an electrochemical and sers study of the gold-thiosulfate
TRANSCRIPT
An Electrochemical and SERS Study of the Gold-Thiosulfate Interface in the Presence of Copper
by
Eric Nicol
A Thesis Presented to
The University of Guelph
In partial fulfillment of requirements for the degree of
Master of Science in
Chemistry
Guelph, Ontario, Canada © Eric Nicol, April 2013
ii
ABSTRACT
AN ELECTROCHEMICAL AND SERS STUDY OF THE GOLD-THIOSULFATE INTERFACE IN THE PRESENCE OF COPPER
Eric Nicol Advisor: University of Guelph, 2013 Prof. Jacek Lipkowski
Complementary electrochemical and spectroscopic techniques were used to characterize
the behavior and composition of the passive layer formed at the gold-thiosulfate interface in the
presence of copper. Raman studies of three different cationic (calcium, ammonium and sodium)
thiosulfate leaching solutions showed that the concentrations of sulfate, thiosulfate, trithionate
and tetrathionate remained constant.
Initial leaching current densities for the three systems were identical, however, significant
differences were noted in the open circuit potentials of these systems. Gold nanorod electrodes
were employed as substrates for Surface Enhanced Raman Spectroscopy (SERS) studies of the
gold-thiosulfate interface. The composition and behavior of the passive layer at the gold-
thiosulfate interface greatly differed from that of the bulk solutions. Higher order polythionate
species were not observed, and significant differences were noted in the behavior of species
common between the three thiosulfate leaching solutions. Passivation levels determined from
SERS indicate that in the presence of copper, the cation associated with thiosulfate may play a
key role in the extent of passivation on the gold surface.
iii
ACKNOWLEDGEMENTS
Many people contributed to the success of this project, and to the maintenance of my sanity
throughout. I would like to thank my supervisor Professor Jacek Lipkowski for being a fantastic
source of knowledge and guidance since I first emailed him asking for a job. His patience,
positive attitude and willingness to discuss challenges were instrumental in the success of this
work.
I would like to express my gratitude to the members of my advisory and examination
committees: Professors Peter Tremaine, Mark Baker, and Abdelaziz Houmam for their time
dedicated to this thesis, as well as Professor Paul Rowntree for chairing both my proposal and
defence.
I would like to recognize the financial support from Barrick Gold Corporation and in
particular Dr. Yeonuk Choi who was actively involved in the project.
I would also like to acknowledge the people who collaborated with the experimental
development and execution of this project: Dr. Janet Baron Gavidia without whom I would have
been lost and confused, and Scott Smith who has done an excellent job of filling Dr. Baron
Gavidia’s shoes. To past and present members of Dr. Lipkowski’s lab group at the University of
Guelph, I would like to express thanks for an unforgettable experience. Especially to Jay Leitch
and Ryan Seenath, who provided invaluable discussion and advice at coffee, and never let my
head swell too large. To my friends Michael Yacyshyn, Dave Sullivan and Christian Carello,
thanks for helping make my successes fantastic and my disasters not so bad.
Lastly, but not least, I would like to give my utmost gratitude to my amazing family and to
Sara Trayes who were cheerleaders when I wanted to quit, and without whom, achievement of
my Masters would not have been possible.
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TABLE OF CONTENTS
Acknowledgements ........................................................................................................................ iii Table of Contents ........................................................................................................................... iv List of Tables ................................................................................................................................. vi List of Figures ................................................................................................................................ vi List of Symbols and Abbreviations ................................................................................................ xi Chapter 1: Introduction ................................................................................................................... 1
1.1 Motivation and Impact: ......................................................................................................... 1 1.2 Objectives: ............................................................................................................................ 3 1.3 Scope: .................................................................................................................................... 4
Chapter 2: Background and Literature Review .............................................................................. 6 2.1 Gold Ore and the Hydrometallurgical Process ..................................................................... 6 2.2 Leaching Mechanism and History ...................................................................................... 12
2.2.1 Cyanidation .................................................................................................................. 13 2.2.2 Environmental Concerns .............................................................................................. 15 2.2.3 Thiosulfate Alternative ................................................................................................ 15
2.3 Thiosulfate Leaching of Gold ............................................................................................. 16 2.3.1 Ammonia-Thiosulfate Leaching .................................................................................. 21 2.3.2 The Role of the Electrolyte on Ammonia-Thiosulfate Leaching ................................. 25
Chapter 3: Experimental Techniques Theory ............................................................................... 28 3.1 Electrochemical Techniques ............................................................................................... 28
3.1.1 Sweep Voltammetry Methods ...................................................................................... 28 3.1.2 Mixed Potential Theory ............................................................................................... 36 3.1.3 Hydrodynamic Methods ............................................................................................... 39
3.2 Spectroscopic Techniques ................................................................................................... 42 3.2.1 Surface Enhanced Raman Spectroscopy ...................................................................... 42 3.2.2 SERS Substrates .......................................................................................................... 49
Chapter 4: Methodology ............................................................................................................... 52 4.1 Reagents .............................................................................................................................. 52 4.2 Cleaning Methods ............................................................................................................... 52 4.3 Experimental Setup and Process ......................................................................................... 54
4.3.1 Preparation of Gold Electrodes .................................................................................... 54 4.3.2 Electrochemical Experiments ...................................................................................... 54
4.3.2.1 Leaching Current Measurements .......................................................................... 54 4.3.2.2 Open Circuit Potential Measurements .................................................................. 55 4.3.2.3 Solution pH Measurements ................................................................................... 56
4.3.3 Preparation of Gold Nanorod Electrodes ..................................................................... 57 4.3.4 Raman Experiments ..................................................................................................... 58
4.3.4.1 SERS ..................................................................................................................... 58 4.3.4.2 Solution Raman ..................................................................................................... 59
Chapter 5: Results and Discussion ................................................................................................ 60 5.1 Preamble ............................................................................................................................. 60 5.2 Characterization of Bulk Solution ...................................................................................... 61
5.2.1 pH Measurements ........................................................................................................ 62
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5.2.2 Solution Raman ............................................................................................................ 64 5.3 Characterization of the Gold-Thiosulfate Interface in the Presence of Copper .................. 71
5.3.1 Initial Characterization ................................................................................................. 71 5.3.2 Leaching Current Measurements ................................................................................. 79 5.3.3 Open Circuit Potential Measurements ......................................................................... 81 5.3.4 Surface Enhanced Raman ............................................................................................ 83
5.4 Summary ........................................................................................................................... 101 Chapter 6: Conclusions and Future Work ................................................................................... 106
6.1 Conclusions ....................................................................................................................... 106 6.2 Future Work ...................................................................................................................... 107
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LIST OF TABLES
Table 5.1: Characteristic Raman active vibrational modes of species expected in the Raman
spectra of the investigated leaching solutions.............................................................................65
Table 5.2: Final change in the normalized peak areas investigated in the calcium, ammonium
and thiosulfate systems.............................................................................................................102
LIST OF FIGURES
Figure 2.1: Flow chart for the processing of gold ore to the pure metal.....................................9
Figure 2.2. Schematic for the electrochemical model for leaching of gold...............................12
Figure 2.3: SERS spectrum of 0.1 M Na2S2O3 at open circuit potential...................................19
Figure 2.4: SERS spectrum of a gold nano-rod electrode submersed in a
0.1M Na2S2O3 + 1 x 10-4M NaOH at open circuit potential.......................................................20
Figure 2.5: Calculated voltammetric current for varying concentrations of ammonia in
thiosulfate leaching solution.......................................................................................................21
Figure 2.6: The effect of copper on gold dissolution in thiosulfate leaching............................23
Figure 2.7: Electrochemical model for the leaching of gold using ammoniacal thiosulfate leach
media...........................................................................................................................................25
Figure 3.1: a) Typical potential time curve for cyclic voltammetry b) Current response
curve............................................................................................................................................28
Figure 3.2: Schematic of the structure of the electrical double layer formed at the metal-
electrolyte interface.....................................................................................................................30
Figure 3.3: Tafel plot of a redox system....................................................................................33
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Figure 3.4: Diagram of the free energy curves of an oxidized and reduced species when a
potential is applied, shifting the system from equilibrium...........................................................35
Figure 3.5: Intersection of the locally linear potential energy surface of an oxidized and reduced
species..........................................................................................................................................35
Figure 3.6: Generic linear sweep voltammogram.......................................................................37
Figure 3.7: Tafel plot a gold electrode in contact with a 0.1 M Na2S2O3 + 0.01 M CuSO4 +
1.1µΜ Ca(OH)2 solution..............................................................................................................37
Figure 3.8: Plot of the current versus the overpotential in a ~40mV range around the mixed
potential for a 0.1 M Na2S2O3 + 0.01 M CuSO4 + 1.1µΜ Ca(OH)2 system................................39
Figure 3.9: a) Profile for a solution with laminar flow b) Profile for a solution with turbulent
flow...............................................................................................................................................40
Figure 3.10: Transitions seen in Raman spectroscopy................................................................45
Figure 3.11: Photon-driven charge-transfer process...................................................................47
Figure 4.1: Cyclic voltammogram of a clean, bare gold RDE in a solution of 0.1 M
NaF...............................................................................................................................................53
Figure 4.2: Schematic for the two cell configurations used.......................................................55
Figure 4.3: Calibration curve for the calculation of pH from the potential measured................56
Figure 5.1: Average pH over a period of 3 hours for calcium, ammonium and sodium thiosulfate
systems (0.1 M S2O3 + 0.01 M CuSO4 and adjusted to a pH of 8.0-8.5) and a blank solution of
pure MilliQ water.........................................................................................................................63
Figure 5.2: Average raw Raman spectra collected over a period of 3 hours for leaching solutions
of: a) 0.1 M CaS2O3 + 0.01 M CuSO4 b) 0.1 M (NH4)2S2O3 + 0.01 M CuSO4 c) 0.1 M Na2S2O3
+ 0.01 M CuSO4...........................................................................................................................66
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Figure 5.3: Deconvoluted spectra of an average of three 0.1 M CaS2O3 + 0.01 M CuSO4
solutions, adjusted to a pH of 8.0-8.5, 20 min after solution preparation...................................67
Figure 5.4: Deconvoluted spectra of an average of three 0.1 M CaS2O3 + 0.01 M CuSO4
solutions, adjusted to a pH of 8.0-8.5, 20 min after solution preparation, in the 900-1000 cm-1
region...........................................................................................................................................68
Figure 5.5: Normalized peak areas for band positions of 1037 cm-1 (tetrathionate = black), 425
cm-1 (trithionate = blue), 980 cm-1 (sulfate = green), 448 cm-1 (thiosulfate = red). a) 0.1 M
CaS2O3 + 0.01 M CuSO4, b) 0.1 M Na2S2O3 + 0.01 M CuSO4, c) 0.1 M (NH4)2S2O3 + 0.01 M
CuSO4...........................................................................................................................................69
Figure 5.6: Linear sweep voltammogram of a 0.1 M Na2S2O3 + 0.01 M CuSO4 solution, pH 8.0-
8.5.................................................................................................................................................71
Figure 5.7: Tafel plot of a 0.1 M Na2S2O3 + 0.01 M CuSO4 solution, pH 8.0-
8.5.................................................................................................................................................72
Figure 5.8: Linear regression of the cathodic reduction of Cu2+ for the determination of the
transfer coefficient........................................................................................................................73
Figure 5.9: Linear regression of the anodic oxidation of Au for the determination of the transfer
coefficient.....................................................................................................................................74
Figure 5.10: Linear regression of data acquired during a linear sweep voltammogram of a 0.1 M
Na2S2O3 + 0.01 M CuSO4 solution, pH 8.0-8.5 solution with a sweep rate of 1 mVs-1..............75
Figure 5.11: Sweep rate dependence of the leaching current measured in a 0.1 M Na2S2O3 +
0.01 M CuSO4 solution, pH 8.0-8.5, at sweep rates of 1, 2, 10, 20, 50 and 100 mVs-1..............77
ix
Figure 5.12: Rotation dependence of the leaching current of a gold electrode in contact with a
0.1 M CaS2O3 + 0.01 M CuSO4 solution, pH 8.0-8.5, at rotation rates of 300, 500, 700 and 1000
RPM..............................................................................................................................................78
Figure 5.13: Average calculated leaching current density for solutions of sodium, calcium and
ammonium thiosulfate solutions (0.1 M S2O3 + 0.01 M CuSO4 and were adjusted to a pH of 8.0-
8.5)................................................................................................................................................80
Figure 5.14: Average open circuit potential of a gold disk RDE as a function of immersion time
in solutions of calcium, sodium and ammonium thiosulfate (0.1 M S2O3 + 0.01 M CuSO4 and
adjusted to a pH of 8.0-8.5)...........................................................................................................82
Figure 5.15: Raw SERS spectra of gold nanorod electrodes exposed to leaching solutions of 0.1
M S2O3 + 0.01 M CuSO4, adjusted to pH 8.0-8.5. Solutions were a) CaTS b) ATS c)
STS................................................................................................................................................85
Figure 5.16: Fitted SERS spectrum of a gold nanorod electrode, after 5 minutes of exposure to a
0.1 M CaS2O3 + 0.01 M CuSO4 leaching solution, with an initial pH of 8.0-
8.5..................................................................................................................................................86
Figure 5.17: Fitted SERS spectrum of a gold nanorod electrode, after 15 minutes of exposure to
a 0.1 M (NH4)2S2O3 + 0.01 M CuSO4 leaching solution, with an initial pH of 8.0-
8.5..................................................................................................................................................88
Figure 5.18: Fitted SERS spectrum of a gold nanorod electrode, after 5 minutes of exposure to a
0.1 M Na2S2O3 + 0.01 M CuSO4 leaching solution, with an initial pH of 8.0-
8.5..................................................................................................................................................90
Figure 5.19: Normalized analytical peak areas for a gold nanorod electrode treated with 0.1 M
CaS2O3 + 0.01 M CuSO4 solution adjusted to pH 8.0-8.5. Peak areas tracked were for band
x
positions of a) 216 cm-1 (green circle), 255 cm-1 and 300 cm-1 (brown square), and b) 400 cm-1
(blue circle), 443 cm-1 (red square), and 610 cm-1 (pink diamond)............................................92
Figure 5.20: Analytical peak areas for a gold nanorod electrode treated with 0.1 M (NH4)2S2O3
+ 0.01 M CuSO4 solution adjusted to pH 8.0-8.5. Peak areas tracked were for band positions of:
a) 216 cm-1 (green circle), 255 cm-1 and 300 cm-1 (brown square), and b) 400 cm-1 (blue circle),
443 cm-1 (red square), and 610 cm-1 (pink diamond)..................................................................96
Figure 5.21: Analytical peak areas for a gold nanorod electrode treated with 0.1 M Na2S2O3 +
0.01 M CuSO4 solution adjusted to pH 8.0-8.5. Peak areas tracked were for band positions of: a)
216 cm-1 (green circle), 255 cm-1 and 300 cm-1 (brown square), and b) 400 cm-1 (blue circle), 443
cm-1 (red square), and 610 cm-1 (pink diamond)..........................................................................99
Figure 5.22: Normalized analytical peak area of a gold nanorod electrode treated with 0.1 M
Na2S2O3 solution adjusted to pH 10............................................................................................104
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LIST OF SYMBOLS AND ABBREVIATIONS
CHAPTER 1:
SERS – Surface enhanced Raman spectroscopy
CHAPTER 2:
REQCM – Rotating electrochemical quartz crystal microbalance
SCE – Saturated calomel electrode
ν – Stretching vibration
δ – Bending vibration
νsym – Symmetric stretching vibration
FSD – Fourier self-deconvolution
CHAPTER 3:
E – Applied potential
E1/E2/E3 – Potential limits
OHP – Outer Helmholtz plane
IHP – Inner Helmholtz plane
νf – Rate of reduction reaction
νb – Rate of oxidation reaction
ic – Cathodic current
ia – Anodic current
CO – Surface concentration of oxidized species
CR – Surface concentration of reduced species
xii
n – number of electrons
A – Electrode area
F – Faraday’s constant
kf – Rate constant of reduction reaction
kb – Rate constant of oxidation reaction
α – Transfer coefficient
k0 – Standard rate constant
E0 – Standard potential
R – Gas constant
T – Temperature
i – Measured current
Eeq – Equilibrium potential
i0 – Exchange current
η – Overpotential
Rct – Charge transfer resistance
iM – Leaching current
EM – Mixed potential
αCu2+ – Transfer coefficient for the reduction of Cu2+
αAu – Transfer coefficient for the oxidation of Au
Re – Reynold’s number
υch – Characteristic velocity of a fluid
l – Characteristic length
ν – Kinematic viscosity
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D – Diffusion coefficient
CO* - Bulk concentration of oxidized species
ω – angular velocity
ik – Current in the absence of mass transfer effects
il.c – Cathodic limiting current
EF – External electric field
E0 – Amplitude of external electric field wave
!ex – Frequency of exciting light
! – Induced dipole moment
!P – Polarizability of the molecule
!0 – Polarizability of the molecule at the equilibrium position
r – Internuclear distance
req – Internuclear distance at the equilibrium position
rm – Maximum internuclear distance
!v – Frequency of vibration of molecule
HOMO – Highest occupied molecular orbital
LUMO – Lowest occupied molecular orbital
Eout - Magnitude of the electric field produced outside a spherical nanoparticle of radius
a
(x, y, z) – Cartesian coordinates
(x, y, z) – Cartesian unit vectors
rd – Radial distribution of electric field
!M – Polarizability of the metal
xiv
εin – Dielectric constant of the nanoparticle
εout – Dielectric constant of the environment surrounding the nanoparticle
E(λ) – Extinction spectrum of a non-spherical nanoparticle
N – Finite number of polarizable elements within a nanoparticle
εi – Imaginary component of the dielectric constant of the nanoparticle
εr – Real component of the dielectric constant of the nanoparticle
λ – Wavelength
a – Radius of nanoparticle
! – Shape factor which accounts for geometries of the nanoparticle, other than spherical
LSPR – Localized surface plasmon resonance
CHAPTER 4:
RDE – Rotating disk electrode
CV – Cyclic voltammogram
RE – Reference electrode
CE – Counter electrode
E – Applied potential
i – Measured current
CHAPTER 5:
SERS – Surface enhanced Raman spectroscopy
FSD – Fourier self deconvolution
! – Rocking vibration
xv
νsym – Symmetric stretching vibration
νasym – asymmetric stretching vibration
δsym – Symmetric bending vibration
δasym – Asymmetric bending vibration
i – Measured current
V – Applied potential
!c – Cathodic transfer coefficient
!a – Anodic transfer coefficient
F – Faraday’s constant
T – Temperature
R – Gas constant
i0 – Exchange current
n – Number of electrons
! - Overpotential
!Cu2+ – Transfer coefficient for the reduction of Cu2+
!Au – Transfer coefficient for the oxidation of Au
nCu2+ - Number of electrons transferred in the reduction of Cu2+
nAu – Number of electrons transferred in the oxidation of Au
iM – Leaching current at the mixed potential
SCE – Saturated calomel electrode
RDE – Rotating disk electrode
ik – Current in the absence of mass transfer effects
il.c – Cathodic limiting current
xvi
D – Diffusion coefficient
A – Area
! – Angular velocity
ν – Kinematic viscosity
CO* - Bulk concentration of oxidized species
CaTS – Calcium thiosulfate
ATS – Ammonium thiosulfate
STS – Sodium thiosulfate
OCP – Open circuit potential
CHAPTER 6:
SERS – Surface enhanced Raman spectroscopy
1
CHAPTER 1: INTRODUCTION
1.1 MOTIVATION AND IMPACT:
Metallic gold has application in a number of industries including the chemical catalysis
and microelectronics fabrication, as a result of exceptional properties such as its catalytic ability,
corrosion resistance, thermal conductivity, and ductility.1
Extraction from its ore is highly dependent on the state of gold within, and its response to
the use of leaching reagents.2 The leaching process involves the addition of a chemical agent to
oxidize the gold, and a complexing agent to bind the Au+ ion and bring it into solution. In most
instances, reduction of molecular oxygen is coupled with the oxidation of gold. The solution
containing the complexed gold ion is later submitted to a process such as electrowinning or
adsorption onto activated carbon to complete extraction of the gold.3
Cyanide is by far the most widely used lixiviant, or leaching reagent, in use today for
extraction of gold from its ores.3-6 However, due to recent concerns over its toxicity and
environmental impact, there has been an increased interest in alternative leaching reagents.3, 6
Stable thiosulfate-gold complexes have been known for over 100 years, but serious
investigations into its use in leaching has only begun in the last two decades. Thiosulfate is
attractive due to its remarkable ability to deal with refractory, preg-robbing and other ores that
may be difficult or non-responsive to cyanide.7 Due to its low toxicity, and comparatively high
initial rate of leaching, thiosulfate is perhaps the most promising alternative to cyanide
leaching.3, 4, 6 This is especially true in the presence of copper and ammonia in the leaching
solution. However, although initial rates of gold leaching from cyanide and thiosulfate solutions
2
are comparable, with extended time leaching, only 60-70% of gold is typically leached in the
presence of thiosulfate.4
The effectiveness of thiosulfate leaching in industry is limited by the formation of a
passive layer on the gold surface, composed of a combination of sulfide and other decomposition
products of the thiosulfate.8, 9 Recent research has shown that the passivating layer consisted of
sulfide and polythionate species10, including thiosulfate, tetrathionate and trithionate.11
The individual addition of both ammonia and copper into thiosulfate leaching solutions
has been shown to increase the rate of leaching and aid in the inhibition of passive layer
formation. Addition of ammonia in to thiosulfate leaching solution results in the formation of an
anionic aurocomplex leading to an increased rate of gold oxidation, providing stability over a
wide pH range, and inhibiting interference from species such as iron oxides, silica, silicates and
carbonates.4, 12 Ammoniacal solutions require the addition of copper due to the inability of O2 to
act as the oxidant.4 Copper, in the absence of ammonia, has also been shown to affect gold
dissolution rates. Zhang and Nicol found that the addition of copper in different forms resulted in
significant increases in the amount of gold dissolved.13
In aqueous thiosulfate solutions containing both copper and ammonia, an increased
leaching rate has been observed, in part due to the formation of an oxidizing copper tetraamine
complex that allows for rapid gold oxidation.6, 14, 15 Although the effects of copper and ammonia
addition on gold dissolution and oxidation are known, the mechanisms through which copper and
ammonia act individually, and in tandem, to affect passive layer formation remain unknown.
The composition of the electrolyte is also known to affect on the leaching rate of gold in
thiosulfate media. Chandra et al.16 observed a significant differences on the ease of oxidation of
gold between sodium, potassium and ammonium thiosulfate solutions. Potassium thiosulfate
3
showed considerable improvement over sodium thiosulfate; however, ammonium was the most
effective cation at increasing the leaching rate. Although explanations for the preference of
heavier alkali metals exist, no explanation for the effect on gold oxidation has yet been provided.
By using a combination of electrochemical techniques and Surface Enhanced Raman
Spectroscopy (SERS), the roles of both Cu2+ and the electrolyte cation in passive layer
composition during thiosulfate leaching of gold can be clarified. Application of the knowledge
gained in the study will ultimately aid in the understanding an improvement of the leaching
process.
Gaining an understanding of the leaching process in the presence of Cu2+ is invaluable
due to its direct application in industrial practices. Ideally, if the function of copper is
understood, optimization of the thiosulfate leaching system could be undertaken in an attempt to
obtain a viable competitor for cyanide leaching; one that is significantly more environmentally
responsible, shows improved gold recovery, and is economically sustainable.
1.2 OBJECTIVES:
Application of thiosulfate leaching of gold in industry requires a complete understanding
of the behavior and composition of the passive layer under applicable conditions. As such, the
role of copper in passive layer formation is vital. The specific objectives of this project are:
• To determine the species composing the passive layer at the gold-thiosulfate interface in
the presence of copper, and determine their behavior through time
• To investigate the effect of the cation in the electrolyte on the composition of the passive
layer at the interface
• To identify which species may be promoting leaching, and which promote passivation
through complementary electrochemical and spectroscopic techniques
4
• To identify dominating reactions at the gold-thiosulfate interface, and determine how
they influence the composition of the passive layer
1.3 SCOPE:
This work consists of 6 chapters. The first chapter is an introduction to the main body of
this thesis, including a brief background of the thiosulfate leaching system, and previous studies
on the effect of copper. The second chapter will include a detailed review of thiosulfate leaching
media, and previous studies for characterization of the passive layer, and its composition. Gold
dissolution rates in the presence and absence of copper will be addressed along with the effect of
ammonia, individually and in combination with copper.
Chapter 3 will summarize the theory behind the experimental techniques used in this
project; electrochemical theory for sweep voltammetry methods and mixed potential theory, as
well as hydrodynamic method theory will be discussed. Fundamentals of both Raman
spectroscopy and Surface Enhanced Raman Spectroscopy (SERS) are examined. Chapter 4
presents the experimental methodology used for all experiments, and the reagents and solutions
used for these investigations.
The fifth chapter presents the results acquired from electrochemical and spectroscopic
methods in calcium, sodium and ammonium thiosulfate leaching media in the presence of
copper. The results are analyzed and discussed in terms of studies of the bulk solution and the
passive layer at the interface, as a function of time. The sixth and final chapter presents the
conclusions of the study as well as suggestions for future work to further the understanding of
the industrial leaching process using thiosulfate media.
5
References
1. Sullivan, A. M.; Kohl, P. A. J. Electrochem. Soc. 1997, 144, 1686-1690. 2. Gupta, C.; Mukherjee, T. Boca Raton 1990, , 127-165. 3. Hilson, G.; Monhemius, A. J. J. Clean. Prod. 2006, 14, 1158-1167. 4. Abbruzzese, C.; Fornari, P.; Massidda, R.; Vegliņ, F.; Ubaldini, S. Hydrometallurgy
1995, 39, 265-276. 5. Zhang, S. C.; Nicol, M. J. J. Appl. Electrochem. 2003, 33, 767-775. 6. Jeffrey, M. I.; Linda, L.; Breuer, P. L.; Chu, C. K. Minerals Eng 2002, 15, 1173-1180. 7. Marsden, J.; House, I. In The chemistry of gold extraction; Society for Mining
Metallurgy: 2006; . 8. Chu, C. K.; Breuer, P. L.; Jeffrey, M. I. Minerals Eng 2003, 16, 265-271. 9. Pedraza, A. M.; Villegas, I.; Freund, P. L.; Chornik, B. J Electroanal Chem 1988, 250,
443-449. 10. Woods, R.; Hope, G. A.; Watling, K. M.; Jeffrey, M. I. J. Electrochem. Soc. 2006, 153,
D105-D113. 11. Baron Gavidia, J. Study of the Gold-Thiosulfate Interface Under Leaching Conditions,
University of Guelph, Guelph, ON, 2010. 12. Aylmore, M. G.; Muir, D. M. Minerals Eng 2001, 14, 135-174. 13. Zhang, S. C.; Nicol, M. J. J. Appl. Electrochem. 2005, 35, 339-345. 14. Aylmore, M. G.; Muir, D. M. Miner Metall Process 2001, 18, 221-227. 15. Senanayake, G. J. Colloid Interface Sci. 2005, 286, 253-257. 16. Chandra, I.; Jeffrey, M. I. Hydrometallurgy 2004, 73, 305-312.
6
CHAPTER 2: BACKGROUND AND LITERATURE REVIEW
2.1 GOLD ORE AND THE HYDROMETALLURGICAL PROCESS
The inert character of gold leads to very limited formation of naturally occurring
compounds within the Earth’s crust. The average gold content (0.005g/t) is far below that of
other metals such as copper and silver (50 g/t and 0.07 g/t respectively), and is mostly found in
residual hydrothermal fluids and metallic and sulfidic sub-phases1. Rocks with a high
concentration of clay, and low concentration of carbonates are considered some of the best
sources, partly due to ease of re-precipitation upon contact with a reducing environment (regions
with high carbonate, carbon or reducing sulfide).1
Native gold has been found to contain gold in concentrations up to 99.8%, however in
most cases gold content is in the range of 85 to 95%, with silver as the main impurity. If the
silver content is between 25 and 55%, the mineral is called electrum1, 2. The density of native
gold is less than that of pure gold (15 000 kg/m3 versus 19 300 kg/m3), and can consequently be
easily recovered with gravity concentration methods when combined with heavier minerals such
as in gangue minerals (i.e. quartz, silicates etc.)1. Gold tellurides are also a widespread form of
native gold, usually accompanied by free gold and sulfide minerals. Low concentrations of gold
are known to occur with bismuth and copper minerals, but such deposits are relatively rare.
Entrapment of ultrafine gold particles can occur within sulfide mineral grain structures,
occurring in concentrations as high as 15 000 g/t in arsenopyrite, for instance.
Gold ore can be classified into 3 main categories: placer deposits, free milling ores and
lastly, refractory ores3. Placers are deposits that have formed as a direct result of weathering and
7
hydraulic transport of gold particles away from the main source1, 3. They require a primary
source of gold (such as quartz-veins, auriferous sulfide deposits or former placers), a long period
of weathering (both chemical and physical) to liberate gold grains from the primary source,
concentration of gold particles by gravity, and a long period of stable bedrock and surface
conditions to allow for a significant concentration to conglomerate1. Leaching is rarely required
for processing of these ores due to the freedom and coarseness of the gold grains3.
Free-milling ores are defined as those which cyanidation (leaching with cyanide) can
extract approximately 95% of the gold, when 80% of the ore is ground to a size less than 75 µm1,
3. These ores contain gold that is finely distributed in a hard rock matrix (typically quartz) and
require thorough grinding in order to release the gold particles before they can be subjected to
subsequent extraction procedures3.
Refractory ores are unresponsive to the cyanide leaching method developed for
processing of free-milling ores. Generally, refractory ores can be grouped into three
subcategories: sulfide ores, carbonaceous ores and telluride ores. Sulfide ores are the largest
group of refractory ores. Gold can be contained within a wide range of host minerals, with pyrite
and arsenopyrite the most common3. Typically these ores contain gold as submicron sized
particles within sulfide grains, which are themselves usually finely distributed within a quartz or
other hard rock matrix1, 3, 4. Carbonaceous ores pose problems due to the preg-robbing effect
(adsorption of gold from solution as leaching is attempted) and formation of a chemical bond
between gold and the carbonaceous material. Gold tellurides, both in the presence and absence of
silver, dissolve very slightly, if at all in cyanide solutions. As the supply of both placer and free-
milling ores is depleted, greater focus is being placed on the extraction and processing of
refractory ores.
8
Ore preparation is the first process in gold extraction, and has a very large impact on the
overall recovery and entrapment of gold within the host mineral1, 3. Figure 2.1 displays a
flowchart description of the overall gold extraction and recovery process. Ore preparation can
involve size reduction, solid-liquid separation, and in some cases where floatation is to be used,
dewatering3. The optimum size is determined by the economics of the process, including reaction
kinetics, reagent consumption and grinding costs for a given processing method. Permeability
and separation efficiency can also play a large role in determining the particulate size1. For each
of the three main extraction methods (flotation, cyanide leaching and oxidative pretreatment
followed by cyanide leaching) particle size can play an important role in process time and
consumption.
In the process of flotation, ground ore, as a finely divided powder, is mixed with water to
form a slurry. Surfactant is added to bind the gold particles, rendering the surface hydrophobic.
Froth containing gold particles is created by introduction of the slurry into an aerated water bath.
The froth rises to the surface of the tank, where it is removed and further concentrated. Froth
flotation can be used for free gold and gold-bearing sulfide minerals. It allows for pre-
concentration, removal of sulfide (producing sulfide-free tailings), and removal of interfering
matrix components such as carbonates or other carbonaceous materials. Differential flotation (a
multi-stage process where specific minerals are targeted for flotation at each step) can also be
used for the separation of gold from host minerals such as pyrite, and arsenopyrite1.
9
Figure 2.1: Flow chart for the processing of gold ore to the pure metal. Adapted from reference5.
Oxidative pretreatment is required for ore that gives poor gold recovery through
conventional leaching methods1. The pretreatment can be performed through chemical oxidation,
dissociation or roasting, biooxidation or acid leaching (usually carried out at high pressure or
temperature in the presence of a strong oxidant)4. Mineralogy plays a large role in determining
Pit/Ore Source
Ore Preparation
Leaching
Solution Collection
Pregnant Pond Barren Pond
Solid/Liquid Separation
Solution Application
Adsorption on Charcoal
Solvent Extraction
Ion Exchange
Crystallization Ionic Precipitation Gas Reduction
Electrochemical Reduction
Electrolytic Reduction
Pure Metal Compound
Metal Compound Metal or its oxide Impure Metal Pure Metal
Leach/Liquor
10
the degree of oxidation used; partial oxidation may suffice to passivate the surfaces of sulphide
minerals in refractory ores, or release gold that is associated with a specific mineral. Complete
oxidation is typically required if gold is finely dispersed within the ore in question, or locked
within sulfide minerals1.
Leaching is used in every hydrometallurgical extraction of gold to produce a gold bearing
solution by dissolving the constituents of the ore1, 5. It is a solid-liquid mass transfer process that
can be carried out at ambient conditions, or at elevated temperatures and pressures; the process
conditions are greatly dependent on the chemical reactions taking place, and their economics5.
Leaching reagents (also called lixiviants) must:
• Dissolve ore minerals rapidly enough for the process to be economical, and ideally be
chemically inert towards gangue minerals
• Be cheap and easily obtainable on the industrial scale
• Ideally, be regenerated in someway during the leaching reaction or process
The leaching process can be applied in four ways: agitated leaching, heap or dump leaching, vat
leaching, and intensive leaching1. Agitated leaching can be used for ore whose particle sizes
don’t allow for passage of the leach solution between the mineral and gold, particularly
submicron sized gold particles. These particles are added to a solvent within an agitated vessel,
where agitation is commonly incurred though gas injection or a rotating impeller. Agitation
ensures that the particles remain finely distributed in solution, increasing the percentage of gold
dissolved5. It is typically used for extraction of gold from unprocessed gangue ore, and may
require further processing, such as solid-liquid separation or the addition of carbon or ion-
exchange resins “in-pulp” for sequestering of the gold.
11
Heap (or dump) leaching is applied to ore which contains gold that be at least partially
liberated without grinding. As such, it is typically only usable with permeable ores 1. In this
process, the leaching solution is sprayed over the top of the ore pile, which rests on a pad with a
slight slope, and is allowed to drain through the pile, oxidizing and solubilizing gold into the
leach solution, which is collected at the bottom of the drainage channels5.
Vat leaching is similar to heap leaching, but the volume of the leaching media or solution
is greatly increased. The ore is contained within an impermeable vat or vessel, and is
successively treated with the leach solution. This can be achieved either through continuous flow,
or on a batch basis. Continuous flow vat leaching typically occurs with upward percolation,
where the leach solution exits through the top, whereas batch leaching generally uses downward
percolation, in which the leach solution exits at the bottom of the vat1.
Intensive leaching is a combination of vat and agitated leaching. The ore is contained
within a closed reactor system where mechanical agitation and high reagent concentrations
ensure high leaching kinetics. The reactors run continuously and have optimal volumes
dependent on the residence time. If necessary, they can also be run at elevated pressures and
temperatures1, 5.
12
2.2 LEACHING MECHANISM AND HISTORY
As stated previously, the leaching of gold involves a solid-to-liquid mass transfer process
across the gold-solution interface. The process is inherently electrochemical in nature,
composed of the anodic oxidation of gold to form a singly charged ion that reacts to form a gold
complex, and cathodic reduction of the oxidant (typically oxygen)1. A diagram of the dissolution
mechanism can be found in Figure 2.2.
Figure 2.2. Schematic for the electrochemical model for leaching of gold (adapted from reference6).
Once oxidized, gold(I) reacts with the lixiviant, or leaching reagent, to form a soluble
gold complex that undergoes mass transport to the bulk solution. Current industrial practices
make use of cyanide as the lixiviant for complexation. Chlorine-chloride leaching was applied
commercially in the 19th century, but use diminished greatly upon introduction of the cyanide
13
process is 18891. Other lixiviants such as ammonia, sulfide, thiocyanate, thiourea and thiosulfate
have been investigated, however, the complexity of these systems limits their application in
industry1.
2.2.1 Cyanidation
The process of cyanidation includes the dissolution of gold into an aerated cyanide
solution to form a dicyano-auro complex. The reaction proceeds via the Elsner equation 3, 7-9:
4Au + 8CN! + O2 + H2O ! 4Au(CN)2! + 4OH! (2.1)
The overall reaction provided by the Elsner equation is the sum of the electrochemical oxidation
of Au0 to Au+, and the reduction of oxygen. If both the anodic and cathodic half reactions are
considered, the system can most accurately be described by Equations 2.2 and 2.3, which
proceed in parallel:
2Au + 4CN! + O2 + 2H2O ! 2Au CN( )2! + H2O2 + 2OH! (2.2)
2Au + 4CN! + H2O2 ! 2Au(CN)2! + 2OH! (2.3)
Returning to Figure 2.2 with the cathodic and anodic reactions in mind, we can see that
the oxidant in the case of cyanidation would be oxygen, which is reduced to hydrogen peroxide
and the hydroxide ion, both of which will undergo mass transport away from the gold surface. In
terms of the anodic reaction, once gold is oxidized from Au0 to Au+, an aurocyanide complex is
formed, which also undergoes mass transport to the bulk solution from the gold-solution
interface1.
The gold dissolution rate in alkaline cyanide solutions is dependent on a wide number of
factors, including cyanide concentration, dissolved oxygen concentration, temperature, pH, and
surface area of exposed gold. Much research has gone into the optimization of the cyanidation
process over the last 100 years since its discovery and implementation. Based on calculated
14
diffusion coefficients for both cyanide and oxygen, as well as observed experimental and
practical values, the optimal molar ratio of CN-: O2 ranges from 4:1 to over 7:1. In practice,
ratios greater than 6:1 are used to ensure that the cyanide concentration is not the rate limiting
factor. 1
Temperature has a strong effect on the gold dissolution rate in aerated alkaline cyanide
leaching solutions. As the temperature is increased, due to increased diffusion rates of reacting
species, the gold dissolution rate also increases up to a maximum of ~85°C. Above this
temperature, oxygen solubility decreases and results in an overall decrease in the gold dissolution
rate1. The pH of the leaching solution is typically kept above 9.4 to avoid hydrolysis of cyanide
and excess consumption, however, in practice the pH conditions are dictated by other process
factors such as the ore composition, and solubility issues. 1
Exposed surface area is directly proportional to the gold dissolution rate. This is directly
related to particle size distribution and liberation characteristics of the ore. In general, with
decreasing particle size, the gold dissolution rate greatly increases, due to a larger amount of
exposed gold1.
15
2.2.2 Environmental Concerns
The presence of copper in ore poses a large problem during the cyanide leaching process,
not only because of the consumption of cyanide due to formation of copper cyanide complexes,
but also due to the increased toxicity of these complexes to birds, animals and fish 3, 9. Recent
concerns over environmental and human health issues involved with cyanide leaching as a result
of collapsed tailings dams in Guyana, and Romania has lead to increased interest in non-cyanide
lixiviants3, 7, 9, 10. One such alternative is thiosulfate, which is not only non-toxic, but also useful
for ores that consume excessive amounts of cyanide or absorb the resulting complex.9
Carbonaceous type ores are susceptible to thiosulfate leaching due to its ability to degrade sulfide
matrices and prevent preg-robbing10, 11. Jeffrey, Breuer and Choo12 compared the relative
leaching rates of chloride, cyanide and thiosulfate system for both pure gold and gold/silver
rotating electrochemical quartz-crystal microbalance (REQCM) electrodes. The authors found
that when comparing the initial rates of leaching, thiosulfate began with a much greater leaching
rate (3.8 × 10-5 mol m-2 s-1) than that of cyanide (0.1 × 10-5 mol m-2 s-1). Also, thiosulfate was
able to leach both pure gold and gold/silver alloy, whereas cyanide was only effective in leaching
of the gold/silver alloy.
2.2.3 Thiosulfate Alternative
Thiosulfate as a chemical has been widely used in both photography and in the
pharmaceutical industry7. It was first proposed for use in leaching of precious metals in the early
1900s in the Von Patera process, wherein gold and silver ores were subjected to a chloridising
roast, then leaching was performed using thiosulfate10. In the late 1970s a process for the
application of ammoniacal thiosulfate to gold ores containing copper and metal sulfides was
developed and patented for commercial use13.
16
2.3 THIOSULFATE LEACHING OF GOLD
Leaching of gold by thiosulfate proceeds through two coupled half reactions: oxidation
of gold (Equation 2.4) and reduction of oxygen (Equation 2.5) 9, 14, 15.
Au + 2S2O32! " Au(S2O3)2
3! + e! (2.4)
O2 + 2H2O + 4e! " 4OH! (2.5)
The overall reaction is reminiscent of the Elsner equation:
4Au + 8S2O32! + O2 + 2H2O " 4Au(S2O3)2
3! + 4OH! (2.6)
The only apparent change in the overall reaction is the lixiviant that acts to sequester the gold in
solution. The anodic reaction in Equation 2.4 has a standard reduction potential of 0.153 V,
while the reduction potential of the cathodic reaction (Equation 2.5) is 0.401 V14. The oxidant is
still O2, and the reaction is typically carried out in alkaline solution due to the acid catalyzed
decomposition of thiosulfate7, 9. Although the initial kinetics of the leaching reaction are rapid,
over extended periods of time the kinetics of this reaction are quite slow mainly due to oxidation
of thiosulfate by oxygen, leading to formation of species such as sulfite, and formation of a
passive layer from decomposition of the thiosulfate ion 6, 16.
The decomposition of thiosulfate under oxidizing conditions can proceed through a
variety of pathways 14, 15, 17:
2S2O32! " S4O6
2! + 2e! (2.7)
S2O32! + 6OH! " 2SO3
2! + 3H2O + 4e! (2.8)
S2O32! + 2OH! " SO4
2! + H2O + S2! (2.9)
Upon formation of these species, a number of additional reactions may occur in solution
including degradation and rearrangement of tetrathionate (S4O62! ) to form trithionateS3O6
2! ,
17
thiosulfate, pentathionate (S5O62! ), and sulfite (Equations 2.10-2.12).17 Subsequent degradation
of trithionate (Equation 2.13) leads to an even more complex and dynamic system. 18-20
S4O62! + S2O3
2! " S5O62! + SO3
2! (2.10)
S4O62! + SO3
2! " S3O62! + S2O3
2! (2.11)
2S4O62! " S3O6
2! + S5O62! (2.12)
S3O62! + H2O " S2O3
2! + SO42! + 2H+ (2.13)
Zhang et al. 17 studied the rearrangement and degradation reactions of both tetrathionate
and trithionate in near-neutral solutions. The presence of excess thiosulfate is known to catalyze
the rearrangement reactions of polythionates (Equation 2.15), which can be described by a
general disproportionation reaction 17:
2SxO62! " Sx!1O6
2! + Sx+1O62! (2.14)
2SxO62! + S2O3
2! H+
OH!" #""$ """ Sx+1O6
2! + SO32! (2.15)
It is important to note here, that unlike other polythionates, trithionate does not rearrange in
accordance with this disproportionation scheme, as dithionate cannot be formed by interaction of
polythionate species21. The direction of Equation 2.15 is highly dependent on the pH of the
solution. At pH > 7, the presence of the sulfite drives the reaction to the left. When the pH is less
than 7, disulfite drives the reaction to the right17. In alkaline solutions, higher order polythionates
show significantly decreased stability. The presence of polythionates higher than hexathionate
are not detected in solutions near neutral pH, or in slightly alkaline solutions17. Pentathionate,
and hexathionate themselves have been detected in solution, however, their degradation is much
more rapid than tetrathionate17.
2S5O62! + 6OH! " 5S2O3
2! + 3H2O (2.16)
18
S6O62! " S5O6
2! + S (2.17)
The presence of these fairly reactive species results in an extremely dynamic and complicated
system. Some, or all, of these species may be responsible for the passive layer on the electrode
surface, effectively inhibiting any further leaching of gold 22.
An investigation into the electro-oxidation of thiosulfate on gold by Pedraza et al. 23
showed that the passivating layer was composed of both elemental sulfur and some form of
oxidized sulfur species. Cyclic voltammetry in a solution of 0.01 M Na2S2O3 and 0.1 M Na2SO4
(the supporting electrolyte) displayed a prominent oxidation peak at 0.44 V versus a saturated
calomel electrode (SCE) that was later assigned to the formation of a sulfide film on the surface
of the gold electrode23. The adsorbed sulfide layer likely formed through the reactions described
by Equations 2.18-2.20: 23
S2O32! + 6OH! " 2SO3
2! + 3H2O + 4e! (2.18)
S2O32! + 2OH! " SO4
2! + H2O + S2! (2.19)
S2! " S + 2e! (2.20)
The second component of the film, oxidized sulfur species, arises not from electro-oxidation of
thiosulfate but from decomposition of thiosulfate at open circuit potential and subsequent
oxidation of these species23, as described by Equations 2.10-2.17.
The hypothesis of Pedraza and co-workers was verified by the experimental results of
Jeffrey et al. 24. Upon application of a triangular potential scan to a gold electrode in contact
with a solution of 0.1 M Na2S2O3, the authors noted an increase in the measured current,
beginning at approximately 0.4 V, with a peak at ~0.73 V. In parallel, a surface enhanced Raman
(SERS) spectrum was collected to track the expected νS-S and δS-S-S bands of elemental sulfur. As
the voltammetric current rose, evidence of peaks corresponding to the expected sulfur stretches
19
appeared in the collected spectrum, beginning at approximately 0.44 V and becoming more
intense as the potential was increased in the positive direction. The increasing intensity of these
bands was believed to correspond to the formation of some form of polymeric sulfur adsorbed on
the gold surface. 24
The presence of oxidation and decomposition products of thiosulfate on the gold surface
was also confirmed by Woods et al. 25 through the collection of a SERS spectrum of 0.1 M
Na2S2O3 at open circuit potential (Figure 2.3).
Figure 2.3: SERS spectrum of 0.1 M Na2S2O3 at open circuit potential. Modified from 25.
At open circuit potential, bands corresponding to the νS-S and νsym(S-O) vibrations of
thiosulfate are seen at 445 cm-1 and 999 cm-1. 25 The authors also noted the presence of bands at
378 cm-1 and 1033 cm-1, which are within 10 cm-1 and 7 cm-1 respectively of the νS-S and νsym(S-O)
vibrations of tetrathionate. Due to this difference in the band position, it was proposed that these
20
bands arose from a polythionate species adsorbed on the gold surface, with a longer chain length
than tetrathionate 25.
Recent work by J. Baron Gavidia proved that tetrathionate was present in the passive
layer forming on the gold surface 6. Time dependent SERS studies of a gold nanorod electrode in
a 0.1 M Na2S2O3 + 1.0 x 10-4 M NaOH solution were conducted at open circuit potential, the
results of which are shown in Figure 2.4. Using the Fourier self-deconvolution method (FSD),
bands at 258, 315, 378 and 1025 cm-1 were identified within the first five minutes of immersion.
These bands were assigned to tetrathionate in view of the fact that adsorption on to the gold
surface would shift the band positions from those typically seen in the solution spectra 6.
Figure 2.4: SERS spectrum of a gold nano-rod electrode submersed in 0.1M Na2S2O3 + 1x 10-4M NaOH
at open circuit potential (modified from 6).
After 630 minutes, the peaks corresponding to tetrathionate were no longer present, but
new bands at 265 and 420 cm-1 arose. These were attributed to the formation of trithionate on the
surface through decomposition of tetrathionate: 6, 19
2S4O62! + 3OH! " 5
2S2O3
2! + S3O62! + 3
2H2O (2.21)
21
2.3.1 Ammonia-Thiosulfate Leaching
The addition of ammonia into thiosulfate leaching solution results in the formation of an
anionic aurocomplex that is stable over a wide pH range. Moreover, ammonia helps to inhibit
interference from other species such as iron oxides, silica, silicates and carbonates, and increases
the rate of gold oxidation 16, 26. This was demonstrated in the work of Breuer and Jeffrey 27.
These authors studied the effect of the concentration of ammonia on gold dissolution rate by use
of a rotating electrochemical quartz crystal microbalance. The results distinctly showed that with
increasing ammonia concentration, the calculated current, and thus the dissolution rate, was
greatly increased (Figure 2.5) 27.
Figure 2.5: Calculated voltammetric current for varying concentrations of ammonia in thiosulfate
leaching solution. Adapted from 27.
The role of ammonia is still not fully understood, but it is believed that it prevents
passivation of the gold by forming a complex with Au+, and subsequently assisting in gold
dissolution13. However, it has been shown that in the absence of thiosulfate, no dissolution of
22
gold into the solution occurs in the presence of ammonia. Breuer and Jeffrey27 proposed that
rather than just ammonia and gold, it is a gold-amine-thiosulfate complex that affects passivation
of the gold surface during the leaching reaction.
Ammonia’s proposed catalytic role is defined by an adsorption/desorption/stabilization
process28. This mechanism corroborates the assumption of Breuer and Jeffrey that it is a gold-
amine-thiosulfate complex that is responsible for enhanced leaching27. The process begins with
the adsorption of ammonia and thiosulfate to the gold surface28:
a) Au(s ) + S2O3
2! + NH3 " Au(S2O3)(NH3)(ads )2!
b) Au(s ) + NH4S2O3! + NH3 " Au(S2O3NH4 )(NH3)(ads )
! (2.22)
Oxidation of the adsorbed species then occurs28:
a) Au(S2O3)(NH3)(ads )
2! " Au(S2O3)(NH3)(ads/aq)! + e!
b) Au(S2O3NH4 )(NH3)(ads )! " Au(S2O3NH4 )(NH3)(ads/aq)
0 + e! (2.23)
Followed by desorption of the now oxidized species28:
a) Au(S2O3)(NH3)(ads/aq)
! + S2O32! " Au(S2O3)2 (aq)
3! + NH3
b) Au(S2O3NH4 )(NH3)(ads/aq)0 + NH4S2O3
! " Au(S2O3)2 (aq)3! + NH3 + 2NH4
+ (2.24)
This mechanism was supported by experimental evidence showing that the rate of
oxidation of gold in ammoniacal thiosulfate solutions was highly dependent on the concentration
product ([M2S2O3][NH3])0.8, where M = Na, NH4+. Also, the concentration of the NH4S2O3
- ion
pair was equal to, or greater than, the concentration of free thiosulfate in solution. 28
Ammonia has also been shown to help in regeneration of thiosulfate from leaching
solutions through ammonolysis of degradation compounds such as trithionate:
a) S3O62! + NH3 ! S3O6 "NH3
2!
b) S3O6 "NH32! # S2O3
2! + SO3NH2! + H+
(2.25)
23
Addition of copper is required in ammonia solutions due to the inability of O2 to act as
the oxidant 16. However, copper is also known to affect gold dissolution without ammonia.
Zhang and Nicol 29 investigated the effect of several different forms, and concentrations of
copper on the gold dissolution reaction (Figure 2.6). The authors found that addition of any form
of copper increased the amount of gold dissolved, but the addition of 0.5 mM CuSO4 resulted in
the highest amount of gold recovered. Based on the work of Rabai and Epstein 30, Zhang and
Nicol 29 proposed that in the absence of ammonia, the oxidation of gold is accomplished through
a copper-thiosulfate-oxygen intermediate, [Cu(S2O3)3O2]5-.
Figure 2.6: The effect of copper on gold dissolution in thiosulfate leaching. Adapted from29.
The intermediate has been observed previously using UV-VIS, but not at the gold-
solution interface. Another suggested role of copper in aiding gold dissolution was the
scavenging of sulfide from the surface of gold, resulting in at least partial removal of the passive
layer. 29
24
One of the best-known methods to limit the formation of the passive layer, and ensure
relatively high gold oxidation rates, is the addition of both ammonia and copper into the solution,
where Cu2+ acts as the oxidant 13, 16, 26, 31. The presence of both of Cu2+ and NH3 is known to
increase the rate of gold dissolution31. Once in solution, Cu2+ and NH3 combine to
form Cu(NH3) 42+, which is believed to act as the oxidant 9, 13, 16, 26, 32. The overall reaction for
the dissolution of gold then becomes13, 15, 26, 33:
Au + 5S2O32! + Cu(NH3)4
2+ " Au(S2O3)23! + 4NH3 + Cu(S2O3)3
5! (2.26)
The reduction potential of the cupric tetra-ammine complex given by Watling 34 is +0.22V, and
can be described by19, 22, 33, 34:
Cu(NH3)42+ + 3S2O3
2! + e! " Cu(S2O3)35! + 4NH3 (2.27)
The reduction of Cu2+ in solution is coupled with the corresponding oxidation of gold, as seen in
Equation 2.4. Once generated, Cu+ can be oxidized back to Cu2+ through coupling to the
reduction of O2 13, 22, 33:
2Cu(S2O3)35! + 8NH3 + 1
2O2 + H2O " 2Cu(NH3)4
2+ + 2OH! + 6S2O32! (2.28)
The proposed mechanism for the dissolution of gold in thiosulfate media containing both
ammonia and copper involves adsorption and desorption of species involved in both the cathodic
and anodic reactions, with a rate-determining redox reaction26, 31. The electrochemical
mechanism is illustrated in Figure 2.7.
Addition of ammonia and copper into thiosulfate leaching solution is known to result in
an increased initial leaching rate, and increased gold recovery over the duration of leaching;
however, the mechanism and exact effect of each additive on the passive layer is still unknown.
25
Figure 2.7: Electrochemical model for the leaching of gold using ammoniacal thiosulfate leach media. Adapted from 26, 31.
2.3.2 The Role of the Electrolyte on Ammonia-Thiosulfate Leaching
The role of the metal cation of the thiosulfate salt was investigated by Chandra et al. 35,
using a polycrystalline solid gold electrode in combination with an REQCM. The authors
measured the gold oxidation polarization curves of the electrode in the presence of ammonium,
sodium and potassium thiosulfate. A significant increase in the current was noted as the metal
cation was changed from sodium to potassium. Two possible explanations were proposed:
i. Heavier cations (such as potassium) would be less hydrated compared to ions with
smaller atomic mass (such as sodium) and therefore would be more likely to be adsorbed
onto the gold surface. 35
26
ii. The dissociation constants for alkali metals with thiosulfate are known to decrease in the
order of Na > K > Rb > Cs, therefore, one would expect potassium ions to more readily
form an ion pair with either gold-thiosulfate, or free thiosulfate.35
However, no explanation yet exists for how either of these occurrences would influence the gold
oxidation reaction. 35
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27
21. Nickless, G. In Inorganic sulphur chemistry; Elsevier Publishing Company: 1968; .
22. Chu, C. K.; Breuer, P. L.; Jeffrey, M. I. Minerals Eng 2003, 16, 265-271. 23. Pedraza, A. M.; Villegas, I.; Freund, P. L.; Chornik, B. J Electroanal Chem 1988,
250, 443-449. 24. Jeffrey, M.; Watling, K.; Hope, G. A.; Woods, R. Minerals Eng 2008, 21, 443-
452. 25. Woods, R.; Hope, G. A.; Watling, K. M.; Jeffrey, M. I. J. Electrochem. Soc. 2006,
153, D105-D113. 26. Aylmore, M. G.; Muir, D. M. Miner Metall Process 2001, 18, 221-227. 27. Breuer, P. L.; Jeffrey, M. I. Hydrometallurgy 2002, 65, 145-157. 28. Senanayake, G. Hydrometallurgy 2005, 77, 287-293. 29. Zhang, S. C.; Nicol, M. J. J. Appl. Electrochem. 2005, 35, 339-345. 30. Rabai, G.; Epstein, I. R. Inorg. Chem. 1992, 31, 3239-3242. 31. Senanayake, G. J. Colloid Interface Sci. 2005, 286, 253-257. 32. Molleman, E.; Dreisinger, D. Hydrometallurgy 2002, 66, 1-21. 33. Arima, H.; Fujita, T.; Yen, W. Miner Metall Process 2003, 20, 81-92. 34. Watling, K. M. Spectroelectrochemical Studies of Surface Species in the
Gold/Thiosulfate System, Griffith University, Australia, 2007. 35. Chandra, I.; Jeffrey, M. I. Hydrometallurgy 2004, 73, 305-312.
28
CHAPTER 3: EXPERIMENTAL TECHNIQUES THEORY
3.1 ELECTROCHEMICAL TECHNIQUES
3.1.1 Sweep Voltammetry Methods
Sweep voltammetry methods can be invaluable for determination of the electrochemical
behavior of a system. A potential varying linearly with time is applied to the system, while the
output current is simultaneously recorded as a function of the applied potential. Figure 3.1a
shows a typical potential-time curve for cyclic voltammetry, and 3.1b the current response curve.
E1
E3
E2
t
E (-)
E0’ E!
A + e- A!-
A + e- A!-
a) b)
Figure 3.1: a) Typical potential time curve for cyclic voltammetry b) Current response curve. Adapted
from Bard and Faulkner1.
Sweep rate is defined as the slope of the potential applied as a function of time, and can
range from 1mV/s to 1000V/s. Starting from potential E1, a voltage ramp is applied to a higher
value, E21. At this point the potential is reversed and cycled back to either the same potential
(E1), or to a different potential, E3, where the sweep ends.
29
In an oxidation reaction or reduction reaction occurring at the surface of an electrode, the
electron transfer across the metal-solution interface produces a current that is defined as Faradaic
current. However, non-Faradaic processes such as charging of the capacitor at the interface may
also occur, resulting in current flow or electrode polarization without a net charge transfer across
the metal electrolyte-interface.
The metal-solution interface behaves as a capacitor. The solution side of the metal
electrolyte interface is composed of several layers, containing all charged species and dipoles.
The simplest model of this region is called the electrical double layer (Figure 3.2). It consists of
the inner and diffuse parts. The inner layer is a region between the metal surface and the outer
Helmholtz plane (OHP) that corresponds to a distance x2. This distance is the point of nearest
approach to the electrode surface, for solvated ions. The inner layer includes solvent molecules
and specifically adsorbed ions. The locations of the specifically adsorbed anions define the
position of the inner Helmholtz plane (IHP), at a distance x1. Beyond the OHP, and extending
into the bulk, is the diffuse layer. Here, solvated ions interact with the charged metal surface only
through long-range electrostatic interactions. Ions in the diffuse layer are treated as point
charges, meaning that their properties are essentially independent of the chemical properties of
the ion.
Current generated through electron transfer of redox reactions follows Faraday’s law:
the amount of a chemical consumed or produced in the electrode reaction is proportional to the
charge that has passed through the interface2. If we consider the reaction
O+ ne! " R (3.1)
30
Figure 3.2: Schematic of the structure of the electrical double layer formed at the metal-electrolyte
interface. Adapted from Bard and Faulkner1.
Where O is the oxidized species and R is the reduced species, the rates of the forward
reaction (reduction) and the reverse reaction (oxidation) are given by:1
! f = k fCO = icnFA
(3.2)
!b = kbCR =ianFA
(3.3)
Where νf is the rate of reduction; νb the rate of oxidation; CO is the surface concentration of the
oxidized species; CR is the surface concentration of the reduced species; n is the number of
electrons transferred in the reaction; A is the area of the electrode and F is Faraday’s constant (96
IHP OHP
Metal
+
+
+
+
-
-
Solvent molecule
x1 x2
31
485.3399 C mol-1). The rate constants of the forward (kf) and reverse reaction (kb) are potential
dependent, as given by Equations 3.4 and 3.5: 1
k f = k0 exp !"nF
RTE ! E0( )#
$%&'( (3.4)
kb = k0 exp
1!"( )nFRT
E ! E0( )#$%
&'(
(3.5)
where α, k0, and E0 are the transfer coefficient, standard rate constant and standard potential,
respectively. The standard rate constant is a measure of the forward or reverse reaction rates
when the system is at equilibrium at standard conditions. Its value indicates the kinetics of a
redox couple; a system with a large standard rate constant will rapidly achieve equilibrium,
whereas a system with a small k0 would be sluggish.
Measured current is the sum of the cathodic and anodic currents:
i = ic ! ia = nFA kfCO (0,t)! kbCR(0,t)"# $% (3.6)
Therefore, by combining the expressions for the rate constants and the potential dependence of
the two reactions, an equation for the current density at any potential can be derived1.
iA= nFk0CO exp
!"nFRT
E ! E0( )#$%
&'(! nFk0CR exp
1!"( )nFRT
E ! E0( )#$%
&'(
(3.7)
If we consider the interface at equilibrium with the solution, in which the bulk concentration of
oxidized species is equal to the surface concentration, and the surface concentration of the
reduced species is equal to the bulk, it follows that E = Eeq , k fCO = kbCR , and k f = kb = k0 . By
definition, the net current at this point is zero, and the electrode adopts a potential based purely
on the bulk concentrations of the oxidized and reduced species. However, this does not mean that
charge transfer across the interface has stopped. The magnitude of the anodic and cathodic
currents at equilibrium are equivalent, such that:
32
nFAk0CO exp!"nF( )RT
(Eeq ! E0 )#
$%&'(= nFAk0CR exp
1!"( )nFRT
(Eeq ! E0 )#
$%&'(
(3.8)
and are defined by the exchange current:
i0 = FAk0CO exp
!"nFRT
Eeq ! E0( )#
$%&'(
(3.9)
Dividing both sides of Equation 3.7 by the exchange current (Equation 3.9), and rearranging for
i, we achieve:
i = nFAk0CO exp!"nFRT
Eeq ! E0( )#
$%&'(exp !"nF
RTEeq ! E( )#
$%&'(! exp
1!"( )nFRT
Eeq ! E( )#$%
&'(
)*+
,-.
(3.10)
The first term in Equation 3.10 is the exchange current as seen in Equation 3.9, and using the
definition of overpotential (! = Eeq " E ), Equation 3.10 can be simplified to:
i = i0 exp!"nFRT
#$%&
'() ! exp
1!"( )nFRT
#$%&
'()
*
+,
-
./ (3.11)
Equation 3.11 is known as the Butler-Volmer equation, and is a description of the current when
the system is out of equilibrium, in the absence of mass transfer effects; that is, when i is < 10%
of the limiting currents. When the overpotential is small, the exponential terms can be estimated
as ex = 1+ x and e! x = 1! x . Using this approximation, Equation 3.11 can then be written as
follows1:
i = !i0nFRT
" (3.12)
The net current is linearly proportional to the overpotential in a narrow range near the
equilibrium potential. The negative reciprocal of the slope in this region has units of resistance,
and is often termed the charge transfer resistance:
33
! "#"i
= Rct =RTi0nF
(3.13)
A plot of log(i) vs. η yields a Tafel plot, a common data form used for analysis of the kinetic
aspects of a redox system. Figure 3.3 displays an example of a Tafel plot for a typical redox
system.
Figure 3.3: Tafel plot of a redox system, modified from Bard and Faulkner.1
When the overpotential is greater than ~60 mV on either the cathodic or anodic side, one
of these contributions becomes negligible. For example, at large negative overpotentials, the
cathodic contribution dominates, i.e. exp !"nFRT
#$%
&'(>> exp (1!" )nF
RT#$%
&'(
, and the anodic term in
Equation 3.11 can be ignored, meaning the Butler-Volmer equation can now be represented as:
i = i0 exp !"nFRT
#$%
&'(
or log i = log i0 !"nFRT
(3.14)
The exchange current can also be determined graphically, as shown in Figure 3.3, by
interpolation of the linear regions at large overpotential, to zero overpotential. The intersection of
the two linear regions gives the log of the exchange current.
-50 -100 -150 -200 50 100 150 200
slope = !"nF2.3RT
slope =1!"( )nF2.3RT
log i0
log |i|
-3.5
-4.5
-5.5
-6.5
!, mV
34
The values of the transfer coefficients can be calculated using the slope of the linear
regions of the Tafel plot. As in Figure 3.3, the slopes of the linear segments at large overpotential
are given by 1!"( )nF2.3RT
and !"nF2.3RT
for the anodic and cathodic branches respectively. From
these expressions, one can calculate the values of the transfer coefficients.
The transfer coefficient is a measure of the symmetry of the intersection of the free
energy curves versus the reaction coordinate for the products and reactants. The reaction
coordinate is variable, for example it could be the torsional angle around a bond, but in
electrochemistry, it is typically considered as the distance of the oxidized species from the
interface. Consider again a one-electron, one-step process in which an oxidized species, O, reacts
with a single electron to form the reduced species, R. The free energy curves of O + e and R are
represented in Figure 3.4. Using the equilibrium potential, Eeq, as a reference point, an applied
potential will shift the O + e energy curve either up or down according toF E ! Eeq( ) . Figure 3.4
shows a positive energy change. Consequently, the activation energy for the oxidation of the
reduced species, R, will also shift by a fraction of the total energy change, 1!!( )F E ! Eeq( ) .
The shift in the activation energy is therefore dependent on the value of the transfer coefficient,
which can range in value from zero to unity depending on the symmetry of the potential energy
curve intersection.
35
Figure 3.4: Diagram of the free energy curves of an oxidized and reduced species when a potential is
applied, shifting the system from equilibrium. On the right is a zoom in of the intersection of the two
potential energy surfaces. Adapted from Bard and Faulkner. 1
If the area of intersection is considered locally linear (Figure 3.5), a definition of the
transfer coefficient can be generated in terms of the geometry at that point.
Figure 3.5: Intersection of the locally linear potential energy surface of an oxidized and reduced species.
Adapted from Bard and Faulkner.1
!
! !
!FE
1!"( )FE
R
O + eE = E
E = 0
Length = x
Reaction Coordinate
36
Using the definition of the angles ϕ and θ:
tan! = "FEx
(3.15)
tan! =1"#( )FE
x (3.16)
The transfer coefficient can be defined as:
! = tan"tan# + tan"
(3.17)
A symmetric intersection of the two potential energy surfaces leads to the angles ϕ and θ being
equivalent, hence α is equal to ½. If the intersection is asymmetric, the value of α will either lie
between 0 and ½ or ½ and 1.
3.1.2 Mixed Potential Theory
The measured current is the sum of both the anodic (oxidation) and cathodic (reduction)
reactions. In the case of gold leaching, this current is the sum of the anodic current generated
from the oxidation of gold, and the cathodic current is generated by the reduction of oxygen or
Cu(II). Figure 3.6 is a representation of typical linear sweep voltammogram of a Au leaching
system. The point at which the curve crosses the x-axis corresponds to a value of zero for the
current. Here, the current contribution from gold oxidation is equivalent in magnitude, but
opposite in sign to the contribution of the reduction of oxygen. The potential at this junction is
defined as the mixed potential. The magnitude of either the cathodic or anodic current at the
mixed potential is known as the leaching current, iM. Using a Tafel-like plot, that is, the
logarithm of the current versus the overpotential, the mixed potential can be easily identified
through extrapolation of the respective currents and their intercept. This is displayed in Figure
3.7.
37
Figure 3.6: Generic linear sweep voltammogram. Dashed lines represent the individual contribution of
the cathodic and anodic currents, while the solid line represents the sum of these two contributions.
Figure 3.7: Tafel plot a gold electrode in contact with a 0.1 M Na2S2O3 + 0.01 M CuSO4 + 1.1µΜ
Ca(OH)2 solution.
Electrode potential
Cur
rent
den
sity
(EM, iM)
-0.26 -0.24 -0.22 -0.20 -0.18 -0.16 -0.14 -0.12 -0.10-2.5
-2.0
-1.5
-1.0
-0.5
0.0
0.5
1.0
1.5
2.0
Log
(i)
Potential / V vs. SCE
Mixed Potential
38
A modified version of the Butler-Volmer equation devised by J. Y. Baron3 can be used to
determine the leaching current. In the case of gold leaching, the exchange current (i0) in the
above equation is equivalent to the leaching current (iM). Because thiosulfate leaching of gold
involves different electron transfer reactions in the cathodic and the anodic processes, the charge
transfer coefficients are different: for copper and gold they are αCu2+ and αAu, respectively. Using
these transfer coefficients, the modified Butler-Volmer equation is:
i = iM exp!"
Cu2+nCu2+F
RT#
$%&
'()! exp
1!"Au( )nAuFRT
#$%&
'()
*
+,
-
./ (3.18)
Because the reaction involves oxidation of gold, and reduction of Cu(II) to Cu(I), the number of
electrons transferred, i.e. n, is equal to 1. When overpotentials are small, the equation can be
simplified to:
i = iM !"Cu2+
! 1!"Au( )( ) #FRT (3.19)
In some instances, the transfer coefficients can be approximated to 0.5, thus allowing a
simplification of the equation. 1 However, in the instance of gold leaching in the presence of Cu,
the transfer coefficients must be calculated as described previously, before the equation can be
further simplified.
By plotting the current versus overpotential in a small potential range (±20mV) on both
sides of the mixed potential, the leaching current can be obtained. A linear fit on the resulting
plot (as show in Figure 3.8) gives the slope, equivalent to the ratio of the current to potential.
39
Figure 3.8: Plot of the current versus the overpotential in a ~40mV range around the mixed potential for
a 0.1 M Na2S2O3 + 0.01 M CuSO4 + 1.1µΜ Ca(OH)2 system.
Combining equations 3.13 and 3.19, and using the calculated slope, the leaching current can be
found:
Rct = ! "#"i
= 1slope
= RTiMF !$
Cu2+ ! 1!$Au( )( )% iM =
slope RT( )F !$
Cu2+ ! 1!$Au( )( ) (3.20)
This method of calculating the leaching current, devised by J. Y. Baron3, will be used to interpret
data obtained from linear and cyclic sweep voltammetry experiments.
3.1.3 Hydrodynamic Methods
Hydrodynamic methods are those involving forced convection of the system, where
either the electrode itself is in motion (such as with a rotating disc electrode), or the solution is
-0.18 -0.17 -0.16 -0.15 -0.14
-1.5
-1.0
-0.5
0.0
0.5
1.0
j / µ
A c
m-2
Potential / V vs. SCE
Δi/Δη
40
forced past a stationary electrode. Techniques using hydrodynamic methods are useful because
the system reaches a steady state fairly rapidly, and effects from double-layer charging are
eliminated, while the contribution from mass transport on the electron transfer reaction is
minimized.
Two types of fluid-flow are typically considered. If the solution can be characterized by
smooth and steady flow in planar layers, the flow is said to be laminar (Figure 3.9a). Flow at the
edges will have the lowest velocity due to friction, while solution at the centre will have the
greatest velocity. Hence a velocity profile for a laminar solution will be parabolic in shape.
Figure 3.9: a) Profile for a solution with laminar flow b) Profile for a solution with turbulent flow.
Adapted from Bard and Faulkner.4
Turbulent flow is characterized by chaotic currents without well-defined layers, only an average
net flow in a particular direction.
The type of flow can be determined through calculation of a factor known as the
Reynold’s number. The Reynold’s number is a dimensionless measure of the ratio of inertial
forces of a fluid, to the viscous forces of the fluid, and is defined by Equation 3.21:
Re = !chlv
(3.21)
Laminar Flow Turbulent Flow
a b
41
Where υch is the characteristic velocity of the fluid, l the characteristic length, and ν the
kinematic viscosity. Although this number is dimensionless, the value is proportional to fluid
velocity, thus a high Reynold’s number implies either high flow rate or electrode rotation rates.
Determination of either laminar, or turbulent flow is based on a critical number, Recr. Below this
value, flow is considered laminar, while above flow is considered turbulent4. Typical values for
Recr lie around 105-106.5
Rotating disk electrodes are the most commonly used electrodes for hydrodynamic
methods. Typically they consist of a round metal disk inserted into a sheath of Teflon, or inert
material. The disk is connected to a motor, which allows rotation at the desired rate.
Three types of reactions are typically encountered during electrochemical hydrodynamic
experiments: those which are mass transport controlled, those which are under both kinetic and
mass transport control, and those which fall only under kinetic control. Reactions under only
kinetic control were discussed previously in terms of the Butler-Volmer model, and thus the
following discussion will only concern systems under mass transport control, or mixed control.
Limiting current for a mass transport controlled reaction is given by the Levich equation:
il ,c = 0.62nFADO2/3! 1/2" #1/6CO
$ (3.22)
Where n is the number of electrons involved in the redox reaction, D is the diffusion coefficient
of the electroactive species, ν is the kinematic viscosity, CO*
is the bulk concentration of the
oxidized species, ω is the angular velocity in s-1, F is the Faraday constant and A is the area of
the electrode. The Levich equation also dictates that the limiting current for the reaction is
proportional to the square root of the angular velocity. Thus a plot of i vs ω1/2 should provide a
straight line with an intercept of zero, if the reaction is indeed under full mass transport control.
42
Deviations from linearity in the above plot are indicative of the system moving from
under pure mass transport control, to that of mixed control. A system whose reaction rate is
limited by both mass transport and kinetic limitations can be described by the Koutecký-Levich
equation:
1i= 1iK
+ 1il ,c
= 1iK
+ 10.62nFADO
2/3! 1/2" #1/6CO$ (3.23)
Where iK is the current in the absence of mass transport effects, i.e. the current that flows only
under a kinetic limitation4.
3.2 SPECTROSCOPIC TECHNIQUES
3.2.1 Surface Enhanced Raman Spectroscopy
Raman spectra arise from the irradiation of a sample with photons in the visible or UV
regions of the electromagnetic spectrum. Scattered radiation may be collected parallel to, or at a
90° angle relative to the incident radiation6, 7. Interaction of the source light with the
polarizability of the molecule is responsible for the generation of an induced dipole moment.
This induced moment results in an excitation that is un-quantized, meaning that the molecule can
be in any one of an infinite number of virtual states between the ground electronic, and first
excited electronic state6.
If the collision between the photon and molecule is elastic, the escaping radiation is of the
same energy as the incident radiation and is defined as Rayleigh scattering. Elastic collisions
occur with the highest probability, thus the Rayleigh band is the central and most intense band in
the Raman spectrum. However, not all collisions are elastic, and thus radiation can be emitted at
either a higher or lower energy than the Rayleigh band6, 7. This is called Raman scattering. The
43
energy lost or gained in the radiation corresponds to the energy difference between the ground
and first excited vibrational energy of the molecule6.
Raman scattering can be explained by consideration of the behavior of a molecule in the
presence of an external electric field (in this instance, produced by the incident radiation):
EF = E0 cos 2!"ext( ) (3.24)
Where νex is the frequency of the exciting light, and E0 is the amplitude of the wave6, 7.
Interaction of the molecule’s electron cloud with the electric field of the incident radiation results
in the formation of an induced dipole moment. The magnitude of the dipole moment is
proportional to the molecule’s polarizability (αP):
µ =!PEF =!PE0 cos 2!"ext( ) (3.25)
In order to be classified as Raman active, the polarizability must vary as a function of the
internuclear distance (r):
!P =!0 + r ! req( ) "!P
"r#$%
&'( (3.26)
Where α0 is the polarizability at the equilibrium position req. The internuclear distance r ! req( ) varies with the frequency of vibration of the molecule, νv:
r ! req( ) = rm cos 2"#vt( ) (3.27)
In the above equation, rm is the maximum internuclear separation relative to the equilibrium
position. Substituting Equation 3.27 into Equation 3.26 we obtain:
!P =!0 +!!P
!r"#$
%&' rm cos 2!"vt( ) (3.28)
The induced dipole moment can then be defined by:
44
µ =!0E0 cos 2!"ext( )+ E0rm !!P
!r"#$
%&' cos 2!"vt( )cos 2!"ext( ) (3.29)
If we apply the trigonometric identity: cos xcos y = cos x + y( ) + cos x ! y( )"# $% 2 , the previous
equation for the induced dipole moment simplifies to:
µ =!0E0 cos 2!"ext( )+ E0rm2
!!P
!r"#$
%&' cos 2! "ex +!v[ ]t( )+ E0rm2
!!P
!r"#$
%&' cos 2! "ex (!v[ ]t( ) (3.30)
The three terms in the above equation represent the two types of scattering observed; the first
term represents oscillation of the dipole at the same frequency as that of the incident radiation,
resulting in Rayleigh scattering.This means that the emitted radiation also matches the frequency
of the incident. Rayleigh scattering accounts for the largest percentage of scattering that occurs.
The second and third terms in Equation 3.30 describe a modulation of the incident frequency by
the vibrational frequency of the bond, resulting in Raman scattering 6. The term containing
!ex "!v( ) represents what is known as Stokes lines, and that containing the term !ex +!v( ) , anti-
Stokes lines. Due to absorption of energy by the molecule, Stokes lines are found at lower
wavenumber than the Rayleigh line, and anti-Stokes lines, due to emission of energy by the
molecule, are found at higher wavenumber than the Rayleigh line 6. Figure 3.10 displays each of
the types of emission seen in Raman spectroscopy, and their origin.
45
Figure 3.10: Transitions seen in Raman spectroscopy. Adapted from6
Although useful to obtain fingerprint identification of molecules, a major limitation of Raman
spectroscopy is its low signal intensity. At most, the intensity of the scattered radiation is 0.001%
of the incident6.
Fleischmann et al. first discovered the benefits of surface enhanced Raman spectroscopy
in 1974 when a large increase in the signal strength of pyridine adsorbed on a roughened silver
electrode was observed 8, 9. At the time, this was attributed to the increase in surface area;
however, Albrecht and Creighton10, and independently Van Duyne and Jeanmaire11, realized that
the increased surface area was too small to cause enhancement over a factor of ~106 for the
predicted scattering cross section for pyridine12, thereby accrediting the extra increase to a
surface enhancement.9
Rayleigh Scattering Raman Scattering
Stokes Shift anti-Stokes Shift
E = h! E = h! + "E
!E
!S = !ex "!v!R = !ex !aS = !ex +!v
0 1 2 3
46
The greatest surface enhancement occurs on roughened metal surfaces of silver, gold,
copper and other coinage metals12; however, enhancement does still occur on metals such as
platinum, aluminum, rhodium and indium9, 12. It is believed that surface enhancement is a
combination of two contributions: an electromagnetic enhancement, and a chemical
enhancement13.
Chemical enhancement is believed to arise from the chemisorption of the adsorbate onto
the metal surface. When a molecule adsorbs onto a metal surface, it undergoes an electronic
rearrangement that allows enhancements for different vibrational modes, even in the ground
electronic state12. Upon chemisorption, it is proposed that a charge-transfer complex is formed
between the metal atoms on a rough surface and the molecule. Because of complex formation,
the energies of the HOMO and LUMO of the adsorbed molecule are close in energy to the Fermi
level of the metal. If the energy difference between the frontier orbitals and the Fermi level is
close to the same frequency of the exciting radiation, resonant enhancement may occur14. A
photon-driven charge transfer can also occur to increase enhancement, as shown in Figure 3.11.
Here, an adsorbed molecule is irradiated with light leading to the formation of an electron hole
pair. Electron tunneling occurs from the metal to the molecule, where vibrational information is
imparted to the electron, if the vibrational frequency matches the residence time within the
molecule. The electron then recombines with the hole, and a Raman photon is emitted with a
frequency modulated by the energy of the vibrational levels of the molecule14.
47
Figure 3.11: Photon-driven charge-transfer process. Adapted from14.
There is some debate in literature as to the importance of the chemical contribution to the
enhancement observed in SERS15, as electromagnetic theory can describe the major features,
including the exceptional intensity for alkali and coinage metals, a dependence on the nano-
structure of the system, the dependence on the aggregation of nanoparticles in the system, and
the polarization dependence16. The electromagnetic contribution to the enhancement arises from
a highly localized concentration of surface plasmons in metal surfaces with nano-sized surface
features17, 18. Plasmons are defined as the collective oscillations of the band of conduction
electrons against the background of the ionic metal cores16. When surface features of the metal
are smaller than the wavelength of the exciting light, dipolar and higher multipolar plasmons can
be excited, or undergo transitions. Many of the higher order multipolar plasmons are non-
radiative, and depending on the size of the feature, or particle, all plasmons other than the dipolar
Electron Hole
CT h!
After t = !
h! " h!v
48
plasmon, can be ignored. If the plasmon is resonant with the incident radiation, these plasmons
are sustained by the local electronic structure of the metal13, 16, 18. The more “free” the conduction
band electrons are in the metal, the sharper and more intense the excitation of dipolar plasmon’s
resonance will be. The excited plasmons may become confined within the nano-sized features,
creating a highly localized electromagnetic field9, 13, 16, 19, that can then interact with the incident
radiation.
The enhancement observed for a metal nanoparticle is highly dependent on a number of
factors, including the dielectric constants of both the environment of the particle and the particle
itself, as well as its aspect ratio9, 13, 17. Equation 3.31 describes the electric field produced outside
a spherical nanoparticle of radius a, when irradiated with z-polarized light of wavelength λ13:
Eout x, y, z( ) = E0 z !!ME0zrd3 !3zrd5 xx + yy + zz( )"
#$
%
&' (3.31)
Where x!, y!, z!( ) are the Cartesian unit vectors, rd the radial distribution, and αM the metal
polarizability given by!M = ga3 , where g is defined as13:
g = ! in " !out! in + 2!out( ) (3.32)
with εin, εout the dielectric constant of the nanoparticle, and the surrounding environment,
respectively. It can be seen from Equation 3.31 that the electric field enhancement decays with
rd-3, meaning there is a finite sensing volume around the nanoparticle, and that surface
enhancement is a near-field effect9, 13, 16, 17, 19.
Using Mie theory, an expression for the extinction spectrum of a non-spherical particle
can be derived:
49
E !( ) = 24"Na3#out
3/2
! ln 10( )# i !( )
# r !( ) + $#out( )2 + # i !( )2%
&''
(
)**
(3.33)
Where N is the finite number of polarizable elements (dipoles) within the nanoparticle, εi and εr
are the imaginary and real components of the dielectric constant for the nanoparticle, and χ is a
shape factor which accounts for geometries other than a spherical particle13. This shape factor is
also highly dependent on the dielectric of the external environment (εout), and so also acts as a
sensitivity factor of the localized surface plasmon resonance (LSPR) extinction spectrum to the
dielectric environment. For a sphere, χ has a value of 2, however, for other aspect ratios, it can
range as high as 20. 13, 20
3.2.2 SERS Substrates
SERS active substrates can be fabricated a number of ways including electrochemically
roughened electrodes and template nanostructures.
Electrochemically roughened electrodes are frequently used throughout the literature as
SERS active substrates. They can be generated through one or more oxidative reductive cycles (a
technique similar to cyclic voltammetry). During the oxidative cycle, the metal combines with a
salt (usually a halide) at the metal surface. When the reductive half cycle is begun the salt
dissociates and the native metal re-deposits on the surface in a non-uniform fashion. Using this
technique, metal clusters are formed with an apparent maximum size of ~200Å9. Although these
surfaces are fairly easily prepared, a common problem is the non-homogeneity of the sample,
subsequently leading to low reproducibility and stability5.
Templated nanostructures, in contrast with electrochemically-roughened surfaces,
provide a reasonably homogenous, and therefore reproducible surface. To achieve such a
reproducible surface, electrochemical reduction of various metals inside nanopores is perhaps
50
one of the best options due to its ability to confine and restrict the geometry of nanostructure to
that of the nanopores. To achieve such a result, anodic aluminum oxide films are commonly
used21. The use of gold and silver nanorods has been found to provide an enhancement up to 109
in a variety of systems19, 22. Pore dimensions, and thus dimensions of the nanorods, can be
controlled by the applied voltage, pH and composition of the electrolyte5.
The presence of a second nanoparticle within 1nm of other particles allows for even
greater enhancement, in the range of 1011. This effect can be attributed to the formation of an
effectively huge capacitive field from the viewpoint of a molecule located at the interstice
between two nanoparticles. As the particles are brought together, the dipoles in each are oriented
in opposite directions. This interaction with the incident light results in a greatly enhanced
Raman signal. 16
References
1. Bard, A. J.; Faulkner, L. R. In Kinetics of Electrode Reactions; Harris, D., Swain, E. and
Aiello, E., Eds.; Electrochemical Methods: Fundamentals and Applications; John Wiley & Sons, Inc.: New York, NY, USA, 2001; Vol. 2nd Ed., pp 87-132.
2. Skoog, D. A.; Holler, F. J.; Crouch, S. R. In Electroanalytical Chemistry; Kiselica, S., Short, M. A., Eds.; Principles of Instrumental Analysis; Thomson Brooks/Cole: Belmont, CA, USA, 2007, 1998; Vol. 6th Ed., pp 627-653.
3. Baron, J.; Szymanski, G.; Lipkowski, J. J Electroanal Chem 2011, 662, 57-63. 4. Bard, A. J.; Faulkner, L. R. In Methods Involving Forced Convection - Hydrodynamic
Methods; Swain, E., Harris, D. and Aiello, E., Eds.; Electrochemical Methods: Fundamentals and Applications; John Wiley & Sons, Inc.: New York, NY, USA, 2001; Vol. 2nd Ed., pp 331-364.
5. Baron Gavidia, J. Study of the Gold-Thiosulfate Interface Under Leaching Conditions, University of Guelph, Guelph, ON, 2010.
6. Skoog, D. A.; Holler, F. J.; Crouch, S. R. In Raman Spectroscopy; Kiselica, S., Short, M. A., Eds.; Principles of Instrumental Analysis; Thomson Brooks/Cole: Belmont CA, USA, 2007,; Vol. 6th Ed., pp 481-495.
7. Banwell, C. N. In Fundamentals of Molecular Spectroscopy; McGraw Hill Book Company (UK) Limited: United Kingdom, Vol. 3rd Ed., pp 338.
8. Fleischmann, M.; Hendra, P.; McQuillan, A. Chemical Physics Letters 1974, 26, 163-166.
9. Moskovits, M. Reviews of Modern Physics 1985, 57, 783-826.
51
10. Albrecht, M. G.; Creighton, J. A. J. Am. Chem. Soc. 1977, 99, 5215-5217. 11. Jeanmaire, D. L.; Van Duyne, R. P. Journal of Electroanalytical Chemistry and
Interfacial Electrochemistry 1977, 84, 1-20. 12. Tian, Z. Q.; Ren, B.; Wu, D. Y. J Phys Chem B 2002, 106, 9463-9483. 13. Stiles, P. L.; Dieringer, J. A.; Shah, N. C.; Van Duyne, R. R. Annual Review of Analytical
Chemistry 2008, 1, 601-626. 14. Brolo, A. G.; Irish, D. E.; Smith, B. D. J. Mol. Struct. 1997, 405, 29-44. 15. Mock, J.; Norton, S.; Chen, S. Y.; Lazarides, A.; Smith, D. Plasmonics 2011, 6, 113-124. 16. Moskovits, M. J. Raman Spectrosc. 2005, 36, 485-496. 17. Liao, Q.; Mu, C.; Xu, D. S.; Ai, X. C.; Yao, J. N.; Zhang, J. P. Langmuir 2009, 25, 4708-
4714. 18. Xu, H.; Käll, M. ChemPhysChem 2003, 4, 1001-1005. 19. Bok, H. M.; Shuford, K. L.; Kim, S.; Kim, S. K.; Park, S. Langmuir 2009, 25, 5266-
5270. 20. Joseph, V.; Matschulat, A.; Polte, J.; Rolf, S.; Emmerling, F.; Kneipp, J. J. Raman
Spectrosc. 2011, . 21. Matefi-Tempfli, S.; Matefi-Tempfli, M.; Vlad, A.; Antohe, V.; Piraux, L. Journal of
Materials Science-Materials in Electronics 2009, 20, 249-254. 22. Orendorff, C. J.; Gearheart, L.; Jana, N. R.; Murphy, C. J. Physical Chemistry Chemical
Physics 2006, 8, 165-170.
52
CHAPTER 4: METHODOLOGY
4.1 REAGENTS
All solutions were prepared using Milli-Q water with a resistivity of 18.2 MΩ cm with a
thiosulfate concentration of 100mM. Calcium thiosulfate solutions were prepared from calcium
thiosulfate (20-30% by weight) aqueous Captor® solutions from Tessenderlo Kerley Ltd. The
pH of the calcium and sodium thiosulfate solutions was adjusted to 8-8.5 using calcium
hydroxide (95%) from Sigma-Aldrich. Ammonium thiosulfate solutions were prepared from
ammonium thiosulfate (99%) from Alfa Aesar, and were pH adjusted using FisherBrand® ACS-
Pur ammonium hydroxide (28-30% NH3 w/w). Sodium thiosulfate pentahydrate (99.5%) salt
was acquired from Acros Organics. Copper was added to all leaching solutions as copper sulfate
pentahydrate (98+%) from Sigma-Aldrich to a concentration of 10 mM.
4.2 CLEANING METHODS
All glassware pieces were cleaned by soaking in a hot acid bath with concentrated H2SO4
and HNO3 in a ratio of 3:1 v/v for 60 min. After soaking in the acid bath the pieces were
thoroughly rinsed with Milli-Q water.
The gold disk rotating disk electrode was polished using 1 µm LECO® Premium
Diamond Suspension, and rinsed by rotation in Milli-Q water and methanol. Prior to use in
leaching current experiments the RDE was subjected to reductive desorption in 0.3 M NaOH
(semi-conductor grade 99.99% trace metals basis from Sigma-Aldrich). The solid polycrystalline
gold electrode was cleaned by flame annealing, followed by rinsing with Milli-Q water and
drying over a flame. Prior to each experiment, the counter electrode was flame annealed, rinsed
53
with Milli-Q water and dried over a flame.
Cleanliness of both the RDE and solid polycrystalline electrodes was checked by
cyclic voltammetry in a solution of 0.1 M NaF (≥99% from Sigma-Aldrich). The system was
considered clean when a cyclic voltammogram (CV) characteristic of a bare polycrystalline gold
electrode was achieved. Figure 4.1 displays a representative CV for either of the two gold
electrodes used. If the shape of the voltammogram matched those of literature1, 2, the electrodes
were considered clean.
Figure 4.1: Cyclic voltammogram of a clean, bare gold RDE in a solution of 0.1 M NaF. The inset graph
displays the double layer region of the gold electrode.
-1.0 -0.5 0.0 0.5 1.0 1.5
-20
-15
-10
-5
0
5
10
15
20
25
-1.0 -0.8 -0.6 -0.4 -0.2 0.0 0.2 0.4 0.6
-2.0
-1.5
-1.0
-0.5
0.0
0.5
1.0
1.5
i / µΑ
cm
-2
Potential / V vs. SCE
54
4.3 EXPERIMENTAL SETUP AND PROCESS
4.3.1 Preparation of Gold Electrodes
The polycrystalline gold electrode used for electrochemical experiments was made in-house
using the procedure outlined by Richer3.
4.3.2 Electrochemical Experiments
Electrochemical experiments were carried out using a HEKA (PG 590, Lambrecht/Pfalz,
Germany) potentiostat/galvanostat connected to an acquisition board from National Instruments
(PCI 6052E). Custom written software written by Professor Dan Bizzotto from the University of
British Columbia, and Professor Ian Burgess from the University of Saskatchewan, was used for
data acquisition.
4.3.2.1 Leaching Current Measurements
A typical three-electrode glass cell was used for the electrochemistry experiments. The
reference electrode (RE) was a saturated calomel electrode (SCE) placed in a separate cell in
order to avoid cross contamination between the investigated solution and the RE. The two cells
were connected with a salt bridge. Leaching current measurements using the RDE were achieved
with a gold coil as the counter electrode (CE), placed in a separate compartment in contact with
the solution through a glass frit. In the case of measurement using the solid polycrystalline
electrode (hanging meniscus configuration), the gold coil CE was submerged into the leaching
solution. Both the solid polycrystalline electrode and the RDE setups were left open to the
atmosphere, without purging of the solution. Figure 4.2 displays schematics of the two cells used
for the RDE and solid gold polycrystalline electrode.
55
Figure 4.2: Schematic for the two cell configurations used.
Linear sweep voltammetry was employed in order to probe the system and calculate the
current as described by J.Y. Baron4.
4.3.2.2 Open Circuit Potential Measurements
Measurement of the open circuit potentials for all three solutions (calcium, sodium and
ammonium thiosulfate with Cu(II)) were performed using the same three electrode setup as in
the leaching current measurements with an RDE (Figure 4.1). The rotating disk electrode was
cleaned according to the procedure described previously. Once cleanliness of the system was
ensured, the RDE was submerged into the leaching solution and the open circuit potential was
recorded for 3 hours using a custom potential-current (E-i) monitor program.
Salt bridge to
SCE
Au working
electrode Au
counter electrode
Au RDE (connected to motor)
Au counter
electrode Salt bridge to
SCE
56
4.3.2.3 Solution pH Measurements
The pH of the leaching solutions was monitored for 3 hours in order to understand possible
reaction pathways at the interface during the leaching reaction. Measurement of the solution
potential was accomplished using an Orion Expandable ionAnalyzer EA 920 pH meter with an
Accumet pH electrode. The electrode was immersed into a 150 mL beaker of the leaching
solution of interest, for 3 hours. Using the E-i monitor program, the potential output from the pH
electrode was recorded. The pH of the solution was calculated through the measurement of the
potential and pH of a set of standard buffer solutions, plotting each on a graph of potential vs. pH
(Figure 4.3). A linear regression was performed, generating a line of best fit with an R2 value of
0.99585. Using the equation of the line, pH = E !0.01836( ) + 6.90614 , the pH was calculated
and plotted as a function of time.
Figure 4.3: Calibration curve for the calculation of pH from the potential measured.
-200 -150 -100 -50 0 50 100 150 2003
4
5
6
7
8
9
10
11
pH
Potential / mV
Equation y = a + b*x
Weight No Weighting
Residual Sum of Squares
0.07405
Pearson's r -0.99862
Adj. R-Square 0.99585Value Standard Error
pHIntercept 6.90614 0.10022Slope -0.01836 6.83401E-4
57
4.3.3 Preparation of Gold Nanorod Electrodes
Gold nanorod electrodes used for SERS leaching experiments were prepared using 13mm
diameter anodized aluminum oxide filters with a 0.1 µm pore size, and a thickness of 60 µm as
templates for the electro-deposition of gold. Nanorods were deposited using two different
template treatments.
In the first method of growth, vapor deposition was employed to deposit a gold film, 70
nm thick, onto the back of the templates for electrical contact to a gold slide. The gold coated
filter was then pressed to the gold slide using a Kel-F® cell, where the uncoated side of the
template faced the open portion of the cell. This assembly was then attached to a potentiostat via
a gold wire and placed in a conventional 3-electrode glass cell, using gold foil as a CE, and a
SCE as the reference.
Uncoated templates were used in the second growth method in order to attach the
nanorods directly attached to the slide. A similar cell design as that for coated filters was used,
with the primary difference being a smaller diameter aperture in the top of the cell, and a greater
distribution of increased pressure on the template and gold slide for minimization of the space
where roughened gold would deposit.
Electrodeposition was carried out by submerging the template in TECHNIC gold solution
(TG-25 RTU), which was de-aerated by passing argon through the solution for 60 minutes. For
growth using coated filters a constant potential of -0.900 V was applied for 3-8 hours. Uncoated
filters required application of a lower potential (-0.800 V) for a significantly longer period of
time (15-20 hours). Once deposition was complete, the electrode was rinsed with Milli-Q water,
and the aluminum template dissolved by submerging the electrode in a 3 M NaOH (99.99% from
Sigma Aldrich) solution, for 3 hours.
58
Nanorod electrodes were cleaned using two different methods, depending on the growth
assembly used. Electrodes grown using coated templates were cleaned by gentle rinsing with
chloroform and a 30 min incubation in the UV/O3 chamber. Uncoated templates produced
electrodes that were mechanically attached to the underlying gold slide, greatly improving the
durability of the gold nanorods. Electrodes grown in such a fashion were cleaned by submerging
the whole cell into piranha solution for 15-30 minutes, followed by thorough rinsing with Milli-
Q water. The nanorod electrodes were then left to soak in Milli-Q water for 60 minutes before
use in experiments
4.3.4 Raman Experiments
A Renishaw Raman Imaging Microscope was used for Raman experiments. A NIR diode
laser with a wavelength of 785 nm and an output power of 300 mW was used for excitation. The
Raman instrument was equipped with a CCD array detector. The spectrometer was calibrated
using the Raman active vibration of silicon at 520 cm-1. A 63x immersion objective from Leica
was used for solution Raman and SERS experiments.
4.3.4.1 SERS
SERS spectra for all systems were collected at 10% power to avoid laser-induced
decomposition of species on the gold surface. For the calcium thiosulfate system the exposure
time was 5 s with 10 accumulations. For time dependent experiments, the static mode was used
to minimize the time required for data acquisition. In order to cover the whole spectral range of
interest, spectra centered at 400 cm-1 and 900 cm-1 were taken for each experiment and then
combined into one spectrum.
Spectra corresponding to the ammonium and sodium thiosulfate systems were collected
using the extended mode with a range of 100-2000 cm-1 and an exposure time of 15 s and 10
59
accumulations. Spectra were collected over a period of 3 hours after exposure of the nanorod
electrode to the leaching solution of interest.
4.3.4.2 Solution Raman
Solution spectra were collected using 100% power in extended mode, with an exposure
time of 15 s and 10 accumulations. Spectra were recorded for 3 hours after the time of solution
preparation.
References
1. Clavilier, J.; Van Huong, C. N. Journal of Electroanalytical Chemistry and Interfacial
Electrochemistry 1977, 80, 101-114. 2. Stolberg, L.; Richer, J.; Lipkowski, J.; Irish, D. Journal of electroanalytical chemistry
and interfacial electrochemistry 1986, 207, 213-234. 3. Richer, J. Measurement of Physical Adsorption of Neutral Organic Species at Solid
Electrodes, University of Guelph, Guelph, ON, Canada, 1985. 4. Baron, J.; Szymanski, G.; Lipkowski, J. J Electroanal Chem 2011, 662, 1, 57-63.
60
CHAPTER 5: RESULTS AND DISCUSSION
5.1 PREAMBLE
Gold leaching in industrial settings is carried out at the open circuit potential of the
system. At this potential, passivation of the gold surface can occur as a result of a number of
complicated reactive pathways that lead to the formation of oxidation and decomposition
products of thiosulfate, such as polythionates and elemental sulfur1-5. Addition of copper and
ammonia to the thiosulfate leaching system has been shown to have significant effects on the
gold leaching rate1-3, 6, 7. However, there is limited literature concerning experimental studies of
the thiosulfate leaching system in the presence of these additives individually. The work
described in this chapter was designed to provide a base understanding of the individual, and
cooperative effects of copper and ammonia on the passive layer.
Characterization of the systems of interest (calcium, sodium and ammonium thiosulfate)
was carried out using both electrochemical and spectroscopic analysis. Using the method devised
by Baron et al.8 the measured leaching currents of a non-passivated surface can be measured with
minimal perturbation to the system. This is an important characteristic of this methodology, as
application of potentials more positive than the mixed potential of the system can lead to electro-
oxidation of thiosulfate5.
Raman spectroscopy is particularly useful in the study of oxygenated inorganic sulfur
species due to their extensive characterization, and distinct intense vibrational modes that allow
for fingerprint identification of a multitude of species in a given sample9-11. This thesis employed
both solution Raman and Surface Enhanced Raman Spectroscopy (SERS) to study the bulk
61
solution and gold-thiosulfate interface in the presence of copper and ammonia to identify species
present at the gold surface.
Band shifting relative to the normal Raman spectrum (aqueous or solid) is expected upon
adsorption of a molecule to a SERS active surface12. The SERS band intensity arises from an
enhancement that is controlled by a number of factors, including: i) electric field localization at
the SERS active surface and ii) the orientation of the polarizability tensor as a result of the
geometry of adsorption. As a general rule, surface vibrations that affect the change in the
polarizability in a direction perpendicular to the surface will undergo the greatest enhancement.
Adsorbates with π electrons are also susceptible to shifts in the vibrational frequency as a result
of a change in the potential of the surface. Thus, the SERS spectra are extremely sensitive to the
open circuit potential of the system12.
Due to the complex composition of the passive layer, a SERS active substrate with very
high enhancement is required to identify various oxygen containing inorganic sulfur species
within the passive layer. The destructive nature of the leaching reaction makes a substrate with
long-term stability a desirable and necessary feature for long term studies. Gold nanorod arrays
have been shown to provide significant enhancement, and long-term stability when used as a
SERS active substrate to study the gold-thiosulfate interface under leaching conditions13, 14.
5.2 CHARACTERIZATION OF BULK SOLUTION
To fully interpret and understand the results acquired in characterizing the passive layer
formed at the gold-thiosulfate interface under leaching conditions, a thorough understanding of
the composition and behavior of the bulk solution must be achieved. Hence, both the behavior
and composition of the solution must be characterized.
62
5.2.1 pH Measurements
All leaching experiments were performed in solutions in equilibrium with ambient air,
and as such are sensitive to the effect of dissolved CO2. The pH of the leaching solution can have
a significant effect on decomposition pathways, and species present. Therefore, it was important
to determine the pH of all three systems over a period of 3 hours by recording the output from a
pH meter. Standard buffer solutions were used to calibrate the readout from the pH meter. The
results after calibration are plotted in Figure 5.1.
It was expected that the pH of the ammonium thiosulfate system would undergo the most
significant change due to possible evaporation of ammonia from the solution. However, the
buffering ability of NH4+ /NH3 can explain the fairly constant pH level in this solution. More
surprising are the behaviors of the calcium and sodium thiosulfate systems. Each of these
solutions displays an initial increase of approximately 0.4 pH units over a period of
approximately 30 min, followed by a drop in pH. After 3 hours the pH of the sodium thiosulfate
solution returns to its initial value (~ pH 8.2), a drop of approximately 0.6 pH units from the
maximum pH (~8.8). A drop of this magnitude is similar to that of the blank system (pure water),
where dissolved carbon dioxide is responsible for the change in pH. The behaviour of the
calcium thiosulfate system is far more drastic. The pH decreases by almost two pH units from its
initial value. A decrease of that magnitude could not solely be caused by dissolved carbon
dioxide.
63
Figure 5.1: Average pH over a period of 3 hours for calcium, ammonium and sodium thiosulfate systems
(0.1 M S2O3 + 0.01 M CuSO4 and adjusted to a pH of 8.0-8.5) and a blank solution of pure MilliQ water.
However, formation of a precipitate was observed in the calcium thiosulfate solution near
the end of the experiment. This could be explained by the precipitation of either CaCO3 or
CaSO4, leading to the observed drop in pH in the calcium thiosulfate solution. Both of these
species may be precipitating from solution; however, based on the solubility product constants of
these two salts (4.96 x 10-9 and 7.10 x 10-5 for calcium carbonate and calcium sulfate,
respectively), and evidence from industrial settings, it is more likely that calcium carbonate is
precipitating.15
0 25 50 75 100 125 150 175 2005.0
5.5
6.0
6.5
7.0
7.5
8.0
8.5
9.0
A
vera
ge p
H o
f Lea
chin
g So
lutio
n
Time / minutes
STS
ATS
CaTSMilliQ Water Blank
64
Significant changes in the pH of the leaching solutions were observed, depending on the
cation of the thiosulfate salt. These changes can greatly affect the composition, and subsequent
efficacy of the leaching process.
5.2.2 Solution Raman
Raman spectra of leaching solutions were collected for a period of 3 hours to correlate
any change in bulk solution species to those observed at the interface in SERS spectra. Many of
the decomposition products of these leaching solutions, such as tetrathionate, trithionate and
sulfite, have well characterized aqueous solution Raman spectra reported in the literature.
Characteristic vibrations for thiosulfate, tetrathionate, trithionate, sulfate and sulfite can be found
in Table 5.1. These vibrations can be used for fingerprint identification in the more complex
spectra of the leaching solutions. Figure 5.2 (a-c) displays the average spectra of 3 separate
Raman experiments for each of the calcium, ammonium and sodium thiosulfate systems,
respectively.
From these spectra, it is evident that there are two main regions of interest, 300 cm-1 to
700 cm-1, and 900 cm-1 to 1200 cm-1. Bands located at ~ 190 and 1640 cm-1 are attributed to
vibrational modes of water, and will not be further discussed in the scope of this chapter. Using
the Fourier Self-Deconvolution (FSD) method described by J. Baron13, 16, each of the regions
were analyzed by deconvoluting the broad bands into individual peaks that could be assigned to
a specific vibration. Fitting of peaks was performed using a mixed Gaussian and Lorentzian band
shape. It was found that the calcium, ammonium and sodium thiosulfate solutions displayed the
same bands, and upon peak fitting, the same component peaks. The deconvoluted spectrum of
the calcium thiosulfate system in the 400 cm-1 region, recorded 20 min after solution preparation,
is presented in Figure 5.3. The upper connected points correspond to the raw experimental
65
envelope, the solid line overlapping the raw data points is the simulated envelope, and the lowest
solid coloured lines represent the deconvoluted peaks.
Table 5.1: Characteristic Raman active vibrational modes of species expected in the Raman spectra of the
investigated leaching solutions10, 17-19.
Species Assignment Raman Shift / cm-1
Thiosulfate (S2O32-)
ρr (SSO) 334
νsym
(SS) 443
δasym
(SO3) 533
δsym
(SO3) 663
νsym
(SO3) 995
νasym
(SO3) 1122
Tetrathionate (S4O62-)
δ (SSinternal
) 260 δ (SS
terminal) 310
δ (SSterminal
) 390 δ
asym (SO
3) 532
δ sym
(SO3) 651
νsym
(SO3) 1040
Trithionate (S3O62-)
δ (SS) 264
δ (SS) 425
δsym
(SO3) 675
νsym
(SO3) 1055
Sulfate (SO42-)
δsym
(OSO) 448
δasym
(OSO) 620
νsym
(SO3) 982
νasym
(SO) 1110
Sulfite (SO32-)
δasym
(OSO) 470
δsym
(OSO) 620
νasym
(SO) 933 ν
sym (SO) 967
66
Figure 5.2: Average raw Raman spectra collected over a period of 3 hours for leaching solutions of: a)
0.1 M CaS2O3 + 0.01 M CuSO4 b) 0.1 M (NH4)2S2O3 + 0.01 M CuSO4 c) 0.1 M Na2S2O3 + 0.01 M
CuSO4. All solutions were adjusted to pH 8.0-8.5.
500 1000 1500 2000
Raman Shift / cm-1
10 min20 min33 min46 min57 min69 min
80 min110 min140 min170 min200 min
500 1000 1500 2000
Raman Shift / cm-1
10 min30 min
42 min
52 min
62 min72 min
92 min112 min
142 min
180 min
500 1000 1500 2000
Raman Shift / cm-1
5 min17 min
28 min39 min53 min
62 min
72 min102 min
142 min
182 min
a) b)
c)
67
Figure 5.3: Deconvoluted spectra of an average of three 0.1 M CaS2O3 + 0.01 M CuSO4 solutions,
adjusted to a pH of 8.0-8.5, 20 min after solution preparation.
Five peaks were identified, at positions of 387, 425, 446, 520, and 666 cm-1. The first
band at 387 cm-1 can be attributed to the presence of tetrathionate in solution10, 13, 17. Strong
bands at 425 and 446 cm-1 correspond to the stretches of trithionate and thiosulfate, respectively.
The assignment of bands at 520 and 666 cm-1 is uncertain, because they could correspond to
multiple oxygenated inorganic sulfur species. The band at 520 cm-1 is very near an asymmetric
stretch of thiosulfate in aqueous solution, 533 cm-1, and an asymmetric stretch of tetrathionate
(532 cm-1).10
300 400 500 600 700 800δ sy
m(O
SO) S
2O2- 3
ν sym
(SS)
S2O
2- 3
δ(SS
) S3O
2- 6
δ asym
(SO
) S2O
2- 3 / δ as
ym(O
SO) S
4O2- 6
Raman Shift / cm-1
δ(SS
term
inal) S
4O2- 6
68
Higher wavenumber sections of the spectra can also contain bands useful for
identification, and thus, cannot be ignored. Figure 5.4 shows the deconvolution of the 900-
1200 cm-1 region in the spectrum recorded 20 minutes after solution preparation.
Figure 5.4: Deconvoluted spectra of an average of three 0.1 M CaS2O3 + 0.01 M CuSO4 solutions,
adjusted to a pH of 8.0-8.5, 20 min after solution preparation, in the 900-1000 cm-1 region.
Four distinct peaks positioned at 980, 997, 1037, and 1130 cm-1 are easily identifiable in
this region of the spectrum. The symmetric stretch of aqueous sulfate matches the band observed
at 980 cm-1 in literature 18. Thiosulfate has a strong peak at 995 cm-1, and a weaker broad peak
near 1122 cm-1. Although there is a slight shift from these literature values, the peaks at 997 and
1130 cm-1 in Figure 5.4 can be assigned to thiosulfate13. The final peak at 1037 cm-1 can be
assigned to a symmetric stretch of tetrathionate10.
900 1000 1100 1200
ν sym(S
O) S
O2- 4
ν asym(S
O) S
2O2- 3
ν sym(S
O) S
2O2- 3
ν sym(S
O) S
4O2- 6
Raman Shift / cm-1
69
FSD was performed on each of the spectra acquired and the deconvoluted bands’
positions were matched with characteristic vibrations of species present in solution. Analytical
areas for each of the species of interest were tracked for the duration of the experiment in each of
the calcium, ammonium and sodium thiosulfate systems (Figure 5.5). Peak areas were
normalized with reference to the water band seen at 1640 cm-1 in each of the spectra in Figure
5.2.
Figure 5.5: Normalized peak areas for band positions of 1037 cm-1 (tetrathionate = black), 425 cm-1
(trithionate = blue), 980 cm-1 (sulfate = green), 448 cm-1 (thiosulfate = red). a) 0.1 M CaS2O3 + 0.01 M
CuSO4, b) 0.1 M Na2S2O3 + 0.01 M CuSO4, c) 0.1 M (NH4)2S2O3 + 0.01 M CuSO4. Solutions were
adjusted to pH 8.0-8.5.
0 50 100 150 200
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
νsym(SO) S4O2-6
νsym(SO) SO2-4
Nor
mal
ized
Ana
lytic
al P
eak
Are
a
Time / minutes
νsym(SS) S2O2-3
δ(SS) S3O2-6
0 50 100 150 2000.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
νsym(SO) SO2-4
νsym(SO) S4O2-6
Nor
mal
ized
Ana
lytic
al P
eak
Are
a
Time / minutes
νsym(SS) S2O2-3
δ(SS) S3O2-6
0 50 100 150 2000.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
νsym(SO) S4O2-6
νsym(SO) SO2-4
Nor
mal
ized
Ana
lytic
al P
eak
Are
a
Time / minutes
νsym(SS) S2O2-3
δ(SS) S3O2-6
a) b)
c)
70
It is evident from the data above that the composition of all three bulk solutions (CaTS,
ATS, and STS) remains fairly constant, with most species maintaining a steady state. Some
variation is observed in the peak area of trithionate in the ammonium thiosulfate system, leading
to a very slight decrease over the final 80 minutes of the experiment. However, the solutions’
compositions maintain a steady state concentration of thiosulfate, tetrathionate, and sulfate.
An important feature of these spectra is the appearance of trithionate early in the
experiment. In previous studies, trithonate was not seen until much later in the experiment13, 20.
Appearance of trithionate within the first 20 minutes after solution preparation is indicative of
catalysis of the decomposition or disproportionation of tetrathionate, likely by Cu2+ in the
solution:
2S4O62! " S3O6
2! + S5O62! (5.1)
4S4O62! + 3OH! " 5S2O3
2! + 2S3O62! + 3H2O (5.2)
In the step-wise reaction sequence for the decomposition of tetrathionate, thiosulfate is
both a product and a catalyst for the reaction. This is one possible reason why the thiosulfate
solutions containing copper maintained a constant thiosulfate concentration. Thiosulfate
solutions without copper can suffer a reduction in thiosulfate concentration over extended
periods of time, depending on the pH of the solution.
71
5.3 CHARACTERIZATION OF THE GOLD-THIOSULFATE INTERFACE IN THE
PRESENCE OF COPPER
5.3.1 Initial Characterization
Using linear sweep voltammetry, kinetic characterization was performed using a gold
electrode in contact with a solution of 0.1 M Na2S2O3 + 0.01 M CuSO4, adjusted to pH 8.0-8.5
using Ca(OH)2. Prior to leaching current measurement, it was necessary to determine the values
for the charge transfer coefficients. A linear sweep voltammogram was collected using a sweep
rate of 1 mVs-1 in a potential range of -0.32 V to 0.2 V, as shown in Figure 5.6.
Figure 5.6: Linear sweep voltammogram of a 0.1 M Na2S2O3 + 0.01 M CuSO4 solution, pH 8.0-8.5.
According to mixed potential theory, the point at which the i-V curve crosses zero
corresponds to the mixed potential (or open circuit potential) of the system. That is, the
magnitude of the current contributed by the cathodic reduction of Cu2+, and that from the
-0.4 -0.3 -0.2 -0.1 0.0 0.1 0.2
-3
-2
-1
0
1
2
3
4
5
Cur
rent
/ µA
Potential / V vs. SCE
72
oxidation of Au, are equal but opposite in sign. Thus, the recorded net current is zero. By
plotting the logarithm of the measured current as a function of applied potential, a Tafel-like plot
can be obtained, as shown in Figure 5.7. The values of the transfer coefficients can be
determined by performing a linear regression analysis of the linear sections of the curve, at large
overpotentials (Figures 5.8 and 5.9).
Figure 5.7: Tafel plot of a 0.1 M Na2S2O3 + 0.01 M CuSO4 solution, pH 8.0-8.5, used for calculation of
the transfer coefficients of the system.
The slopes of the cathodic and anodic linear regions are given by Equations 5.3 and 5.4,
respectively:
slope = !" cF2.3RT
(5.3)
slope =1!" a( )F2.3RT
(5.4)
-0.4 -0.3 -0.2 -0.1 0.0 0.1 0.2
-3.0
-2.5
-2.0
-1.5
-1.0
-0.5
0.0
0.5
1.0AnodicLinear
Log(i)
Potential / V vs. SCE
CathodicLinear
73
Using the slope calculated through the linear regressions in Figures 5.3 and 5.4, the transfer
coefficients for the reduction of Cu2+ and oxidation of Au were found to be 0.35 ± 0.02 and 0.78
± 0.05, respectively.
Figure 5.8: Linear regression of the cathodic reduction of Cu2+ for the determination of the transfer
coefficient. A slope of -6.10 was calculated.
Quantification of the transfer coefficients is necessary for the calculation of leaching
currents in the systems of interest. The Butler-Volmer equation discussed in Chapter 3 can
provide an accurate description of the current-overpotential curve for a simple electron transfer
reaction:
i = i0 exp!" cnFRT
#$%&
'() ! exp
1!" a( )nFRT
#$%&
'()
*
+,
-
./ (5.5)
-0.32 -0.30 -0.28 -0.26 -0.24 -0.22 -0.20-0.3
-0.2
-0.1
0.0
0.1
0.2
0.3
0.4
0.5
Log
(i)
Potential / V vs. SCE
74
A modified version is required here since the values of the transfer coefficients are not the same
for the cathodic and anodic reactions:
i = iM exp!"
Cu2+nCu2+F
RT#
$%&
'()! exp
1!"Au( )nAuFRT
#$%&
'()
*
+,
-
./ (5.6)
Figure 5.9: Linear regression of the anodic oxidation of Au for the determination of the transfer
coefficient. A slope of 4.75 was calculated.
The term i0 in Equation 5.5 is equivalent to the leaching current, iM, in the case of the
gold leaching reaction, as in Equation 5.6. The terms nCu2+ and nAu in Equation 5.6 both refer to
one electron transfer redox reactions, the rengeneration of Cu2+ and oxidation of gold,
thus nCu2+ = nAu =1 . Because the potential region of interest is small, Equation 5.6 can be
simplified to:
-0.08 -0.06 -0.04 -0.02 0.00
-0.2
0.0
Log
(i)
Potential / V vs. SCE
75
i = iM !"Cu2+
! 1!" Au( )( ) nFRT # (5.7)
A plot of the current density response in a ~ 40 mV region around the mixed potential
should have linear characteristics (Figure 5.10). Application of a linear regression to the data
provided a slope with which to calculate the value of the leaching current. For the data shown in
Figure 5.10, a leaching current density of 2.14 µA cm-2 was obtained for a temperature of 293 K,
assuming a one-electron transfer reaction.
Figure 5.10: Linear regression of data acquired during a linear sweep voltammogram of a 0.1 M Na2S2O3
+ 0.01 M CuSO4 solution, pH 8.0-8.5 solution with a sweep rate of 1 mVs-1.
This method of calculating leaching current densities was used throughout this thesis for
analysis of all systems of interest. The current efficiency of thiosulfate leaching media has been
shown to be close to 100% up to 0.08 V vs SCE20, thus the calculation of the current density in a
-0.16 -0.14 -0.12
-1
0
1
j / µ
Αcm
-2
Potential / V vs. SCE
76
small potential range around the mixed potential can be directly related to the gold leaching
current density.
To gain an understanding of the kinetic behavior of the systems of interest, it is necessary
to observe under which sweep rates the system is under mass transport or kinetic control. The
thickness of the diffuse layer is inversely proportional to the square root of the sweep rate. At
low sweep rates, the diffuse layer extends further from the electrode, greatly decreasing the flux
to the electrode surface, and thus producing a relatively low current. Higher sweep rates reduce
the thickness of the diffuse layer, leading to increased current. Current produced by a system
under mass transport control (such as at low sweep rates) are proportional to the square root of
the sweep rate. Deviations from such a relationship indicate that the system is moving under
mixed control, that is, contributions to the current arise from both diffusion to the electrode
surface and kinetic limitations of the electron transfer reaction. To determine if the systems of
interest were susceptible to kinetic control, the leaching currents at 6 sweep rates (1, 2, 10, 20, 50
and 100 mVs-1) were measured and plotted as a function of the square root of the sweep rate
(Figure 5.11).
The linear behavior at low sweep rates in Figure 5.11 indicates that the system is under
mass transport control under these conditions. As the sweep rate increases to 20 mVs-1, deviation
from the line is observed. Hence, at this sweep rate, and higher, the system moves to mixed
control. Since this probing of the system did not involve forced convection of the solution, it
should be possible through use of an RDE to bring the system under full kinetic control.
Measurement of the leaching current while under full kinetic control is required in order to have
an accurate understanding of the gold leaching reaction rate, without the contribution of diffusion
effects.
77
Figure 5.11: Sweep rate dependence of the leaching current measured in a 0.1 M Na2S2O3 + 0.01 M
CuSO4 solution, pH 8.0-8.5, at sweep rates of 1, 2, 10, 20, 50 and 100 mVs-1.
Therefore, leaching current densities were also measured as a function of the angular
velocity of a gold disk RDE, using rotation rates of 300, 500, 700 and 1000 RPM. Due to rapid
passivation of the gold surface as a result of the forced convection in the solution, the potential
range of the linear sweep was decreased to a 120 mV region (-220 mV to -100 mV). The
dependence of the leaching current density was plotted as a function of the square root of the
angular velocity (calculated from the rotation rate), as shown in Figure 5.12.
Within experimental error, the data displayed below are independent of the angular
velocity (and thus rotation rate). According to the Koutecky-Levich equation discussed in
Chapter 3, at each sweep rate a maximum, or limiting current, exists.
0 1 2 3 4 5 6 7 8 9 10 110
2
4
6
8
10
j m / µA
cm-2
ν1/2
/ mV1/2 s-1/2
78
Figure 5.12: Rotation dependence of the leaching current of a gold electrode in contact with a 0.1 M
CaS2O3 + 0.01 M CuSO4 solution, pH 8.0-8.5, at rotation rates of 300, 500, 700 and 1000 RPM. Data was
recorded using a sweep rate of 5 mV s-1.
This limiting current is given by the term iK in Equation 5.8, and describes the current in
the absence of mass transfer effects:
1i= 1iK
+ 1il ,c
= 1iK
+ 10.62nFADO
2/3! 1/2" #1/6CO* (5.8)
When iK is small, such that:
1iK
>> 10.62nFADO
2/3! 1/2" #1/6CO* (5.9)
6 8 100.0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
4.0
4.5
5.0
5.5
6.0
j M / µA
cm-2
ω1/2
79
and mass transfer is efficient enough to keep the surface concentration of species equal to that of
the bulk, a plot of i vs. ω1/2 will show that i is independent of the rotation rate, as seen in Figure
5.12 with respect to the leaching current densities. Thus, it can be concluded that with rotation
rates as low as 300 RPM, kinetic control of the system can be achieved. Since measurement of
the leaching current in the solutions of interest needs to be carried out in a regime where no mass
transfer effects are present, a rotation rate of 300 RPM and sweep rate of 5 mVs-1 were chosen
for characterization of the systems of interest.
5.3.2 Leaching Current Measurements
Leaching current densities of all three systems of interest (calcium, sodium and
ammonium thiosulfate) were calculated using a set solution of 0.1 M S2O32- + 0.01 M CuSO4, pH
adjusted to 8.0-8.5. Measurement of the current response of a gold disk RDE was carried out
using the 3-electrode cell described previously. The rotation rate was set at 300 RPM, and the
potential was scanned between -220 mV and -100 mV using a sweep rate of 5 mVs-1. The
average calculated leaching current density for calcium thiosulfate (CaTS), sodium thiosulfate
(STS) and ammonium thiosulfate (ATS) are shown in Figure 5.13.
The data show that there is no variation, within experimental error, in the leaching current
density measured in all 3 systems. These results show a different effect of the cation on the
leaching reaction to those presented by Chandra et al.21, who noted an increase in the gold
oxidation polarization curves upon changing the alkali metal cation from sodium to potassium.
Based on these results, it was expected that upon comparison of the leaching current in the
sodium and calcium thiosulfate systems, the leaching current would be greater in the calcium
system.
80
Figure 5.13: Average calculated leaching current density for solutions of sodium, calcium and
ammonium thiosulfate solutions (0.1 M S2O3 + 0.01 M CuSO4 and were adjusted to a pH of 8.0-8.5).
The disagreement with literature can be attributed to the introduction of copper into the
systems studied in this work. In both the calcium thiosulfate and sodium thiosulfate systems,
there is a lack of ammonia in the leaching solutions. Ammonia is normally required to stabilize
the Cu2+ ion as the copper tetraamine complex, which is the oxidant in the leaching system1-3, 22,
23. In an ammonia free solution, Zhang and Nicol6 proposed that Cu2+ may form a reactive
intermediate with thiosulfate and oxygen that can react with free thiosulfate, producing the
Cu(S2O3)35! complex and tetrathionate through Equation 5.10:
S2O3( )3 Cu !O2"# $%5&
+ 4S2O32& + 2H2O' 2S4O6
2& + Cu(S2O3)35& + 4OH& (5.10)
STS CaTS ATS0.0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
4.0
4.5
5.0
5.5
j M / µA
cm-2
Thiosulfate Salt
81
The formation of complexes of copper and thiosulfate is well documented in literature. The
Cu(S2O3)35! or Cu(S2O3)2
3! complexes are the most thermodynamically stable copper complexes
in the concentration, pH and potential ranges employed in this work2, 3, 24-26. The presence of this
copper-thiosulfate redox couple significantly changes the properties of the solutions and explains
the differences between the work of Chandra et al.21, and the results presented in this chapter.
5.3.3 Open Circuit Potential Measurements
The open circuit potential (OCP) of a stationary gold disk RDE submerged in the
solutions of interest was tracked for an immersion period of 3 hours. The resulting curves as a
function of time can be seen in Figure 5.9. At t = 0, the OCP of the system is 28.6 mV, 18.9 mV,
and -3.6 mV for STS, CaTS and ATS, respectively. The high positive values for these systems
may be the result of adsorption of thiosulfate to the clean electrode surface. Each system then
decays rapidly to more negative potentials, which may indicate that the gold leaching reaction is
favored during this time. In the calcium and ammonium thiosulfate leaching solutions, the
potential curve reaches a negative limit, and then rises in an asymptotic fashion. The OCP of the
sodium thiosulfate system, while it behaves similarly to the calcium and ammonium thiosulfate
systems in the first 15 minutes of immersion, differs in the rate at which the potential rises
toward positive potentials. Rather than rising asymptotically after reaching a minimum, the
potential curve rises in a quasi-linear fashion.
The relative magnitudes for the negative limits for the OCPs match the results of the
oxidation polarization curve results of Chandra et al.21. That is, a progressively more negative
limit is reached as the cation is changed from sodium, to calcium, to ammonium.
82
Figure 5.14: Average open circuit potential of a gold disk RDE as a function of immersion time in
solutions of calcium, sodium and ammonium thiosulfate (0.1 M S2O3 + 0.01 M CuSO4 and adjusted to a
pH of 8.0-8.5).
In order to replicate the conditions of the SERS leaching tests, measurement of the OCP
for all systems was performed while holding the disk electrode stationary. Under these
conditions, diffusion from the bulk solution to the electrode surface may directly affect the
potentials measured. The differences between the limiting values of the OCP may be a result of
ion pairing with thiosulfate in solution, as proposed by Chandra et al.21, or by the formation of
adsorbed mixed complexes at the gold-thiosulfate interface. Senanayake27 has shown that the
oxidation of gold is highly dependent on the concentration product [M2S2O3][NH3] (where M =
Na or NH4+), and that the concentration of the NH4S2O3
- ion pair was greater than or equal to the
concentration of free thiosulfate ions.
0 25 50 75 100 125 150 175 200
-0.125
-0.100
-0.075
-0.050
-0.025
0.000
0.025
0.050
Ave
rage
Ope
n C
ircui
t Pot
entia
l / V
vs.
SCE
Time / minutes
ATS
CaTSSTS
83
A greater percentage of available copper ions would be present as an oxidizing cupric
tetraamine complex in the ammonium thiosulfate leaching system. This should increase the rate
of gold electrode oxidation. In the sodium and calcium thiosulfate leaching systems, the stability
and subsequent effect of these ion pairs could vary greatly depending on the identity of the
cation, as no ammonia is present to stabilize the copper ions. The stability of the ion pairing
between the alkali metal cation and thiosulfate increases as one moves down the rows of the
periodic table21. Therefore, the heavier mass and increased charge of calcium could result in a
much more stable ion pair with thiosulfate than sodium. These differences in ion pairing may
lead to vastly different solvation and diffusion rates in each system; however, further work is
required to confirm these assumptions.
Typically, more positive open circuit potentials are indicative of passivation of the gold
surface. Thus, Figure 5.14 shows that in the first 15 minutes of exposure, the leaching process is
favored over passivation. At low exposure times, both the calcium and ammonium thiosulfate
solutions leach at a greater rate than the sodium thiosulfate leaching system. However, at
exposure times greater than 15 minutes, passivation occurs in these solutions. Differences in the
rate of passivation in the three systems could indicate that the identity of the cation may
influence the level and rate of passivation of the gold surface.
5.3.4 Surface Enhanced Raman
To characterize the effect of copper on the composition and properties of the gold-
thiosulfate interface, gold nanorod electrodes were employed as SERS active surfaces. Leaching
solution was introduced into a spectro-electrochemical cell with an assembled nanorod electrode.
Spectra were collected in situ during a 3 hour gold leaching experiment.
84
Figure 5.15 shows spectra collected for each of the calcium, ammonium and sodium
thiosulfate systems, as a function of exposure time. Upon comparison, it is evident that the SERS
spectra of species at the gold-thiosulfate interface are different from the Raman spectra of the
bulk solution (Figures 5.2-5.4). Unlike the bulk solution, the passive layer at the interface
undergoes significant changes in composition from the initial spectra.
Fourier Self-Deconvolution was used to identify peak positions, and peak fitting was
performed using a mixed Gaussian and Lorentzian band shape, as with the bulk solutions.
Fitted spectra of a gold nanorod electrode, 5 minutes after exposure to the calcium
thiosulfate leaching solution, can be seen in Figure 5.16.
Six bands are visible at positions of approximately 255, 400, 520, 650, 950 and 1010 cm-
1. Upon peak fitting, 10 individual peaks were identified. The broad band at 255 cm-1 was
separated into three peaks at positions of 216, 255, and 285 cm-1. The peak at 216 cm-1 was
assigned to the δ(S-S-S) of S8, while the peaks at 255 and 285 cm-1 could be attributed to either
ν(Cu-S) or ν(Au-S) 28. Although the solution spectra of both trithionate and tetrathionate display
vibrations near 260 cm-1, assignment of the 255 cm-1 peak to these species was avoided due to
the absence of vibrations definitively corresponding to tetrathionate (390, 1040, and 1233 cm-1),
and the lack of correlation in the intensity changes of the trithionate peaks (refer to Table 5.1).
However, the peak at 285 cm-1 is almost a direct match for NaHS on gold as shown by Jeffrey et
al.28
85
Figure 5.15: Raw SERS spectra of gold nanorod electrodes exposed to leaching solutions of 0.1 M S2O3
+ 0.01 M CuSO4, adjusted to pH 8.0-8.5. Solutions were a) CaTS b) ATS c) STS.
a)
b)
c)
86
Identification of the 255 and 285 cm-1 peaks as either ν(Cu-S) or ν(Au-S) is not possible
without further analysis of the samples with a complementary technique, due to the overlap
between the vibrational modes.
Figure 5.16: Fitted SERS spectrum of a gold nanorod electrode, after 5 minutes of exposure to a 0.1 M
CaS2O3 + 0.01 M CuSO4 leaching solution, with an initial pH of 8.0-8.5.
A peak centered at 405 cm-1 was assigned to the δ(S-S) mode of trithionate, although the
peak center was shifted 15 cm-1 lower than the corresponding vibration in aqueous solution. Such
a shift could be the result of adsorption to the gold surface13. A weak shoulder at 443 cm-1
directly correlated to the symmetric S-S vibration of the thiosulfate ion (νsym(S-S)). The weak band
at 530 cm-1 could either be assigned to the δasym(S-O) of the thiosulfate ion or the δasym(O-S-O) of the
tetrathionate ion. The asymmetric band near 650 cm-1 was composed of two strong peaks
centered near 620 and 660 cm-1. The first peak at 620 cm-1 was assigned to a δasym(O-S-O) vibration
of either the sulfate or sulfite ion18, 19. The peak at 660 cm-1 was assigned to the δsym(S-O) of the
thiosulfate ion or the δsym(O-S-O) of the trithionate ion.
100 200 300 400 500 600 700
δ(SS
) S3O
2- 6
ν sym(S
S) S
2O2- 3
δ asym
(SO
) S2O
2- 3 / δ as
ym(O
SO) S
4O2- 6
δ asym
(OSO
) SO
2- 4 / ν sy
m S
O2- 3
δ sym(S
O) S
2O2- 3
/ δ sy
m(O
SO) S
3O2- 6
Raman Shift / cm-1
δ(S-
S-S)
S8
ν(A
u-S)
/ ν(
Cu-S
)
700 800 900 1000 1100
ν sym S
O2- 3
ν sym(S
O) S
2O2- 3
Raman Shift / cm-1
87
The final region of interest, from 900 to 1100 cm-1, displayed two bands at 950 and 1010
cm-1. In conjunction with the peak at 620 cm-1, its most likely that the peak at 950 cm-1
corresponds to the νsym(S-O) of sulfite at the interface rather than sulfate19. As was seen in the bulk
solution, the most intense peak for sulfate should be found near 980 cm-1, but is distinctly lacking
in this spectrum18. Thus, it can be concluded that sulfate is not a species responsible for
passivation of the interface, at least at low exposure times. Based on literature, the strong peak at
1010 cm-1 can be assigned to the gold-thiosulfate complex, Au(S2O3)23! , with the downward shift
in peak position from the aqueous spectrum likely resulting from adsorption to the gold
surface29.
Throughout the duration of the experiment, peaks shifted to both higher and lower
wavenumbers from their initial positions. A shift to higher wavenumber could be the result of a
change in the ionic strength of the solution, or through a change in the dielectric constant of the
SERS active surface as a result of the leaching process30. For example, at longer exposure times,
the 285 cm-1 peak shifted to higher wavenumber centered closer 305 cm-1. However, it is well
documented that this shift occurs as a result of an increase in coverage of the gold surface20, a
process that significantly changes the local dielectric environment of the nanorod substrate31.
Peak shifting was also accompanied by the appearance and disappearance of bands. As
exposure time increased, a strong peak at 469 cm-1 developed, along with a weaker band near
140 cm-1. Both of these peaks may be assigned as vibrations of elemental sulfur (S8); specifically
as ν(S-S) and δ (S-S-S) modes, respectively. 28
88
Deconvolution and peak fitting was also carried out on the ammonium thiosulfate system;
the fitted spectrum of a gold nanorod electrode after 15 minutes of exposure can be seen in
Figure 5.17.
Figure 5.17: Fitted SERS spectrum of a gold nanorod electrode, after 15 minutes of exposure to a 0.1 M
(NH4)2S2O3 + 0.01 M CuSO4 leaching solution, with an initial pH of 8.0-8.5.
Below 670 cm-1, peak positions were similar to the initial spectrum of the calcium
thiosulfate system. Peaks were identified with positions of 216, 255, 300, 400, 443, 530, 610 and
640 cm-1. Assignment of these bands was the same as in the calcium thiosulfate system,
indicating that the passive layer after 15 minutes contained polymeric sulfur, a variety of either
Au-S or Cu-S bonds, trithionate, thiosulfate and either sulfate or sulfite. An additional band at
690 cm-1 was noted, which was tentatively assigned as a vibration of dithionate. Although
literature states that dithionate may not be formed from the interaction of polythionates32, it can
be generated via oxidation of hydrogen sulfite (bisulfite) or sulfite by Cu(II) 3, 32. Since the
leaching solution itself constitutes an oxidizing environment, and the presence of sulfite is
possible, as indicated by the spectra, it is possible that dithionate was generated in this system.
100 200 300 400 500 600 700
δ sym(S
O) S
2O2- 3
/ δ sy
m(O
SO) S
3O2- 6
δ asym
(OSO
) SO
2- 4 / ν sy
m S
O2- 3
δ asym
(SO
) S2O
2- 3 / δ as
ym(O
SO) S
4O2- 6
ν sym(S
S) S
2O2- 3
δ(SS
) S3O
2- 6
ν(A
u-S)
/ ν(
Cu-
S)
δ(SS
S) S
8
Raman Shift / cm-1
900 1000 1100 1200 1300 1400
ν sym(C
O2)N
H2C
OO
-
ν asym
(CO
) CO
2- 3 /
δas
ym(N
H2)
NH
+ 4
ν asym
(SO
) S3O
2- 6 /
HSO
- 3
ν sym(S
O) A
u-S 2O
2- 3
ν SO
3NH
- 2
ν sym(S
O) S
2O2- 3
Raman Shift / cm-1
ν SO
2 / S
O3N
H- 2
89
Significant differences were noted in the SERS spectra of the ammonium and calcium
thiosulfate systems, above 1000 cm-1. An asymmetric band centered near 1000 cm-1 was
composed of two peaks positioned at 998 and 1011 cm-1 that were assigned to the νsym(S-O) of
thiosulfate ion, and a vibration of the gold-thiosulfate complex, respectively. Five additional
peaks were observed between 1200 and 1500 cm-1 at positions of 1237, 1267, 1351, 1404 and
1443 cm-1. The first peak at 1237 cm-1 was assigned as either a νasym(S-O) of trithionate or bisulfite.
Since both of these species have corresponding vibrations in the rest of the spectrum, further
identification is not possible. Peaks located at 1267 and 1351 cm-1 matched literature values for
vibrations of the sulfamate ion, produced through reaction of trithionate with ammonia33:
S3O62! + 2NH3 ! S2O3
2! + SO3NH2! + NH4
+ (5.11)
Such a reaction pathway indicates that decomposition of trithionate at the interface was
occurring, while regenerating thiosulfate. The remaining peaks at 1404 and 1443 cm-1 were
assigned as symmetric vibrations of the carbamate, and carbonate ions, respectively.
The final system analyzed via deconvolution and peak fitting was the sodium thiosulfate
system with copper. The fitted spectrum of a gold nanorod electrode after five minutes of
exposure to the leaching solution can be found in Figure 5.18.
In the first five minutes of exposure, 7 peaks were identified at positions of 256, 300,
391, 443, 525, 613 and 1004 cm-1. As in the previous two systems, these bands were assigned to
ν(Cu-S) or ν(Au-S), δ(S-S) mode of trithionate, νsym(S-S) of thiosulfate, δasym(S-O) of the thiosulfate ion or
the δasym(O-S-O) of the tetrathionate ion, δasym(O-S-O) vibration of either the sulfate or sulfite ion and
νsym(S-O) of thiosulfate ion, respectively. An interesting note here is the absence of elemental
sulfur, which was observed in the initial spectra of both the calcium and ammonium thiosulfate
systems.
90
Figure 5.18: Fitted SERS spectrum of a gold nanorod electrode, after 5 minutes of exposure to a 0.1 M
Na2S2O3 + 0.01 M CuSO4 leaching solution, with an initial pH of 8.0-8.5.
In each of the three thiosulfate systems investigated, six peaks were common amongst the
Raman spectra; vibrational modes corresponding to δ(S-S-S) of S8 (~ 216 cm-1), ν(Cu-S) or ν(Au-S)
vibrations (at ~ 255 and 300 cm-1), δ(S-S) mode of trithionate (~ 400 cm-1), νsym(S-S) of thiosulfate
and the δasym(O-S-O) vibration of either the sulfate or sulfite ion were identified in the three systems
of interest, and their peak areas tracked as a function of exposure time of the gold nanorod
electrodes to the respective leaching solutions. Changes in the peak areas are directly related to
changes in the composition of the passive layer.
Analysis of the peaks’ area as a function of time was performed by normalizing the peak
area at time ‘t’, with respect to the initial value of the area. Thus, plotted values can be viewed as
a percent change. This normalization allows for a semi-quantitative comparison of the rate of
change in area between the 6 peaks of interest, and a comparison of changes in band intensity
100 200 300 400 500 600 700
Raman Shift / cm-1
ν(A
u-S)
/ ν(
Cu-
S)
δ(SS
) S3O
2- 6
ν sym(S
S) S
2O2- 3
δ asym
(SO
) S2O
2- 3 / δ as
ym(O
SO) S
4O2- 6
δ asym
(OSO
) SO
2- 4 / ν sy
m S
O2- 3
900 1000 1100 1200 1300 1400 1500
Raman Shift / cm-1
ν sym(S
O) S
2O2- 3
91
between the 3 electrolytes. Data for the calcium, ammonium, and sodium thiosulfate systems are
presented in Figures 5.19 to 5.21.
Two panels were used to present the peak analysis data; panel (a) groups data that display
large changes in concentration in the first 100 minutes of exposure, and panel (b) contains data
that show small changes in concentration, either in the first 100 minutes of exposure, or over the
duration of the experiment. Panel (b) in Figures 5.19-5.21 also display the open circuit potential,
discussed and shown previously in Figure 5.14, for each of the 3 leaching systems.
92
Figure 5.19: Normalized analytical peak areas for a gold nanorod electrode treated with 0.1 M CaS2O3 +
0.01 M CuSO4 solution adjusted to pH 8.0-8.5. Peak areas tracked were for band positions of a) 216 cm-1
(green circle), 255 cm-1 and 300 cm-1 (brown square), and b) 400 cm-1 (blue circle), 443 cm-1 (red square),
and 610 cm-1 (pink diamond).
0 50 100 150 200 250 3000
4
8
12
16
20
24
28
32
Nor
mal
ized
Pea
k A
rea
Time / minutes
ν(Au-S/Cu-S) [255 cm-1 + 305 cm-1]
δ(SSS) S8
0 50 100 150 200 250 3000
4
8
12
16
20
Nor
mal
ized
Pea
k A
rea
Time / minutes
δ(SS) S3O2-6
νsym(SS) S2O2-3
νsym SO2-3
-0.12
-0.10
-0.08
-0.06
-0.04
-0.02
0.00
Ave
rage
Ope
n C
ircui
t Pot
entia
l / V
a)
b)
93
Over the first hour of exposure in the calcium thiosulfate system, the amount of elemental
sulfur at the interface significantly increased to a maximum, approximately 16 times the initial
concentration. From 60 minutes to 120 minutes, almost 100% of the sulfur at the interface was
removed. Deposition was noted through the remainder of the experiment, resulting in a final
concentration approximately 2.5 times the initial concentration.
Absolute identification of the peaks at 255 cm-1 and 305 cm-1 as either Au-S or Cu-S
interactions was not possible based solely upon the SERS spectra acquired. However, each of
these bands contains information as to the level of passivation at the interface. Thus, the sum of
the areas of both of these peaks should represent the total number of Au-S or Cu-S interactions
making up the passive layer. The change in the normalized area of this sum provides an
indication of the level of passivation.
The change in the normalized area of the Au-S/Cu-S interactions paralleled that of
elemental sulfur at the interface, especially in the first hour of exposure. A maximum value of
approximately 24 times the initial concentration was reached after 90 minutes of exposure. From
approximately 120-300 minutes the normalized area decayed significantly, reaching a final value
approximately 8 times the initial. The large spike at 90 minutes of exposure was most likely the
result of a sudden increase in the surface enhancement from the gold nanorod substrate. The
leaching process is a destructive process, and the sudden increase in enhancement may arise
from a temporary ‘lightning rod’ effect resulting from the collapse of nanorods, bringing them
within 1nm of each other. 34, 35
A significant decrease in the amount of trithionate at the interface occurred in the first
hour of exposure, with almost 100% of the initial value removed either through decomposition or
desorption of trithionate from the gold surface. A spike in the amount of trithionate at the
94
interface was observed after 90 minutes. By 120 minutes, the concentration of trithionate had
returned to a level on par with the initial surface concentration; the final concentration at the
interface was approximately the same level as the initial.
Thiosulfate at the interface experienced an increase in the surface concentration (to
approximately 4 times its initial value) during the first hour of exposure. This may be a result of
the decomposition of trithionate36, 37:
2S3O62! + 6OH! ! S2O3
2! + 4SO32! + 3H2O (5.12)
At the same time, decomposition of thiosulfate to form elemental sulfur was occurring,
providing a competitive pathway between production and decomposition of thiosulfate at the
interface2:
2S2O32! + H2O ! 2SO4
2! + 4S + OH! (5.13)
3S2O32! + 6OH! ! 4SO3
2! + 2S2! + 3H2O (5.14)
S2O32! ! SO3
2! + S0 (5.15)
The thiosulfate concentration at the interface over the remaining time of the experiment
was observed to increase greatly, with a final concentration approximately 20 times the initial.
The behavior of sulfite in the first 90 minutes of exposure was similar to that of
thiosulfate; an increase to approximately 4 times the initial concentration was observed. This
production is further support of Equation 5.12. Sulfite at the interface experienced a significant
decay from 90 minutes to approximately 260 minutes of exposure; such an effect could be
produced by oxidation of sulfite by Cu(II) to form sulfate23:
2Cu2+ + SO32! + OH! ! 2Cu2+ + SO4
2! + H2O (5.16)
However, sulfate was only observed in solution and not at the interface. Therefore, if
sulfate is being formed as a by-product of sulfite oxidation, or other decomposition reactions, it
95
must be rapidly transported to the bulk upon formation. The final concentration of sulfite was
approximately the same as the initial value.
Figure 5.19b also displays the open circuit potential of the calcium thiosulfate system as a
function of exposure time. While the time scale of the SERS leaching experiment and open
circuit potential do not directly overlap, some correlation can be drawn between the species
present and the rate of passivation. The OCP potential curve displays a minimum, followed by an
increase between 15 and 60 minutes of exposure. During this time, an increase in elemental
sulfur was observed, along with a decrease in the concentration of trithionate; removal and
deposition of these species may be responsible for the curvature in this region of the potential
plot. Although almost all trithionate was removed in the first hour, the concentration of elemental
sulfur increases to approximately 16 times its initial concentration. Hence, the open circuit
potential experienced a steep rise as the surface became extensively passivated. Between 90 and
120 minutes, a plateau is observed in the potential curve. This corresponds to the achievement of
a steady state concentration of elemental sulfur, and trithionate. Such a result could indicate that
the rate of passivation depends on a combination of changes in the surface concentration of both
trithionate and elemental sulfur in this system.
Peak area analysis of the interfacial species for the ammonium thiosulfate system can be
found in Figure 5.20 (a) and (b). The overall changes in band intensity were much smaller in the
ammonium thiosulfate system than the corresponding changes in the calcium thiosulfate system.
The data are presented in two panels that plot the same normalized band intensities as panels (a)
and (b) in Figure 5.19.
96
Figure 5.20: Analytical peak areas for a gold nanorod electrode treated with 0.1 M (NH4)2S2O3 + 0.01 M
CuSO4 solution adjusted to pH 8.0-8.5. Peak areas tracked were for band positions of: a) 216 cm-1 (green
circle), 255 cm-1 and 300 cm-1 (brown square), and b) 400 cm-1 (blue circle), 443 cm-1 (red square), and
610 cm-1 (pink diamond).
0 20 40 60 80 100 120 140 160 180 2000
4
8
Nor
mal
ized
Pea
k A
rea
Time / minutes
ν(Au-S/Cu-S) [255 cm-1+ 305 cm-1]
δ(SSS) S8
0 20 40 60 80 100 120 140 160 180 2000
1
2
3
4
Nor
mal
ized
Pea
k A
rea
Time / minutes
δ(SS) S3O2-6
νsym(SS) S2O2-3
δasym(OSO) SO2-4 / νsym SO2-
3
-0.12
-0.10
-0.08
-0.06
-0.04
-0.02
0.00
Ave
rage
Ope
n C
ircui
t Pot
entia
l / V
a)
b)
97
More scatter was observed in the ATS data (Figure 5.20), however, the scale is expanded
which may contribute to the random appearance of the data. This scatter may be due to
fluctuations in the surface enhancement as a result of physical damage to the nanorods caused by
dissolution during the experiment.
The concentration of elemental sulfur at the gold-thiosulfate interface in the ammonium
thiosulfate system increases significantly over the course of the experiment, achieving a
maximum value approximately 9 times the initial concentration. Removal of sulfur from the
interface was observed over some periods of the experiment. This resulted in a final
concentration approximately 4 times the initial.
The normalized area for Au-S/Cu-S interactions at the surface increased over the duration
of the experiment, reaching a maximum value of 3 times the initial area. The final normalized
area at the interface is approximately 2 times the initial value.
Allowing for scatter, the data in Figure 5.20 (b) show a general increase in all peak
intensities over the course of the experiment, with the exception of the δ(S-S) mode of trithionate.
Over the first hour of the experiment the concentration of trithionate at the surface decayed. A
large spike was observed after 90 minutes, as in the calcium thiosulfate system. Again, this spike
may arise from a sudden increase in enhancement arising from a temporary lightning rod effect34,
35. Decay of trithionate at the interface continued from 90 minutes onward, resulting in a final
concentration one-tenth the initial. The final values for thiosulfate and sulfite at the interface
were 3 times, and on par with the initial values, respectively.
The open circuit potential of the ammonium thiosulfate system (Figure 5.20 b) displays a
much slower rate of passivation at longer exposure times than the calcium thiosulfate system.
This may be the result of constant removal of trithionate from the interface, and alternating
98
periods of deposition and removal of elemental sulfur at the interface; a slower, and smaller,
change in the amount of all adsorbed species may also contribute. Because neither trithionate nor
sulfur was completely removed, or achieved a steady state, passivation of the surface continued;
hence, the OCP of the system moved consistently towards more positive potentials.
Decay of trithionate at the interface may be the result of interaction with ammonia,
resulting in the production of sulfamate (Equation 5.11), which was observed at the interface at
early exposure times. This reaction also produces thiosulfate, which may contribute to the
periodic behavior observed for this species.
The composition of the passive layer formed on the gold surface in the sodium thiosulfate
system displayed much different behavior over the duration of the experiment than in the
previous two systems. Normalized peak area analysis for the sodium thiosulfate system can be
found in Figure 5.21 (a) and (b). Data in both panels displayed a sharp increase at 80 minutes of
exposure time. Assuming that this increase is due to a temporary lightning rod effect, as in the
calcium and ammonium thiosulfate systems, the data in Figure 5.21 (b) show a continuous quasi-
linear increase of all band intensities with time. The number of Au-S/Cu-S interactions at the
interface reached a final value approximately 18 times the initial. Elemental sulfur at the
interface also experienced a large increase, reaching a final value approximately 9 times the
initial. Final values for thiosulfate, trithionate were approximately two times the initial
concentration, and sulfite experienced a 3-fold increase over the duration of the experiment.
99
Figure 5.21: Analytical peak areas for a gold nanorod electrode treated with 0.1 M Na2S2O3 + 0.01 M
CuSO4 solution adjusted to pH 8.0-8.5. Peak areas tracked were for band positions of: a) 216 cm-1 (green
circle), 255 cm-1 and 300 cm-1 (brown square), and b) 400 cm-1 (blue circle), 443 cm-1 (red square), and
610 cm-1 (pink diamond).
0 20 40 60 80 100 120 140 160 180 2000
4
8
12
16
20
Nor
mal
ized
Pea
k A
rea
Time / minutes
ν(Au-S/Cu-S) [255 cm-1 + 305 cm-1]
δ(SSS) S8
0 20 40 60 80 100 120 140 160 180 2000
4
8
12
Nor
mal
ized
Pea
k A
rea
Time / minutes
νsym(SS) S2O2-3
δ(SS) S3O2-6
δasym(OSO) SO2-4 / νsym SO2-
3
-0.12
-0.08
-0.04
0.00
0.04
Ave
rage
Ope
n C
ircui
t Pot
entia
l / V
a)
b)
100
The sodium thiosulfate system showed the highest level of passivation at early exposure
times, as shown by the negative limit in the open circuit potential of the system in Figure 5.14
(the OCP is also displayed in Figure 5.21 (b)). However, it also displayed the slowest rate of
change over the duration of the leaching experiment; a linear increase in the OCP was observed
from 25-180 minutes of exposure. A gradual linear increase was also observed in the behavior of
trithionate, sulfite, thiosulfate, and up to 160 minutes, sulfur. Because no single species displays
behavior solely related to the trend in the open circuit potential, it can be concluded that in the
sodium thiosulfate system the passive layer is complex, and no single species can be assigned
responsibility for passivation of the gold surface.
It has been proposed that Cu(II) may aid in gold dissolution through the removal of sulfur
contaminants from the surface by scavenging sulfide, or through interaction with sulfur chains on
the surface6. If, in the work presented in this chapter, the peak observed at ~305 cm-1 in all 3
systems corresponds to a Au-S vibration28, then the peak at ~255 cm-1 could correspond to Cu-S
vibrations. Increases observed in the area of the Cu-S peak may have been a result of interaction
of Cu(II) with sulfur on the gold surface, and the observed decrease in the calcium and
ammonium thiosulfate systems could correspond to subsequent diffusion from the interface to
the bulk solution. This is possible, as leaching solutions left for several hours after preparation
reproducibly precipitated black copper sulfides.
Peaks corresponding to characteristic vibrations of tetrathionate were not observed in the
SERS spectra collected, thus, it can be concluded that in the presence of Cu(II) in these systems,
tetrathionate is not present at the interface, or if it is formed, undergoes rapid decomposition.
101
5.4 SUMMARY
The composition and behavior of species in the bulk leaching solution, and at the gold
thiosulfate interface in the presence of copper, were characterized through complementary
electrochemical and spectroscopic techniques. The bulk solutions for the calcium, ammonium
and sodium thiosulfate systems all maintained steady state concentrations of thiosulfate,
tetrathionate, sulfate, and trithionate. These systems were unaffected by the changes in pH
caused by carbon dioxide absorption and precipitation of calcium carbonate (or calcium sulfate).
Solubility values for each of these salts indicate that it is most likely calcium carbonate that is
precipitating from solution.
The gold-thiosulfate interface showed large differences between the three systems.
Tetrathionate and sulfate were found to be absent in each system at the interface; however,
elemental sulfur, and sulfite were present. The negative limits of the open circuit potentials
indicated that the level of passivation was most significant in the sodium thiosulfate system,
followed by calcium and ammonium thiosulfate, respectively. This result was also reflected in
the SERS results. Table 5.2 displays the final change in the normalized peak area of the six
bands analyzed in this thesis (to δ(S-S-S) of S8 [~ 216 cm-1], ν(Cu-S) or ν(Au-S) vibrations [at ~ 255
and 300 cm-1], δ(S-S) mode of trithionate [~ 400 cm-1], νsym(S-S) of thiosulfate and the δasym(O-S-O)
vibration of either the sulfate or sulfite ion), for the calcium, ammonium, and sodium thiosulfate
systems.
102
Table 5.2: Final change in the normalized peak areas investigated in the calcium, ammonium and
thiosulfate systems.
Overall Change in the Normalized Peak Area
Peak Calcium Thiosulfate Ammonium Thiosulfate Sodium Thiosulfate
δ(S-S-S) of S8 [~ 216 cm-1] 2.6 4.6 9.4
ν(Cu-S) / ν(Au-S) [~ 255 and 300 cm-1] 8.0 2.1 17.8
δ(S-S) of S3O62-
[~ 400 cm-1] 1.2 0.1 2.4
νsym(S-S) of S2O32-
[~ 445 cm-1] 19.8 3.2 2.4
δasym(O-S-O) of SO32-/SO4
2- [~ 610 cm-1] 1.4 0.8 3.1
The overall change in the normalized areas are representative of change in the passive
layer at the gold surface over the duration of the experiment. In terms of the relative increase
from the initial concentration at the surface, the sodium thiosulfate system experienced the
greatest overall increase in the amount of Au-S or Cu-S interactions at the interface, followed by
the calcium, and ammonium thiosulfate systems, respectively. A higher level of passivation (or
larger number of Au-S/Cu-S interactions at the surface) should result in a lower leaching rate.
The results in Table 5.2 match what is expected based on the results from literature21. The
calcium thiosulfate and sodium thiosulfate systems show significant differences in the behavior
of thiosulfate at the interface. An increase to 2.4 times the initial concentration was noted in the
sodium thiosulfate system; however, a massive increase in calcium thiosulfate system was
observed, with the final concentration almost 20 times the initial. Such a large concentration of
thiosulfate at the interface should lead to promotion of leaching.
103
Large amounts of sulfur were present at the interface at early exposure times in both the
calcium and ammonium systems; however, each of these two systems displayed a significant
ability for removal of sulfur from the interface (final values for the concentration at the interface
for the calcium and ammonium systems are 2.6 and 4.6 times the initial concentrations,
respectively). Comparably, trithionate was observed to decompose in both of these systems as
well. Almost 100% removal occurred over the course of the experiment in the ammonium
thiosulfate system (the final surface concentration was one-tenth the initial), while the calcium
thiosulfate system displayed decomposition only at early exposure times (final surface
concentration was 1.2 times the initial). The combined presence of both elemental sulfur and
trithionate at the interface in calcium thiosulfate may explain the higher level of passivation
when compared with the ammonium thiosulfate system, where elemental sulfur appears to be the
main component of the passive layer. The increased concentration of Cu(II) as a result of
stabilization by NH3 in the ammonium thiosulfate solution should lead to increased leaching, and
minimize degradation of thiosulfate.
Previous studies13, 20 of the gold-thiosulfate interface in the absence of copper displayed
evidence of thiosulfate and tetrathionate in the passive layer at early exposure times. Extended
exposure showed that tetrathionate at the interface degraded, forming trithionate and thiosulfate;
such behavior reflects known solution reactions. Normalized peak area analysis of the work by
Baron13 can be found in Figure 5.22.
104
Figure 5.22: Normalized analytical peak area of a gold nanorod electrode treated with 0.1 M Na2S2O3
solution adjusted to pH 10. Modified from13.
In the presence of copper, an enhanced degradation rate of higher polythionate species
was noted. Also, elemental sulfur was identified at early exposure times at the interface in some
systems. Thus, copper appears to be aiding in removal of polythionates from the interface, but
may also be contributing to enhanced degradation of thiosulfate to form elemental sulfur, a
species which seems to contribute to passivation. Introduction of copper is both beneficial and
detrimental to the leaching system, as leaching appears to be promoted, as well as passivation.
References
1. Abbruzzese, C.; Fornari, P.; Massidda, R.; Vegliņ, F.; Ubaldini, S. Hydrometallurgy
1995, 39, 265-276. 2. Aylmore, M. G.; Muir, D. M. Miner Metall Process 2001, 18, 221-227. 3. Aylmore, M. G.; Muir, D. M. Minerals Eng 2001, 14, 135-174. 4. Byerley, J. J.; Fouda, S. A.; Rempel, G. L. J.Chem.Soc., Dalton Trans. 1973, , 889-893. 5. Pedraza, A. M.; Villegas, I.; Freund, P. L.; Chornik, B. J Electroanal Chem 1988, 250,
443-449. 6. Zhang, S. C.; Nicol, M. J. J. Appl. Electrochem. 2005, 35, 339-345.
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7. Breuer, P. L.; Jeffrey, M. I. Hydrometallurgy 2002, 65, 145-157. 8. Baron, J.; Szymanski, G.; Lipkowski, J. J Electroanal Chem 2011, 662, 57-63. 9. Meyer, B.; Ospina, M. Phosphorus Sulfur and Silicon and the Related Elements 1982, 14,
23-36. 10. Sato, S.; Higuchi, S.; Tanaka, S. Appl. Spectrosc. 1985, 39, 822-827. 11. Haigh, J.; Hendra, P.; Rowlands, A.; Degen, I.; Newman, G. Spectrochim. Acta, Pt. A:
Mol. Spectrosc. 1993, 49, 723-725. 12. Fanigliulo, A.; Bozzini, B. Trans. Inst. Met. Finish. 2002, 80, 132-136. 13. Baron Gavidia, J. Study of the Gold-Thiosulfate Interface Under Leaching Conditions,
University of Guelph, Guelph, ON, 2010. 14. Antohe, V. A.; Radu, A.; Matefi-Tempfli, M.; Attout, A.; Yunus, S.; Bertrand, P.; Dutu,
C. A.; Vlad, A.; Melinte, S.; Matefi-Tempfli, S.; Piraux, L. Appl. Phys. Lett. 2009, 94, 073118.
15. Chang, J. C. In Section 8: Analytical Chemistry; Lide, D. R., Ed.; CRC Handbook of Chemistry and Physics; CRC Press: Boca Raton, 1990; pp 8-39.
16. Surewicz, W. K.; Mantsch, H. H. Biochim. Biophys. Acta 1988, 952. 17. Meyer, B.; Ospina, M. Abstracts of Papers of the American Chemical Society 1982, 183,
56-INOR. 18. Daly, F. P.; Brown, C. W.; Kester, D. R. J. Phys. Chem. 1972, 76, 3664-3668. 19. Evans, J. C.; Bernstein, H. J. Canadian Journal of Chemistry-Revue Canadienne De
Chimie 1955, 33, 1270-1272. 20. Jeffrey, M.; Watling, K.; Hope, G. A.; Woods, R. Minerals Eng 2008, 21, 443-452. 21. Chandra, I.; Jeffrey, M. I. Hydrometallurgy 2004, 73, 305-312. 22. Jeffrey, M. I.; Linda, L.; Breuer, P. L.; Chu, C. K. Minerals Eng 2002, 15, 1173-1180. 23. Molleman, E.; Dreisinger, D. Hydrometallurgy 2002, 66, 1-21. 24. Arima, H.; Fujita, T.; Yen, W. Miner Metall Process 2003, 20, 81-92. 25. Wan, R. Y.; LeVier, K. M. Int. J. Miner. Process. 2003, 72, 311-322. 26. Senanayake, G. Hydrometallurgy 2004, 75, 55-75. 27. Senanayake, G. Hydrometallurgy 2005, 77, 287-293. 28. Jeffrey, M. I.; Watling, K.; Hope, G. A.; Woods, R. Minerals Eng 2008, 21, 443-452. 29. Watling, K.; Hope, G. A.; Jeffrey, M. I.; Woods, R. ECS Transactions 2006, 2, 121-132. 30. Moskovits, M. J. Raman Spectrosc. 2005, 36, 485-496. 31. Jain, P. K.; Huang, X.; El-Sayed, I. H.; El-Sayed, M. A. Plasmonics 2007, 2, 107-118. 32. Nickless, G. In Inorganic sulphur chemistry; Elsevier Publishing Company: 1968; . 33. Naito, K.; Hayata, H.; Mochizuki, M. Journal of Inorganic and Nuclear Chemistry 1975,
37, 1453-1457. 34. Tian, Z. Q.; Ren, B.; Wu, D. Y. J Phys Chem B 2002, 106, 9463-9483. 35. Stiles, P. L.; Dieringer, J. A.; Shah, N. C.; Van Duyne, R. R. Annual Review of Analytical
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106
CHAPTER 6: CONCLUSIONS AND FUTURE WORK
6.1 CONCLUSIONS
Overall this thesis has described the composition and behavior of the bulk solutions, and
the passive layer that forms on the gold surface, in three different thiosulfate-leach solutions
containing copper. In an attempt to identify passivating species at the interface, the open circuit
potentials and pH of all three systems have been characterized along with the kinetics of the
electron transfer reaction. This work was supported by NSERC and Barrick Gold Inc.
Beginning with the characterization of the bulk solutions of interest (calcium, ammonium
and sodium thiosulfate with Cu(II)), some conclusions can be drawn. Firstly, the identity of the
cation in thiosulfate salt has no effect on the composition of the bulk solution; all three solutions
displayed identical compositions. In each solution thiosulfate, trithionate, tetrathionate and
sulfate were identified. Secondly, the time dependent concentration of these species is stable, and
independent of pH. Each of the three solutions displayed steady state concentrations of
thiosulfate, tetrathionate, trithionate, and sulfate over a period of 3 hours. This is likely an effect
of Cu(II) in solution, as it is known to catalyze decomposition of both thiosulfate and
polythionates1, leading to regeneration of both.
At the gold-thiosulfate interface, in the presence of copper, little correlation to the
characteristics of the bulk solution was observed. Species identified at the interface included
elemental sulfur, trithionate, thiosulfate, and sulfite. Polythionates higher than trithionate were
not observed. The absence of higher polythionates is likely an effect of the local pH of the
interface. The cation of the thiosulfate salt was found to have no effect on the initial leaching rate
107
of gold; however, significant differences were noted in the open circuit potential curves of a gold
disc electrode exposed to the calcium, ammonium and sodium thiosulfate solutions. This trend in
the open circuit potentials indicates that there are differing levels and rates of passivation
between the three systems. The time dependent composition of the passive layer was drastically
different between calcium, ammonium and sodium thiosulfate systems. Both the calcium and
ammonium systems displayed the ability to remove trithionate from the interface, as well as
elemental sulfur. This was not seen in the sodium system, which displayed a slower rate of
passivation, once the negative limit was reached. However, Cu(II) has been proposed to
scavenge sulfur contaminants from the gold surface1; this interaction maybe responsible for the
removal elemental sulfur observed. Thus, more work is needed to confirm whether the behavior
of elemental sulfur at the interface is an effect of the cation, or copper.
One common feature between the systems was the behavior of trithionate at the interface.
In all 3 systems, a large spike in the concentration of trithionate is observed between 75-90
minutes exposure of the gold electrode to leaching solution. Rapid removal or decomposition of
trithionate was then observed, returning the concentration at the interface to near pre-spike
values. This large spike is likely the result of a temporary lightning rod effect created as a result
of dissolution of the gold nanorod electrode.
6.2 FUTURE WORK
Potential roles of both copper and the cation of the thiosulfate salt have been proposed in
the scope of this thesis. However, further comparative studies are required before definitive
conclusions can be drawn. In order to determine if the cation of the thiosulfate salt is having a
significant effect on the time dependent composition of the passive layer, electrochemical
measurements of the leaching current as a function of time should be performed. If the cation
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effect arises from the extent or rate of passivation, a significant difference in leaching currents
should be noted for a non-passivated and sequentially passivated electrode surface.
Although open circuit potentials for the calcium, ammonium and sodium thiosulfate
systems in the presence of copper have been recorded, a polished gold disk electrode was
employed. This “smooth” gold electrode had a significantly different geometry than the gold
nanorod electrode employed in the SERS studies. Differing geometries could lead to different
rates of mass transport, and subsequently different rates of adsorption. Thus, measurement of the
open circuit potential of the gold nanorod electrode directly would allow for correlation of SERS
results and the open circuit potentials of the systems. It would allow for tracking peak shifts
within the SERS spectrum, giving insight into the adsorption and desorption characteristics of
each species. As well, SERS studies for all three salts must be performed in the absence of
copper in the leaching solution. These studies have been executed on the sodium thiosulfate
system, however, characterization of the ammonium and calcium thiosulfate systems must be
completed. This will clarify the interaction of copper with adsorbed elemental sulfur. If Cu(II) is
unable to scavenge sulfur or sulfide from the surface, a constant buildup would be expected in
the copper free experiments, in all solutions.
Addition of ammonia to each of the leaching solutions, in the absence of copper would
allow for an understanding of the individual effect of this additive. SERS results may indicate
whether ammonia/ammonium is adsorbing at the interface. Once complete, SERS studies of the
leaching solutions with both copper and ammonia can be performed. In this way, the individual
and tandem role of both additives can be identified.
Extending both SERS and solution Raman studies beyond the exposure limit within this
work would greatly enhance the industrial applicability of the results, as industrial leaching
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processes are extremely long-scale (eg. up to 90 days). Identifying the end products at the
interface, and bulk solution would allow for greater understanding and optimization of the leach
system.
References
1. Zhang, S. C.; Nicol, M. J. J. Appl. Electrochem. 2005, 35, 339-345.