposition(6). In fact, in situ methylation of inorganic Hg(II) is considered to be the primary source of MeHg in most aquatic settings. The contribution of the external MeHg input (e.g., wastewater discharge, surface runoff, groundwater infiltration) to the bioaccumulation of MeHg in aquatic food webs is relatively minor(8-10, 45). In situ production of MeHg may be mediated by microorganism(11, 46) or sunlight(47, 48), and occur in sediments(8) or water columns(49). Mass-balances for MeHg in both coastal marine and freshwater settings have shown that the sediment, where there is little to no sunlight penetration, is the most significant reservoir for MeHg(8, 50). The production of MeHg is mainly driven by the anaerobic bacteria that thrive in surface sediments. Many bacteria responsible for mercury methylation degrade MeHg as well(51). In addition, a variety of other bacterial

strains that cohabitate with the mercury methylators can demethylate MeHg via cellular detoxification(52) or metabolic mechanism(53). Thus, MeHg concentration and net accumulation in natural sediments are governed by the balance of simultaneously occurred microbial production and degradation of MeHg.

MeHg that accumulates in the surface sediments establishes a concentration gradient of MeHg between sediment pore waters and overlaying waters, which drives the diffusive flux of MeHg from the sediment to upper water bodies(54). This process can be enhanced by advective mixing at the sediment-water interface(55, 56) or bioturbation by benthic organisms(57). Once MeHg is transported into the photic zone, sunlight induced degradation is believed to take place and remove up to 80% of the

MeHg flux (9, 10). Thus, photodegradation is the predominant pathway for the decomposition of MeHg and maintains the MeHg concentrations at low levels in surface waters.

1.3 Microbial formation of methylmercury in natural sediments

1.3.1 Mechanisms of microbial mercury methylation

1.3.1.1 Mercury methylating bacteria

Microbial Hg(II) methylation was first demonstrated in 1969(46) and it is now widely accepted that in aquatic sediments the primary Hg(II) methylators are sulfate- reducing bacteria (SRB), a group of obligate anaerobic bacteria that utilize sulfate as the

terminal electron acceptor for respiration(58). Although the possibility of Hg(II) methylation by other microbes (i.e., iron-reducing bacteria and methanogens) has been recently reported(59-61), the dominant role of SRB in Hg(II) methylation is supported by extensive evidence from pure culture and sediment microcosm experiments. Key results from these works include the following: (1) a variety of SRB cultures can effectively convert Hg(II) to MeHg(51, 62); (2) the addition of molybdate, a specific inhibitor of sulfate reduction suppressed Hg(II) methylation in sediment samples(58); (3) Hg(II) methylation rates were observed to correlate with sulfate-reduction rates(63) or with the distribution of SRB populations in sediment(64); and (4) the experimental addition of sulfate enhanced MeHg production in sediment cores(65).

To date, mercury methylation capacity has been tested in fewer than 50 bacterial strains (Figure1.1). However, it appears that the ability to methylate mercury is not shared by all SRB(66). The order Desulfovibriales has been the most extensively examined, with about half of the species having the ability to produce MeHg. Mercury-methylating

SRB are also found within the Desulfobacterales order as well as the Desulfuromonales order. The capacity for mercury methylation was found to depend on the strain rather than species, genus or metabolic group(62).

As shown in Figure 1.1, the vast majority of the Hg methylators also have the ability to demethylate MeHg. Biotic demethylation occurs through both reductive and oxidative pathways(66). Reductive demethylation is mediated by the mer operon system

and result in the formation of CH 4 and either inorganic Hg(II) or Hg 0. In oxidative demethylation, MeHg is degraded to inorganic Hg(II), CO 2, and small amounts of CH 4 as a cometabolic by-product of methylotrophic metabolism(66). Reductive demethylation is a well known mechanism for mercury resistance and detoxification.

This process occurs in both aerobic and anaerobic environments(67), and is favored at high concentrations of mercury(52). Oxidative demethylation, on the other hand, mostly occurs in anaerobic sediments, and has been detected in both pristine and mercury contaminated areas(53, 68, 69). Many mercury methylating bacteria including all obligatory anaerobic microorganisms, do not have the mer sequence in their genomes, suggesting that they most likely degrade MeHg through an oxidative pathway.

1.3.1.2 Mercury uptake prior to methylation

Microbial methylation of mercury has been broadly recognized as an intracellular reaction(70, 71), and thus transport of inorganic mercury from the microbe’s surroundings through the cellular membrane is the first step leading to mercury methylation. In general, microorganisms can take up metals either by a metabolism- independent, passive process or a metabolism-dependent, active process. Passive processes would probably dominate in low nutrient environments since active uptake requires energy, and microorganisms normally lack specific uptake systems for metals(72).

Desulfobacterium autotrophicum (ABR92474)*

Desulfomicrobium baculatum (AM419440) DESULFOVIBRIONACEAE Desulfovibrio termitidis (AF418142) Desulfolubus propionicus (BAK69215) Desulfobacterium dehalogenas (AAL07562)

Desulfomaculum nigrificans (DSM574) ▲ PEPTOCOCCACEAE Desulfomicrobium escambiense (BAB55555) DESULFOMICROBIACEAE Desulfomicrobium sp. (ADR28) Desulfovibrio africanus (BAB55563) DESULFOVIBRIONACEAE Desulfovibrio gigas (AAB41001) Salinibacter ruber (ABC45686)* RHODOTHERMACEAE

Geobacter sulfurreducens PCA (NP954464) GEOBACTERACEAE Thermus thermophilus HB8 (YP144419) Aeropyrum pernix K1 (NP157957) DESULFOVIBRIONACEAE Desulfovibrio vulgaris (AF418179) Thermoproteus tenax (CAF18529) THERMOPROTEACEAE Sulfolubus solfataricus (NP344015) SULFOLOBACEAE Haloarcula marismortui (YP134400) HALOBACTERIACEAE Thermosulfovibrio yellowstonii 0.2 (ACI22200)

Figure 1.1: Phylogeny of bacterial strains capable of methylating and demethylating mercury for those strains with available 16S rRNA gene sequences.

Strains denoted with symbols are microorganisms with only methylation (*) or only demethylation ( ▲) capabilities. The phylogenetic tree was generated using the neighbor-joining analysis. All accession numbers are indicated.

The speciation of mercury has long been considered as an important determinant of its biological uptake. In Escherichia coli (E. coli ) and marine diatom organisms cultured in a chloride gradient and exposed to Hg(II), a maximum reduction of cell viability

(presumably reflecting mercury uptake and toxicity) was achieved when HgCl 2 complexes were the dominant Hg(II) species in the culture solutions(73, 74). The large uptake of Hg(II) in conditions with high concentrations of HgCl 2, but not in conditions where HgCl 3- or HgCl 42- dominated Hg speciation, would suggest passive diffusive uptake of lipophilic, neutrally-charged forms of Hg(II)(73, 74). These findings have been 11

extrapolated to the process of mercury methylation, in which anaerobic bacteria were also presumed to take up mercury by a passive diffusion mechanism(13, 75, 76).

Mercury uptake by bacteria could also occur through a facilitated diffusion mechanism. Facilitated uptake of dissolved Hg(II) has been proposed by studies involving Gram-negative bacteria that are not known to methylate mercury(77, 78). The details of the facilitated uptake mechanism remain unclear, but in an engineered strain of Vibrio anguillarum pRB28, a small change in pH (from 7.3 to 6.3) resulted in a large increase in the uptake of inorganic Hg(II)(78). This effect of pH on uptake could not be explained by a model of passive diffusion of uncharged Hg(II) species, and likely occurred through a facilitated mechanism(78).

Researchers have also suggested that microbial uptake of mercury can involve an active transport mechanism. This conclusion is based on studies performed with both methylating and non-methylating microorganisms(16, 17, 79). Mercury uptake was reduced when the incubations were performed at decreased temperatures and in the presence of Na +/K + ATPase inhibitors, indicating a metabolic influence on the mercury uptake which cannot be explained simply by diffusion into the cells(17, 79). Moreover, recent studies by Schaefer et al .(16, 17) suggested that mercury uptake by two mercury methylating bacteria ( Geobacter sulfurreducens and Desulfovibrio desulfuricans ) was energy dependent, either through an electrogenic or ATP-driven mechanism. Complexation with certain types of thiols, such as cysteine, enhanced the uptake and subsequent

methylation of inorganic Hg(II). However, Schaefer et al.(16, 17) postulated that the involvement of a specific amino acid membrane transporter was unlikely, and that mercury uptake may involve a non-specific transporter for trace metal nutrients, resulting in the accidental uptake of Hg(II).

1.3.1.3 Biochemical mechanisms of mercury methylation

The major obstacles for advancing our knowledge of microbial mercury methylation include the absence of identified genetic systems responsible for mercury methylation and lack of clear correlations between taxonomy of methylating microbes and methylation rates. The biochemical reactions causing methylation of mercury are considered to be intracellular, followed by a rapid efflux or diffusion of MeHg outside the cell(80). The ability to produce MeHg is constitutive rather than induced by exposure to mercury(80). Despite many previous research efforts, there is still little understanding on the biochemical pathways of MeHg formation in SRB. Since MeHg production is primarily associated with the activity of sulfate-reducing organisms, it has been proposed that the organism's ability to methylate mercury is most likely associated with substrate specificity of its enzymes(71). For some microorganisms methylmercury production could be linked to a specific methyl-transferase pathway, to a Hg-specific uptake pathway, or to the biochemistry of mercury binding within the cell(80).

The main biochemical agents that can induce a spontaneous transfer of a methyl group include: S-adenosylmethionine, N 5-methyl-tetrahydrofolate derivatives, and 13

methylcobalamin(70). Among these, methylcobalamin has been identified as an agent that can facilitate the transfer of methyl groups (CH 3-) to mercuric ions (Hg 2+ ). The importance of methylcobalamin has been shown in studies conducted with the sulfate- reducing bacterium Desulfovibrio desulfuricans LS, a dissimilatory incomplete oxidizer of short-chain fatty acids(71, 81). For this organism, the proposed mechanism of carbon flow involves the transfer of a methyl group to methyl-tetrahydrofolate via methylcobalamin (MeB 12 ), with the methyl group originating from C-3 of Ser (by serine hydroxymethyl transferase) or from formate. Next, the enzymatic acetyl-coenzyme A

(CoA) pathway initiates a methyl group transfer from methyl-tetrahydrofolate to a corrinoid membrane protein and enzymatic transmethylation to Hg 2+ . The consensus is that mercury methylation is most likely an incidental rather than a primary function of methylcobalamin. Paradoxically, methylcobalamin was identified in some acetogens and methanogens(82), which were previously excluded from the groups of organisms considered to be mercury methylators.

Although mercury methylation occurred mainly through the acetyl-CoA pathway in Desulfovibrio desulfuricans LS(71, 81), experiments involving other incomplete oxidizing SRB strains have demonstrated that inhibitors of the acetyl–CoA pathway

(such as chloroform and cobalt limitation) had no effect on MeHg production(83, 84).

Furthermore, SRB strains that are complete oxidizers (i.e., microbes that consume acetate to generate CO 2) have shown a correlation between mercury methylation and the

abundance of CO dehydrogenase (CODH), a key enzyme in acetyl-CoA pathway(83).

These inconsistencies highlight the gaps in our knowledge of the biochemical process responsible for MeHg production, and underscore the point that the methylation of mercury within SRB may involve more than one mechanism.

1.3.2 Chemical speciation and reactivity of inorganic mercury in natural sediments

1.3.2.1 Chemical speciation of inorganic mercury in the aquatic environment

In the aquatic environment, inorganic Hg(II) generally persists in the form of mercury-ligand complexes (e.g., Hg 2+ complexes with chloride, inorganic sulfide, or dissolved organic matter), or Hg(II) is associated with particles (mercury-bearing minerals or Hg 2+ adsorbed to particle surfaces). The relative partitioning of inorganic

Hg(II) in various dissolved and particulate forms will govern the overall mobility of Hg and the availability of Hg to methylating microorganisms in anaerobic settings. In the past, researchers have attempted to differentiate the speciation of Hg(II) based on size separation (i.e. filtration with a particular pore size or molecular weight cutoff) or metal- ligand complexation (from experimentally determined thermodynamic binding strengths of ‘dissolved’ Hg complexes).

In a surface water or sediment porewater, the fraction of Hg(II) in the particulate phase has traditionally been operationally defined by the amount of mercury that is retained by a 0.2 or 0.45 μm filter. Smaller particles (i.e. < 0.2 or < 0.45 μm) suspended in

water (i.e. colloids) are capable of passing through these filters. Therefore, ultrafiltration devices have been applied to surface waters to quantify the amount of colloidal mercury relative to dissolved and particulate forms(85-87). Previous work has shown that in 0.4

μm-filtered water, the percentage of mercury that is captured by a 10-kDa ultrafiltration device (i.e. the ‘colloidal fraction’) varied from 20 to 80%, where the lower end of this range occurred in saline water and the higher end occurred in freshwater(85-87). This result is consistent with the coagulation of colloidal particles in saline water to form aggregates larger than 0.4 μm. In the porewater of anoxic sediment where methylation primarily occurs, the proportion of mercury in the colloidal fraction is not as well documented. However, the presence of colloidal Hg could be expected in light of evidence showing that nanoparticulate forms of HgS can persist as by-products of mineral precipitation occurring in the presence of dissolved organic matter (DOM)(18,

88-90).

The species of Hg(II) present in the aquatic environment can also be differentiated based on binding strength of dissolved Hg(II)-ligand complexes. Trace metal complexation has been studied extensively in the past using a wide variety of methods that include electrochemical, competitive ligand exchange, and chromatographic approaches(91). When all of these techniques were applied to streams, rivers, estuaries and municipal wastewater effluent(92-95), the results generally demonstrated that the stability constants for Hg-ligand complexes resembled those for

Hg-thiol complexes. Presumably, dissolved Hg(II) in these settings is specifically binding to thiol functional groups associated with dissolved organic matter (DOM), as demonstrated in spectroscopic studies(96-98). In some settings, particularly wastewater effluent(92), the experimentally-determined stability constants for Hg(II) species were greater than expected for Hg-thiol complexes, suggesting that ‘stronger’ Hg(II) complexes (such as Hg-sulfide clusters or nanoparticles) were present in these settings(99).

1.3.2.2 Current model for predicting methylation potential

In natural sediments, methylmercury levels rarely correlate to the amount of total mercury and only a small portion of the total inorganic Hg(II) is likely to be available to the mercury methylating bacteria. To this end, bioavailability models for mercury have been devised to link the geochemical speciation of inorganic Hg(II) to

MeHg production in sediments and other anaerobic settings. The most established approach for modeling mercury bioavailability assumes that uptake occurs through a passive diffusion mechanism(12). Thus, only neutrally-charged forms of dissolved Hg(II) can be taken up by methylating microorganisms, and the concentration of the

‘bioavailable’ forms of Hg(II) can be estimated from thermodynamic equilibrium models of Hg(II) complexes(12). This modeling approach has been utilized by several others in attempts to draw correlations between observed MeHg concentrations in environmental samples and the calculated concentrations of neutrally-charged forms of dissolved 17

Hg(II)(13-15, 100). However, careful examination of the underlying assumptions reveals that this model is inherently flawed.

First of all, the outcome of this model heavily relies on the accuracy of the stability constants for the formation of Hg(II)-ligand complexes. However, much uncertainty arise from the HgS 0(aq) (or HgOHSH (aq) ) complex, a species that has been proposed to be the available form of mercury for microbial methylation(75). The intrinsic solubility of HgS 0(aq) ( Ksp1 = 10 -10 for the reaction: HgS (s) + H 2O ⇔ HgS 0(aq) ) was not experimentally determined, but was instead extrapolated from the formation constants of other metal complexes (ZnS and CdS)(101). The original study that proposed the ZnS 0 and CdS 0 species acknowledged that these compounds were likely to consist of colloidal forms of metal sulfides rather than mononuclear aqueous complexes(101). A more recent study(102) postulated that the formation of HgS 0(aq) from HgS (s) should be represented by the smaller value estimated by Dryssen and Wedborg ( Ksp1 = 10 -22.3 ). Moreover, aqueous

HgS 0 remains as an unknown chemical species as it has never been directly measured in any study, rendering it difficult to prove the viability of a formation reaction and equilibrium constant.

Another source of uncertainty for equilibrium calculations is the solubility products Ks0 for minerals such as metacinnabar and cinnabar, which vary by orders of magnitude in a database for critically selected stability constants (log Ks0 = -38 ± 2)(103).

The values of these constants determine the saturation index of Hg-sulfide, a parameter

for predicting the potential of particle formation. The development of the neutral mercury-sulfide bioavailability model required the selection of relatively large solubility constants for HgS (s) ( Ks0 = 10 -36.5 ) so that the transition of the predominant dissolved Hg- sulfide species (e.g. HgS 0(aq) versus HgS(HS) -) could be matched to field data showing a decrease of MeHg with an increase of sulfide concentration(75). However, this selection leads to the prediction that in sediment pore waters, dissolved Hg(II) is understaturated

(with respect to metacinnabar HgS (s) ) and thus Hg-sulfide particles is not thermodynamically favored in most sediment pore waters(12), a notion that conflicts with direct observations of HgS (s) in sediments(104).

The second questionable assumption of this model is that the mercury speciation in pore waters can be represented by chemical equilibrium. Previous measurements of mercury speciation (whether they involve metal-ligand stability or characterizations of particulate vs. dissolved) rarely fully considered the heterogeneity of mercury species in natural aquatic systems, and that there is no clearly defined cutoff that distinguishes a dissolved molecule from a colloidal particle(105). Rather, the constituents that comprise a surface water or sediment porewater include a continuum of species: from dissolved molecules to polynuclear clusters, amorphous nanoparticles, and larger (perhaps crystalline) particles (Figure 1.2). This mixture of compounds would not be predicted from equilibrium mineral solubility and likely represent intermediates of metal-ligand complexation, mineral dissolution, and precipitation processes at non-equilibrium.

Several studies have pointed to the importance of rate-limited processes (e.g. HgS (s) dissolution, precipitation, mass transfer across depth) for influencing Hg geochemistry in sulfidic settings(18, 55, 88, 105, 106). Taken together, the current model that utilizes equilibrium speciation to predict the reactivity of inorganic Hg(II) in microbial methylation needs to be re-examined.

Stable NOM-capped Heterogeneous polynuclear clusters amorphous nanoparticles

bioavailable to

dissociation methylating bacteria

ripening or aging sorption cluster formation in presence of NOM of NOM dissolved Hg(II) complexes

not bioavailable dissolution

Stable colloid and Aggregated colloids nanoparticle suspension and nanoparticles

Figure 1.2: The transformations of mercuric sulfides in sediments. These processes involve a diverse collection of species, many of which are intermediates of metal-ligand complexation reactions and precipitation and dissolution of HgS (s) . Modified from Aiken et al. (105).

1.3.2.3 Hg-sulfide-DOM speciation

Although the current model completely neglects the role of DOM in microbial mercury methylation, field and experimental data have demonstrated correlations between organic carbon concentration and MeHg production(107-109). DOM could contribute to Hg bioavailability and methylation potential in two ways. First,

complexation of Hg(II) by DOM is thought to decrease the amount of Hg(II) available to the methylating bacteria due to the difficulty for the large macromolecular and hydrophilic Hg-DOM complexes to diffuse through the cell membranes(110). On the other hand, in most settings MeHg concentration was observed to increase with organic carbon content in sediments(107, 111). This positive correlation is typically attributed to a stimulating effect of the labile carbon on microbial growth. However, neither of these two theories fully captures the inter-related roles of DOM and sulfide for inorganic

Hg(II) speciation and methylation. Exceptions to the correlation between MeHg and

DOM have been reported(112) in which the co-existence of sulfide and DOM appeared to yield a favorable geochemical environment for microbial Hg(II) uptake(112).

Therefore, the explicit mechanism through which DOM contributes to mercury methylation needs to be investigated in conjunction with other environmental variables, especially sulfide.

Thermodynamic equilibrium calculations generally show that Hg 2+ preferentially coordinates to inorganic sulfides over organic thiols associated with DOM due to the higher abundance and/or affinity of sulfides. However, DOM can influence Hg(II) speciation in other ways, particularly if the transformations involving Hg are at a non- equilibrium status. DOM is known to enhance the dissolution rate of cinnabar and inhibit the precipitation rate of metacinnabar(88, 113). The occurrence of an aqueous ternary DOM-Hg-sulfide complex is a possible explanation(114). However, more recent

studies have demonstrated that DOM plays a significant role in slowing the growth and aggregation of HgS nanoparticles as they precipitate in aqueous suspension(115, 116).

These nanoparticles are likely to consist of a metacinnabar-like material (in terms of Hg-

S coordination structure) that result in amorphous or nanocrystalline Hg-S-DOM nanoparticles(89, 90).

Discrete nanoparticles of HgS have been detected directly in soil and sediment(117-119). However, these examples were highly contaminated settings (mining and industrial sites where mercury-enriched materials were actively processed).

Methods to directly detect nanoparticles (e.g. electron microscopy) generally require high concentrations of the target element in the sample (e.g. greater than parts-per- million). In most settings where mercury concentrations in sediments are much more dilute, nanoscale mercuric sulfides will likely comprise of a mixture of metal sulfides, such as Hg sorbed to mackinawite-like (FeS) nanoparticles(120).

1.3.2.4 Reactivity of nanoparticles

Recent research on mercury-sulfide-organic matter chemistry has indicated that nanoparticles of Hg-S-DOM need to be considered in speciation models(18, 89, 90). The role of these particles for bioavailability and methylation potential remains unknown

(and is a goal of this dissertation). Nanoscale particles are expected to behave differently than the compositionally identical, larger materials due to the high specific surface areas and unique reactivity of materials at the nanoscale(121, 122). Indeed, the defining 22

characteristics of nanoparticles are not only the small size (i.e., at least one dimension smaller than 100 nm) but also size-specific reactivity exhibited by the nanomaterials(123).

Nano-specific reactivity is generally observed in particles smaller than 30 nm and stems from the relatively large specific surface area and crystal lattice imperfections in a material with a large proportion of atoms on the surface. Nano-specific reactivity may include increased sorption capacity (normalized to surface area), enhanced transport, and faster rates of dissolution(121, 123).

The reactivity of nanoparticles can lend them to unique pathways for uptake into organisms, and at the very least, will influence the microbial bioavailability of the metal constituents of the nanoparticle. Clues toward understanding the importance of nano-

HgS for microbial uptake and methylation could be gained from more established research on biouptake of iron originating from nanoscale iron oxides. For example, in microbial iron reduction, nanosized iron oxide colloids exhibited up to 100 times greater iron transformation rates than their respective bulk minerals(124). This observation was attributed to the enhanced solubility(125) and larger mineral particle-bacteria contact for nanoparticulate Fe(III)(126, 127). Moreover, microscopic analysis revealed that iron oxide nanoparticles can penetrate the outer membrane of iron reducing bacteria,

Shewanella putrefaciens , without collapsing the cells(126), and this bacterium tended to dissolve Fe(III) at the bacteria-mineral interface(128).

In previous studies, the reactivity of nanoparticles has almost always been examined in relatively simple and non-environmentally relevant matrices (e.g., experimental solutions, pure bacterial cultures)(124, 129). The information that we gain from this is valuable but not sufficient for constructing a predictive model for real environmental settings. In natural aquatic systems, nanoparticles commonly exist as aggregates(121). Environmental factors, including pH, salinity, sulfide content and DOM content are expected to influence the structure of the nanoparticle aggregates(18, 130,

131); this in turn has been shown to affect the reactivity of nanoparticles in biogeochemical processes (e.g., dissolution and bio-reduction)(129, 132-134). Therefore, in order to understand the potential role of nanoparticulate Hg-S-DOM species in microbial methylation, experimental efforts under environmentally relevant conditions are warranted.

1.4 Photodegradation of methylmercury in natural waters

1.4.1 Mechanisms of sunlight induced degradation

Photodegradation refers to the degradation processes that are triggered by absorption of photons. In natural waters, photodegradation is driven by sunlight irradiation. Sunlight induced degradation of various organic compounds, including pharmaceuticals(135-137), pesticides(138, 139), and phenol(140) has been extensively documented. In general, photodegradation occurs via two types of pathways: direct photolysis and indirect photolysis.

1.4.1.1 Direct photolysis

Direct photolysis is a process in which a substrate of interest is excited to a higher energy state by the absorption of photons and consequently decomposes to other substances. In this process, not all excited molecules undergo chemical reactions that result in new products. Instead, a portion of the excited molecules may drop to their ground state while releasing energy in the form of heat or light, or transfer their energy to other compounds that are in the same matrix. The efficiency of direct photolysis in contaminant degradation is quantified by quantum yield, the ratio of the amount of molecules degraded during photolysis and the amount of photons absorbed by the reaction matrix.

There are two prerequisites for direct photolysis to occur and become an important removal mechanism of contaminants in natural waters. First, the contaminants need to reside in a location that receives adequate amounts of sunlight irradiation(141). At each depth in the water column, the sunlight spectrum and intensity is dependent on the turbidity of water and the presence of other light absorbing compounds or organisms. Generally, radiation with shorter wavelength is attenuated more rapidly in the water column relative to longer-wavelength radiation. As a result, direct photolysis driven by the high energy component of sunlight (i.e., ultraviolet (UV)

A and B) may be negligible at water depth greater than several centimeters(142, 143).

Second, the absorption spectra of the contaminants need to overlap with the spectra of

the available sunlight irradiation(144). Solar radiation below 290 nm is significantly blocked by the stratospheric ozone layer, and thus sunlight reaching the Earth’s surface contains mainly visible and infrared light, with a small amount of UV light.

Contaminant compounds that strongly absorb radiation at wavelengths corresponding to UV-B (290-320 nm), UV-A (320-400 nm) and visible light (400-750 nm) are known to readily decompose via direct photolysis in natural waters(137, 140). However, many compounds lack chemical structures that enable them to absorb sunlight, or the energy of the available sunlight is too limited to induce degradation processes. MeHg follows this second example. Quantum mechanical calculations have shown that breaking the C-

Hg bond in MeHg molecules requires energy higher than 4.4 eV(19) which corresponds to UV light with a wavelength shorter than 280 nm. Therefore, the sunlight degradation of MeHg likely occurs through indirect pathways.

1.4.1.2 Indirect photolysis

During indirect photodegradation of a contaminant, a different light absorbing compound, often referred as photosensitizer, transfers the energy from a photochemically excited state to the substrate of interest by directly reacting with the substrate. Alternatively, the excited photosensitizer can generate a suite of reactive intermediates that subsequently decompose the substrate. These reactive intermediates are transient species in natural waters and their production follows two types of mechanisms. The type I mechanism involves electron transfer from the excited 26

photosensitizer to the dissolved oxygen molecules and results in the production of

• • • hydroxyl radical ( OH), superoxide radical anion (HO 2 /O 2 -), hydrogen peroxide (H 2O2),

• peroxyl radical (ROO ), etc . The type II mechanism involves energy transfer from the excited photosensitizer to the dissolved oxygen molecules in a collision-dependent process and typically results in the production of singlet oxygen ( 1O2).

The reactive intermediates, including the excited photosensitizer molecules, are widely prevalent in surface waters(145, 146) and their concentrations reflect the balance between rapid generation and consumption processes. Among all the reactive intermediates, hydroxyl radical and singlet oxygen have received most of the attention, due to their important role in photodegradation of environmentally relevant compounds.

In particular, previous photodemethylation studies have demonstrated that indirect photolysis of MeHg can occur through both hydroxyl radical and singlet oxygen pathways(147, 148).

Hydroxyl radical is a very strong and non-selective oxidant that reacts with most contaminants at near diffusion-controlled rates (~10 9-10 10 M -1s-1)(146). The primary quenchers of hydroxyl radical include DOM, carbonate and bicarbonate(149, 150), which are ubiquitous in natural waters. As a result, the abundance of hydroxyl radical is extremely low (10 -16 ~10 -18 M) in pristine waters but can be elevated by orders of magnitude due to human activities(146, 151, 152). For instance, the photolysis of nitrate that is released from agriculture application is a major source of hydroxyl radical in

freshwater systems(152). In another example, in surface waters impacted by acidic mine drainage, Fenton reactions likely dominate the production of hydroxyl radicals(153).

Singlet oxygen is the first excited state of molecular oxygen and preferentially reacts with electron-rich compounds(154). In the aquatic environment, singlet oxygen is mainly consumed by water molecules through physical quenching(155). Compared with hydroxyl radical, singlet oxygen has lower oxidative reactivity but higher selectivity and longer life time in water. The steady-state concentrations of singlet oxygen in natural waters are in the range of 10 -12 ~10 -14 M, significantly larger than concentrations of hydroxyl radical(151, 156).

The kinetics of indirect photodegradation in natural waters is generally determined by three factors: 1) sunlight intensity and penetration; 2) the reactivity between the substrate of interest and reactive intermediates; 3) the chemical characteristics of natural waters, such as pH, particle content, dissolved oxygen, and the type and concentration of photosensitizers. To understand the mechanisms of indirect photodegradation, numerous studies have performed laboratory or in situ experiments to answer the following questions: 1) What are the reactive intermediates responsible for the degradation? 2) What is the source of these reactive intermediates? 3) What is the reactivity between the reactive intermediates and the target contaminants? 4) How do environmental parameters (i.e., solar irradiation, water chemistry characteristics) affect the degradation kinetics? In these studies, Electron paramagnetic resonance (EPR)

spectroscopy with the application of spin-trapping agents to stabilize the unpaired electrons in the reactive intermediates is one of the most commonly employed methods for identifying and quantifying the reactive intermediates(139, 157, 158). Researchers have also utilized chemical probes and quenchers that preferentially react with the reactive intermediates at a known reaction rate(151, 159-161).

1.4.2 Photochemistry of dissolved organic matter

In natural waters, DOM is the most abundant light absorbing compound and present at a wide range of concentration levels (0.2-200 mg-C/L)(162). DOM consists of the decomposed plant and animal material and is a mixture of organic molecules with ill-defined structures. The main fraction of aquatic DOM is humic substances, which comprise two types of macromolecular hydrophobic acids, humic acid and fulvic acid.

The highly heterogeneous nature of DOM dictates the complexity of its role in aquatic photodegradation processes.

The chromophoric moieties in DOM readily absorb sunlight and induce indirect photolysis by generating a variety of reactive intermediates(163). First, upon absorbing sunlight, the DOM molecule is promoted to its excited triplet state ( 3DOM *) and this excited species can directly react with contaminants and lead to degradation(164). A number of studies have demonstrated the significant role of DOM in the photodegradation of pharmaceutical and pesticide compounds(165-167). This process may involve energy, electron or hydrogen atom transfer between 3DOM * and the

contaminant molecules. The degradation kinetics appeared to be correlated with the extent of adsorption of the organic contaminant onto DOM macromolecules(165-167).

Second, 3DOM * can react with dissolved oxygen and generate singlet oxygen via the type

II mechanism. DOM is known to be the primary photosensitizer responsible for singlet oxygen formation in natural waters(168, 169). Singlet oxygen is thought to exhibit a heterogeneous spatial distribution around photosensitized humic acid molecules in aqueous solutions. The 1O2 concentration appears to be several orders of magnitude higher at the origin of 1O2 generation (i.e., humic acid molecules) than in the bulk solution (Figure 1.3)(169, 170). These findings suggest that the compounds in close proximity to DOM (e.g., hydrophobic organics, metals complexed by DOM) are expected to be exposed to elevated levels of 1O2, relative to bulk solution. Moreover,

• DOM has been reported to be a source of other reactive intermediates, such as OH(171-

• 174), O2 -(175) and H 2O2(175-177). In particular, H 2O2 that is generated from photolysis of

• DOM can react with dissolved ferrous iron (Fe(II)) to produce OH. This reaction is commonly referred as photo-Fenton reaction and is considered as a potentially

• important source of OH in sunlit waters(171-173).

Figure 1.3: 1O2 distribution around a 1.3-nm raidus DOM macromolecule. in H 2O (solid line), D 2O ( 1O2 enhancer; dashed line) and H 2O with 1 mM sodium azide

(1O2 quencher; dotted line)(169).

The inhibitory effects of DOM have also been observed in photodegradation studies. Three mechanisms have been proposed to account for this phenomenon: 1)

• DOM is a scavenger of photogenerated reactive intermediates, such as OH(139, 149,

150); 2) DOM competes with other (possibly more efficient) photosensitizers for sunlight irradiation and the presence of DOM retards direct photolysis by screening the reactive components of sunlight(178-180); and 3) In some cases, DOM decelerates photodegradation through quenching of the transformation intermediates of the contaminants(181, 182).

To further complicate matters, DOM compounds that are from different sources have exhibited distinct behavior in photodegradation processes(183, 184). Studies have shown that fulvic acids were more efficient in generating 3DOM *(165) and 1O2(185) than

humic acids, and autochthonous fulvic acids were more photochemically reactive compared to allochthonous fulvic acids(184). Recent research has investigated the properties of DOM responsible for such discrepancies. The ability of DOM to generate

1O2 appeared to correlate with its molecular weight(157), fluorescent emission(185) and

UV absorbance properties(186). The ratio of absorbance at 254 nm and 365 nm may be a good indicator of 1O2 and H2O2 production from photolysis of DOM(186).

1.4.3 Chemical speciation and reactivity of methylmecury in natural waters

Apart from the complex role of DOM in photochemistry as discussed in section

1.4.2, a unique role which DOM may play in MeHg photodegradation is that DOM can strongly chelate MeHg. As a result, in some settings, DOM significantly influences the speciation of MeHg through the formation of MeHg-DOM complexes. In natural waters, the MeHg cation (CH 3Hg +) is predominantly complexed by dissolved ligands such as chloride (Cl -), hydroxide (OH -), and DOM. The thermodynamic stability of CH 3Hg + complexes with thiol functional groups in DOM are orders of magnitude greater than

CH 3Hg + complexes with other ligands, as indicated by the larger stability constants for these complexes. Thus, CH3Hg + preferentially associates with DOM in natural waters(187, 188). In addition to the affinity of CH3Hg + for the ligands, MeHg speciation is also determined by the concentration of each ligand. Figure 1.4 shows the computed concentrations of dissolved CH 3Hg-ligand complexes as a function of chloride concentration under environmentally relevant conditions. CH 3HgCl complexes are 32

expected to be the main MeHg species in coastal marine waters with chloride concentration greater than 0.1 M. In contrast, most freshwater systems contain enough

DOM (>1 mg-C/L) and low chloride such that CH 3Hg + are bound to thiol ligands associated with DOM(189, 190).

1.0E-12 13.5 KCH3HgNOM = 10

8.0E-13 CH 3HgCl CH 3Hg-NOM 6.0E-13

4.0E-13

concentration,M 2.0E-13 CH 3HgOH 0.0E+00 -3.0 -2.5 -2.0 -1.5 -1.0 -0.5 - log [Cl ], M Figure 1.4: MeHg speciation calculated for a range of chloride concentrations. Speciation for water containing 10 -12 M total MeHg, 0.001 M to 0.7 M Cl -, 10 -10 mol DOM binding site per L (corresponding to 1 mg/L DOM), and pH 7. Seawater contains ~0.5 M Cl - (191).

Recent field study in a marine environment revealed that MeHg photodegradation is relatively slow in sea water compared to microbial processes(22), while other studies demonstrated consistently higher photodegradation rates of MeHg across freshwater systems with different water chemistry characteristics (e.g., DOC, pH)(9, 10, 20, 21). In theoretical computation studies, Ni et al. concluded that coordination of CH3Hg + ions to thiol-containing ligands increased the electron density at the Hg-C bond of MeHg. The shift of negative charge towards the methyl group increased its vulnerability to attack by an electrophile (e.g., OH and 1O2) and

consequently increased the reactivity of MeHg towards degradation (i.e., breaking the

C-Hg bond)(23). Although this theory has not yet been proven experimentally, it provides a plausible explanation for the difference observed in the kinetics of MeHg photodegradation between freshwater and marine systems. Here, this dissertation will address this potential mechanism of degradation via DOM activation by sunlight and specificity for MeHg-thiol complexes.

1.5 Research objectives

The overriding goal of this research was to understand how MeHg is generated and degraded in the aquatic environment. The main focus of this dissertation is to establish a relationship between the geochemical speciation of mercury and its availability in microbial methylation and sunlight degradation. The specific objectives include: 1) Demonstrate that in natural sediments, the methylation potential of inorganic

Hg(II) is determined by the kinetically-hindered mercury sulfide precipitation and mineralization processes, rather than equilibrium pore water chemistry; and 2)

Investigate the degradation pathway of MeHg in sunlit natural waters and the influence of MeHg speciation on the degradation kinetics.

The first objective is addressed in Chapter 2 and 3. To achieve this objective, three forms of mercury sulfides, including dissolved Hg(NO 3)2 and Na 2S, humic-coated

HgS nanoparticles and microscale crystalline HgS particles, were selected to represent different aging states of mercury in natural sediments. The methylation potential of 34

these mercury sulfide species were examined using pure cultures of SRB (Chapter 2) and sediment slurry microcosms (Chapter 3).

In Chapter 2, the methylation of mercury derived from nano-HgS aged for different time periods (16 h to 1 week) was also assessed. To explore the mechanism of microbial methylation of nano-HgS, the structure of HgS nanoparticles and bulk crystalline particles were characterized using a suite of techniques, including dynamic light scattering (DLS), BET N 2 adsorption, transmission electron microscopy (TEM), energy dispersive X-ray (EDX) spectroscopy, selected area electron diffraction (SAED), synchrotron X-ray diffraction (XRD) and X-ray photoelectron spectrometry (XPS).

Mercury fractionation method (filtration and centrifugation) coupled with TEM-EDX were applied to investigate the mechanism of mercury delivery to the methylating bacteria. The speciation of mercury in the culture media during methylation was analyzed using a previously developed competitive ligand exchange-solid phase extraction method.

In Chapter 3, natural sediments and waters from both freshwater and saline settings were utilized in sediment slurry microcosm experiments to investigate the

‘aging’ effect on the bioavailability of mercury under more environmentally relevant conditions. In these experiments, environmental variables (e.g., TOC, sulfate, salinity, sulfide, iron, pH) were monitored concurrently with net MeHg production. The speciation of mercury in sediment slurries was assessed using a centrifugation and

ultracentrifugation method. These results were analyzed to understand how environmental conditions influence the speciation of mercury and the consequence of these changes for MeHg production. In addition, saturation indices for sulfide mineral phases, including metacinnabar (β-HgS (s) ), mackinawite (FeS (s) ) and pyrite (FeS 2(s) ), in filtered (<0.2 μm) pore waters were calculated to evaluate the potential occurrence of nanoparticles during mercury methylation.

The second objective regarding the mechanism of MeHg photodegradation is addressed in Chapter 4. Both natural water and simulated water, natural sunlight and

UV-A radiation were utilized in this chapter. First, photodegradation experiments were conducted with MeHg and DOM at various concentration ratios to demonstrate the ability of DOM to induce and enhance MeHg photodegradation. Furthermore, selective quenchers were added to those reaction matrices to identify the photo-generated reactive intermediates responsible for MeHg degradation. In particular, singlet oxygen quenchers with different hydrophobicities were employed to assess how the heterogeneous distribution of singlet oxygen around photo-sensitized DOM influenced

MeHg decomposition. Finally, the second order rate constants for the degradation (via photo-generated reactive intermediates) of environmentally relevant MeHg-ligand complexes were quantified using competition kinetics. Glutathione (GSH) and mercaptoacetate (MA) were selected as surrogates of thiol-containing ligands that were expected to bind to MeHg ions in DOM-containing fresh water. Cl - and OH - ligands

were relevant for saline or low-DOM water. The effect of thiol-binding on the reactivity of MeHg molecules was also investigated in the direct photolysis experiments with UV-

C radiation.

The final chapter (Chapter 5) summarizes the key findings from this dissertation work and emphasizes the importance of mercury speciation for understanding the mechanisms of microbial formation and photodegradation of MeHg. The implications of this research to mercury biogeochemistry, nanogeoscience and DOM photochemistry are also discussed in this chapter.

Chapter 2. Methylation of mercury by bacteria exposed to dissolved, nanoparticulate, and microparticulate mercuric sulfides

This chapter was published in Zhang T., Kim B., Levard C., Reinsch B.C., Lowry

G.V., Deshusses M.A., Hsu-Kim H. (In press). Environmental Science & Technology .

American Chemical Society.

2.1 Introduction

Methylmercury (MeHg) production in the aquatic environment is primarily mediated by anaerobic bacteria, particularly sulfate reducing bacteria (SRB) (58, 62).

Microbial mercury methylation has been extensively studied for decades (46).

Nevertheless, our knowledge of the mechanisms of this process is rather limited. The speciation (and subsequent bioavailability to methylating bacteria (192)) of inorganic divalent mercury (Hg(II)) is controlled by inorganic sulfide (S(-II)) and dissolved natural organic matter (NOM) that can strongly bind Hg 2+ (12, 193, 194).

In the past, mercury bioavailability and methylation potential has been predicted based on chemical equilibrium models of dissolved mercury complexes in porewater

(nominally defined by 0.2 or 0.45-µm filters) (12). This approach presumes that mercury uptake is controlled by hydrophobic partitioning of dissolved neutrally-charged Hg- sulfide species that passively diffuse through the cell membranes of methylating bacteria.

Thus, the methylation potential from this approach is estimated from the concentrations of neutrally-charged Hg sulfide species calculated at chemical equilibrium. Our previous research has shown that this model is flawed because it must invoke an unknown species, HgS 0(aq) that may represent nanoparticulate HgS rather than a mononuclear aqueous complex (18). Here, we hypothesize that mercury speciation in anaerobic settings represent a mixture of dissolved, nanoparticulate, and bulk scale forms of HgS whose concentrations and bioavailability cannot be represented by conventional equilibrium models.

Mercuric sulfide (HgS) nanoparticles are known to exist in nature (117-119) and can account for a portion of mercury passing through 0.2 or 0.45-µm filters (195-197), an operational cut-off often used to separate ‘dissolved’ from particulate. In filtered porewater, mercury is often supersaturated with respect to HgS (s) (18). As with other minerals, the precipitation of HgS (s) involves nanoparticles that can be prevented from growing or aggregating by NOM (18, 198). Nanoparticles of HgS and other metal sulfides have been observed in anaerobic sediments and other settings where active precipitation was occurring (117-119, 199, 200), yet the role of nanoscale HgS for biomethylation has not been elucidated. As a result, accurate prediction of MeHg production and accumulation in the environment remains elusive.

Nanoparticles are not simply smaller versions of larger particles. Indeed, nanoparticles exhibit unique reactivity due to the high surface area-to-mass ratio of

nanoscale materials and resulting alterations in lattice structure and surface chemistry

(121). Conventional models of metal bioavailability use a chemical equilibrium speciation approach that presumes all particles (nano or otherwise) to be unavailable to microbes, yet this approach neglects properties such as enhanced solubility and greater deposition of nanoscale materials directly onto cell surfaces (126, 201, 202). Thus, nanoscale-specific reactivity of mercury may be contributing to its bioavailability and methylation potential in the environment.

In this study, we examined the methylation potential of various forms of mercuric sulfides by exposing fermentatively cultured SRB strains, Desulfobulbus propionicus 1pr3 and Desulfovibrio desulfuricans ND132, to three forms of mercury: dissolved Hg(NO 3)2 freshly mixed with Na 2S (dissolved Hg+S exposure), humic- stabilized HgS nanoparticles, and commercially-purchased HgS microparticles. These forms of mercury represented three different aging states of mercury in sulfidic sediments. We also characterized the structure of the HgS nanoparticles and assessed the speciation of mercury in the culture media during the incubation experiments.

2.2 Materials and Methods

2.2.1 Microorganisms and Culture Conditions

Desulfobulbus propionicus 1pr3 (ATCC 33891) and Desulfovibrio desulfuricans

ND132 (C. Gilmour, Smithsonian Environmental Research Center) were utilized as the test microorganisms. These strains were cultured in Hungate tubes (Bellco Glass) placed

in an anaerobic chamber. Cell growth was monitored by optical density (OD 660 ) and protein content (203). The bacterial cultures were maintained between experiments on sulfate-containing medium. Prior to mercury methylation bioassays, the cultures were transferred three times in fermentative media that contained 20 mM pyruvate (for 1pr3) or 40 mM fumarate (for ND132) as the organic carbon source, 0.15 mM Ti-nitrilotriacetic acid (NTA) as the reductant and 10 mg/L resazurin as the redox indicator, according to previous methods (193, 204).

2.2.2 HgS Particle Preparation

The Hg stock solution consisted of Hg(NO 3)2 dissolved in 0.1 N HNO 3. Na 2S stocks were prepared by dissolving freshly washed and dried crystals of Na 2S·9H 2O

(Fisher Scientific) in N 2-purged water and were utilized within 4 h of preparation. HgS nanoparticles were synthesized by dissolving 50 μM Hg(NO 3)2 and 50 μM Na 2S with 10 mg-C L -1 Suwannee River humic acid (SRHA, International Humic Substances Society) in a solution of 0.1 M NaNO 3 and 4 mM sodium 4-(2-hydroxyethyl) piperazine-1- ethanesulfonate (HEPES) (pH 7.5, double-filtered to <0.1 μm). Depending on the experiments, the Hg-S-NOM nanoparticle stock solution was allowed to age for 16 hours to 1 week at room temperature prior to use in the methylation experiments. A microparticulate HgS stock suspension was prepared by adding a commercial metacinnabar powder (β-HgS, Alfa Aesar) into nanopure-filtered water (>18 MΩ-cm).

This suspension was mixed end-over-end prior to taking an aliquot for the experiments.

2.2.3 HgS Particle Characterization

The average hydrodynamic diameter of HgS nanoparticles and microparticles were analyzed by light-intensity weighted dynamic light scattering (DLS) (Malvern

Zetasizer NS). The diameters of the monomers within the aggregates were analyzed by transmission electron microscopy (TEM). BET surface areas of HgS nanoparticles and microparticles were determined using the BET N 2 adsorption technique (Beckman

Coulter SA3100 Surface Area Analyzer). The geometric surface areas of HgS particles were calculated from the individual particle size obtained from TEM images. The geometric surface area calculations assumed spherical particles with a density of 7.71 g cm -3 (205).

The crystallographic structure of HgS nanoparticles and microparticles was analyzed by synchrotron X-ray diffraction (XRD) performed at the Stanford Synchrotron

Radiation Laboratory (SSRL) BL 11-3. The average crystallite diameter D of nano-HgS was estimated from the broadening of the X-ray diffraction peaks by the Scherrer formula (206):

D = K λ (2.1) βcos θ

where K is the constant of proportionality ( K=0.9), λ is the x-ray wavelength

( λ =0.0977 nm), β is the full width at half the maximum intensity in radians (FWHM) and

θ is the Bragg angle.

The elemental composition of the HgS particles was analyzed by X-ray photoelectron spectrometry (XPS) using a PHI VersaProbe Scanning XPS Microprobe.

Additional details on the preparation of samples for these analyses are provided in the

SI section.

2.2.4 Mercury Methylation Bioassay

The bacterial cultures were pre-grown in a fermentative medium and incubated until exponential growth phase (19 h for D. propionicus 1pr3 and 67 h for D. desulfuricans

ND132) prior to dosing with mercury. In the dissolved Hg+S exposure, Hg(NO 3)2 and

Na 2S were added into the test cultures separately. While we refer to this exposure as

“dissolved Hg+S” because of the initial method of dosing, HgS was supersaturated in these cultures and likely consisted of early-stage precipitation products (i.e., HgS clusters and nanoparticles). The cultures were also exposed to humic-associated HgS nanoparticles, representing an intermediate stage of heterogeneous HgS precipitation, and microscale crystalline HgS, representing a mercury-bearing mineral encountered in soil and sediments (118, 207). In the dissolved and micro-HgS treatments, SRHA, NaNO 3 and HEPES were also added to the test cultures to account for the chemical carryover from the HgS nanoparticle stock in the nano-HgS treatment. The concentration of spiked mercury sulfides was 6 to 20 nM in micro-HgS treatment and 1 to 5 nM in all the other treatments. The cultures were continuously mixed end-over-end and stored in an anaerobic chamber during incubation (1 to 10 days). All mercury methylation bioassays

were conducted in the dark at room temperature (25-27°C). All the bacterial cultures were buffered with 19.2 mM 3-(N-morpholino)propanesulfonic acid (MOPS) at pH 7.0-

7.3.

At each time point, triplicate vials were sacrificed and subsampled for measurements of total protein and mercury concentration. After subsampling, the remaining cultures were preserved by adding 0.4% (v/v) concentrated hydrochloric acid

(HCl) (trace metal grade) and stored at 4°C prior to MeHg analysis. Two sets of controls were incubated under the same conditions including: 1) abiotic control consisting of uninoculated media amended with Hg(NO 3)2 and Na 2S; 2) killed control consisting of autoclaved (121°C, 30 min) cultures amended with Hg(NO 3)2 and Na 2S after the autoclave step. MeHg concentrations in all control samples were below the detection limit (≤ 8 pM MeHg) and significantly lower than MeHg in viable cultures amended with Hg(NO 3)2 and Na 2S.

2.2.5 Chemical Analysis

MeHg concentration was quantified by distillation, aqueous phase ethylation, gas chromatographic separation, and atomic fluorescence spectrometry (Tekran 2600)

(208). Samples for total mercury analysis was first digested with 2 to 4% (v/v) BrCl for at least 12 h and analyzed by SnCl 2 reduction, gold amalgamation, and cold vapor atomic fluorescence spectrometry (209).

2.2.6 Mercury Fractionation by Filtration

In mercury methylation bioassays, total mercury in a subset of D. propionicus

1pr3 cultures was fractionated using filtration. Separate test cultures were filtered with either 0.22 μm polycarbonate (GE Osmonics Labstore) or 0.02 μm aluminum oxide

(Whatman) syringe filters. Total mercury concentration in the filtrates was quantified and represented different mercury species: 1) <0.02 μm fraction considered the nominally “dissolved” mercury and likely consisted of aqueous mononuclear mercury complexes (e.g., Hg(OH) x2-x, Hg(HS) x2-x) and possibly polynuclear mercury sulfide clusters; 2) 0.02 to 0.22 μm fraction which contained colloidal mercury (e.g., Hg-S-NOM nanoparticles); and 3) >0.22 μm fraction which contained cell- and/or large particle- associated mercury. Filtration experiments were also performed with bacteria-free media amended with the three forms of mercury (dissolved Hg+S, nano-HgS, and micro-HgS). These solutions were stored at room temperature and filtered with 0.02 μm and 0.22 μm filters at multiple time points up to 1 day.

Centrifugation and ultracentrifugation were also used to separate dissolved and nanoparticulate species. Further details are provided in the SI section.

2.2.7 TEM Analysis of HgS-Amended Cultures

After 14-h exposure to HgS, cells from 1pr3 cultures were separated by centrifugation, washed with 10 mM phosphate buffered saline (PBS, pH 7.4), and resuspended in a fixative solution containing 4% (v/v) formaldehyde and 2%

glutaraldehyde. After storing for 4 h in the fixative, the cells were washed with high purity deionized water (>18 MΩ-cm). This suspension was deposited on a carbon-coated copper grid (200 mesh) and imaged by FEI Tecnai TEM operating at 80 keV (Figure 2.2b,

2.2d, 2.2f) and a JEOL 2000 FX TEM operating at 200 keV with an energy dispersive X- ray (EDX) spectrometer for the element analysis. (Figure A7).

2.2.8 Mercury Ligand Exchange Reactivity

A subset of mercury-treated media solution (no bacteria) was analyzed for chemical speciation of mercury using a previously developed competitive ligand exchange-solid phase extraction (CLE-SPE) method (18, 92, 99). This technique separates labile mercury species from strongly complexed (i.e., inert) mercury species based on the chemical reactivity of mercury in the presence of a competing ligand: glutathione (GSH) or diethyl dithiocarbamate (DEDC). GSH and DEDC are both thiol-containing compounds and form strong hydrophilic complexes (HgH 2(GSH) 22-) or hydrophobic complexes (Hg(DEDC) 20) with mercury that can be differentiated from the original Hg- sulfide or Hg-NOM species using C 18 -resin solid phase extraction. In the past, researchers have used CLE-SPE to quantify stability constants (assuming that reactions reach equilibrium). Here, we do not imply equilibrium but simply use this method to quantify the reactivity of the Hg for thiol-ligand exchange.

The bacteria-free media containing mercury (either dissolved Hg+S, nano-HgS, or micro-HgS) were sampled at the beginning and end of a one-day holding period at

room temperature. These samples were divided into three aliquots that were amended with either 0.1 mM GSH, 0.1 mM DEDC, or no additional thiol. After 1 h of reaction time, the samples were filtered through a C 18 -resin packed column. The hydrophobic fraction was defined by mercury retained by the resin, while the hydrophilic fraction was defined as mercury passing through the C 18 -filter.The concentration of chemically ‘labile’

Hg was quantified from the difference of hydrophilic Hg concentrations in the thiol- amended sample and in the control (i.e. no thiol added).

2.3 Results and Discussion

2.3.1 Methylation of Mercury Sulfides

The net production of MeHg in the cultures varied depending on the type of HgS added (Figure 2.1). For each SRB strain, the cultures exposed to dissolved Hg(NO 3)2 and

Na 2S (and likely to be precipitating HgS in situ ) demonstrated the highest net MeHg production. MeHg production was observed to a lesser extent in the nanoparticle exposures. In cultures exposed to HgS microparticles, MeHg concentration was less than

8 pM and similar to the autoclaved and abiotic controls. Furthermore, MeHg production by cultures exposed to nanoparticles depended on the age of the nano-HgS stock solutions. The cultures methylated 6% to 10% of the total mercury derived from nanoparticles aged for 16 hours, while cultures methylated a smaller fraction (2% to 4%) if exposed to older nanoparticles (aged 3 days or 1 week) (Figure 2.1a and 2.1b).

Consistent results were obtained in replicate experiments employing higher mercury doses (5 nM) and longer incubation time (up to 10 days) (Figure A1).

Overall, the results demonstrated that the methylation potential of mercury introduced as HgS nanoparticles was greater than bulk scale HgS particles. These results were not due to differences in cell growth, as the optical density (OD 660 ) and protein content were identical in all HgS exposures (Figure A2). In the D. propionicus 1pr3 cultures, OD 660 increased from 0.27 to 0.31 (0.0050 h -1) and total protein concentration increased from 9.5 to 14.1 μg/mL (0.017 h -1) after one day of incubation. In the D. desulfuricans ND132 cultures, OD 660 increased from 0.23 to 0.35 (0.017 h -1) and total protein concentration increased from 26.1 to 49.5 μg/mL (0.027 h -1) after one day of incubation.

The differences of methylation between the nanoparticle and microparticle exposures were likely due to geochemical Hg speciation rather than growth rates of the cultures. The diameter of the nanoparticles was smaller (3 to 4 nm) and specific surface area was larger (220 to 260 m 2 g -1) compared to the microparticles (>500 nm, 2.5 m 2 g -1)

(Table 2.1, Figure A3 and A4). While the surface composition of Hg and S was similar for the nano- and micro-HgS, as shown by X-ray photoelectron spectroscopy (Figure A5a and A5b), the degree of crystallinity varied between nano- and microparticles. X-ray diffraction and electron diffraction data suggested that the nanoparticles were poorly crystalline compared to the HgS microparticles (Figure A4 and A5c).

While nanoparticles generally have high specific surface areas relative to their bulk scale analogs, they can also exhibit unique reactivity due to lattice or surface imperfections that occur with nanoscale particles (121, 122). Here, we provide two lines of evidence to demonstrate that biomethylation of HgS nanoparticles did not depend simply on surface area. First, as the HgS nanoparticles were allowed to age for 16 hours and 3 days prior to exposure to D. propionicus 1pr3, their methylation potential was considerably reduced (Figure 2.1a) while their size and specific surface area remained similar (Table 2.1). Statistically significant differences in geometric surface area and monomer diameter were observed only with the 1-week old nanoparticles.

Second, nano-HgS was more reactive per unit surface area relative to micro-HgS.

MeHg generated from nano-HgS with 5.3×10 -5 m 2 L-1 surface area was 3 times greater than MeHg generated from micro-HgS with 11×10 -5 m 2 L-1 surface area (Figure 2.1c), corresponding to a production of MeHg per m 2 of material that was 6 times higher for nano-HgS compared to micro-HgS (1.13 µmol m -2 and 0.18 µmol m -2, respectively). The reduced availability of nano-HgS during aging may be due to the structural changes occurring with amorphous nanoparticles or cluster/particles at the small size range (1-2 nm). These changes would not be reflected in the diameter and specific surface area measurements from TEM or XRD analyses. Due to sample preparation and analysis requirements, these methods are not quantitative reflections of all forms of mercury in

Figure 2.1: Net MeHg production in SRB cultures exposed to different forms of mercuric sulfides. Methylation by (a) D. propionicus 1pr3 and (b) D. desulfuricans ND132 cultures that were exposed to 1 nM dissolved Hg(NO 3)2 and Na 2S, 1 nM humic-HgS nanoparticles, and 6 nM HgS microparticles. The HgS nanoparticle stock solution was stored at room temperature for 16 h and a longer period (3 days in Figure 2.1a and 1 week in Figure 2.1b) prior to amending to cultures. (c) Methylation by D. propionicus 1pr3 cultures that were exposed to the similar geometric surface area of HgS nano- and microparticles: 1 nM (5×10 -5 m 2 L-1) HgS nanoparticles aged for 16 h, 56 nM HgS microparticles (3×10 -5 m 2 L-1), and 227 nM HgS microparticles (11×10 -5 m 2 L-1). Autoclaved cultures or abiotic culture media were amended with 1 nM dissolved

Hg(NO 3)2 and Na 2S. All cultures received the same humic acid concentration (0.2 μg-C L-1). The error bars represent ±1 s.d. of duplicate samples for the controls and triplicate samples in all other experiments.

the nano-HgS stock solution, and the data are likely to be biased towards more crystalline particles.

The net production of methylmercury was relatively fast in the first few hours and slow after this initial time period (Figure 2.1). This deceleration of MeHg production could not be explained by microbial growth, as we observed a steady increase of cell density throughout the one-day mercury methylation experiments (Figure A2). Similar trends were observed in other mercury methylation studies using the same SRB strains

(193, 204) and estuarine sediment slurries (63). These results are possibly due to the saturation of enzymes and/or depletion of certain compounds (e.g. methyl donors) that were required for mercury methylation. Furthermore, inorganic Hg speciation may have shifted after the first few hours towards less bioavailable forms for the bacteria. The declining net methylation rate may also be explained by the contribution of a reverse process (i.e. methylmercury degradation) balancing overall methylmercury concentrations in the cultures. Desulfovibrio desulfuricans ND132 and Desulfobulbus propionicus strains are known to simultaneously generate and degrade MeHg (51, 204,

210). We performed experiments with the 1pr3 strain exposed to methylmercury chloride and observed MeHg degradation in these cultures (Figure A6).

Table 2.1: Average size and surface area of humic-HgS nanoparticles and HgS microparticles utilized in methylation bioassays. Diameters and geometric surface areas were compared to the 16-h HgS nanoparticles using an unpaired two-tailed t-test. Values that are statistically different (p<0.01) from the 16-h nanoparticles are indicated by an asterisk (*).

Monomer Crystallite 2 -1 Hydrodynamic Surface area (m g ) HgS particles diameter diameter diameter (nm) (a) (nm) (b) (nm) (c) BET (d) Geometric (e)

Nanoparticles 25.8 ± 2.9 3.2 ± 0.8 5.7 ± 0.1 (n=3) 47.9 264 ± 72 (n=110) (aged for 16 h) (n=10) (n=110) Nanoparticles 27.6 ± 3.0 3.3 ± 0.9 5.0 ± 0.3 250 ± 63 (aged for 3 ND (n=5, p=0.27) (n=110,p=0.14) (n=3,p=0.012) (n=110,p=0.13) days) Nanoparticles 28.3 ± 4.9 3.6 ± 0.7* 5.7 ± 0.1 224 ± 47* (aged for 1 ND (n=4, p=0.25) (n=110, p=10 -5) (n=3, p=0.89) (n=110,p=2×10 -6) week) 1457 ± 435* 530 ± 367* 2.5 ± 1.8* Microparticles NA 2.5 (n=7, p=2×10 -8) (n=78, p=10 -25 ) (n=78, p=10 -68 ) (a) Quantified by light-intensity weighted dynamic light scattering. (b) Estimated from individual monomers observed in TEM images (Figure A4). (c) Estimated from the broadening of the X-ray diffraction peak widths by the Scherrer formula (Figure A5c). (d) BET surface area quantified by N 2-gas adsorption. (e) Geometric surface areas (based on approximation of spherical monomers) were calculated from the size of individual particles in TEM images. NA: Not available ND: Not determined

The net production of methylmercury was relatively fast in the first few hours

and slow after this initial time period (Figure 2.1). This deceleration of MeHg production

could not be explained by microbial growth, as we observed a steady increase of cell

density throughout the one-day mercury methylation experiments (Figure A2). Similar

trends were observed in other mercury methylation studies using the same SRB strains

210). We performed experiments with the 1pr3 strain exposed to methylmercury chloride and observed MeHg degradation in these cultures (Figure A6).

2.3.2 Mercury Fractionation in Methylating Cultures

In cultures exposed to the three forms of mercury, we fractionated the mercury into nominally dissolved mercury (<0.02 μm), colloidal mercury (between 0.02 and 0.22

μm), and particulate or cell-associated mercury (>0.22 μm) using filters with two different pore sizes (0.02 and 0.22 μm) (Figure 2.2). Bacteria-free media that were amended with dissolved HgNO 3, nanoparticulate HgS, and microparticulate HgS was also filtered in the same manner (Figure 2.3). The results indicated that the two filters could be used to distinguish these forms of mercury (Figure 2.3a).We also examined the cultures with transmission electron microscopy (TEM) to further differentiate mercury associated with cells from mercury associated with large aggregates of HgS particles. In

the nano-HgS exposures, the filtration results showed that the amount of mercury in the

>0.22 m fraction increased over incubation time (Figure 2.2c). In TEM images, on the other hand, large aggregates of HgS particles (>0.22 m), as seen in the micro-HgS treated cultures (Figure 2.2f and Figure A7c), were not observed in the nanoparticle exposures (Figure 2.2d and Figure A7b).

While mercury was quantified in the >0.22 µm fraction, the nanoparticles were likely too dilute to be observed in bacterial cultures that contained a complex mixture of particles (Figure A7b). The nanoparticles could also be dissolving into solution, as indicated by a small increase of dissolved mercury (from 0.073 nM to 0.21 nM in one day) in bacteria-free media amended with 1 nM nano-HgS (Figure 2.3c). This concentration range is greater than would be expected from the equilibrium solubility of β-HgS (s) (K sp =

10 -38.7±2 for the reaction: Hg 2+ + HS - = HgS (s) + H +) (103). Using Hg-sulfide equilibrium equations described in our previous paper (18), we calculated that in our samples with 1 nM Hg(II) and 1 nM S(-II) at pH 7.5, dissolved Hg concentration at equilibrium with

HgS (s) should be 10 -9 to 10 -5 nM, depending on the solubility product for HgS (s) . In media containing both bacteria and nanoparticles, the percent of total mercury in the >0.22 µm fraction was 60% after 1 day and greater than bacteria-free media containing nanoparticles (8%) (Figure 2.2c and Figure 2.3c), indicating that the nanoparticles were either depositing onto cells or releasing dissolved mercury that was immediately adsorbed to or taken up by the cells. Micro-HgS was less accessible for biomethylation,

possibly due to the minimal or slow mercury dissolution (<0.02 nM mercury dissolved in the bacteria free experiments, Figure 2.3d).

= 10 -38.7±2 for the reaction: Hg 2+ + HS- = HgS (s) + H +) (103). Using Hg-sulfide equilibrium equations described in our previous paper (18), we calculated that in our samples with 1 nM Hg(II) and 1 nM S(-II) at pH 7.5, dissolved Hg concentration at equilibrium with

HgS(s) should be 10 -9 to 10 -5 nM, depending on the solubility product for HgS(s). In media containing both bacteria and nanoparticles, the percent of total mercury in the

>0.22 µm fraction was 60% after 1 day and greater than bacteria-free media containing nanoparticles (8%) (Figure 2.2c and Figure 2.3c), indicating that the nanoparticles were either depositing onto cells or releasing dissolved mercury that was immediately adsorbed to or taken up by the cells. Micro-HgS was less accessible for biomethylation, possibly due to the minimal or slow mercury dissolution (<0.02 nM mercury dissolved in the bacteria free experiments, Figure 2.3d).

Figure 2.2: Percentages of mercury in solution after filtration of D. propionicus 1pr3 cultures and TEM images of the cultures.

Cultures were exposed to 1 nM dissolved Hg(NO 3)2 and Na 2S (a and b); 1 nM humic- HgS nanoparticles (aged for 16 h, c and d); and 6 nM HgS microparticles (e and f). “Hg<0.02 μm” represented the fraction of total mercury that passed through 0.02-μm filters. “0.02 μm

a. b.

c. d.

Figure 2.3: Filtration of mercury-amended bacteria-free media (for culturing D. propionicus 1pr3).

(a) Medium solution was amended with 1 nM Hg(NO 3)2, 1 nM humic-HgS nanoparticles (aged for 16 h) or 6 nM HgS microparticles, and filtered immediately (less than 10 min) after mercury addition. Filtration of separate samples at different time points after they were amended with different Hg-sulfide species, including (b)

1 nM dissolved Hg(NO 3)2 and Na 2S, (c) 1 nM humic-HgS nanoparticles (aged for 16 h), and (d) 6 nM HgS microparticles. “Hg<0.02 μm” represented the fraction of total mercury that passed through 0.02-μm filters. “0.02 μm

In the dissolved Hg+S and nano-HgS exposures, the amount of mercury in the

>0.22 µm fraction was similar (Figure 2.2a and 2.2c), yet these treatments exhibited markedly different MeHg production (Figure A8). While HgS clusters and nanoparticles were likely forming in the cultures receiving dissolved Hg(II) and S(-II), net MeHg production was faster in the dissolved mercury exposure than the nanoparticle exposure.

These results agree with recent studies that suggested transmembrane mercury uptake as the rate-limiting step of intracellular mercury methylation (51, 211) and imply that

HgS nanoparticles are not as bioavailable as their precursors (e.g., dissolved mercury- sulfide complexes and clusters). Similar patterns of mercury size fractionation were observed in replicate cultures processed by centrifugation and ultracentrifugation

(Figure A9).

2.3.3 Mercury Speciation by Competitive Ligand Exchange

We applied competitive ligand exchange-solid phase extraction (92, 99) to further examine the ligand exchange reactivity of mercury in bacteria-free media. In this method, labile mercury species were replaced by Hg-thiol complexes: either hydrophilic mercury-glutathione (GSH) complexes or hydrophobic Hg-diethyldithiocarbamate

(DEDC) complexes. Labile mercury was quantified by the change of mercury in the hydrophilic fraction (defined as mercury passing through a C18 -resin filter). In the micro-

HgS exposure, mercury speciation remained unchanged after addition of GSH or DEDC

(Figure 2.4), indicating that mercury was largely inert. In the media containing nano- 58

HgS, the amount of hydrophilic mercury increased after 1 day, and this fraction was mostly removed by DEDC ligand exchange (Figure 2.4b), indicating the presence of labile mercury. However, in the dissolved Hg+S exposure, the changes in the hydrophilic mercury fraction after addition of GSH and DEDC were both larger than in the nano-HgS treatment. This pattern of decreasing thiol-exchange reactivity between the dissolved, nanoparticulate and microparticulate mercury corresponded to decreasing methylation rates (Figure 2.1). Mercury is believed to bind to bacterial cells through thiol-containing ligands on the membrane surfaces (212), and these complexes may enter the cells as favorable substrates for methylation (211) . Hence, the labile mercury quantified by thiol ligand exchange could signify the available fraction of mercury for microbial uptake and methylation.

2.3.4 Environmental Implications

Our overall results have shown that mercury bioavailability (and methylation potential) is not adequately represented by equilibrium speciation of aqueous dissolved mercury (defined by a 0.2 or 0.4-µm filter) (12). Our previous work (18) has suggested that HgS (aq) 0 (a form of dissolved mercury presumed to be bioavailable (12)) may represent HgS nanoparticles rather than a mononuclear aqueous mercury-sulfide complex.

Figure 2.4: Hydrophilic mercury in HgS-amended media after competitive ligand exchange with GSH or DEDC and C 18 solid-phase extraction. Uninoculated medium solutions (for culturing D. propionicus 1pr3) were spiked with

1 nM dissolved Hg(NO 3)2, 1 nM dissolved Hg(NO 3)2 and Na 2S, 1 nM humic-HgS nanoparticles (aged for 16 h), and 6 nM HgS microparticles. Ligand exchange reactions were performed by amending aliquots of these samples with GSH or DEDC at two time points after the mercury amendment: (a) Immediately (less than 10 min). (b) 23 h. GSH and DEDC were mixed in the samples for 1 h and then filtered through a C 18 resin. All solutions received the same humic acid concentration (0.2 μg-C/L). Error bars represent ±1 s.d. for duplicate samples.

Bacteria are not known to directly take up nanoparticles without compromising their membrane integrity. Moreover, previous works have indicated that bacteria can take up metal constituents of nanoparticles through the dissolution of nanoparticles that accumulated at cell surfaces (126, 201, 202). Indeed, a mechanism of direct uptake of nanoparticles is not necessary for explaining the data presented here. Instead, our results suggest that bioavailability is related to the kinetics of reactions (rates of cluster formation, crystal ripening, dissolution, etc.).

Therefore, a single entity to represent nanoparticulate or colloidal HgS is overly simplistic. The bioavailability of mercury depends on the evolving nanoscale properties of mercury compounds that fall in the fraction typically designated as dissolved and colloidal (less than 0.2 or 0.45 µm). This conclusion could help explain observations that mercury recently deposited to surface waters from the atmosphere (as weak HgCl 2 complexes) is more readily transformed to MeHg than older mercury that persists mainly as crystalline HgS (s) in historically contaminated sediments (109, 213).

Although the occurrence of nanoparticulate or colloidal HgS has been suggested in a number of studies (18, 89, 90, 99, 117-119), our investigation is the first to explore the potential of HgS nanoparticles to serve as an important, but previously unrecognized source of bioavailable mercury for methylating bacteria. Overall, our results point to a new approach that should consider reaction mechanisms and Hg transformation kinetics for modeling mercury bioavailability. Such models could facilitate prediction

and mitigation of MeHg hotspots in the aquatic environment. Given that mineral nanoparticles are ubiquitous in the environment (121), the importance of nanoscale processes for trace metal bioavailability and toxicity has yet to be fully realized. Our findings provide a new approach that may be applied to other metal-sulfide nanoparticles (e.g., ZnS, CuS, FeS) and their potential roles in biogeochemical metal cycling.

Chapter 3. Net methylation of mercury in estuarine sediment slurry microcosms amended with dissolved, nanoparticulate, and microparticulate mercuric sulfides

3.1 Introduction

The major route for human exposure to neurotoxic methylmercury (MeHg) is dietary consumption of contaminated fish(3). MeHg production in aquatic ecosystems mainly occurs through methylation of inorganic mercury by anaerobic bacteria in sediments (46, 214). However, production rate of MeHg appears to vary across aquatic ecosystems, sometimes by orders of magnitude(192, 215, 216). Although numerous studies have been conducted to investigate the microbiological(60, 217-219) and geochemical factors(100, 112, 220, 221) that control MeHg production in natural sediments, great uncertainty remains regarding the biochemical pathway of MeHg production and the identity of inorganic mercury species available for methylation.

Our previous study has indicated that the availability of mercury for methylation decreased during aging of mercury in sediments, a process in which mercury is expected to transform from dissolved Hg-sulfides to nanoparticulate and microparticulate forms of HgS. Our results demonstrated that nanoparticulate HgS resulted in significantly higher MeHg production relative to micro-crystalline HgS. Dissolved Hg species (i.e.,

Hg that passed through 0.02 μm filters or remained in the supernatant after ultracentrifugation) was most susceptible to methylation by pure cultures of sulfate- reducing bacteria (SRB) grown under fermentative conditions (222).While this previous 63

work highlighted the importance of Hg-sulfide speciation for MeHg production and the previously unrecognized role for nanoparticles, natural sediments comprise of much more complex biological and chemical conditions than represented by the pure bacterial cultures in our previous work. For example, natural sediments often contain a diverse group of microorganisms that are expected to simultaneously generate and degrade

MeHg(51, 61, 223). Non-methylating microbes may also compete with methylating bacteria for labile carbon for growth, resulting in limitations to MeHg production rates.

The geochemistry of mercury in sediments is also different from our previous pure culture experiments. Sediments comprise of a mixture of the mineral particles or particles coated with natural organic matter that can scavenge Hg from pore water via adsorption or complexation(224, 225). Therefore, the results from pure culture studies may not directly apply to real sediments, and the aging effect on the availability of Hg for methylation needs to be examined under more environmentally relevant conditions.

In this study, we conducted sediment slurry microcosm experiments to investigate whether ‘aging’ of mercuric sulfides decreased MeHg production, as we observed in pure cultures of SRB. In these slurry experiments, we also sought to understand how environmental variables (e.g., salinity, organic carbon) could influence the speciation of mercury between various sulfide phases (e.g. dissolved, nanoparticulate, microparticulate) and the consequence of these changes for MeHg production. For our slurries, we selected sediments and water at two locations in the San

Francisco Bay estuary (California, U.S.A.) to represent sediments in freshwater and saline settings. The slurries were amended with one of three forms of mercury: dissolved

Hg(NO 3)2 freshly mixed with Na 2S, HgS nanoparticles, and HgS microparticles. In these

HgS treatments, we compared the net MeHg production and other water quality parameters relative for Hg speciation.

3.2 Materials and Methods

3.2.1 Sediment and water collection

Sediment samples were obtained from two sites at San Francisco estuary, CA in

August 2011, as part of a periodic sampling of the San Francisco Estuary Institute

Regional Monitoring Program in the San Francisco Bay-Delta region. Site BG30 is a freshwater location in the San Joaquin River (38.023° N, 121.808° W), and site LSB129S is a saline water located at lower south bay area near the city of San Jose (37.487° N,

122.101° W). Triplicate surficial (0-5 cm) sediment samples were collected with Van Veen sediment grab samplers and packed into acid-cleaned polyethylene jars with Teflon- lined caps. The sediment samples were covered with a thin layer of overlaying water, and the sample jars were sealed with no headspace. At the same sampling sites, surface water was also collected using acid-cleaned polyethylene jugs filled to capacity.

Sediment and water samples were transported on ice to the laboratory and stored at 4°C.

A portion of each sediment samples was analyzed for total mercury, methylmercury, acid volatile sulfide (AVS) and water content. Surface water samples were analyzed for

pH, total organic carbon (TOC), sulfate (SO 42-), ferrous iron (Fe(II)), total Fe, and total mercury (Hg) content.

3.2.2 Pore water characterization

Pore waters were extracted from the sediment samples by centrifugation at 3000 g for 20 min and analyzed for total mercury concentration. Total mercury in pore waters was fractionated using a previously developed centrifugation method(222). The bulk- scale particulate mercury was first separated from pore waters by centrifugation at 6,700 g for 5 min. The mercury in the supernatant was further fractionated by ultracentrifugation at 370,000 g for 1 h. After ultracentrifugation, the mercury remaining in the supernatant was considered to be nominally dissolved, while mercury in the pellet after ultracentrifugation consisted of colloidal mercury. The applicability of this procedure for separating dissolved, nanoparticulate and microparticulate mercury was demonstrated by centrifugation and ultracentrifugation of simulated natural water that were treated by these different forms of mercury (Figure 3.1a).

Pore water samples were also filtered through 0.2 μm nylon syringe filters in an anaerobic chamber. A portion of the filtered pore water was analyzed for dissolved Fe(II) immediately. Aliquots for TOC and SO 42- measurements were refrigerated prior to analysis. Aliquots for analysis of AVS were preserved with 0.01 N zinc sulfate (ZnSO 4) and 0.01 N potassium hydroxide (KOH, trace metal grade) at 4°C. Additional samples were collected for analysis of major cations (e.g., Na, Mg, K, Ca, Fe, Mn, etc. ).

3.2.3 Sediment slurry preparation

For slurry preparation, sediment samples were homogenized, apportioned into

acid-cleaned serum bottles, and mixed with N 2-purged surface water in an anaerobic chamber. A redox indicator, resazurin, was added to a final concentration of 2 mg/L. The serum bottles were then capped with butyl rubber stoppers and crimped with aluminum seals. The slurries were incubated in dark at room temperature (20-22°C) to deplete residual oxygen. The slurries were not utilized for mercury methylation experiments until resazurin became clear.

The preparation of the sediment slurries involved two experimental procedures that involved differences in the solid to water ratio and pre-incubation period (prior to

Hg addition). For Experiment 1, 20 g (wet weight) sediment was mixed with 140 mL water and incubated for 2.5 days prior to mercury spikes. For Experiment 2, 50 g (wet weight) sediment was mixed with 120 mL water and an external carbon source (10 mM sodium pyruvate). These slurries were incubated for 7.5 days prior to mercury amendments.

3.2.4 HgS particle preparation

The mercury stock solution consisted of Hg(NO 3)2 dissolved in 0.1 N HNO 3. Na 2S stocks were prepared by dissolving freshly washed and dried crystals of Na 2S·9H 2O in

N2-purged water and were utilized within 20 h of preparation. HgS nanoparticles were synthesized by dissolving 50 μM Hg(NO 3)2 and 50 μM Na 2S with 10 mg-C L -1 Suwannee

River humic acid (SRHA, International Humic Substances Society) in a solution of 0.1 M

NaNO 3 and 4 mM sodium 4-(2-hydroxyethyl) piperazine-1-ethanesulfonate (HEPES)

(pH 7.5, filtered to <0.1 μm). The Hg-S-NOM nanoparticle stock solution was allowed to age for 16 h at room temperature prior to use in the methylation experiments. A microparticulate HgS stock suspension was prepared by adding a commercial metacinnabar powder (β-HgS, Alfa Aesar) into nanopure-filtered water (>18 MΩ-cm).

This suspension was mixed end-over-end prior to taking an aliquot for the experiments.

Our previous data indicated that the nanoparticulate HgS comprised primarily of metacinnabar-like particles with an average diameter of 3.2 ± 0.8 nm (based on TEM analysis) and geometric surface area of 264 ± 72 m 2 g -1. The microparticulate HgS was actually a mixture of metacinnabar and cinnabar particles (based on X-ray diffraction) with an average diameter of 530 ± 367 nm (based on TEM analysis) and geometric surface area of 2.5 ± 1.8 m 2 g -1(222).

3.2.5 Mercury methylation experiments

After pre-incubation, sediment slurries were spiked with three forms of mercury sulfides, including dissolved Hg+S (i.e., Hg(NO 3)2 and Na2S added from their respective stock solutions), HgS nanoparticles and HgS microparticles. Three sets of controls were also incubated in the methylation experiments: (1) A mercury blank consisting of sediment slurries without the mercury addition; (2) Amendment with 20 mM sodium molybdate, a specific inhibitor of sulfate reduction, added one day prior to dosing with

Hg(NO 3)2 and Na 2S; (3) autoclaved control consisting of autoclaved (121 °C, 30 min) slurries amended with Hg(NO 3)2 and Na 2S after the autoclave step. In the dissolved

Hg+S exposure, micro-HgS exposure and the three controls, SRHA, NaNO 3 and HEPES were added to the slurries to account for the chemical carryover from the HgS nanoparticle stock in the nano-HgS treatment. The spiked mercury sulfide in each slurry sample was 200 nmol in micro-HgS treatment and 50 nmol in all other treatments. All the slurries were incubated statically for up to 7 days in the dark at room temperature

(20-22°C).

At each time point, replicate serum bottles (n=2-3) were sacrificed for chemical analysis. First, the slurries were mixed end-over-end, and 1 mL of gas was collected from the headspace using a gas-tight syringe for analysis of volatile Hg (mainly Hg 0).

The gas samples were then injected into nanopure-filtered water (>18 MΩ-cm) containing 2% (v/v) BrCl. These samples were stored in gas-tight glass vials and equilibrated for > 3 days under room temperature prior to total mercury analysis. The highest Hg 0 concentration detected from these samples accounted for < 0.1% of total Hg spike. Thus, Hg(II) reduction to gaseous Hg 0 was not significant in our experiments.

After collection of the headspace sample, liquid aliquots were collected from the water overlaying the sediment, filtered through 0.2 μm nylon syringe filters and preserved for chemical analysis, including total Hg, AVS, Fe(II), TOC, major anions (e.g.,

SO 42- and Cl -) and major cations (e.g., Na, Mg, K, Ca, Fe, Mn, etc. ). After pore water

sampling, aliquots of sediment-water mixture were fractionated using the

(ultra)centrifugation method, as described in section 3.2.2, to separate particulate, colloidal, and dissolved forms of Hg. The remainder of the slurries were frozen (-80°C) until analysis for whole sediment AVS, MeHg, and total mercury content. The pH of the slurries and the dissolved Fe(II) concentration were measured immediately after sample collection. All sample collection was conducted in an anaerobic chamber (Coy Lab

Products).

3.2.6 Chemical analysis

MeHg concentration in the slurry samples was analyzed using a modified version of a previous method (226). MeHg was extracted from an aliquot of the slurries

(containing ~1 g wet sediment) by adding 5 mL of a potassium bromide (KBr) extraction solution, 1 mL of 1 M CuSO 4, and 10 mL dichloromethane (CH 2Cl2) into the slurry samples The KBr extraction solution consisted of 18% KBr (w/v) in 5% concentrated 36

M H 2SO 4 (v/v, trace metal grade) and 0.02% NH 2OH•HCl (w/v). The sediment-CH 2Cl2 mixture was held under static conditions for 1 h and then continuously mixed for 1 h.

The CH 2Cl2 phase of this mixture was then separated from the sediments. MeHg that partitioned in the CH 2Cl2 was back-extracted into nanopure-filtered water while heating at 70°C for 2 h. At the end of the heating, CH 2Cl2 was purged from the sample with ultrapure N 2 for 5 min. MeHg concentration in the aqueous phase was quantified by aqueous phase ethylation, gas chromatographic separation, and atomic fluorescence

spectrometry (Brooks Rand Model III)(208). Sediment MeHg concentrations were corrected for extraction efficiency using a standard reference material (CC580) and reported on a dry weight basis.

For total mercury analysis, water samples were first digested with 0.5-2% (v/v)

BrCl and analyzed by SnCl 2 reduction, gold amalgamation, and cold vapor atomic fluorescence spectrometry(209). Total mercury in sediment slurries was measured by atomic absorption spectrometry (Milestone DMA-80)(227).

Acid volatile sulfide (AVS) in water and slurry samples was quantified using acid leaching with 1 N HCl, volatilization of H 2S and subsequent trapping in 0.5 N

NaOH, followed by colorimetric detection of sulfide by the Cline method(228, 229).

Dissolved Fe(II) concentration was determined using a phenanthroline colorimetric method(230). SO 42- and Cl - were analyzed on a Dionex ICS-2000 ion chromatograph equipped with an AS18 analytical column, ASRS 300 suppressor and KOH eluent generator (Dionex, Sunnyvale, CA). TOC were measured on a TOC-V CPH total organic carbon analyzer (Shimadzu, Kyoto, Japan). Water samples for analysis of major cations were treated with 0.5% (v/v) hydrochloric acid (HCl, trace metal grade) and 2% (v/v) nitric acid (HNO 3, trace metal grade), and analyzed using inductively coupled plasma mass spectroscopy (ICP-MS, Agilent Technologies, Santa Clara, CA). The water content in sediment samples was determined by overnight drying of 4-5 g wet sediments at

100°C.

3.2.7 Equilibrium speciation calculations

Hg and Fe(II) speciation calculations were conducted using MINEQL + software

(version 4.5, Environmental Research Software)(231) to determine saturation indices for metacinnabar (HgS (s) ), mackinawite (FeS (s) ), pyrite (FeS 2(s) ). The input parameters included measured pH, total Hg, Fe(II), AVS, Cl -, and thiol concentration in the water samples (slurries filtered to <0.2 μm). The thiol concentration was estimated from the

TOC measurements of the water samples from the slurries without carbon addition, assuming 73 μmol thiol per gram TOC(113). The results were corrected for temperature

(21°C) and the ionic strength estimated from the concentrations of the major cations (Na,

Mg, K and Ca). The formation constants used for calculation were from MINEQL+ database or literature, and summarized in Table B1.

3.3 Results and Discussion

3.3.1 Surface water and sediment characteristics

BG30 and LSB129S represented typical conditions in freshwater and saline estuary environment, respectively. Total Hg content in both sediment and water samples were relatively low (Table 3.1) and fell within the range of previous measurements for the San Francisco Bay-Delta region (232), indicating that these sampling areas were not recently impacted by mercury contamination. Although the total sediment Hg from the two sites were close to each other, the fraction of MeHg in total Hg was much higher for LSB129S (0.90%), relative to BG30 (0.11%). The percentage

of total mercury in sediments present as MeHg has been proposed to be a proxy for net

MeHg production rate(107, 215, 233-235). Thus, our results suggested that mercury methylation was likely more active in the sediment at LSB129S, which was possibly due to the high sulfate abundance typically encountered in the marine environment that can result in enhanced productivity of SRB. Elevated AVS concentration in pore water and the whole sediment from LSB129S (an order of magnitude higher than BG30, Table 3.1) also suggested a greater extent of sulfate reduction at this site.

In addition to microbial activity, MeHg production is also dependent on the geochemical speciation of inorganic mercury(12, 222). The mercury speciation in pore waters was analyzed using a centrifugation and ultracentrifugation method. Pore waters were first extracted from BG30 and LSB129S sediments using centrifugation at 3,000 g for 20min, and then further centrifuged at two higher speeds (6,700 g and 370,000 g) to remove large and colloidal particles from the aqueous phase, respectively. Simulated water that was amended with dissolved HgNO 3, nanoparticulate HgS and microparticulate HgS was processed using this method and the three forms of mercury were effectively distinguished (Figure 3.1a).

In the sediment samples, we utilized this procedure to make operationally- defined distinctions between particulate, colloidal, and dissolved Hg in the sediments and to semi-quantitatively assess the distribution of Hg between colloidal and dissolved phases in the porewater. The results from the Hg fractionation in pore waters showed

that the majority of the mercury was associated with bulk-scale particles (i.e., the fraction removed from pore waters after centrifugation at 6,700 g for 5 min, Figure 3.1b).

The colloidal fraction (i.e., the difference between supernatant Hg after centrifugation at

6,700 g for 5 min and supernatant Hg after centrifugation at 370,000 g for 1 h) was significant in BG30 pore water but negligible in LSB129S pore water (Figure 3.1b). This result is consistent with the aggregation of colloidal particles under high salinity condition to form large aggregates that settled from solution after ‘light’ centrifuging (18,

130). Similar patterns have been observed in mercury fractionation experiments with surface waters(85-87). In a separate aliquot of the pore water sample, a portion of the colloidal Hg (as defined above) in BG30 pore water was also able to pass through the 0.2

µm filters. This type of filtration is often used to determine dissolved forms of mercury in water, and our data demonstrate that a significant portion of this Hg is colloidal

(Figure 3.1b).

3.3.2 Net MeHg production in sediment slurry microcosms

The net production of MeHg in sediment slurries varied depending on the procedure of slurry preparation (Figure 3.2). In Experiment 1, the slurries were prepared with a lower sediment-water ratio and pre-incubated without an additional carbon substrate. Here, microbial activity appeared to be limited, especially in comparison to slurries from Experiment 2 that contained a higher sediment-to-water ratio, longer pre- incubation period, and external carbon addition (pyruvate).

Table 3.1: Surface water and sediment characteristics

Site BG30 LSB129S Sample Surface water a Pore water b Sediment c Surface water a Pore water b Sediment c pH 7.62 7.43 ND 7.87 7.66 ND [TOC] (mg/L) 2.8 ± 0.5 20.0 ± 0.2 ND 4.1 ± 0.1 16.4 ± 0.1 ND [SO 42-] (mM) 0.69 ± 0.02 0.40 ± 0.09 ND 22.3 ± 0.1 20.0 ± 0.5 ND AVS (μmol g -1) ND 0.001 ± 0.0002 1.34 ± 0.32 ND 0.021 ± 0.004 13.51 ± 1.54 [Fe(II)] (mM) < DL d 0.10 ± 0.02 ND < DL d 0.11 ± 0.02 ND [Total Fe] (mM) 0.006 ± 0.0001 0.10 ± 0.02 ND 0.011 ± 0.0004 0.13 ± 0.02 ND [Total Hg] (pmol g -1) 0.020 ± 0.008 0.034 ± 0.013 1943 ± 104 0.011 ± 0.001 0.031 ± 0.0005 1477 ± 57 [MeHg] (pmol g -1) ND ND 2.19 ± 0.65 ND ND 13.3 ± 1.5 a 75 The values represent mean ± standard deviation of duplicate samples (w/o 0.2 μm filtration).

bThe values represent mean ± standard deviation of triplicate samples (w/ 0.2 μm filtration). cThe values represent mean ± standard deviation of triplicate samples and are expressed on a dry weight (dw) basis. d< DL: values below the detection limit of Fe(II), 0.004 mM. ND: not detected

Although the slurries in the two experiments received the same amount of Hg spike,the Hg loading per gram sediment was higher in Experiment 1 due to the lower sediment mass. The net MeHg production was reported as the difference between the starting and final MeHg concentration in the slurry microcosms during a 2.8-day incubation in Experiment 1 and 2.2-day incubation in Experiment 2 (Figure 3.2). Despite the longer incubation and higher inorganic Hg loading per sediment mass, net MeHg production in Experiment 1 was very limited, probably due to the suppressed bacterial activity caused by the lack of carbon source. The availability of labile organic carbon is known to be an important factor in determining microbial MeHg production(194, 219).

In Experiment 1, TOC did not appear to decrease during the incubation (Figure B1a and

B1b), possible due to its minimal availability for bacterial metabolism.

In Experiment 2, samples were also collected after the 2.2-day incubation until up to 7 days (Figure 3.3). Net MeHg production in autoclaved slurries was below the detection limit (0.2 pmol g -1 (dw)), indicating that MeHg was generated from biological activity. Addition of molybdate, a specific metabolic inhibitor of SRB, decreased the

MeHg production by 92.9 ± 0.6% in BG30 slurries and 92.7 ± 1.1% in LSB129S slurries.

This result suggested that SRB were the major mercury methylators in this experiment.

The dominant role of SRB in MeHg sedimentary accumulation has been confirmed by many other studies(63-65, 214).

Figure 3.1: Mercury fractionation using centrifugation and ultracentrifugation.

(a) Simulated water, consisting of 5 mM NaHCO 3, 50 mM NaCl and 1 mg-C/L Pony Lake fulvic acid (PLFA, International Humic Substances Society), was amended with

1 nM dissolved Hg(NO 3)2, HgS nanoparticles or HgS microparticles, and (ultra)centrifuged immediately (less than 10 min) after mercury addition. (b) Pore water was first collected from the supernatant after centrifugation of the sediments at 3,000 g for 20 min. This pore water was further centrifuged at 6,700 g for 5 min or 370,000 for 1 h, or filtered with 0.2-μm filters. “Hg < 0.02 μm” represents the fraction of total mercury that passed through 0.02-μm filters. The error bars represent 1 s.d. for duplicate samples.

)

-1 180 Experiment 1: BG30 150 Experiment 1: LSB129S Experiment 2: BG30 120 Experiment 2: LSB129S 90

Net MeHg production (pmol g (pmol production MeHg Net 0 Dissolved Hg+S Nano-HgS Micro-HgS

Figure 3.2: Net MeHg production in sediment slurries 2.2-2.8 days after amendment of

50 nmol mercury (either dissolved Hg(NO 3)2 + Na 2S, HgS nanoparticles or HgS microparticles). The MeHg concentration was normalized to the dry sediment mass. In Experiment 1, slurries consisted of 20 g wet sediment and 140 mL surface water, and pre-incubated for 2.5 days prior to Hg addition. In Experiment 2, slurries consisted of 50 g wet sediment, 120 mL surface water and 10 mM sodium pyruvate, and pre-incubated for 7.5 days prior to Hg addition. The error bars represent 1 s.d. of replicate samples (n=2- 3). The net MeHg production from micro-HgS exposed LSB129S slurries in Experiment 1 was below the detection limit (0.2 pmol g -1).

Most of the sulfate in BG30 samples was consumed prior to the mercury amendments in Experiment 2. Hence, sulfate concentration in the overlaying water was near or below the detection limit (~0.01 mM), which became an inhibitory factor of microbial activity and subsequent MeHg production in these slurries. As a result, TOC degradation appeared to be insignificant (Figure B1c) and less MeHg was generated in

BG30 samples relative to LSB129S samples (Figure 3.3). Moreover, the net MeHg production in BG30 slurries amended with Hg(NO 3)2+Na 2S and nano-HgS was almost identical (Figure 3.3a), which deviates from our previous observation that dissolved

Hg(NO 3)2+Na 2S was more available to methylation than nano-HgS(222). This discrepancy is likely because the previous experiments were conducted using actively

growing bacteria and Hg availability was the limiting factor of MeHg production(222).

However, in the present study, bacterial activity was so low (Figure B2a) that the available Hg was in excess in both Hg(NO 3)2+Na 2S and nano-HgS treatments.

It should be noted that the MeHg concentration in Hg(NO 3)2+Na 2S and nano-HgS treated BG30 slurries at the end of the 7-day incubation was significantly higher than that at the time zero ( p < 0.05, Figure 3.3a). This result suggests that active sulfate reduction may not be a prerequisite for Hg methylation. Pure cultures of SRB have been shown to efficiently methylate mercury under fermentative conditions(193, 236) and comparable MeHg production rates have been detected from natural sediments with low sulfate content relative to those with high sulfate content(215). Therefore, the relationship between mercury methylation and sulfate reduction in SRB may need to be re-examined, particularly in low salinity environment.

In Experiment 2 that included pyruvate addition, sulfate reduction (Figure B2b),

TOC degradation (Figure B1d) and mercury methylation (Figure 3.3b) all actively occurred in LSB129S slurries. The net MeHg production in the slurries amended with

Hg(NO 3)2+Na 2S was found to be the largest of the three types of Hg treatments, similar to our previous experiments with SRB in pure culture (222). Likewise, the methylation potential of nano-HgS was higher than that of micro-HgS (Figure 3.3b). The difference in

MeHg production among the three HgS treatments can not be explained by the activity

Figure 3.3: Net MeHg production in the slurry microcosms exposed to 2 nmol g -1 (dw) dissolved Hg(NO 3)2 and Na 2S, HgS nanoparticles or HgS microparticles in Experiment 2. The slurries were prepared with sediments from site (a) BG30 and (b) LSB129 S. The MeHg concentration was normalized to the dry sediment mass in each serum bottle. Incubation time represents the time after Hg amendments. The error bars represent 1 s.d. for duplicate samples in the test groups. Single replicate of slurries were incubated for the controls.

of SRB,as sulfate concentration decreased at similar rates in all HgS exposures (Figure

B2b). Instead, the availability of inorganic mercury most likely controlled the MeHg production in LSB129S slurries. Furthermore, the net MeHg production during the 7-day incubation in the slurries treated with Hg(NO 3)2+Na 2S was 170 ± 26 pmol g -1 (dw) for

LSB129S and 63 ± 6 pmol g -1 (dw) for BG30, while the sulfate reduction was at least 40- fold greater in LSB129 samples (2.7 mM), relative to BG30 samples (< 0.068 mM). The maximum extent of sulfate reduction in BG30 samples (0.068 mM) was estimated by assuming that the sulfate measured at time zero was completely depleted during the 7- day incubation (Figure B2a). Normalizing the net MeHg production to the sulfate reduction revealed that the mercury species present in the BG30 slurries may be more available for methylation than LSB slurries.

3.3.3 Mercury speciation in sediment slurry microcosms

In Experiment 2, the Hg spikes applied in the Hg(NO3)2+Na 2S and nano-HgS treatments (~2 nmol g -1(dw)) only increased the total Hg concentration to 2 to 2.5 times the concentration in the original sediments (Table 3.1). However, the MeHg production from the newly added Hg (as dissolved Hg-S or HgS nanoparticles) was 18-48 times greater compared to “Hg blank controls” that contained mercury from the original sediment material (Figure 3.3). The availability of native Hg in the sediments resembled the HgS microparticles (Figure 3.3), which can be explained by the pore water fractionation results showing that Hg mainly existed as bulk-scale particles in BG30 and 81

LSB129S samples (Figure 3.1b). This result is also consistent with the previously reported ‘aging’ effects of mercury in sediments that reduce bioavailability over time(109, 213).

To further explore the relationship between mercury speciation and its methylation potential, total aqueous mercury in the slurries amended with the three forms of HgS or lacking Hg addition was fractionated using centrifugation (6,700 g for 5 min) and ultracentrifugation (370,000 g for 1h) in Experiment 2. In both types of sediment slurries (BG30 and LSB129S), the addition of Hg(NO 3)2+Na 2S and nano-HgS appeared to increase the dissolved Hg content of the mixtures relative to the control with no Hg added. For each time point (0 and 7 days) in both BG30 and LSB129S samples, the concentration of dissolved Hg was larger than the dissolved Hg in the slurries exposed to microparticulate HgS and the Hg blank control (Figure 3.4). In our previous research with pure cultures, the dissolved Hg was the most susceptible Hg species to microbial methylation (222). Thus, in these sediment slurries, the higher concentration of dissolved Hg may, at least partly, explain the greater amounts of MeHg production in slurries that were amended with Hg(NO3)2+Na 2S and nano-HgS relative to the other slurries (Figure 3.3).

However, the amount of dissolved Hg alone does not explain the observation that the dissolved Hg fractions in LSB129S samples exposed to Hg(NO 3)2+Na 2S and nano-HgS were similar (Figure 3.4c and 3.4d), yet the net MeHg production in these two

treatments differed significantly (p < 0.05, Figure 3.3b). This discrepancy is probably caused by the mineral particles in the sediments which were not present in the pure bacterial cultures that we utilized for the previous study(222). In Experiment 2, the majority of mercury added as Hg(NO 3)2 or nanoparticles (~345 nmol Hg per L of water in slurries) was immediately bound to large particles. Upon Hg addition in these two treatments, only 0.3-1.3% of the total Hg spike remained in the supernatant after centrifugation at 6,700 g for 5min (Figure 3.4a and 3.4c). Bacteria tend to attach to the surface of mineral particles via sorption or biofilm formation(237). Thus, association with sedimentary particles may localize Hg to the methylating bacteria and create a

‘micro-environment’ with chemical conditions different from the bulk samples.

Moreover, it is possible that a proportion of the dissolved Hg species are not available for microbial methylation. In other words, only the portion of mercury that can be readily taken up by the methylating bacteria contributes to the MeHg production.

Other researchers have postulated that Hg-thiol complexes are the dominant form of Hg species on the bacterial surface(212), and this form of Hg can be taken up into the cells through an active membrane transport mechanism(17, 211). Additionally, recent studies have demonstrated that Hg uptake was the rate-limiting step toward methylation(51,

211), and the microbial MeHg production correlated with the Hg fraction that was readily available for thiol ligand exchange(222).

a. b.

c. d.

Figure 3.4: Hg fractionation in sediment slurries amended with 2 nmol g -1 (dw) (345 nmol Hg per L of water in the slurries) dissolved Hg(NO 3)2 and Na 2S, HgS nanoparticles or HgS microparticles in Experiment 2. The slurries were prepared with sediments from site BG30 (a and b) and LSB129S (c and d). (Ultra)centrifugation was performed immediately (a and c) and 7 days (b and d) after mercury amendments. ‘Dissolved Hg’ represents the total Hg remained in the aqueous phase after ultracentrifugation at 370,000 g for 1 h. ‘Colloidal Hg’ represents the concentration difference of supernatant Hg after centrifugation at 6,700 g for 5 min and ultracentrifugation at 370,000 g for 1 h. The error bars represent 1 s.d. for duplicate samples in the test groups. Single replicate of slurries were incubated for the Hg blank control.

A substantial amount of colloidal Hg was detected in the Hg(NO 3)2+Na 2S and nano-HgS exposed BG30 slurries and this fraction increased during the 7-day incubation period, while the colloidal Hg in the LSB129 samples was little to none (Figure 3.4). This result can be explained by the previous observation that HgS nanoparticles are stable in the presence of natural organic matter, but subject to aggregation in high salinity waters(18).

In sedimentary environments, other sulfide-complexing metals, such as Zn, Fe and Cu, often coexist at much higher levels relative to Hg. In particular, due to the high abundance of iron, Fe(II) likely controls the sulfide speciation and thus indirectly influences the speciation and bioavailability of inorganic Hg in pore waters. In BG30 and

LSB129S sediments, Fe(II) was detected at micromolar quantities (Table 3.1). Also, we observed a surficial black layer inside the LSB129S slurry microcosms, indicative of FeS precipitates.

We attempted to model the equilibrium speciation of Hg, Fe, and S in the filtered porewater from our study. In these calculations, we utilized the concentration measurements of filtered (< 0.2 μm) water from the HgS-treated slurries (Figure 3.5 and

3.6). In the calculations, we assumed that these concentrations represented ‘dissolved’ species, and used the results to determine the saturation indices for mineral phases, including metacinnabar (HgS (s) ), mackinawite (FeS (s) ) and pyrite (FeS 2(s) ). The solubility product Ks0 for metacinnabar varies by 4 orders of magnitude in databases of critically

selected stability constants(103). Here, we utilized two values at the high and low end of this range (10 -36 and 10 -40 ) for our speciation calculations.

In speciation calculations for the BG30 slurries, the HgS (s) saturation index was greater than zero in slurries treated with Hg(NO 3)2+Na 2S and nano-HgS for both values of Ks0 (Figure 3.5). This result indicated that HgS (s) was above saturation and precipitation of metacinnabar was thermodynamically favored in these filtered (<0.2 μm) samples. In the micro-HgS exposed and non-spiked BG30 slurries, oversaturation of

HgS (s) could also occur, particularly at the end of the incubation (Figure 3.5b).

Nanoparticles are small enough to pass through the 0.2 μm filters, although a portion may adsorb or deposit on the filter membrane surface. In contrast, the speciation calculations indicated lower (negative, if Ks0 = 10 -36 ) saturation indices of metacinnabar in

LSB129S samples (Figure 3.6), suggested a lower thermodynamic driving force for HgS precipitation.

The saturation indices for pyrite were 10 to 11 orders of magnitude greater than

zero in all the samples (Figure 3.5 and 3.6), indicating a high potential for FeS 2(s) formation if S 0 was present in the sediments. The increase of AVS content in the slurries did not account for all the sulfate reduced during the 7-day incubation (Figure B3), which may be attributable to the formation of reduced sulfur species such as pyrite that are not detected by AVS. Moreover, all BG30 samples were oversaturated with respect to mackinawite, while the saturation indices of FeS (s) was near or below zero in most

LSB129S samples. Overall, these calculations indicated that the measured colloidal Hg fraction in the slurry samples was not simply HgS. Instead, mercury was probably associated with other metal sulfides (e.g. FeS) via co-precipitation or Hg adsorption to the surface of FeS nanoparticles(238, 239). The role of these colloidal Hg-S-Fe species for microbial methylation potential is unclear and needs to be studied further.

3.3.4 Environmental implications

Our study has demonstrated that MeHg production in natural sediments was largely governed by two factors, the activity of the methylating bacteria and the availability of inorganic mercury. Environmental variables, such as TOC, sulfate, salinity and iron influence MeHg production by regulating one or both of these factors.

SRB was predominantly responsible for Hg methylation in our study and their activity could be limited by the supply of sulfate and/or labile organic carbon.

Nevertheless, in the presence of abundant carbon source (Experiment 2), MeHg production occurred to a comparable extent under sulfate-limited conditions (BG30) relative to sulfate-rich conditions (LSB129S) (Figure 3.3). There are two possible explanations for this phenomenon. First, mercury methylation may not be necessarily coupled to sulfate reduction activity. The relationship between MeHg generation and sulfate reduction in SRB may be more complex than previously recognized. Second, despite the sufficient source of sulfate in the LSB129S slurries, MeHg production was

Figure 3.5: Calculated saturation indices for the water samples withdrawn from the sediment (BG30) slurries. -1 Slurries were treated with 2 nmol g (dw) dissolved Hg(NO 3)2 and Na 2S, HgS nanoparticles or HgS microparticles. Saturation indices of metacinnabar (β-HgS(s) ), mackinawite (FeS (s) ) and pyrite (FeS 2(s) ) were calculated using measured pH, Hg, Fe(II), AVS, Cl - and DOC concentrations in 0.2-μm filtered water samples collected at two time points during the methylation experiments: (a) Immediately (less than 10 min) and (b) 7 days.

Figure 3.6: Calculated saturation indices for the water samples withdrawn from the sediment (LSB129S) slurries. -1 Slurries were treated with 2 nmol g (dw) dissolved Hg(NO 3)2 and Na 2S, HgS nanoparticles or HgS microparticles. Saturation indices of metacinnabar (β-HgS (s) ), mackinawite (FeS (s) ) and pyrite (FeS 2(s) ) were calculated using measured pH, Hg, Fe(II), AVS and Cl - concentrations in 0.2-μm filtered water samples collected at two time points during the methylation experiments: (a) Immediately (less than 10 min) and (b) 7 days.

limited by mercury availability in the marine sediment due to enhanced colloidal aggregation under high ionic strength conditions.

Our results also confirmed the previous pure culture study(222) that aging of mercury from dissolved HgS to nanoparticulate HgS and then crystalline micro-HgS, decreased its availability for microbial methylation. However, the chemistry in natural sediments is more complex than that found in pure cultures of bacteria. Partitioning of

Hg to bulk-scale mineral particles and colloids (especially FeS) may considerably influence the Hg speciation and MeHg production. As a result, dissolved Hg (i.e., Hg remained in the supernatant after ultracentrifugation) in sediment pore water may not accurately represent the available fraction of Hg for methylation. Additional approaches targeting the Hg species that are likely taken up by the methylating bacteria (e.g., Hg- thiol complexes) are needed to assess the Hg availability in natural samples.

Chapter 4. Photolytic degradation of methylmercury enhanced by binding to natural organic ligands

This chapter was published in Zhang T. and Hsu-Kim H. (2010). Nature

Geoscience , 3(7), 473-476. Reproduced with permission from Nature Geoscience.

4.1 Introduction

Demethylation of methylmercury (MeHg) in natural waters occurs through both biological(240, 241) and photochemical processes. In sunlit waters, the photochemical process is the dominant removal pathway for methylmercury and has been estimated to account for as much as 80% of MeHg flux out of lakes(9, 242). While sunlight is known to induce MeHg degradation, rates vary among different types of water.

Photodegradation is rapid in remote lakes in North America and the Arctic(9, 20, 21), but relatively slow in seawater, especially compared to microbial processes(22). The reason for such variability is unknown because the photodemethylation process is poorly understood.

MeHg degradation is likely to be induced by the ultraviolet spectrum (<400 nm)

• of sunlight(20). Hydroxyl radicals ( OH), which are produced via photo-Fenton reactions or by photolysis of nitrate, have been proposed as the reactive intermediates responsible

• for photodegradation(243, 244). However, with previously reported OH demethylation rates ( k ~2 to 9 × 10 9 M -1 s -1)(243-245), significant MeHg degradation by this pathway

• would require large steady state OH concentrations that would occur only under specific conditions: agriculturally-impacted waters (with high nitrate and low dissolved organic carbon(9, 245, 246)) or in acidic waters (pH ≤ 6) with sufficient soluble Fe to

• drive photo-Fenton processes(247). While production of OH is possible at neutral pH(248), production rates are highly sensitive to pH shifts, dissolved natural organic matter (NOM) content, and filtration of samples (in which removal of particulate Fe-

• oxides would decrease OH production rates). Such patterns are inconsistent with previous field studies on MeHg photodegradation in which pH variations (from 6 to 7.5) and filtration of the sample did not change photodegradation rates(9, 21, 242, 249). Such

• discrepancies point to another photoreactive intermediate (in addition to OH) that can induce MeHg degradation in surface waters.

MeHg degradation can also occur through singlet oxygen ( 1O2)(250), which is generated by sunlight sensitization of dissolved NOM(169, 251). In this study, we examined the mechanism of MeHg photodegradation by exposing MeHg to natural sunlight and UV-A radiation in the presence of NOM and selective quenchers for photo- generated reactive intermediates. We also demonstrated that the reaction rates between

1O2 and MeHg were sufficient to account for sunlight-induced degradation. However, these rates applied only to certain types of MeHg species in water: CH 3Hg-thiol complexes with NOM.

4.2 Materials and Methods

4.2.1 Materials

Nanopure-filtered water (>18 M Ω-cm) was used to prepare all stock solutions.

Ultrahigh purity grade nitrogen was used for degassing water. Chemical stock solutions were stored in borosilicate glass containers at 4°C. All containers were acid cleaned by overnight soaking in 1 N HCl followed by 3-times rinses with Nanopure water.

Monomethylmercury chloride stock solutions contained either 1000 mg-Hg/L

(Alfa Aesar) or 1 mg-Hg/L (Brooksrand) that were preserved with 0.5% (v/v) glacial acetic acid and 0.2% (v/v) HCl. The humic acid stock solution was prepared by dissolving Suwannee River humic acid standard (SRHA, International Humic

Substances Society) in nanopure water. The pH was adjusted to 7.5 by trace metal grade

HCl and NaOH (Fisher Scientific). The SRHA solution was then filtered with 0.2 μm nylon syringe filters (VWR) and total organic carbon (TOC) concentration was determined in the filtered stock by combustion catalytic oxidation/infrared spectroscopy

(Shimadzhu TOC Analyzer). Stock solutions of glutathione (GSH, Sigma-Aldrich) and mercaptoacetate (MA, Sigma-Aldrich) were prepared in N 2-purged water and were utilized within 48 hours of preparation. Stock solutions of nitrobenzene (NB,

Mallinckrodt) and 2-chlorophenol (2CP, Alfa Aesar) were prepared by diluting the original liquid stock in nanopure water. A working solution of 10 g/L H 2O2 was diluted from a 30% H 2O2 (Fisher Biotech) stock and utilized within 2 hours of preparation.

Catalase (from bovine liver, Sigma-Aldrich) stocks were prepared by diluting to 0.2 mg/mL in 1 mM phosphate buffer (pH=7.3). Aliquots of the catalase stock were stored at

-20 °C and utilized within 8 hr of thawing. Reagents for thiol detection consisted of 2.0 mM 2,2’-dithiobis (5-nitropyridine) (DTNP, Fluka) dissolved in acetonitrile, 0.5 M sodium acetate buffer (pH = 6), and 7.5 mM tributylammonium (TBA, J.T. Baker) dissolved in an acetate buffer (pH=3.5).

4.2.2 Sample matrix

Methylmercury (MeHg) photodegradation experiments were conducted in both simulated water and seawater. The simulated water consisted of Nanopure water

(Barnstead) buffered to pH 7.3-7.4 with potassium phosphate (1-10 mM). The seawater matrix comprised of filtered seawater from the Gulf of Mexico (Sigma Aldrich). TOC concentration of this water was measured at 2.8 mg-C/L. The phosphate buffer or seawater matrix was then amended with MeHg and a ligand: either GSH, MA, SR humic acid, or Cl -. In buffered solutions without additional ligands, methylmercury speciation was predominantly in the form of CH 3HgOH and CH 3HgHPO 4- complexes, depending on pH and total phosphate concentration.

4.2.3 Photodegradation experiments

Sunlight exposure experiments were conducted on the roof of Hudson Hall at

Duke University, Durham, NC (36°0’0” N, 78°56’28” W). Solar fluence was calculated based on the hourly sunlight intensity data recorded by the National Oceanic and

Atmospheric Agency’s weather station (Site ID #NC Durham 11W) in the Blackwood

Division of the Duke Forest (located approximately 14 km away). Sunlight exposure experiments were performed twice: October 20, 2008 and December 2-4, 2008, during which the average temperature at noon was 15.2°C and 8.9°C, respectively.

Singlet oxygen experiments were performed by exposing the samples containing rose bengal to artificial UV-A radiation at room temperature (25-26°C). The photolysis chamber consisted of four parallel low-pressure fluorescent UV lamps (blacklight blue,

Atlantic Ultraviolet, 15W), which have an output spectrum ranging from 320 to 400 nm and a peak intensity at 365 nm. UV-A and sunlight photodegradation experiments were performed in FEP Teflon bottles. These bottles were expected to transmit approximately

50-80% of radiation in the UV-A range (315-400 nm)(249, 252, 253). Other factors such as lamp intensity (which can dim over time) and distance from the lamp also affected fluence during the photochemical experiments. Thus, data for comparing MeHg degradation over time (e.g., data plotted on the same figure) were acquired by placing the samples side-by-side on one platform at the same time in the UV reactor.

2CP was applied as the probe compound to detect the formation of 1O2 (254, 255) and was added in the same reaction matrix with MeHg- ligand complexes for quantification of rate constants by competition kinetics. During each sunlight and UV-A photodegradation experiment, temperature was controlled by partially submerging the

Teflon bottles in a deionized water bath at ambient temperature. The deionized water

was replaced every couple hours during the photolysis experiment. The temperature of the water was monitored in another Teflon bottle that was placed next to the other samples. The fluctuation of temperature during irradiation was less than ±1 °C.

••• Direct photodegradation by UV-C radiation (λ=254 nm) and OH-induced photodegradation experiments were conducted in a bench-scale reactor with quasi- collimated beam system. The reactor contained four low-pressure mercury vapor lamps

(USHIO, 15 W). Incident UV irradiance was measured by a UV radiometer (International

Light Inc., Model 1700/SED 240/W). Fluence was defined as the average irradiance multiplied by the exposure time. For this system, fluence was calculated as the radiation emitted at 254 nm. Each sample was placed in a 300-mL borosilicate glass dish, covered with a quartz disc, and continuously mixed with a stir bar during UV-C exposure. H2O2

••• was added as a photochemical generator of OH (256), and NB was applied as the probe

••• compound to detect the formation of OH (257). Catalase was added to samples immediately after UV-C irradiation to quench the excess H 2O2 prior to detection (256).

No additional temperature controls were employed since the temperature did not fluctuate more than 1°C during the course of UV-C exposure.

4.2.4 MeHg decomposition by superoxide

••• We performed additional experiments with superoxide (O 2 -) generated from an

••• enzyme reaction to determine if MeHg degradation can occur with O 2 -. The samples consisted of 10 mM potassium phosphate (buffered to pH 7.3) spiked with 15 nM MeHg

and 4 μM GSH. Superoxide was generated at room temperature (25-26°C) by the reaction of 2.5 mM xanthine with 0.15 units/mL xanthine oxidase (XO) in the sample solution. A probe compound 3’-(258)-bis(4-methoxy-6-nitro)benzenesulfonic acid

••• hydrate (XTT) was employed to monitor the formation of O2 - (according to previous

••• published methods (258)). Upon reaction with O 2 -, XTT is reduced to a compound that has an absorbance maximum at 470 nm. Prior to MeHg and GSH analysis, 16 units/mL superoxide dismutase (SOD) was added to samples immediately after sample collection

••• to quench any additional generation of O 2 -.

MeHg recoveries were tested by preparing a separate set of samples with all reagents added (XO, XTT, SOD in phosphate buffer) prior to the MeHg spike. The average recovery of MeHg in this sample was only 75.7% (±0.6%, n=2), possibly due to a matrix effect caused by the combination of XO, XTT, and SOD reagents. Thus, all MeHg

••• concentrations in the O 2 --induced degradation samples were corrected for this matrix effect.

4.2.5 Chemical analyses

Methylmercury concentration was quantified in the water samples using a standard method (259, 260) involving aqueous phase ethylation, purge and trap on

Tenax resin, gas chromatographic separation, and atomic fluorescence spectrometry

(Tekran 2600). Aliquots of samples were taken prior to photodegradation experiments and analyzed to check the recoveries of MeHg from each reaction matrix. The average

recovery of MeHg was 98.3% (±7.6%, n=56). No strong matrix effect was observed during ethylation and MeHg analysis steps. Thus, a pre-separation or distillation step was not necessary for MeHg quantification in the samples.

Total Hg concentration was monitored in the photolyzed samples by atomic absorption spectrometry (Milsteone DMA-80) (261). The average recovery of total Hg was 94.6% (±6.6%, n=39) after the photodegradation experiments, indicating that substantial loss of volatile Hg 0 was not occurring during photolytic exposure.

GSH and MA concentration was quantified using published methods (262). In summary, thiol compounds in sample aliquots were derivatized by DTNP and quantified within 48 hours of derivatization using reverse phase high performance liquid chromatography (RP-HPLC) with detection by UV absorbance. Thiol concentration decreased rapidly due to reactive oxygen species generated during the photolytic exposure. However, during the entire duration of the experiment, thiol concentration always exceeded total Hg concentration (Figure C5). Therefore, we assumed that all MeHg species was complexed by thiol ligands during the experiment.

NB was quantified by a previously developed isocratic elution RP-HPLC method

(256). 2CP was quantified by a gradient elution RP-HPLC (263). All HPLC measurements were carried out on a Varian HPLC workstation fitted with an Alltech C 18 column and a photodiode array detector. Aliquots of samples were taken prior to photodegradation experiments and analyzed to check the recoveries of the HPLC

methods. The average recoveries were 97.1% (±8.2%, n=31) for thiols, 103.2% (±7.3%, n=24) for 2CP and 104.6% (±8.0%, n=25) for NB.

Methylmercury concentration data were reported as average ±1 standard deviation of duplicate or triplicate measurements. Error bars for GSH, MA, NB, and 2CP concentrations represent ±1 standard error calculated from the linear regression of the external calibration employing a 95% confidence level. In all figures, error bars that were smaller than the symbols were not plotted.

4.2.6 Quantification of rate constants

Direct UV-C photodegradation rates were measured by quantifying the first

order fluence-based rate constant, kD, Fluence , for MeHg-ligand complexes (CH 3HgL) corresponding to the following reaction:

+ν →kD CH3 HgL h products (4.1)

The value for kD, Fluence was obtained by linear regression of logarithmic MeHg concentration ratios determined as a function of fluence: [ ] CH3 HgL ln t = −k ⋅ Fluence (4.2) []CH HgL D 3 0

The time-dependent first order rate constant for photodecomposition, k ,tD , can

be estimated from kD, Fluence by considering the incident irradiance IUV of the UV lamps:

= ⋅ kD, t k D , Fluence I UV (4.3)

For our direct UV degradation experiments, IUV was measured by a UV radiometer prior to each experiment. For the time frame of our experiments, we 99

observed direct UV degradation of MeHg only with UV-C irradiation and not with UV-

••• A irradiation. Thus, in calculating indirect photodegradation rates (i.e., by 1O2 and OH)

••• we needed to consider direct UV degradation only in the experiments with OH and not in the experiments with 1O2.

••• For 1O2 and OH-induced degradation experiments, competition kinetics were applied to determine the second-order rate constants for decomposition of CH 3HgL

••• complexes. Probe compounds (either 2CP for 1O2 or NB for OH) were added to reaction matrices. Simultaneous degradation of MeHg and the probe compound was monitored.

Competition kinetics in the singlet oxygen experiments are summarized in the following reactions:

k1 +1 →O2, CH 3 HgL CH3 HgL O 2 products (4.4a)

+1 →kr 2CP O2 products (4.4b)

− ' +1 →kr 2CP O2 products (4.4c)

where 2CP and 2CP - are the protonated and deprotonated forms of 2- chorophenol, and the reaction rate constants are kr = 3.82×10 5 M -1s-1 and kr’ = 2.05×10 8 M -

1s-1 (reference (255)). The overall second-order degradation rate constant for 2- chlorophenol, is equal to:

− []2CP 2CP  = + '   k1 kr k (4.5) O2 ,2 CP [] r [] 2CPtot 2 CP tot

The second-order rate constants for CH 3HgL degradation by singlet oxygen was calculated by combining the rate laws for CH 3HgL and 2CP degradation (264):

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[CH HgL] k1 [2 CP ] 3 t =O2, CH 3 HgL ⋅ tot t ln[] ln [] (4.6) CH3 HgL k1 2 CP tot 0 O2 , 2 CP 0

The rate constant k 1 was calculated from the slope of data plotted as 2 ,CHO 3HgL ln[ MeHg ]/[ MeHg ]0 versus ln[ 2CP tot ]/[ 2CP tot ]0.

In the kinetics experiment with hydroxyl radicals, nitrobenzene (NB) was used as the probe compound. The competition reactions are summarized by the following:

+ν →kD CH3 HgL h products (4.7a)

k• +• →OH, CH3 HgL CH3 HgL OH products (4.7b)

k• NB+• OH →OH, NB products (4.7c)

• 9 -1 -1 where the rate constant for Eq. (7c) is k OH ,NB = 2.43×10 M s (reference (256)).

From Equations (4.7a), (4.7b) and (4.7c), the rate laws for MeHg and NB degradation were derived as the following:

d[ CH HgL ] 3 =− ⋅•  ⋅[][] −⋅ k•  OH  CH3 HgL kD , t CH 3 HgL (4.8a) dt OH, CH3 HgL

d[ NB ] =− ⋅•  ⋅ [] k•  OH  NB (4.8b) dt OH, NB

••• The second order rate constant k • for CH 3HgL degradation by OH OH ,CH 3HgL radicals was calculated by integrating Equations (8a) and (8b) to the following:

[CH HgL ] t ln 3 t =−k ⋅• OH  dt −⋅ k t • OH, CH HgL ∫   D, t (4.9a) []CH HgL 3 0 3 0

[NB ] t t = − ⋅ •  ln k• ∫  OH  dt (4.9b) []NB OH, NB 0 0

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••• The steady state approximation was made for OH radicals. As a result,

Equations (4.9a) and (4.9b) were combined to solve for k • : OH ,CH 3HgL [CH HgL ] −ln3 t / t − k []CH HgL D, t =3 0 ⋅ k• k • (4.10) OH, CH3 HgL []NB OH, NB −lnt / t []NB 0

The rate constant k • was solved from the data by calculating OH ,CH 3HgL [ ] [ ] CH3 HgL NB −lnt / t and −lnt / t from the slope of MeHg and NB concentration data []CH HgL []NB 3 0 0 plotted as ln C/C 0 versus exposure time.

4.2.7 Dark controls and direct photolysis

Dark controls were performed using replicate samples in bottles that were wrapped in foil and placed under the same conditions as the photolyzed samples (i.e., exposed to sunlight, or placed under UV-A or UV-C lamps). As shown in Figures C3, more than 97.2% of MeHg and 99.7% of 2CP were recovered during the entire course of the experiment, indicating that these compounds were not degraded by rose bengal and other sample constituents in the dark. The recoveries of GSH and MA in the dark controls were >80.5%, indicating that some oxidation of the thiols were occurring under

••• ambient (i.e., oxic) laboratory conditions. Dark controls for OH-induced degradation experiments comprised of water samples that contained H 2O2, another type of reactive oxygen species (ROS). As shown in Figure C6, MeHg degradation was not observed

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after 65 min of exposure to H 2O2. Degradation of GSH was observed at a rate similar to

••• that observed during UV-C irradiation (i.e., OH generation).

Dark controls were also conducted for water samples that contained SR humic acid as the source of photochemical reactive species for MeHg decomposition. These samples contained 0.5 nM to 4000 nM MeHg, 2.8 mg-C/L SRHA in a 10 mM phosphate buffer (pH=7.3-7.4) and were wrapped in foil prior to placement in the UV-A batch reactor. The recovery of MeHg during more than 100 hr of UV-A exposure was 96.8%

(±9.1%, n=6).

MeHg degradation by direct photolysis (i.e., no ROS generated) was also tested in simulated water samples containing MeHg. UV-C decomposition of CH 3Hg-GSH complexes was observed (Figure 4.4). However, direct photolysis by UV-A irradiation or sunlight was not observed. The recovery of MeHg was 102.5% (±10.1%, n=4) after 18.5 hr of sunlight exposure (Figure 4.1a) and 103.8% (±6.5%, n=4) after 50 hr of UV-A exposure

(data not shown).

4.3 Results and Discussion

4.3.1 MeHg photodegradation by NOM

In the sunlight degradation experiments, simulated freshwater containing MeHg,

SRHA, and a phosphate buffer (pH 7 to 7.4) was exposed to natural sunlight over 4 days in October and December 2008 on a building roof top in Durham, North Carolina.

During the exposure period, MeHg did not degrade in appreciable amounts in samples

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that contained only the phosphate buffer (Figure 4.1a). MeHg degradation was observed only when humic acid was present and when the MeHg concentration was low (15 nM) relative to humic acid. The different rates observed between the 15 nM and 1500 nM

MeHg treatments could be caused by differences in complexation between MeHg and thiols associated with the humic. In the sample with 2 mg-C/L humic, the reduced-sulfur concentration from the humic acid was estimated at 150 nM (assuming 7.3×10 -5 moles reduced-S per g C(113)) and greater than total MeHg (15 nM). In contrast, the sample with 1.1 mg-C/L humic and 1500 nM MeHg contained 80 nM reduced-S, less than MeHg concentration.

We performed similar photodegradation experiments using a bench scale UV-A reactor (λ=365 nm) to show that degradation rates depended on the concentration of

MeHg relative to humic acid. MeHg degradation occurred only if humic acid was present (Figure C1). Furthermore, degradation was observed only at MeHg concentrations (0.5 nM and 125 nM) that were less than the estimated reduced-S content of the humic (200 nM) (Figure 4.1b). We observed similar trends demonstrating the concentration dependence for MeHg degradation in samples formulated in seawater and irradiated by artificial UV-A (Figure C1).

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Figure 4.1: Concentration dependence of MeHg photodegradation rates. a) MeHg degradation by sunlight in water containing Suwannee River humic acid (SRHA) and phosphate (1 mM, pH = 7).The high MeHg treatment (1500 nM) was exposed in October 2008 (15.2 °°°C average noon temperature). The low MeHg treatment (15 nM) was exposed in December 2008 (8.9 °°°C). Error bars represent ±1 s.d. for replicate measurements (n=2-3) of the same sample. b) MeHg degradation by artificial UV-A ( λλλ=365 nm) in water containing SRHA (2.8 mg-C/L) and phosphate (10 mM, pH 7.4). Data points represent the average MeHg concentration (±1 s.d.) for duplicate samples. 105

Sunlight irradiation of chromophoric NOM is known to generate reactive

• intermediates such as OH, superoxide, and 1O2(265). In additional photodegradation experiments, we utilized selective probes to identify the reactive intermediate that was responsible for MeHg decomposition (Figure 4.2). The addition of β-carotene and sodium azide (NaN 3), known to be sensitive to 1O2, resulted in slower MeHg degradation kinetics (Figure 4.2a). Experiments were also repeated in 98.5% deuterated water (D 2O), a solvent that is ~13 times slower than H 2O at quenching singlet oxygen(169,

266) and thus, increases steady state 1O2 levels. The use of D 2O appeared to increase the rate of decomposition; however, this enhancement was smaller than expected. The

• addition of isopropyl alcohol (a probe for OH(267)) and isoprene (a probe for photosensitized triplet state NOM(267, 268)) did not appreciably change photodegradation rates (Figure 4.2b). MeHg degradation was not observed in dark controls: the recovery of MeHg was 92% (± 8.3%) after 120 hr. Additional dark experiments with enzyme-generated superoxide indicated that superoxide was not capable of decomposing MeHg within this time frame (Figure C2). Collectively these results suggested that 1O2 was the reactive intermediate responsible for photodegradation in our experiments and not other types of intermediates such as •OH, photosensitized NOM, and superoxide.

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Figure 4.2: MeHg degradation via generation of 1O2 from photosensitized humic acid. MeHg was degraded by UV-A in simulated water containing MeHg (0.5 nM), SRHA (2.8 mg-C/L), and phosphate (10 mM, pH 7.3). a) Replicate samples were amended with either D 2O (a singlet oxygen enhancer), NaN 3 or β-carotene (singlet oxygen quenchers); b) No difference was observed in replicates amended with isoprene (a ••• quencher for triplet state dissolved NOM) or isopropyl alcohol (quencher for OH). Error bars represent ±1 s.d. for replicate measurements (n=2-3).

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The experiments with β-carotene, NaN 3, and D 2O demonstrated the consequence of a heterogeneous distribution of 1O2 in humic acid solutions and the binding of MeHg to NOM. According to Latch and McNeill(169), singlet oxygen is generated in the light- absorbing regions of dissolved NOM and quickly quenched by water. Therefore, the steady state 1O2 concentration can be orders of magnitude greater in close proximity to dissolved NOM molecules than in bulk solution(169). In our experiments, this concentration gradient may explain why hydrophilic modifiers such as NaN 3 and D 2O had less effect on MeHg degradation rates than would be predicted based on known reaction rates between the modifiers and 1O2 (Table C1). In contrast, β-carotene is a hydrophobic modifier that sorbed to NOM macromolecules and subsequently inhibited

MeHg degradation to an extent that was predicted (Table C1). These results indicated that MeHg was sorbed to or complexed by humic molecules and, consequently, reacted with 1O2 molecules in the hydrophobic regions of the NOM.

4.3.2 Speciation-dependent MeHg photodegradation rate

In our experiments water composition was an important factor controlling degradation kinetics. Rates depended on the relative concentrations of MeHg and humic acid (Figure 4.1). Direct interactions between MeHg and NOM (such as metal-ligand complexation) may be critical to photodemethylation kinetics.

We tested the importance of ligand complexation by quantifying second-order

1 + - rate constants k 1 for reactions between O2 and CH 3Hg complexes with Cl , 2 ,CHO 3HgL

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glutathione (GSH) or mercaptoacetate (MA). GSH and MA were selected as surrogates of thiol-containing ligands that are expected to bind to CH 3Hg + ions in NOM-containing freshwater(216, 269). Singlet oxygen was generated in these experiments by UV-A sensitization of rose bengal. Dark controls were also performed to ensure that the reactants did not cause degradation (Supplementary Figure C3). Additional experiments with amendments (D 2O or NaN 3) and with Eosin-Y (instead of rose bengal) were performed to demonstrate that singlet oxygen, and not sensitized rose bengal, was the reactive intermediate for the degradation experiments (Figure C4). Reaction rate constants were quantified by competition kinetics in which the simultaneous degradation of 2-chlorophenol was monitored in the same solution.

In water containing GSH and MA, we observed faster MeHg decomposition by

1O2 than in water containing Cl - or the phosphate buffer alone (Figure 4.3). We observed similar trends in experiments with seawater (Figure C5). These results indicated that

1O2–induced degradation was faster for CH 3Hg-GSH and CH 3Hg-MA complexes than

for CH 3HgCl and CH 3HgOH/ CH 3HgHPO 4- complexes (corresponding to PO 4-buffered water with no additional ligands).

Complexation of sulfhydryl groups by mercury was expected to slow thiol oxidation by reactive oxygen species(270). However, the total thiol concentration was in excess of total Hg in the samples. Thus, the thiols were decomposing simultaneously during the reaction (Figure C6), producing compounds such as organic sulfur radicals

109

that may be capable of inducing demethylation. We tested the reactivity of thiol oxidation products by performing dark oxidation experiments with hydrogen peroxide

(H 2O2), which oxidized GSH and MA to form sulfhydryl radical intermediates(271).

Decomposition of CH 3Hg-GSH complexes was not observed, even as GSH (which was in excess of MeHg) was oxidized by H 2O2 (Figure C7). Therefore, we concluded that sulfhydryl radicals were not reactive intermediates that caused degradation.

0.8 0

0.6

50 mM ClCl-- 0.4 - 2- OH/HPO4-OH /HPO 4 [MeHg]/[MeHg] 0.2 100 µM MA 75 µM GSH 0 0 2 4 6 8 10 UV-A exposure time (hr)

Figure 4.3: Effect of ligand complexation on singlet oxygen-induced of MeHg.

MeHg (350 nM) was complexed with GSH, MA, Cl -, or OH -/HPO 42- in water containing 2-chorophenol (40 μM) and phosphate (10 mM, pH=7.3). 1O2 was generated by UV-A (λ=365 nm) irradiation of rose bengal. Lines indicate model predictions from rate constants. Error bars represent ±1 s.d. for replicate measurements (n=2-3).

Second-order reaction rates constants (Table 4.1) calculated from the 1O2

3 3 experiments indicated that k 1 ,CHO HgL values for CH Hg-GSH and CH Hg-MA complexes 2 3 were 2 to 4 times greater than the rate constant measured for the CH 3HgOH/

CH 3HgHPO 4- mixture. Degradation of CH 3HgCl was not detected for the time scale of

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our experiments (34 hr). Thus, the CH 3HgCl rate constant with 1O2 was at least an order

• of magnitude slower. For the OH degradation pathway, rate constants did not appear to differ between CH 3Hg-thiol complexes and CH 3HgOH/CH 3HgHPO 4- mixed speciation (Table 4.1).

Table 4.1: Degradation rate constants for methylmercury complexes.

Values represent average (± 1 standard deviation) of duplicate photochemical experiments at 25-26 ºC.

• 1O2-induced OH-induced Direct UV (λ=254 MeHg species degradation degradation nm) degradation -1 -1 -1 -1 k 1 (M s ) k • (M s ) D 2 2 ,CHO 3HgL OH ,CH 3HgL k (cm /mJ)

CH 3HgCl <0.32×10 6 (b) NA NA

CH 3HgOH, 1.9 (±0.04)×10 6 1.9 (±0.09)×10 9 <0.23×10 -4 (c) CH 3HgHPO 4- mix (a)

CH 3Hg-GSH 7.4 (±0.2)×10 6 1.9 (±0.04)×10 9 2.2 (±0.02)×10 -4

CH 3Hg-MA 4.9 (±0.5)×10 6 NA NA

(a) The initial methylmercury speciation (in phosphate buffer) was 49% CH 3HgOH and 51% CH 3HgHPO 4- in the 1O2-induced degradation experiments (pH = 7.3), and 90% • CH 3HgOH and 10% CH 3HgHPO 4- in the OH-induced and direct UV-C degradation experiments (pH = 7.4). (b) Calculated limit of quantification based on >95% of the initial MeHg measured in solution after exposure to UV-generated 1O2 for the maximum duration of the experiment (34 hr). (c) Calculated limit of quantification based on >94% of the initial MeHg measured in solution after the maximum UV fluence (2600 mJ/cm 2) for the experiment. NA: Not available

The increased reactivity between CH 3Hg-thiol complexes and 1O2 may be caused by shifts in carbon-mercury bond energies induced by thiol complexation. In theoretical studies, Ni et al.(23) demonstrated that coordination of a reduced sulfur residue (-SH) to

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CH 3Hg + resulted in greater electronegativity at the carbon atom. Tossel(19) also utilized quantum mechanical calculations to show that sulfhydrl coordination decreased electron transition energies to a range where irradiation by low UV wavelengths ( λ< 270 nm) could induce demethylation.

We confirmed these theoretical calculations by performing direct photolysis experiments with UV-C radiation ( λ = 254 nm) to demonstrate that thiol complexation alters electron transition energies of the CH 3Hg-ligand molecule. Degradation of MeHg by UV-C was observed only if methylmercury was in the form of CH 3Hg-GSH complexes (Figure 4.4).

2.5

2.0

1.5 /mJ) 2 (cm 4 1.0 *10 D k 0.5

0.0 0.0 0.5 1.0 1.5 2.0 15 20 25 30 [GSH] :[MeHg] 0 0 Figure 4.4: Effect of ligand complexation on direct UV-C degradation of MeHg. Reaction matrices contained MeHg at varying initial ratios of MeHg and GSH (4 μM initial MeHg, 1 mM phosphate, pH 7.4). Solid line indicates a linear regression of data for [GSH] 0:[MeHg] 0 ≤ 1.

Degradation was not observed in the phosphate buffer alone, where methylmercury consisted of 90% CH 3HgOH and 10% CH 3HgHPO 4-. Moreover, the

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observed decomposition rates increased as the GSH:MeHg molar ratio increased from zero to 1, and remained constant for all mixtures where GSH was in excess of MeHg.

GSH did not decompose during direct UV-C exposure (91.9% ± 1.4% GSH recovered after 2600 mJ/cm 2 fluence). These results showed that GSH complexation altered the reactivity of the MeHg molecule to make it more susceptible to photodegradation.

4.3.3 Environmental Implications

Speciation-dependent degradation rates can explain the variability observed between freshwater lakes and seawater. MeHg speciation differs greatly between seawater and freshwater due to ligand competition between Cl - and thiolate (R-S-) functional groups associated with NOM. CH 3HgCl complexes are likely dominating methylmercury speciation in coastal marine waters with Cl - concentration greater than

0.1 M (Figure C8). In contrast, most freshwater systems contain enough dissolved NOM

(>1 mg-C/L) and low Cl - such that CH 3Hg + ions are bound to thiol ligands associated with NOM(269, 272, 273).

In sunlit surface water containing NOM, the steady state concentration of 1O2 can be as high as 10 -13 to 10 -12 M (169). Using the CH 3Hg-thiol rate constants in Table 4.1, we calculated a first order photodegradaton rate ranging from 0.04 to 0.6 d -1, values that are close to rates quantified in lakes(9, 21, 249). Overall, our results indicate that ecosystem mass balances for MeHg must consider water composition and MeHg speciation when estimating losses from sunlight decomposition. With this new understanding of how

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photodemethylation occurs, we can improve efforts to predict mercury cycling in the aquatic environment and prevent bioaccumulation of MeHg in food webs.

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Chapter 5. Conclusions

5.1 Summary

The overarching goal of this dissertation research was to investigate the mechanisms of microbial MeHg formation and MeHg photodegradation in the natural aquatic environment. Most importantly, our study focused on elucidating what mercury species were actively participating in these transformation reactions, and how environmental conditions influenced these processes. To that end, we investigated: 1)

How the kinetically-limited mercury sulfide mineralization reactions control the availability of inorganic mercury to methylating bacteria in sediments (Chapters 2 and 3); and 2) How MeHg is degraded by natural sunlight in surface waters and the role of

MeHg speciation for influencing degradation kinetics (Chapter 4).

In Chapter 2, the methylation potential of three forms of mercury sulfides (i.e., dissolved Hg(NO 3)2 and Na 2S, humic-coated HgS nanoparticles and microscale crystalline HgS particles) that represent different aging states of mercury in natural sediments was examined using pure cultures of SRB. The bacteria exposed to dissolved

Hg-sulfide generated the most MeHg, and the methylation potential of nanoparticulate mercury was considerably higher than microparticulate mercury. This result may be explained by the dissolution of mercury sulfides and their reactivity in thiol ligand exchange. In cultures amended with dissolved and nanoparticulate mercury, mercury

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likely partitioned onto the bacterial cells and released dissolved Hg species that were susceptible to thiol ligand exchange. In contrast, HgS microparticles were inert to thiol ligand exchange and did not dissolve appreciably during incubation. Moreover, in bacterial cultures exposed to nano-HgS, MeHg production decreased with an increase in the age of nanoparticle stock solutions. The HgS nanoparticles comprised of smaller crystallite diameter higher specific surface area, and more disordered structure (lower degree of crystallinity) than the microparticulate HgS. However, no significant structural changes were detected by DLS, TEM, XRD or XPS during aging of HgS nanoparticles.

In Chapter 3, we extended the concepts of the pure culture experiments (in

Chapter 2) and performed methylation experiments in sediment slurry microcosms to better simulate environmental conditions. Natural sediments collected from both freshwater and coastal marine systems were utilized in sediment slurry microcosm experiments to quantify the methylation potential of dissolved Hg-sulfide, HgS nanoparticles and HgS microparticles in relation to various environmental parameters.

As can be seen in Chapter 3, environmental variables, including sulfate, labile carbon, salinity, sulfide and iron, influenced the net MeHg production by regulating the activity of methylating bacteria or the availability of inorganic mercury. SRB appeared to be the principle Hg methylators in the sediments and their productivity was determined by the availability of both labile carbon (electron donor) and sulfate (electron acceptor). Net

MeHg production in nano-HgS amended slurries was consistently greater relative to

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micro-HgS amended slurries in all experiments (freshwater or marine sediment, carbon- limited, sulfate-limited or microbial activity-abundant condition), possibly due to the higher concentration of dissolved mercury species (defined as the mercury remained in the aqueous phase after ultracentrifugation at 370,000 g for 1 h) released from HgS nanoparticles. Dissolved Hg-sulfide appeared to be more available for methylation compared to nano-HgS in the slurries with high sulfate reduction activity, although

MeHg production from these two treatments was indistinguishable under sulfate- limited conditions. The bulk-scale particles in natural sediments (such as FeS or bacterial cells) were likely scavenging mercury added to the system. We hypothesize that such processes would substantially affect mercury speciation in the slurries and the subsequent bioavailability for methylation. In addition, filtered pore water (<0.2 μm) in the freshwater sediment treated with dissolved and nano-HgS was supersaturated with respect to metacinnabar (β-HgS (s) ), mackinawite (FeS (s) ) and pyrite (FeS 2(s) ), suggesting that colloidal/nanoparticulate mercury was likely to co-exist with these particles through co-precipitation with and/or adsorption to Fe-sulfides. In contrast, colloidal mercury fraction was minimal in marine sediment pore water, which may be attributable to the high extend of colloidal aggregation in high salinity environment.

In Chapter 4, we reported our findings regarding the mechanism of sunlight induced degradation of MeHg in surface waters. The degradation kinetics of MeHg was evaluated in natural water and simulated water samples with various concentration

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ratios of DOM:MeHg and in the presence of selective quenchers for photo-generated reactive intermediates. The results demonstrated that photosensitization of DOM by natural sunlight or UV-A radiation induced MeHg degradation via generating 1O2.

Furthermore, significant degradation of MeHg only occurred when the abundance of thiol functional groups in DOM was high enough to complex with all MeHg molecules.

This is likely because thiol-binding can effectively decrease the activation energy of the

Hg-C bond breaking process, and thus increase the reactivity of MeHg toward decomposition. Indeed, the degradation rate of MeHg-thiol complexes appeared to be significantly greater than that of other MeHg complexes (CH 3HgCl or CH 3HgOH) in both 1O2 and UV-C radiation induced degradation experiments. Our study emphasizes that the role of DOM in MeHg sunlight degradation is critical and complex. First, DOM is the photosensitizer that induces MeHg degradation by generating reactive intermediates ( 1O2) upon sunlight irradiation. Second, DOM is capable of binding

CH 3Hg + ions (via thiol complexation) and resulting in elevated exposure to 1O2 within

hydrophobic portions of the DOM macromolecule (where 1O2 concentration is expected to be greater than in bulk solution (169)).

Overall, this dissertation work has demonstrated that not all mercury species are equally available for the transformation processes between inorganic Hg(II) and MeHg.

The occurrence and abundance of the available mercury species is determined by a variety of environmental factors. In natural sediments, where MeHg is generated from

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microbial methylation, DOM-coated HgS nanoparticles may persist and contribute available mercury species by releasing chemically labile mercury. In surface waters,

MeHg complexed with thiol ligands (e.g., DOM) is readily available for sunlight degradation, likely through a pathway involving 1O2 generated from DOM photosensitization.

5.2 Implications and future research

Considering the complexity of the mercury biogeochemical cycle, the intention of this dissertation is not to suggest a specific mechanism and exclude the others. Instead, this study aims to bridge the knowledge gap between the speciation and availability/reactivity of mercury in the processes controlling MeHg concentrations in aquatic ecosystems. Several implications arise from the research presented in this dissertation and are summarized below.

5.2.1 Towards an understanding of mercury methylation potential

The results in Chapter 2 point out that the equilibrium speciation of ‘dissolved’ aqueous mercury (defined by a 0.2-µm filter) cannot accurately predict the production of

MeHg because this mercury fraction (<0.2 µm) contains a continuum of mercury species that include dissolved molecules, polynuclear clusters, and amorphous nanoparticles.

These different forms of mercury exhibited distinct methylation potential. The transformation reactions involving these mercury species, such as cluster formation, monomer aggregation and crystal ripening, are often times kinetically-hindered in the 119

presence of DOM(105). Therefore, the bioavailability and methylation potential of mercury is most likely related to the ‘slow’ kinetics of these processes that control the relative abundance of various mercury species (i.e., those falling through a 0.2-µm filter), rather than the equilibrium chemistry. Future modeling efforts for predicting mercury bioavailability will need to consider the rate of transformations involving mercury species. Such an approach would require a series of rate constants for the geochemical reactions that dictate the concentration of the available forms of inorganic mercury for microbial methylation.

Chapter 2 and 3 provide the first documentation of HgS nanoparticles serving as an important, but previously unrecognized source of available mercury for biomethylation. The enhanced methylation of the nanoparticles (relative to bulk-scale

HgS particles) cannot be simply explained by their greater surface area, and is likely caused by the unique reactivity at the nanoscale. The exact mechanism for microbial methylation of HgS nanoparticles was not thoroughly elucidated with the available data on my dissertation. In particular, my research did not provide a clear reason for the decreased availability of nano-HgS during aging. Future research is needed to address two main questions. First, the properties that change during the aging of mercury sulfides (particularly at the nanoscale) should be indentified and related to methylation potential. The structural changes of HgS nanoparticles aged for different length of time could not be reflected in the DLS, TEM or XRD results in the present study, probably

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because these methods are probing the more crystalline phases in the nano-HgS stock solution and are less sensitive towards amorphous materials. Additional techniques that better probe short scale atomic structure, such as small angle X-ray scattering (SAXS) and X-ray absorption spectroscopy (XAS), should be applied and complemented with bulk scale mercury speciation analysis (e.g., CLE-SPE) to assess the evolving properties of HgS during aging. Secondly, the environmental conditions, including pH, salinity

(ionic strength), the concentration and type of DOM and metal:sulfide ratio, are known to control the formation and stability of metal sulfide nanoparticles(18, 131, 274-276).

The influence of these factors on the properties of mercury sulfide that change during aging and result in lower MeHg production needs to be investigated. The information from these works will possibly link the measurable environmental parameters to the amount of bioavailable mercury in sediments and ultimately to MeHg production.

The understanding of how the geochemical speciation determines the availability of inorganic mercury for methylation will be tremendously improved if future research in the field of microbiology and molecular biology can fully address the microbiological mechanism of mercury methylation. For example, the mechanism of mercury uptake by the methylating bacteria is poorly understood, even though bacterial uptake is believed to be the rate-limiting step of intracellular mercury methylation(51). A few recent studies have proposed that methylating bacteria may preferentially take up Hg(II)-thiol complexes through an active transport system(17, 211). Our results in Chapter 2 agree

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with this observation by showing that the susceptibility of mercury to thiol ligand exchange correlated with the methylation potential of these mercury species. If future studies further corroborate the notion that methylating microbes take up mercury through an active transport mechanism, our research then provides a methodology (i.e.,

CLE-SPE) to quantify the availability of inorganic mercury for microbial uptake and subsequent methylation.

Additionally, the biochemical pathway and the enzymes that are responsible for mercury methylation are currently unknown. Although sulfate reduction rate has been widely used as a measure of microbial activity in mercury methylation studies(63, 218) due to the dominant role of SRB in sedimentary MeHg production(58, 64), our study, among others(193, 215, 236), has shown that MeHg production can occur to a significant extent under sulfate-limited conditions, suggesting that mercury methylation may not be directly coupled with sulfate reduction activity. Thus, further research is needed to determine a better measure of biological activity pertaining to MeHg production. This information will help design the experiments for assessing the bioavailability of inorganic mercury by ensuring that mercury methylation is not limited by microbial activity (e.g., enzyme, methyl donor molecules) and thus MeHg production solely reflects the availability of inorganic mercury.

Because of the complexity of environmental samples, the results from the pure culture study (Chapter 2) should be applied to real sediment system with caution.

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Additional factors need to be taken into account when assessing mercury methylation in natural sediments. In Chapter 3, our results have shown that mercury has a high propensity to partition onto bulk sediment particles. The importance of this process for bioavailability would largely depend on the relative distribution of mercury and methylating bacteria. For instance, mercury bound to sediment particles may be directly available for the methylating bacteria that are also attached to the sediment solid phase.

Thus, the biogeochemical reactions occurring at the microbe-sediment particle interfaces

(e.g., adsorption, complexation, dissolution, precipitation, aggregation, etc. ) may be important for determining MeHg production in real sediments. Analytical tools with both high spatial resolution and chemical sensitivity are required to investigate these interfacial processes, since these reactions involve nano-scale particles (as demonstrated in Chapter 2 and 3) and mercury often occurs at relatively low concentrations in natural samples. High-resolution electron microscopy, synchrotron-based X-ray microscopy, and microprobe mapping have been utilized to examine the distribution of mercury and other trace elements and to identify the ‘hot spots’ of these elements in biological samples(117, 277, 278). These techniques can also be coupled with metal speciation analysis, including X-ray absorption spectroscopy, X-ray diffraction, electron diffraction, and energy dispersive X-ray spectroscopy, and have shown great promise in elucidating the mechanisms of nanoparticle-microbe transformation processes(279).

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5.2.2 Other metal sulfide nanoparticles

In the sediment environment, an array of sulfide-complexing metals, such as Fe,

Mn, Zn and Cu, often exist and possibly affect mercury speciation. In particular, iron sulfides are the most abundant sulfide minerals and thus tend to control the fate of other metals(239), including mercury. Sulfate reduction and iron reduction are both commonly encountered electron-accepting processes in anaerobic sediments. Therefore, the products of these reactions (i.e., S 2- and Fe 2+ ) are expected to precipitate in sediments as FeS (s) . Amorphous mackinawite is the initial product of FeS precipitation. This mineral phase slowly transforms into well-crystalline iron-sulfide minerals, such as pyrite (FeS 2(s) ) under ambient conditions(280). FeS (s) is known to scavenge metals from aqueous phase and renders them less available for biogeochemical reactions via three mechanisms: surface adsorption, co-precipitation and metathesis (surface metal exchange)(239).

In the case of mercury, amendment of wetland sediment with Fe(II) has been suggested as a strategy for preventing MeHg accumulation, based on the idea that Fe(II) can reduce the solubility (and subsequent availability for microbial methylation) of mercury by decreasing the concentration of dissolved sulfide(100, 281). This appeared to be true at high Fe(II) doses, but it has also been repeatedly shown that low doses of Fe(II) enhanced mercury methylation(100, 220, 281), which can not be explained by equilibrium speciation. Instead, this phenomenon may be due to the ‘slow’ kinetics of

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FeS precipitation, and as a result, the stable occurrence of polynuclear cluster and nanoscale particles of FeS(282-284) (as suggested in Chapter 3).

Future research is needed to prove the promoting effect of nano-FeS on mercury methylation and more importantly, to elucidate why this occurs. A few studies have been conducted to investigate the interaction between Hg 2+ and FeS nanoparticles(238,

285, 286). However, the influence of FeS-Hg interaction on the availability of mercury for microbial methylation has yet to be evaluated. Moreover, pre-synthesized FeS nanoparticles were utilized in these studies, which do not necessarily simulate all the environmental conditions. For example, instead of reacting with pre-formed nanoparticulate FeS, Hg 2+ may co-precipitate with Fe 2+ and S 2- in natural sediments, especially during redox oscillation at the sediment-water interface, where mercury methylation often occurs. As suggested in our study (Chapter 2 and 3), the initial precipitation products of HgS (e.g., polynuclear clusters) are likely more bioavailable than the nanoparticles. Thus, the clusters formed during the initial stage of co- precipitation of Fe 2+ , S2- and Hg 2+ may be a potentially important source of available mercury for biomethylation.

It is interesting to note that the proposed mechanism of the interaction between

Hg 2+ and nano-FeS changed from Hg adsorption onto nano-FeS surface to surface precipitation of metacinnabar as the concentration ratio of Hg 2+ to nano-FeS increased(238). Also, natural organic matter (NOM) appeared to compete with

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disordered mackinawite for binding with Hg 2+ through thiol complexation(286). These parameters (Fe, S, Hg and NOM content) may significantly vary in natural settings and thus affect reactions involving Fe, S and Hg. Therefore, future research should focus on:

1) characterizing the structure (e.g., size, crystallinity, surface chemical composition, etc. ) of the stable products of Hg, S and Fe co-precipitation; 2) quantifying the methylation potential of these products in relation to their structural characteristics; 3) assessing how environmental conditions (e.g., redox oscillation, pH, salinity, NOM, Fe, S and Hg content) affect the structure and methylation potential of these materials.

5.2.3 Implications for MeHg photodegradation

The mechanism of MeHg sunlight degradation proposed in this dissertation

(Chapter 4) challenges the long-standing view that the kinetics of MeHg photodegradation in surface waters is not influenced by water chemistry characteristics other than those affect light intensity and MeHg concentration. We demonstrate for the first time that the speciation of dissolved MeHg is a factor for photodegradation and thiol complexation of MeHg enhances its reactivity toward photodemethylation in natural waters. Our results explain the field observation that MeHg photodegradation occurs to a lesser extent in sea water, where CH 3HgCl complexes dominate MeHg speciation, while MeHg photodegradation is relatively rapid in freshwater lakes where

MeHg-thiol complexes control MeHg speciation(191).

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Our study also demonstrates that the role of DOM in MeHg sunlight degradation is of critical importance because DOM is the photosensitizer compound that generates reactive intermediates upon sunlight irradiation. Moreover, DOM enhances the reactivity of MeHg via thiol-complexation. Suwannee river humic acid (SRHA) was selected as a model DOM in this dissertation because it has been well characterized(113) and extensively used in studies of mercury-DOM interaction(18, 89, 114). However,

DOM is highly heterogeneous and the structure varies widely between differing aquatic systems. Depending on the source and chemical composition, DOM may behave very differently as a photosensitizer and as a metal binding ligand. For example, a variety of

• • reactive intermediates(164, 169, 171, 175, 177), including 1O2, OH, 3DOM * O2 - and H2O2, are found to be generated from photosensitization of some types of DOM, while other types of DOM can serve as a scavenger of photo-generated reactive intermediates, such

• as OH(139, 149, 150). Also, MeHg binding capacities vary between different types of humic substances. Therefore, we suggest that a diverse group of DOM be examined further for their role in MeHg photodegradation.

In addition, the quantifiable properties of DOM that possibly determine its ability to complex MeHg and generating reactive intermediates should be characterized and related to the kinetics of photo-induced MeHg decomposition. These properties may include reduced sulfur content(187, 287), molecular weight(157), UV absorbance(186) and fluorescence excitation-emission matrices(185) as suggested by

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other photodegradation studies, and should be incorporated into the MeHg photodegradation model together with the parameters that have already been discussed in the previous research (e.g., temperature, solar radiation, light penetration and pH).

5.2.4 Photochemistry in the aquatic environment

Another implication of our MeHg photodegradation study is that the role of 1O2 in sunlight-induced degradation may have been long underestimated. In the previous studies, a hydrophilic probe (furfuryl alcohol, FFA) has been used to quantify the

abundance of 1O2 in water where H 2O molecules are the primary quenchers of 1O2(151,

161). As a result, this approach has neglected the notion that 1O2 concentration may be unevenly distributed in DOM-containing water and that hydrophobic regions of DOM macromolecules may act as sites or regions for elevated 1O2 concentrations (relative to the bulk water)(169). With this underestimate of 1O2 concentration, researchers have presumed that 1O2 is not significant for photodegradation of compounds Other reactive

• intermediates, particularly OH, are presumed to be more important. Our study is the first to provide an example where the heterogeneous distribution of 1O2 in DOM is critical for photochemical degradation kinetics.

The effect of the heterogeneous distribution of 1O2 around DOM is not limited to the fate of MeHg but may also be relevant to photodegradation of other compounds that associate with DOM. According to Latch and McNeill, the 1O2 concentration at the origin of generation is roughly 1000-fold higher than in bulk solution(169, 251). Even if a

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substance that is only 10% associated with DOM, the contribution of 1O2 degradation would be underestimated by 100 times using the 1O2 concentration measured with hydrophilic probe. This could be the case with photodegradation of a variety of substances that associate with DOM through hydrophobic interaction, electrostatic interaction or ligand complexation.

In fact, the binding with DOM appeared to control the kinetics of photodegradation of halogenated organic contaminants(288-291) and nitroxides(292) in recent studies. In particular, intra-DOM 1O2 reaction has been demonstrated to be the major pathway of sunlight degradation of a polychlorinated biphenyl compound(288).

Additionally, association with DOM may also enhance the photoinactivation of pathogenic microorganisms by bringing them in close proximity to high concentrations of 1O2. This has been proven to be true in a photodegradation study using MS2 coliphase as a surrogate for enteric viruses(293). Hence, intra-DOM 1O2 reactions may be important photochemical mechanisms for a variety of substances in the environment.

Future research should investigate the kinetics of this process by utilizing proper 1O2 probes that accurately mimic the interactions between DOM and the compound of interest . The potential application of intra-DOM 1O2 reactions for water and wastewater treatment should also be assessed.

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5.2.5 Predictive model for methylmercury

The prevailing theme throughout this dissertation was the role of chemical speciation (of mercury) for reactivity in biological and photochemical transformation processes. The knowledge gained from this research will significantly improve our understanding of how MeHg is generated and degraded in aquatic ecosystems. The results from this dissertation and proposed future research will contribute toward establishing a quantitative model that link the kinetics of MeHg formation and degradation to environmental conditions. This model will include speciation-dependent rate constants and measurable environmental parameters that are known to determine mercury speciation. This model could enable an accurate prediction of MeHg accumulation in response to ecosystem alterations, such as accidental mercury contamination (e.g., oil spill, coal ash spill, acid mine drainage inflow, municipal wastewater inflow, etc. ), acid deposition, wetland restoration and sediment remediation.

This information could also facilitate the development of effective cleanup strategies and proper policies for the regulation of mercury sources, and ultimately reducing the human exposure to MeHg.

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Appendix A Supporting information for Chapter 2

SI Materials and Methods

Methylmercury Degradation Bioassay. The cultures of D. propionicus 1pr3 were pre- grown in a fermentative medium and incubated until exponential growth phase prior to dosing with methylmercury chloride (MeHgCl). All methylmercury degradation bioassays were conducted under the same conditions as in the mercury methylation bioassays. Abiotic control consisted of uninoculated media amended with MeHgCl.

Characterization of HgS Particles. The average hydrodynamic diameter of HgS nanoparticles and microparticles were analyzed by light-intensity weighted dynamic light scattering (DLS) (Malvern Zetasizer NS). The diameters of the monomers within the aggregates were analyzed by transmission electron microscopy (TEM). Samples for

DLS were directly quantified in their respective stock suspensions. Samples for TEM were prepared by depositing droplets of the particle stocks on a carbon-coated copper grid (200 mesh), wiping the excess liquid with a lint-free tissue, and allowing the grid to air dry under protective cover. Micro- and nano-HgS particles (aged for 16 h or 1 week) were examined by a FEI Technai TEM at 200 keV and characterized with selected area electron diffraction (SAED) to assess crystal structure of the particles (Figure A4a-f).

Images of HgS nanoparticles (aged for 3 days) were captured on a Hitachi HF2000 TEM

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operating at 200 keV and analyzed for elemental composition by energy dispersive X- ray (EDX) spectroscopy (Oxford Inca EDX system) (Figure A4g and h).

For X-ray diffraction (XRD) analysis, the HgS nanoparticle stocks were filtered through 0.025 μm pore size mixed cellulose ester membranes (Millipore). The material that deposited on the filter membrane was analyzed for crystallographic structure by synchrotron XRD performed at the Stanford Synchrotron Radiation Laboratory (SSRL)

BL 11-3. Samples were placed in the path of a 12,700 eV beam measuring 0.150 mm 2 and the diffraction pattern was captured using a MAR 345 detector set at a distance of 149.9 mm away from the sample, as determined through calibration with a LaB 6 standard. The diffraction image was integrated and peak matched in a manner similar to Reinsch et al.

(2010) (294). The resulting spectra indicated that the commercial HgS microparticles consisted of a mixture of metacinnabar and cinnabar (approximately 50-50 proportion based on Rietveld analysis of XRD data in Figure A5c). The spectra for the humic-HgS nanoparticles indicated metacinnabar-like structure, but peaks were broader and less defined than the micro-HgS spectrum, signifying that nano-HgS consisted of smaller crystallite size and less crystallinity (Figure A5c).

The HgS particles that deposited on the 0.025 μm mixed cellulose ester filters

(Millipore) were also analyzed for elemental composition by X-ray photoelectron spectrometry (XPS) (Figure A5a and b). A PHI VersaProbe Scanning XPS Microprobe was used for these measurements. XPS data were calibrated using the binding energy of

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C(1s) (285.00 eV) as the internal standard. In the XPS spectra, binding energy and peak width of the Hg(4f) and S(2p) transitions were comparable for nano-HgS and micro-HgS, indicating similar local structure.

Centrifugation and Ultracentrifugation of HgS-Amended Cultures. In mercury methylation bioassays, total mercury in D. propionicus 1pr3 cultures was also fractionated using centrifugation and ultracentrifugation (Figure A9). HgS-treated bacteria were first separated from the bulk medium by centrifugation at 6,700 g for 5 min. The pellets (containing mercury associated with bacterial cells and/or large particle aggregates) were washed with 10 mM PBS (pH 7.4) by centrifugation and resuspended in nanopure-filtered water (>18 MΩ-cm). The cells within the washed pellets were then lysed through four freeze-thaw cycles plus 15-min sonication (90 W). This procedure was adequate to fully compromise the cell membrane integrity as determined from the

LIVE/DEAD® viability assay (Invitrogen). The mercury released from the lysed cells was collected from the supernatant after centrifugation at 10,800 g for 30 min and the cell debris- and/or large particle-associated mercury was collected from the pellets. Total mercury in the supernatant of the centrifuged (6,700 g, 5 min) sample was further fractionated by ultracentrifugation at 370,000 g for 2 h. After ultracentrifugation, we considered the mercury remaining in the supernatant to be nominally dissolved, while mercury in the pellet after ultracentrifugation consisted of nanoparticulate/colloidal mercury. The applicability of this procedure for separating dissolved, nanoparticulate

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and microparticulate mercury was demonstrated by (ultra)centrifugation of bacteria-free media that were treated by these different forms of mercury (Figure A9d).

Figure A 1: Net MeHg production in D. propionicus 1pr3 cultures exposed to different forms of mercuric sulfides.

Test cultures were exposed to 5 nM dissolved Hg(NO 3)2 and Na 2S, 5 nM humic-HgS nanoparticles (aged for 16 h), and 20 nM HgS microparticles during (a) 1-day time course experiments and (b) 10-day time course experiments. Abiotic controls were uninoculated medium solutions amended with 5 nM dissolved Hg(NO 3)2 and Na2S. The error bars represent ±1 s.d. for duplicate samples.

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Figure A 2: Bacterial growth in SRB cultures exposed to different forms of mercuric sulfides.

Total protein content and optical density (OD 660 ) of (a) D. propionicus 1pr3 cultures and (b) D. desulfuricans ND132 cultures.

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Figure A 3: Intensity-weighted size distribution of HgS particles measured by dynamic light scattering. (a ) HgS nanoparticles (aged for 16 h); (b) HgS nanoparticles (aged for 3 days); (c) HgS nanoparticles (aged for 1 week); (d) Commercial HgS microparticles.

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a. b.

c. d.

e. f.

g. h.

Figure A 4: TEM and SAED patterns of HgS nanoparticles. Aged for 16 h (a,b), HgS nanoparticles aged for 1 week (c,d), and commercial HgS microparticles (e,f). The EDX spectra for HgS nanoparticles aged for 3 days (g) is shown in (h). EDX spectrum indicated the presence of Hg and S in the nano-HgS sample. In SAED patterns, defined rings indicate long range crystal structure.

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a. b.

Figure A 5: X-ray photoelectron spectroscopy (XPS) and X-ray diffraction (XRD) results of HgS particles. XPS data corresponding to (a) Hg(4f) and (b) S(2p) electrons for the HgS nanoparticles (aged for 16 h, 3 days or 1 week) and HgS microparticles used in methylation experiments. XPS spectra were also collected for pure minerals of metacinnabar and cinnabar. (c) XRD spectra of the same samples. Dotted lines correspond to expected peak positions for pure cinnabar and pure metacinnabar. The spectra for humic-HgS nanoparticles indicated metacinnabar-like structure. HgS microparticles consisted of a mixture of metacinnabar and cinnabar. 138

1000

800 MeHg in 1pr3 culture

600 Total Hg in 1pr3 culture MeHg in abiotic control 400 Total Hg in abiotic control

200 Mercury Mercury concentration (pM) 0 0 10 20 30 40 50 Incubation time (h) Figure A 6: MeHg degradation in D. propionicus 1pr3 cultures. Test cultures were exposed to 1 nM methylmercury chloride (MeHgCl) during 2-day time course experiments. Abiotic controls were uninoculated medium solutions amended with 1 nM MeHgCl. The error bars represent ±1 s.d. for duplicate samples.

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Figure A 7: TEM images and EDX spectra of D. propionicus 1pr3 cultures.

Cultures were exposed to (a) 1 nM dissolved Hg(NO 3)2 and Na 2S, (b) 1 nM humic-HgS nanoparticles (aged for 16 h), and (c) 6 nM HgS microparticles. Test cultures for TEM imaging were collected 14 h after exposure to HgS. Elemental composition of the particles around and/or on the bacterial cells (in black circles) was determined by EDX. The Cu, C, O and Si peaks are from the sample grid. Hg-containing particles were not observed in the cultures exposed to dissolved and nanoparticulate mercury; however, other nanoparticles comprising of elements in the culture media were present, highlighting the heterogeneous nature of the bacteria cultures.

140

80 Dissolved Hg+S 60 HgS nanoparticles HgS microparticles 40 Autoclaved control Abiotic control 20

0 Methylmercury Methylmercury concentration (pM) 0 5 10 15 20

Incubation time (h) Figure A 8: Net MeHg production in D. propionicus 1pr3 cultures exposed to different forms of mercuric sulfides.

Test cultures were exposed to 1 nM dissolved Hg(NO 3)2 and Na 2S, 1 nM humic-HgS nanoparticles (aged for 16 h), and 6 nM HgS microparticles. The error bars represent ±1 s.d. for duplicate samples. Results of Hg fractionation by filtration and (ultra)centrifugation of these samples are shown in Figure 2.2 and Figure A9, respectively.

141

a. b.

c. d.

Figure A 9: Centrifugation and ultracentrifugation of mercury-amended D. propionicus 1pr3 cultures. Cultures were treated with different Hg-sulfide species, including (a) 1 nM dissolved

Hg(NO 3)2 and Na 2S, (b) 1 nM humic-HgS nanoparticles (aged for 16 h), and (c) 6 nM HgS microparticles. Cultures were incubated at room temperature for up to 1 day after mercury addition. At each time point, cultures were first centrifuged at 6,700 g for 5 min. The pellets were lysed through freeze-thaw cycles plus sonication and then centrifuged at 10,800 g for 30 min. The supernatant (after 6,700 g for 5 min) was ultracentrifuged at 370,000 g for 2 h. (d) Bacteria-free media were amended with 1 nM

Hg(NO 3)2, 1 nM humic-HgS nanoparticles (aged for 16 h) or 6 nM HgS microparticles, and (ultra)centrifuged immediately (less than 10 min) after mercury addition. The error bars represent ±1 s.d. for duplicate samples. 142

Appendix B Supporting information for Chapter 3

Table B 1: Stability constants utilized to calculate saturation indices for HgS (s)

(metacinnabar), FeS (s) (mackinawite) and FeS 2(s) (pyrite).

log K (I = 0 M, 25°C) Reference H2S ⇔ HS - + H + -7.02 (231) HS- ⇔ S 2- + H + -17.3 (231)

HgS (s) + H + ⇔ Hg 2+ + HS - log Ks0 = -38 ± 2 (103)

Hg 2+ + HS - ⇔ HgSH + 30.2 (12) Hg 2+ + 2HS - ⇔ Hg(SH) 20 37.7 (295) Hg 2+ + 2HS - ⇔ HgHS 2- + H + 31.5 (295) Hg 2+ + 2HS - ⇔ HgS 22- + 2H + 23.2 (295)

Hg 2+ + RS 2- ⇔ Hg(RS 2) log KDOM = 28.7 (aquatic humic) (296) RS 22- + H + ⇔ RS 2H- 8.4 (297) RS 2H- + H + ⇔ RS 2H2 8.4 (297)

Hg 2+ + H 2O ⇔ HgOH + + H + -3.4 (103) Hg 2+ + 2H 2O ⇔ Hg(OH) 20 + 2H + -6.2 (103) Hg 2+ + 3H 2O ⇔ Hg(OH) 3- + 3H + -21.1 (103)

Hg 2+ + Cl - ⇔ HgCl + 7.3 (103) Hg 2+ + 2Cl - ⇔ Hg(Cl) 20 14.0 (103) Hg 2+ + 3Cl - ⇔ Hg(Cl) 3- 15.0 (103)

Hg 2+ + Cl - + H 2O ⇔ HgOHCl 0 + H + 4.2 (103)

Fe 2+ + HS - ⇔ FeS (s), mackinawite + H + 3.6 (231) Fe 2+ + HS - + S 0(s) ⇔ FeS 2(s), pyrite + H + 14.2 (283)

Fe 2+ + HS - ⇔ Fe(HS) + 4.34 (298) Fe 2+ + 2HS - ⇔ Fe(HS) 2(aq) 8.95 (231) Fe 2+ + 3HS - ⇔ Fe(HS) 3- 10.99 (231)

Fe 2+ + H 2O ⇔ FeOH + + H + -9.40 (231) Fe 2+ + 2H 2O ⇔ Fe(OH) 2(aq) + 2H + -20.49 (231) Fe 2+ + 3H 2O ⇔ Fe(OH) 3- + 3H + -28.99 (231)

Fe 2+ + Cl - ⇔ FeCl + -0.20 (231) 143

a. b.

c. d.

Figure B 1: Total organic carbon (TOC) concentrations in water samples withdrawn -1 from the sediment slurries exposed to 2 nmol g (dw) dissolved Hg(NO 3)2 and Na 2S, HgS nanoparticles or HgS microparticles. The slurries were prepared with sediments from site BG30 (a and c) and LSB129S (b and d). In Experiment 1 (a and b), slurries consisted of 20 g wet sediment and 140 mL surface water, and incubated for 2.5 days prior to Hg addition. In Experiment 2 (c and d), slurries consisted of 50 g wet sediment, 120 mL surface water and 10 mM sodium pyruvate, and incubated for 7.5 days prior to Hg addition. The error bars represent 1 s.d. of replicate samples (n=2-3). Single replicate of slurries were incubated for the control.

144

Figure B 2: Sulfate reduction in slurry microcosms exposed to 2 nmol g -1 (dw) dissolved Hg(NO 3)2 and Na 2S, HgS nanoparticles or HgS microparticles in Experiment 2. The slurries were prepared with sediments from site (a) BG30 and (b) LSB129 S. The error bars represent 1 s.d. for duplicate samples in the test groups. Single replicate of slurries were incubated for the controls. After 2.2-day incubation, sulfate was only detected in the samples collected from molybdate-treated and autoclaved BG30 slurries (a).

145

Sulfate reduction 700 mol) AVS production µ 600

500

400

300

200

100

0 Sulfate reduction or AVS production ( production AVS or reduction Sulfate Dissolved Hg+S Nano-HgS Micro-HgS Hg blank

Figure B 3: Sulfate reduction and AVS production in sediment (LSB129S) slurries -1 exposed to 2 nmol g (dw) dissolved Hg(NO 3)2 and Na 2S, HgS nanoparticles or HgS microparticles (Experiment 2). The sulfate reduction and AVS production were calculated from the difference in pore water sulfate and total AVS concentrations between day 0 and day 7 after mercury addition to the slurries, respectively. The error bars represent 1 s.d. for duplicate samples in the test groups. Single replicate of slurries were incubated for the Hg blank control.

146

Appendix C Supporting information for Chapter 4

Table C 1: Effects of D 2O, sodium azide, and β-carotene on the steady state 1O2 concentration during UV-A irradiation of simulated water containing Suwannee River humic acid. The samples consisted of 0.5 nM MeHg, 2.8 mg-C/L SRHA and 10 mM phosphate

(pH=7.3) amended with D 2O, NaN 3 or β-carotene.

0.1 mM 98.5% D 2O 0.2 M NaN 3 β-carotene [1 ] [1 ] O2 ss , with amendment O2 ss , no amendment calculated 11.0 0.00328 0.294 [1 ] [1 ] O2 ss , with amendment O2 ss , no amendment measured 1.13 0.138 0.393

The ‘calculated’ steady-state concentration of 1O2 was estimated using a previously published approach (169, 267, 299) that is summarized by the following:

k 1  = f O2  ss k+ k[] Q solv q (11) where kf is the zero-order formation rate constant of 1O2, ksolv is the rate constant between

1O2 and the solvent (either H 2O or D 2O), and kq is the reaction rate between 1O2 and the quencher Q (either NaN 3 or β-carotene). The following rate constants were utilized for the calculation: k=2.5 × 105 s -1 , k=1.9 × 104 s -1 , solv, H2 O solv, D2 O

8 -1 -1 9 -1 -1 k=3.8 (± 0.4) × 10 M s , and kβ =6 × 10 M s (references (169, 267, 299, q, NaN 3 q, - carotene

300)). The value for kf was assumed to be the same with and without the amendment.

The measured steady state 1O2 concentration was quantified from the methylmercury degradation data assuming pseudo-first order kinetics: d[ CH Hg ] 3 = − []1  kCH Hg CH3 Hg O 2  (12) dt 3 ss k = k [1O ] (13) obs CH 3Hg 2 ss

147

where the observed rate constant kobs was calculated from the slope of the linear regression of experimental data plotted as ln [MeHg]/[MeHg] 0 versus time. The rate constant k was assumed to be the constant with and without the amendment. CH3 Hg

Table C 2: Stability constants utilized for methylmercury speciation calculations in MINEQL+ (I= 0 M, 25 °C). All values were obtained from the NIST/Smith and Martell database (301,), unless otherwise noted.

log K + - 0 CH 3Hg + OH ⇔ CH 3HgOH 9.45 + - 0 CH 3Hg + Cl ⇔ CH 3HgCl 5.44 + 2- - CH 3Hg + HPO 4 ⇔ CH 3HgHPO 4 5.12

L2- = mercaptoacetate log K + 2- - CH 3Hg + L ⇔ CH 3HgL 17.3 L2- + H + ⇔ HL - 10.6 - + HL + H ⇔ H 2L 3.6

L3- = glutathione log K + 3- 2- CH 3Hg + L ⇔ CH 3HgL 16.7 2- + - CH 3HgL + H ⇔ CH 3HgHL 9.4 - + CH 3HgHL + H ⇔ CH 3HgH 2L 3.7 L3- + H + ⇔ HL 2- 8.88 2- + - HL + H ⇔ H 2L 17.12 - + H2L + H ⇔ H 3L 20.40 + + H3L + H ⇔ H 4L 22.48

Natural organic matter (NOM) (a) log K + CH 3Hg + NOM ⇔ CH 3Hg-NOM 13.5, 14.0, 14.5

(a) CH 3Hg-NOM binding constants represent the range of values reported by

Hintelmann et al. (272) and Amirabahman et al. (269) for aquatic humic substances that contain binding site density of 0.001-0.24 nmol per mg NOM.

148

b. c. Sea water

Figure C 1: Suwannee River humic acid (SRHA)-induced photodegradation of MeHg. a) UV-A degradation of MeHg dissolved in nanopure water amended with SRHA. b) Degradation of MeHg dissolved in 10 mM phosphate buffer (pH 7.3) amended with SRHA. c) Degradation of MeHg dissolved in seawater (containing 2.8 mg/L organic carbon) and seawater spiked with additional DOC (in the form of Suwannee River humic acid) (pH=7.8). Error bars represent ±1 s.d. for replicate measurements (n=2-3).

149

1 0 0.8

0.6

0.4 [MeHg]/[MeHg] 15 nM MeHg alone 0.2 15 nM MeHg with 4 µM GSH

0 0 2 4 6 8 Reaction time (hr)

b. c.

Figure C 2: Reaction between MeHg and superoxide (O2-).

MeHg was exposed to O2- generated by xanthine-xanthine oxidase (XO) in water containing 10 mM phosphate (pH 7.3), 50 μM XTT, and 15 nM initial MeHg with or without 4 μM GSH. a) Degradation of MeHg was not observed. Error bars represent

±1 s.d. for duplicate measurements. b) The probe compound XTT reacted with O 2- to generate a reaction product with an absorbance at 470 nm. c) GSH was observed to degrade as O 2- was generated in the sample with 4 μM initial GSH. The steady state superoxide concentration in these experiments was approximately 3 ×××10 -5 M (assuming

2nd order rate constant = 8.6 ×××10 4 M-1s-1 between XTT and O 2- (302)). Given this steady state concentration, the 2 nd order rate constant between MeHg and O 2- is less than 0.03 M-1s-1, a rate that is too slow to account for MeHg photodegradation in our experiments (e.g., Figure 4.1).

150

Figure C 3: Recovery of reactants in dark controls corresponding to competition kinetics experiments for 1O2 degradation of CH 3Hg-thiol complexes. a) CH 3Hg-GSH sample consisted of 350 nM MeHg, 100 µµµM GSH, 40 µµµM 2CP in a phosphate buffer (10 mM, pH 7.3) and 15 µµµM rose bengal. b) CH 3Hg-MA sample consisted of 350 nM MeHg, 100 µµµM MA, 40 µµµM 2CP in a phosphate buffer (10 mM, pH 7.3) and 15 µµµM rose bengal. All sample bottles for dark controls were wrapped with foil before placement under UV-A lamps.

151

0.8 0

0.6

No amendments

[MeHg]/[MeHg] 0.4 Eosin-Y 0.1 M NaN3NaN 0.2 3 47.9% D2O 47.9% D 2O

0 0 2 4 6 8 UV-A exposure time (hr)

Figure C 4: MeHg degradation by 1O2 generated by UV-A irradiation of rose bengal or Eosin-Y. UV-A degradation of MeHg in simulated water containing 350 nM MeHg, 200 μM GSH, 60 μM 2CP, 15 μM rose bengal and 10 mM phosphate (pH=7.3). The singlet oxygen enhancer D 2O and quencher NaN 3 were amended to replicate samples. In one sample, 15 μM Eosin-Y instead of rose bengal was used as the 1O2 generator.

152

Seawater 1

0.8 0

0.6

0.4 No amendments [MeHg]/[MeHg] 0.2 100 µM MA 100 µM GSH 0 0 10 20 30 UV-A exposure time (hr)

Figure C 5: Effect of ligand complexation on 1O2-induced degradation of MeHg in seawater.

1O2 was generated by UV-A irradiation (λ=365 nm) of 15 μM rose bengal in seawater amended with 350 nM MeHg and either GSH or MA (pH=7.4). The residual organic carbon in the seawater matrix was 2.8 mg-C/L. Error bars represent ±1 s.d. for replicate measurements (n=2-3).

153

a. b.

Figure C 6: Simultaneous degradation of methylmercury and thiols during reactions ••• with 1O2 and OH. a) Degradation of glutathione (GSH) and MeHg by 1O2. b) Degradation of mercaptoacetate (MA) and MeHg by 1O2. Singlet oxygen was generated by UV-A sensitization of rose bengal (15 µµµM) in samples containing 0.35 µµµM MeHg, a thiol (either 75 µµµM GSH or 100 µµµM MA) and 40 µµµM 2CP dissolved in a phosphate buffer (10 ••• mM, pH=7.3). c) Degradation of GSH and MeHg by OH generated from UV-C ( λλλ=254 nm) irradiation of H 2O2 (0.18 mM). Samples initially contained 0.35 µµµM MeHg, 12.5 µµµM GSH and 4 µµµM NB dissolved in phosphate buffer (1 mM, pH=7.4). In all cases,

154

thiol concentrations exceeded MeHg concentration during the entire duration of the photodegradation reactions.

0.8 0 0.6 C/C

0.4 MeHg GSH 0.2 NB

0 0 20 40 60 Reaction time (mins)

Figure C 7: Recovery of reactants for dark experiments with hydrogen peroxide. The samples consisted of 2 µµµM MeHg, 10 µµµM glutathione (GSH), 4 µµµM nitrobenzene

(NB), and 1.5 mM H2O2 dissolved in 1 mM phosphate buffer (pH=7.4).

155

Figure C 8: MeHg speciation calculated for a range of chloride concentrations. Speciation for water containing 10 -12 M total MeHg, 0.001 M to 0.7 M Cl -, 10 -10 mol NOM binding site per L (corresponding to 1 mg/L NOM), and pH 7. Calculations were performed with MINEQL+ using constants in Table C1 and assuming three different values for the CH 3Hg-NOM binding constant K : a) 10 13.5 ; b) 10 14.0 ; c) CH3 Hg- NOM

10 14.5 . Seawater contains ~0.5 M Cl -. Thus, CH 3HgCl complexes are expected to dominate MeHg speciation.

156

References

1. Mergler D, Anderson HA, Chan LHM, Mahaffey KR, Murray M, Sakamoto M, et al. Methylmercury exposure and health effects in humans: A worldwide concern. Ambio. 2007;36(1):3-11.

2. Mahaffey KR, Clickner RP, Bodurow CC. Blood organic mercury and dietary mercury intake: National Health and Nutrition Examination Survey, 1999 and 2000. Environmental Health Perspectives. 2004;112(5):562-70.

3. Grandjean P, Satoh H, Murata K, Eto K. Adverse Effects of Methyl-mercury: Environmental Health Research Implications. Environmental Health Perspectives. 2010;118(8):1137-45.

4. Hrenchuk LE, Blanchfield PJ, Paterson MJ, Hintelmann HH. Dietary and Waterborne Mercury Accumulation by Yellow Perch: A Field Experiment. Environmental Science & Technology. 2012;46(1):509-16.

5. Gehrke GE, Blum JD, Slotton DG, Greenfield BK. Mercury Isotopes Link Mercury in San Francisco Bay Forage Fish to Surface Sediments. Environmental Science & Technology. 2011;45(4):1264-70.

6. Engstrom DR. Fish respond when the mercury rises. Proceedings of the National Academy of Sciences of the United States of America. 2007;104(42):16394-5.

7. Benoit JM, Gilmour CC, Heyes A, Mason RP, Miller CL. Geochemical and biological controls over methylmercury production and degradation in aquatic ecosystems. In: Cai Y, Braids OC, editors. Biogeochemistry of Environmentally Important Trace Elements2003. p. 262-97.

8. Hollweg TA, Gilmour CC, Mason RP. Mercury and methylmercury cycling in sediments of the mid-Atlantic continental shelf and slope. Limnology and Oceanography. 2010;55(6):2703-22.

9. Sellers P, Kelly CA, Rudd JWM, MacHutchon AR. Photodegradation of methylmercury in lakes. Nature. 1996;380:694-7.

10. Sellers P, Kelly CA, Rudd JWM. Fluxes of methylmercury to the water column of a drainage lake: The relative importance of internal and external sources. Limnol Oceanogr. 2001;46(3):623-31.

157

11. Compeau GC, Bartha R. Sulfate-reducing bacteria: Principal methylators of mercury in anoxic estuarine sediment Applied and environmental microbiology. 1985;50:498-502.

12. Benoit JM, Gilmour CC, Mason RP, Heyes A. Sulfide controls on mercury speciation and bioavailability to methylating bacteria in sediment pore waters. Environ Sci Technol. 1999;33(6):951-7.

13. Drott A, Lambertsson L, Bjorn E, Skyllberg U. Importance of dissolved neutral mercury sulfides for methyl mercury production in contaminated sediments. Environmental Science & Technology. 2007;41(7):2270-6.

14. Benoit JM, Gilmour CC, Mason RP. The influence of sulfide on solid phase mercury bioavailability for methylation by pure cultures of Desulfobulbus propionicus (1pr3). Environmental Science & Technology. 2001;35(1):127-32.

15. Mehrotra AS, Horne AJ, Sedlak DL. Reduction of net mercury methylation by iron in Desulfobulbus propionicus (1pr3) cultures: implications for engineered wetlands. Environmental Science & Technology. 2003;37(13):3018-23.

16. Schaefer JK, Morel FMM. High methylation rates of mercury bound to cysteine by Geobacter sulfurreducens. Nature Geoscience. 2009;2(2):123-6.

17. Schaefer JK, Rocks SS, Zheng W, Liang LY, Gu BH, Morel FMM. Active transport, substrate specificity, and methylation of Hg(II) in anaerobic bacteria. Proceedings of the National Academy of Sciences of the United States of America. 2011;108(21):8714-9.

18. Deonarine A, Hsu-Kim H. Precipitation of mercuric sulfide nanoparticles in NOM-containing water: implications for the natural environment. Environ Sci Technol. 2009;43(7):2368-73.

19. Tossell JA. Theoretical study of the photodecomposition of methyl Hg complexes. Journal of Physical Chemistry A. 1998;102(20):3587-91.

20. Lehnherr I, St Louis VL. Importance of ultraviolet radiation in the photodemethylation of methylmercury in freshwater ecosystems. Environ Sci Technol. 2009;43(15):5692-998.

21. Hammerschmidt CR, Fitzgerald WF. Photodecomposition of methylmercury in an arctic Alaskan lake. Environmental Science & Technology. 2006;40(4):1212-6.

158

22. Whalin L, Kim EH, Mason R. Factors influencing the oxidation, reduction, methylation and demethylation of mercury species in coastal waters. Marine Chemistry. 2007;107(3):278-94.

23. Ni B, Kramer JR, Bell RA, Werstiuk NH. Protonolysis of the Hg-C bond of chloromethylmercury and dimethylmercury. A DFT and QTAIM study. Journal of Physical Chemistry A. 2006;110(30):9451-8.

24. Lindberg S, Bullock R, Ebinghaus R, Engstrom D, Feng XB, Fitzgerald W, et al. A synthesis of progress and uncertainties in attributing the sources of mercury in deposition. Ambio. 2007;36(1):19-32.

25. Selin NE. Science and strategies to reduce mercury risks: a critical review. Journal of Environmental Monitoring. 2011;13(9):2389-99.

26. Morel FMM, Kraepiel AML, Amyot M. The chemical cycle and bioaccumulation of mercury. Annual review of ecology and systematics. 1998;29:543-66.

27. Simmons-Willis TA, Koh AS, Clarkson TW, Ballatori N. Transport of a neurotoxicant by molecular mimicry: the methylmercury-L-cysteine complex is a substrate for human L-type large neutral amino acid transporter (LAT) 1 and LAT2. Biochem J. 2002;367:239-46.

28. Yin ZB, Jiang HY, Syversen T, Rocha JBT, Farina M, Aschner M. The methylmercury-L-cysteine conjugate is a substrate for the L-type large neutral amino acid transporter. Journal of Neurochemistry. 2008;107(4):1083-90.

29. Kerper LE, Ballatori N, Clarkson TW. Methylmercury transport across the blood- brain-barrier by an amino-acid carrier. American Journal of Physiology. 1992;262(5):R761-R5.

30. Kajiwara Y, Yasutake A, Adachi T, Hirayama K. Methylmercury transport across the placenta via neutral amino acid carrier. Archives of Toxicology. 1996;70(5):310-4.

31. Clifton JC. Mercury exposure and public health. Pediatr Clin N Am. 2007;54(2):237-+.

32. Farina M, Rocha JBT, Aschner M. Mechanisms of methylmercury-induced neurotoxicity: Evidence from experimental studies. Life Sci. 2011;89(15-16):555-63.

33. Farina M, Aschner M, Rocha JBT. Oxidative stress in MeHg-induced neurotoxicity. Toxicology and Applied Pharmacology. 2011;256(3):405-17.

159

34. Clarkson TW. The toxicology of mercury. Critical Reviews in Clinical Laboratory Sciences. 1997;34(4):369-403.

35. Roman HA, Walsh TL, Coull BA, Dewailly E, Guallar E, Hattis D, et al. Evaluation of the Cardiovascular Effects of Methylmercury Exposures: Current Evidence Supports Development of a Dose-Response Function for Regulatory Benefits Analysis. Environmental Health Perspectives. 2011;119(5):607-14.

36. Clarkson TW, Magos L, Myers GJ. The toxicology of mercury - Current exposures and clinical manifestations. N Engl J Med. 2003;349(18):1731-7.

37. Renner R. Mercury woes appear to grow. Environmental Science & Technology. 2004;38(8):144A-A.

38. Hight SC, Cheng J. Determination of methylmercury and estimation of total mercury in seafood using high performance liquid chromatography (HPLC) and inductively coupled plasma-mass spectrometry (ICP-MS): Method development and validation. Anal Chim Acta. 2006;567(2):160-72.

39. Gong YS, Krabbenhoft DP, Ren LP, Egelandsdal B, Richards MP. Mercury Distribution and Lipid Oxidation in Fish Muscle: Effects of Washing and Isoelectric Protein Precipitation. J Agric Food Chem. 2011;59(20):11050-7.

40. Krystek P, Ritsema R. Mercury speciation in thawed out and refrozen fish samples by gas chromatography coupled to inductively coupled plasma mass spectrometry and atomic fluorescence spectroscopy. Anal Bioanal Chem. 2005;381(2):354-9.

41. Wiener JG, Krabbenhoft DP, Heinz GH, Scheuhammer AM. Ecotoxicology of mercury. In: Hoffman DJ, Rattner BA, G.A. Burton J, Cairns J, editors. Handbook of Ecotoxicology (2nd ed). Boca Raton, FL: Lewis Publishers; 2003. p. 409-63

42. Clarkson TW. The three modern faces of mercury. Environmental Health Perspectives. 2002;110:11-23.

43. U.S. EPA (U.S. Environmental Protection Agency). 2011. Biennial National Listing of Fish Advisories. 2010. EPA-820-F-11-014. Available: http://water.epa.gov/scitech/swguidance/fishshellfish/fishadvisories/upload/tech nical_factsheet_2010.pdf [accessed 26 Feb 2012].

44. Pirrone N, Cinnirella S, Feng X, Finkelman RB, Friedli HR, Leaner J, et al. Global mercury emissions to the atmosphere from anthropogenic and natural sources. Atmos Chem Phys. 2010;10(13):5951-64. 160

45. Yee D, McKee LJ, Oram JJ. A regional mass balance of methylmercury in San Francisco Bay, California, USA. Environmental Toxicology and Chemistry. 2011;30(1):88-96.

46. Jensen S, Jernelov A. Biological methylation of mercury in aquatic organisms. Nature. 1969;223(5207):753-4.

47. Siciliano SD, O'Driscoll NJ, Tordon R, Hill J, Beauchamp S, Lean DRS. Abiotic production of methylmercury by solar radiation. Environmental Science & Technology. 2005;39(4):1071-7.

48. James H W. Review of possible paths for abiotic methylation of mercury(II) in the aquatic environment. Chemosphere. 1993;26(11):2063-77.

49. Lehnherr I, St Louis VL, Hintelmann H, Kirk JL. Methylation of inorganic mercury in polar marine waters. Nature Geoscience. 2011;4(5):298-302.

50. Hammerschmidt CR, Fitzgerald WF, Lamborg CH, Balcom PH, Tseng CM. Biogeochemical cycling of methylmercury in lakes and tundra watersheds of Arctic Alaska. Environmental Science & Technology. 2006;40(4):1204-11.

51. Bridou R, Monperrus M, Gonzalez PR, Guyoneaud R, Amouroux D. Simultaneous determination of mercury methylation and demethylation capacities of various sulfate-reducing bacteria using species-specific isotopic tracers. Environ Toxicol Chem. 2011;30(2):337-44.

52. Schaefer JK, Yagi J, Reinfelder JR, Cardona T, Ellickson KM, Tel-Or S, et al. Role of the bacterial organomercury lyase (MerB) in controlling methylmercury accumulation in mercury-contaminated natural waters. Environmental Science & Technology. 2004;38(16):4304-11.

53. Oremland RS, Miller LG, Dowdle P, Connell T, Barkay T. Methylmercury oxidative-degradation potentials in contaminated and pristine sediments of the Carson River, Nevada. Applied and Environmental Microbiology. 1995;61(7):2745-53.

54. Rolfhus KR, Sakamoto HE, Cleckner LB, Stoor RW, Babiarz CL, Back RC, et al. Distribution and fluxes of total and methylmercury in Lake Superior. Environmental Science & Technology. 2003;37(5):865-72.

55. Merritt KA, Amirbahman A. Methylmercury cycling in estuarine sediment pore waters (Penobscot River estuary, Maine, USA). Limnology and Oceanography. 2008;53(3):1064-75. 161

56. Kim EH, Mason RP, Porter ET, Soulen HL. The impact of resuspension on sediment mercury dynamics, and methylmercury production and fate: A mesocosm study. Marine Chemistry. 2006;102(3-4):300-15.

57. Abi-Ghanem C, Nakhle K, Khalaf G, Cossa D. Mercury Distribution and Methylmercury Mobility in the Sediments of Three Sites on the Lebanese Coast, Eastern Mediterranean. Archives of Environmental Contamination and Toxicology. 2011;60(3):394-405.

58. Compeau GC, Bartha R. Sulfate-reducing bacteria - principal methylators of mercury in anoxic estuarine sediment. Appl Environ Microbiol 1985;50(2):498- 502.

59. Hamelin S. AM, BarkayT., YWang Y., Planas D.,. Methanogens: Principal Methylators of Mercury in Lake Periphyton. Environmental Science & Technology. 2011;45:7693-700.

60. Fleming EJ, Mack EE, Green PG, Nelson DC. Mercury methylation from unexpected sources: Molybdate-inhibited freshwater sediments and an iron- reducing bacterium. Applied and Environmental Microbiology. 2006;72(1):457-64.

61. Kerin EJ, Gilmour CC, Roden E, Suzuki MT, Coates JD, Mason RP. Mercury methylation by dissimilatory iron-reducing bacteria. Applied and Environmental Microbiology. 2006;72(12):7919-21.

62. Ranchou-Peyruse M, Monperrus M, Bridou R, Duran R, Amouroux D, Salvado JC, et al. Overview of mercury methylation capacities among anaerobic bacteria including representatives of the sulphate-reducers: implications for environmental studies. Geomicrobiol J. 2009;26(1):1-8.

63. King JK, Saunders FM, Lee RF, Jahnke RA. Coupling mercury methylation rates to sulfate reduction rates in marine sediments. Environ Toxicol Chem. 1999;18(7):1362-9.

64. Devereux R, Winfrey MR, Winfrey J, Stahl DA. Depth profile of sulfate-reducing bacterial ribosomal RNA and mercury methylation in an estuarine sediment. Fems Microbiology Ecology. 1996;20(1):23-31.

65. Gilmour CC, Henry EA, Mitchell R. Sulfate stimulation of mercury methylation in fresh-water sediments. Environmental Science & Technology. 1992;26(11):2281-7.

162

66. Barkay T, Wagner-Dobler I. Microbial transformations of mercury: Potentials, challenges, and achievements in controlling mercury toxicity in the environment. In: Laskin AI, Bennett JW, Gadd GM, editors. Advances in Applied Microbiology, Vol 572005. p. 1-52.

67. Barkay T, Kritee K, Boyd E, Geesey G. A thermophilic bacterial origin and subsequent constraints by redox, light and salinity on the evolution of the microbial mercuric reductase. Environmental Microbiology. 2010;12(11):2904-17.

68. Marvin-DiPasquale M, Agee J, McGowan C, Oremland RS, Thomas M, Krabbenhoft D, et al. Methyl-mercury degradation pathways: A comparison among three mercury-impacted ecosystems. Environmental Science & Technology. 2000;34(23):4908-16.

69. Marvin-DiPasquale M, Agee JL. Microbial mercury cycling in sediments of the San Francisco Bay-Delta. Estuaries. 2003;26(6):1517-28.

70. Berman M, Chase T, Bartha R. Carbon flow in mercury biomethylation by Desulfovibrio desulfuricans. Applied and Environmental Microbiology. 1990;56(1):298-300.

71. Choi SC, Chase T, Bartha R. Metabolic pathways leading to mercury methylation in Desulfovibrio desulfuricans LS. Applied and Environmental Microbiology. 1994;60(11):4072-7.

72. Haferburg G. KE. Microbes and metals: interactions in the environment. Journal of Basic Microbiology. 2007;47:453-67.

73. Najera I. LCC, Kohbodi G.A., Jay J.A. Effect of chemical speciation on toxicity of mercury to Escherichia coli biofilms and planktonic cells. Environ Sci Technol. 2005;39:3116-20.

74. Mason RP, Reinfelder JR, Morel FMM. Uptake, toxicity, and trophic transfer of mercury in a coastal diatom. Environ Sci Technol. 1996;30(6):1835-45.

75. Benoit JM, Gilmour CC, Mason RP, Heyes A. Sulfide controls on mercury speciation and bioavailability to methylating bacteria in sediment pore waters. Environ Sci Technol. 1999;33(6):951-7.

76. Jay JA, Murray KJ, Gilmour CC, Mason RP, Morel FMM, Roberts AL, et al. Mercury methylation by Desulfovibrio desulfuricans ND132 in the presence of polysulfides. Applied and Environmental Microbiology. 2002;68(11):5741-5.

163

77. Golding GR, Kelly CA, Sparling R, Loewen PC, Rudd JWM, Barkay T. Evidence for facilitated uptake of Hg(II) by Vibrio anguillarum and Escherichia coli under anaerobic and aerobic conditions. Limnology and Oceanography. 2002;47(4):967- 75.

78. Kelly CA, Rudd JWM, Holoka MH. Effect of pH on mercury uptake by an aquatic bacterium: Implications for Hg cycling. Environmental Science & Technology. 2003;37(13):2941-6.

79. Najera I. LCC, Kohbodi G.A., Jay J.A. Effect of chemical speciation on toxicity of mercury to Escherichia coli biofilms and planktonic cells. Environmental science and technology. 2005;39:3116-20.

80. Gilmour CC, Elias DA, Kucken AM, Brown SD, Palumbo AV, C.W. S, et al. Sulfate-reducing bacterium Desulfovibrio desulfuricans ND132 as a model for understanding bacterial mercury methylation. Applied and Environmental Microbiology. 2011(77):3938-51.

81. Choi SC, Bartha R. Cobalamin-mediated mercury methylation by Desulfovibrio desulfuricans LS. Applied and Environmental Microbiology. 1993;59(1):290-5.

82. Stupperich E, Krautler B. Pseudo vitamin-B12 or 5-hydroxybenzimidazolyl- cobamide are the corrinoids found in methanogenic bacteria. Arch Microbiol. 1988;149(3):268-71.

83. Ekstrom EB, Morel FMM. Cobalt limitation of growth and mercury methylation in sulfate-reducing bacteria. Environmental Science & Technology. 2008;42(1):93- 9.

84. Ekstrom EB, Morel FMM, Benoit JM. Mercury methylation independent of the acetyl-coenzyme a pathway in sulfate-reducing bacteria. Applied and Environmental Microbiology. 2003;69(9):5414-22.

85. Choe KY, Gill GA, Lehman R. Distribution of particulate, colloidal, and dissolved mercury in San Francisco Bay estuary. 1. Total mercury. Limnology and Oceanography. 2003;48(4):1535-46.

86. Babiarz CL, Hurley JP, Hoffmann SR, Andren AW, Shafer MM, Armstrong DE. Partitioning of total mercury and methylmercury to the colloidal phase in freshwaters. Environmental Science & Technology. 2001;35(24):4773-82.

164

87. Lee S, Han S, Gill GA. Estuarine mixing behavior of colloidal organic carbon and colloidal mercury in Galveston Bay, Texas. Journal of Environmental Monitoring. 2011;13(6):1703-8.

88. Ravichandran M, Aiken GR, Ryan JN, Reddy MM. Inhibition of precipitation and aggregation of metacinnabar (mercuric sulfide) by dissolved organic matter isolated from the Florida Everglades. Environmental Science & Technology. 1999;33(9):1418-23.

89. Slowey AJ. Rate of formation and dissolution of mercury sulfide nanoparticles: The dual role of natural organic matter. Geochim Cosmochim Acta. 2010;74(16):4693-708.

90. Gerbig CA, Kim CS, Stegemeier JP, Ryan JN, Aiken GR. Formation of Nanocolloidal Metacinnabar in Mercury-DOM-Sulfide Systems. Environ Sci Technol. 2011;45(21):9180-7.

91. Pesavento M, Alberti G, Biesuz R. Analytical methods for determination of free metal ion concentration, labile species fraction and metal complexation capacity of environmental waters: A review. Analytica Chimica Acta. 2009;631(2):129-41.

92. Hsu H, Sedlak DL. Strong Hg(II) complexation in municipal wastewater effluent and surface waters. Environ Sci Technol. 2003;37(12):2743-9.

93. Han SH, Gill GA. Determination of mercury complexation in coastal and estuarine waters using competitive ligand exchange method. Environmental Science & Technology. 2005;39(17):6607-15.

94. Black FJ, Bruland KW, Flegal AR. Competing ligand exchange-solid phase extraction method for the determination of the complexation of dissolved inorganic mercury(II) in natural waters. Analytica Chimica Acta. 2007;598(2):318- 33.

95. Han SH, Gill GA, Lehman RD, Choe KY. Complexation of mercury by dissolved organic matter in surface waters of Galveston Bay, Texas. Marine Chemistry. 2006;98(2-4):156-66.

96. Hesterberg D, Chou JW, Hutchison KJ, Sayers DE. Bonding of Hg(II) to reduced organic, sulfur in humic acid as affected by S/Hg ratio. Environmental Science & Technology. 2001;35(13):2741-5.

165

97. Nagy KL, Manceau A, Gasper JD, Ryan JN, Aiken GR. Metallothionein-Like Multinuclear Clusters of Mercury(II) and Sulfur in Peat. Environmental Science & Technology. 2011;45(17):7298-306.

98. Xia K, Skyllberg UL, Bleam WF, Bloom PR, Nater EA, Helmke PA. X-ray absorption spectroscopic evidence for the complexation of Hg(II) by reduced sulfur in soil humic substances. Environmental Science & Technology. 1999;33(2):257-61.

99. Hsu-Kim H, Sedlak DL. Similarities between inorganic sulfide and the strong Hg(II) - Complexing ligands in municipal wastewater effluent. Environ Sci Technol. 2005;39(11):4035-41.

100. Mehrotra AS, Sedlak DL. Decrease in net mercury methylation rates following iron amendment to anoxic wetland sediment slurries. Environmental Science & Technology. 2005;39(8):2564-70.

101. Dyrssen D, Wedborg M. The sulfure-mercury(II) system in natural waters. Water Air and Soil Pollution. 1991;56:507-19.

102. Skyllberg U. Competition among thiols and inorganic sulfides and polysulfides for Hg and MeHg in wetland soils and sediments under suboxic conditions: Illumination of controversies and implications for MeHg net production. Journal of Geophysical Research-Biogeosciences. 2008;113.

103. Smith RD, Martell AE. NIST Critical Stability Constants of Metal Complexes Database v. 2.0. 1993.

104. Wolfenden S, Charnock JM, Hilton J, Livens FR, Vaughan DJ. Sulfide species as a sink for mercury in lake sediments. Environmental Science & Technology. 2005;39(17):6644-8.

105. Aiken GR, Hsu-Kim H, Ryan JN. Influence of dissolved organic matter on the environmental fate of metals, nanoparticles, and colloids. Environmental Science & Technology. 2011;45(8):3196-201.

106. Ravichandran M, Aiken GR, Reddy MM, Ryan JN. Enhanced dissolution of cinnabar (mercuric sulfide) by dissolved organic matter isolated from the Florida Everglades. Environmental Science & Technology. 1998;32(21):3305-11.

107. Lambertsson L, Nilsson M. Organic material: The primary control on mercury methylation and ambient methyl mercury concentrations in estuarine sediments. Environmental Science & Technology. 2006;40(6):1822-9. 166

108. Ullrich SM, Tanton TW, Abdrashitova SA. Mercury in the aquatic environment: A review of factors affecting methylation Critical reviews in environmental science and technology. 2001;31(3):241-93.

109. Munthe J, Bodaly RA, Branfireun BA, Driscoll CT, Gilmour CC, Harris R, et al. Recovery of mercury-contaminated fisheries. Ambio. 2007;36(1):33-44.

110. Hammerschmidt CR, Fitzgerald WF, Balcom PH, Visscher PT. Organic matter and sulfide inhibit methylmercury production in sediments of New York/New Jersey Harbor. Marine Chemistry. 2008;109(1-2):165-82.

111. Furutani A, Rudd JWM. Measurement of mercury methylation in lake water and sediment samples. Applied and Environmental Microbiology. 1980;40(4):770-6.

112. Sunderland EM, Gobas F, Branfireun BA, Heyes A. Environmental controls on the speciation and distribution of mercury in coastal sediments. Marine Chemistry. 2006;102(1-2):111-23.

113. Waples JS, Nagy KL, Aiken GR, Ryan JN. Dissolution of cinnabar (HgS) in the presence of natural organic matter. Geochimica et Cosmochimica Acta. 2005;69(6):1575-88.

114. Miller CL, Mason RP, Gilmour CC, Heyes A. Influence of dissolved organic matter on the complexation of mercury under sulfidic conditions. Environmental Toxicology and Chemistry. 2007;26(4):624-33.

115. Deonarine A, Hsu-Kim H. Precipitation of mercuric sulfide nanoparticles in NOM-containing water: Implications for the natural environment. Environ Sci Technol. 2009;43(7):2368-73.

116. Slowey AJ. Rate of formation and dissolution of mercury sulfide nanoparticles: The dual role of natural organic matter. Geochim Cosmochim Acta. 2010;74(16):4693-708.

117. Patty C, Barnett B, Mooney B, Kahn A, Levy S, Liu YJ, et al. Using X-ray microscopy and Hg L-3 XAMES to study Hg binding in the rhizosphere of Spartina cordgrass. Environ Sci Technol. 2009;43(19):7397-402.

118. Barnett MO, Harris LA, Turner RR, Stevenson RJ, Henson TJ, Melton RC, et al. Formation of mercuric sulfide in soil. Environ Sci Technol. 1997;31(11):3037-43.

167

119. Lowry GV, Shaw S, Kim CS, Rytuba JJ, Brown GE. Macroscopic and microscopic observations of particle-facilitated mercury transport from new idria and sulphur bank mercury mine tailings. Environ Sci Technol. 2004;38(19):5101-11.

120. Jeong HY, Klaue B, Blum JD, Hayes KF. Sorption of mercuric ion by synthetic nanocrystalline mackinawite (FeS). Environ Sci Technol. 2007;41(22):7699-705.

121. Hochella MF, Lower SK, Maurice PA, Penn RL, Sahai N, Sparks DL, et al. Nanominerals, mineral nanoparticles, and earth systems. Science. 2008;319(5870):1631-5.

122. Waychunas GA, Zhang HZ. Structure, chemistry, and properties of mineral nanoparticles. Elements. 2008;4(6):381-7.

123. Auffan M, Rose J, Bottero JY, Lowry GV, Jolivet JP, Wiesner MR. Towards a definition of inorganic nanoparticles from an environmental, health and safety perspective. Nature Nanotechnology. 2009;4(10):634-41.

124. Bosch J, Heister K, Hofmann T, Meckenstock RU. Nanosized Iron Oxide Colloids Strongly Enhance Microbial Iron Reduction. Applied and Environmental Microbiology. 2010;76(1):184-9.

125. Bonneville S, Van Cappellen P, Behrends T. Microbial reduction of iron(III) oxyhydroxides: effects of mineral solubility and availability. Chemical Geology. 2004;212(3-4):255-68.

126. Glasauer S, Langley S, Beveridge TJ. Sorption of Fe (hydr)oxides to the surface of Shewanella putrefaciens: Cell-bound fine-grained minerals are not always formed de novo. Applied and Environmental Microbiology. 2001;67(12):5544-50.

127. Bonneville S, Behrends T, Van Cappellen P, Hyacinthe C, Roling WFM. Reduction of Fe(III) colloids by Shewanella putrefaciens: A kinetic model. Geochimica et Cosmochimica Acta. 2006;70(23):5842-54.

128. Grantham MC, Dove PM, DiChristina TJ. Microbially catalyzed dissolution of iron and aluminum oxyhydroxide mineral surface coatings. Geochimica et Cosmochimica Acta. 1997;61(21):4467-77.

129. Liu J, Aruguete DM, Murayama M, Hochella MF. Influence of Size and Aggregation on the Reactivity of an Environmentally and Industrially Relevant Manomaterial (PbS). Environmental Science & Technology. 2009;43(21):8178-83.

168

130. Gondikas AP, Jang EK, Hsu-Kim H. Influence of amino acids cysteine and serine on aggregation kinetics of zinc and mercury sulfide colloids. Journal of Colloid and Interface Science. 2010;347(2):167-71.

131. Mullaugh KM, Luther GW. Spectroscopic determination of the size of cadmium sulfide nanoparticles formed under environmentally relevant conditions. Journal of Environmental Monitoring. 2010;12(4):890-7.

132. Liu J, Aruguete DA, Jinschek JR, Rimstidt JD, Hochella MF. The non-oxidative dissolution of galena nanocrystals: Insights into mineral dissolution rates as a function of grain size, shape, and aggregation state. Geochimica et Cosmochimica Acta. 2008;72(24):5984-96.

133. Cutting RS, Coker VS, Fellowes JW, Lloyd JR, Vaughan DJ. Mineralogical and morphological constraints on the reduction of Fe(III) minerals by Geobacter sulfurreducens. Geochimica et Cosmochimica Acta. 2009;73(14):4004-22.

134. Hotze EM, Phenrat T, Lowry GV. Nanoparticle Aggregation: Challenges to Understanding Transport and Reactivity in the Environment. Journal of Environmental Quality. 2010;39(6):1909-24.

135. Boreen AL, Arnold WA, McNeill K. Photodegradation of pharmaceuticals in the aquatic environment: A review. Aquatic Sciences. 2003;65(4):320-41.

136. Peuravuori J, Pihlaja K. Phototransformations of selected pharmaceuticals under low-energy UVA-vis and powerful UVB-UVA irradiations in aqueous solutions- the role of natural dissolved organic chromophoric material. Analytical and Bioanalytical Chemistry. 2009;394(6):1621-36.

137. Packer JL, Werner JJ, Latch DE, McNeill K, Arnold WA. Photochemical fate of pharmaceuticals in the environment: Naproxen, diclofenac, clofibric acid, and ibuprofen. Aquatic Sciences. 2003;65(4):342-51.

138. Burrows HD, Canle M, Santaballa JA, Steenken S. Reaction pathways and mechanisms of photodegradation of pesticides. Journal of Photochemistry and Photobiology B-Biology. 2002;67(2):71-108.

139. Garbin JR, Milori D, Simoes ML, da Silva WTL, Neto LM. Influence of humic substances on the photolysis of aqueous pesticide residues. Chemosphere. 2007;66(9):1692-8.

169

140. Rayne S, Forest K, Friesen KJ. Mechanistic aspects regarding the direct aqueous environmental photochemistry of phenol and its simple halogenated derivatives. A review. Environment International. 2009;35(2):425-37.

141. Loiselle SA, Bracchini L, Cozar A, Dattilo AM, Tognazzi A, Rossi C. Variability in photobleaching yields and their related impacts on optical conditions in subtropical lakes. Journal of Photochemistry and Photobiology B-Biology. 2009;95(2):129-37.

142. Morris DP, Zagarese H, Williamson CE, Balseiro EG, Hargreaves BR, Modenutti B, et al. The attentuation of solar UV radiation in lakes and the role of dissolved organic carbon. Limnology and Oceanography. 1995;40(8):1381-91.

143. Hines NA, Brezonik PL. Mercury inputs and outputs at a small lake in northern Minnesota. Biogeochemistry. 2007;84(3):265-84.

144. Liu QT, Williams HE. Kinetics and degradation products for direct photolysis of beta-blockers in water. Environmental Science & Technology. 2007;41(3):803-10.

145. Cooper WJ, Zika RG, Petasne RG, Fischer AM. Sunlight-induced photochemistry of humic substances in natural-waters - major reactive species. Acs Symposium Series. 1989;219:333-62.

146. Mill T. Predicting photoreaction rates in surface waters. Chemosphere. 1999;38(6):1379-90.

147. Suda I, Suda M, Hirayama K. Degradation of methyl and ethyl mercury by singlet oxygen generated from sea-water exposed to sunlight or ultraviolet-light. Archives of Toxicology. 1993;67(5):365-8.

148. Chen J, Pehkonen SO, Lin CJ. Degradation of monomethylmercury chloride by hydroxyl radicals in simulated natural waters. Water Research. 2003;37(10):2496- 504.

149. Goldstone JV, Pullin MJ, Bertilsson S, Voelker BM. Reactions of hydroxyl radical with humic substances: Bleaching, mineralization, and production of bioavailable carbon substrates. Environmental Science & Technology. 2002;36(3):364-72.

150. Westerhoff P, Mezyk SP, Cooper WJ, Minakata D. Electron pulse radiolysis determination of hydroxyl radical rate constants with Suwannee river fulvic acid and other dissolved organic matter isolates. Environmental Science & Technology. 2007;41(13):4640-6. 170

151. al Housari F, Vione D, Chiron S, Barbati S. Reactive photoinduced species in estuarine waters. Characterization of hydroxyl radical, singlet oxygen and dissolved organic matter triplet state in natural oxidation processes. Photochemical & Photobiological Sciences. 2010;9(1):78-86.

152. Brezonik PL, Fulkerson-Brekken J. Nitrate-induced photolysis in natural waters: Controls on concentrations of hydroxyl radical photo-intermediates by natural scavenging agents. Environmental Science & Technology. 1998;32(19):3004-10.

153. Allen JM, Lucas S, Allen SK. Formation of hydroxyl radical ((OH)-O-center dot) in illuminated surface waters contaminated with acidic mine drainage. Environmental Toxicology and Chemistry. 1996;15(2):107-13.

154. Punjabi PB, Kabra BV, Pitliya RL, Vaidya VK, Ameta SC. The chemistry of singlet molecular oxygen. Journal of the Indian Chemical Society. 2001;78(4):175-84.

155. DeRosa MC, Crutchley RJ. Photosensitized singlet oxygen and its applications. Coordination Chemistry Reviews. 2002;233:351-71.

156. Zepp RG, Wolfe NL, Baughman GL, Hollis RC. Singlet oxygen in natural waters. Nature. 1977;267(5610):421-3.

157. Sandvik SLH, Bilski P, Pakulski JD, Chignell CF, Coffin RB. Photogeneration of singlet oxygen and free radicals in dissolved organic matter isolated from the Mississippi and Atchafalaya River plumes. Marine Chemistry. 2000;69(1-2):139- 52.

158. Kochany J, Bolton JR. Mechanism of photodegradation of aqueous organic pollutants. 1. EPR spin-trapping technique for the determination of .OH radical rate constants in the photooxidation of chlorophenols following the photolysis of H2O2. Journal of Physical Chemistry. 1991;95(13):5116-20.

159. Minella M, Romeo F, Vione D, Maurino V, Minero C. Low to negligible photoactivity of lake-water matter in the size range from 0.1 to 5 mu m. Chemosphere. 2011;83(11):1480-5.

160. Chin YP, Miller PL, Zeng LK, Cawley K, Weavers LK. Photosensitized degradation of bisphenol a by dissolved organic matter. Environmental Science & Technology. 2004;38(22):5888-94.

161. Vione D, Bagnus D, Maurino V, Minero C. Quantification of singlet oxygen and hydroxyl radicals upon UV irradiation of surface water. Environmental Chemistry Letters. 2010;8(2):193-8. 171

162. Westerhoff P, Anning D. Concentrations and characteristics of organic carbon in surface water in Arizona: influence of urbanization. Journal of Hydrology. 2000;236(3-4):202-22.

163. O'Sullivan DW, Neale PJ, Coffin RB, Boyd TJ, Osburn SL. Photochemical production of hydrogen peroxide and methylhydroperoxide in coastal waters. Marine Chemistry. 2005;97(1-2):14-33.

164. Canonica S. Oxidation of aquatic organic contaminants induced by excited triplet states. Chimia. 2007;61(10):641-4.

165. Chen Y, Hu C, Hu XX, Qu JH. Indirect Photodegradation of Amine Drugs in Aqueous Solution under Simulated Sunlight. Environmental Science & Technology. 2009;43(8):2760-5.

166. Xu HM, Cooper WJ, Jung J, Song WH. Photosensitized degradation of amoxicillin in natural organic matter isolate solutions. Water Research. 2011;45(2):632-8.

167. Gerecke AC, Canonica S, Muller SR, Scharer M, Schwarzenbach RP. Quantification of dissolved natural organic matter (DOM) mediated phototransformation of phenylurea herbicides in lakes. Environmental Science & Technology. 2001;35(19):3915-23.

168. Haag WR, Hoigne J, Gassman E, Braun AM. Singlet oxygen in surface waters. 2. Quantum yields of its production by some natural humic materials as a function of wavelength. Chemosphere. 1984;13(5-6):641-50.

169. Latch DE, McNeill K. Microheterogeneity of singlet oxygen distributions in irradiated humic acid solutions. Science. 2006;311(5768):1743-7.

170. Grandbois M, Latch DE, McNeill K. Microheterogeneous Concentrations of Singlet oxygen in Natural Organic Matter Isolate Solutions. Environ Sci Technol. 2008;42(24):9184-90.

171. Southworth BA, Voelker BM. Hydroxyl radical production via the photo-Fenton reaction in the presence of fulvic acid. Environmental Science & Technology. 2003;37(6):1130-6.

172. Vione D, Falletti G, Maurino V, Minero C, Pelizzetti E, Malandrino M, et al. Sources and sinks of hydroxyl radicals upon irradiation of natural water samples. Environmental Science & Technology. 2006;40(12):3775-81.

172

173. Vermilyea AW, Voelker BM. Photo-Fenton Reaction at Near Neutral pH. Environmental Science & Technology. 2009;43(18):6927-33.

174. Vaughan PP, Blough NV. Photochemical Formation of Hydroxyl Radical by Constituents of Natural Waters. Environ Sci Technol. 1998;32:2947-53.

175. Garg S, Rose AL, Waite TD. Photochemical production of superoxide and hydrogen peroxide from natural organic matter. Geochimica et Cosmochimica Acta. 2011;75(15):4310-20.

176. Draper WM, Crosby DG. The photochemical generation of hydrogen-peroxide in natural-waters. Archives of Environmental Contamination and Toxicology. 1983;12(1):121-6.

177. Scully NM, McQueen DJ, Lean DRS, Cooper WJ. Hydrogen peroxide formation: The interaction of ultraviolet radiation and dissolved organic carbon in lake waters along a 43-75 degrees N gradient. Limnology and Oceanography. 1996;41(3):540-8.

178. Leifer A. The Kinetics of Environmental Aquatic Photochemistry - Theory and Practice; American Chemical Society: Washington, DC, 1988; p 304.

179. Espy R, Pelton E, Opseth A, Kasprisin J, Nienow AM. Photodegradation of the Herbicide Imazethapyr in Aqueous Solution: Effects of Wavelength, pH, and Natural Organic Matter (NOM) and Analysis of Photoproducts. Journal of Agricultural and Food Chemistry. 2011;59(13):7277-85.

180. Canonica S, Laubscher HU. Inhibitory effect of dissolved organic matter on triplet-induced oxidation of aquatic contaminants. Photochemical & Photobiological Sciences. 2008;7(5):547-51.

181. Wenk J, von Gunten U, Canonica S. Effect of Dissolved Organic Matter on the Transformation of Contaminants Induced by Excited Triplet States and the Hydroxyl Radical. Environmental Science & Technology. 2011;45(4):1334-40.

182. Walse SS, Morgan SL, Kong L, Ferry JL. Role of dissolved organic matter, nitrate, and bicarbonate in the photolysis of aqueous fipronil. Environmental Science & Technology. 2004;38(14):3908-15.

183. Grannas AM, Cory RM, Miller PL, Chin YP, McKnight DM. The role of dissolved organic matter in arctic surface waters in the photolysis of hexachlorobenzene and lindane. Journal of Geophysical Research-Biogeosciences. 2012;117.

173

184. Guerard JJ, Miller PL, Trouts TD, Chin YP. The role of fulvic acid composition in the photosensitized degradation of aquatic contaminants. Aquatic Sciences. 2009;71(2):160-9.

185. Coelho C, Guyot G, ter Halle A, Cavani L, Ciavatta C, Richard C. Photoreactivity of humic substances: relationship between fluorescence and singlet oxygen production. Environmental Chemistry Letters. 2011;9(3):447-51.

186. Dalrymple RM, Carfagno AK, Sharpless CM. Correlations between Dissolved Organic Matter Optical Properties and Quantum Yields of Singlet Oxygen and Hydrogen Peroxide. Environmental Science & Technology. 2010;44(15):5824-9.

187. Yoon SJ, Diener LM, Bloom PR, Nater EA, Bleam WF. X-ray absorption studies of CH3Hg+-binding sites in humic substances. Geochimica et Cosmochimica Acta. 2005;69(5):1111-21.

188. Skyllberg U, Qian J, Frech A. Combined XANES and EXAFS study on the bonding of methyl mercury to thiol groups in soil and aquatic organic matter. Physica Scripta. 2005;T115:894-6.

189. Amirbahman A, Reid AL, Haines TA, Kahl JS, Arnold C. Association of methylmercury with dissolved humic acids. Environ Sci Technol. 2002;36:690-5.

190. Hintelmann H, Welbourn PM, Evans RD. Measurement of complexation of methylmercury (II) compounds by freshwater humic substances using equilibrium dialysis. Environ Sci Technol. 1997;31:489-95.

191. Zhang T, Hsu-Kim H. Photolytic degradation of methylmercury enhanced by binding to natural organic ligands. Nature Geoscience. 2010;3(7):473-6.

192. Merritt KA, Amirbahman A. Mercury methylation dynamics in estuarine and coastal marine environments - A critical review. Earth Sci Rev. 2009;96(1-2):54-66.

193. Benoit JM, Gilmour CC, Mason RP. Aspects of bioavailability of mercury for methylation in pure cultures of Desulfobulbus propionicus (1pr3). Appl Environ Microbiol. 2001;67(1):51-8.

194. Drott A, Lambertsson L, Bjorn E, Skyllberg U. Importance of dissolved neutral mercury sulfides for methyl mercury production in contaminated sediments. Environ Sci Technol. 2007;41(7):2270-6.

174

195. Diaz X, Johnson WP, Fernandez D, Naftz DL. Size and elemental distributions of nano- to micro-particulates in the geochemically-stratified Great Salt Lake. Appl Geochem. 2009;24(9):1653-65.

196. Guentzel JL, Powell RT, Landing WM, Mason RP. Mercury associated with colloidal material in an estuarine and an open-ocean environment. Mar Chem. 1996;55(1-2):177-88.

197. Stordal MC, Gill GA, Wen LS, Santschi PH. Mercury phase speciation in the surface waters of three Texas estuaries: importance of colloidal forms. Limnol Oceanogr. 1996;41(1):52-61.

198. Aiken GR, Hsu-Kim H, Ryan JN. Influence of dissolved organic matter on the environmental fate of metals, nanoparticles, and colloids. Environ Sci Technol. 2011;45:3196-201.

199. Labrenz M, Druschel GK, Thomsen-Ebert T, Gilbert B, Welch SA, Kemner KM, et al. Formation of sphalerite (ZnS) deposits in natural biofilms of sulfate-reducing bacteria. Science. 2000;290(5497):1744-7.

200. Moreau JW, Weber PK, Martin MC, Gilbert B, Hutcheon ID, Banfield JF. Extracellular proteins limit the dispersal of biogenic nanoparticles. Science. 2007;316(5831):1600-3.

201. Bosch J, Heister K, Hofmann T, Meckenstock RU. Nanosized iron oxide colloids strongly enhance microbial iron reduction. Appl Environ Microbiol. 2010;76(1):184-9.

202. Dehner CA, Barton L, Maurice PA, Dubois JL. Size-dependent bioavailability of hematite (alpha-Fe 2O3) nanoparticles to a common aerobic bacterium. Environ Sci Technol. 2011;45(3):977-83.

203. Lowry OH, Rosebrough NJ, Farr AL, Randall RJ. Protein measurement with the folin phenol reagent. J Biol Chem. 1951;193(1):265-75.

204. Gilmour CC, Elias DA, Kucken AM, Brown SD, Palumbo AV, Schadt CW, et al. The sulfate-reducing bacterium Desulfovibrio desulfuricans ND132 as a model for understanding bacterial mercury methylation. Appl Environ Microbiol. 2011:In press.

205. Institute of Experimental Mineralogy RAoS. Crystallographic and Crystallochemical Database for Minerals and their Structural Analogues. Created 1997, Updated 2009( http://database.iem.ac.ru/mincryst/index.php ). 175

206. Birks LS, Friedman H. Particle size determination from X-ray line broadening. J Appl Phys. 1946;17(8):687-91.

207. Kim CS, Bloom NS, Rytuba JJ, Brown GE. Mercury speciation by X-ray absorption fine structure spectroscopy and sequential chemical extractions: A comparison of speciation methods. Environ Sci Technol. 2003;37(22):5102-8.

208. U.S. Environmental Protection Agency. Method 1630: Methyl Mercury in water by Distillation, Aqueous Ethylation, Purge and Trap, and CVAFS. Washington, DC. 2001.

209. U.S. Environmental Protection Agency. Method 1631, Revision D: Mercury in Water by Oxidation, Purge and Trap, and Cold Vapor Atomic Fluorescence Spectroscopy. Washington, DC. 2001.

210. Pak KR, Bartha R. Mercury methylation and demethylation in anoxic lake sediments and by strictly anaerobic bacteria. Appl Environ Microbiol. 1998;64(3):1013-7.

211. Schaefer JK, Morel FMM. High methylation rates of mercury bound to cysteine by Geobacter sulfurreducens. Nature Geosci. 2009;2(2):123-6.

212. Mishra B, Fein J, Yee N, Beverridge T, Myneni S. Hg(II) adsorption and speciation on bacterial surfaces. Geochim Cosmochim Acta. 2010;74(12):A713-A.

213. Harris RC, Rudd JWM, Amyot M, Babiarz CL, Beaty KG, Blanchfield PJ, et al. Whole-ecosystem study shows rapid fish-mercury response to changes in mercury deposition. Proc Nat Acad Sci USA. 2007;104(42):16586-91.

214. Compeau GC, Bartha R. Sulfate-reducing bacteria - principal methylators of mercury in anoxic estuarine sediment. Appl Environ Microbiol. 1985;50(2):498- 502.

215. Drott A, Lambertsson L, Bjorn E, Skyllberg U. Do potential methylation rates reflect accumulated methyl mercury in contaminated sediments? Environmental Science & Technology. 2008;42(1):153-8.

216. Hintelmann H, Keppel-Jones K, Evans RD. Constants of mercury methylation and demethylation rates in sediments and comparison of tracer and ambient mercury availability. Environ Toxicol Chem. 2000;19(9):2204-11.

176

217. Warner KA, Roden EE, Bonzongo JC. Microbial mercury transformation in anoxic freshwater sediments under iron-reducing and other electron-accepting conditions. Environmental Science & Technology. 2003;37(10):2159-65.

218. King JK, Kostka JE, Frischer ME, Saunders FM, Jahnke RA. A quantitative relationship that remonstrates mercury methylation rates in marine sediments are based on the community composition and activity of sulfate-reducing bacteria. Environmental Science & Technology. 2001;35(12):2491-6.

219. King JK, Kostka JE, Frischer ME, Saunders FM. Sulfate-reducing bacteria methylate mercury at variable rates in pure culture and in marine sediments. Applied and Environmental Microbiology. 2000;66(6):2430-7.

220. Han S, Obraztsova A, Pretto P, Deheyn DD, Gieskes J, Tebo BM. Sulfide and iron control on mercury speciation in anoxic estuarine sediment slurries. Marine Chemistry. 2008;111(3-4):214-20.

221. Hammerschmidt CR, Fitzgerald WF. Geochemical controls on the production and distribution of methylmercury in near-shore marine sediments. Environmental Science & Technology. 2004;38(5):1487-95.

222. Zhang T, Kim B, Levard C, Reinsch BC, Lowry GV, Deshusses MA, et al. Methylation of mercury by bacteria exposed to dissolved, nanoparticulate, and microparticulate mercuric sulfides. Environmental Science & Technology. In press.

223. Oremland RS, Culbertson CW, Winfrey MR. Methylmercury decomposition in sediments and bacterial cultures - Involvement of methanogens and sulfate reducers in oxidative demethylation. Applied and Environmental Microbiology. 1991;57(1):130-7.

224. Behra P, Bonnissel-Gissinger P, Alnot M, Revel R, Ehrhardt JJ. XPS and XAS study of the sorption of Hg(II) onto pyrite. Langmuir. 2001;17(13):3970-9.

225. Andrews JC. Mercury speciation in the environment using X-ray absorption spectroscopy. Recent Developments in Mercury Science2006. p. 1-35.

226. Bloom NS, Colman JA, Barber L. Artifact formation of methyl mercury during aqueous distillation and alternative techniques for the extraction of methyl mercury from environmental samples. Fresenius Journal of Analytical Chemistry. 1997;358(3):371-7.

177

227. U.S. Environmental Protection Agency. Method 7473: Mercury in Solids and Solutions by Thermal Decomposition, Amalgamation, and Atomic Absorption Spectrophotometry. 1998.

228. Allen HE, Fu GM, Deng BL. Analysis of acid-volatile sulfide (AVS) and simultaneously extracted metals (SEM) for the estimation of potential toxicity in aquatic sediments. Environmental Toxicology and Chemistry. 1993;12(8):1441-53.

229. Cline JD. Spectrophotometric determination of hydrogen sulfide in natural waters. Limnology and Oceanography. 1969;14(3):454-&.

230. Olson RV, Ellis RJ. Iron. In: Page, A.L., et al. (Eds.), Methods of Soil Analysis. Part 2. Chemical and Microbiological Properties. 2nd ed. Agronomy no. 9. American Society of Agronomy, Soil Science Society of America, Madison, WI, pp. 301-312. 1982.

231. Schecher WD. MINEQL+: A Chemical Equilibrium Program for Personal Computers, 4.5 ed.; Environmental Research Software, Hallowell, ME, 2001.

232. Institute SFE. San Francisco Estuary Institute Regional Monitoring Program data; San Francisco Estuary Institute: San Francisco, CA; accessed at http://www.sfei.org on Feb 19, 2012.

233. Heyes A, Mason RP, Kim EH, Sunderland E. Mercury methylation in estuaries: Insights from using measuring rates using stable mercury isotopes. Marine Chemistry. 2006;102(1-2):134-47.

234. Hammerschmidt CR, Fitzgerald WF. Methylmercury cycling in sediments on the continental shelf of southern New England. Geochimica et Cosmochimica Acta. 2006;70(4):918-30.

235. Sunderland EM, Gobas F, Heyes A, Branfireun BA, Bayer AK, Cranston RE, et al. Speciation and bioavailability of mercury in well-mixed estuarine sediments. Marine Chemistry. 2004;90(1-4):91-105.

236. Gilmour CC, Elias DA, Kucken AM, Brown SD, Palumbo AV, Schadt CW, et al. Sulfate-Reducing Bacterium Desulfovibrio desulfuricans ND132 as a Model for Understanding Bacterial Mercury Methylation. Applied and Environmental Microbiology. 2011;77(12):3938-51.

237. Long GY, Zhu PT, Shen Y, Tong MP. Influence of Extracellular Polymeric Substances (EPS) on Deposition Kinetics of Bacteria. Environmental Science & Technology. 2009;43(7):2308-14. 178

238. Jeong HY, Klaue B, Blum JD, Hayes KF. Sorption of mercuric ion by synthetic manocrystalline mackinawite (FeS). Environmental Science & Technology. 2007;41(22):7699-705.

239. Morse JW, Luther GW. Chemical influences on trace metal-sulfide interactions in anoxic sediments. Geochimica et Cosmochimica Acta. 1999;63(19-20):3373-8.

240. Marvin-Dipasquale MC, Oremland RS. Bacterial methylmercury degradation in Florida Everglades peat sediment. Environ Sci Technol. 1998;32(17):2556-63.

241. Oremland RS, Culbertson CW, Winfrey MR. Methylmercury decomposition in sediments and bacterial cultures: Involvement of methanogens and sulfate reducers in oxidative demethylation. Appl Environ Microbiol. 1991;57(1):130-7.

242. Sellers P, Kelly CA, Rudd JWM. Fluxes of methylmercury to the water column of a drainage lake: The relative importance of internal and external sources. Limnol Oceanogr. 2001;46(3):623-31.

243. Hoigne J, Bader H. Ozone and hydroxyl radical initiated oxidations of organic and organometallic trace impurities in water. In: Brinckman FE, Bellama JM, editors. Organometals and organo-metalloids; Occurence and fate in the environment Washington, DC: ACS Symposium Series 82. American Chemical Society; 1978. p. 292-313.

244. Chen J, Pehkonen SO, Lin C-J. Degradation of monomethylmercury chloride by hydroxyl radicals in simulated natural waters. Water Res. 2003;37:2496-504.

245. Hoigne J. Comment on "degradation of monomethylmercury chloride by hydroxyl radicals in simulated natural waters". Water Research. 2004;38(14- 15):3470-1.

246. Zepp RG, Hoigne J, Bader H. Nitrate-induced photooxidation of trace organic- chemicals in water. Environmental Science & Technology. 1987;21(5):443-50.

247. Southworth BA, Voelker BM. Hydroxyl Radical Production via the Photo-Fenton in the Presence of Fulvic Acid. Environ Sci Technol. 2003;37:1130-6.

248. Vermilyea AW, Voelker BM. Photo-Fenton Reaction at Near Netural pH. Environ Sci Technol. 2009;43:6927-33.

249. Lehnherr I, St Louis VL. Importance of ultraviolet radiation in the photodemethylation of methylmercury in freshwater ecosystems. Environ Sci Technol. 2009;43(15):5692-998.

179

250. Suda I, Suda M, Hirayama K. Degradation of methyl and ethyl mercury by singlet oxygen generated from sea water exposed to sunlight or ultraviolet light. Archives of Toxicology. 1993;67:365-8.

251. Grandbois M, Latch DE, McNeill K. Microheterogeneous Concentrations of Singlet Oxygen in Natural Organic Matter Isolate Solutions. Environmental Science & Technology. 2008;42(24):9184-90.

252. Amyot M, Gill GA, Morel FMM. Production and loss of dissolved gaseous mercury in coastal seawater. Environ Sci Technol. 1997;31:3606-11.

253. Black FJ, Conaway CH, Flegal AR. Stability of dimethyl mercury in seawater and its conversion to monomethyl mercury. Environ Sci Technol. 2009;43(11):4056-62.

254. Gryglik D, Miller JS, Ledakowicz S. Singlet molecular oxygen application for 2- chlorophenol removal. Journal of Hazardous Materials. 2007;146(3):502-7.

255. Miller JS. Rose bengal-sensitized photooxidation of 2-chlorophenol in water using solar simulated light. Water Research. 2005;39:412-22.

256. Bandy J, Shemer H, Linden KG. Impact of lamp choice and H 2O2 dose on photodegradation of nitrobenzene. Environmental Engineering Science. 2009;26:973-80.

257. Hoigne J. Inter-calibration of OH radical sources and water quality parameters. Water Science and Technology. 1997;35:1-8.

258. Ukeda H, Maeda S, Ishii T, Sawamura M. Spectrophotometric Assay for Superoxide Dismutase Based on Tetrazolium Salt 3'-{1-[(Phenylamino)- carbonyl]-3,4-tetrazolium}-bis(4-methoxy-6-nitro)benzenesulfonic Acid Hydrate Reduction by Xanthine–Xanthine Oxidase. ANALYTICAL BIOCHEMISTRY. 1997;251:206-9.

259. Bloom NS. Determination of picogram levels of methylmercury by aqueous phase ethylation, followed by cryogenic gas chromatography with cold vapor atomic fluorescence detection. Can J Fish Aquat Sci. 1989;46:1131-40.

260. U.S. Environmental Protection Agency. Method 1630; Methyl Mercury in water by Distillation, Aqueous Ethylation, Purge and Trap, and CVAFS. Washington, DC; 2001 Contract No.: Document Number|.

180

261. U.S. Environmental Protection Agency. Method 7473 Mercury in Solids and Solutions by Thermal Decomposition, Amalgamation, and Atomic Absorption Spectrophotometry; 1998 Contract No.: Document Number|.

262. Vairavamurthy A, Mopper K. Field methods for determination of traces of thiols in natural waters. Anal Chim Acta. 1990;236:363-70.

263. Chae S, Hotze EM, Wiesner M. Evaluation of the oxidation of organic compounds by aqueous suspensions of photosensitized hydroxylated-C60 fullerene aggregates. Environmental Science & Technology. 2009;43:6208-13.

264. Pereira VJ, Weinberg HS, Linden KG, Singer PC. UV Degradation Kinetics and Modeling of Pharmaceutical Compounds in Laboratory Grade and Surface Water via Direct and Indirect Photolysis at 254 nm. Environmental Science & Technology. 2007;41:1682-8.

265. Brezonik PL. Chemical Kinetics and Process Dynamics in Aquatic Systems. Boca Raton, FL: CRC Press; 1993.

266. Schmidt R, Afshari E. Collisional deactivation of O-2(1-Delta-G) by solvent molecules - compartive experiments with O-16(2) and O-18(2). Ber Bunsen-Ges Phys Chem Chem Phys. 1992;96(6):788-94.

267. Boreen AL, Edhlund BL, Cotner JB, McNeill K. Indirect photodegradation of dissolved free amino acids: the contribution of singlet oxygen and the differential reactivity of DOM from various sources. Environmental Science & Technology. 2008;42:5492-8.

268. Zepp RG, Schlotzhauer PF, Sink RM. Photosensitized transformations involving electronic-energy transfer in natural waters: Role of humic substances. Environ Sci Technol. 1985;19(1):74-81.

269. Amirbahman A, Reid AL, Haines TA, Kahl JS, Arnold C. Association of methylmercury with dissolved humic acids. Environmental Science & Technology. 2002;36(4):690-5.

270. Hsu-Kim H. Stability of Metal-Glutathione Complexes During Oxidation by Hydrogen Peroxide and Cu(II)-Catalysis. Environ Sci Technol. 2007;41:2338-42.

271. Abedinzadeh Z, Gardes-Albert M, Ferradini C. Kinetic study of the oxidation mechanism of glutathione by hydrogen peroxide in neutral aqueous solution. Can J Chem. 1989;67:124701255.

181

272. Hintelmann H, Welbourn PM, Evans RD. Measurement of complexation of methylmercury(II) compounds by freshwater humic substances using equilibrium dialysis. Environ Sci Technol. 1997;31(2):489-95.

273. Karlsson T, Skyllberg U. Bonding of ppb levels of methyl mercury to reduced sulfur groups in soil organic matter. Environ Sci Technol. 2003;37(21):4912-8.

274. Horzempa LM, Helz GR. Controls on the stability of sulfide sols-colloidal covellite as an example. Geochimica et Cosmochimica Acta. 1979;43(10):1645-50.

275. Deonarine A, Lau BLT, Aiken GR, Ryan JN, Hsu-Kim H. Effects of Humic Substances on Precipitation and Aggregation of Zinc Sulfide Nanoparticles. Environmental Science & Technology. 2011;45(8):3217-23.

276. Lau BLT, Hsu-Kim H. Precipitation and growth of zinc sulfide nanoparticles in the presence of thiol-containing natural organic ligands. Environmental Science & Technology. 2008;42(19):7236-41.

277. Kim B, Park CS, Murayama M, Hochella MF. Discovery and Characterization of Silver Sulfide Nanoparticles in Final Sewage Sludge Products. Environmental Science & Technology. 2010;44(19):7509-14.

278. Kemner KM, Kelly SD, Lai B, Maser J, O'Loughlin EJ, Sholto-Douglas D, et al. Elemental and redox analysis of single bacterial cells by X-ray microbeam analysis. Science. 2004;306(5696):686-7.

279. Suzuki Y, Kelly SD, Kemner KM, Banfield JF. Radionuclide contamination - Nanometre-size products of uranium bioreduction. Nature. 2002;419(6903):134-.

280. Rickard D, Luther GW. Kinetics of pyrite formation by the H2S oxidation of iron(II) monosulfide in aqueous solutions between 25 and 125 degrees C: The mechanism. Geochimica et Cosmochimica Acta. 1997;61(1):135-47.

281. Ulrich PD, Sedlak DL. Impact of Iron Amendment on Net Methylmercury Export from Tidal Wetland Microcosms. Environmental Science & Technology. 2010;44(19):7659-65.

282. Rozan TF, Lassman ME, Ridge DP, Luther GW. Evidence for iron, copper and zinc complexation as multinuclear sulphide clusters in oxic rivers. Nature. 2000;406(6798):879-82.

182

283. Luther GW, Rozan TF, Taillefert M, Nuzzio DB, Di Meo C, Shank TM, et al. Chemical speciation drives hydrothermal vent ecology. Nature. 2001;410(6830):813-6.

284. Luther GW, Ferdelman TG. Voltammetric characterization of iron (II) sulfide complexes in laboratory solutions and in marine waters and porewaters. Environmental Science & Technology. 1993;27(6):1154-63.

285. Jeong HY, Sun K, Hayes KF. Microscopic and Spectroscopic Characterization of Hg(II) Immobilization by Mackinawite (FeS). Environmental Science & Technology. 2010;44(19):7476-83.

286. Skyllberg U, Drott A. Competition between Disordered Iron Sulfide and Natural Organic Matter Associated Thiols for Mercury(II)-An EXAFS Study. Environmental Science & Technology. 2010;44(4):1254-9.

287. Qian J, Skyllberg U, Frech W, Bleam WF, Bloom PR, Petit PE. Bonding of methyl mercury to reduced sulfur groups in soil and stream organic matter as determined by X-ray absorption spectroscopy and binding affinity studies. Geochimica et Cosmochimica Acta. 2002;66(22):3873-85.

288. Chen L, Tang X, Shen C, Chen C, Chen Y. Photosensitized degradation of 2,4 ',5- trichlorobiphenyl (PCB 31) by dissolved organic matter. Journal of Hazardous Materials. 2012;201:1-6.

289. Grannas AM, Cory RM, Miller PL, Chin Y-P, McKnight DM. The role of dissolved organic matter in arctic surface waters in the photolysis of hexachlorobenzene and lindane. Journal of Geophysical Research-Biogeosciences. 2012;117.

290. Burns SE, Hassett JP, Rossi MV. Binding effects on humic-mediated photoreaction: Intrahumic dechlorination of mirex in water. Environmental Science & Technology. 1996;30(10):2934-41.

291. Burns SE, Hassett JP, Rossi MV. Mechanistic implications of the intrahumic dechlorination of mirex. Environmental Science & Technology. 1997;31(5):1365-71.

292. Blough NV. Electron-paramagnetic resonance measurements of photochemical radical production in humic substances. 1. Effects of O-2 and charge on radical scavenging by nitroxides. Environmental Science & Technology. 1988;22(1):77-82.

183

293. Kohn T, Grandbois M, McNeill K, Nelson KL. Association with natural organic matter enhances the sunlight-mediated inactivation of MS2 coliphage by singlet oxygen. Environmental Science & Technology. 2007;41(13):4626-32.

294. Reinsch BC, Forsberg B, Penn RL, Kim CS, Lowry GV. Chemical Transformations during Aging of Zerovalent Iron Nanoparticles in the Presence of Common Groundwater Dissolved Constituents. Environmental Science & Technology. 2010;44(9):3455-61.

295. Dyrssen D, Wedborg M. The sulfur-mercury (II) system in natural-waters. Water Air and Soil Pollution. 1991;56:507-19.

296. Haitzer M, Aiken GR, Ryan JN. Binding of mercury(II) to aquatic humic substances: Influence of pH and source of humic substances. Environmental Science & Technology. 2003;37(11):2436-41.

297. Khwaja AR, Bloom PR, Brezonik PL. Binding constants of divalent mercury (Hg2+) in soil humic acids and soil organic matter. Environmental Science & Technology. 2006;40(3):844-9.

298. Wei DW, Osseoasare K. Formation of iron monosulfide - A spectrophotometric study of the reaction between ferrous and sulfide ions in aqueous-solutions. Journal of Colloid and Interface Science. 1995;174(2):273-82.

299. Latch DE, Stender BL, Packer JL, Arnold WA, McNeill K. Photochemical Fate of Pharmaceuticals in the Environment: Cimetidine and Ranitidine. Environmental Science & Technology. 2003;37:3342-50.

300. Schmidt R. Deactivation of O 2 (1△g) Singlet Oxygen by Carotenoids: Internal

Conversion of Excited Encounter Complexes. J Phy Chem A. 2004;108:5509-13.

301. National Institute of Standards and Technology. Critical Stability Constants of Metal Complexes Database. v. 2.0 ed. Database NSR, editor1993.

302. Sutherland MW, Learmonth BA. Tetrazolium dyes MTS and XTT provide new quatitative assays for superoxide and superoxide dismutase. Free Radical Research. 1997;27:283-9.

184

Biography

Tong Zhang was born in 1981 in Tianjin, China. Ms. Zhang obtained a Bachelor and Master of Science in Environmental Science and Engineering at Tsinghua University,

Beijing, China (2006), followed by a Doctor of Philosophy in Civil and Environmental

Engineering at Duke University, Durham, North Carolina (2012).

Published work to date includes the following journal articles: Methylation of mercury by bacteria exposed to dissolved, nanoparticulate, and microparticulate mercuric sulfides,

Environmental Science &Technology (2012) in press; Photolytic degradation of methylmercury enhanced by binding to natural organic ligands, Nature Geoscience (2010) 3(7), pp 473-476.

Awards include the American Chemical Society Ellen C. Gonter Environmental

Chemistry Award (2012) and the Duke University Department of Civil and

Environmental Engineering Senol Utku Award for student publications (2011).

185