Determining Active Oxidant 

Species Reacting with 

Organophosphate Pesticides 

in Chlorinated Drinking Water

R E S E A R C H A N D D E V E L O P M E N T
EPA/600/R-06/103 

September 2006 

Determining Active Oxidant Species Reacting with 

Organophosphate Pesticides in Chlorinated Drinking 

Water 

By 

Stephen E. Duirk, Daniel P. Cherney†, Christopher J. TarrI, and Timothy W. Collette 

National Exposure Research Laboratory 

Ecosystems Research Division 

Athens, GA 

†

National Research Council 

I

Student Services Authority 

U.S. Environmental Protection Agency 

Office of Research and Development 

Washington D.C. 20460 
NOTICE 

This work has been subject to external peer and administrative review, and has been approved 

for publication as an EPA document. Mention of trade names or commercial products does not 

constitute endorsement or recommendation for use. 

ii 
ABSTRACT 

Chlorpyrifos (CP) is an organophosphate (OP) pesticide that was used as a model 

compound to investigate the transformation of OP pesticides at low pH and in the presence of 

bromide and natural organic matter (NOM) under drinking water treatment conditions. Raman 

spectroscopy was used to determine which active chlorine species was responsible for the rapid 

oxidation of CP below neutral pH. Over the pH range of 2-11, three active chlorine species were 

identified with Raman bands of 725, 711, and 538 cm-1. The first two bands were easily 

assigned to HOCl and OCl- respectively. The 538 cm-1 band was identified as Cl2(aq) after 

bubbling chlorine gas through phosphate buffered water at pH 2. Either at low pH or in water 

treatment plants that use direct injection of Cl2(g), molecular chlorine rapidly reacts with CP 

transforming it to chlorpyrifos oxon (CPO). In the presence of bromide, the loss of CP was 

found to be accelerated when aqueous chlorine was added. It was found that bromide acts as a 

catalyst in the oxidation of CP to CPO via the formation of hypobromous acid (HOBr) at 

concentrations relevant to drinking water treatment. Also, NOM was found to not inhibit the 

oxidation of CP to CPO in the presence of free chlorine under the experimental conditions in this 

study. A previously developed screening-level model used to predict OP pesticide degradation 

pathways was modified to model the effects of both bromide and NOM. The model adequately 

described the oxidation of CP to CPO by free chlorine as well as the stability of CPO over the 

course of the experiment. This work demonstrated the applicability of screening-level models to 

predict the fate of OP pesticides under drinking water treatment conditions. 

iii 
ACKNOWLEDGEMENT 

The authors would like to thank Jimmy Avants his technical assistance. Also, we would like to 

thank Dr. Wayne Garrison, Dr. Jackson Ellington, and Dr. John Kenneke for their consultation 

and expertise. 

iv 
TABLE OF CONTENTS 

ABSTRACT...................................................................................................................................iii 

ACKNOWLEDGEMENT ............................................................................................................. iv

LIST OF TABLES......................................................................................................................... vi 

LIST OF FIGURES ...................................................................................................................... vii 

EXECUTIVE SUMMARY ............................................................................................................ x 

1 INTRODUCTION .................................................................................................................. 1

2 EXPERIMENTAL PROCEDURES....................................................................................... 7 

2.1 Materials ......................................................................................................................... 7 

2.2 Methods........................................................................................................................... 8 

2.2.1 Spectroscopic Studies of Active Chlorine Species ................................................. 8 

2.2.2 Chlorpyrifos Kinetic Experiments with Bromide and NOM................................ 10 

2.3 Instrumentation ............................................................................................................. 11 

3 RESULTS AND DISCUSSION........................................................................................... 15 

3.1 Spectroscopic Studies of Active Chlorine Species ....................................................... 15 

3.2 Loss of CP in the Presence of Bromide and Free Chlorine .......................................... 24 

3.3 Loss of CP in the Presence of Free Chlorine and NOM............................................... 29 

4 CONCLUSIONS................................................................................................................... 32 

5 REFERENCES ..................................................................................................................... 34 

TABLES ....................................................................................................................................... 38 

FIGURES...................................................................................................................................... 42 

v 
LIST OF TABLES 

Table 1 NOM and source water characteristics for the ACC water collected prior to 

chlorination. ............................................................................................................. 39 

Table 2 The slopes of kobs vs. [Br-] for each pH and [HOCl] concentration......................... 40 

Table 3 Stoichiometric equations and rate coefficients used in the chlorpyrifos degradation 

pathway model......................................................................................................... 41 

vi 
LIST OF FIGURES 

Figure 1 Observed first order loss rate coefficients for chlorpyrifos over the pH range of 2-11 

in the presence of ECS. [CP]o = 0.5 μM, [Buffer]T = 10 mM, Temperature = 

25±1°C, and [HOCl]T = 10 μM. .............................................................................. 43 

Figure 2 (A.) Speciation plot of the 26.2 mM commercial free chlorine solution (the ECS 

solution) generated by monitoring the normalized areas of the Raman bands from 

OCl­ at 711 cm-1 (>), HOCl at 725 cm-1() and the low pH species at 538 cm-1 

(‚). (B.) Background-subtracted Raman spectra of ECS solutions. The OCl­ band 

is observed in the top two spectra and the HOCl band is observed in the bottom 

three spectra. An additional band for the low pH form appears as the pH is lowered 

below 5..................................................................................................................... 44 

Figure 3 Speciation plots of commercial free chlorine solutions (ECS solutions) generated by 

monitoring the normalized areas of the Raman bands from OCl­ at 711 cm-1 (>, 

dash-dot model line), HOCl at 725 cm-1 (, dashed model line) and the low pH 

species at 538 cm-1 (‚, solid model line). Vertical lines indicate the apparent pKas 

at 13.1 mM. The pKa for the protonation of OCl­ to HOCl occurs at 7.50 for all 

three concentrations. The apparent pK at lower pH varies from (from top to 

bottom): 2.11, 1.84, and 1.25................................................................................... 45 

Figure 4 Raman spectrum acquired from (A) a buffered aqueous solution at pH 1.45 through 

which chlorine gas was bubbled (BCS) and (B) a pH 1.63 solution containing 13.1 

mM commercial free chlorine (ECS). The band at 538 cm-1 is due to the presence 

of Cl2 (aq) in each solution. ..................................................................................... 46 

Figure 5 Raman spectra comparing 13.1 mM free chlorine solutions with (A) equimolar 

chloride ion at pH 0.75 (ECS) and (B) free chlorine solution with reduced chloride 

ion at pH 0.80 (LCS). The spectrum of the solution containing equimolar chloride 

ion has a much larger band at 538 cm-1 (due to aqueous Cl2) than the solution with 

greatly reduced chloride concentration.................................................................... 47 

Figure 6 Speciation profile of the LCS free chlorine solution with reduced chloride ion 

concentration generated by monitoring the normalized areas of the Raman bands 

from HOCl at 725 cm-1 () and OCl­ at 711 cm-1 (>). The predominant species in 

solution at low pH remains HOCl throughout, and a significant drop in its total area 

is not observed as pH decreases............................................................................... 48 

Figure 7 Observed first order loss rate coefficients for chlorpyrifos over the pH range of 2-11 

in the presence of LCS. [CP]o = 0.5 μM, [Buffer]T = 10 mM, Temperature = 

25±1°C, and [HOCl]T = 10 μM. .............................................................................. 49 

vii 
Figure 8 Observed first order rate coefficients for CP loss in the presence of increasing 

bromide concentrations at [HOCl]T = 10 μM over the pH range of 7-9. [CP]o = 0.5 

μM, [Buffer]T = 10 mM, and Temperature = 25 oC................................................ 50 

Figure 9 Observed first order rate coefficients for CP loss in the presence of increasing 

bromide concentrations at [HOCl]T = 50 μM at pH 8 and 9. [CP]o = 0.5 μM, 

[Buffer]T = 10 mM, and Temperature = 25 oC. ....................................................... 51 

Figure 10 Two possible oxidation pathways for CP at high and low [HOCl]:[Br-] ratios. The 

pathway on the left shows the predominant pathway when [HOCl][Br-] and the 

pathway on the right shows the predominant pathway when [HOCl]:[Br-] 

approaches 1............................................................................................................. 52 

Figure 11 Loss of CP in the presence of hypobromous acid at pH 6.5. [CP]o = 0.42 μM, 

[HOBr]T = 10 μM, [Buffer]T = 10 mM, and Temperature = 25 oC. ........................ 53 

Figure 12 CP degradation in the presence of free chlorine at pH 8.0 in the absence of bromide. 

[CP]o = 0.39 μM, [HOCl]T = 10 μM, [PO4]T = 10 mM, and Temperature = 25 o

C. 

Lines represent model results................................................................................... 54 

Figure 13 CP degradation in the presence of free chlorine at pH 8.0 in the absence of bromide. 

[CP]o = 0.42 μM, [HOCl]T = 25 μM, [PO4]T = 10 mM, and Temperature = 25 oC. 

Lines represent model results................................................................................... 55 

Figure 14 CP degradation in the presence of free chlorine at pH 8.0 in the absence of bromide. 

[CP]o = 0.42 μM, [HOCl]T = 50 μM, [PO4]T = 10 mM, and Temperature = 25 oC. 

Lines represent model results................................................................................... 56 

Figure 15 CP degradation in the presence of free chlorine and bromide at pH 8.0. [CP]o = 

0.51 μM, [HOCl]T = 10 μM, [Br-] = 1 μM, [PO4]T = 10 mM, and Temperature = 25 o

C. Lines represent model results............................................................................ 57 

Figure 16 CP degradation in the presence of free chlorine and bromide at pH 8.0. [CP]o = 

0.48 μM, [HOCl]T = 25 μM, [Br-] = 1 μM, [PO4]T = 10 mM, and Temperature = 25 o

C. Lines represent model results............................................................................ 58 

Figure 17 CP degradation in the presence of free chlorine and bromide at pH 8.0. [CP]o = 

0.49 μM, [HOCl]T = 50 μM, [Br-

] = 1 μM, [PO4]T = 10 mM, and Temperature = 25 o

C. Lines represent model results............................................................................ 59 

Figure 18 CP degradation in the presence of free chlorine and bromide at pH 8.0. [CP]o = 

0.35 μM, [HOCl]T = 50 μM, [Br-

] = 1 μM, [PO4]T = 10 mM, and Temperature = 25 o

C. Lines represent model results............................................................................ 60 

viii 
Figure 19 CP degradation in the presence of free chlorine and bromide at pH 8.0. [CP]o = 

0.43 μM, [HOCl]T = 50 μM, [Br-] = 5 μM, [PO4]T = 10 mM, and Temperature = 25 o

C. Lines represent model results............................................................................ 61 

Figure 20 CP degradation in the presence of free chlorine and bromide at pH 8.0. [CP]o = 

0.39 μM, [HOCl]T = 50 μM, [Br-

] = 10 μM, [PO4]T = 10 mM, and Temperature = 

25 oC. Lines represent model results....................................................................... 62 

Figure 21 Chlorine demand of ACC water over the pH range of 7-9. [HOCl]T = 50 μM, 

[Buffer]T = 10 mM, [DOC] = 1.1 mg-C/L, and Temperature = 25 oC. ................... 63 

Figure 22 CP degradation in the presence of free chlorine and ACC water over the pH range 

of 7-9. [CP]o ≈ 0.5 μM, [HOCl]T = 50 μM, [DOC] = 1.1 mg-C/L, [Buffer]T = 10 

mM, and Temperature = 25 oC. Lines represent model results. ............................. 64 

Figure 23 CP degradation in the presence of free chlorine and ACC water at pH 7. [CP]o = 

0.46 μM, [HOCl]T = 50 μM, [DOC] = 1.1 mg-C/L, [Buffer]T = 10 mM, and 

Temperature = 25 oC. Lines represent model results.............................................. 65 

Figure 24 CP degradation in the presence of free chlorine and ACC water at pH 7.5. [CP]o = 

0.35 μM, [HOCl]T = 50 μM, [DOC] = 1.1 mg-C/L, [Buffer]T = 10 mM, and 

Temperature = 25 oC. Lines represent model results.............................................. 66 

Figure 25 CP degradation in the presence of free chlorine and ACC water at pH 8.0. [CP]o = 

0.42 μM, [HOCl]T = 50 μM, [DOC] = 1.1 mg-C/L, [Buffer]T = 10 mM, and 

Temperature = 25 oC. Lines represent model results.............................................. 67 

Figure 26 CP degradation in the presence of free chlorine and ACC water at pH 8.5. [CP]o = 

0.35 μM, [HOCl]T = 50 μM, [DOC] = 1.1 mg-C/L, [Buffer]T = 10 mM, and 

Temperature = 25 o

C. Lines represent model results.............................................. 68 

Figure 27 CP degradation in the presence of free chlorine and ACC water at pH 9.0. [CP]o = 

0.5 μM, [HOCl]T = 50 μM, [DOC] = 1.1 mg-C/L, [Buffer]T = 10 mM, and 

Temperature = 25 o

C. Lines represent model results.............................................. 69 

ix 
EXECUTIVE SUMMARY 

Environmental regulations require that all relevant routes of human exposure to 

anthropogenic chemicals be considered in risk assessments. Community water systems (CWSs) 

serve approximately 95% of the US population and potable water is considered a relevant route 

of exposure to anthropogenic chemicals. There is available monitoring data for important 

pesticides and toxic chemicals in drinking water sources (both surface and ground water). 

However, there is very little monitoring data for these chemicals or their degradates in finished 

drinking water. Limited experimental studies show that some chemicals are partially removed 

by physical water treatment processes (e.g., filtration, flocculation, etc.), and some are 

transformed by reactions that occur during chemical treatment (e.g., disinfection and softening). 

Transformation products of some contaminants have been shown to be more toxic than the 

parent compound. 

This report is in partial fulfillment of the National Exposure Research Laboratory Task # 

16608, “Fate of Pesticides and Toxic Chemicals During Drinking Water Treatment”, under Goal 

4, GPRA Objective 4.5, and GPRA Sub-objective 4.5.2. The goals of this research task are to: 1) 

provide chemical-specific information on the effects of water treatment for high-priority 

pollutants, 2) provide physicochemical parameters for transformation products, and 3) develop 

predictive models for forecasting treatment effects that cross chemical class and treatment 

conditions. 

Drinking water sources are a complex matrix of inorganic anions and cations and natural 

organic matter (NOM). Because of this complexity, it is exceedingly difficult to predicting the 

potential transformation of anthropogenic compounds under conditions that simulate drinking 

water treatment. Therefore, it is necessary to identify specific components in drinking water 

matrices that could significantly influence the transformation of anthropogenic chemicals (e.g., 

x 
organophosphate (OP) pesticides). Also, the effect of these native aqueous constituents could be 

dependent on other drinking water plant operational parameters such as pH adjustment or 

application of disinfectants. 

Results of experiments reported here demonstrate that the change in risk associated with 

the transformation of OP pesticides due to drinking water chlorination practices is a very 

complex issue. For example, pH of the finished water can significantly affect the rate of OP 

pesticide transformation. Previous work has shown that oxidation of OP pesticides by free 

chlorine near neutral pH leads very rapidly to a more toxic transformation product (i.e., oxons). 

Also, OP pesticides are susceptible to chlorine-assisted hydrolysis above neutral pH, which leads 

to a less-toxic transformation product. Below pH of 6, the oxidation of OPs can proceed much 

more rapidly than expected. This phenomenon has been observed by others in the drinking water 

community examining the loss of other anthropogenic compounds in the presence of free 

chlorine. Therefore, we chose to investigate this phenomenon in order to determine its relevance 

to drinking water treatment. Additionally, two native aquatic constituents (i.e., bromide and 

NOM) were investigated to determine their effect on the transformation of OP pesticides. It was 

hypothesized that bromide could accelerate the rate of OP transformation, while the presence of 

NOM was expected to decrease the rate of OP transformation due to the fact that NOM could act 

as a sink for active oxidants commonly used in drinking water treatment. 

Raman spectroscopy was used to determine the presence of active chlorine species over 

the pH range of 1-11. Over the pH range 5-11, unique Raman bands were identified and 

associated with hypochlorous acid (HOCl) and hypochlorite ion (OCl-

) with Raman shifts of 725 

and 711 cm-1 respectively. As the pH was lowered below pH 5, in-growth of a band at 538 cm-1 

appeared while the band at 725 cm-1 disappeared. This low-pH active chlorine species was 

xi 
identified as Cl2(aq). Molecular chlorine in the aqueous phase is a more potent oxidant than 

HOCl. The relevance of this to the drinking water industry is that many plants use chlorine gas 

to chlorinate potable water. Cl2(g) does hydrolyze to HOCl and OCl- quickly near neutral pH; 

however, its selective and rapid reactivity with not only OP pesticides but other anthropogenic 

chemicals is a factor in determining the risk associated with drinking potable water contaminated 

with anthropogenic chemicals. 

Bromide and NOM were investigated to determine their effect on the transformation rate 

of OP pesticides. Bromide was found to significantly increase the rate of chlorpyrifos (CP) 

transformation. This is due to the oxidation of Br- by HOCl resulting in the formation of 

hypobromous acid (HOBr), which is a more potent oxidant than HOCl. Since the formation of 

HOBr is pH dependent, a screening-level model that forecasts the concentrations of all reaction 

products as a function of pH, chlorine dose, OP pesticide concentration, bromide concentration, 

and time after chlorine dosing was developed. Using this simple model, the concentrations of 

CP, chlorpyrifos oxon (CPO), and 3,5,6-trichloro-2-pyridinol (TCP) can be adequately described 

under a variety of scenarios that are similar to drinking water treatment conditions using an 

apparent rate coefficient describing the reaction between HOBr and CP. In the presence of 

NOM, the model was found to predict the transformation of CP to CPO as well as the 

degradation of CPO to TCP in the presence of free chlorine. 

The work reported here demonstrates the “proof-of-concept” towards the development of 

a comprehensive modeling tool for OP fate in drinking water treatment plants and distribution 

systems. Following the reported work, we have now begun to apply the experiments and models 

to other OP pesticides that were judiciously selected to reflect the full range of chemical 

structure variability in this class. Our objective is to develop predictive models that associate 

xii 
relative rates of reactivity to structural variability. This will allow decision makers to rank and 

prioritize chemicals found in drinking water sources according to potential risk. Also, we are 

currently in the process of investigating chloramination disinfection practices on the 

transformation of OP pesticides. 

xiii 
1 INTRODUCTION 

The Food Quality Protection Act of 1996 (FQPA) requires that all pesticide chemical 

residuals in or on food be considered for anticipated human exposure. Drinking water is 

considered a potential pathway for dietary exposure, but there is reliable monitoring data for only 

the source water. For example, the United States Geological Survey (USGS) completed a 

national reconnaissance survey known as NAWQA (National Water Quality Assessment) to help 

define human exposure to various contaminants (1). For the NAWQA survey, 90 pesticides (and 

some selected metabolites) were chosen as target chemicals to monitor in US drinking water 

sources. However, there is a relative dearth of information on occurrence of pesticide residuals 

and pesticide metabolites in finished drinking water. Two surveys have been conducted for a 

few community water systems examining pesticide concentrations in the source and finished 

drinking water (2,3). Neither of these studies thoroughly examined the effect of each treatment 

process on a single slug of water, hence only the influent and effluent of each treatment facility 

could be qualitatively discussed with respect to overall removal efficacy. Also, these studies did 

not account for the treatment plant hydraulic retention time, thus it was not possible to ensure 

that influent and effluent samples were properly paired. Clearly, more thorough occurrence 

studies that include monitoring drinking water for both pesticides and their expected 

transformation products are needed. 

Chlorination is the most commonly used chemical disinfection process for community 

water systems (4) and is known to react with numerous pesticides. For example, four s-triazines 

were found to degrade in the presence of free chlorine (HOCl + OCl-

) (5,6). Atrazine was also 

found to be significantly degraded by ozone (7); however, subsequent chlorination of the 

ozonated effluent had very little affect on the concentration of residual atrazine or its ozone 

degradation products (2). Also, some carbamate pesticides have been shown to react with free 

1
chlorine while other members of this pesticide class were found to be stable in chlorinated water. 

For example, carbaryl and propoxur do not react with free chlorine; but aldicarb, methomyl, and 

thiobencarb do exhibit significant reactivity (8-10). These findings demonstrate that free 

chlorine reactivity with different members in a specific class of pesticides can vary significantly 

due to chemical structure variations. Therefore, it is prudent to study the fate and transformation 

pathways of entire chemical classes, using class members that exhibit systematic structural 

variations and employing carefully selected experimentation and numerical modeling. 

When chlorine reacts with the phosphorothioate subgroup of organophosphate (OP) 

pesticides, the thiophosphate functionality (P=S) can be oxidized to its corresponding oxon 

(P=O) (11-13). The resulting oxons are typically more potent than the parent as an inhibitor of 

acetlycholinesterase, an enzyme necessary for proper function of the nervous system (13). Duirk 

and Collette, (14) elucidated the fate of chlorpyrifos (CP) and its transformation products over 

the pH range of 6-11. They were able to model the loss of CP and chlorpyrifos oxon (CPO) to 

the stable end-product of 3,5,6-trichloro-2-pyridinol (TCP) over this pH range in buffered 

deionized water systems. However, other factors such as pH below 6, bromide ion, and natural 

organic matter can significantly influence the transformation rate of CP. 

Several recent papers in drinking water literature have reported free chlorine oxidation 

phenomena that stray from expected results as the pH was lowered in the course of kinetic 

experiments with various anthropogenic chemicals (15-17). Specifically, a dramatic increase in 

reaction rate was observed in solutions when the pH was lowered below approximately 5. The 

same trend has been observed while studying the free chlorine oxidation of CP, and a kinetic plot 

depicting this phenomenon is displayed in Figure 1. Initially, the observed reaction rate was 

assumed to be relatively constant as the pH was lowered below the point where HOCl was the 

2
principal species and the OCl- concentration was insignificant. This result was expected because 

HOCl is generally assumed to be the predominate oxidant in free chlorine solutions. However, 

since the observed rate coefficient increased dramatically as pH was dropped below 5, it was 

clear that there was another chlorine species in solution, more potent as an oxidant than HOCl. 

At least two competing hypotheses have been put forward by various researchers 

regarding the identity of the species responsible for this rapid oxidation. Specifically, some 

researchers have proposed H2OCl+ as the oxidant (18), and other researchers have proposed 

Cl2(aq) (17). In literature, there are no thorough studies designed to rigorously determine the 

identity and speciation profile of this potent oxidant. It appears that the H2OCl+ species was first 

proposed as an active species at low pH in 1962 by Arotsky and Symons (19) in their discussion 

of halogen cations. Although they did not measure the pKa for H2OCl+, it was estimated to be 

between -3 and - 4 (19). Some publications have discussed the existence of the H2OCl

+ cation 

(20,21) as part of a reverse hydrolysis reaction. Despite the low proposed pKa, many 

chlorination (15,16,18,22,23) and other oxidation (24) studies have either proposed or discussed 

H2OCl+ as the reactive, oxidative species in low pH aqueous solutions. There have been a few 

publications disagreeing with the hypotheses involving this species, but these are much fewer in 

number and are found in the inorganic chemistry literature (25,26). 

The presence of bromide ion (Br-

) in drinking water sources has been of interest to the 

drinking water industry for several reasons. Bromide ion can be oxidized by ozone (O3) 

resulting in the formation of bromate ion (BrO3

-

) (27). Bromate is a suspected carcinogen and is 

regulated under Stage 1 of the Disinfection/Disinfection Byproducts (D/DBP) Rule at a 

maximum contaminant level (MCL) of 0.01 mg/L. Bromide can also be easily oxidized to 

hypobromous acid in the presence of free chlorine (28), it is a stronger oxidant that HOCl, 

3
HOCl + Br­ HOBr + Cl- kHOCl,Br = 5.6 x 106

 M-1h-1 (1) 

HOBr +H2O OBr- + H3O+ pKa = 8.8 (2) 

and the pKa of hypobromous acid is 8.8 (28). In the presence of excess chlorine, all bromide 

present will be oxidized to active bromine (29). Over the pH range of 6.5-9, which is the pH 

range of the National Secondary Drinking Water Standards, most of the active bromine present 

will be in the HOBr form (i.e., the active oxidant). Therefore, it is believed that the presence of 

bromide under drinking water treatment conditions could accelerate the transformation of 

anthropogenic chemicals. For example, the rate coefficient for HOBr reacting with pyrene was 

found to be 2 orders of magnitude faster than HOCl in aqueous solutions (30). In the presence of 

natural organic matter (NOM), both brominated and chlorinated oxidants react with NOM 

resulting in the formation of disinfection byproducts (DBPs) (31,32). The brominated DBPs are 

believed to be more carcinogenic than their chlorinated analogs (33). Westerhoff et al., (34) 

conducted a nationwide survey of 101 drinking water sources and found that the average 

bromide concentration was ≈ 1.25 μM. Since bromide is a relatively ubiquitous and 90% of 

community drinking water systems use chlorine for both primary and secondary disinfection, the 

presence of bromide ion in drinking water sources could affect the rate of organophosphorus 

(OP) pesticide transformation in chlorinated drinking water. 

The presence of NOM in drinking water sources has also been a primary concern for the 

drinking water industry. The physical-chemical properties of NOM are a function of decaying 

organic matter from both terrestrial and aquatic sources, environmental conditions, and 

microorganisms present in the watershed (35). Hence, the composition and concentration of 

NOM in source waters fluctuates from source to source and often varies within each source 

depending on conditions such as rainfall, change in seasons, and temperature. When chlorine 

4
reacts with NOM, it often results in the formation of DBPs (33). The electron density (i.e., 

aromaticity) of NOM is often expressed as specific ultraviolet absorbance (SUVA), which is the 

ratio of ultraviolet absorbance of the NOM at a specific wavelength divided by the total organic 

carbon concentration (TOC) in mg-C/L. This specific physical-chemical property of NOM has 

been found to correlate to reactivity of NOM with free chlorine and serve as a reliable surrogate 

for DBP formation (36,37). Since both chlorinated and brominated oxidants react with NOM 

(38), the presence of NOM can potentially act as a sink for active oxidants in potable water. 

Thus, its presence can impact the rate of OP pesticide transformation under drinking water 

treatment conditions. 

One purpose of this study was to identify active oxidants under drinking water treatment 

conditions. Toward this end, Raman spectroscopy was used to determine active chlorine species 

present over the pH range of 1-11 in a solution where active chlorine to chloride ion molar ratios 

were equivalent (equal-molar chlorine solution: ECS). Also, a low chloride ion active chlorine 

solution (low-chloride chlorine solution: LCS) was prepared to assist in the identification of 

active chlorine species as a function of pH. Kinetic studies with CP in the presence of both ECS 

and LCS were conducted to investigate CP transformation below pH 6. Bromide ion was 

investigated to determine the effect it has on transformation of CP in the pH range of 7-9. Also, 

the influence of NOM on the transformation of CP was investigated over the pH range of 7-9. 

The work reported here serves to further elucidate the OP transformation pathways under 

drinking water treatment conditions. The long-term objective is the development of a 

comprehensive modeling tool for OP fate in drinking water treatment plants and distribution 

systems. This will allow decision makers to rank and prioritize chemicals found in drinking 

water sources according to potential risk. The goal of these experiments was to incorporate these 

5
important matrix effects into our models so that the models can be applied to actual drinking 

water treatment scenarios. 

6
2 EXPERIMENTAL PROCEDURES 

2.1 Materials 

Chlorpyrifos (99.5%), chlorpyrifos oxon (98.7%), and 3,5,6-trichloro-2-pyridinol (99%) 

were purchased from ChemService (West Chester, PA). Commercial 10-13% sodium 

hypochlorite (NaOCl), purchased from Aldrich (Milwaukee, WI), contained equal-molar 

amounts of OCl- and Cl-

. Aqueous stock solutions and experiments utilized laboratory prepared 

deionized water (18 MΩ cm-1) from a Barnstead ROPure InfinityTM/NANOPure TM system 

(Barnstead-Thermolyne Corp., Dubuque, IA). Chlorine gas (99.9% purity) was purchased from 

Matheson Tri-Gas (Montgomeryville, PA). Filters with pore size of 0.45 μm were purchased 

from Millipore (Billerica, MA). Phosphate and carbonate salts used for buffer solutions were 

dissolved in deionized water and filtered through a 0.45 μm filter, which was pre-rinsed with 

deionized water. The pH for the experiments was adjusted with either 1 N H2SO4 or NaOH 

unless otherwise noted. All other organic and inorganic chemicals were certified ACS reagent 

grade and used without further purification. The glassware and polytetrafluoroethylene (PTFE) 

septa used in this study were soaked in a concentrated free chlorine solution for 24 hours, rinsed 

with deionized water, and dried at 105 o

C prior to use. All chlorination experiments were 

conducted at constant temperature (25±1°C). 

Water was collected form the Athens-Clarke County (ACC) water treatment plant and 

used in some experiments. The ACC water treatment facility uses conventional physicalchemical surface water treatment (i.e., coagulation, flocculation, sedimentation, and dual media 

filtration) prior to chlorination. Water was collected prior to chlorination and the NOM and 

source water characteristics are shown in Table 1. The SUVA254 of the ACC effluent was 1.61 

L/mg-m, which indicates a low aquatic humic substance concentration in this water (36). 

7
2.2 Methods 

2.2.1 Spectroscopic Studies of Active Chlorine Species 

For speciation studies, Raman spectra were collected of free-chlorine solutions as a 

function of pH (at least 20 points) over the range of approximately 1 to 12. The commercial 

sodium hypochlorite solution (ECS) was diluted with deionized water to a stock solution with a 

free chlorine concentration of 262.0 mM to more closely match that of the low chloride solution 

(LCS) that was prepared in our lab. Total free chlorine concentration was again determined via 

Standard Method 4500-Cl F DPD-FAS titrimetric method (39). The final concentrations of free 

chlorine in the ECS solutions monitored with Raman spectroscopy were 26.2, 13.1 and 6.6 mM. 

The pH of 50 mM phosphate buffer was raised with 5 N sodium hydroxide and lowered with a 2 

N sulfuric acid solution. The cuvette used for Raman work was relieved of chlorine demand by 

soaking it in a concentrated free-chlorine solution overnight and rinsing it several times with 

deionized water. Samples for Raman analysis were prepared by adding an appropriate amount of 

buffer directly to the cuvette, followed by the addition of free-chlorine solution. The cuvette was 

then rapidly capped, and inverted several times for mixing. Minimal headspace was left in the 

capped cuvette to minimize the possible loss of a volatile species in a low pH solution. (There 

was ~ 100 μL of gas volume between the solution and the cap.) The filled cuvette was then 

placed into a sample holder and the Raman spectra were acquired. Unless otherwise noted, the 

time for each acquisition was 5 s. To improve the signal-to-noise ratio in the spectrum, 200 of 

these 5 s acquisitions were co-added (i.e., averaged). After acquisition of the Raman spectra, the 

samples were transferred into a vial that had been relieved of chlorine demand and the final 

solution pH was recorded. 

Some of our experiments were carried out with a free-chlorine solution containing a low 

concentration of chloride ion, LCS. Our method for preparing this solution, which was adapted 

8
from Reinhard and Stumm (40), involves complexing and removing excess chloride ion from 

hypochlorous acid/hypochlorite preparations. First, all glassware was soaked in a concentrated 

free-chlorine solution to satisfy any chlorine demand. The glassware was rinsed with deionized 

water and dried at 105 oC prior to use. To a 250 mL graduated cylinder, sodium hydroxide 

(2.5mL of a 1 N solution) was added to deionized water for a final volume of 250 mL. Mercuric 

oxide (8.33g) was added to the graduated cylinder and the solution was stirred for 10 min. 

Chlorine gas was then bubbled through the HgO/NaOH/water solution for 3 min and this 

solution was stirred for 1 h. After the stirring time had elapsed, the solution was filtered through 

a 0.45 μm filter that had been pre-rinsed with 1.5 L of deionized water. The filtrate was then 

distilled at 90°C into sodium hydroxide (2.5 mL of a 0.1 N solution). The distillate was stored at 

4°C until further used. 

The amount of chloride ion in the LCS was determined in a multi-step process. First, 

using the Standard Method 4500-Cl F DPD-FAS titrimetric procedure (39), we determined that 

the total free-chlorine concentration was 202.5 mM. (At the alkaline pH of this solution, the 

total free-chlorine concentration is essentially equal to the OClconcentration.) Next, an aliquot 

of the LCS was reduced with excess sodium sulfite, thereby converting all species of chlorine to 

chloride ion. Then, using ion chromatograph with chemical suppression, we determined that the 

chloride ion concentration in the reduced aliquot was 220.0 mM. Therefore, by difference, we 

determined that the chloride ion concentration in the LCS was 17.5 mM. This method of 

producing a low chloride ion solution was effective – the free chlorine to chloride ion molar ratio 

in our prepared LCS was 11.6:1, whereas in the commercial ECS the ratio was 1:1. 

Many studies have been published on the absorption of chlorine into water (41) and on 

the kinetics of HOCl production at various temperatures (42-44) and ionic strengths (44). This 

9
reaction has been shown to proceed more rapidly at both higher temperature and higher ionic 

strength, (42-47) and experimental evidence has been obtained with several buffers arguing that 

the reaction involves base-assisted hydrolysis (48-50). A low pH free-chlorine solution was 

prepared by bubbling chlorine gas through a solution of 50 mM phosphate buffer for 

approximately 20 sec. The initial pH of the solution was 1.8 and the final pH was 1.4 and the 

color of the solution changed from clear to yellow. We report experiments here using this 

solution, which was assigned the acronym BCS to denote Bubbled Chlorine Solution 

2.2.2 Chlorpyrifos Kinetic Experiments with Bromide and NOM 

For all CP oxidation experiments, CP was spiked by adding 0.5 mL of a 4 mM stock 

solution in ethyl acetate into an empty 4 L borosilicate glass Erlenmeyer flask. A gentle flow of 

nitrogen gas was used to evaporate the ethyl acetate and then 4 L of deionized or ACC water was 

added to the flask. The solution was slowly stirred and allowed to dissolve for 12 hours resulting 

in an aqueous concentration of 0.5 μM. Experiments studying the effect of bromide, bromide 

was added to the aqueous system prior to chlorination. CP chlorination kinetic experiments were 

conducted under pseudo-first-order conditions: total chlorine to chlorpyrifos molar ratios of 

20:1, 50:1, and 100:1 were used. Chlorine was added to solutions under rapid mix conditions 

achieved with a magnetic stir plate and a PTFE coated stir bar. All kinetic experiments used to 

estimate rate coefficients measured at least 87% loss of parent compound. 

Above pH 8, 10 mM carbonate [CO3]T buffer was used to maintain pH. The purchased 

free chlorine solution was first diluted to 250 mM then added to the aqueous system containing 

0.5 μM chlorpyrifos and carbonate buffer in a 2 L Erlenmeyer flask. Fourteen aliquots from the 

large 2 L reactor were then placed into 128 mL amber reaction vessels with PTFE septa and 

stored in the dark. In the pH range of 6.0-7.5, the rate of CP loss in the presence of free chlorine 

10
was very fast. Therefore, fourteen 100 mL aliquots of the 2 L aqueous system containing 10 mM 

phosphate buffer, [PO4]T, and 0.5 μM CP were placed in 250 mL amber Erlenmeyer flasks. 

Then each flask was individually dosed with the appropriate amount of chlorine. 

At each discrete sampling interval, two reaction vessels were sacrificed in their entirety. 

One vessel was used to determine total free chlorine concentration ([HOCl]T = [HOCl] +[OCl-]) 

via Standard Method 4500-Cl F DPD-FAS titrimetric method (39). In the other vessel, free 

chlorine residuals were quenched and the pH of a 100 mL sample was then adjusted to 2 for 

analysis of CP and its degradation products. 

Chlorine residuals were quenched with sodium sulfite in 20% excess of the initial free 

chlorine concentration. Both sodium sulfite and ascorbic acid were tested as quenching agents to 

determine if they affected the stability or recovery of CP, CPO, or TCP. Control studies 

compared the recoveries of all three analytes in water without chlorine present after two hours at 

room temperature with and without a quenching agent present. Sodium sulfite had no apparent 

effect on the recovery of any of the three analytes. However, ascorbic acid reduced the recovery 

of CP by about 5.1% and increased the recovery of CPO by about 7.6%, as compared to both the 

sodium sulfite quenched samples and the samples with no quenching agent present. 

Furthermore, ascorbic acid has previously been found to react with organophosphate pesticides 

(51). Therefore, ascorbic acid was not used to quench chlorine residuals in this work. 

2.3 Instrumentation 

The Raman spectra were acquired with a Kaiser Optical Systems (Ann Arbor, MI) 

HoloProbe, using 532 nm laser excitation from a CoherentTM (Santa Clara, CA) Verdi V Series 

frequency-doubled YVO4 laser. This type of Raman instrument has been fully described 

elsewhere (52,53). Briefly, laser light was coupled to a remote probe head via a 1.9 m long 

fiber-optic cable. Raman scattering and luminescence were removed from the excitation optical 

11
train via a holographic grating and spatial filter within the probe head. This filtered laser light 

was brought to a focus with a 10x microscope objective (Olympus) approximately in the center 

of a standard quartz cuvette with a screw cap. The power of the laser light at the cuvette was 

approximately 220 mW. Raman scattered photons from the solution were collected by the 

microscope objective along the same path as the excitation laser beam (i.e. 180° backscattering 

geometry), and were coupled to a separate 1.9 m fiber-optic cable for delivery to the f/1.8 axial 

transmissive-type spectrograph. The probe head contained a holographic notch filter placed 

along the collection path such that only inelastically scattered light was inputted into the fiber. 

Inside the spectrograph the scattered light was passed through another holographic notch filter to 

further remove elastically scattered light returning from the probe head. The Raman scattered 

light was then focused through a 50 μm slit and directed through a holographic grating with 

double transmission layers. These layers permit two regions of the dispersed Raman spectrum to 

be stacked and simultaneously focused onto the back-illuminated CCD detector (Andor 

Technology model DU420). This configuration permitted the acquisition of the entire Raman 

spectrum with a useable Stokes Raman shift of about 4450 to 100 cm-1 in a single exposure with 

about 5 cm-1 spectral resolution. The CCD detector was thermoelectrically cooled to -65°C and 

contained 1024x 256 pixels. 

A Kaiser Optical Systems, Inc. HoloLab calibration accessory was used to correct both 

the frequency and intensity of the Raman spectra. Briefly, each spectrum was frequencycorrected based on previously measured lines from a neon bulb. Also, each spectrum was 

intensity-corrected based on a previously measured NIST-traceable, broadband light source (i.e., 

a tungsten bulb) that was powered by a current regulated power source for consistent spectral 

output. 

12
Analysis of a Raman spectrum required the subtraction of a 50 mM phosphate buffer 

spectra that was the same pH of the chlorine-containing solutions. The pH of the buffer was no 

more than a few hundredths of a pH unit different from that of the associated solution. The 

bands in the spectrum were then analyzed by fitting a Gaussian-Lorentzian mixed curve in 

GRAMS/32 (Thermo Galactic). Since the HOCl and OCl- bands overlapped with each other, the 

center frequency and band width for these two species were determined with spectra that were 

not near the pKa of hypochlorous acid. These two parameters were then held constant 

throughout the analysis and the band areas were determined for each spectral acquisition. In 

addition, the band associated with atmospheric nitrogen was averaged and used, as has been 

previously described (54) to correct the band areas corresponding to the chlorine species for any 

changes in instrumental conditions (e.g., alignment, laser power fluctuations, etc.). 

We sought to use these individual Raman band areas to construct chemical speciation 

plots. For our purpose, a speciation plot is useful when it depicts the fractions of the total 

analytical concentration that belong to the various species in solution as a function of pH. 

Unfortunately, individual band areas could not be used directly to construct such plots because 

we could not assume that Raman scattering cross sections of the various species were equal. 

However, at the highest analytical concentration of free chlorine in ECS solutions examined by 

Raman spectroscopy (26.2 mM), it was possible by varying pH to obtain three individual Raman 

spectra, each of which uniquely contained appreciable intensity for only one of the three 

observed species. From these three spectra, we estimated an apparent Raman scattering cross 

section for each of the three species. Then, using these apparent cross sections, we calculated the 

intensity that we would observe for a given species at a given analytical concentration if 100% of 

the analytical concentration existed in that form. Individual Raman band area collected at the 

13
various pH’s for a given analytical concentration were then divided by this quantity in order to 

“normalize” them. These normalized band intensities, which ranged from nominal values of 0 to 

1 were used in speciation plots to approximate the fractions of the total analytical concentration 

that belonged to the various species in solution as a function of pH. 

CP and its degradation products were extracted from water using C-18 solid phase 

extraction cartridges purchased from Supleco (Bellefonte, PA). The 100 mL sample was preadjusted to pH ≤ 2 to increase the recovery of TCP (pKa = 4.55) on the solid phase adsorbent 

(55). Then, the sample was spiked with 1 μM of phenthorate (internal standard), mixed 

thoroughly by hand for two minutes, passed through the SPE cartridge at an approximate flowrate of 7 mL/min, and eluted with 3 mL of ethyl acetate. Quantification for each analyte was 

compared to eight extracted standards over the concentration range of 0.01 to 1 μM. A HewlettPackard 6890 GC equipped with a 5973 MSD was used to analyze CP, CPO, and TCP. GC 

conditions were as follows: 30-m Restek Rtx-200 column with a 0.25-mm ID and 0.5-μm film 

thickness. The temperature profile was: 100 °C for 5 minutes, 100 to 250 °C at 10 °C/minute, 

and then held at 250 °C for 25 minutes. Mass balances of 80% or greater were obtained for each 

experiment. 

UV-Vis spectra were acquired with a Shimadzu 1700 UV-Vis spectrometer (Shimadzu 

Scientific, Columbia, MD). Total organic carbon (TOC) was measured using a Shimadzu TOC 

5000 (Shimadzu Scientific, Columbia, MD) and calibrated according to Standard Methods 505A 

(39). Ion chromatography was performed on a Metrohm (Houston, TX) MIC-2 ion 

chromatograph with chemical suppression. All pH measurements were obtained with an Orion 

940 pH meter using a Ross combination electrode from Fisher Scientific (Pittsburgh, PA). 

14
3 RESULTS AND DISCUSSION 

3.1 Spectroscopic Studies of Active Chlorine Species 

As discussed earlier, there was a large increase in the rate of aqueous free chlorine 

oxidation for chlorpyrifos when the pH was decreased below ~ 5. This behavior, which is 

similar to that recently reported with other anthropogenic chemicals (15,17,56,57), suggests the 

presence of an oxidant other than OCl-

 and HOCl, and more potent than either, that is formed as 

the pH decreases below 5. To firmly establish the existence of this third species, and to 

determine whether it is Cl2 (aq) or H2OCl+

, Raman spectra of the aqueous free chlorine ECS 

solution (26.2 mM) were collected over a pH range from approximately 1 to 12. Across this pH 

range, three Raman bands appeared in the spectra due to various chlorine species in solution. 

Specifically, a band at 711 cm-1 was observed to increase in intensity as the pH was increased 

above 6, and its intensity was relatively constant above pH of ~ 9.5. A second band at 725 cm-1 

was observed to increase in intensity as the pH was lowered from 9 to 6, and its intensity 

decreased as the pH was lowered below ~ 5. Finally, a third band at 538 cm-1 was observed to 

increase in intensity as the pH was lowered below 5. The normalized areas of these bands are 

plotted over the entire pH range in Figure 2A. 

Figure 2B shows spectra of chlorine acid species in solution at four discrete pHs (11.96, 

7.46, 4.59, and 2.73) with the aqueous phosphate buffer subtracted. Bands in the figure are 

labeled according to the species to which they are attributed. The band at 711 cm-1 is confidently 

attributed uniquely to the OCl- species and the band at 725 cm-1 uniquely to the HOCl species, in 

agreement with an earlier study (58). The in-growth of a Raman band at 538 cm-1 as the pH is 

lowered below 5 supports the presence of a third chlorine species. We believe this band at 538 

cm-1 can be attributed uniquely to the species that is responsible for the unexpectedly rapid 

15
oxidation of anthropogenic chemicals that is observed in free chlorine solutions as pH decreases 

below 5. 

Raman spectroscopy has been demonstrated as an effective method for determining the 

equilibrium constants of organic species in aqueous solution (54). By monitoring monotonic 

changes in spectra, the amount of each species can be quantified at each pH studied. This 

general method is applied to the current experiment to model chlorine species in aqueous 

solution. The pKa for the protonation of OCl-, described by: 

HOCl OCl-+ H3O+ (3) 

can be readily approximated by monitoring the Raman-active vibrational bands of the species 

HOCl and OCl- (at 725 and 711 cm-1, respectively) as a function of pH. Specifically, for a 

simple isolated protonation event of this type, the pKa can be approximated as the pH at which 

the normalized intensities of the HOCl and OCl- bands are equal because, according to the 

Henderson-Hasslebach equation, when [HOCl] = [OCl-], the 

pH = -log K1 = pK1 (4) 

and 

K1 = [OCl-] [H3O+

]/[HOCl] (5) 

(Although it would be accurate in this case to label the equilibrium constant as Ka, for this work 

the more generic K1 and K2 to denote constants for the equilibria discussed at high and low pH, 

respectively. For simplicity, concentration-based equilibrium constants will be presented. Using 

the extended Debye-Hhckel equation (59) the calculated activity coefficients for all charged 

species discussed herein and have determined that, at the concentrations used in this work , 

activity-based equilibrium constants do not differ significantly from those based on 

concentration. Specifically, for determination of pKs, the differences were calculated to be 

16
between 0.7 and 1.2% for the high pH event and 1.3 to 6.4% for the low pH event for all 

equilibrium cases and concentrations considered. For these calculations, activity coefficients of 

all neutral species were assumed to be 1. 

Furthermore, the species arising at a lower pH can be described by either equation 6 or 7, 

as discussed in the Introduction: 

H2OCl+ H3O+ + HOCl (6) 

H2O + Cl2 H3O+ + HOCl + Cl- (7)

If equation 6 is correct, then the low-pH species is formed from a simple protonation event like 

that described in equation 3. On the other hand, if equation 7 is correct, the equilibrium is more 

complicated, involving species other than a proton and a protonated/deprotonated pair. 

A convenient way to distinguish between the two situations described by equations 6 and 

7 was to monitor the relative concentrations of HOCl and the low-pH form as a function of pH as 

the absolute concentration of total chlorine was varied. This was achieved by preparing 

speciation plots such as that shown in Figure 2A for several chlorine concentrations. If equation 

6 was the proper description, the pH at which the normalized intensities of the bands for the lowpH form and HOCl are equal would approximate the pKa and this observed pKa would not vary 

as a function of total free chlorine. Specifically, for equation 6 (as with equation 3), when 

[HOCl] = [H2OCl+], the 

pH = -log K2

H 2OCl + = pK2 H 2OCl+ (8) 

K 

where 

H 2OCl +

 = [HOCl] [H3O+] / [H2OCl+] (9) 2

17
Alternately, if equation 7 was the proper description, the pH at which the normalized 

intensity of the bands were equal would change as a function of total free chlorine. Specifically, 

for equation 7, when [HOCl] = [Cl2], the 

pH = -log KCl 2 + log [Cl-] = p KCl 2 + log [Cl-] (10) 2 2 

K

where 

Cl 2

 = [HOCl] [Cl-] [H3O+] / [Cl2] (11) 2

The commercial reagents used to prepare free-chlorine solutions (ECS) were equimolar in HOCl 

and Cl-. So, as the free chlorine concentration is increased, [Cl-] will increase. Thus, from 

equation 10, we would expect the pH at which the normalized intensity of the bands are equal to 

increase as the free chlorine concentration increases in ECS solutions if equation 7 is the proper 

description of the low pH behavior. 

Full speciation plots, using normalized Raman band intensities, for three different 

commercial total free chlorine solutions (ECS solutions) are displayed in Figure 3. In order to 

determine if the apparent pK for either equilibrium changed as a function of chlorine 

concentration, the normalized band intensities were fit using standard equations that describe the 

concentration profiles of polyprotic acids (60). To present these equations, we assigned K2 as 

the equilibrium constant for the low pH event described by equation 9 or 11 and K1 for the 

higher pH event described by equation 5. For these generalized equations, the amount of each 

component (αi) is described in terms of the total amount of the acid or base present. In this case: 

(60) 

α = [lowpHform] (12) 0 CT 

18
[HOCl ]

α = (13) 1 CT 

[

OCl−

]

α =

C (13) 2 

T 

where CT = [low pH form] + [HOCl] + [OCl-

]. It is important to note that the sum of the αi 

components must be equal to 1 under every condition. The numerical value of each αi 

component was obtained by utilizing the following equations: (60) 

α = [H 3O+

]

2 

(14) 0 2

[H 3O+] + K2[H 3O+] + K2 K1 

α1 = 2 

K 2[H 3O+] (15)

[H 3O+] + K2[H 3O+] + K2 K1 

α = K2 K1 (16) 2 2 

[H + + 3O+] K2[H 3O+] K2 K1 

The model was fit to the data by using a least-squares analysis and allowing the values 

for K2 and K1 to vary in order to minimize the sum of the squared residuals. The fit of the model 

in the top graph of Figure 3 most closely matches the experimental data because it was used to 

determine the apparent Raman scattering cross-section for each species. Subsequent plots of 

experimental data for 13.1 mM and 6.6 mM solutions deviate slightly from the model because 

the spectral correction with the nitrogen band (as described in the Experimental section) cannot 

perfectly encompass all variations in the instrument over long periods of time. (Note that all of 

the data in Figure 3 was collected over a time period of 60 days). Furthermore, the greatest 

variation between experimental data and the model is seen in the amounts of HOCl and OCl-

. 

19
This is likely due to these bands not being completely resolved from one another, limiting the 

accuracy of the curve fitting routine. 

The vertical line on the right in Figure 3 indicates the apparent pKa (i.e. pK1) for the 

protonation of OCl- that occurs between pH 7 and 8 for each concentration of free chlorine. This 

line indicates that the pKa of hypochlorous acid is approximately 7.50, which agrees very well 

with the previously published result of a pKa near 7.5 (44). Indeed, the pKa for each 

concentration was identical to the significant digits that we can accurately report. On the other 

hand, it is evident, as shown by the line on the left in Figure 3, that the apparent pK for the low 

pH equilibrium does not occur at the same pH regardless of concentration. The apparent pK 

increases from 1.25 to 2.11 as the concentration of free chlorine in solution is increased. The 

inconsistent location of the pK in this equilibrium strongly suggests that a protonation event is 

not the source of the change and that equation 7 is the proper description of the low-pH behavior. 

Indeed, the direction of the change – an increase in apparent pK as the free chlorine 

concentration is increased – was predicted (in equation 8) when equation 7 was the proper 

description. Furthermore, the magnitude of the change in apparent pK on going from 6.6mM to 

26.2 mM – over 65% – was much larger than could be explained by changes in activity 

coefficients over this concentration range. Indeed, changes in activity coefficients alone would 

result in a change in apparent pK over this concentration range of less than 5%. 

To further investigate the hypothesis that Cl2 was the correct assignment for the low-pH 

form of aqueous free chlorine, a Raman spectrum was acquired of a solution (BCS) that had been 

prepared by bubbling chlorine gas through an acidic 50mM phosphate buffered solution for 

approximately 20 s. The pH of the BCS solution was intentionally kept low during preparation 

for comparison to the diluted, acidic ECS solutions that were monitored earlier. The Raman 

20
spectrum of the acidic BCS solution is shown as trace A in Figure 4. The band at 538 cm-1, as 

previously observed in low-pH ECS solutions prepared from HOCl, was very prominent in the 

BCS spectrum and is much larger than the observed HOCl vibration (at 725 cm-1). The Raman 

spectrum of a diluted ECS solution at approximately the same pH is also shown, as trace B, in 

Figure 4. Our hypothesis that Cl2 is the correct assignment of the low pH form of chlorine in 

solution is strengthened because the band at 538 cm-1 is coincident in both spectra. The observed 

disparity in relative band areas is simply indicative that Cl2 in the BCS solution is much more 

concentrated relative to HOCl than in the ECS solution. 

If equation 7 were, indeed, the correct description of the low-pH behavior of aqueous free 

chlorine, then minimizing the chloride ion (Cl-) concentration in solution would diminish the 

formation of the potent oxidant as the pH is reduced. As described above, our commercial freechlorine solutions (ECS) contained an equimolar amount of HOCl and Cl-

. However, we were 

able to prepare a stock low chloride solution (LCS) at alkaline pH using the method adapted 

from Reinhard and Stumm (40) as described in the Experimental Section. This LCS stock 

solution was diluted to 13.1 mM in free chlorine, which resulted in a Cl- concentration of 1.12 

mM. For the ECS solution at this free-chlorine concentration, the concentration of Cl- (i.e. 13.1 

mM) was more than ten times higher. 

Figure 5 displays the Raman spectrum acquired from the LCS at pH 0.80 (trace B) and of 

the ECS at a similar pH (trace A) with the same concentration of free chlorine (i.e., 13.1 mM). It 

is easily seen that the ratio of band intensities in the Raman spectra at 538 to 725 cm-1 is much 

larger for the ECS solution, which contained equimolar chloride ion. Figure 6 displays the 

normalized band areas for the Raman bands located at 711 cm-1 (OCl-

) and 725 cm-1 (HOCl) as a 

function of pH for the 13.1 mM LCS sample The normalized band area of the HOCl Raman band 

21
at 725 cm-1 does not exhibit a clear and consistent decrease as the pH of the solution is dropped 

below 5 as was observed in Figure 3 with the ECS solutions. Even at a pH below 0, there was 

not a significant decrease in the HOCl Raman band at 725 cm-1, nor is there any significant 

Raman scattering observed in the 538 cm-1 region of the spectrum. (Note that a normalization 

procedure that was simpler than that described in the Experimental Section was used to construct 

the speciation plot in Figure 6. Because only two species were observed and both species could 

be completely isolated one from the other within the pH range sampled, it was acceptable to 

normalize each individual band area by the maximum area observed for that band.) Without 

chloride ion present in solution, a reverse hydrolysis reaction cannot take place if equation 7 is 

the correct description of the low-pH behavior. However, if equation 6 were the correct 

description, then lowering the chloride ion concentration should have no effect on the 

equilibrium towards H2OCl+

 (equation 5). Hence all of the spectral observations lead to the 

conclusion that the Raman band at 538 cm-1 results from scattering by Cl2(aq). 

Having established that equation 7 is the correct description of the low pH event, it is 

possible to use data from Figure 3 to estimate the concentration-based equilibrium constant 

K2

Cl2. By rearranging equation 10, we establish that, when [HOCl] = [Cl2]: 

pKCl 2 = pH – log [Cl-] (17) 2

The three speciation plots in Figure 3, and the assumed chloride ion concentrations, provide three 

independent estimates of K2

Cl2. The mean (± sample standard deviation) of these three estimates 

was 2.56 (± 1.01) x 10-4 M2

. However, note that this equilibrium constant should be considered 

an estimate only. For example, it was based on only three determinations. Also, the Clconcentration was not confirmed by measurements in these ECS solutions. Instead, it was 

assumed to be equal (as stated for the commercial product) to the known free-chlorine 

22
concentration. Furthermore, it is a concentration-based equilibrium constant. However, we note 

that our value is in reasonable agreement with the concentration-based value previously reported 

(42) of 6 x 10-4 M2

 for solutions with an ionic strength nearly 10 times greater than present in the 

solutions that we used in these experiments. 

CP loss was observed in the presence of both ECS and LCS chlorine solutions over the 

pH range of 2-11. In the presence of ECS (Figure 1), the observed first order loss of CP 

increased as pH decreased from 11 to 2. The insert in Figure 1 shows the loss of CP over the pH 

range of 6-11. Duirk and Collette, (14) found over this pH range, that HOCl was responsible for 

the oxidation of CP to CPO and that OCl- accelerates the hydrolysis of both CP and CPO 

chlorine assisted hydrolysis. Below pH 5, the observed CP loss rate greatly increased as pH 

decreased from 5 to 2. This increase in the observed loss rate of CP was then attributed to the 

formation of Cl2(aq). In the presence of LCS (Figure 7), the observed first order rate of CP loss 

did increase as pH decreased from 11 to 2. However, the observed rate of CP loss in the 

presence of LCS at low pH was not nearly as significant as it was in the presence of ECS. At pH 

of 2, the rate of loss of CP in LCS was about 2 orders of magnitude smaller than that observed 

for the ECS. However, this increase in the rate of CP loss with LCS is significant, and, in part, 

could be explained via the following proposed pathway. When HOCl reacts with CP, chloride 

ion is one of the products as well as a reduced sulfur species (S), equation 18. In the presence of 

excess free chlorine, 

HOCl + CP CPO + Cl-

+ S + H+ (18) 

4HOCl + S SO4

2- + 4Cl- + 4H+ (19) 

HOCl + H+

 + Cl­ Cl2(aq) + H2O (20) 

Cl2(aq) + CP products (21) 

23
S will be rapidly oxidized to sulfate (SO4

2-) resulting in four more moles of chloride for each 

mole of S oxidized (61), equation 19. With the presence of additional chloride ions in solution at 

low pH, the formation potential of Cl2(aq) is increased greatly. Therefore, the observed rate of 

CP loss in the presence of the LCS free chlorine solution increases due to the fact that aqueous 

molecular chlorine can be formed under these reaction conditions. 

3.2 Loss of CP in the Presence of Bromide and Free Chlorine 

The presence of bromide in drinking water sources has long been a concern of the water 

treatment industry because it can result in enhanced formation of potentially-hazardous DBPs 

under drinking water treatment conditions. Since free chlorine can rapidly oxidize bromide to 

hypobromous acid (HOBr) (28), it was thought that its presence might act as a catalyst in the 

transformation of CP during drinking water chlorination. 

Initial experiments to determine if bromide does accelerate the rate of CP loss were 

conducted over the pH range of 7-9, [CP] = 0.5 μM, [HOCl]T = 10 and 50 μM, and increasing 

bromide concentrations from 0-10 μM. At [HOCl]T = 10 μM (Figure 8), plotting kobs versus the 

bromide concentration was found to be linearly proportional to increasing bromide 

concentration. This was also observed at pH 8 and 9 at the higher chlorine concentration of 50 

μM (Figure 9). Experiments were also conducted at pH 7 and 50 μM chlorine concentration; 

however, the reaction rate was too rapid to accurately quantify an observed rate of CP loss. 

By comparing the slope of the regression lines at constant chlorine dose and varying pH 

(Table 2), the kobs in the presence of bromide and free chlorine was found to be pH dependent. 

The pH of the solution determines the speciation of [HOCl]T between HOCl and OCl-

. 

Experiments conducted at lower pH have a greater percentage of [HOCl]T present as HOCl 

species. Since the pKa of HOCl (7.5) is close to the pH range over which experiments are 

conducted, there is a large variation in the speciation of HOCl as pH increased from 7 to 9. At 

24
pH 7, 8, and 9, the percentage of free chlorine present as HOCl species is approximately 75%, 

23%, and 3% respectively. Since HOCl is primarily responsible for the formation of HOBr, 

lowering the pH increased the rate of HOBr formation and subsequently the kobs for CP loss. 

The formation of HOBr, due to HOCl oxidizing Br-, could be responsible for the increase 

in kobs with respect to increasing bromide concentration. At the lower bromide concentrations 

(0.1 and 1.0 μM) and chlorine concentration of 10 μM, CP is competitive with bromide for 

oxidation by HOCl. Table 3 shows the rate coefficients for HOCl reacting with both CP and 

bromide. The reaction rate coefficient for HOCl with bromide is approximately 3 times faster 

than HOCl with CP. As bromide concentrations increased to 5.0 and 10.0 μM, active bromine 

becomes present in greater abundance than CP. Therefore, HOCl is more likely to oxidize 

bromide than CP when chlorine to bromide molar ratios approach 1. At low bromide 

concentrations and in the presence of excess free chlorine, the oxidation of CP is most likely 

carried out by both HOCl and HOBr. Therefore, bromide appears to act as a catalyst in the 

oxidation of CP by free chlorine. 

While all observed rates were faster with a 50 μM chlorine concentration as compared to 

the analogous experiment with [HOCl]T = 10.0 μM, the magnitude of bromide-catalyzed CP 

oxidation on kobs compared to the control experiments was greater at the lower chlorine dose for 

pH 8 and 9. At the 10 μM chlorine concentration, addition of 10 μM of bromide induced a ~20 

fold and ~30 fold increase in kobs for pH 9 and 8 respectively, whereas addition of 10 μM of 

bromide at the 50 μM chlorine dose only increased the kobs ~4 fold and 10 fold at pH 9 and 8 

respectively. This indicates that the magnitude of bromide catalyzed oxidation of CP is 

controlled by the ratio of HOCl to bromide as well as pH. When a large excess of HOCl exists, 

it is the primary oxidant responsible for the loss of CP. However, when [HOCl]:[bromide] molar 

25
ratios approach unity, HOBr becomes the primary oxidant. These two reaction scenarios are 

depicted schematically in Figure 10. 

The relationship between kobs and bromide concentration shows that increasing bromide 

is linearly proportional to the observed rate of CP loss in the presence of free chlorine. This 

result indicates that HOBr is likely the dominant oxidant responsible for CP degradation at 

higher bromide concentrations. The linearity of the trend also indicates that bromide acts as a 

catalyst in the oxidation reaction between HOCl and CP. Bromide is a catalyst in this pathway 

because it does not appear in the balanced equation for the reaction and after being reduced by 

chlorpyrifos, HOBr yields H+

 and Br-. The regenerated bromide ion can then be re-oxidized by 

HOCl and participate in additional reactions with CP. The slope of the regression line for kobs 

vs. [Br-] concentration was observed to be dependent on the concentration of free chlorine, pH, 

and the concentration of bromide (Table 2). All three of these parameters affect the rate of 

HOBr formation. 

A model was developed to predict the loss of CP in the presence of free chlorine and 

bromide. From the stoichiometric equations in Table 3, a system of ordinary differential 

equations (ODEs), equations 22-28, were incorporated into a model developed to predict CP loss 

pathways in buffered deionized water (14). Key assumptions in this model are that HOBr will 

only react with CP resulting in CPO, and that all HOBr that reacts with CP results in 

d[HOCl] T − − = −k [HOCl][CP] − k [OCl ][CP] − k [OCl ][CPO] dt HOCl,CP OCl,CP OCl,CPO (22) 

− 4 * k HOCl,S[HOCl][S] − k HOCl,Br [HOCl][Br − ] 

d[HOBr]T = k HOCl,Br [HOCl][Br − ] − k HOBr,CP [HOBr][CP] (23) dt 

d[Br-] = −k HOCl,Br [HOCl][Br − ] − k HOBr,CP [HOBr][CP] (24) dt 

26
d[S] = k HOCl,CP [HOCl][CP] + k HOBr,CP [HOBr][CP] − k HOCl,S[HOCl][S] (25) dt 

d[CP] −

dt = −k HOCl,CP [HOCl][CP] − k h,CP [CP] − k OCl,CP [OCl ][CP] (26) 

− k HOBr,CP [HOBr][CP] 

d[CPO] −

dt = k HOCl,CP [HOCl][CP] − k h,CPO [CPO] − k OCl,CPO [OCl ][CPO] (27) 

+ k HOBr,CP [HOBr][CP] 

d[TCP] = k h,CP [CP] + k OCl,CP [OCl − ][CP] + k h,CPO [CPO] + k OCl,CPO [OCl − ][CPO] (28) dt 

the regeneration of bromide in the aqueous system. Also, that oxidation of the reduced sulfur 

species (S) can be adequately predicted by the rate coefficient determined for HOCl reaction 

with sulfite (SO3

2-) (61). ScientistTM, an ODE solver by Micromath (Salt Lake City, UT), was 

used to fit the rate coefficient for CP oxidation by hypobromous acid (kHOBr,CP). ScientistTM uses 

a modified Powell algorithm to minimize the unweighted sum of the squares of the residual error 

between the predicted and experimentally observed values to estimate specific parameters in the 

model. 

To aid in parameter estimation, the loss of CP in the presence of HOBr was examined at 

pH 6.5 (Figure 11). At this pH, over 99% of the active bromine present will be in the HOBr 

form due to its pKa of 8.8 (28). The reaction of HOBr with CP was expected to be 1 to 3 orders 

of magnitude faster than HOCl with CP. This has been observed by others examining the 

reactivity of HOCl and HOBr with pyrene, NOM, and sulfite (30,38,61,62). The reaction was 

essentially complete after only 10 seconds. The second order rate coefficient for this reaction 

was determined by nonlinear regression analysis using Sigma Plot version 8.0 (Point Richmond, 

CA). The kHOBr,CP was found to be 1.14 x 109

 M-1h-1, which is three orders of magnitude greater 

than HOCl oxidation of CP (kHOCl,CP = 1.72 x 106

 M-1h-1). 

27
The system of ODEs in equations 22-28 were used to model the loss of CP in the 

presence of free chlorine and bromide. In Figures 12-14, model results are shown for CP loss as 

well as the formation of CPO and TCP at pH 8 and [HOCl]T = 10, 25 and 50 μM in the absence 

of bromide. The model adequately predicts loss of CP and the formation of CPO and TCP. The 

only significant change from the model presented here verses the original model by Duirk and 

Collette (14) is the incorporation of oxidation of the reduced sulfur species, equation 25. Since 

no rate coefficients exist in literature for the oxidation of S by HOCl, it was assumed that the rate 

coefficient for HOCl reacting with SO3

2- was sufficiently large enough to predict the loss of free 

chlorine due to the reaction with S. Under these experimental conditions, the model appears 

quite capable of predicting CP degradation pathways in the presence of free chlorine. 

At pH of 8, two different sets of experiments were conducted to determine if the 

experimentally determined rate coefficient for kHOBr,CP was adequate to describe the loss of CP in 

the presence of bromide and free chlorine. The first set of experiments was conducted at a 

constant bromide concentration of 1 μM, with increasing free chlorine concentrations from 10 to 

50 μM. The second set of experiments was conducted at constant free chlorine concentration of 

50 μM, with increasing bromide concentrations from 1 to 10 μM. Using the experimentally 

determined kHOBr,CP, CP loss was found to be significantly over-predicted for both sets of 

experiments. This could have been due to the fact that all active bromine being formed was 

assumed to only react with CP. However, HOBr does react with reduced sulfur species, such as 

sulfite, more rapidly than HOCl (61,62). 

In the current model, all reduced sulfur species are assumed to be oxidized by HOCl and 

all active bromine reacts with CP. Since reaction rate coefficients for HOCl and HOBr with all 

the reduced sulfur species from S to SO3

2- are not available in literature, it is not possible to 

28
proportion the amount of HOBr that could potentially react with either CP or one of the reduced 

sulfur species. Therefore, an apparent second order rate coefficient (kapp_HOBr,CP) was determined 

by pooling these data sets to predict the loss of CP in the presence of free chlorine and bromide 

for these experimental conditions (Figures 15-20). The kapp_HOBr,CP was found to be 6.01(±0.40) 

x 107 M-1h-1 (95% confidence interval shown in parentheses), which is two orders of magnitude 

less than the experimentally determined kHOBr,CP. Although the apparent rate coefficient is 

significantly smaller than the experimentally determined rate coefficient, it was able to 

adequately predict the loss of CP and the formation of CPO and TCP under these experimental 

conditions using this value for the apparent rate coefficient. Also, it still falls within the 

expected range of 1 to 3 orders of magnitude faster than its the chlorinated analog (61,62). The 

reaction with the HOBr and CP may have been too fast to be adequately captured by the 

experimental procedure used to investigate HOCl reactions with CP below pH of 8. This might 

explain, in part, the discrepancy between the experimentally determined and the apparent rate 

coefficients. In any event, these results indicate that active bromine pathways in the presence of 

OP pesticides under drinking water treatment conditions still needs to be investigated further. 

3.3 Loss of CP in the Presence of Free Chlorine and NOM 

The presence of NOM could potentially decrease the rate of CP degradation by acting as 

a sink for oxidants used in drinking water treatment. Conventional physical-chemical water 

treatment processes such as coagulation and lime softening remove a portion of the hydrophobic 

fraction of the NOM (63,64). More advanced water treatment processes like granular activated 

carbon (GAC) and membranes are capable of removing more hydrophilic NOM fractions (65). 

However, few community water systems employ these more advanced technologies. The ACC 

water treatment plant uses only conventional surface water treatment processes. Therefore, only 

a portion of the hydrophobic NOM fraction will be removed. Figure 21 shows the chlorine 

29
demand studies conducted over the pH range of 7-9. Over the first half hour, approximately 12% 

of the free chlorine initially present is consumed by the NOM. After the initial chlorine demand 

of the NOM is satisfied, the loss of free chlorine attributable to reaction with NOM is relatively 

slow (kobs = 0.012 h-1) over the pH range of 7-9. These results suggest that the NOM present in 

the ACC water may affect the rate of CP transformation over the first half hour in the presence of 

free chlorine. 

To examine the effect of NOM on the transformation of CP, experiments were conducted 

over the pH range of 7-9 and in the presence of 50 μM free chlorine. Figure 22 shows the 

observed first order rate of CP loss with and without the presence of NOM. In the presence of 

ACC NOM, the observed loss of CP was very similar to the control experiments. It appears that 

the conventional treatment performed at the ACC removed a portion of the hydrophobic NOM 

fraction that is usually associated with significant chlorine demand (36). Therefore, it might 

possible for the current model (developed with experiments in laboratory water) to accurately 

predict the transformation of CP in this low humic water. 

Figures 23-27 show the experimental and model results for the loss of CP and formation 

of CPO and TCP in ACC water for a variety of experimental conditions. At pH 7 and 7.5 

(Figures 23 and 24), the model adequately predicted CP loss and the formation of CPO and TCP. 

However, at these pHs, the model did slightly over-predict CP loss. This could have been due to 

bromide being present at 0.35 μM, but this low bromide concentration was not expected to 

significantly affect the rate of CP transformation (See Figures 8 and 9). On the other hand, the 

model performed very well predicting the formation and stability of CPO, the more toxic 

transformation product, as well as TCP. Over the pH range of 8-9 (Figures 25-27), the model 

adequately predicted the loss of CP over the experimental time frame. At pH values greater than 

30
7.5, CPO becomes susceptible to chlorine-assisted hydrolysis due to the greater abundance of the 

hypochlorite ion. The model not only predicted CPO formation, but it also showed that CPO 

was slowly being transformed to TCP over the course of the experiment at pH 8-9. This appears 

to validate the ability of the model to adequately predict the loss of CP in the presence of free 

chlorine and NOM, as well as the ability of the model to predict the formation and stability of 

CPO in the presence of free chlorine. 

31
4 CONCLUSIONS 

The goal of this work was to determine the active oxidant species responsible for the 

degradation of CP under drinking water treatment conditions. Using Raman spectroscopy, the 

third species in free chlorine solutions at low pH was identified, which is responsible for the very 

rapid oxidation of CP at low pH. The Raman band at 538 cm-1 appeared as the pH of the ECS 

free chlorine solution was lowered and indicated the presence of a third species in addition to the 

hypochlorite ion and hypochlorous acid. The speciation profiles of ECS solutions at different 

concentrations showing variation in the low-pH “apparent pKa,” and disappearance of the band 

at 538 cm-1 in LCS solutions indicated that the third species in solution is Cl2(aq) and not 

H2OCl+ as has been frequently assumed. 

In the presence of bromide, the transformation of CP was found to be accelerated upon 

chlorination. HOCl rapidly oxidizes bromide to HOBr, which is a stronger oxidant. Once HOBr 

reacts with CP, bromide is then regenerated in the aqueous system (i.e., bromide is a catalyst 

increasing the rate of CP transformation). In the presence of HOBr, the reaction rate coefficient 

of CP with HOBr was found to be three orders of magnitude greater than the rate coefficient 

found for HOCl with CP. However, the experimentally determined rate coefficient was found to 

be too large to accurately predict the loss of CP and the formation of CPO and TCP under the 

experimental conditions used in this study. An apparent second order rate coefficient 

(kapp_HOBr,CP), which was two orders of magnitude less than the experimentally determined rate 

coefficient, was found to adequately fit the experimental results. This was assumed to be due to 

HOBr reacting with not only CP but with reduced sulfur species evolving from the initial 

oxidation of CP by either HOCl or HOBr. Future work will be done to further elucidate active 

bromine pathways in the presence of free chlorine and CP. 

32
The role of NOM on the transformation of CP was investigated using Athens-Clarke 

County water, which was collected prior to chlorination. Under the experimental conditions 

used in this study, the established model, initially developed to predict the transformation 

pathways of CP in laboratory water, was found to adequately predict the loss of CP and the 

formation of CPO and TCP in the ACC water. It was assumed that due to the low humic content 

of the ACC water, the chlorine demand that was exerted by the NOM was not sufficient to 

significantly affect the rate of CP transformation. Also, the model was able to predict the 

formation and stability of CPO over the entire pH range. In waters that are low in aquatic humic 

substances, the model appears to adequately predict CP transformation pathways. 

The work presented here shows the application of a model, developed to predict the loss 

of CP in buffered aqueous systems, to predict the loss of CP over a greater range of conditions 

that are more closely associated with drinking water treatment. However, the need still exists to 

systematically investigate chlorination of OP pesticides as a class and determine if reactivity 

with aqueous chlorine is related in a predictable way to OP chemical structure. If a relationship 

between reactivity and pesticide structure exists, the ability to predict the fate of OP pesticides in 

the presence of free chlorine would greatly aid regulators in assuring that tolerances to OP 

pesticide exposure from drinking water are not exceeded. After developing predictive tools for 

the entire class of OP pesticides, we plan to address other important classes of pesticides and 

toxic chemicals. 

33
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37
TABLES 

38
Table 1 NOM and source water characteristics for the ACC water collected prior to 

chlorination. 

Source 

ACC 

SUVA254

(L/m-mg) 

1.61 

SUVA280 

(L/m-mg) 

1.07 

a

[NH3]T

(μM) 

2.06 

[Cl-] 

(μM) 

93.46 

[NO2 

-

] 

(μM) 

0.04 

[Br-] 

(μM) 

0.35 

b

[PO4]T

(μM) 

c

ND 

[SO4 

2-] 

(μM) 

8.30 a

[NH3]T = [NH4

+

] + [NH3]

b

[PO4]T = [H3PO4] + [H2PO4

-

] + [HPO4

2-] + [PO43-] c

ND = Non-detect 

39
Table 2 The slopes of kobs vs. [Br-

] for each pH and [HOCl] concentration. 

pH [HOCl] (μM) kobs (h-1) vs. [Br-

] 

7 10 10.3 

8 10 2.7 

9 10 0.4 

8 50 5.3 

9 50 0.6 

40
Table 3 Stoichiometric equations and rate coefficients used in the chlorpyrifos degradation 

pathway model. 

Reaction Stoichiometry Rate/Equilibrium Coefficient Reference 

(25 °C) 

1 kr + − kHOCl,CP = 1.72 ×106 M-1h-1

HOCl + CP ⎯⎯→CPO + H + Cl + S 

2 k kHOCl,S = 2.74 ×1012 M-1h-1

HOCl, S + − 2­ 4HOCl + S ⎯⎯⎯⎯→4H + 4Cl + SO4 

3 − kHOCl,Br − kHOCl,Br = 5.6 × 106

 M-1h-1

HOCl + Br ⎯⎯⎯⎯→HOBr + Cl 

4 kHOBr,CP + − kHOBr,CP = see text HOBr + CP ⎯⎯⎯⎯→CPO + H + Cl 

5 kh,

⎯CP kh,CP = kN,CP + kB,CP[OH-

] CP ⎯⎯ →TCP 

kN,CP = 3.72×10-4 h-1 

kB,CP = 37.0 M-1h-1 

6 kh,CPO

⎯⎯ kh,CPO = kN,CPO + kB,CPO[OH-

] CPO ⎯⎯ →TCP 

kN,CPO = 2.13 ×10-3 h-1 

kB,CPO = 230 M-1h-1 

7 kOCl,CP kOCl,CP = 990 M-1h-1

CP + OCl- ⎯⎯⎯⎯→TCP + Cl8 - kOCl,CPO - kOCl,CPO = 1340 M-1h-1

CPO + OCl ⎯⎯⎯⎯→TCP + Cl 

9 HOCl H+

 + OCl- pKa = 7.5 

(14) 

(61) 

(28) 

This work

(66) 

(14) 

(14) 

(14) 

(59) 

41
FIGURES 

42
kobs (h -1 ) 

0 

500 

1000 

1500 

2000 

pH 

6 7 8 9 10 11 12 

k obs (h -1 ) 

0 

2 

4 

6 

8 

10 

0 2 4 6 8 10

pH 

Figure 1 Observed first order loss rate coefficients for chlorpyrifos over the pH range of 2­

11 in the presence of ECS. [CP]o = 0.5 μM, [Buffer]T = 10 mM, Temperature = 

25±1°C, and [HOCl]T = 10 μM. 

43

12 
N

o

r

malized Band Area

N

o

r

malized Band Area

A

1

0.5

0

02468 10 12

pH

OCl- pH B

11.96

Inte

n

sit

y Inte

n

sit

y

7.46

HOCl

4.59

low pH 

form

800 750 700 650 600 550 500 450

Raman Shift (cm-1 Raman Shift (cm ) -1

2.73

Figure 2 (A.) Speciation plot of the 26.2 mM commercial free chlorine solution (the ECS 

solution) generated by monitoring the normalized areas of the Raman bands from 

OCl- at 711 cm-1 (>), HOCl at 725 cm-1() and the low pH species at 538 cm-1 

(‚). (B.) Background-subtracted Raman spectra of ECS solutions. The OCl- band 

is observed in the top two spectra and the HOCl band is observed in the bottom 

three spectra. An additional band for the low pH form appears as the pH is lowered 

below 5. 

44
N

ormalized B

and Area 

26.2 mM 

13.1 mM 

6.6 mM 

1 

0.5 

0 

1 

0.5 

0 

1 

0.5 

0 

N

ormalized B

and Area

26.2 mM

13.1 mM

6.6 mM

1 

0.5

0 

1 

0.5

0 

1 

0.5

0 

0 2 4 6 8 10 12

pH

Figure 3 Speciation plots of commercial free chlorine solutions (ECS solutions) generated 

by monitoring the normalized areas of the Raman bands from OCl- at 711 cm-1 (>, 

dash-dot model line), HOCl at 725 cm-1 (, dashed model line) and the low pH 

species at 538 cm-1 (‚, solid model line). Vertical lines indicate the apparent pKas 

at 13.1 mM. The pKa for the protonation of OCl- to HOCl occurs at 7.50 for all 

three concentrations. The apparent pK at lower pH varies from (from top to 

bottom): 2.11, 1.84, and 1.25. 

45
Intensity 

(A) 

(B) 

Intensit

y

(A)

(B)

800 750 700 650 600 550 500 450

Raman Shift (cm-1 Raman Shift (cm ) -1

Figure 4 Raman spectrum acquired from (A) a buffered aqueous solution at pH 1.45 through 

which chlorine gas was bubbled (BCS) and (B) a pH 1.63 solution containing 13.1 

mM commercial free chlorine (ECS). The band at 538 cm-1 is due to the presence 

of Cl2 (aq) in each solution. 

46
Intensity

(A) 

(B) 

(A)

(B)

800 750 700 650 600 550 500 450

Raman Shift (cm-1 Raman Shift (cm ) -1

Figure 5 Raman spectra comparing 13.1 mM free chlorine solutions with (A) equimolar 

chloride ion at pH 0.75 (ECS) and (B) free chlorine solution with reduced chloride 

ion at pH 0.80 (LCS). The spectrum of the solution containing equimolar chloride 

ion has a much larger band at 538 cm-1 (due to aqueous Cl2) than the solution with 

greatly reduced chloride concentration. 

47
0 

0.5 

1 

-2 0 2 4 6 8 10 12 

pH

Nor

malized Peak Area 

0

0.5

1

-2 0 2 4 6 8 10 12

pH

Nor

malized Peak Area

Figure 6 Speciation profile of the LCS free chlorine solution with reduced chloride ion 

concentration generated by monitoring the normalized areas of the Raman bands 

from HOCl at 725 cm-1 () and OCl- at 711 cm-1 (>). The predominant species in 

solution at low pH remains HOCl throughout, and a significant drop in its total area 

is not observed as pH decreases. 

48
kobs (h -1 ) 

0 

10 

20 

30 

40 

0 2 4 6 8 10 12

pH 

Figure 7 Observed first order loss rate coefficients for chlorpyrifos over the pH range of 2­

11 in the presence of LCS. [CP]o = 0.5 μM, [Buffer]T = 10 mM, Temperature = 

25±1°C, and [HOCl]T = 10 μM. 

49
kobs (h -1 ) 

0 

25 

50 

75 

100 

125 

pH 7 

pH 8 

pH 9 

0 2 4 6 8 10 12

Bromide (μM) 

Figure 8 Observed first order rate coefficients for CP loss in the presence of increasing 

bromide concentrations at [HOCl]T = 10 μM over the pH range of 7-9. [CP]o = 0.5 

μM, [Buffer]T = 10 mM, and Temperature = 25 o

C. 

50
kobs (h-1) 

70 

pH 8

60 pH 9

50

40

30

20

10

0

0 2 4 6 8 10

Bromide (μM)

Figure 9 Observed first order rate coefficients for CP loss in the presence of increasing 

bromide concentrations at [HOCl]T = 50 μM at pH 8 and 9. [CP]o = 0.5 μM, 

[Buffer]T = 10 mM, and Temperature = 25 o

C. 

51

12 
CP 

HOCl BrHOBr 

CP 

BrCPO ClFigure 10 Two possible oxidation pathways for CP at high and low [HOCl]:[Br-

] ratios. The 

pathway on the left shows the predominant pathway when [HOCl][Br-

] and the 

pathway on the right shows the predominant pathway when [HOCl]:[Br-

] 

approaches 1. 

52
CP (

μM) 

0.0 

0.1 

0.2 

0.3 

0.4 

0.5 

0 2 4 6 8 10 

Time (seconds) 

Figure 11 Loss of CP in the presence of hypobromous acid at pH 6.5. [CP]o = 0.42 μM, 

[HOBr]T = 10 μM, [Buffer]T = 10 mM, and Temperature = 25 o

C. 

53

12 
Co

n

c

e

ntr

atio

n (

μM) 

0.6 

CP 

CPO 0.5 

TCP 

0.4 

0.3 

0.2 

0.1 

0.0 

0.0 0.5 1.0 1.5 2.0 2.5 

Time (hours) 

Figure 12 CP degradation in the presence of free chlorine at pH 8.0 in the absence of 

bromide. [CP]o = 0.39 μM, [HOCl]T = 10 μM, [PO4]T = 10 mM, and Temperature 

= 25 o

C. Lines represent model results. 

54
Co

n

c

e

ntr

atio

n (

μM) 

0.6 

CP 

CPO 0.5 

TCP 

0.4 

0.3 

0.2 

0.1 

0.0 

0.00 0.25 0.50 0.75 1.00 1.25 

Time (hours) 

Figure 13 CP degradation in the presence of free chlorine at pH 8.0 in the absence of 

bromide. [CP]o = 0.42 μM, [HOCl]T = 25 μM, [PO4]T = 10 mM, and Temperature 

= 25 o

C. Lines represent model results. 

55
Co

n

c

e

ntr

atio

n (

μM) 

0.6 

CP 

CPO 0.5 

TCP 

0.4 

0.3 

0.2 

0.1 

0.0 

0.00 0.25 0.50 0.75 

Time (hours) 

Figure 14 CP degradation in the presence of free chlorine at pH 8.0 in the absence of 

bromide. [CP]o = 0.42 μM, [HOCl]T = 50 μM, [PO4]T = 10 mM, and Temperature 

= 25 o

C. Lines represent model results. 

56
Co

n

c

e

ntr

atio

n (

μM) 

0.0 

0.1 

0.2 

0.3 

0.4 

0.5 

0.6 

CP 

CPO 

TCP 

0.0 0.1 0.2 0.3 0.4 

Time (hours) 

Figure 15 CP degradation in the presence of free chlorine and bromide at pH 8.0. [CP]o = 

0.51 μM, [HOCl]T = 10 μM, [Br-

] = 1 μM, [PO4]T = 10 mM, and Temperature = 25 o

C. Lines represent model results. 

57
Co

n

c

e

ntr

atio

n (

μM) 

0.0 

0.1 

0.2 

0.3 

0.4 

0.5 

0.6 

CP 

CPO 

TCP 

0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 

Time (hours) 

Figure 16 CP degradation in the presence of free chlorine and bromide at pH 8.0. [CP]o = 

0.48 μM, [HOCl]T = 25 μM, [Br-

] = 1 μM, [PO4]T = 10 mM, and Temperature = 25 o

C. Lines represent model results. 

58
Co

n

c

e

ntr

atio

n (

μM) 

0.0 

0.1 

0.2 

0.3 

0.4 

0.5 

0.6 

CP 

CPO 

TCP 

0.00 0.02 0.04 0.06 0.08 

Time (hours) 

Figure 17 CP degradation in the presence of free chlorine and bromide at pH 8.0. [CP]o = 

0.49 μM, [HOCl]T = 50 μM, [Br-

] = 1 μM, [PO4]T = 10 mM, and Temperature = 25 o

C. Lines represent model results. 

59
Co

n

c

e

ntr

atio

n (

μM) 

0.0 

0.1 

0.2 

0.3 

0.4 

0.5 

0.6 

CP 

CPO 

TCP 

0.0 0.1 0.2 0.3 0.4 

Time (hours) 

Figure 18 CP degradation in the presence of free chlorine and bromide at pH 8.0. [CP]o = 

0.35 μM, [HOCl]T = 50 μM, [Br-

] = 1 μM, [PO4]T = 10 mM, and Temperature = 25 o

C. Lines represent model results. 

60
Co

n

c

e

ntr

atio

n (

μM) 

0.0 

0.1 

0.2 

0.3 

0.4 

0.5 

0.6 

CP 

CPO 

TCP 

0.000 0.025 0.050 0.075 0.100 0.125 0.150 

Time (hours) 

Figure 19 CP degradation in the presence of free chlorine and bromide at pH 8.0. [CP]o = 

0.43 μM, [HOCl]T = 50 μM, [Br-

] = 5 μM, [PO4]T = 10 mM, and Temperature = 25 o

C. Lines represent model results. 

61
Co

n

c

e

ntr

atio

n (

μM) 

0.0 

0.1 

0.2 

0.3 

0.4 

0.5 

0.6 

CP 

CPO 

TCP 

0.00 0.02 0.04 0.06 0.08 

Time (hours) 

Figure 20 CP degradation in the presence of free chlorine and bromide at pH 8.0. [CP]o = 

0.39 μM, [HOCl]T = 50 μM, [Br-

] = 10 μM, [PO4]T = 10 mM, and Temperature = 

25 o

C. Lines represent model results. 

62
[HOCl]T (mm) 

0.00 

0.01 

0.02 

0.03 

0.04 

0.05 

0.06 

pH 7 

pH 8 

pH 9 

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 

Time (hours) 

Figure 21 Chlorine demand of ACC water over the pH range of 7-9. [HOCl]T = 50 μM, 

[Buffer]T = 10 mM, [DOC] = 1.1 mg-C/L, and Temperature = 25 o

C. 

63
Figure 22 CP degradation in the presence of free chlorine and ACC water over the pH range 

of 7-9. [CP]o ≈ 0.5 μM, [HOCl]T = 50 μM, [DOC] = 1.1 mg-C/L, [Buffer]T = 10 

mM, and Temperature = 25 o

C. Lines represent model results. 

30 

35 

Control 

ACCF 

25 

20 

(h 

-1 ) k 

15

obs

10 

5 

6.5 

0 

7.0 7.5 

pH 

8.0 8.5 9.0 9.5 

64
Co

n

c

e

ntr

atio

n (

μM) 

0.0 

0.1 

0.2 

0.3 

0.4 

0.5 

0.6 

CP 

CPO 

TCP 

0.0 0.1 0.2 0.3 0.4 0.5 0.6 

Time (hours) 

Figure 23 CP degradation in the presence of free chlorine and ACC water at pH 7. [CP]o = 

0.46 μM, [HOCl]T = 50 μM, [DOC] = 1.1 mg-C/L, [Buffer]T = 10 mM, and 

Temperature = 25 o

C. Lines represent model results. 

65
Y Data 

0.0 

0.1 

0.2 

0.3 

0.4 

0.5 

0.6 

CP 

CPO 

TCP 

0.00 0.05 0.10 0.15 0.20 0.25 0.30 0.35 

Time (hours) 

Figure 24 CP degradation in the presence of free chlorine and ACC water at pH 7.5. [CP]o = 

0.35 μM, [HOCl]T = 50 μM, [DOC] = 1.1 mg-C/L, [Buffer]T = 10 mM, and 

Temperature = 25 o

C. Lines represent model results. 

66
Co

n

c

e

ntr

atio

n (

μM) 

0.0 

0.1 

0.2 

0.3 

0.4 

0.5 

0.6 

CP 

CPO 

TCP 

0.00 0.25 0.50 0.75 1.00 1.25 

Time (hours) 

Figure 25 CP degradation in the presence of free chlorine and ACC water at pH 8.0. [CP]o = 

0.42 μM, [HOCl]T = 50 μM, [DOC] = 1.1 mg-C/L, [Buffer]T = 10 mM, and 

Temperature = 25 o

C. Lines represent model results. 

67
0.5 

0.6 

CP 

CPO 

TCP 

μM)

0.4 

Co

n

c

e

ntr

atio

n (

0.3 

0.2 

0.1 

0.0 

0.0 

0.5 

Time (hours) 

1.0 1.5 2.0 2.5 

Figure 26 CP degradation in the presence of free chlorine and ACC water at pH 8.5. [CP]o = 

0.35 μM, [HOCl]T = 50 μM, [DOC] = 1.1 mg-C/L, [Buffer]T = 10 mM, and 

Temperature = 25 o

C. Lines represent model results. 

68
Co

n

c

e

ntr

atio

n (

μM) 

0.6 

CP 

CPO 

0.5 TCP 

0.4 

0.3 

0.2 

0.1 

0.0 

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 

Time (hours) 

Figure 27 CP degradation in the presence of free chlorine and ACC water at pH 9.0. [CP]o = 

0.5 μM, [HOCl]T = 50 μM, [DOC] = 1.1 mg-C/L, [Buffer]T = 10 mM, and 

Temperature = 25 o

C. Lines represent model results. 

69