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

ii

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

iii

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.

iv

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.

v

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.

vi

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

vii

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

viii

5.1 Introduction .......................................................................................................... 112

5.2 Conclusions .......................................................................................................... 112

5.3 Recommendations ................................................................................................ 115

REFERENCES .......................................................................................................... 116

ix

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

xi

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

xii

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







4.7 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


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



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


xiii

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


4.14 THM and HAA formation in reaction mixtures containing iopamidol and



intervals........................................................................................................................88

4.15 THM and HAA formation in reaction mixtures containing iopamidol and



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.



xiv

4.25 TOCl formation in chlorinated Cleveland source water as a function of pH.



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.



4.28 TOI degradation in chloraminated Cleveland source water as a function of pH.



4.29 TOCl formation in chloraminated Akron source water as a function of pH.



4.30 TOCl formation in chloraminated Barberton source water as a function of pH.



4.31 TOCl formation in chloraminated Cleveland source water as a function of pH.



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

http://www.whocc.no/atc_ddd_index/?code=V08A

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).

http://www.epa.state.oh.us/dsw/tmdl/CuyahogaRiver.aspx

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-tribromophenol 15 60.31 68.28



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




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




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,


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,


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



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



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




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



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.


= 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.



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.



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.



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 =


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



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 =


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



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



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

REFERENCES

Abia, L., Armesto, X.L., Canle, L.M., Garcia, M.V. and Santaballa, J.A.

(1998). Oxidation of aliphatic amines by aqueous chlorine. Tetrahedron 54: 521–530

Adams, C.D. (2009). Pharmaceuticals. In: Contaminants of emerging

environmental concern (Eds.: Bhandari, A., Surampali, R.Y., Adams, C.D.,

Champagne, P., Ong, S.K., Tyagi, R.D. and Zhang, T.C.). American Society of Civil

Engineers, US

Adam, L.C. and Gordon, G. (1999). Hypochlorite ion decomposi-tion: effects

of temperature, ionic strength, and chloride ion. Inorg. Chem. 38: 1299-1304.

Agus, E., Voutchkov, N., Sedlak, D.L., (2009). Disinfection by-products and

their potential impact on the quality of water produced by desalination systems: A

literature review. Desalination 237, 214-237.

Akin, E.W., Hoff, J.C. and Lippy, E.C. (1982). Waterborne outbreak control:

which disinfectant? Environ. Health Perspect. 46: 7 – 12

Allan,J. D., and Castillo, M. M. (2007). Stream Ecology: Structure and

function of running waters. Springer, Dordrecht, The Netherlands

American Water Works Association Research Foundation (AWWARF),

(1987). Current Methodology for the control of algae in surface water. Research

report, AWWA, Denver, CO

APHA, AWWA and WEF, (1995). Micro liquid-liquid extraction gas

chromatographic method. In Standard methods for in the examination of water wand

wastewater. American Public Healthh Association (APHA), American Water Works

Association (AWWA) and Water Environment federation (WEF).

Armesto, X.L., Canle, L.M., Garcia, M.V., Losada, M. and Santaballa, J.A.,

(1994). Chlorination of dipeptides by hypochlorous acid in aqueous solution. Gazz.

Chim. Ital. 124: 519–523

AWWA (2008). Committee report: Disinfection survey, Part 1–recent

changes, current practices and water quality. J. AWWA 100 (10):76–90

AWWA (2000). Committee report: Disinfection at large and medium size

systems. Journal AWWA, 92 (5): 32 – 43

Bahr, C., Schumacher, J., Ernst, M., Luck, F., Heinzmann, B. and Jekel, M.

(2007). SUVA as a control parameter for the effective ozonation of organic pollutants

in secondary effluent. Water Sci. Technol. 55 (12): 267–274.

117

Baribeau, H., Pre´vost, M., Desjardins, R. and Lafrance, P. (2001). Changes in

chlorine and DOX concentrations in distribution systems. J. of the AWWA 93: 102–

114.

Batt, A. L., Kim, S. and Aga, D. (2006). Enhanced biodegradation of

iopromide and trimethoprim in nitrifying sludge. Environ. Sci. Technol. 40 (23):

7367–7373.

Bell, R.P. and Gelles, E. (1951). The halogen cations in aqueous solution.

Journal of Chemical Society 73: 2734 – 2740

Bellar, T. A., Lichtenberg, J. J., and roner, R. C. (1974). “The occurrence of

organohalides in chlorinated drinking water.” J of AWWA. 66(12): 703-706.

Benfield, E. F. (2006). Decomposition of Leaf Material in F. R. Hauer and G.

A. Lamberti, editors. Methods in Stream Ecology. Academic Press, Burlington, MA,

USA.

Benjamin, M.M. and Lawler, D.F. (2013). Water quality engineering: Physical

and chemical treatment processes. John Wiley and Sons Inc.

Betts, K. (1998). Growing concern about disinfection by-products. Environ.

Sci. and Technol. 546A-548A.

Bichsel, Y. (2000). Behaviour of iodine species in oxidative process during

drinking water treatment. Doctoral dissertation, Swiss Federal Institute of Tech. Diss.

ETH No 13429

Bichsel Y. and von Gunten U. (2000). Formation of iodo-trihalomethanes

during disinfection and oxidation of iodide-containing waters. Environ. Sci.

Technol 34: 2784–2791

Bichsel Y. and von Gunten U. (1999a). Hypoiodous acid: Kinetics of the

buffer-catalysed disproportionation. Wat. Res. 34 (12): 3197 – 3203.

Bichsel Y. and von Gunten U. (1999b). Oxidation of iodide and hypoiodous

acid in the disinfection of natural waters. Environ. Sci. Technol. 33: 4040 – 4045.

Brezonik, P.L. (1994). Chemical kinetics and Process Dynamics in Aquatic

Systems, Lewis Publishers Inc. Boca Raton, FL, USA.

Bruchet A. and Duguet, J.P. (2004). Role of oxidants and disinfectants on the

removal, masking and generation of tastes and odours. Water Sci. and Techn. 49(9):

297- 306

Bryant, E. A., Fulton, G. P., and Budd, G. C. (1992). Disinfection Alternatives

for Safe Drinking Water. Hazen and Sawyer. Non Nostrand Reinhold: New York.

Bull, R.J., Birnbaum, L.S., Cantor, K.P., Rose, J.B., Butterworth, B.E.,

Pegram, R. and Tuomisto, J. (1995). Water chlorination: essential process or cancer

hazard? Fundam. Appl. Toxicol. 28 (2): 155-166.

118

Bull, R.J. et al. (1991) Health effects of disinfectants and disinfection by-

products. Denver, CO, American Water Works Association

Burgot, J. (2012). Ionic equilibra in analytical chemistry. Springer Science and

Business media, NY, USA

Busetti, F., Linge, K. L., Blythe, J. W. and Heitz, A. (2008). Rapid analysis of

iodinated X-ray contrast media in secondary and tertiary wastewater by direct

injection liquid chromatography-tandem mass spectrometry. J Chromatogr A. 1213

(2): 200–208.

Carballa, M., Omil, F., Ternes, T. A. and Lema, J. M.(2007). Fate of

pharmaceuticals and personal care products (PPCPs) during anaerobic digestion of

sewage sludge. Water Res. 41 (10): 2139–2150.

Cherney, D.P., Duirk, S.E., Tarr, J.C. and Colette, T.W. (2006). Monitoring

the speciation of aqueous free chlorine from pH 1 to 12 with Raman spectroscopy to

determine the identity of potent low pH oxidant. Appl. Spectroscopy. 60: 764–772.

Choi, J. and Valentine, R.L. (2001). Formation of N-nitrosodimethylamine

(NDMA) in Chloraminated Water: New disinfection by-product. Proceedings the

221st National Meeting (Environmental Division), San Diego, CA, April 1-5, 2001,

Vol 41, No. 1, pg 8-11.

Clough, P.N. and Starke, H.C. (1985). A review of the aqueous chemistry and

partitioning of inorganic iodine under LWR severe accident conditions. European

Applied Research Reports: Nuclear. Sci. and Tech. 6: 631 – 776.

Cowman, G. A. and Singer, P. C. (1996). Effect of bromide ion on haloacetic

acid speciation resulting from chlorination and chloramination of aquatic humic

substances. Environ. Sci. Technol. 30: 16-24.

Crittenden, J. C., Trussell, R. R., Hand, D. W., Howe, K. J. and

Tchobanoglous, G. (2012). Water Treatment: Principles and Design. John Wiley &

Sons Inc

Daughton, C.G. and Ternes, T.A. (1999). Pharmaceuticals and Personal Care

Products in the Environment: Agents of Subtle Change? Environmental Health

Perspect. 107: 907-938.

Deborde, M. and von Gunten, U. (2008). Reactions of chlorine with inorganic

and organic compounds during water treatment - kinetics and mechanisms: A critical

review. Water Research, 42(1-2): 13-51

Deborde, M., Rabouan, S., Gallard, H. and Legube, B. (2004). Aqueous

chlorination kinetics of some endocrine disruptors. Environ. Sci. Technol. 38: 5577–

5583.

Diehl, A.C., Speitel, G.E., Symons, J.M., Krasner, S.W., Hwang, C.J. and

Barrett, S.E. (2000). DBP formation during chloramination. J of AWWA 92: 76–90.

119

Dodd, M.C., Shah, A.D., von Gunten, U. and Huang, C.H. (2005). Interactions

of fluoroquinolone antibacterial agents with aqueous chlorine: reaction kinetics,

mechanisms, and transformation pathways. Environ. Sci. Technol. 39: 7065–7076.

Dore´, M. (1989). Chimie des oxydants et traitement des eaux, Edition

Technique et Documentation. Lavoisier, Paris.

Dressman, R.C. and Stevens, A.A.(1983). Analysis of organohalides in water

– an evaluation update. J. Am Water Works Assoc. 75:431 – 434.

Drewes ,J.E., Fox, P. and Jekel, M. (2001). Occurrence of iodinated X-ray

contrast media in domestic effluents and their fate during indirect potable reuse. J.

Environ. Sci. Health Part AToxic/Hazard. Subst. Environ. Eng. 36(9):1633-1645

Duirk, S.E., Lindell, C., Cornelison, C.C., Kormos, J., Ternes, T.A., Attende-

Ramos, M., Osiol, J., Wagner, E.D., Plewa, M.J. and Richardson, S.D. (2011).

Formation of toxic iodinated disinfection by-products from compounds used in

medical imaging. Environ. Sci. and Tech. 45(16): 6845 – 6854.

Duirk, S.E., Gombert, B., Croue, J.-P. and Valentine, R.L. (2005).Modeling

monochloramine loss in the presence of natural organic matter. Water Research 39

(14), 3418 – 3431

Duirk, S.E., Gombert, B., Choi, J., and Valentine, R.L. (2002).

Monochoramine loss in the presence of humic acid. J. Environ. Monitor. 4 (1), 85–89.

Elliott, S., Lead, J.R. and Baker, A. (2006). Characterisation of the

fluorescence from freshwater, planktonic bacteria. Water Research 40:2075-2083

Fabian, I. and Gordon G. (1997). The kinetics and mechanism of the chlorine

dioxide iodide ion reaction. Inorganic Chemistry, 36(12): 2494-2497.

Fuge, R. and Johnson C.C. (1986). The geochemistry of iodine - a review.

Environmental Geochemistry and Health. 8(2): 31-54.

Gallard, H., Leclercq, A. and Croue´, J.P. (2004). Chlorination of bisphenol a:

kinetics and byproducts formation. Chemosphere 56: 465–473.

Gallard, H. and von Gunten, U. (2002). Chlorination of phenols: kinetics and

formation of chloroform. Environ. Sci. Technol. 36: 884–890

Gang, D., Clevenger, T.E. and Banerji, S.K. (2003). Relationship of chlorine

decay and THMs formation to NOM Size. J. Hazard. Mater. 96(1): 1-12.

Garland, J.A., Elzerman, A.W. and Penkett,, S.A. (1980). The mechanism for

dry deposition of ozone to seawater surfaces. J. Geophys. 85(C12): 7488 – 7492

Gartiser, S.; Brinker, L.; Erbe, T.; Kummerer, K.; Willmund, R. (1996).

Contamination of hospital wastewater with hazardous compounds as defined by 7a

WHG. Acta Hydrochim. Hydrobiol. 24 (2): 90–97.

Gates, D. (1998). The Chlorine Dioxide Handbook. AWWA, Denver, CO.

120

Gerritsen, C.M. and Margerum, D.W. (1990). Non- metal kinetics:

hypochlorite and hypochlorous acid reactions with cyanide. Inorg. Chem. 29: 2757–

2762.

Gordon, G. and Bubnis, B. (2000). Sodium hypochlorite speciations, in Proc.

of the AWWA and Water Quality Technology Conference, Denver, CO, USA, June

11-15th

, 2000.

Gottardi. W. (1983). Iodine and iodine compounds. In: Disinfection,

Sterilization, and Preservation. Ed.: S.S. Block; Lea & Febiger, Philadelphia,

Pennsylvania: 83-196.

Gottardi, W. (1981). The formation of iodate as a reason for the decrease of

efficiency iodine containing disinfectant. Zentralblatt fur Bakterologie Hyg 1 Abt

Orig B 172: 498–507

Grassi, M., Kaykioglu, G., Belgiorno, V. and Lofrano, G. (2012). Emerging

Compounds removal from Wastewater. Springer Briefs in Green Chemistry for

Sustainability. 10: 978-994.

Greenwood, N.N. and Earnshaw, A. (1984). Chemistry of the elements.

Pergamom Press, Oxford.

Halling-Sprensen, B., Nielsen, S.N., Lanzky, P.F., Ingerslev, F., Lηzhρft,

H.C.H. and Jρgensen, S.E. (1998). Occurrence, fate and effects of pharmaceuticals

substances in the environment – a review. Chemosphere 36: 357 – 393.

Hansson R.C., Henderson M.J., Jack R., and Taylor R.D. (1987). Iodoform

taste complaints in chloramination. Water Res. 21: 1265–1271

Harrington-Brook, K., Doerr, C.L. and Moore, M.M. (1998). Mutagenicity of

three disinfection by-products: di- and trichloroacetic acid and chloral hydrate in

L5178Y/T /− (-)3.7.2C mouse lymphoma cells. Mutat Res. 413: 265–276

Hirsch, R., Ternes, T. A., Lindart, A., Haberer, K. and Wilken, R.D. (2000). A

sensitive method for the determination of iodine containing diagnostic agents in

aqueous matrices using LC-electrospray tandem-MS detection. Fresenius J Anal

Chem.366 (8): 835–841.

Hoff, J.C. and Geldreich, E.E. (1981). Comparison of the biocidal efficiency

of alternative disinfectants. J. Am. Water Works Assoc. 73, 40–44

Hoigne´, J. (1998). Chemistry of aqueous ozone and transformation of

pollutants by ozonation and advances oxidation processes. In: Hubrec, J. (Ed.), The

Handbook of Environmental Chemistry Quality and Treatment of Drinking Water.

Springer, Berlin.

Hoigne, J. and Bader, H. (1994). Kinetics of reactions of chlorine dioxide

(ClO2) in water – I. Rate constants for inorganic and organic compounds. Water Res.

28: 45–55.

121

Hua, G. and Reckhow, D.A. (2006). Determination of TOCl, TOBr, and TOI

in drinking water by pyrolysis and off-line ion chromatography. Analytical and

Bioanalytical Chemistry 384: 495–504.

Hua, G., Reckhow, D.A. and Kim, J. (2006). Effect of Bromide and Iodide

Ions on the Formation and Speciation of Disinfection by-products during

Chlorination. Environ. Sci & Technol. 40: 3050-3056.

Hudson, N., Baker, A. and Reynolds, D. (2007). Fluorescence analysis of

dissolved organic matter in natural, waste and polluted waters- A review. River

Research and Applications 23:631-649

Jafvert, C.T. and Valentine, R.L. (1992). Reaction scheme for the chlorination

of ammoniacal water. Environ. Sci. Technol. 26: 577–586.

Jekel, M.R. and Roberts, P.V. (1980). Total Organic Halogen as a Parameter

for the Characterization of Reclaimed Waters: Measurement, Occurrence, Formation,

and Removal. Environmental Science and Technology 14(8): 970-975.

Jensen, J.N., Johnson, J.D., Aubin, J. and Christman, R.F. (1985). Effect of

monochloramine on isolated fulvic acid. Org. Geochem. 8 (1), 71–76.

Johnson, D.W. and Margerum, D.W. (1991). Non-metal redox kinetics: a

reexamination of the mechanism of the reaction between hypochlorite and nitrite ions.

Inorg. Chem. 30: 4845–4851.

Johnson, J.D. and Jensen, J.N. (1986). THM and TOX formation: routes, rates

and precursors. J. Am. Water Works Assoc. 78: 156 – 162.

Jolley, R. L. and Carpenter, J. H. (1983). A review of the chemistry and

environmental fate of reactive oxidant species in chlorinated water. In Water

Chlorination: Environmental Impacts and Health Effects. Vol. 4. Book 1. Eds: Jolley,

R.L., Brungs, W.A., Cotruvo, J.A., Gumming, R.B., Mattice, J.S. and Jacobs, V.A.

Ann Arbor Science Publishers.

Kargalioglu, Y., McMillan, B.J., Minear, R.A. and Plewa, M.J.

(2002).Analysis of the cytotoxicity and mutagenicity of drinking water disinfection

by-products in salmonella typhimurium.Teratogenesis Carcinogenesis and

Mutagenesis 22 (2): 113-128

Kavanaugh, M.C., Trussel, A.R., Cromer, J., and Trussel, R.R. (1980). An

empirical kinetic model of trihalomethane formation: application to meet the proposed

THM standard. J. AWWA 72 (10):578.

Kirkmeyer, G.J., Martel, K., Thompson, G., and Radder, L. (1993).

Optimizing Chloramine treatment. Prepared for the AWWARF, Denver, CO.

Knocke, R., van Benschoten,J.E., Kearney, M., Soborski, A. and Reckho,

D.A. (1990). Alternate oxidants for the removal of soluble iron and manganese. AA

Research Foundation, Denver, CO.

122

Kogevinas, M., Villanueva, C. M., Font-Ribera, L., Liviac, D., Bustamante,

M., Espinoza, F. and Marcos, R. (2010). Genotoxic effects in swimmers exposed to

disinfection by-products in indoor swimming pools. Environ. Health Perspectives

118 (11): 1531-1537

Koivusalo, M., Jaakkola, J. J., Vartiainen, T., Hakulinen, T., Karjalainen, S.,

Pukkala, E. And Tuomisto, J. (1994). Drinking water mutagenicity and

gastrointestinal and urinary tract cancers: an ecological study in Finland. Am. J.

Public Health 84 (8) 1223-1228.

Koplin, D.W., Furlong, E.T., Meyer, M.T., Thurman, E.M., Zaugg, S.D.,

Barber, L.B., and Buxton, H.T. (2002). Pharmaceuticals, hormones and other organic

wastewater contaminants in US streams, 1999-2000: A national reconnaissance.

Environ. Sci. & Technol. 38 (23): 6377-6384

Kormos, J. L., Schulz, M., Kohler, H.-P. E. and Ternes, T. A. (2010).

Biotransformation of selected iodinated X-ray contrast media and characterization of

microbial transformation pathways. Environ. Sci. Technol. 44: 4998–5007

Kormos, J. L., Schulz, M., Wagner, M. and Ternes, T. A. (2009). Multistep

approach for the structural identification of biotransformation products of iodinated

X-ray contrast media by liquid chromatography/hybrid triple quadrupole linear ion

trap mass spectrometry and 1H and

13C nuclear magnetic resonance. Anal. Chem. 81

(22): 9216–9224.

Krasner, S.W. (2009). The formation and control of emerging disinfection by-

products of health concern. Philosophical Transactions of the Royal Society a-

Mathematical Physical and Engineering Sciences 367(1904), 4077-4095.

Krasner, S.W. (1999). Chemistry of disinfection by-product formation. In:

formation & control of disinfection byproducts in drinking water. Singer, PC (ed.)

AWWA, Denver, CO

Krasner S.W., Weinberg H.S., Richardson S.D., Pastor S.J., Chinn R.,

Sclimenti M.J., Onstad G.D., Thruston A.D., Jr (2006). Occurrence of a new

generation of disinfection byproducts. Environ. Sci. Technol. 40: 7175–7185

Krasner, S.W., Croue´, J.-P., Buffle, J. and Perdue, E.M. (1996). Three

Approaches for Characterizing NOM. J. American Water Works Association, 88(6):

66–79.

Krause, W. and Schneider, P.W. (2002). Optical, Ultrasound, X-ray and

Radiopharmaceutical Imaging. In: Merbach AE, Tóth É, editors. The Chemistry of

Contrast Agents in Medical Magnetic Resonance Imaging. p 107-150

Kristiana, I., Gallard, H., Jol, C. And Croué, J-P. (2009). The formation of

halogen-specific TOX from chlorination and chloramination of natural organic matter

isolates. Water Research 43: 4177 – 4186.

Kühn W. and Sontheimer H. (1973). Einige Untersuchungen zu Bestimmung

von organischen Chlorverbindungen auf Aktivkohlen. Vom Wasser 41: 65–79.

http://www.ecs.umass.edu/eve/secure/DBPs/Occurrence/Krasner%202009%20DBP%20review.pdf

http://www.ecs.umass.edu/eve/secure/DBPs/Occurrence/Krasner%202009%20DBP%20review.pdf

123

Kumar, K. and Margerum, D.W. (1987). Kinetics and mechanism of general

acid-assisted oxidation of bromide by hypochlorite and hypochlorous acid. Inorg.

Chem. 26: 2706-2711.

Kumar, K., Day, R.A and Margerum, D.W. (1986). Atom transfer redox

kinetics: general acid assisted oxidation of iodide by chloramines and hypochlorite.

Inorg. Chem. 25(24): 4344 – 4350

Kummerer, K., Erbe, T., Gartiser, S and Brinker, L. (1998). AOX- emissions

from hospitals into municipal waste water. Chemosphere 36 (11): 2437–2445.

Legube, B. (2003). Ozonation By-products. The Handbook of Environmental

Chemistry, vol. 5 (Part G), pp. 95–116.

Leigh, G.J., Ed. (1990). Nomenclature of inorganic chemistry:

Recommendations. Blackbcll Scientific Publications. Oxford.

Li, C., Benjamin, M.M. and Korshin, G.V. (2002). The relationship between

TOX formation and spectral changes accompanying chlorination of pre-concentrated

or fractionated NOM. Water Research 36: 3265–3272.

Loffler, D., Rombke, J., Meller, M. and Ternes, T. A. (2005). Environmental

fate of pharmaceuticals in water/sediment systems. Environ. Sci. Technol. 39 (14):

5209–5218.

Luong, T. V., Peters, C. J. and Perry, R. (1982). Influence of bromide and

ammonia upon the formation of trihalomethanes under water treatment conditions.

Environ. Sci. Technol. 16: 473-479.

Magazinovic, R.S., Nicholson, B.C., Mulcahy, D.E. and Davey, D.E. (2004).

Bromide levels in natural waters: its relationship to levels of both chloride and total

dissolved solids and the implications for water treatment. Chemosphere 57, 329-335.

McArdell, C.S., Kovalova, L., Eugster, J., Hagenbuch, M., Wittmer, A. and

Siergrist, H. (2010). Elimination of pharmaceuticals from hospital wastewater in a

pilot membrane bioreactor with PAC or ozone post-treatment. Conference

proceeding: SETAC Europe, 20th

Annual meeting, 23 – 27 May, 2010, Seville, Spain

McGuire, M.J. (2006). Eight revolutions in the history of US drinking water

disinfection. Journal AWWA, 98(3): 123 – 150

Mitch, W. A. and Sedlak, D. L. (2002). Factors affecting the formation of

NDMA during chlorination. Environ. Sci. Technol., 36: 588–595.

Moran, J.E., Oktay, S.D. and Santschi, P.H. (2002). Sources of iodine and

iodine-129 in rivers. Water Resource Res. 38(8) Art no. 1149

Morris, R. D., Audet, A. M., Angelillo, I. F., Chalmers, T. C. and Mosteller, F.

(1992). Chlorination, chlorination by-products, and cancer: a meta-analysis. Am. J.

Public Health 82 (7): 955-963.

124

Morris, J.C. (1986). Aqueous chlorine in treatment of water supplies. In: Ram,

N.M., Calabrese, E.J., Christmas, R.F. (Eds.), Organic Carcinogens in Drinking

Water: Detection, Treatment and Risk Assessment. Wiley, New York, pp. 33–54.

Morris, J.C. and Isaac, R.A. (1983). A critical review of kinetic and

thermodynamic constants for aqueous chlorine-ammonia system. In: Jolleys, R.L.,

Brungs, W.A., Cotruvo, J.A., Cumming, R.B., Mattice, J.S., Jacobs, V.A. (Eds.),

Water Chlorination: Environmental Impact and Health Effects, vol. 4. Ann Arbor

Science Publishers, Michigan, pp. 49–62.

Morris, J.C. (1978). The chemistry of aqueous chlorine in relation to water

chlorination. In: Jolleys, R.L. (Ed.), Water Chlorination: Environmental Impact and

Health Effects, vol. 1. Ann Arbor Science Publishers, Michigan, pp. 21–35

Myers, O.E. and Kenedy, J.W. (1950). The kinetics of iodine-iodate isotopic

exchange reaction. J.Am. Chem. Soc. 72: 897 – 906.

Nagy, J.C., Kumar, K. And Margerum, D.E. (1988). Non-metal redox kinetics:

oxidation of iodide by hypochlorous acid and by nitrogen trichloride measured by the

pulse-accelerated flow method. Inorg. Chem. 27(16): 2773 – 2780

NEFCO (2011). Wolf Creek watershed plan – phase I. Northeast Ohio Four

County Regional Planning and Developmental Organization (NEFCO) Draft Report.

Nguyen, M., Westerhoff, P., Baker, L.,Hu, Q., Esparza-Soto, M., and

Sommerfeld, M. (2005). Characteristics and Reactivity of Algae-Produced Dissolved

Organic Carbon. Journal of Environmental Engineering: 1574 – 1782.

Nieuwenhuijsen, M. J., Toledano, M. B., Eaton, N. E., Elliott, P. and Fawell,

J. (2000). Chlorination disinfection by-products in water and their association with

adverse reproductive outcomes: a review. Occup. Environ. Med. 57: 73–85.

Nriagu, J.S. and Simmons, M.S. (1994). Oxidants in the environment. John

Wiley and Sons, New York.

Odeh, I.N., Francisco, J.S. and Margerum, D.W. (2002). New pathways for

chlorine dioxide decomposition in basic solution. Inorg. Chem. 41, 6500–6506.

Ohio Department of Natural Resources. Lake Erie Watershed. Available

online (01/07/14):

http://ohiodnr.com/Portals/13/Atlas_Maps_GIS/coastalatlas2/CH3_watershed.pdf

Oppel, J., Broll, G., Lo¨ffler, D., Meller, M., Ro¨mbke, J. and Ternes, T.A.

(2004). Leaching behaviour of pharmaceuticals in soil-testing systems: a part of an

environmental risk assessment for groundwater protection. Sci. Total Environ. 328 (1-

3): 265–273.

Oleksy-Frenzel, J., Wischnack, S. and Jekel, M. (2000). Application of ion-

chromatography for the determination of the organic-group parameters AOCl, AOBr

and AOI in water. Fresenius Journal of Analytical Chemistry 366(1):89-94.

http://ohiodnr.com/Portals/13/Atlas_Maps_GIS/coastalatlas2/CH3_watershed.pdf

125

Peck, A. (2006). Analytical methods for the determination of persistent

ingredients of personal care products in environmental matrices. Anal Bioanal. Chem.

386: 907 – 939.

Pérez, S. and Barceló, D (2007). Fate and occurrence of x-ray contrast media

in the environment. Anal Bioana Chem. 387(4): 1235 – 1246

Perez, S., Eichhorn, P., Celiz, M.D., and Aga, D.S. (2006). Structural

characterization of metabolites of the X-ray contrast agent iopromide in activated

sludge using ion trap mass spectrometry. Analytical Chemistry 78(6):1866-1874.

Pelhybridge, A.D. and Prue, J.E. (1967). Equilibria in aqueous solutions of

iodic acid. Transactions ol the Faraday Society. 63: 2019-2033.

Plewa, M. J., Wagner, E. D. and Jazwierska, P. (2004). Halonitromethane

drinking water disinfection byproducts: chemical characterization and mammalian

cell cytotoxicity and genotoxicity. Environmental Sci. & Technol. 38(1): 62-68.

Plewa, M.J., Kargalioglu, Y., Vankerk, D., Minear, R.A., and Wagner, E. D.

(2002).Mammalian cell cytotoxicity and genotoxicity analysis of drinking water

disinfection by-products. Environmental and Molecular Mutagenesis 40(2): 134-142

Post, G.B., Atherholt, T.B., and Cohn, P.D. (2011). Health and Aesthetic

aspects of drinking water. In Water quality and treatment (Editor: Edzwald, J.K.).

McGraw Hill, NY

Pourmoghaddas, H. and Stevens, A.A. (1995). Relationship between

trihalomethanes and haloacetic acids with total organic halogen during chlorination.

Wat. Res. 29: 2059–2062

Pourmoghaddas, H., Stevens, A. A., Kinman, R. N., Dressman, R. C., Moore,

L. A. and Ireland, J. C. (1993). Effect of bromide ion on formation of HAAs during

chlorination. J.sAm. Water WorksAssoc. 85: 82-87.

Putschew, A., Miehe, U., Tellez, A. S., and Jekel, M. (2007). Ozonation and

reductive deiodination of iopromide to reduce the environmental burden of iodinated

X-ray contrast media. Water Sci.Technol. 56 (11): 159–165.

Putschew, A. and Jekel, M. (2006). Iodinated X-ray contrast media. In

Organic Pollutants in the Water Cycle, Reemtsma, T.; Jekel, M., Eds.; Wiley-VCH:

Weinheim, Germany. pp 87 – 98.

Putschew, A. and Jekel, M.(2001). Iodierte R€ontgenkontrastmittel im

anthropogen beeinflussten Wasserkreislauf. Vom Wasser 97: 103–114.

Putschew, A., Schittko, S. and Jekel, M. (2001). Quantification of triiodinated

benzene derivatives and X-ray contrast media in water samples by liquid

chromatography-electrospray tandem mass spectrometry. J. Chromatogr., A 930 (1-

2): 127–134.

126

Putschew, A., Wischnack, S., and Jekel, M (2000). Occurrence of

triiodoinated X-ray contrast agents in the aquatic environment. Sci. Total Environ.

255 (1): 129–134.

Qiang, Z. And Adams, C. (2004). Determination of monochloramine

formation rate constants with stopped-flow spectrometry. Environ. Sci. Technol. 38:

1435–1444

Rebenne, LM., Gonzalez, A.C. and Olson, T.M. (1996). Aqueous chlorination

kinetics and mechanism of substituted dihydrobenzenes. Environ. Sci. Technol. 30:

2235–2242.

Reckhow, D.A. and Singer, P.C. (1984). Removal of Organic Halide

Precursors by Pre-ozonation and Alum Coagulation. Journal of AWWA 76 (4): 151-

157.

Richardson, S.D. (2011) Disinfection Byproducts: Formation and Occurrence

in Drinking Water. In: The Encyclopedia of Environmental Health. Nriagu, J.O. (ed),

pp. 110-136, Elsevier, Burlington, MA

Richardson, S. D. (2009) Water analysis: Emerging contaminants and current

issues. Anal. Chem. 81 (12): 4645–4677.

Richardson, S.D. (2003). Disinfection by-products and other emerging

contaminants in drinking water. Trends in Analytical Chemistry 22: 666–684.

Richardson, S. (1998). Drinking water disinfection by-products. In

Encyclopedia of environmental analysis and remediation. New York, NY: John Wiley

& Sons, Inc

Richardson, S. D., Fasano, F., Ellington, J.J., Crumley, F.G., Buettner,

K.M., Evans, J.J., Blount, B.C., Silva, L.K., Waite, T.J., Luther, G.W., Mckague,

A.B., Miltner, R.J., Wagner, E.D. and Plewa, M.J. (2008). Occurrence and

mammalian cell toxicity of iodinated disinfection by-products in drinking

water. Environ. Sci. Technol.42, 8330–8338

Richardson, S.D., Plewa, M.J., Wagner, E.D., Schoeny, R.,and DeMarini,

D.M. (2007). Occurrence, genotoxicity, and carcinogenicity of regulated and

emerging disinfection by-products in drinking water: A review and roadmap for

research. Mutation Research/Reviews in Mutation Research 636, 178-242

Richardson, S.D., Simmmons, J.E. and Rice, G. (2002). Disinfection by-

product: the next generation. Environ. Sci. Technol. 36 (9): 198 – 205

Roberson, J. A. (2008). The evolution of disinfection byproduct regulations:

past, present, and future. In Disinfection By-Products in Drinking Water: Occurrence,

Formation, Health Effects, and Control. Karanfil, T., Krasner, S. W., Westerhoff, P.,

Xie Y. (eds). American Chemical Society: Washington, D. C.

Roberts, J.D. and Caserio, M.C. (1968). Chimie Organique Moderne.

Ediscience, Paris.

http://www.ncbi.nlm.nih.gov/pubmed?term=Fasano%20F%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Ellington%20JJ%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Crumley%20FG%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Buettner%20KM%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Buettner%20KM%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Evans%20JJ%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Blount%20BC%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Silva%20LK%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Waite%20TJ%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Luther%20GW%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Mckague%20AB%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Mckague%20AB%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Miltner%20RJ%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Wagner%20ED%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

http://www.ncbi.nlm.nih.gov/pubmed?term=Plewa%20MJ%5BAuthor%5D&cauthor=true&cauthor_uid=19068814

127

Rook, J. J. (1974) Formation of haloforms during chlorination of natural

water. Water Treatment and Examination. 23 (3): 234-243.

Sacher, F., Lange, F. T., Brauch, H.-J., and Blankenhorn, I. (2001).

Pharmaceuticals in groundwaters: Analytical methods and results of a monitoring

program in Baden-Wu¨rttemberg, Germany. J. Chromatogr. A 938 (12): 199–210.

Schlesinger, W.H. (1997). Biogeochemistry: An analysis of global change.

Academic Press

Schulz, M., Lo¨ffler, D., Wagner, M., and Ternes, T. A. (2008).

Transformation of the X-ray contrast medium iopromide in soil and biological

wastewater treatment. Environ. Sci. Technol. 42 (19): 7207–7217.

Seitz, W., Weber, W. H., Jiang, J.-Q., Lloyd, B. J., Maier, M., Maier, D. and

Schulz, W. (2006a). Monitoring of iodinated X-ray contrast media in surface water.

Chemosphere 64 (8): 1318–1324.

Seitz, W., Jiang, J.-Q., Weber, W. H., Lloyd, B. J., Maier, M. and Maier, D

(2006b). Removal of iodinated X-ray contrast media during drinking water treatment.

Environ Chem. 3 (1): 35–39.

Sharma, V.K. (2008). Oxidative transformations of environmental

pharmaceuticals by Cl2, ClO2, O3 and Fe(VI): Kinetics assessment. Chemosphere 73:

1379 – 1386.

Simmons, J.E., Richardson, S.D., Speth, T.F., Miltner, R.J., Rice, G., Schenck,

K.M., Hunter III, E.S., and Teuschler, L.K. (2002). Development of a research

strategy for integrated technology-based toxicological and chemical evaluation of

complex mixtures of drinking water disinfection byproducts. Environ. Health

Perspect. 110: 1013–1024.

Singer, P.C. (1994). Control of disinfection by-products in drinking water.

Journal of Environmental Engineering 120: 727 – 744.

Singer, P.C., Weinberg, H.S., Brophy, K., Liang, L., Roberts, M., Grisstede, I.,

Krasner, S., Baribeau, H., Arora, H. and Najm, H. (2002). Relative dominance of

haloacetic acids and trihalomethanes in treated drinking water. AWWA and

AWWARF, Denver, CO.

Singer, P.C. and Reckhow, D.A. (1999). Chemical oxidation. In Water quality

and treatment. Letterman R.D. technical editor, AWWA, McGraw-Hill, New York,

NY

Speck, U. and Hübner-Steiner, H. (1999). Pharmakologie und Toxikologie,

Eds. Oberdisse, Hackenthal, Kuschinsky, 2. Auflage, Springer-Verlag Berlin,

Heidelberg, NY

Sprehe, M., Geissen, S. U. and Vogelpohl, A. (2001). Photochemical

oxidation of iodized X-ray contrast media in hospital wastewater. Wat. Sci. Technol.

44 (5): 317–323.

128

Steger-Hartmann, T., Länge, R. and Schweinfurth, H. (2000). Iodinated x-ray

contrast media in the aquatic environment – fate and effects. Symposia paper

presented before the Division of Environmental Chemistry, ACS, San Francisco, CA.

March 26 – 30, 2000).

Steger-Hartmann, T., Länge, R. and Schweinfurth, H. (1999). Environmental

risk assessment for the widely used iodinated x-ray contrast agent iopromide

(Ultravist). Ecotox. Environ. Saf. 42: 274 – 281

Suberkroopp, K., and Klug, M.J. (1976). Fungi and bacteria associated with

leaves during processing in a woodland stream. Ecology 57:707-719

Świetlik, J. and Sikorska, E. (2005). Characterization of natural organic matter

fractions by high pressure size-exclusion chromatography, specific UV absorbance

and total luminescence spectroscopy. Polish Journal of Environmental Studies, 15 (1):

145-153

Symons, J.M., Xia, R., Speitel, G.E., Diehl, A.C., Hwang, C.J., Krasner, S.W.

and Barrett, S.E. (1998). Factors affecting disinfection by-product formation during

chloramination. AWWA Research Foundation, USA.

Symons, J. M., Krasner, S. W., Simms, L. A. and Sclimenti, M. (1993).

Measurement of THM and precursor concentrations revisited: the effect of bromide

ion. J. Am. Water Works Assoc. 85: 51-62.

Ternes, T. A., Bonerz, M., Hermann, N., Teiser, B., and Andersen, H.

R.(2007). Irrigation of treated wastewater in Braunschweig, Germany: An option to

remove pharmaceuticals and musk fragrances. Chemosphere 66 (5): 894–904.

Ternes, T. A., Stu¨ber, J., Herrman, N., McDowell, D., Ried, A., Kampmann,

M. and Teiser, B. (2003). Ozonation: a tool for removal of pharmaceuticals, contrast

media and musk fragrances from wastewater. Water Res. 37 (8): 1976–1982.

Ternes, T. A. and Hirsch, R. (2000). Occurrence and behavior of X-ray

contrast media in sewage facilities and the aquatic environment. Environ. Sci.

Technol. 34 (13): 2741–2748.

Ternes, T. A. (1998). Occurrence of drugs in German sewage treatment plants

and rivers. Water Res. 32 (11): 3245–3260.

Thomas, T.R., Pence, D.T. and Hasty R.A. (1980). The disproportionation of

hypoiodous acid. Journal of Inorganic & Nuclear Chemistry. 42: 183-186.

Truesdale. V.W. ( 1997). Kinetics of disproportionation of hypoiodous acid at

high pH, with an extrapolation to rainwater. Journal of the Chemical Society, Faraday

Transactions, 93(10): 1909-1914.

Ueno, H., Moto, T., Sayato, Y. and Nakamuro, K. (1996). Disinfection by-

products in the chlorination of organic nitrogen compounds: By-products from

kynurenine. Chemosphere 33 (8): 1425–1433.

129

United State Environmental Protection Agency (US EPA), (2013). Enhanced

coagulation and enhanced precipitative softening guidance manual. BiblioGov, USA.

United State Environmental Protection Agency (US EPA), (2006). National

Primary Drinking Water Regulations: Stage 2 Disinfectants and Disinfection By-

products Rule.

http://water.epa.gov/lawsregs/rulesregs/sdwa/stage2/regulations.cfm#prepub

United State Environmental Protection Agency (US EPA), (2005). Rule

factsheet: stage 2 disinfection and disinfection by-product. EPA 815-F-05-003

United State Environmental Protection Agency (US EPA), (2000). The history

of drinking water treatment. EPA-816-F-00-006.

http://www.epa.gov/safewater/consumer/pdf/hist.pdf Accessed on 01/28/2014

United State Environmental Protection Agency (US EPA), (1999).

Alternative Disinfectants and Oxidants Guidance Manual. EPA 815-R-99-014

USEPA (1997). Research plan for microbial pathogens and disinfection by-

products in Drinking Water. EPA-600-R-97-122

Urbansky, E.T., Cooper, B.T. and Margerum D.W. (1997). Disproportionation

kinetics of hypoiodous acid as catalyzed and suppressed by acetic acid-acetate buffer.

Inorganic Chemistry 36: 1338-1344.

Valero, F. And Arbós, R. (2010). Desalination of brackish river water using

Electrodialysis Reversal (EDR): Control of the THMs formation in the Barcelona (NE

Spain) area. Desalination 253, 170-174.

van Benschoten J. E., Lin W. and Knocke W. R. (1992). Kinetic modeling of

manganese(II) oxidation by chlorine dioxide and potassium permanganate. Env Sci

and Technol. 26: 1326-1333.

van Hoof, F. (1992). Identifying and characterizing effects of disinfection by-

products, in E.A. Bryant, G.P. Fulton and G.C. Budd (Eds.), Disinfection Alternatives

for Safe Drinking Water, van Nostrand Reinhold, NY, USA.

Vikesland P. J., Ozekin K. and Valentine R. L. (2000). Monochloramine decay

in model and distribution system waters. Wat Res. 35 (7): 1766 – 1776.

Vikesland P. J., Ozekin K. and Valentine R. L. (1998). Effect of natural

organic matter on monochloramine decomposition: pathway elucidation through the

use of mass

and redox balances. Environ. Sci. Tech. 32(10): 1409–1416.

von Gunten, U. (2003). Ozonation of drinking water: Part I. Disinfection and

by-product formation in presence of bromide, iodide or chlorine. Water Res. 37:

1443-1467.

Waller, K., Swan, S., Windham, G. C. and Fenster, L. (2001). Influence of

exposure assessment methods on risk estimates in an epidemiologic study of

http://water.epa.gov/lawsregs/rulesregs/sdwa/stage2/regulations.cfm#prepub

http://www.epa.gov/safewater/consumer/pdf/hist.pdf%20%20Accessed%20on%2001/28/2014

130

trihalomethane exposure and spontaneous abortion. J. Expo. Anal. Environ.

Epidemiol. 11: 522–531.

Weinberg H, (1999). Disinfection by-products in drinking water: the analytical

challenge. Analytical Chemistry 71(23): 801- 808.

Weinberg H.S., Krasner S.W., Richardson S.D., Thruston A.D., Jr (2002). The

occurrence of disinfection by-products (DBPs) of health concern in drinking water:

results of a nationwide DBP occurrence study. EPA/600/R-02/068.Athens, GA: US

EPA

Weissbrodt, D.,Kovalova, L., Ort, C., Pazhepurackel, V., Moser, R., Jollender,

J., Siegrist, H. and McArdell, C.S. (2009). Mass flows of x-ray contrast media and

cytostatics in hospital wastewater. Environ. Sci. Tech 255(1): 129 – 134

Westerhoff, P., Chao, P. and Mash, H. (2004). Reactivity of natural organic

matter with aqueous chlorine and bromine. Water Res. 38: 1502-1513.

Wolfe, R.L., Ward, N.R., Olson, B.H. (1984). Inorganic chloramines as water

disinfectants: a review. J. Am. Water Works Assoc. 76: 74–88

Wilcox, P. and Williamson, S. (1986). Mutagenic activity of concentrated

drinking water samples. Environ Health Perspect 69:141–149.

Wong, G.T.F. (1991). The marine geochemistry of iodine. Rev Aqua Sci. 4(1):

45-73.

Wren, J.C, Paquette. J., Sunder, S. and Ford, B.L. (1986). Iodine chemistry in

the +1 oxidation state II. A Raman and UV/visible spectroscopic study of the

disproportionation of hypoiodite in basic solutions. Canadian Journal of Chemistry,

64(12): 2284-2296.

Wu, W.W., Chadik, P.A., Davis, W.M., Delfino, J.J. and Powell, D.H. (2000).

The effect of structural characteristics of humic substances on disinfection by-product

formation in chlorination. In: Barrett, S.E., Krasner, S.W., Amy, G.L. (Eds.), Natural

Organic Matter and Disinfection By-Products: Characterisation and Control in

Drinking Water. American Chemical Society, New York, pp. 109–121.

Ziegler M, Schulze-Karal C, Steiof M, Rüden H. Reduzierung der AOX-

Fracht von Krankenhäusern durch Minimierung des Eintrags iodorganischer

Röntgenkontrastmittel _Abstract in English.. Korrespondenz Abwasser 1997;

44:1404 – 1408.

Zwiener, C. And Richardson, S.D. (2005). Analysis of disinfection by-

products in drinking water by LC-MS and related MS techniques. Trac-trends in

Analytical Chemistry 24(7): 613 – 621

Documents

total organic halogen formation in the presence of iopamidol and chlorinated oxidants with and