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The formation, fate and transformation of nitromethane in potable reuse processes

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The formation, fate and transformation of nitromethane in potable reuse processes

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Content Copyright 2021 Jiaming Lily Shi
The Formation, Fate and Transformation of Nitromethane in Potable Reuse Processes
by
Jiaming Lily Shi
A Dissertation Presented to the
FACULTY OF THE USC GRADUATE SCHOOL
UNIVERISITY OF SOUTHERN CALIFORNIA In Partial
Fulfillment of the
Requirements for the Degree
DOCTOR OF PHILOSOPHY
ENGINEERING (ENVIRONMENTAL ENGINEERING)
December, 2021
ii

TABLE OF CONTENTS
LIST OF TABLES AND SCHEMES -------------------------------------------------------------------- iv
LIST OF FIGURES ----------------------------------------------------------------------------------------- v
ABSTRACT ----------------------------------------------------------------------------------------------- viii
1. INTRODUCTION --------------------------------------------------------------------------------------- 1
1.1 BACKGROUND AND MOTIVATION ----------------------------------------------------------------- 1
1.2 OBJECTIVES AND SCOPE OF WORK ---------------------------------------------------------------- 6
2. NITROMETHANE AS AN OZONE BYPRODUCT AND KEY INTERMEDIATE TO
HALONITROMETHANES IN WASTEWATER, AND N-METHYLAMINE DRUGS AS
POTENT NITROMETHANE PRECURSORS ---------------------------------------------------------- 9
ABSTRACT --------------------------------------------------------------------------------------------------- 9
2.1 OVERVIEW AND BACKGROUND ------------------------------------------------------------------ 10
2.2 MATERIALS AND METHODS ---------------------------------------------------------------------- 13
2.2.1 Sample collection ------------------------------------------------------------------------------ 13
2.2.2 Material and Reagents ---------------------------------------------------------------------- 14
2.2.3 Analytical Methods -------------------------------------------------------------------------- 14
2.2.4 Batch Ozonation Experiments ------------------------------------------------------------- 16
2.2.5 Standard Addition Experiments ----------------------------------------------------------- 17
2.3 RESULTS AND DISCUSSION ----------------------------------------------------------------------- 18
2.3.1 Nitromethane Formation in Wastewater Effluent --------------------------------------- 18
2.3.2 Nitromethane as the Precursor of Chloropicrin in Ozonated Wastewater Effluent. 23
2.3.3 N-Methylamines as Precursors of Nitromethane and Chloropicrin. ------------------ 26
2.3.4 Reaction Mechanism and Intermediates. ------------------------------------------------- 30
2.3.5 Relationship between Precursor Structure and Nitromethane Yield ------------------ 36
2.4 IMPLICATIONS ------------------------------------------------------------------------------------- 40
3. FORMATION AND PERSISTENCE OF NITROMETHANE IN OZONE-BASED WATER
REUSE TREATMENT PROCESSES ------------------------------------------------------------------- 42
ABSTRACT ------------------------------------------------------------------------------------------------- 42
3.1 INTRODUCTION ------------------------------------------------------------------------------------ 43
3.2 METHODS AND MATERIALS ---------------------------------------------------------------------- 45
3.2.1 Sample collection ------------------------------------------------------------------------------ 45
3.2.2 Material and Reagents ------------------------------------------------------------------------ 46
3.2.3 Chlorination Experiments -------------------------------------------------------------------- 47
3.2.4 RO Experiments ------------------------------------------------------------------------------ 47
3.2.5 UV and UV/AOP Experiments ------------------------------------------------------------- 48
3.2.6 Analytical Method ---------------------------------------------------------------------------- 49
3.3 RESULTS AND DISCUSSIONS ---------------------------------------------------------------------- 50
3.3.1 Formation of Nitromethane in Reuse Treatment Trains --------------------------------- 50
3.3.2 Nitromethane Removal Efficiency by Bench-Scale RO --------------------------------- 52
3.3.3 Direct and indirect photodegradation of nitromethane. ---------------------------------- 53
iii

3.3.4 Direct and Indirect Photolysis Calculation ------------------------------------------------ 62
3.3.5 Secondary disinfection ----------------------------------------------------------------------- 67
3.4 IMPLICATIONS -------------------------------------------------------------------------------------- 72
4. CHLORAMINATION OF NITROMETHANE --------------------------------------------------- 74
4.1 INTRODUCTION ------------------------------------------------------------------------------------ 74
4.2 MATERIALS AND METHODS ---------------------------------------------------------------------- 75
4.2.1 Materials and Reagents ---------------------------------------------------------------------- 75
4.2.2 Chloramination Experiments. --------------------------------------------------------------- 76
4.2.3 Deprotonation Rate Experiments ----------------------------------------------------------- 77
4.2.4 Nitromethane Chloramination Kinetics Experiments ------------------------------------ 78
4.3 RESULTS AND DISCUSSION ----------------------------------------------------------------------- 80
4.3.1 Reaction Mechanism ------------------------------------------------------------------------- 80
4.3.2 Kinetic Modeling ------------------------------------------------------------------------------ 85
4.4 CONCLUSIONS ------------------------------------------------------------------------------------- 87
5. CONCLUSIONS ---------------------------------------------------------------------------------------- 88
REFERENCES ---------------------------------------------------------------------------------------------- 92

iv

List of Tables and Schemes
Table 1. Literature summary on halonitromethane removal by O3/BAC ..................................... 11
Table 2. Basic water quality data. ................................................................................................. 13
Table 3. Retention times and quantification ions used for halonitromethanes and
nitromethane ................................................................................................................. 15
Table 4: Water quality table .......................................................................................................... 46
Table 5: Manufacturer’s Data for Membranes Tested in an RO Bench-Scale System ................. 47
Table 6: Percent change the concentration of nitromethane after each treatment step,
calculated relative to the concentration after the prior step. ......................................... 51
Table 7: RO Rejection Rates of Nitromethane in DI water .......................................................... 53

Scheme 1: N-methylamine ozonation scheme: single electron oxidation pathway. ..................... 31
Scheme 2: N-methylamine ozonation scheme: self-dimerization pathway. ................................. 31
Scheme 3: N-methylamine ozonation scheme: dissolved oxygen pathway. ................................. 31
Scheme 4: N-methylamine ozonation paired electron reaction pathway scheme ......................... 36
Scheme 5: Two possible pathways of methamphetamine intermediate ozonation toward
nitrated products. .......................................................................................................... 39
Scheme 6: Deprotonation of nitromethane to nitromethyl anion and corresponding nitronate
resonance structure. ...................................................................................................... 55
Scheme 7. Proposed Reaction Scheme of the Nitrohexane Reaction with Hydroxyl Radical
and Dissolved Oxygen Leading to the Release of Nitromethane ................................. 59
Scheme 8: Reaction mechanism of nitromethane and monochloramine/dichloramine ................ 85
Scheme 9: Reactions involve total nitromethane. ......................................................................... 85
v

List of Figures

Figure 1: Application of ozone in wastewater reclamation ............................................................. 3
Figure 2: Nitromethane formation from ozonation of Water B with nitrite-corrected ozone
doses. Error bars represent the range of experimental duplicates. Unless indicated,
pH = 7.0, [PO4] TOT = 10 mM, [t-BuOH] = 10 mM, trxn = 24 h, and T = 22 ± 1 º .......... 19
Figure 3: Nitromethane formation from ozonation of the filtered influent to two anonymous
wastewater reuse plants (i.e., secondary treated wastewater effluent) and the effluent
of one wastewater treatment plant. Panels A and D are from Water A, panels B and E
from Water C. Error bars represent the range of experimental duplicates. Unless
indicated, pH = 7.0, [PO4] TOT = 10 mM, [t-BuOH] = 10 mM, trxn = 24h,
and T = 22 ± 1º C ............................................................................................................ 20
Figure 4: Ozonation of ephedrine to putative nitrone intermediate (calculated m/z = 180.1019;
found = 180.1017; error = 1.1 ppm) and nitromethane. [Epinephrine]0 = 15 μM,
[O3]0 = 0.68 mM, [PO4] TOT = 10 mM, pH = 6, [t-BuOH] = 20 mM. ............................. 21
Figure 5: Comparison of nitromethane formation between bioreactor effluent and secondary
treated wastewater effluent. Experimental details: pH = 7, [PO4] TOT = 10 mM,
trxn = 24 h, and T = 22 ± 1 ºC. ........................................................................................ 23
Figure 6: Chloropicrin formation following standard addition of nitromethane to ozonated
wastewater effluent samples. Yellow squares represent nitromethane formation from
directchlorination without ozonation. Red circles represent nitromethane formation
from ozonation. Red circles represent nitromethane formation from ozonation with
subsequent chlorination. Panels A, B and C represent experiments with Waters A,B
and C respectively. Error bars represent the range of experimental duplicates.
Experimental conditions: pH 7, [PO4] TOT = 10 mM, [t-BuOH] = 10 mM,
[HOCl] TOT,0 = 50 mg/L for A, 375 mg/L for B, 100 mg/L for C, trxn = 24 h,
and T=22 ± 1 ºC .............................................................................................................. 25
Figure 7: Nitromethane formation from ozonation of wastewater as a function of initial nitrite
concentration. Experimental details: pH = 7, [O3]0 = 0.7 mg/mg DOC, [PO4] TOT = 10
mM, trxn = 24 h, and T = 22 ± 1 ºC. ................................................................................. 26
Figure 8: Molar yields (calculated relative to starting concentration of amine) of nitromethane
and chloropicrin from ten N-methylamines (5 μM) treated with 50 µM ozone followed
by free chlorine (50 mg/L = 704 μM). Nitromethane yields are reported prior to
chlorination. Experimental details: pH = 7.0, [PO4] TOT = 10 mM, trxn = 24 h. Error
bars represent the standard deviation of experimental triplicates. .................................. 27
Figure 9: Nitromethane formation from ephedrine into Water B. Error bars represent the
standard deviation of experimental triplicates. Experimental details: pH = 7.0,
[PO4] TOT = 10 mM, trxn = 24 h, T = 22 ± 1 ºC, ozone dose: 1 mg O3/mg DOC ............. 29
Figure 10: Molar yields (calculated relative to starting concentration of amine) of nitromethane
and chloropicrin from three N-methylamines (5 μM) treated with 50 μM ozone
followed by free chlorine (2.5 mg/L). Nitromethane yields are reported prior to
chlorination. Experimental details: pH = 7.0, [PO4] TOT = 10 mM, trxn = 24 h. Error
bars represent the standard deviation of experimental triplicates. .................................. 29
Figure 11: Nitromethane formation from ephedrine with different concentration of tert-butanol
Experimental details: pH = 7, [PO4] TOT = 10 mM, trxn = 24 h, and T = 22 ± 1 ºC. ........ 32
vi

Figure 12: Nitromethane formation from different ephedrine starting concentrations.
Experimental details: pH = 7, [PO4] TOT = 10 mM, [t-BuOH] = 10 mM, trxn = 24 h,
and T = 22 ± 1 ºC. ........................................................................................................... 32
Figure 13: Comparison of nitromethane molar yield after ozonation between benzaldehyde N-
methyloxime (nitrone intermediate) and benzymethylamine. Experimental details:
pH = 7, [PO4] TOT = 10 mM, [t-BuOH] = 10 mM, trxn = 24 h, and T = 22 ± 1 ºC ........... 33
Figure 14: Ozonation of N-benzylmethylamine to its corresponding nitrone intermediate
(benzaldehyde N-methyloxime) and nitromethane. Panel A: Concentration profile
with respect to time. Panel B: Mass balance with respect to specific ozone dose.
Experimental details: [N-benzylmethylamine]0 = 3 μM, [t-BuOH] = 20 mM, [O3]0 =
0.86 mM, [PO4] TOT = 10 mM, pH= 5.9. Error bars represent the range of
experimental duplicates. .................................................................................................. 35
Figure 15: Molar yields of nitromethane from ozonation of four N,N-alkylmethylamines with
increasing alkyl substitution. Experimental details: [t-BuOH] = 10 mM,
[PO4] TOT = 10 mM, pH = 7, trxn = 24 h. Error bars represent the standard deviation of
experimental triplicates. .................................................................................................. 38
Figure 16: Molar yields of nitromethane from ozonation of methylethylamine and 2-
(methylamino)ethanol. Experimental details: [t-BuOH] = 10 mM, [PO4] TOT = 10 mM,
pH = 7, trxn = 24 h. Error bars represent the standard deviation of experimental
triplicates. ........................................................................................................................ 38
Figure 17: Chromatographic and mass spectral evidence for intermediate produced by ozonation
of methamphetamine at 15 molar equivalent doses, including extracted ion
chromatograms for m/z values of 164.1081, and 150.0920. ........................................... 40
Figure 18: Nitromethane concentration in the effluent of each treatment step (indicated on x-axis
labels) through four reuse plant treatment processes. Panels A1−A3 present three
sequential sampling of plant A, panel B corresponding to plant B, panel C for Plant C,
and panel D for plant D. .................................................................................................. 51
Figure 19: UV molar extinction coefficient of nitromethane in DI water adjusted to pH 7, 9, or
12 with 10 mM phosphate buffer (pH 7 and 12) or borate buffer (pH 9). Spectra
obtained at 20 mg/L (pH 7), 10 mg/L (pH 9), and 5 mg/L (pH 12) to account for
differences in UV absorbance as a function of pH and maintain A < 1.0. ...................... 55
Figure 20 Direct photolysis of nitromethane at pH 7, 9, and 12. .................................................. 55
Figure 21: Apparent rate constant of direct photodegradation of nitromethane at pH 12. ........... 56
Figure 22: Competition kinetics between nitromethane and the hydroxyl radical in deionized
water at pH 7 (buffered with 10 mM phosphate) and pH 12 (no buffer). Experimental
details: [nitromethane]0 = 5 μM, [acetate]0 = 500 μg/L, [benzoate]0 = 50 μg/L,
[H2O2]0 = 0.2 mM, and T = 22 ± 1 °C. ............................................................................ 57
Figure 23. Nitromethane formation from nitrohexane during UV/H2O2 treatment. Experimental
details: [Nitrohexane]0 = 500 μM, pH = 7, [PO4] TOT = 10 mM, [H2O2]0 = 0.2 mM,
T = 22 ºC. ........................................................................................................................ 58
Figure 24. Panel A: Initial phase of nitromethane formation from nitrohexane with UV/H2O2
[Nitrohexane]0 = 1000 μM, pH = 7, [PO4] TOT = 10 mM, [H2O2]0 = 0.2mM, T = 22 ºC.
Panel B: Initial rate experiment to determine reaction rate order for nitromethane
formation from nitrohexane. Experimental details: [nitrohexane]0 = 400, 500, 750,
1000 and 1500 μM; pH = 7, [PO4] TOT = 10 mM, [H2O2]0 = 0.2mM, T = 22 ºC. ........... 59
vii

Figure 25: Nitromethane formation from different nitroalkane during UV/H2O2 treatment.
Experimental details: [Nitroalkane]0 = 500 μM, pH = 7, [PO4] TOT = 10 mM,
[H2O2]0 = 0.2 mM, T = 22 ºC. ......................................................................................... 60
Figure 26: Chlorine residual measured via DPD colorimetry after 48-hour chlorination in
samples from plant A, B and D (Panels A, B and C). ..................................................... 68
Figure 27: Halonitromethane concentration in samples collected from plant E (panel A), which
employed 2 mg/L free chlorine addition prior to RO, and halonitromethane formation
after chlorine addition (10 mg/L) to 40 mL of the samples collected from plants A, B,
and D (panels B, C, and D). Species not depicted were below detection limits (0.1
μg/L). Experimental conditions: pH = 7, 10 mM phosphate buffer, t = 48 h,
and T = 22 °C. ................................................................................................................. 69
Figure 28: Molar yield of chloropicrin from nitromethane after chlorination in samples from
plant A, B and D. ............................................................................................................. 70
Figure 29: Linear fitting between nitromethane and chloropicrin in samples from plant D after
chlorination. [HOCl]0 = 10 mg/L, quenched by sodium thiosulfate at a 1:1 molar ratio
after 48 h. Experimental conditions: [PO4] tot = 10mM, pH 7, T = 22 °C. ...................... 70
Figure 30: Comparison of nitromethane concentration in samples from plant D and chloropicrin
concentration in the same samples after chlorination. [HOCl]0 = 10 mg/L, quenched by
sodium thiosulfate at a 1:1 molar ratio after 48 h. Experimental conditions:
[PO4] tot = 10 mM, pH 7, and T = 22 °C. ........................................................................ 71
Figure 31. Predicted concentrations of monochloramine and dichloramine. Data generated
using rate constants from Jafvert and Valentine (1992) and Kintecus software package.
......................................................................................................................................... 79
Figure 32: Concentration profile of chloraminating of nitromethane to its products with respect
to time. Experiment details: Reaction time = 76 hr, [Nitromethane]0 = 2500 µg/L,
[Monochloramine]0 = 10 mM, For pH 7, 8 and 12, [PO4] TOT = 10 mM; For pH 9,
[Borate] = 10mM; For pH 10, [CO3
2-
] = 10 mM. Error bars represent the range of
experimental triplicates. .................................................................................................. 80
Figure 33: Chloramination of nitromethane to its corresponding chlorinated products and nitrate
at pH 5, 8 and 10, concentration profile with respect to time. Experimental details: For
pH 5, [HOCl] = 2.5 mM, [NH4Cl] = 2 mM, [Acetate] TOT = 10 mM. For pH 8,
[PO4] TOT = 2.5 mM, [HOCl] = 1.5 mM, [NH4Cl] = 2 mM. For pH 10, [CO3] TOT = 2.5
mM, [HOCl] = 1.5 mM, [NH4Cl] = 2 mM. Error bars represent the standard deviation
of experimental duplicates. .............................................................................................. 83
Figure 34: Speciation of chloramines and nitromethane with respect to pH. ............................... 84
Figure 35: Pseudo first order rate of chloramination of nitromethane at pH 9.14 and 10.31,
Experimental details: For pH 9.14, [Monochloramine]0 = 1 mM, [Borate] TOT = 10 mM.
For pH 10.31, [CO3] TOT = 10 mM, [Monochloramine] = 1 mM, Error bars represent
the standard deviation of experimental duplicates. ......................................................... 87
Figure 36: Nitromethane being the link between N-methlyamine type of precursors in
wastewater effluent to halonitromethanes formation during O3-chlorine processes ....... 89

viii

Abstract

As water resources become increasingly scarce, potable reuse of wastewater increasingly
attractive. An increasing number of treatment processes in used in full- and pilot- scale water
reclamation plants involve ozonation. Ozone is a highly effective oxidant and disinfectant.
However, previous research has demonstrated that ozonation of wastewater overall drastically
increases the formation potential of chloropicrin, a highly toxic disinfection byproduct, during
subsequent chlorination. Chloropicrin is synthesized by chlorinating nitromethane, suggesting
that nitromethane may be the immediate precursor of chloropicrin, although nitromethane is
unlikely to occur naturally in wastewater.
In this work we demonstrated that wastewater ozonation forms nitromethane, which is
the key intermediate toward halonitromethanes (HNMs), including chloropicrin. Ozonation of
wastewater effluent was shown to form abundant nitromethane. The formation pathway was
studied, and a category of stimulant and prescription drugs featuring N-methylamine functional
groups, including ephedrine and methamphetamine, which have previously been shown to exist
in wastewater, were found to ubiquitously formed nitromethane, typically at >50% yield, in the
presence of ozone.
The fate of nitromethane through water reuse treatment trains was characterized by
analyzing samples from five reuse operations employing ozone. Nitromethane was poorly
rejected by reverse osmosis (RO), and not removed by, and in some cases, increased by
ultraviolet/advanced oxidation processes (UV/AOP). In contrast, biological activated carbon
removed most nitromethane.
ix

The transformation of nitromethane to other products through secondary disinfection was
investigated. When free chlorine was added to the system, chloropicrin formation was
consistently observed. When chloramines were added to the system, other products such as
nitrate, monochloronitromethane and dichloronitromethane were found, and the reaction
mechanism was studied.
These results indicate that nitromethane presents a unique hazard to direct potable reuse
systems, due to its ubiquitous formation during wastewater ozonation, poor removal by RO and
UV/AOP, and facile conversion into genotoxic halonitromethanes upon secondary disinfection.
x

Preface

The motivation of this research work has originally started with the concern for the safety
of potable reuse water. I have always wanted to devote my time to mitigating human health risk
caused by emerging contaminants in drinking water. I have a strong belief that by combining
advance knowledge and real-life application together, we can improve our current water system
step by step.
This lengthy work was impossible to achieve without the help of my department, USC
Viterbi School of Engineering, Sonny Astani Department of Civil and Environmental
Engineering and my amazing research group. I am also forever grateful for the support I received
from my family.

1

CHAPTER 1.

1. Introduction

1.1 Background and Motivation
Fresh water scarcity is a severe global concern [1] for many reasons, such as population
growth, increasing demand [2], and decreasing water source availability. In California, due to
climate change, droughts are happening more frequently. The annual rainfall during an extreme
drought w recently was only 40-60% of the historical average value [3]. Previous research stated
that more than four billion people have already experienced water shortages [4] and the problem
continues to worsen. Due to this increasing stress, wastewater reclamation has been as a
considered potential way to augment water supplies. Wastewater reclamation refers to treating
wastewater to a degree that it can be safely used for drinking water, either directly, or indirectly
after passing through an environmental barrier. Indirect potable reuse is discharging highly
treated wastewater into an environmental buffer like groundwater aquifers or lakes, to be later
treated at a drinking water facility. Direct potable reuse refers to direct transmission of recycled
wastewater to the influent of a drinking water treatment plant, which demands higher water
quality for assurance of safety, but is attractive because direct potable reuse can save energy and
infrastructure expenditures for transporting the water. Indirect potable reuse has already been
implemented around the US, especially in Southern California. There are more than 20 full-scale
indirect potable reuse plants operating in the United States [5].
The standard treatment configuration for indirect potable reuse relies heavily on
membrane processes, such as microfiltration, nanofiltration, and/or reserve osmosis (RO). To
prevent membrane biofouling by microorganisms in wastewater, a disinfectant is often added
2

prior to membrane processes for water reuse. Because RO membrane can be degraded by free
chlorine [6], chloramines (i.e., monochloramine and dichloramine) are a common alternative to
chlorine, and additionally more effectively prevent aftergrowth of bacteria [7]. However, shortly
after chloramine use gained popularity, soon it was discovered that chloramine use is strongly
associated with formation of NDMA, a potent carcinogen [8], which led to the temporary closure
and subsequent retrofit of the largest water reuse plant in the world, in Orange County, CA [9].
To minimize NDMA in recycled wastewater, most modern operations use an ultraviolet light
advanced oxidation process (UV/AOP) to generate highly reactive hydroxyl radicals from UV
photolysis of hydrogen peroxide [10], which can oxidize NDMA formed during pre-disinfection.
Yet more recent research has shown that NDMA can be re-formed post UV/AOP [11]. One
alternative disinfectant to chloramines, ozone, has more recently gained attention for water reuse,
because the only known major ozone byproduct (and only ozone byproduct under regulation) is
bromate, which forms from ozone oxidation of bromide. Because low levels of bromide are
expected in wastewater effluent, using ozone as the pre-oxidant rather than chloramines is
potentially attractive, and has been selected at two full-scale water reuse facilities in California to
date.

3

Figure 1: Application of ozone in wastewater reclamation
Ozone is a strong oxidant and disinfectant that can effectively inactivate bacteria, viruses
and most micropollutants [12, 13]. Figure 1 demonstrates two common water reuse
configurations featuring ozone. Pre-ozonation can significantly reduce formation of regulated
disinfection byproducts (DBPs) like haloacetic acids and trihalomethanes formation during
subsequent chlorination [14]. Additionally, ozone is often used upstream of biological activated
carbon or other forms of biofiltration because of its ability to make dissolved organic matter
biodegradable. The lower energy consumption of O3-BAC and its reasonably good performance
for organic carbon removal has made it an attractive alternative to reserve osmosis that has
gained significant attention in recent years [15].
Despite several advantages, use of ozone for wastewater disinfection is associated with
its own problems. As an electrophile, ozone tends to attack chemical moieties with high electron
density. Besides bromate, there are many byproducts formation associated with ozone, yet all
remain unregulated. Potentially toxic oxidative byproducts such as ketones, aldehyde and
carboxylic acids increase in concentration during ozonation [16]. The formation of carbonyl
compounds and carboxylic acids is attributed to the fact that ozone tends to react with
compounds through oxygen addition pathways. The transformation products from ozonation of
4

various micropollutants were investigated by many studies. One of the well-known ozone
reactions is the Criegee mechanism of the bond cleavage of olefin groups [17]. The reaction is
achieved through the formation of an ozonide which subsequently degrades into carbonyl
products. This type of reaction can also happen to aromatic compounds which results in the
formation of ring opening products such as aliphatic carbonyl compounds. Phenol and phenolic
derived compounds also tend to have high ozone reactivity [18]. The major transformation
products from ozonating phenolic compounds are hydroquinones, catechol and ring open
products like muconic acid [19].
Ozone also tends to react with organic nitrogen compounds to form nitrogenous DBPs.
Amine containing compounds can react with ozone through an ozonolysis pathway to form
amine oxides which can be further oxidized [20]. Nitrogenous DBPs play a much more
important role than carbonous DBPs in overall toxicity [21]. Previous research has demonstrated
that ozone can decrease the formation potential of certain classes of nitrogenous DBPs, but the
formation of halonitromethanes is found to be strongly promoted by ozonation when followed by
subsequent chlorination or chloramination [22]. Many ozonation byproducts are small and polar
which makes them hard to remove by subsequent treatment using conventional physical water
treatment techniques, and the same applies to halonitromethanes. Halonitromethanes are the
compound of interest for this thesis due to their association with wastewater ozonation, and high
genotoxicity and cytotoxicity [23]. The most commonly detected halonitromethane,
trichloronitromethane (chloropicrin) was used as chemical warfare agents in World War I [24].
Since chloropicrin was first discovered in ozonated water in the 1980s [25] significant effort has
been expended to determine the formation mechanism of chloropicrin during water treatment,
and to identify the source of halonitromethane precursors in drinking water and wastewater. It
5

was recently demonstrated that the nitrogen source for chloropicrin formation during treatment
of convention water sources (e.g., rivers, lakes) comes from dissolved organic nitrogen,
especially secondary amines [26], but nothing is known about the precursor functional groups in
wastewater, nor about what specific precursor molecules to which its formation may be
attributed.
Another complication with ozone is that it decays fast in water which leaves no residual,
and disinfectant residuals are legally required for drinking water in the United States [27].
Therefore, following ozonation, often secondary disinfectants are used to meet the regulation of
disinfectant residual during drinking water distribution. For direct potable reuse, a secondary
disinfectant must be added because the water is directly distributed to household either as part of
the reuse process or during subsequent drinking water treatment. For indirect potable reuse, after
discharging into the aquifer, the water at some point will be treated again in drinking water
treatment facility and secondary disinfection will happen then. The most common secondary
disinfectants are free chlorine or chloramines. Post ozonation, subsequent chlorination or
chloramination is inevitable, and this combination is historically associated with chloropicrin
formation [28].
To fully access the risk of ozone usage in wastewater reclamation, mechanistic studies are
needed provide more information related to the formation of unwanted byproducts. More
importantly, the fate and transport of the ozone byproducts through different treatment process
can provide the information to plant designers and operators to optimize their system and reduce
public health risk.

6

1.2 Objectives and Scope of Work
The goals of this work are to first demonstrate formation of the key intermediate toward
chloropicrin during ozonation of wastewater, and to identify the precursors present in wastewater
effluent that cause the halonitromethane formation during ozonation and subsequent chlorination.
Based on previous research, we learned that one class of chemical that can serve as an efficient
precursor to halonitromethanes are secondary N-methylamines [26], because ozonation of
secondary amines leads to the formation of primary nitro compounds [20, 29]. Then with
subsequent chlorination, the nitro compound can be transformed into halonitromethanes. Many
stimulant drugs and antidepressants carry this functional group and are often detected in
wastewater effluent due to incomplete removal [30]. Therefore, we hypothesized that ozonation of
N-methylphenethylamine drugs in wastewater might lead to formation of nitromethane, which is
known to be an efficient precursor to chloropicrin during chlorination [31].
For the first part of this thesis, nitromethane formation was quantified during wastewater
ozonation, to establish that nitromethane is the linkage between N-methylamine type precursors in
wastewater and subsequent formation of chloropicrin during secondary disinfection, and the
importance of nitromethane as the precursor of chloropicrin was determined. Additionally, we are
going show that like simple aliphatic N-methylamines, N-methyl phenethylamines can also be
converted into nitromethane and determine the yield and the reaction pathway.
Nitromethane is a small neutral compound, and it is also very polar, all the above makes it
likely to pass through reverse osmosis. Nitromethane removal rates by advanced water treatment
processes applied in potable reuse have never been characterized because of the existing analytical
method were not sensitive enough to detect it at low concentrations. Therefore, the next objective
of this study was to profile the fate of nitromethane across water treatment processes in real full-
7

scale and pilot-scale potable reuse treatment plants. If nitromethane is insufficiently removed by
the advanced treatment processes of reverse osmosis and UV/AOP, its presence in final recycled
water is all but assured, likely leading to elevated chloropicrin formation during treatment of
recycled wastewater for drinking water production.
Finally, because chloropicrin and all other halonitromethanes, rather than nitromethane,
are the problematic compounds relevant to public health risk, the mechanism of halonitromethanes
formation during from nitromethane during secondary disinfection should also be understood. The
kinetics and reaction mechanism of free chlorine reactions with nitromethane to form chloropicrin
are well understood [31], and in fact this pathway is used during the industrial synthesis of
chloropicrin for use as an agricultural fumigant [32]. But because chlorination of nitromethane in
wastewater will occur in the context of relatively high levels of ammonia background, adding free
chlorine (bleach) will lead to rapid formation of inorganic chloramines. Chloramines refers to the
combination of monochloramine (NH2Cl) and dichloramine (NHCl2), which exist in an
equilibrium depending on pH and the chlorine:ammonia ratio. Chloramines are gaining popularity
because they can suppress chlorinated DBPs formation compared to free chlorine, and this helps
utility to meet drinking regulations. Therefore, either direct or in situ chloramination of
nitromethane can happen downstream of potable reuse. We believe that two forms of chloramines
will have different rate constants with nitromethane. In general, chloramines are slower
halogenating reagents than free chlorine, so the kinetic model might be rather different as the one
for free chlorine, in which deprotonation of nitromethane is the rate-limiting step below pH 12.
Also, due to having reactive nitrogen in chloramine species, the reaction mechanism can be
different as well. The objective of this portion of the thesis is to measure the rate constant of
8

chloramination of nitromethane under different condition and set up a kinetic model, also to
understand the mechanistic role of monochloramine and dichloramine respectively in this process.
Overall, this thesis focuses on the formation, fate, and transformation of nitromethane as an
ozonation byproduct in potable reuse processes, and to provide crucial information for
halonitromethane minimization during water reuse treatment design, to mitigate potential public
health risk.

9

CHAPTER 2.

2. Nitromethane as an Ozone Byproduct and Key Intermediate to
Halonitromethanes in Wastewater, and N-methylamine Drugs as
Potent Nitromethane Precursors

This chapter is published in the Environmental science and technology [33].

Abstract
Potable reuse of wastewater is expanding, and ozonation for water reuse is becoming more
common, either as a preoxidant before membranes or as part of ozone/biological activated
carbon (O3/BAC) systems. However, previous research has demonstrated that ozone drastically
increases the formation potential of genotoxic halonitromethanes (HNMs), including during
O3/BAC. Chloropicrin, the most common HNM, is synthesized by chlorinating nitromethane,
suggesting that nitromethane may be the immediate precursor of chloropicrin, although
nitromethane is unlikely to occur naturally in wastewater. We hypothesized that wastewater
ozonation forms nitromethane, which would be the key intermediate toward HNMs.
Ozonation of wastewater effluent was shown to form abundant nitromethane, which explained
the majority (in one case all) subsequent chloropicrin formation. Next, we investigated a
suspected category of nitromethane precursor: stimulant drugs such as ephedrine and
methamphetamine, and certain antidepressants. These drugs all feature N-methylamine
functional groups, and certain N-alkylamines have been shown to produce primary nitroalkanes
upon ozonation. Ozonation of N-methylamine drugs ubiquitously formed nitromethane, typically
10

at >50% yield. Subsequent chlorination converted nitromethane to chloropicrin. The reaction
mechanism was investigated to understand the variation in nitromethane yield between different
precursors. These results suggest that nitromethane fate during reuse and nitromethane control
should be investigated.

2.1 Overview and background
Intensifying water scarcity has led to increased interest in potable reuse (USEPA, 2007;
NRC, 2012). Water reuse processes frequently employ ozone, either as pretreatment before
microfiltration and reverse osmosis (RO) to prevent membrane biofouling [10], or followed by
biologically activated carbon (O3/BAC) as an alternative to RO [34]. O3/BAC is growing in
popularity as it is generally less energy intensive than RO and does not generate brine [15, 34].
Ozone is an effective oxidant for most pharmaceutical compounds in wastewater and other trace
contaminants in wastewater [35-38] and has been implemented in Europe to oxidize
micropollutants prior to wastewater discharge into the environment [39]. It generally reduces the
formation of regulated disinfection byproducts such as trihalomethanes and haloacetic acids
during subsequent chlorination [14]. DBP prevention is potentially a major benefit of ozone, as
DBPs are thought to be the major drivers of human health risk in recycled wastewater [40].
However, O3/BAC has been recently shown to increase the calculated toxicity of the final
effluent relative to traditional RO-based treatment trains

[41, 42], and ozone oxidizes bromide to
bromate, which can limit its use in high-bromide waters [36], suggesting a need to better
understand formation and control of ozone byproducts suggesting a need to better understand
formation and control of ozone byproducts.
One class of byproducts long associated with drinking water ozonation are
halonitromethanes (HNMs) [25, 26, 28, 43-45], which are highly genotoxic

[23, 46]. HNMs have
11

more recently been observed as byproducts in O3/BAC systems with post-chlorination or
chloramination: in systems treating conventional drinking water sources with O3/BAC, ozone
typically increases chloropicrin formation potential by several fold, followed by a modest
decrease across the BAC step, for a net increase compared to the influent

[47-49]. Table 1
contains a summary of halonitromethane data in O3/BAC systems found in the literature.
Table 1. Literature summary on halonitromethane removal by O 3/BAC

Compared to drinking water, less has been reported about HNM formation by ozone in
wastewater, likely because application of O3/BAC for water reuse is a relatively new
development

[15]. The ozone step of O3/BAC has been recently shown to increase subsequent
chloropicrin formation in tertiary effluent, but chloropicrin precursor removal by BAC appears to
be inconsistent, leading to inconsistent net effects by O3/BAC on chloropicrin formation
potential in wastewater effluent

[41, 42].
Despite substantial research on the topic, the precursors and formation pathway of
halonitromethanes remain somewhat unclear, particularly in wastewater. The origin of the
nitrogen in halonitromethanes formed in wastewater has been shown to be dissolved organic
Reference Water Source
Reactor
Type Removal Efficiency
Chu et al.,
2012
5

Surface Water (Lake
Taihu)
pilot
-188% after ozonation, +50% after BAC
stepwise
Chu et al.,
2015
6

Surface Water (Lake
Taihu)
pilot 29.5%-56.9%
Krasner et al.,
2012
7

Surface Water full
-226% after ozonation, +42–61% after BAC
stepwise
Selbes et al.,
2016
8

Surface Water laboratory
Biofiltration decreased HNM FP to 8–11 μg/L,
25% reduction
Selbes et al.,
2017
9

Surface Water full/pilot
Biofiltration removed some HNM precursors,
which were within the range of 25–60%
reduction
Chuang et al.,
2017
10

Secondary
Wastewater Effluent
laboratory
Ozonation increased TCNM formation at 2-5-
fold, BAC with 15 min EBCT reduced TCNM at
95%
Chuang et al.,
2019
11

Secondary
Wastewater Effluent
pilot
O3/BAC treatment increased chloropicrin
formation up to 5-fold to tertiary effluent
12

nitrogen (DON) rather than inorganic nitrogen

[50]. The remaining question is the origin of the
nitro group, which is not commonly found in natural biomolecules or DON

[51]. Recently, it was
demonstrated with model compounds that ozonation could oxidize amines to nitro compounds,
which could serve as chloropicrin precursors, and that chloropicrin formation in surface water
was attributable to primary and/or secondary amines

[26]. In particular, ozonation of secondary
N-methylamines was shown to produce nitromethane [26], which has been used industrially as
the precursor to synthetic chloropicrin (Vanderbilt et al., 1939). Nitromethane is relatively strong
carbon acid (pKA = 10.2) with potential carcinogenicity
[52]
and its conjugate base carbanion can
be successively chlorinated to form chloropicrin at 100% yield with specific-specific rate
constant between HOCl and the nitromethyl anion of k = 9500 M
-1
s
-1
[31]. Thus, we
hypothesized that nitromethane may be a singular common intermediate between
halonitromethane precursors and halonitromethanes in ozonated wastewater.
Because a few model N-methylamines (dimethylamine, sarcosine, N-
methylethanolamine) have been shown to form nitromethane at high yield

[26], we suspected
that molecules previously observed in wastewater effluent with N-methylamine functional
groups might be efficient nitromethane precursors. These include illicit stimulants (e.g.,
methamphetamine) and certain antidepressants (e.g., sertraline). Detection of illicit and
antidepressant drugs in wastewater effluent and surface water is common because they are
incompletely removed by conventional wastewater treatment process

[53, 54]. For instance,
ephedrine has been detected at 500 – 3500 ng/L (median = 1400 ng/L) in wastewater and 1.8 –
145 ng/L (median = 25 ng/L) in surface water

[30, 54]. MDMA has been found in wastewater
effluent at 30 - 110 ng/L (median = 94 ng/L) [55], and methamphetamine detected at 10 ng/L –
2000 ng/L (median = 150 ng/L) in wastewater

[30]. The antidepressant drugs sertraline (Zoloft)
13

and fluoxetine (Prozac), which bear N-methylamine groups have also been found in multiple
wastewater treatment plants and downstream surface water, indicating incomplete removal by
wastewater treatment

[56].
The objectives of this study were to (1) determine whether ozonation of wastewater
effluent leads to nitromethane formation, (2) determine whether nitromethane is the key (and
possibly exclusive) intermediate toward halonitromethanes in ozonated wastewater, (3)
investigate whether N-methylamine drugs can be ozonated to nitromethane at high yields and
then form chloropicrin during subsequent chlorination, and (4) investigate the formation pathway
of nitromethane from N-methylamine precursors.
2.2 Materials and Methods
2.2.1 Sample collection
Secondary wastewater effluent samples were obtained from three anonymous wastewater
utilities in Southern California and the effluent of a bench scale anaerobic membrane bioreactor
fed by synthetic nutrients containing potato starch, milk powder, yeast, soy oil, fatty acids, urea
and various metal ions and inorganic salts

[57]. Samples were collected in polycarbonate jugs
pre-rinsed with methanol and deionized water, filtered (0.7 µM glass fiber), and stored at 4 ºC
until use. Water quality data are available in Table 2.
Table 2. Basic water quality data.
Wastewater
treatment plant
A
Wastewater
treatment plant
B
Wastewater
treatment plant
C
Anaerobic
membrane
bioreactor effluent
pH 6.04 8.10 7.82 6.97
DOC (mg/L) 9.0 16.0 11.6* 4.3
COD (mg/L) n/a n/a 42.0 20.0
UV254 0.15 0.24 0.24 0.12
14

2.2.2 Material and Reagents
Model precursors and analytical standards were purchased at the highest available purity
from Sigma-Aldrich (St. Louis, MO). The nitrone analogue of N-benzylmethylamine was
purchased from Sigma-Aldrich under the name benzaldehyde N-methyloxime. Free chlorine
working stocks (~200 mM) were prepared by diluting reagent grade sodium hypochlorite
(available chlorine 4.00-4.99%, Sigma-Aldrich) into deionized water and standardizing
spectrophotometrically (e292 = 365 M
-1
cm
-1
) [58].
2.2.3 Analytical Methods
Nitromethane and chloropicrin were measured via gas chromatography/triple quadrupole
mass spectrometry (GC-MS/MS) (Agilent 7890B/7010, Santa Clara, CA). 5 mL aqueous samples
in 20 mL headspace vials were salted out with approximately 2 - 3 g sodium sulfate and introduced
into the GC inlet with a static headspace sampler (Agilent 7697A). Samples were first equilibrated
with moderate shaking for 20 minutes at 85 ºC, then pressured with helium into a 1 mL headspace
sample loop and injected onto a DB-1701 column (60 m ´ 0.25 mm ´ 0.25 µm) at 150 ºC. The
inlet was operated in split mode with a split ratio of 10:1 at 120 °C. Helium was used as the carrier
gas with a constant flow of 1.2 mL/min. The GC oven temperature was held at 40 ºC for 1 min and
then increased to 150 ºC at a rate of 10 ºC/min and finally raised to 250 ºC at a rate of 100 ºC/min
with a holding time of 1 min. The mass spectrometry interface temperature was set at 250 ºC. Ions
SUVA254 16.6 14.72 21.1 27.8
Ammonia
(mg/L)
1.78 46.3 46.2 66.5
Nitrate (mg/L) 11.4 0.57 0.30 75%) nitromethane. Bench-scale
experiments were conducted to verify low removal by RO in clean systems and with wastewater
effluent and to quantify the kinetics of direct and indirect photolysis of nitromethane in
UV/AOP. An explanation for increasing nitromethane concentration during AOP is proposed.
These results indicate that nitromethane presents a unique hazard to direct potable reuse systems,
due to its ubiquitous formation during wastewater ozonation, poor removal by RO and UV/AOP,
and facile conversion into genotoxic halonitromethanes upon chlorine addition.
43

3.1 Introduction
Ozonation is widely used in wastewater reclamation processes, both as an oxidant for
removing micropollutants that cannot be efficiently removed by conventional treatment
processes [35, 36, 39] and as a disinfectant [80], either prior to ecological discharge or upstream
of reuse processes. During potable reuse of wastewater, advanced treatment processes including
mem- brane filtration (microfiltration (MF) or ultrafiltration, reserve osmosis (RO)), ultraviolet
advanced oxidation process (UV/ AOP), and ozonation/biofiltration (O
3
/BAC) are employed to
ensure high water quality [81]. The most commonly used treatment train is MF-RO-UV/AOP
[15]. In these treatment trains, a disinfectant, most often chlorine, is added upstream of
membranes to prevent biofouling, but ozone is an effective alternative [82]. Preozonation of
wastewater can also decrease trihalomethane and haloacetic acid formation potential during
downstream secondary disinfection [14], and rapid ozone decay leaves no residual to damage
membranes. The application of biofiltration or biological activated carbon (BAC) post ozonation
instead of MF and RO [83, 84] is gaining popularity because ozone is capable of increasing the
biodegradability of organic matter [85] to enable reduction of total organic carbon (TOC) to
below regulatory limits for reuse while avoiding the cost, energy input, and brine disposal
associated with RO (California Water Board, 2019). Recent research indicates that ozone
oxidizes N-methyl- amines present in secondary wastewater effluent and generates abundant
nitromethane (up to ∼35 μg/L) and that nitro- methane is the key intermediate toward
halonitromethanes during subsequent chlorination [33]. Halonitromethanes are highly genotoxic
[46, 86, 87] and have been observed in disinfected wastewater [43] and recycled wastewater,
particularly in systems using ozone as a primary disinfectant or preoxidant, followed by chlorine
or chloramines as a final disinfection step [42, 88, 89]. In one study on wastewater reuse with
44

ozone/BAC, ozonation was found to increase the halonitromethane formation potential of
wastewater by 8−10 fold, but the identity of the halonitromethane precursors remained unknown
[41]. The same study found effective removal of halonitromethane precursors by the BAC step
after ozonation, but the treatment efficacy of RO and AOP for halonitromethane precursors was
not assessed.
Nitromethane formed by ozone is potentially of serious public health concern if it persists
through advanced treatment into the final disinfection step of direct potable reuse (DPR), but
nitromethane removal rates by advanced treatment processes have never been characterized, and
molecular properties of nitromethane (small, polar, and uncharged) suggest that it may persist
through RO [90]. Additionally, we hypothesized that nitromethane may persist through UV/AOP
due to anticipated low molar absorptivity at 254 nm, and relatively low reactivity with hydroxyl
radical due to the lack of C−C or C−N double bond insertion sites,1 and the electron-
withdrawing nitro group geminal to abstractable hydrogens. While nitromethane may not be a
concern in indirect potable reuse (IPR) plants, which are far more common at present, if
nitromethane persists through advanced treatment, it will have implications for future DPR
operations.
The objectives of this study were to (1) quantify nitromethane formation by ozonation at
full- and pilot-scale reuse plants, (2) investigate nitromethane fate through reuse treatment trains,
(3) understand the removal mechanisms and efficiency of advanced treatment processes at the
bench scale, and (4) quantify halonitromethane formation from nitro- methane during secondary
disinfection of finished recycled water, including bench-scale chlorination of effluent from IPR
plants to anticipate nitromethane transformation in future DPR scenarios.

45

3.2 Methods and Materials
3.2.1 Sample collection
Water samples were obtained from five full- or pilot-scale wastewater reuse plants
employing ozone as the primary disinfectant, at all available sampling points along the treatment
trains. Plants A and B are full-scale reuse plants with the same treatment configuration: O
3
-MF-
RO-UV/AOP followed by decarbonation. Plant C is a pilot- scale reuse plant employing
ozonation followed by biological activated carbon (O
3
/BAC) only. Plant D is also a pilot-scale
plant, employing sedimentation upstream of O
3
/BAC (EBCT = 12 min), followed by granular
activated carbon (GAC) and UV/AOP. Plant E is a demonstration plant for direct potable reuse
employing O
3
-BAC-RO-UV/AOP. Chlorine was added as a secondary disinfectant at plant E
after BAC. Samples were collected in 40 mL headspace free amber vials capped with Teflon-
faced septa and shipped on ice overnight to the laboratory, where they were stored at 4 °C until
analysis. Vials for samples from plant E to be collected downstream from chlorination were
augmented with 37.8 μL of a 5 g/L sodium thiosulfate stock prior to sampling, to quench
chlorine at the time of sampling. Simultaneously, two large-volume samples were collected in 1
L HDPE containers from the influent of the same plants for bench-scale testing and shipped and
stored as described above. Water quality data were acquired from plant influents after filtration
(0.7 μm glass fiber baked at 400 °C for 4 h). Chemical oxygen demand (COD), ammonia, nitrate,
and nitrite were measured using Hach TNT kits (Loveland, CO). All available water quality data
are provided in Table 4.

46

Table 4: Water quality table

3.2.2 Material and Reagents
Analytical standards of nitro- methane and chloropicrin were purchased at the highest
available purity from Sigma-Aldrich (St. Louis, MO). Standards for chloronitromethane and
dichloronitromethane at 99.9% purity were purchased from CanSyn Chem. Corp. (Toronto,
Canada). Sodium hypochlorite solutions (4−6%) purchased from Sigma-Aldrich were diluted
into Milli-Q water to make free chlorine working stocks for chlorination experiments and were
standardized by UV spectroscopy (ε
292
= 365 M
−1
cm
−1
) [58].

47

3.2.3 Chlorination Experiments
To assess conversion of nitromethane to chloropicrin in real recycled wastewater, free
chlorine was added to samples from the treatment trains of Plants A, B, and D at a uniform dose
of 10 mg/L and then quenched by sodium thiosulfate after 48 h at a 1:1 molar ratio relative to the
initial chlorine concentration. Chlorine residual was measured by adding 0.5 mL of DPD and 0.5
mL of phosphate buffer (1:1 NaH
2
PO
4
/Na
2
HPO
4
, pH 6.2, 10 mM) into 10 mL of the samples, and
then UV absorbance was measured at 515 nm and compared to a previously constructed standard
curve prepared with free chlorine in deionized (DI) water [91].
3.2.4 RO Experiments
Nitromethane rejection was evaluated using three commercially available thin-film
composite RO membranes (Table 5). Membrane coupons were tested in a flat-sheet RO module
(Sterlitech Corporation, Kent, WA) with an active membrane area of 0.014 m
2
. Spacers (1.65 and
0.43 mm thick) were used in the feed and permeate channels to provide support and increase
mixing.
Table 5: Manufacturer’s Data for Membranes Tested in an RO Bench-Scale System

To quantify nitromethane rejection, a ∼40 μg/L nitromethane feed solution was prepared
using reagent-grade nitromethane and Milli-Q water. This synthetic solution was tested on all
membranes.
The feed solution was circulated through the RO system at a rate of 0.5 L/min with the
RO concentrate recycling back into the feed tank. A high-pressure positive displacement pump
(Wanner Engineering Inc. Minneapolis, MN) was used to recirculate the feed solution. The
48

pressure was manually controlled with a needle valve at the feed outlet of the module and was
measured with a pressure transducer (Cole Parmer Instrument, Vernon Hills, IL) at the feed inlet
and the feed outlet of the module. Prior to each experiment, membrane coupons were compacted
with DI water at 2070 kPa (300 psi) for 12 h to ensure adequate time for equilibration. Once
equilibrated, the pressure was gradually increased at a rate of 35 kPa/min (5 psi/min) until the
system reached 690 kPa (100 psi). Experiments remained at 690 kPa for 1 h at a water flux of 25
L/m
2
h. Three 20 mL permeate samples were collected every 20 min and the concentration of
nitromethane was measured as described in the Analytical Methods section. Nitromethane
rejection (R in %) was calculated using where C
f
and C
p
are concentrations of the feed and
permeate solutions in μg/L.
𝑅 =
(𝐶
!
−𝐶
"
)
𝐶
!
×100%
3.2.5 UV and UV/AOP Experiments
UV experiments were conducted in a lab-scale UV chamber consisting of a UV LED
lamp (4 W 254/365 nm UV lamp, Analytik Jena) operated in short-wave (254 nm) mode inside
of a metal box with a circular aperture to produce a semicollimated beam, similar to previous
work.25 UV lamp fluence was quantified with iodide/iodate actinometry. A solution of 0.1 M
KIO
3
/0.6 M KI and 0.01 M Na
2
B
4
O
7
·10 H
2
O was irradiated in a 90 cm crystallization dish, the
absorbance of the solution was measured at 325 nm periodically [92], and the computed surface
irradiance was 0.53 mJ/ (cm
2
s). To minimize evaporation, the crystallization dish was capped
with a UV-transparent quartz disc during experiments. The pH was controlled with a 10 mM
phosphate (pH 7 and 12) or borate (pH 9) buffer. UV spectra were collected prior to photolysis
experiments to determine the peak absorbance wavelength and the extinction coefficient at 254
nm with a Cary 60 spectrophotometer (Agilent, Santa Clara, CA). For indirect photolysis
49

experiments, a hydrogen peroxide stock standardized by UV spectroscopy (ε
254
= 18.6 M
-1
cm
-1
)
[93] was added to the solution at 0.2 mM. The nitromethane reaction rate constant with hydroxyl
radical was determined via competition kinetics with benzoate (pH 12) or acetate (pH 7) [94].
Benzoate and acetate were used at 50 and 500 μg/L, respectively (k
benzoate,•OH
= 5.90 × 10
8
M
−1
s
−1

and k
acetate,•OH
= 7.9 × 10
7
M
−1
s
−1
) [94, 95]. Acetate was used at pH 7 and benzoate at pH 12
because nitromethane reacted with hydroxyl radical much more quickly at pH 12 than at pH 7, so
a fast-reacting probe was needed at pH 12 and a slower reacting probe at pH 7.
3.2.6 Analytical Method
Nitromethane and halonitromethanes were measured via static headspace gas
chromatography/triple quadrupole mass spectrometry (GC/MS) as described previously[33].
More details are provided in Chapter 2.2.3 and Table 3. Benzoate and acetate were measured via
high-pressure liquid chromatography (LC) with UV detection (Agilent 1260). Benzoate samples
were injected onto a C18 column (100 mm × 4.6 mm × 2.7 μm; Poroshell, Agilent) with an
isocratic mobile phase consisting of 30% acetonitrile and 70% aqueous formic acid (0.1%) at a
flow rate of 1.5 mL/ min. Acetate samples were injected onto an HPX-87H ion exclusion column
(300 × 7.8 mm) (Aminex, Bio-rad) with a flow rate of 0.5 mL/min and an isocratic mobile phase
of 30% acetonitrile and 70% 10 mM sulfuric acid. The column temperature was maintained at 60
°C with an integrated column heater. The detection wavelengths for benzoate and acetate were
250 and 206 nm, respectively. Intermediates from UV/AOP experiments were detected by liquid
chromatography ion mobility−time-of-flight mass spectrometry (LC− IM−QTOF) (Agilent
1290/6560). The samples were injected onto an Agilent Eclipse Plus C18 column (50 mm × 2.1
mm × 1.8 μm) and separated with a gradient mobile phase. The initial mobile phase composition
was 10% acetonitrile and 90% water, and at 1 min the mobile phase composition began a ramp to
90% acetonitrile and 10% water over 4 min, followed by 2 min isocratic at 90% acetonitrile and
50

10% water. The flow rate was 0.4 mL/min. Analytes were ionized with positive electrospray
ionization and detected in scan mode with an m/z range of 30−350.
3.3 Results and Discussions
3.3.1 Formation of Nitromethane in Reuse Treatment Trains
The influent to each of the five full- and pilot-scale plants, which was conventional
secondary effluent in all cases, contained little nitromethane (0.61−1.43 μg/L), consistent with
recent bench-scale findings that nitromethane is a byproduct caused by the treatment process
rather than a source water pollutant [33]. After ozonation, the concentration of nitromethane
increased at all five plants, ranging from 2.5 to 10 μg/L (Figure 18). Plant A was sampled three
times to assess the temporal variability of nitromethane formation at a single site (Figure
18A1,A2,A3). During the second sampling (Figure 18A2), lower nitromethane formation (2.2
μg/L) (Figure 18A2) was observed than during the other events (4.6−6.9 μg/L) (Figure
18A1,A2,A3), potentially indicating variability in the concentration of nitromethane precursors
in secondary effluent.
Nitromethane concentrations decreased after microfiltration (MF) (generally 10−20%,
but in one case 63%) (Figure 18A1,A2,A3,B), and the reason for removal by MF is unclear.
Because nitromethane is highly water soluble, removal by filtration was not expected, and the
cause may be worth investigating in the future. Nitromethane removal by RO was consistently
poor (range: 29−37%) (Figure 18A1,A2,A3,B), in contrast to most trace organic pollutants,
which are well removed by RO [96]. Nitromethane is small (MW = 61 Da), uncharged, and
highly polar (dipole moment = 3.46 D) [97], all of which are associated with poor rejection by
RO [90, 98]. Nitromethane rejection by RO was further investigated at the bench scale, as
described below. The percent change in concentration after each step is provided in Table 6.
51

Table 6: Percent change the concentration of nitromethane after each treatment step, calculated relative to the
concentration after the prior step.

Figure 18: Nitromethane concentration in the effluent of each treatment step (indicated on x-axis labels) through four
reuse plant treatment processes. Panels A1−A3 present three sequential sampling of plant A, panel B corresponding
to plant B, panel C for Plant C, and panel D for plant D.

UV/AOP was expected to degrade nitromethane, as with other low-molecular-weight
compounds of concern that can pass RO (e.g., 1,4-dioxane, NDMA) [88, 99]. However, nitro-
52

methane failed to decrease across the AOP step of any of the AOP plants sampled (Figure 18).
Surprisingly, UV/AOP increased nitromethane concentrations at four out of five evaluated plants
using UV/AOP and in all three sampling events of plant A, for reasons investigated as described
below.
In contrast to RO and UV/AOP, Figure 18C, D indicates that BAC can effectively
remove nitromethane. Nitromethane removal efficiency by BAC was 77% at plant C and 95% at
plant D. Previous research has demonstrated the bioremediation of nitro-substituted explosives,
with the associated formation of inorganic nitrogenous byproducts [100-102]. A similar process
may explain the nitromethane removal observed by BAC in plants C and D and is worth further
investigation. Plant D has an additional GAC column immediately after BAC. GAC was not
expected to sorb nitromethane due to its hydrophilicity (log K
ow
= −0.35) [103]. However, the
GAC column employed in plant D showed a similar removal efficacy as the BAC unit,
suggesting that bacterial washout from the BAC column may have colonized the GAC column,
converting the GAC column to become functionally a second BAC column.

3.3.2 Nitromethane Removal Efficiency by Bench-Scale RO
Nitromethane rejection was investigated with bench-scale RO to compare with the full-
scale results. Nitromethane rejections for the three membranes with a feed solution of ∼40 μg/L
nitromethane in DI water are shown in Table 7. Nitromethane rejection ranged from 20 to 24%.

53

Table 7: RO Rejection Rates of Nitromethane in DI water

The rejection of an RO membrane for an organic solute depends on the membrane
material, molecular weight cutoff, surface charge, and surface morphology, as well as solute
properties (i.e., structure, molecular weight, molecular geometry, pKa, and K
ow
) [98]. The low
molecular weight of nitromethane (61.04 g/mol), its neutral charge (with a pKa of 10.2 at a
feedwater pH of 8.37), and its low log K
ow
(−0.35) are consistent with poor rejection by RO
membranes [104-106].
The bench-scale results generally agree with the data collected from the full-scale RO
processes, which show nitromethane rejections of 37, 29, 31, and 34% for plants A, B, C, and D,
respectively, suggesting that poor nitromethane rejection is generally low during RO, both in
clean and fouling environments. The higher rejections observed at the full scale are likely due to
fouling deposition on the RO membranes. Increased rejection of low-molecular-weight solutes in
the presence of a fouling layer has been previously attributed to more dominant size exclusion
caused by narrower pores in the fouling layer [101-103, 107].
3.3.3 Direct and indirect photodegradation of nitromethane.
To understand the poor performance of UV/ AOP for nitromethane removal, lab-scale
UV and UV/AOP experiments were performed to quantify direct and indirect photolysis rates of
nitromethane. First, UV spectra were obtained from nitromethane at pH values well below and
above the pKa value (to obtain the spectra of nitromethane and the nitromethyl anion, its
54

conjugate base), to estimate the peak wavelength and photon absorption efficiency during UV
treatment for direct photodegradation.
At pH 7, at which nitromethane is negligibly deprotonated, the UV absorption of
nitromethane was limited, with a molar extinction coefficient indistinguishable from 0 at 254 nm
(Figure 19); the emission wavelength of the low-pressure mercury lamps have been used most
commonly in water reuse processes. At a higher pH, stronger UV absorption was observed,
consistent with stronger UV absorbance by the nitromethyl anion. At pH 9, the apparent molar
extinction coefficient at 254 nm increased to 23.4 M
−1
cm
−1
, and at pH 12, at which the majority
of the nitromethane is present as the nitromethyl anion, an absorbance peak at 233 nm emerged
(Figure 19), with a molar extinction coefficient of 6083 M
−1
cm
−1
. The new absorbance peak can
possibly be attributed to the nitromethyl anion possessing a nitronate resonance structure, which
has an additional C−C π-bond compared to nitromethane (Scheme 6), potentially making the
anion a better UV chromophore than neutral nitromethane. At pH 7 and 9, the direct
photodegradation of nitromethane was negligible over 100−250 min, corresponding to a UV
fluence of 840−2050 mJ/cm
2
(Figure 20). At pH 12, nitromethane decayed with an observed
first-order rate constant of k = 1.8 × 10
−4
s
−1
(Figure 21), corresponding to an incident fluence-
based rate constant of 1.91 × 10−3 cm
2
/mJ, resulting in an expected 73% direct photodegradation
at a typical UV/AOP incident fluence of 1000 mJ/cm
2
if the AOP were carried out at pH 12
(fluence calculations shown in section 3.3.4). However, because UV/AOP is usually performed
on neutral to slightly acidic RO permeate, negligible direct photolysis is expected in real
systems.
55

Figure 19: UV molar extinction coefficient of nitromethane in DI water adjusted to pH 7, 9, or 12 with 10 mM
phosphate buffer (pH 7 and 12) or borate buffer (pH 9). Spectra obtained at 20 mg/L (pH 7), 10 mg/L (pH 9), and 5
mg/L (pH 12) to account for differences in UV absorbance as a function of pH and maintain A < 1.0.

Scheme 6: Deprotonation of nitromethane to nitromethyl anion and corresponding nitronate resonance structure.

Figure 20 Direct photolysis of nitromethane at pH 7, 9, and 12.
56

Figure 21: Apparent rate constant of direct photodegradation of nitromethane at pH 12.
To evaluate the effectiveness of indirect photodegradation during UV/AOP, competition
kinetics was used to determine the second-order rate constant of nitromethane with hydroxyl
radical. The indirect photodegradation rate increased with respect to pH, implying that hydroxyl
radical reacts more quickly with the nitromethyl anion than nitromethane (Figure 22). The
apparent second-order rate constants were 1.63 × 10
8
M
−1
s
−1
at pH 7 and 6.62×10
8
M
−1
s
−1
at pH
12 (Section 3.3.4). The corresponding second-order rate constants for nitro- methane and the
nitromethanyl anion were calculated to be 1.63 × 10
8
and 6.70 × 10
8
M
−1
s
−1
, respectively,
indicating relatively weak selectivity for the more electron-rich anion by the highly reactive
hydroxyl radical, consistent with the reactivity/selectivity principle [108]. The extent of
nitromethane oxidation by hydroxyl radical under typical AOP conditions was estimated by
calculating the required hydroxyl radical exposure to achieve 0.5−log oxidation of 1,4-dioxane,
as previously,6 and was found to be 5.4% (Section 3.3.4).
57

Figure 22: Competition kinetics between nitromethane and the hydroxyl radical in deionized water at pH 7 (buffered
with 10 mM phosphate) and pH 12 (no buffer). Experimental details: [nitromethane] 0 = 5 μM, [acetate] 0 = 500 μg/L,
[benzoate] 0 = 50 μg/L, [H 2O 2] 0 = 0.2 mM, and T = 22 ± 1 °C.
Poor removal by direct and indirect photolysis at neutral pH explains why full-scale
UV/AOP is not effective for removing nitromethane, but it does not explain the unexpected
increase of nitromethane post UV/AOP. We hypothesized that this increase can be explained by
the indirect photodegradation of longer chain nitroalkanes by hydroxyl radical to release
nitromethane. It was recently demonstrated that secondary amines in water can be ozonated to
form various nitro- alkanes [72]. Based on existing research on hydroxyl radical reaction
mechanisms [109-112], we hypothesized that hydroxyl radical can fragment higher-molecular-
weight nitroalkanes formed during wastewater ozonation to nitromethane rather than directly
oxidizing amines to nitromethane because hydroxyl radical is not involved in established
pathways for producing nitro compounds with ozone [33, 70].
58

Figure 23. Nitromethane formation from nitrohexane during UV/H2O2 treatment. Experimental details: [Nitrohexane] 0
= 500 μM, pH = 7, [PO 4] TOT = 10 mM, [H 2O 2] 0 = 0.2 mM, T = 22 ºC.
To test the hypothesis that fragmentation of long-chain nitroalkanes by AOP can release
nitromethane, 1-nitrohexane was treated with UV/AOP. Nitromethane formed quickly after UV
irradiation began in the presence of nitrohexane and hydrogen peroxide (Figure 23), indicating
that nitromethane formation by fragmentation of larger nitroalkanes during UV/ AOP is possible
and could potentially explain the increase in nitromethane concentrations during full-scale
UV/AOP. Previous research on hydroxyl radical reaction mechanisms with alkanes suggested the
formation of peroxyl radicals by hydrogen abstraction followed by a reaction between the
corresponding carbon-centered radical and dissolved oxygen (Scheme 7), and the resulting
peroxy radicals were reported to dimerize to a tetroxide.1 Initial rate experiments were
conducted to evaluate the reaction order, and the rate order results indicate that the rate-limiting
step is the first order in nitrohexane (Figure 24), indicating that if dimerization occurs, it is not
rate limiting. Reaction intermediates leading to nitromethane from nitrohexane were sought
using LC−MS−QTOF with positive and negative ionization. Two intermediates depicted in
Scheme 7 were found, though at small abundance based on the peak area. The masses of the
observed intermediates correspond to either the tetroxide or an artifactual dimer of the peroxy
59

radical (C12H24N2O8) labeled A in Scheme 7 and the conjugate acid alcohol of the alkoxide
product labeled C in Scheme 7 (C6H13NO3), both formed during the proposed tetroxide
decomposition.

Scheme 7. Proposed Reaction Scheme of the Nitrohexane Reaction with Hydroxyl Radical and Dissolved Oxygen
Leading to the Release of Nitromethane

Figure 24. Panel A: Initial phase of nitromethane formation from nitrohexane with UV/H2O2 [Nitrohexane] 0 = 1000
μM, pH = 7, [PO 4] TOT = 10 mM, [H 2O 2] 0 = 0.2mM, T = 22 ºC. Panel B: Initial rate experiment to determine reaction
rate order for nitromethane formation from nitrohexane. Experimental details: [nitrohexane] 0 = 400, 500, 750, 1000
and 1500 μM; pH = 7, [PO 4] TOT = 10 mM, [H 2O 2] 0 = 0.2mM, T = 22 ºC.
60

Figure 25: Nitromethane formation from different nitroalkane during UV/H 2O 2 treatment. Experimental details:
[Nitroalkane] 0 = 500 μM, pH = 7, [PO 4] TOT = 10 mM, [H 2O 2] 0 = 0.2 mM, T = 22 ºC.
We suspect that the ketone product (shown as B in Scheme 7) from the tetroxide decomposition
reaction could hydrolyze to release the nitromethyl anion. Previous research has demonstrated
the analogous formation of chloroform from hydrolyzing 1,1,1-trichloroacetone [113]. Because
the pKa of nitromethane is even lower than that of chloroform (10.2 vs 15.7) [114], the
nitromethyl anion should be a better leaving group than the trichloromethyl anion during ketone
hydrolysis. Finally, the nitromethyl anion released by hydrolysis would be rapidly protonated,
leading to the nitromethane formation we observed. The observed low yields of nitromethane
relative to nitrohexane depletion (<1%) could potentially be explained by the low selectivity of
hydroxyl radical for the hydrogens on the 2-carbon of nitrohexane. These hydrogens comprise
only 2/13 possible reaction sites with hydroxyl radical and may be less attractive targets due to
their proximity to the electron- withdrawing nitro group than hydrogens farther from the nitro
group.
61

Nitrohexane was chosen as a model nitroalkane in AOP experiments for analytical convenience
and is expected to be well removed by RO, hence contributing negligibly to nitromethane
formation by AOP. To demonstrate that the generation of nitromethane during AOP can be
generalized to other nitroalkanes, additional nitroalkanes were subjected to UV/AOP treatment
(Figure 25). Shorter nitroalkanes produced a higher nitromethane yield, consistent with our
explanation for low nitromethane yield from nitrohexane, except for nitroethane. The low yield
of nitromethane from nitroethane can potentially be explained by the low stability of the primary
alkyl radical that would be produced upon the abstraction of a hydrogen atom from the 2-carbon.
Primary carbon radicals are known to be less stable than secondary carbon radicals [115]. In
addition to producing nitromethane at a higher yield, shorter nitroalkanes will likely also be
rejected by RO at lower rates, suggesting that small nitro compounds may contribute to the
nitromethane increase observed after UV/ AOP.

62

3.3.4 Direct and Indirect Photolysis Calculation

63

64

65

66

67

3.3.5 Secondary disinfection
Due to the rapid decay of ozone in water and regulations requiring drinking water to
carry a disinfectant residual during distribution [116], secondary disinfectants are likely to be
added to any recycled water post treatment in a direct potable reuse scenario [117].The most
68

common secondary disinfectant is chlorine [118]. Plant E is a direct potable reuse demonstration
plant, at which chlorine is added post RO at 2 mg/L. Halonitromethanes and nitro- methane were
measured directly from plant samples without subsequent treatment. Chlorine residual was
measured via DPD colorimetry [119] and is reported in Figure 26. Figure 27A shows that before
chlorine addition, nitromethane was formed by ozone and then effectively (94%) removed by
BAC in agreement with the previous results (Figure 18). Speciation swiftly shifted after chlorine
addition (post RO) toward the more toxic halogenated nitromethanes. The molar sum of the
measured halonitromethanes post UV/AOP corresponds to 93.6% of the nitromethane present
before chlorine addition, consistent with nitromethane serving as the primary precursor to
halonitromethanes, as observed previously in secondary effluent [33].

Figure 26: Chlorine residual measured via DPD colorimetry after 48-hour chlorination in samples from plant A, B and
D (Panels A, B and C).

69

Figure 27: Halonitromethane concentration in samples collected from plant E (panel A), which employed 2 mg/L free
chlorine addition prior to RO, and halonitromethane formation after chlorine addition (10 mg/L) to 40 mL of the
samples collected from plants A, B, and D (panels B, C, and D). Species not depicted were below detection limits (0.1
μg/L). Experimental conditions: pH = 7, 10 mM phosphate buffer, t = 48 h, and T = 22 °C.
Because all sampled plants besides plant E were indirect potable reuse plants, final
disinfectants were not added to the final effluent before plant discharge. To more broadly
evaluate the potential of nitromethane formed by ozonation to transform into halonitromethanes
in recycled water, chlorine was added to the effluent from each treatment step of plants A, B, and
D. Final chlorine residuals are shown in Figure 26. The major halonitromethane species formed
was generally chloropicrin; however, dichloronitromethane was also detected in samples from
plants A and B (Figure 27 B, C), indicating incomplete conversion of nitromethane to
chloropicrin under the conditions evaluated. Chloropicrin was the only halonitro- methane
detected in samples from plant D (Figure 27 D). Overall, chloropicrin was detected at 0.2−1 μg/L
in all three plants’ final product water after laboratory chlorination, highlighting the potential risk
of using ozone and chlorination in the direct potable reuse process, especially without a BAC
step.

70

Figure 28: Molar yield of chloropicrin from nitromethane after chlorination in samples from plant A, B and D.

Figure 29: Linear fitting between nitromethane and chloropicrin in samples from plant D after chlorination. [HOCl] 0 =
10 mg/L, quenched by sodium thiosulfate at a 1:1 molar ratio after 48 h. Experimental conditions: [PO 4] tot = 10mM,
pH 7, T = 22 °C.

71

Figure 30: Comparison of nitromethane concentration in samples from plant D and chloropicrin concentration in the
same samples after chlorination. [HOCl] 0 = 10 mg/L, quenched by sodium thiosulfate at a 1:1 molar ratio after 48 h.
Experimental conditions: [PO 4] tot = 10 mM, pH 7, and T = 22 °C.
The molar yield of chloropicrin from nitromethane in chlorinated samples from plants A,
B, and D ranged from 2 to 40% (Figure 28), which is smaller than that observed previously in DI
water,13,57 likely due to scavenging of chlorine by other species present in wastewater. The
concentration trends of chloropicrin across the treatment trains generally mirrored the
concentration profile of nitromethane prior to chlorination (data from plant D shown in Figure
30), and although the yield was not 100%, the trend is consistent with our hypothesis that
halonitromethanes as a class of byproducts generally derive from halogenation of nitromethane.
Chloropicrin formation was positively correlated to nitromethane concentrations prior to
chlorination (Figure 29).
Incomplete conversion of nitromethane to chloropicrin has potentially harmful
toxicological consequences, as monochloronitromethane and dichloronitromethane have been
shown to be more toxic than chloropicrin [23]. For plants A and D, the highest concentration of
chloropicrin was in samples taken immediately downstream from ozone addition, consistent with
greater nitromethane concentrations present in the ozonation effluent (Figure 28). The reason
72

behind the variation in molar yields between plants and treatment steps remains unclear. One
possibility is that elevated ammonia concentrations in wastewater effluent lead to in situ
chloramination rather than chlorination, but unfortunately ammonia concentration data from the
treatment train are unavailable. Chloramines are generally slower halogenating agents than free
chlorine; therefore, the rate constant between chloramines and nitro- methane is expected to be
smaller than that between chlorine and nitromethane. The kinetics of nitromethane chlorination
by chloramines merits further research.
3.4 Implications
We found that nitromethane is a ubiquitous wastewater ozonation byproduct that can
persist through common advanced treatment processes in full- and pilot-scale wastewater reuse
plants. Due its size and polarity, nitromethane is not effectively removed by membrane-based
treatment, and at neutral pH, direct and indirect photolysis by UV/AOP fail to remove
nitromethane and even produce it in some cases. In contrast, biological activated carbon showed
promise for removing nitromethane. During subsequent chlorination in real and simulated direct
potable reuse scenarios, nitromethane was transformed into halonitromethanes, including
intermediate species between nitro- methane and chloropicrin, which are highly genotoxic.
Halonitromethanes have been routinely documented in the final effluent of potable reuse systems
using chlorine for secondary disinfection [41, 42, 88, 89]. This study connects halonitromethane
formation to nitromethane formed far upstream of final disinfection and provides field- and
bench-scale evidence and explanations for which treatment strategies are likely to be effective at
mitigating halonitromethane formation during potable reuse. These observations are likely to be
most relevant to direct potable reuse scenarios because conventional drinking water processes
may not remove nitromethane prior to chlorination. In indirect potable reuse scenarios,
73

nitromethane may be biodegraded in the environmental buffer (i.e., in groundwater or a surface
water reservoir). While halonitromethane concentrations in reuse effluent are typically at least an
order of magnitude lower than more commonly measured DBPs (e.g., THMs at 0.6−38 μg/L and
HAAs at 0.1−10 μg/L compared to chloropicrin at 0.02−2.5 μg/L) [120], their high relative
toxicity (e.g., chloropicrin is 33 times more cytotoxic than trichloroacetic acid) [86] merits
concern even at low concentrations, and previous work has suggested that halonitromethanes
contribute a significant portion to the total toxicity of potable reuse water [42].

Acknowledgement
The authors acknowledge financial support from the T.F. Yen Fellowship (to J.L.S.) and the U.S.
National Science Foundation (Award No. CBET-1944810). The authors thank Dr. Emily Marron
of Stanford University for a useful suggestion on hydroxyl radical exposure estimation, and the
five anonymous participating reuse facilities

74

CHAPTER 4.

4. Chloramination of Nitromethane
4.1 Introduction
Many studies have shown that the sequence of ozonation-chlorination strongly increases
the formation of halonitromethanes [47], in particular chloropicrin. In our previous work, we
have demonstrated that nitromethane is an ozonation byproduct and it is the key and possible
sore precursor to chloropicrin during subsequent chlorination [33]. Nitromethane can persist
through advance treatment processes such reserve osmosis and be detected in the plant final
effluent water [79]. Because ozone does not leave a disinfectant residual, in order to meet
drinking water regulations, a secondary disinfectant must to be added [121]. The most widely
used secondary disinfectant is free chlorine (HOCl + OCl
-
; i.e., bleach), but free chlorine is
strongly associated with formation of many regulated disinfection byproducts of public health
concern [122]. As an alternative to free chlorine, chloramines (i.e., monochloramine and
dichloramine) are gaining popularity because chloramines can suppress the formation of
regulated trihalomethanes (THM) and haloacetic acids (HAA) relative to free chlorine while
maintaining adequate disinfection [123].
Chloramines and free chlorine are halogenating reagents which have different chemical
properties. In general, chloramines are weaker oxidants than free chlorine and react with
nucleophiles more slowly. Chlorination of nitromethane by free chlorine has been well-studied
[31]. Hypochlorite catalyzes the deprotonation of nitromethane to the corresponding
nitromethayl anion, and nitromethyl anion is oxidized by hypochlorous acid to form chloropicrin.
75

Monochloronitromethane and dichloronitromethane are rapidly formed as short-lived
intermediates which are quickly halogenated to chloropicrin as a stable final product [31].
In our previous study, samples collected from potable reuse plants showed lower than
expected level of chloropicrin formation relative to their levels of nitromethane, and also had
detectable level of monochloronitromethane and dichloronitromethane [79]. We suspected that,
due to the high background ammonia in non-nitrified wastewater, in situ chloramination was
happening when free chlorine was added (i.e., chlorination of ammonia to form monochloramine
and dichloramine). The reaction mechanism and kinetics of nitromethane chloramination remain
unclear.
The objectives of this study were to (1) determine the reaction mechanism of
chloramination nitromethane. (2) determine whether monochloramine and dichloramine play
different role when reacting with nitromethane, and (3) measure the rate constants for each
reaction step and develop a model to predict nitromethane chloramination rates under a range of
conditions.
4.2 Materials and Methods
4.2.1 Materials and Reagents
Analytical standards of nitromethane and d3-nitromethane were purchased at the highest
available purity from Sigma-Aldrich (St. Louis, MO). Standards for chloronitromethane and
dichloronitromethane at 99.9% purity were purchased from CanSyn Chem. Corp. (Toronto,
Canada). An analytical standard for chloropicrin (1000 µg/mL in methanol) was purchased from
Fisher Scientific. Sodium hypochlorite solutions (4−6%) was purchased from Sigma-Aldrich and
was standardized by UV spectroscopy (ε292 = 365 M
−1
cm
−1
). All other chemicals were
purchased from Sigma-Aldrich.
76

4.2.2 Chloramination Experiments.
To measure the yield of nitromethane chloramination products, monochloramine,
nitromethane, 1,2-dibrompropane (as an internal standard) and buffer were added to a 25 mL
amber glass vial. 1,2-dibromopropane was used as the internal standard. For different pH, buffer
was selected differently. For pH 5, acetate was used, for pH 7, 8 and 12, phosphate buffer was
used, for pH 9, borate buffer was used and for pH 10, carbonate was used. Buffer concentration
was set at 10 mM for all samples, and the initial concentration of nitromethane and
monochloramine were 2500 µg/L and 0.5 mM respectively. Monochloramine stocks were
prepared by titrating an equal volume of 48 mM ammonium chloride into 40 mM free chlorine
solution, the pH of the monochloramine stock was adjusted to above 8.5 to prevent dichloramine
formation. Monochloramine was standardized by UV-spectrophotometry.
After 67 hour of reaction time, sodium thiosulfate was added at a 1:1 molar ratio relative
to initial chloramine concentration to quench the excess chloramine. 5 mL of the sample was
withdrawn and acidified to pH 1, then 5 mL of MtBE was added. After shaking for one minute, 1
mL of the organic layer was transferred into a GC vial, and injected to a GC/MS (Agilent
7890/7010) to analyze nitromethane, monochloronitromethane, dichloronitromethane and
chloropicrin. Nitrate was measured by using Hach Nitrate TNTplus Vial test (0.23-13.5 mg/L
NO3
-
/N), when below detection limit, 100 µL of a standard solution (13.8 mg/L NO3
-
/N) was
added into 5 mL sample. Because breakpoint chlorination can cause nitrate formation as well,
blank controls were made in parallel with the reaction mixture. The nitrate formation for each
time point was determined by subtracting the background nitrate measured from the blank
control.

77

4.2.3 Deprotonation Rate Experiments
Nitromethane is a weak acid (pKa = 10.2) that can exchange protons with water. The rate
of deprotonation was measured by measuring the rate of deuterium labeled nitromethane losing
deuterium in water (i.e., exchanging D for H). Due to deprotonation, d3-nitromethane will
gradually lose the deuterium and exchange for regular proton because the concentration of
deuterium is much smaller than the concentration of regular protons. By measuring the loss of
labeled nitromethane, the deprotonation rate can be determined. Deprotonation reactions are
general base catalyzed; in order to maintain a steady pH, phosphate buffer was used. Because
phosphate buffer can increase the deprotonation rate through general base catalysis, three
different buffer concentrations were used to extrapolate back to a hypothetical zero-buffer rate.
By plotting the measured rate constant against buffer concentration and extrapolate the liner
fitting of the data to y-axis, the true deprotonation rate constant of d3-nitromethane in water can
be determined.
Reaction mixtures were prepared in a 250 mL glass beakers. The initial concentration of
d3-nitromethane was 1 mg/L, and 1,2-dibromopropane was used as the internal standard. The
ionic strength of the solution was fixed at 20 mM with sodium chloride. For each time point, 5
mL of the sample was withdrawn and immediately extracted with 5 mL of MTBE. After shaking
for one minute, the organic layer was transferred into a 15 mL centrifuge tube. Magnesium
sulfate was used for de-watering the sample. After the de-watering step, sample was injected into
GC/MS. Four different multiple reaction transitions were monitored, which are 64 to 46, 63 to
46, 62 to 46, and 61 to 46, for tri-, di-, mono-, and unlabeled nitromethane, respectively. The
observed rate constant was computed based on the 64 to 46 transitions.

78

4.2.4 Nitromethane Chloramination Kinetics Experiments

To assess the conversion of nitromethane to chloropicrin by chloramines, free chlorine
and ammonium chloride were added to the reactor to achieve in situ chloramination. To prevent
products escaping through volatilization, a complete sealed syringe was used as the reactor.
Glass syringes were purchased from Fisher Scientific and were cleaned and baked at 400 °C
prior to using. To maintain a steady pH, buffer was added, and the buffer concentration was set
at low level due to the general acid catalyzed disproportionation reaction of monochloramine to
dichloramine and ammonia. At pH 5, 2.5 mM of sodium hypochlorite and 2 mM ammonium
chloride were added with 10 mM of acetate buffer. At pH 8 and 10, 1.5 mM sodium hypochlorite
and 2 mM of ammonium chloride with 2.5 mM buffer were added into the reactor. 5 mL of
sample was withdrawn each time, 1:1 molar ratio of sodium thiosulfate was added to quench the
chloramine. then HCl was used to adjust the pH below 1 to ensure all chlorinated nitromethane
species remained in the neutral form prior to extraction. To prevent excess quencher reducing the
chlorinated species, liquid-liquid extraction was done immediately after quenching. After
shaking for 1 min with 5 mL of methyl tert-butyl ether (MtBE), 1 mL of the sample was
transferred into a GC vial. Nitromethane and chlorinated species were measured on the GC/MS
as described below. Nitrate was measured at each time point with a Hach TNT plus nitrate kit.
The chloramine speciation with respect to time was predicted by chemical kinetics modeling as
shown in Figure 31. In order to extract rate constant at pH 8 and 10, experiments were operated
under pseudo first order condition, the concentration change of monochloramine was controlled
within 5%.
79

Figure 31. Predicted concentrations of monochloramine and dichloramine. Data generated using rate constants from
Jafvert and Valentine (1992) and Kintecus software package.

4.2.5 Analytical Methods

Nitromethane and chlorinated nitromethane samples were injected onto a DB-1701
column (60 m´ 0.25 mm ´ 0.25 μm) The inlet was operated in split mode with a split ratio of
10:1 at 150 °C. Helium was used as the carrier gas with a constant flow of 1.2 mL/min. The GC
oven temperature was held at 40 ºC for 1 min and then increased to 150 ºC at a rate of 10 ºC/min
and finally raised to 250 ºC at a rate of 100 ºC/min with a holding time of 1 min. The mass
spectrometry interface temperature was set at 250 ºC. Ions were generated with an electron
ionization (EI) source kept at 230 ºC, and both quadrupoles were kept at 150 ºC. Mass spectra of
nitromethane and chloropicrin were acquired in multiple reaction mode with m/z transitions of
61 to 46 (nitromethane), and 117 to 82 (chloropicrin) at collision energies at 5 eV and 30 eV,
respectively. Monochloronitromethane and dichloronitromethane were measured with the m/z
transitions of 48.9 to 48.9 and 82.9 to 46.9. 1,2-dibromopropane was used as the internal
standard and was measured with the transition of 121 to 121.

0 50000 100000 150000 200000
0.0000
0.0005
0.0010
0.0015
time (s)
Concentration (M)
NH
2
Cl
NHCl
2
pH 5.1
0 50000 100000 150000 200000 250000
0.0000
0.0005
0.0010
0.0015
0.0020
time (s)
Concentration (M)
NH
2
Cl
NHCl
2
pH 8.3
0 50000 100000 150000 200000 250000
0.0000
0.0005
0.0010
0.0015
0.0020
time (s)
Concentration (M)
NH
2
Cl
NHCl
2
pH 9.9
80

4.3 Results and Discussion
4.3.1 Reaction Mechanism

Figure 32: Concentration profile of chloraminating of nitromethane to its products with respect to time. Experiment
details: Reaction time = 76 hr, [Nitromethane] 0 = 2500 µg/L, [Monochloramine] 0 = 10 mM, For pH 7, 8 and 12, [PO 4]
TOT = 10 mM; For pH 9, [Borate] = 10mM; For pH 10, [CO3
2-
] = 10 mM. Error bars represent the range of
experimental triplicates.
To better understand the reaction mechanism, the yields of monochloronitromethane,
dichloronitromethane and chloropicrin during chloramination of nitromethane were measured at
a range of pH values. After 67 hours of reaction, at neutral pH, there was a significant amount of
nitromethane remained unreacted, while at higher pH, less nitromethane was unreacted which
indicates the overall reaction rate of chloramination of nitromethane increase with respect to pH.
This is consistent with the chlorination of nitromethane; the rate limiting step of nitromethane
chlorination by free chlorine is the deprotonation of nitromethane [31].
By accounting only nitromethane, monochloronitromethane, dichloronitromethane and
chloropicrin, there was a significant amount of mass balance missing, especially at pH 7, 8 and 9.
At higher pH, mass balance recovered. This indicates that at neutral pH, there are potentially
other products forming during the process besides chlorinated nitromethanes. We suspected that
7 8 9 10 11 12
0
10
20
30
40
pH
Concentration (µM)
Chloronitromethane
Dichloronitromethane
Nitromethane
Chloropicrin
Nitrate
mass balance
81

one potential product was nitrate, and by measuring nitrate, 90% of the mass balance was
recovered at pH 7 and 8.
The yield of chloropicrin decreased with respect to pH, meanwhile the yield of
monochloronitromethane and dichloronitromethane increased significantly under basic
conditions. Overall, the total halonitromethane yield increase at higher pH. Previously, it was
believed that the initial chlorine addition was always the rate limiting step, and that the
secondary and tertiary chlorine additions were considered nearly instantaneous [31]. This
contradicts with the result shown in Figure 32. The pKas of monochloronitromethane and
dichloronitromethane are 7.05 and 6.23 respectively. At basic pH, for both
monochloronitromethane and dichloronitromethane, the predominant species would be in the
deprotonated form. By having more chlorine substitution, the strong electron withdrawing effect
will decrease the electron density of the anion, makes it harder to oxidize. At high pH, more
nitromethyl anion is available for halogenation. But instead of chloropicrin,
dichloronitromethane and monochloronitromethane are the final reaction products due to the
decreased electron density of the corresponding halonitromethyl anions.
In Figure 33, different products formation were shown with respect to time under acidic,
neutral, and basic conditions. At pH 5, initially the predominant chloramine species was
monochloramine, then as time proceeds, the dichloramine concentration increases and eventually
reached a steady state. At pH 8 and 10, condition was set to ensure monochloramine can stabilize
and the formation of dichloramine is minimized (Figure 31).
At pH 5, the reaction is very slow. A small amount of nitrate formation was observed in
the beginning, but the concentration of nitrate did not increase with respect to time. The major
product is chloropicrin, and chloropicrin concentration increases after monochloramine depletion
82

which indicates halogenation can be achieved by dichloramine. The mass balance remained
relatively constant.
At pH 8, nitromethane reacted much faster with chloramine. As shown in Figure 33, the
concentration of monochloronitromethane and dichloronitromethane first increased and then
decreased. Chloropicrin, as the final product of halogenation, increases and reaches to steady
level. Nitrate is the main product from the reaction. Mass balance maintained constant
throughout the process which indicates nitrate formation and halogenation are the only two
reactions taken place at pH 8.
At pH 10, low level of nitrate formation was observed as shown in Figure 33. Chlorinated
nitromethane are the major products. Monochloronitromethane was formed quickly, and the
concentration increases with respect to time. High amounts of dichloronitromethane were also
formed.l; while a lower level of chloropicrin was detected compared to pH 5 and 8.
83

Figure 33: Chloramination of nitromethane to its corresponding chlorinated products and nitrate at pH 5, 8 and 10,
concentration profile with respect to time. Experimental details: For pH 5, [HOCl] = 2.5 mM, [NH 4Cl] = 2 mM, [Acetate]
TOT = 10 mM. For pH 8, [PO 4] TOT = 2.5 mM, [HOCl] = 1.5 mM, [NH 4Cl] = 2 mM. For pH 10, [CO 3] TOT = 2.5 mM,
[HOCl] = 1.5 mM, [NH 4Cl] = 2 mM. Error bars represent the standard deviation of experimental duplicates.
0 2000 4000 6000
0
20
40
60
0
1
2
3
4
5
time (min)
Concentration (mM)
TCNM/MCNM/DCNM (mM)
Nitromethane
Nitrate
Chloropicrin
Monochloromethane
Dichloronitromethane
mass balance
pH 5.11
0 1000 2000 3000
0
10
20
30
40
50
0
2
4
6
8
time (min)
Concentration (mM)
TCNM/DCNM/MCNM (mM)
Nitromethane
Nitrate
Chloropicrin
Monochloromethane
Dichloronitromethane
mass balance
pH 8.35
0 1000 2000 3000
0
10
20
30
0
2
4
6
time (min)
Concentration (mM)
TCNM/DCNM/MCNM (mM)
Nitromethane
Nitrate
Chloropicrin
Monochloromethane
Dichloronitromethane
mass balance
pH 9.9
84

Figure 34: Speciation of chloramines and nitromethane with respect to pH.
A speciation diagram of chloramines and nitromethane was made with respect to pH
(Figure 34). Because nitrate formation peaked at pH 8 and more nitrate was formed under
conditions favoring monochloramine stabilization, we concluded that nitrate is formed is from
the reaction between monochloramine and neutral nitromethane.
Based on the results, the following mechanism (Scheme 8) was proposed. Instead of
acting as an oxidant, monochloramine may also be react as a nucleophile, having a lone pair of
electrons on the nitrogen [124]. SN2 nucleophilic substitution would lead to the nitro group
leaving, which can potentially explain nitrate formation. Chlorinated products are resulting from
the reaction between nitromethyl anion with monochloramine and dichloramine. At acidic to
neutral pH, same as chlorination, deprotonation of nitromethane is the rate limiting step. When
the pH exceeds the pKa, more nitromethyl anion will be available and overall chlorination rate
becomes faster. Unlike what has been reported in previous research [31], we believe increased
chlorine substitution will decrease the reaction rate constant due to the strong electron
withdrawing effect.
5 6 7 8 9 10 11 12
0
20
40
60
80
100
pH
percent (%)
Monochloramine
Dichloramine
Nitromethane
Nitromethyl anion
85

Scheme 8: Reaction mechanism of nitromethane and monochloramine/dichloramine
4.3.2 Kinetic Modeling

To further understand nitromethane reactions with chloramines, a rate model was
constructed. Total nitromethane concentration was measured with respect to time because the
analytical method cannot distinguish neutral nitromethane and nitromethyl anion. Reactions
involving nitromethane and nitromethyl anion are listed in Scheme 9. Based on the sum rule of
derivatives, the rate of total nitromethane changing equals to the summation of the rate of
nitromethane and the rate of nitromethyl anion changing over time.

Scheme 9: Reactions involve total nitromethane.
The rate expressions for nitromethane and nitromethyl anion were written as following.
At basic pH, deprotonation and protonation reaction can be assumed at equilibrium. The pKa
value of nitromethane is 10.2 [125]. When pH reaches above 9, the amount of dichloramine
H
3
C
N
O
O
H
2
C
N
O
O
H
Cl
NH
2
H
3
C
N
Cl
H
H
+ NO
2
-
further oxidation
NO
3
-
NH
2
Cl
C
H
N
O
O
k
d
k
1
k
2
Cl
k
3
NHCl
2
NH
2
Cl
NHCl
2
C
N
O
O
Cl
Cl
k
4
NH
2
Cl
NHCl
2
C
N
O
O
Cl
Cl
Cl
H
2
O
H
2
C
N
O
O
OH
+ Cl
-
k
5
OH
-
+ CH
3
NO
2
k
d
k
p
CH
2
NO
2
+ H
+
CH
3
NO
2
+ NH
2
Cl NO
3
-
k
1
CH
2
NO
2
+ NH
2
Cl
k
2
CH
2
ClNO
2
CH
2
NO
2
+ NHCl
2
CH
2
ClNO
2
k
3
86

present in the system is so low that it can be considered negligible, hence k3 will not affect the
experiment observed rate constant.
[𝐶𝐻
#
𝑁𝑂
$
]
%
=[𝐶𝐻
#
𝑁𝑂
$
]+[𝐶𝐻
$
&
𝑁𝑂
$
]

𝑑[𝐶𝐻
#
𝑁𝑂
$
]
%
𝑑𝑡
=
𝑑[𝐶𝐻
#
𝑁𝑂
$
]
𝑑𝑡
+
𝑑[𝐶𝐻
$
&
𝑁𝑂
$
]
𝑑𝑡

𝑑[𝐶𝐻
#
𝑁𝑂
$
]
𝑑𝑡
= −𝑘
'
[𝑂𝐻
&
][𝐶𝐻
#
𝑁𝑂
$
]+ 𝑘
"
[𝐻
(
][𝐶𝐻
$
&
𝑁𝑂
$
]
−𝑘
)
[𝑁𝐻
$
𝐶𝑙][𝐶𝐻
#
𝑁𝑂
$
]

𝑑[𝐶𝐻
$
&
𝑁𝑂
$
]
𝑑𝑡
= 𝑘
'
[𝑂𝐻
&
][𝐶𝐻
#
𝑁𝑂
$
]− 𝑘
"
[𝐻
(
][𝐶𝐻
$
&
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$
]

−𝑘
$
[𝑁𝐻
$
𝐶𝑙][𝐶𝐻
$
&
𝑁𝑂
$
]− 𝑘
#
[𝑁𝐻𝐶𝑙
$
][𝐶𝐻
$
&
𝑁𝑂
$
]

[𝐶𝐻
#
𝑁𝑂
$
]=𝛼
*
[𝐶𝐻
#
𝑁𝑂
$
]
%

[𝐶𝐻
#
𝑁𝑂
$
]=𝛼
)
[𝐶𝐻
#
𝑁𝑂
$
]
%

𝛼
*
=
[𝐻
(
]
[𝐻
(
]+𝐾
+

𝛼
)
=1− 𝛼
*

𝑑[𝐶𝐻
#
𝑁𝑂
$
]
%
𝑑𝑡
= −𝑘
)
𝛼
*
[𝑁𝐻
$
𝐶𝑙][𝐶𝐻
#
𝑁𝑂
$
]
%
− 𝑘
$
𝛼
)
[𝑁𝐻
$
𝐶𝑙][𝐶𝐻
#
𝑁𝑂
$
]
%

− 𝑘
#
𝛼
)
[𝑁𝐻
$
𝐶𝑙][𝐶𝐻
#
𝑁𝑂
$
]
%

'[-.
!
/0
"
]
#
'2
= −(𝑘
)
𝛼
*
[𝑁𝐻
$
𝐶𝑙]+ 𝑘
$
𝛼
)
[𝑁𝐻
$
𝐶𝑙]+ 𝑘
#
𝛼
)
[𝑁𝐻
$
𝐶𝑙])[𝐶𝐻
#
𝑁𝑂
$
]
%

𝑘
3
=𝑘
)
𝛼
*
[𝑁𝐻
$
𝐶𝑙]+𝑘
$
𝛼
)
[𝑁𝐻
$
𝐶𝑙]+𝑘
#
𝛼
)
[𝑁𝐻𝐶𝑙
$
]

𝑙𝑛:
𝐶
𝐶
*
; =−𝑘
3
𝑡

To calculate the rate constant k1 and k2 shown in the scheme 9, the observed rates from
two set of experiments with different condition were used (Figure 35). Based on the pH and the
observed rate constant, k1 was calculated to be 0.0147 M
-1
s
-1
and k2 is 0.00335 M
-1
s
-1
. This
indicates that SN2 nucleophilic substitution is faster than halogenation, which agrees with our
87

result showing at neutral pH nitrate is the major product. More work needs to be done to better
corporate deprotonation reaction of nitromethane into our current model.

Figure 35: Pseudo first order rate of chloramination of nitromethane at pH 9.14 and 10.31, Experimental details: For
pH 9.14, [Monochloramine] 0 = 1 mM, [Borate] TOT = 10 mM. For pH 10.31, [CO 3] TOT = 10 mM, [Monochloramine] = 1
mM, Error bars represent the standard deviation of experimental duplicates.
4.4 Conclusions
The reaction mechanism and kinetics of nitromethane chloramination is strong affected
by pH because of the speciation of nitromethane and chloramines. At acidic pH, deprotonation of
nitromethane is the rate limiting step, nitromethyl anion react with dichloramine to form
chloropicrin. At neutral pH, nitromethane and monochloramine react through SN2 nucleophilic
substitution to form nitrate, halonitromethanes are the minor products due to less nitromethyl
anion available. At basic pH, deprotonation is significant faster, more nitromethyl anion is
available. Chlorination of nitromethane is the predominant reaction. However, at high pH, the
secondary and the tertiary chlorine addition reaction slow down so monochloronitromethane and
dichloronitromethane are the major product at reaction time relevant to water treatment.

0 500 1000 1500 2000
-1.0
-0.5
0.0
time (min)
ln(C/C
0
)
pH 9.14
k = 0.00048
0 1000 2000 3000
-2
-1
0
time (min)
ln(C/C
0
)
pH 10.31
k = 0.0008348
88

CHAPTER 5.

5. Conclusions

This dissertation sought to identify toxic byproducts formation during potable reuse processes to
mitigate public health risk. The overall objective was apporached from three fronts:
1. Identifying the key precursor of the extremely toxic nitrogenous disinfection byproducts,
halonitromethanes. The principal chemical formation pathway of halonitromethanes from
a class of precursors, N-methylamines, that are likely to be present in wastewater effluent.
2. Investigating the fate of an ozone byproduct and key intermediate toward
halonitromethanes, nitromethane, through water treatment trains that are currently
implemented in direct and indirect potable reuse plants. It was demonstrated the
nitromethane can persist through most of the exisiting treatment units and experiments to
rationalize the reason behind the concentration changes at each different treatment step
were performed and interpreted.
3. Understanding the transformation pathway of nitromethane upon secondary disinfection,
filling in the knowledge gap on reactions with chloramines, which is another commonly
used secondary disinfectant, and how final transformation products formation is affected
by chloramination versus more conventional chlorination.

In chapter 2, we have successfully identified the one key intermediate for
halonitromethane formation which is very unusual in disinfection byproducts research
(Figure 36). Above 70% of the halonitromethane formation in ozonated water can be
attributed to nitromethane, and ozonation of wastewater effluent leads to significant amount
89

of nitromethane formation. Even though the main precursors led to nitromethane formation
aren’t immediately clear, we have demonstrated that a class of chemicals can serve as good
candidates and likely contributed a major portion of the overall nitromethane formation.
Ozonation of N-methylamine type of precursors all can form nitromethane from modest to
high yield. The yield varation is determined by the pka of the 𝛼-carbon next to the methyl
group carbon because the key step for the key nitrone intermediate formation is a
deprotonation step. The more acidic the carbon is, the easier the deprotionation at neutral pH,
which explains higher yield. With this information, yield of unknown precursors can be
predicted.

Figure 36: Nitromethane being the link between N-methlyamine type of precursors in wastewater effluent to
halonitromethanes formation during O 3-chlorine processes

In Chapter 3, we showed that nitromethane is a ubitiquous byproduct in every plant
that employed aqueous ozonation, and it can persisit through most of the advanced treatment
90

units at each plant. Membrane processes such as microfiltration and reserve osmosis are not
efficient at removing nitromethane due to nitromethane being neutral, polar and small in size.
Generally reserve osmosis demonstrated rejection ranged from 20% to 25%, and the rejection
is consisitent with different feed concentration and different membranes. UV/AOP fails to
degrade nitromethane and can produce additional nitromethane by degrading longer chain
nitro compounds. Biological actived carbon shows sufficient removal of nitromethane which
indicates bioremediation of nitromethane can be achieved by certain microbial communities.
Upon secondary disinfection, toxic transformation products of nitromethane were detected in
samples such as monochloronitromethane, dichloronitromethane and chloropicrin. However,
the overall yields for halonitromethanes were lower than expected, based on previously
research on nitromethane chlorination, which led to our efforts in chapter 4 to characterized
nitromethane reactions with chloramines.
In chapter 4, the reaction kinetics and mechansim of nitromethane chloramination
were studied, because chloramines (monochloramine and dichloramine) are becoming
popular alternative secondary disinfectant to chlorine. Similar to chlorination of
nitromethane, the formation of halonitromethanes such as monochloronitromethane,
dichloronitromethane and chloropicrin was resulted from nitromethyl anion reacting with
monochloramine and dichloramine. However, at neutral pH, an unexpected SN2 nucleophilic
subsitution was the predominant pathway, producing nitrate as the main transformation
product. This indicates that by using chloramines as the secondary disinfectant, at neutral pH,
the overall toxicity of disinfected water may be reduced, because nitrate is not a potent
carcinogen, unlike halonitromethanes. Meanwhile, at basic pH, the predominant
91

transformaiton products of chloriminating nitromethane are monochloronitromethane and
dichloronitromethane, which can cause public health concern.
The research described in this dissertation contains information that can provide
guidelines for potable reuse operation on how to mitigate risk. First, it raised attention to a
newly-identified ozone byproduct, nitromethane, and the negative health impact associated
with nitromethane transformation upon seconadry disinfection. It also provided the
knowledge that can help engineers to tailor the current potable reuse treatment technology on
how to minimize the concentration of nitromethane, and subsequent halonitromethane
formation in the final effluent for potable reuse plants.

92

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Abstract (if available)

Abstract As water resources become increasingly scarce, potable reuse of wastewater increasingly attractive. An increasing number of treatment processes in used in full- and pilot-scale water reclamation plants involve ozonation. Ozone is a highly effective oxidant and disinfectant. However, previous research has demonstrated that ozonation of wastewater overall drastically increases the formation potential of chloropicrin, a highly toxic disinfection byproduct, during subsequent chlorination. Chloropicrin is synthesized by chlorinating nitromethane, suggesting that nitromethane may be the immediate precursor of chloropicrin, although nitromethane is unlikely to occur naturally in wastewater. ❧ In this work we demonstrated that wastewater ozonation forms nitromethane, which is the key intermediate toward halonitromethanes (HNMs), including chloropicrin. Ozonation of wastewater effluent was shown to form abundant nitromethane. The formation pathway was studied, and a category of stimulant and prescription drugs featuring N-methylamine functional groups, including ephedrine and methamphetamine, which have previously been shown to exist in wastewater, were found to ubiquitously formed nitromethane, typically at >50% yield, in the presence of ozone. ❧ The fate of nitromethane through water reuse treatment trains was characterized by analyzing samples from five reuse operations employing ozone. Nitromethane was poorly rejected by reverse osmosis (RO), and not removed by, and in some cases, increased by ultraviolet/advanced oxidation processes (UV/AOP). In contrast, biological activated carbon removed most nitromethane. ❧ The transformation of nitromethane to other products through secondary disinfection was investigated. When free chlorine was added to the system, chloropicrin formation was consistently observed. When chloramines were added to the system, other products such as nitrate, monochloronitromethane and dichloronitromethane were found, and the reaction mechanism was studied. ❧ These results indicate that nitromethane presents a unique hazard to direct potable reuse systems, due to its ubiquitous formation during wastewater ozonation, poor removal by RO and UV/AOP, and facile conversion into genotoxic halonitromethanes upon secondary disinfection.

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Asset Metadata

Creator Shi, Jiaming Lily (author)

Core Title The formation, fate and transformation of nitromethane in potable reuse processes

School Viterbi School of Engineering

Degree Doctor of Philosophy

Degree Program Engineering (Environmental Engineering)

Publication Date 11/19/2021

Defense Date 11/19/2021

Publisher University of Southern California (original), University of Southern California. Libraries (digital)

Tag chloropicrin,nitromethane,OAI-PMH Harvest,ozone,potable reuse

Format application/pdf (imt)

Language English

Contributor Electronically uploaded by the author (provenance)

Advisor McCurry, Daniel (committee chair), Childress, Amy (committee member), Fokin, Valery (committee member), Smith, Adam (committee member)

Creator Email jiamings@usc.edu,lilyshi0530@gmail.com

Permanent Link (DOI) https://doi.org/10.25549/usctheses-oUC17138271

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