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This is to certify that the
thesis entitled

FATE OF PHARMACEUTICALS IN DRINKING WATER
UTILITIES

presented by

REBECCA HULLMAN

has been accepted towards fulfillment
of the requirements for the

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FATE OF PHARMACEUTICALS IN DRINKING WATER UTILITIES
By

Rebecca Hullman

A THESIS

Submitted to
Michigan State University
in partial fulfillment of the requirements
for the degree of

MASTER OF SCIENCE

Environmental Engineering

2009

ABSTRACT

FATE OF PHARMACEUTICALS IN DRINKING WATER UTILITIES
By
Rebecca Hullman

Studies suggest that drinking water treatment is effective at decreasing the amount
of pharmaceuticals present in source water. Primary treatment processes such as
coagulation and sedimentation are not effective on their own, while a combination of
filtration and disinfection processes is effective at reducing the concentrations of most
pharmaceuticals. The type of oxidant used also determines the degree at which a
pharmaceutical is removed from source water. Ozone or the combination of ultraviolet
(UV) light with hydrogen peroxide seem to be the most effective oxidation processes.
Chlorine oxidation is known to produce chlorinated byproducts when it reacts with
organic compounds. In bench-scale studies, the reaction of the pharmaceutical,
acetaminophen with free chlorine was studied. Acetaminophen reacts with free chlorine
to produce the byproduct 1,4-benzoquinone. Results indicate that acetaminophen is most
reactive with free chlorine at pH 9.0 and least reactive at pH 6.0. As pH increased,
degradation of acetaminophen also increased. The formation of 1,4-benzoquinone was
also affected by pH and reached a maximum of 68.7% of the initial acetaminophen
concentration when the pH was at 6.0, the molar ratio at 1,275, and after a contact time of
30 minutes. At all pH values the rate of degradation of acetaminophen was slowest at a

molar ratio of about 100, and the highest at a molar ratio of about 10,000.

COPyright by
REBECCA HULLMAN

2009

ACKNOWLEDGEMENTS

This research would not have been possible without the contribution of many
people at Michigan State University. First, I would like to express my gratitude to Irene
Xagoraraki who suggested the topic for my research, and who guided me throughout the
entire process. She provided me with excellent opportunities for growth as a researcher
and an environmental professional. Her support and encouragement were invaluable to
the completion of this research.

I acknowledge Tom Voice for his advice on the direction of the research and his
technical expertise. He continually challenged me to assess the real purpose behind this
research and to work towards a significant scientific contribution to the field. I am also
grateful towards Phani Mantha who patiently helped me set up a mathematical model. I
gratefiilly thank Hui Li and Wenlu Song for their assistance with instrumentation and

method development. Finally, I thank my family for their encouragement over the years.

iv

TABLE OF CONTENTS

LIST OF TABLES .................................................................................... vi
LIST OF FIGURES ................................................................................. vii
INTRODUCTION .................................................................................... 1
CHAPTER 1

LITERATURE REVIEW AND SYNTHESIS ................................................... 2
Summary ............................................................................................... 2
Occurrence and Sources .............................................................................. 3
Fate of Pharmaceuticals in Drinking Water Utilities ............................................. 9
Oxidation of Pharmaceuticals ..................................................................... 21
Conclusions and Recommendations .............................................................. 38
CHAPTER 2

LABORATORY EXPERIMENTS ............................................................... 41
Summary ............................................................................................. 41
Primary Experiments ............................................................................... 42
Additional Experiments and Continuing Research ............................................. 64
Conclusions .......................................................................................... 76
APPENDIX ........................................................................................... 78
REFERENCES ...................................................................................... 79

LIST OF TABLES

Table 1. Overall removal of pharmaceuticals during full-scale drinking water
treatment ................................................................................... 13

Table 2. Removal of pharmaceuticals during individual process in full-scale drinking
water utilities .............................................................................. 14

Table 3. Comparison of the reactivity of different oxidants for selected
pharmaceuticals ........................................................................... 28

Table 4. Effect of pH during chlorination of pharmaceuticals by sodium hypochlorite..32

Table 5. Byproduct formation during pharmaceutical oxidation ............................ 36
Table 6. Acetaminophen occurrence in the environment ..................................... 44
Table 7. Expected molar ratios of chlorine to acetaminophen ............................... 46

Table 8. Summary of results and comparison with literature data involving reaction of

acetaminophen with sodium hypochlorite ............................................ 62
Table 9. Conditions tested in this study ......................................................... 65
Table 10. Acetaminophen and chlorine species composition at each pH tested ............ 70

vi

LIST OF FIGURES

Figure 1. Source of pharmaceuticals in water .................................................... 7
Figure 2. Degradation of acetaminophen and formation of 1,4—benzoquinone (pH 6.0,

MR = 99) .................................................................................. 51
Figure 3. Degradation of acetaminophen and formation of 1,4-benzoquinone (pH 6.0,

MR = 1,275) .............................................................................. 51
Figure 4. Degradation of acetaminophen and formation of 1,4-benzoquinone (pH 6.0,

MR = 10,386) ............................................................................ 52
Figure 5. Effect of pH and contact time on acetaminophen (MR = 106 :t 6) ............... 53
Figure 6.. Effect of pH and contact time on acetaminophen (MR = 1,417 i 285).... ......53
Figure 7. Effect of pH and contact time on acetaminophen (MR = 9,789 i 1,430). . ......54
Figure 8. Effect of pH and contact time on 1,4-benzoquinone formation

(MR = 106 i 6) .......................................................................... 57
Figure 9. Effect of pH and contact time on 1,4-benzoquinone formation

(MR= 1,417d: 285) ...................................................................... 58
Figure 10. Effect of pH and contact time on 1,4-benzoquinone formation

(MR = 9,789 i 1,430) ................................................................... 59
Figure 11. Effect of pH on acetaminophen degradation and 1,4-benzoquinone formation

(MR = 107 :l: 5) ........................................................................ 69

Figure 12. Effect of initial acetaminophen concentration on acetaminophen degradation

Figure 13

and 1,4-benzoquinone formation (MR = 121 :h 18, pH 7.5) ..................... 72

. Effect of initial acetaminophen concentration on acetaminophen degradation

and 1,4-benzoquinone formation (MR = 1,159 :I: 164, pH 6.0) .................. 73

Figure 14. Effect molar ratio on acetaminophen degradation and 1,4-benzoquinone

formation ................................................................................ 75

vii

INTRODUCTION

“We forget that the water cycle and the life cycle are one.”

- Jacques Cousteau

The research presented below addresses the detection of pharmaceuticals in
environmental waters and the fate of these pharmaceuticals during drinking water
treatment. The results of full- and bench-scale studies are presented to evaluate the
effectiveness of primary treatment (i.e. coagulation, flocculation, sedimentation),
filtration, and disinfection on pharmaceutical removal from source water. Disinfection,
which has the potential to produce oxidation byproducts, is evaluated in more detail than
primary treatment and filtration. The fate of pharmaceuticals in drinking water utilities is
affected by the individual pharmaceutical’s chemical structure and properties, the type
and combination of treatment process used. Factors affecting oxidation of
pharmaceuticals include the water pH, contact time, oxidant dosage, initial
pharmaceutical concentration, water matrix, and the individual pharmaceutical’s
chemical structure and properties. Experiments on the chlorination of acetaminophen
provide a case study on how a pharmaceutical might be transformed during disinfection
of drinking water. Current risk assessments suggest that risk of human exposure to
pharmaceuticals via drinking water is very low, but more research is needed on mixture
toxicity of low levels of pharmaceuticals, and the effects of pharmaceuticals on microbial
systems. Overall, the research presented below gives the reader an idea of the complexity

and uncertainty of the issue of pharmaceuticals in the environment.

CHAPTER 1

LITERATURE REVIEW AND SYNTHESIS

Summary

Pharmaceuticals have recently been identified as emerging contaminants in
environmental waters. The presence of pharmaceuticals in wastewater effluent and
source water for drinking water utilities may have implications on human and ecological
health. Water and wastewater treatment processes are capable of transforming
pharmaceuticals, but each compound responds differently to treatment, making it difficult
to predict the fate of pharmaceuticals as a whole. Drinking water oxidants such as
chlorine, ozone, and advanced UV oxidation have been shown to degrade and transform a
wide variety of pharmaceuticals. A review of literature suggests that conventional water
treatment may effectively remove the majority of tested pharmaceuticals. Certain
compounds, such as carbamazepine, appear to undergo treatment without transformation,
thus requiring further study.

Chlorine and ozone are shown to react with various pharmaceuticals, and often
times oxidation results in the formation of byproducts that may or may not be harmful to
humans and the environment. In most cases, UV irradiation does not degrade
pharmaceuticals unless an advanced oxidation process is applied. Ozone appears to be
the most effective at oxidizing different classes of pharmaceuticals present in waters.
Predictive quantitative structure activity relationship (QSAR) models may be key to
determining which pharmaceuticals persist during drinking water treatment. However,
the current state of knowledge does not provide clear answers regarding potential human

risk of exposure to pharmaceuticals via water consumption.

Occurrence and Sources

Pharmaceutical contamination of aquatic systems is an emerging issue in
environmental science and engineering. Pharmaceuticals are biologically active
compounds that may produce adverse effects in humans, animals, and plants. Increasing
numbers of studies reveal that pharmaceuticals are found in natural waters, wastewater,
and drinking water. When water undergoes treatment, such as oxidation, pharmaceuticals
may be transformed into compounds that are less or more harmful to humans and the
environment than the parent compound. The oxidation rates and oxidation byproducts of
pharmaceuticals will largely depend on their concentrations in the environmental waters.
In the United States there is more consumption of new pharmaceuticals and high-strength
or long—acting formulations than in other countries [Danzon and Furukawa, 2008]. This
suggests that more persistent and/or newer man-made compounds may occur in US.

waters.

0. Occurrence of Pharmaceuticals in Wastewater, Drinking Water and Surface Waters
The presence of pharmaceuticals in raw and treated wastewater has been well
documented with concentrations averaging from less than 10 ug/L in finished wastewater
to greater than 100 ug/L in raw wastewater to [Batt et al., 2006; Boyd et al., 2003; Brun
et al., 2006; Buser et al., 1999; Buser et al., 1998; Carballa et al., 2004; Castiglioni et al. ,
2006; Castiglioni et al., 2005; Clara et al., 2005; Ellis, 2006; Ferrari et al., 2003; Giger et
al., 2003; Gobel et al., 2005; Golet et al., 2003; Golet et al., 2002; Gomez et al., 2007;

Gros et al., 2006; Gross et al., 2004; Han et al., 2006; Heberer, 2002; J elicic and Ahel,

2003; Jones et al., 2005; Karthikeyan and Meyer, 2006; Kummerer et al., 1997; Lee et
al., 2003; Lindberg et al., 2005; Lindqvist et al., 2002; McArdell et al., 2003; Metcalfe et
al., 2003; Petrovic et al.,2003; Quintana et al., 2005; Radjenovic et al., 2007; Renew and
Huang, 2004; Santos et al., 2007; Sedlak and Pinkston, 2001; Tauxe-Wuersch et al.,
2005; Temes, 1998; Thomas and Foster, 2004; Tixier et al., 2003; Vieno et al., 2005;
Wennmalm and Gunnarsson, 2005; Yang and Carlson, 2004; Zuccato et al., 2006].
Examples of pharmaceuticals that have been detected in wastewater are common
prescription and veterinary drugs such as beta-blockers (e. g. metoprolol, propranolol),
analgesics (e. g. ibuprofen, naproxen), and antibiotics (e. g. erythromycin, trimethoprim,
ciprofloxacin, tetracylcline, clindomycin, sulfonamides, tetracycline, fluoroquinolone,
macrolides, and trimethoprim). For example, Santos et al., (2007) detected ibuprofen in
wastewater influent and effluent at concentrations of 12.1 — 373 and 0.78 — 48.2 rig/L,
respectively. Naproxen was also detected at 1.1 — 27.4 pg/L in wastewater influent and
0.22 — 4 .3 rig/L in effluent . Gomez et al. (2007) detected acetaminophen, codeine,
diclofenac, and ibuprofen in wastewater influent at mean concentrations of 134, 5.2, 1.5,
and 84 rig/L, and in wastewater effluent at mean concentrations of 0.22, 3.7, 0.9, and 7.1
rig/L, respectively.

Although most of the published literature focuses on occurrence in sewage
effluents, pharmaceuticals have also been detected in raw and treated drinking water,
mostly at levels less than 1 ug/L [Boyd et al., 2003; Heberer, 2002; Loraine and
Pettigrove, 2006; M011 et al., 2001; Perret et al., 2006; Petrovic et al., 2003; Rodriguez-
Mozaz et al., 2004; Stackelberg et al., 2004; Stolker et al., 2004; Temes etal., 2002;

Vieno et al., 2005; Zuhlke et al., 2004]. For example, Loraine and Pettigrove (2006)

detected ibuprofen in finished drinking water at a mean concentration of 0.93 rig/L. In
Italy, three sulfonamide antibiotics were detected in store-bought mineral waters at
concentrations ranging from 0.009 to 0.080 rig/L [Perret et al., 2006]. In 2004, the US
Geological Survey conducted a study on the fate of 106 contaminants throughout a
conventional drinking water treatment plant [Stackelberg et al., 2004]. Pharmaceuticals
present in raw water samples included carbamazepine, trimethoprim, erythromycin-HZO,
acetaminophen, codeine, and sulfamethoxazole. Carbamazepine was detected at a
maximum concentration of 0.258 rig/L in finished water samples [Stackelberg et al. ,
2004]

The presence of pharmaceuticals has been confirmed in surface waters (rivers and
lakes) as well, with concentrations generally less than 1 rig/L [Abuin et al., 2006; Batt et
al., 2006; Bound and Voulvoulis, 2006; Boyd etal., 2005; Boyd et al., 2003; Buser et al.,
1999; Buser et al., 1998; Calamari et al., 2003; Cha et al., 2006; Giger et al., 2003; Gros
et al., 2006; Hirsch et al., 1999; Hua et al., 2006; Kolpin et al., 2002; Kosjek et al., 2005;
McArdell et al., 2003; Perret et al., 2006; Rabiet et al., 2006; Rodriguez-Mozaz et al.,
2004; Stackelberg et al., 2004; Stolker et al., 2004; Stumpf et al., 1999; Temes, 1998;
Vanderford et al., 2003; Wennmalm and Gunnarsson, 2005; Wiegel et al., 2004; Yang
and Carlson, 2004; Zuccato et al., 2006]. For example, Bound and Voulvoulis (2006)
reported concentrations of ibuprofen and acetaminophen up to 3 and 0.56 pg/L,
respectively. Rabiet et al. (2006) detected acetaminophen, carbamazepine, and
diclofenac in surface water at concentrations of 0.011 — 0.072, 0.024 —- 0.056, and 0.001 —
0.033, respectively. The US. Geological Survey conducted the first major study of

pharmaceuticals and personal care products in surface waters during 1999 and 2000

[Kolpin et al., 2002]. The study focused on locations likely to be contaminated by
wastewater effluent. Kolpin et al. (2002) detected pharmaceuticals, hormones, and other
organic contaminants in 80% of 139 streams sampled. The most frequently detected
pharmaceuticals by Kolpin et al. in US. streams included trimethoprim, acetaminophen,
erythromycin, estriol, lincomycin, and sufamethoxazole at detection frequencies of 27.4,
23.8, 21.5, 21.4, 19.2 and 19.0% and at median concentrations of 0.013, 0.11, 0.1, 0.019,
0.06, and 0.07 rig/L, respectively. Other pharmaceuticals that have been found in surface

waters are naproxen, clofibric acid, ketoprofen and clarithromycin.

b. Sources of Pharmaceuticals

There are several proposed pathways responsible for pharmaceutical
contamination of water. Figure 1 shows possible pathways for pharmaceuticals to enter
the water supply. Pharmaceuticals taken by humans in daily life are partially metabolized
and excreted into wastewater. If not transformed during treatment, a portion of these
contaminants can be released in wastewater treatment plant effluents and discharged in
receiving water bodies.

Pharmaceuticals will either be transformed or remain unchanged during treatment
by wastewater treatment processes, then released into surface water. Pharmaceuticals,
originating from wastewater or manure, may be present in surface water or ground water
used as a source for raw drinking water. Once inside the drinking water treatment
facility, a pharmaceutical may again be transformed during treatment or it may remain
unchanged. Upon distribution of treated drinking water, there is a possibility of human

exposure to unchanged or transformed pharmaceuticals.

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Animal production is another point source responsible for the wide dissemination
of pharmaceuticals in water, particularly from industrialized animal agriculture, or
Concentrated Animal Feeding Operations (CAFOs). Industrialized animal agriculture is
used to provide livestock for human consumption, typically with animals living in dense
populations. Antibiotics and hormones are commonly administered to livestock for
disease control and added into feed at sub—therapeutic levels to improve the feeding
efficiency, growth rate, and animal health. About 70% (12,600 tons) of antibiotics
produced for medical and agricultural purposes in the US. are used for non-therapeutic
treatment of livestock [Mellon et al., 2001]. As much as 75% of the antibiotics used for
animal growth are excreted back into manure [Campagnolo et al., 2002], of which the
US. produces 133 million tons of dry weight per year [Burkholder et al., 2007]. The
most common growth promoters of livestock are tetracycline, chlorotetracycline, and
bacitracen [J indal et al., 2006; Mackie et al., 2006]. Other antibiotics commonly used in
animal agriculture include macrolide, sulfonamide, and B-lactam antibiotics [Huang et
al., 2001]. All of the antibiotics listed above are also used in human therapy, which
instigated a ban on growth promoting antibiotics by the European Union in order to
preserve the effectiveness of human antibiotics [Union of Concerned Scientists, 2006].
After a period of storage manure is often land-applied as plant fertilizer, and as
consequence the pharmaceuticals are discharged to surface and ground water via surface
runoff and percolation. Studies indicate application of swine manure to crops results in
temporary increases of antibiotic resistant bacteria (Sengelov et al., 2003; Halling-
Sorensen et al., 2005). Other potential sources of pharmaceuticals include biosolid

application sites and landfill sites.

c. Conclusions

The presence of pharmaceuticals in wastewater, drinking water and surface water
is well documented. Commonly detected pharmaceuticals include ibuprofen,
acetaminophen, carbamazepine, and various antibiotics. Primary sources of
pharmaceuticals in surface water are likely from human usage and land application of
animal manure. When pharmaceuticals present in surface water are not removed by
drinking water treatment, human exposure to very low levels of one or more

pharmaceutically-active compounds may occur.

Fate of Pharmaceuticals in Drinking Water Utilities

The fate of pharmaceuticals during water treatment depends on the compound’s
physicochemical properties and the types and sequence of processes used in the treatment
plant. Physical processes such as settling, flocculation, and filtration show varying
removal efficiencies of pharmaceuticals (Gobel et al., 2007; Nakada et al., 2007;
Stackelberg et al., 2007; Peng et al., 2006; Clara et al., 2005; Temes et al., 2002).
Disinfection processes, such as chlorine oxidation, are expected to degrade or transform
pharmaceuticals since a reactive oxidative substance is added to the water. Most
pharmaceuticals are organic substances which are expected to be highly susceptible to
oxidation. Full-scale studies indicate that an appropriate combination of processes will
effectively reduce the amount of pharmaceuticals leaving a drinking water treatment
plant [Nakada et al., 2007; Stackelberg et al., 2007; Boyd et al., 2003; Temes et al.,

2002}

(1. Removal of Pharmaceuticals during F ull-Scale Drinking Water Treatment

Table 1 and Table 2 summarize removal of pharmaceuticals during full-scale
drinking water treatment. Table 1 provides overall removal, while Table 2 provides
removal during individual treatment processes. Pharmaceuticals that are not detected
after treatment are often referred to as ‘removed’ by the treatment.

Few studies have looked at removal of pharmaceuticals during full-scale drinking
water treatment. Full-scale studies are valuable because they provide information
regarding the effectiveness of individual processes on pharmaceutical removal, as well as
concentrations in source and finished water. Stackelberg et al. (2007) sampled water
from a drinking water treatment plant (DWTP) in the United States that utilized
screening, clarification, primary chlorine disinfection, sand/GAC filtration, and
secondary chlorine disinfection. The DWTP treated 235 million liters per day. The
degree of removal by each process varied for each organic compound investigated.
Forty-five organic compounds were detected in the source water, including
pharmaceuticals, flame retardants, plasticizers, pesticides, and others. Overall, GAC
filtration resulted in an average percent removal of 53% for all compounds, disinfection
resulted in 32% removal, and clarification resulted in 15% removal. Only the
pharmaceuticals acetaminophen, carbamazepine, erythromycin, and sulfamethoxazole
were detected in at least 25% of source water samples. After primary disinfection,
sulfamethoxazole and acetaminophen were not detected, while carbamazepine and
erythromycin were removed by 20% and 92%, respectively. The initial concentrations
before primary disinfection were 30, 15, 191, and 10 ng/L, respectively. The

combination of clarification, chlorine disinfection, and sand/GAC filtration was effective

10

at reducing the studied pharmaceuticals. These results were similar to a previous study
conducted by Stackelberg et al. (2004) at the same DWTP. The utility was located in a
very urbanized area, with at least 50 wastewater treatment plants discharging into streams
or tributaries supplying source water to the DWTP. In the original study, the treatment
process train followed screening, the addition of powdered activated carbon (PAC),
coagulation, primary disinfection, flocculation, sedimentation, sand/GAC filtration, and
secondary disinfection. Several changes were made to the treatment process before the
second study, including discontinuing the addition of PAC, adding microsand to the
clarification process, and reversing the order of clarification and primary disinfection.
Similar to the second study, carbamazepine was the only investigated pharmaceutical that
was not entirely removed by conventional drinking water treatment.

Boyd et al. (2003) took samples at various stages of treatment from two different
drinking water treatment plants. One plant in Louisiana, USA used PAC addition,
coagulation, flocculation, sedimentation, chlorination, filtration, and a final
chloramination. Only naproxen was detected at 68 ng/L, but was completely removed
after the combination of chlorination, filtration, and chloramination. The second plant
was located in Windsor, Ontario and used ozonation, coagulation, flocculation,
sedimentation, filtration, and a final chlorination. Clofibric acid and naproxen were
detected in the influent at 103 and 63 ng/L , respectively. Neither of the compounds were
detected in the final effluent.

Temes et al. (2002) also monitored selected pharmaceuticals during drinking
water treatment. Samples were obtained from two different DWTPs. One plant used pre-

ozonation, flocculation, main ozonation, multiple-layer filtration, and GAC filtration for

11

the treatment train, while the other plant used sedimentation, flocculation, GAC filtration,
underground passage, bank filtration, and slow sand filtration. The study found the
combination of ozonation and GAC filtration effective at reducing carbamazepine,
diclofenac, and clofibric acid, with all three compounds removed by greater than 90% in

the first treatment plant. The initial concentration of these compounds ranged from 10 to

180 ng/L.

12

 

 

TABLE 1. Overall Removal of Pharmaceuticals during Full—Scale Drinking Water Treatment

 

 

 

 

 

 

 

 

 

 

 

 

 

Effluent
Influent . Overall .
Concentration Concentrati Removal Treatment Process Train
Pharmaceutical (n I/L on (0,) Reference
5 ) (ng/L) m
Acetamino hen 15a 033 98 Screening, clarification. chlorination, sand/C AC filtration. secondary Stackelberg et al. 20079
p chlorination 7
1913 293 85 Screening, clarification, chlorination, sand/GAC filtration, secondary Stackelberg et al. 20079
ab 8], 36.73 chlorination, St k lb ’ t 1. 2 4
200 126-5 Screening, PAC and sulfuric acid addition, coagulation, chlorination, ac e erg e a 7 00
. . flocculation. sedimentation GAC filtration secondary chlorination
> ' ’ , . ’ . T ‘ t 1., 2 2
Carbamazepine 80 BDL 99 C DWTPI: Pre-ozonation, flocculation, main ozonation, layer and GAC emes e a 00
8O — 180b BDL 7 30b 83 , 100 filtration
DWTPZ: Sedimentation, flocculation, GAC, underground passage, bank
and sand filtration
80b BDL 100‘: DWTP]: Pre—ozonation, flocculation, main ozonation, layer and GAC Temes et al., 2002
Bezafibrate ND filtration . . .
DWTP2: Sedimentation, flocculation, GAC, underground passage, bank
and sand filtration
b BDL C DWTPI: Pre—ozonation. flocculation, main ozonation, layer and GAC Temes et al., 2002
~10 100
Clofibric Acid b BDL 00° filtrauon . . .
“0’15 1 DWTPZ: Sedimentation, flocculation, GAC, underground passage, bank
and sand filtration
~35b BDL 1000 DWTPI: Pre—ozonation, flocculation, main ozonation, layer and GAC Temes et al., 2002
- BDL C filtration
Drclofenac ~65 100 DWTPZ: Sedimentation, flocculation, GAC, underground passage, bank
and sand filtration
_ a ND 100 Screening, clarification, chlorination, sand/GAC filtration, seconda C
Erythromycul 10 hlorination ry SiaCkSIberg Ct al.7 2007
63-65 ND 100 DWTPI: PAC addition, coagulation/flocculahon/sedimentation,
N chlorination, filtration, chloramination Boyd et al., 2003
aproxe“ 64 ND 100 DWTPZ: Ozonation, o ' v e ' v 1' . ~ filtration,
chlorination
a 100 Screenin , clarification, chlorination, sand/GAC filtrat'on, . d c
Sumamethoxazole 30 ND Chlorinafon ‘ mo“ ary Stackelberg et al., 2007

 

 

 

 

 

 

 

 

Note: ND: pharmaceutical not detected in sample; NM: not measured; NA: not available; BDL: below detection limits.

8 Value represents average.
b Approximated from figure.
C Removal calculated by authors.

 

 

 

 

TABLE 2. Removal of Pharmaceuticals during Individual Processes in Full—Scale Drinking Water Utilities

 

 

 

 

 

 

 

 

 

 

 

Primary Treatment
Influent _ . .
(Coagulation. Filtration Disinfection
_ Concentration _ . ‘
Pharmaceutical ( flocculation, settlmg) Reference
n0/L)
a ng/L % Removal rig/L % Type ng/L "/6 Oxidant
Removal Removal
b 4 r b \
Acetaminophen 15a 6" 60 Id 98 Sand/GA ND‘l ' 100 N300 Stackelberg et al., 2007L
C .
b . . .
191a 1868 3b 4a 97 Sand/GA 149d . 20 b Naoa Stackelberg et al., 2007L
C ‘ NaOCl
~ b ’7 b '. l 1 . ,
Carbamazepine 200?Mi 82.3 “1 59.98 106 0d GAC 126.5“ . 0db o StaCkelberg mil-~20“
a d , zone T 7 9
C - ~ 1 ) emes et al., -00_
arbamazepme 80—180d 100 w 0 to 14b >90b GAC NA : >99L
C b ' I
ar amazepme 180d 0 _
d
10
Bezafibrate 80d 73d 8.8b BDL >95b GAC NM . NM Ozone Temes et al., 2002
d b
Clofibric Acid 9 - 10d 10d 0b 5 50 GAC NA i 776 Ozone Temes Ct 31'» 2002
Diclofenac 35 _ 65d 60d 77b BDL >95b GAC NA ; >996 Ozone Temes et al., 2002
. b b « b
Erythmmycin 10a 53“ 47 ND8 100 Sand/GA 0.4‘il I 92 N800 Stackelberg eta1., 2007C
C ,
Naproxen 63—65 63-68 0b NA NA NA ND6 100*) NaOCl Boyd et al., 2003
b b .
SUIfamethovaOI 30“ 20a 33 NDa NA Sand/GA ND81 100 NaOC1 Stackelberg et al., 2007L
e C

 

 

 

 

 

 

 

 

 

 

 

 

Note: Percent removals were calculated based on pharmaceutical concentration before and after each treatment process (eg. primary treatment,
filtration, disinfection). ND: pharmaceutical not detected in sample; NM: not measured; NA: not available; BDL: below detection limit.

a Value represents average.

b Authors calculated removals from given concentrations.

6 Filtration followed primary disinfection.

d Approximated from f1 gure.

6 Removal after main ozonation

 

 

 

This section will review the following processes: lime softening, coagulation,
sedimentation, filtration, adsorption, and membrane filtration. Oxidation is considered an
important process that can effectively degrade or transform pharmaceuticals, and is

reviewed in the next section.

b. Lime Softening, Coagulation, Sedimentation

Several studies have investigated the effect of coagulation/flocculation, lime
softening, and sedimentation on removal of pharmaceuticals. Sedimentation by
coagulation/flocculation and lime softening are not expected to remove pharmaceuticals
due to the hydrophilic nature of most pharmaceuticals. Temes et al. (2002) observed
little or no removal of carbamazepine, diclofenac, and clofibric acid after coagulation
with iron chloride during full-scale drinking water treatment. Similarly, Stackelberg et
al. (2007) saw incomplete removal of acetaminophen, carbamazepine, erythromycin, and
sulfamethoxazole during coagulation with iron chloride at a DWTP. Clarification
resulted in percent removals of 3%, 33%, 47%, and 60% for carbamazepine,
sulfamethoxazole, erythromycin, and acetaminophen, respectively (Stackelberg et al.,
2007).

In a bench scale study by Adams et al. (2002), the coagulants, aluminum sulfate
and ferric sulfate, were added at doses of 20 - 107 mg/L and 35 - 169 mg/L, with initial
antibiotic concentrations of 50 rig/L. Initially the samples were mixed at 100 rpm for 1
min, 30 rpm for 20 min, followed by 3 hours of settling time. For lime softening, lime
and soda ash dosages of 232 and 191 mg/L, respectively, were applied in the same way as

coagulants, and the pH was adjusted to approximately 11.3. No significant removal of

15

antibiotics occurred during coagulation or lime softening processes. Similar results were
obtained by Westerhoff et al. (2005) in coagulation and lime softening experiments
involving 22 different pharmaceuticals. In a pilot-scale study, the average removal of 13
different pharmaceuticals was 3% during coagulation with fenic sulfate (Vieno et al.,
2007). Choi et al. (2008) conducted coagulation experiments on tetracycline antibiotics
using poly-aluminum chloride, added at dosages of 5 - 60 mg/L and mixed for 5 minutes.
Removal increased with higher dosage from 5 - 40 mg/L, but declined when a dosage
greater than 40 mg/L was added. The authors explained this by charge neutralization of
the antibiotics at high coagulant dosages, which caused a re-stabilization of antibiotics,
thus reducing removal efficiency. These studies indicate chemical precipitation

processes are not effective for removing pharmaceuticals.

c. Filtration

i. Adsorptive

Adsorption by activated carbon is commonly used in drinking water treatment for
removal of organic chemicals, taste and odor compounds, and NOM, either in powdered
(PAC) or granular form (GAC). The adsorption of organic chemicals to activated carbon
depends on the chemical’s structure, polarity, hydrophobicity, porosity of media,
activated carbon dosage, and presence of NOM (Choi et al., 2008; Mestre et al., 2007;
Vieno et al., 2007; Westerhoff et al., 2005; Yoon et al., 2003; Yu et al., 2003).

Adsorption by GAC may be effective for removing pharmaceuticals with high octanol-

water partitioning coefficients (kow), such as estradiol, ethinylestradiol, and fluoxetine

(log kow 2 4.0). In a pilot study by Vieno et al. (2007) most pharmaceuticals were

16

removed after a two-step GAC filtration process except those having high hydrophillicity
(atenolol, sotalol, ciprofloxacin). Choi et al. (2008) found the use of a coal-based GAC
column more effective than coagulation, with greater than 68% of tetracycline antibiotics
removed. Although not statistically confirmed, coal-based carbon exhibited slightly
better removal than coconut-based carbon.

PAC addition may also be effective, but is often only used seasonally to control
taste and odor compounds and NOM. Another issue that must be considered is disposal
of used PAC in landfills. Adams et a1. (2002) found a PAC dosage of 50 mg/L effective
for removing all antibiotics studied from both pure water and surface water. Westerhoff
et al. (2005) also found that increasing PAC dosage resulted in greater removal, with a
dosage of 20 mg/L removing greater than 80% of pharmaceuticals studied. The
researchers observed trends of protonated bases having greater removal than other
compounds, particularly compounds with deprotonated functional groups having low kOW
values. Yoon et al. (2003) investigated the adsorption of l7B-estradiol and l70t-ethynyl
estradiol onto six different PAC brands, tested at 5 and 15 mg/L doses and a 4 hour
contact time. Hormones were spiked at 100 nM. For the model water, greater than 99%
of the hormones were removed by all brands except for one, which had the lowest point
of zero charge value. Removals were lower in raw drinking water because of the
presence of NOM. Snyder et al. (2006) studied spiked pharmaceutical removal from
natural waters in bench and pilot scale experiments. Pharmaceutical removal depended
on PAC dose, contact time, and structure/behavior of each compound. The majority of
compounds tested were removed by greater than 90% with a PAC dose of 35 mg/L and 5

hour contact time.

17

In a study by Yu et al. (2003), adsorption of naproxen and carbamazepine to
bituminous carbon and coconut shell were compared using Freundlich isotherms. Results
indicate carabamazepine adsorbs to carbon to a greater extent than naproxen. Naproxen
showed greater adsorption affinity for bituminous carbon than coconut shell, while
carbamazepine did not display a significant difference. Authors attributed differences
between the compounds to naproxen being present in anionic form during the
experiments, while carabamazepine remained uncharged.

Mestre et al. (2007) studied adsorption of ibuprofen to activated carbon prepared
from waste cork powder. One type, CAC, was made using chemical activation with
potassium carbonate. The second type, CPAC, was made by physical activation of CAC
using steam. Surface chemistry characterization based on point of charge determination
indicated CAC had a more acidic surface than CPAC. Adsorption isotherms showed
CPAC had a higher adsorption capacity due to more developed porosity. Ibuprofen,
which appeared to follow the Langmuir isotherm model, adsorbed to CPAC at higher
rates and to a greater capacity than to CAC, likely due to the more developed porosity of

CPAC (Mestre et al., 2007). Temperature did not seem to affect adsorption, while

increase in pH led to decreased removal. The pKa of ibuprofen is 4.91, so at a pH greater

than 5, electrostatic repulsion occurs between the ibuprofen anion and the activated
carbon surface.

The adsorption behavior of acetaminophen and nalidixic acid onto polar and
nonpolar adsorbents was investigated using neutral polymeric resins as polar adsorbents
and powdered activated carbon as a nonpolar adsorbent (Suntisukaseam et al., 2007). It

was found that nalidixic acid, a quinolone antibiotic with a carboxylic functional group,

18

adsorbed to all media better than acetaminophen. Acetaminophen has a much greater

water solubility, lower Kow and higher polarity than nalidixic acid, so it tends to stay in

solution. Adsorption onto polymeric resins was found to follow the Langmuir isotherm,
while activated carbon adsorption followed the Freundlich isotherm. Greater polarity of
the polymeric resins resulted in greater amounts of pharmaceuticals adsorbed per area of
adsorbent. For both pharmaceuticals, activated carbon resulted in higher sorption than

the polymeric resins.

ii. Sand

Removal by sand filtration is related to a compound’s hydrophobic or hydrophilic
tendencies. For example, Nakada et al. (2007) observed that after sand filtration,
hydrophilic compounds, such as carbamazepine, sulfapyridine, sulfamethoxazole, and
estriol, had lower removal efficiencies (<50%) than hydrophobic compounds, such as
ibuprofen, which were removed by >80% in some cases. At a drinking water treatment
plant studied by Stackelberg et al. (2007), sand/GAC filtration was employed in between
primary and secondary chlorine disinfection, and completely removed the remaining
erythromycin, and the remaining carbamazepine by 97%. Temes et al. (2002) took
samples after GAC filtration in full-scale drinking water treatment: diclofenac and
bezafibrate were removed by greater than 95%, carbamazepine by more than 75%, and

clofibric acid by 20%.

19

iii. Membrane

Verliefde et al. (2007) studied removal of 20 pharmaceuticals from surface water
by a bench-scale nanofiltration (NF) membrane (1500 l/h feed flow, 10% recovery) and
by 3 NF membrane with subsequent GAC column (80% recovery). Results indicate
removal of pharmaceuticals by NF depends on the charge of the pharmaceutical.
Rejection values were lowest for positively charged pharmaceuticals and highest for
negatively charged pharmaceuticals. Neutral compounds exhibited intermediate removal,
likely due to negatively charged membrane surface, or hydrophobicity of neutral
compounds (Verliefde et al., 2007). Higher removals were achieved when NF was
combined with GAC filtration. Yoon et al. (2006) also investigated removal of
pharmaceuticals by membrane filtration, by NF as well as by ultrafiltration (U F).
According to results, NF membranes retained pharmaceuticals based on hydrophobicity
and size of particle, while UF membranes retained pharmaceuticals mainly based on
hydrophobicity alone.

Snyder et al. (2006) investigated rejection of pharmaceuticals by a partially fouled
UF membrane, and found the majority of compounds in the feed water were not rejected.
Compounds that were well-removed included the steroid hormones such as estradiol,
estrone, ethinylestradiol, and progesterone. In a pilot RO system, Snyder et al. (2006)
found that all pharmaceuticals spiked into feed water were well-rejected by both virgin
membranes and fouled membranes. However, pharmaceuticals are then present in
greater concentrations in the retentate water, thus requiring further treatment prior to

discharge.

20

d. Conclusions

Drinking water treatment consists of a train of physicochemical processes, which
may or may not remove or transform pharmaceuticals into less reactive compounds.
Certain combinations of processes are more effective than others, and bench-scale studies
indicate that many factors influence the degree at which a pharmaceutical is transformed.
Coagulant dosage, type or brand of activated carbon used, and the compound’s chemical
properties are just a few parameters to consider when planning a treatment train that will
minimize the concentration of pharmaceuticals in finished drinking water. Lime
softening, coagulation, and sedimentation processes are not effective at removing
pharmaceuticals. Adsorptive, sand, and membrane filtration are effective at removing

some pharmaceuticals, depending on the compound’s chemical structure and properties.

Oxidation of Pharmaceuticals

Oxidation processes are routinely applied at wastewater and drinking water
utilities for disinfection and taste and odor control. Most pharmaceuticals are susceptible
to oxidation because they are polar organic compounds containing reactive functional
groups. Appendix 1 provides the chemical formula, molecular weight, octanol-water
partitioning coefficient, acid-dissociation constant, and structure for various
pharmaceuticals. Data and structures in this table were obtained from the chemical
database, DrugBank (W ishart et al., 2006). Chemical oxidants typically used in water
treatment include chlorine, chloramines, and ozone. UV based oxidation processes can

also be used. Oxidation occurs when there is a net loss of electrons from an atom.

21

Oxidants, such as chlorine, are electrophilic molecules that attack the electron-rich areas

of a nucleophile.

a. Oxidation by Chlorine-Containing Oxidants

Chlorine is capable of oxidizing many different inorganic and organic compounds
(Deborde and von Gunten, 2008). Typically, chlorine is added to water in the form of
sodium hypochlorite (N aOCl), or chlorine gas, which exists primarily as hypochlorite ion
(OCl') and hypochlorous acid (HOCl) in solution. Collectively the two species are

referred to as free available chlorine, and the concentration of free chlorine is expressed

as mg/L C12. The chlorine species that is dominant in solution depends on the pH,

temperature, and total concentration of chlorine [AWWA, 1999]. The pKa of free

chlorine is 7.5 with HOCl predominating at pH levels less than 7.5, and the weaker
disinfectant, OCl-, predominating at a pH greater than 7.5. Major parameters that affect
oxidation pathways are the pharmaceutical molecular structure, pharmaceutical
concentration, chlorine concentration, contact time, pH, temperature, and water
composition.

Depending on the molecular structure of the pharmaceutical, chlorine will react
differently with different types of compounds. In some cases a pharmaceutical may not
react with chlorine at all. For example, Glassmeyer and Shoemaker (2005) investigated
the fate of pharmaceuticals during chlorination via bench top experiments. Chlorinated
and non-chlorinated mass chromatograms and spectra for pharmaceuticals were
compared to determine the effect of chlorination on each compound. The results

indicated that each pharmaceutical fell into one of three groups: unchanged by

22

chlorination, chlorinated, or changed by chlorination but did not appear to become
chlorinated. For the conditions tested, the pharmaceuticals that were unchanged were
aspirin, and 60t-methyl-170t-hydroxy progesterone acetate. Amoxicillin, cephalexin,
cimetidine, trimethoprim, diltiazem, and warfarin were changed by chlorination but did
not appear to become chlorinated. Only acetaminophen and gemfibrozil showed definite
signs of chlorination.

In some cases, pharmaceuticals react rapidly and are transformed by free chlorine,
such as acetaminophen, naproxen, triclosan, diclofenac, and estrogens [Bedner and
Maccrehan, 2006; Boyd et al., 2005; Westerhoff et al., 2005]. In other cases,
pharmaceuticals exhibit little or no reactivity and do not undergo transformation with free
chlorine, such as meprobamate, ibuprofen, ketoprofen, carbamazepine, and erythromycin
[Gibs et al., 2007; Pinkston and Sedlak, 2004; Westerhoff et al., 2005].

Even though the majority of chlorination studies have used sodium hypochlorite
as the disinfectant, the effectiveness of different forms of chlorine in water treatment (i.e.
chlorine dioxide, monochloramine) has also been assessed. Chamberlain and Adams
(2006) compared free chlorine with monochloramine as an oxidant and found that
monochloramine is less effective at removing pharmaceuticals from drinking water than
free chlorine. However, monochloramine is thought to form less halogenated byproducts
than free chlorine [Chamberlain and Adams, 2006; Richardson, 2003]. In the
Chamberlain and Adams study (2006), the only antibiotic that showed a high degree of
reactivity with monochloramine was carbadox. In a study by Pinkston and Sedlak

(2004), none of the pharmaceuticals tested reacted fast enough with monochloramine to

23

be transformed. Also, the synthetic hormone, Hot-ethinylestradiol, was not transformed
by monochloramine at typical water treatment doses of a few mg/L (Lee et al., 2008).
Few studies have been published that examine the reactivity of pharmaceuticals
with chlorine dioxide. In a study by Huber et al. (2005b), chlorine dioxide was applied to
groundwater, drinking water, or surface water spiked with various pharmaceuticals.
Chlorine dioxide doses ranged from 0.1 to 11.5 mg/L, and pharmaceuticals were spiked
at concentrations less than or equal to 1 rig/L. Pharmaceutical solutions were in contact
with chlorine dioxide up to 180 minutes, and the pH was kept at ambient levels (7.2-7.9).
Some compounds were readily oxidized (diclofenac, l7B-estradiol, estrone, 17a-
ethinylestradiol, phenazone, sulfonamide antibiotics), while some compounds did not
react at all (bezafibrate, carbamazepine, diazepam, fenoprofen, ibuprofen, ketoprofen).
The study showed that chlorine dioxide is effective at oxidizing only specific classes of
compounds such as sulfonamides and estrogens. Lee et al. (2008) also reported

significant transformation of l7a—ethinylestradiol by chlorine dioxide.

b. Ozonation of Pharmaceuticals

The use of ozonation in drinking water treatment is more common in Europe and
Asia than in the United States, due to relatively high costs in the United States. Ozone
reacts more rapidly with pharmaceuticals than chlorine because of its high reactivity with
most organic functional groups [Adams et al., 2002; Alum et al., 2004; Huber et al.,
2005a; Nakada et al., 2007; Westerhoff et al., 2005]. For example, Westerhoff et al.
(2005) applied 2.5 to 4 mg/L ozone to a wide variety of pharmaceuticals at

concentrations ranging from 10-250 ng/L. After 10 minutes in contact with ozone, most

24

of the compounds were oxidized by greater than 80%. The degradation efficiency of
ozone is greater when it is applied in combination with hydrogen peroxide, which helps
stimulate the generation of OH radicals [Zweiner and F rimmel, 2000]. For example,
when ozone was applied directly to water spiked with 2 [Lg/L each of clofibric acid,
diclofenac, and ibuprofen at a dose of 1 mg/L, and the compounds were allowed to react
for 10 minutes, only diclofenac was degraded readily [Zweiner and F rimmel, 2000].
However, when hydrogen peroxide was used in combination with ozone, the degradation

efficiency of all three compounds increased.

c. UV Irradiation of Pharmaceuticals

Ultraviolet irradiation is used in both drinking water and wastewater treatment.
UV processes are often combined with hydrogen peroxide and are used to oxidize taste
and odor-causing chemicals and to disinfect water [Rosenfeldt and Linden, 2004].
Typically, low pressure high-output lamps or medium pressure lamps are used in large-
scale water disinfection (Linden and Kullman, 2007). A comparison of monochromatic
low-pressure (LP) UV lamps with polychromatic medium pressure (MP) lamps was done

by Rosenfeldt and Linden (2004) on endocrine disruptors, including the hormones ethinyl

estradiol and estradiol. At the same fluence of 1,000 mi cm-2, the MP lamps resulted in

greater degradation of the hormones than the LP lamps. This is due to the low
probability of radiation absorption by the hormones using the monochromatic LP lamp
which emits at 254 nm (Linden and Kullman, 2007). The polychromatic MP lamp
radiation covers wavelengths between 200 and 300 nm, allowing molecules to reach

excited states and become unstable.

25

The addition of hydrogen peroxide (H202) to UV irradiation results in the
photolysis of H202 into OH radicals, one of the strongest known oxidants [Vogna et al.,

2004]. Direct UV irradiation in the absence of H202 has been found to be a less effective

oxidant than the combined process [Adams et al., 2002; Rosenfeldt and Linden, 2004;
Rosenfeldt et al., 2007; Vogna et al., 2004]. Adams et al. (2002) applied UV radiation to

water spiked with 50 [lg/L carbadox, trimethoprim, and sulfonamide antibiotics. Their

study showed that typical UV dosages, about 30 mJ-cm'z, used during disinfection of

drinking water were ineffective at removing all compounds. Even at doses 100 times
greater the antibiotics were not completely removed. Adams et al. (2002) attributed this
inefficiency to weak ultraviolet absorption of the specific compounds. Linden and

Kullmen (2007) studied degradation of 17B-estradiol and l7a-ethinyl estradiol by direct

UV and UV/ H202 in terms of both chemical degradation and destruction of estrogenic

activity. Results indicated the oxidation products of the estrogens were less
estrogenically active than the parent compounds. Another interesting finding was that
more oxidation of these compounds occurred at environmentally-relevant concentrations
and at lower UV doses than suggested by previous research which used higher

concentrations and dosages.

d. Comparison of Different Oxidants
The type of oxidant used in water treatment plays a major role in the fate of
pharmaceuticals. As shown in Table 5, different oxidants react differently with the same

pharmaceutical compounds. In general, studies have shown that pharmaceuticals exhibit

26

mixed reactivity with free chlorine [Chamberlain and Adams, 2006; Deborde et al., 2004;
Gibs et al., 2007; Glassmeyer and Shoemaker, 2005; Qiang et al., 2006; Westerhoff et
al., 2005]. The same is true for chlorine dioxide and monochloramine [Chamberlain and
Adams, 2006; Huber et al., 2005b; Qiang et al., 2006]. Ozone reacts rapidly with most
pharmaceuticals except for some acidic pharmaceuticals, such as clofibric acid,
gemfibrozil, ibuprofen, bezafibrate, and meprobamate [Huber et al., 20053; Temes et al.,
2002; Zweiner and Frimmel, 2000]. Most acidic pharmaceuticals contain a carboxylic
acid group, which may not react well with ozone.

UV irradiation of pharmaceuticals seems to only be effective in the presence of
hydrogen peroxide [Adams et al., 2002; Rosenfeldt and Linden, 2004; Rosenfeldt et al.,
2007; Vogna et al., 2004]. In a full-scale observation, Renew and Huang (2004) sampled
effluents from two wastewater treatment plants for antibiotics. One treatment plant used
chlorine for disinfection and the other used UV as the disinfectant. The results showed

that chlorination removed significantly more antibiotics than UV disinfection.

27

 

 

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28

e. Mechanisms of Oxidation

Reactions between chlorine and organic molecules may occur by oxidation of
functional groups, addition to double bonds, and electrophilic substitution [Brezonik,
1994]. The reactivity of an organic contaminant with chlorine depends largely on its
functional groups [Westerhoff et al., 2005], although the reactivity of the oxidant itself is
also important. Snyder et al. ( 2003) makes two generalizations about the reactivity of
oxidants with pharmaceuticals: protonated forms of pharmaceuticals are less reactive than
dissociated acidic forms, and aliphatic pharmaceuticals are less reactive than aromatic
compounds. The order of reactivity (highest to lowest) for aromatic or aliphatic
compounds is: thiols > amines > hydroxyl > carboxyl [Snyder et al., 2003]. Many
pharmaceuticals contain such functional groups as well as phenolic groups. Two of the
best documented reactions of HOCl/OCI' is the reaction of free chlorine with phenols,
and chlorine addition to primary and secondary amines [Pinkston and Sedlak, 2004]. In
the first reaction, HOCl reacts with the phenolate anion to yield ortho- and para-

substituted chloro-phenols. In the reactions with amines, HOCl reacts with the

unprotonated amine. For both cases, the maximum reaction rates occur between the pka

of HOCl (7.5) and that of the phenol or amine [Pinkston and Sedlak, 2004]. Hormones
such as estrone, B-estradiol, and l7a—ethinylestradiol contain a phenolic ring and exhibit
rapid reactivity with chlorine, suggesting that other compounds containing phenolic rings
will undergo rapid elimination during chlorination as well [Deborde et al., 2004].

Dodd and Huang (2004) investigated the transformation of the antibiotic,

sulfamethoxazole, during chlorination. Results showed that chlorine attacks the aniline-

29

nitrogen of sulfamethoxazole directly leading to the halogenation of the aniline moiety,
yielding a ring-chlorinated product. They also found that the sulfonamide moiety of
sulfamethoxazole is ruptured when free chlorine is in excess, resulting in the formation of
3—amino-5-methylisoxazole and N-chloro-p- benzoquinoneimine. Huber et al. (2005b)
studied the sulfonamide antibiotics sulfamethoxazole, sulfamethazine, sulfapyridine, and
sulfathiazole, which all contain the aniline moiety. Although the compounds have similar

structures, they exhibited varying reactivity with chlorine dioxide, which the authors

attributed to differences in speciation. For example, sulfapyridine has a pka of 8.4 so

during the experiments it remained in its neutral form and was less reactive than

sulfamethoxazole, which has a pkgl of 5.7.

In ozonation, pharmaceuticals can be oxidized by one of two mechanisms, either
by ozone directly, or by OH' radicals. Ozone spontaneously decomposes in water by
chain reactions involving different free radical species [AWWA, 1999]. When applied
directly, a significant amount of ozone may be lost due to decomposition in water
[Staehelin and Holgne, 1982]. Ozone tends to react selectively with functional groups,
while OH' radicals are nonselective toward organic compounds [Calamari et al., 2003;
Haag and Yao, 1992]. Westerhoff et al. (2005) found that functional groups affect how
efficiently a pharmaceutical is oxidized by ozone. For example, ibuprofen, which
contains a propionate moiety, was not as efficiently oxidized as compounds that contain
an amine functional group, such as diclofenac. Westerhoff et al. (2005) also compared
the structural properties of compounds that reacted rapidly with both chlorine and ozone

and found that those compounds usually had hydroxyl or amine functional groups along

with low pka values. Evidence from both bench-scale and full-scale studies suggest

3O

pharmaceuticals with a double carbon bond (C=C) or an activated aromatic ring are more
susceptible to ozonation than pharmaceuticals with amide structures [Nakada et al.,2007].
UV or UV/HzOz processes are capable of oxidizing organic contaminants, but the
effectiveness of the processes depends on the absorption spectra of the molecule. Adding
hydrogen peroxide (H202) helps to lower the UV dose required [Rosenfeldt and Linden,
2004]. The effectiveness of UV processes also depends on the pressure of the lamp used
and the ability of absorbed radiation by the pharmaceutical to reach an excited state
[Rosenfeldt and Linden, 2004]. Other factors that influence UV efficiency are

background absorbance and scavenging capacity of the water matrix.

fl Factors Affecting Oxidation Rates

The degradation of pharmaceuticals during oxidation of wastewater or drinking
water depends primarily on the pH of the water, the amount of oxidant present, the
concentration of pharmaceutical, the available contact time, and matrix water. Table 4
shows pH effects on pharmaceuticals during chlorination studies involving sodium
hypochlorite. Boyd et al. (2005) found the transformation of naproxen during
chlorination was dependent on pH, contact time, and chlorine dose. The rate and extent
of the reaction increased when a higher chlorine dose was applied and favored low pH

levels.

31

TABLE 4. Effect of pH During Chlorination of Pharmaceuticals by Sodium

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Hypochlorite
Pharmaceutical pH Studied pH Effect Source
6, 7.5, 9 . . Xagoraraki et al., 2008;
Acetaminophen 5.5, ambient Effects :5? ‘“ “Ch Westerhoff et al., 2005;
5 — 10 y Pinkston & Sedlak, 2004
Atenolol 5 — 10 Ifigrmedlr}? Pinkston & Sedlak, 2004
6.1, 7.6, 9.1 .
Carbadox No effect observed Chamberlain & Adams, 2006
. Decreased with
Carbamazepine 5. 5,ambien t increased PH Westerhoff et al., 2005
Diazepam 55’ ambient Decreased With Westerhoff et al., 2005
increased pH
Diclofenac 5.5, ambient No effect observed Westerhoff et al., 2005;
Dilantin 5.5, ambient No effect observed Westerhoff et al., 2005;
. 6.1, 7.6, 9.1 Chamberlain & Adams, 2006;
Erythromycm 5.5, ambient No effect observed Westerhoff er al., 2005
Estradiol 3.5 - 12 Greatest between pH Deborde er al., 2004;
5.5, ambient 8 and 10 Westerhoff er al., 2005
Estriol 3.5 — 12 Greatest between pH Deborde et al., 2004;
5.5, ambient 8 and 10 Westerhoff et al., 2005
Es trone 3.5 — 12 Greatest between pH Deborde et al., 2004;
5.5, ambient 8 and 10 Westerhoff er al., 2005
Ethin lestradiol 3.5 — 12 Greatest between pH Deborde er al., 2004;

y 5.5. ambient 8 and 10 Westerhoff et al., 2005
Fluoxetine 5.5, ambient No effect observed Westerhoff et al., 2005
Gemfibrozil 5 — 10 Decreased with Pinkston & Sedlak, 2004;

5.5, ambient increased pH Westerhoff er al., 2005
5 - 10 Pinkston & Sedlak, 2004;
Ibupmfe“ 5.5, ambient N0 ”women“ Westerhoff et al., 2005
Indometacine 5 - 10 Decreased “m" Pinkston & Sedlak, 2004
increased pH
Metoprolol 5 ‘ 1° [immeased ‘23 Pinkston & Sedlak, 2004
5, 7, 9 . Boyd et al., 2005;
Naproxen 5 — 10 xix wgh Pinkston & Sedlak, 2004;
5.5, ambient p Westerhoff et al., 2005
Propanolol 5 _ 10 Increased With Pinkston & Sedlak, 2004
increasmg pH
Sulfamerazine 6'1’ 7'6’ 9 1322:2222 :1?) Chamberlain & Adams, 2006;
Sulfamethazine 6.1, 7.6, 9 Decreased with Chamberlain & Adams, 2006;
6.6, 7.6, 8.6 increased pH Qiang er al., 2006
6.1, 7.6, 9 Chamberlain & Adams, 2006;
Sulfamethoxazole 4 — 9 Decreased with Dodd & Huang, 2004;
6.6, 7.6. 8.6 increased pH Qiang et al., 2006;
5.5, ambient Westerhoff er al., 2005
Sulfathiazole 6'1’ 76’ 9 12:32:: :11? Chamberlain & Adams, 2006;
Trimethoprim 5.5, ambient No effect observed Westerhoff er al., 2005

 

32

 

Pinkston and Sedlak (2004) found that the reaction rates of pharmaceuticals with
free chlorine increased when the pH was reduced from 10 to 7, since free chlorine exists
predominantly as HOCl below a pH of 7.5, and HOCl is more reactive than OCl'. Dodd
and Huang (2004) also found that reactions of sulfamethoxazole with chlorine were pH
dependent. The reactivity of sulfamethoxazole with free chlorine decreased substantially
with increasing pH. Similar results were observed by Chamberlain and Adams (2006),
when they studied the removal of six different sulfonamides with chlorine; sulfonamides
were readily removed from drinking water at neutral pH, but were only partially removed
at pH 9. Pinkston and Sedlak (2004) observed reactions of aromatic ether- and amine-
containing pharmaceuticals with free chlorine where they found that the reaction rates
tended to increase with decreasing pH. Acidic pharmaceuticals, such as gemfibrozil,
ketoprofen, and ibuprofen, exhibit low reactivity with chlorine-containing oxidants [Gibs
et al., 2007; Glassmeyer and Shoemaker, 2005; Pinkston and Sedlak, 2004; Westerhoff et
al., 2005]. These pharmaceuticals all have acid dissociation constants less than 5, but
were studied at pH levels greater than or equal to 5.0. For acidic pharmaceuticals, the
greatest reactivity with oxidants may fall at lower pH levels.

The extent that a pharmaceutical reacts with an oxidant is also largely dependent
on the oxidant dose and the concentration of pharmaceutical present. When the molar
ratio of oxidant to pharmaceutical is greater, a higher degree of reactivity is expected.
For instance, Huber et al. (2005a) observed that the amount of oxidation of a
pharmaceutical increased with greater ozone doses. Also, Chamberlain and Adams
(2006) applied either 0.1 or 1.0 mg/L free chlorine to distilled water spiked with

macrolide antibiotics. After two hours, the average removal of antibiotics for a 0.1 mg/L

33

C12 dose was only 28%, while the average removal for a 1.0 mg/L C12 dose was 85%.

Adams et al. (2002) found that UV doses typical to water disinfection were ineffective at

removing several different antibiotics. In order to achieve 50 — 80% removal of the

antibiotics, the UV dose had to be increased to 3,000 mJ/cmz. This dose was 100 times

greater than those typically used in disinfection.

Oxidation rates and the formation of intermediate products are affected by the
length of time a pharmaceutical is in contact with an oxidant as well. Chamberlain and
Adams (2006) observed that antibiotics were removed to a greater extent as the contact
time with chlorine increased, however, the time required for adequate removal varied for
each antibiotic. Boyd et al. (2005) found that when naproxen was chlorinated,
intermediate products increased over contact time, and then decreased after one hour.

Another factor affecting degradation rates of pharmaceuticals is the water matrix.
In a study done by Adams et al., removal times of antibiotics by free chlorine were
shorter in distilled water than in surface water suggesting natural organic matter in
surface water decreases reactivity (2002). Huber et al. (2005b) demonstrated that when
chlorine dioxide is used as a disinfectant for surface waters containing a considerable
amount of organic matter, the pharmaceuticals and the water matrix compete for chlorine
dioxide, thereby increasing the chlorine dioxide dose necessary to oxidize the
pharmaceuticals. When ozone is used as an oxidant, the oxidation rates of OH radicals
with organic compounds may be limited by the rate of OH generation and by competing
OH scavengers such as dissolved organic carbon [Haag and Yao, 1992]. Adams et al.
(2002) found that reactions of ozone with antibiotics were faster in distilled water than in

surface water.

34

g. Formation of Byproducts

Often, pharmaceuticals are said to be ‘removed’ by oxidation processes, when in
fact they are only transformed into different compounds. It is known that chlorine reacts
with dissolved natural organic matter to produce disinfection byproducts including
trihalomethanes and haloaccetic acids. Deborde and von Gunten (2008) reported that 600
different disinfection byproducts have been identified. Ozonation may also produce
bromate when it is applied to waters that contain bromide, and chlorine dioxide reacts
with organic compounds to form chlorite. Monochloramine as a disinfectant can lead to
the formation of nitrosamines, while UV irradiation may convert nitrate to nitrite.

Pharmaceuticals may react with chlorine during water treatment to form
byproducts with confirmed toxicity to humans [Bedner and Maccrehan, 2006; Korshin et
al., 2006], or byproducts of unknown or no toxicity to humans [Andreozzi et al., 2005;
Bedner and Maccrehan, 2006; Dodd and Huang, 2004; Gould and Richards, 1984;
McDowell et al., 2005; Pinkston and Sedlak, 2004]. Similarly, ozonation of drinking
water may result in the formation of byproducts as well [Alum et al. , 2004; Andreozzi et
al., 2005]. The formation of oxidation byproducts that cause adverse health effects is of
particular concern in regards to human health.

Table 5 presents a summary of oxidation byproducts of pharmaceuticals and
whether or not the byproducts are known to be toxic to humans. As shown in the table,
the toxicity of most of the byproducts is relatively unknown, indicating more research in
this area is needed. Bedner and Maccrehan (2006) discovered two byproducts of known
toxicity to humans resulting from the chlorination of the common pain reliever,

acetaminophen. Phenolic compounds, such as acetaminophen, react to form monochloro-

35

and dichloro- intermediate products from the chlorine substitution at the ortho and para
positions of the ring [Brezonik, 1994]. The antibiotic, sulfamethoxazole contains 2m
aniline ring, which rearranges to ortho-chlorinated intermediates in the presence of free
chlorine [Dodd and Huang, 2004]. Depending on oxidation conditions, finished water
may contain compounds with confirmed toxicity to humans or compounds of unknown
toxicity. Further chemical, toxicological, and risk assessment studies are needed to

determine the potential for human and environmental risk.

TABLE 5. Byproduct Formation during Pharmaceutical Oxidation

 

 

 

 

 

 

Byproduct
Oxidant Pharamceutical Byprodch Toxicity Reference
NaOCl Acetaminophen N-Acetyl-p-benzoquinone imine Toxic Bedner &
1 ,4-Benzoquinone Toxic Maccrehan,
Chloro-4-acetamindophenol Unknown 2006
Dichloro-4-acetamidophenol Unknown
NaOCl Estrogen Chloroacetic acids Toxic Korshin et
al., 2006
NaOCl Sulfamethoxazole 3-amino-5-methylisoxazole Not Toxic Dodd and
N-chloro—p-benzoquinoneimine Possibly Huang,
2004
Ozone Amoxicillin 2—amino—2-(p—hydroxyphenyl) - Unknown Andreozzi
acetic acid et al., 2005
1-(2-benzaldehyde)-4-hydro-(1H, 3
Ozone Carbamazepine H))-quinazoline—2-one Unknown
l-(2-benzaldehyde)-(l H, 3H)- McDowell
quinazoline-2,4-dione Unknown et al. , 2005
l-(2-benzoic acid)-(1 H, 3H)-
quinazoline-2,4-dione Unknown

 

 

 

h. Quantitative Structure-Property Relationships (QSPRs)

To study how each pharmaceutical behaves during water treatment would require
a great deal of time and money. Since each pharmaceutical compound has different
properties (see Appendix 2), and a large number of pharmaceuticals exist. The use of

QSPRs for modeling contaminants reduces the amount of experimentation required to

36

understand contaminant fate and transport by making predictions based on a compound’s
structure and chemical properties (Papa and Gramatica, 2008; Lei and Snyder, 2007; Hu
et al., 2000; Dai et al., 1999). One of the primary assumptions of QSPR modeling is that
the chemical and physical properties of a compound depend on its structure (Dai et al.,
1999). Chemical properties and structures are then used to form mathematical
relationships that allow for prediction of contaminant fate in the environment. Statistical
testing and comparison with available data determine if the developed equations are
predictive of reaction rates or removal mechanisms. An approach to QSPR modeling of
compounds with similar structures usually involves molecular connectivity characteristics
and quantum chemical descriptors (Dai et al., 1999).

Different molecular properties or descriptors may be used to predict removal by
separate treatment processes. For instance, common predictors used for correlation with

oxidation rate constants are oxidation potentials, o constants in the Hammett equation,

and the energy of the highest occupied molecular orbital (SHOMO) (Hu et al., 2000). To

model contaminant removal by free chlorine and ozone, Lei and Snyder (2007) used
various geometrical properties, solubility, Henry’s law constant, hexadecane/gas partition
coefficient, octanol/ gas partition coefficient, and water/ gas partition coefficient. The
study also evaluated dipole moment, electron affinity, ionization potential, molecular
weight, octanol-water partition coefficient, polarizability, and number and types of
functional groups. Lei and Snyder (2007) were able to determine the most important
parameters for a QSPR ozonation model based on statistical testing. They found
contaminant removal correlated with weakly polar surface area, number of metabolites, 1:

surface area, number of reactive functional groups and electron affinity, with the most

37

important parameter being weakly polar surface area. In the same study, Lei and Snyder
found functional groups were significant in determining reactivity with chlorine. For a
more detailed approach to kinetic modeling, refer to Crittenden’s study on contaminant

degradation by advanced oxidation processes (Crittenden et al., 1999).

i. Conclusions

The oxidation of a pharmaceutical is affected by the pharmaceutical’s functional groups,
oxidant dose, water matrix, pH of the system, and the initial pharmaceutical
concentration. Ozone is capable of oxidizing more functional groups than chlorine, and
this is demonstrated through comparison of the studies involving pharmaceutical
oxidation by ozone or chlorine. UV is only effective at oxidizing pharmaceuticals if
hydrogen peroxide is used in combination. To avoid costly and time-consuming
experiments, QSARs may be employed to predict the oxidation potential of a

pharmaceutical.

Conclusions and Recommendations

Drinking water utilities may be required to treat source water for pharmaceuticals
if federal or state agencies decide to regulate these compounds. Currently, there is some
debate within the scientific community over the relevance of pharmaceuticals in drinking
water and whether they pose significant risks to human health. A question that remains
is, does the ability to detect substances in water mean that it is really a human health risk?
Have humans been consuming water containing pharmaceuticals for decades without

realizing? Humans have used medicines since ancient times, but only recently have

38

detection methods been developed. Instruments today are able to detect pharmaceuticals
at ng/L concentrations, but without sufficient toxicology studies involving relevant
mixtures of organic compounds present in drinking water, it is very difficult to assess the
risk these compounds pose to humans.

Generalized predictions for the removal or transformation of pharmaceuticals in
water and wastewater utilities cannot be made since their fate depends on multiple
parameters. The efficiency of an oxidant and its effects on the fate of pharmaceuticals
depend on the molecular structure of the pharmaceutical, concentration of
pharmaceutical, pH of the water, oxidant dose, contact time, and water matrix. For
example, some compounds exhibit greater reactivity at lower pH levels and some at
higher pH levels. Conventional studies of fate, transformation and removal of
pharmaceuticals are important, but it is not possible to evaluate each and every
pharmaceutical under all treatment processes and all operating conditions. Modeling
efforts that take into account the molecular structure of pharmaceuticals, such as
quantitative structure activity relationships (QSARs), and kinetic parameters, may
provide significant results in this area.

The focus of research can be narrowed to more persistent and/or unpredictable
compounds. These include dilantin, diclofenac, carbamazepine, fluoxetine, and
ibuprofen. More research on oxidation byproducts of endocrine disruptors is also
needed.

Implementation of a risk assessment is necessary to determine acceptable levels
of pharmaceuticals in environmental waters and drinking water. Once levels are

determined, preventative measures may be taken to reduce or remove pharmaceuticals.

39

More research is needed on pharmaceutical transport as it applies to the land application
of biosolids. Also, the increasing development of new pharmaceuticals requires a method
for predicting and identifying compounds that are not currently being monitored, or a
way to predict risk of new compounds before they enter environmental waters. It may
also be possible to predict mixtures of drugs that occur by region based on drug
prescriptions and sales. Furthermore, there is a need for research on long term effects on
humans and organisms of exposure to multiple pharmaceuticals, metabolites, and

oxidation byproducts.

40

CHAPTER 2

LABORATORY EXPERIMENTS

Summary

To investigate further the oxidation of pharmaceuticals, bench-scale experiments
involving the chlorination of acetaminophen were conducted. In total, 23 experiments
were performed at room temperature using sodium hypochlorite. The initial objectives of
the experiments were to study how pH, molar ratio of oxidant to pharmaceutical, and
initial pharmaceutical concentration effected degradation and the formation of an
oxidation byproduct. Later objectives were to determine reaction mechanisms that
dominate during the chlorination of acetaminophen by developing equations that model
the reaction kinetics. Acetaminophen was chosen because it is commonly used, and has
been documented to produce 1,4-benzoquinone during chlorination.

The first nine experiments were conducted at pH levels of 6.0, 7.5, and 9.0 and
the molar ratio of chlorine to acetaminophen was varied from about 100 to 10,000.
Results indicate that acetaminophen is highly susceptible to oxidation by free chlorine,
but greatest degradation occurs at pH 9.0 and a molar ratio of 10,000. The production of
1,4-benzoquinone was also effected by pH and molar ration, with the maximum amount
of 1,4-benzoquinone produced at pH 6.0 and a molar ratio of 1,275.

The next fourteen experiments were conducted at pH levels of 6.0, 7.5, 9.0 and
11.0 and the initial acetaminophen concentration was lowered from 2,000 rig/L to 200

and 1,000 ug/L. Overall, acetaminophen reacts with chlorine to the greatest extent

41

between pH 7.5 and 9.0, and the least at pH 6.0 and 11.0. Lowering the initial
acetaminophen concentration reduced the amount of degradation that occurred.
Preliminary risk assessments were also performed using literature toxicity data
which suggest a relatively low level of ecological and human health risk. Currently,
research on this subject by another graduate student is using the data from the
acetaminophen experiments to develop a set of equations that can be computer modeled

to predict reaction rates at different conditions.

Primary Experiments
a. Abstract

The concern over pharmaceuticals and their toxicity in wastewater and drinking
water has grown over the past decade. In this study, the common analgesic,
acetaminophen, was chlorinated with sodium hypochlorite to determine the effects of pH
and chlorine—to-pharmaceutical molar ratios on the degradation of acetaminophen and the
formation of the toxic byproduct 1,4-benzoquinone. Reactions were studied for pH 6.0,
7.5 and 9.0 at average molar ratios of 106 i 6, 1,417 i 285, and 9,789 i 1,430 over a
period of 100 minutes. The degradation of acetaminophen and the formation of 1,4-
benzoquinone were monitored using liquid chromatography tandem mass spectrometry
(LC/MS/MS). Results indicate that acetaminophen is most reactive with free chlorine at
pH 9.0 and least reactive at pH 6.0. As pH increased, degradation of acetaminophen also
increased. The formation of 1,4-benzoquinone was also affected by pH and reached a
maximum of 68.7% of the initial acetaminophen concentration when the pH was at 6.0,

the molar ratio at 1,275, and after a contact time of 30 minutes. At all pH values the rate

42

of degradation of acetaminophen was slowest at a molar ratio of about 100, and the

highest at a molar ratio of about 10,000.

b. Introduction
Acetaminophen, also known as paracetamol, is a very common over-the-counter

analgesic used for fever, headaches and other minor pain. In 2002, the US produced 3.6 x

109 g of acetaminophen (Bedner and Maccrehan, 2006). In 1998, it was estimated that

3.2 x 109 tablets were consumed in the UK alone (Bessems and Vermeulen, 2001).

Acetaminophen has been detected in surface waters, wastewater, and drinking
water. In the 2002 survey of 139 US. streams, Kolpin et al. detected acetaminophen in
roughly 25% of the samples tested at a median concentration of 0.11 rig/L (2002). The
median concentration of acetaminophen detected in surface waters is 0.055 i 0.051 pig/L
(Bound & Voulvoulis, 2006; Gros et al., 2006; Stackelberg et al., 2004; Wiegel et al.,
2004; Boyd & Furlong, 2002; Kolpin et al., 2002). In raw wastewater, acetaminophen
was detected at a median concentration of 48 i 75 ug/L (Gomez et al., 2007; Gros et al.,
2006; Han et al., 2006). The median concentration of acetaminophen detected in finished
wastewater is 0.76 i 0.96 ug/L (Gomez et al., 2007 ; Radjenovic et al., 2007; Bound &
Voulvoulis, 2006; Brun et al., 2006; Gros et al., 2006; Han et al., 2006). Table 6
summarizes the occurrence of acetaminophen and the range concentrations detected in

water samples.

43

TABLE 6. Acetaminophen Occurrence in the Environment (jig/L)

 

 

 

Surface Raw Finished Finished
Water Wastewater Wastewater Dringng Water
<01 (8) 0 - 0.26 (6) 0 - 6.0 (12) Detected at
0.11 (10) 29 - 246 (l) 0 - 0.16 (6) low ppb levels
0 - 0.026 (9) 0.13 - 26 (5) 1.9 (4) (1 l)
0-0.25 (5) 0-4.3 (1)
0.052 - 0.56 (3) 0 - 5.99 (5)
0.007 — 0.066 (7) 0.048 - 0.418 (2)
0.079 — 0.22 (3)

 

 

(1) Gomez et al., 2007 ; (2) Radjenovic et al., 2007; (3) Bound and Voulvoulis, 2006; (4) Brun et al., 2006;
(5) Gros et al., 2006; (6) Han et al., 2006; (7) Wiegel et al., 2004; (8) Stackelberg et al., 2004; (9) Boyd &
Furlong, 2002; (10) Kolpin et al., 2002; (l l) Moll et al., 2001; (12) Temes et al., 1998.

The widespread use of acetaminophen raises the concern of whether or not the
compound persists during treatment of wastewater and drinking water. One of the most
common treatment processes in water utilities is chlorination. Chlorine is a strong
electrophile, and although chlorination targets the inactivation of microorganisms,
chlorine may also react with chemical compounds present in water, such as
acetaminophen and other pharmaceuticals. It has been shown that acetaminophen reacts
with chlorine (Pinkston and Sedlak (2004); Glassmeyer and Shoemaker (2005);
Westerhoff et al. (2005); Bedner and Maccrehan (2006); Gibs et al. (2007).

Bedner and Maccrehan (2006) reported that during chlorination of
acetaminophen, 11 different chlorination products were observed, including the toxic
substances N-acetyl-p-benzoquinone imine (NAPQI) and 1,4-benzoquinone. Although

these toxicants may exist only at very low levels in drinking water and wastewater, their

presence along with multiple other pharmaceuticals deserves further consideration.

The objective of this study is to investigate the effect of pH on oxidation of
acetaminophen and production of 1,4—benzoquinone during chlorination, at three molar
ratios representative of chlorine-to-acetaminophen molar ratios observed in water and
wastewater utilities. The actual concentrations of acetaminophen that were tested are
higher than those observed in drinking water supplies and wastewater effluents and
therefore the chlorine chorine doses were also higher than that those applied in practice in

water and wastewater utilities.

0. Methods

i. Chlorine to Acetaminophen Molar Ratios

Minimum and maximum acetaminophen concentrations detected in raw
wastewater, finished wastewater, and raw drinking water (Table 6) were combined with
typical chlorine doses (5-50 mg/L for wastewater and 0.1-10 for drinking water) in order
to determine target molar ratios for the experiments. Based on the range of calculated
molar ratios in Table 7, the target molar ratios chosen for the experiments were 100,

1,000, and 10,000.

45

TABLE 7. Expected Molar Ratios of Chlorine to Acetaminophen

 

 

 

 

 

Type of Chlorine Acetaminophen Acetaminophen
Water (mg/L) (jg/L) Source MOI“ Ratios
Surface Waters
0.1 0.007 — 0.56 384 — 30, 416
'3" 10 0.007 — 0.56 . 5mm “’3‘” 38, 363 — 3,041,650
Drinking 0 l 0 048 6 0 Frmshed Wastewater
Water ' ' ' ' Finished 3'5 ’ 4’436
10 0.048 — 6.0 W 3,549 - 443, 574
astewatcr
Finished
Finished 5 0.043 —6.0 Wastewater 171’7774;:22221i;§770
Wastewater 50 0.048 - 6.0 Finished ’ ’ ’
Wastewater
Raw 5 0.13 — 246 Raw Wastewater 43 — 81, 891
Wastewater 50 0.13 - 246 Raw Wastewater 433 — 818, 906

 

 

 

 

 

 

 

ii. Chlorination Experiments

The day prior to conducting each batch chlorination experiment, 6 beakers were
filled with 1500 mL distilled water, covered and allowed to reach room temperature
overnight. A chlorine stock was prepared by adding 25 mL of 4-6 % sodium hypochlorite
to 500 mL DI water. This stock was adjusted to a pH of 6, 7.5 or 9 with hydrochloric
acid or sodium hydroxide. The concentration of free chlorine in the stock was
determined using the DPD Ferrous Titrimetric Method (4500-Cl F.). Next, ascorbic acid
was weighed out (1-2 g per 1 L sample) and set aside. A sodium bicarbonate buffer stock
was prepared by adding 4.2 g sodium bicarbonate to 500 mL DI water in a volumetric
flask.

The experiment started with recording the temperature of the water in the 6
beakers. Then, 15 mL of sodium bicarbonate buffer stock was added to each beaker and
the pH was adjusted to 6, 7.5, or 9. A predetermined amount of a 1000 mg/L
acetaminophen solution was added to one beaker while stirring. Next, the appropriate

amount of primary chlorine stock was added to the same beaker to achieve the desired

46

amount of primary chlorine stock was added to the same beaker to achieve the desired
molar ratio of chlorine to acetaminophen. The solution was then transferred to a 1 L
amber bottle, headspace free, and a timer was started. The remaining solution was
divided into two 130-mL amber bottles, headspace free, and set aside. At the end of the
desired contact time, ascorbic acid was added to the 1 L bottle to stop the reaction of
chlorine with acetaminophen. Also, the chlorine residual, temperature, and pH were
measured using the two 130 mL samples at the end of the contact time. The addition of
pharmaceutical and chlorine was repeated for the remaining 5 beakers with contact times
of 0, 3, 10, 30, and 100 minutes, and one replicate contact time randomly chosen for each
batch. The l L samples were stored in a refrigerator at 4 degrees Celsius until analysis on

a LC/MS/MS.

iii. Analytical Methods for Acetaminophen

AcetaminOphen was purchased from Sigma Aldrich, >99% purity. The initial
acetaminophen stock solution was prepared at 1000mg/L in methanol and the stock was
then diluted to 10 mg/L. A standard curve of 100, 200, 500, 1,000, and 2,000 rig/L
(pg/L) was used for measuring acetaminophen concentrations. Samples were analyzed
using liquid chromatography/ tandem mass spectrometry (LC/MS/MS) which a
Shimadzu Prominence Liquid Chromatograph (LC-20AD) with autosampler (SIL-20A)
and an Applied Biosystems API 3200 tandem mass spectrometer. An electrospray
ionization (ESI) interface was used as the ionization source. The retention time on a
Phenomenex Luna C18 (150 x 4.60 mm 3 micron) column was from 6.7 to 7.3 minutes.

For the mobile phases, eluent A was 0.1% formic acid/5% methanol/95% distilled water

47

and eluent B was 01% formic acid/5% distilled water/95% methanol. Isocratic flow was
used at a rate of 0.10 mL/min eluent A, and 0.17 mL/min eluent B. A volume of 10 uL
of sample was injected through the LC/MS/MS system. An ESI (+) MRM scan produced
a precursor ion for acetaminophen at 152 amu and two product ions at 110 and 93 amu
using unit resolution. The tandem MS parameters used for analysis of acetaminophen
were as follows: curtain gas (CUR) of 12 psi, collision gas (CAD) of 2 psi, ionspray
voltage of 5400 V, nebulizer gas (GSl) of 55 psi, auxiliary gas (G82) of 40 psi, and
temperature of probe at 650°C. The standard curves used for quantification of

acetaminophen were linear with R-squared values of 0.995 or greater.

iv. Analytical Methods for 1, 4-Benzoquinone

1, 4-Benzoquinone was purchased from Acros Organics, >98% purity. A 1000
mg/L stock was made by adding 0.100g 1,4-benzoquinone to 100 mL distilled water.
This stock was diluted to 10 mg/L, from which standard curve dilutions were made. The
standard curves consisted of 50, 100, 200, 500, and 1000 ug/L. Samples were analyzed
using liquid chromatography/ tandem mass spectrometry (LC/MS/MS). An atmospheric
pressure chemical ionization (APCI) interface was used as the ionization source. The
retention time on a Phenomenex Luna C18 (150 x 4.60 mm 3 micron) column was from
6.5 to 6.9 minutes. For the mobile phases, eluent A was 0.1% formic acid/5%
methanol/95% distilled water and eluent B was 01% formic acid/ 5% distilled water/95%
methanol. Isocratic flow was used at a rate of 0. 10 mL/min eluent A, and 0.17 mL/min
eluent B. 10 uL of sample was injected through the LC/MS/MS system. An APCI (-)

ion scan produced a precursor ion of 108.1 amu. A product ion for 1,4-benzoquinone

48

could not be identified. The MS parameters used for analysis of benzoquinone were as
follows: curtain gas (CUR) of 10 psi, collision gas (CAD) of -5 psi, nebulizer gas (GSl)
of 40 psi, auxiliary gas of 50 psi, and the temperature of the probe at 675°C. The
standard curves used to quantitate 1,4-benzoquinone were linear with R-squared values

greater than 0.995.

d. Results and Discussion

i. Overall Results

Nine chlorination experiments were performed at five contact times each, at room
temperature (21.8 i 2.2 0C), using buffered water. For each experiment, one contact time
was run in duplicate to confirm the accuracy of the results. The average relative standard
deviation (defined as the standard deviation divided by the mean times 100) in
acetaminophen concentrations between replicate samples was 9.52%. The average
relative standard deviation in 1,4-benzoquinone concentrations between replicate samples
was 2.98%.

The experiments were conducted at pH 6.0, 7.5 and 9.0, and at molar ratios of
chlorine/acetaminophen of 106 d: 6, 1,417 i 285, and 10,454 i 1,430. Each experiment
produced an acetaminophen degradation curve and a 1,4-benzoquinone formation curve.
Examples of acetaminophen and 1,4-benzoquinone curves are shown in Figures 2, 3, and
4. At a free chlorine dose of 100 mg/L and at pH 6.0 with an initial acetaminophen
concentration of 2150 ug/L (molar ratio = 99), acetaminophen was only slightly
degraded, and 1,4-benzoquinone was produced at a slow rate (Figure 2). When the molar

ratio of chlorine-to-acetaminophen was increased to 1,745 and the pH was kept at 6.0,

49

produced in comparison to the 99 molar ratio (Figure 3). When the molar ratio was
further increased to 10,386 and the pH was kept at 6.0, acetaminophen was degraded
rapidly over time and 1,4-benzoquinone was produced within the first three minutes of
chlorination, after which 1,4-benzoquinone began to slowly degrade over time as well
(Figure 4). At the highest chlorine to pharmaceutical molar ratio (Figure 4), more
chlorine was available to attack both acetaminophen and 1,4-benzoquinone over time.
Similar graphs were produced for all conditions. The graphs showed that degradation of
acetaminophen and production of 1,4-benzoquinone were dependent on pH and chlorine-
to-acetaminophen molar ratios.

Rate constants were estimated for the degradation of acetaminophen and the

formation of 1,4-benzoquinone. For example, in Figure 3, the rate constant for

acetamin0phen degradation is -0.044 min.1 and the rate of formation of 1,4-
benzoquinone (0-30 min) is 0.077 min-1 using an exponential equation. In Figure 4, the
rate constant for degradation of acetaminophen increased to -0.595 min.1 and the rate of

formation of 1,4—benzoquinone (3-100 min) is 0.045 min'l. The rate constants of

acetaminophen degradation were not equal to the rate constants of 1,4—benzoquinone
formation signifying the possible formations of other intermediate products during the
chlorination of acetaminophen. This study the focuses on the formation of 1,4-

benzoquinone because it is known to be toxic.

50

 

NEW
D

 

 

 

1000 ‘
5(1) - benoquinone
OED-“Ci r 3--.-f...........‘ ........... I .......... '1?
0 20 40 60 80 100
Tim (min)

Figure 2. Degradation of acetaminophen and formation of 1,4-benzoquinone at pH 6.0
over a lOO-minute contact time, with a free chlorine dose of 100 mg/L, initial
acetaminophen concentration of 2150 ug/L (molar ratio = 99).

2500

 

Conc., ug/l.

 

 

 

 

Figure 3. Degradation of acetaminophen and formation of 1,4—benzoquinone at pH 6.0
over a 100-minute contact time, with a free chlorine dose of 1,000 mg/L, initial
acetaminophen concentration of 1670 rig/L (molar ratio = 1,275).

51

 

 

 

 

Tim: (min)

Figure 4. Degradation of acetaminophen and formation of 1,4-benzoquinone at pH 6.0
over a lOO-minute contact time, with a free chlorine dose of 10,000 mg/L, initial
acetaminophen concentration of 2050 ug/L (molar ratio = 10,386).

ii. Efi'ect of pH on Acetaminophen Degradation
Figures 5, 6, and 7 show the effect of pH on acetaminophen degradation at molar

ratios of chlorine/acetaminophen of 106 i 6, 1,417 1 285, and 10,454 r 1,430

respectively. In the figures, C/Co is the remaining concentration of acetaminophen after

chlorination divided by the initial acetaminophen concentration. At the beginning of the

experiments (0 min contact time), C/Co equals one since no reaction has occurred. As

acetaminophen is allowed to react with chlorine, C/Co decreases since the concentration

of acetaminophen decreases. Figure 5 reveals that at a molar ratio of 106 t 6, the
acetaminophen degradation rate was lowest at pH 6.0 and highest at pH 9.0. Estimated

rate constants for the degradation of acetaminophen by free chlorine at an average molar

ratio of 106 i 6 are -0.00279, -0.0257, and -0.0577 min.1 for pH 6.0, 7.5, and 9.0,

52

respectively. For pH 6.0, 7.5, and 9.0, the percent degradation of acetaminophen after

100 minutes was 24.7, 87.0, and 94.8%, respectively.

 

 

A
v1.1

 

 

Time (nrin)

Figure 5. Effect of pH and contact time on acetaminophen for an average initial
acetaminophen concentration (Co) of 2013.33 1 119.30 [Lg/L and 100 mg/L free chlorine
dose (chlorine/acetaminophen molar ratio: 106 i 6).

 

C/Co

 

 

 

 

Figure 6. Effect of pH and contact time on acetaminophen degradation for an average
initial acetaminophen concentration (Co) of 1540 1- 278. 75 jug/L and 1,000 mg/L free
chlorine dose (chlorine/acetaminophen molar ratio = 1,417 i 285).

53

 

 

 

 

 

+pH6.0

--13--pH7.5

-~O--pH9.0
l ‘ I l l l I I 1
0102030405060708090100

Titre (nin)

Figure 7. Effect of pH and contact time on acetaminophen degradation at an average
initial acetaminophen concentration (Co) of 2036.67 :1: 270.25 [Lg/L and 10,000mg/L free
chlorine dose (chlorine/acetaminophen molar ratio = 9789 i 1430).

The same trends are observed for molar ratios of 1,417 i 285 and 9,789 1 1,430

(Figures 5 and 6). For an average molar ratio of 1,417 :t 285, the rate constants for

acetaminophen degradation are -0.0435, -0.400, and -l.23 min.1 for pH 6.0, 7.5, and 9.0,

respectively. The percent degradation of acetaminophen after 100 minutes was 99.2% for

pH 6.0 and 100% for pH 7.5 and 9.0. For an average molar ratio of 9,789 1 1,430, the

estimated rate constant for pH 6.0 is -0.791 min.1 and -O.333 min"1 for pH 7.5 and 9.0.

The percent degradation of acetaminophen after 100 minutes was 100% at all pH levels.

Free chlorine, which consists of hypochlorous acid (HOCl) and hypochlorite ion

(OCl-), has a pka of 7.5. Since the stronger oxidant, HOCl, is the dominant chlorine

species at pH less than 7.5, it was expected that the acetaminophen degradation rate

would be greater at pH 6.0 than at pH 9.0. At pH 9.0, the dominant chlorine species is

54

OCl', a weaker oxidant, so it was surprising to observe acetaminophen react to a greater

extent at a higher pH. The results suggest that the acetaminophen molecule is more

reactive with oc1' than with HOC].

Acetaminophen contains a phenolic functional group as well as an amide group.
Amides do not react rapidly with free chlorine so the main site of the reaction is most
likely the phenol. Pinkston and Sedlak (2004) stated that the main site of the reaction
between acetaminophen and free chlorine would likely be between HOCl and the

phenolic functional group. Depending on pH, two forms of acetaminophen are present,

the protonated form (ROH) and the phenolate form (R0)

The ionization constant of acetaminophen is pka= 9.5 (Dasmalchi et al., 1995).

According to Pinkston and Sedlak (2004) in reactions with free chlorine, phenols tend to

exhibit a maximum reaction rate between the pka of free chlorine (pka= 7.5) and that of

the phenol. When determining rate constants, Pinkston and Sedlak did not consider
reactions with OCl- because previous research proved substituted phenols do not react
fast enough with hypochlorite. However, the results of this study indicate that
acetaminophen reacts with hypochlorite at a faster rate than it reacts with hypochlorous

acid.

iii. Effect of pH on Benzoquinone Production
Figures 8, 9, and 10 show the effect of pH on 1,4-benzoquinone formation at

molar ratios of chlorine/acetaminophen of 106 1- 6, 1,417 i 285, and 10,454 i 1,430

respectively. In these figures, CB/Co is the 1,4-benzoquinone molar concentration formed

55

from the chlorination of acetaminophen divided by the initial acetaminophen molar
concentration. For an average molar ratio of 106 1 6, 1,4-benzoquinone was produced up
to 11.8, 50.5, and 24.4 % of the initial acetaminophen concentration for pH 6.0, 7.5, and
9.0, respectively. For an average molar ratio of 1,417 :1: 285, 1,4-benzoquinone was
produced up to 68.7, 54.3, and 33.1 % of the initial acetaminophen concentration for pH
6.0, 7.5, and 9.0, respectively. For an average molar ratio of 9,789 :9: 1,430, 1,4-
benzoquinone was only produced for pH 6.0 up to 38% of the initial acetaminophen
concentration.

The effect of pH on 1,4-benzoquinone formation is more complex than the effect
of pH on acetaminophen degradation since1,4-benzoquinone also reacts with free
chlorine over time, resulting in its own degradation. Bedner and Maccrehan (2006)
observed increasing 1,4-benzoquinone production over a 60 minute contact time with
free chlorine at pH 7, reaching a maximum concentration of 2.5 umol/L, or 25% of the
initial acetaminophen concentration. This is lower than the 50% 1,4-benzoquinone
production observed in this study for pH 7.5 and average molar ratio 106 i 6. However,
their study involved a much lower molar ratio of chlorine to acetaminophen (5.7) than

those used in this study (minimum of 100).

56

 

CB/CO

 

 

 

 

 

40506070 8090100

Tint: (min)

Figure 8. Effect of pH and contact time on 1,4-benzoqiunone formation (CB) at an

average initial acetaminophen concentration (C0) of 2013.33 1 119.30 jig/L and 100
mg/L free chlorine dose (chlorine/acetaminophen molar ratio = 106 :1: 6).

57

 

1.0

+pH 6.0
--B-- pH 7.5

 

 

 

 

Figure 9. Effect of pH and contact time on 1,4-benzoquinone formation (CB) at an

average initial acetaminophen concentration (C0) of 1540 :t 278. 75 rig/L and 1,000
mg/L free chlorine dose (chlorine/acetaminophen molar ratio = 1,417 :1: 285).

58

 

 

 

 

 

 

1.0
08 _ +pH6.0
--13--pH7.5
o 0.6 _ ---<>-- pH9.0
U
E
U
r l l I r fl

 

0 10 20 30 40 50 60 70 80 90 100

True (min)

Figure 10. Effect of pH and contact time on 1,4-benzoquinone formation (CB) at an

average initial acetaminophen concentration (Co) of 2036.67 1 270. 25 ug/L and 10,000
mg/L free chlorine dose (chlorine/acetaminophen molar ratio = 9789 1 1430).

iv. Effect of Molar Ratios

At all pH values the rate of degradation of acetaminophen was slowest at a molar
ratio of about 100, and highest at a molar ratio of about 10,000. At pH 6.0 and a molar
ratio of 100 acetaminophen is removed by only 25% after a 100 minute contact time,
while at a molar ratio of 10,000, acetaminophen is removed entirely after 100 minutes. A
similar trend was observed for pH 7.5 and 9.0.

At pH 6.0, 1,4-benzoquinone formation was greatest at the intermediate molar
ratio of about 1,000. At pH 7.5 and 9.0, 1,4-benzoquinone formation increased steadily

at the molar ratio of about 100. At the molar ratio of about 1,000, 1,4—benzoquinone was

59

produced rapidly after 3 minutes, but then quickly dropped. For the molar ratio of about

10,000, 1,4-benzoquinone was detected only at pH 6.0.

v. Comparison with Literature Data

The results of this study are compared with the results from previous
acetaminophen chlorination work (Gibs et al., 2007; Bedner and Maccrehan, 2006;
Glassmeyer and Shoemaker, 2005; Westerhoff et al., 2005; Pinkston and Sedlak, 2004).
Table 8 summarizes experimental conditions and basic conclusions obtained from
published studies and this study. Table 8 presents chemical dose, acetaminophen
concentration, molar ratio, pH, and contact time conditions. All studies concluded that
acetaminophen was degraded when reacting with chlorine.

One study identified chlorinated byproducts. The transformation of
acetaminophen during chlorination into the toxic byproduct, 1,4-benzoquinone, first
observed by Bedner and Maccrehan (2006), was confirmed in this study, and the effects
of pH and molar ratios were further explored. Also, molar ratio conditions representing
the whole range of ratios observed in water and wastewater engineering practices are
only tested in this study.

Pinkston and Sedlak (2004) found that acetaminophen was significantly
transformed by free chlorine with a half-life of 5.2 minutes. Rate constants were studied
over a similar pH range, 5 - 10, and at a molar ratio of chlorine/acetaminophen of 30,
which is lower than the molar ratios used in this study. Glassmeyer and Shoemaker
(2005) found that acetaminophen showed definite signs of chlorination when

acetaminophen-spiked water was dosed with 28.75 mg/L of free chlorine and allowed to

60

react for 48 hours. The primary product of the reaction was a singly chlorinated
acetaminophen molecule, but other chlorinated products were also observed, and some of
the acetaminophen did not react. Glassmeyer and Shoemaker (2005) did not attempt to

test the effects of pH or molar ratio.

61

 

 

TABLE 8. Summary of Results and Comparison with Literature Data Involving the Reaction of Acetaminophen with Sodium Hypochlorite

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Chemical Acetaminophen Molar Contact Acetaminophen 1,4-Benzoquinone Reference
Dose Concentration Ratio pH Time Degradation Formation
1,4-benzoquinone accounted for 25% of the initial
57 umol/L 1 umollL 57 7.0 2—90 min 88% of acetaminophen (10 pmol/L initial) was acetaminophen Bedner & Maccrehan,
10 umol/L 5.7 transformed in 1 hour concentration after 1 hour. (2006)
Glassmeyer &
405 umol/L NA NA NA 48 hrs Definite signs of chlorination N/A Shoemaker, (2005)
53.5 pmol/L 1.98 pmol/L 27 5.5 24 hr. >95% oxidized N/A Westerhoff et al., (2005)
49.3 pmol/L 25
600 umol/L 20 nmol/L 30 5—10 5 days Acetaminophen was significatly transformed N/A Pinkston and Sedlak,
(2004)
Completely degraded after 1 day in contact with
16.9 umol/L 0.0033 umol/L 5121 8 1—10 days free chlorine N/A Gibs et al., (2007)
Benzoquinone was produced up to 1 1.8% of the initial
1,408.45 Acetaminophen was removed by 6.51 and acetaminophen after 100 min, no amount was detected
umol/L 14.22 pmol/L 99 6 0—100 min 24.7% after 30 and 100 min, respectively. before 100 min This study
1,408.45 Acetaminophen was removed by 56.4 and Benzoquinone was produced up to 41.1 and 50.5% of the
umol/L 12.97 pmol/L 109 75 0—100 min 87.0% after 30 and 100 min, respectively. initial acetaminophen after 30 and 100 min, respectively. This Study
1,408.45 Acetaminophen was removed by 80.0 and Benzoquinone was produced up to 24.4 and 17.7% ofthe
umol/L 12.77 umol/L 1 10 9 0—100 min 94.8% after 30 and 100 min, respectively. initial acetaminophen after 30 and 100 min, respectively. This study
14,0845 Acetaminophen was removed by 76.8 and Benzoquinone was produced up to 68.7 and 57.3% ofthe
umol/L 1 105 pmol/L 1,275 6 0—100 min 99.2% after 30 and 100 min, respectively. initial acetaminophen after 30 and 100 min, respectively. This study
Benzoquinone was produced up to 54.3% of the initial
14,084, 5 Acetaminophen was removed by 100% after 30 acetaminophen after 3 min, no amount was detected after
umol/L 807 pmol/L 1,745 7.5 0—100 min and 100 min. 100 min. This study
Benzoquinone was produced up to 33.1% of the initial '
14,0845 Acetaminophen was removed by 100% after 30 acetaminophen after 3 min., no amount was detected after
umoUL 1 1.44 pmol/L 1,231 9 0-100 min and 100 min. 100 min. This study
Benzoquinone was produced up to 38.0% of the initial
1407845 Acetaminophen was removed by 100% after 30 acetaminophen after 3 min., no amount was detected after
umol/L 13.56 pmol/L 10,386 6 0-100 min and 100 min. 100 min. This study
140,845 Acetaminophen was removed by 100% after 30
umol/L 15.21 umol/L 9,257 7.5 0—100 min and 100 min. Benzoquinone was not detected in any of the samples. This study
140,845 Acetaminophen was removed by 100% after 30
umol/L 1 1.51 pmol/L 12,237 9 0-100 min and 100 min. Benzoquinone was not detected in any of the samples. This study

 

 

62

 

 

 

Westerhoff et a1. (2005) added 3.5—3.8 mg/L free chlorine to surface water spiked
with 1.98 umol/L of acetaminOphen at pH 5.5. After a 24 hour contact time, greater than
95% of acetaminophen was oxidized, indicating a high degree of reactivity with chlorine.
The effects of pH and molar ratio were not tested in this study either, but different water
matrices were studied.

Bedner and Maccrehan (2006) monitored the reaction of free chlorine with
acetaminophen over 90 minutes at pH 7 and at molar ratios of 5.7 and 57. At an initial
concentration of 1 umol/L, acetaminophen reacted to a greater extent than at an initial
concentration of 10 umol/L due to the greater molar ratio of free chlorine to
acetaminophen (molar ratio of 57) at the lower initial concentration. Similarly, in this
study degradation of acetaminophen increased with increasing chlorine-to-acetaminophen
molar ratios. In this study a wide range of molar ratios was evaluated. Bedner and
Maccrehan also found that acetaminophen exhibited a high degree of reactivity with
hypochlorite at a neutral pH.

Gibs et al. (2007) analyzed chlorinated samples of drinking water over 10 days
and found that acetaminophen reacted completely with residual chlorine within one day

at pH 8 and at a molar ratio of free chlorine to acetaminophen of 5121.

e. Conclusions

This study found that acetaminophen was degraded and transformed by free
chlorine. The rate of degradation was affected by pH and chlorine/acetaminophen molar
ratio. The highest degradation rates were observed at pH 9.0, and the lowest degradation

rates were observed at pH 6.0. Acetaminophen degradation was also greatest at molar

63

ratios of approximately 10,000, and lowest at molar ratios of approximately 100.
Acetaminophen was transformed by free chlorine to form the toxic byproduct, 1,4 —
benzoquinone. The formation of 1,4 — benzoquinone was also affected by pH, and
chlorine/acetaminophen molar ratio. Depending on the applied molar ratios, up to 68.7 %
of acetaminophen may be converted into the toxic byproduct, 1,4— benzoquinone. The
effects of varying pH and chlorine/acetaminophen molar ratios are significant in drinking
water and wastewater treatment plants since those parameters can be slightly modified by
water utility professionals. Water chlorination conditions affect the potential of
acetaminophen transformation to toxic 1,4-benzoquinone and further affect the potential

of human exposure to toxic substances.

Additional Experiments and Continuing Research
a. Additional Experiments

For the additional 14 experiments, initial acetaminophen concentrations of 200,
1,000, and 2,000ug/L were tested. These concentrations are not typically observed in
environmental waters, but were necessary to quantify the formation of 1,4-benzoquinone.
Therefore, chlorine doses used in the experiments are also greater than those typical of
water treatment plants. To achieve the desired molar ratios of chlorine to acetaminophen,
chlorine doses of 10, 50, 100, 500, 1,000, and 10,000 were applied. The pH levels tested
were 6, 7.5, 9, and 11. Table 9 shows the conditions tested in each experiment.
Chlorination experiments were conducted in buffered, distilled water using 1.5 L beakers

(method details are given in Xagoraraki et al., 2008).

64

TABLE 9. Conditions tested in this study.

 

 

 

 

 

Initial Chlorine Molar Ratio
Acetaminophen Dose Chlorine/Acetaminophen pH
Concentration (rig/L) (mg/L)
10 99.5 1 3.04 6, 7.5, 9
188 1 42 100 995 1 121 6, 7.5, 9
1,000 15,705 1 489 6, 7.5, 9
50 134 7.5
908 1 158 500 1,044 6
100 107 1 5 ll (6, 7.5, 9)
1,950 1 352 1,000 1,325 1 285 11 (6, 7.5, 9)
10,000 9,722 1 1,727 11 (6, 7.5, 9)

 

 

 

 

 

Note: pH levels in parentheses were tested in previous study, Xagoraraki et al., 2008.

b. Analytical Procedures

Samples were analyzed using liquid chromatography/tandem mass spectrometry
(LC/MS/MS). The instrumentation consisted of a Shimadzu Prominence Liquid
Chromatograph (LC-20AD) with autosampler (SIL-20A) and an Applied Biosystems API
3200 tandem mass spectrometer. A Phenomenex Luna C18 (150 x 4.60 mm 3 micron)

column was used. Refer to previous publication for specific analytical method details

(Xagoraraki et al., 2008).

0. Determination of pka

Knowledge of the acid dissociation constant, or pka, of acetaminophen is
important for predicting speciation at different pH levels. Since previous studies on the
determination of acetaminophen pk,l could not be obtained, the reported pka of 9.38 was

verified by spectrophotometric method. This method consisted of preparing two 100
ppm solutions of acetaminophen in distilled water. One solution was adjusted to pH 3

with hydrochloric acid and the other solution was adjusted to pH 12 with sodium

65

hydroxide. Each solution was diluted to 1, 5, 10, and 25 ppm, and absorbance
measurements for each dilution were obtained at wavelengths (A) 241 and 255.5 nm. A

plot of absorbance versus acetaminophen concentration produced linear absorbance

curves for each pH, and the slopes of the lines (a1, a2) were used to calculate ratios of

ROH (protonated acetaminophen) and R0- (unprotonated acetaminophen) at known pH

values. The mixture composition of ROH and RO- was calculated from measured

absorbancies (A1, A2) at 711:241 nm and 12:2555 nm,

A1 = 31(ROH) + a1(RO-) (1)

A2 = a2(ROH) + 32(R0-) (2)

For acetaminophen solutions ranging from pH 7-11, a pH 9.47 solution resulted in

approximately equal amounts of ROH and RO-, verifying the reported pka of 9.38

(Dasmalchi et al., 1995).

d. Overall Results

Fourteen chlorination experiments were performed at five contact times each, at
room temperature using buffered, distilled water. Each experiment included a duplicate
contact time to confirm the accuracy of the results. The average relative standard
deviation, defined as the standard deviation divided by the mean times 100, in
acetaminophen concentrations between replicate samples was 10.9%. The average

relative standard deviation in 1,4-benzoquinone concentrations between replicate samples

66

was 28.3%. This value is high due to the low concentrations of 1,4-benzoquinone
produced and the poor sensitivity of the instrumentation at low concentrations.
Furthermore, only five of the fourteen experiments could be analyzed for 1,4-
benzoquinone due to the low initial acetaminophen concentration of the other nine
experiments. For initial and final temperature, the average relative standard deviation
between replicates was 0.2 and 0.3%, respectively. For initial and final pH, the average
relative standard deviation between replicates was 0.5 and 0.7%, respectively. The
average relative standard deviation for final measured chlorine in replicates was 16%,
indicating some variability in the rate of chlorine consumption during the experiments.

One experiment was performed without the addition of acetaminophen at pH 7.5
and chlorine dose 1,000 mg/L. This was to ensure that chlorine was not reacting with any
other compounds than acetaminophen during the original experiments. Chlorine was
added to each beaker at a dose of 1,000 mg/L, and the pH was adjusted to 7.5. Samples
were collected after 0, 3, 10, 30, and 100 minutes and analyzed for residual chlorine
concentration. The experiment where no acetaminophen was added resulted in no
reduction of residual chlorine, indicating that chlorine was not reacting with any other
compounds.

The experiments were conducted at pH 6.0, 7.5, 9.0, and 11.0. At pH 11.0, molar
ratios of 109, 1,049, and 8,035 were tested only at an initial acetaminophen concentration
of 2,000 pig/L, to supplement previous experiments (see Xagoraraki et al., 2008). At an
initial acetaminophen concentration of 200 [Lg/L, molar ratios of 99.5 1 3, 995 1 121, and
15,705 1 489 were tested for pH 6.0, 7.5, and 9.0. Two experiments were performed at

an initial acetaminophen concentration of 1,000 pig/L. One was performed at pH 7.5,

67

molar ratio 134, and the other at pH 6.0, molar ratio 1,044. Each experiment produced an
acetaminophen degradation curve and a 1,4-benzoquinone formation/degradation curve.

1,4-benzoquinone is targeted in this study since it is a toxic compound.

e. Efiect of pH

Figure 11a shows the effect of pH on acetaminophen degradation at an average

molar ratio of chlorine to acetaminophen of 107 1 5. In the figure, C/Co is the
concentration of acetaminophen after each contact time with free chlorine divided by the
initial acetaminophen concentration. Before free chlorine is added, C/Co equals one

since no reaction has occurred. Figure 11a shows that acetaminophen degradation is
faster at pH 7.5 and 9.0 than at pH 6.0 and 11.0.

Figure 11b shows the effect of pH on detected 1,4-benzoquinone concentration at

an average molar ratio of chlorine to acetaminophen of 107 1 5. In the figure, CB/Co is

the concentration of 1,4-benzoquinone after each contact time of acetaminophen with free

chlorine divided by the initial acetaminophen concentration. Before free chlorine is

added CB/Co equals zero since no 1,4-benzoquinone has formed. Figure .1 1b shows that

the highest 1,4-benzoquinone concentration is detected at pH 9.

68

 

0.90 4 \ c
0.80 — ~
0.70 -
0.60 —
g 0.50 -
0.40 .
0.30 -
0.20 —

 

0.10 -

 

 

 

Cam I I I I I I I T I 1

Trrne (trim)

(a)

 

1.00
0.90 a + pH 6.0
0.80 J -I— pH 7.5
0.70 A + pH 9.0

050 a +pH 11.0

g 0.50 4

0.40 r

l.

 

 

 

 

 

 

 

 

 

Tim: (nin)

0))

FIGURE 11. (a) Effect of pH on acetaminophen degradation and (b) 1,4-benzoquinone
formation for an average initial acetaminophen concentration of 13.2 1 0.68 umol/L
(average molar ratio 107 1 5).

69

Different species of free chlorine and acetaminophen are expected to be dominant

in solution at different pH levels. Free chlorine, which consists of hypochlorous acid

(HOCl) and hypochlorite ion (OCl-), has a pka of 7.5. Therefore, at pH 6.0 it is assumed

that majority of chlorine is in HOCl form, and at pH 9.0 and 11.0 the majority of chlorine

is in OCl- form. Similarly, acetaminophen consists of a protonated form (ROH) and an

unprotonated form (RO-) that are in equilibrium at the pka of 9.4. Therefore, it is

assumed that at pH 6.0 and 7.5 the majority of acetaminophen is in the ROH form, and at
pH 11.0 the majority of acetaminophen is in the RO- form. Table 10 shows which
species of acetaminophen and chlorine are dominant at the pH levels studied. The lowest
extent of acetaminophen degradation is observed at pH 6.0, when HOCl reacts with
ROH, and at pH 11, when OCl- reacts with RO-. The highest degradation is observed at
pH 7.5 and pH 9.0, when different forms of acetaminophen and chlorine coexist. Similar

results are observed at all molar ratios.

TABLE 10. Acetaminophen and chlorine species composition at each pH tested.

 

52H 6.0 7.5 9.0 11.0
Dominant
Chlorine Species HOCI (HOCI + 00-) OCl- ocr-
Dormmm ROH ROH (ROH + RO-) RO-

Acetaminophen Species

Pseudo first-order

l 1 I

0.044 min‘1 0.4 min‘ 1.23 min“ 0.06 min'

 

Reaction Rate

*
Reaction rates are for experiments at molar ratio 1,325 1 285 and initial acetaminophen
concentration of 1,663 1 334 [Lg/L.

70

f. Effect of Initial Concentration

Figures 12 and 13 show the effect of initial acetaminophen concentration on
acetaminophen degradation and 1,4-benzoquinone formation over contact time with free
chlorine. Figures 12a and 13a show that slightly less degradation of acetaminophen
occurs at the lower initial concentration of 1,000 [lg/L than at 2,000 rig/L. Figures 12b

and 13b show that more 1,4-benzoquinone is formed at the higher initial acetaminophen
concentration of 2,000 [Lg/L than at 1,000 [Lg/L.

71

 

 

 

 

 

 

 

 

 

 

 

1.2

”l +2000uglL

—B-1000ug/L

CB/CO

 

 

 

 

 

 

 

 

 

 

(1))

FIGURE 12. (a) Effect of initial acetaminophen concentration on acetaminophen
degradation and (b) 1,4-benzoquinone formation (average molar ratio 121 1 18, pH 7.5).

72

 

 

 

 

 

 

 

 

 

 

 

1.2

1.1— +2000ug/L
1..

0.9 “ -B—1000ug/L

0.8 i

CB/CO

 

 

 

 

 

 

 

 

0))

FIGURE 13. (3) Effect of initial acetaminophen concentration on acetaminophen
degradation and (b) 1,4-benzoquinone formation (average molar ratio 1,159 1 164, pH
6.0).

73

 

g. Efiect of Molar Ratio and Contact Time

Acetaminophen underwent greater degradation with increasing oxidant dose and
increasing contact time with free chlorine. Figure 14a shows the effect of molar ratio of
chlorine to acetaminophen on acetaminophen degradation at pH 6.0. Degradation
increases with increasing molar ratio. Similar trends are observed at all pH values.

Figure 14b shows the effect of molar ratio of chlorine to acetaminophen on 1,4-
benzoquinone formation at pH 6.0. At the lower molar ratio of 109, 1,4-benzoquinone
formed after ten minutes then remained at a steady concentration. When the molar ratio
increased to 1,049, 1,4-benzoquinone peaked at three minutes then began to degrade with
longer contact time with free chlorine. At the highest molar ratio of 8,035, a small
amount of 1,4-benzoquinone formed after three minutes, after which none could be

detected.

74

 

 

 

   
 

 

 

 

 

 

 

 

 

 

1.0
0.9 a
0.8 a
0.7 ~
+MR99
0‘5 1 -B—MR 1.275
g 0.5 ~ —A—MR10.386
0.4 ~
0.3 -
0.2 a
0.1 -
0.0
0 10 20 30 40 50 60 70 80 90 100
The (nin)
(a)
1.0
+MR 109
-S—MR1,275
0.8 ~
-A—MR9,257

 

 

 

 

 

 

 

 

0))
FIGURE 14. Effect of molar ratio on (a) acetaminophen degradation and (b) 1,4-

benzoquinone formation for an average initial acetaminophen concentration (Co) of 12.9
1 1.7 umol/L (pH 6).

75

 

The greatest amount of 1,4-benzoquinone formed at pH 6, molar ratio 1,275, with
69% of acetaminophen converted to 1,4-benzoquinone (initial acetaminophen
concentration 2,000ug/L). At pH 7.5, molar ratio 109 and molar ratio 1,745, 50.5% and
54.3% of acetaminophen was converted to 1,4-benzoquinone, respectively. Significant
amounts of 1,4-benzoquinone also formed at pH 6.0, molar ratio 10,386, with 38% of
acetaminophen converted to 1,4-benzoquinone, and at pH 9, molar ratio 1,231, where

33.1% was converted.

h. Continuing Research

The experimental results are being used to evaluate the reaction rate order of the
reaction between acetaminophen and chlorine. The reaction mechanisms are being
investigated by calculating reaction rates from a set of equations. The equations are
being evaluated using a computer modeling program. Also, these reaction rates are being
compared to those reported in previous literature for both chlorination and oxidation
studies. Preliminary results suggest that the reaction is second-order, and that ozonation

is expected to oxidize acetaminophen more efficiently than chlorination.

Conclusions

The concern over pharmaceuticals in drinking water is spurring research at all
levels of drinking water treatment processes. The above reported experiments evaluated
the chlorination of acetaminophen by free chlorine. Results indicate at which conditions
greatest degradation of acetaminophen will occur, as well as the conditions at which the

greatest amount of 1,4-benzoquinone may be produced as a byproduct. Although the

76

experiments only evaluated one compound, the research may be applied to other
compounds in terms of the effect pH, molar ratios, and initial pharmaceutical
concentration have on a pharmaceutical. Furthermore, the continuing research on
reaction rates and mechanisms may be used as a model for investigating other

pharmaceuticals.

77

 

*
Appendix 1. Pharmaceutical Chemical Properties

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Pharmaceutical Formula ”weir/gym“ Log Kow pKa
Acetaminophen C8H9N02 151.2 0.46 9.4
Atenolol C 14H22N203 266.3 0.5 -
Carbadox C17H15N306 262.2 1.3 -
Carbamazepine C 15H 12N20 236.3 2.5 (2
Diazepam C16H13C1N20 284.8 2.9 3.4
Diclofenac C 14Hl 1C12N02 296.2 3.9 4.2
Dilantin C15H12N202 252.3 2.2 8.3
Erythromycin C37H67N013 734 3.1 8.8
Estradiol C18H2402 272.2 4.0 10.4
Estriol C 18H2402 288 .4 2.5 10.4
Estrone C18H2202 270.2 3.1 10.3
Ethinylestradiol C2OH2402 296.2 4.7 10.5
Fluoxetine C 17H18F3N0 309.3 4.6 -
Gemfibrozil C 15H2203 250.2 3.4 9.4
Ibuprofen C 13H 1 802 206.1 3.6 4.5
Indometacine C 19H16C1NO4 357.8 3.4 4.5
Metoprolol C 15H25N03 267 .4 1.6 -
Naproxen C 14H 1403 230.1 2.8 4.2
Progesterone C21H3002 314.2 3.9 -
Propanolol C 1 6H2 1 N02 259.3 3 -
Sulfamerazine C1 1H12N4OZS 264.3 0.14 -
Sulfamethazine C12H14N402S 278.3 0.89 -
Sulfamethoxazole ClOH 1 1N303S 253.1 2.1 2.1
Sulfathiazole C9H10N402S 270.3 0.9 -
Trimethoprim C 14H 1 8N403 290.1 0.6 6.3

 

*
Chemical data obtained from www.DrugBank.com (W ishart et al., 2006)

78

 

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