EVALUATING THE EFFICIENCY OF CYCLIC VOLTAMMETRY USING MODIFIED AND UNMODIFIED GLASSY CARBON ELECTRODES FOR THE ANALYSIS OF

SULFUR CONTAINING COMPOUNDS IN WATER

ASHLEY LYNN RYAN

A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE

DEGREE OF MASTER OF ENVIRONMENTAL SCIENCE

SCHOOL OF GRADUATE STUDIES NIPISSING UNIVERSITY NORTH BAY, ONTARIO

JULY, 2017

©ASHLEY LYNN RYAN, 2017

ii

Signature Page

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Abstract

When low oxidation state sulfur compounds such as sulfides dissolve in water, they create an

undesirable rotten egg smell. A prolonged exposure to concentrations of about 0.01 mg L1- of

H2S has been known to have negative effects on human gastro-intestinal, and nervous

systems. Concentrations of dissolved low oxidation state sulfur-based compounds in the

range of 0.01 to 0.5 mg L1- in aquatic systems have been shown to have chronic and acute

toxic effects in aquatic organisms. The ability to qualitatively and quantitatively determine

sulfur-based compounds in water, and further be able to speciate these compounds in water

would be of great significance. Sulfur-based compounds tend to disproportionate in aqueous

solutions to form compounds, the identity of which is dependent on the pH, temperature, and

ionic strength of the solution. The work that was investigated in this study involved the

detection of simple sulfides (H2S, HS- and S2-) in water through use of cyclic voltammetry

(CV), iodometric, and colorimetric techniques. In addition, the analysis of polysulfide and

polythionate compounds using these techniques were also investigated in this study. The

sodium salts of disulfide (Na2S2), trisulfide (Na2S3), and tetrasulfide (Na2S4) were used in

this study as representatives of polysulfides. These compounds were not available

commercially and were therefore synthesized. With regards to the CV technique, the bare

glassy carbon electrode (BGCE) and the vanadium oxide modified glassy carbon electrode

(MGCE) were investigated for their suitability to analyze the polysulfides and thiosulfate.

The results obtained in this work demonstrated a potential for use of the CV technique for the

qualitative analysis of the above class of compounds. It was however found necessary to first

reduce the above type of sulfur compounds to simple sulfide (S2-) before they could

successfully be analyzed using the iodometric and colorimetric techniques. Even though the

latter two techniques were successfully used for the quantitative analysis of sulfide, the CV

technique did not show much success. This was primarily as a result of the unpredictable

magnitudes of the various peak currents that resulted. In addition, the background currents

especially when the MGCE was used were found to interfere with the peaks associated with

sulfide. The colorimetric method and iodometric methods were found to have a great

agreement to each other (F (1,10) = 0.418, p = 0.532) for the determination for reduced sulfur-

based compounds.

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Keywords

Electrochemical Analysis, Sulfide, Polysulfide, Glassy Carbon Electrode, Vanadium Oxide,

Electrode Modification

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Acknowledgments

I would like to thank my supervisor Dr. Stephen Kariuki for offering me the opportunity to

take on this project, and for his continued guidance and support throughout. All of his advice

and encouragement was greatly appreciated. I would also like to thank my committee

members, Dr. Mukund Jha and Dr. Samuel Mugo, for reviewing the enclosed research, for

their advice and Dr. Jeffery Dech for chairing my defense.

Exclusive thank you goes to Dr. Lesley Lovett-Doust and Dr. Jeffery Dech for their endless

candor, knowledge and guidance through the editing process. Thank you goes out to anyone

else who was brave enough to read this document. It was very much appreciated. I would

also like to thank Dr. Krys Chutko for collecting one of the water samples that was used to

test the methodologies outlined in this work.

Recognition also goes out to the Instrumental Laboratory in Lakehead University (LUIL) for

the work on the surface characterizations using SEM-EDX technology, and ICP analyses for

the water sample. To iCAMP though Canadore College for the preliminary exploration of

the electrode surfaces through SEM. Also, thanks to Willam Zhe from Laurentian University

for the SEM and EDX work done there as well.

I am thankful for my family, friends, and colleagues for their advice, guidance and continued

fortitude throughout this endeavor. I am especially thankful to my spouse, Tommi Jo, for

being an ever patient and supportive partner while I took the time to complete this work.

During the course of this project I spent a sleepless night watching a re-run of Back to the

Future. I am reminded of the antics of Dr. Emmett Brown (Christopher Lloyd) and can relate

a few quotes from the movies to this work. It was challenging to pick just one.

I end with this:

“You're not thinking fourth dimensionally!” – Dr. Emmet Brown

Cheers

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Table of Contents

Signature Page .................................................................................................................... ii

Abstract .............................................................................................................................. iii

Acknowledgments............................................................................................................... v

Table of Contents ............................................................................................................... vi

List of Tables ..................................................................................................................... ix

List of Figures .................................................................................................................... xi

List of Abbreviations and Symbols................................................................................... xv

Overview and Research Objectives .................................................................................... 1

Chapter 1 ........................................................................................................................... 14

Sulfide (HS-, H2S and S2-) ................................................................................................. 14

1.1 INTRODUCTION ................................................................................................ 14

1.2 MATERIALS AND METHODS .......................................................................... 15

1.2.1 Reagents and Solutions ............................................................................. 15

1.2.2 Standardization of 25 mM Thiosulfate and 25 mM Iodine Solutions ...... 17

1.2.3 Standardization of Sulfide Using Iodometric Titration ............................ 18

1.2.4 Electrochemical Analysis of Sulfide ......................................................... 18

1.3 RESULTS ............................................................................................................. 21

1.3.1 Standardization of Iodine and Thiosulfate Solutions ................................ 21

1.3.2 Standardization of Sulfide Solutions using Iodometric Titration ............. 21

1.3.3 Electrochemical Analysis of Sulfide ......................................................... 22

1.4 CONCLUSIONS................................................................................................... 49

Chapter 2 ........................................................................................................................... 54

Analysis of Polysulfides and Polythionates ...................................................................... 54

2.1 INTRODUCTION ................................................................................................ 54

vii

2.2 MATERIALS AND METHODS .......................................................................... 56

2.2.1 Making Na-polysulfides............................................................................ 56

2.2.2 N,N-Diethyl-p-phenylenediamine (DEPD) Analysis of Sulfide ............... 59

2.2.3 Reduction of Na-polysulfides and polythionates Using Chromium as the

Reducing Agent ........................................................................................ 59

2.3 RESULTS ............................................................................................................. 61

2.3.1 Making Na-polysulfides............................................................................ 61

2.3.2 Reduction and Analysis of Na-polysulfides and Polythionates ................ 63

2.4 CONCLUSIONS................................................................................................... 66

Chapter 3 ........................................................................................................................... 68

Electrochemical Analysis of Na-polysulfides and Polythionates ..................................... 68

3.1 INTRODUCTION ................................................................................................ 68

3.2 MATERIALS AND METHODS .......................................................................... 70

3.2.1 Equipment ................................................................................................. 70

3.2.2 Dissolution of Na-polysulfides and Thiosulfate Ions ............................... 70

3.3 RESULTS ............................................................................................................. 71

3.3.1 Electrochemical Analysis of Na-polysulfides and Polythionates at BGCE

and MGCE ................................................................................................ 71

3.4 CONCLUSIONS................................................................................................... 80

Chapter 4 ........................................................................................................................... 82

Analysis of Low Oxidation-State Sulfur Species in Water .............................................. 82

4.1 INTRODUCTION ................................................................................................ 82

4.2 MATERIALS AND METHODS .......................................................................... 83

4.2.1 Water Source and Collection Parameters ................................................. 83

4.2.2 Reduction and DEPD Analysis of Lake Water Sample ............................ 84

4.2.3 Electrochemical Analysis of Lake Water Samples ................................... 84

4.3 RESULTS ............................................................................................................. 85

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4.3.1 Reduction and DEPD Analysis of Lake Water Sample ............................ 85

4.3.2 Electrochemical Analysis of Lake Water ................................................. 90

4.4 CONCLUSIONS................................................................................................... 91

General Conclusions ......................................................................................................... 93

Appendix A ....................................................................................................................... 96

Appendix B ....................................................................................................................... 99

References ....................................................................................................................... 102

Curriculum Vitae: Ashley Ryan (nee Marcellus) ........................................................... 133

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List of Tables

Table 1. Standardized values (mM) for thiosulfate (S2O32-) and iodine (I2) used in iodometric

determination of sulfide. Values are the mean of n = 5. ......................................................... 21

Table 2. Mean potential positions and peak current (µA) values in the anodic segment for a 1

mM solution of HS- through the 2nd to the 10th CV cycle. ...................................................... 27

Table 3. Peak currents at the cathodic peak at -1208 mV for CV cycles 1 through 10 for a 1

mM HS- solution. .................................................................................................................... 29

Table 4. Mean of six replicate runs and ± SE of current responses (Ipc) for CV analysis of

sulfide obtained at the BGCE at the Epc of -1263 mV. ........................................................... 31

Table 5. Mean CV anodic currents (Ipa) for the sulfide analysis at the BGCE at potentials of

0, +565, and +1142 mV. ......................................................................................................... 31

Table 6. Peak current, potential and scan number for 0.36 mM HS- run over MGCE at a scan

rate of 100 mV s1-. .................................................................................................................. 46

Table 7. Mean ± SE for anodic current response Ipa (µA) of various concentrations of HS- at

MGCE; at a mean anodic potential Epa of +200 mV, (n = 6; R2 = 0.937). ............................. 48

Table 8. Mean ± SE for cathodic current response Ipc (µA) of various concentrations of HS- at

MGCE at a mean cathodic potential Epc of -570 mV, (n = 6; R2 = 0.907). ............................ 49

Table 9. Temperature profile for the preparation of Na-polysulfides (Rosen and Tegman,

1971). ...................................................................................................................................... 58

Table 10. Comparison mean % yields ± SE of reduced Na-polysulfides from two separate

batch products, summer 2015 and winter 2016. ..................................................................... 63

Table 11. Mean and SE ± for various concentrations of sulfide used to develop a standard

curve for DEPD analyses (n=11) ............................................................................................ 64

Table 12. Mean percent yields ± SE of the reduced polysulfides after applying heat at 60 °C

and reducing the solution for 4 hours (n = 5). ........................................................................ 66

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Table 13. Weights (g) and concentrations (mM) of Na-Polysulfides used for CV analyses .. 71

Table 14. Anodic segment peak potential (mV) and peak current (µA) for CV analysis of Sn2-

ions (1 mM) at BGCE (n = 3). ................................................................................................ 73

Table 15. Cathodic segment peak potential (mV) and peak current (µA) for CV analysis of

Sn2- ions (1 mM) at BGCE (n = 3). ......................................................................................... 73

Table 16. Inorganic metal concentrations (ppm) from site IG9 for water sample collected in

October of 2015. ..................................................................................................................... 87

Table 17. Sediment concentrations (ppm), of the metals (and sulfur) associated with smelting

of nickel ores. Mean values (mg/kg ± standard error) (Chase et al. 2016, unpublished results).

................................................................................................................................................. 87

Table 18. Comparison of percent recoveries and recovered concentrations (mM) of sulfide

spiked water samples for filtered and unfiltered samples. ...................................................... 88

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List of Figures

Figure 1: Electrochemical cell setup for chemical analysis. (A) From left to right: Reference

electrode (Ag/AgCl) shaded grey, glassy carbon electrode (GCE) shaded black and platinum

auxiliary electrode (Pt) represented as light grey coil. (B) The BASi setup for

electrochemical analysis, following the same electrode order from left to right as is the case

in A.......................................................................................................................................... 19

Figure 2. Fraction dissociation plot for sulfide as a function of pH. Where indications of

sulfide species in solution are H2S(aq) (♦), HS- (■) and S2- (▲). ............................................. 23

Figure 3. Overlapped voltammograms showing five CV cycles of HS- (1 mM), with the

supporting electrolyte at BGCE with a scan rate of 100 mV s-1. ............................................ 25

Figure 4. Overlapped voltammograms showing 10 – 15 CV cycles of 1 mM HS- solution at

the BGCE with supporting electrolyte solution (dashed line), at scan rate of 100 mV s-1. .... 25

Figure 5. Epa positions (mV) at -6 (♦), +434 (▲), and +1158 (■) with peak current values

(Ipa) µA at CV cycles 2 and 5-10 for 1 mM HS-. .................................................................... 26

Figure 6. Epc position (mV) at -1208 with peak current values (Ipc) µA at CV cycles 2 and 5-

10 for 1 mM HS-. .................................................................................................................... 28

Figure 7. CV analysis for the various HS- concentrations at a mean Epc = -1263 mV at BGCE.

Error bars represent ± S.E. of six replicate runs (R2 = 0.988). ............................................... 30

Figure 8. CV anodic currents (Ipa) for various HS- concentrations at Epa = 0 mV (♦) R2 =

0.628, +565 mV (■) R2 = 0.926, and +1142 mV (▲) R2 = 0.897 on BGCE. Error bars

represent ± S.E. of six replicate runs. ..................................................................................... 32

Figure 9. BGCE SEM images A) is polished and electrochemically pretreated in 0.1 M

phosphate buffer, B) BGCE after 10 CV cycles of 1 mM HS-, in 0.1 M potassium phosphate

buffer, chevrons indicate some product build-up from sulfide electrochemical oxidation. ... 33

Figure 10. EDX spectra of BGCE pretreated by running CV scans of the BGCE in phosphate

buffer from -200 to +800 mV at 100 mV s1-; where C = carbon. ........................................... 34

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Figure 11. EDX spectra of BGCE after 10 CV cycles (-1600 to +1600 mV) in 1 mM HS-

solution made in 0.1 M phosphate buffer. Where C = carbon and O = oxygen. ................... 34

Figure 12. SEM of BGCE obtained after scanning 10 CV cycles (-1600 to +1600 mV) in a 1

mM HS- solution prepared in 0.1 M phosphate buffer. Chevron indicates a deposition onto

electrode surface. .................................................................................................................... 35

Figure 13. SEM images of BGCE (A) and MGCE (B), showing the surface characterizations

between the BGCE and MGCE. ............................................................................................. 36

Figure 14. The fifth cycle of a CV output of 0.36 mM concentration of HS- (solid line) at

MGCE, and the supporting electrolyte solution (dashed line) at a scan rate of 100 mV s1-. .. 38

Figure 15. SEM image of MGCE showing foreign deposits on the film modification after 10

CV cycles of 1 mM HS- at 100 mV s1- (A) and MGCE showing 10 CV cycles in the

supporting electrolyte solution at 100 mV s1- (B). .................................................................. 38

Figure 16: EDX spectra of MGCE after 10 CV cycles (-1600 to +1600 mV) of 1 mM HS-

solution made in 0.1 M phosphate buffer. Where C, O, S and Cu represent carbon, oxygen,

sulfur and copper respectively. ............................................................................................... 39

Figure 17: EDX spectra of MGCE after 10 CV (-1600 to +1600 mV) cycles in supporting

electrolyte solution (0.1 M phosphate buffer), where C = carbon. ......................................... 40

Figure 18. SEM image of MGCE (left), showing sites of interest for EDX along with

supporting EDX data expressed as weight % (right) for V2O5 modifier. ............................... 40

Figure 19. SEM image of MGCE after 10 CV cycles in 1 mM HS- solution, showing sites of

deposits on surface (left) and EDX data expressed as weight % (right) for MGCE after 10 CV

cycles in 1 mM HS- solution. .................................................................................................. 41

Figure 20. CV analysis of supporting electrolyte solution on MGCE from cycles 1 to 10, scan

rate 100 mV s1-; the numbers indicate the scan number of the CV cycle. .............................. 43

Figure 21. CV graph at MGCE of HS- solution (0.44 mM); scan rate of 100 mV s1- showing

cycles 2, 5, and 10 as well as the blank (dashed line). ............................................................ 44

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Figure 22. Peak current values (Ipa1) µA at CV cycles 1 to 10 for 0.36 mM HS- (♦), compared

to CV cycle numbers 1 to 10, and compared to the supporting electrolyte solution (▲) on

MGCE. .................................................................................................................................... 45

Figure 23. Peak current values (Ipa2) µA at CV cycles 1 to 10 for 0.36 mM HS- (♦), compared

to CV cycle numbers 1 to 10, and compared to the supporting electrolyte solution (▲) on

MGCE. .................................................................................................................................... 45

Figure 24. Peak current values (Ipc) µA at CV cycles 1 to 10 for 0.36 mM HS- (♦), compared

to CV cycle numbers 1 to 10, and compared to the supporting electrolyte solution (▲) on

MGCE. .................................................................................................................................... 46

Figure 25. CV analysis increasing towards more positive potentials Ipa (µA) for various,

concentrations of HS- at Epa = +200 mV on MGCE after blank corrections. Error bars

represent ± S.E. of six replicate runs (R2 = 0.937). ................................................................ 48

Figure 26. CV analysis increasing towards more negative potentials Ipc (µA) for various

concentrations of HS- at Epc = -570 mV on MGCE after blank corrections. Error bars

represent ± S.E. of six replicate runs (R2 = 0.907). ................................................................ 49

Figure 27. Pictures of the synthesized Na-polysulfides: (A) Na2S2; (B) Na2S3; (C) Na2S4 .... 59

Figure 28. Purge and trap system set-up for the reduction of aqueous Na-polysulfides (Sn2-)

and thiosulfate (S2O32-): flask (A) (Sn

2- or S2O32-); flasks B and C (traps for reduced sulfide).

................................................................................................................................................. 61

Figure 29. Calibration curve of sulfide obtained using the DEPD method (R2 = 0.999). ...... 64

Figure 30. CV graph at a potential scan rate of 100 mV s-1 for S22- [1.39 mM], S3

2- [1.31 mM]

and S42- [1.22 mM] at BGCE. With supporting electrolyte solution as the blank (dashed

line). ........................................................................................................................................ 72

Figure 31. CV graph of the comparison of simple sulfide ion with di-, tri- and tetrasulfide

ions at a potential scan rate of 100 mV s-1 for HS- [1 mM], S22- [1.39 mM], S3

2- [1.31 mM]

and S42- [1.22 mM] at BGCE. With the supporting electrolyte solution as the blank (dashed

line). ........................................................................................................................................ 74

xiv

Figure 32. CV graph of the comparison of quantitative amounts of S2O32- at a potential scan

rate of 100 mV s-1 for S2O32- [1.06 mM], and 2 x the amount of S2O3

2- [2.12 mM] at BGCE.

With the supporting electrolyte solution as the blank (dashed line). ...................................... 75

Figure 33: CV graph of the comparison of polysulfide ions with thiosulfate ion at a scan rate

of 100 mV s-1 for S2O32- [1.06 mM] (dotted line), S2

2- [1.39 mM], S3

2- [1.31 mM] and S42-

[1.22 mM] at the BGCE. With supporting electrolyte solution as the blank (dashed line). .. 76

Figure 34. CV graph of the comparison of thiosulfate ion, sulfide ion and a combination of

thiosulfate ion and sulfide ion.at a potential scan rate of 100 mV s-1 for S2O32- [1.06 mM],

sulfide [1 mM], and mixture of HS- + S2O32- at BGCE. With the supporting electrolyte

solution as the blank (dashed line). ......................................................................................... 77

Figure 35. CV graph of the comparison of thiosulfate ion, trisulfide ion and a combination of

thiosulfate ion and trisulfide ion at a potential scan rate of 100 mV s-1 for S32- [1.31 mM] and

a mixture of S32-

[1.31 mM] and S2O32- [1.06 mM], and S2O3

2- [1.06 mM] at the BGCE. With

supporting electrolyte solution as the blank (dashed line). ..................................................... 78

Figure 36. CV graph at potential scan rate at 100 mV s1- at the MGCE for S22-, S3

2-, S42-,

S2O32-, and HS- (dotted line). With the blank as the supporting electrolyte solution (dashed

line). ........................................................................................................................................ 80

xv

List of Abbreviations and Symbols

Aux Auxillary Electrode

BASi Bioanalytical Systems Inc.

BGCE Bare Glassy Carbon Electrode

cm Centimeter

CTFE Chlorotrifluoroethylene

CDR Colour Developing Reagent

Cu Copper

DEPD N,N-Diethyl-p-phenylenediamine

EC Electrochemical Cell

GCE Glassy Carbon Electrode

g Grams

HDPE High-density polyethylene

Fe Iron

FeS Iron Sulfide

PbS Lead Sulfide

HgS Mercury Sulfide

μ Micro

μA Micro Amps

μM Micromolar

mM Millimolar

mL Milliters

mV Millivolts

MOE Ministry of Environment

MnS Manganese Sulfide

MGCE Modified Glassy Carbon Electrode

M Molar

N2 Nitrogen Gas

N Normal

I Peak Current

Ipa Peak Current Anodic (oxidation) μA

Ipc Peak Current Cathodic (reduction) μA

Epa Peak Potential Anodic (oxidation) mV

Epc Peak Potential Cathodic (reduction) mV

s1- Per second

Pt Platinum

MΩcm Resistivity measure unit MegOhms

centimeter

ν Scan rate (mV s1-)

Ag/AgCl Silver/Silver chloride

Ag2S Silver Sulfide

S Sulfur

V Volts

Zn Zinc

ZnS Zinc Sulfide

1

Overview and Research Objectives

Sulfur in the Environment

Interactions between the abiotic factors of the environment and the living organisms of

the biosphere are accompanied by a continuous cycling of matter in nature. Different

species of living organisms assimilate substances needed for their growth and to support

life. Living and non-living organisms emit by-products of metabolism, complex

minerals, and organic compounds of chemical elements in the form of non-assimilated

food or dead biomasses into the environment. This evolution to the biosphere formed a

stable connection of global biogeochemical cycles. With well-known and most important

biogeochemical cycles being: carbon, nitrogen, oxygen, phosphorous, sulfur and water.

Land and water ecosystems play an import role in the dynamics of biogeochemical

cycles, which have been agitated due to anthropogenic processes. The impact from these

anthropogenic processes has caused unpredicted changes to climate, increased

greenhouse gas production, decreases in biodiversity, progressive desertification, as well

as other factors (Krapivin and Varatsos, 2008; Ciais et al., 2014). Biogeochemical cycles

naturally aim at employing an equilibrium state in order to balance the cycling of the

elements between land and surface compartments for example carbon, nitrogen,

phosphorous, and sulfur. Impact of human activity to these biogeochemical cycles sees

most of the resources taken from nature. Those resources that were once taken from

nature are returned as waste products, which is more often than not poisonous or

unsustainable. These anthropogenic practices, thus, tip the equilibrium scales that should

exist between both the biosphere and humankind.

The nitrogen, sulfur and carbon biogeochemical cycles play important roles in the

regulation of many biological, chemical and geochemical processes. Carbon is an

essential element to life forms so too is nitrogen, phosphorus and sulfur. For example,

three of the essential Amino Acids (for humans) contain sulfur –these are cysteine,

cystine, and methionine (Canfield & Raiswell, 1999; Sievert et al., 2007). Amino acids

are proteins essential for functional metabolisms in the body. Cystine is formed by the

2

oxidation of two cysteine molecules. Cystine is found in the skeleton, hair, nails,

connective tissues, and is required to form glutathione. Other well-known sulfur

containing compounds in biological systems are the tripeptide glutathione, and many

important (protein) enzymes, coenzymes, vitamins, and hormones. Microorganisms can

use inorganic forms of sulfur compounds, such as sulfate, to process energy referred to as

assimilation (Andreae, 1990; Canfield & Raiswell, 1999; Sievert et al., 2007; Tang et al.,

2009; Knöller & Schubert, 2010). Due to an ever increasing impact from anthropogenic

activities such as the production and combustion of fossil fuels, this includes peat, coal,

oil and natural gas. These activities have increased the availability of sulfur compounds

in natural and managed ecosystems.

Numerous chemical and biological processes contribute to various transformations of

sulfur from one form to another through oxidation-reduction reactions. Sulfur based

compounds can occur in a variety of oxidation states, with -2 (sulfide), 0 (elemental

sulfur) and +6 (sulfate) being the most common (Kuhn et al., 1983; Luther III, 1985;

Andreae, 1990; Kariuki et al., 2001; Keller-Lehmann et al., 2006; Sievert et al., 2007;

Tang et al., 2009). Of the sulfur-based compounds, the more stable form sulfate can

function as an electron acceptor in metabolic pathways to be utilized by microorganisms,

and can be converted to sulfide (Sievert et al., 2007; Tang et al., 2009; Fazzini et al.,

2013). On the other hand, reduced sulfur compounds, like sulfides, can serve as electron

donors converting those compounds to elemental sulfur and sulfate (Friedrich et al.,

2001; Rohwerder and Sand, 2007; Tang et al., 2009). Some microorganisms that utilize

this type of conversion are sulfur reducing bacteria. Sulfur reducing bacteria represents a

diverse group of anaerobes and aerobes thriving in both oxic and anoxic environments

The biochemical reactions involved in the oxidation and reduction of sulfate to sulfide

and vice versa are elaborate. Different groups of sulfur reducing bacteria utilize these

pathways in different ways, depending on their environmental conditions, to suit their

needs. According to authors Friedrich et al. (2001), Friedrich et al. (2005), Rohwerder

and Sand (2007), Mohapatra et al. (2008), Ghosh and Dam (2009), and Tang et al.,

(2009), the metabolic processes of sulfur reducing bacteria have been extensively

reviewed. In brief, sulfate reduction occurs though two pathways, assimilatory and

3

dissimilatory. Both pathways require activation of sulfate by adenosine triphosphate

(ATP). The attachment of sulfate to ATP results in the formation of adenosine

phosphosulfate (APS) then is catalyzed by the enzyme ATP sulphurylase. The sulfate

portion of the APS is further reduced to sulfite by the enzyme APS reductase. In the

assimilatory pathway sulfate is used to generate reduced sulfur compounds, using organic

compounds, for the synthesis of amino acids and proteins. The assimilatory pathway

does not excrete sulfide directly. Rather, the developed sulfide is incorporated into

organic sulfur compounds, such as dimethyl sulfide and dimethyl sulfoxide. In the

dissimilatory pathway the reduction of sulfate or sulfur is converted to inorganic sulfide,

such as hydrogen sulfide, by available anaerobic sulfate and the resulting sulfide

produced is thus excreted into the surrounding environment.

Some sulfur-based compounds which are of particular interest include: hydrogen sulfide

(H2S), hydrosulfide (HS-), sulfide (S2-), elemental sulfur (S0), thiols (R-SH), sulfate

(SO42-), sulfite (SO3

2-), polythionates (S2On2-) and polysulfides (Sn

2-) (Ciglenečki &

Ćosovíc, 1997). In anoxic environments, there may be an accumulation of reduced sulfur

species such as, hydrogen sulfide, metal sulfides, polysulfides, thiosulfates and elemental

sulfur (Brouwer and Murphy, 1995). Sulfur species formed depend largely on

temperature, pH, ionic strength, the source of the sulfur species and oxidizing agent

present (Boulegue et al., 1982; Koh, 1990; Petre and Larachi, 2006; Kaasalainen and

Stefánsson, 2011; Pan et al., 2013). As outlined by Holmer and Storkholm (2001) and

Mohapatra et al. (2009), the oxidation metabolism of inorganic sulfur compounds starts

with sulfide oxidase which catalyzes the oxidation of hydrogen sulfide to elemental

sulfur. Sulfur oxidizing enzymes oxidize elemental sulfur to sulfite. Further oxidation of

sulfite to sulfate is catalyzed by sulfite oxidase. During the oxidation of sulfide to the

various sulfur compounds, sulfite and elemental sulfur can react chemically together to

form thiosulfate. Inorganic forms of sulfur-based compounds can be found in oxic,

anoxic, fresh and marine aquatic systems (Al-Farawati & van den Berg, 1999; Dutta et

al., 2010).

Sulfur based compounds are released naturally from living organisms, some examples

include metabolic excretions from bacteria, decomposition of organic plant matter under

4

anaerobic conditions, and discharge from agriculture livestock. Several anthropogenic

processes such as wastewaters, petroleum refineries, burning of fossil fuels, and paper

mills also produce sulfur-based compounds. Releases of high concentrations of reduced

sulfur compounds, like sulfide, can seriously disrupt ecosystems by producing elevated

levels of sulfur-based contamination downstream (Witter and Jones, 1997; Ateya et al.,

2007). Brouwer and Murphy (1995), reported that hydrogen sulfide in concentrations

down to 0.015 mg L1- cause chronic toxicity to aquatic organisms, with levels of 0.5 mg

L1- exhibiting acute toxicity to rainbow trout, and salmon in a Spanish fish farm (Ortiz et

al., 1993). Hydrogen sulfide may be produced from the decomposition of organic matter

underground, such as decaying plant material, or by chemical reduction of sulfate by

sulfate-reducing bacteria. Hydrogen sulfide can be found in deep or shallow wells. It is

often present in areas underlain by shale or sandstone, near coal or peat deposits, and near

oil fields. Hydrogen sulfide can also occur naturally in groundwater systems. Some of

the sources of hydrogen sulfide in ground water are located near farm areas and swamps

in Ontario, as well as areas in Nova Scotia where there are sulfide bearing mineral

deposits and out west through Saskatchewan in the gas fields (Edwards et al., 2011). The

level of contamination from the hydrogen sulfide, in the ground water, depends on the

level of the water table. Dissolved sulfide in sewage and paper mill waste water effluent

were reported to be released at concentrations of 47 mg L1- and 51 mg L1- respectively

(Dutta et al., 2008 and 2010). Another example of discharge into the environment is the

effluent from the un-hairing process that occurs during tanning of animal hides.

Discharges from tanning process effluents have reported sulfide concentrations in the

range of 700 – 2000 mg L1- (Font et al., 1996). The resulting untreated waste water, from

tanneries, has also been reported to have dissolved sulfide concentrations of 20 mg L1-

(Font et al., 1996).

Development of reliable methods for the detection of various sulfur-based compounds

has become an important goal for analytical and environmental chemists (Lawrence et

al., 2000; O’Reilly et al., 2001; Khudaish & Al-Hinai 2006; Dutta et al., 2009; Titova et

al., 2009; Paim and Straditto 2010). The importance of the detection of sulfur-based

compounds has increased due to the increased amount of occupational exposure to

hydrogen sulfide. Sulfur content in crude oils is 0.3 to 0.8 wt. % and the hydrogen

5

sulfide content in natural gas ranges from 0.01 to 30 wt. % (Lawrence et al., 2000(a)).

Further increases to the extraction of oil and natural gas products to meet future supply

and demands will cause further increases to hydrogen sulfide concentrations in both oil

and natural gas. Clinical cases involving sulfide poisoning can range from 0.03 to 1 mg

L1-; with lethal doses ranging from 0.3 to 3 mg L1- (Lawrence et al., 2000(a); Ardelean et

al., 2014). As a weak acid, hydrogen sulfide has corrosive properties that can lead to

damage to infrastructure in municipal and city sewer systems. Zhang et al. (2008)

reported that at levels of 0.1 – 0.5 mg L1- of sulfide content in waste water will have

minor effects to infrastructure. However, at sulfide concentrations of > 2.0 mg L1-, severe

corrosion problems are prominent. Sulfide present in waste waters from sewer systems,

combined with the enclosed environment related to sewer systems, the exposure to levels

of sulfide that induce corrosion to infrastructure pose a risk of toxicity to sewer system

workers. Several authors such as Lawrence et al. (2000a), Lawrence et al., (2000b),

Lawrence et al. (2004), Ateya et al. (2007), Lawrence et al. (2007), Titova et al. (2009),

Edwards et al. (2011), Pikaar et al. (2011), and Hu & Mutus (2013), have reported some

clinical cases that have arisen from the exposure to sulfide. These authors have also

reported toxic and chronic effects with continued exposure. Exposures to sulfide, even at

concentrations as low as 0.01 mg L1-, have been reported to have neurotoxic effects as

well as causing chronic side effects to the skin, eyes, circulatory system, digestive and

respiratory tracts (Lawrence et al., 2000(a); Lawrence et al., 2000(b); Khudaish & Al-

Hinai 2006; Hughes et al., 2009; Titova et al., 2009; Hu & Mutus, 2013). The risks

associated with sulfide, not only to environments that are close to wastewater effluents

but also to surrounding occupational hazards mean that, continued investigation is

essential to the understanding of the environmental processes undergone by sulfur-based

compounds.

Current Detection Methods for Sulfur-Based Compounds

Sulfide analyses are well represented across a variety of analytical methods. The sulfide

anion can be rather versatile, in that, the anion can form other intermediate species of

sulfur based compounds to which they can be determined by an assortment of analytical

methods. A few of those intermediate sulfur based compounds that can form from the

6

sulfide anion are thiosulfate, sulfite, sulfate, and acid volatile sulfides. Some major

analytical methods for the detection of the sulfide anion, and other sulfur-species include

titration (Pawlak & Pawlak, 1999; Ciesielski & Zakrzewski, 2006), spectroscopy (Silva et

al., 2001; Kariuki et al., 2008), liquid chromatography (O’Reilly et al., 2001; Kamyshny

et al., 2006), ion chromatography (Miura et al., 2005; Stefansson et al., 2007), gas

chromatography (Wardencki, 1998; Kristiana et al., 2010) and electrochemistry (García-

Calzada et al., 1999; Rozan et al., 2000b; Giovanelli et al., 2003; Cheng et al., 2005;

Manova et al., 2007; Titova et al., 2009; Piam & Stradiotto, 2010; Manan et al., 2011;

Huang et al., 2012). For the detection of simple sulfide, the methodologies vary over a

broad instrumental platform, with each method posing its own set of benefits and

limitations.

In a review published by Lawrence et al. (2000a), the authors have presented methods for

analysis of the sulfur-species and their respective limitations. Iodometric titration is

reported to be reliable for the standardization of pure, laboratory-generated sulfide

samples for calibration since the overall concentration of dissolved total sulfide is ≥ 1.0

mg L1-. The authors suggest that titrimetric methods have significant limitations in terms

of sensitivity and selectivity when used to analyze environmental samples. In the

presence of interfering ions like thiosulfate, sulfite, and organic compounds, both

dissolved and solid iodometric titration methods can be hindered (Clesceri et al., 1998).

These interferences cause a reduction in sensitivity towards the overall detection of the

sulfide ion in environmental samples, due to the formation of side reactions whose

products have detection limits that range from 0.2 to 4.0 mg L1- (Lawrence et al., 2000a).

Spectroscopic Determinations of Sulfur Based Compounds

There are a variety of spectroscopic detection methods known for analyzing sulfides.

The most common analysis quantifies the concentration of sulfide with the Methylene

Blue complex, detected via the UV/visible spectrum. This particular reaction involves

aqueous samples of sulfide to react with N,N-dimethyl-p-phenylenediamine, through

oxidative coupling in the presence of ferric ions under acidic conditions, which forms a

pentacyclic phenothiazine dye, blue coloration referred to as methylene blue that has an

7

absorbance maximum at 670 nm (see Scheme 1) (Cline, 1969; Spaziani et al., 1997;

Lawrence et al., 2000a; Lawrence et al., 2000c; Silva et al., 2001; Kariuki et al., 2008).

The methylene blue method is reported to be consistently sensitive, selective and simple.

For example, wastewater samples analyzed for sulfide using methylene blue have upper

and lower detection limits of 3.2 x 10-2 and1.0 x 10-4 mg L1- respectively (Lawrence et

al., 2000a). Although, the Methylene Blue approach has been popular, the sample

turbidity and the formation of H2S when S2- is in contact with oxygen tend to limit the

accurate evaluation of sulfide in aqueous media (Lawrence et al., 2000c). Also, reducing

agents cause interference with the methylene blue reaction, for example, thiosulfate in

concentrations of 10 mg L1- may prevent or impair the formation of the blue coloration

(Clesceri et al., 1998). When analyzing natural water samples with a complex matrix,

combined with low sulfide concentrations, interferences from dissolved organic material

can occur. The absorbance shoulder of dissolved organic material can spread over the

wavelength of 668 nm, could falsely increase the methylene blue signal (Tang &

Santschi, 2000). Other chromogens that have been mentioned to be analogous to

methylene blue, towards the detection of sulfide, are: resazurin, aszurea, thionine and

toluidine blue (Lawrence et al., 2000a). Other spectroscopic methods, use the

fluorimetric protocols with 2,7-dichlorofluorescien. This method suffers from problems

of poor selectivity in the presence of thiol (RSH) compounds, since the latter reacts in a

similar way as the sulfide anion does (Lawrence et al., 2000a).

Scheme1

Chromatographic Determinations of Sulfur Based Compounds

Chromatographic separations, like those of capillary electrophoresis (Font et al., 1996;

Petre & Larachi, 2006), gas (Wardencki, 1998; Kristiana et al., 2010), liquid (Tang &

Santschi, 2000; Pan et al., 2011), and ion chromatography (Casella et al., 2000; Ohira &

Toda, 2006) hold advantages for complex media. Chromatographic methods can offer

8

better resolution and separation from other sulfur based compounds present in the

sample. For capillary electrophoresis, earlier work by Font et al. (1996), observed good

linearity for sulfide in concentration ranges from 0.5 to 10 mg L1- at pH of 10.5, using

internal and external calibration standards. The authors reported a limit of detection for

sulfide of 0.2 mg L1-. They also observed that thiosulfate, nitrate, and nitrite could be

resolved using this technique and these compounds did not cause any interferences with

the resolution of the sulfide anion. However, in work performed by Petre and Larachi

(2006), the authors used capillary electrophoresis to resolve inorganic sulfur bearing

species. These species included sulfate, sulfite, polythionates, polysulfides, and sulfide

anions. Over a pH range of 8.2 to 12.2 that the separations were tested, a pH of 9.5 was

found suitable for the separation of thiosulfate, sulfate, sulfide, tetrathionate and sulfite at

a sulfide concentration of 6 mg L1- and the other sulfur species at concentrations of 10 mg

L1-. They also observed that polysulfide distribution depended largely on pH. In addition,

they noted that the polysulfide ions were unstable over the pH range of 8.2 to 12.2. A

complete quantitative determination of these ions was not possible with this method.

Separations and quantification of sulfide is possible with capillary electrophoresis, for

examining environmental effluents from leather, paper processing and mining water with

detection ranges from 5.0 x 10-4 mg L1- to 20 mg L-1 (Lawrence et al., 2000a). However,

some sulfur containing compounds like polysulfides appear more challenging to separate

and quantify due to their rapid dissociation with changes in pH of the solution being

analyzed. With liquid and gas chromatography separation analyses, these techniques are

susceptible to common interferences such as hydroxyl ions and metal ions that may be

present in the sample. Also, many liquid and gas chromatography separations employ

much of the emphasis on the organic forms of sulfide, such as dimethyl sulfide, rather

than the inorganic forms of sulfide (Wardencki, 1998; Kristiana et al., 2010). However,

there has been success in separating sulfide, sulfite, sulfate and thiosulfate anions using

ion-pair chromatography. This technique requires the conversion of sulfide to

thiocyanate and sulfite to sulfate in order to be analyzed. In concentrations of up to 10 -

mM, the common anions listed above, did not interfere with the determination of sulfur

anions. The recoveries of these anions from hot-spring waters were reported to be 99.5 to

101.3 % respectively (Miura & Kawaoi, 2000; Miura et al., 2005). Pre-column

9

derivatizations based on methylene blue and florescent labeled monobromoimane have

been utilized towards the detection of sulfide anion through liquid chromatography. Ion

chromatography, using gel sorbents or solid lead (II) based chromate columns suffer

impairment from hydroxyl ions as well as metal ions that could be found in complex

sample matrices, such as, calcium, magnesium and iron (Lawrence et al., 2000a). The

technique also lacks selectivity with regards to other redox species of sulfur anions.

Electrochemical Determinations of Sulfur Based Compounds

Electrochemical detection methods have become popular in the analysis of sulfur

compounds (Ciglenečki & Ćosović, 1997; García-Calzada et al., 1999; Lawrence et al.,

2000b; Chadwell et al., 2001; Lawrence et al., 2002; Diligin et al., 2012; Huang et al.,

2012; Hu & Mutus, 2013). The responses of these methods are based on the oxidation or

reduction of an electroactive species, like those of sulfur-based species. Unlike some

other analytical methods, the electrochemical based methods can provide information

about reaction kinetics, chemical behavior of the electroactive species in solution, and

information regarding possible adsorbed products to the surface of the working electrode

(Kissinger and Heineman, 1996; Scholz, 2010; Zanello et al., 2012). The

electrochemical cell contains a reference (saturated calomel (SCE) or silver/silver

chloride (Ag/AgCl)), an auxiliary, and a working electrode. Working electrodes like

those of inert metals such as, gold, silver, and platinum offer favorable electron transfer

kinetics and a wide potential range (Kissinger and Heineman, 1996; Scholz, 2010;

Zanello et al., 2012; Li and Miao, 2013). However, the cathodic potential is restricted

due to the low hydrogen overvoltage, which forms surface oxides and hydrogen layers

leading to high background currents affecting the kinetics of the electroactive species at

the working electrode. Mercury electrodes offer high hydrogen overvoltage and an

extended cathodic potential window (Kissinger and Heineman, 1996). The most

commonly used mercury electrode is the hanging mercury drop electrode, which possess

a highly reproducible, renewable, smooth surface (Kissinger and Heineman, 1996;

Umiker et al., 2002; Scholz, 2010; Li and Miao, 2013). The hanging mercury drop

electrode does not require cleaning before an experiment since it has a self-renewing

surface, but the limited anodic range and toxicity around handling and disposal has

10

researchers looking towards alternative working electrode materials that have comparable

chemical function to that of mercury. (Li and Miao, 2013). Carbon based electrodes have

soft surface properties, in other words, the surface can be easily polished compared to

metal electrodes like silver and platinum. Carbon surfaces can be easily renewed for

electron exchange during polishing and cleaning procedures (Kissinger and Heineman,

1996; Scholz, 2010; Zanello et al., 2012). Carbon electrodes have a complex underlying

microstructure, broad potential window, low background current, which allows for the

formation of a wider variety of surface bonds and functional groups (McCreery, 2008;

Scholz, 2010; Zanello et al., 2012). The cost of carbon based electrodes is quite low

compared to that of traditional metal electrodes, which makes them a favorable material

for electrochemical analyses (Li and Miao, 2013). Common carbon based electrodes

include pyrolytic graphite, glassy carbon, carbon paste, carbon-fiber, nanotubes, and

carbon composite electrodes (Kissinger and Heineman, 1996; Scholz, 2010; Zanello et

al., 2012; Li and Miao, 2013).

Electrochemical techniques include polarography (Umiker et al., 2002), cathodic

stripping voltammetry (Ciglenečki & Ćosović, 1997) anodic stripping voltammetry

(Huang et al., 2012), and cyclic voltammetry (Lawrence et al., 2004; Lawrence et al.,

2007). Electrochemical methods can provide fast, sensitive detection while probing

reaction mechanisms of electroactive species. Differential pulse polarography has been

used for many years as a direct method for sulfur compound speciation (Rozan et al.,

2000). When compared to direct current and normal pulse polarography, differential

pulse polarography better resolves multi component systems and exhibits closely spaced

half-wave potentials especially when both oxidized and reduced species of the redox

couple are present (Umiker et al., 2002). When utilizing the hanging mercury drop

electrode for sulfide analysis, the cathodic reaction at the mercury electrode (scheme 2)

shows that sulfide oxidizes the mercury through the formation of mercury sulfide.

Similar sulfur based species like those of sulfite, thiols, polythionates, and polysulfides

have a similar reaction with mercury (Rozan et al., 2000; Umiker et al., 2002). Umiker

et al (2002) conducted differential pulse polarography to quantify sulfur species in soil

samples. The authors reported that the concentrations for sulfite, thiosulfate, cysteine and

sulfide were 0.580, 3.16, 1.58, and 0.296 µM, respectively.

11

𝐻𝑆− + 𝐻𝑔 ↔ 𝐻𝑔𝑆 + 𝐻+ + 2𝑒−

Scheme 2

Cathodic striping voltammetry can also be used to determine sulfur species in the

presence of sulfide. In a study done by Ciglenečki & Ćosović (1997), the authors

measured sulfur based compounds in anoxic sea water using the mechanism outlined in

scheme 2, but utilizing the reverse reaction when the potentials are moved to more

negative regions and the resulting current of mercury (II) to elemental mercury is

measured. The authors report that sulfide and elemental sulfur concentrations were 42

µM, with thiosulfate concentrations being 8 µM in sea water samples. They also reported

that sulfide and thiosulfate had good linearity in the range of 0.01 to 1 µM and 1 to 100

µM respectively. Huang et al (2012) investigated the use of bismuth film modified

electrodes towards the detection of sulfide in water samples using anodic stripping

voltammetry. The use of anodic stripping voltammetry is based on the selective reaction

between cadmium (Cd2+) and sulfide to form cadmium sulfide precipitate and then

stripping the Cd2+ from the formed sulfide deposit. The process leads to a current which

is proportional to the Cd2+ concertation in the sample. Bismuth, as an electrode modifier,

has been reported towards electrochemical detection of metal ions (Wang, 2005). Huang

and others go on to report that bismuth films have the same advantages of mercury, but

with less toxicity. In their study they compared the anodic stripping voltammetric

method to the classic methylene blue reaction. The initial analysis of the water samples

yielded no detectable limits for either method, but upon spiking the water samples with

sulfide, recoveries of 101 ± 0.583 µM L1- for the anodic stripping method resulted. The

methylene blue reaction showed no detectable levels of sulfide, even after spiking except

in two of the ten water samples they analyzed. Good linearity and reproducibility was

also reported towards the detection of sulfide, as well as little to no matrix interferences.

Lawrence et al. (2007) used cyclic voltammetry and various carbon based substrates,

more specifically edge-plane pyrolytic graphite, to examine the direct oxidation of sulfide

in a river water sample. The authors compared the techniques of cyclic and square wave

voltammetry, reporting a linear range of sulfide being detected 5 – 6 µM and 10 – 60 µM

12

respectively. Lawrence et al. (2007) chose to use cyclic voltammetry due to its lower

linear range for detectable sulfides. These authors further report recoveries of sulfide

spiked river water of 104 %. The authors also noted that the approach they used would

suffer when determining sulfide in matrices that contain oxidizable electroactive species.

Another group from Lawrence et al. (2004) used carbon nanotubes to modify glassy

carbon electrodes for the detection of sulfide. Although these authors did not apply their

modifications to environmental samples, they had good linear range for sulfide using bare

glassy carbon and modified glassy carbon with carbon nanotube electrodes, in the range

of 12.5 – 50 µM and 1.25 – 112.5 µM respectively. Many of the techniques summarized

here show adequate selectivity with some sensitivity, and vice versa towards the

determination of sulfide in various matrices, with sample handling and clean up measures

being required in most cases. To increase the comprehensive assessment of detecting

sulfide and other sulfur based compounds it appears that combining the various analytical

techniques, could perhaps enhance the selectivity and sensitivity of detection of sulfur

based compounds.

This work will attempt to examine if electrochemical analyses, such as cyclic

voltammetry, with the use of bare and modified glassy carbon electrodes will be efficient

towards possible quantification of lab-generated sulfide samples. Chemically modified

electrodes function to immobilize molecules with specific functions on the electrode

surface by either physical or chemical means (Kissinger and Heineman, 1996; Li &

Miao, 2013). The use of glassy carbon electrodes is favored more over most traditional

metal electrodes like gold, silver, platinum (Kapusta et al., 1983; Mohtadi et al., 2005) or

mercury (Umiker et al., 2002). This study will apply cyclic voltammetry using

chemically modified and unmodified glassy carbon electrodes against lab-generated

sulfide, along with polysulfides, namely, disulfide, trisulfide and tetrasulfide. Also, a

representative of the polythionates, such as thiosulfate, will be examined to explore

whether these compounds can be distinguished individually and as part of a mixture with

simple sulfide. The use of cyclic voltammetry will also be compared to some of the

classical methodological approaches, for example, iodometric titrations and ethylene

blue, analogous to methylene blue, for its efficiency in the determination of sulfides,

polysulfides and polythionates. Lastly, this work will be applied to a natural water

13

sample in order to test cyclic voltammetry, iodometric titration and ethylene blue

reactions against environmentally relevant concentrations of sulfur-based compounds

such as the ones mentioned above.

14

Chapter 1

Sulfide (HS-, H2S and S2-)

1.1 INTRODUCTION

Electrochemical techniques measure the oxidation and the reduction signals of the sulfur

compounds at the surface of a working electrode. Modifying the working electrode

surface with a metal oxide, such as vanadium oxide (V2O5), has been shown to enhance

the detection of inorganic sulfide species like those of HS-, S2-, and H2S (Park et al.,

1998; Park et al., 2002; D’Elia et al., 2004; Khudaish and Al-Hinai, 2006). Khudaish

and Al-Hinai (2006) have outlined the electrochemical deposition of vanadium oxide

films on glassy carbon electrodes according to a procedure described by D’Elia et al.

(2004). Vanadium compounds are not soluble in water and lead to a slow

electrochemical response. Once they are deposited on the electrode surface the catalytic

activity, towards the complexation with the sulfide ion in solution, are increased and can

be used in aqueous solutions (Li et al., 1996; D’Elia et al., 2004; Khudauish and Al-

Haini, 2006; Salimi et al., 2006). The V2O5 film deposition was realized through the

electrochemical oxidation of the vanadyl species, VO2+, that was obtained by scanning

potentials of a solution containing VO2+ from 0 to + 2000 mV versus Ag/AgCl. The

suggested reaction which occurs during this process is shown in Scheme 3 (Khudaish and

Al-Hinai, 2006). Other metal oxide catalysts for the oxidation of hydrogen sulfide have

been explored, such as, bismuth-molybdenum oxide (Li and Cheng, 1966), iron-

antimonate and iron-tin (Li et al., 1997); aluminum oxide, titanium dioxide, and iron (III)

oxide (Park et al., 1998). The authors that indicated that these metal oxides can oxidize

hydrogen sulfide did not propose any mechanisms details for the oxidation process.

However, Li and Cheng (1997) have described a two-step mechanism towards the

oxidation of hydrogen sulfide to elemental sulfur using mixed oxides of bismuth and

molybdate. These authors describe that, first, hydrogen sulfide reacts with the oxygen of

the metal oxide catalyst to form elemental sulfur, and the metal oxide is then partially

reduced. Subsequently, the partially reduced metal oxide is re-oxidized by the oxygen

present in the reaction mixture (Scheme 3).

15

2𝑉𝑂2+ + 3𝐻2𝑂 → 𝑉2𝑂5 + 6𝐻+ + 2𝑒−

Scheme 3

Voltammetric measurements can provide insight into chemical mechanisms in regards to

reaction mechanisms, kinetics, electron transfer, reversibility or irreversibility of a

reaction, and the behavior of a species in solution (Batchelor-McAuley et al., 2015). The

cyclic voltammetric technique allows for fast and easy accumulation of data. This

technique employs the use of various substrates for a working electrode, with the most

common being metal based electrodes (gold, silver, platinum or mercury). With the

enhanced awareness of environmental protection and health concerns towards human

exposure, the applications of mercury electrodes ought to be reduced due to their high

toxicities (Huang et al., 2012). Carbon based electrodes such as glassy carbon, carbon

paste, carbon fiber, or carbon composite electrodes are more advantageous over

conventional metallic based electrodes. The working surfaces of glassy carbon electrodes

are readily recharged by mechanical polishing, and can be modified by applying

electrochemical pretreatments (Wang & Hutchins, 1985; Pocard et al., 1992; McCreery,

2008). This chapter outlines the electrochemical analysis of sulfide using cyclic

voltammetry, with the bare glassy carbon and also with the modified glassy carbon

working electrodes.

1.2 MATERIALS AND METHODS

1.2.1 Reagents and Solutions

A 0.1 M potassium phosphate buffer solution (pH = 10.3) was prepared and used to make

the sulfide stock solutions. This buffer solution also served as a supporting electrolyte

for all the electrochemical experiments that were carried out in this work. Once the

sulfide was dissolved in the buffer solution, the sulfide stock solution (1.2.1.1) was

standardized against 25 mM standardized iodine and thiosulfate solutions. The

standardization of the sulfide stock solution was required since the sulfide solution can

easily be oxidized when exposed to air. Therefore, to ensure the concentrations of sulfide

16

used during the electrochemical and colorimetric experiments were as accurate as

possible, iodometric titrations were conducted. The procedures about how the solutions

used for the iodometric titration and for the phosphate buffer preparation are summarized

in Appendix A. These include the preparation of 6 M hydrochloric acid, 25 mM sodium

thiosulfate, 2 mM potassium bi-iodate, 25 mM iodine and 2 % (w/v) starch solution.

1.2.1.1 Stock Sulfide Solution

Working in a Polymer Series 100, Baxter Glove Box, equipped with a dual purge

nitrogen flow control, as well as a humidity and temperature control capability (Terra

Universal Inc.), a sulfide stock solution of 106 ± 1 mM (n = 9) was prepared by weighing

and dissolving 0.158 g of anhydrous sodium sulfide with a purity of ≥ 90.0 % (Acros

Organics, Fisher Scientific) into 20-mL of pre-purged, 0.1 M phosphate buffer with a pH

of 10.3, contained in a 20 mL, 28 x 61 mm borosilicate glass scintillation vial with

polyethylene caps (Fisherbrand, Fisher Scientific). The solution was well sealed and

stored at a temperature of 4 ºC, until it was required for analysis.

1.2.1.2 Vanadium Tetroxide (V2O4) Solution

A 0.01 M of V2O4 was prepared in a 250-mL volumetric flask by dissolving 0.415 g of

V2O4 (99.9 %, SIGMA) in 1.0 M phosphoric acid (H3PO4) (85 %, Fisher Scientific). The

V2O4 solution is reported to have been used for the modification of a glassy carbon

electrode surface (D’Elia et al., 2004; Khudaish and Al-Hinai, 2006). These authors

report that they used the modified glassy carbon electrode for the electrochemical

detection of sulfide. The vanadyl species present in solution. According to formal

reduction potentials (Harris, 2010), is VO2+.

In the present study, the VO2+ solution was to be used for the modification of the glassy

carbon electrode surface for the electrochemical analysis of not only simple sulfide, but

also of other sulfur species. The process used for the electrochemical deposition of the

V2O5 film is described in section 1.2.4.4.

17

1.2.2 Standardization of 25 mM Thiosulfate and 25 mM Iodine

Solutions

1.2.2.1 Standardization of Thiosulfate Solution against Bi-Iodate Solution

Sodium thiosulfate (Na2S2O3) is a common secondary standard which therefore, requires

standardization through the use of a primary standard. A 25 mM S2O32-

solution,

prepared as described in Appendix A, was standardized with a solution of potassium bi-

iodate (KIO3), also prepared as outlined in Appendix A. In the standardization of S2O32-,

a weighed amount of KIO3 was reacted with excess potassium iodide (KI) under acidic

conditions. As shown in reaction R – 1.1 (Appendix B), the reaction releases iodine (I2)

which then reacts with the S2O32-

solution being standardized (Appendix B, R – 1.2). The

summary of how the standardization was carried out is as follows. A 20-mL aliquot of 2-

mM KIO3 solution, prepared as shown in Appendix A, was pipetted into a 150-mL

solution containing 2 g KI. The resulting solution was acidified with 2 drops of

concentrated sulfuric acid (ACS Grade, Caledon Labs) and then diluted with ultra-pure

water to a volume of 200-mL. The 200-mL solution was then titrated with 25 mM S2O32-

prepared as specified in Appendix A. A few drops of 2 % (w/v) starch solution prepared

as described in Appendix A, was used as the indicator of the end point for the titration.

The reactions used to calculate the concentration of the thiosulfate ion are outlined in

Appendix B.

1.2.2.2 Standardization of Iodine Solution against Standardized Thiosulfate Solution:

A 20-mL aliquot of the standardized S2O32- (25 mM) was added to a 250-mL Erlenmeyer

flask. Two drops of 2 % (w/v) starch solution was added to the S2O32-solution. The

resulting solution was titrated with iodine solution (Appendix A) to a blue-starch

complex. The balanced reaction between thiosulfate and iodine is shown in Appendix B,

reaction R – 1.3. The molarity of the iodine solution was calculated using equation Eq.

1.7 (Appendix B). This process was repeated for a total of three replicates.

18

1.2.3 Standardization of Sulfide Using Iodometric Titration

A 25-mL portion of the standardized iodine solution (Appendix A) was transferred into a

250-mL Erlenmeyer flask. The iodine solution was acidified by pipetting 2-mL of 6 M

hydrochloric acid (Appendix A), and was swirled to mix. A 0.5-mL aliquot of the sulfide

stock solution (section 1.2.1.1) was then pipetted into the iodine/acid mixture. To limit

exposure to air, the sulfide solution was discharged below the surface of the iodine/acid

mixture. As the reaction R – 1.4 Appendix B shows, the added sulfide solution reacted

with I2 which had been added in excess. The unreacted I2 was then back-titrated with

standardized S2O32- solution (Reaction R – 1.5, Appendix B). For this titration, a

standardized S2O32- solution (Appendix A) was titrated into the iodine solution mixture

until a pale straw colour was achieved, after which a few drops of 2 % (w/v) starch

solution (appendix A) was added. The addition of the starch solution led to the formation

of a blue starch complex coloration which indicated the presence of some unreacted I2.

Further titration with the S2O32- led to the disappearance of the blue coloration. This

marked the end point of the titration. The reactions R – 1.4 and R – 1.5 (Appendix B)

were used in the computation of the sulfide in solution.

1.2.4 Electrochemical Analysis of Sulfide

1.2.4.1 Equipment

All electrochemical measurements were conducted using the BASi work station with

Epsilon USB software (Bioanaytical System Inc.). The three-electrode cell consists of a

C3 glass cell vial with a dimension of 50 mm x 59 mm, that houses the supporting

electrolyte solutions, and electrodes used for analysis. A glassy carbon electrode with a

solvent resistant coating of chlorotrifluoroethylene with a working surface diameter of

3.0 mm (area = 0.017 cm2); was used as the working electrode. A coiled Pt wire, 23-cm

in length was used as the auxiliary electrode. An Ag/AgCl, filled with 3.0 M NaCl, was

used as the reference electrode. All electrodes were supplied through BASi, and were

assembled as shown in Figure 1.

19

A)

B)

Figure 1: Electrochemical cell setup for chemical analysis. (A) From left to right:

Reference electrode (Ag/AgCl) shaded grey, glassy carbon electrode (GCE) shaded black

and platinum auxiliary electrode (Pt) represented as light grey coil. (B) The BASi setup

for electrochemical analysis, following the same electrode order from left to right as is

the case in A.

1.2.4.2 Cleaning the Electrodes

Between experiments, the glassy carbon electrodes were cleaned using various grades of

diamond and alumina polish over their respective polishing pads. Surfaces were rinsed

well with ultra-pure water, and sonicated (FS20, 3.0 qt; Fisher Scientific) for 5 minutes to

ensure polishing particulates were removed. The Ag/AgCl reference electrode and the Pt

auxiliary wire were rinsed with ultra-pure water between experiments.

1.2.4.3 Electrochemical Pretreatment of Glassy Carbon Electrode

Using the electrochemical cell, described in section 1.2.4.1, all glassy carbon electrodes

were pretreated electrochemically using the 0.1 M phosphate buffer solution (Appendix

A) as the supporting electrolyte solution. To do this, CV was used to cycle the glassy

carbon electrode potential from -200 mV to +800 mV at 100 mV s-1 for 25 cycles. This

procedure, according to a number of authors, Nagaoka and Yoshino, 1986; Kamau, 1988;

Yang and Lin, 1994; McCreery and Cline, 1996 (Ch. 10); Dekanski et al., 2001; Kiema et

al., 2003; Zhao et al., 2008 and McCreery, 2008, helps to remove any residue on the

surface of an electrode that may have remained during the electrode polishing process.

20

1.2.4.4 Electrochemical Deposition of V2O5 Film for Modified Glassy Carbon

Electrode (MGCE):

The electrochemical deposition of V2O5 onto glassy carbon electrodes has been

recommended by Khudaish & Al-Hinai (2006) and D’Elia et al. (2004). The VO2+

solution (section 1.2.1.2) was added to the electrochemical cell so that there was 2 cm of

head space at the top of the cell. The pretreated working glassy carbon, reference and

auxiliary wire electrodes were then placed in the cell. Using CV, the electrochemical

deposition of the V2O5 film was coated onto the surface of the glassy carbon electrode by

cycling the electrode potential from 0 to +2000 mV, at a scan rate of 50 mV s-1, for 20

cycles. The MGCE was removed and rinsed gently with ultrapure water, and stored in

the supporting electrolyte solution (Appendix A) until needed. The MGCE’s were used

for the electrochemical experiments the same day that they were modified.

1.2.4.5 Electrochemical Analysis Using Cyclic Voltammetry (CV)

Electrochemical measurements of the sulfide solution (section 1.2.1.1) using CV were

conducted by cycling the electrode potential from -1600 mV to +1600 mV, at a scan rate

of 100 mV s-1 for 10 cycles. The peak potential (mV), height (µA), and area (µC) were

recorded for all 10 cycles. Six replicate runs were carried out for each experiment over

bare glassy carbon electrodes (BGCE) and MGCE. All the potential measurements were

carried out using the set-up described in section 1.2.4.1, and the electrochemical

parameters outlined above were also used for the experiments involving the polysulfides

and polythionates; which are outlined in chapter 3.

1.2.4.6 Characterization of Bare and Modified Glassy Carbon Electrode Surfaces

Working electrode surface characterizations on the BGCE and vanadium oxide MGCE

were conducted through an external laboratory using a Hitachi SU-70 Schottky Field

Emission Scanning Electron Microscope coupled with Energy Dispersive Spectrometer

(Oxford Aztec 80 mm/124 eV EDX).

21

1.3 RESULTS

1.3.1 Standardization of Iodine and Thiosulfate Solutions

As shown in Table 1, the standardized concentrations for the iodine and thiosulfate

solutions as prepared in Appendix A were 25 ± 0.2 mM and 25 ± 0.1 mM respectively.

The small standard error points to the fact that iodine and thiosulfate were of very high

purity. As indicated in reaction R – 1.3 (Appendix B), iodometric titration involves the

reduction of iodine with the thiosulfate ion. Prior to using the sulfide solution for

electrochemical analyses, the sulfide solution needed to be standardized using iodometric

titration.

Table 1. Standardized values (mM) for thiosulfate (S2O32-) and iodine (I2) used in

iodometric determination of sulfide. Values are the mean of n = 5.

Solution Expected Conc. (mM) Mean Conc. (mM)

S2O32- 25 25 ± 0.1

I2 25 25 ± 0.2

1.3.2 Standardization of Sulfide Solutions using Iodometric

Titration

As described above, the standardization of sulfide requires the addition of the iodine

solution to be in excess. Since the iodine is present in excess, the iodine oxidizes the

entire sulfide that is present in solution (Appendix B, R – 1.4). The excess iodine is

determined by back titrating with standardized thiosulfate (Appendix B, R – 1.3). The

concentration of the sulfide in the stock solution (section 1.2.1.1) was found to be 106 ± 1

mM (n = 9). The difference between the calculated, based on the original weight of Na2S

prior to dissolution, and the experimental values determined iodometrically was 4 ± 1

mM (n = 9). It should be noted that there was some precipitate formed in solution during

the titration process. This was most likely as a result of the formation of elemental sulfur

during the oxidation of sulfide by iodine (Appendix B, R – 1.4). Also, according to

Pawlak & Pawlak (1999), some of the concentration differences between the calculated

22

and the titrated values could be attributed to the oxidation of sulfide to sulfate during

transfer of the sulfide from the basic medium (pH ≥ 10) to the acidified iodine solution.

Also, the preparation of the sulfide stock solution (section 1.2.1.1) which involved the

dissolution of a small weight of the anhydrous Na2S into a small volume of the buffer

(20-mL) could also have generated the significant difference between the calculated and

experimental values as observed above.

1.3.3 Electrochemical Analysis of Sulfide

1.3.3.1 Electrochemical Analysis of Sulfide using BGCE

According to the sulfur fractional dissociation of sulfide plot (Figure 2), drawn as a

function of pH, the main species present in the sulfide containing solution having a pH of

10.3 is primarily HS-. A small amount of S2- may also have been present. The acid base

equilibria for the dissociation of sulfide in solution are outlined in reactions R – 1.5 and R

– 1.6, in Appendix B. The Ka1 and Ka2 values used in the plotting of Figure 2 were 9.1 x

10-8 and 1.1 x 10-12 (Harris, 2010).

As the HS- becomes oxidized it is possible that more than one oxidation product is

formed at the electrode surface (Kuhn et al., 1983). As shown in Appendix B, the

oxidation products of sulfide may include S0, S2O32-, S4O6

2-, SO32-, and SO4

2-. It has been

reported that at the pH ranges of 7.5 – 11, sulfide oxidation may result in the formation of

SO32-, S2O3

2-, and SO42- (Zhang and Millero, 1993).

23

Figure 2. Fraction dissociation plot for sulfide as a function of pH. Where indications of

sulfide species in solution are H2S(aq) (♦), HS- (■) and S2- (▲).

Figure 3 shows six overlapped voltammograms, obtained at the BGCE. Five of the

voltammograms are for 1 mM HS- solution and the other one for the supporting

electrolyte solution. As indicated in Figure 3, the voltammograms were obtained by

scanning potential from -1600 to +1600 mV, the potential at which it was switched back

to -1600 mV. The results in Figure 3 show that more than one oxidation products were

obtained while anodically scanning the potential. On the other hand, one cathodic

product appears to have formed while scanning the potential in reverse. There was no

presence of electroactive species in the blank under these analysis conditions.

The purpose for conducting several CV cycles for each solution was to monitor any

changes that may have occurred at the surface of the electrode. Specifically, the potential

was switched back and forth ten times between -1600 mV and +1600 mV for each

solution to provide information on probable adsorption onto the surface of the electrode.

In order to produce a more distributed spread between the CV cycles, only the even

cycles are represented in Figure 3. The Epa appears at three distinct potentials, with their

0

10

20

30

40

50

60

70

80

90

100

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

% C

om

po

siti

on

of

sulf

ide

sp

eci

es

pH Units

H2S ♦HS- ■ S2- ▲

24

mean values being -6, +434, and +1158 mV vs. Ag/AgCl respectively. There is also a

peak present in the Epc at a mean value of -1208 mV vs. Ag/AgCl. Peak current values,

at the indicated potentials in the Epa and Epc segments, exhibit increases from cycles 1 to

10. However, when the peak currents alone are examined, as shown in Figures 4 and 5,

versus the cycle number, there is more or less a steady current response from cycles 5 to

10 in most cases. More cycles, were examined, up to 15, but it became clear that there

were no significant increases in terms of the CV outputs after the 10th cycle (Figure 4).

Since this was the case, the 10th cycle was selected to represent the CV analysis for each

of the proceeding experiments using BGCE.

25

Figure 3. Overlapped voltammograms showing five CV cycles of HS- (1 mM), with the

supporting electrolyte at BGCE with a scan rate of 100 mV s-1.

Figure 4. Overlapped voltammograms showing 10 – 15 CV cycles of 1 mM HS- solution

at the BGCE with supporting electrolyte solution (dashed line), at scan rate of

100 mV s-1.

-160

-100

-40

20

80

140

200

260

320

380

-1600 -1200 -800 -400 0 400 800 1200 1600

I /

µA

E / mV vs. Ag/AgCl

Blank

2

4

6

810

-200

-125

-50

25

100

175

250

325

400

-1600 -1200 -800 -400 0 400 800 1200 1600

I / µA

E mV vs. Ag/AgCl

Blank

10-15

26

As shown in Figure 5, peak current responses (Ipa), moving towards more positive

potentials, at the mean potential positions listed above from Figure 3 exhibit an increase

in their values from cycle 2 to cycle 5. However, from cycles 5 to 10 the peak current

plateaus at potential positions -6 and +434 mV (Figure 3), with a mean Ipa = 15 ± 0.45 µA

and 57 ± 0.55 µA, respectively. The peak current at the potential position +1158 mV

displays a gradual increase from cycle 5 to cycle 9 with the peak current dropping off by

11 µA in cycle 10 (Table 2). The peak(s) through this potential region get masked by the

electrochemical response to the presence of water molecules in solution as the CV cycles

increase. Khudaish and Al-Hiani (2006) have suggested that this region is close to the

electrolysis of water, and one should use caution when interpreting data in this region.

Figure 5. Epa positions (mV) at -6 (♦), +434 (▲), and +1158 (■) with peak current values

(Ipa) µA at CV cycles 2 and 5-10 for 1 mM HS-.

0

10

20

30

40

50

60

70

80

90

100

2 5 6 7 8 9 10

I pa

/ µ

A

CV cycle number

27

Table 2. Mean potential positions and peak current (µA) values in the anodic segment for

a 1 mM solution of HS- through the 2nd to the 10th CV cycle.

Potential Positions Epa (mV)

-6 +434 +1158

CV Cycle No. Peak Current Ipa (µA)

2 0 45 46 3 8 51 68

4 12 51 74

5 13 56 79

6 15 56 82

7 15 56 84

8 15 57 89

9 16 57 92

10 16 57 81

Outlined in Figure 6 are the peak current responses (Ipc) at the potential, Epc, of -1208

mV. The results indicate that from CV cycle 2 to cycle 5, there is an increase in peak

current from 52 µA to 97 µA in the 2nd and 5th cycles respectively. The peak current then

becomes steady at a value of 106 µA from the 6th to the 10th cycle (Table 3).

As stated above, the results shown in Figure 3 and summarized in Table 2 and Table 3

indicate that when 1 mM HS- is scanned anodically, three peaks result whose peak

currents appear to become steady from the 6th CV cycle. The anodic peaks appear at – 6,

+ 434, and at +1158 mV. The cathodic peak appears at -1208 mV. A number of authors

such as Kissinger and Heineman (1996), Scholz (2010) and Zanello et al. (2012) have

pointed out that for a reversible redox reaction, the difference between the anodic peak

and the corresponding cathodic peak potentials (∆Ep) should be 59

𝑛 mV, where n

represents the number of electrons involved in the redox reaction. The minimum

separation between the anodic peak potential and that of the only cathodic peak potential

(Figure 3) is 1214 mV. The redox reactions that occurred during the CV of sulfide

cannot therefore be considered reversible. According to Kissinger and Heineman (1996),

electrochemical irreversibility is caused by slow electron exchange of the electroactive

species at the working electrode. In some cases, irreversibility may be, according to the

authors, be a result of diffusing away the product formed at the working electrode into the

bulk of the solution before the reverse scan, at the particular potential of interest, is

28

reached. These authors also state that the irreversibility of a redox system may be

influenced by smaller rate constants. They also go on to state that irreversibility

influences the peak current ratio, in that smaller peak current values may be observed in

the reverse scan than in the forward scan. Based on these facts, there was no evidence

from the data that was obtained (Figure 6 and Table 3), of an indication of reversibility

through this oxidation-reduction couple. Also, the anodic peak current values were 16,

57 and 92 µA after 10 CV cycles (Table 2), the anodic peak currents showed no increase

in value from CV cycles 10-15 (Figure 4). This phenomenon could suggest a quasi-

reversible redox system (Zanello et al., 2012).

Figure 6. Epc position (mV) at -1208 with peak current values (Ipc) µA at CV cycles 2 and

5-10 for 1 mM HS-.

0

20

40

60

80

100

120

2 5 6 7 8 9 10

I pc

/ µ

A

CV cycle number

29

Table 3. Cathodic peak curretns at -1208 mV for CV cycles 1 through 10 for a 1 mM HS-

solution.

Potential Position Epc (mV)

-1208

CV Cycle No. Peak Current Ipc (µA)

1 52 2 52

3 76

4 89

5 97

6 104

7 106

8 106

9 106

10 106

Depicted below in Figure 7 and Table 4, are the results of the CV analysis of the various

concentrations of sulfide at -1263 mV on the BGCE. There is a good positive linear

correlation towards increasing concentrations of sulfide (R2 = 0.914). However, the

precision of the current signal decreases as the concentration of sulfide increases. Table

4 outlines the mean peak current response as a function of the sulfide concentration

solution. At lower sulfide concentrations (0.6 mM), the standard error is decreased over

the upper concentrations examined (2.2 mM), where the standard error is 4 times greater

than that of the lower end of the concentration range examined. Three to six replicates

were examined in order to narrow the observed error over the concentration range of

sulfide, and to provide adequate information to assess the reproducibility of the

experiment. However, even with the increase to the number of replicates, the error

decreased only slightly. Table 5 and Figure 8 illustrate the CV analysis and the peak

current response of the various concentrations of sulfide in solution with potentials

moving to more positive regions. Of the three oxidative peaks, the one at + 565 mV has

the best linear correlation between the sulfide concentration and the current response.

The R2 value at this potential is 0.926 while as at the other two peaks, 0 mV and +1142

mV, the R2 values are 0.628 and 0.897, respectively. There are 3 oxidative products that

correspond to the potentials of 0, +565 and +1142 mV. Based on the half-reaction and

standard formal potentials outlined by Bouroushain (2010), the possible oxidative

products formed could be elemental sulfur, sulfate, sulfite and thiosulfate. As reported by

30

authors Lawrence et al. (2000b) and Dutta et al. (2009), the error that is outlined in

Tables 4 and 5 could be attributed to the mechanical polishing that is required between

experiments in order to remove any adhered oxidative products, such as elemental sulfur,

that may be on the surface of the electrode after analysis. Polishing is known to re-

activate the surface of an electrode (Kamau, 1988; Chen and McCreery, 1996; Kiema et

al., 2003; McCreery, 2008). Even though every effort was made to polish the electrodes

that were used during the CV analyses of the HS- solutions as consistently as possible,

there was no guarantee that every residue on the surface of any given electrode was

removed. Further, since it was not possible to consistently use only one electrode

through the CV experiments, it is likely that imperfections in the various electrodes could

have contributed to some of the standard error values that were observed.

Figure 7. CV analysis for the various HS- concentrations at a mean Epc = -1263 mV at

BGCE. Error bars represent ± S.E. of six replicate runs (R2 = 0.988).

0

50

100

150

200

250

300

0 0.25 0.5 0.75 1 1.25 1.5 1.75 2 2.25 2.5

Ipc

/ µ

A

Bulk Sulfide Concentration (mM)

31

Table 4. Mean of six replicate runs and ± SE of current responses (Ipc) for CV analysis of

sulfide obtained at the BGCE at the Epc of -1263 mV.

Bulk Conc. Sulfide (mM) Mean Ipc (µA)

0.6 33 ± 8

1.1 69 ± 16

1.4 95 ± 22

1.6 124 ± 26

1.9 218 ± 20

2.2 243 ± 36

Table 5. Mean CV anodic currents (Ipa) for the sulfide analysis at the BGCE at potentials

of 0, +565, and +1142 mV. Mean (µA) ± SE

Conc. Sulfide (mM) Ipa(1) (0 mV) Ipa(2) (565 mV) Ipa(3) (1142 mV)

0.6 5 ± 1 27 ± 5 39 ± 8

1.1 12 ± 3 46 ± 8 50 ± 8

1.4 21 ± 3 63 ± 7 77 ± 7

1.6 27 ± 8 73 ± 12 80 ± 13

1.9 68 ± 14 71 ± 14 ―

2.2 70 ± 13 106 ± 22 ―

2.5 41 ± 8 98 ± 13 ―

32

Figure 8. CV anodic currents (Ipa) for various HS- concentrations at Epa = 0 mV (♦) R2 =

0.628, +565 mV (■) R2 = 0.926, and +1142 mV (▲) R2 = 0.897 on BGCE. Error bars

represent ± S.E. of six replicate runs.

Figure 9 shows scanning electron microscopy (SEM) images of the BGCE surfaces for a

polished and electrochemically pretreated electrode (image A) and after running 10 CV

cycles (-1600 to +1600 mV) of a 1 mM HS- solution on the polished and

electrochemically pretreated electrode (image B). As indicated in section 1.2.4.3 the

pretreatment of the electrode was accomplished by running CV scans by varying the

potential of the working electrode from –200 to +800 mV at 100 mV s1- for 25 cycles.

The chevrons, shown in Figure 9 (B), indicate small deposits present on the surface of the

BGCE. Energy-dispersive X-ray (EDX) was conducted in conjunction to SEM work

(Figure 9). However, the only spectral band present in both samples was carbon for the

results attained though EDX (Figure 10 and Figure 11). Based on the EDX shown in

Figure 9, there appears to have been insufficient sulfur build-up on the electrode surface

for detection. The SEM image in Figure 9 (B) shows that there was some deposition that

had built up on the surface of the electrode. The SEM and EDX work performed was

conducted at an external laboratory. The time it took for the sample preparation to be

done and the impending analysis to be completed was substantial and could have affected

0

20

40

60

80

100

120

140

0 0.25 0.5 0.75 1 1.25 1.5 1.75 2 2.25 2.5

I pa

/µA

Bulk concentration of sulfide (mM)

33

the chemistry of any adsorbed material on the surface of the electrode. Figure 12 shows

another SEM image of a BGCE after 10 CV cycles (-1600 to +1600 mV) of a 1 mM HS-

solution made up in 0.1 M phosphate buffer. The SEM image, taken using Joel JCM-

6000 (Hoskin Scientific), was obtained on the same day that the CV cycles of 1 mM HS-

solution were analyzed on BGCE. Although it was observed that there was more

deposition on the surface of the electrode, as opposed to the image observed in Figure 9

(B), the composition of the surface could not be determined due to the limitation of the

SEM utilized for the image in Figure 12. The SEM utilized for the acquisition of the

image in Figure 12 was limited due to the lack of an alternative detection source (EDX)

for appropriate characterization of the deposits observed. The images attained, using that

particular SEM (Joel JCM-6000), were for exploratory purposes. Due to this limitation

the presence of sulfur or any other product that may have been deposited on the BGCE

could not be confirmed.

Figure 9. BGCE SEM images A) is polished and electrochemically pretreated in 0.1 M

phosphate buffer, B) BGCE after 10 CV cycles of 1 mM HS-, in 0.1 M potassium

phosphate buffer, chevrons indicate some product build-up from sulfide electrochemical

oxidation.

34

Figure 10. EDX spectra of BGCE pretreated by running CV cycles of the BGCE in 0.1 M

phosphate buffer from -200 to +800 mV at 100 mV s1-; where C = carbon.

Figure 11. EDX spectra of BGCE after 10 CV cycles (-1600 to +1600 mV) in 1 mM HS-

solution made in 0.1 M phosphate buffer. Where C = carbon and O = oxygen.

0

250

500

750

1000

1250

1500

1750

2000

2250

2500

0 2.5 5 7.5 10 12.5

Inte

nsi

ty (

Co

un

ts)

keV

C

0

1200

2400

3600

4800

6000

7200

8400

9600

10800

12000

0 2.5 5 7.5 10 12.5

Inte

nsi

ty (

Co

un

ts)

keV

C

O

35

Figure 12. SEM of BGCE obtained after scanning 10 CV cycles (-1600 to +1600 mV) in

a 1 mM HS- solution prepared in 0.1 M phosphate buffer. Chevron indicates deposition

onto electrode surface.

1.3.3.2 Cyclic Voltammetric Analysis of Sulfide Solutions at MGCE

As explained in section 1.2.4.4, the modified electrode surface film was prepared by the

electrochemical deposition of an oxide of vanadium on the GCE surface. This was

intended to enhance the detection of sulfide at lower concentrations. With repetitive

potential cycling between 0 to +2000 mV, for 20 cycles at a scan rate of 50 mV s1-, the

current density and peak potential in the anodic segment decrease. This decay is

attributed to the consumption of VO2+ to form V2O5 at the electrode surface (Weckjuysen

and Keller, 2003; D’Elia et al., 2004; Khudauish and Al-Haini, 2006). Weckjuysen and

Keller (2003) and Harris (2010) have indicated that the reaction responsible for the

formation of the V2O5 film from VO2+ is given by the Reaction R – 1.7, shown below.

2𝑉𝑂2+ + 3𝐻2𝑂 → 𝑉2𝑂5 + 6𝐻+ + 2𝑒− (𝑅 − 1.7)

The deposition of the film can be observed below in Figure 13 (B) where the MGCE and

the BGCE Figure 13 (A) are compared. As shown below in Figure 13 (A) the BGCE has

pores on the surface which according to McCreery (2008) are formed during fabrication.

The MGCE SEM image shown in Figure 13 (B) shows a uniform coverage over the

36

glassy carbon surface which is consistent with literature that presents characterization of

the electrodeposition of V2O5 films onto GCE (D’Elia et al., 2004).

Figure 13. SEM images of BGCE (A) and MGCE (B), showing the surface

characterizations between the BGCE and MGCE.

Figure 14 summarizes the comparison of the blank (supporting electrolyte solution) and a

0.36 mM sulfide solution at MGCE. There are peak current responses in both the Ipa/Ipc

segments of the supporting electrolyte solution, but the current density shows a marked

increase in the presence of sulfide solution. This overlap in current was corrected for by,

taking a blank measurement before each set of experiments with MGCE, taking the

current responses (Ipa/Ipc), from those blank measurements, and then subtracting the blank

current responses from the current responses of the sulfide solutions along the potential

regions of interest (Epa/Epc). In the case of Figure 14, the blank current response was 37 ±

4 µA at 132 ± 18 mV, and the current response measured for 0.36 mM sulfide was 70 ± 3

µA at 145 ±18 mV. The actual current response for 0.36 mM of sulfide taken at 145 ± 18

mV was 33 ± 3 µA with background correction applied. In contrast, the supporting

electrolyte response in the case of the BGCE (Figure 3) does not interfere with the

current responses from the sulfide present in solution as it does in the case of the MGCE.

EDX and SEM analyses of the surface of the MGCE after 10 CV cycles in just the

supporting electrolyte solution did not show any accumulation of foreign products on the

film. Also, contrary to the observation of D’Elia et al. (2004), the analyses did not show

the presence of vanadium bands. However, when the sulfide was added and 10 CV

37

cycles were conducted between -1600 and +1600 mV, there was a presence of deposits

on the MGCE surface (Figure 15 and Figure 18). EDX analyses of these electrode

surfaces revealed the presence of copper, sulfur, carbon and oxygen bands (Figure 16).

As noted earlier, even though the CV of the supporting electrolyte has significant residual

currents, the bands found on MGCE did not show any interfering products (Figure 17 and

Figure 19) that could explain the signal observed in Figure 14. Since the EDX and the

SEM with just the supporting electrolyte solution on MGCE did not show any traces of

deposition to the surface related to the supporting electrolyte solution, that would be

indicate that they interaction between the deposited V2O5 film and sulfide is responsible

for the deposits on the MGCE shown in Figure 15. Khudaish and Al-Hinai (2006)

proposed the following electrochemical mechanism for the oxidation V2O4 to soluble V5+

as shown in reactions R – 1.8 and 1.9. The proposed mechanisms (Khudaish and Al-

Hinai, 2006) for the reactions happening during CV cycles at MGCE suggest that the

vanadium oxide modifier is a self-regenerating film, which could be the cause of the

segment peaks observed in the anodic and cathodic segments (Figure 14) at Epa/Epc =

+303, +1125 and -641 mV respectively. When the MGCE is in the presence of sulfide,

the catalytic response from the MGCE increases the peak current in the overlapped

potential regions where the blank is present.

𝑉2𝑂4 + 2𝑂𝐻− → 2𝑉𝑂3− + 2𝐻+ + 2𝑒− (𝑅 1.8)

2𝑉𝑂3− + 𝐻2𝑂 → 𝑉2𝑂5 + 2𝑂𝐻− (𝑅 1.9)

38

Figure 14. The fifth cycle of a CV output of 0.36 mM concentration of HS- (solid line) at

MGCE, and the supporting electrolyte solution (dashed line) at a scan rate of 100 mV s1-.

Figure 15. SEM image of MGCE showing foreign deposits on the film modification after

10 CV cycles of 1 mM HS- at 100 mV s1- (A) and MGCE showing 10 CV cycles in the

supporting electrolyte solution at 100 mV s1- (B).

-80

-60

-40

-20

0

20

40

60

80

100

120

-1600 -1200 -800 -400 0 400 800 1200 1600

I /

µA

E/ mV vs. Ag/AgCl

39

Figure 16: EDX spectra of MGCE after 10 CV cycles (-1600 to +1600 mV) of 1 mM HS-

solution made in 0.1 M phosphate buffer. Where C, O, S and Cu represent carbon,

oxygen, sulfur and copper respectively.

0

250

500

750

1000

1250

1500

1750

2000

2250

2500

0 1 2 3 4 5 6 7 8 9 10 11 12

Inte

nsi

ty (

Co

un

ts)

keV

Cu

Cu

Cu

SO

C

40

Figure 17: EDX spectra of MGCE after 10 CV (-1600 to +1600 mV) cycles in supporting

electrolyte solution (0.1 M phosphate buffer), where C = carbon.

Figure 18. SEM image of MGCE (left), showing sites of interest for EDX along with

supporting EDX data expressed as weight % (right) for V2O5 modifier.

0

250

500

750

1000

1250

1500

1750

2000

2250

2500

0 1 2 3 4 5 6 7 8 9 10 11 12

Inte

nsi

ty (

Co

un

ts)

keV

C

41

Figure 19. SEM image of MGCE after 10 CV cycles in 1 mM HS- solution, showing sites

of deposits on surface (left) and EDX data expressed as weight % (right) for MGCE after

10 CV cycles in 1 mM HS- solution.

Figure 20 shows 10 overlapped CV cycles of the supporting electrolyte obtained using

MGCE. A progressive decay in the anodic current density from cycle 1 to 10 at an Epa of

+1200 mV is observed. According to authors Barrado et al. (1997), Weckjuysen &

Keller (2003), D’Elia et al. (2004), and Khudaish & Al-Hinai (2006), the likely cause of

this decay is the electroxidation of V2O5 to V2O4. This occurs when the V2O5 undergoes

an electron transfer in the presence of water molecules. This current decay at +1200 mV

can also be observed in Figure 21 which contains an overlap of 3 CV cycles of 0.44 mM

sulfide solution at the MGCE. This process, for the response of the V2O5 film towards

the electrochemical response to sulfide, is outlined below as reaction R – 1.10 as

proposed by McCleverty & Meyer (2004) and Khudaish & Al-Hinai (2006). According

to D’Elia et al., (2004), the presence of this decay in current density has been attributed

to the presence of dissolved oxygen. These authors pointed out that at potentials greater

than +1230 mV, the oxidation of water occurs. This in turn could enhance the

electrooxidation of V2O5 as depicted in reaction R – 1.10. Figure 20 and Figure 21 show

that between 0 and +200 mV, there is an anodic peak that develops at the MGCE for both

the supporting electrolyte alone and the 0.44 mM sulfide solution. However, the anodic

peak current in 0 to +200 mV region is more enhanced in the presence of the sulfide ion.

42

According to McCleverty & Meyer (2004) and Khudaish & Al-Hinai (2006), the

variation in current density in the Epa region 0 to +200 mV, in the presence of sulfide, is

attributed to the oxidation of V2O5·HS- which creates elemental sulfur (Reaction R –

1.10). As later CV cycles occur, further accumulation of elemental sulfur on the surface

may cause the active regions of the vanadyl film to become fouled (Figure 15 A).

𝑉2𝑂5 ∙ 𝐻𝑆− → 𝑉2𝑂4 + 𝑆 + 𝑂𝐻− (𝑅 − 1.10)

It is possible that the build-up of elemental sulfur on the MGCE surface could reduce the

availability of the reactive sites of the electrode surface (Figure 16) for continuous HS-

oxidation over a series of CV cycles. The presence of elemental sulfur was supported by

the SEM and EDX data provided from Figure 15 (A) and Figure 16 for the analysis of

sulfide at MGCE. In addition, the Epc shifted to a more negative region with each

successive CV cycle. The Epc shown in Figure 17 has an average Epc of -540 mV. When

compared to the Epc obtained with the BGCE (-1208 mV), the catalytic activity of the

V2O5 used in the MGCE may be considered responsible for a potential shift of

approximately 668 mV. These changes in current densities over the MGCE when in the

presence of the sulfide ion are also illustrated in Figure 22 to Figure 24 as well as in

Table 6. It should be noted that in Table 6 for both the Epa1 and Epc there is significant

peak shift, where the shift is 456 mV towards more positive potentials from CV cycles 1

to 10, and a 247 mV shift towards more negative potentials from CV cycles 1 to 10,

respectively. This peak shift was also observed in the MGCE CV cycles with just the

supporting electrolyte solution, where the peaks shifted 466 mV towards more positive

potentials and 325 mV towards more negative potentials for Epa1 and Epc, respectively.

As shown in Figure 22 to Figure 24, there is little separation in the current response

between the supporting electrolyte solution at the MGCE compared to the current signal

when in the presence of sulfide. However, as shown in Figure 22, the best separation in

the current response between the sulfide signal and just the supporting electrolyte signal

at the MGCE occurs at the 4th CV cycle. As a result of this, the 4th CV cycle was used

for the sulfide solution analysis at the MGCE Figure 25 and Table 7.

43

Figure 20. CV analysis of supporting electrolyte solution on MGCE from cycles 1 to 10,

scan rate 100 mV s1-; the numbers indicate the scan number of the CV cycle.

-200

-140

-80

-20

40

100

160

220

280

340

400

-1600 -1200 -800 -400 0 400 800 1200 1600

I /

µA

E / mV vs. Ag/AgCl

1

2

3

4

}5-10

44

Figure 21. CV graph at MGCE of HS- solution (0.44 mM); scan rate of 100 mV s1-

showing cycles 2, 5, and 10 as well as the blank (dashed line).

-200

-150

-100

-50

0

50

100

150

200

250

300

350

-1600 -1200 -800 -400 0 400 800 1200 1600

I / μ

A

E / mV vs. Ag/AgCl

2

10

5 & blank

45

Figure 22. Peak current values (Ipa1) µA at CV cycles 1 to 10 for 0.36 mM HS- (♦),

compared to CV cycle numbers 1 to 10, and compared to the supporting electrolyte

solution (▲) on MGCE.

Figure 23. Peak current values (Ipa2) µA at CV cycles 1 to 10 for 0.36 mM HS- (♦),

compared to CV cycle numbers 1 to 10, and compared to the supporting electrolyte

solution (▲) on MGCE.

0

20

40

60

80

100

0 1 2 3 4 5 6 7 8 9 10

I pa

1/

µA

CV Cycle No.

0

50

100

150

200

250

300

0 1 2 3 4 5 6 7 8 9 10

I pa

2/

µA

CV Cycle No.

46

Figure 24. Peak current values (Ipc) µA at CV cycles 1 to 10 for 0.36 mM HS- (♦),

compared to CV cycle numbers 1 to 10, and compared to the supporting electrolyte

solution (▲) on MGCE.

Table 6. Peak current, potential and scan number for 0.36 mM HS- run over MGCE at a

scan rate of 100 mV s1-.

Scan No. Ipa1 (µA) Epa1 (mV) Ipa2 (µA) Epa2 (mV) Ipc (µA) Epc (mV)

1 57 -250 274 1235 156 -423

2 45 -125 217 1234 128 -448

3 62 -45 135 1184 95 -492

4 77 64 76 1107 73 -543

5 71 141 41 1107 58 -594

6 57 185 25 1120 48 -620

7 50 211 18 1120 43 -664

8 43 224 14 1117 39 -643

9 38 233 14 1142 35 -668

10 37 206 12 1090 34 -670

Six concentration levels were examined in an attempt to quantify various sulfide

concentrations at the MGCE. The sulfide concentrations examined at the MGCE were:

0.67, 0.44, 0.36, 0.31, 0.22 and 0.13 mM. The sulfide concentrations used for the

calibration curve in the case of MGCE were lower than when BGCE was used (Figure 7

and Figure 8). According to Khudaish and Al-Hiani (2006), there was a significant

change in the concentrations of the sulfide solutions used in the BGCE analyses when

0

40

80

120

160

200

0 1 2 3 4 5 6 7 8 9 10

I pc

/ µA

CV Cycle No.

47

compared to that of the MGCE analyses. This change in the sulfide concentrations for

the MGCE analyses was in the part from the enhanced catalytic activity provided by the

V2O5 modifier film on the surface of the electrode. Figure 25 shows that there is a good

linear correlation between the peak current at +200 mV and the various concentrations of

sulfide in the anodic segment. These results indicate that the oxidative products of HS-

slightly increase linearly with increasing concentration of sulfide (Table 7). The results

in Figure 26 and Table 8 show that in the cathodic segment, at a potential of -570 mV, is

a linear increase in the peak current as the sulfide concentration increases during the

reduction of HS- at the MGCE. However, the standard error results (Table 8) for the HS-

concentrations 0.13, and 0.31 mM raises some concerns. For those two concentrations

(0.13 and 0.31 mM), after background corrections were conducted, had values of 0.7 ± 3

and 1 ± 4 µA. The standard error is larger than the measurement attained, which would

suggest that since these two measurements are less than their standard errors, it would be

difficult to differentiate them from background noise and are therefore unreliable. The

lower peak current values observed in the cathodic segment provide evidence that the

redox system being analyzed at MGCE is irreversible, since electrochemical

irreversibility influences the peak current ratio and that the more irreversible a couple the

smaller the peak current values will be on the reverse segment (Tables 7 and 8)

(Kissinger and Heineman, 1996; Zanello et al., 2012). Although, direct proportionality

between the peak current at -570 mV and the sulfide concentrations is weak, the signal

for detection for the reduced HS- concentrations in the cathodic segment are detected at

values ≥ 0.44 mM. Nonetheless, there is a catalytic increase in current response at the

MGCE of solutions containing sulfide when compared to the supporting electrolyte

solution without sulfide (Figure 14).

48

Figure 25. CV analysis increasing towards more positive potentials Ipa (µA) for various,

concentrations of HS- at Epa = +200 mV on MGCE after blank corrections. Error bars

represent ± S.E. of six replicate runs (R2 = 0.937).

Table 7. Mean ± SE for anodic current response Ipa (µA) of various concentrations of HS-

at MGCE; at a mean anodic potential Epa of +200 mV, (n = 6; R2 = 0.937).

Bulk Conc. Sulfide (mM) Mean Ipa (µA)

0.13 35 ± 3

0.22 36 ± 4

0.31 40 ± 6

0.36 52 ± 6

0.67 54 ± 4

25

30

35

40

45

50

55

60

0.00 0.13 0.26 0.39 0.52 0.65 0.78

I pa

/ µ

A

Bulk Sulfide Concentration (mM)

49

Figure 26. CV analysis increasing towards more negative potentials Ipc (µA) for various

concentrations of HS- at Epc = -570 mV on MGCE after blank corrections. Error bars

represent ± S.E. of six replicate runs (R2 = 0.907).

Table 8. Mean ± SE for cathodic current response Ipc (µA) of various concentrations of

HS- at MGCE at a mean cathodic potential Epc of -570 mV, (n = 6; R2 = 0.907).

Bulk Conc. Sulfide (mM) Mean (Ipc)

0.13 0.7 ± 3

0.31 1 ± 4

0.36 9 ± 4

0.44 13 ± 3

0.67 23 ± 3

1.4 CONCLUSIONS

As mentioned above fifteen CV cycles were initially examined in order to optimize an

appropriate number of cycles for the determination of sulfide at BGCE and MGCE. Only

ten CV cycles were used for each of the six replicates examined at each concentration

level listed for BGCE and MGCE. The reason for choosing 10 CV cycles was the

observation that after the 10th CV cycle there were no increases to the CV outputs (Figure

4). As illustrated in the SEM images in Figure 12 and Figure 15 (A), it is believed that

surface fouling on both the BGCE and MGCE may be attributed to the buildup of

-5

0

5

10

15

20

25

30

0.00 0.13 0.26 0.39 0.52 0.65 0.78

I pc

/ µ

A

Bulk Sulfide Concentration (mM)

50

elemental sulfur on the working surface of the electrode. Such surface fouling leads to a

progressively reduced active working electrode surface as a series of the CV scans

proceed. This effect could, in turn, cause a progressive decrease in the output current.

As the CV output presented in Figure 3 shows, there are three oxidative products that

appear at three distinct potentials, namely at 0, 565, and 1142 mV. The possible

oxidative products that may have formed during the CV analysis of sulfide at these

potentials are elemental sulfur, sulfate and thiosulfate. Kinetic studies conducted on the

oxidation of sulfide in aqueous solutions by several authors (Avrahami & Goulding,

1968; O’Brien & Birkner, 1977; Millero, 1986; Kotronarou & Hoffmann, 1991) have

proposed mechanisms for the formation of sulfur intermediates based on the presence of

oxygen in sulfide solutions being analyzed. In a study done by Avrahami and Goulding

(1968) that examined the oxidation of sulfide in water in the pH range of 11-13, the

disappearance of sulfide was found to follow first order kinetics, with sulfate and

thiosulfate being the products of oxidation. The authors further suggest that the oxidation

of sulfide forms sulfite in the presence of oxygen. The sulfite is then further oxidized to

sulfate. These authors have presented possible mechanisms for the intermediate

oxidative products of sulfur, formed from sulfide. O’Brien and Birkner (1977) also

examined the oxidation of sulfide in water, through the pH range of 7.5-11, and they

suggested that oxygen as well as the pH of the sulfide solution influence the formation of

sulfur species during oxidation. Based on their kinetic study, the authors proposed that

with low sulfide to oxygen ratios the sulfur products formed would be sulfite, thiosulfate

and sulfate. However, when the ratio of sulfide to oxygen is high the formed products are

elemental sulfur and polysulfides. In the work presented in this chapter the potentials

were cycled from -1600 to +1600 mV for 10 cycles. As reported by Harris (2010), the

electrolysis of water occurs at potentials of -1230 mV. Since the CV analysis conducted

in this work included potentials where water molecules are oxidized to oxygen, this could

potentially lead to the formation of sulfur intermediates during the anodic sweeping. The

speciation of the specific sulfur intermediates that form may be realized through further

investigations of the chemistry occurring at the surface of the electrode.

As observed from the results obtained in this chapter, there was a fair amount of

variability among the experiments performed. Increasing the number of replicates and

51

performing Dixon’s Q test to recognize and remove outliers did not seem to help reduce

the variability of sequential replicate values observed here. The results in this chapter

further indicate that the variability was reduced when CV analysis of HS- was conducted

over MGCE (Table 8 & Table 9), compared to the variability that resulted from the

analysis of sulfide at BGCE (Table 4 & Table 5).

It is clear that the higher the concentration of HS- being measured, the greater the

variability of replicate measurements (S.E. increases with increasing concentration of

sulfide). It is possible that once the sulfide is in solution, any exposure to air could cause

rapid oxidation of the HS-. Even though precautions were taken to ensure that the

exposure to oxygen was limited, i.e. significant purging of solutions with nitrogen was

maintained, HS- solution was discharged below the surface of the electrochemical cell

solution, and all procedures were carried out under a blanket of nitrogen, the variable

readings from the same sample suggest that there could have been slight exposure to air.

It has been reported that the presence of oxygen can cause the oxidation of HS-, which

leads to for formation of intermediate sulfur-based compounds such as S0, SO32-, SO4

2-

and S2O32- (Avrahami and Golding, 1968; O’Brien and Birkner, 1977; Millero, 1986;

Kotronarou and Hoffmann, 1991). Secondly, the variation of current responses could

have occurred when more than one GCE was inevitably used. Even if it were possible to

use just only one electrode, slight differences in the polishing/cleaning procedures

between analyses could also lead to current output variations.

The results obtained in this chapter indicate that the MGCE has an enhanced detectability

of HS- ion at lower concentrations than those over the BGCE. Using a MGCE, as more

elemental sulfur gets deposited onto the active vanadium film, there is decline in the

current observed over the series of ten CV cycles (Figure 21). The shift in potential and

the reduced current response after four of the ten cycles examined suggests that useable

HS- analysis could be carried out with fewer cycles than is the case for the analyses

carried out using BGCE. A major disadvantage realized when using MGCE is that the

current responses for the blank sample occurs at the same potential as do those of the HS-

solutions. This could limit the use of this approach to measure low detectable levels of

sulfide. Even though, the lowest sulfide concentration detected using MGCE was 0.13

52

mM, the observed standard error results indicated at concentrations of 0.13 and 0.31 mM

(Table 8), for the cathodic segment, these values would not be considered quantifiable.

To achieve the electrodeposition of V2O5 on the GCE, twenty CV cycles were required to

achieve the film coverage over the GCE surface. There may have been microabrasions

on the surface, like those observed in Figure 9 (A), which could have prevented uniform

deposition of the V2O5 film on the GCE surface, a phenomenon that could reduce the

precision of the results. Although the V2O5 film has previously been described as a stable

product (Barrado et al., 1997; Park et al., 2002, Weckhuysen & Keller, 2003; D’Elia et

al., 2004) the present study shows that there is a significant decay in current density in

the supporting electrolyte solution as well as in the HS- solutions. A possible explanation

for this is that the V2O5 may be undergoing secondary reactions over the sequence of CV

cycles. When the results using the BGCE and the MGCE are compared in terms of their

ability to detect sulfide, the MGCE allows lower levels of sulfide to be detected.

However, during the CV analysis at the MGCE there appears to be current output

stability after 10 cycles. When the HS- solution was added to the electrochemical cell

and CV cycles were resumed, the MGCE surface seemed to show inactivity towards HS-

oxidation. The proposed theory by Khudaish and Al-Hiani (2006) does not seem to

support the regeneration of this modification film. It has been reported elsewhere

(Barrado et al., 1997) that the V2O5 may undergo side reactions with the supporting

electrolyte as well as with any present oxygen during the CV cycles when the potential is

swept from -1600 to +1600 mV. Perhaps as a result of deposition of elemental sulfur on

the GCE surface, the accumulation of these adsorptive products can eventually deactivate

the surface of the working electrode requiring renewal (polishing) (Kissinger and

Heineman, 1996). Zanello et al. (2012), suggested that the MGCE acts like a membrane

which regenerates under different ionic environment thereby causing a shift in the

potential around a particular peak potential, at the working electrode surface. This might

help explain the potential shifts observed in Table 6.

53

Overall, the results obtained from the BGCE and the MGCE experiments casts doubts on

the reliability of these methods for the routine quantitative analysis of sulfide. The

methods nevertheless seem to provide useful qualitative information about the presence

of sulfide in a sample, and it may be possible to modify the apparatus to address the

sources of drift in measured current density and surface changes in the electrodes

themselves as further measurements are made.

54

Chapter 2

Analysis of Polysulfides and Polythionates

2.1 INTRODUCTION

Polysulfides and polythionates are a unique class of sulfur compounds, known to

participate in many environmentally significant processes due to their reactivity as strong

nucleophiles, and reducing agents (Chadwell et al., 2001; Kamyshny et al., 2004;

Kamyshny et al., 2006; Kamyshny et al., 2008 and Kristiana et al., 2010). For example,

polysulfides are important geochemically, and have a high affinity for transition metal

ions. They form sulfide minerals like pyrite (iron sulfide) (Chadwell et al., 2001 and

Kristiana et al., 2010). Polysulfides also participate in the sulfurization of organic matter

in aquatic systems, and are precursors to the formation of volatile organic sulfur

compounds such as dimethyl disulfide (Gun et al., 2000; Kamyshny Jr. et al., 2008 and

Kristiana et al., 2010). Previous reports have stated that polysulfides are present in oxic

and anoxic environments (Kariuki et al., 2001 and Kristina et al., 2010). Sulfur chain

lengths of up to S182- have been reported (Chadwell et al., 2001) being present in

biofilms, and in drinking water distribution systems (Kristiana et al., 2010). The concern

with the presence of sulfides and polysulfides in drinking water systems is that these

species consume disinfectants, dissolved oxygen and react with metal ions. These

processes can produce insoluble metal sulfides and can cause taste and odor problems

(Kristiana et al., 2010). However, they are challenging to analyze, due to their low

concentrations in natural environments (10 – 20 µM), thermal instability, and

susceptibility to oxidation and transformation (Chadwell et al., 2001; Kariuki et al., 2001;

Kamyshy Jr. et al., 2004; Kamyshny Jr. et al., 2006; Kamyshny Jr. et al., 2008 and

Kristiana et al., 2010). The dissociation of polysulfides to polythionates, sulfur and

sulfide depends on the pH, temperature and ionic strength of an aqueous solution.

Along with sulfide and polysulfides, polythionates, such as the thiosulfate ions are also

present in natural environments. These compounds have a similar chemical structure and

their metabolism appears closely related. Polythionates are important intermediate

55

species in the redox transformations of sulfur compounds in many environments and in

the metabolism of sulfur-oxidizing and sulfur-reducing microorganisms (Koh, 1990;

Druschel et al., 2003; Mohapatra et al., 2008). The versatility of thiosulfate is that it can

be oxidized to sulfate or tetrathionate, reduced back to sulfide, or disproportionate into

both sulfide and sulfate, through the metabolism of sulfate reducing bacteria (Ciglenečki

and Ćosović, 1997). Controlling and understanding the distribution of sulfur-base

compounds could aid in the reduction of hydrogen sulfide production from these redox

processes, which in turn could be useful in the management of odors released in anoxic

environments and water treatment processes as well as managing outbreaks of taste and

odour problems in drinking water supplies.

The availability of polysulfide compounds commercially is rare and when available, the

compounds are largely costly. These compounds were therefore synthesized as needed in

the present study. The methods used to synthesize the sodium salts of disulfide,

trisulfide, and tetrasulfide from sodium sulfide and the elemental sulfur are described in

this chapter. The analysis of the synthesized Na-polysulfide compounds using the N,N-

diethyl-p-phenylenediamine (DEPD) colorimetric method, and also the iodometric

titration method was carried out as described in section 1.3.2. Using these two methods

to directly analyze polysulfides and polythionates would be the preferred option. If the

polysulfide ions cannot be directly analyzed using DEPD, they would be reduced to

simple sulfides (S2-) using chromium. The reduced Sn2- ions will be analyzed using

DEPD and iodometric titrations. The DEPD and iodometric titrimetric methods, also

used to analyze thiosulfate ions, are also described in this chapter. The thiosulfate ion

was used as a representative of the class of compounds known as polythionates. An

attempt on how these two methods (DEPD and iodometric) were used to quantify the

polysulfides and thiosulfate directly without transforming them first is also described in

the chapter. Ultimately, it became necessary to reduce these compounds to simple sulfide

before being quantified with these two techniques.

56

2.2 MATERIALS AND METHODS

2.2.1 Making Na-polysulfides

2.2.1.1 Equipment and Preparation

Commercially available polysulfide compounds are only obtainable in the tetrasulfide

form. Even though available, the product is pricy. A 5-g quantity of sodium tetrasulfide

currently costs $500 and even then, its purity is not guaranteed to that of analytical grade.

A preparative method outlined by Rozen and Tegman (1971) was used to make the

polysulfide compounds of interest for this study. Due to the high reactivity of sodium

sulfide (Na2S) with air, the weighing, mixing and transfer of the reactants used in the

making of the polysulfides into the reaction vials was carried out in a Polymer Series 100,

Baxter Glove Box, equipped with a dual purge nitrogen flow control, for humidity, and

temperature monitoring (Terra Universal Inc.). Stoichiometric amounts of elemental

sulfur (S0) and anhydrous Na2S were weighed out on an analytical balance (Ohaus

Explorer Pro; model no. EP114C), in a nitrogen purged environment. The equations 2.1

(a & b), 2.2 (c & d), and 2.3 (e & f) (Appendix B) illustrate the proportions of Na2S and S

that were used to synthesize sodium disulfide (Na2S2), sodium trisulfide (Na2S3), and

sodium tetrasulfide (Na2S4). For example, to make a polysulfide containing an n-sulfur

chain (Sn2-), the molar ratio of Na2S to S0 would be 1:(n – 1).

The relative humidity of the glove box in which the anhydrous Na2S and elemental sulfur

were prepped was set at a range of 20-30%. Ambient temperature in the glove box was

maintained. Mortars and pestles, scoopulas, glass tubes, and glass enclosure vials, used

in this procedure were dried in an oven at 150°C for 12 hours, removed, wrapped in

aluminum foil, and cooled to ambient temperature before using. The glass tubes, and

glass enclosure vials were purged with nitrogen gas, and their openings were sealed with

parafilm before introducing them in the glove box. The glove box was initiated and the

main chamber was evacuated. Manipulations of weighing or mixing the Na2S or S did not

occur until the atmosphere inside the main chamber reached 30% relative humidity or

lower.

57

2.2.1.2 Polysulfide Mixing Process

Stoichiometric amounts of Na2S were weighed out and ground using a mortar and pestle.

A stoichiometric amount of S0 was weighed and added to the same mortar containing the

Na2S. The two reactants were mixed, using the pestle, until completely homogenized.

The homogeneity of the Na2S and S reactants was required to ensure that there would be

uniform melting, and re-crystallization of the final product. Stoichiometric portions of

the individually identified polysulfide mixtures were placed into 18 x 150 mm

borosilicate glass tubes (Baxter Scientific Products, T1290-9A), and sealed with parafilm,

before removing them from the glove box via the air lock port. This was to ensure that

any exposure to oxygen prior to evacuating the nitrogen from the tube during the sealing

process was minimized as much as practically possible.

2.2.1.3 Sealing the Glass Tubes for Final Preparation

Using a water aspirator to create a vacuum, the parafilm-sealed tube containing the

reaction mixture was quickly un-wrapped and attached to the vacuum tubing. Making

sure the sealing from the vacuum tube to the sample tube was air tight. The reactant tube

was evacuated for 60 minutes. Using a hand torch cylinder (MG9 - 14.1 oz. MAP-Pro,

Worthington Cylinders U.S.A.), the tube was sealed by gently softening the upper portion

of the glass tube. This was done by carefully rotating the glass tube around the flame

until the tube walls collapsed, creating a sealed environment for the reaction mixture.

2.2.1.4 Preparation and Retrieval of Na-polysulfides

Outlined in Table 9 is the temperature profile (Rosen and Tegman, 1971) that was used to

form the Na-based polysulfides. The Na-polysulfide reactant mixture was conducted in

an Isotemp Muffle Furnace (Fisher Scientific) using the temperature profile outlined

below in Table 9. Once the reaction was complete, each of the polysulfide products was

removed from its respective glass tube and placed in the mortar for crushing and grinding

to a fine powder. Before being used, the mortars, pestles, and the glass storage vials had

been baked in an oven at 150 °C for 12 - 24 hours and cooled to room temperature. The

58

glass vials into which the synthesized polysulfides were placed had been purged with

nitrogen prior to transferring them into the glove box.

Table 9. Temperature profile for the preparation of Na-polysulfides (Rosen and Tegman,

1971).

Temperature

Segment

Reaction

Temperature (°C)

Reaction

Time

(hours)

Comments

Segment 1 230 10 – 12

Solid state conversion, 80-90 % of

reaction occurs at this segment

Segment 2 300 – 490 ½

Liquid state reaction, complete

conversion

Segment 3 205 1 - 10

Tempering period, recrystallization

occurs

Retrieval of the Na-polysulfides from the glass tubes in which they were synthesized was

conducted in the Baxter Glove Box, under the same conditions as those used for the

initial mixing of the reactants as outlined above. To detach the synthesized polysulfides

from the walls of the glass tubes in which the reactions were carried out, the glass tubes

were carefully tapped with an object of moderate weight. Each glass tube was then

gradually broken apart, starting from the top of the seal and moving downward towards

the polysulfide product. Once the Na-polysulfide was removed from the glass tube it was

placed in a mortar, and a pestle was used to crush and grind it to a powder. The latter

was then transferred into a 12 x 75 mm glass storage vial, with threaded screw cap

(Figure 27). The glass vials containing the Na-polysulfides were packed under nitrogen,

in zipper-lock bags, and placed in the fridge until needed for further analyses.

The final Na-polysulfides were sent to an external laboratory for X-Ray Diffraction

(XRD) (Pananalyical Expert Pro, Pixcel Detector) analysis for purity. The system was

operated at 45 kV, 40 mA, in which the incidence angle spanned from 6° to 140°2θ,

under argon. A copper anode was used in this analysis.

59

Figure 27. Pictures of the synthesized Na-polysulfides: (A) Na2S2; (B) Na2S3; (C) Na2S4

2.2.2 N,N-Diethyl-p-phenylenediamine (DEPD) Analysis of Sulfide

The spectrophotometric method for the analysis of sulfide using DEPD was adopted from

another study by Kariuki et al. (2008). Into a 25-mL volumetric flask, a small amount of

the alkaline ultra-pure water, which had been previously purged with nitrogen, was added

to a 25-mL volumetric flask. A 250-μL aliquot of the 106 mM sulfide stock solution was

then added to the flask, along with 2-mL of the color developing reagent, prepared as

described in Appendix A. The alkaline ultra-pure water was used to top up the flask to

the 25-mL total volume. The 1 mM sulfide solution was left to react for a 30-minute

period. The resulting solution was used to create 25, 15, 10, 5 and 2.5 µM of sulfide. The

DEPD spectrophotometric method is analogous to the Methylene Blue method whose

main reaction is shown below. The diluted solutions were immediately analyzed using a

UV-VIS spectrophotometer (GENESYS, Thermo Scientific) at a fixed wavelength of

670 nm, using a rectangular quartz glass cuvette with a path length of 1 cm (Agilent

Technologies). A blank was prepared by using nitrogen-purged ultra-pure water.

2.2.3 Reduction of Na-polysulfides and polythionates Using

Chromium as the Reducing Agent

The set-up for the reduction of the polysulfides and the thiosulfate ion is shown in Figure

28. The reduction for these compounds was adopted from a procedure reported by

Kariuki et al. (2008). In brief, a purge and trap set-up consisting three 125-mL

60

Erlenmeyer flasks were connected in series. A 50 mL aliquot solution of the sulfur-

containing compound being reduced was placed in the flask labeled A. If thiosulfate was

being reduced, the solution was made up in water. The polysulfides were made up in

0.1 M NaOH solution. To minimize air oxidation of the polysulfides, their solutions were

made up in the Baxter glove box under a nitrogen rich environment. Before removal

from the glove box, the solutions were sealed with a rubber stopper and parafilm prior to

removal from glove box for reduction.

The 0.1 M NaOH solution was pre-purged with nitrogen gas prior to using with

polysulfides. This was to limit any prior exposure to oxygen in order to control the rapid

oxidation of the polysulfides. The thiosulfate solution was made up in pre-purged ultra-

pure water. The thiosulfate is a very stable product and the need for high pH solutions, as

is the case for simple sulfides and polysulfides, was not required for this compound. In

each of flasks (B and C) 60 mL 0.1 M NaOH was placed. Three, two gas-tube stoppers

were connected with Tygon tubing and placed on the top of each of the flasks (Figure

28). The third flask (C) in the sequence (see Figure 28) was intended to function as an

extra ‘catch flask’. If recovery of polysulfides from flask (B) were to be <100 %,

following reduction, some of that reduced solution might escape into flask (C) and be

trapped there. Nitrogen gas was connected to the inlet source tube for reaction flask (A),

and a gentle stream of gas flow was initiated. Flasks B and C were connected first,

wrapped in parafilm, and quickly, the third stopper was connected to flask A. All three

flasks were purged for 30 minutes. With the nitrogen gas still flowing, approximately

1.50 g sample of chromium metal was weighed, added to reaction flask (A), and then 7

mL of concentrated hydrochloric acid was also added to the same flask. The gas-tube

stopper was quickly placed on the top of the flask, immediately sealed with parafilm, and

continuously purged for 4 hours, at 60 °C while the reaction completed. After the

reaction was complete, working under a blanket of nitrogen gas, flasks B and C were

removed and the reduced solutions were poured, individually, into a 60 mL Teflon

storage bottle (Nalagene®, Fisher Scientific). The DEPD reaction, as outlined in section

2.2.3, was conducted on reduced solutions captured in flask B. These reduced solutions

were also used for electrochemical experiments, which will be described in Chapter 3.

61

Figure 28. Purge and trap system set-up for the reduction of aqueous Na-polysulfides

(Sn2-) and thiosulfate (S2O3

2-): flask (A) (Sn2- or S2O3

2-); flasks B and C (traps for reduced

sulfide).

2.3 RESULTS

2.3.1 Making Na-polysulfides

As described in the procedure section, 2.2.1 the synthesis the Na-polysulfides involved

several stages. One of the stages involved the sealing of the reaction mixture in

borosilicate glass tubes, before placing the tubes in a muffle furnace. Extra care was

required to ensure that the sealing of the glass tubes was perfect. This appeared to have

been a delicate step because only about 50% of sealing attempts ended up being

successfully done. Also, due to the low moisture environment inside the glove box, the

glass tubes became highly charged with static electricity. When the Na2S and S

homogenized mixtures were transferred from the mortar to the glass tube, some of the

mixture tended to adhere to the side of the tube rather than settling at the bottom of the

tube. An attempt was made to circumvent this problem by having the powdered mixture

delivered through a rolled 15.24 x 15.24 cm weigh paper (Fisherbrand, Fisher Scientific).

This helped deliver most of the mixture to the bottom of the tube, rather than sticking to

the tube walls. Any of the mixture retained on the walls would not mix with the rest of

the reactant mixture that did make it to the bottom of the tube. That could have reduced

the yields below 100 % of the final Na-polysulfide product. The colors of the Na2S2,

62

Na2S3, and Na2S4 synthesized in our lab were yellow, orange or yellow-orange, and olive

green, respectively (Figure 27). The colors for Na2S2 and Na2S4 agreed with those

reported by Rosen and Tegan (1971). When the polysulfides were ground into a powder-

like form some of the physical characteristics between the different polysulfides were

quite evident. This was the case when the trisulfides were extracted from their glass; they

did not grind into a dry-like powder, but more crystalline and seemed to have more

moisture. This observation, for the trisulfide, did not seem to be affected by the humidity

levels in the lab environment or by the relative humidity levels inside the glovebox. The

disulfide and tetrasulfide compounds ground easily into a dry, powder-like form.

The standardization of the polysulfides through iodometric titration (Section 1.3.2) and

the DEPD method (Section 2.2.2), without first reducing them to S2-, was unsuccessful. It

is believed that the white cloudy precipitate that formed during the use of the two

standardization methods impaired the ability of the apparatus to accurately measure and

standardize the aqueous Na-polysulfides. Specifically, in the DEPD experiment, the

cloudy precipitate blocked the incident light in the spectrophotometer from passing

through the cuvette. To circumvent the problem of precipitate formation, the aqueous

Na-polysulfides needed to be reduced to the simple sulfide ion (HS- or S2-), as described

in section 2.2.3, prior to analysis using DEPD. Such a reduction of the polysulfides to

sulfide also allowed the use of the iodometric titration for the standardization of the

polysulfides.

A comparison of two separate batches of Na-polysulfides following reduction is outlined

below in Table 10. As the results indicate, the purity of the disulfide, trisulfide, and

tetrasulfide prepared in January 2016 was 94%, 87%, and 89%, respectively. This purity

level of the polysulfides was better than of the ones prepared in August 2015 by 9%, 7%,

and 5% for disulfide, trisulfide, and tetrasulfide, respectively. There is a possibility that

the relative humidity at the time of synthesizing the polysulfides could have impacted the

purity of the synthesized polysulfides. There is usually higher humidity in the lab in the

month of August than in January.

63

Table 10. Comparison mean % yields ± SE of reduced Na-polysulfides from two separate

batch products, summer 2015 and winter 2016.

Polysulfide August 2015 January 2016

Na2S2 85 ± 7 94 ± 3

Na2S3 80 ± 3 87 ± 7

Na2S4 84 ± 19 89 ± 5

A typical relative humidity range in August is anywhere between 45 - 55 %; compared to

January relative humidity levels which tend to be between 18 - 20%. An increased level

of relative humidity in the lab environment may lead to an increased level of moisture in

the air. Sodium sulfide and the Na-polysulfides are very hygroscopic. As reported by

Petri and Larachi (2006), excess moisture could lead to loss of these types of compounds

due to oxidation to hydrogen sulfide as a gas. Steps were taken to ensure humidity levels

were reduced as far as possible by running a de-humidifier and split air-conditioning unit.

However, the two external units were only able to reduce the humidity to about 45 % in

the lab space. As a result, even working inside the glove box and manipulating the

samples during the month of August occurred in an environment where the relative

humidity could only be reduced to 30 - 35 %. This is much higher than measurements

made in January when the relative humidity was between 20 - 25 %.

The Na-polysulfides were analyzed using XRD, under conditions as outlined above.

Unfortunately, due to the low signal to noise ratio, the final products could not be

quantified. There was also the presence of some impurities such as disodium sulfite

found within these samples. Some of these impurities did not coincide with the oxidative

products described in appendix B or listed in Table 12. This may suggest the existence of

side reactions during the synthesis of the polysulfides at high temperatures. The

quantification of the side products was not determinable using XRD.

2.3.2 Reduction and Analysis of Na-polysulfides and Polythionates

According to a number of authors (Cline, 1969; Lawrence et al., 2000; Kariuki et al.,

2008; Reese et al., 2011), sulfide, the product of reduction of sulfur containing

compounds with chromium, gets protonated in an acidic medium to give off H2S gas

(Reaction R-2.4). The H2S gas generated, when collected in an alkaline solution, reverts

64

back to S2− which in turn reacts with DEPD to give an intense blue solution. The

intensity of the blue colour, from the DEPD reaction, is proportional to the concentration

of S2− in the solution being analysed.

𝑆2−(𝑎𝑞) + 2𝐻+(𝑎𝑞) → 𝐻2𝑆(𝑔) (𝑅 − 2.4)

Figure 29. Calibration curve of sulfide obtained using the DEPD method (R2 = 0.999).

Table 11. Mean and SE ± for various concentrations of sulfide used to develop a standard

curve for DEPD analyses (n=11)

Concentration (µM) Mean Absorbance ± SE

25 1.5 ± 0.07

15 0.91 ± 0.04

10 0.63 ± 0.03

5.0 0.32 ± 0.02

2.5 0.17 ± 0.01

Figure 29 shows the calibration curve obtained by plotting the concentrations of sulfide

in the DEPD reaction versus absorbance at a wavelength of 670 nm. A correlation with

an R2 value of 0.999 was obtained for the sulfide analysis through the DEPD method. To

account for possible matrix affects, a method blank was prepared using ultra-pure water

and 0.1 M NaOH. Both were separately put through the reduction process and run

through DEPD and iodometric titration. Both the ultra-pure water and 0.1 M NaOH

0

0.2

0.4

0.6

0.8

1

1.2

1.4

1.6

0 3 6 9 12 15 18 21 24 27

Ab

sorb

an

ce a

t 6

70

nm

Concentration of Sulfide (µM)

65

showed no presence of sulfide after analysis through the DEPD and iodometric titration

methods.

Earlier attempts to reduce the aqueous Na-polysulfides, specifically the tetrasulfide

compound, were not very promising. When the reduction was run at room temperature

for 1.5 hours, the yield was ≤ 80 % conversion of the aqueous disulfide and trisulfide

polysulfide compounds to their reduced forms of Sn2-. There was ≤ 60 % conversion of

the tetrasulfide compound under these conditions. After the reduction time had ended,

and the solutions from flask B and C were recovered, an odor of H2S was apparent from

reaction flask A. The presence of this characteristic “rotten egg” odor indicated that not

all of the aqueous Na-polysulfide had been captured in flask B and converted to their

simple sulfide form. This was confirmed after completing the DEPD analysis (section

2.2.3) when the recovery values of Sn2- were shown to have been ≤ 80 % for the disulfide

and trisulfide compounds, and ≤ 60 % for the tetrasulfide. As noted by Kariuki et al.

(2008), the increase in chain length of the aqueous Na-polysulfide makes the reduction of

these compounds more difficult under these conditions. The method was modified in

terms of applying heat to 60 °C, and the reduction of the polysulfides was set up to occur

over a 4-hour period rather than the original 1.5 hours. The heat source was turned off

after 3.5 hours and the N2 gas flow was increased slightly as the reaction solution cooled

to room temperature. This was done to expel any remaining reduced Sn2- from reaction

flask A before retrieval of the solutions for further analyses. The additional heat

treatment and reaction time caused the % yields of recovered Sn2- to increase by 19, 12

and 40% for disulfide, trisulfide and tetrasulfide, respectively. Table 12 summarizes the

% yields of the reduced Sn2- products after applying heat at 60 °C and reducing the

solution for 4 hours. With this modification of the procedure there was no smell of H2S

from reaction flask A after applying the reduction treatment. As indicated below in Table

12, there are some remaining challenges in terms of getting closer to the desirable level of

100 % conversion of reduced Sn2- of the trisulfide and tetrasulfide.

66

Table 12. Mean percent yields ± SE of the reduced polysulfides after applying heat at 60

°C and reducing the solution for 4 hours (n = 5).

Na-Polysulfide Mean ± SE

Na2S2 [S22-] 93 ± 5

Na2S3 [S32-] 85 ± 4

Na2S4 [S42-] 83 ± 5

The percentage purities of the synthesized polysulfides using the DEPD method were in

agreement with those obtained through the iodometric titrations. In order to directly

compare the two sets of results, an Analysis of Variance (ANOVA) was carried out to

compare the effect of each method on reduced sulfide compounds. There was no

significant difference between DEPD and iodometric titration in terms of the

determination of reduced sulfide compounds (F (1,10) = 0.418, p = 0.532). This supports

the conclusion that the DEPD and iodometric titration methods are in agreement, and that

the DEPD method provides a reliable means for determining the concentrations of

reduced sulfur compounds.

2.4 CONCLUSIONS

The several steps required to make the Na-polysulfides can be a source of significant

errors, which in turn may lead to incomplete product formation. Through several batch

adjustments attempts were made to optimize successful product formation. The Na-

polysulfides made in January 2016 have proved the best products to date. The early

round of testing of the final Na-polysulfides in August 2015 was challenging; for

example, when the Na-polysulfides were added to the supporting electrolyte solution

(Appendix A) at a pH of 10.3, a fine black to grey-ish precipitate formed instantly. It has

been suggested that the polysulfide molar distribution is pH dependent, with tetrasulfide

being dominant at pH ≈ 8.0 with trace amounts of trisulfide; and at a pH ≥12.0 trisulfide

disulfide becoming the predominant species (Petre and Larachi, 2006). Also, when

testing the aqueous Na-polysulfides in alkaline media (pH > 12.0) against DEPD method,

with no reduction analysis done, the polysulfide ions in the acidic environment

precipitated a white cloudy substance that only allowed for a slight blue coloration to

form. Sonne and Dasgupta (1991), showed that where H2Sx has a value of x > 1, the

67

compounds are unstable in acidic solutions, thereby producing elemental sulfur and

hydrogen sulfide. It is likely that this was the case when the Na-polysulfide prior to the

reduction process were subjected to the DEPD and iodometric titration methods.

An improvement that can be made to the study of reduced aqueous Na-polysulfides is to

vary the pH of the trap solution (Flask B) or vary the pH during the initial mixing of the

Na-polysulfides prior to conducting the reduction process. Heating the reaction flask (A)

increased the rate of conversion (Table 12). Another consideration could be the use of

other reducing agents besides chromium that would have the potential to increase the

conversion of Na-polysulfides to simple sulfide (HS- or S2-).

The exact purity of the Na-polysulfides could not be determined though XRD analysis.

Some of the deviation from the high purity of the synthesized polysulfides as determined

through the XRD analysis could have been as a result of a significant time delay from

when the samples were synthesized to the time when they were analyzed through XRD. It

is likely that side reactions could have been occurring during storage of the samples. In

future, the ideal situation would be the XRD analysis of the polysulfides shortly after

their synthesis.

A drawback was recognized when using the DEPD (section 2.2.3) and iodometric

titrimetric (Section 1.2.3) methods for the reduced Sn2- or SnO2

2- solutions. The two

methods do not allow differentiation between the different reduced Sn2- or SnO2

2- species

if they are all present in solution at the same time. A further investigation of the methods

of analysis for the polysulfides and polythionates might lead to the speciation of these

compounds. In the following chapter, electrochemical analyses will be used to

investigate whether these aqueous Na-polysulfides and polythionates can be readily

differentiated.

68

Chapter 3

Electrochemical Analysis of Na-polysulfides and Polythionates

3.1 INTRODUCTION

Direct determination of polysulfides and polythionates still remains a challenge. This is

partially due to their high redox reactivity in the natural systems 10 - 20 µM (Kamyshny,

2006; Chadwell et al., 2001). Polysulfide ions are subject to autoxidation when present

in solution. This autoxidation reaction tends to be rapid, when in solution, but when the

polysulfides are in a solid state the reaction is slower (Kamyshny, 2006; Steudel, 2003;

O’Reilly et al., 2001; Sonne and Dasgupta, 1991). Most of the previous attempts to

distinguish and separate polysulfide and polythionates in environmental samples use

capillary electrophoresis (Petre and Larachi, 2006), ion chromatography (Miura et al.,

2005), titrimetric methods (Kamyshny et al., 2004), and electrochemical analyses (Manan

et al., 2011; Kariuki et al., 2001; Rozan et al., 2000b). Ion chromatography and capillary

electrophoresis often require pH changes, which could alter the speciation of sulfur based

compounds present in samples. Derivatization using substances like methyl iodide are

sometimes required for analysis of samples using the above techniques. The uses of

modifiers, such as a chromate electrolyte in a solution of hexamethonium bromide, have

also been used towards the determination of polysulfides and polythionates in samples.

These changes to pH, derivatizations, and added modifiers can alter the relative

abundance of the sulfur-based species originally present in a given environmental sample.

For example, S2O32- under acidic conditions will decompose to elemental sulfur and

sulfite (SO32-) (Ciglenečki and Ćosović, 1997). This distribution change seems to alter

the sulfur-based compound speciation in the original sample. Kamyshny et al. (2004),

reviewed the various published attempts of the polysulfide speciation, concluding that the

speciation and distribution of polysulfides and polythionates seem to vary widely from

lab to lab. These authors also pointed out that not one lab used the same technique or set

of techniques to determine sulfur based compounds in a consistent manner. However,

these authors also identified the thermodynamic constants (pKn) for polysulfide

disproportionation which could be useful in understanding polysulfide and polythionate

69

kinetics in solution. They also developed a new approach to determining the polysulfide

speciation in aqueous media. Their results show a narrower range of disproportionation

to polysulfide speciation by accounting for variations of the pH of the solutions that the

polysulfide was dissolved in. Their results support the findings of Petre and Larachi

(2006), where they classify the abundance of the polysulfide species present in solution

over a pH range from 8.0 to 12.0.

Electrochemical analyses for the detection of polysulfides using mercury drop electrodes

(Chadwell et al. 2001; Rozan et al. 2001) showed that there is an interaction between

polysulfides and bisulfides with divalent cations such as iron, nickel, cobalt, copper and

zinc. Both groups of authors concluded that a metal complex was formed through the

decomposition of the polysulfide to bisulfide in solution, rather than through direct

complexation with the polysulfide ion. This might suggest that in characterizing

environmental samples, one should be aware that the matrix of the solution affects the

determination of polysulfide concentrations. Also, other authors who have used the

polarographic electroanalytical techniques for the analysis of sulfur compounds have

observed a significant interference from higher order polythionates (Steudel 2003;

Kariuki et al. 2001; Rozan et al. 2000).

This chapter reports on the electrochemical analyses of aqueous solutions of the Na-

polysulfides (Na2S2, Na2S3 and Na2S4). In addition, the analysis of thiosulfate as a

representative of the class of polythionate compounds is also presented. Cyclic

voltammetry using both the modified and the unmodified glassy carbon electrodes used

for the characterization of the polysulfides and polythionate in alkaline media is also

discussed. The electrochemical analyses for the polysulfides and polythionate presented

here will demonstrate that the mixture is best analyzed when the sulfur compounds get

converted to simple sulfide.

70

3.2 MATERIALS AND METHODS

3.2.1 Equipment

Electrochemical measurements were conducted using the BASi work station with Epsilon

USB software (Bioanaytical System Inc.). The three-electrode cell consists of a C3 glass

cell vial 50 mm x 59 mm housing the supporting electrolyte solutions, and electrodes for

analysis. The working electrode was a glassy carbon electrode with a working surface

diameter of 3.0 mm (area = 0.017 cm2). The auxiliary electrode used was a 23-cm long

coiled platinum wire. An Ag/AgCl probe was used as the reference electrode. All

electrodes were supplied by BASi. In each case the electrochemical measurements

(section 1.2.4.5) of the Na-polysulfides, the polythionate, and reduced polysulfides and

the polythionate solutions were conducted using cyclic voltammetry. Cyclic

voltammetric cycles were operated by scanning the electrode potential from -1600 mV to

+1600 mV, at a scan rate of 100 mVs-1 for 10 cycles. The peak potential (mV), height

(µA), and area (µC) were recorded for all 10 cycles. Three sets at 10 CV cycles per

replicate were analyzed for each experiment. MGCE experiments were also carried out

on the above solutions using the procedures outlined in section 1.2.4.4.

3.2.2 Dissolution of Na-polysulfides and Thiosulfate Ions

3.2.2.1 Na-Polysulfides

The final products selected for electrochemical analysis of the Na-polysulfides that were

synthesized as outlined in section 2.2.2.5 were those that generated the highest percent

yield following the chromium-reduction process described in section 2.2.4. The weights

of the various Na-polysulfides that were used are shown in Table 13. Each compound

was dissolved in 20 mL 0.1 M NaOH as described in section 2.2.3.1, stored in a 30 mL

Teflon storage bottle (Nalgene®, Fisher Scientific), and placed in the fridge at 4 °C until

required. Due to the instability of the Na-polysulfides, once they were dissolved, they

had to be used within 48 hours. All procedures described here were carried out inside the

Baxter Glovebox as described in section 1.2.1.2.

71

Table 13. Weights (g) and concentrations (mM) of Na-Polysulfides used for CV analyses

Na-Polysulfide Weight (g) Concentration (mM) in 20 mL 0.1 M

NaOH Na2S2 0.149 67.8 Na2S3 0.175 61.4 Na2S4 0.175 50.1

3.2.2.2 Polythionates (S2O32-)

The thiosulfate ion was used to represent the polythionates group of sulfur compounds.

For the thiosulfate ion, 0.539 g of sodium thiosulfate was weighed and dissolved in 20

mL of ultra-pure water, placed in a 20 mL glass scintillation vial (Fisherbrand, Fisher

Scientific), and stored in the fridge at 4 oC until required later for analyses. The

thiosulfate concentration in this solution was 109 mM. The thiosulfate ion is not readily

oxidized by exposure to air, so the preparation (weighing and dissolution process) was

carried out on the bench top rather than in the glove box.

3.3 RESULTS

3.3.1 Electrochemical Analysis of Na-polysulfides and Polythionates at

BGCE and MGCE

Cyclic voltammetry was performed using each of the BGCE and MGCE for the aqueous

Na-polysulfides (Na2S2, Na2S3, Na2S4), the polythionate (thiosulfate), as well as for the

reduced polysulfide solutions (section 2.2.4). For the BGCE experiments, the three Na-

polysulfides analyzed are shown in Figure 27. The Na2S2 showed no peak current

response through the potential scan from -1600 to +1600 mV. Indeed even when the

Na2S2 concentration was doubled relative to that shown in Figure 30 (1.39 mM), no

current response was observed.

72

Figure 30. CV graph at a potential scan rate of 100 mV s-1 for S22- [1.39 mM], S3

2- [1.31

mM] and S42- [1.22 mM] at BGCE. With supporting electrolyte solution as the blank

(dashed line).

It is unclear why there is no CV response for the S22- ion. Some researchers have

suggested that at extremely high pH values (≥ 12.0) the disulfide ion becomes the

predominant species (Petre et al., 2006 and Steudel, 2003). However, in this

electrochemical analysis the pH of the supporting electrolyte solution was 10.3, which is

significantly below the pH of 12.0 that they were considering. Neither our own work, nor

other published reports clearly address whether the disulfide ion breaks down to other

sulfur species when the solution is at a pH < 12.0. It was, however, observed that after

the analysis on the BASi equipment, the electrochemical cell containing the disulfide ion

released the characteristic odor of H2S. This may indicate that a reduced volatile sulfide

was present, but perhaps not in quantities that could be detected by this particular

analytical method. It is worth noting that humans are particularly sensitive to the odor of

H2S, and some people can detect the gas at levels as low as 0.47 ppb. There were,

however, clear current responses for both S32- and S4

2- ions (Figure 30), occurring in the

-100

-75

-50

-25

0

25

50

75

100

125

150

175

-1600 -1200 -800 -400 0 400 800 1200 1600

I /

µA

E / mV vs. Ag/AgCl

blank

S22-

S42-

S32-

73

anodic and cathodic segments. The values for their potential position(s) and current

response(s) are summarized in Table 14 and Table 15 with the values for HS- ion added

for comparison.

Table 14. Anodic segment peak potential (mV) and peak current (µA) for CV analysis of

Sn2- ions (1 mM) at BGCE (n = 3).

Na-Polysulfide Epa (mV) Ipa (µA)

S22- ― ―

S32- -8, +1050 17, 30

S42- -15, +1055, +1339 21, 28, 18

HS- -10, +412, +1114 18, 69, 87

Table 15. Cathodic segment peak potential (mV) and peak current (µA) for CV analysis

of Sn2- ions (1 mM) at BGCE (n = 3).

Na-Polysulfide Epc (mV) Ipc (µA)

S22- ― ―

S32- -1135 43

S42- -1100 71

HS- -1200 115

In Appendix A, the table for oxidation states of various sulfur compounds, provides some

insight as to which species could be produced through the oxidation of S32- and S4

2- ions.

Some of the potential positions for both the anodic and cathodic segments shown in

Figure 32 match the CV analyses conducted on sulfide in Figure 3. However, there is a

less marked rise in the current baseline in the tri- and tetrasulfide CV graphs at the Epa

positions at ≈ +1060 and +1340 mV, compared to the response observed for the sulfide

CV (Figure 31). The current responses in the tri- and tetrasulfide CV graphs are easier to

integrate at potentials of ≈ +1060 to +1340 mV due to the less marked rise in baseline

current in these regions of high potential. The current responses of HS- ion at potentials

of +1600 mV are twice as high as the current response of the tri- and tetrasulfides at a

potential of +1600 mV (Figure 31). All of the concentrations examined for the ions of

Sn2- and HS- comparison were close to 1 mM.

74

Figure 31. CV graph of the comparison of simple sulfide ion with di-, tri- and tetrasulfide

ions at a potential scan rate of 100 mV s-1 for HS- [1 mM], S22- [1.39 mM], S3

2- [1.31

mM] and S42- [1.22 mM] at BGCE. With the supporting electrolyte solution as the blank

(dashed line).

When thiosulfate was analyzed at the BGCE there were only two peak current responses

in the anodic segment of the CV trace. There were no peak(s) present in the reductive

segment (Figure 32). A 1.1 mM solution of S2O32- in the electrochemical cell gave

potential positions Epa = +1037 and +1415 mV with a current response of Ipa = 42 and 59

µA, respectively. A check was done to see if the S2O32-

could, on its own, be

quantitatively determined on a BGCE (Figure 33). This was, indeed, found to be the

case. However, when thiosulfate was combined with sulfide, the potentials at which

oxidation of the two species appeared were indistinguishable (Figure 34). Overall the

results shown in Figure 32 and Figure 33 suggest that thiosulfate can be detected on its

own, but when it is mixed with sulfide its response cannot be readily distinguished from

that of the HS- ion. In contrast with that result, Figure 32 and Figure 35 show that when

-150

-100

-50

0

50

100

150

200

250

300

350

-1600 -1200 -800 -400 0 400 800 1200 1600

I /

µA

E / mV vs. Ag/AgCl

HS-

blank

S22-

S32-

S42-

75

S2O32- is mixed in with a Sn

2- such as S32-, the S2O3

2- can still be individually

distinguished from the baseline current.

Figure 32. CV graph of the comparison of quantitative amounts of S2O32- at a potential

scan rate of 100 mV s-1 for S2O32- [1.06 mM], and 2 x the amount of S2O3

2- [2.12 mM] at

BGCE. With the supporting electrolyte solution as the blank (dashed line).

-50

-25

0

25

50

75

100

125

150

175

200

225

250

275

-1600 -1200 -800 -400 0 400 800 1200 1600

I /

µA

E / mV vs. Ag/AgCl

blank

S2O32-

(2X) S2O32-

76

Figure 33: CV graph of the comparison of polysulfide ions with thiosulfate ion at a scan

rate of 100 mV s-1 for S2O32- [1.06 mM] (dotted line), S2

2- [1.39 mM], S3

2- [1.31 mM] and

S42- [1.22 mM] at the BGCE. With supporting electrolyte solution as the blank (dashed

line).

-100

-75

-50

-25

0

25

50

75

100

125

-1600 -1200 -800 -400 0 400 800 1200 1600

I /

µA

E / mV vs. Ag/AgCl

blank

S2O32-

S22-

S42-

S32-

77

Figure 34. CV graph of the comparison of thiosulfate ion, sulfide ion and a combination

of thiosulfate ion and sulfide ion. At a potential scan rate of 100 mV s-1 for S2O32- [1.06

mM], sulfide [1 mM], and mixture of HS- + S2O32- at BGCE. With the supporting

electrolyte solution as the blank (dashed line).

According to Steudel (2003) polysulfide anions are subject to autoxidation in the

presence of molecular oxygen. Reaction R – 3.1 shows the S42- example. In the presence

of molecular oxygen, polysulfides autoxidize to form thiosulfate and sulfur (R – 3.1).

Even though measures had been taken in this experiment to exclude oxygen, the

potentials that are used in this work can help further explain the source of oxygen that

may readily react with polysulfides, as shown in reaction R – 3.1. Anodic oxidation of

sulfide can yield products like elemental sulfur, polysulfides, thiosulfates and sulfates

(Al-Kharafi et al. 2010). This may explain the two peaks observed at potentials of +1050

and +1339 mV respectively (Figure 30 and Table 14).

𝑆4 2− +

3

2 𝑂2 → 𝑆2𝑂3

2− + 1

4 𝑆8 (𝑅 − 3.1)

This proposed mechanism and explanation is further supported by the results of Lessner

et al. (1993), where spectroscopic and chemical investigations revealed that in alkaline

-150

-100

-50

0

50

100

150

200

250

300

350

400

450

-1600 -1200 -800 -400 0 400 800 1200 1600

I / µ

A

E mV Vs. Ag/AgCl

(HS- + S2O32-)

HS-

S2O32-

blank

78

pH ranges, the polysulfide and bisulfide species predominate. This can be represented by

the following reaction (R – 3.2):

4𝑆42− + 8𝑂𝐻− + 𝐻2𝑂 ↔ 3𝑆2𝑂3

2− + 10𝐻𝑆− (𝑅 − 3.2)

According to the reaction R – 3.1, as the CV cycles are continuously run on the BGCE, in

the presence of sulfur-based compounds, elemental sulfur builds up on the working

surface of the electrode, thereby reducing the area of activity. It should be noted that

thiosulfate on its own does not exhibit a reductive segment. This may suggest that in the

case of thiosulfate there is no reduction to form HS- and that the thiosulfate ion is

oxidized to tetrathionate (S4O62-). According to O’Reilly et al. (2001) and Koh (1990),

there is no production of elemental sulfur in the electroxidation process that could impair

the working surface of the glassy carbon electrode.

Figure 35. CV graph of the comparison of thiosulfate ion, trisulfide ion and a

combination of thiosulfate ion and trisulfide ion at a potential scan rate of 100 mV s-1 for

S32- [1.31 mM] and a mixture of S3

2- [1.31 mM] plus S2O3

2- [1.06 mM], and S2O32- [1.06

mM] at the BGCE. With supporting electrolyte solution as the blank (dashed line).

-60

-20

20

60

100

140

180

220

-1600 -1200 -800 -400 0 400 800 1200 1600

I /

µA

E / mV vs. Ag/AgCl

blank

S2O32-

S32-

S32- + S2O3

2-

79

As illustrated in Figure 36, the analysis of the Sn2- ions at MGCE’s exhibit similar

behavior to that of HS-. After background corrections were adjusted at Epa ≈ +83 mV

(LOD = 28 µA at MGCE), the HS-, S22-, S3

2-, S42- ions have a current response of 50 µA,

22 µA, 35 µA, and 43 µA respectively. It should be noted that there is a shift in potential

to the left of the Sn2- and S2O3

2- ions relative to the blank reference at an Epa of +249 mV.

However, there is no apparent shift in potential for the current responses located at a

mean Epa of +1138 mV (n = 5) relative to the blank at a Epa of +1107 mV. Based on the

proposed formation of a complex of V2O5 ·HS- at MGCE (Khudaish and Al-Hinai 2006)

(Equation R – 3.3), the interactions of the Sn2- with the MGCE surface may be explained.

𝑉2𝑂5 ∙ 𝐻𝑆− → 𝑉2𝑂4 + 𝑆 + 𝑂𝐻− (𝑅 − 3.3)

With regard to the activity of the Sn2- ions at the MGCE, the present study may suggest

that the reaction that may have occurred can be described by the reaction R – 3.4. This

hypothesis is based on the mechanism proposed by Khudaish and Al-Hinai (2006), and

shown as R – 3.3.

𝐻2𝑂 + 𝑉2𝑂5 ∙ 𝑆𝑛2− → 𝑉2𝑂4 + 𝑛𝑆 + 𝑂𝐻− (𝑅 − 3.4)

Figure 36 clearly shows that it would be very difficult to attribute the peak responses

from a series of CV scans on an MGCE, separately for HS-, S2O32-, and Sn

2- ions present

in a mixture. This observed overlap with all these sulfur anions poses a problem for the

application of this modification agent towards environmental and industrial samples.

Samples that come from environment and industrial waste basins would be more likely to

contain a mixture of these substances rather than pure solutions that were individually

analyzed in Figure 36.

80

Figure 36. CV graph at potential scan rate at 100 mV s1- at the MGCE for S22-, S3

2-, S42-,

S2O32-, and HS- (dotted line). With the blank as the supporting electrolyte solution

(dashed line).

3.4 CONCLUSIONS

At the BGCE, S32- and S4

2- ions can be detected, but they do not show unique, distinctive,

or characteristic potential positions during CV cycles. According to Lessner et al.

(1993), thiosulfate can be a side product formed from the dissociation of polysulfide ions

in solution. Thiosulfate, when mixed with a polysulfide, appears to show the same

potential positions, between +1000 to +1300 mV (Figure 32). The thiosulfate ion can be

distinguished when it is in a pure solution rather than being part of a mixture, and it is

unique in that it has no reductive segment. The CV results of the thiosulfate ion at the

BGCE suggests that under the CV conditions described above, only the oxidation peak,

and not the reduction peak, is observed. This observation also agrees with that of several

other authors (Zanello et al., 2012; Scholz, 2010; Kissinger and Heineman, 1996). As

observed in this chapter, there are several potential positions that overlap in the CV

-100

-75

-50

-25

0

25

50

75

100

125

150

175

200

-1600 -1200 -800 -400 0 400 800 1200 1600

I / µA

E / mV vs. Ag/AgCl

S22-

S32-

S42-

S2O32-

HS-

blank

81

graphs for the S32- and S4

2- ions. There are several potential positions that overlap for the

S32- and S4

2- ion CV graphs. This overlap in potential positions makes these two

compounds hard to distinguish, which would pose a problem when analyzing

environmental samples since these samples would likely contain a mixture of these

compounds. The problem is particularly apparent when Sn2- ions are combined with HS-

(Figure 31 and Figure 34). There is a catalytic response to Sn2- and S2O3

2- ions at

concentrations of about 1 mM, but their potential positions did not vary enough during

CV analyses to allow for distinction between these two classes of sulfur species. The

interference of the supporting electrolyte solution during the experiment at MGCE would

suggest that at concentrations lower than 1 mM for the Sn2- and S2O3

2-, the current

response of these ions would be unquantifiable. However, this phenomenon was not

observed in reactions carried out using BGCE. As was the case with the BGCE, the

potential positions and current responses of the Sn2-, HS-, and S2O3

2- ions were very

similar and therefore, indistinguishable under the CV conditions used in this work. The

unexpected fluxes in terms of the increases and decreases in the current recorded over a

series of ten CV cycles at the MGCE (Figures 22 to 24) may need further investigation.

Earlier studies have successfully determined quantifiable polysulfides in solution

mixtures in combination with thiosulfate and sulfide ions (Kariuki et al. 2001; Rozan et

al. 2000). However, these analyses were carried out using a mercury drop electrode,

which provides a fresh working surface with each sequential drop. With the mercury

drop electrode, there was no need for manual manipulation or electrode cleaning as was

the case with the glassy carbon electrodes used in this study, where the surface had to be

polished between analyses.

Although the polysulfides could not be distinguished from each other using the BGCE

and MGCE, valuable qualitative data was gathered regarding the individual ions and how

they might interact during CV analyses.

82

Chapter 4

Analysis of Low Oxidation-State Sulfur Species in Water

4.1 INTRODUCTION

Sulfur compounds are prevalent in environmental and industrial processes. Substances

such as sulfide, thiosulfate, polysulfides, sulfate and sulfite can be generated by natural

processes (biological or geochemical) or by anthropogenic activities such as petroleum-

based or mining industries (Casella et al. 2002). Identification and quantification of these

sulfur compounds is important due to the health risks associated with exposure to

hydrogen sulfide and the effects these compounds have in aquatic systems (Kariuki et al.

2008; Lawrence et al. 2007; Lawrence et al. 2002; Witter and Jones 1998; Zhang and

Millero 1993). As previously stated, the levels of dissolved H2S at 0.01 mg L1- have been

reported to have a chronic toxicity to aquatic organisms (Brouwer and Murphy, 1995).

Levels of dissolved H2S approaching 0.5 mg L1- exhibit acute toxicity to rainbow trout

and salmon (Ortiz et al., 1993). The sulfide anion can be rather versatile in that it forms

intermediate species of sulfur based compounds. These intermediate sulfur based

compounds include but are not limited to: thiosulfate, sulfite, sulfate, and acid volatile

sulfides. According to Gun et al. (2000), polysulfides are an integral part of the sulfur

cycle and they are required for the formation of volatile sulfur compounds, which are a

group of organic volatile sulfur species such as dimethyl sulfide. Also, the polysulfide

ions have the ability to disassociate to form reduced inorganic sulfur species like those of

HS- and H2S during the organic decomposition of R-SH groups prompted by sulfur

reducing bacteria (Gosh and Dam, 2009; Kamyshny et al., 2008; Kamyshny et al., 2006;

Friedrich et al., 2001). The standard detection methods in use were outlined in an

extensive review by Lawrence et al. (2000) and were briefly summarized in the

background information. Among those detection methods, electrochemical analyses have

become more popular for the detection of environmental contaminants such as sulfide

species. Examples of electrochemical methods used for sulfide determination are

potentiometry (using an ion-specific electrode), voltammetry and amperometry (Ardelean

et al., 2014).

83

In this chapter, a water sample collected from Callander Bay, Ontario was examined

using the methods described in the earlier chapters. The electrochemical, iodometric

titration and ethylene blue analyses employed earlier have set the frame work for testing

the efficiency of these methodologies towards environmental samples to determine if

sulfide, polysulfides and polythionates can be detected at environmentally relevant

concentrations.

4.2 MATERIALS AND METHODS

4.2.1 Water Source and Collection Parameters

Water was collected from a site in the middle of Callander Bay of Lake Nipissing,

Ontario. This site has been identified by the Ontario Ministry of the Environment

(OMOE) as site IG9 (Easting: 624587 Northing: 5119449) where the water is 9 – 10 m

deep (OMOE Oct 6, 2015). The temperature of the water, dissolved oxygen

concentration (DO) and the pH of the water were 13.8 °C, 75.5% (7.88 mg/L), and 6.5,

respectively. The total watershed area for Lake Nipissing, including Callander Bay, is

296.12 km2 of which 83% is the Wasi watershed (Karst-Riddoch 2011). Samples for two

one-litre high-density polyethylene (HDPE) bottles pre-purged with N2 gas and sealed,

were collected using a Van Dorn water sampler to a depth of 10 m. The water sample

was filled to the very top of the HDPE bottle to ensure no headspace was left. The two

1-L HDPE sample bottles were returned to the lab and stored in the fridge (4 ºC) for no

longer than 48 hours before being repackaged and frozen until needed for further

analyses. The contents of the bottles were then split up into 15 individual 125-mL HDPE

bottles. In the Baxter Glove Box, under an atmosphere of nitrogen, the individual bottles

were sealed and frozen for subsequent analyses. Three 50-mL sub-samples of the water

were portioned into 125-mL flasks to be used for reduction and DEPD analyses. Those

samples were carefully sealed and stored in a fridge (4 ºC) until the analyses could be

completed. The three 50-mL samples were also prepared in the Baxter Glove box as

described above. A subsample of the lake water, from site IG9, was filtered (0.45µm)

and sent out for Ion-coupled Plasma Atomic Emission Spectroscopy (ICP-AES) analysis

at an external lab for trace inorganic metal analysis.

84

4.2.2 Reduction and DEPD Analysis of Lake Water Sample

Using the procedures for the reduction of sulfur compounds with chromium as described

in section 2.2.4, three replicates of the unfiltered water samples were first reduced

without any spiking of a known concentration of sulfide. Also, using the same

procedure, another triplicate set of untreated water samples was then spiked with a

known concentration of the sulfide stock solution and was reduced with chromium

(section 1.2.1.2). Similarly, another set of water samples was acidified without the

addition of chromium and analyzed using the same procedure as the above two sets of

water samples. The last two sets of water samples were being tested in order to detect the

presence of any acid-labile sulfides in the water sample. The procedure above was

repeated on filtered lake water samples. The filtration of the water samples was carried

out through a 0.45-µm membrane. All the three sets of replicates for the filtered and

unfiltered water samples were tested using DEPD, following the procedure outlined in

section 2.2.3, to detect the presence of sulfide in the lake water. A control group, using

ultra-pure water, with combinations of chromium and sulfide was analyzed using this

procedure to determine if there were any matrix interferences.

4.2.3 Electrochemical Analysis of Lake Water Samples

Electrochemical measurements were conducted using the BASi work station with Epsilon

USB software (Bioanaytical System Inc.). The three-electrode cell consists of a C3 glass

cell vial 50 mm x 59 mm, housing the supporting electrolyte solutions, and electrodes for

analysis. A glassy carbon, with a working surface diameter of 3.0 mm (area = 0.017

cm2); was used as the working electrode. The auxiliary electrode, Pt wire, in the form of

a 23-cm long coil was used and the reference electrode was a Ag/AgCl. All electrodes

were supplied by BASi. The electrochemical set up for the analysis using BASi is shown

in Figure 1. Electrochemical measurements (section 1.2.4.5), conducted using CV, where

the electrode potential was cycled from -1600 mV to +1600 mV at a scan rate of 100 mV

s-1 for 10 cycles. The peak potential (mV), height (µA), and area (µC) were recorded for

all 10 cycles. Three sets at 10 CV cycles per replicate were performed for each

experiment for BGCE and MGCE. The unfiltered lake water sample was examined using

85

the procedures outlined above; also the sulfide spiked lake water sample was examined.

The reduced lake water sample described in section 4.2.1.2 was also assessed.

4.3 RESULTS

4.3.1 Reduction and DEPD Analysis of Lake Water Sample

The reduction of the lake water samples with chromium and their analysis using the

DEPD methodology was carried out on a couple of water samples, chosen to represent a

freshwater sample that was collected from Lake Nipissing. The present study was

designed to test current methodologies outlined in the earlier chapters in order to

determine whether these methods would be able to detect sulfur compounds at

environmentally relevant concentrations. It was also an opportunity to see if individual

sulfur compounds could be distinguished and to see if natural samples showed any form

of matrix interferences.

The first subsample of lake water was reduced without being spiked with simple sulfide.

Two reduction procedures were carried out, the first without the addition of chromium,

and the second with chromium being added to the water sample. Also, filtered and

unfiltered samples were analyzed. When these reductions were completed, there were no

detectable traces of sulfide observed using the DEPD method (section 2.2.2). It is worth

noting that when the lake water samples were reduced without the addition of chromium,

a white precipitate formed at the bottom of the reaction flask (A), shown in Figure 26.

From the results indicated in Table 16 and Table 17 for water and sediments collected

from site IG9, there are very few of the divalent cations that are found in the water

sample (Table 16). However, as indicated in Table 17 provided by Chase et al. (2016),

there are several of the key divalent cations in the sediment samples from site IG9.

According to Kaasalainen and Stefánsson, (2011); Kariuki et al., (2008); Rozan et al.,

(2000) and Janssen et al., (1999), upon acidification of samples with metallic sulfides

such as, FeS2, PbS, Ag2S, MnS and HgS, there is a production of a white precipitate

forming during the reduction processes. Acidification of a sulfur containing sample

promotes the dissociation of any acid-volatile sulfide complexes present in the water

86

sample (Brouwer and Murphy, 1995). Such metal sulfide complexes have been shown to

not only increase the rate of oxidation of sulfide but also to change the distribution

products formed during sulfide-oxidation (Zhang and Millero, 1993). It has been

reported that metal sulfide complexes tend to occur for the bisulfide (HS-) of the divalent

cations like iron, cobalt, nickel, copper, and zinc (Chadwell et al., 2001; Al-Farawati and

van den Berg, 1999; Zhang and Millero, 1994). Since the pH range of most aquatic

systems falls in the slightly acidic to slightly basic range, and according to the fractional

dissociation for sulfide as a function of pH (Figure 2), the HS- ion is the most dominant

and readily available species. In Table 17 the presence of inorganic metals found in fresh

water sediments have been shown to form metal sulfide complexes, and are responsible

for controlling the solubility of trace metals in anoxic waters, and are also thought to be

important for the stabilization of H2S in oxic surface waters (Al-Farawati and van den

Berg, 1999; Zhang and Millero, 1994). Further examination of Table 17 indicates that

the values for iron and sulfur are above the limits set out in the guidelines by

Environment Canada for sediment samples. It has been suggested by Rozan et al., (2000)

that in natural water samples, iron sulfides are the most abundant, and that iron sulfides

diffuse from the anoxic sediments into the water column. Although the iron content in

the water sample, from IG9, was below the limit of detection along with the other metals

(Table 16), there was a likelihood of presence of sulfur in the form of sulfide complexed

with the metals. The presence of sulfur, found in the water sample at IG9, is a possible

cause for the white precipitate that formed at the bottom of reaction flask (A) (Figure 26)

after carrying out the chromium-oxidation reaction A possible reaction that may have

occurred to generate the precipitate is shown Reaction R – 4.1 (Bouroushain, 2010).

𝐻2𝑆(𝑔) → 𝑆(𝑠) + 2𝐻+ + 2𝑒− (𝑅 − 4.1)

87

Table 16. Inorganic metal concentrations (ppm) from site IG9 for water sample collected

in October of 2015.

Site Aluminum Barium Copper Iron Lead Magnesium Nickel Zinc Sulfur

IG9 ― 0.012 0.005 ― ― 1.99 ― ― 1.59

*CCME 0.005 ― 0.002 0.300 0.007 ― 0.025 0.030 ―

*CCME: Canadian Council of Ministers of the Environment, water quality data table for

the protection of aquatic life. (―) indicates no data or below limit of detection.

Table 17. Sediment concentrations (ppm), of the metals (and sulfur) associated with

smelting of nickel ores. Mean values (mg/kg ± standard error) (Chase et al. 2016,

unpublished results).

*ECG indicates Environment Canada Guidelines for sediment for each element (also

expressed in mg/kg)

According to several authors Kaasalainen and Stefánsson (2011); Kariuki et al. (2008);

Rozan et al. (2000) and Janssen et al. (1999), the addition of chromium to reaction flask

(A), in combination with HCl, promotes the dissociation of the metal sulfide complexes

(PbS, MnS, HgS, etc.). This in turn increases the recovery of simple sulfides in the

sample being tested (Rozan et al. 2000). However, Cainfield et al. (1986) and Luther III

et al. (1985) had earlier suggested that the addition of chromium would not promote the

dissociation of metal sulfide complexes of PbS, MnS, HgS, etc. In the present study,

after the chromium was added to the water sample and reduced again, there was still no

detectable sulfide present. Initially, aliquots of 1, 5 and 10 mL of the water samples

treated with chromium were analyzed through the DEPD measurement. However, since

there was no detectable sulfide in either of these sample volumes, a 40-mL aliquot of a

water sample treated with chromium was analyzed for the presence of sulfide through the

DEPD method. Despite the increased sample volume for both with and without added

chromium, the DEPD analyses still did not detect any sulfides in the water sample.

However, when the water sample was spiked with a known concentration of sulfide and

analyzed for sulfide recovery both with and without chromium, there were positive

Site Iron Nickel Copper Cobalt Arsenic Lead Zinc Sulfur

IG9 32,987

(629) 47

(0.5) 40

(0.3) 15

(0.2) 6

(0.4) 48

(0.7) 166 (1)

1038 (18)

*ECG 20,000 16 16 22 6 31 120 1000

88

results for the detection of sulfides using the DEPD and the iodometric titration methods,

as outlined below in Table 18.

Table 18. Comparison of percent recoveries and recovered concentrations (mM) of

sulfide spiked water samples for filtered and unfiltered samples. Experiment

Sample and Treatment DEPD (mM) % Recovery Iodometry (mM) % Recovery

Duchesnay Creek (UF) + S2- +Cr 3 85 3 78

IG9 (UF) + S2- + Cr 0.2 4 BLD ―

Duchesnay Creek (UF) + S2- 3 67 2 47

IG9 (UF) + S2- 0.1 4 BLD ―

Duchesnay Creek (F) + S2- + Cr 4 108 4 105

IG9 (F) + S2- + Cr 5 113 3 82

Duchesnay Creek (F) + S2- 3 72 2 53

IG9 (F) + S2- 4 98 3 74

Type I water + S2- + Cr 4 91 3 90

Type I water + S2- 2 56 2 41

UF = Unfiltered; F = Filtered (0.45 µm); BLD = Below Limit of Detection

The results in Table 18 indicate that the addition of chromium to the reaction flask A

(Figure 28) enhances the overall recovery of sulfides in the natural water samples

collected from the two different sites. This observation holds both for the filtered and

unfiltered natural water samples. The results in Table 18 also indicate that there are

matrix effects encumbered in natural water samples. According to the results, the matrix

effects were more enhanced in the unfiltered water samples since the sulfide recovery,

when compared to that of the filtered water samples, was lower. As stated above, even

though the addition of chromium to unfiltered water samples enhanced the recovery of

sulfides, the recovery of the sulfides in natural water samples is enhanced when the water

sample is filtered. This increase in sulfide recovery is in agreement with the results of

Mylon et al. (2002), Kariuki et al. (2008) and Rozan et al. (2000) who showed that the

addition of chromium to acidified water samples containing sulfides could cause the

transition metal sulfide complexes, if contained in the water samples, to release the

chromium-labile sulfides.

89

The collection of the IG9 water sample (Table 18) was done in early October of 2015,

and initial recoveries of unfiltered water samples yielded very low percent recoveries of

sulfide. Even after the sample’s treatment with chromium, low percent recoveries were

still observed. Based on the chemistry of typical field samples, and as described in

section 4.2.1, collection was after the fall overturn event, when stratification of water at

different temperatures and densities breaks down and water from the hypolimnion

(deeper zone) mixes freely with water in the epilimnion (Dodson, 2005). As a result, at

the depth of 10 m where the sample was collected, the water sample would have been

relatively well-oxygenated, despite the fact that earlier, during summer stratification, the

deep water region can become severely anoxic. Since the initial IG9 water sample was

unfiltered, and the fall-overturn had occurred, there is a good possibility that the

sediment-water interface has mixed. This mixing alters the chemistry at the sediment-

water interface, which would stir up fine colloid suspensions, chemotrophs, and bacteria;

thus could impair the conversion of simple sulfides in a natural water sample (Table 16

and Table 17). However, the conversion of sulfides was greatly increased when the water

sample was filtered prior to reduction in both the treatments applied.

The results in Table 18 indicate that there were elevated percentage yield recoveries with

some water samples that were treated with chromium over those that did not get the

chromium treatment. This could suggest that some sulfur-based intermediates such as

polysulfides, thiosulfate or tetrathionate may have been present. Chadwell et al., 2001

suggested that polysulfides can be formed by the oxidation of H2S by O2, iron III, and by

manganese III and IV oxides in the absence of O2 or by micro biotic processes. Some of

this micro biota can be removed if the water sample is filtered prior to treatment with

chromium, thus limiting the activity of their use of sulfur intermediates for assimilation

and dissimilation. These sulfur-based intermediates would normally not be detected

using DEPD, unless they first get reduced with chromium (Kariuki et al. 2008). If there

were any aqueous simple sulfides (H2S, S2- or HS-) in the field samples, those compounds

would have become oxidized by the fall turnover in the lake by air. Therefore, these

oxidized sulfur species would not have been detected as simple sulfides (Sorokin 2011;

Kuwabara et al. 1999) except in the case where the samples were first treated with

chromium (Morse and Luther III, 1999). For example, if the polysulfides were present in

90

the water samples, they could have been oxidized to thiosulfate, and without the

chromium reduction of the sample, these compounds would not have been detected.

Also, since the condition of the lake was oxic at this time of year, it is possible that the

metal-sulfide complexes may have formed, in the presence of oxygen (Al-Farawati and

van den Berg 1999; Morse and Luther 1999; Zhang and Millero 1994). These metal-

sulfide complexes would also go undetected if one were simply using the DEPD method

to analyze for just acid labile sulfides. At the time of sampling, the reported pH value for

the water sample was 6.5. According to Rozan et al. (1999), metal sulfide complexes are

likely to be present in natural water samples at this pH. In order to preserve the water

samples prior to analysis they were separated and frozen. O’Reilly et al. (2001) have

reported that freezing the samples failed to prevent the oxidation of sulfide. Since the

water samples were slowly frozen at -20 ºC and not flash frozen using liquid nitrogen,

there is a possibility that oxidation could still have been occurring during the freezing

process. The presence of naturally occurring thiosulfate in the water samples analyzed

through treatment of the samples with chromium could cause these compounds to

become detectable using DEPD, and this in turn could have inflated the observed percent

recovery of sulfide.

4.3.2 Electrochemical Analysis of Lake Water

Even though there was no detectable sulfide found in the lake water samples prior to their

being reacted with chromium, the lake water was still compared against the

electrochemical parameters described above. There was no detectable current response in

terms of the potential scan window of -1600 to +1600 mV on BGCE and MGCE. The

same results held for the chromium-reduced lake water samples in combination with the

3.63 mM spiked sulfide to the water sample prior to addition of chromium. As well,

there was no detectable current signal for the water samples spiked with 3.63 mM sulfide,

and with no chromium added to it. Using a larger amount of the chromium-reduced water

sample in the EC analysis still provided no current response. More CV cycles were run,

going up to 20 cycles, using the BGCE only, since the MGCE showed physical

breakdown of the film after 10 CV cycles. Nevertheless, no current response was seen.

At this point it is not clear why this was the case.

91

4.4 CONCLUSIONS

For the original, unfiltered and filtered water sample, after going through the reduction

process without being spiked with sulfide, the sulfide was not readily detected through

DEPD or CV analyses. This was also the case with the water sample being reduced with

and without chromium. When the unfiltered water sample was spiked with sulfide and

reduced, with and without chromium there was sulfide detected (4 mM) through DEPD

reaction. However, when the water sample was filtered, spiked with sulfide, and reduced

with and without chromium the amount of sulfide detected was close to and above 100 %

recovery (Table 18) for DEPD reaction at site IG9. When using the filtered, sulfide

spiked lake water samples, and applying the electrochemical conditions there was no

detectable sulfides present over BGCE and MGCE. This should have been the case since

the aliquot added to test this procedure exceeded the lowest detectable amount of sulfide

at BGCE (≥ 0.57 mM). The exceedingly high percent recovery values for these

chromium-reduced water samples (Table 18), when compared to the results from the

original water sample, prior to spiking with sulfide, from the reductions done raises some

concerns. As indicated in Tables 16 and 17, there is a presence of transition metal ions in

both the water and in the sediments at site IG9. Metal-sulfide complexes may have

formed, in the presence of oxygen during the change in conditions of the lake from

anoxic to oxic (Al-Farawati and van den Berg 1999; Morse and Luther 1999; Zhang and

Millero 1994). According to Morse and Luther III (1999), these metal-sulfide ion

complexes are disassociated when in the presence of chromium during the reduction

procedure. The interactions between the metal ions in the water and the sediment, at the

pH measurement taken during sample collection (pH = 6.5) could be the cause for the

higher than expected recoveries noted in Table 18. Also, it should be noted that the lower

recoveries observed in Table 18 for the unfiltered water sample could suggest that

bacterial activity is promoting the dissociation of the sulfide ion (Gosh and Dam, 2009).

An observation is made that although CV could be applied to the analysis of lab grade

solutions, it did not seem to show promise for the analysis of the natural lake water

sample where the sulfide is present at very low concentrations. Even after the reduced

lake water sample was analyzed in terms of the presence of sulfide in the sample there

92

was no detectable sulfide observed during CV analyses. This was the case for both BGCE

and MGCE. Sulfide concentration in the environmental samples, determined using some

electrochemical techniques, is reported to be higher than in the samples examined in this

study. For example, Dutta et al. (2008 and 2010) reported dissolved sulfide levels in

sewage and paper mill waste water to have been 47 mg L1- and 51 mg L1-, respectively.

Also, Font et al. (1996) reported sulfide levels in waste water downstream from a tannery

to be 20 mg L1-. For future work it is possible that different electrochemical techniques

may hold more promise for the investigation of sulfide species in water samples at low

concentrations.

93

General Conclusions

In a recent review of the value of voltammetry, Batchelor-McAuley et al. (2015)

suggested that CV could be used along with other electrochemical-based techniques to

perform the quantitative analysis of sulfide. However, based on the present study, CV

using glassy carbon electrodes appears to be inefficient as a technique for the

quantification of sulfide at environmentally relevant concentrations using BGCE and

vanadium oxide MGCE. The addition of an electrode modifier, such as vanadium oxide,

enhances the catalytic response at lower sulfide concentrations when compared to that of

the concentrations analyzed using BGCE. However, when analyzing solutions with the

vanadium oxide modifier, the supporting electrolyte solution directly overlaps with the

current response related to electrooxidation of sulfide. This causes a concern, since the

interference from the supporting electrolyte solution does not allow for the detection limit

of sulfide to be lower than 0.13 mM using the CV parameters used in this study in

combination with the vanadium oxide modifier. However, the observed standard error

results (Table 8) for the concentration of 0.13 mM at MGCE, would indicate that levels

this low would be unquantifiable leading to further limitations using CV at MGCE.

Although, the SEM results, for the MGCE, indicated that there was a homogenous

coating on the electrode surface, the EDX results did indicate the presence of any

vanadium elements for that film covering. The BGCE and MGCE analyses, performed in

this study, are limited in their application towards environmentally relevant

concentrations of sulfide. Also, the considerable variation shown in measurements made

using both the BGCE and MGCE electrodes suggest that sulfide is not easily quantified

using this approach.

Methods using hanging mercury drop electrodes with polarographic, square-wave or

stripping voltammetric techniques have been shown to be more reliable for the

quantification of sulfide in solution (Dilgin et al. 2012; Huang et al. 2012; Lawrence

2006; Cheng et al. 2005; García-Calzada et al. 1999). Mercury drop electrodes have a

unique advantage over glassy carbon electrodes, as they provide a fresh working surface

with every drop of mercury. This guarantees a clean working surface with no surface

imperfections. The issue with using mercury is that it is toxic, so other electrode

94

materials are being examined in order to circumvent concerns over the use of mercury

electrodes. Other electrode materials such as gold (Waite et al., 2006), platinum

(Kapusta et al. 1983) and boron doped diamond (Lawrence et al., 2002) electrodes have

been shown to detect sulfide in solution. In contrast, the use of glassy carbon electrodes

requires that the working surface be manually polished to remove deposits from previous

analyses. These repeated manual manipulations of the glassy carbon surface will leave

grooves and uneven surfaces, which in turn may differentially accumulate particles

during subsequent runs accounting for the variation observed among the replicates.

Refinement of polishing procedures, combined with electrochemical pre-treatment prior

to analyses may help identify electrodes that may not have been properly polished

between analyses.

Any attempt to use CV to measure individual substances in mixtures of various sulfur-

based compounds, such as polysulfides and polythionates in solution, was also found to

have been challenging. Although some qualitative information for each group of

compounds was gained when they were analyzed separately, trisulfide and tetrasulfide

showed very similar CV responses using BGCE and the disulfide showed no detectable

signal under the conditions that the CV was run. When analyzed in a mixture, the sulfide,

polysulfides and polythionates are not readily distinguished from each other with the CV

protocols used in this study. This also seems to be true for the DEPD experiments. These

experiments showed good percent recoveries of simple sulfide and detectable levels of

chromium-labile sulfides, but the DEPD method could not distinguish combinations of

these sulfur-based compounds in solution. The DEPD-colorimetric method used in this

study is very similar to the methylene blue method commonly used in other studies. In

this case, the total sulfide available in solution in the form of acid-labile, metal sulfides,

polysulfides and polythionates in a water sample can be quantified when the sample has

been reduced with chromium under the conditions described earlier (Kamyshny et al.

2008).

Clearly, sulfide was not readily detected in the lake water sample. Even after reducing

the water sample with chromium according to the procedures outlined earlier, no sulfide

was detected using either the DEPD or the CV methods. When the sample was spiked

95

with a known concentration of sulfide, sulfide was detected and measurable. However,

the water sample must first be filtered prior to analyses in order to reduce any matrix

interferences that may occur. Reduction of the sulfide-containing water samples at a

temperature of 60ºC, with chromium metal as a reducing agent, allowed for the total

recovery of sulfide. Cyclic voltammetry of the lake water sample and the chromium-

reduced lake water sample provided no qualitative or quantitative information on sulfide

concentrations using either BGCE or MGCE electrodes. A further investigation of the

techniques presented in this study, as well as others not presented may cast more light

into how sulfur compounds in a sample can be detected.

This study has provided a useful assessment of the potential of CV, using BGCE and

vanadium oxide MGCE as a method for the qualification and quantification of lab

generated sulfur based compounds in solutions. It is now clear that when these sulfur

based compounds are combined as a mixture, there is a lack of speciation between the

compounds, thereby making them not easily distinguished from one another. When

examining environmental samples, being able to have speciation of the various sulfur

compounds, contained as a mixture in the natural samples would be a great contribution

to the community at large. The chemistry of sulfur in natural systems is complex, and

highly dependent on oxidation states and the availability of molecular oxygen. In

addition, there is some likelihood that chemical changes may easily develop during

transportation and storage, especially in raw unfiltered water samples that still contain

numerous active microorganisms. Other factors that may have affected the results of the

field analyses would include the concentrations of metal ions in the water sample, the

sediment, and the time of year the sample was taken. A thorough study of chemistry of

the vanadium oxide electrode modifiers may provide the real key to the usefulness of the

CV technique for the analysis of sulfur species in natural water and other environmental

samples. Such a study could include reducing the forward switching potential to a value

less than where water oxidizes to oxygen. Exploration of other metal oxide modifiers

besides vanadium films, such as bismuth (Huang et al., 2012), and an investigation of

other electrochemical techniques for the analysis of sulfur compounds in natural samples

could turn out to be very informative.

96

Appendices

Appendix A

Supporting Electrolyte Solution:

0.1 M Potassium Phosphate Buffer

A potassium phosphate buffer solution was prepared by combining solutions of 0.075 M

dibasic potassium phosphate with 0.025 M monobasic potassium phosphate. The buffer

solution was brought to a pH of 10.3 using sodium hydroxide (NaOH) (RICCA

Chemicals). The phosphate buffer solution was filtered through a 47.00 mm, 0.45 µm

mixed cellulose membrane (Millipore Corp.) using a glass vacuum filtration apparatus,

and water aspirator. The mono and dibasic phosphate salts were reagent grade, and

required filtration to remove any particulates that would otherwise be unwanted. After

filtration, the solution was purged for 1 hour using nitrogen gas, by inserting a piece of

Teflon tubing (ID 3.2 mm x OD 6.4 mm x wall1.6 mm) to the bottom of the glass storage

bottle that was connected to the regulator supplying the gas. Before the tube was fully

removed, once the purging was complete, the tube was brought above the liquid level in

the bottle to insert a blanket of N2 gas in the head space of the storage bottle before the

buffer solution was capped for storage until use.

Solutions for Iodometric Titrations:

6 M Hydrochloric Acid Solution

Using a 100-mL volumetric flask, 50-mL of concentrated hydrochloric acid (trace metal

grade, Fisher Scientific) was added to 30-mL of ultra-pure water; then filled to the mark,

on the flask, to make a total volume of 100-mL. This solution was used for standardizing

the sulfide stock solution during the titration procedure with 25 mM iodine and

thiosulfate.

97

Standard Thiosulfate Solution 25 mM

Using a 1000-mL graduated cylinder, approximately 600-mL of the ultrapure water was

transferred into 1000-mL Erlenmeyer flask. Sodium thiosulfate pentahydrate

(Na2S2O3•5H2O) (≥ 98.5 %, Anachemia, ACS grade) weighing 6.21 ± 0.05 g was

dissolved into the ultrapure water, contained in the Erlenmeyer flask. Next, 0.4 ± 0.01 g

of NaOH was added to the Na2S2O3•5H2O solution and the contents of the Erlenmeyer

flask were brought to a total volume of 1000-mL. The NaOH was added as a

preservative. Once prepared, the thiosulfate solution was poured into an amber storage

bottle. The solution was expected to be stable for 6 months (Clesceri, Greenberg, and

Eaton, 1998). The thiosulfate solution was standardized against potassium bi-iodate

solution.

Potassium Bi-Iodate 2 mM solution

A sample of potassium bi-iodate (KIO3) (SIGMA, analytical reagent ≥ 99.8 %), weighing

0.814 ± 0.005 g was transferred to a 1000-mL Erlenmeyer flask. The KIO3 was dissolved

in about 600-mL of ultrapure water. Once dissolved the KIO3 solution was brought to a

total volume of 1000-mL, mixed well, and then the solution was transferred into an

amber storage bottle. The solution was expected to be stable for 6 months (Clesceri,

Greenberg, and Eaton, 1998).

Iodine Solution 25 mM

Using a 1000-mL graduated cylinder, approximately 600-mL of ultrapure water was

added to a 1000-mL Erlenmeyer flask. A sample of potassium iodide (KI) (Fisher, ≥ 99.0

%) weighing 24.1 ± 0.05 g was and added to the flask containing the ultrapure water.

Once the KI was dissolved completely, 3.2 ± 0.02 g of resublimed iodine solid (I2)

(Fisher, ACS grade) was weighed, added to the KI solution, and dissolved. Once the

iodine solid was completely dissolved, the remaining ultra-pure water was added to reach

a total volume of 1000-mL. The iodine solution was transferred to an amber bottle for

storage, and was expected to be stable for 6 months (Clesceri, Greenberg, and Eaton,

1998). The iodine solution was standardized against 25 mM thiosulfate solution.

98

Starch Solution 2 % (w/v)

Using a 250-mL beaker, 200-mL of ultra-pure water was heated thoroughly on a hotplate.

Some of hot water from the beaker was poured into a 100-mL volumetric flask, and the

flask was placed back on the hotplate. A sample of 2.00 g soluble starch (Lab Grade,

Fisher Scientific), was weighed, added to the hot water, and dissolved. The flask was

capped, and mixed well. The remaining hot water was poured to the mark of the

volumetric flask to reach total volume of 100-mL. The starch solution was expected to

keep stable for a period of 1 month (Clesceri, Greenberg, and Eaton, 1998). This

solution is used to check for the presence of iodine during iodometric titrations.

Solutions for DEPD Procedure

0.1 M Sodium Hydroxide (NaOH) Solution

Using a 1000-mL glass storage bottle, ultra-pure water (18.2 MΩcm) was added to within

a few inches of the top. Approximately 4.00 g of NaOH pellets were weighed and placed

in the ultra-pure water, a stir bar was added, and the solution was placed on a hot-plate

stirrer, with no heat applied, and set to 800 rpm. While the solution was being stirred, it

was also purged for 1 hour using N2 gas (99.998 %) in order to remove any presence of

oxygen from the solution. The pH was confirmed using a pH meter with an ATC pH

combination electrode (Accumet AB150 pH/mV meter and pH/ATC electrode, Fisher

Scientific).

Colour Developing Reagent

Colour developing reagent was made by weighing out 4.3 g of N,N- Diethyl-p-

phenylenediamine oxalate salt 96 % (DEPD) (Fisher Scientific), and 6.4 g iron (III)

chloride hexahydrate 97 – 102 % (FeCl3·6H2O) (ACS grade, Sigma). These two

compounds were added to 100-mL of 50% (v/v) HCl solution, placed in a glass storage

bottle, and stored at 4°C.

99

Appendix B

3𝐻2𝑆𝑂4(𝑎𝑞) + 𝐾𝐼𝑂3(𝑎𝑞) + 5𝐾𝐼(𝑎𝑞) ↔ 3𝐼2(𝑎𝑞) + 3𝐻2𝑂(𝑙) + 3𝐾2𝑆𝑂4(𝑎𝑞) (R − 1.1)

2𝑁𝑎2𝑆2𝑂3(𝑎𝑞) + 𝐼2(𝑎𝑞) → 2𝑁𝑎𝐼(𝑎𝑞) + 𝑁𝑎2𝑆4𝑂6(𝑎𝑞) (𝑅 − 1.2)

ɳ𝐾𝐼𝑂3 = 𝐶 × 𝑉(𝐿) (𝐸𝑞. 1.3)

ɳ𝐼2 = 3 × ɳ𝐾𝐼𝑂3

(𝐸𝑞. 1.4)

ɳ𝑁𝑎2𝑆2𝑂3 = 2 × ɳ𝐼2

(𝐸𝑞. 1.5)

𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 𝑆2𝑂3 = (ɳ𝑁𝑎2𝑆2𝑂3

𝑉𝑜𝑙. 𝑁𝑎2𝑆2𝑂3(𝐿)

) × 2 (𝐸𝑞. 1.6)

2𝑆2𝑂32− + 𝐼2 → 𝑆4𝑂6

2− + 2𝐼− (R − 1.3)

𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 𝐼2 = (𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 𝑜𝑓 𝑆2𝑂3) ∙(𝑉𝑜𝑙𝑁𝑎2𝑆2𝑂3 (𝐿))

𝑉𝑜𝑙. 𝐼2 𝑡𝑖𝑡𝑟𝑎𝑡𝑒𝑑 (𝐿) (𝐸𝑞. 1.7)

𝐼2 + 𝑆2− → 𝑆 + 2𝐼− (𝑅 − 1. 4)

𝐼2 + 2𝑆2𝑂3 → 𝑆4𝑂62− + 2𝐼− (𝑅 − 1.5)

The following equation was used to calculate the concentration of sulfide in mg/L (Eq.

1.8):

𝑚𝑔

𝐿 𝑆2− =

[(𝐴 × 𝐵) − (𝐶 × 𝐷)] × 32000

𝑉𝑜𝑙. 𝑜𝑓 𝑆2− 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝑚𝐿) (𝐸𝑞. 1. 8)

Where:

A = mL iodine solution

B = molarity of iodine solution

C = mL thiosulfate solution

D = molarity of thiosulfate solution

100

𝐻𝑆− ↔ 𝐻+ + 𝑆2− (𝑅 − 1.5)

𝑆2− ↔ 𝑆 + 2𝑒− (𝑅 − 1.6)

Oxidation states of various sulfur compounds.

Species S2- HS- S S2O32- S4O6

2- SO32- SO4

2-

Oxidation No. -2 -1 0 +2 +2.5 +4 +6

Stoichiometric Equations that show the amount of Na2S and S required for

making Na2S2, Na2S3, and Na2S4

𝑵𝒂𝟐𝑺 + 𝑺 → 𝑵𝒂𝟐𝑺𝟐 (𝑅 − 2.1)

The amount of Na2S and S0 required making 1.00 g Na2S2 are given in equations (2.1 a)

and (2.1 b):

1.00 𝑔 𝑁𝑎2𝑆2 × 1 𝑚𝑜𝑙 𝑁𝑎2𝑆2

110.11 𝑔 ×

1 𝑚𝑜𝑙 𝑁𝑎2𝑆

1 𝑚𝑜𝑙 𝑁𝑎2𝑆2 ×

78.04 𝑔 𝑁𝑎2𝑆

1 𝑚𝑜𝑙 𝑁𝑎2𝑆= 0.7087 𝑔 𝑁𝑎2𝑆 (2.1 𝑎)

1.00 𝑔 𝑁𝑎2𝑆2 × 1 𝑚𝑜𝑙 𝑁𝑎2𝑆2

110.11 𝑔 ×

1 𝑚𝑜𝑙 𝑆

1 𝑚𝑜𝑙 𝑁𝑎2𝑆2 ×

32.06 𝑔 𝑆

1 𝑚𝑜𝑙 𝑆= 0.2912 𝑔 𝑆 (2.1 𝑏)

𝑵𝒂𝟐𝑺 + 𝟐𝑺 → 𝑵𝒂𝟐𝑺𝟑 (𝑅 − 2.2)

The amount of Na2S and S0 required making 1.00 g Na2S3 are given in equations (2.2 c)

and (2.2 d):

1.00 𝑔 𝑁𝑎2𝑆3 × 1 𝑚𝑜𝑙 𝑁𝑎2𝑆3

142.17 𝑔 ×

1 𝑚𝑜𝑙 𝑁𝑎2𝑆

1 𝑚𝑜𝑙 𝑁𝑎2𝑆3 ×

78.04 𝑔 𝑁𝑎2𝑆

1 𝑚𝑜𝑙 𝑁𝑎2𝑆= 0.5489 𝑔 𝑁𝑎2𝑆 (2.2 𝑐)

1.00 𝑔 𝑁𝑎2𝑆3 × 1 𝑚𝑜𝑙 𝑁𝑎2𝑆3

142.17 𝑔 ×

2 𝑚𝑜𝑙 𝑆

1 𝑚𝑜𝑙 𝑁𝑎2𝑆3 ×

32.06 𝑔 𝑆

1 𝑚𝑜𝑙 𝑆= 0.4510 𝑔 𝑆 (2.2 𝑑)

𝑵𝒂𝟐𝑺 + 𝟑𝑺 → 𝑵𝒂𝟐𝑺𝟒 (𝑅 − 2.3)

The amount of Na2S and S0 required making 1.00 g Na2S4 are given in equations (2.3 e)

and (2.3 f):

101

1.00 𝑔 𝑁𝑎2𝑆4 × 1 𝑚𝑜𝑙 𝑁𝑎2𝑆4

174.24 𝑔 ×

1 𝑚𝑜𝑙 𝑁𝑎2𝑆

1 𝑚𝑜𝑙 𝑁𝑎2𝑆4 ×

78.04 𝑔 𝑁𝑎2𝑆

1 𝑚𝑜𝑙 𝑁𝑎2𝑆= 0.4479 𝑔 𝑁𝑎2𝑆 (2.3 𝑒)

1.00 𝑔 𝑁𝑎2𝑆4 × 1 𝑚𝑜𝑙 𝑁𝑎2𝑆4

174.24 𝑔 ×

3 𝑚𝑜𝑙 𝑆

1 𝑚𝑜𝑙 𝑁𝑎2𝑆4 ×

32.06 𝑔 𝑆

1 𝑚𝑜𝑙 𝑆= 0.5520 𝑔 𝑆 (2.3 𝑓)

102

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Curriculum Vitae: Ashley Ryan (nee Marcellus)

SUMMARY OF QUALIFICATIONS

Central Analytical Facility (CAF) Technologist – Nipissing University

The CAF is a unique part of Nipissing University’s Research Facility. A dedicated

technologist supervises, coordinates, and performs its day-to-day operations, providing

training and support to students and faculty researchers, external clients in academe, and

government or industry clients. Six pieces of equipment requiring specialized technical

support include: High Pressure Liquid Chromatograph (HPLC), Gas Chromatograph

Mass Spectrometer (GC-MS), Fourier-Transformation Infra-Red Spectrometer (FTIR),

Nuclear Magnetic Resonance Spectrometer (NMR), Atomic Absorption Spectrometer

(AAS), Electrochemical Analyzer, and a Microplate Reader. Two additional pieces of

equipment are the Confocal Microscope and Electron Microscope. These microscopes

are also maintained and operated by an individual with specialized technical training.

This suite of research instruments is integral to the research and scientific training

missions of the University and represents a significant component of the infrastructure

and technical capacity for scientific research at NU.

The CAF technologist ensures the functioning and integrity of all equipment in the

research facility. Each instrument involves a unique set of tasks and functions that need

to be performed on a regular basis. The skills required to maintain the equipment in the

CAF must be constantly updated, and they continue to evolve as the equipment ages.

Among the many duties of the CAF technologist, one important role is as a liaison,

facilitating interactions among students, faculty, and administrators in support of the

success of the research programs at this institution. The CAF technologist is involved in

the collection of data, but also is responsible for the efficient scheduling and optimal use

of equipment. It is critical that research deadlines (especially for thesis student

researchers) are met with all parties working together smoothly.

Over the past seven years as the CAF technologist I have worked diligently to support

research initiatives at Nipissing University. I am a versatile, hard-working, innovative

individual with a practical hands-on approach; I always strive for excellence. I have a

strong work ethic, and dedicate myself to ensuring that a job is done correctly. Applying

my analytical and problem solving skills, I routinely identify challenges and develop

effective solutions so that researchers will be able to collect meaningful, high quality

results.

PROFILE SUMMARY

A versatile, hard-working, innovative individual with a practical hands-on approach who

always strives toward excellence. Ability to collect, analyze and interpret data and

quickly grasp complex issues. Excellent interpersonal, analytical and problem solving

134

skills, promptly identifying problems and developing effective solutions. Proven ability

to complete projects to highest standard and meet deadlines with meticulous attention to

detail.

EDUCATION Masters of Environmental Science

Nipissing University, North Bay

2012-Present

Research Focus: Evaluating the efficiency of cyclic voltammetry using modified and

unmodified glassy carbon electrodes for the analysis of sulfur containing compound in

water.

Expected graduation date: October 2017

Honors Bachelor of Science Biology

Nipissing University, North Bay

2003-2008

Thesis Project: Characterization of senescence patterns of Eriophorum vaginatum

Related course work: General and Analytical Chemistry, Microbiology, Plant

Physiology and Fresh Water Biology.

Equipment: Olympus FV1000 Laser Scanning Confocal Microscope

ON THE JOB TRAINING

Bruker Fourier NMR 300

Bruker Ltd. Dec 2012

Supervised installation of NMR equipment into the CAF

Onsite training provided by Bruker NMR Specialists for Operation

and Maintenance

Maintenance: Filling of liquid Nitrogen (every week); Filling of

liquid Helium (every 4-5 months) to ensure stability of the supercoil

magnet. Updating shim files, field and phase values for optimal

spectral performance

Operation: Acquisition of proton and carbon NMR spectra, high

temperature NMR parameters, recognizing bad spectra and what can

be done to correct it. Troubleshooting issues with shimming the

magnetic pull to optimize the spectra. Inputting parameters for new

solvents. How to prepare samples for NMR analysis.

Since the installation of the NMR in CAF the equipment has yielded

several peer-reviewed journals from 2013 to present. These are listed

on Dr. Mukund Jha’s webpage. Data collected from NMR have been

collected by myself, BSc Honours students and research interns

135

Solid Phase Extraction (SPE) Seminar

Waters Ltd. Mississauga, ON

Sept 2012

Introduction to importance to SPE to analytical analyses

Sample preparation

Isolation of analyte to match choice of bedding in SPE

Optimization for maximum analyte recovery

Troubleshooting issues with SPE

SFR: Electron Microscope Commissioning

Nipissing University, North Bay May 2010

On Site training of Philips CM 10 Transmission Electron Microscope

provided by Peter Maloney from SFR

Operation: alignment of electron beam, optimizing viewing

conditions based on specific sample, alignment of column for

optimum performance, etc.

Maintenance and care: Removing parts of the column and cleaning

the appropriate parts, water cooling maintenance, filter changes, water

changes to reduce bacterial growth within the smaller pipes of the

microscope, etc.

Training received allowed me to train other students in basic operation

of the EM scope; this has supported the research of graduate students,

honors students and faculty

John Dolan HPLC Troubleshooting/Diagnostics Work Shop

Marriot Hotel, Vaughan June 2010

Troubleshooting principles

o Measurements and basic practices

Performance Qualification

o Quality/validity of results as well as accurate reproducibility

Column Physics/Chemistry

o Why and how columns die; problems with the stationary

phase; how and when to avoid contamination (samples/mobile

phase)

Pumps and Autosamplers

o Maintenance to pump heads and check valves; degassing

solvents for mobile phase; what type of tubing is suitable for

each type of analyses; how to recognize failure in frits;

blockages and leaks.

Detectors

o UV detectors noise and drift in baseline; wavelength selection

and other types of detectors available (PDA-Photo Diode

Array)

Quantification

o Measuring peak area/height; issues with calibration curves and

reproducibility issues.

136

Olympus Confocal FV1000 Laser Scanning Confocal Microscope

Nipissing University, North Bay August 2007

Received onsite training from Vince Varallo (Olympus Canada)

Operation: Image acquisition control for laser optimization and

operation, acquisition setting for scan rate function and position of

specimen for optimum performance, recognizing spectral bleed and

how to correct for it, etc.

Maintenance: Recognizing signs of degradation in signal, alignments

and replacements of bulbs/diodes; proper cleaning of parts for

operation i.e. objective pieces, condensers, etc.

Web Seminars: Analytical Content

One Hour in Length, North Bay

Started viewing in March 2010 to Present (2016) Wide variety of topic are covered over these presentations

Topics include but are not limited to: troubleshooting equipment issues, samples prep,

appropriate column choices, new techniques/technologies, etc.

EXPERIENCE Central Analytical Facility Technologist

Nipissing University, North Bay

March 2009-Present

Recognizing, troubleshooting, fixing issues with all equipment in CAF, and other labs Preparing reagents and solutions on a regular basis for ongoing projects Analyzing samples for researchers, professors and collaborators Writing test reports for professors, and collaborators Setting up instruments based on methods and protocols for CAF users Maintaining all equipment in CAF, Confocal Microscope, Electron Microscope, water

systems, equipment in teaching labs Keeping records pertaining to replacement of parts either for maintenance or repair Keeping track of financials for the CAF, maintenance, and supplies budget Training students to operate instruments Writing and updating Standard Operation Procedures for each instrument Developing new analytical methods or protocols to meet research requests Developing quality control procedures for specific analytical methods

Plant Growth Facilities Intern

Nipissing University, North Bay

June 2008-March 2009

Maintained cleanliness of Greenhouse

Contained pest outbreak, using appropriate measures

Assisted with ongoing research projects conducted in the Greenhouse

Assisted students on their current projects

Learned to operate mechanical components as well as the Argus operation program that

provides climate and irrigation control

137

Research Assistant Chemistry

Biology Department

Nipissing University, North Bay

May 2005-Aug 2005

Followed existing research protocols

Validate quality of data, elements and editing

Collaborate with senior researchers in order to ensure accurate research procedures and results

Conducted assays using mine tailings and acidic solutions to isolate individual elements

PUBLISHED WORK , AWARDS, COMMUNTIY SERVICE

Rao V.K., Kaswan P, Shelke G.M., Ryan A., Jha M and Kumar A. 2015. Iodine-Mediated, Microwave-

Assisted Synthesis of 1-Arylnaph-thofurans via Cyclization of 1-(1′-Arylvinyl)-2-naphthols. Synthesis.

Vol 47, pp. 3990 – 3996

Jha, M., Edmunds M., Lund K., and Ryan A. 2014. A new route to the versatile synthesis of thiopyrano

[2,3-b:6,5-b’] diindoles via 2-(alkylthio)-indole-3-carbaldehydes. Tetrahedron Letters. Vol. 55; pp.

5691-5694.

Mirza, R. S., Laraby, C. A., & Marcellus, A. 2013. Knowing Your Behaviour: The importance of

Behavioural Assays in the Characterisation of Chemical Alarm Cues in Fishes and Amphibians.

In Chemical Signals in Vertebrates 12 (pp. 295-308). Springer New York.

Jha, M., Enaohwo, O. & Marcellus, A. 2009. Chemoselective S-benzylation of indoline-2-thiones using

benzyl alcohols. Tetrahedron Letters. Vol. 50(51); pp. 7184-7187.

Certificate of appreciation for Science Day (Feb 22, 2014); volunteered knowledge about

importance of water that engaged the youths of Scouts Canada and Girl Guides of Canada in

fun hands-on lab activities.

Image of Distinction Category for Nikon Small Worlds 2009 Competition; two images of the

four were placed in this category

Judge in North Bay Regional Science Fair (April 2015, April 2016)

Judge for North Bay Regional Robotics Competition (March 2016)