TOTAL ORGANIC HALOGEN FORMATION IN THE PRESENCE OF

IOPAMIDOL AND CHLORINATED OXIDANTS WITH AND WITHOUT

NATURAL ORGANIC MATTER.

A Thesis

Presented to

The Graduate Faculty of The University of Akron

In Partial Fulfilment

of the Requirements for the Degree

Master of Science

Nana Osei Bonsu Ackerson

May, 2014

TOTAL ORGANIC HALOGEN FORMATION IN THE PRESENCE OF

IOPAMIDOL AND CHLORINATED OXIDANTS WITH AND WITHOUT

NATURAL ORGANIC MATTER.

Nana Osei Bonsu Ackerson

Thesis

Approved: Accepted:

______Advisor Department Chair Dr. Stephen E. Duirk Dr. Wieslaw Binienda

______Committee member Dean of the College Dr. Christopher C. Miller Dr. George K. Haritos

______Committee member Dean of Graduate School Dr. Lan Zhang Dr. George R. Newkome

______Date

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ABSTRACT

The objectives of this study were to investigate the transformation of ICM as a function of pH (6.5 to 9.5) and time (up to either 72 or 168 hr) in the presence of chlorinated oxidants. Total organic iodide (TOI) loss was used as a surrogate for the

ICM. Experiments were performed with and without natural organic matter (NOM).

Degradation of TOI in the absence of NOM was carried out at low and high concentrations of iopamidol and aqueous chlorine. Also, the effect of NOM variation on iodate formation was investigated.

The TOI degradation and iodate formation at low reactant and buffer concentrations were greatest at pH 7.5 and least at pH 9.5. TOI degradation followed observed first-order kinetics at all pH except pH 6.5, which exhibited bi-phasic degradation kinetics. Iodate formation did not follow either first or second order observed formation and was the predominant iodine-containing species after 24 hr.

Furthermore, disinfection by-products (DBPs) formed at pH 6.5 – 8.5 were chloroform, trichloroacetic acid and chlorodiiodomethane. In the presence of monochloramine and in the absence of NOM, the loss of TOI was insignificant and no iodate formation was observed.

At high concentrations of iopamidol and aqueous chlorine, TOI loss and iodate formation at pH 6.5 and 8.5 was rapid for the first 24 hr and ceased afterwards. The formation of total organic chloride (TOCl) was initially observed at 6 hr and 2 hr for pH 6.5 and 8.5 respectively. Also, chloroform, dichloroiodomethane, chlorodiiodomethane, dichloroacetic acid and trichloroacetic acid was observed.

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About 99% of the remaining TOI formed at each discrete time was contained in unidentified iopamidol transformation products.

When TOI was monitored in the presence of NOM and aqueous chlorine, source waters from Akron, Barberton and Cleveland respectively recorded 68 to 74%,

62 to 72% and 68 to 77% loss of TOI. However, no iodate was formed in any of the source water experiments. No significant degradation of TOI was observed in the presence of NOM and monochloramine. Iodate was not formed in varying NOM concentrations in Barberton and Cleveland source waters.

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ACKNOWLEDGEMENTS

I thank my God Almighty who has brought me this far and blessed me with wisdom and understanding in my academic pursuits. I would like to express my profound appreciation to my advisor, Dr. Stephen E. Duirk for his guidance, assistance and time. His patience, constructive criticisms and dedication were vital to the success of this thesis. Also, my gratitude goes to Dr. Christopher C. Miller and

Dr. Lan Zhang for their time, advice and insightful comments. To my laboratory colleagues, both graduate and undergraduate students, I say thank you for your unflinching support. My sincere thanks to all and sundry who supported me in diverse ways. Finally, my deepest appreciation goes to my wife Irene Ackerson and my daughter Nana Onomaa Ackerson for their prayers, love, support, patience and understanding during my busy schedules and throughout my studies.

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TABLE OF CONTENTS Page

LISTS OF TABLES ...... ix

LIST OF FIGURES ...... x

CHAPTER

I INTRODUCTION ...... 1

1.1 Background ...... 1

1.2 Problem Statement ...... 5

1.3 Specific Objectives ...... 7

II LITERATURE REVIEW ...... 9

2.1 Introduction ...... 9

2.2 Iodinated X-ray Contrast Media ...... 9

2.2.1 Occurrence and Concentration of ICM in Water and Wastewater ...... 11

2.2.2 Transformation of ICM ...... 13

2.3 Reactions of Chlorinated Oxidants Used in Water Treatment ...... 14

2.3.1 Chlorine ...... 14

2.3.2 Chlorine Dioxide ...... 17

2.3.3 Chloramines ...... 18

2.4 Chemistry and Reactions of Iodine ...... 19

2.5 Total Organic Halogen Formation ...... 22

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2.6 Toxicity of Halogenated Disinfection By-Products ...... 26

III MATERIALS AND METHODS ...... 29

3.1 Chemicals and Reagents ...... 29

3.2 Source Water Characterization ...... 30

3.3 Experimental Methods ...... 38

3.3.1 Experiments with Deionized Water ...... 38

3.3.2 Experiments with Source Waters ...... 42

3.4 Analytical Procedures ...... 44

3.4.1 Total Organic Halogen ...... 44

3.4.2 Disinfection By-product ...... 45

3.5 Analyses of TOX, Iodate and Iodide ...... 47

3.6 Analyses of DBPs ...... 55

IV RESULTS AND DISCUSSION ...... 71

4.1 Introduction ...... 71

4.2 Transformation of Iopamidol in the Absence of NOM ...... 71

4.2.1 Transformation at Low Concentration ...... 71

4.2.2 Transformation at High Concentration ...... 83

4.3 Transformation of Iopamidol in the Presence of Chlorine and NOM ...... 96

4.4 Transformation of Iopamidol in the Presence of Monochloramine and NOM .... 104

4.5 Iodate Formation as a Function of Dissolved Organic Carbon ...... 110

V CONCLUSIONS AND RECOMMENDATIONS ...... 112

vii

5.1 Introduction ...... 112

5.2 Conclusions ...... 112

5.3 Recommendations ...... 115

REFERENCES ...... 116

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LISTS OF TABLES

Table Page

2.1 Aqueous Iodine species...... 20

2.2 Reactions forming TOX and iodate...... 25

3.1 Source water characteristics from Akron, Barberton and Cleveland water...... 32

3.2 Florescence EEM regions proposed by Chen et al. (2003)...... 34

3.3 Florescence regions for Akron, Barberton and Cleveland source waters for 1 mg/L C...... 34 3.4 Comparison of recovery at 4°C and room temperature using 2,4,6- trichlorophenol, 2,4,6-tribromophenol and 4-iodophenol. [TCP] = 25 – 100 μM, [TBP] = 5 – 15 μM, [IPh] = 5 – 15 μM...... 41

3.5 Oven temperature programming for THMs and HANs analysis on GC/μECD.....55

3.6 Oven temperature programming for HAAs analysis on GC/μECD...... 56

3.7 Limit of quantification for the detection of DBPs...... 70

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LIST OF FIGURES

Figure Page

2.1 The chemical structures of ICM of common usage in hospitals...... 11

2.2 TOX formation and oxidation products...... 24

2.3 Iodo-DBP formation pathway (Adapted from Duirk et al., 2011)...... 26

3.1 Fluorescence excitation-emission spectrum of Akron source water. [DOC] = 5.57 mg/L, SUVA254 = 2.27 L/mg.m...... 35

3.2 Fluorescence excitation-emission spectrum of Barberton source water. [DOC] = 4.47 mg/L, SUVA254 = 4.31 L/mg.m...... 36

3.3 Fluorescence excitation-emission spectrum of Cleveland source water. [DOC] = 2.51 mg/L, SUVA254 = 1.17 L/mg.m...... 37

3.4 Modified schematic diagram of the TOX gas absorption system...... 45

3.5 Gradient profile for the analysis of Total organic halogen...... 49

3.6 Calibration curve for Chloride using 2,4,6-trichlorophenol. [Cl-] = 0 – 250 μM..50

3.7 Calibration curve for Iodide using 4-iodophenol. [I-] = 0 – 50 μM...... 51

3.8 Calibration curve for Bromide using 4-iodophenol. [Br-] = 0 – 50 μM...... 52

3.9 Calibration curve for Iodide using KI. [I-] = 0 – 100 μM...... 53

3.10 Gradient profile for the analysis of iodate...... 54

3.11 Calibration curve for Iodate using NaIO3. [IO3-] = 0 – 20 μM...... 54

3.12 Calibration curve for CHCl3using chloroform. [CHCl3] = 0 – 1000 nM...... 57

3.13 Calibration curve for CHBr2Cl using dibromochloromethane. [CHBr2Cl] = 0 – 300 nM...... 57

3.14 Calibration curve for CHBrI2 using bromodiiodomethane. [CHBrI2] = 0 – 125 nM...... 58

3.15 Calibration curve for CHClI2 using chlorodiiodomethane. [CHClI2] = 0 – 250 nM...... 58

x

3.16 Calibration curve for CHBr2I using dibromoiodomethane. [CHBr2I] = 0 – 250 nM...... 59

3.17 Calibration curve for CHBrClI using bromochloroiodomethane. [CHBrClI] = 0 – 250 nM...... 59

3.18 Calibration curve for CHBr3 using bromoform. [CHBr3] = 0 – 500 nM...... 60

3.19 Calibration curve for CHCl2I using dichloroiodomethane. [CHCl2I] = 0 – 500 nM...... 60

3.20 Calibration curve for CHCl2Br using bromodichloromethane. [CHCl2Br] = 0 – 400 nM...... 61

3.21 Calibration curve for CHI3 using iodoform. [CHI3] = 0 – 50 nM...... 61

3.22 Calibration curve for CAN using chloroacetonitrile. [CAN] = 0 – 500 nM...... 62

3.23 Calibration curve for DCAN using dichloroacetonitrile. [DCAN] = 0 – 500 nM...... 62

3.24 Calibration curve for TCAN using trichloroacetonitrile. [TCAN] = 0 – 125 nM...... 63

3.25 Calibration curve for BAN using bromoacetonitrile. [BAN] = 0 – 125 nM...... 63

3.26 Calibration curve for DBAN using dibromoacetonitrile. [DBAN] = 0 – 250 nM...... 64

3.27 Calibration curve for BCAN using bromochloroacetonitrile. [BCAN]=0–250 nM...... 64

3.28 Calibration curve for IAN using iodoacetonitrile. [IAN] = 0 – 31 nM...... 65

3.29 Calibration curve for CAA using chloroacetic acid. [CAA] = 0 – 250 nM...... 65

3.30 Calibration curve for DCAA using dichloroacetic acid. [DCAA] = 0 – 500 nM...... 66

3.31 Calibration curve for TCAA using trichloroacetic acid. [TCAA] = 0 – 250 nM...... 66

3.32 Calibration curve for BCAA using bromochloroacetic acid. [BCAA]=0–250 nM...... 67

3.33 Calibration curve for BDCAA using bromodichloroacetic acid. [BDCAA] = 0 – 250 nM...... 67

3.34 Calibration curve for BAA using bromoacetic acid. [BAA] = 0 – 1000 nM...... 68

3.35 Calibration curve for DBAA using dibromoacetic acid. [DBAA] = 0 – 500 nM...... 68

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3.36 Calibration curve for IAA using iodoacetic acid. [IAA] = 0 – 125 nM...... 69

4.1 TOI degradation as a function of pH in reaction mixtures containing iopamidol and aqueous chlorine ([Cl]T=100 μM, [iopamidol]=5 μM, [Buffer]=1 mM, and temperature= 25°C). Error bars represent 95% confidence intervals...... 72

4.2 Observed pseudo-first order loss of TOI as a function of pH. [Cl2]T = 100 μM, [Iopamidol] = 5 μM, [Buffer]T = 1 mM, Temperature = 25°C...... 73

4.3 Iodate formation as a function of pH in reaction mixtures containing iopamidol and aqueous chlorine. [Cl]T=100 μM, [iopamidol]=5 μM, [Buffer]=1 mM, and temperature= 25°C. Error bars represent 95% confidence intervals...... 75

4.4 THM and HAA formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 6.5. [Cl2]T = 100 μM, [Iopamidol] = 5 μM, [Buffer]T = 4 mM, and temperature = 25C. Error bars represent 95% confidence intervals...... 77

4.5 THM and HAA formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 7.5. [Cl2]T = 100 μM, [Iopamidol] = 5 μM, [Buffer]T = 4 mM, and temperature = 25C. Error bars represent 95% confidence intervals...... 78

4.6 THM and HAA formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 8.5. [Cl2]T = 100 μM, [Iopamidol] = 5 μM, [Buffer]T = 4 mM, and temperature = 25C. Error bars represent 95% confidence intervals...... 79

4.7 TOCl formation as a function of pH in reaction mixtures containing iopamidol and aqueous chlorine. [Cl2]T = 100 μM, [Iopamidol] = 5 μM, [Buffer]T = 4 mM, and temperature = 25C. Error bars represent 95% confidence intervals...... 81

4.8 TOI loss as a function of pH in reaction mixtures containing iopamidol and monochloramine. [NH2Cl] = 100μM, [iopamidol] = 5μM, [Buffer] = 1mM, and temperature = 25°C. Error bars represent 95% confidence intervals ………...…..…82

4.9 Iodide formation as a function of pH in reaction mixtures containing iopamidol and monochloramine. [NH2Cl] = 100μM, [iopamidol] = 5μM, [Buffer] = 1mM, and temperature = 25°C. Error bars represent 95% confidence intervals...... 83

- - 4.10 TOI, I , and IO3 mass balance in reaction mixtures containing iopamidol and aqueous chlorine at pH 6.5. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Error bars represent 95% confidence intervals...... 84

- - 4.11 TOI, I , and IO3 mass balance in reaction mixtures containing iopamidol and aqueous chlorine at pH 8.5. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Error bars represent 95% confidence intervals...... 85

4.12 TOCl formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 6.5. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Error bars represent 95% confidence intervals...... 86

xii

4.13 TOCl formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 8.5. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Error bars represent 95% confidence intervals...... 87

4.14 THM and HAA formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 6.5. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Error bars represent 95% confidence intervals...... 88

4.15 THM and HAA formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 8.5. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Error bars represent 95% confidence intervals...... 89

4.16 Proportion of iodinated DBPs in TOI at pH 6.5 at (a) 12 hr (b) 24 hr (c) 48 hr and (d) 72 hr. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Unknown T.P. is the unknown transformation products (remaining TOI)………………….…………………………………………...... ……92

4.17 Proportion of iodinated DBPs in TOI at pH 8.5 at (a) 12 hr (b) 24 hr (c) 48 hr and (d) 72 hr. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Unknown T.P. is the unknown transformation products (remaining TOI)….…………………………………………………………………..93

4.18 Proportion of chlorinated DBPs in TOCl at pH 6.5 at (a) 12 hr (b) 24 hr (c) 48 hr and (d) 72 hr. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C..…………………………………………………………………94

4.19 Proportion of chlorinated DBPs in TOCl at pH 8.5 at (a) 2 hr (b) 6 hr (c) 12 hr (d) 24 hr (e) 48 hr and (f) 72 hr. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C…………..……………………………………...95

4.20 TOI loss in chlorinated Akron source water as a function of pH. [Cl2]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 5.57 mg/L. Error bars represent 95% confidence intervals...... 97

4.21 TOI loss in chlorinated Barberton source water as a function of pH. [Cl2]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 4.47 mg/L. Error bars represent 95% confidence intervals...... 98

4.22 TOI loss in chlorinated Cleveland source water as a function of pH. [Cl2]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 2.51 mg/L. Error bars represent 95% confidence intervals...... 99

4.23 TOCl formation in chlorinated Akron source water as a function of pH. [Cl2]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 5.57 mg/L. Error bars represent 95% confidence intervals...... 101

4.24 TOCl formation in chlorinated Barberton source water as a function of pH. [Cl2]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 4.47 mg/L. Error bars represent 95% confidence intervals...... 102

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4.25 TOCl formation in chlorinated Cleveland source water as a function of pH. [Cl2]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 2.51 mg/L. Error bars represent 95% confidence intervals...... 103

4.26 TOI degradation in chloraminated Akron source water as a function of pH. [NH2Cl]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 5.57 mg/L. Error bars represent 95% confidence intervals...... 105

4.27 TOI degradation in chloraminated Barberton source water as a function of pH. [NH2Cl]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 4.47 mg/L. Error bars represent 95% confidence intervals...... 106

4.28 TOI degradation in chloraminated Cleveland source water as a function of pH. [NH2Cl]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 2.51 mg/L. Error bars represent 95% confidence intervals...... 107

4.29 TOCl formation in chloraminated Akron source water as a function of pH. [NH2Cl]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 5.57 mg/L. Error bars represent 95% confidence intervals...... 108

4.30 TOCl formation in chloraminated Barberton source water as a function of pH. [NH2Cl]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 4.47 mg/L. Error bars represent 95% confidence intervals...... 109

4.31 TOCl formation in chloraminated Cleveland source water as a function of pH. [NH2Cl]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 2.51 mg/L. Error bars represent 95% confidence intervals...... 110

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CHAPTER I

INTRODUCTION

1.1 Background

The quality of drinking water is vital for public health and safety. Although the quantity of water available for human consumption is small relative to the total volume of water on earth, water resources are continuously subjected anthropogenic contamination from point and non-point sources. Contaminants from both chemical and microbiological sources in drinking water pose threats to public health or cause undesirable aesthetic properties (Post et al., 2011). In order for water to be safe for human consumption, it requires a certain level treatment. Early treatment of water focused on aesthetic qualities which included taste, odour and turbidity. In the late

19th and early 20th century, drinking water quality further focused on disease-causing microbes in public water supplies (US EPA, 2000). Drinking water disinfection has been one of the major practices, which has been used since the 19th century (Zwiener and Richardson, 2005), to control microbial pathogens (biological contamination) responsible for the outbreak of waterborne diseases and protect public health. In the

United States of America (USA), disinfection has significantly reduced outbreaks of typhoid fever and cholera (McGuire, 2006).

Disinfection has been accomplished with disinfectants like chlorine, chloramines, chlorine dioxide, ozone and ultraviolet radiation. Apart from their use as disinfectants, the chemical oxidants can be used for the oxidation of taste and

1 odour compounds, micropollutant removal or transformation, and to improve coagulation of surface water (Bruchet and Duguet, 2004; Legube, 2003; von Gunten,

2003; Hoigné, 1998; Morris, 1986; Wolfe et al., 1984; Hoff and Gelderich, 1981). A survey conducted in 1998 and repeated in 2007 showed that many water utilities in

USA use multiple disinfectants (AWWA, 2008; 2000). Disinfection has contributed to the decline of waterborne diseases; however, the use of chemical disinfectants has lead to the formation of disinfection by-products (DBPs) (Richardson, 1998).

Specific DBPs have been linked to cancer of the bladder, stomach, pancreas, kidney and rectum (Bull et al., 1995; Koivusalo et al., 1994; Morris et al., 1992.

Trihalomethanes (THMs) in chlorinated drinking water were discovered in the

1970s (Bryant et al., 1992; Rook, 1974; Bellar et al., 1974). The US Environmental

Protection Agency (USEPA) passed the Safe Drinking Water Act (SDWA) and the

Total Trihalomethane (TTHM) Rule in 1974 and 1979 correspondingly in response to the chloroform report (Roberson, 2008). Also, the Information Collection Rule (ICR) required a collection of broad spectrum of water quality and treatment data including

THMs and haloacetic acids (HAAs) in both water treatment plants and distribution systems (Singer et al., 2002). THMs and HAAS accounted for more than 50% on weight basis of the chlorination by-products (USEPA, 1997). In 1998 and 2006, the stage 1 and stage 2 disinfectants and disinfection by-product (D/DBP) rules respectively (which superseded the TTHM rule) were enacted. The stage 1 D/DBP rule was based on a running annual average, which considered the results from all monitoring points. The Stage 2 D/DBP rule is based on locational running annual average. The maximum contaminant level (MCL) for TTHMs and five HAAs are respectively 80 μg/L and 60 μg/L. In addition, two inorganic DBPs, bromate and

2 chlorite, are regulated at MCLs of 0.01 mg/L and 1.0 mg/L respectively (US EPA,

2005; 1999).

DBPs are chemical compounds formed due to the reaction between disinfectants/oxidants and certain water matrix components (called precursors) or micropollutants (Krasner et al., 2006; Richardson, 2005; Plewa et al., 2004; Simmon et al., 2002; Cancho et al., 2000). There are over 600 identified DBPs (Richardson,

2011). Some of the major classes of DBPs are THMs, haloacetic acids (HAAs), haloacetonitriles (HANs), haloketones (HK), halonitromethanes (HNMs), haloacetamides (HAMs), haloacetaldehydes (HALs), cynogen halides (CNX), N- nitrosamines, oxyhalides, carboxylic acids, halogenated furanones, and halobenzoquinones (Richardson, 2011; Hrudey, 2009; Krasner et al., 1989; Backland et al., 1988; Bieber and Trehy, 1983; Miller and Uden, 1983; Christman et al., 1983;

Rook, 1974; Bellar et al., 1974). The speciation of DBPs may depend on the presence bromide and iodide in water matrix (Hua et al., 2006; Bichsel and von Gunten, 2000;

Krasner, 1999). Iodide is vital because its occurrence in source waters can result in the formation of iodo-DBPs which are known to be highly genotoxic and cytotoxic, with iodoacetic acid being the most genotoxic DBP (Richardson et al., 2008; Plewa et al., 2004).

There are different species of iodine that can be found in water – each exhibiting different mobility, bioavailability and chemical behaviour in the environment. The iodine species include iodide, iodate and organo-iodine (Hansen et al., 2011; Gilfedder et al., 2009). Hu et al. (2005) noted that both organic and inorganic iodine species have different hydrophilic and biophilic properties. Iodide naturally exists in seawater and is suspected to have a world-wide average concentration of 30 μg/L (Yokota et al., 2004). Iodide can also exist in freshwater as

3 well as in rain water (Gilfedder et al., 2009, 2008; Schwehr and Santschi, 2003).

Iodine concentration ranging from 0.5 to 212 μg/L has been detected in major U.S.,

Canadian and European rivers (Moran et al., 2002). Iodide is rapidly oxidised to HOI in the presence of chlorine (Nagy et al., 1988), chloramines (Kumar et al., 1986) and ozone (Garland et al., 1980). Chlorine and ozone can further oxidise HOI to form

- IO3 , the preferred sink for iodide (Bichsel and von Gunten, 1999b). Nonetheless,

HOI can react with NOM to form iodinated DBPs (Richardson et al., 2008; Krasner et al., 2006). Concentrations of natural iodide in source waters are reported to be very low or below detection limits in some cases to the extent that formation of iodo-DBPs was difficult to account for by the natural iodide (Richardson et al., 2008). Other possible source of iodide that contributes to iodo-DBP formation is iodinated x-ray contrast media (ICM) (Duirk et al., 2011).

ICM are large molecular (> 600 Da) triiodobenzoic acid pharmaceuticals which are used for imaging soft tissues, internal organs and blood vessels.

Administered dose to humans can be up to 200 g per diagnostic session (Pérez and

Barceló, 2007). Due to their highly hydrophilic property, they are resistant to human metabolism and are thus excreted through urine and feces un-metabolised within 24 hr

(Weissbrodt et al., 2009; Pérez and Barceló, 2007; Steger-Hartmann et al., 2000).

Therefore, they are detected at high concentrations in domestic and clinical wastewaters, surface waters (Weissbrodt et al., 2009; Busetti et al., 2008; Putschew et al., 2007; Seitz et al., 2006a; Ternes and Hirsch, 2000; Hirsch et al., 2000; Ternes,

1998), groundwater and bank filtrate (Schulz et al., 2008; Ternes et al., 2007; Sacher et al., 2001), soil leachates (Oppel et al., 2004) and drinking water supplies (Seitz et al., 2006b). ICM removal is negligible during conventional surface water treatment

(i.e. coagulation, flocculation, sedimentation, and filtration). On the contrary,

4 removal with activated carbon filtration has been achieved (Carballa et al., 2007;

Seitz et al., 2006b). Advanced oxidation process and ozonation have been found not to be effective at ICM removal (Bahr et al., 2007; Putschew et al., 2007; Seitz et al.,

2006a; Ternes et al., 2003). In municipal sewage treatment plants, ICM have been partly transformed/sorbed during nitrification with high sludge retention time (Schulz et al., 2008; Carballa et al., 2007; Batt et al., 2006). About 27 transformation products (TP) of iohexol, iomeprol and iopamidol (all ICM) were identified by

Kormos et al. (2009). Furthermore, 46 biotransformation products of iopromide, iohexol, iomeprol and iopamidol from aerobic soil-water and sediment water systems were detected by Kormos et al. (2010) and Schulz et al. (2008).

1.2 Problem Statement

ICM are known to be primary contributors to the total organic halogen (TOX) burden in clinical wastewater (Gartiser et al., 1996). Also, ICM are contributors of more than 90% of the adsorbable organic iodide in wastewater and surface water

(Putschew and Jekel, 2006; Putschew and Jekel, 2001; Putschew et al., 2001; Sprehe et al., 2001; Kummerer et al., 1998; Gartiser et al., 1996). Pharmaceuticals, oestrogen, textile dyes, personal care products, alkylphenol surfactants, diesel fuel, pesticides and UV filters are contaminants that form DBPs (Richardson, 2009). Due to the potential reactivity of contaminants containing activated aromatic benzene moieties with chlorine and other oxidants, Duirk et al. (2011) investigated ICMs as a potential source of iodine in iodo-DBPs found in chlorinated and chloraminated drinking waters. They observed iodo-acid and iodo-THMs were formed in 72 hr of reaction time due to the reaction of chlorinated oxidants with ICM in the presence of natural organic matter (NOM).

5

An occurrence study conducted in 23 cities revealed that up to 10.2 μg/L of iodo-THMs and 1.7 μg/L of iodo-acids were detected in chlorinated and chloraminated drinking water in USA and Canada (Richardson et al., 2008). Further re-examination of the source water in 10 out of the 23 cities showed that 4 out of 5 commonly used ICM were detected (Duirk et al., 2011). The detected ICM were iopamidol, iohexol, iopromide and diatrizoate. Iopamidol, the most predominant detected ICM, was sampled in 6 out of 10 treatment plants with concentrations up to

2700 ng/L.

Halogenated organic DBPs have been quantified as individual class species.

They can be quantified as total organic halogen. TOX is a group parameter which is an indication of the total amount of organic bound halogen in water (Dressman and

Stevens, 1983; Jekel and Roberts, 1980; Kuhn and Sontheimer, 1973). Specific TOX parameters include total organic chloride (TOCl), total organic bromide (TOBr) and total organic iodide (TOI). In chlorination and chloramination of natural water treatment, THMs and HAAs together account for about 50% and less than 20% of the

TOX respectively (Richardson, 2003; Li et al., 2002; Reckhow and Singer, 1984).

TOBr and TOI are formed when bromide and iodide are respectively in the natural water. Recent studies by Hua and Reckhow (2006) and Hua et al. (2006) reported the formation of TOCl, TOBr and TOI from chlorination of natural waters. Kristiana et al (2009) studied the formation and distribution of halogen-specific TOX in chlorination and chloramination of NOM isolates in the presence of bromide and iodide. Prior to the studies above, there had been studies on formation and behaviour of TOX from chlorination (Li et al., 2002; Baribeau et al., 2001; Pourmoghasddas and

Stevens, 1995) and chloramination (Wu et al., 2003; Diehl et al., 2000; Symons et al.,

1998) of water samples and NOM isolates. However, there have been no studies on

6 the transformation of TOI in the presence of NOM and chlorinated oxidants. Since

ICM are widespread at elevated levels in rivers and streams (Putschew et al., 2000), this research monitored the degradation of TOI (i.e. iopamidol, an ICM) in chlorinated and chloraminated oxidants in the presence and absence of NOM. Also mass balance on iodide, known DBPs and unidentified TOX were investigated.

1.3 Specific Objectives

In this research, the following specific objectives were considered:

1. Investigated the transformation of TOI (i.e., iopamidol) as a function of time

and pH in the presence of chlorinated oxidants. These experiments were

conducted in 2 phases with laboratory prepared deionized water. The study

was carried out at low concentrations of iopamidol, buffer solutions, and

chlorinated oxidants (detailed experimental procedures are in chapter 3) and at

high concentrations of iopmaidol, buffer solutions and chlorine. The

degradation of iopamidol and iopamidol transformation products was

monitored as TOI at pH 6.5 to 9.5 and reaction time of 0 to 72 hr. In addition

the formation of TOCl, iodate and iodide as well as THMs, HANs and HAAs

were accessed.

2. Investigated the transformation of iopamidol as a function of time and pH in

the presence of NOM with chlorinated oxidants. Three source waters were

obtained from the drinking water treatment plants intakes at the cities of

Akron, Barberton and Cleveland’s Garett Morgan Treatment Plant and were

filtered through 0.45 μm nylon membrane filter to remove particulate NOM.

The experiments were carried out at pH 6.5 to 8.5 and reaction time of 0 to 72

at 25°C. TOI loss as well as TOCl and iodate formation were monitored as

7

function of time and pH using aqueous chlorine. In addition the loss of TOI

and formation of TOCl in the presence of monochloramine were investigated.

3. Accessed the impact of dissolved organic carbon (DOC) variations on iodate

formation in the presence of aqueous chlorine as a function of pH. Barberton

and Cleveland source waters were used to conduct this study at pH 6.5 to 8.5.

8

CHAPTER II

LITERATURE REVIEW

2.1 Introduction

Bromide and both organic and inorganic iodide are found in source water matrix. Bromide in drinking water sources are due to contributions from seawater intrusion, geologic sources, seawater desalination, mining tailings, chemical production, sewage and industrial effluent s (Valero and Arbós, 2010; Richardson et al., 2007; Magazinovic et al., 2004; von Gunten, 2003). Also seawater intrusion, seawater desalination and dissolution of geologic sources are key contributing factors to iodide concentrations in drinking water sources (Agus et al., 2009; Hua et al.,

2006; von Gunten, 2003). The addition of chlorinated oxidants (aqueous chlorine and chloramines) to the water for the purpose of achieving microbial inactivation, though successful, has resulted in the formation of disinfection by-products (DBPs) in the presence of precursors like natural organic matter (NOM) and halides. In this chapter a review of literature is carried out on iodinated contrast media and their transformation, chlorinated oxidants used in water treatment, the chemistry of iodine, total organic halogen and toxicity of DBPs

2.2 Iodinated X-ray Contrast Media

Iodinated x-ray contrast media (ICM) are derivatives of 2,4,6-triiodobenzoic acid. They have recorded enormous usage in the medical sector especially radiology

9 for x-ray diagnostic imaging of soft tissues like organs, veins and blood vessels. They are large molecules (> 600 Da) with approximately 3.5x106 kg/year global consumption. Out of the total global consumption, Germany alone uses approximately 5x105 kg/year (Steger-Hartmann et al., 1999). Speck and Hübner-

Steiner (1999) reported that about 100 g of ICM is administered for each medical examination and up to 200 g per diagnosis (Perez and Barcelo, 2007). Steger-

Hartmann et al. (2000) reported that the estimated amount of ICM consumed in the

United States (US) in 1999 was 1330 t (1.33 x 106 kg).

The side chains of the ICM are comprised of hydroxyl, carboxyl and amide moieties (Figure 2.1) to impart elevated polarity and aqueous solubility (Krause and

Schneider, 2002). Due to their biological and chemical stability and inertness, they are excreted from the body unmetabolised within a day (Perez et al., 2006). They are resistant to conventional wastewater and drinking water treatment processes with 10% removal recorded (Drews et al., 2001; Ternes and Hirsch, 2000; Hirsch et al., 2000).

Also, samples have be detected in relatively high concentrations (>1 μg/L) in aqueous environments like creeks, rivers, effluents of wastewater and surface water in the world (Drews et al., 2001; Ternes and Hirsch, 2000; Hirsch et al., 2000; Putschew et at., 2000).

According to the World Health Organisation (WHO) Collaborating Centre for

Drug Statistics Methodology, there are over 35 ICMs and these can be categorized as water soluble, nephrotropic, high osmolar ICM; water soluble, nephrotropic, low osmolar ICM; water soluble, hepatotropic ICM; and non-water soluble ICM.

Examples of water soluble, nephrotropic, high osmolar ICM include diatrizoic acid, metrizoic acid, iodamide, iotalamic acid, ioglicic acid, acetrizoic acid, iocarmic acid, methiodal, and diodone. Metrizamide, iohexol, ioxaglic acid, iopamidol, iopromide,

10 iotrolan, ioversol, iopentol, iodixanol, iomeprol, iobitridol and ioxilan are examples of water soluble, nephrotropic, low osmolar ICM. Iodoxamic acid, iotroxic acid, ioglycamic acid, adiopiodone, iobenzamic acid, iopanoic acid, iocetamic acid, sodium iopodate, tyropnoic acid and calcium iopodate are in the water soluble, hepatotropic

ICM category. Some of the non-water soluble ICM are ethyl ester of iodised fatty acid, iopydol, propyliodone, iofendylate (http://www.whocc.no). The five ICMs were chosen due their frequent occurrence and detection in water sources and wastewater.

Figure 2.1: The chemical structures of ICM of common usage in hospitals

2.2.1 Occurrence and Concentration of ICM in Water and Wastewater

The increasing usage of ICM is alarming because Gartiser et al. (1996) in their research identified ICM compound as main contributors to the burden of total adsorbable organo-iodine (AOI) in clinical wastewater. ICM have been detected in wastewater from medical imaging facilities (Ziegler et al., 1997; Gartiser et al.,

1996). Of the 69 pharmaceutical agents that were suspected to be in wastewater from hospitals, McArdell et al (2010) detect 52 active pharmaceutical agents with

11 iopamidol recording the highest concentration in the mg/L range. This confirms the high administration of iopamidol in the medical imagining facilities. When the wastewater from the medical imagining facilities was treated with membrane bioreactor, many of the pharmaceuticals including ICMs (iopamidol, diatrizoate, ioxitalamic acid, iomeprol and iopromide) were removed to less than 20%. However, when powdered activated carbon (PAC) was used for treatment, about 70% removal was achieved for iopamidol although about 900 μg/L of iopamidol remained in the effluent. The concentration of diatrizoate, ioxitalamic acid, and iomeprol remaining in the wastewater after treatment with PAC was above 100 μg/L while the concentration of iopromide was 8.5 μg/L. This was not unexpected due to the high polarity of the ICMs resulting in ineffective adsorption (Steger-Hartmann et al.,

1999). In Berlin, Oleksy-Frenzel et al. (2000) found concentrations of ICM up to 100

μg I/L in municipal treatment plant effluents.

In Germany, where extensive research has been conducted on ICM, the most detected ICM in sewage effluent were diatrizoate and iopromide with maximum concentrations of 15 and 21 μg/L respectively (Putschew et al., 2001). In addition, iopamidol, iomeprol and iohexol have been found in sewage effluent (Putschew et al.,

2001). Also Ternes and Hirsch (2000) detected iopamidol, diatrizoate, iothalamic acid, ioxithalamic acid, iomeprol and iopromide in the influents of municipal wastewater treatment plant (WWTP) at almost the same concentration. Iopromide recorded the highest concentration of 7.5 μg/L. Since municipal and sewage treatment plants discharged their effluents into rivers and creeks, a median iopamidol and diatrizoate concentrations of 0.49 μg/L and 0.23 μg/L respectively has been detected in the receiving water bodies (Ternes and Hirsch, 2000).

12

2.2.2 Transformation of ICM

The fate of contaminants of emerging environmental concern including pharmaceuticals in source water and wastewater includes partitioning and transformation (Adams, 2009). Partitioning processes include adsorption (onto activated carbon) and membrane separation (Grassi et al., 2012). The transformation processes include aerobic and anaerobic biodegradation, hydrolysis, and chemical oxidation with chlorine, ozone, or advanced oxidation (Adams, 2009). Detection of these compounds at lower concentrations (ng/L) is accomplished with sophisticated and sensitive analytical methods and instruments (Peck, 2006; Kolpin et al., 2002;

Daughton and Ternes, 1999; Halling-Sprensen et al., 1998).

The transformations of ICM have been investigated by many researchers.

Despite their relative stability in humans ICM are subject to chemical and biological transformation in the environment (Schulz et al. 2008; Batt et al., 2006; Loffler et al.,

2005; Putschew et al., 2000). Schulz et al. (2008) studied the biotransformation and the transformation products of iopromide in water/soil systems. Using high performance liquid chromatography-ultraviolet (HPLC-UV) and liquid chromatography (LC) tandem mass spectrometry (MS) they identified 12 transformation products.

In another research, Kormos et al (2010) investigated the biotransformation products of ICM in aerobic soil-water and river sediment-water batch systems in

Germany. The soils used were loamy sand soil with organic matter content of 2.3% and the upper ploughed agricultural soil layer with 0.9% organic matter content. The soil had been irrigated and treated with secondary treated wastewater effluent and sludge for about 50 years. The groundwater used was collected from a deep well.

HPLC-UV and LC/MS were employed in the identification of the transformation

13 products. There was no biotransformed product detected with diatrizoate. However, at neutral pH, 11, 15 and 8 biotransformation products were detected from the transformation of iohexol, iomeprol and iopamidol respectively.

2.3 Reactions of Chlorinated Oxidants Used in Water Treatment

Chemical oxidation processes are used in water treatment to oxidise pollutants of concern to their terminal end product (CO2 and H2O) or to intermediate products that are more readily biodegradable or of less toxicological effect (Nriagu and

Simmons, 1994). Chemical oxidants are used to transform organic compounds into harmless forms and oxidise insoluble inorganic metals for precipitation (Crittenden et al., 2012). Also chemical oxidants have found tremendous use in the disinfection of drinking water as well as wastewater. In addition, oxidation processes are used for the removal of taste and odour compounds like geosmin and 2-methylisoborneol

(MIB) (AWWARF, 1987). Some oxidants used in drinking water treatment include chlorine, chlorine dioxide, chloramines (monochloramine, dichloramine and trichloramine), ozone, hydrogen per oxide and potassium per manganate (Crittenden et al., 2012; Nriagu and Simmons, 1994). Due to their reactivity, chemical oxidants can form potentially harmful by-product (Krasner et al., 2006; Simmons et al., 2002;

Bichsel and von Gunten, 2000). This research focuses on chlorinated oxidants.

2.3.1 Chlorine

Chlorine is a widely used oxidant which can exist in the gaseous (Cl2), liquid

(NaOCl) or solid (Ca(OCl)2) states (AWWA, 2011). Gaseous chlorine rapidly hydrolyses to form hypochlorous acid (HOCl) (eq. 2.1) which can also dissociate to form hypochlorite (OCl-) (eq. 2.2). The total concentration of aqueous chlorine

14 existing as HOCl or OCl- is known as the free chlorine. The speciation of chlorine is dependent on the pH of the solution. The main chlorine species found at a pH range of 6 to 9 under typical water treatment conditions are HOCl and OCl- (Deborde and von Gunten, 2008).

Morris (1978) indicated that hypochlorous acid is the major reactive form during water treatment since the other species of chlorine are present in low or insufficient concentrations for significant reaction. At low pH Cherney et al (2006) also argued that Cl2(aq) is the most probable reactive chlorine species.

Chlorine reacts with both organic and inorganic compounds. In their paper,

Deborde and von Gunten (2008) reported that most kinetics of the oxidation reactions of chlorine with inorganic and organic compounds followed a second order reaction – first order with respect to total chlorine ([HOCl]T) and first order with respect to total compound ([X]T) as shown in eq. 2.3. There are significant variations in the reactivity of HOCl and OCl- for any given compound. With reference to other authors

(Armesto et al., 1994a; Rebenne et al., 1996; Abia et al., 1998; Gallard and von

Gunten, 2002; Gallard et al., 2004; Deborde et al., 2004; Dodd et al., 2005) Deborde and von Gunten (2008) further indicated that for chlorination reactions the apparent second order rate constant is pH dependent.

kapp is the apparent second order rate constant

- - [HOCl]T = [HOCl] + [OCl ] and [X]T = [HX] + [X ]

15

When ammonia is present in water, chlorination results in the formation of monochloramine (NH2Cl), dichloramine (NHCl2) and trichloramine (NCl3) (Qiang and Adams, 2004; Jafvert and Valentine, 1992; Morris and Isaac, 1983). Deborde and von Gunten (2008) indicated that as the number of atoms of chlorine on the nitrogen increased, the reactivity of chlorine decreased – a confirmation of the presumed initial mechanism of electrophilic attack of HOCl on the nitrogen (Jafvert and Valentine,

1992; Morris, 1978).

In addition, HOCl is the predominant reactive species in the presence of other halides. Also HOCl oxidises other inorganic compounds like sulphite, cyanide and nitrite via an electrophilic attack of HOCl (Johnson and Margerum , 1991; Gerritsen andMargerum, 1990). Deborde and von Gunten (2008) concluded that weak variations of nucleophilicity of the inorganic compounds induces strong changes in

HOCl reactivity and a high sensitivity of chlorine reactivity with regard to the nucleophilic character can be anticipated.

The dominance of HOCl species is also evident in its reactions with organic compounds. The possible pathways reactions include oxidation, addition and electrophilic substitution. As a result of its high selectivity, HOCl has a restricted reactivity to limited site (Deborde and von Gunten 2008). Furthermore HOCl form a more oxidised or chlorinated compound due to its capability to induce modifications in the parent molecular structure (Dore, 1989). Chlorine reacts with aromatic compounds and other moieties bound to the aromatic ring by electrophilic substitution with an initial reaction occurring primarily in ortho or para position to a substituent

(Deborde and von Gunten, 2008; Roberts and Caserio, 1968). The influence of the substituents on the aromatic ring on substitution reaction cannot be overemphasised.

Deborde and von Gunten (2008) explained that faster substitution reaction is due to

16 the properties of the electron donor of the substituent that increases the charge density of the aromatic ring.

2.3.2 Chlorine Dioxide

Chlorine dioxide (ClO2) is a stable free radical and potent oxidant (Sharma,

2008). It has slow decomposition in neutral aqueous solution (Odeh et al., 2002) but accelerated degradation in basic solution (Sharma, 2008). Gates (1998) explained that the use of ClO2 is restricted to high quality water with low dosage (1.0 to 1.4 mg/L in

USA). It goes through a wide variety of redox reactions with organic matter to form oxidized organics and reduced chlorine species (Singer and Reckhow, 1999). ClO2 also rapidly oxidizes inorganic species like Fe (II) and Mn (II), but this is limited by the presence of organic matter due to competition between the organics and metals for the oxidants as well as the possible formation of metal-organic complexes (Knocke et al., 1990; van Benschoten et al., 1992). ClO2 reactivity with both organic and inorganic compounds also follows a second order reaction – first order with respect to

ClO2 and first order with respect to the organic or inorganic compound (Hoigne and

Bader, 1994). Aromatic compounds, hydrocarbons, carbohydrates, aldehydes, acetone and primary and secondary amine compounds are unreactive with ClO2.

ClO2 reacts selectively with phenols and the reaction is influenced by pH (Sharma,

2008; Hoigne and Bader, 1994). Compared with free chlorine, ClO2 has much lower tendency to produce chlorinated DBPs especially THMs and HAAs (Benjamin and

Lawler, 2013).

17

2.3.3 Chloramines

Chloramines have found use in water treatment for disinfection purpose.

Chloramine disinfectant is produced by substitution reaction between free chlorine and NH3. As stated above chloramines includes NH2Cl, NHCl2 and NCl3. In typical water treatment conditions, NH2Cl is the predominant species over the drinking water pH range of 6.5 – 8.5 (Vikesland et al., 1998). Although NH2Cl has the same oxidising capacity as free chlorine, it is a weaker disinfectant (Wolfe et al., 1984).

On the contrary, it has been shown that NH2Cl is unstable at neutral pH even in the absence of organic and inorganic compounds. Also it undergoes a series of reactions known as auto-decomposition that result in the oxidation of ammonia and reduction of active chlorine (Jafvert and Valentine, 1992). These reactions to a large extent depend on the pH of the solution and the chlorine to ammonia nitrogen ratio – larger ratio results in faster oxidation of ammonia (Vikesland et al., 2000).

A monochloramine concentration of 0.5 – 2 mg/L has been detected in water supply systems where monochloramine was used as the primary disinfectant or to provide chlorine residual in the distribution system (Bull et al., 1991). Vikesland and

Valentine (2000) showed that NH2Cl reacted in solution with Fe (II) through a direct interaction between molecular monochloramine and aqueous ferrous iron. They further indicated that they are autocatalytic reactions since the iron oxide product of the aqueous-phase reaction sped up the overall reaction kinetics enabling the formation of highly reactive ferrous iron surface complex. NH2Cl also react with dimethylamine (NMA), an organic substance found in water, to produce N- nitrosodimethylamine (NDMA) by the oxidation of NMA to unsymmetrical

18 dimethylhydrazine as an intermediate and further oxidation to NDMA (Mitch and

Sedlak, 2002; Choi and Valentine, 2001).

2.4 Chemistry and Reactions of Iodine

Iodine (I) is a halogen with atomic number 53. Elemental iodine is slightly soluble in water (1.18 x 10-3 mol/L at 25°C) (Burgot, 2012) but solubility increases

- with the addition of alkali iodide to form triiodide (I3 ) and polyiodides. Iodine is freely soluble in organic solvent (Burgot, 2012). Aqueous iodine species known are elemental iodine (I2) is used as a disinfectant and has proven to be effective and economical (Gottardi, 1983). Iodine as a disinfectant is often applied in drinking water disinfection in emergency situations like floods and earthquakes (Bichsel,

2000). Despite its effectiveness in disinfection, it has drawbacks comparable to disinfectant like chlorine.

A system comprised of iodine and water can undergo different equilibria (eq.

2.4 – 2.11) (Clough and Starke, 1985). The aqueous iodine species are shown in table

2.1. In the aqueous system the equilibrium is influenced by pH and iodide ions.

Equations 2.4 to 2.10 have fast reaction rate, that is, they occur instantaneously.

However disproportionation of HOI to form iodate is relatively slow with a rate highly influenced by pH and iodide concentration (Gottardi, 1981). Iodine is hydrolysed to form HOI and I- (eq. 2.4). High pH values results in the dissociation of

HOI (eq. 2.5) to OI- (pKa = 10.4) (Bell and Gelles, 1951). Further disproportionation reaction of HOI (eq. 2.11) forms iodate and iodide with the equilibrium shifting more to the right at environmental conditions (pH ≥ 6, total iodine < 2 μM) (Bichsel, 2000).

In addition, iodic acid is formed by the protonation of iodate (Pethybridge and Prue,

19

- - 1967) while the electrochemical oxidation of IO3 on PbO2 anode forms IO4

(Greenwood and Earnshaw, 1984)

I2 H2O HOI I H 2.4

HOI OI H , p a = 10.4 2.5

I2 I I3 2.6

I3 I2 I5 2.7

2 2I3 I6 2.8

OI I H2O HI2O OH (2.9)

HI2O I2O H 2.10

3HOI IO3 2I 3H (2.11)

Table 2.1: Aqueous Iodine species

Chemical Formula Customary name IUPAC name (Leigh, 1990) Valance

I- Iodide Iodide (-1) -I

I2 Iodine Diiodine 0

- I3 Triiodide Triiodide(-1) -1/3

HOI Hypoiodous acid Hydrogen oxoiodate +I

OI- Hypoiodite Oxoiodate(-1) +I

- IO2 Iodite Dioxoiodate(-1) +III

HIO3 Iodic acid Hydrogen trioxoiodate +V

- IO3 Iodate Trioxoiodate(l-) +V

- IO4 Periodate Tetroxoiodate (-1) +VII

Adapted from Bichsel, 2000

A series of reaction mechanism (eq. 2.12 and 2.13) is used to describe equation 2.11 (overall reaction). The rate limiting step is either equation 2.14 or 2.15.

At pH > 5, the reaction in eq. 2.11 is forced to the right (Bichsel and von Gunten,

20

1999b). The equilibrium constant for the overall reaction is 6x10-11 (Myers and

Kennedy, 1950). The kinetics of the reaction is second order with respect to [HOI]T

- ([HOI]T = [HOI] + [OI ]) (Urbansky et al., 1997; Truedale, 1997; Wren et al., 1986;

Thomas et al., 1980).

The concentration of total iodine in water resources is usually in between 0.5 –

10 μg/L. However groundwater can show concentration in excess of 50 μg/L (Wong,

1991; Fuge and Johnson, 1986). The main species of iodine in fresh waters are I- and

- IO3 . During drinking water treatment both organic and inorganic iodide present in the water can be oxidised by chlorinated (aqueous chlorine, ClO2 and chloramines) and non-chlorinated (ozone) oxidants. The oxidation of I- in iodide-containing waters rapidly forms HOI as the first product in the presence of ozone (Garland et al., 1980), chloramines (Kumar et al., 1986) and chlorine (Nagy et al., 1988). Nonetheless the

- - chemistry of oxidation of I by ClO2 is different. ClO2 oxidises I to I radical (Fabian and Gordon, 1997).

Bichsel and von Gunten (1999b) investigated the stoichiometry of the reaction of HOCl/OCl- with I- at a pH range of 5.3 – 8.7 and a molar ratio of [HOCl]:[I-] = 4:1.

- - The first oxidation step from I to HOI occurred very fast. Formation of IO3 was

- - - measured together with the sum of [HOCl], [OCl ] and [HOI] as I3 (in excess of I ) by spectrophotometry. Every mole of I- reacted with 3.0±0.1 moles of HOCl/OCl- to

- produce 0.99±0.02 mol of IO3 (eq. 2.14 – 2.15). From their research it was assumed

- that no stable intermediate nor IO4 was formed.

21

In addition, Bichsel and von Gunten (1999) determined the rate constant for

- the oxidation of HOI by NH2Cl by measuring IO3 formation in a pH range of 7.2 –

8.5 in the presence of 0.005 – 1.0 mM NH2Cl and 0.1 μM HOI. NH2Cl is already known to oxidise I- to HOI in a relatively fast, pH-dependent reaction (Kumar et al.,

- 1986). Within the first 77 hr IO3 was detected (representing < 25% [HOI]0). The calculated maximum rate constant (kNH2Cl+HOI), if HOI was the reactive species was

2x10-3 M-1s-1. However, if OI- was the reactive species the maximum rate constant

-1 -1 (kNH2Cl+OI-) was 3 M s .

2.5 Total Organic Halogen Formation

Total organic halogen (TOX) in environmental analysis is a measure that represents the total amount of organically bound halogen in waters (Dressman and

Stevens, 1983; Jekel and Roberts, 1980; Kuhn and Sontheimer, 1973). It has been adopted as a surrogate measurement for the total halogenated disinfection by-products

(DBP) in drinking water formed from the reaction between chemical disinfectants and natural organic matter (NOM) (Reckhow and Singer, 1984; Luong et al., 1982). The halogen specific fractions of TOX include total organic chloride (TOCl), total organic bromide (TOBr) and total organic iodide (TOI).

The formation of DBP is initiated by the addition of a disinfectant to the water treatment train. Disinfectants used are chlorine, chloramines, ozone and ultraviolet

(UV) light. The reaction of free chlorine in water with water constituents can be described in four general pathways: oxidation, addition, substitution and catalysed or light decomposition (Gang et al., 2003; Johnson and Jensen, 1986). In addition and substitution reactions chlorine is added or substituted into the NOM molecular structure that produces chlorinated organic intermediates with further decomposition

22 resulting in DBP formation (van Hoof, 1992). If compounds containing double bonds are present in the water chlorine addition reaction with water is too slow unless double bonds are activated by substituent group (Brezonik, 1994). Brezonik (1994) further indicated that substitution reactions with chlorine are typically electrophilic.

Also Gordon and Bubnis (2000) reported of the slow process of the decomposition of OCl- in basic solution. The decomposition involves chlorite ion (eq.

2.16) as an intermediate (Adam and Gordon, 1999). Hypochlorite in a decomposition that is catalysed by transition metal ions like Ni(II), Cu(II) and Fe(II) (Gordon and

Bubnis, 2000) results in the formation of O2 (eq. 2.17).

In oxidation reaction, the molecule/compound being oxidised by chlorine donates two electrons to Cl+ radical to form Cl- (Gang et al., 2003). Oxidation reactions account for more than 90% of the chlorine demand in natural waters while the other chlorine reactions account for the remainder (Jolley and Carpenter, 1983). If bromide and iodide are present in the water matrix, a proposed mechanism is the transfer of Cl+ from HOCl to the halide (X-) to form an intermediate (XCl) which as a result of the hydrolysis produces OX- (Johnson and Margerum, 1991; Kumar and

Margerum, 1986). Thus bromide and iodide are rapidly oxidised to HOBr and HOI

- respectively. HOI is further oxidised to IO3 in the absence of NOM (Bichsel and von

Gunten, 1999a). On the contrary, in the presence of NOM, the active oxidants have the ability to react with NOM to form brominated and iodinated DBP in a way similar to HOCl (Bichsel and von Gunten, 2000; Symons et al., 1993; Rook, 1974) (fig 2.2).

Two major classes of DBPs are THMs and HAAs. About a total of 10 and 19 halogenated THMs and HAAs respectively can be formed during chlorination of

23 drinking water in the presence of bromide and iodide (Hua et al., 2006). THMs and

HAAs account for about 50% of the TOX formed during chlorination (Reckhow and

Singer, 1984). The presence of bromide in natural waters shifts the speciation of

THMs and HAAs from chlorinated to brominated species (Cowman and Singer, 1996;

Symons et al., 1993; Pourmoghaddas et al., 1993; Luong et al., 1982) due to the efficient substitution characteristics of HOBr (Westerhoff et al., 2004).

Figure 2.2: TOX formation and oxidation products

In their research to determine the effect of bromide on TOX formation at pH

7, Hua et al. (2006) found that TOCl concentration gradually decreased while TOBr gradually increased as the bromide concentration increased. The authors also found that increasing the concentration of iodide did not significantly change the concentration of TOCl for iodide concentration of 0 – 2 μM. However TOCl decreased sharply at iodide concentrations of 10 μM and 30 μM. The TOI

24 concentration also increased with increasing iodide concentration. To determine the effect of chlorine dose on I-THMs formation, Hua et al. (2006) spiked the source water (Tulsa water) with 2 μM iodide and chlorine concentration range of 0.5 to 5 mg/L at pH 7 for 48 hr. TOCl increased almost linearly with increasing chlorine concentration. TOBr increased with increasing chlorine up to 3 mg/L. At chlorine concentration of 0.5 mg/L TOI peaked and gradually decreased to 3 mg/L. However

- increasing chlorine concentration saw a significant increase in IO3 from 0.5 to 3 mg/L. The observations were attributed to the reactions shown in table 2.2.

Table 2.2: Reactions forming TOX and iodate

Equation Reaction Rate constant Reference 8 -1 -1 2.18 k1 = 4.3 x 10 M s Nagy et al., 1988 -1 -1 2.19 k2 = 1550 M s Kumar and Margerum, 1987

-1 -1 2.20 k3 = 8.2 M s Bichsel and von

Gunten, 1999b

-1 -1 2.21 k4 = 52 M s Bichsel and von

Gunten, 1999b -1 -1 2.22 HOCl + NOM Products K5 = 00.7 – 5 M s Westerhoff et al., 2004 -1 -1 2.23 HOBr + NOM Products K6 =15 – 167 M s Westerhoff et al., 2004 -1 -1 2.24 HOI + NOM Products (TOI) K7 =0.1 – 0.4 M s Bichsel and von Gunten, 2000

Although natural iodide in source waters was believed to be the primary source of iodo-DBPs (Bichsel and von Gunten, 2000; 1999a), Duirk et al. (2011) showed that organically bound iodide, in the presence of chlorinated oxidants can be released from the aromatic ring of an aromatic compound to be incorporated into the

25

NOM to form iodo-DBPs. In the absence of NOM, iodate was formed. The proposed reaction pathway is shown in fig 2.3.

Figure 2.3: Iodo-DBP formation pathway. (Adapted from Duirk et al., 2011)

2.6 Toxicity of Halogenated Disinfection By-Products

Chlorination of water supplies was introduced to inactivate harmful pathogenic microorganisms in the water to protect public health from risk of infection. While the goal of chlorination of water was successful (Akin et al., 1982)

DBPs were formed. The formation of DBP was not known until the early 1970s when

Rook (1974) reported of the formation of chloroform (Bryant et al., 1992; Bellar et al., 1974). More than 600 DBP have been identified in drinking water (Richardson et al., 2007). The types of DBPs formed are dependent on source water, pH, temperature, type of disinfectant used and the treatment processes (Krasner, 2009;

Richardson et al., 2007; Ueno et al., 1996). Majority of BDPs formed due to water disinfection have yet to be chemically defined (Richardson et al., 2002; Weinberg,

1999). Due to public health concerns the United States Environmental Protection

26

Agency (US EPA) regulates 11 DBP – they include 4 THMs, 5 HAAs, bromate and chlorite. The four THMs are chloroform (CHCl3), bromodichloromethane (CHBrCl2), dibromochloromethane (CHBr2Cl) and bromoform (CHBr3). Also the regulated

HAAs include monochloroacetic acid (MCAA), dichloroacetic acid (DCAA), trichloroacetic acid (TCAA), bromoacetic acid (BAA) and dibromoacetic acid

(DBAA). Each has been assigned a maximum contaminant level (Weinberg et al.,

2002). Nonetheless, the focus has been on THMs and HAAs as the most prevalent in drinking water and as the surrogates for other DBPs (US EPA, 2006). Drinking water

DBP represents a class of environmentally hazardous chemicals with long term health effects (Betts, 1998; Richardson, 1998).

Studies in epidemiology have linked elevated risk of cancer of the bladder, stomach, pancreas, kidney and rectum as well as Hodgkin’s and non-Hodgkin’s lymphoma to the consumption of chlorinated water (Bull et al., 1995; Koivusalo et al., 1994; Morris et al., 1992). Also Waller et al. (2001) and Nieuwenhuijsen et al.

(2000) have linked the increase in risk of spontaneous abortions and birth defects in human to DBP. Studies have further shown that concentrated extracts of drinking water samples were toxic in many in vivo and in vitro bioassays (Wilcox and

Williamson, 1986).

The genotoxicity of the regulated THMs has been studied (Kogevinas, et al.,

2010; Kargalioglu, et al., 2002). Kargaliouglu et al. (2002) observed that in strains of salmonella, in the presence of the enzyme, glutathione S-transferase theta (GSTT1-1), bromodichloromethane, dibromochloromethane and bromoform induced genotoxicity.

Richardson et al. (2007) also noted that bromodichloromethane, dicbromochloromethane and bromoform have no genotoxic induction response except in the presence of GSTT1-1. A study conducted by Plewa et al. (2002), focused on

27 mammalian cell cytoxicity and genotoxicity of brominated and chlorinated HAAs in

Chinese Hamster Ovary (CHO), they detected that BAA was the most genotoxic and cytotoxic. The brominated HAAs were more cytotoxic and genotoxic than the chlorinated analogues. DCAA and TCAA have been found to be mutagenic in mouse lymphoma cells (Harrington-Brook et al., 1998)

Later it was shown in a study to measure five iodo acids and two THMs in chlorinated and chloraminated drinking waters from 23 cities in United States of

America and Canada that iodinated DBPs are highly genotoxic and cytotoxic – iodoacetic acid was the most identified genotoxic DBP in mammalian cell

(Richardson et al., 2008). The iodo-THMs were less cytotoxic than the iodo-acids except for iodoform. Iodoacetic acid is highly cytotoxic and more genotoxic in mammalian cells than bromoacetic acid (Plewa et al., 2004). Furthermore, iodo-

THMS are expected to be more toxic than their brominated and chlorinated analogues. Duirk et al. (2011) in addition confirmed the cytotoxicity and genotoxicity of iodo-DBP in mammalian cells after dosing chlorinated or chloraminated Athens-

Clark County source water with iopamidol at pH 7.5. From their study the rank of iodo-DBP in descending order of cytotoxicity in chlorinated source water spiked with iopamidol was iodoacetic acid (IAA) > chlorodiiodomethane >dichloroiodomethane > iodoform > bromochloroiodomethane. The same ranking for chloraminated water was IAA > chlorodiiodomethane > dichloroiodomethane. Also they noted that iodo-

DBP induced the highest genotoxicity.

28

CHAPTER III

MATERIALS AND METHODS

3.1 Chemicals and Reagents

2, 4, 6 trichlorophenol (98%), 4-iodophenol (99%) and NaI (99%) were purchased from Sigma Aldrich (St. Louis, MO, USA). Also 2, 4, 6 tribromophenol

(98%) was purchased from Acros Organics (NJ, USA). Iopamidol was purchased from U.S. Pharmacopeia (Rockville, MD, USA). In addition, NaCl (99%) was purchased from EMD chemicals (Gibbstown, NJ, USA). NaBr (99.5%) and KI

(99.5%) were purchased from Fisher Scientific (NJ, USA). Commercial 10-15% sodium hypochlorite (NaOCl) which contained equimolar amounts of OCl- and Cl- was purchased from Sigma Aldrich (St. Louis, MO, USA). The standard soutions used for the disinfection by-product (DBP) included: iodoacetic acid and iodoform from Sigma Aldrich (St. Louis, MO, USA), haloacetic acid mix (containing various concentrations in methyl tert-butyl ether (MtBE) of monochloro-, monobromo-, dichloro-, trichloro-, bromochloro-, dibromo-, bromodichloro-, chlorodibromo-, and tribromoacetic acid) from Restek (Bellefonte, PA, USA), trihalomethane mix

(including chloroform, bromoform, bromodichloromethane, and dibromochloromethane) purchased through Chem Service (West Chester, PA, USA), iodo-THMs (dichloroiodo-, dibromoiodo-, bromochloroiodo-, chlorodiiodo-, and bromodiiodomethane) purchased from CanSyn Chem Corporation (Toronto, ON,

Canada), Chloro-, dichloro-, and trichloroacetonitrile purchased from Chem Service

29

(West Chester, PA, USA) bromoacetonitrile and dibromoacetonitrile purchased from

Arcos Organics (Geel, Belguim) and bromochloroacetonitrile and iodoacetonitrile purchased from Crescent Chemical and Alfa Aesar (Ward Hill, MA, USA) respectively. All DBPs were purchased at the highest possible purities. All other organic and inorganic chemicals used were certified American Chemical Society

(ACS) reagent grade and were used without further purification.

Deionized water prepared from a Barnstead ROPure Infinity/NANOPure system (Barnstead-Thermolyne Corp. Dubuque, IA, USA) was used to generate deionized water (18.2 MΩ.cm-1) for the experiments. Experimental pH was monitored with Orion 5 star pH meter equipped with Ross ultra combination electrode

(Thermo Fisher Scientific, Pittsburgh, PA, USA) and pH adjustments for the experiments were achieved with 0.1 N H2SO4 and 0.1 N NaOH. All glasswares and polytetrafluoroethylene (PTFE) were soaked in a chlorine bath or base bath for 24 hours, rinsed with large amount of deionized water and dried before use.

3.2 Source Water Characterization

Source waters for the experiments were sampled from the intake of Akron,

Barberton and Cleveland drinking water treatment plants in Ohio, USA. The Akron water treatment plant receives water from the Upper Cuyahoga River through three impounding reservoirs: East branch Reservoir, Wendell R. Ladue Reservoir, and Lake

Rockwell (Franklin, Portage County). Also water is taken directly from Lake Erie for treatment at the Garret Morgan Water Treatment Plant on the near Westside of

Cleveland. Water from the Upper Wolf Creek forms the Barberton reservoir which serves the Barberton water treatment plant (Norton, OH).

30

The Cuyahoga river watershed which is located in northeastern Ohio drains a total of 812 square miles (2103 km2) and flows through 6 counties. Akron, Cleveland and some of its suburb, Cuyahoga Falls and Kent are major municipalities partially or fully in the watershed. At the downstream of Cuyahoga Falls, the river turns abruptly northward and flows in a wide, deep preglacial valley to Cleveland and its mouth in

Lake Erie. Agricultural land uses like cultivated crops and forest are located on the eastern portion of the watershed whiles urban development, with some forest and pockets of hay and pasture lands are predominantly in the western portion of the watershed (http://www.epa.state.oh.us/dsw/tmdl/CuyahogaRiver.aspx).

The Upper Wolf Creek is a small headwater tributary to Tuscarawas River.

According to NEFCO (2011), “The Creek originates from Medina County and flows east into Summit County before forming the Barberton reservoir in the City of Norton and Copley Township.” The creek has ten tributaries, with all ten tributaries flowing into the Barberton Reservoir. Adjacent to the creek are forest, wetlands, shrub and/or old field lands. The watershed is bedeviled with developmental works as a result of its close proximity to Akron (east), Medina (west), Wadsworth (south) and Cleveland

(north) (NEFCO, 2011).

Lake Erie in Ohio covers 11,649 square miles (30,171 km2). About 72% of this land is agricultural or open space, 20% is wooded while slightly more than 2% remains wetland. Also other 4% accounts for the developed and urban environment use (includes industrial, commercial, residential, quarries, transportation and institutional). Inland lakes and rivers cover 1%. Dominant land use in the basin is crop agriculture. Due to its intensive land use, Lake Erie receives large loads of sediments, nutrients and pesticides to the surface waters (Ohio Department of Natural

Resources).

31

Source water characteristics from the three water treatment plants in the three cities are shown in table 3.1. Total organic carbon (TOC) concentrations were measured using Shimadzu TOC analyzer (Shimadzu Scientific, Columbia, MD, USA) and calibrated according to Standard Method 505A (APHA et al, 1992). The ultraviolet absorbance at 254 nm (UV254) and spectral characteristics of the NOM were measured with Shimadzu UV 1601 UV visible spectrophotometer in accordance with Standard Method 5910B (APHA et al, 1998). The specific ultraviolet absorbance at 254 nm (SUVA254) was calculated from the relation:

. DBP formation has been linked to water characteristics like

SUVA254, bromide concentration and DOC concentration (Njam et al., 1994).

Table 3.1: Source water characteristics from Akron, Barberton and Cleveland water

Akron source Barberton Cleveland

water source water source water

DOC (mg/L C) 5.57 4.47 2.51

Bromide (µM) 1.6 2.0 < 0.5

Iodide (µM) < 0.5 < 0.5 < 0.5

-1 UV254 (cm ) 0.121 0.132 0.029

SUVA254 (L/mg-m) 2.17 3.08 1.17

The source waters were further characterized using florescence spectroscopy, which yielded the excitation-emission matrix (EEM) spectra. Parlanti et al. (2000) used ratios of florescence EEM peak intensities to track NOM changes in natural water. The preparation of the samples for the florescence spectra detection followed the method developed by Chen et al. (2003) with slight modifications. The water

32 samples were acidified with sulfuric acid to lower the pH to 2.75 – 3.25 to remove inorganic carbon. Also, the samples were diluted to a final DOC of 1 mg/L with 0.01

M KCl to allow direct comparisons of fluorescence intensities (Nguyen et al., 2005).

The EEM florescence spectra were obtained with an F-7000 FL fluorescence spectrophotometer (Hitachi Hi-Tech, Tokyo, Japan). The spectrophotometer uses xenon lamp as its light source. The excitation slit as well as emission slit were set to a band-pass of 10 nm. The spectra of the source water samples were measured at successive emission spectra at 2 nm intervals across the range 290 to 550 nm and using excitation wavelengths spaced at 5 nm from 204 to 404 nm. The resulting spectra were then merged into the EEM and constructed using SigmaPlot 12.0 (SPSS

Inc.) to generate contour maps of the fluorescence intensity with the regional integration (Figures 3.1 - 3.3). Florescence regional integration (FRI) was proposed by Chen et al. (2003) to quantify multiple broad-shaped EEM peaks. The FRI is a quantitative technique which integrates volume under EEM region (Table 3.2). In addition, the FRI technique has been used to quantitatively analyze all wavelength- dependent florescence intensity data from EEM spectra (Marhuenda-Egea et al.,

2007). The five distinctive regions proposed by Chen et al. (2003) are indicated in table 3.2. The five regions found in the NOM EEM of the three source waters are also shown in table 3.3.

33

Table 3.2: Florescence EEM regions proposed by Chen et al. (2003)

Excitation Emission Range

Regions Representation Range (nm) (nm)

I Aromatic 200 – 250 280 – 330

II Aromatic protein-like 200 – 250 330 – 380

III Fulvic acids 200 – 250 380 – 550

IV Soluble microbial by-products 250 – 400 280 – 380

V Humic acids 250 – 400 380 – 550

Table 3.3: Florescence regions for Akron, Barberton and Cleveland source waters for 1 mg/L C

Akron Barbertion Cleveland Fluorescence Regions % % %

Aromatics (I) 1.9 14.7 1.8 8.3 2.4 22.7

Aromatic Protein-Like (II) 3.4 25.9 5.5 26.0 3.3 31.1

Fulvics (III) 5.1 38.5 9.1 42.7 3.0 28.7

Microbial (IV) 1.1 8.5 2.0 9.5 1.2 11.2

Humics (V) 1.6 12.3 2.8 13.4 0.7 6.3

Total 13.1 100 21.2 100 10.5 100

34

400 0 10 20 30 40 350 50 60

300 IV V

Excitation (nm) Excitation

250 III I II

300 320 340 360 380 400 420 440 460 480 500 520 540

Emission (nm)

Figure 3.1: Fluorescence excitation-emission spectrum of Akron source water. [DOC] = 5.57 mg/L, SUVA254 = 2.27 L/mg.m

35

400 0 5 380 10 15 360 20 25 30 340

320

300 V IV

Excitation (nm) Excitation 280

260

240 I II III 220

300 320 340 360 380 400 420 440 460 480 500 520 540

Emission (nm)

Figure 3.2: Fluorescence excitation-emission spectrum of Barberton source water. [DOC] = 4.47 mg/L, SUVA254 = 4.31 L/mg.m

36

400 0 380 5 10 360 15 20 25 340

320 V 300 IV

Excitation (nm) Excitation 280

260

240

I II III 220

300 320 340 360 380 400 420 440 460 480 500 520 540 Emission (nm)

Figure 3.3: Fluorescence excitation-emission spectrum of Cleveland source water. [DOC] = 2.51 mg/L, SUVA254 = 1.17 L/mg.m

The emission and excitation matrices (EEM) shows the emission and excitation spectra of the three source waters. It is evident from the fluorescence excitation-emission spectra that the waters from Akron and Barberton water treatment plants recorded the highest percentage of volume in the region III (fulvic acid) whiles source water from Cleveland treatment plant had the highest percentage volume in region II (aromatic protein-like). Source water from Cleveland is lower in fulvics and humics comparable to source water from Akron and Barberton. Humic acid, a

37 category of humic substances results from degradation of plant materials by biological and natural chemical processes in terrestrial and aquatic environment (Hudson et al.,

2007). The humic and fulvic acids in the source waters vary due to the vegetation near the watershed, algal concentration in the water and possibly the season of the year (Singer, 1994; Kavanaugh et al., 1980). Aromatic protein in the source waters may be of bacterial origin (Elliot et al., 2006), possibly enzymes a particular microbial community use to break down leaf litter (Allan and Castillo, 2007;

Benfield, 2006; Suberkropp and Klug, 1976) from the forest and other vegetation around the entire watershed. The high aromatic proteins in Cleveland source water may also be attributed to algae bloom since Lake Erie is noted for the toxic algae bloom (personal communication). Leaf litter is a vital source of fulvic acid

(Schlesinger, 1997) which is likely to be a contribution factor to the high fulvic in

Akron and Barberton source waters. The plant materials in the watershed may be from the farms, forests or other vegetative cover. All the source waters have almost equal percentage of volume of soluble microbial by-products which is low.

3.3 Experimental Methods

The experimental procedures were categorised into two – experiments using deionized water (without NOM) and experiments using source waters which contain

NOM.

3.3.1 Experiments with Deionized Water

Controlled laboratory experiment was conducted using deionized water at pH of 6.5, 7.5, 8.5, 9.0 and 9.5. Five 500 mL Erlenmeyer flask (batch reactor) were filled will deionized water. A total of 1 mM aqueous buffer solution was added to each

38 batch reactor as well as 5 μM iopamidol. The aqueous buffer solutions added to the batch reactors were phosphate for pH 6.5 and 7.5, borate for pH 8.5 and carbonate buffer pH 9.0 and 9.5. Using a magnetic stir plate and a PTFE-coated stir bar, under rapid mix, 100 μM of aqueous chlorine was added to the aqueous solution. To ensure uniform mix, the reactants were allowed to mix for 3 min. The samples were afterward transferred into 40 mL amber vials and16 mL amber vials with PTFE septa.

The samples in the 40 mL and 16 mL amber vials were used TOX and iodate analyses respectively. They were stored at 25±1°C in an incubator for reaction times of 0, 6,

12, 24, 48 and 72 hours. Similar experiments, following the same experimental protocol were carried out using monochloramine as the oxidant. The procedures for the preparation of monochloramine are described below. At the end of each reaction time, residual oxidant in each of the samples in the 40 mL and 16 mL amber vials was quenched with 120 μM aqueous sulphite solution and resorcinol for TOX extraction and iodate analysis respectively. The TOX sample was further acidified to pH 2 with nitric acid prior to concentration on the activated carbon columns.

Similarly, three 1000 mL Erlenmeyer flasks were filled with deionized, 4 mM buffer and 5 μM iopamidol. They were rapidly stirred on magnetic stir plate using

PTFE-coated stir bar. About 100 μM of aqueous chlorine was added to each

(representing pH 6.5, 7.5 and 8.5) and the reactants were allowed to uniformly mix for

3 min. Six aliquots from each batch reactor were transferred into 128 mL amber, headspace free with PTFE septa and stored in an incubator at 25±1°C for reaction times of 0, 6, 12, 24, 48 and 72 hr. Also, oxidant residual was quenched in each sample transferred into the 128 mL amber bottle with 120 μM aqueous sulphite solution and analysed for DBPs.

39

Furthermore, experiments were carried out using deionized water to investigate the degradation of TOI and the formation of TOCl, iodate, iodide, THMs,

HAAs and haloacetonitriles (HANs). These were done at higher concentrations of aqueous chlorine, iopamidol and buffer in the absence of NOM. The molar concentration ratio of total chlorine and iopamidol were maintained at a ratio of 20:1 respectively. The experiments were executed at pH 6.5 and 8.5 using 200 mM phosphate buffer, 1.29 mM iopamidol and 25.7 mM aqueous chlorine. Aqueous solutions containing iopamidol and buffer were prepared in a 125 mL Erlenmeyer flask. Using a magnetic stir plate and a PTFE-coated stir bar, under rapid mix, aqueous chlorine was added to the aqueous solution. The reactants were allowed to mix for 3 minutes. Samples were transferred into eight 10 mL amber vials with PTFE septa and stored at 25±1°C in the incubator for reaction times of 0, 1, 2, 6, 12, 24, 48, and 72 hours. Samples were taken at the end of the reaction for analysis. Since the concentrations were very high for effective adsorption in the activated carbon and avoid overloading of columns in the gas chromatography/electron capture detector

(GC/ECD) system, aliquots (3.9 mL) of the samples were transferred into a 1 L

Erlenmeyer flask and diluted to 1 L using deionized water. The diluted sample was put on a magnetic stir plate and using a PTFE coated stir bar, mixed under rapid condition for 3 minutes. After the uniform mix, 10 mL each of the diluted sample were transferred into two 16 mL amber vials. One was quenched with aqueous sulphite solution (120% of the diluted aqueous chlorine concentration) and analyzed for iodide while the other was quenched with 120 μM resorcinol solution for subsequent analysis of iodate formed on the ion chromatography system.

Furthermore, 30 mL of the diluted sample was transferred into a 40 mL amber vial and quenched with 120 μM aqueous sulphite solution for TOX analysis. Nitric acid

40 was added to decrease the pH to 2. Samples were stored at 4°C for 30 minutes before analytical procedures were carried out on the TOX analyzing module and ion chromatography system. The extractions of samples were carried at 4°C since earlier comparison experiments using the 2,4,6-trichlorophenol (TCP), 2,4,6-tribromophenol

(TBP) and 4-iodophenol (IPh) standards to determine recovery of halides resulted in better recovery at 4°C than room temperature (table 3.4). In addition, 100 mL of the diluted samples were transferred into 128 mL amber bottle, quenched with 120 μM aqueous sulphite solution for THMs, HAAs and HANs analyses.

Table 3.4: Comparison of recovery at 4°C and room temperature using 2,4,6- trichlorophenol, 2,4,6-tribromophenol and 4-iodophenol. [TCP] = 25 – 100 μM, [TBP] = 5 – 15 μM, [IPh] = 5 – 15 μM

Recovery

Standard Concentration (μM) Room temperature 4°C

2,4,6-trichlorophenol 100 55.83 66.90

2,4,6-trichlorophenol 50 96.69 118.96

2,4,6-trichlorophenol 25 158.30 162.69

2,4,6-tribromophenol 15 60.31 68.28

2,4,6-tribromophenol 10 60.76 81.99

2,4,6-tribromophenol 5 68.03 83.57

4-iodophenol 15 87.97 95.18

4-iodophenol 10 100.15 110.57

4-iodophenol 5 117.57 110.67

41

3.3.2 Experiments with Source Waters

Also, source waters collected from the Akron, Barberton and Cleveland drinking water treatment plants were filtered through 0.45 μm Whatman nylon membrane filters (Whatman, West Chester, PA, USA) and stored at 4°C prior to use.

Chlorination and chloramination kinetic experiments were conducted under a pseudo first order conditions using [Cl2]T:[iopamidol] = 20:1. Also to determine the effect of

NOM concentration on the iodate formation, the concentrations of NOM in Barberton and Cleveland source waters were decreased by ½ and ¼ by diluting with deionized water.

Aqueous solutions for each of the source waters were prepared in batch reactors. For each of the source waters, aqueous solutions containing NOM, iopamidol and buffer were prepared in a 250 mL Erlenmeyer flask. Buffer was used to maintain the pH of the solution. About 1 mM of phosphate buffer (for pH 6.5 and

7.5) and borate buffer (for pH 8.5) were used to maintain the pH. The lower concentration of the buffer was used to mitigate interferences in the IC chromatograms. Under rapid mix condition, using a magnetic stir plate and a PTFE- coated stir bar, relatively high concentration of aqueous chlorine was added to the aqueous solution at the requisite [Cl2]T:[iopamidol] ratio. The relatively high concentration of aqueous chlorine was used to ensure that excess disinfectant was present in the aqueous mixture throughout the duration of the experiment. Prior to the addition of aqueous chlorine, the chlorine concentration was checked using ferrous ammonium sulphate (FAS)/N, N′-diphenyl-p-phenylenediamine (DPD) titration

(APHA et al., 2005). Stirring was maintained for about 3 min. Aliquots of the aqueous solution were transferred into five 40 mL amber vials with PTFE septa and

42 stored headspace free at 25±1°C in an incubator for a reaction time of 0, 6, 24, 48 and

72 h.

In a similar experimental protocol as the above, pre-formed monochloramine

(that is zero minute of free aqueous chlorine contact time) was used to avoid the artefacts caused by the reactions of excess free chlorine that may briefly exist when forming monochloramine in-situ (Duirk et al, 2005). Pre-formed monochloramine solution was prepared by mixing 5.64 mM ammonium chloride with 3.7 mM hypochlorous acid to achieve a Cl/N molar ratio of 0.7 in a 10 mM carbonate buffer solution. The solution under rapidly mixed condition on a magnetic stir plate using a

PTFE stir bar at a pH 8.5 was allowed to react and reach equilibrium for 30 min. A higher pH (8.5) was used to minimise monochloramine decomposition and to ensure monochloramine remains the active species (Symons et al, 1998) in the aqueous solution. The concentration of the preformed monochloramine was checked with UV visible spectrophotometer and FAS/DPD titration (APHA et al., 2005).

Analytical experimental triplicates were carried to observe the TOI loss and iodate formation in all the source waters for pH of 6.5, 7.5 and 8.5 for 0, 6, 24, 48 and

72 hours. In addition, TOCl formation was observed. At each reaction time, samples were taken from the incubator and the residual chlorine was quenched with aqueous sodium sulphite solution (120% of the initial total chlorine concentration) to measure the TOI and TOCl concentrations formed. The samples were further acidified to pH 2 with nitric acid (70% ACS grade). On the contrary, since iodate is directly oxidised by sulphite at pH above 4 (Rabai and Beck, 1987), resorcinol (120% of the initial total chlorine concentration) was used to quench residual chlorine concentration. Similar chemical procedures were used to quench excess monochloramine in the chloramination experiments.

43

3.4 Analytical Procedures

Two chemical analytical procedures were adopted. One procedure was used to detect halogen-specific TOX, iodide and iodate in the sample. The other analytical procedure was used to detect DBPs in the samples.

3.4.1 Total Organic Halogen

The analytical method developed by Hua and Reckhow (2000) with slight modifications was adopted to analyse the halogen-specific TOX. Using the TOX-100 adsorption module from Cosa Instruments/Mitsubishi (Horseblock Road, NY, USA),

30 mL of each acidified sample was concentrated on a pre-packed granular activated carbon (GAC) column (Cosa Instruments/Mitsubishi, Horseblock Road, NY, USA) through adsorption at an extraction flow rate of 3.3 mL/min. The inorganic halides in the column that will interfere with the results were washed with 15 mL KNO3 solution

- (1000 mg NO3 /L at pH 2) at extraction rate of 3.3 mL/min. The GAC column was placed in a sample quartz boat and automatically introduced into the combustion chamber of the TOX-100 analyzer (Cosa Instruments/Mitsubishi, Horseblock Road,

NY, USA). Furthermore using oxygen as the carrier gas, the GAC was combusted for

15 min at a temperature of 900°C. Using a customized coarse diffuser, the off-gas

(hydrogen halides) was absorbed into a 20 mL phosphate solution (fig. 3.4). Some portions of the 20 mL phosphate solution were used to rinse the diffuser to ensure full recovery of halides.

44

Figure 3.4: Modified schematic diagram of the TOX gas absorption system

3.4.2 Disinfection By-product

THMs, HANs and HAAs analysis were carried out using micro liquid-liquid extraction with MtBE at acidic pH. The THMs which were concurrently analysed included bromodichloromethane (CHBrCl2), dibromochloromethane (CHBr2Cl), chloroform (CHCl3), dichloroiodomethane (CHCl2I), bromochloroiodomethane

(CHBrClI), bromoform (CHBr3), dibromoiodomethane (CHBr2I), chlorodiiodomethane (CHClI2), bromodiiodomethane (CHBrI2), and iodoform (CHI3).

Also chloroactonitrile (CAN), trichloroacetontrile (TCAN), dichloroacetonitrile

(DCAN), bromochloroacetonitrile (BCAN), dibromoacetonitrile (DBAN) bromoacetonitrile (BAN), and iodoacetontrile (IAN) were the HAN compounds analysed.

THMs and HANs were extracted using the US EPA method 551.1 (Munch and Hautman, 1998) with slight modifications. After samples were quenched with aqueous sulphite solution, the sample was acidified with 5 mL of concentrated

45 sulphuric acid. About 3 mL of MtBE and 10 μL of 123.9 mM of 1,2-dibromopropane internal standard were transferred into the acidified sample to achieve approximately

12.4 μM internal standard in the sample. MtBE was used to extract non-dissociated acidic compounds (APHA AWWA and WEF, 1995). In addition, 30 g of anhydrous sodium sulphate salt (dried at 100°C) was added to decrease the activity of inorganic compounds and increase the activity of the organic compounds – to increase partitioning of the DBPs from the aqueous phase to the MtBE (US EPA, 2013), which increases extraction efficiency. The samples in the 128 mL amber bottles were capped with polyseal cone-lined cap, hand-shaken for a minute and then shaken on the wrist action shaker (Burrell Scientific, Pittsburgh, PA, USA) for 30 minutes.

After the mechanical shake, the sample was left to settle for 3 minutes in a 100-ml volumetric flask. A disposable Pasteur pipette was used to transfer at least 1.5 mL

MtBE extract into a 2 mL GC autosampler vial through another Pasteur pipette filled with glass wool and dried anhydrous sodium sulphate salt to dry out water from the organic extract. The extracted sample was then split – 0.5 mL used for derivatization with diazomethane for HAAs analysis and the remaining used for THMs and HANs analyses. The extracted samples were stored in the freezer. Vials were finally placed in the GC autosampler for injection into the GC.

HAAs were measured using a modified US EPA method 552.1 (Hodgeson and

Becker, 1992) which uses liquid-liquid extraction with MTBE, derivatization with diazomethane and analysis with GC/MS. The HAA compounds analysed were comprised of chloroacetic acid (CAA), bromoacetic acid (BAA), dichloroacetic acid

(DCAA), trichloroacetic acid (TCAA), iodoacetic acid (IAA), bromochloroacetic acid

(BCAA), bromodichloroacetic acid (BDCAA) and dibromoacetic acid (DBAA).

Aliquot of the extracted sample was methylated with diazomethane for the production

46 of methyl ester or other derivatives for gas chromatographic separation (APHA,

AWWA and WEF, 1995). Diazomethane was generated by adding 0.367 g diazald and 1 mL carbitol (2-[2-ethoxyethoxy] ethanol) to the inner tube of the diazomethane generator. Also 3 mL of MtBE was added to the outer tube of the diazomethane generator. The two parts of the generator were assembled and the lower part of the outer tube was immersed in ice bath to ensure an isothermal condition of 0°C was maintained. After equilibrating to 0°C, 1.5 mL of KOH (37%) was slowly injected

(dropwise) into the generator through the septum to initiate the reaction. The apparatus was shaken gently by hand to ensure uniform mixture of reactants in the inner tube while avoiding spill into the outer tube. When the solution in the outer tube becomes yellow it is an indication of excess diazomethane. The apparatus with the solution was left to stand for 50 minutes, after which the tube was opened to destroy unreacted diazomethane with activated silica. After preparing the diazomethane, about 0.5 mL of the extracted sample was transferred into another GC autosampler vial and 250 μL of the diazomethane added to it. The sample stood for

15 minutes to allow adequate methylation of the HAAs, and then 1 – 3 grains of activated silica were added to the sample to destroy any excess diazomethane.

3.5 Analyses of TOX, Iodate and Iodide

Detection of halogen specific TOX, iodate and iodide were achieved with

Dionex ICS-3000 ion chromatograph system (Dionex Corporation, Sunnyvale, CA,

USA) with conductivity detector and an ASRS®300 4 mm anion self-regenerating suppressor. For the detection of the halides (TOX measured as halides and inorganic iodide), AS20 analytical column (4 x 250 mm) and guard column (Dionex

Corporation, Sunnyvale, CA, USA) with KOH as the mobile phase were employed.

47

The flow rate of the mobile phase was 1 mL/min. Figure 3.5 shows the gradient profile of the method used for the determination of the halides.

Organic compounds (2,4,6-trichlorophenol, 2,4,6-tribromophenol and 4- iodophenol) were used for standardization test to determine the recovery of Cl-, Br-, and I-. The phenol standards were run through an activated carbon cartridge, combusted in a TOX analyzer and analyzed on the ICS-3000 using the method developed. Inorganic halogenated compounds of the same concentrations as the phenols were analyzed on the ICS-3000 to determine the recovery of the phenols after combustion in the TOX analyzer. The phenols were then used to generate calibration curves (figures 3.6 – 3.8) to further determine the concentrations of specific halogens in the samples. Also standard solutions of potassium iodide at concentrations of 0 to

50 μM were run on the ICS 3000 and a calibration curve was developed (figure 3.9) to determine the concentrations of iodide in the sample.

Absorbed combusted source water samples were delivered by AS50 autosampler (Dionex Corporation, Sunnyvale, CA, USA) and a volume of 500 μL was injected. For the automatic control of ICS module and data analysis

(chromatograms), the Chromeleon software by Dionex Corporation (Sunnyvale, CA,

USA) was used. The area of the integrated chromatograms, measured for each halide, was fitted into the equation of the calibration curve and the molar concentrations of

TOI and TOCl were calculated as μM I- and μM Cl- respectively.

48

60

50

40

30

20

Eluent Concentration (mM) Eluent Concentration 10

0 0 5 10 15 20

Time (min)

Figure 3.5: Gradient profile for the analysis of Total organic halogen

49

300

250

200

M)

150 y= 5.6135*x-5.250 R2=0.9964

Concentration ( Concentration 100 Chloride

50

0 0 10 20 30 40 50

Area ( S*min)

Figure 3.6: Calibration curve for Chloride using 2,4,6-trichlorophenol. [Cl-] = 0 – 250 μM

50

60

50

40

M)

30 y= 5.1236*x+1.2031 R2=0.9988

Concentration ( Concentration 20 Iodide

10

0 0 2 4 6 8 10

Area ( S*min)

Figure 3.7: Calibration curve for Iodide using 4-iodophenol. [I-] = 0 – 50 μM

51

60

y = 6.4756x + 2.0445 50 R2 = 0.9917

40

M)



30

Concentration ( Concentration 20

10

0 0 2 4 6 8

Area (S.min)

Figure 3.8: Calibration curve for Bromide using 4-iodophenol. [Br-] = 0 – 50 μM

52

100 y = 3.995x + 1.1775 R² = 0.9995

80

M)

 60

40

Concentration ( Concentration

20

0 0 5 10 15 20 25 30

Area (S.min)

Figure 3.9: Calibration curve for Iodide using KI. [I-] = 0 – 100 μM

For iodate detection, the AS18 analytical column (4 x 250 mm) and guard column (Dionex Corporation, Sunnyvale, CA, USA) with KOH as the mobile phase were also used. The flow rate of the mobile phase was 1 mL/min. The gradient profile of the method used for the detection of iodate is shown in figure 3.10. To determine the iodate formed in the sample, sodium iodate (at concentrations of 0 to 50

μM) was used for standardization test. The concentrations of the standard were run directly in the ICS system and a graph of concentration of standards versus area of respective chromatograms was used as calibration curve (figure 3.11). The concentration of iodate formed in the reaction was calculated from the equation of the calibration curve.

53

35

30

25

20

15

10

Eluent Concentration (mM) Eluent Concentration

5

0 0 5 10 15 20

Time (min)

Figure 3.10: Gradient profile for the analysis of iodate.

25

y = 5.4459x 20 R² = 0.9997

M)  15

10

Concentration ( Concentration

5

0 0 1 2 3 4

Area (S.min)

- Figure 3.11: Calibration curve for Iodate using NaIO3. [IO3 ] = 0 – 20 μM 54

3.6 Analyses of DBPs

The extracted and derivatized samples were analyzed with 7890A GC system equipped with 63Ni microelectron capture detector (μECD) (Agilent Technologies,

Santa Clara, CA, USA). A Restek 13638-127 GC column (Restek Corporation,

Bellefonte, PA, USA) was connected from the injector to the μECD to achieve separation of analytes. The column conditions were as follows: length 30 m, internal diameter 0.25 mm, film thickness 0.5 μm and flow rate 1 mL/min. Samples were delivered by 7693 autosampler (Agilent Technologies, Santa Clara, CA, USA).

Splitless injections were achieved by injecting 1 μL of the sample into the column.

The temperature of the μECD was 250°C and the make-up gas was ultrahigh purity nitrogen gas with flow rate of 19 mL/min. The carrier gas employed was helium gas

(ultrahigh purity). There were two oven temperature programming used – one for analysis of THMs and HANs (Table 3.5) and the other for HAAs analysis (Table 3.6).

Table 3.5: Oven temperature programming for THMs and HANs analysis on GC/μECD

Rate (°C/min) Temperature (°C) Hold time (min) Run time (min)

Initial 50 10 10

Ramp 1 2.5 65 0 16

Ramp 2 5 85 0 20

Ramp 3 7.5 205 0 36

Ramp 4 10 280 0 43.6

55

Table 3.6: Oven temperature programming for HAAs analysis on GC/μECD

Rate (°C/min) Temperature (°C) Hold time (min) Run time (min)

Initial 50 10 10

Ramp 1 0.25 50.5 5 17

Ramp 2 0.25 52 5 28

Ramp 3 0.25 62.5 0 70

Ramp 4 35 280 0 76.214

THMs, HANs and HAAs standard solutions were prepared using deionized water. The known concentrations of the THMs and HANs standards were extracted using the extraction procedure described above using 10 μL of 123.9 mM 1,2- dibromopropane internal standard to achieve about 12.4 μM internal standard in the sample. The HAAs of known concentrations were also derivatized with diazomethane after extraction using the same volume and concentration of 1,2- dibromopropane as internal standard. All standards were analyzed with 7890A GC system equipped with μECD using their respective methods. A calibration curve of concentration of the standard versus the relative response of the standard solution to the internal standard was developed to calculate the concentrations of the DBPs formed in the samples. The relative response of standard to the internal standard is referred to in the calibration curve as response ratio (shown on the abscissa). The calibration curves for all the standard solutions are shown in figures 3.12 to 3.36. The concentration of the specific DBP was calculated from the equation of the line of best fit of the corresponding standard curve. The limits of quantification (LOQ) for the

DBPs are shown in table 3.7.

56

1200

1000

800

] (nM)

3 600 y = 1471.4x R² = 0.9785

[CHCl 400

200

0 0.0 0.2 0.4 0.6 0.8 1.0

Response ratio

Figure 3.12: Calibration curve for CHCl3using chloroform. [CHCl3] = 0 – 1000 nM

350

300

250

200

Cl] (nM)

2 150 y = 109.38x

[CHBr R² = 0.9974 100

50

0 0.0 0.5 1.0 1.5 2.0 2.5 3.0

Response ratio

Figure 3.13: Calibration curve for CHBr2Cl using dibromochloromethane. [CHBr2Cl] = 0 – 300 nM

57

160

140

120

100

] (nM)

2 80

[CHBrI 60 y = 135.95x R² = 0.9925 40

20

0 0.0 0.2 0.4 0.6 0.8 1.0

Response ratio

Figure 3.14: Calibration curve for CHBrI2 using bromodiiodomethane. [CHBrI2] = 0 – 125 nM

300

250

200

] (nM)

2 150 y = 266.06x

[CHClI R² = 0.9993 100

50

0 0.0 0.2 0.4 0.6 0.8 1.0

Response ratio

Figure 3.15: Calibration curve for CHClI2 using chlorodiiodomethane. [CHClI2] = 0 – 250 nM

58

300

250

200

I] (nM) I]

2 150 y = 1152.7x

[CHBr R² = 0.9941 100

50

0 0.00 0.05 0.10 0.15 0.20 0.25

Response ratio

Figure 3.16: Calibration curve for CHBr2I using dibromoiodomethane. [CHBr2I] = 0 – 250 nM

300

250

200

150 y = 1595.3x R² = 0.998

[CHBrClI] (nM) 100

50

0 0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 0.16 0.18

Response ratio

Figure 3.17: Calibration curve for CHBrClI using bromochloroiodomethane. [CHBrClI] = 0 – 250 nM

59

600

500

400

] (nM)

3 300 y = 204.28x R² = 0.9968

[CHBr 200

100

0 0.0 0.5 1.0 1.5 2.0 2.5 3.0

Response ratio

Figure 3.18: Calibration curve for CHBr3 using bromoform. [CHBr3] = 0 – 500 nM

600

500

400

y = 2013.1x

I] (nM) I]

2 300 R² = 0.9985

[CHCl 200

100

0 0.00 0.05 0.10 0.15 0.20 0.25 0.30

Response ratio

Figure 3.19: Calibration curve for CHCl2I using dichloroiodomethane. [CHCl2I] = 0 – 500 nM

60

400

300

Br] (nM) 200

2 y = 158.05x R² = 0.9965

[CHCl

100

0 0.0 0.5 1.0 1.5 2.0 2.5 3.0

Response Ratio

Figure 3.20: Calibration curve for CHCl2Br using bromodichloromethane. [CHCl2Br] = 0 – 400 nM

60

50

40

y = 68.338x

] (nM)

3 30 R² = 0.9933

[CHI 20

10

0 0.0 0.2 0.4 0.6 0.8 1.0

Response ratio

Figure 3.21: Calibration curve for CHI3 using iodoform. [CHI3] = 0 – 50 nM

61

600

500

400

300

[CAN] (nM) y = 478.7x 200 R² = 0.9988

100

0 0.0 0.2 0.4 0.6 0.8 1.0

Response ratio

Figure 3.22: Calibration curve for CAN using chloroacetonitrile. [CAN] = 0 – 500 nM

600

500

400

300 y = 2852.9x

[DCAN] (nM) 200 R² = 0.9991

100

0 0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 0.16 0.18 0.20

Response ratio

Figure 3.23: Calibration curve for DCAN using dichloroacetonitrile. [DCAN] = 0 – 500 nM

62

140

120

100

80

60

[TCAN] (nM) y = 160.4x R² = 0.9631 40

20

0 0.0 0.2 0.4 0.6 0.8 1.0

Response ratio

Figure 3.24: Calibration curve for TCAN using trichloroacetonitrile. [TCAN] = 0 – 125 nM

140

120

100

80

y = 129.6x 60 R² = 0.9987

[BAN] (nM)

40

20

0 0.0 0.2 0.4 0.6 0.8 1.0 1.2

Response ratio

Figure 3.25: Calibration curve for BAN using bromoacetonitrile. [BAN] = 0 – 125 nM

63

300

250

200

150 y = 346.63x R² = 0.9977

[DBAN] (nM) 100

50

0 0.0 0.2 0.4 0.6 0.8

Response ratio

Figure 3.26: Calibration curve for DBAN using dibromoacetonitrile. [DBAN] = 0 – 250 nM

300

250

200

150 y = 372.69x R² = 0.9931

[BCAN] (nM) 100

50

0 0.0 0.2 0.4 0.6 0.8

Response ratio

Figure 3.27: Calibration curve for BCAN using bromochloroacetonitrile. [BCAN]=0– 250 nM

64

35

30

25

20

y = 33.062x 15 R² = 0.9946

[IAN] (nM) [IAN]

10

5

0 0.0 0.2 0.4 0.6 0.8 1.0 1.2

Response ratio

Figure 3.28: Calibration curve for IAN using iodoacetonitrile. [IAN] = 0 – 31 nM

300

250

200

150 y = 773.91x

[CAA] (nM) R² = 0.9997 100

50

0 0.00 0.05 0.10 0.15 0.20 0.25 0.30 0.35

Response ratio

Figure 3.29: Calibration curve for CAA using chloroacetic acid. [CAA] = 0 – 250 nM

65

600

500

400

300

y = 376.03x

[DCAA] (nM) 200 R² = 0.9693

100

0 0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6

Response ratio

Figure 3.30: Calibration curve for DCAA using dichloroacetic acid. [DCAA] = 0 – 500 nM

300

250

200

150

[TCAA] (nM) y = 100.41x 100 R² = 0.9986

50

0 0.0 0.5 1.0 1.5 2.0 2.5 3.0

Response ratio

Figure 3.31: Calibration curve for TCAA using trichloroacetic acid. [TCAA] = 0 – 250 nM

66

300

250

200

150 y = 105.74x R² = 0.9969

[BCAA] (nM) 100

50

0 0.0 0.5 1.0 1.5 2.0 2.5 3.0

Response ratio

Figure 3.32: Calibration curve for BCAA using bromochloroacetic acid. [BCAA]=0– 250 nM

300

250

200

150 y = 85.951x R² = 0.9982 [BDCAA] (nM) 100

50

0 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5

Response ratio

Figure 3.33: Calibration curve for BDCAA using bromodichloroacetic acid. [BDCAA] = 0 –250 nM

67

1200

1000

800

600 y = 1378.1x [BAA] (nM) R² = 0.991 400

200

0 0.0 0.2 0.4 0.6 0.8

Response ratio

Figure 3.34: Calibration curve for BAA using bromoacetic acid. [BAA] = 0 – 1000 nM

600

500

400

300

y = 139.48x

[DBAA] (nM) 200 R² = 0.9949

100

0 0 1 2 3 4

Response ratio Figure 3.35: Calibration curve for DBAA using dibromoacetic acid. [DBAA] = 0 – 500 nM

68

140

120

100

80

y = 83.588x 60 R² = 0.9896

[IAA] (nM) [IAA]

40

20

0 0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8

Response ratio

Figure 3.36: Calibration curve for IAA using iodoacetic acid. [IAA] = 0 – 125 nM

69

Table 3.7: Limit of quantification for the detection of DBPs

DBPs Limit of quantification (nM)

CHCl3 1.0

CH BrCl2 1.0

CH Br2Cl 1.0 CHClBrI 1.0

CHCl2I 0.2

CHClI2 0.2

CHBr3 1.0

CHBr2I 0.2

CHBrI2 0.2

CHI3 0.2 CAN 1.0 TCAN 1.0 DCAN 1.0 BAN 1.0 BCAN 1.0 DBAN 1.0 IAN 0.2 CAA 1.0 BAA 1.0 DCAA 1.0 TCAA 1.0 IAA 0.2 BCAA 1.0 BDCAA 1.0 DBAA 1.0

70

CHAPTER IV

RESULTS AND DISCUSSION

4.1 Introduction

This chapter focuses on the transformation of iopamidol in the presence of chlorinated oxidants (aqueous chlorine and monochloramine) and the absence of

NOM using deionized water where TOI was used as surrogate for iopamidol. This complements the study by Pushpita Kumkum in 2013. The measured rates of degradation of TOI and the formation of iodate in the absence of NOM as a function of pH were assessed. Also, the transformation of iopamidol in the presence of chlorinated oxidants and NOM were investigated using three source waters.

4.2 Transformation of Iopamidol in the Absence of NOM

Transformation of iopamidol was monitored at both low and high concentrations of reactants and buffer. Iopamidol degradation was monitored as TOI loss. Chlorine incorporation was also monitored as TOCl. Other parameters investigated were iodate, iodide and DBPs formed.

4.2.1 Transformation at Low Concentration

The degradation of iopamidol in the absence of NOM in excess aqueous chlorine was conducted at pH of 6.5 to 9.5 as a function of time (figure 4.1). The loss of TOI with aqueous chlorine in deionized water resulted in a great degradation of

71

TOI. TOI decreased with respect to time. The greatest degradation of TOI was observed at pH 7.5 whiles the least was observed at 9.5. This is also evident in the

-6 observed rate constants (kobs) (figure 4.2). The kobs for pH 7.5 and 9.5 are 3.97 x 10 s-1 (0.0143 hr-1) and 2.89 x 10-6 s-1 (0.0104 hr-1) respectively. At the end of the 72-

- hour reaction, TOI exhibited the same degradation at pH 8.5 and 9.0 (kobs = 3.83 x 10

6 s-1 {0.0138 hr-1}). The loss of TOI in the presence of aqueous chlorine follows a pseudo first order reaction.

25

pH 6.5 pH 7.5 20 pH 8.5 pH 9.0 pH 9.5 15

M)



TOI ( TOI 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.1: TOI degradation as a function of pH in reaction mixtures containing iopamidol and aqueous chlorine [Cl]T=100 μM, [iopamidol]=5 μM, [Buffer]=1 mM, and temperature= 25°C. Error bars represent 95% confidence intervals.

It has already been proposed that OCl- may primarily initiate the transformation of iopamidol (Duirk et al., 2011). This however does not follow the known conventional iodide oxidation pathway (Bichsel and von Gunten, 2000;

72

Bichsel and von Gunten, 1999b). OCl-, a strong nucleophile, is thought to have initially attacked one of the amide side chains of the iopamidol. This results in the formation of a primary amine transformation product. On the contrary, the reaction mechanism culminating in the cleavage of iodide from the benzene ring is yet to be entirely understood. The loss of TOI decreased with increasing pH. From figure 4.2, all the observed pseudo first order rate line almost approximate to zero at the ln

([TOI]t/[TOI]0) intercept except at pH 6.5. There is an observed biphasic behaviour exhibited at pH 6.5 due to its positive intercept at the ln ([TOI]t/[TOI]0) axis. This may imply that there is not enough OCl- to initiate the degradation of iopamidol in the rate limiting reaction to transform iopamidol to its initial amine transformation product. At pH 6.5 there is about 90% HOCl species present. Therefore, both HOCl and OCl- species may have participated in the degradation of iopamidol at pH 6.5.

0.0

-0.2

-0.4

)

0 -0.6

/[TOI]

t -0.8

ln ([TOI] -1.0 pH 6.5 y = -0.0117x + 0.0436, R² = 0.9848 pH 7.5 y = -0.0143x - 0.0296, R² = 0.9766 -1.2 pH 8.5 y = -0.0138x - 0.0658, R² = 0.9807 pH 9.0 y = -0.0138x - 0.0907, R² = 0.9802 pH 9.5 y = -0.0104x - 0.0502, R² = 0.9679 -1.4

0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75 Time (hr)

Figure 4.2: Observed pseudo-first order loss of TOI as a function of pH. [Cl2]T = 100 μM, [Iopamidol] = 5 μM, [Buffer]T = 1 mM, Temperature = 25°C

73

Iodide is known to be rapidly oxidized to HOI in the presence of aqueous chlorine (Bichsel and von Gunten, 1999a; Nagy et al., 1988). HOI can

- - disproportionate to IO3 and I . However, Bichsel and von Gunten (1999a) argued that in the presence of excess aqueous chlorine (1 – 10 μg/L HOI, pH 6 – 8, [CO3]T = 0 –

5 mM), HOI/OI- disproportionation is too slow to be of importance to the fate of HOI.

Therefore the fate of HOI will be its reaction with NOM to form TOI or further

- oxidation to form IO3 (Hua et al., 2006; Bichsel and von Gunten, 1999a). In the reaction of iopamidol with aqueous chlorine, the iodine on the aromatic ring will be

- oxidised to HOI and HOI will either be oxidised to form IO3 or may be incorporated back into iopamidol transformation products to form TOI. Duirk et al. (2011) proposed the possibility of iopamidol being the source of iodine in iodo-DBPs since iodo-DBPs were not detected in control experiment (raw source water in the presence of aqueous chlorine) which was confirmed in this study.

The formation of iodate was monitored in the experiment and it was found that iodate was formed at all pH (figure 4.3). The formation of iodate was found to increase with respect to increase in time for all pH. Iodate formation was observed to be greatest at pH 7.5 and lowest at pH 9.5. Iodate formation was approximately the same at pH 8.5 and 9.0. Using the same experimental condition (except reaction time up to 48 hr), Duirk et al. (2011) found that iodate formation was highest at pH 7.5.

Since iodate formation is due to chlorine oxidation of HOI, it is expected that either

HOCl or OCl- oxidises HOI. If OCl- oxidises HOI, iodate formation is expected to increase with increasing pH (6.5 – 9.5). However, if the active chlorine species is

HOCl, the converse will exist.

74

25

pH 6.5 20 pH 7.5 pH 8.5 pH 9.0 pH 9.5 15

M)



Iodate ( Iodate 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75 Time (hr)

Figure 4.3: Iodate formation as a function of pH in reaction mixtures containing iopamidol and aqueous chlorine. [Cl]T=100 μM, [iopamidol]=5 μM, [Buffer]=1 mM, and temperature= 25°C. Error bars represent 95% confidence intervals.

From the result, iodate formation decreased from pH 7.5 to 9.5 – an indication

- that HOCl may be the species oxidising HOI to IO3 . On the contrary, at pH 6.5

- (about 90% HOCl is available as free chlorine), IO3 formation decreased. This was also observed in the research by Duirk et al. (2011). This may be due to the low proportion of OCl- (approximately 10%) to initiate the reaction (cleavage of amide group on the aromatic ring) at pH 6.5. Also, it may be possible that both species of chlorine (HOCl and OCl-) were involved in the oxidation of HOI to iodate. In their study, Bichsel and von Gunten (2000) concluded that HOCl and OCl- were the kinetically (pseudo first order reaction) dominating species in the oxidation of HOI to

- IO3 at pH 5.3 to 6.4 and pH 8.2 to 8.9 respectively. Also, the rate constant for the first order reaction for OCl- (52±5 M-1.s-1) was significantly higher than the rate

75

-1 -1 - constant for HOCl (8.2±0.8 M .s ). The formation kinetics of IO3 in this study did not follow either first or second order reaction (not shown).

To determine the DBPs formed at low concentration, experiments were carried out at pH 6.5, 7.5 and 8.5 using 5 μM iopamidol and 100 μM aqueous chlorine in the absence of NOM for reaction time up to 72 hr. The predominant DBPs formed at these pH were CHCl3 and TCAA (figures 4.4 to 4.6). The formations of these DBPs were observed from 6 to 72 hr. Also, relatively small concentrations of CHClI2 were observed after 12 hr. CHCl3 and TCAA increased significantly from 6 to 72 hr. At pH 6.5, chloroform was predominant DBP up to 12 hr. Afterwards, trichloroacetic acid dominated to 72 hr. On the contrary, almost equal concentrations of CHCl3 and

TCAA were observed at 0 to 12 hr at pH 7.5. TCAA then became the main species.

The trend at pH 8.5 was quite different from the above – approximately equal TCAA and CHCl3 were observed at each discrete sampling time. The formation of CHCl3 increased with increasing pH. However, the increasing order of TCAA formation with pH was pH 6.5 < 8.5 < 7.5. The formation of CHClI2 was almost equal at all pH.

Iodinated DBPs were not formed until 12 hr sampling time. At 0 and 6 hr, all the TOI formed at all pH may be unknown iopamidol transformation products. The proportions of TOI formed as CHClI2 at pH 6.5, 7.5 and 8.5 were less than 0.2% at all sampling time of 12, 24, 48 and 72 hr. Therefore more than 99% of the remaining

TOI formed were unknown transformation products with known toxicity.

76

400

CHCl3 300 CHClI2 TCAA

200

Concentration (nM) Concentration 100

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.4: THM and HAA formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 6.5. [Cl2]T = 100 μM, [Iopamidol] = 5 μM, [Buffer]T = 4 mM, and temperature = 25C. Error bars represent 95% confidence intervals.

77

400

CHCl3 CHClI2 300 TCAA

200

Concentration (nM) Concentration 100

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.5: THM and HAA formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 7.5. [Cl2]T = 100 μM, [Iopamidol] = 5 μM, [Buffer]T = 4 mM, and temperature = 25C. Error bars represent 95% confidence intervals.

78

400

CHCl3 CHClI2 300 TCAA

200

Concentration (nM) Concentration 100

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.6: THM and HAA formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 8.5. [Cl2]T = 100 μM, [Iopamidol] = 5 μM, [Buffer]T = 4 mM, and temperature = 25C. Error bars represent 95% confidence intervals.

The chlorinated DBPs formed were CHCl3, CHClI2 and TCAA. The proportions of TOCl formed as chlorodiiodomethane at all pH were less than 0.05% at all sample times. Nonetheless, the proportions of TCAA and CHCl3 formed were greater than 2%. The predominant chlorinated DBP formed at pH 6.5 at sampling time 6 and 12 hr was CHCl3, which were 4% and 5% respectively. However, TCAA was major chlorinated DBP at 24, 48 and 72 hr. In all, the total proportions of chlorinated DBPs formed relative to TOCl (figure 4.7) at sampling time 6, 12, 24, 48 and 72 hr were approximately 4, 7, 17, 5 and 11% respectively. The proportions of

TCAA formed at pH 7.5 were greater than CHCl3 at all sample times. Proportions of

TCAA formed at 6, 12, 24, 48 and 72 hr were approximately 3, 8, 16, 15 and 6%

79 correspondingly. The total proportions of all chlorinated DBPs formed at pH 7.5 were 6, 18, 23, 25 and 14% at sampling time of 6, 12, 24, 48 and 72 hr respectively.

These proportions were higher than the proportions formed at pH 6.5 and 8.5.

Approximately equal proportions of TCAA and CHCl3 were formed at pH 8.5. At pH

8.5, about 6, 3, 9, 7 and 13% represented the total chlorinated DBPs formed at 6, 12,

24, 48 and 72 hr respectively. In conclusion more chlorinated DBPs were formed than iodinated DBPs.

Also, the degradation of iopamidol was investigated in the presence of monochloramine at pH 6.5 to 9.0 for up to 168 hr reaction time. There was no observed significant degradation of iopamidol (TOI) over the 168 hr (figure 4.8).

Monochloramine is known to react with iopamidol to form iodo-DBPs in aqueous solutions containing iopamidol and NOM (Duirk et al., 2011). Therefore, the iodide on the benzene ring may be oxidised to HOI (Bichsel and von Gunten, 1999b). HOI has been shown to be stable in the presence of NH2Cl and in the absence of other reactants (Bichsel and von Gunten, 1999a). Thus, the formation of iodate in the presence of NH2Cl is implausible. This may explain why iodate was not formed in the presence of NH2Cl. When sulphite was used to quench the reaction, it was

- 2- 2- expected that HOI will be reduced to I whiles SO3 will be oxidised to SO4 .

Substantial concentrations of iodide were quantified (figure 4.9) which was relatively constant from 6 hr to 168 hr. HOI/I- may appear to be in pseudo-steady state with the iopamidol transformation products and iopamidol.

80

25 pH 6.5 pH 7.5 20 pH 8.5 pH 9.0 pH 9.5 15

M)



TOCl ( TOCl 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.7: TOCl formation as a function of pH in reaction mixtures containing iopamidol and aqueous chlorine. [Cl2]T = 100 μM, [Iopamidol] = 5 μM, [Buffer]T = 4 mM, and temperature = 25C. Error bars represent 95% confidence intervals.

81

20

15

M)  10

TOI ( TOI pH 6.5 pH 7.5 5 pH 8.5 pH 9

0 0 25 50 75 100 125 150 175

Time (hr)

Figure 4.8: TOI loss as a function of pH in reaction mixtures containing iopamidol and monochloramine. [NH2Cl] = 100 μM, [iopamidol] = 5 μM, [Buffer] = 1 mM, and temperature = 25°C. Error bars represent 95% confidence intervals.

82

5

pH 6.5 pH 7.5 4 pH 8.5 pH 9

3

M)



] (

- [I 2

1

0 0 25 50 75 100 125 150 175

Time (hr)

Figure 4.9: Iodide formation as a function of pH in reaction mixtures containing iopamidol and monochloramine. [NH2Cl] = 100μM, [iopamidol] = 5μM, [Buffer] = 1mM, and temperature = 25°C. Error bars represent 95% confidence intervals.

4.2.2 Transformation at High Concentration

The transformation of iopamidol in the absence of NOM was carried out in chlorinated deionized water at high concentrations of reactants and buffer at pH 6.5 and 8.5. The concentrations of iopamidol, aqueous chlorine and buffer were 1.29 mM, 25.7 mM and 200 mM respectively. The degradation of iopamidol (TOI) was fast for the first 24 hr at both pH (figures 4.10 – 4.11). After 24 hr, TOI loss ceased at pH 6.5 and only slightly continued to degrade at pH 8.5. The same was observed in the formation of iodate at both pH. Iodide formation remained constant but concentrations were very low. After approximately 24 hours of reaction, the predominant iodine species at pH 6.5 and pH 8.5 was iodate. The formation of TOCl

83 was observed after 6 hr at pH 6.5 and 2 hr at pH 8.5 and continued until the discrete sample at 24 hr (figure 4.12 – 4.13). After 24 hr, TOCl formation remained fairly constant at pH 8.5 but there was marginal increase of approximately 17% at pH 6.5.

7000 TOI 6000 Iodide Iodate [I] 5000 T

M)

 4000

3000

Concentration ( Concentration 2000

1000

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

- - Figure 4.10: TOI, I , and IO3 mass balance in reaction mixtures containing iopamidol and aqueous chlorine at pH 6.5. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and tempeerature = 25C. Error bars represent 95% confidence intervals.

84

6000 TOI Iodide 5000 Iodate [I]T 4000

M)



3000

Concentration ( Concentration 2000

1000

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

- - Figure 4.11: TOI, I , and IO3 mass balance in reaction mixtures containing iopamidol and aqueous chlorine at pH 8.5. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Error bars represent 95% confidence intervals.

85

2500

2000

1500

M)



TOCl ( 1000

500

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.12: TOCl formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 6.5. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Error bars represent 95% confidence intervals.

86

3000

2500

2000

M)

 1500

TOCl ( 1000

500

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.13: TOCl formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 8.5. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Error bars represent 95% confidence intervals.

Although all reactions (TOI loss, and iodate and iodide formation) stopped at

24 hr, other reactions continued resulting in TOCl formation – that may be the cause of the marginal increment in TOCl formation (chlorine incorporation). Thus the chlorinated DBPs formed (figures 4.14 – 4.15) may have accounted for the observed pattern in TOCl formation. Formation of iodinated DBP was relatively low. At both pH, CHCl2I and CHClI2 were observed after 12 hr and 24 hr respectively.

87

2e+5 CHCl3 CHCl2I CHClI2 2e+5 DCAA TCAA

1e+5

Concentration (nM) Concentration 5e+4

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.14: THM and HAA formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 6.5. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Error bars represent 95% confidence intervals.

88

2e+5 CHCl3 CHCl2I CHClI2 2e+5 DCAA TCAA

1e+5

Concentration (nM) Concentration 5e+4

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.15: THM and HAA formation in reaction mixtures containing iopamidol and aqueous chlorine at pH 8.5. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Error bars represent 95% confidence intervals.

Total organic halogen is comprised of the halogenated DBPs and the unknown

TOX. The unknown TOX may be the unknown iopamidol transformation products.

At each discrete sample time, the DBPs were normalized to the amount and type of halogen contained within the chemical structure and the percent of TOCl and/or TOI it accounted for. The only iodinated DBPs formed at pH 6.5 and 8.5 were CHCl2I and

CHClI2. At the initial reaction time 100% TOI was observed at both pH 6.5 and 8.5.

After 1 hr, no degradation of TOI was observed at pH 6.5. Nevertheless, there was

TOI loss after 1 hr at pH 6.5, that is, the amount of TOI formed as reaction time increased gradually decreased. The TOI remaining at 2, 6 and 12 hr were 97, 87 and

67%. There was significant degradation of TOI at these sample times. However, the loss of TOI from 24 to 72 hr was insignificant at pH 6.5. Approximately 43, 42 and

89

42% TOI remained at sample times 24, 48 and 72 hr respectively. Therefore at the end of the 72-hr reaction time, about 58% of the initial TOI had degraded. In contrast, TOI degraded from 1 hr to 72 hr at pH 8.5. TOI remaining at 1, 2, 6 and 12 hr at pH 8.5 was respectively 88, 80, 66 and 49%. Degradation of TOI after 12 hr was slow. At 24, 48 and 72, the remaining TOI was 41, 36 and 34%. In all, approximately 66% of the initial TOI degraded at pH 8.5 at the end of the 72-hr reaction time.

The formation of CHCl2I at both pH was observed from 12 hr to 72 hr while

CHClI2 was observed from 24 hr to 72 hr. The relative proportions of the iodinated

DBPs (I-DBPs) were small relative to the TOI (figures 4.16 and 4.17). Although the percentage increase of I-DBPs as a function of time was significant, the proportion of

I-DBPs formed was infinitesimally small relative to unkown TOI. This implies about

99% of the TOI formed was not identified and thus the relative toxicity of these unknown transformation products (unknown T.P.) cannot be confirmed. The formation of the I-DBPs increased with increasing time and increasing pH.

On the contrary, the formation of chlorinated DBPs (Cl-DBPs) was significant

(figures 4.18 and 4.19) relative to the TOCl formed. At pH 6.5 CHCl3 recorded the highest relative proportion of Cl-DBPs followed by TCAA for the reaction times.

More chloroform was formed than unknown TOCl at 12 hr. However, there was remarkable decrease and increase in CHCl3 and unknown TOCl respectively at 24 hr.

The proportion of both CHCl2I and CHClI2 increased with increasing time at pH 6.5.

At 48 hr approximately equal proportions of chloroform and TCAA were formed.

Also at pH 6.5, a decreasing pattern of Cl-DBPs formation was observed at all discrete times – trichlorinated DBPs > dichlorinated DBPs > monochlorinated DBPs.

At pH 8.5, more unknown TOCl was formed although Cl-DBPs were also formed.

90

This may be due to the low concentration of HOCl at pH 8.5. Only the trichlorinated

DBPs were formed at 2 and 6 hr. Thus their formations were rapid. At 12 to 48 hr, chloroform was the predominant Cl-DBPs but more DCAA was formed than TCAA.

Also there was high increase in CHCl2I. Approximately equal proportion of CHCl2I and TCAA were formed at 48 hr. At 72 hr the trichlorinated DBPs were the predominant Cl-DBPs formed.

It has been proposed that OCl- is the initial reactive species on one of the amide groups on the aromatic ring (Duirk et al., 2011). It is therefore expected that the oxidants cleaves the C-N bond which will result in NH2 bonded to the aromatic ring. HOCl is a strong oxidant with reduction potential of 1.49 V. Since the oxidant is in high concentration and in excess and iopamidol is also in high concentration, there will be enough collision of molecules. In a more oxidising environment aniline

(R-NH2) is expected to be oxidised to azobenzene (R–N=N–R) (Schwarzenbach et al., 2002). Because the iopamidol assumes an aniline structure, it is possible that a dimer with N=N is formed.

91

a b

CHCl2I CHCl2I Unknown T.P. CHClI2 Unknown T.P.

c d

CHCl2I CHCl2I CHClI2 CHClI2 Unknown T.P. Unknown T.P.

Figure 4.16: Proportion of iodinated DBPs in TOI at pH 6.5 at (a) 12 hr (b) 24 hr (c) 48 hr and (d) 72 hr. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Unknown T.P. is the unknown transformation products (remaining TOI).

92

a b

CHCl2I CHCl2I CHClI2 Unknown T.P. Unknown T.P.

c d

CHCl2I CHCl2I CHClI2 CHClI2 Unknown T.P. Unknown T.P.

Figure 4.17: Proportion of iodinated DBPs in TOI at pH 8.5 at (a) 12 hr (b) 24 hr (c) 48 hr and (d) 72 hr. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. Unknown T.P. is the unknown transformation products (remaining TOI).

93

a b

CHCl3 CHCl2I CHClI2 DCAA CHCl3 TCAA CHCl2I U.T.P CHClI2 DCAA TCAA U.T.P

c d

CHCl3 CHCl3 CHCl2I CHCl2I CHClI2 CHClI2 DCAA DCAA TCAA TCAA U.T.P U.T.P

Figure 4.18: Proportion of chlorinated DBPs in TOCl at pH 6.5 at (a) 12 hr (b) 24 hr (c) 48 hr and (d) 72 hr. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. U.T.P. is the unknown transformation products (remaining TOI).

94

a b

CHCl3 CHCl3 TCAA TCAA U.T.P U.T.P

d c

CHCl3 CHCl3 CHCl2I CHCl2I DCAA CHClI2 TCAA DCAA TCAA U.T.P U.T.P

f e

CHCl3 CHCl3 CHCl2I CHCl2I CHClI2 CHClI2 DCAA DCAA TCAA TCAA U.T.P U.T.P

Figure 4.19: Proportion of chlorinated DBPs in TOCl at pH 8.5 at (a) 2 hr (b) 6 hr (c) 12 hr (d) 24 hr (e) 48 hr and (f) 72 hr. [Cl2]T = 25.7 mM, [Iopamidol] = 1.29 mM, [Buffer]T = 200 mM, and temperature = 25C. U.T.P. is the unknown transformation products (remaining TOI).

95

4.3 Transformation of Iopamidol in the Presence of Chlorine and NOM

Experiments were carried out with source waters from the intake of Akron,

Barberton and Cleveland Water Treatment Plants using iopamidol in the presence of excess aqueous chlorine. These experiments were conducted to measure the degradation of TOI as a function of time and pH in the presence of NOM. In addition, the experiments were to determine the stability of TOI in the presence of

NOM. Also the formation of TOCl and iodate over 72-hour time period as a function of pH was also monitored.

The degradation of TOI followed almost the same degradation pattern for all the source waters (Figures 4.20 – 4.22). The loss of TOI ranged from 68% to 74% in

Akron source water, 62% to 72% in Barberton source water and 68% to 77% in

Cleveland source water. This may be due to the rapid oxidation of iodide on the aromatic ring to HOI (Nagy et al., 1988) which was subsequently substituted into the natural organic matter in the source waters (Kristina et al., 2009; Richardson et al.,

2007; Bichsel and von Gunten, 2000) forming TOI. There was approximately the same magnitude of degradation of TOI at the end of 72 hr in the three source waters.

The least degradation was evident at pH 6.5 whiles pH 7.5 and 8.5 were almost the same. TOX has been used as a surrogate measurement for the total halogenated DBPs formed from the reaction between chemical disinfectants and NOM (Stevens et al.,

1985; Reckhow and Singer, 1984). THMs and HAAs account for approximately 50% of TOX in chlorination of natural water (Kristina et al., 2009; Reckhow and Singer,

1984). The low degradation of TOI may imply higher formation of iodinated DBPs which are known to be highly genotoxic and cytotoxic (Richardson et al., 2008;

Plewa et al., 2004). Duirk et al (2011) indicated that the iopamidol was involved in the formation of iodo-DBPs along with NOM.

96

25

pH 6.5 pH 7.5 20 pH 8.5

15

M)



TOI ( TOI 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.20: TOI loss in chlorinated Akron source water as a function of pH. [Cl2]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 5.57 mg/L. Error bars represent 95% confidence intervals.

97

25

pH 6.5 20 pH 7.5 pH 8.5

15

M)



TOI ( TOI 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.21: TOI loss in chlorinated Barberton source water as a function of pH. [Cl2]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 4.47 mg/L. Error bars represent 95% confidence intervals.

98

25 pH 6.5 pH 7.5 20 pH 8.5

15

M)



TOI ( TOI 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.22: TOI loss in chlorinated Cleveland source water as a function of pH. [Cl2]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 2.51 mg/L. Error bars represent 95% confidence intervals.

Since the concentration of Br- was low in Akron and Barberton source waters and below detection limit in Cleveland source waters, TOBr was not detected.

Bromide in source waters are rapidly oxidised to HOBr (Hua et al., 2006) and are incorporated into THMs during chlorination (Rook, 1974). The speciation of THMs and HAAs will shift from chlorinated species to mixed species and finally to fully brominated species if Br- is in relatively high concentration (Cowman and Singer,

1996; Pourmoghaddas et al., 1993) since HOBr is more efficient at substitution while

HOCl is more effective at oxidation (Cowman and Singer, 1996; Symons et al., 1993;

Long et al., 1982).

99

Degradation of TOI in Akron and Barberton sources water were almost complete in 24 hr. It can be seen from the TOCl formation in the two source water that the incorporation of chlorine almost plateaued from 24 to 72 hr (figure 4.23 –

4.24). This was however different in the Cleveland source water as the loss of TOI was steady at 48 hr. It is evident in the TOCl formation (figure 4.25). In both Akron and Barberton waters, about 20% of the initial aqueous chlorine was incorporated into the reaction while almost 10% incorporation was observed in Cleveland water. The difference in chlorine incorporation may be due to the presence of relatively high activated aromatic structures in the NOM structure in Akron and Barberton waters which are very reactive with chlorine (Reckhow and Singer, 1985; de Laat et al.,

1982; Norwood et al., 1980). In addition, the high percentage volume of humic acids shown in the EEM of Akron and Barberton source waters may have contributed to the relatively high TOCl because aquatic humic substance consumes more chlorine and forms more TOX (Reckhow et al., 1990).

The relatively low degradation of TOI in the source waters may be partly due to relatively high SUVA254 values of the source waters. SUVA254 is used to characterise aromaticity and molecular weight distribution of NOM and significant correlations have been observed between aromaticity and DBP formation (Wu et al.,

2000; Croué et al., 2000; Reckhow et al., 1990; Singer and Chang, 1989; Edzwald et al., 1985). Also, there has been reported linkage between UV254 and the aromatic and unsaturated components of NOM (Traina et al., 1990). UV254 has been used to predict the formation of THMs and HAAs in chlorinated source waters (Singer and

Reckhow, 1999; Owen et al., 1998). Electrophilic reaction of aqueous chlorine with

NOM will produce DBPs due to electron-rich sites on the NOM molecule (Singer and

Reckhow, 1999). From table 3.1, the UV254 values for the source water were high.

100

25

20

M) 15

 pH 6.5 pH 7.5 TOCl ( TOCl 10 pH 8.5

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.23: TOCl formation in chlorinated Akron source water as a function of pH. [Cl2]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 5.57 mg/L. Error bars represent 95% confidence intervals.

101

25

20

15 pH 6.5

M)

 pH 7.5 pH 8.5

TOCl ( TOCl 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.24: TOCl formation in chlorinated Barberton source water as a function of pH. [Cl2]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 4.47 mg/L. Error bars represent 95% confidence intervals.

102

25

pH 6.5 20 pH 7.5 pH 8.5

15

M)



TOCl ( TOCl 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.25: TOCl formation in chlorinated Cleveland source water as a function of pH. [Cl2]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 2.51 mg/L. Error bars represent 95% confidence intervals.

In all the source waters, neither iodate nor iodide was detected in the samples after chlorination reaction for 72 hr. This may be as a result of HOI reacting with

NOM to reform TOI and unidentified TOI or iodo-DBPs. Iodide, a surrogate for

HOI, was not detected due to its reaction with NOM after iodide oxidation. There may be other iodide species or iopamidol transformation products formed which may have not been adsorbed on the activated carbon cartridge. The sum of the unadsorbed iodine species and TOI will satisfy the mass balance of iodine.

103

4.4 Transformation of Iopamidol in the Presence of Monochloramine and NOM

Pre-formed monochloramine was also used to monitor the degradation of iopamidol in the three source waters. In the presence of monochloramine, the degradation of TOI was almost negligible in the three source waters (figure 4.26 –

4.28). The concentration of TOI formed during chloramination was higher than chlorination. Hua and Reckhow (2006) made the same observation when they spiked surface water with inorganic iodide. It has been observed by other researchers

(Richardson et al., 2008; Krasner et al., 2006; Weinberg et al., 2002) that levels of iodinated THMs formation are higher in chloramination than chlorination. Formation of iodo-THMs is highest when chloramines are used with addition of ammonia before chlorine addition (Bichsel and von Gunten, 2000; Hansson et al., 1987). Also no idoate was detected in the source waters.

Iodide is rapidly oxidised to HOI in the presence of monochloramine (Kumar

- et al., 1986). However, monochloramine does not oxidise HOI to IO3 (Bichsel and von Gunten, 1999b). On the other hand, the HOI formed reacts with NOM to form iodo-DBPs. Duirk et al. (2011) again proposed that OCl- may be the primary reactive species in the monochloramine reaction. The formation of TOCl was relatively low

(figures 4.29 – 31) compared with TOCl formed in the chlorinated source waters. The formation of TOCl was highest at pH 6.5 followed by pH 7.5 and 8.5 in that order.

Thus HOCl may be the active oxidant after the initial hydrolysis of monochloramine to form HOCl and NH3 (Vikesland et al., 2001). The low TOCl formed may be due the hydrolysis of NH2Cl. The HOCl formed is low in concentration and will slowly react with NOM to form Chlorinated DBPs (Duirk et al, 2005; Vikesland et al., 1998;

Cowman and Singer, 1996; Jensen et al., 1985). Also monochloramine can directly react with NOM (Duirk et al., 2002). Furthermore, the loss of monochloramine could

104 have been due to its autodecomposition (Vikesland et al., 2001; Vikesland et al.,

1998). On the contrary, autodecomposition of monochloramine does not result in the formation of DBP (Duirk et al., 2002); therefore DBP formation via autodecomposition mechanism is not plausible. Kirkmeyer et al. (1993) also confirmed that the use of monochloramine for disinfection resulted in lower levels of total chlorinated by-products (measured by total organic halides).

25

pH 6.5 pH 7.5 20 pH 8.5

15

M)



TOI ( TOI 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.26: TOI degradation in chloraminated Akron source water as a function of pH. [NH2Cl]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 5.57 mg/L. Error bars represent 95% confidence intervals.

105

25

pH 6.5 20 pH 7.5 pH 8.5

15

M)



TOI ( TOI 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.27: TOI degradation in chloraminated Barberton source water as a function of pH. [NH2Cl]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 4.47 mg/L. Error bars represent 95% confidence intervals.

106

25 pH 6.5 pH 7.5 20 pH 8.5

15

M)



TOI ( TOI 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.28: TOI degradation in chloraminated Cleveland source water as a function of pH. [NH2Cl]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 2.51 mg/L. Error bars represent 95% confidence intervals.

107

25

pH 6.5 20 pH 7.5 pH 8.5

15

M)



TOCl ( TOCl 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.29: TOCl formation in chloraminated Akron source water as a function of pH. [NH2Cl]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 5.57 mg/L. Error bars represent 95% confidence intervals.

108

25

20 pH 6.5 pH 7.5 pH 8.5 15

M)



TOCl ( TOCl 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.30: TOCl formation in chloraminated Barberton source water as a function of pH. [NH2Cl]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 4.47 mg/L. Error bars represent 95% confidence intervals.

109

25

pH 6.5 pH 7.5 20 pH 8.5

15

M)



TOCl ( TOCl 10

5

0 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75

Time (hr)

Figure 4.31: TOCl formation in chloraminated Cleveland source water as a function of pH. [NH2Cl]T = 100 µM, [Iopamidol] = 5 µM, [Buffer]T = 1 mM, temperature = 25C, DOC = 2.51 mg/L. Error bars represent 95% confidence intervals.

4.5 Iodate Formation as a Function of Dissolved Organic Carbon

Iodinated DBPs are more carcinogenic than their bromine and chlorine analogues. Iodoacetic acid is the most genotoxic DBP identified to date (Richardson et al., 2008; Plewa et al., 2004). Due to this toxicological effect, the preferred sink for source water iodide in drinking water is iodate. Iodate can be reduced to iodide in vivo and in vitro (Taurog et al., 1966) – this is innocuous in quantities usually found in drinking water (Hua et al., 2006). Therefore, the DOC of Barberton and Cleveland source waters were reduced to investigate its effect on the formation of iodate in the presence of aqueous chlorine and NOM.

110

The DOC of Barberton source water was decreased to 2.235 mg/L and 1.118 mg/L while that of Cleveland source water was decreased to 1.255 mg/L and 0.628 mg/L with deionized water. About 5 μM and 100 μM iopadimol and aqueous chlorine were added to the samples respectively and stored at 25°C in the dark for 72 hr. At the end of the reaction time, the samples were quenched with 120 μM resorcinol solution and 120 μM aqueous sulphite solution to analyse for iodate and iodide respectively in the IC system. Neither iodate nor iodide was detected in the sample. Humic substances comprised of large molecular weight compounds are suspected to be precursors for THM formation potential (THMFP) and make up more than half the mass of DOC in water (Sweitlik and Sikorska, 2005). Although DOC fractionation was not carried, it is suspected that the high percentage volume of humic substances reacted with HOI to form TOI. Consequently, there was not enough concentration of HOI to be oxidised to form iodate, which is proceeds slowly (Bichsel and von Gunten, 1999b).

111

CHAPTER V

CONCLUSIONS AND RECOMMENDATIONS

5.1 Introduction

The study investigated the reaction of iopamidol with chlorinated oxidants in the absence of NOM as a function of time and pH. Degradation of iopamidol was monitored as TOI, which also includes iopamidol transformation products. Formation of iodate, TOCl and DBPs was investigated. The oxidants used were aqueous chlorine and monchloramine. Similar experiments were conducted at pH 6.5 and 8.5 using aqueous chlorine at high concentrations to investigate the loss of TOI and the formation of iodate, iodide, TOCl and DBPs. In addition, the degradation of iopamidol was monitored in the presence of NOM and chlorinated oxidants (aqueous chlorine and monochlramine). Three source waters from Akron, Barberton and

Cleveland Water Treatment Plants were used for the experiments. The loss of TOI as a function of time and pH was investigated. Furthermore formation of TOCl and iodate was investigated. Finally, formation of iodate in the presence of aqueous chlorine and NOM as a function of pH and DOC was studied.

5.2 Conclusions

1. In the absence of NOM at low reactant and buffer concentrations, the

degradation of iopamidol (TOI) was greatest at pH 7.5 and least at pH 9.5 in

112

aqueous chlorine. Approximately the same degradation was observed at pH

8.5 and 9.0. The degradation of iopamidol followed pseudo first-order

reaction kinetics at all pH except at pH 6.5, which exhibited a bi-phasic

behaviour. Since the maximum observed rate of TOI loss was at pH 7.5, it

was assumed both HOCl and OCl- participated in the degradation of iopamidol

and iopamidol transformation products.

2. At low reactant and buffer concentrations, iodate was formed and the

formation was greatest at pH 7.5 and least at 9.5. Both pH 8.5 and 9.0

exhibited the same formation pattern. Formation of iodate did not follow

either first or second order observed degradation.

3. Disinfection by-products formed at low reactant and buffer concentrations in

the absence of NOM were chloroform, trichloroacetic acid and

chlorodiiodomethane. All the DBPs were observed at pH 6.5, 7.5 and 8.5.

The formation of CHCl3 and TCAA were observed initially at 6 hr while

formation of CHClI2 was observed at 12 hr. Formation of CHCl3 increased

with increasing pH. There was however no observed difference in formation

of CHClI2 with pH. The proportions of chlorinated DBPs formed were higher

than the iodinated DBPs.

4. When high concentrations of reactants and buffer were used, degradation of

iopamidol was rapid up to 24 hr but remained fairly constant from 24 to 72 hr

at both pH 6.5 and 8.5. Also, iodate showed rapid formation from 0 to 24 hr

and the reaction stopped afterwards, that is, formation of iodate remained

fairly constant. In addition, TOCl was formed at both pH 6.5 and 8.5 after 6

hr and 2 hr respectively. DBPs formed at pH 6.5 and 8.5 were chloroform,

dichloroiodomethane, chlorodiiodomethane, trichloroacetic acid and

113

dichloroacetic acid. Higher concentrations of THMs were formed at pH 8.5

comparable to pH 6.5.

5. In the absence of NOM, insignificant degradation of iopamidol was observed

in the presence of monochloramine over the pH range of 6.5 – 8.5 under

similar experimental conditions as low concentration experiments using

aqueous chlorine. Iodate formation was not observed; however, iodide was

measured at a 2 µM pseudo steady-state concentration over the 168 hr

experiment.

6. In the presence of NOM and aqueous chlorine, TOI exhibited almost the same

degradation rate and pattern in all the three source waters at pH 6.5 to 8.5.

Degradation of TOI ranged from 62 to 77% in all the three source waters after

the 72-hr sample time. No iodate formation was observed. About 20% of the

initial aqueous chlorine (represented as TOCl) was incorporated into the

reaction in Akron and Barberton source waters while approximately 10%

aqueous chlorine was incorporated into the reaction in the Cleveland source

water.

7. Iopamidol showed no degradation in the presence of NOM and

monochloramine for pH 6.5 to 8.5 and reaction time 0 to 72 hr. Almost all the

iodide was incorporated into the NOM to form TOI. As expected, iodate was

not formed.

8. The decrease in dissolved organic carbon (DOC) in Barberton and Cleveland

source waters did not result in the formation of iodate at pH 6.5, 7.5 and 8.5

for 72 hr. The DOC of source waters were diluted to half and one quarter their

initial DOC concentration.

114

5.3 Recommendations

1. There should be a study that will compare the degradation of ICM (iopamidol)

in prechlorination followed by the addition of ammonia in the presence of

NOM with preformed monochloramine (done in this research) using the same

experimental condition used in this study.

2. A study should be carried out to investigate the optimum concentration of

dissolved organic matter for iodate formation in the presence of aqueous

chlorine and NOM using the same source waters. The same experimental

conditions should be used.

3. An investigation into the speciation of total organic halogen (TOX) in source

water spiked with varying concentrations of bromide while maintaining all

other conditions used in this research should be carried out.

115

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