Mercury Stabilization Using Thiosulfate Or Selenosulfate

Home , Mercury selenide, Mercury sulfide, Selenide, Sodium hypochlorite

MERCURY STABILIZATION USING THIOSULFATE OR SELENOSULFATE

Zizheng Zhou

B.A.Sc, The University of British Columbia, 2011

A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF

THE REQUIREMENTS FOR THE DEGREE OF

MASTER OF APPLIED SCIENCE

The Faculty of Graduate Studies

(Materials Engineering)

THE UNIVERSITY OF BRITISH COLUMBIA

(Vancouver)

April, 2013

ABSTRACT

Mercury is often found associated with gold and silver minerals in ore bodies. It is recovered as liquid elemental mercury in several stages including carbon adsorption, carbon elution, electrowinning and retorting. Thus a great amount of mercury is produced as a by-product in gold mines. The Mercury Export Ban Act of 2008 prohibits conveying, selling and distributing elemental mercury by federal agencies in United States. It also bans the export of elemental mercury starting January 1, 2013. As a result, a long-term mercury management plan is required by gold mining companies that generate liquid mercury as a by-product.

This thesis will develop a process to effectively convert elemental mercury into much more stable mercury sulfide and mercury selenide for safe disposal. The process consists of 1) extraction of elemental mercury into solution to form aqueous mercury (II) and 2) mercury precipitation as mercury sulfide or mercury selenide.

Elemental mercury can be effectively extracted by using hypochlorite solution in acidic environment to form aqueous mercury (II) chloride. The effect of different parameters on the extent and rate of mercury extraction were studied, such as pH, temperature, stirring speed and hypochlorite concentration. Results show that near complete extraction can be achieved within 8 hours by using excess sodium hypochlorite at pH 4 with a fast stirring speed of 1000RPM.

Mercury precipitation was achieved by using thiosulfate and selenosulfate solution. In thiosulfate precipitation, cinnabar, metacinnabar or a mixture of both can be obtained depending on the experimental conditions. Elevated temperatures, acidic environment and high reagent concentrations favour the precipitation reaction. Complete mercury removal can be achieved within 4 hours. However, it appears that the less stable metacinnabar tends to form when the precipitation rate increases.

Selenosulfate solution can be produced by dissolving elemental selenium in sulfite solution at elevated temperature. Precipitation of mercury selenide using selenosulfate

reagent was found to be very effective. The precipitation rate proved to be extremely fast, and the formed precipitates have been confirmed to be tiemannite (HgSe) in all experiments.

Finally, Solid Waste Disposal Characterization (SWDC) experiments were conducted to examine the mobility of the formed mercury sulfide and mercury selenide. The results show that none of the formed precipitates exceed the Ultimate Treatment Standard (UTS) limit.

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TABLE OF CONTENTS

ABSTRACT...... ii

TABLE OF CONTENTS...... iv

LIST OF TABLES...... ix

LIST OF FIGURES ...... x

ACKNOWLEDGEMENTS...... xiv

DEDICATION...... xv

1 INTRODUCTION ...... 1

2 LITERATURE REVIEW ...... 3

2.1 Mercury Generation in Mining Industry and Its Impact...... 3

2.1.1 Mercury in Small-Scale Gold Mining...... 3

2.1.2 Mercury in Modern Gold Mining and Non-Ferrous Mining ...... 4

2.1.3 Impact of Mining Activities on Mercury Emissions...... 5

2.2 Legislative Background ...... 6

2.3 Overview of Existing Mercury Stabilization Technologies...... 7

2.3.1 Stabilization of Mercury as Mercury Sulfide or Mercury Selenide...... 7

2.3.1.1 DELA Process, German...... 8

2.3.1.2 Bethlehem Apparatus...... 8

2.3.1.3 STMI Process, France...... 8 iv

2.3.1.4 CENIM Milling Process, Spain ...... 9

2.3.1.5 Synthesis of Mercury Sulfide by Shaking ...... 9

2.3.1.6 Wet Process...... 9

2.3.1.7 Stabilization of Mercury as Mercury Selenide ...... 9

2.3.2 Mercury Stabilization via Amalgamation...... 10

2.3.2.1 Amalgamation with Copper...... 10

2.3.2.2 Amalgamation with Zinc ...... 10

2.3.3 Stabilization of Mercury into a Stable and Insoluble Matrix...... 10

2.3.3.1 ATG Stabilization Process...... 10

2.3.3.2 Sulfur Polymer Cement Process ...... 11

2.3.3.3 Magnesia Binder ...... 11

2.3.3.4 Solidification with Hydraulic Cement and Sulfur Polymer Cement ...... 11

2.4 Aqueous Chemistry of Mercury...... 12

2.4.1 Hg - H2O Chemistry...... 12

2.4.2 Hg - Cl - H2O Chemistry ...... 15

2.4.2.1 Interaction between Mercury and Hypochlorite ...... 19

2.4.3 Hg - S - H2O Chemistry...... 20

2.4.3.1 Thiosulfate Chemistry...... 22

2.4.3.2 Interaction between Mercury and Thiosulfate ...... 23

2.4.4 Hg - Se - H2O Chemistry ...... 26

2.5 Research Objectives...... 30

3 EXPERIMENTAL METHODS...... 31

3.1 Mercury Leaching Experiments...... 31

3.1.1 Hypochlorite Leaching Experiments ...... 32

3.1.2 Hydrogen Peroxide Leaching Experiments ...... 35

3.1.3 Cyanidation Experiments...... 36

3.2 Mercury Precipitation Experiments ...... 36

3.2.1 Thiosulfate Precipitation Experiments...... 36

3.2.2 Selenosulfate Precipitation Experiments ...... 39

3.2.3 Selenious Acid Precipitation Experiments...... 40

3.3 Selenium Dissolution Experiments...... 40

3.4 Solid Waste Disposal Characterization...... 42

4 RESULTS AND DISCUSSION...... 44

4.1 Hypochlorite Leaching Experiments ...... 44

4.1.1 Effect of pH...... 44

4.1.2 Effect of Stirring Speed ...... 46

4.1.3 Effect of Hypochlorite Concentration...... 47

4.1.4 Effect of Temperature...... 49

4.1.5 Randomness and Errors ...... 51

4.2 Other Types of Mercury Leaching Experiments ...... 52

4.2.1 Hydrogen Peroxide Leaching Experiments ...... 52

4.2.2 Cyanidation Experiments...... 53

4.2.3 Summary...... 54

4.3 Thiosulfate Precipitation Experiments...... 54

4.3.1 Preliminary Study ...... 55

4.3.2 Effect of pH...... 55

4.3.3 Effect of Temperature...... 59

4.3.4 Effect of "Seeding" ...... 62

4.4 Selenium Dissolution Experiments...... 63

4.5 Selenosulfate Precipitation Experiments ...... 65

4.5.1 Preliminary Experiment...... 66

4.5.2 Effect of Temperature...... 66

4.5.3 Effect of Selenosulfate Concentration ...... 69

4.6 Selenious Acid Precipitation Experiments...... 70

4.7 Solid Waste Disposal Characterization Experiments ...... 70

5. CONCLUSION...... 71

6. REFERENCES ...... 76

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LIST OF TABLES

Table 1 Global average mercury content in non-ferrous ores [36]...... 3

Table 2 Test conditions of hypochlorite leaching experiments ...... 34

Table 3 Test conditions of thiosulfate precipitation experiments...... 38

Table 4 Test conditions of selenosulfate precipitation experiments...... 40

Table 5 Experimental conditions for selenium dissolution experiments...... 42

Table 6 Solid waste disposal characterization experiments...... 43

Table 7 Comparison of theoretical and measured hypochlorite consumption...... 49

Table 8 Thermodynamic data for the oxidation reaction (HSC Database)...... 51

Table 9 Thermodynamic data for the overall leaching reaction (HSC Database) ...... 51

Table 10 Hydrogen peroxide leaching results ...... 53

Table 11 Results of cyanidation experiments ...... 54

Table 12 Preliminary selenium dissolution experiments done by Ullah (2012)...... 64

Table 13 Results for selenium dissolution experiments ...... 65

Table 14 Effect of temperature on selenosulfate precipitation experiments

...... 66

Table 15 Solid waste disposal characterization experiment results...... 70

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LIST OF FIGURES

Figure 1 Pourbaix diagram of Hg-H2O system at 25oC. [Hg] = 1 molal, Pressure = 1atm. (HSC 6.1) ...... 12

Figure 2 Pourbaix diagram of Hg-H2O system at 298K. [Hg] = 10-3 molal, Pressure = 1atm. (HSC 6.1) ...... 13

Figure 3 Pourbaix diagram of Hg-H2O system at 298K. [Hg] = 10-6 molal, Pressure = 1atm. (HSC 6.1) ...... 13

Figure 4 Pourbaix diagram for the Cl-H2O system at 298K. [Cl] = 1Molal, Pressure = 1atm. (HSC) ...... 16

Figure 5 Pourbaix diagram for the Hg-Cl-H2O system at 298K. [Hg] = 1Molal, [Cl] = 1Molal, Pressure = 1atm. (HSC)...... 17

-3 Figure 6 Pourbaix diagram for the Hg-Cl-H2O system at 298K. [Hg] = 10 Molal, [Cl] = 1Molal, Pressure = 1atm. (HSC)...... 18

-3 Figure 7 Pourbaix diagram for the Hg-Cl-H2O system at 298K. [Hg] = 10 Molal, [Cl] = 10-3Molal, Pressure = 1atm. (HSC) ...... 18

- Figure 8 Pourbaix diagram for the Hg-Cl-H2O system at 298K. [Hg] = 1Molal, [Cl] = 10 3Molal, Pressure = 1atm. (HSC) ...... 19

Figure 9 The structure of metacinnabar [6] ...... 20

Figure 10 Eh-pH diagram for the Hg-S-H2O system at 298K with activities of dissolved mercury and sulfur of 0.1 and 1, respectively. [24]...... 21

Figure 11 Eh-pH diagram for the Hg-S-H2O system at 298K with activities of dissolved mercury and sulfur of 10-6 and 1, respectively. [24]...... 22

Figure 12 Metastable Eh-pH diagram for the Hg-S-H2O system at 298K with activities of dissolved mercury and sulfur of 0.1 and 1, respectively; S(VI) species excluded from calculation [24] ...... 24

Figure 13 Metastable Eh-pH diagram for the Hg-S-H2O system at 298K with activities of dissolved mercury and sulfur of 10-6 and 1, respectively; S(VI) species excluded from calculation [24] ...... 25

Figure 14 Metastable Eh-pH diagram for the Hg-S-H2O system at 298K with activities of dissolved mercury and sulfur of 0.1 and 1, respectively; all sulfur-oxy species excluded from calculation [24]...... 26

Figure 15 Structure of mercury selenide...... 27

Figure 16 Pourbaix diagram for the Se-H2O system at 298K. [Se] = 1Molal, Pressure = 1atm. (HSC) ...... 27

Figure 17 Pourbaix diagram for the Hg-Se-H2O system at 298K. [Hg] = 1Molal, [Se] = 1Molal, Pressure = 1atm. (HSC)...... 28

-3 Figure 18 Pourbaix diagram for the Hg-Se-H2O system at 298K. [Hg] = 10 Molal, [Se] = 1Molal, Pressure = 1atm. (HSC)...... 28

-3 Figure 19 Pourbaix diagram for the Hg-Se-H2O system at 298K. [Hg] = 10 Molal, [Se] = 10--3Molal, Pressure = 1atm. (HSC)...... 29

Figure 20 Pourbaix diagram for the Hg-Se-H2O system at 298K. [Hg] = 1Molal, [Se] = 10--3Molal, Pressure = 1atm. (HSC)...... 29

Figure 21 Experimental setup for mercury leaching experiments. pH was monitored by the pH and temperature probe, and controlled by the pH controller; a balance was used to monitor the weight change of the acid/basic solution; temperature was monitored by a thermometer and controlled by a waterbath which is not in this Figure; Stirring was achieved by a magnetic stirrer. The system was sealed by using rubber stoppers...... 32

Figure 22 Experimental setup for selenium dissolution experiments...... 41

Figure 23 Effect of pH on hypochlorite leaching of mercury. Test Conditions: 20oC, 500RPM, hypochlorite concentration 10 times as much as stoichiometrically required.. 45

Figure 24 Effect of stirring speed on hypochlorite leaching of mercury. Test conditions: 20oC, pH = 4, hypochlorite concentration 10 times as much as stoichiometrically required...... 46

Figure 25 Effect of hypochlorite concentration on mercury leaching. Test conditions: 20oC, pH = 4, 1000 RPM, hypochlorite concentration N times as much as stoichiometrically required as indicated...... 47

Figure 26 Measured hypochlorite consumption in mercury leaching ...... 48

Figure 27 Effect of temperature on hypochlorite leaching of mercury. Test conditions: pH = 4, 500RPM, hypochlorite concentration 10 times as much as stoichiometrically required...... 50

Figure 28 Effect of Randomness on hypochlorite leaching of mercury. Test conditions: temperature at 20oC, pH = 4, 500RPM, hypochlorite concentration 10 times as much as stoichiometrically required...... 52

Figure 29 Effect of pH on precipitation rate in thiosulfate precipitation experiments. Test conditions: 80oC, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required...... 56

Figure 30 XRD pattern for the formed precipitates at test conditions: pH = 2, 80oC, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required..... 57

Figure 31 XRD pattern for the formed precipitates at test conditions: pH = 4, 80oC, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required..... 58

Figure 32 XRD pattern for the formed precipitates at test conditions: pH = 5, 80oC, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required..... 58

Figure 33 XRD pattern for the formed precipitates at test conditions: pH = 6, 80oC, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required..... 59

Figure 34 Effect of temperature on mercury precipitation at pH 2. Test conditions: pH = 2, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required..... 60

Figure 35 Effect of temperature on mercury precipitation at pH 4. Test conditions: pH = 4, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required..... 60

Figure 36 Effect of temperature on mercury precipitation at pH 5. Test conditions: pH = 5, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required..... 61

Figure 37 Effect of temperature on mercury precipitation at pH 6. Test conditions: pH = 6, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required..... 61

Figure 38 Effect of seeding at pH 5. Test conditions: 60oC, pH = 5, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required...... 62

Figure 39 Effect of seeding at pH 6. Test conditions: 60oC, pH = 5, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required...... 63

Figure 40 XRD pattern of formed precipitates at 10oC ...... 67

Figure 41 XRD pattern of formed precipitates at 40oC ...... 67

Figure 42 XRD pattern of formed precipitates at 60oC ...... 68

Figure 43 XRD pattern of formed precipitates at 80oC ...... 68

Figure 44 Effect of selenosulfate concentration on mercury precipitation experiments, Test conditions: 20oC, 500RPM, selenosulfate concentration 10 times as much as stoichiometrically required...... 69

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Figure 45 Flowsheet of mercury stabilization process via thiosulfate precipitation ...... 72

Figure 46 Flowsheet of mercury stabilization process via selenosulfate precipitation .... 74

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ACKNOWLEDGEMENTS

I want to specially thank my supervisor Dr. David Dreisinger for all the kind guidance and encouragement he has offered to me.

I want to thank Dr. Berend Wassink for all the kind assistance and support.

Finally, I want to thank all my family and my friends for their faith and support.

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DEDICATION

I lovingly dedicate this thesis to my wife, Henglin Jin, for her endless love and support through the completion of this project.

1 INTRODUCTION

As one of the first metals to be mined in history, mercury, also well known as quicksilver and hydrargyrum and has been found in Egyptian tombs dating back to 1500 B.C. Mercury has had an active role in ancient civilizations. Ancient Chinese, Greeks and Romans widely used mercury as almost everything from medicine to talismans. However, its toxicity started to be recognized when mercury mining became associated with human illness beginning as tremor and progressing to severe mental damage [1]. Mercury was first commercialized as early as 2700 B.C. after recovery at the Almadén mines in Spain [7]. However, mercury production became industrialized and globalized in 1554 due to the development of the "Patio" amalgamation process, in which mercury is used to extract silver from ores [2, 7].

Elemental mercury is a silvery, extremely dense liquid. It has an atomic number of 80, and an atomic weight of 200.59 g/mol. Mercury is a Group IIB element. At 25oC, the density of mercury is 13,534 kg/m3. At atmospheric pressure, its freezing point is -38.85oC, and its boiling point is 356.6oC. Mercury has extremely high surface tension, which gives it very unique rheological behaviour [3]. Mercury has a high electric conductivity and also a very good germicidal ability [5].

Elemental mercury, all inorganic mercury compounds and most organic mercury compounds are highly toxic to human beings by ingestion, inhalation and skin absorption. After being absorbed into the human body, mercury can attack and accumulate in body tissues, particularly in the brain and kidneys [4].

Despite its toxicity, mercury has various applications due to its unique properties, especially in metallurgical extraction (for example, gold and silver extraction), and chlorine and alkali production. Mercury was also employed in dental amalgams, catalysts, thermometers, barometers, manometers, electrical apparatus, mercury vapour lamps, mirror coatings, and as a coolant and neutron absorber in nuclear power plants [3].

Mercury does not occur in nature in the native form. The most common mineral of mercury in nature is Cinnabar (red HgS). Other mineral sources can also be found in

nature, including metacinnabar (black HgS), living stone (HgSb4S7), coloradite (HgTe),

tiemannite (HgSe), and calomel (Hg2Cl2). Mercury is estimated to have a concentration of 0.08 mg/kg in the earth's crust [9]. Mercury vapour is mainly generated from volcanic emissions and evaporation from oceans. Typically, the mercury concentration in the atmosphere ranges from 2-4 ng/m3 in uncontaminated areas, increases to about 20 ng/m3 in urban areas, and can reach up to 18 μg/m3 near some active volcanoes [4].

This thesis will focus on developing a process for potential long-term mercury management in the gold mining industry, which mainly consists of two major components: 1) leaching elemental mercury into aqueous solution, and 2) precipitation of mercury as a stable mercury compound.

This thesis consists of five chapters. The second chapter provides a chemical and legislative background that lies behind this project. It also reviews current mercury stabilization technologies. The third chapter describes all the experimental methods adopted in this work in detail, including the experimental procedures and chemicals used. The fourth chapter presents all the experimental results and discussions. The last chapter draws conclusions of this work and proposes two processes to stabilize elemental mercury.

2 LITERATURE REVIEW

This chapter will first look into mercury generation in the mining industry, then go through the legislative background of mercury regulation, and finally provide an overview of the aqueous chemistry of mercury and current mercury stabilization technologies.

2.1 Mercury Generation in Mining Industry and Its Impact

Mercury is closely associated with the mining industry, especially the gold mining industry, no matter if it is in the small-scale gold mines in developing regions, or in the large-scale modern gold mining facilities in developed regions.

Mercury can also be typically found coexisting with other non-ferrous metals including copper, lead, zinc and silver in ore bodies. Table 1 below shows a global average mercury concentration in non-ferrous ores.

Table 1 Global average mercury content in non-ferrous ores [36]

Ore Type Average Mercury Content in Ore Unit Copper 5 - 10 g Hg / t Cu Lead 3 - 44 g Hg / t Pb Zinc 7 - 87 g Hg / t Zn Gold and Silver 0.1 - 200 g Hg / t ore

2.1.1 Mercury in Small-Scale Gold Mining

Mercury has the unique property to readily form amalgams with precious metals [6, 35] This property has been widely utilized to concentrate or extract gold and silver from low concentration ores. [7] After crushing the ore, mercury will be applied to contact the gold and silver minerals via several different methods in order to form amalgams. Mercury amalgams can be separated by washing with water due to its high density. Then the

amalgams will be heated in a retort device to remove mercury, leaving behind gold and silver of relatively high purity.

Due to the fact that mercury amalgamation will generate large quantities of highly toxic mercury waste and emit gaseous mercury into the atmosphere, this technology has already been prohibited in most of the developed countries. However, since the 1970s, the method of using mercury amalgamation to extract gold has been widely applied by small-scale gold mines, or artisanal gold mines, in numerous developing countries and regions, such as Brazil, China, Southeastern Asia, and some African countries. [8, 15, 16] This is mainly due to the inexpensive, convenient and fast nature of the amalgamation process and most importantly, the lack of regulations in these regions.

When a mercury retort device is not used, mercury loss into the environment can reach up to more than half of the initially applied amount. According to Veiga, it can be reasonably estimated that approximately one ton of mercury would be released into the environment when one ton of gold is produced [19]. Obviously, this will cause severe long-term damage to the ecosystem and raise serious global concerns regarding mercury pollution.

2.1.2 Mercury in Modern Gold Mining and Non-Ferrous Mining

As mentioned before, in the gold mining industry in developed countries like the United States, the mercury amalgamation process has already been prohibited. Mercury is involved mainly as a by-product in gold mining industry and other non-ferrous metals industry in the US.

In the gold cyanidation process, cyanide is utilized to extract gold in a basic environment in the presence of oxygen. The cyano gold complex is then selectively loaded onto activated carbon, eluted, and finally recovered through electrowinning or other refining processes. However, mercury(II) can also react with cyanide to form highly soluble cyano mercury complexes.

2+ - 0 Hg + 2CN → Hg(CN)2

2+ - 2- Hg + 4CN → Hg(CN)4

As a result, mercury is also loaded onto activated carbon along with gold. Then the cyano mercury complex is eluted with gold and electrowon to the metallic state as a gold- mercury alloy. Retorting is practiced to separate the mercury by volatilization and condensation. The mercury is ultimately collected as liquid elemental mercury. Furthermore, mercury can also be recovered through several other hot processes such as in roasting and autoclaves via mercury controlling devices. Mercury is then stored on site until it can be delivered to other commercial facilities for purification and preparation for further sale. [17]

Mercury is also involved in the non-ferrous mining industry. In pyrometallurgy processes such as smelting and roasting, the high temperatures cause mercury to vaporize and therefore present to the off gas. [36] Then mercury can be recovered via mercury control technologies. In the most common case, mercuric chloride is sprayed in the scrubber cell for the roaster to form non-volatile mercurous chloride precipitates:

0 HgCl2 + Hg → Hg2Cl2

The formed mercurous chloride is then treated in a mercury recycler to convert mercurous chloride to elemental mercury.

2.1.3 Impact of Mining Activities on Mercury Emissions

Due to its volatility and toxicity, mercury emissions are always an environmental concern. Therefore, research has been conducted in this field [7, 8, 12, 15, 16, 18, 19]. It is estimated that mercury emissions from natural sources are approximately 2,500 tonnes annually, while anthropogenic mercury emissions can reach up to 4,000 tonnes annually. [8, 12]

Gold mining activities are contributing an increasing amount of mercury in recent decades. It can be summarized that gold mining activities are responsible for between 10% to 20% of the annual anthropogenic mercury [8, 12]. It is possible that the mercury

amalgamation processes in small-scale gold mining facilities all over the world is one major contributor. In developing countries such as Brazil, where numerous small-scale gold mines are still in operation in the Amazon River area, gold mining activities are responsible for about 2/3 of the total mercury emissions in the country. [8]

The dominant mercury species that is directly released into the environment is gaseous elemental mercury, rather than mercury salts. However, due to a series of natural or atmospheric chemical reactions, nearly all elemental mercury will be converted into mercury(II) in the ecosystem. Unfortunately, elemental mercury is relatively mobile and has a life time of 1 to 2 years in the atmosphere, which is much higher in comparison with mercury (II) salts (several days). Therefore, it is not difficult for elemental mercury to be transported to remote areas. [10] Thus, serious concerns have been raised regarding global mercury pollution due the high mobility of elemental mercury.

2.2 Legislative Background

On October 14, 2008, the Mercury Export Ban Act (MEBA) was signed into law prohibiting "conveying, selling, or distributing elemental mercury", or exporting elemental mercury from the United States, unless qualifies for the exemptions as stated in the act. The act thus requires the Department of Energy to establish a long-term management and storage of the elemental mercury produced within the country. The main goal of this act is to significantly lower the elemental mercury availability in the global market, which will then impact the small-scale gold mining facilities and other industries utilizing mercury.

As mentioned before, elemental mercury is currently produced as a by-product in gold mines and it is stored on site for sale or secondary treatment. The act requires gold mines and other mercury producing facilities to develop long-term mercury stabilization processes and storage plans.

2.3 Overview of Existing Mercury Stabilization Technologies

The purpose of mercury stabilization is to convert mercury to its stable compounds with low mercury leachability and low mercury vapour pressure in order to meet standards set by relevant regulations. In this section current available mercury stabilization technologies will be briefly reviewed.

Generally three mechanisms are behind current mercury stabilization processes:

1. Stabilization of mercury as mercury sulfide or mercury selenide. The mercury stabilization technology presented in this paper also lies in this category. 2. Stabilization of mercury as amalgams. 3. Stabilization of mercury into a stable and insoluble matrix.

The final mercury containing product in the mercury stabilization processes shall undergo toxic characteristic leaching procedure (TCLP), which is designed to simulate landfill condition to examine the mobility of the analytes in wastes. TCLP experiments consist of preparing samples for leaching, leaching, preparing leachate solution for analysis and leachate analysis. The goal is to examine whether the concentration of the analyte in the final leachate passes the regulated standard. If so, the waste shall be deemed toxic. The RCRA limit for mercury is 0.2 mg/L, and the universal treatment standard (UTS) limit for mercury is 0.025 mg/L. If a mercury waste meets the RCRA limit in TCLP, the waste can be considered as "non-hazardous". If a mercury waste meets the UTS limit in TCLP, the waste can be disposed in landfills.

2.3.1 Stabilization of Mercury as Mercury Sulfide or Mercury Selenide

As mentioned before, mercury sulfide and mercury selenide both are physically and chemically stable compounds. Therefore, they are the desired products for numerous mercury stabilization processes.

2.3.1.1 DELA Process, German

DELA process was developed by the German DELA GmbH to stabilize elemental mercury as mercury sulfide. In this process, elemental mercury and sulfur are mixed in a heated vacuum; oxygen is absent. The reaction temperature is kept above 580oC which is higher than the boiling points for both elements (356.6oC for mercury and 444.6oC for sulfur) to ensure the reaction occurs in gaseous phase:

Hg(g) + S(g) → HgS(g)

Excess amount of sulfur is added prior to the addition of mercury. The mercury sulfide is then collected by condensation, and is a mixture of metacinnabar and cinnabar. [35]

2.3.1.2 Bethlehem Apparatus

This process was developed by Bethlehem Apparatus Co., USA. The idea is the reaction between elemental mercury and sulfur in gaseous phase at high temperature. The unique point of this process is that the formed mercury sulfide is mixed with patented polymers to produce pellets of the size of 7 x 7 mm. It was confirmed that the product has the same physical and chemical properties as cinnabar. [35, 37]

2.3.1.3 STMI Process, France

The direct interaction between elemental mercury and sulfur is still adopted in the mercury stabilization process developed by STMI:

Hg(l) + S(s) → HgS(s)

Unlike the DELA process, the STMI process takes place at a much lower temperature of 60 to 80oC. The molar ratio of sulfur and mercury is controlled between 1 to 3. Mixing is achieved by rotating the reactor at 50 RPM. The final product of this process was confirmed to be a mixture of metacinnabar and sulfur. [35]

2.3.1.4 CENIM Milling Process, Spain

In this process, formation of mercury sulfide is achieved by reacting elemental mercury and sulfur in a ball mill. Milling can be done for 15 minutes to 3 hours at 400RPM and room temperature. The final product mainly consists of metacinnabar. However, other species can also be formed depending on the time of milling. A longer milling time will result in the formation of undesirable mercury oxide. [35]

2.3.1.5 Synthesis of Mercury Sulfide by Shaking

Similar to the milling process described above, mercury sulfide is synthesized by applying work to elemental mercury and sulfur. Elemental mercury and powdered sulfur are shaken in a paint shaker with steel milling balls. The shaking is carried out longitudinally for one hour, and then transversely for one more hour. The formed product was determined to be mercury sulfide. However, whether cinnabar or metacinnabar is formed was not stated in the literature. [35]

2.3.1.6 Wet Process

Unlike previous stabilization technologies, elemental mercury can also be stabilized in cold and aqueous media. In a typical process, mercury is first dissolved in an oxidizing acid (e.g. concentrated nitric acid) to form soluble mercury(II). Then mercury is precipitated by the addition of sulfide to form metacinnabar.

0 + - 2+ 3Hg + 8H + 2NO3 → 3Hg + 2NO(g) + 4H2O

2+ 2- Hg + S → HgS(s)

2.3.1.7 Stabilization of Mercury as Mercury Selenide

In this process, mercury containing waste is treated to form mercury selenide. In a closed system, waste is first heated in a rotary furnace to vaporize all mercury at a controlled temperature of 600 to 850oC. Enough selenium is then added to ensure complete mercury conversion into mercury selenide. Finally after the mercury free waste is separate from

the system, the gaseous phase is then cooled down to obtain solid mercury selenide. Again, the process should be conducted without oxygen in the system. [35]

2.3.2 Mercury Stabilization via Amalgamation

As mentioned before, mercury readily forms amalgams with heavy metals just by physically contacting the metals. However, the formed solid amalgams have lower stability and higher solubility in comparison with mercury sulfide and mercury selenide.

2.3.2.1 Amalgamation with Copper

Copper amalgamation can be done by mixing mercury with fine copper powder which is first washed with nitric acid. The ratio between mercury and copper is controlled so that the mixture contains 65% of mercury by weight. The mixture is then milled for a total of 90 minutes, then hardened and later crushed into powder if required. [35]

2.3.2.2 Amalgamation with Zinc

Just like copper amalgamation, zinc amalgams can also be formed by mixing mercury with fine zinc powder which should also be washed with nitric acid. The mixture should contain 45% of mercury by weight and it should be milled for a total of two hours. [35]

2.3.3 Stabilization of Mercury into a Stable and Insoluble Matrix

2.3.3.1 ATG Stabilization Process

The stabilization process was designed by the Allied Technology Group (ATG) to stabilize both solid and liquid mercury containing waste. Mercury waste is first mixed with a sulfur containing immobilizing agent to stabilize mercury. Then clay or cement is added to solidify the product. The mercury-containing waste load in the final product can reach up to 70% by weight. [35]

2.3.3.2 Sulfur Polymer Cement Process

In this process, a sulfur polymer cement (SPC) was developed to form a sulfur polymer matrix which can encapsulate mercury sulfide. Mercury or mercury containing waste is first mixed with SPC at elevated temperature for 4 to 8 hours. A chemical stabilizer such as sodium sulfide may also be added to ensure all mercury compounds are converted into mercury sulfide. The obtained mixture will then be heated to 120 to 150oC until a molten paste is obtained. Additional SPC is applied to increase the viscosity of the molten product and also to make sure that all mercury is converted into the form of mercury sulfide. Finally the molten product can be cast to obtain the final product. It has been confirmed by XRD that both cinnabar and metacinnabar can be observed within the sulfur polymer matrix. It should be noted that the process should be conducted under inert atmosphere or in a vacuum to prevent formation of mercury oxide. [35, 37]

2.3.3.3 Magnesia Binder

In this process developed by Dolomatrix, Australia, toxic waste is first homogenized in water, followed by the addition of Dolocrete® reagents consisting of magnesium oxide and a proprietary mixure of additives. The resultant slurry can then be cast into desirable shapes after complete mixing. The final product has been confirmed to be very stable under TCLP (<0.01 mg/L). [35]

2.3.3.4 Solidification with Hydraulic Cement and Sulfur Polymer Cement

This process was developed by the Southwest Research Institute, USA to treat a broad range of toxic wastes. Wastes are first combined with water to form a 50 wt% slurry. Then it is mixed with pozzolana lime, kiln dust, a hydraulic cement, and other unspecified additives. The mixture is cured at room temperature first and then at 180oC for another 8 hours. The formed product is crushed to pebbles with a particle size between 1/4 to 1/2 inch. They are then combined with sulfur polymer cement, additional pozzolana and sand. After curing, the final concrete-like product is formed with high compression strength. [35]

2.4 Aqueous Chemistry of Mercury

The aqueous chemistry of mercury will be reviewed in this section. The interaction between mercury and hypochlorite, thiosulfate and selenium will especially be discussed in detail.

2.4.1 Hg - H2O Chemistry

-3 The Pourbaix diagrams for the Hg-H2O system at solute concentrations of 1 molal, 10 molal, and 10-6 molal at 25oC are shown in Figures 1 - 3. The diagrams were completed using the HSC 6.1 software with the thermodynamic data in their database.

Figure 1 Pourbaix diagram of Hg-H2O system at 25oC. [Hg] = 1 molal, Pressure = 1atm. (HSC 6.1)

Figure 2 Pourbaix diagram of Hg-H2O system at 298K. [Hg] = 10-3 molal, Pressure = 1atm. (HSC 6.1)

Figure 3 Pourbaix diagram of Hg-H2O system at 298K. [Hg] = 10-6 molal, Pressure = 1atm. (HSC 6.1)

According to the Pourbaix diagrams, metallic mercury is largely stable in the presence of water within the whole pH range between 0 - 14. Thermodynamically, mercury can be oxidized to mercury oxide solids by oxidants like oxygen. However, the kinetics of such a reaction is quite slow. In fact, metallic mercury is very stable when exposed to air or oxygen. It can only be corroded by oxygen very slowly in the presence of moisture and form red mercury oxide [23].

Mercury oxide has a zig-zag structure of O - Hg - O chains. It can appear in red and yellow forms. They both share the exact same structures. The difference in appearance is solely due to the difference in particle size. Red mercury oxide can be produced via hot method by heating the elements at approximately 350oC, while yellow mercury oxide can be produced by precipitating mercury (II) from alkaline solution [6]. Mercury oxide is not a very stable compound in the presence of other elements, and therefore it is very rare in nature.

As shown in the Pourbaix diagrams, mercury is stable in neutral and basic solutions. It is quite stable in most acids too. It can only be dissolved in concentrated sulfuric acid and nitric acid. Mercury has two oxidation states: mercury(I) and mercury(II). Mercuric salts can be formed when mercury is dissolved in excess amount of concentrated nitric acid, or under hot conditions in diluted nitric acid or concentrated sulfuric acid [6, 23]:

Hg + 4HNO3 ↔ Hg(NO3)2 + 2NO2 + 2H2O

On the other hand, in dilute nitric acid at room temperature, mercury can also be slowly dissolved to form mercurous nitrate:

6Hg + 8HNO3 ↔ 3Hg2(NO3)2 +2NO + 4H2O

Unlike other monovalent metal cations such as cuprous ions, mercurous ions exist in the 2+ form of Hg2 in solution. This is due to the fact that mercurous ions tend to form covalent bonds instead of ionic bonds. As a result, Hg+ ions are dimerized and Hg+ − Hg+ 2+ 2+ dimers are formed [22]. Hg2 (aq) can be disproportioned into Hg (aq) and metallic Hg:

2+ 2+ Hg2 ↔ Hg + Hg

The disproportion reaction can occur with the addition of OH- or HS- ions [24]. Thus 2+ Hg2 ions only have a relatively small domain of stability.

According to Figures 2 and 3, when the mercury concentration decreases, the domains of 2+ 2+ stability of both Hg (aq) and Hg2 (aq) enlarge. It should be noted that when the mercury concentration is reduced to as low as 10-6M, instead of HgO(s), aqueous

Hg(OH)2(aq) becomes the stable form in oxidizing and basic conditions. Moreover, in oxidizing environment, Hg(OH)+(aq) appears to be thermodynamically stable at pH value of around 3.5.

It is worth mentioning that even though mercury cannot directly react with elemental carbon [6], the mercury atom has the ability to replace the hydrogen atom from organic compounds to form Hg − C bonds [22]. The inertness of the Hg − C bonds gives organic compounds of mercury, such as methyl mercury, a very strong ability to bioaccumulate in the food chain [25]. These are among the most toxic of mercury compounds and are readily formed in nature.

2.4.2 Hg - Cl - H2O Chemistry

Mercury has the ability to associate with halogens to form halides. White mercuric

chloride (HgCl2) has a structure of linear Cl - Hg - Cl molecules. It is odorless, volatile and soluble in water. However, the covalent bonding between the mercury and chlorine

atoms ensures that mercuric chloride is dissolved in the solution in form of HgCl2 molecules instead of mercuric and chloride ions [6]. In chloride solutions, mercuric chloride can further associate with chloride anions to form the tetrahedral complex 2- (HgCl4) .

On the other hand, mercurous chloride (Hg2Cl2), or dimercury dichloride, or calomel, is also a white, odorless chemical compound. It also has a linear structure. However, unlike mercuric chloride, mercurous chloride is virtually insoluble in water.

o Shown below in Figure 4 is the Eh pH diagram for the Cl-H2O system at 25 C. It can be seen that chloride can be oxidized to perchlorate in oxidizing environment through the entire pH range of 0 ~ 14. This can affect the domains of stability of mercury chloride

species in Hg-Cl-H2O systems (Figures 5 ~ 8).

Figure 4 Pourbaix diagram for the Cl-H2O system at 298K. [Cl] = 1Molal, Pressure = 1atm. (HSC)

o Shown in Figures 5 - 8 are the Pourbaix diagrams for the Hg-Cl-H2O system at 25 C with dissolved mercury and chlorine concentrations of 1M and 10-3M. In comparison with the

Hg-H2O system, Figure 5 indicates that the presence of chloride ions greatly extends the 2- domain of stability of aqueous mercury(II) by forming HgCl2(a) and HgCl4 (a). As 2+ 2- mentioned above, the transition between Hg , HgCl2(a) and HgCl4 (a) is due to the oxidation/reduction of the chlorine species (Figure 5). The domain of stability of mercury(I) is also extended by forming calomel in the presence of chloride.

-3 2- When decreasing the mercury activity to 10 , the domain of stability of HgCl4 (aq) further extends downwards and to higher pH, and overlays most of the domain of

stability of Hg2Cl2(s) due to the overwhelming presence of excess chloride (Figure 6). However, if chloride concentration is drastically reduced to a comparable level (10-3M), 2- HgCl4 (aq) can no longer be observed on the diagram (Figure 7). Aqueous mercury(II)

compound in this case becomes HgCl2(a) with a shrinking domain of stability. However, as suggested in Figure 8, when the presence of chloride species is much less than the mercury level, solid mercuric chloride and mercurous chloride is formed instead of the aqueous species, and as expected, the domain of stability of mercury chlorine compounds further shrinks.

Figure 5 Pourbaix diagram for the Hg-Cl-H2O system at 298K. [Hg] = 1Molal, [Cl] = 1Molal, Pressure = 1atm. (HSC)

-3 Figure 6 Pourbaix diagram for the Hg-Cl-H2O system at 298K. [Hg] = 10 Molal, [Cl] = 1Molal, Pressure = 1atm. (HSC)

-3 -3 Figure 7 Pourbaix diagram for the Hg-Cl-H2O system at 298K. [Hg] = 10 Molal, [Cl] = 10 Molal, Pressure = 1atm. (HSC)

-3 Figure 8 Pourbaix diagram for the Hg-Cl-H2O system at 298K. [Hg] = 1Molal, [Cl] = 10 Molal, Pressure = 1atm. (HSC)

2.4.2.1 Interaction between Mercury and Hypochlorite

Hypochlorite (OCl-) is a strong oxidant. However, hypochlorite is not stable. It can easily decompose into chloride and oxygen. Hypochlorite can react with acid and produce chlorine.

+ - HOCl + H + Cl → Cl2 + H2O

- Hypochlorite can easily oxidize iodide (I ) to iodine (I2), and therefore can be titrated accurately using thiosulfate based on the reactions below.

- - + - OCl + 2I + 2H → I2 + Cl + H2O

2- - 2- I2 + 2S2O3 → 2I + S4O6

Research has already been done to study the ability of hypochlorite solution to dissolve elemental mercury [26-30]. It has been reported that elemental mercury can be absorbed by the hypochlorite solution based on the following reaction:

- + - Hg + OCl + 2H + Cl ↔ HgCl2(a) + H2O

According to Liu et al. (2010), a lower pH will result in a faster mercury extraction. Therefore, it is assumed that the active chlorine concentration in the solution plays a significant role in the oxidation of elemental mercury. As expected, a higher hypochlorite concentration will also lead to a faster extraction rate of mercury. However, temperature is a less decisive factor. It is found that temperature has a slight detrimental impact on mercury extraction possibly due to loss of volatile chlorine gas.

2.4.3 Hg - S - H2O Chemistry

Mercury sulfide has two stable allotropic forms: the most stable red hexagonal cinnabar (α form); and the less stable black cubic metacinnabar (β form) (Figure 9) [6, 9]. Cinnabar is the naturally existing form of mercury sulfide due to its stability. In comparison, the less stable black metacinnabar is rarely found in the nature. In the laboratory, metacinnabar is the better known mercury sulfide form, as it can be produced

by the reaction between aqueous mercury(II) with H2S [6]:

2+ + Hg + H2S → HgS + 2H

Figure 9 The structure of metacinnabar [6]

Mercury sulfide is a very stable compound [6, 24]. Thus it is one of the desired stabilized products for mercury, preferably as cinnabar.

2+ 2- Cinnabar: HgS(s) = Hg (aq) + S (aq) pKsp = 56.4

2+ 2- Metacinnabar: HgS(s) = Hg (aq) + S (aq) pKsp = 51.8

The Hg-S-H2O system is complicated. Researchers [24] have provided a thorough study of this system. The Hg-S-H2O system was calculated with unit activity for dissolved sulfur species and an activity of 0.1 and 10-6 for dissolved mercury (Figures 10 and 11).

Figure 10 Eh-pH diagram for the Hg-S-H2O system at 298K with activities of dissolved mercury and sulfur of 0.1 and 1, respectively. [24]

Figure 11 Eh-pH diagram for the Hg-S-H2O system at 298K with activities of dissolved mercury and sulfur of 10-6 and 1, respectively. [24]

2.4.3.1 Thiosulfate Chemistry

2- Thiosulfate (S2O3 ) is an unstable anion and can be prepared by dissolving elemental sulfur in boiling sulfite solution [13].

2- 2- S + SO3 → S2O3

It is well know that thiosulfate is not stable in acid, forming elemental sulfur and sulfur dioxide:

2- + 0 S2O3 + 2H → S + SO2(g) + H2O

2- + 0 - S2O3 + H → S + HSO3

Thiosulfate can be oxidized to tetrathionate or sulfate depending on the oxidizing ability of the oxidants. A well known example of thiosulfate being oxidized to tetrathionate is the oxidation by iodine. This property plays a significant role in the analytical chemistry

field, for instance the aforementioned titration to determine the hypochlorite concentration.

2- - 2- I2 + 2S2O3 → 2I + S4O6

Thiosulfate has the ability to form complexes with numerous metals including Cu(I), Cd(II), Bi(III), Hg(II), Ag(I) and Au(I) [13]. This property is utilized in the extractive metallurgy field. For example, thiosulfate can be used as an alternative reagent to extract gold [34].

2- 3- - 4Au + 8S2O3 + O2 + 2H2O = 4Au(S2O3)2 + 4OH

2.4.3.2 Interaction between Mercury and Thiosulfate

In the literature and through preliminary experimental work done by Ullah at the University of British Columbia, the complexation of mercury(II) and thiosulfate anion has been proved to be effective [11, 13, 34]:

2+ 2- 2- Hg + 2S2O3 ↔ Hg(S2O3)2

2+ 2- 4- Hg + 3S2O3 ↔ Hg(S2O3)3

However, mercury thiosulfate complexes do not appear in the Pourbaix diagrams shown above. According to the Pourbaix diagrams, mercury sulfide is stable in reducing conditions, but it can be oxidized to form metallic metal and S(VI) species. However, it is believed that such reactions can hardly occur kinetically due to the strong interaction between sulfur and mercury, and the large amount of activation energy involved [24]. Thus, presented in Figure 12 and 13 are another two Eh-pH diagrams calculated under the same conditions with all S(VI) species excluded. As shown in Figure 12, the domain m- of stability of mercury sulfide extends. Meanwhile, Hg(S2O3)n complexes have a small domain of stability when the mercury sulfide is oxidized below pH of 7. On the other hand, when pH is greater than 7, mercury sulfide can be theoretically oxidized to metallic

mercury and S(V) or S(IV) species, even though this is kinetically extremely hard to achieve.

Figure 12 Metastable Eh-pH diagram for the Hg-S-H2O system at 298K with activities of dissolved mercury and sulfur of 0.1 and 1, respectively; S(VI) species excluded from calculation [24]

Figure 13 Metastable Eh-pH diagram for the Hg-S-H2O system at 298K with activities of dissolved mercury and sulfur of 10-6 and 1, respectively; S(VI) species excluded from calculation [24]

In order to further accurately predict the chemical behaviour of mercury in sulfide solutions, researchers [24] have excluded all the sulfur-oxy species and recalculated another Eh-pH diagram, which is shown in Figure 14. It is obvious that the domain of mercury sulfide greatly extends. Polysulfide anions are found within mercury sulfide's domain when the pH is greater than 8.

Figure 14 Metastable Eh-pH diagram for the Hg-S-H2O system at 298K with activities of dissolved mercury and sulfur of 0.1 and 1, respectively; all sulfur-oxy species excluded from calculation [24]

2.4.4 Hg - Se - H2O Chemistry

Mercury selenide is also an extremely stable and unreactive substance. It can be prepared directly from the elements in the gas phase [6, 35]. Tiemannite is the natural existing mineral of mercury(II) selenide. It has a cubic structure as shown below in Figure 15, in which mercury and selenide atoms are tetrahedrally coordinated.

The Pourbaix Diagram of Se-H2O system is shown below in Figure 16. Based on the diagram, selenium can either be oxidized or reduced through the entire pH range.

Figure 15 Structure of mercury selenide

Figure 16 Pourbaix diagram for the Se-H2O system at 298K. [Se] = 1Molal, Pressure = 1atm. (HSC)

The Hg - Se - H2O system is less complicated in comparison with the mercury sulfur system. The Pourbaix diagrams with mercury and selenium activity of 1 and 10-3 are shown below in Figures 17 - 20.

Figure 17 Pourbaix diagram for the Hg-Se-H2O system at 298K. [Hg] = 1Molal, [Se] = 1Molal, Pressure = 1atm. (HSC)

-3 Figure 18 Pourbaix diagram for the Hg-Se-H2O system at 298K. [Hg] = 10 Molal, [Se] = 1Molal, Pressure = 1atm. (HSC)

-3 --3 Figure 19 Pourbaix diagram for the Hg-Se-H2O system at 298K. [Hg] = 10 Molal, [Se] = 10 Molal, Pressure = 1atm. (HSC)

--3 Figure 20 Pourbaix diagram for the Hg-Se-H2O system at 298K. [Hg] = 1Molal, [Se] = 10 Molal, Pressure = 1atm. (HSC)

It can be observed that mercury selenide is the only mercury-selenium species shown on the Pourbaix diagrams. The domain of stability of mercury selenide is not affected much

by the concentration of mercury or selenium in the system. It can also be observed that if the domain of stability of mercury selenide is eliminated from the diagrams, the

remaining lines are consistent with those of the Hg - H2O system.

Selenosulfate solution can be prepared by dissolving selenium powder in hot sulfite solution with vigorous agitation [38]:

2- 2- Se + SO3 → SeSO3

However, no literature can be found regarding the interaction between mercury and selenosulfate. It can only be presumed that selenosulfate will behave similar to thiosulfate in terms of its interaction with mercury(II) salts.

2.5 Research Objectives

A process shall be developed to stabilize elemental mercury as a stable product which will not exceed the UTS limit in TCLP experiments. The idea is to first leach elemental mercury into solution as aqueous mercury (II). Then, thiosulfate or selenosulfate can be used to precipitate mercury (II) as mercury sulfide or mercury selenide. In order to achieve this goal, the following subjects are required to be studied:

1. Examination of the leaching behaviour of elemental mercury: The effectiveness of different leaching reagents including sodium hypochlorite, hydrogen peroxide and cyanide will be investigated. At the same time, factors that can have an impact on the extraction rate will also be studied, including temperature, pH, stirring speed, and reagent concentration. 2. Examination of the precipitation behaviour of mercury (II) salts: Thiosulfate, selenosulfate and selenious acid will be applied to precipitate mercury (II). The formed precipitates shall be examined by XRD and TCLP.

3 EXPERIMENTAL METHODS

This chapter will provide details regarding experimental apparatus, applied chemical reagents, and analytical methods.

3.1 Mercury Leaching Experiments

Different approaches to extract elemental mercury into aqueous solution have been explored, including hypochlorite leaching, hydrogen peroxide leaching and cyanidation.

Figure 21 below shows the experimental setup for the leaching experiment. The reactor is a glass jacketed vessel (from Kontes Glass) and has a total volume of 1 litre. A thermometer (alcohol thermometer from Fisher Scientific) was used to monitor the solution temperature. The temperature of the solution was controlled by a 6-litre circulating water bath (Cole-Parmer polystat heated circulating bath with analog control) attached to the reactors. The pH of the solution was monitored and controlled by using a pH controller (Cole-Parmer pH controller) and a pump (Masterflex L/S variable speed modular drive with Masterflex standard pump head). The reactor was sealed with a rubber stopper. The addition of pH adjustment reagent was recorded by a balance with an accuracy of 0.1 g (Sartorius M-Power Toploader balance). Stirring was achieved by using a magnetic stirrer (Corning Digital Stirrer 5 inch x 7 inch, with a 2 inch magnetic stirring bar). Samples were taken by using syringes and stored in sealed glass vials to avoid oxidation by air.

It should be noted that for those experiments requiring high stirring speed (above 500 RPM), the glass jacketed vessel will not function well since it has a V-shaped bottom. Another type of reactor was used, which was a plastic, flat-bottom reactor. All the other experimental apparatus was the same as described above.

Thermometer pH Probe

Temperature Probe

pH Controller

Water In Water Out

Pumping Device

Magnetic Stirrer Acid/Basic Solution

Balance

3.1.1 Hypochlorite Leaching Experiments

0.5 g of mercury was used in hypochlorite leaching experiments. Hypochlorite solution was prepared by diluting a more concentrated sodium hypochlorite solution (Ricca

Chemical 2.5% (w/w) sodium hypochlorite solution) using de-ionized water. pH of the reaction was adjusted by hydrochloric acid solution (1 mole/L) and sodium hydroxide solution (0.1 molar). Experiments were all started with 500 mL of solution. Elemental mercury was added into the solution after the target temperature and pH was reached.

The majority of the experiments were conducted for a total of 8 hours. During each experiment, samples of size of 10 mL or 4 mL were taken at 15 min, 30 min, 1 hr, 1.5 hr, 2 hr, 3 hr, 4 hr, 6 hr and 8 hr. Samples were diluted and then analyzed by Atomic Absorption Spectrometry (AAS) to determine the concentration of mercury in solution.

Preliminary experiments were first conducted to validate the capability of hypochlorite to extract mercury. More experiments were then conducted to investigate the effect of different parameters on hypochlorite leaching, including pH, temperature, stirring speed, and hypochlorite concentration. Detailed experimental conditions are provided in Table 2.

Table 2 Test conditions of hypochlorite leaching experiments

Sodium ClO- : Hg Mercury Hypochlorite Molar Stirring Dosage Concentration Ratio pH Temperature Speed g/L M oC RPM 1 0.05 10 1 20 500 1 0.05 10 2 20 500 1 0.05 10 3 20 500 Effect of pH 1 0.05 10 4 20 500 1 0.05 10 5 20 500 1 0.05 10 6 20 500 1 0.05 10 4 20 400 Effect of 1 0.05 10 4 20 600 Stirring 1 0.05 10 4 20 800 Speed 1 0.05 10 4 20 1000 1 0.01 2 4 20 1000 Effect of 1 0.02 4 4 20 1000 Hypochlorite 1 0.03 6 4 20 1000 Concentration 1 0.04 8 4 20 1000 1 0.05 10 4 20 1000 1 0.05 10 4 20 500 Effect of 1 0.05 10 4 30 500 Temperature 1 0.05 10 4 40 500 1 0.05 10 4 50 500

It should be noted that in some experiments, samples were titrated in order to determine the consumption of hypochlorite. The titration procedures are described below:

1. Starch indicator solution was prepared by dissolving soluble starch in hot de- ionized water until the solution was clear.

2. Thiosulfate titration solution was prepared by dissolving sodium thiosulfate powder (Alfa Aesar Chemical sodium thiosulfate, 99%, anhydrous) in de-ionized water. 3. Iodide solution was prepared by dissolving potassium iodide powder (Fisher Chemical, Certified ACS) in de-ionized water. 4. 1 mL of hypochlorite containing sample was added into 20 mL of potassium iodide solution (Excess amount of iodide was required). 5. Several drops of concentrated hydrochloric acid were added into the solution right before titration. 6. Titration began by adding thiosulfate solution slowly into the solution. 7. When the colour of the solution faded to pale yellow, several drops of starch indicator were added into the solution. 8. The addition of thiosulfate was stopped when the solution became clear and colourless. 9. Finally the volume of added thiosulfate was recorded.

3.1.2 Hydrogen Peroxide Leaching Experiments

In hydrogen peroxide leaching experiments, 0.5 g of mercury was used in each test. Hydrogen peroxide solution was prepared by diluting a more concentrated hydrogen peroxide solution (Fisher Chemical hydrogen peroxide ACS 30%) using de-ionized water. The molar ratio between mercury and hydrogen peroxide was controlled at 1 : 5. The temperature of the solution was controlled at 20oC and 50oC. Hydrochloric acid was added into the solution to adjust the pH to 1.2. The stirring speed was controlled at 500 RPM. Experiments were all started with an initial 500 mL of solution. Mercury was added into the solution after the target pH was reached.

Experiments were conducted for a total of 24 hours. During each experiment, samples of size 10 mL were taken at 10 min, 20 min, 40 min, 1 hr, 1.5 hr, 2 hr, 3 hr, 4 hr, 6 hr, 21 hr, and 24 hr. Samples were analyzed through AAS to determine the concentration of mercury in solution.

3.1.3 Cyanidation Experiments

Air was blown into the bottom of the reactor by using a pumping device (Masterflex L/S variable speed modular drive with Masterflex standard pump head). 0.5 g of mercury was used in cyanidation experiments. Cyanide solution was prepared by dissolving sodium cyanide powder (Fisher Chemical sodium cyanide ACS) using sodium hydroxide solution. The starting cyanide concentrations were 2.5 g/L and 5 g/L. The temperature of the solution was controlled at 20oC. The pH of the solution was controlled at 12. The stirring speed was controlled at 500 RPM. Mercury was added into the solution after the target pH was reached.

Experiments were conducted for as long as 1 or 2 days. Samples were taken throughout the experiments, and then analyzed by AAS to determine the aqueous mercury concentration.

3.2 Mercury Precipitation Experiments

Three approaches have been explored to precipitate mercury as its stable forms, i.e. cinnabar and mercury selenide, including using thiosulfate, selenosulfate, and selenous acid. Mercury precipitation experiments have exactly the same experimental setup as mercury leaching experiments shown in Figure 21. Temperature and pH of the solution were carefully controlled during the experiments. Samples of 4 mL were taken by syringe, filtered using syringe filters (Fisher Brand 0.22μm PVDF membrane), and then stored in sealed glass vials. At the end of the experiment, precipitates were separated and collected by filtration. They were dried in an oven, and then characterized by using X-Ray Diffraction (XRD).

3.2.1 Thiosulfate Precipitation Experiments

Mercury was added into the reactor in the form of mercury (II) salts: mercuric oxide (Alfa Aesar Chemical mercury (II) oxide, red, 99%) or mercuric chloride (Fisher Chemical mercuric chloride, certified ACS). Thiosulfate solution was prepared by dissolving sodium thiosulfate powder (Alfa Aesar Chemical sodium thiosulfate, 99%,

anhydrous) using de-ionized water. The pH of the reaction was adjusted by hydrochloric acid solution (approximately 1 molar) and sodium hydroxide solution (0.1 molar). Experiments were all started with an initial 500 mL of solution. Mercury salt was added into the solution after the target temperature and pH were reached. Stirring speed was controlled at 500 RPM for all the experiments.

The majority of the thiosulfate precipitation experiments were conducted for a total of 22 hours. During these experiments, samples were taken at 30 min, 1 hr, 2 hr, 4 hr, 6 hr, 8 hr, and 22 hr. Samples were filtered, diluted and then analyzed using AAS.

Experiments were conducted to study the effect of different parameters on thiosulfate precipitation, including pH, temperature, thiosulfate concentration, and seeding. Detailed experimental conditions are listed below in Table 3.

Table 3 Test conditions of thiosulfate precipitation experiments

2- Sodium S2O3 : "Seed" : Mercury Thiosulfate Hg Molar pH Temperature Hg Molar Dosage Concentration Ratio Ratio g/L M oC 2 0.1 10 2 80 N/A 2 0.1 10 4 80 N/A Effect of pH 2 0.1 10 6 80 N/A 2 0.1 10 8 80 N/A 2 0.1 10 10 80 N/A 2 0.1 10 2 20 N/A Effect of 2 0.1 10 2 40 N/A Temperature 2 0.1 10 2 60 N/A 2 0.1 10 2 80 N/A 2 0.1 10 4 20 N/A Effect of 2 0.1 10 4 40 N/A Temperature 2 0.1 10 4 60 N/A 2 0.1 10 4 80 N/A 2 0.1 10 5 20 N/A Effect of 2 0.1 10 5 40 N/A Temperature 2 0.1 10 5 60 N/A 2 0.1 10 5 80 N/A 2 0.1 10 6 20 N/A Effect of 2 0.1 10 6 40 N/A Temperature 2 0.1 10 6 60 N/A at pH 6 2 0.1 10 6 80 N/A 2 0.1 10 5 60 1 Effect of 2 0.1 10 5 60 5 "Seeding" 2 0.1 10 6 60 1 2 0.1 10 6 60 5

3.2.2 Selenosulfate Precipitation Experiments

Sodium selenosulfate was prepared by dissolving selenium powder in sodium sulfite solution. Details are provided in Section 3.3. Mercury salt (mercuric oxide or mercuric chloride) was added into the selenosulfate solution after the target temperature was reached. It should be noted that due to the fact that the precipitation reaction was completed extremely fast, pH was adjusted to 7 by using sodium hydroxide solution and then not controlled after mercury salt was added. Stirring speed was controlled at 500 RPM for all the experiments.

Experiments were started with 500 mL solution for a maximum duration of 4 hours. Samples were taken at 10 min, 20 min, 40 min, 1 hr, 1.5 hr, 2 hr and 4 hr. Samples were filtered, diluted, and then analyzed through AAS to determine the concentration of mercury in solution.

Experiments were conducted to study the effect of temperature and selenosulfate concentration on the precipitation results. Detailed experimental conditions are provided in Table 4 below.

Table 4 Test conditions of selenosulfate precipitation experiments

2- Sodium SeSO3 : Mercury Selenosulfate Hg Molar Stirring Dosage Concentration Ratio pH Temperature Speed g/L M oC RPM 2 0.1 10 Basic 10 500 2 0.1 10 Basic 20 500 Effect of 2 0.1 10 Basic 40 500 Temperature 2 0.1 10 Basic 60 500 2 0.1 10 Basic 80 500 2 0.02 2 Basic 20 500 Effect of 2 0.04 4 Basic 20 500 Selenosulfate 2 0.06 6 Basic 20 500 Concentration 2 0.08 8 Basic 20 500 2 0.1 10 Basic 20 500

3.2.3 Selenious Acid Precipitation Experiments

In selenious acid precipitation experiments, mercuric chloride was first added into the solution as the source of mercury (II). Selenious acid was added into the solution when mercuric chloride was completely dissolved. Hydrochloric acid was added into the solution to control the pH at 2. Stirring speed was controlled at 500 RPM. Experiments were conducted at 20oC and 80oC. The initial molar ratio between mercury and selenious acid was 1 : 10. Experiments were conducted for one day. Samples were taken for AAS to examine the aqueous concentration of mercury.

3.3 Selenium Dissolution Experiments

Selenium dissolution experiments were conducted to obtain selenosulfate solution for selenosulfate precipitation experiments. Figure 22 below shows the experimental setup for the selenium dissolution experiments. Temperature was roughly controlled at 90 ~ 100oC. Stirring and heating was achieved by using a magnetic stirring hot plate (Cole 40

Parmer). pH of the reaction was neither adjusted nor controlled throughout the experiments. Nitrogen was blown into the solution to de-aerate the reaction environment.

Thermometer

Heating Magnetic Stirrer

Figure 22 Experimental setup for selenium dissolution experiments

Sulfite solution was prepared by dissolving sodium sulfite powder (Fisher Chemical sodium sulfite certified ACS) using de-ionized water. After the sulfite solution was heated to 90oC, selenium powder (Alfa Aesar Chemical selenium powder, 325 mesh, 99.5%) was added into the reactor. Stirring speed was controlled at above 1000 RPM.

After the experiment was completed, the solution was cooled down to room temperature and then transferred into a 500 mL volumetric flask.

For those experiments whose goal was to prepare selenosulfate solution for precipitation experiments, experiments took place in a 500 mL volumetric flask. Experiments were stopped when clear solution was obtained. Then the solution was transferred into a glass bottle. Nitrogen was used to remove the air inside the bottle. It was then stored in a fridge for further use. In such experiments, the molar ratio between sulfite and selenium was controlled to be four to achieve complete selenium dissolution.

Several experiments were conducted to study the effect of sulfite concentration on selenium dissolution. Such experiments took place in a 500 mL beaker. They were run for a total of 2 hours. One sample of 10 mL was taken at the end of the experiment. It was then filtered and analyzed by ICP to determine the concentration of selenium in solution. The detailed experimental conditions are listed below in Table 5.

Table 5 Experimental conditions for selenium dissolution experiments

2- Se:SO3 Selenium Added Sodium Sulfite Added Molar Ratio g mole g mole 4 1.974 0.025 12.604 0.1 2 1.974 0.025 6.302 0.05 1.5 1.974 0.025 4.7265 0.0375 1.25 1.974 0.025 3.939 0.03125 1 1.974 0.025 3.151 0.025

3.4 Solid Waste Disposal Characterization

Solid waste disposal characterization experiments were conducted to examine the mobility of mercury in the formed precipitates. However, since the experimental procedure was not exactly following the standard TCLP, the results can not be used for regulatory purposes.

The extraction fluid was prepared by diluting the mixture of 5.7mL anhydrous glacial acetic acid and 64.3mL 1.00M sodium hydroxide solution to 1000mL using de-ionized water. The pH of the extraction fluid was measured to be 4.90. In each experiment, 0.5 to 1 gram of the as received formed precipitates were used, and extraction fluid was added to ensure a 20 : 1 liquid to solid mass ratio. Experiments were all conducted at room temperature. Stirring speed was controlled at 50 RPM using a stainless steel overhead stirrer.

Each experiment was conducted in a beaker at room temperature for a total of 18 hours. A plastic cover was used to minimize the effect of water evaporation. Filtration was done by using syringe filters at the end of the experiment. The filtered solution was then examined by the Cold Vapour Atomic Absorption Spectrometry (CVAAS).

Listed below are the precipitation experiments whose precipitates were examined in the solid waste disposal characterization experiments.

Table 6 Solid waste disposal characterization experiments

Test No. Precipitation Conditions 1 Cinnabar (Certified ACS) 2 Thiosulfate ppt, 80oC, pH2 3 Thiosulfate ppt, 80oC, pH4 4 Thiosulfate ppt, 80oC, pH6 5 Selenosulfate ppt, 10oC 6 Selenosulfate ppt, 40oC 7 Selenosulfate ppt, 60oC 8 Selenosulfate ppt, 80oC

4 RESULTS AND DISCUSSION

In this chapter, experimental results will be reported, analyzed and discussed in detail.

4.1 Hypochlorite Leaching Experiments

In hypochlorite leaching experiments, several series of experiments were conducted to study the effect of pH, temperature, stirring speed, hypochlorite concentration, and chloride concentration on the mercury extraction rate. The goal was to find the optimum leaching conditions to extract mercury effectively and efficiently.

4.1.1 Effect of pH

First of all, the effect of pH on hypochlorite leaching was studied. As mentioned before, hypochlorite anion can react with protons to produce chlorine and water. Then chlorine oxidizes mercury to soluble mercury (II) chloride.

- + - ClO + 2H + Cl ↔ Cl2 + H2O

Cl2 + Hg → HgCl2(a)

According to the reaction above, it was expected that mercury leaching in hypochlorite media would be favoured in acidic conditions, as excessive protons will lead to the formation of chlorine molecules which will act as the direct oxidant of mercury. The results of mercury extraction rate over time are reported in Figure 23 below.

Effect of pH on Hypochlorite Leaching of Mercury

100%

90%

80%

70% pH=1 60% pH=2 pH=3 50% pH=4 40% pH=5 30% pH=6 % Mercury Extracted Mercury % 20%

10%

0% 0 100 200 300 400 500 600 Time (min)

Figure 23 Effect of pH on hypochlorite leaching of mercury. Test Conditions: 20oC, 500RPM, hypochlorite concentration 10 times as much as stoichiometrically required.

During each experiment, the pH of the reaction tended to increase due to the consumption of protons and thus hydrochloric acid solution was used to control the pH. According to Figure 23, mercury leaching was not successful when pH was too high. Only less than 20% of mercury was extracted into the solution when pH was above 5, even when the hypochlorite concentration was substantially high. However, mercury leaching was fastest when pH was 4, instead of the expected 1. In fact, for the experiments conducted under pH 1, 2, 3 and 4, the results were quite irregular. It is observed that following the fastest extraction at pH 4 was pH 2, 1, and 3. This is probably due to the fact that since only half gram of mercury was dosed in each experiment, liquid mercury droplet suspension and breakage could vary significantly from time to time, especially when the stirring speed was set at 500 RPM, which was not fast enough to break the mercury droplets completely.

4.1.2 Effect of Stirring Speed

Stirring speed is also a very important factor on mercury leaching. Mercury is expected to be extracted faster at a higher stirring speed. This is due to the fact that a higher stirring speed will provide a higher surface area of reaction, which will lead to faster kinetics. Experimental results are provided below in Figure 24.

Effect of Stirring Speed on Hypochlorite Leaching of Mercury

100%

90%

80%

70%

60% 400RPM 600RPM 50% z 800RPM 40% 1000RPM

30% % Mercury Extracted % Mercury

20%

10%

0% 0 100 200 300 400 500 600 Time (min)

Figure 24 Effect of stirring speed on hypochlorite leaching of mercury. Test conditions: 20oC, pH = 4, hypochlorite concentration 10 times as much as stoichiometrically required.

As expected, a higher stirring speed leads to a faster extraction speed. Especially, there is a jump from the 800 RPM curve to 1000 RPM curve. At 800 RPM, mercury extraction reached approximately 75% after 8 hours, while it only took less than 3 hours when the stirring speed was set at 1000 RPM. This proves the significant role of agitation in mercury extraction.

4.1.3 Effect of Hypochlorite Concentration

Obviously, a higher hypochlorite concentration will yield a faster extraction rate. As shown in Figure 25, experimental results match the expectation.

Effect of Hypochlorite Concentration on Mercury Leaching

100% 90% 80% 70% 2 Times 60% 4 Times 50% 6 Times 40% 8 Times 10 Times 30% % Mercury Extracted Mercury % 20% 10% 0% 0 100 200 300 400 500 600 Time (min)

Figure 25 Effect of hypochlorite concentration on mercury leaching. Test conditions: 20oC, pH = 4, 1000 RPM, hypochlorite concentration N times as much as stoichiometrically required as indicated.

In this series of experiments, samples were titrated to measure the concentration of hypochlorite, so that hypochlorite consumption can be determined in each experiment. The results are summarized below in Figure 26.

Hypochlorite Consumption

0.05 0.045 0.04

0.035 2 Times 0.03 4 Times 0.025 6 Times 0.02 8 Times 0.015 10 Times 0.01

Hypochlorite Concentration0.005 (M) 0 0 100 200 300 400 500 600 Time (min)

Figure 26 Measured hypochlorite consumption in mercury leaching

It should be noted that in each experiment the measured hypochlorite consumption is higher than its theoretical value calculated based on the amount of extracted mercury provided in Figure 25. Shown below in Table 7 is the comparison of final hypochlorite consumption between its theoretical value and measured value. The difference between the theoretical and measured value is mainly due to the loss of chlorine gas during the experiments. It can be observed that the chlorine loss is more severe when the initial hypochlorite concentration is higher.

Table 7 Comparison of theoretical and measured hypochlorite consumption

Initial Hypochlorite Consumption Initial NaClO : Hypochlorite Hg Molar Ratio Theoretical Measured Differences Concentration M M M M % 0.01 2 0.002704 0.002788 0.000084 3.11% 0.02 4 0.003838 0.005483 0.001645 42.86% 0.03 6 0.003169 0.007587 0.004418 139.41% 0.04 8 0.003982 0.009131 0.005150 129.34% 0.05 10 0.004918 0.012784 0.007867 159.97%

4.1.4 Effect of Temperature

The effect of temperature was also investigated by conducting experiments at 20oC, 30oC, 40oC and 50oC. Results are shown in Figure 27.

Effect of Temperature on Hypochlorite Leaching of Mercury

100%

90%

80%

70%

60% 20C 30C 50% 40C 40% 50C 30% % Mercury Extracted 20%

10%

0% 0 100 200 300 400 500 600 Time (min)

Figure 27 Effect of temperature on hypochlorite leaching of mercury. Test conditions: pH = 4, 500RPM, hypochlorite concentration 10 times as much as stoichiometrically required.

From Figure 27, it is observed that high temperature is detrimental to the hypochlorite leaching process at temperature range of 20oC - 50oC. Tables 8 and 9 below give thermodynamic data for the leaching and oxidation reaction.

Oxidation Reaction: Hg + Cl2 → HgCl2(a)

- + - Overall Leaching Reaction: Hg + ClO + 2H + Cl → HgCl2(a) + H2O

From Tables 8 and 9, it is observed that the enthalpy change (ΔH) for both reactions is negative, which implies that both reactions are exothermic. Thus a raise in temperature will result in the equilibrium to shift backwards. However, the equilibrium constant (K) in each case is so large that back reaction is unlikely. Thus, the increasing hypochlorite loss with higher temperature might be the reason why a higher temperature leads to a lower extraction rate. However, further study is required in order to fully understand the kinetic behaviour of mercury leaching in hypochlorite media.

Table 8 Thermodynamic data for the oxidation reaction (HSC Database)

Hg + Cl2(a) = HgCl2(a) T ΔH ΔS ΔG K Log(K) oC kJ J/K kJ 20 -215.444 -98.314 -186.623 1.803E+033 33.256 30 -212.931 -89.882 -185.683 9.934E+031 31.997 40 -210.824 -83.041 -184.820 6.781E+030 30.831 50 -208.984 -77.255 -184.019 5.594E+029 29.748

Table 9 Thermodynamic data for the overall leaching reaction (HSC Database)

- + - ClO (a) + 2H (a) + Cl (a) + Hg = HgCl2(a) + H2O T ΔH ΔS ΔG K Log(K) oC kJ J/K kJ 20 -252.308 -11.807 -248.847 2.209E+044 44.344 30 -246.401 8.012 -248.830 7.560E+042 42.879 40 -240.950 25.706 -249.000 3.449E+041 41.538 50 -235.749 42.056 -249.339 2.028E+040 40.307

4.1.5 Randomness and Errors

As mentioned above in Section 4.1.1, the breakage of mercury droplets can vary significantly, especially when a small dosage of mercury is applied in leaching experiments. Therefore, randomness had an important role in hypochlorite leaching. Shown below in Figure 28 is the result of two experiments sharing the exact same experimental conditions. It can be observed that even under identical experimental conditions, the extraction rate can vary significantly. The final extraction of the two experiments has a difference of about 25%. The effect of randomness can also be observed in the experimental results in the other sections. It implied that the results were compromised by the varying surface area of liquid mercury under the stirring conditions

applied. This is a difficult problem experimentally. A sealed flask with a magnetic stirrer was used for mixing. It is recommended that future work employ more standard mixing systems.

Randomness

1 0.9 0.8 0.7 0.6 Test 1 0.5 Test 2 0.4 0.3

% Mercury Extracted % Mercury 0.2 0.1 0 0 100 200 300 400 500 600 Time (min)

Figure 28 Effect of Randomness on hypochlorite leaching of mercury. Test conditions: temperature at 20oC, pH = 4, 500RPM, hypochlorite concentration 10 times as much as stoichiometrically required.

4.2 Other Types of Mercury Leaching Experiments

As mentioned in the previous chapter, another two approaches have been explored to extract elemental mercury, which are hydrogen peroxide leaching and cyanidation.

4.2.1 Hydrogen Peroxide Leaching Experiments

Hydrogen peroxide was also used to approach mercury extraction based on the following reaction.

+ - H2O2 + Hg + 2H + 2Cl → HgCl2(a) + 2H2O

Preliminary experiments were conducted to test the feasibility of this method. The idea is to provide experimental conditions that favour the reaction to assure that mercury will be extracted if possible. Detailed experimental conditions and procedures are provided in the previous chapter. Results are summarized below in Table 10.

Table 10 Hydrogen peroxide leaching results

H2O2 : Hg Temperature Test Duration Hg Extraction Molar Ratio oC hr % 5 20 24 0.63% 5 50 8 0.82%

It can obviously be concluded that mercury extraction was not successful at all even when the experimental conditions seem to favour the leaching reaction. Therefore, it can be concluded that hydrogen peroxide is not an effective oxidant for elemental mercury.

4.2.2 Cyanidation Experiments

Cyanide was also used to extract elemental mercury based on the following reaction:

- - Hg + 2CN + 1/2O2 + H2O → Hg(CN)2(a) + 2OH

As mentioned in the previous chapter, a high concentration of sodium cyanide solution was kept in preliminary experiments at 2.5 g/L and 5 g/L. Furthermore, air was blown into the solution in one of the experiments. Results are shown in Table 11.

Table 11 Results of cyanidation experiments

Mercury Sodium Cyanide Concentration Air Mercury Dosage Duration Extracted g/L Molar g/L Molar Days % 2.5 0.051 No 1 0.00499 2 2.83% 5 0.102 No 1 0.00499 2 1.68% 5 0.102 Yes 1 0.00499 1 4.36%

In each experiment the pH was controlled at 12, but only very little mercury was extracted in the end. It was not practical to raise the temperature in cyanidation to improve the kinetics, due to the increasing possibility of hydrogen cyanide and mercury volatilization. Therefore, based on the results of preliminary experiments, cyanidation is not an effective method to extract mercury into aqueous solution.

4.2.3 Summary

As described above, even under favourable experimental conditions, only poor extraction rate can be achieved for both extraction methods. Therefore, it can be concluded that neither hydrogen peroxide nor cyanide can be used to extract mercury effectively.

4.3 Thiosulfate Precipitation Experiments

Thiosulfate can be used as a source of sulfur to precipitate mercury (II) as mercury sulfide based on the reactions described below:

2- 2- - HgO + 2S2O3 + H2O → Hg(S2O3)2 + 2OH

2- 2- 2- + Hg(S2O3)2 + H2O → HgS + SO4 + S2O3 + 2H

Or,

2+ 2- 2- Hg + 2S2O3 + H2O → Hg(S2O3)2

2- 2- 2- + Hg(S2O3)2 + H2O → HgS + SO4 + S2O3 + 2H 54

The desired product is the stable allotropic form of mercury sulfide, cinnabar. Ullah (2012) did some preliminary experiments on mercury precipitation in sodium thiosulfate media. This section will briefly review Ullah's work and provide further details about the effect of different factors on thiosulfate precipitation, such as pH, temperature, thiosulfate concentration and "seeding".

4.3.1 Preliminary Study

Ullah (2012) investigated mercury precipitation mainly at pH 5 and 6 with a relatively high initial mercury concentration of 10 g/L. According to his work, precipitates produced at pH 5 and 6 at 80oC were mainly cinnabar, with a small portion of metacinnabar (less than 10%).

The production of the less desirable metacinnabar could be eliminated when the precipitation took place at room temperature. Unfortunately, it takes an unacceptably long time (weeks) for near complete mercury removal to occur. In his work, Ullah also studied the effect of thiosulfate concentration on mercury precipitation, with the conclusion that a higher thiosulfate concentration leads to a faster and more complete removal of mercury.

Therefore, the goal here is to find the proper experimental conditions to gain pure cinnabar production within a reasonable period of time, meanwhile achieving near complete mercury removal. It is worth mentioning that a lower mercury concentration of 2 g/L is investigated in this work.

4.3.2 Effect of pH

Effect of pH was first investigated. The series of experiments were conducted at 80oC to achieve relatively fast precipitation rate. Results are shown in Figure 29.

Effect of pH on Precipitation Rate

100%

90% 80% 70% pH=2 60% pH=4 50% pH=6

40% pH=8 pH=10 30% 20%

Percent Mercury in Solution(%) Mercury Percent 10% 0% 0 200 400 600 800 1000 1200 1400 Time (min)

Figure 29 Effect of pH on precipitation rate in thiosulfate precipitation experiments. Test conditions: 80oC, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required.

It is obvious that the precipitation rate is much faster when the environment is more acidic. This is due to the fact that thiosulfate anions are not stable in acidic conditions. With the presence of protons, thiosulfate anions tend to break down more easily and provide a reactive sulfur source for mercury sulfide formation. Complete mercury removal was reached within 4 hours of reaction when pH was 4 or lower. When the pH was raised to 6, the precipitation rate slowed down significantly, but near complete mercury removal was still achieved within one day of reaction.

During each experiment, it was observed that final precipitates were all in black colour. It should be noted that for the experiment conducted at pH 2, it was observed that white precipitates were formed immediately after the pH was adjusted to 2. This is because of the fact that at low pH, thiosulfate will dissociate and precipitate elemental sulfur which can appear as a fine white solid. This was confirmed by the X-Ray pattern.

The precipitates formed at pH 2, 4, 5 and 6 were collected and examined by X-Ray Diffraction. From the X-Ray pattern, it can be concluded that at pH 2, the formed 56

precipitate is a combination of metacinnabar and sulfur (Figure 30). At pH 4, the formed precipitate is pure metacinnabar (Figure 31). Cinnabar starts to form when the pH increases to 5 (Figure 32). At pH 6, the formed precipitate is still a mixture of metacinnabar and cinnabar (Figure 33). Comparing with pH 5, at pH 6, the characteristic peaks of cinnabar become much stronger, which suggests a higher composition of cinnabar. In fact, at pH 6, the characteristic peaks of cinnabar seem to be even higher than those of metacinnabar. The precipitates formed at pH 8 and 10 were not examined because there were not enough precipitates collected at the end of these two experiments.

Figure 30 XRD pattern for the formed precipitates at test conditions: pH = 2, 80oC, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required.

Figure 31 XRD pattern for the formed precipitates at test conditions: pH = 4, 80oC, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required.

Figure 32 XRD pattern for the formed precipitates at test conditions: pH = 5, 80oC, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required. 58

Figure 33 XRD pattern for the formed precipitates at test conditions: pH = 6, 80oC, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required.

4.3.3 Effect of Temperature

Temperature functions as a very important factor in mercury precipitation in thiosulfate media. As mentioned before in section 4.2.1, mercury precipitation is extremely slow at room temperature, which is unacceptable. Therefore, several series of experiments were conducted to investigate the effect of temperature at pH of 2, 4, 5 and 6. The results are provided below in Figure 34 - 37.

Effect of Temperature at pH 2

100% 90%

80%

70%

60% 20C 40C 50% 60C 40% 80C 30% 20% Percent Mercury in Solution (%) 10% 0% 0 200 400 600 800 1000 1200 1400 Time (min)

Figure 34 Effect of temperature on mercury precipitation at pH 2. Test conditions: pH = 2, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required.

Effect of Temperature at pH 4

100%

90%

80%

70%

60% 20C 40C 50% 60C 40% 80C 30%

20% Percent Mercury inSolution (%) 10%

0% 0 200 400 600 800 1000 1200 1400 Time (min)

Figure 35 Effect of temperature on mercury precipitation at pH 4. Test conditions: pH = 4, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required. 60

Effect of Temperature at pH 5

100%

90%

80%

70%

60% 20C 40C 50% 60C 40% 80C 30%

20% Percent Mercury inSolution (%) 10%

0% 0 200 400 600 800 1000 1200 1400 Time (min)

Figure 36 Effect of temperature on mercury precipitation at pH 5. Test conditions: pH = 5, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required.

Effect of Temperature at pH 6

100%

90% 80%

70%

60% 20C 40C 50% 60C 40% 80C 30%

20% Percent Mercuryin Solution (%) 10%

0% 0 200 400 600 800 1000 1200 1400 Time (min)

Figure 37 Effect of temperature on mercury precipitation at pH 6. Test conditions: pH = 6, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required. 61

It is obvious that the precipitation rate significantly increases with temperature. It is observed that when pH is equal or below 4, temperature of 60oC can still yield a fast enough precipitation rate. When pH increases to 5 and 6, however, only at near complete mercury removal can only occur at 80oC.

4.3.4 Effect of "Seeding"

In this series of experiments, fine cinnabar powder was added into the solution initially as "seed", with the purpose of increasing the precipitation rate at lower temperatures. However, due to the fact that mercury precipitation was extremely slow at room temperature, experiments were conducted to study if "seeding" can accelerate the precipitation rate at pH 5 and 6 and temperature of 60oC.

Results are provided below in Figures 38 and 39. It can be observed that the precipitation rate at pH 5, 60oC was enhanced by seeding. However, seeding has no impact on the precipitation rate at pH 6, 60oC.

Effect of Seeding at pH 5

100%

90%

80%

70%

60% pH5 No Seed 50% pH5 with Seed 40%

30%

20% Percent Mercury in Solution (%) 10%

0% 0 200 400 600 800 1000 1200 1400 Time (min)

Figure 38 Effect of seeding at pH 5. Test conditions: 60oC, pH = 5, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required. 62

Effect of Seeding at pH 6

100%

90%

80%

70%

60% pH6 No Seed 50% pH6 with Seed 40%

30%

20% Percent Mercury in Solution (%) 10%

0% 0 200 400 600 800 1000 1200 1400 Time (min)

Figure 39 Effect of seeding at pH 6. Test conditions: 60oC, pH = 5, 500RPM, thiosulfate concentration 10 times as much as stoichiometrically required.

4.4 Selenium Dissolution Experiments

The goal of selenium dissolution experiments is to prepare selenosulfate solution for the selenosulfate precipitation experiments in Section 4.5. Selenium was dissolved in sulfite solution according to the following reaction.

2- 2- Se + SO3 ↔ SeSO3

Ullah (2012) has conducted some experiments to investigate selenium dissolution behaviour. Experimental conditions and results are provided below in Table 12.

Table 124 Preliminary selenium dissolution experiments done by Ullah (2012)

Sodium Selenium Sulfite Selenium Time Temperature Dissolved g g hr oC % 1 0.05 4 50 40% 1 0.05 8 50 45% 1 0.05 24 60 70% 1 0.01 24 70 68% 1.5 0.05 24 60 62% 2 0.1 48 60 60%

The detailed experimental conditions, however, were not specified by the author, for example, the total volume of the solution, the stirring speed and how the determination of selenium concentration was carried out. From the table, it is observed that selenium dissolution was not successfully achieved under the experimental conditions above. The tendency of higher sulfite concentration and higher temperature leading to a better selenium dissolution can still be observed.

As mentioned in Chapter 3, most experiments were conducted in a 500 mL volumetric flask with very fast stirring (over 1000 RPM) at temperature between 90 to 100oC. In these experiments, the initial molar ratio between sulfite and selenium was controlled at 4.

It was observed that selenium was completely dissolved within 30 minutes and a clear solution of pale yellow colour was produced in each experiment. Then the solution was stored at low temperature without the presence of oxygen for further use.

Another series of experiments were conducted to study the effect of sulfite concentration on selenium dissolution. Each experiment was conducted for a total of 2 hours, and then analyzed for selenium concentration in solution. Experimental conditions for the series of experiments were provided in Table 13.

Table 13 Results for selenium dissolution experiments

% Selenium 2- SO3 : Se Measured Selenium Concentration Dissolved Molar Ratio Mole/L g/L % 4 0.0479 3.78 96.2% 2 0.0361 2.852 72.6% 1.5 0.0193 1.521 38.7%

It can be observed from the experiments that when sodium sulfite was 4 times as much as required stoichiometrically, selenium powder can be dissolved very fast within 10 minutes. When sodium sulfite was 2 times as much, selenium powder was dissolved within 30 minutes. When sodium sulfite dosage was decreased to as much as 1.5 times or below, selenium dissolution became harder. Complete dissolution could not be achieved within 2 hours of reaction. After the solution was filtered, transferred into a glass bottle, and stored in a fridge, massive red selenium precipitates were formed within several hours. Even for the "2 Times" solution, red precipitates were observed after one day of storage, so the solution has to be filtered again in order to obtain clear samples for analysis.

4.5 Selenosulfate Precipitation Experiments

The inspiration of using selenosulfate to precipitate mercury as mercury selenide comes from the success of thiosulfate precipitation experiments. The possible precipitation chemistry are described below:

2- 2- - HgO + 2SeSO3 + H2O → Hg(SeSO3)2 + 2OH

2- 2- 2- + Hg(SeSO3)2 + H2O → HgSe + SO4 + SeSO3 + 2H

However, since selenosulfate salts are apparently unstable at lower temperature, they cannot be purchased elsewhere. Therefore, selenosulfate solutions were prepared in the laboratory as described in the previous section. Furthermore, the effect of temperature

and selenosulfate concentration on mercury precipitation was investigated and the results are presented below.

4.5.1 Preliminary Experiment

One experiment was conducted to test the feasibility of precipitating mercury using aqueous selenosulfate solution. In this experiment, sodium selenosulfate solution was prepared 10 times more concentrated than required stoichiometrically. The reaction took place at 80oC to achieve fast kinetics. Experiments were conducted for a total of 4 hours. Results show that mercury is precipitating extremely rapidly. Atomic Absorption Spectrometry showed that mercury was precipitated completely within the first 10 minutes of the experiment, which proved the feasibility of this method.

4.5.2 Effect of Temperature

The effect of temperature was first investigated. In all the experiments, large amounts of black precipitates were formed immediately after mercury salts were added. As summarized in Table 14, the kinetics of the reaction was very fast even when it was at room temperature. The formed precipitates were all confirmed to be pure tiemannite by using X-Ray Diffraction (Figures 40 - 42), except at 80oC, a great amount of selenium was also precipitated along with tiemannite (Figure 43).

Table 14 Effect of temperature on selenosulfate precipitation experiments

Temperature Time when more than 97% of mercury is precipitated oC min 80 < 10 60 < 10 40 < 10 20 < 20 10 < 120

Figure 40 XRD pattern of formed precipitates at 10oC

Figure 41 XRD pattern of formed precipitates at 40oC

Figure 42 XRD pattern of formed precipitates at 60oC

Figure 43 XRD pattern of formed precipitates at 80oC

4.5.3 Effect of Selenosulfate Concentration

The effect of selenosulfate concentration on mercury precipitation was also investigated. It was observed during the experiments that the formed mercury selenide precipitates became much more fine when selenosulfate concentration was lower. For instance, samples taken at 10 minutes and 20 minutes could not be effectively filtered (pore size of 0.22 μm) in the "2 Times" experiment because the precipitates were too fine.

Effect of Selenosulfate Concentration

100.00%

90.00%

80.00%

70.00%

60.00% 2X 4X 50.00% 6X 40.00% 8X 30.00%

20.00% Percent Mercury in Solution (%) Solution in Mercury Percent 10.00%

0.00% 0 20 40 60 80 100 120 140 Time (min)

Figure 44 Effect of selenosulfate concentration on mercury precipitation experiments, Test conditions: 20oC, 500RPM, selenosulfate concentration 10 times as much as stoichiometrically required.

According to Figure 44, mercury precipitation using selenosulfate was proved to be very effective and efficient even when a relatively low selenosulfate concentration was used at room temperature.

4.6 Selenious Acid Precipitation Experiments

There were no precipitates formed at the end of each experiment at 20oC and 80oC. The solution analysis via AAS also proved that no mercury was precipitated at all. Therefore, it can be concluded that selenious acid cannot function as a selenium source to precipitate mercury, at least under the conditions tested.

4.7 Solid Waste Disposal Characterization Experiments

Shown below in Table 15 are the results for the solid waste disposal characterization experiments. It can be observed that all samples were below the Universal Treatment Standard of mercury (0.025 ppm or 25 ppb), especially the mercury selenide precipitates presented excellent immobility.

Table 15 Solid waste disposal characterization experiment results

Test No. Precipitation Conditions Aqueous Hg Concentration (ppb) 1 Cinnabar (Certified ACS) 19 2 Thiosulfate ppt, 80oC, pH2 9 3 Thiosulfate ppt, 80oC, pH4 10 4 Thiosulfate ppt, 80oC, pH6 8 5 Selenosulfate ppt, 10oC 2 6 Selenosulfate ppt, 40oC 1 7 Selenosulfate ppt, 60oC <1 8 Selenosulfate ppt, 80oC <1

5. CONCLUSION

The following conclusion can be drawn based on the research in this work:

1. Elemental mercury leaching can be achieved by using hypochlorite solution. A higher agitation speed, higher hypochlorite concentration and lower pH favour the mercury extraction. 2. Alternative leaching of elemental mercury by using hydrogen peroxide and cyanide were found to be ineffective. 3. Thiosulfate can be used to precipitate mercury as mercury sulfide. A higher temperature, lower pH and higher thiosulfate concentration can lead to a faster precipitation. However, a higher pH and lower temperature favours the formation of cinnabar. 4. Selenium can be completely dissolved in excessive sulfite solution at above 90oC with vigorous agitation. 5. Selenosulfate can be used to precipitate mercury as mercury selenide. The precipitation speed was found to be very fast under nearly all experimental conditions. The formed precipitates were confirmed to be tiemannite. 6. Solid waste disposal characterization experiments confirmed that none of the formed precipitates exceeded the UTS limit for mercury. 7. Selenious acid was not effective to precipitate mercury as mercury selenide. 8. A mercury stabilization process via thiosulfate precipitation has been proposed, which is shown below in Figure 45. First, elemental mercury is leached in excessive hypochlorite solution with vigorous agitation. Then, the leachate undergoes a pH adjustment by adding sodium hydroxide to precipitates mercury (II) as mercuric oxide. After solid liquid separation, the aqueous solution containing sodium hypochlorite can be recycled back to the hypochlorite leaching stage, while the solid mercuric oxide is brought to a second leaching process using sodium thiosulfate solution. The leachate containing mercury thiosulfate complex will finally undergo a precipitation stage to produce mercury sulfide.

The filtered liquid solution can also be partially recycled to the thiosulfate leaching stage.

Sodium Elemental Hypochlorite Mercury Solution Hydrochloric Acid

Leaching Tailing Treatment

Sodium Hydroxide Leachate

pH Adjustment

Solid Liquid

Separation Liquid solution

Sodium Thiosulfate Mercury Oxide

Leaching

Tailing Treatment Hydrochloric Acid

Precipitation

Solid Liquid Separation Liquid solution

Mercury Sulfide

Figure 45 Flowsheet of mercury stabilization process via thiosulfate precipitation

9. Shown below in Figure 46 is the mercury stabilization process via selenosulfate precipitation. Just like the previous process, hypochlorite solution is first applied to leach elemental mercury into aqueous solution, and mercuric oxide is obtained by pH adjustment. Then selenosulfate solution is prepared by dissolving elemental selenium in excessive sulfite solution at elevated temperature with vigorous agitation. It is then introduced to the second leaching process to dissolve mercuric oxide and precipitate mercury selenide. The waste liquid solution containing sulfite can be partially recycled back to selenium dissolution.

Sodium Elemental Hypochlorite Mercury Solution

Hydrochloric Acid

Leaching Tailing Treatment

Elemental Sodium Hydroxide Selenium Leachate Sodium pH Adjustment Sulfite Solution

Solid Liquid Leaching Separation Liquid solution

Leachate Mercury Oxide

Leaching

Leachate

Precipitation

Liquid solution Solid Liquid Separation

Tailing Treatment Mercury Selenide

Figure 46 Flowsheet of mercury stabilization process via selenosulfate precipitation

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