mercury stabilization using thiosulfate or selenosulfate - Open ...

MERCURY STABILIZATION USING THIOSULFATE OR

SELENOSULFATE

by

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

in

The Faculty of Graduate Studies

(Materials Engineering)

THE UNIVERSITY OF BRITISH COLUMBIA

(Vancouver)

April, 2013

© Zizheng Zhou, 2013

ii

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

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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

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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

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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

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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


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

4.4 Selenium Dissolution Experiments......................................................................... 63

4.5 Selenosulfate Precipitation Experiments ................................................................ 65

4.5.1 Preliminary Experiment .................................................................................. 66


4.5.3 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

ix

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

Figure 6 Pourbaix diagram for the Hg-Cl-H2O system at 298K. [Hg] = 10-3Molal, [Cl] =


Figure 7 Pourbaix diagram for the Hg-Cl-H2O system at 298K. [Hg] = 10-3Molal, [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

x

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


dissolved mercury and sulfur of 10-6 and 1, respectively; S(VI) species excluded from

calculation [24] ................................................................................................................. 25


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] =


Figure 18 Pourbaix diagram for the Hg-Se-H2O system at 298K. [Hg] = 10-3Molal, [Se] =


Figure 19 Pourbaix diagram for the Hg-Se-H2O system at 298K. [Hg] = 10-3Molal, [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

xi

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


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



xii





Figure 34 Effect of temperature on mercury precipitation at pH 2. Test conditions: pH = 2,








Figure 38 Effect of seeding at pH 5. Test conditions: 60oC, pH = 5, 500RPM, thiosulfate

concentration 10 times as much as stoichiometrically required. ...................................... 62


concentration 10 times as much as stoichiometrically required. ...................................... 63

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




Figure 44 Effect of selenosulfate concentration on mercury precipitation experiments,

Test conditions: 20oC, 500RPM, selenosulfate concentration 10 times as much as


xiii

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

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

xiv

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.

xv

DEDICATION

I lovingly dedicate this thesis to my wife, Henglin Jin, for her endless love and support

through the completion of this project.

1

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].

2

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.

3

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

4

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.

Hg2+ + 2CN- → Hg(CN)20

5

Hg2+ + 4CN- → Hg(CN)42-

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:

HgCl2 + Hg0 → 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

6

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.

7

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.

8

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]

9

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.

3Hg0 + 8H+ + 2NO3- → 3Hg2+ + 2NO(g) + 4H2O

Hg2+ + S2- → 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

10

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]

11

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]

12

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

The Pourbaix diagrams for the Hg-H2O system at solute concentrations of 1 molal, 10-3

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)

13

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)

14

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

form of Hg22+ 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+

dimers are formed [22]. Hg22+(aq) can be disproportioned into Hg2+(aq) and metallic Hg:

15

Hg22+ ↔ Hg2+ + Hg

The disproportion reaction can occur with the addition of OH- or HS- ions [24]. Thus

Hg22+ ions only have a relatively small domain of stability.

According to Figures 2 and 3, when the mercury concentration decreases, the domains of

stability of both Hg2+(aq) and Hg22+(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

(HgCl4)2-.

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.

16

Shown below in Figure 4 is the Eh pH diagram for the Cl-H2O system at 25oC. 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)

Shown in Figures 5 - 8 are the Pourbaix diagrams for the Hg-Cl-H2O system at 25oC 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

domain of stability of aqueous mercury(II) by forming HgCl2(a) and HgCl42-(a). As

mentioned above, the transition between Hg2+, HgCl2(a) and HgCl42-(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.

When decreasing the mercury activity to 10-3, the domain of stability of HgCl42-(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),

HgCl42-(aq) can no longer be observed on the diagram (Figure 7). Aqueous mercury(II)

17

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)

18

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


Figure 7 Pourbaix diagram for the Hg-Cl-H2O system at 298K. [Hg] = 10-3Molal, [Cl] = 10-3Molal,


19

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


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

I2 + 2S2O32- → 2I- + S4O6

2-

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:

20

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]:

Hg2+ + H2S → HgS + 2H+

Figure 9 The structure of metacinnabar [6]

21

Mercury sulfide is a very stable compound [6, 24]. Thus it is one of the desired stabilized

products for mercury, preferably as cinnabar.

Cinnabar: HgS(s) = Hg2+(aq) + S2-

(aq) pKsp = 56.4

Metacinnabar: HgS(s) = Hg2+(aq) + S2-

(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]

22

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

Thiosulfate (S2O32-) is an unstable anion and can be prepared by dissolving elemental

sulfur in boiling sulfite solution [13].

S + SO32- → S2O3

2-

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

dioxide:

S2O32- + 2H+ → S0 + SO2(g) + H2O

S2O32- + H+ → S0 + 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

23

field, for instance the aforementioned titration to determine the hypochlorite

concentration.

I2 + 2S2O32- → 2I- + S4O6

2-

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].

4Au + 8S2O32- + O2 + 2H2O = 4Au(S2O3)2

3- + 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]:

Hg2+ + 2S2O32- ↔ Hg(S2O3)2

2-

Hg2+ + 3S2O32- ↔ Hg(S2O3)3

4-

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

of stability of mercury sulfide extends. Meanwhile, Hg(S2O3)nm- 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

24

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]

25


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.

26


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.

27

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.

28

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


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


29

Figure 19 Pourbaix diagram for the Hg-Se-H2O system at 298K. [Hg] = 10-3Molal, [Se] = 10--3Molal,


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


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

30

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]:

Se + SO32- → SeSO3

2-

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.

31

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.

32

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.

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

pH Probe Thermometer

Temperature Probe

pH Controller

Acid/Basic Solution

Balance

Pumping Device

Water Out Water In

Magnetic Stirrer

33

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.

34

Table 2 Test conditions of hypochlorite leaching experiments

Mercury

Dosage

Sodium

Hypochlorite

Concentration

ClO- : Hg

Molar

Ratio pH Temperature

Stirring

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

1 0.05 10 4 20 500

1 0.05 10 5 20 500

Effect of pH

1 0.05 10 6 20 500

1 0.05 10 4 20 400

1 0.05 10 4 20 600

1 0.05 10 4 20 800

Effect of

Stirring

Speed 1 0.05 10 4 20 1000

1 0.01 2 4 20 1000

1 0.02 4 4 20 1000

1 0.03 6 4 20 1000

1 0.04 8 4 20 1000

Effect of

Hypochlorite

Concentration

1 0.05 10 4 20 1000

1 0.05 10 4 20 500

1 0.05 10 4 30 500

1 0.05 10 4 40 500

Effect of

Temperature

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.

35

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.

36

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%,

37

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.

38

Table 3 Test conditions of thiosulfate precipitation experiments

Mercury

Dosage

Sodium

Thiosulfate

Concentration

S2O32- :

Hg Molar

Ratio

pH Temperature

"Seed" :

Hg Molar

Ratio

g/L M oC

2 0.1 10 2 80 N/A

2 0.1 10 4 80 N/A

2 0.1 10 6 80 N/A

2 0.1 10 8 80 N/A

Effect of pH

2 0.1 10 10 80 N/A

2 0.1 10 2 20 N/A

2 0.1 10 2 40 N/A

2 0.1 10 2 60 N/A

Effect of

Temperature

2 0.1 10 2 80 N/A

2 0.1 10 4 20 N/A

2 0.1 10 4 40 N/A

2 0.1 10 4 60 N/A

Effect of

Temperature

2 0.1 10 4 80 N/A

2 0.1 10 5 20 N/A

2 0.1 10 5 40 N/A

2 0.1 10 5 60 N/A

Effect of

Temperature

2 0.1 10 5 80 N/A

2 0.1 10 6 20 N/A

2 0.1 10 6 40 N/A

2 0.1 10 6 60 N/A

Effect of

Temperature

at pH 6 2 0.1 10 6 80 N/A

2 0.1 10 5 60 1

2 0.1 10 5 60 5

2 0.1 10 6 60 1

Effect of

"Seeding"

2 0.1 10 6 60 5

39

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.

40

Table 4 Test conditions of selenosulfate precipitation experiments

Mercury

Dosage

Sodium

Selenosulfate

Concentration

SeSO32- :

Hg Molar

Ratio pH Temperature

Stirring

Speed

g/L M oC RPM

2 0.1 10 Basic 10 500

2 0.1 10 Basic 20 500

2 0.1 10 Basic 40 500

2 0.1 10 Basic 60 500

Effect of

Temperature

2 0.1 10 Basic 80 500

2 0.02 2 Basic 20 500

2 0.04 4 Basic 20 500

2 0.06 6 Basic 20 500

2 0.08 8 Basic 20 500

Effect of

Selenosulfate

Concentration

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

41

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.

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.

Heating

Magnetic Stirrer

Thermometer

42

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

Selenium Added Sodium Sulfite Added Se:SO32-

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.

43

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



5 Selenosulfate ppt, 10oC




44

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.

45

Effect of pH on Hypochlorite Leaching of Mercury

0%

10%

20%

30%

40%

50%

60%

70%

80%

90%

100%

0 100 200 300 400 500 600

Time (min)

% M

erc

ury

Ex

tra

cte

d

pH=1

pH=2

pH=3

pH=4

pH=5

pH=6

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.

46

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

0%

10%

20%

30%

40%

50%

60%

70%

80%

90%

100%

0 100 200 300 400 500 600

Time (min)

% M

ercu

ry E

xtra

cted

400RPM

600RPM

800RPM

1000RPM

z

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.

47

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

0%

10%

20%

30%

40%

50%

60%

70%

80%

90%

100%

0 100 200 300 400 500 600

Time (min)

% M

erc

ury

Ex

tra

cte

d

2 Times

4 Times

6 Times

8 Times

10 Times

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.

48

Hypochlorite Consumption

0

0.005

0.01

0.015

0.02

0.025

0.03

0.035

0.04

0.045

0.05

0 100 200 300 400 500 600

Time (min)

Hyp

och

lori

te C

on

cen

trat

ion

(M

)

2 Times

4 Times

6 Times

8 Times

10 Times

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.

49

Table 7 Comparison of theoretical and measured hypochlorite consumption

Hypochlorite Consumption Initial

Hypochlorite

Concentration

Initial NaClO :

Hg Molar Ratio Theoretical Measured Differences

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.

50

Effect of Temperature on Hypochlorite Leaching of Mercury

0%

10%

20%

30%

40%

50%

60%

70%

80%

90%

100%

0 100 200 300 400 500 600

Time (min)

% M

ercu

ry E

xtra

cted

20C

30C

40C

50C

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.

51

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

52

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

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

0 100 200 300 400 500 600

Time (min)

% M

ercu

ry E

xtra

cted

Test 1

Test 2

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

53

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

Molar Ratio Temperature Test Duration Hg Extraction

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.

54

Table 11 Results of cyanidation experiments

Sodium Cyanide Concentration Air Mercury Dosage Duration

Mercury

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:

HgO + 2S2O32- + H2O → Hg(S2O3)2

2- + 2OH-

Hg(S2O3)22- + H2O → HgS + SO4

2- + S2O32- + 2H+

Or,

Hg2+ + 2S2O32- + H2O → Hg(S2O3)2

2-

Hg(S2O3)22- + H2O → HgS + SO4

2- + S2O32- + 2H+

55

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.

56

Effect of pH on Precipitation Rate

0%

10%

20%

30%

40%

50%

60%

70%

80%

90%

100%

0 200 400 600 800 1000 1200 1400

Time (min)

Per

cen

t M

ercu

ry i

n S

olu

tio

n (

%)

pH=2

pH=4

pH=6

pH=8

pH=10

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

57

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.

58





59




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.

60

Effect of Temperature at pH 2

0%

10%

20%

30%

40%

50%

60%

70%

80%

90%

100%

0 200 400 600 800 1000 1200 1400

Time (min)

Pe

rce

nt

Me

rcu

ry in

So

luti

on

(%

)

20C

40C

60C

80C

Figure 34 Effect of temperature on mercury precipitation at pH 2. Test conditions: pH = 2, 500RPM,



0%

10%

20%

30%

40%

50%

60%

70%

80%

90%

100%

0 200 400 600 800 1000 1200 1400

Time (min)

Per

cen

t M

erc

ury

in S

olu

tio

n (

%)

20C

40C

60C

80C



61


0%

10%

20%

30%

40%

50%

60%

70%

80%

90%

100%

0 200 400 600 800 1000 1200 1400

Time (min)

Pe

rce

nt

Me

rcu

ry in

So

luti

on

(%

)

20C

40C

60C

80C




0%

10%

20%

30%

40%

50%

60%

70%

80%

90%

100%

0 200 400 600 800 1000 1200 1400

Time (min)

Pe

rce

nt

Me

rcu

ry in

So

luti

on

(%

)

20C

40C

60C

80C



62

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

0%

10%

20%

30%

40%

50%

60%

70%

80%

90%

100%

0 200 400 600 800 1000 1200 1400

Time (min)

Per

cen

t M

ercu

ry i

n S

olu

tio

n (

%)

pH5 No Seed

pH5 with Seed


concentration 10 times as much as stoichiometrically required.

63

Effect of Seeding at pH 6

0%

10%

20%

30%

40%

50%

60%

70%

80%

90%

100%

0 200 400 600 800 1000 1200 1400

Time (min)

Per

cen

t M

ercu

ry i

n S

olu

tio

n (

%)

pH6 No Seed

pH6 with Seed


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.

Se + SO32- ↔ SeSO3

2-

Ullah (2012) has conducted some experiments to investigate selenium dissolution

behaviour. Experimental conditions and results are provided below in Table 12.

64

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

Sodium

Sulfite Selenium Time Temperature

Selenium

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.

65

Table 13 Results for selenium dissolution experiments

Measured Selenium Concentration

% Selenium

Dissolved SO32- : Se

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:

HgO + 2SeSO32- + H2O → Hg(SeSO3)2

2- + 2OH-

Hg(SeSO3)22- + H2O → HgSe + SO4

2- + SeSO32- + 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

66

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.


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

67

Figure 40 XRD pattern of formed precipitates at 10oC


68



69

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

0.00%

10.00%

20.00%

30.00%

40.00%

50.00%

60.00%

70.00%

80.00%

90.00%

100.00%

0 20 40 60 80 100 120 140

Time (min)

Pe

rce

nt

Me

rcu

ry in

So

luti

on

(%

)

2X

4X

6X

8X

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.

70

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



5 Selenosulfate ppt, 10oC 2

6 Selenosulfate ppt, 40oC 1

7 Selenosulfate ppt, 60oC <1

8 Selenosulfate ppt, 80oC <1

71

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.

72

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

leaching stage.

73

Figure 45 Flowsheet of mercury stabilization process via thiosulfate precipitation

Hydrochloric Acid

Elemental

Mercury

pH Adjustment

Leachate

Sodium Hydroxide

Solid Liquid

Separation

Leaching

Leaching

Mercury

Oxide

Sodium Thiosulfate

Precipitation

Hydrochloric Acid

Mercury Sulfide

Liquid solution

Tailing Treatment

Tailing Treatment

Sodium

Hypochlorite

Solution

Solid Liquid

Separation Liquid solution

74

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.

75

Figure 46 Flowsheet of mercury stabilization process via selenosulfate precipitation

Hydrochloric Acid

Elemental

Mercury

pH Adjustment

Leachate

Sodium Hydroxide

Solid Liquid

Separation

Leaching

Leaching

Mercury

Oxide

Precipitation

Mercury Selenide

Liquid solution

Tailing Treatment

Sodium

Hypochlorite

Solution

Leaching

Solid Liquid

Separation

Tailing Treatment

Liquid solution

Elemental

Selenium

Sodium

Sulfite

Solution

Leachate

Leachate

76

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