total organic halogen formation in the presence of iopamidol and chlorinated oxidants with and
Transcript of total organic halogen formation in the presence of iopamidol and chlorinated oxidants with and
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
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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
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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
x
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
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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 = 25C. 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 = 25C. 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 = 25C. 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 = 25C. 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 = 25C. 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 = 25C. 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 = 25C. Error bars represent 95% confidence intervals...........................86
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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 = 25C. 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 = 25C. 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 = 25C. 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 = 25C. 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 = 25C. 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 = 25C..…………………………………………………………………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 = 25C…………..……………………………………...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 = 25C, 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 = 25C, 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 = 25C, 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 = 25C, 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 = 25C, 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 = 25C, 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 = 25C,
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 = 25C,
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 = 25C,
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 = 25C,
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 = 25C,
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 = 25C,
DOC = 2.51 mg/L. Error bars represent 95% confidence intervals...........................110
1
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
2
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
3
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
4
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,
5
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).
6
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
7
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
8
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.
9
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
10
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,
11
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
12
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).
13
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
14
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
15
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
-]
16
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
17
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).
18
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
19
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,
20
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
2I3 I6
2 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,
21
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.
22
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-1
s-1
. However, if OI- was the reactive species the maximum rate constant
(kNH2Cl+OI-) was 3 M-1
s-1
.
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
23
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
24
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
25
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
2.18 k1 = 4.3 x 108 M
-1s
-1 Nagy et al., 1988
2.19 k2 = 1550 M-1
s-1
Kumar and
Margerum, 1987
2.20
k3 = 8.2 M-1
s-1
Bichsel and von
Gunten, 1999b
2.21
k4 = 52 M-1
s-1
Bichsel and von
Gunten, 1999b
2.22 HOCl + NOM Products K5 = 00.7 – 5 M-1
s-1
Westerhoff et al.,
2004
2.23 HOBr + NOM Products K6 =15 – 167 M-1
s-1
Westerhoff et al.,
2004
2.24 HOI + NOM Products (TOI) K7 =0.1 – 0.4 M-1
s-1
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
26
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
27
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
28
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.
29
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
30
(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).
31
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).
32
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
water
Barberton
source water
Cleveland
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
UV254 (cm-1
) 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
33
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.
34
Table 3.2: Florescence EEM regions proposed by Chen et al. (2003)
Regions Representation
Excitation
Range (nm)
Emission Range
(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
Fluorescence Regions
Akron Barbertion Cleveland
% % %
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
35
Figure 3.1: Fluorescence excitation-emission spectrum of Akron source water.
[DOC] = 5.57 mg/L, SUVA254 = 2.27 L/mg.m
Emission (nm)
300 320 340 360 380 400 420 440 460 480 500 520 540
Ex
cita
tion (
nm
)
250
300
350
400
0
10
20
30
40
50
60
I IIIII
IV V
36
Figure 3.2: Fluorescence excitation-emission spectrum of Barberton source water.
[DOC] = 4.47 mg/L, SUVA254 = 4.31 L/mg.m
Emission (nm)
300 320 340 360 380 400 420 440 460 480 500 520 540
Ex
cita
tion (
nm
)
220
240
260
280
300
320
340
360
380
4000
5
10
15
20
25
30
I II III
IVV
37
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
Emission (nm)
300 320 340 360 380 400 420 440 460 480 500 520 540
Ex
cita
tion (
nm
)
220
240
260
280
300
320
340
360
380
400
0
5
10
15
20
25
I II III
IV
V
38
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
39
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.
40
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
41
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
Standard Concentration (μM)
Recovery
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
42
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
43
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.
44
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.
45
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
46
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
47
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.
48
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.
49
Figure 3.5: Gradient profile for the analysis of Total organic halogen
Time (min)
0 5 10 15 20
Elu
ent
Conce
ntr
atio
n (
mM
)
0
10
20
30
40
50
60
50
Figure 3.6: Calibration curve for Chloride using 2,4,6-trichlorophenol. [Cl-] = 0 – 250
μM
y= 5.6135*x-5.250
R2=0.9964
Area ( S*min)
0 10 20 30 40 50
Co
ncen
trati
on (
M)
0
50
100
150
200
250
300
Chloride
51
Figure 3.7: Calibration curve for Iodide using 4-iodophenol. [I-] = 0 – 50 μM
y= 5.1236*x+1.2031
R2=0.9988
Area ( S*min)
0 2 4 6 8 10
Conce
ntr
atio
n (
M)
0
10
20
30
40
50
60
Iodide
52
Figure 3.8: Calibration curve for Bromide using 4-iodophenol. [Br-] = 0 – 50 μM
Area (S.min)
0 2 4 6 8
Co
nce
ntr
atio
n (
M)
0
10
20
30
40
50
60
y = 6.4756x + 2.0445
R2 = 0.9917
53
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.
Area (S.min)
0 5 10 15 20 25 30
Con
centr
atio
n (
M)
0
20
40
60
80
100 y = 3.995x + 1.1775
R² = 0.9995
54
Figure 3.10: Gradient profile for the analysis of iodate.
Figure 3.11: Calibration curve for Iodate using NaIO3. [IO3-] = 0 – 20 μM
Time (min)
0 5 10 15 20
Elu
ent
Conce
ntr
atio
n (
mM
)
0
5
10
15
20
25
30
35
Area (S.min)
0 1 2 3 4
Conce
ntr
atio
n (
M)
0
5
10
15
20
25
y = 5.4459x
R² = 0.9997
55
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
56
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.
57
Figure 3.12: Calibration curve for CHCl3using chloroform. [CHCl3] = 0 – 1000 nM
Figure 3.13: Calibration curve for CHBr2Cl using dibromochloromethane. [CHBr2Cl]
= 0 – 300 nM
Response ratio
0.0 0.2 0.4 0.6 0.8 1.0
[CH
Cl 3
] (n
M)
0
200
400
600
800
1000
1200
y = 1471.4x
R² = 0.9785
Response ratio
0.0 0.5 1.0 1.5 2.0 2.5 3.0
[CH
Br 2
Cl]
(nM
)
0
50
100
150
200
250
300
350
y = 109.38x
R² = 0.9974
58
Figure 3.14: Calibration curve for CHBrI2 using bromodiiodomethane. [CHBrI2] = 0 –
125 nM
Figure 3.15: Calibration curve for CHClI2 using chlorodiiodomethane. [CHClI2] = 0 –
250 nM
Response ratio
0.0 0.2 0.4 0.6 0.8 1.0
[CH
BrI
2]
(nM
)
0
20
40
60
80
100
120
140
160
y = 135.95x
R² = 0.9925
Response ratio
0.0 0.2 0.4 0.6 0.8 1.0
[CH
ClI
2] (
nM)
0
50
100
150
200
250
300
y = 266.06x
R² = 0.9993
59
Figure 3.16: Calibration curve for CHBr2I using dibromoiodomethane. [CHBr2I] = 0 –
250 nM
Figure 3.17: Calibration curve for CHBrClI using bromochloroiodomethane.
[CHBrClI] = 0 – 250 nM
Response ratio
0.00 0.05 0.10 0.15 0.20 0.25
[CH
Br 2
I] (
nM)
0
50
100
150
200
250
300
y = 1152.7x
R² = 0.9941
Response ratio
0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 0.16 0.18
[CH
BrC
lI]
(nM
)
0
50
100
150
200
250
300
y = 1595.3x
R² = 0.998
60
Figure 3.18: Calibration curve for CHBr3 using bromoform. [CHBr3] = 0 – 500 nM
Figure 3.19: Calibration curve for CHCl2I using dichloroiodomethane. [CHCl2I] = 0 –
500 nM
Response ratio
0.0 0.5 1.0 1.5 2.0 2.5 3.0
[CH
Br 3
] (n
M)
0
100
200
300
400
500
600
y = 204.28x
R² = 0.9968
Response ratio
0.00 0.05 0.10 0.15 0.20 0.25 0.30
[CH
Cl 2
I] (
nM
)
0
100
200
300
400
500
600
y = 2013.1x
R² = 0.9985
61
Figure 3.20: Calibration curve for CHCl2Br using bromodichloromethane. [CHCl2Br]
= 0 – 400 nM
Figure 3.21: Calibration curve for CHI3 using iodoform. [CHI3] = 0 – 50 nM
Response Ratio
0.0 0.5 1.0 1.5 2.0 2.5 3.0
[CH
Cl 2
Br]
(nM
)
0
100
200
300
400
y = 158.05x
R² = 0.9965
Response ratio
0.0 0.2 0.4 0.6 0.8 1.0
[CH
I 3]
(nM
)
0
10
20
30
40
50
60
y = 68.338x
R² = 0.9933
62
Figure 3.22: Calibration curve for CAN using chloroacetonitrile. [CAN] = 0 – 500 nM
Figure 3.23: Calibration curve for DCAN using dichloroacetonitrile. [DCAN] = 0 –
500 nM
Response ratio
0.0 0.2 0.4 0.6 0.8 1.0
[CA
N]
(nM
)
0
100
200
300
400
500
600
y = 478.7x
R² = 0.9988
Response ratio
0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 0.16 0.18 0.20
[DC
AN
] (n
M)
0
100
200
300
400
500
600
y = 2852.9x
R² = 0.9991
63
Figure 3.24: Calibration curve for TCAN using trichloroacetonitrile. [TCAN] = 0 –
125 nM
Figure 3.25: Calibration curve for BAN using bromoacetonitrile. [BAN] = 0 – 125
nM
Response ratio
0.0 0.2 0.4 0.6 0.8 1.0
[TC
AN
] (n
M)
0
20
40
60
80
100
120
140
y = 160.4x
R² = 0.9631
Response ratio
0.0 0.2 0.4 0.6 0.8 1.0 1.2
[BA
N]
(nM
)
0
20
40
60
80
100
120
140
y = 129.6x
R² = 0.9987
64
Figure 3.26: Calibration curve for DBAN using dibromoacetonitrile. [DBAN] = 0 –
250 nM
Figure 3.27: Calibration curve for BCAN using bromochloroacetonitrile. [BCAN]=0–
250 nM
Response ratio
0.0 0.2 0.4 0.6 0.8
[DB
AN
] (n
M)
0
50
100
150
200
250
300
y = 346.63x
R² = 0.9977
Response ratio
0.0 0.2 0.4 0.6 0.8
[BC
AN
] (n
M)
0
50
100
150
200
250
300
y = 372.69x
R² = 0.9931
65
Figure 3.28: Calibration curve for IAN using iodoacetonitrile. [IAN] = 0 – 31 nM
Figure 3.29: Calibration curve for CAA using chloroacetic acid. [CAA] = 0 – 250 nM
Response ratio
0.0 0.2 0.4 0.6 0.8 1.0 1.2
[IA
N]
(nM
)
0
5
10
15
20
25
30
35
y = 33.062x
R² = 0.9946
Response ratio
0.00 0.05 0.10 0.15 0.20 0.25 0.30 0.35
[CA
A]
(nM
)
0
50
100
150
200
250
300
y = 773.91x
R² = 0.9997
66
Figure 3.30: Calibration curve for DCAA using dichloroacetic acid. [DCAA] = 0 –
500 nM
Figure 3.31: Calibration curve for TCAA using trichloroacetic acid. [TCAA] = 0 –
250 nM
Response ratio
0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6
[DC
AA
] (n
M)
0
100
200
300
400
500
600
y = 376.03x
R² = 0.9693
Response ratio
0.0 0.5 1.0 1.5 2.0 2.5 3.0
[TC
AA
] (n
M)
0
50
100
150
200
250
300
y = 100.41x
R² = 0.9986
67
Figure 3.32: Calibration curve for BCAA using bromochloroacetic acid. [BCAA]=0–
250 nM
Figure 3.33: Calibration curve for BDCAA using bromodichloroacetic acid.
[BDCAA] = 0 –250 nM
Response ratio
0.0 0.5 1.0 1.5 2.0 2.5 3.0
[BC
AA
] (n
M)
0
50
100
150
200
250
300
y = 105.74x
R² = 0.9969
Response ratio
0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5
[BD
CA
A]
(nM
)
0
50
100
150
200
250
300
y = 85.951x
R² = 0.9982
68
Figure 3.34: Calibration curve for BAA using bromoacetic acid. [BAA] = 0 – 1000
nM
Figure 3.35: Calibration curve for DBAA using dibromoacetic acid. [DBAA] = 0 –
500 nM
Response ratio
0.0 0.2 0.4 0.6 0.8
[BA
A]
(nM
)
0
200
400
600
800
1000
1200
y = 1378.1x
R² = 0.991
Response ratio
0 1 2 3 4
[DB
AA
] (n
M)
0
100
200
300
400
500
600
y = 139.48xR² = 0.9949
69
Figure 3.36: Calibration curve for IAA using iodoacetic acid. [IAA] = 0 – 125 nM
Response ratio
0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8
[IA
A]
(nM
)
0
20
40
60
80
100
120
140
y = 83.588x
R² = 0.9896
70
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
71
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
72
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
observed rate constants (kobs) (figure 4.2). The kobs for pH 7.5 and 9.5 are 3.97 x 10-6
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.
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;
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
I (
M)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
pH 9.0
pH 9.5
73
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.
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
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
ln (
[TO
I]t/
[TO
I]0)
-1.4
-1.2
-1.0
-0.8
-0.6
-0.4
-0.2
0.0
pH 6.5 y = -0.0117x + 0.0436, R² = 0.9848pH 7.5 y = -0.0143x - 0.0296, R² = 0.9766
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
74
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.
75
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
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
Iodat
e (
M)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
pH 9.0
pH 9.5
76
constant for HOCl (8.2±0.8 M-1
.s-1
). 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.
77
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 = 25C. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
Conce
ntr
atio
n (
nM
)
0
100
200
300
400
CHCl3
CHClI2
TCAA
78
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 = 25C. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
Conce
ntr
atio
n (
nM
)
0
100
200
300
400
CHCl3
CHClI2
TCAA
79
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 = 25C. 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%
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
Co
nce
ntr
atio
n (
nM
)
0
100
200
300
400
CHCl3
CHClI2
TCAA
80
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
expected that HOI will be reduced to I- whiles SO3
2- will be oxidised to SO4
2-.
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.
81
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 = 25C. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
Cl
( M
)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
pH 9.0
pH 9.5
82
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.
Time (hr)
0 25 50 75 100 125 150 175
TO
I (
M)
0
5
10
15
20
pH 6.5
pH 7.5
pH 8.5
pH 9
83
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
Time (hr)
0 25 50 75 100 125 150 175
[I- ]
( M
)
0
1
2
3
4
5
pH 6.5
pH 7.5
pH 8.5
pH 9
84
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.
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 = 25C. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
Conce
ntr
atio
n (
M)
0
1000
2000
3000
4000
5000
6000
7000
TOI
Iodide
Iodate
[I]T
85
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 = 25C. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
Conce
ntr
atio
n (
M)
0
1000
2000
3000
4000
5000
6000
TOI
Iodide
Iodate
[I]T
86
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 = 25C. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
Cl
( M
)
0
500
1000
1500
2000
2500
87
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 = 25C. 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.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
Cl
( M
)
0
500
1000
1500
2000
2500
3000
88
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 = 25C. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
Conce
ntr
atio
n (
nM
)
0
5e+4
1e+5
2e+5
2e+5
CHCl3
CHCl2I
CHClI2
DCAA
TCAA
89
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 = 25C. 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
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
Conce
ntr
atio
n (
nM
)
0
5e+4
1e+5
2e+5
2e+5
CHCl3
CHCl2I
CHClI2
DCAA
TCAA
90
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.
91
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.
92
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 = 25C. Unknown T.P. is the unknown transformation products
(remaining TOI).
CHCl2IUnknown T.P.
CHCl2I
CHClI2
Unknown T.P.
CHCl2I
CHClI2
Unknown T.P.
CHCl2I
CHClI2
Unknown T.P.
a b
c d
93
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 = 25C. Unknown T.P. is the unknown transformation products
(remaining TOI).
CHCl2I
Unknown T.P.
CHCl2I
CHClI2
Unknown T.P.
CHCl2I
CHClI2
Unknown T.P.
CHCl2I
CHClI2
Unknown T.P.
a b
c d
94
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 = 25C. U.T.P. is the unknown transformation products
(remaining TOI).
CHCl3CHCl2ICHClI2DCAATCAAU.T.P
CHCl3CHCl2ICHClI2DCAATCAAU.T.P
CHCl3CHCl2ICHClI2DCAATCAAU.T.P
CHCl3CHCl2ICHClI2DCAATCAAU.T.P
a b
c d
95
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 = 25C. U.T.P. is the unknown transformation
products (remaining TOI).
CHCl3
TCAA
U.T.P
CHCl3
TCAA
U.T.P
CHCl3
CHCl2I
DCAA
TCAA
U.T.P
CHCl3CHCl2ICHClI2DCAATCAAU.T.P
CHCl3CHCl2ICHClI2DCAATCAAU.T.P
CHCl3
CHCl2I
CHClI2
DCAA
TCAA
U.T.P
a b
c d
e f
96
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.
97
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 = 25C, DOC
= 5.57 mg/L. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
I (
M)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
98
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 = 25C, DOC
= 4.47 mg/L. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
I (
M)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
99
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 = 25C, 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).
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
I (
M)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
100
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.
101
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 = 25C, DOC
= 5.57 mg/L. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
Cl
( M
)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
102
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 = 25C,
DOC = 4.47 mg/L. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
Cl
( M
)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
103
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 = 25C,
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.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
Cl
( M
)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
104
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
105
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).
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 =
25C, DOC = 5.57 mg/L. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
I (
M)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
106
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 =
25C, DOC = 4.47 mg/L. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
I (
M)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
107
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 =
25C, DOC = 2.51 mg/L. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
I (
M)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
108
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 =
25C, DOC = 5.57 mg/L. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
Cl
( M
)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
109
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 =
25C, DOC = 4.47 mg/L. Error bars represent 95% confidence intervals.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
Cl
( M
)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
110
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 =
25C, 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.
Time (hr)
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75
TO
Cl
( M
)
0
5
10
15
20
25
pH 6.5
pH 7.5
pH 8.5
111
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).
112
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
113
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
114
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.
115
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.
116
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