EPA/600/R-08/089 | September 2008 | www.epa.gov/athens 

Fate of High Priority Pesticides 

During Drinking Water 

Treatment 

Ecosystems Research Division, Athens, GA 30605 

National Exposure Research Laboratory 

Office of Research and Development 
EPA/600/R-08/089 

September 2008 

 Fate of High Priority Pesticides During 

 Drinking Water Treatment 

by 

Stephen E. Duirk, †

Lisa M. Desetto, and †Gary M. Davis 

National Exposure Research Laboratory 

Ecosystems Research Division 

Athens, GA 

†

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. 

i
ABSTRACT 

The fate of organophosphorus (OP) pesticides in the presence of chlorinated oxidants was 

investigated under drinking water treatment conditions. In the presence of aqueous chlorine, 

intrinsic rate coefficients were found for the reaction of hypochlorous acid (kHOCl,OP) and 

hypochlorite ion (kOCl,OP) with eight OP pesticides. The reaction of hypochlorous acid (HOCl) 

with each OP pesticide was relatively rapid at near neutral pH, kHOCl,OP = 0.86 – 3.56 x 106 M-1h1

. The reaction of HOCl with OP pesticides occurs at the thiophosphate (P=S) moiety resulting 

in the formation of the corresponding oxon (P=O), which is more toxic than the parent OP 

pesticide. Hypochlorite ion (OCl-

) was found not to oxidize the pesticide but act like a 

nucleophile accelerating hydrolysis, kOCl,OP = 37.3 – 15,908.9 M-1h-1. Both the kHOCl,OP and the 

kOCl,OP were found to correlate well with molecular descriptors within each subgroup of OP 

pesticide class. The most commonly detected OP pesticides in drinking water sources were then 

investigated in the presence of monochloramine (NH2Cl). Monochloramine was found not to be 

very reactive with OP pesticides, kNH2Cl,OP = 10.6 – 21.4 M-1h-1. Dichloramine (NHCl2) was 

found to be two orders of magnitude more reactive with the OP pesticides investigated than 

monochloramine, kNHCl2,OP = 1995.0 – 2931.9 M-1h-1. The reactivity of the three chlorinated 

oxidants was then found to correlate with half-wave potentials (E1/2) for each OP pesticide 

respectively. A model was developed to predict the transformation of OP pesticides in the 

presence of chlorinated oxidants. With hydrolysis rate coefficients, the transformation of OP 

pesticides under drinking water treatment condition was found to be adequately predicted over 

the pH range of 6.5-9. The structure-activity relationships and mechanistic models developed 

here could be used by regulators to determine if drinking potable water contaminated with OP 

pesticides represents significant risk to a receiving population serviced by a community water 

system. 

ii
ACKNOWLEDGEMENT 

The authors would like to thank Mr. Jimmy Avants and Dr. Lidia Samarkina their technical 

assistance, Dr. Paul Winget for calculating highest occupied molecular orbital (EHOMO) energies 

for the organophosphorus pesticides, Dr. Wayne Garrison and Dr. Jackson Ellington for their 

consultation and expertise, as well as Dr. Craig Adams at the Missouri University of Science and 

Technology and Dr. David Sedlak at the University of California, Berkeley for agreeing to 

review this report. Finally, the authors would like to thank Dr. Mac Long and Dr. Eric Weber for 

their continued support of this work at the Ecosystems Research Division in Athens, GA. 

iii
TABLE OF CONTENTS 

LIST OF TABLES.......................................................................................................................... v 

LIST OF FIGURES ....................................................................................................................... vi 

EXECUTIVE SUMMARY ........................................................................................................... ix 

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

2 EXPERIMENTAL PROCEDURES....................................................................................... 5 

2.1 Materials ......................................................................................................................... 5 

2.2 Methods........................................................................................................................... 5 

2.2.1 Chlorination and Hydrolysis of OP Pesticides........................................................ 5 

2.2.2 Chloramination of OP Pesticides............................................................................ 7 

2.3 OP Pesticide Analysis..................................................................................................... 8 

2.4 Numerical Calculations and Parameter Estimation ........................................................ 8 

3 RESULTS AND DISCUSSION............................................................................................. 9 

3.1 Hydrolysis of select OP pesticides.................................................................................. 9 

3.2 OP Pesticide Degradation in the Presence of Aqueous Chlorine ................................. 11 

3.3 OP Oxon Chlorine-Assisted Hydrolysis....................................................................... 16 

3.4 OP Pesticide-Chlorine Model Validation ..................................................................... 17 

3.5 OP Pesticide Degradation in the Presence of Monochloramine ................................... 17 

4 CONCLUSION..................................................................................................................... 24 

5 REFERENCES ..................................................................................................................... 27 

TABLES ....................................................................................................................................... 31 

FIGURES...................................................................................................................................... 39 

iv
LIST OF TABLES 

Table 1 Stoichiometric equations and coefficients in the monochloramine autodecomposition 

model........................................................................................................................... 32 

Table 2 Structures and some chemical properties of select OP Pesticides (48). ..................... 33 

Table 3 Neutral and alkaline hydrolysis rate coefficients for 8 OP pesticides and 3 oxon 

transformation products. 95% confidence intervals shown in parentheses. .............. 34 

Table 4 Stoichiometric equations used in the chlorine-OP pesticide transformation pathway 

model........................................................................................................................... 35 

Table 5 Oxidation and chlorine-assisted hydrolysis rate coefficients for 8 OP pesticides and 3 

oxon transformation products. 95% confidence intervals shown in parentheses. ..... 36 

Table 6 Differential equations for the monochloramine autodecomposition model. .............. 37 

Table 7 Intrinsic rate coefficients for monochloramine and dichloramine reacting with CP, 

DZ, and MA. 95% confidence intervals shown in parentheses. ................................ 38 

v
LIST OF FIGURES 

Figure 1 Hydrolytic behavior of chlorethoxyfos (CE) over the pH range of 3-11. Insert: slope 

represents the alkaline hydrolysis rate coefficient. [CE]o = 0.5 μM, [Buffer]T = 10 

mM, and Temperature = 25±1°C. Error bars represent 95% confidence intervals.... 40 

Figure 2 Hydrolytic behavior of tebupirimfos (TE) over the pH range of 3-10. Insert: slope 

represents the alkaline hydrolysis rate coefficient. [TE]o = 0.5 μM, [Buffer]T = 10 

mM, and Temperature = 25±1°C. Error bars represent 95% confidence intervals.... 41 

Figure 3 Hydrolytic behavior of methidathion (ME) over the pH range of 3-11. Insert: slope 

represents the alkaline hydrolysis rate coefficient. [ME]o = 0.5 μM, [Buffer]T = 10 

mM, and Temperature = 25±1°C. Error bars represent 95% confidence intervals.... 42 

Figure 4 Hydrolytic behavior of phosmet (PM) over the pH range of 2-10. Insert: slope 

represents the alkaline hydrolysis rate coefficient. [PM]o = 0.5 μM, [Buffer]T = 10 

mM, and Temperature = 25±1°C. Error bars represent 95% confidence intervals.... 43 

Figure 5 Observed first-order rate of CE loss in the presence of aqueous chlorine at pH 6.5. 

[CE]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate............................................................ 44 

Figure 6 Observed first-order rate of DZ loss in the presence of aqueous chlorine at pH 6.5. 

[DZ]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate............................................................ 45 

Figure 7 Observed first-order rate of MA loss in the presence of aqueous chlorine at pH 6.5. 

[MA]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate............................................................ 46 

Figure 8 Observed first-order rate of ME loss in the presence of aqueous chlorine at pH 6.5. 

[ME]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate............................................................ 47 

Figure 9 Observed first-order rate of PA loss in the presence of aqueous chlorine at pH 6.5. 

[PA]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate............................................................ 48 

Figure 10 Observed first-order rate of PM loss in the presence of aqueous chlorine at pH 6.5. 

[PM]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate............................................................ 49 

Figure 11 Observed first-order rate of TE loss in the presence of aqueous chlorine at pH 6.5. 

[TE]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate............................................................ 50 

vi
Figure 12 The reaction order of chlorine with 8 OP pesticides at pH 6.5. [OP]o = 0.5 μM, 

[PO4]T = 10 mM, Temperature = 25±1°C, and [HOCl]T = 10-100 μM...................... 51 

Figure 13 Second-order apparent rate coefficient for the OP pesticides at pH 6.5. [OP]o = 0.5 

μM, [PO4]T = 10 mM, Temperature = 25±1°C, and [HOCl]T = 0-100 μM. Error bars 

represent 95% confidence intervals. ........................................................................... 52 

Figure 14 The pH dependency of the first-order observed rate coefficients for the OP pesticides. 

[OP]o = 0.5 μM, [HOCl]T = 25 μM, [Buffer]T = 10 mM, and Temperature = 25 °C. 

Error bars represent 95% confidence intervals. .......................................................... 53 

Figure 15 Second-order apparent rate coefficient for the OP pesticides at pH 9. [OP]o = 0.5 μM, 

[CO3]T = 10 mM, Temperature = 25±1°C, and [HOCl]T = 0-100 μM. Error bars 

represent 95% confidence intervals. ........................................................................... 54 

Figure 16 Relationship between the kHOCl,OP with EHOMO as a function of OP pesticide subgroup. 

Error bars about the regression line represent 95% confidence intervals................... 55 

Figure 17 Relationship between the kOCl,OP with alkaline hydrolysis rate coefficient as a function 

of OP pesticide subgroup. Error bars about the regression line represent 95% 

confidence intervals. ................................................................................................... 56 

Figure 18 Second-order apparent rate coefficient of 3 OP pesticide oxons at pH 9. [OP]o = 0.5 

μM, [CO3]T = 10 mM, Temperature = 25±1°C, and [HOCl]T = 0-200 μM. Error bars 

represent 95% confidence intervals for both data points and regression lines. .......... 57 

Figure 19 Experimental and model results for DZ loss in the presence of chlorine at pH 7.0. 

[DZ]o = 0.6 μM, [HOCl]T = 20 μM, [PO4]T = 10 mM, and Temperature = 25±1°C. 

Error bars represent 95% confidence intervals. Lines represent model results. ........ 58 

Figure 20 Experimental and model results for PA loss in the presence of chlorine at pH 8.0. 

[PA]o = 0.57 μM, [HOCl]T = 25 μM, [PO4]T = 10 mM, and Temperature = 25±1°C. 

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

Figure 21 First-order observed loss of CP and DZ in the presence of monochloramine as a 

function of pH. [NH2Cl] = 50 μM, [OP]o = 0.5 μM, Cl/N = 0.7 mol/mol, [Buffer]T = 

10 mM, Temperature = 25±1°C, and pH 6.5-9.0. Error bars represent 95% 

confidence intervals. ................................................................................................... 60 

Figure 22 First-order observed loss of MA in the presence of monochloramine as a function of 

pH. [NH2Cl] = 50 μM, [MA]o = 0.5 μM, Cl/N = 0.7 mol/mol, [Buffer]T = 10 mM, 

Temperature = 25±1°C, and pH 6.5-9.0. Error bars represent 95% confidence 

intervals....................................................................................................................... 61 

vii
viii

Figure 23 Observed first-order observed loss of three OP pesticides in the presence of 

monochloramine at pH 8.5. [NH2Cl] = 50 μM, [OP]o = 0.5 μM, Cl/N = 0-0.7 

mol/mol, [H3BO3]T = 10 mM, and Temperature = 25±1°C. Error bars represent 95% 

confidence intervals. ................................................................................................... 62 

Figure 24 Experimental and model results for CP loss in the presence of monochloramine. 

[CP]o = 0.5 μM, [NH2Cl]o = 50 μM, Cl/N = 0.7 mol/mol, [Buffer]T = 10 mM, pH 6.5-

9, and Temperature = 25±1°C. Error bars represent 95% confidence intervals. Lines 

represent model results. .............................................................................................. 63 

Figure 25 Experimental and model results for DZ loss in the presence of monochloramine. 

[DZ]o = 0.5 μM, [NH2Cl]o = 50 μM, Cl/N = 0.7 mol/mol, [Buffer]T = 10 mM, pH 6.5-

9, and Temperature = 25±1°C. Error bars represent 95% confidence intervals. Lines 

represent model results. .............................................................................................. 64 

Figure 26 Experimental and model results for MA loss in the presence of monochloramine. 

[MA]o = 0.5 μM, [NH2Cl]o = 50 μM, Cl/N = 0.7 mol/mol, [Buffer]T = 10 mM, pH 

6.5-8.5, and Temperature = 25±1°C. Error bars represent 95% confidence intervals. 

Lines represent model results...................................................................................... 65 

Figure 27 Structure-activity relationship relating the reaction of monochloramine, dichloramine, 

and hypochlorous acid with CP, DZ, and MA. Piela and Wrona (49) found half-wave 

potentials for monochloramine (-0.2 V), dichloramine (0.4), and hypochlorous acid 

(0.8 V). These values were then converted using the Nernst equation to the intensive 

property of the log ratio of the activities of the reduced products to the oxidized forms 

of the oxidant. ............................................................................................................. 66 
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 high priority 

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 transformation products 

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 # 

ERD08103-1, “Fate of Pesticides and Toxic Chemicals During Drinking Water Treatment”, 

under Safe Pesticides/Safe Products (SP2) Long-Term Goal 1.4.1. The goals of this research 

task are to: 1) provide chemical-specific information on the effects of water treatment for highpriority pollutants, 2) provide physicochemical parameters for transformation products, and 3) 

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

treatment conditions. 

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

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

distribution systems. Our objective was to develop predictive models that associate intrinsic 

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. For this purpose, eight OP 

ix
pesticides and three oxon transformation products were investigated. The 8 OP pesticides 

chosen for this study contained the thiophosphate moiety (P=S) from the phosphorothioate 

(chlorethoxyfos (CE), chlorpyrifos (CP), diazinon (DZ), parathion (PA), and tebupirimfos (TE)) 

and the phosphorodithioate (malathion (MA), methidathion (ME), and phosmet (PM)) 

subgroups. The three oxon (i.e., P=O) products were from the phosphate subgroup (diazoxon 

(DZO), malaoxon (MAO), and paraoxon (PAO)). It was originally thought that by investigating 

a class of pesticides that structure-activity relationships could easily be developed. However, 

reactivity with chlorinated oxidants could only be correlated with subgroups and not across the 

class of OP pesticides. 

The phosphorodithioate subgroup was found to hydrolyze much more rapidly than the 

phosphate or the phosphorothioate subgroups, and oxons are more susceptible to alkaline 

hydrolysis when compared to the corresponding thionates. Also, ethyl esters appear to hinder 

nucleophilic attack by hydroxide ions (i.e., alkaline hydrolysis) compared to a structurally 

similar pesticide with methyl ester linkages at the phosphorus atom; however, there is a need for 

more hydrolysis studies to be conducted at environmentally relevant conditions to build 

structure-activity relationships in order to predict intrinsic hydrolysis rate coefficients. 

In the presence of chlorine, intrinsic rate coefficients for both hypochlorous acid (HOCl) 

and hypochlorite (OCl-

) were calculated, kHOCl,OP and kOCl,OP respectively. Since multiple 

reaction pathways occur simultaneously over the pH range of 6.5-9, a chlorine-OP pesticide 

reaction model previously developed for CP was used. The reaction with HOCl and the OP 

pesticides resulted in the rapid formation of oxons, which are more toxic than the parent OP 

pesticide. Hypochlorite accelerated the hydrolysis of each OP pesticide but did not result in 

oxon formation. The energy of the highest occupied molecular orbital (EHOMO) showed that 

x
xi

kHOCl,OP correlated within the two subgroups. The alkaline hydrolysis rate (kB,OP) was found to 

correlate with kOCl,OP within each OP subgroup, which also implies that the reaction mechanism 

of OCl- and OH- are similar. Therefore, the cross correlation demonstrates that OCl- does act as 

a nucleophile accelerating the hydrolysis of OP pesticides. Oxons did not undergo further 

oxidation by aqueous chlorine, but they were susceptible to chlorine assisted hydrolysis. 

Three of the most commonly detected OP pesticides (CP, DZ, and MA) in drinking water 

supplies were investigated in the presence of chloramines. The loss of all three pesticides under 

chloramination conditions was found to be highly pH dependent. At pH 6.5, the primary 

degradation pathway for OP pesticides was due to reaction with dichloramine and hypochlorous 

acid, which are formed during monochloramine autodecomposition. The direct reaction of 

monochloramine with each OP pesticide was found to be a minimal loss pathway. The order of 

reactivity for the three chlorinated oxidants was HOCl  NHCl2  NH2Cl. The reactivity of the 

three chlorinated oxidants was found to correlate with half-wave potentials (E1/2) for each OP 

pesticide. 

The stability of OP pesticides was investigated under conditions similar to drinking water 

treatment. The critical reaction parameters could not be correlated across the OP pesticide class, 

but correlations were made within subgroups indicating that minute structure differences 

significantly impact reactivity. The structure-activity relationships developed here could allow 

for the prediction of critical reaction parameters of OP pesticides in the presence of chlorinated 

oxidants. Along with the presented model, regulators could access the potential risk associated 

when consuming potable water contaminated with OP pesticides. 
1 INTRODUCTION 

The United States Geological Survey (USGS) conducts a national reconnaissance survey 

known as NAWQA (National Water-Quality Assessment) Program to help define the effect 

contaminants have on drinking water supplies and aquatic ecosystems (1). 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. 

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. When assessing potential pesticide exposure due to drinking potable water, all 

potential transformation pathways need to be addressed. Under drinking water treatment 

conditions, hydrolysis and chemical oxidation are the most relevant transformation pathways for 

the class of organophosphorus (OP) pesticides (4). The OP pesticides were chosen for their 

wide-spread use and measured concentrations in drinking water supplies (1, 3, 5). However, 

there is little quantitative rate coefficient information for either pathway. 

Only 6 OP pesticides have been studied in-depth to reveal their hydrolytic behavior at 

acidic, neutral, or alkaline conditions. Malathion has been extensively studied over the pH range 

1
of 1-10. In the pH range of 1-4, hydrolysis occurs at an ether linkage resulting in the formation 

of malathion monoacid (6). Above pH of 4, nucleophilic attack at the tetrahedral phosphorus 

atom by hydroxide ion was found to be twice as fast as carboxyl ester hydrolysis (7). Malathion 

appears to be the most stable at pH 4, with an estimated half-life greater than one year. 

Diazinon, diazoxon, parathion, and paraoxon have been studied over the pH range of 3.1-10.4 

(8). All four of these OP pesticides were found to have an acidic, neutral, and alkaline 

component to their hydrolytic behavior. Also, chlorpyrifos and its corresponding oxon were 

studied over the pH range of 3-11, which neutral and alkaline hydrolysis rate coefficients were 

determined for both pesticides (4, 9). However, OP pesticide hydrolysis is very dependent on 

structure allowing for rate coefficient comparisons only to be made within each subgroup (i.e., 

phosphorothioate, phosphorodithioate, phosphorothiolate, and phosphate) (8). 

Chlorination is the most commonly used chemical disinfection process for community 

water systems (10), and it is known to react with numerous pesticides. For example, four striazines were found to degrade in the presence of aqueous chlorine ([HOCl]T = HOCl + OCl-

) 

(11, 12). Atrazine was also found to be significantly degraded by ozone (13); however, 

subsequent chlorination of the ozonated effluent had very little effect on the concentration of 

residual atrazine or its ozone degradation products (2). Also, some carbamate pesticides have 

been shown to react with 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 chlorine; but 

aldicarb, methomyl, and thiobencarb do exhibit significant reactivity (14-16). These findings 

demonstrate that 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 

2
systematic structural variations and employing carefully selected experimentation and numerical 

modeling. 

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

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

(P=O) (17-19). The resulting oxons are typically more potent than the parent as an inhibitor of 

acetlycholinesterase, an enzyme necessary for regulating nerve impulse transmission between 

nerve fibers (19). Duirk and Collette, (4) 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, as well as in the presence of naturally occurring 

aqueous constituents such as bromide and natural organic matter (NOM) (20). 

Chloramines are a common secondary disinfectant alternative for many drinking water 

utilities. In some cases, monochloramine is used for both primary and secondary disinfection 

when excessively high background ammonia concentration in the source water do not allow for 

breakpoint chlorination (10). However, very little is known about the reactions of 

monochloramine with anthropogenic chemicals. In the presence of aromatic-ether and aminecontaining pharmaceuticals, aqueous chlorine was found to react more rapidly than 

monochloramine (21). Also, monochloramine was found to react with diuron, a phenyl urea 

herbicide, resulting in the formation of N-nitrosodimethylamine (NDMA) (22), which is a known 

disinfection byproduct of chloramination (22, 23). 

However, it is difficult to directly observe monochloramine reactions with anthropogenic 

compounds due to its own autodecomposition, i.e., a series of reactions ultimately resulting in 

monochloramine loss (Table 1) (24). This mechanism is highly pH dependent and allows for 

3
multiple oxidants to coexist (i.e., monochloramine (NH2Cl), dichloramine (NHCl2), and HOCl), 

which confounds the interpretation of observational results. Using this well understood 

autodecomposition reaction mechanism, the reactions of monochloramine with natural organic 

matter resulting in chlorinated and brominated disinfection byproduct formation (DBP) were 

resolved (25-27). This same technique has been used to investigate triclosan reactivity in 

chloraminated water (28). It was found that dichloramine was approximately three orders of 

magnitude more reactive with triclosan than monochloramine, but aqueous chlorine was found to 

be three orders of magnitude faster than dichloramine. Similar results have been see with sulfite 

and cyanide (29). Therefore, reaction pathway models can be used as a tool to propose and 

elucidate reaction pathways in complex systems with parallel reactions. 

The purpose of this study was to further elucidate the kinetics and transformation 

pathways of OP pesticides in the presence of chlorinated oxidants. The 8 OP pesticides chosen 

in this study contained the thiophosphate moiety from the phosphorothioate (chlorethoxyfos 

(CE), chlorpyrifos (CP), diazinon (DZ), parathion (PA), and tebupirimfos (TE)) and the 

phosphorothioate (malathion (MA), methidathion (ME), and phosmet (PM)) subgroups (Table 

2). Hydrolysis neutral and alkaline rate coefficients were determined for CE, ME, PM, and TE. 

Since all these OP pesticides can form oxons in the presence of chlorine, reaction rate 

coefficients HOCl and OCl-

 were correlated to molecular descriptors. For both correlations, 

subgroup differences were observed. CP, DZ, and MA were investigated in the presence of 

monochloramine over the pH range of 6.5-9. These three OP pesticides were chosen due to their 

frequency of detection in drinking water supplies (3). 

4
2 EXPERIMENTAL PROCEDURES 

2.1 Materials 

Chlorpyrifos (CP), chlorpyrifos oxon (CPO), 3,5,6-trichloro-2-pyridinol (TCP), parathion 

(PA), paraoxon (PAO), 4-nitrophenol (n-Ph), diazinon (DZ), diazoxon (DZO), 2-isopropyl-6-

methyl-4-pyrimidinol (IMP), malathion (MA), malaoxon (MAO), methidathion (ME), 

tebupirimfos (TE), phosmet (PM) and chlorethoxyfos (CE) 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). 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. 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 prior to 

use. All chlorination, hydrolysis, and chloramination experiments were conducted at constant 

temperature (25±1°C) and pH was maintained to ±0.1 of the initial pH for the duration of each 

experiment. 

2.2 Methods 

2.2.1 Chlorination and Hydrolysis of OP Pesticides 

For all OP oxidation and hydrolysis experiments, each OP was spiked by adding 1 mL of 

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

5
of nitrogen gas was used to evaporate the ethyl acetate and 4 L of deionized water was then 

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

in an aqueous OP concentration of 1 μM. 

OP chlorination kinetic experiments were conducted under pseudo first-order conditions: 

total chlorine, [HOCl]T, to OP molar ratios of 20:1, 50:1, 100:1, and 200:1. Chlorine was added 

to solutions under rapid mix conditions achieved with a magnetic stir plate and a PTFE coated 

stir bar. All chlorination experiments were conducted until at least 87% loss in the pesticide 

initial concentration was achieved. Above pH 8, 10 mM carbonate [CO3]T buffer was used to 

maintain pH. The commercial aqueous chlorine solution was first diluted to 250 mM and then 

added to the aqueous system containing 0.5 μM OP pesticide and carbonate buffer in a 2 L 

Erlenmeyer flask. Thirteen aliquots from the large 2 L reactor were then placed into 120 mL 

amber reaction vessels with PTFE septa and stored in the dark. In the pH range of 6.5-8, the rate 

of OP loss in the presence of chlorine was very fast. Therefore, thirteen 100 mL aliquots of the 2 

L aqueous system containing 10 mM phosphate buffer, [PO4]T, and 0.5 μM OP were placed in 

250 mL amber Erlenmeyer flasks. Each flask was individually dosed with chlorine. 

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

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

]) via 

Standard Method 4500-Cl F DPD-FAS titrimetric method (30). Free chlorine residuals were 

quenched in the second reaction vessel with sodium sulfite in 20% excess of the initial free 

chlorine concentration and the pH of the 100 mL sample was then adjusted to 7. Sulfite had no 

effect on the recovery of any of the analytes as previously shown (4). 

Hydrolysis experiments were conducted in a similar manner to the chlorination 

experiments. Phosphate buffer was used over the pH range of 2-8; while carbonate buffer was 

6
used in the experiments at alkaline pH. Since hydrolysis or most of the OP pesticides was 

relatively slow, these experiments were conducted in 2-L Erlenmeyer flaks with 10 mM of 

buffer. However, the hydrolysis rate of PM was extremely fast at pH 9-10. The experimental 

setup used to study the chlorination of OP pesticides at neutral pH was modified in order to study 

PM hydrolysis at alkaline pH. Instead of spiking a 100 mL solution under rapid mix conditions 

with a minuscule amount of chlorine, 98 mL of a [PM]o = 0.5 μM solution was spiked with 2 mL 

of concentrated carbonate buffer. The concentrated buffer solution was the proper proportion of 

bicarbonate/carbonate, which immediate adjusted the 100 mL sample to the target pH of 9, 9.5, 

or 10. The rapid rate of PM hydrolysis at alkaline pH was then stopped by immediately 

adjusting the pH to 7 with the appropriate amount of sulfuric acid. Hydrolysis experiments were 

conducted until at least 70% loss in the initial pesticide concentration was achieved. 

2.2.2 Chloramination of OP Pesticides 

Monochloramine kinetic experiments utilized additions of preformed monochloramine to 

avoid potential artifacts caused by reactions of excess free chlorine that may briefly exist if 

monochloramine was formed in-situ. The stock solution was prepared by mixing 5.64 mM 

ammonia with 3.7 mM hypochlorous acid to achieve the desired 0.7 Cl/N molar ratio. The 

solution was aged for 30 minutes in 10 mM bicarbonate buffer, pH 8.5, prior to use in any of the 

experiments. Kinetic experiments were performed in a 2 liter Erlenmeyer flask in which 

monochloramine ([NH2Cl]o = 0.05 mM) was added to a buffer ([Buffer]T = 10.0 mM) solution 

containing [OP]o = 0.5 μM under rapid mix conditions, which was then poured into sixteen 128 

mL batch reactors. At each sample interval, two reactors were used in their entirety and 

monochloramine, pH, as well as the pesticide and its known degradation products were measured 

for at least 140 hours. 

7
2.3 OP Pesticide Analysis 

Parent OP, oxon products, and other degradation products were extracted using C-18 

solid phase extraction cartridges purchased from Supleco (Bellefonte, PA). Prior to extraction, 

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 flow-rate 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 Hewlett-Packard 6890 GC 

equipped with a 5973 MSD was used to analyze the OPs and their degradation products. 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. 

2.4 Numerical Calculations and Parameter Estimation 

ScientistTM by Micromath (Salt Lake City, UT) was used to solve the systems of stiff 

ordinary differential equations. The model was used to estimate rate coefficients using nonlinear regression analysis techniques. Scientist 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. Final parameter estimates were 

based on global pooling of all experimental data for each OP pesticide. All molecular 

descriptors were calculated using GAMESSPLUS software. 

8
3 RESULTS AND DISCUSSION 

3.1 Hydrolysis of select OP pesticides 

Understanding the hydrolytic behavior of water soluble pesticides, such as the OP 

pesticides, is paramount in order to determine their fate in the environment or engineered 

systems. The phosphorothioate subgroup examined contained five pesticides. Of these, very 

little information was known about CE or TE. CE was first investigated at pH 3 and 7 in order to 

determine the acidic and neutral hydrolysis rate coefficients. The hydrolytic behavior was 

similar to CP over this pH range and CE exhibited no statistical difference in the first order 

observed rate coefficient at pH 3 or 7 (Figure 1). At alkaline pH, hydroxide ion appears to 

significantly accelerate CE hydrolysis (Figure 1, insert). TE was also investigated over the pH 

range of 3-10 with only slightly different results (Figure 2). At pH 3 and 7, the observed firstorder rate coefficients of TE hydrolysis were statistically similar. However, the observed rate of 

TE loss at pH 10 was found to be statistically larger than the observed rate at either pH 3 or 7. 

The presence of hydroxide ion at pH 10 appears to slightly increase the observed rate of TE loss. 

Since the observed hydrolysis rate of TE at pH 7 was only minimally different when compared to 

the observed hydrolysis rate at pH 10, the second order alkaline hydrolysis rate coefficient was 

estimated from these two pHs (Figure 2, insert). 

The neutral hydrolysis rates for both TE and CE were found to be very similar with CP, 

DZ, and PA. Neutral hydrolysis rates for the phosphorothioate subgroup ranged from 1.56 x 10-4

– 1.10 x 10-3 h-1. However, the range in second-order rate coefficients for alkaline hydrolysis 

was significantly greater but with most of the OP pesticides relatively close in magnitude (Table 

3). CP, DZ, PA, and TE second order hydrolysis rate coefficients ranged from 6.0-37.1 M-1 h-1. 

However, CE was found to be 2 orders of magnitude faster than CP. Hydroxide ion attacks the 

tetrahedral phosphorus atom via an SN2 nucleophilic substitution releasing the best leaving group 

9
and resulting in the formation of diethyl thiophosphate. CP, DZ, PA, and TE all have aryl 

functional groups as leaving groups. CE has an alkyl leaving group with four chlorine atoms. 

The highly electronegative chlorine atoms may draw the electrons associated with the phosphoether linkage creating a slightly more partial positive charge at the tetrahedral phosphorous atom, 

which would allow CE to be more susceptible to SN2 attack by hydroxide ion. CP also has 

chlorine atoms associated with the aryl structure and is more susceptible to nucleophilic attack 

than the other phosphorothioates; however, the aryl structure appears to significantly negate the 

electron withdrawing effects of the three chlorine atoms. The oxons of DZ, PA, and CP (i.e., 

phosphate subgroup) have all previously been examined (4, 8). The alkaline hydrolysis rate for 

the oxons in the phosphate subgroup is an order of magnitude faster than their corresponding 

phosphorothioates (Table 3), which has been previously reported (4, 8, 31). This demonstrates 

that the alkaline hydrolysis rate of OP pesticides can be significantly different with minor 

changes in molecular structure. 

In the phosphorodithioate subgroup, MA was the only pesticide in this subgroup in 

which its hydrolytic behavior has been well characterized (6, 7). The observed rate coefficients 

for ME at pH 3 and 7 were statistically indistinguishable (Figure 3). At pH 10 and above, ME 

alkaline hydrolysis was very rapid (Figure 3, insert). PM was found to be relatively stable at pH 

2; however, its hydrolysis increased rapidly with increasing hydroxide ion concentration (Figure 

4). Alkaline hydrolysis of PM was found to be very rapid above pH 8 (Figure 4, insert). For the 

phosphorodithioate subgroup, ME was the only OP pesticide in this subgroup found to be stable 

at neutral pH. Neutral hydrolysis for the phosphorodithioate subgroup was then operationally 

defined as the pH where each pesticide was the most stable. By this definition, the neutral 

hydrolysis rate coefficients are comparable to the phosphorothioate and phosphate subgroups. 

10
Alkaline hydrolysis was generally found to be more rapid than the other two subgroups. All the 

phosphorodithioate pesticides in Table 2 have methyl esters linkages at the tetrahedral 

phosphorus atom. Alkaline hydrolysis may be faster due to the fact that SN2 nucleophilic attack 

by the hydroxide ion is hindered more by the ethyl esters than the methyl esters. Chlorpyrifos 

and parathion alkaline hydrolysis rates have been found to be slower when compared to 

chlorpyrifos-methyl and parathion-methyl (8). Also, the thiol ester could be more susceptible to 

nucleophilic attack by hydroxide ion than phenyl esters. This makes it difficult to directly 

compare the phosphorodithioates to the phosphorothioates at environmentally relevant 

conditions. 

3.2 OP Pesticide Degradation in the Presence of Aqueous Chlorine 

In the presence of aqueous chlorine, the oxidation of CP was found to be first order 

resulting in an overall second order reaction (4). Therefore, observed loss of each OP pesticide 

in the presence of aqueous chlorine was assumed to be first order with respect to the OP 

pesticide. If ln([OP]/[OP]o) versus time (t) plots are linear, then this assumption would be valid 

when there is a molar excess of chlorine. The observed first-order rate coefficients (kobs) for all 

the OP pesticides were determined from these plots via the slope of the regression line as shown 

in the following expression. 

tk

[OP]

[OP] ln obs

o

−= (1) 

Under pseudo first-order chlorination conditions, all the OP pesticides from CE-TE exhibited a 

first order dependency with respect to the OP pesticide in the presence of excess aqueous 

chlorine at pH 6.5 (Figures 5-11). The reaction order of the aqueous chlorine reacting with each 

OP pesticide was determined by plotting the log of kobs versus the log of the initial chlorine 

11
concentration at pH 6.5 (Figure 12). Since the slope of the regression line is approximately 1 for 

all OPs, this indicates that the loss of OP pesticides can be described as a second-order reaction. 

The apparent loss of each OP pesticide in the presence of aqueous chlorine at a specific 

pH could then be described by the following rate expression where kapp is the apparent secondorder rate coefficient at a specific pH (equation 2). The observed first-order rate coefficients at 

[HOCl]k [OP] dt

d[OP] −= app T (2) 

each pH were assumed to linearly increase with increasing free chlorine concentration (equation 

3). The kapp for each OP pesticide was determined by plotting kobs versus the initial 

obs app T = [HOCl]kk (3) 

total aqueous chlorine concentration. Figure 13 shows that kobs increased linearly with 

increasing chlorine concentration at pH of 6.5. Diazinon and phosmet degraded the fastest at this 

pH, while the rest of the OP pesticides exhibited similar reactivity with aqueous chlorine. Since 

approximately 91% of the active chlorine is in the HOCl form, the apparent rate coefficient will 

be very close to the intrinsic rate coefficient of hypochlorous acid reacting with each OP 

pesticide. As pH increased from 6.5 to 9, kobs decreased for all the OP pesticides (Figure 14). 

When the pH of the aqueous system shifts from neutral to alkaline pH, the chlorine species also 

shifts from HOCl, which has a pKa of 7.5 (32), to OCl-

. The decrease in the kobs would be 

expected if hypochlorous acid is the dominant reacting species resulting in oxon formation. 

Oxons for all the parent OP pesticides are not commercially available; therefore, transformation 

products were identified by their mass spectra. For all the OP pesticides, their corresponding 

oxons were the only transformation products identified by NIST mass spectra database at pH 6.5. 

HOCl appears to be the active chlorine species over the pH range investigated responsible for OP 

oxidation and oxon formation. 

12
The transformation of OP pesticides can become relatively complicated at alkaline pH. 

Several transformation pathways occur simultaneously resulting in unique transformation 

products for each pathway. At pH 9, approximately 97% of the active chlorine will be in the 

hypochlorite form. However, all the OP pesticides degraded rapidly in the presence of aqueous 

chlorine (Figure 15). Phosmet degraded faster than diazinon due to PM being more susceptible 

to alkaline hydrolysis. Just like at pH 6.5, the observed rate coefficients increased linearly with 

increasing chlorine concentration. However, oxon products were not necessarily the major 

transformation product. CP and its corresponding oxon were both found to be susceptible to 

chlorine-assisted hydrolysis (4). Since tetrahedral phosphorus atoms are known to be more 

susceptible to hydrolysis by supernucleophiles (33), and hypochlorite is a supernucleophile, 

intrinsic rate coefficients for both the chlorine-assisted hydrolysis and the HOCl-oxidation 

pathways need to be determined for each OP pesticide. 

Degradation pathway models have been shown to be an effective tool determining rate 

coefficients in complex systems. Previously, researchers have used the monochloramine 

autodecomposition model to determine rate coefficients for the reactions of monochloramine and 

dichloramine with triclosan and NOM (27, 28). More applicable to this reaction system, Duirk 

and Collette (4) developed a degradation pathway model for CP in order to determine the 

intrinsic rate coefficients for both HOCl and OCl-

. Knowing that the pH dependency on the rate 

of OP loss is due to chlorine speciation, pooled data at pH 6.5 and 9 will be used to parameterize 

the intrinsic rate coefficients for HOCl (kHOCl,OP) and OCl- (kOCl,OP) reacting with each OP 

pesticide using the following system of ordinary differential equations. In the following 

equations, kOCl,OPO is the intrinsic rate coefficient for hypochlorite assisting in the hydrolysis of 

13
 k [OCl ][OPO]

k5 [HOCl][OP] ][OP][OClk dt

d[HOCl]

OPOOCl,

OPHOCl, OPOCl,

T

−

− −= − − (4) 

k [HOCl][OP] ][OP][OClk[OP]k dt

d[OP]

OPHOCl, CPh, OPOCl,

− −= − − (5) 

k [HOCl][OP] k k[OPO] [OCl ][OPO] dt

d[OPO]

OPHOCl, OPOh, OPOOCl,

− = − − (6) 

 k [OCl ][OPO]

k][OP][OClk[OP]k [OPO] dt

d[OPH]

OPOOCl,

OPh, OPOCl, OPOh,

−

− = + + +

 (7) 

the OPO (oxon transformation product), and kh,OP and kh,OPO are the rate coefficients for both OP 

and OPO hydrolysis, and OPH represents the hydrolysis product. Rate equations were written 

based on the stoichiometric equations in Table 4. For most of the OP pesticides, only equations 

4 and 5 are applicable due to the inability to accurately quantify the oxon product. 

The kHOCl,OP were found to cluster in the range of 1.7-2.2 X 106 M-1h-1, with CE and DZ 

being notable exceptions (Table 5). The kOCl,OP appear to show differences between the 

subgroups. For the phosphorothioate subgroup, kOCl,OP appears to be an order of magnitude 

faster than kB,OP (Tables 3 and 5). However, the OCl-

 does not appear to increase the rate of 

hydrolysis as significantly for the phosphorodithioate subgroup. This could be due to the fact 

that they are generally found to be extremely susceptible to nucleophilic attack from hydroxide 

ion. DZ is the only OP pesticide in which an intrinsic rate coefficient has been previously 

reported (17), kHOCl,DZ = 4.68 x 105

 M-1h-1 and kOCl,DZ =972 M-1h-1. Although the hypochlorite 

intrinsic rate coefficients were similar, HOCl rate coefficients were found to be an order of 

magnitude different. The difference can be explain because kHOCl,DZ was approximated from 

data gathered only over the pH range of 9.5-11, and it was thought that both HOCl and OCl14
resulted in oxon formation. Since parallel reaction pathways can occur at any pH, the most 

accurate method to parameterize reaction rate coefficients of OP pesticides in the presence of 

chlorine would be to use a comprehensive transformation pathway model and data sets relevant 

to drinking water treatment conditions. 

Frontier molecular orbital theory has been used to correlate oxidation rate coefficients 

with the easily calculated energy of highest occupied molecular orbital (EHOMO) (34). EHOMO was 

initially thought to be able to correlate all 8 OP pesticides; however, subgroup differences were 

quickly unveiled (Figure 16). The phosphorodithioate subgroup was found to have a slightly 

higher potential than the phosphorothioate subgroup. This is most likely due to the sulfur 

linkage and the methyl esters at the tetrahedral phosphorus atom allowing them to be more easily 

oxidized by chlorine than the phosphorothioate subgroup, which primarily had ethyl and phenyl 

esters. EHOMO was found to be a good molecular descriptor describing the oxidation of OP 

pesticides within each subgroup. 

Other molecular descriptors were needed in order to correlate the reactivity of 

hypochlorite ion with each OP pesticide. Cross correlations use rate coefficients as molecular 

descriptors with a well understood reaction mechanism (i.e., SN2 alkaline hydrolysis) and infer a 

mechanistically similar reaction for a different reactant (i.e., chlorine-assisted hydrolysis) (35). 

Therefore, correlating kOCl,OP with alkaline hydrolysis rate coefficients for all 8 OP pesticides 

would then confirm that OCl-

 acts as a nucleophile accelerating OP pesticide hydrolysis. 

Subgroup differences were evident as the phosphorothioates generally hydrolyze slower at 

alkaline pH than the phosphorodithioates (Figure 17); however, chlorine appeared to have a 

greater effect accelerating OP pesticide hydrolysis as indicated by the slope of the regression line 

being an order of magnitude greater then the slope for the phosphorodithioates. Even with the 

15
differences between the two OP pesticide subgroups, this cross correlation validates that OClacts like a nucleophile attacking the tetrahedral phosphorus atom accelerating the hydrolysis of 

OP pesticides. However, it is important to note that HOCl still rapidly transforms OP pesticides 

to their corresponding oxon at pH 9. Oxon stability in the presence of aqueous chlorine still 

needs to be investigated. 

3.3 OP Oxon Chlorine-Assisted Hydrolysis 

Chlorine assisted hydrolysis was first observed by Edwards et al., (33) investigating the 

factors that determine the reactivity of nucleophiles, which are basicity, polarizability, and the 

presence of unshared pairs of electrons on the adjacent atom to the nucleophilic atom (i.e., the 

alpha effect). The three lone-pairs of electrons on both the chlorine and oxygen atom enhance 

hypochlorite’s nucleophilicity towards specific moieties such as a tetrahedral phosphorus (i.e., 

phosphoesters) and carbonyls. Hypochlorite is considered to be a supernucleophile towards 

phosphoesters because of these repulsions (36). At pH 9, DZO, MAO, and PAO loss was 

observed in the presence of aqueous chlorine (Figure 18). With subgroup differences aside, the 

loss of each oxon increased linearly proportional with increasing chlorine concentrations. DZO 

and PAO were significantly more stable than MAO at pH 9 because the phosphorothiolates are 

more susceptible to alkaline hydrolysis (8). 

The rate coefficients for chlorine-assisted hydrolysis for all four oxons were found only 

to be slightly larger than their corresponding parents (Table 5). This was first reported with CP 

and CPO in the presence of chlorine and bromine (4, 20). This is due to the structure of the 

nucleophile, OCl-

, and not minor structural difference between a phosphorothioate and phosphate 

analogs (36). Since tetrahedral phosphate moieties are susceptible to SN2 attack by OCl-

 (33), 

the difference in the partial positive charge at the tetrahedral phosphorus atom between (P=S) 

and (P=O) does not significantly influence the rate of nucleophilic attack. Therefore, chlorine16
assisted hydrolysis rates for oxons can then be assumed to be approximately the same as their 

corresponding parents. This could aid risk assessors estimate exposure to the more toxic oxon 

products as well as direct future exposure research activities related to oxon stability in drinking 

water distribution systems. 

3.4 OP Pesticide-Chlorine Model Validation 

With the available rate coefficients determined here and in cited literature, OP pesticide 

transformation pathways in the presence of aqueous chlorine can be predicted under drinking 

water conditions. Diazinon (DZ) and parathion (PA) were chosen because their corresponding 

oxons and hydrolysis products are commercially available. Diazinon was found to react the 

fastest among the OP pesticides selected (Table 5). At a pH and chlorine concentration typical 

of drinking water treatment, DZ was rapidly transformed to diazoxon (DZO) (Figure 19). Also, 

DZO is stable in the presence of chlorine for over 48 hours at neutral pH (18). At pH 8, PA was 

also rapidly transformed to paraoxon (PAO) (Figure 20). PAO was also found to be relatively 

stable at alkaline pH, as shown previously (Figure 18). This model has already been shown to 

adequately predict CP transformation and degradation of CPO in the presence of natural 

occurring aqueous constituents (i.e., bromide and natural organic matter) (20). Regulators could 

potentially use this model to access potential exposure to OP pesticides and their more toxic 

oxon products using the structure-activity relationships in Figures 16 and 17 to estimate reaction 

rate coefficients across the phosphorothioate and phosphorodithioate subgroups. 

3.5 OP Pesticide Degradation in the Presence of Monochloramine 

Observing the direct reaction of monochloramine with analytes in solution can be very 

difficult at environmentally relevant concentrations. Previous work elucidating the reaction 

mechanism of monochloramine with bromide, cyanide, or nitrite was performed at high 

millimolar concentrations of the reactants (29, 37, 38). Due to the autodecomposition 

17
mechanism of monochloramine (24), it is nearly impossible to observe the direction reaction of 

monochloramine with specific analytes when at environmentally relevant concentrations. 

Therefore, the monochloramine autodecomposition model can be used as a tool to validate 

proposed transformation pathways with comprehensive data sets over a range of relevant 

conditions. 

The observed loss of each OP pesticide was assumed to be first-order with respect to each 

OP. Monochloramine, dichloramine, and hypochlorous acid will exist simultaneously and vary 

in concentration as a function of pH. Therefore, the kobs presented in Figures 21 and 22 should 

be viewed as a function of total oxidant in the aqueous system. The observed first-order rate 

coefficients for CP and DZ both decreased as pH increased from 6.5 to 9 (Figure 21). This was 

found to be very similar to the reaction of chlorine with both CP and DZ (Figure 14). MA 

followed the same trend initially until approximately pH 7.5 when alkaline hydrolysis became 

the dominant degradation pathway (Figure 22). However, it appears that dichloramine and 

hypochlorous acid are primarily responsible for the loss of the OP pesticides at pH 6.5 and 7. As 

pH grows more alkaline, the direct reaction between monochloramine and the OP may be 

inconsequential compared to alkaline hydrolysis. 

OP pesticide degradation pathways in the presence of monochloramine will be intimately 

intertwined with monochloramine autodecomposition. Since alkaline hydrolysis of MA was a 

significant loss pathway, further discussion on OP pesticide degradation pathways will primarily 

refer to CP and DZ. As pH decreased so did the observed rate of OP pesticide loss. The initial 

monochloramine concentration was constant for all the experiments; therefore, the direct reaction 

of monochloramine with OP pesticides appears to be minimal. However, near neutral pH three 

possible oxidants are formed due to monochloramine autodecomposition that could be 

18
responsible for the loss of OP pesticides: hypochlorous acid (HOCl), dichloramine (NHCl2), and 

monochlorammonium ion (NH3Cl+

). Hypochlorous acid forms primarily due to the hydrolysis 

of monochloramine (Table 1 reaction 2), and it does react with all three OP pesticides. However, 

HOCl can either react with ammonia to reform monochloramine or with monochloramine to 

form dichloramine (Table 1 reaction 3). Dichloramine formation is also due to monochloramine 

disproportionation, which is acid catalyzed Table 1 reaction 5. This acid-catalyzed reaction 

involves the formation of monochlorammonium ion (39). 

NH2Cl +H3O+ NH3Cl+ H2O K = 28 M-1 (8) 

Even though phosphate and carbonate salts participate as proton donors at high concentrations, 

they are several orders of magnitude less effective in the formation of monochlorammonium ion 

at environmentally relevant concentrations. Therefore, the primary dichloramine formation 

pathway will be due to HOCl reacting with NH2Cl. This can be significant because dichloramine 

has been found to be more reactive with cyanide, sulfite, and triclosan than monochloramine (28, 

29, 38). The following are proposed additional reactions of chlorinated oxidants with OP 

pesticides in the presence of monochloramine. 

NH2Cl + OP products kNH2Cl,OP (9) 

NHCl2 + OP products kNHCL2,OP (10) 

Since hypochlorous acid is a product of monochloramine autodecomposition, all known OP 

degradation pathways will also be included in the partial rate expressions, equations 11-17, 

which are added to the full monochloramine autodecomposition model (Table 6). 

k [NH kCl][OP] [NHCl ][OP] dt

Cl]d[NH

2OPNH2Cl, OPNHCL2, 2

2 −= + (11) 

k [NHCl ][OP] dt

]d[NHCl

OPNHCl2, 2

2 −= (12) 

19
 k [OCl ][OPO]

k [HOCl][OP] ][OP][OClk dt

d[HOCl]

OPOOCl,

OPHOCl, OPOCl,

T

−

− −= − − (13) 

k [NH Cl][OP] dt

]d[NH

NHC2Cl, 2OP

T3 = (14) 

k [HOCl][OP] ][OP][OClk[OP]k dt

d[OP]

OPHOCl, CPh, OPOCl,

− −= − − (15) 

k [HOCl][OP] k k[OPO] [OCl ][OPO] dt

d[OPO]

OPHOCl, OPOh, OPOOCl,

− = − − (16) 

 k [OCl ][OPO]

k][OP][OClk[OP]k [OPO] dt

d[OPH]

OPOOCl,

OPh, OPOCl, OPOh,

−

− = + + +

 (17) 

Additional assumptions are that in the monochloramine-OP pesticide reaction model the 

reaction of monochloramine and dichloramine with each OP pesticide do not lead to oxon 

formation. Also, that the reactions of monochloramine and dichloramine are 1-to-1 elemental 

stoichiometric reactions. The reaction of dichloramine and OP pesticides results in the 

regeneration of monochloramine, which is similar to dichloramine reacting with inorganics (29, 

38). The direct reaction of monochloramine with OP pesticides results in ammonia formation. 

Finally, rate coefficients determined using the monochloramine-OP pesticide reaction model 

exhibit no pH dependency. 

Two rate coefficients need to be determined for both monochloramine and dichloramine 

with these OP pesticides. Since dichloramine rapidly decays above neutral pH (40), experiments 

to determine kNH2Cl,OP were conducted in the presence of excess ammonia at pH 8.5 to promote 

monochloramine stability and reformation (41). In the presence of excess ammonia ([NH3]T = 

0.02-0.45 or Cl/N = 0.7-0.1 mol/mol), the rate of OP pesticide degradation decreased as 

ammonia concentrations increases (Figure 23). Due to MA being susceptible to alkaline 

20
hydrolysis, MA loss at pH 8.5 was significantly faster than CP and DZ. For each pesticide, these 

data sets were pooled for each OP pesticide and an intrinsic rate coefficient for NH2Cl with each 

OP pesticide was determined using nonlinear regression analysis. The rate coefficients found for 

the direct reaction of monochloramine with each OP pesticide were relatively small in magnitude 

(Table 7) compared to kHOCl,OP (Table 5). This is not unusual because another study has reported 

similar results with triclosan (28). Knowing the kHOCl,OP of each OP pesticide, the remaining 

data for each OP pesticide was pooled in order to determine kNHCl2,OP. Dichloramine was found 

to be two orders of magnitude great than monochloramine with each OP pesticide (Table 7), 

which is consistent with previously reported results (29, 38). Dichloramine appears to play a 

significant role in the loss of OP pesticides near neutral pH. 

Using the final parameter estimates in Table 7, the monochloramine-OP pesticide 

reaction model was used to predict the loss of each OP pesticide in the presence of 

monochloramine over the pH range of 6.5-9. Hydrolysis was not a major loss pathway for either 

CP or DZ over this range; therefore, oxidation was the primary degradation pathway for both 

pesticides in the presence of monochloramine (Figures 24 and 25). The loss of CP and DZ in the 

presence of monochloramine was found to be relatively fast at pH 6.5 due to the formation of 

dichloramine and hypochlorous, which then react with the OP pesticides. As pH increased to 9, 

the rate of CP and DZ loss in the presence of monochloramine decreased due to: 1) the rapid 

disproportionation of dichloramine above neutral pH (42), 2) speciation of ammonium to 

ammonia which rapidly reacts with hypochlorous acid reforming monochloramine, 3) 

hypochlorous acid speciation to hypochlorite which does not participate in monochloramine 

reformation, and 4) monochloramine not being as reactive with either pesticide compared to 

other chlorinated oxidants. Since hydrolysis of MA is a significant degradation pathway at pH 

21
7.5 and above, modeling MA hydrolysis was very important in order to correctly parameterize 

kNH2Cl,OP and kNHCl2,OP. As shown in Figure 26, the model was found very capable of predicting 

both monochloramine and MA concentrations. The excellent correlation between the 

experimental results and the model predictions indicate that the fundamental reactions 

responsible for the loss of both monochloramine and the OP pesticides were incorporated into 

the monochloramine-OP pesticide reaction model. 

Unlike the chlorination experiments, only a select few OP pesticides were used in the 

chloramination experiments. Since CP and DZ are from the phosphorothioate subgroup and MA 

is from the phosphorodithioate subgroup, it is not possible to correlated the reactivity of 

dichloramine and monochloramine to EHOMO as with kHOCl,OP. However, the reactivity of the 

three chlorinated oxidants can be correlated to each OP pesticide using half-wave potentials 

(E1/2). Half-wave potentials have been shown to be a good descriptor for the electrophilicties of 

HOCl, NH2Cl, and NHCl2 with organic and inorganic compounds (28); however, using halfwave potentials to correlate structure to reactivity is not as robust as thermodynamic parameters 

(43). Figure 27 shows a good correlation for the three oxidants with the three pesticides. It is 

interesting to note that the slope for each pesticide was found to be the same (slope = 0.30) with 

just slight differences in the intercept. This appears to indicate that degree of electrophilicity of 

each oxidant with each OP pesticide may be proportional across the OP pesticide class and not 

dependent on subgroup. Therefore, estimates for the reactivity of monochloramine and 

dichloramine with other OP pesticides could be made from knowing only kHOCl,OP (Table 5), or 

the kNH2Cl,OP and kNHCL2,OP for CP or MA could be used to quickly estimate the fate of other OP 

pesticides in the presence on monochloramine.. 

22
Although the reaction rate coefficients of the chlorinated oxidants were found to correlate 

with each OP, they appear not to react via the same mechanism (i.e., reaction center). Product 

studies have shown that only the HOCl reaction pathway resulted in oxon formation, and that the 

resulting oxons also appear not to be stable in the presence of monochloramine. The reaction of 

monochloramine and dichloramine do not appear to result in oxon formation; however, they may 

react with each OP pesticide resulting in a difference transformation product while not mitigating 

toxicity. Additional work is needed in order to understand the fate of OP pesticides and their 

oxon transformation products in the presence of monochloramine. 

23
4 CONCLUSION 

Eight OP pesticides and three of the oxon transformation products were investigated in 

the presence of chlorinated oxidants. It was originally thought that by investigating a class of 

pesticides that structure-activity relationships could easily be developed. However, it quickly 

became apparent that subgroups could be correlated within each OP pesticide class; however, 

correlations could not be made across the chemical class. This became evident looking at 

hydrolysis rate coefficients for the phosphorothioate, phosphate, and the phosphorodithioate 

subgroups. Chlorination of all eight pesticides led to a similar conclusion. Chloramination of 

three of the most commonly detected OP pesticides provided some insight into how chlorinated 

oxidants react with OP pesticides. Overall, OP pesticide structure was found to dictate reactivity 

in all three systems. 

Hydrolysis rate coefficients were determined for 4 pesticides due to lack of reported 

intrinsic rate coefficients. Pesticides from the phosphorodithioate group hydrolyzed rapidly 

under alkaline conditions. Also, oxons were found to be more susceptible to alkaline hydrolysis 

compared to the corresponding thionates due to the oxygen being more electronegative than 

sulfur, which facilitates the nucleophilic attack by hydroxide ion at the slightly more positively 

charged phosphorus atom. Also, ethyl esters appear to hinder nucleophilic attack by hydroxide 

ion compared to a structurally similar pesticide with methyl esters at the phosphorus atom; 

however, no direct comparison can be inferred for the difference in hydrolysis rates between the 

phosphorothioate and phosphorodithioate subgroups investigated here. Therefore, more 

hydrolysis studies need to be conducted at environmentally relevant conditions to determine the 

hydrolytic behavior of OP pesticides, which intrinsic hydrolysis rate coefficients could then be 

predicted via structure-activity relationships. 

24
In the presence of aqueous chlorine, eight OP pesticides were investigated and intrinsic 

rate coefficients for both HOCl and OCl-

 were calculated. Since multiple reaction pathways 

occur simultaneously over the pH range of 6.5-9, a chlorine-OP pesticide reaction model 

previously developed for CP was used. Intrinsic rate coefficients were first found for HOCl at 

pH 6.5 then OCl-

 at pH 9. The reaction of HOCl with OP pesticides results in the rapid 

formation of the corresponding oxon. Hypochlorite acts as a nucleophile accelerating the 

hydrolysis of the OP pesticide and not resulting in oxon formation. When the intrinsic rate 

coefficients were correlated to molecular descriptors, subgroup differences were readily 

apparent. The energy of the highest occupied molecular orbital (EHOMO) showed that kHOCl,OP

correlated within the two subgroups. The phosphorothioates were found to be lower in energy, 

and the phosphorodithioates were more reactive as a group as indicated by the greater slope of 

the regression line. The alkaline hydrolysis rate (kB,OP) was found to correlate to kOCl,OP within 

each OP subgroup, which also implies that the mechanism of OCl-

 and OH- are similar. 

Therefore, OCl- does attach the tetrahedral phosphorus as a nucleophile accelerating the 

hydrolysis of all eight OP pesticides and three oxons. 

Three of the most commonly detected OP pesticides (CP, DZ, and MA) in drinking water 

supplies were examined in presence of chloramines. The loss of all three pesticides under 

chloramination conditions was found to be highly pH dependent. At pH 6.5, the primary 

degradation pathway for OP pesticides was due to reaction with dichloramine and hypochlorous 

acid, which are formed during monochloramine autodecomposition. The direct reaction of 

monochloramine with each OP pesticide was found to be a minimal degradation pathway, which 

was confirmed when parameterizing kNH2Cl,OP for each OP pesticide in the presence of excess 

ammonia. The order of reactivity for the three chlorinated oxidants with each OP pesticide was 

25
HOCl  NHCl2  NH2Cl. The reactivity of the three chlorinated oxidants correlated with halfwave potentials (E1/2) for each OP pesticide. 

The stability of OP pesticides was investigated under conditions similar to drinking water 

treatment. The eight pesticides in this study showed that reaction parameters could not be 

correlated across the OP pesticide class, but correlations could be made within subgroups 

indicating that minute structure differences can significantly impact reactivity. Structure-activity 

relationships developed here can allow for the prediction of critical reaction parameters of OP 

pesticides in the presence of chlorinated oxidants. Along with the presented model, regulators 

could access the potential risk associated when consuming potable water contaminated with OP 

pesticides. 

26
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30

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E255-E265. 
TABLES 

31
Table 1 Stoichiometric equations and coefficients in the monochloramine 

autodecomposition model. 

Reaction Stoichiometry Rate/Equilibrium 

Coefficients (25 oC) 

Reference 

1 HOCl + NH3 NH2Cl + H2O k1= 1.5 x 1010M-1h-1 (44) 

2 NH2Cl + H2O HOCl + NH3 k2=7.6 x 10-2 h-1 (44) 

3 HOCl + NH2Cl NHCl2 + H2O k3= 1.0 x 106M-1h-1 (45) 

4 NHCl2 + H2O HOCl + NH2Cl k4= 2.3 x 10-3 h-1 (45) 

5 NH2Cl + NH2Cl NHCl2 NH3 ak5= pH dependent (40) 

6 NHCl2 + NH3 NH2Cl + NH2Cl k6= 2.16 x 108M-2h-1 (46) 

7 NH2Cl + NHCl2 N2+ 3H++3Cl- k7= 55.0 M-1h-1 (47) 

8 NHCl2 + H2O bI + 2HCl k8= 4.0 x 105M-1h-1 (42) 

9 b

I + NHCl2 HOCl+N2+HCl k9= 1.0 x 108M-1h-1 (47) 

10 bI + NH2Cl N2+H2O+ HCl k10= 3.0 x 107M-1h-1 (47) 

11 HOCl H+

 + OCl- pKa = 7.5 (32) 

12 NH4

+ H+

 + NH3 pKa = 9.3 (32) 

13 H2CO3 H+

 + HCO3

- pKa = 6.3 (32) 

14 HCO3

- H+

 + CO3

2- pKa = 10.3 (32) 

Note: ak5 = [H+] + [H2CO3] + [HCO3

-

]: = 2.5 x 107 M-2h-1, = 4.0 x

103

 M-2h-1, and = 800 M-2h-1 at 25 oC. bI: reactive intermediate. 

+ H k

k

COH 32 k -

3 HCO k + H k COH 32 k

-

3 HCO

32
Table 2 Structures and some chemical properties of select OP Pesticides (48). 

Structure and Name Log Kow Vapor Pressure 

(mm Hg) 

Water 

Solubility 

At 20°C 

(mg/L) 

Chlorethoxyfos (CE) 4.59 8.25 x 10-4 2.0 x 100

Chlorpyrifos (CP) 5.11 1.85 x 10-5 2.0 x 100

Diazinon (DZ) 3.81 1.2 x 10-2 4.0 x 101

Malathion (MA) 2.36 7.9 x 10-6 1.4 x 102

Methidathion (ME) 2.5 1.88 x 10-6 2.0 x 102

Parathion (PA) 3.83 9.7 x 10-6 6.5 x 100

Phosmet (PM) 2.95 6.4 x 10-1 2.5 x 101

Tebupirimfos (TE) a

NA 3.75 x 10-5 5.5 x 100

a

NA means not available. 

33
Table 3 Neutral and alkaline hydrolysis rate coefficients for 8 OP pesticides and 3 oxon 

transformation products. 95% confidence intervals shown in parentheses. 

OP Pesticide kN,OP at 25 °C 

(h-1) 

kB,OP at 25 °C 

(M-1h-1) 

References 

Chlorethoxyfos (CE) 4.68(±0.69) x 10-4 1.25(±0.30) x 103 this work 

Chlorpyrifos (CP) 

Chlorpyrifos oxon (CPO) 

3.72 x 10-4

2.13 x 10-3

37.1 

230.2 

(9) 

(4) 

Diazinon (DZ) 

Diazoxon (DZO) 

1.56 x 10-4 

9.99 x 10-4

18.9 

165.6 

(8) 

Malathion (MA) a

7.92 x 10-5 1.98 x 103 (6) 

Methidathion (ME) a

1.42(±1.01) x 10-3 2.22(±0.04) x 102 this work 

Parathion (PA) 

Paraoxon (PAO) 

2.66 x 10-4

2.00 x 10-4

4.3 

46.1 

(8) 

Phosmet (PM) b

5.55(±0.11) x 10-4 2.73(±0.08) x 105 this work 

Tebupirimfos (TE) 1.10(±0.0.01) x 10-3 6.0(±0.1) this work 

a

MA hydrolysis rate coefficient at pH 4 

b

PM hydrolysis rate coefficient at pH 2 

34
Table 4 Stoichiometric equations used in the chlorine-OP pesticide transformation pathway 

model. 

Reaction Stoichiometry Rate/Equilibrium 

Coefficient (25 °C) 

Reference 

1 -2 SO5Cl5HOPO 4

k 5HOCl OP OPHOCl, ⎯+ ⎯ → + − +++⎯⎯

kHOCl,OP = Table 5 this work 

2 OPH k OP ⎯⎯ →⎯OPh, kh,OP = kN,OP 

 + kB,OP[OH-

] 

kN,OP = Table 3 

kB,OP = Table 3 

3 OPH k OPO ⎯⎯ → OPOh,

⎯⎯ kh,OPO = kN,OPO 

 + kB,OPO[OH-

] 

kN,OPO = Table 3 

kB,OPO = Table 3 

4 OPH k OCl OP - ⎯+ ⎯ →⎯⎯OPOCl, kOCl,OP = Table 5 this work 

5 OPH k OCl OPO - ⎯+ ⎯ → OPO OCl,

⎯⎯ kOCl,OPO = Table 5 this work 

6 HOCl H+

 + OCl- pKa = 7.5 (HOCl/OCl-) (32) 

35
Table 5 Oxidation and chlorine-assisted hydrolysis rate coefficients for 8 OP pesticides and 

3 oxon transformation products. 95% confidence intervals shown in parentheses. 

OP pesticide kHOCl,OP at 25 °C 

(M-1h-1) 

kOCl,OP or kOCl,OPO 

at 25 °C (M-1h-1) 

References 

Chlorethoxyfos (CE) 0.86(±0.18) x 106 15,900±2100 this work 

Chlorpyrifos (CP) 

Chlorpyrifos oxon (CPO) 

1.72 x 106 990 

1340 

(4) 

Diazinon (DZ) 

Diazoxon (DZO) 

3.56(±0.65) x 106 627±30 

914.1±54.2 

this work 

Malathion (MA) 

Malaoxon (MAO) 

1.72(±0.36) x 106 382±26 

565±99 

this work 

Methidathion (ME) 1.89(±0.12) x 106 252±47 this work 

Parathion (PA) 

Paraoxon (PAO) 

2.20(±0.53) x 106 37±10 

48±10 

this work 

Phosmet (PM) 2.84(±0.80) x 106 1000±100 this work 

Tebupirimfos (TE) 1.76(±0.43) x 106 71±13 this work 

36
Table 6 Differential equations for the monochloramine autodecomposition model. 

1)

d[NH Cl]

dt k [NH ][HOCl]- k [NH Cl]- k [HOCl][NH Cl] + k [NHCl ]-

 k [NH Cl] k [NHCl ][NH ]- k [NHCl ][NH Cl]- k [I][NH Cl]

2

13 22 3 2 4 2

5 2

2

6 2 3 7 2 2 10 2

=

2 +

2)

d[NHCl ]

dt k [HOCl][NH Cl] k [NHCl ] + k [NH Cl] k [H ][NH ][NHCl ]

 k [NH Cl][NHCl ] k [NHCl ][OH ] k [I][NHCl ]

2

3 242 5 2

2

6 3

72 28 2 9 2

=− −

− −

+

−

2

3)

d[TOTOCl]

dt -k [HOCl][NH ] k [NH Cl]- k [HOCl][NH Cl] k [NHCl ]

 k [I][NHCl ] 

1 3 2 23 2 4

9 2

= + + 2 +

4) d[TOTNH ]

dt -k [HOCl][NH ] k [NH Cl] k [NH Cl] - k [NHCl ][NH ] 3

1 3 2 2 5 2

2 = ++ 6 2 3

5) d[Cl ]

dt k [NH Cl] k [NHCl ] k [I][NHCl ] k [I][NH Cl]

-

7 2

2 = ++ + 3 2 8 2 9 2 10 2

6) d[I]

dt

k [NHCl ][OH ] k [I][NHCl ] k [I][NH Cl] 8 2

- = −+ 9 210 2

7) d[N ]

dt k [NH Cl][NHCl ] k [I][NHCl ] k [I][NH Cl] 2 = ++ 7 2 2 9 2 10 2

8)

d[pH]

dt k [NH Cl][NHCl ] k [I][NHCl ] + 3k [I][NH Cl]

 k [NH ][HOCl]( NH OCl ) k OCl NH )[NH Cl]

 k [NH Cl][HOCl]( OCl k [NHCl ]( OCl

 k [NH Cl] NH k [H ][NH ][NHCl ]( NH /

7 2 2 9 2 10 2

1 3 o 4 1 21 o 4 2

3 2 1 4 21

5 2

2

o4 6

+

3 2o4

= − + +

−+ −

−+ +

− +

−

− −

(

(

) )

( ) ))

3 3

α α αα

α α

α α

+

β

37
38

Table 7 Intrinsic rate coefficients for monochloramine and dichloramine reacting with CP, 

DZ, and MA. 95% confidence intervals shown in parentheses. 

OP Pesticide kNH2Cl,OP

(M-1h-1) 

kNHCl2,OP

(M-1h-1) 

Chlorpyrifos (CP) 11.2(±1.2) 2700(±100) 

Diazinon (DZ) 21.4(±1.9) 2930(±120) 

Malathion (MA) 10.6(±0.6) 2000(±200) 
FIGURES 

39
pH

0 2 4 6 8 10 12

kobs (h-1)

0.00

0.25

0.50

0.75

1.00

1.25

[OH-] (M)

0 300x10-6 600x10-6 900x10-6

kobs (h-1)

0.00

0.25

0.50

0.75

1.00

1.25

y(x) = 1252.9(x) -0.0252 r2 = 0.99

Figure 1 Hydrolytic behavior of chlorethoxyfos (CE) over the pH range of 3-11. Insert: 

slope represents the alkaline hydrolysis rate coefficient. [CE]o = 0.5 μM, [Buffer]T

= 10 mM, and Temperature = 25±1°C. Error bars represent 95% confidence 

intervals. 

40
pH

0 2 4 6 8 10 12

kobs (h-1)

0.000

0.001

0.002

0.003

0.004

0.005

0.006

[OH-] (M)

0 25x10-6 50x10-6 75x10-6 100x10-6 125x10-6

kobs (h-1) 0.000

0.001

0.002

y(x) = 6.0 1(x) - 1.05 x 10-3 

r

2

 = 1.00

Figure 2 Hydrolytic behavior of tebupirimfos (TE) over the pH range of 3-10. Insert: slope 

represents the alkaline hydrolysis rate coefficient. [TE]o = 0.5 μM, [Buffer]T = 10 

mM, and Temperature = 25±1°C. Error bars represent 95% confidence intervals. 

41
pH

0 2 4 6 8 10 12

kobs (h-1)

0.00

0.05

0.10

0.15

0.20

0.25

[OH-] (M)

0 300x10-6 600x10-6 900x10-6

kobs (h-1)

0.00

0.05

0.10

0.15

0.20

0.25

y(x) = 222.51(x) -0.01 r2 = 1.00

Figure 3 Hydrolytic behavior of methidathion (ME) over the pH range of 3-11. Insert: slope 

represents the alkaline hydrolysis rate coefficient. [ME]o = 0.5 μM, [Buffer]T = 10 

mM, and Temperature = 25±1°C. Error bars represent 95% confidence intervals. 

42
pH

0 2 4 6 8 10 12

kobs (h-1)

0

5

10

15

20

25

30

[OH-] (M)

0 25x10-6 50x10-6 75x10-6 100x10-6 125x10-6

kobs (h-1)

0

5

10

15

20

25

30

y(x) = 2.73 x 105(x) - 6.36 x 10-2 

r

2

 = 1.00

Figure 4 Hydrolytic behavior of phosmet (PM) over the pH range of 2-10. Insert: slope 

represents the alkaline hydrolysis rate coefficient. [PM]o = 0.5 μM, [Buffer]T = 10 

mM, and Temperature = 25±1°C. Error bars represent 95% confidence intervals. 

43
Time (hours)

0.00 0.25 0.50 0.75

ln([CE]/[CE]o)

-3.0

-2.5

-2.0

-1.5

-1.0

-0.5

0.0

[HOCl]T = 10 μM

[HOCl]T = 25 μM

[HOCl]T = 50 μM

[HOCl]T = 100 μM

r

2

 = 0.98

r

2

 = 0.98

r

2

 = 0.97

r

2

 = 0.99

Figure 5 Observed first-order rate of CE loss in the presence of aqueous chlorine at pH 6.5. 

[CE]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate. 

44
Time (hours)

0.000 0.025 0.050 0.075 0.100 0.125 0.150

ln([DZ]/[DZ]o)

-6

-5

-4

-3

-2

-1

0

[HOCl]T = 10 μM

[HOCl]T = 25 μM

[HOCl]T = 50 μM

[HOCl]T = 100 μM

r

2

 = 0.99

r

2

 = 0.99 r

2

 = 0.99

r

2

 = 0.99

Figure 6 Observed first-order rate of DZ loss in the presence of aqueous chlorine at pH 6.5. 

[DZ]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate. 

45
Time (hours)

0.00 0.05 0.10 0.15 0.20 0.25 0.30

ln([MA]/[MA]o)

-6

-5

-4

-3

-2

-1

0

[HOCl]T = 10 μM

[HOCl]T = 25 μM

[HOCl]T = 50 μM

[HOCl]T = 100 μM r

2

 = 0.99

r

2

 = 0.99

r

2

 = 0.99

r

2

 = 0.99

Figure 7 Observed first-order rate of MA loss in the presence of aqueous chlorine at pH 6.5. 

[MA]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate. 

46
Time (hours)

0.00 0.05 0.10 0.15 0.20 0.25

ln([ME]/[ME]o)

-3.0

-2.5

-2.0

-1.5

-1.0

-0.5

0.0

[HOCl]T = 10 μM

[HOCl]T = 25 μM

[HOCl]T = 50 μM

[HOCl]T = 100 μM

r

2

 = 0.99

r

2

 = 0.99

r

2

 = 0.99

r

2

 = 0.99

Figure 8 Observed first-order rate of ME loss in the presence of aqueous chlorine at pH 6.5. 

[ME]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate. 

47
Time (hours)

0.0 0.1 0.2 0.3 0.4 0.5

ln([PA]/[PA]o)

-7

-6

-5

-4

-3

-2

-1

0

[HOCl]T = 10 μM

[HOCl]T = 25 μM

[HOCl]T = 50 μM

[HOCl]T = 100 μM

r

2

 = 0.99

r

2

 = 0.99

r

2

 = 0.99

r

2

 = 0.99

Figure 9 Observed first-order rate of PA loss in the presence of aqueous chlorine at pH 6.5. 

[PA]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate. 

48
Time (hours)

0.00 0.05 0.10 0.15 0.20 0.25

ln([PM]/[PM]o)

-3.0

-2.5

-2.0

-1.5

-1.0

-0.5

0.0

[HOCl]T = 10 μM

[HOCl]T = 25 μM

[HOCl]T = 50 μM

[HOCl]T = 100 μM

r

2

 = 0.99

r

2

 = 0.99

r

2

 = 0.99

r

2

 = 0.99

Figure 10 Observed first-order rate of PM loss in the presence of aqueous chlorine at pH 6.5. 

[PM]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate. 

49
Time (hours)

0.00 0.05 0.10 0.15 0.20 0.25 0.30 0.35

ln([TE]/[TE]

o)

-3.0

-2.5

-2.0

-1.5

-1.0

-0.5

0.0

[HOCl]T = 10 μM

[HOCl]T = 25 μM

[HOCl]T = 50 μM

[HOCl]T = 100 μM

r

2

 = 0.99

r

2

 = 0.99

r

2

 = 0.99

r

2

 = 0.99

Figure 11 Observed first-order rate of TE loss in the presence of aqueous chlorine at pH 6.5. 

[TE]o = 0.5 μM, [HOCl]o = 10-100 μM, [PO4]T = 10 mM, and Temperature = 

25±1°C. Experiments performed in triplicate. 

50
-log[HOCl]T

3.8 4.0 4.2 4.4 4.6 4.8 5.0 5.2

-log(kobs)

-3.0

-2.5

-2.0

-1.5

-1.0

-0.5

0.0

CP 

CE

DZ

MA

ME

PA

PM

TE

slopeavg = 1.05+0.09

Figure 12 The reaction order of chlorine with 8 OP pesticides at pH 6.5. [OP]o = 0.5 μM, 

[PO4]T = 10 mM, Temperature = 25±1°C, and [HOCl]T = 10-100 μM. 

51
[HOCl]T (μM)

0 25 50 75 100 125

kobs (h-1)

0

50

100

150

200

250

300

350

CP 

CE 

DZ 

MA 

ME 

PA 

PM

TE

Figure 13 Second-order apparent rate coefficient for the OP pesticides at pH 6.5. [OP]o = 0.5 

μM, [PO4]T = 10 mM, Temperature = 25±1°C, and [HOCl]T = 0-100 μM. Error 

bars represent 95% confidence intervals. 

52
pH

6.0 6.5 7.0 7.5 8.0 8.5 9.0 9.5

kobs (h-1)

0

25

50

75

100

125

CP 

CE 

DZ 

MA 

ME 

PA 

PM 

TE 

Figure 14 The pH dependency of the first-order observed rate coefficients for the OP 

pesticides. [OP]o = 0.5 μM, [HOCl]T = 25 μM, [Buffer]T = 10 mM, and 

Temperature = 25 °C. Error bars represent 95% confidence intervals. 

53
[HOCl]T (μM)

0 25 50 75 100 125

kobs (h-1)

0.0

2.5

5.0

7.5

10.0

12.5

15.0

CP 

CE 

DZ 

MA 

ME 

PA 

PM

TE

Figure 15 Second-order apparent rate coefficient for the OP pesticides at pH 9. [OP]o = 0.5 

μM, [CO3]T = 10 mM, Temperature = 25±1°C, and [HOCl]T = 0-100 μM. Error 

bars represent 95% confidence intervals. 

54
Ε HOMO-HFMIDI (eV)

-0.37 -0.36 -0.35 -0.34 -0.33

ln(kHOCl,OP)

13.0

13.5

14.0

14.5

15.0

15.5

Phosphorothioate

Phosphorodithioate

y(x) = 67.78(x) + 38.54

r

2

 = 0.95

y(x) = 95.22(x) + 47.09

r

2

 = 0.99

Figure 16 Relationship between the kHOCl,OP with EHOMO as a function of OP pesticide 

subgroup. Error bars about the regression line represent 95% confidence intervals. 

55
log (kB,OP)

0123456

log (kOCl-,OP)

0

1

2

3

4

5

Phosphorothioate

Phosphorodithioate y(x) = 1.00(x) + 1.22

r

2

 = 0.96

y(x) = 0.19(x) + 1.96

r

2

 = 0.99

Figure 17 Relationship between the kOCl,OP with alkaline hydrolysis rate coefficient as a 

function of OP pesticide subgroup. Error bars about the regression line represent 

95% confidence intervals. 

56
[HOCl]T x 10-6 (M)

0 50 100 150 200 250

kobs (h-1)

0.0

0.1

0.2

0.3

0.4

DZO

MAO 

PAO 

y(x) = 565.29(x) + 0.21 r2 = 0.99

y(x) = 914.13(x) + 0.00 r2

 = 0.99

y(x) = 48.35(x) + 0.00 r2

 = 0.99

Figure 18 Second-order apparent rate coefficient of 3 OP pesticide oxons at pH 9. [OP]o = 

0.5 μM, [CO3]T = 10 mM, Temperature = 25±1°C, and [HOCl]T = 0-200 μM. 

Error bars represent 95% confidence intervals for both data points and regression 

lines. 

57
Time (hours)

0.00 0.05 0.10 0.15 0.20 0.25 0.30

Concentration (

μM)

0.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

DZ

DZO

DZH

Figure 19 Experimental and model results for DZ loss in the presence of chlorine at pH 7.0. 

[DZ]o = 0.6 μM, [HOCl]T = 20 μM, [PO4]T = 10 mM, and Temperature = 25±1°C. 

Error bars represent 95% confidence intervals. Lines represent model results. 

58
Time (hours)

0.0 0.1 0.2 0.3 0.4 0.5 0.6

Concentration (

μM)

0.0

0.1

0.2

0.3

0.4

0.5

0.6

PA

PAO

PAH

Figure 20 Experimental and model results for PA loss in the presence of chlorine at pH 8.0. 

[PA]o = 0.57 μM, [HOCl]T = 25 μM, [PO4]T = 10 mM, and Temperature = 25±1°C. 

Lines represent model results. 

59
pH

6789 10

kobs (h-1)

0.000

0.005

0.010

0.015

0.020

0.025

CP 

DZ 

CP Control 

DZ Control 

Figure 21 First-order observed loss of CP and DZ in the presence of monochloramine as a 

function of pH. [NH2Cl] = 50 μM, [OP]o = 0.5 μM, Cl/N = 0.7 mol/mol, [Buffer]T

= 10 mM, Temperature = 25±1°C, and pH 6.5-9.0. Error bars represent 95% 

confidence intervals. 

60
pH

6.0 6.5 7.0 7.5 8.0 8.5 9.0

kobs (h-1)

0.00

0.01

0.02

0.03

0.04

MA

MA Control 

Figure 22 First-order observed loss of MA in the presence of monochloramine as a function 

of pH. [NH2Cl] = 50 μM, [MA]o = 0.5 μM, Cl/N = 0.7 mol/mol, [Buffer]T = 10 

mM, Temperature = 25±1°C, and pH 6.5-9.0. Error bars represent 95% confidence 

intervals. 

61
Cl/N (mol/mol)

0.00 0.25 0.50 0.75

kobs (h-1)

0

500x10-6

1x10-3

2x10-3

25x10-3

50x10-3

75x10-3

CP 

DZ 

MA

0.5x10-3

Figure 23 Observed first-order observed loss of three OP pesticides in the presence of 

monochloramine at pH 8.5. [NH2Cl] = 50 μM, [OP]o = 0.5 μM, Cl/N = 0-0.7 

mol/mol, [H3BO3]T = 10 mM, and Temperature = 25±1°C. Error bars represent 

95% confidence intervals. 

62
Time (hours)

0 50 100 150 200 250 300 350

Concentration (

μM)

0.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

10.0

20.0

30.0

40.0

50.0

60.0

pH 6.5, CP

pH 7,CP

pH 8, CP

pH 9, CP

pH 6.5, NH2Cl

pH 7, NH2Cl

pH 8, NH2Cl

pH 9, NH2Cl

Figure 24 Experimental and model results for CP loss in the presence of monochloramine. 

[CP]o = 0.5 μM, [NH2Cl]o = 50 μM, Cl/N = 0.7 mol/mol, [Buffer]T = 10 mM, pH 

6.5-9, and Temperature = 25±1°C. Error bars represent 95% confidence intervals. 

Lines represent model results. 

63
Time (hours)

0 100 200 300 400

Concentration (

μM)

0.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

10.0

20.0

30.0

40.0

50.0

60.0

pH 6.5, DZ

pH 7, DZ

pH 8, DZ

pH 9, DZ

pH 6.5, NH2Cl

pH 7, NH2Cl

pH 8, NH2Cl

pH 9, NH2Cl

Figure 25 Experimental and model results for DZ loss in the presence of monochloramine. 

[DZ]o = 0.5 μM, [NH2Cl]o = 50 μM, Cl/N = 0.7 mol/mol, [Buffer]T = 10 mM, pH 

6.5-9, and Temperature = 25±1°C. Error bars represent 95% confidence intervals. 

Lines represent model results. 

64
Time (hours)

0 50 100 150 200 250 300 350

Concentration (

μM)

0.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

10.0

20.0

30.0

40.0

50.0

60.0

pH 6.5, MA

pH 7, MA

pH 8, MA

pH 8.5, MA

pH 6.5, NH2Cl

pH 7, NH2Cl

pH 8, NH2Cl

pH 8.5, NH2Cl

Figure 26 Experimental and model results for MA loss in the presence of monochloramine. 

[MA]o = 0.5 μM, [NH2Cl]o = 50 μM, Cl/N = 0.7 mol/mol, [Buffer]T = 10 mM, pH 

6.5-8.5, and Temperature = 25±1°C. Error bars represent 95% confidence 

intervals. Lines represent model results. 

65
66

E1/2 (V vs. SCE)/0.059 V

-5.0 -2.5 0.0 2.5 5.0 7.5 10.0 12.5 15.0

log(kint)

0

1

2

3

4

5

6

7

CP: y(x) = 0.30(x) + 1.88 r2 = 0.97

DZ: y(x) = 0.30(x) + 2.09 r 2 = 0.96

MA: y(x) = 0.30(x) + 1.81 r2 = 0.97 

Figure 27 Structure-activity relationship relating the reaction of monochloramine, 

dichloramine, and hypochlorous acid with CP, DZ, and MA. Piela and Wrona (49) 

found half-wave potentials for monochloramine (-0.2 V), dichloramine (0.4), and 

hypochlorous acid (0.8 V). These values were then converted using the Nernst 

equation to the intensive property of the log ratio of the activities of the reduced 

products to the oxidized forms of the oxidant. 
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