Balancing Bromate Formation, Organics Oxidation

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Balancing Bromate Formation, Organics Oxidation, and Pathogen Inactivation: The Impact of Bromate Suppression Techniques on Ozonation System Performance in Reuse Waters

Peter Hamilton Buehlmann

Thesis submitted to the faculty of the Virginia Polytechnic Institute and State University in partial fulfillment of the requirements for the degree of

Master of Science In Environmental Engineering

John T. Novak Amy J. Pruden-Bagchi Charles B. Bott

August 5th, 2019 Blacksburg, VA

Keywords: Ozone, Bromate, Monochloramine, Chlorine-Ammonia Process, Advanced Oxidation

Balancing Bromate Formation, Organics Oxidation, and Pathogen Inactivation: The Impact of Bromate Suppression Techniques on Ozonation System Performance in Reuse Waters

Peter Hamilton Buehlmann

Academic Abstract

Ozonation is an integral process in ozone-biofiltration treatment systems and is beginning to be widely adopted worldwide for water reuse applications. Ozone is effective for pathogenic inactivation and organics oxidation: both increasing assimilable organic carbon for biofiltration and eliminating trace organic contaminants which may pose a threat to human health. However, ozone can also form disinfection byproducts such as bromate from the oxidation of naturally occurring anion bromide. Bromate is a known human carcinogen and is regulated by the EU, WHO, and USEPA to a maximum limit of 10µg/L. In waters high in bromide, especially above 100µg/L, bromate formation becomes a major concern. In the secondary wastewater effluent studied, bromide concentration may exceed 500µg/L. Several bromate suppression techniques have been devised in previous work, including free ammonia addition, monochloramination, and the chlorine-ammonia process. While free ammonia addition was not found to adequately reduce bromate formation below the required MCL, monochloramine addition and the chlorine-ammonia process were found to be effective. However, the impact of these chemical suppression techniques on organics oxidation and disinfection has not been fully studied. This study explored the impact of these bromate suppression techniques at a wide range of ozone doses on bromate formation, pathogenic inactivation, ozone-refractory organics oxidation through the surrogate 1,4-dioxane, and N-nitrosodimethylamine (NDMA) formation. Additionally, bromate suppression mechanisms of monochloramine were explored further through a variety of different water quality parameters, such as through hydroxyl radical exposure and ultraviolet absorption spectrum measurements, which were correlated and utilized to develop a hydroxyl radical exposure predictive model.

Balancing Bromate Formation, Organics Oxidation, and Pathogen Inactivation: The Impact of Bromate Suppression Techniques on Ozonation System Performance in Reuse Waters

Peter Hamilton Buehlmann

General Abstract

Ozone is a powerful oxidant used in water treatment in order to degrade contaminants of emerging concern into less harmful moieties and to inactivate pathogens. Upon application to process water, ozone quickly reacts with constituents in the water to form hydroxyl radicals: the most powerful oxidant in water treatment. These hydroxyl radicals, though with extremely short half-lives, are able to degrade ozone-recalcitrant organics, such as 1,4-dioxane through a process called advanced oxidation. Ozone itself also has the capability of inactivating a multitude of pathogenic organisms, including viruses Giardia and Cryptosporidium parvum when specific contacts times are met. However, ozone does have the potential to form disinfection byproducts such as N- nitrosodimethylamine (NDMA) and bromate. NDMA, though not currently regulated by the United States’ Environmental Protection Agency (USEPA), has a drinking water health advisory limit of 10ng/L in the State of California. Bromate, on the other hand, is a known human carcinogen regulated to 10µg/L by the USEPA. Formed within the ozone system from the naturally occurring ion bromide, bromate can be limited through various chemical treatments such as ammonia addition, pH adjustment, monochloramination, and the chlorine-ammonia process. To date, these methods of bromate suppression have not been comprehensively studied in terms of bromate suppression as well as disinfection and organics oxidation in water reuse systems. The purpose of this research was to minimize bromate formation while ensuring NDMA formation was minimized, and disinfection and organics oxidation were maximized. Through this study, system efficiencies were improved and water quality for future generations will be improved.

Acknowledgements

I would like to thank my committee: Dr. Amy Pruden, Dr. Charles Bott, and Dr. John Novak for their knowledge, support, and guidance throughout these four years of research. To the Hampton Roads Sanitation District and Dr. Charles Bott for both funding and the equipment to make all of this research possible, as well as the incredible opportunity to be involved in such an innovative and ambitious project as the Sustainable Water Initiative for Tomorrow. To everyone involved in the SWIFT pilot and research center, you have all contributed greatly to my success. There are so many of you – too many to list – but I would like to highlight the SWIFT research team, plant operators and maintenance operators, and the hard-working people at the Central Environmental Laboratory and the Technical Services Division. Lastly, I would like to thank my family and my friends from all over the world for being with me through this grand adventure.

Contents 1. Introduction ...... 1 1.1 Project Motivation and Objectives ...... 5 2. Literature Review ...... 7 2.1 Ozonation in Water Treatment ...... 7 2.2 Reactions of Ozone in Water ...... 7 2.3 Ozone Decomposition in Water ...... 8 2.4 Reactions of ozone and hydroxyl radicals with organics ...... 9 2.5 Hydroxyl Radical Measurement Methods ...... 11 2.6 Absorption spectrum of ozone-treated process water ...... 12 2.7 Disinfection ...... 12 2.8 Chlorinated and brominated disinfection byproducts (Haloacetic Acids and Trihalomethanes) ...... 13 2.9 NDMA Formation ...... 13 2.10 Bromate Formation ...... 14 2.11 Bromate Suppression ...... 15 2.11.1 pH suppression ...... 15 2.11.2 Bromate Formation Mitigation: Free Ammonia Addition ...... 16 2.11.3 Bromate Formation Mitigation: Chlorine-Ammonia Process ...... 17 2.11.4 Bromate Formation Mitigation: Monochloramination ...... 18 2.12 Effects of preoxidation on ozone processes ...... 18 2.13 Ozone control strategies ...... 19 3. Manuscript 1: Balancing bromate formation, organics oxidation, and pathogen inactivation: the impact of bromate suppression techniques on ozonation system performance in reuse waters ...... 20 3.1 Abstract ...... 20 3.2 Introduction ...... 21 3.2.1 Background of Ozone ...... 21 3.2.2 Reactions of Ozone in Water ...... 21 3.2.3 Ozone Decomposition Kinetics ...... 22 3.2.7 OH• Exposure Measurement Methods ...... 23 3.2.5 Bromate Formation ...... 24 3.2.6 Bromate Suppression Methods ...... 26

3.2.7 Hypothesis and Objectives ...... 27 3.3 Materials and Methods ...... 29 3.3.2 Preformed Monochloramines ...... 29 3.3.3 Ozone Demand Free Water ...... 30 3.3.4 Sample Water ...... 31 3.3.5 Analytical Methods ...... 32 3.3.6 Experimental Setup ...... 33 3.4 Results and Discussion ...... 34 3.4.1 Bromate Formation with Increasing Ozone Dose for Different Suppression Methods ...... 34 3.4.2 Ozone Exposure Calculations ...... 35 3.4.3 Impact of Suppression Techniques on Bromate Formation and Ozone Exposure ...... 37 3.4.4 Impact of Chemical Addition on Disinfection Credits...... 41 3.4.5 Hydroxyl Radical Exposure ...... 45 3.4.6 Impact of dose and suppression techniques on organics oxidation: UV spectral scans ...... 50 3.4.7 Impact of dose and suppression techniques on NDMA formation ...... 55 3.5 Conclusions ...... 57 References ...... 59 4. Engineering Significance ...... 62 Appendix A ...... 64

List of Tables

Table 1: Factors affecting bromate formation ...... 25 Table 2: Sample water collected from flocculation/sedimentation process effluent ...... 31 Table 3: Linear regression parameters for the percent decrease in 1,4-dioxane versus the percent decrease in UV254 absorbance ...... 53

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List of Figures

Figure 1: Bromate formation pathways. Adapted from Buffle et al, 2004 ...... 14 Figure 2: Ammonia suppression of bromate: the formation of bromamine and its' degradation via ozone. Adapted from Buffle et al., 2004 ...... 16 Figure 3: The chlorine-ammonia process, adapted from Buffle et al., 2004...... 18 Figure 4: Actions of ozone and hydroxyl radicals (OH•) upon application to process water...... 22 Figure 5: The bromate formation pathway, adapted from Buffle et al., 2004...... 25 Figure 6: Suppression of bromate formation through free ammonia addition, adapted from Buffle et al., 2004...... 27 Figure 7: The chlorine-ammonia process’s suppression of bromate. Adapted from Buffle et al., 2004. ... 28 Figure 8: Bench-scale Preformed Monochloramine System ...... 30 Figure 9: Bench-scale ozonation apparatus used for experimentation...... 33 Figure 10: Impact of ozone dose (in terms of the ozone dose to TOC ratio) versus bromate formed. Figure (b) shows bromate concentrations near the maximum contaminant level of 10 µg/L. Free ammonia addition increased bromate formation, whereas both monochloramine (MC) and the chlorine-ammonia process (Cl2-

NH3) process inhibited bromate formation...... 35 Figure 11: Illustration depicting the first order decay of aqueous ozone and C*t used for calculating disinfection credits...... 36 Figure 12: (a) Impact of increasing ozone dose on ozone exposures under different bromate suppression techniques, including (b) free ammonia addition, (c) preformed monochloramine addition, and (d) the chlorine-ammonia process with a 1 minute free-chlorine contact time prior to ammonia addition...... 38 Figure 13: Bromate formation as a function of ozone exposure when utilizing differenct bromate suppression techniques, where ozone exposure is represented as total ozone exposure (TOE) rather than a C*t value...... 40 Figure 14: C*t values calculated along an ozone decay curve. Insert shows the derivation for the time at which maximum C*t occurs (1/k)...... 41 Figure 15: Log removal values calculated from LT2SWTR for (a & b) virus, (c) Giardia, and (d)

Cryptosporidium parvum as a function of applied ozone dose in terms of O3:TOC ...... 43 Figure 16: Bromate formation as a function of Virus Log Removal, calculated using values from the LT2SWTR...... 44

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Figure 17: OH* exposures as a function of ozone dose for different bromate suppression techniques: (a) all methods, (b) free ammonia addition, (c) preformed monochloramine addition, and (d) the chlorine- ammonia process with 1 minute free chlorine contact time...... 46 Figure 18: Bromate formation as a function of hydroxyl radical exposure ...... 47

Figure 19: Impact of monochloramine concentrations on hydroxyl radical exposures for various O3:TOC ratios in (a) preformed monochloramine addition and (b) the chlorine-ammonia process...... 48

Figure 20: (a) RCT (ratio of hydroxyl radical exposure to ozone exposure) as a function of ozone dose for the bromate suppression techniques studied: (b) free ammonia addition, (c) preformed monochloramine addition, and (d) the chlorine-ammonia process with a one minute free chlorine contact time...... 49 Figure 21: Changes in UV absorbance at 254nm as compared to the background sample as a function of ozone dose in mg O3:mg TOC for a variety of chemical suppression techniques, including (a) all experiments, (b) free ammonia addition (NH4), (c) preformed monochloramine addition (NH2Cl), and (d) the chlorine-ammonia process (Cl2-NH4) with a one minute free chlorine contact time ...... 51 Figure 22: Percent decreases in 1,4-dioxane concentrations versus the percent decreases in UV254 absorbance for different bromate suppression techniques, subjected to various ozone doses...... 52 Figure 23: (a) Validation of the OH˙ exposure model using ∆UV254 when compared to measured OH˙ exposures through 1,4-dioxane. (b) The impact of bromate suppression technique on the accuracy of the OH˙ exposure model ...... 54 Figure 24: (a) Impact of various ozone doses on NDMA formation for the various bromate suppression techniques studied, including (b) free ammonia addition, (c) preformed monochloramine addition, and (d) the chlorine-ammonia process ...... 55 Figure 25: Absorbance spectrum of the monochloramine stock solution used for 1mg/L preformed monochloramine testing...... 64 Figure 26: Absorbance spectrum of the monochloramine stock solution used for 3mg/L preformed monochloramine testing...... 64 Figure 27: Absorbance spectrum of the monochloramine stock solution used for 5mg/L preformed monochloramine testing...... 65 Figure 28: Absorbance ozonated water samples from 240nm to 330nm for the different chemical suppression techniques studied. (a) no bromate suppression, (b) 0.4mg/L-N free ammonia, (c) 0.8mg/L-N free ammonia, (d) 1mg/L preformed monochloramine, (e) 3mg/L preformed monochloramine, (f) 5mg/L preformed monochloramine, (g) 1mg/L chlorine-ammonia process, (h) 3mg/L chlorine-ammonia process, and (i) 5mg/L chlorine-ammonia process...... 66

1. Introduction

As the demand for water has increased in the past half century, so too has the bid for more sustainable, cost efficient, and reliable treatment technologies. Thus, there has been an emergence of non-potable, indirect and direct potable reuse from treated wastewater sources. While non- potable reuse has had applications for many years in the agricultural industry, potable reuse has gained recent attention due to growing urban centers in coastal regions and diminishing freshwater supplies. California’s water reclamation programs, for instance, have more than tripled their production since the 1970’s.

Multiple water-related issues are troubling the Hampton Roads region of eastern Virginia. The Chesapeake Bay, with a fragile ecosystem of oysters, is very susceptible to algal blooms. Therefore, wastewater treatment facilities discharging into this water body are facing much stricter nutrient permits in the near future. Simultaneously, the Potomac aquifer, stretching from North Carolina to Maryland, is in a multiple-trillion-gallon deficit due to overuse in the past century. This is directly causing land subsidence and saltwater intrusion into the aquifer. Additionally, groundwater scarcity may cause a reduction in withdrawal limits for industrial users within the next century, causing significant economic harm to the mid-Atlantic region. Hampton Roads Sanitation District’s Sustainable Water Initiative for Tomorrow (SWIFT) aims to address these issues simultaneously by treating wastewater effluent to drinking water standards for the purposes of groundwater augmentation, thus eliminating nutrient discharge into the Chesapeake Bay.

Traditionally, indirect and direct potable water reuse has centered on the concept of “full advanced treatment” (FAT), consisting of multiple membrane-based treatment processes. These treatment processes include a reverse osmosis system, associated with high capital, operational, and maintenance costs. Additionally, this method of treatment creates a concentrated stream of dissolved minerals, salts, and pollutants. This brine stream is exceedingly difficult to properly dispose at inland facilities, and the pollutants in wastewater effluent are concentrated in this stream rather than being removed or destroyed. Lastly, the quality of water produced by reverse osmosis is often incompatible with natural soil chemistry, requiring further treatment for groundwater augmentation.

An emerging alternative to this treatment method is the “carbon-based” approach. This treatment train centers on ozone (O3) and biologically active carbon (BAC) filtration steps. These processes provide multiple barriers for contaminants. Pathogens, micropollutants and bulk total organic carbon (TOC) are inactivated or degraded chemically within the ozone system and physically or biologically within the BAC.

Ozone, when dissolved into an aqueous system, has two primary routes of oxidation: direct ozonation with aqueous ozone gas, and advanced oxidation with hydroxyl radicals (OH•). These hydroxyl radicals form relatively slowly from the autocatalytic degradation of ozone, but rather quickly in waters with constituents such as dissolved organics or metals. These OH• are necessary for the degradation of ozone-refractory contaminants such as 1,4-dioxane but are not primarily responsible for biological inactivation and bulk organic oxidation. Due to the increased organics content in wastewater applications, the OH• exposures in wastewater effluents during ozonation may be orders of magnitude higher than in drinking water applications and are often considered inherent advanced oxidation processes (AOPs).

While ozonation is an excellent treatment technology for both the inactivation of pathogens and the advanced oxidation of organic constituents, it does include a drawback of having the potential of forming bromate, a suspected human carcinogen. Due to this risk, the World Health Organization (WHO), United States Environmental Protection Agency (USEPA), and others have adopted a 10ug/L maximum contaminant level (MCL) for drinking water. Bromate formation is of primary concern in waters containing high background bromide concentrations.

To curtail the formation of this carcinogen, operational changes must be made in order to limit bromate formation. The simplest method to reduce bromate formation is to reduce ozone dose, thus limiting O3 and OH• exposures. This will lead to less organic oxidation and pathogen inactivation, however. Ozone dose can be controlled and described in many ways, with applied dose, or the amount of ozone applied to the water in mg/L. This value does not account for the ozone not dissolved in the process water. Transferred ozone dose (TOD) is equal to the product of transfer efficiency and applied ozone dose.

TOD is an absolute value of dose which does not factor in water quality parameters. Due to the impact of organic carbon on ozone reaction kinetics and the variety of organic carbon content between municipalities, ozone dose can be normalized to TOC in the form of an ozone to TOC ratio (O3:TOC). This parameter is useful in comparing ozone doses between different water qualities but does not capture the reactivity of the organic constituents. An indirect method for controlling ozone dose based on ozone reaction kinetics with organic matter is through residual control. This technique measures aqueous ozone residual at a specific contact time within an ozonation system, factoring in some decay kinetics. Another method to control ozone through indirect organic oxidation measurement is through the reduction in UV absorbance at 254nm (UV254). Though useful, these methods cannot fully elucidate ozone reaction kinetics, and more advanced methods for evaluating ozone control are required for optimal system performance. Controlling and classifying an ozonation system via ozone and OH• exposures would be ideal, as full reaction kinetics would be known and may be modeled.

Multiple chemical suppression techniques have also been developed to maintain drinking water quality. The oldest practiced method is pH suppression, which shifts the equilibrium of hypobromite (OBr-) to hypobromous acid (HOBr), a key intermediate in the bromate formation pathway. HOBr has a considerably slower reaction rate with ozone than OBr-, oftentimes regarded as insignificant during the timeframe of ozone contactors. Whereas this technique is considered the best available treatment technology, high alkalinity water can render this alternative cost prohibitive. Therefore, other methods have been developed such as free ammonia addition, which acts to suppress HOBr as a bromamine. This compound forms in a manner similar to its chloramine counterpart and is an intermediate product which does not directly react into bromate. This strategy is limited in effectiveness as several characteristics must be met to achieve adequate suppression: pH must be relatively neutral or acidic to prevent OBr- from forming bromate in lieu of bromamine, - and bromide must be oxidized by either OH• or O3 to HOBr/OBr prior to amine formation, increasing the probability of a bromine molecule “missing” the intermediate HOBr/OBr- step and forming bromate. The chlorine-ammonia process aims to preoxidize bromide into hypobromous acid, followed by quenching free chlorine residual with free ammonia, forming both bromamine and chloramine. Theoretically, this minimizes the probability of bromine molecules shortcutting the pathway to bromate, as ozone and OH• are absent during bromamine formation.

Monochloramine, a byproduct of this technique, was found to have excellent bromate suppression capabilities as well, though the exact mechanism is not known. It is hypothesized that monochloramine scavenges OH•, thus slowing the formation of bromate through the OH•-driven pathway. Another possibility is monochloramine is creating either bromamine or an as of yet unknown intermediate product, similar in function to bromamine masking. The chlorine-ammonia process also has the added benefit of preoxidizing organics in a water matrix, reducing instantaneous ozone demand (IOD) and the OH• exposure associated with this phase. However, this may lead to decreases in a system’s advanced oxidative capabilities.

Along with bromate formation, N-nitrosodimethylamine (NDMA) formation is of chief concern. NDMA is currently has a health advisory limit of 10ng/L in California. NDMA is traditionally of concern during the chloramination of drinking water, but NDMA is also generated within the ozonation process. When coupling chloramination and ozonation, NDMA becomes an obvious concern. Not much research has been conducted on NDMA formation with this combination in wastewater sources. NDMA may be a potential drawback to the implementation of chloramine and chlorine-ammonia based bromate suppression techniques and needs further exploration.

1.1 Project Motivation and Objectives

In the Hampton Roads region of eastern Virginia, water scarcity is a hidden, yet imminent, threat to the region’s economic stability. Over the past two centuries, a multiple-trillion-gallon water deficit has accumulated within the Potomac Aquifer, which stretches from North Carolina to Maryland, resulting from overconsumption. This issue has led to issues including land subsidence and saltwater intrusion. Hampton Roads Sanitation District’s (HRSD) Sustainable Water Initiative for Tomorrow (SWIFT) aims to alleviate this negative trend by augmenting the Potomac Aquifer with 100 MGD with tertiary-treated wastewater effluent. The advanced treatment train, consisting of flocculation/sedimentation, ozonation, biofiltration, granular activated carbon, and ultraviolet disinfection, will ensure effluent parameters meeting or exceeding drinking water standards.

Pilot-scale testing began at HRSD’s York River Treatment Plant (YRTP) in June 2016, consisting of the carbon-based treatment train, as well as a membrane-based treatment train comprising ultrafiltration, reverse osmosis, and UV-advanced oxidation with peroxidation. After 6 months of continuous operation, the carbon-based treatment system was selected for full-scale implementation, and a 1 million gallon per day (MGD) demonstration facility was constructed.

Central to this process, the ozonation step serves to oxidize organics to increase carbon assimilability to be picked up by the BAC, degrade trace organic constituents into less harmful moieties, as well as to serve as a disinfection step in a multi-barrier process. Prior to startup of the pilot-scale treatment system, bromate formation was of great concern due to high influent bromide concentrations (300-600µg/L). This concern was realized within the first few weeks of pilot operation, with bromate concentrations often exceeding 15x the maximum contaminant level. Free ammonia addition (0.3mg/L), coupled with ozone dose limitations, was able to reduce bromate formation to 50µg/L, though still five times the maximum contaminant level (MCL) of 10µg/L. pH suppression was not considered due to cost limitations associated with a high expected acid demand in secondary effluent with elevated alkalinity (120mg/L as CaCO3).

Alternative mitigation technologies were explored, including the chlorine-ammonia process and chloramination. Monochloramination formation through the ammonia-chlorine process was

5 attempted, but a stable residual of chloramine was not attainable at the pilot scale. Additionally, in-situ formed monochloramine and chlorine-ammonia processes were initially avoided due to concerns of NDMA and disinfection by-product (DBP) formation, but further experimentation is required to either confirm or deny these hypotheses. Preformed monochloramine was selected for bromate mitigation due to the ability to adequately maintain safe bromate levels while still allowing for high enough ozone doses for proper disinfection.

The main objective of this project was to explore the efficacy of bromate suppression techniques in high bromide source waters subjected to an ozonation process. This information is vital to the further development of SWIFT facilities throughout the region, where bromide concentrations can be found as high as 1,300µg/L. Additionally, parameters such as organics degradation and pathogen inactivation were investigated to ensure maximum system efficiency. By maximizing organics degradation, granular activated carbon (GAC) reactivation frequencies could be reduced, saving cost. Maximizing pathogen inactivation through the ozonation system contributes heavily to the philosophy of a multi-barrier treatment process and can eliminate the need for further disinfection downstream.

2. Literature Review

2.1 Ozonation in Water Treatment

Ozone has been widely used in drinking water disinfection since the 1970’s and has become increasingly prevalent in the field of water reuse for disinfection and oxidation. Over the past two decades, this process has seen a renaissance as an oxidation step for the purposes of increasing organic assimilability in biologically active filtration and degrading contaminants of emerging concern into less harmful moieties (Volk and Lechevallier 2002). Ozonation has multiple advantages when compared with other advanced oxidation processes for these applications including low cost, replicability, and limited production of chlorinated and brominated disinfection by products (DBPs). Bromate is a key DBP unique to the ozonation system with an EPA, EU and WHO MCL of 10µg/L. For this reason, it is imperative to balance ozone oxidation and disinfection with bromate formation.

2.2 Reactions of Ozone in Water

Upon application to process water, aqueous ozone reacts to form products such as hydroxyl radicals (OH•) through autocatalytic decay or through reactions with organic and inorganic constituents in the water matrix. OH• are short lived molecules with an extremely high oxidative reactivity, oxidizing compounds quickly (k ≈ 108-10 M-1s-1) (Rosario-Ortiz et al. 2008). OH•, the most reactive species in water and wastewater treatment, is the primary actor in advanced oxidation processes (AOPs). Direct reactions between ozone and organics allow for large increases in biodegradable organic matter (BOM) and can generate large amounts of OH•. These reactions mostly occur within the first few seconds of the process and this phase is referred to as instantaneous ozone decay (IOD). Generating radicals through ozone-organic matter reactions (OH• initiation) during IOD in turn allow for a greater number of organics to be oxidized by OH•, some of which are ozone-refractory. These organics can either generate further radicals (OH• propagation) or end the radical chain reactions (OH• termination) through OH• scavenging. Following fast initial reactions, ozone degradation rates slow significantly to a first-order decay

7 rate on the order of minutes to hours. Ozone decay in water treatment is often regarded as “biphasic” in nature due to these two phases.

2.3 Ozone Decomposition in Water

During the initial decomposition of ozone, electron rich organic moieties (ERMs) can react very quickly with aqueous ozone. Dissolved natural organic matter (DNOM) can react with ozone yielding two primary products: an oxidized DNOM molecule alone, or a radicalized DNOM - molecule with ozonide radical (O3 •) as described in Section 2.4. These ozonide radials lead to the decomposition of O3 into OH•. Because of these reactions, O3 residuals during the initial phase decrease rapidly, especially in wastewaters when compared to natural waters. Understanding the trend of this decay is paramount to comprehending oxidation reactions during this stage, as both

O3 and OH• concentrations change rapidly. While decomposition within IOD may at first appear to fit a linear trend, sources have found that both ozone and OH• exposures decrease exponentially. When ozone decay within the first phase is modelled with first order regression, as described in Buffle et al., 2006, the resulting decay constant versus reaction time best fits a power function. This can be attributed to fast second order reactions having a large impact on the pseudo-first order model. It can also be stated that, in wastewater effluents, there is no clear distinction between IOD and phase 2 (Buffle et al. 2006b).

Due to OH• proliferation reactions during IOD, ozonation is often considered an inherent AOP. Quantifying these transients OH• concentrations is important for determining oxidation characteristics. By integrating OH• concentrations with respect to time, an exposure value may be calculated. This value is crucial for calculations relating to organics degradation. It is also important to factor in the primary oxidant’s effect: the exposure of ozone. As depicted in Equation

1, the removal of a compound P is affected by the exposures of both OH• and O3 as well as the respective reaction rates. Relative exposures shift as the reaction progresses and the impact of each of the oxidant’s changes. To track these variations, RCT may be used. RCT is the ratio of OH• exposure to O3 exposure as expressed in equation 2 (Elovitz and Von Gunten 1999). RCT is not a constant value: during IOD, it is a described as being an order of magnitude higher than during first order decay. RCT is also a valuable tool for calculating the relative oxidation of a given

8 compound by either O3 or OH•. Knowing the RCT value, the fraction of oxidation by OH• may be calculated using Equation 3, simplified in Equation 4 (Elovitz and Von Gunten 1999).

푃 ln [ ] = −{퐾푂퐻• ∗ ∫[푂퐻 •]푑푡} + {퐾푂3 ∗ ∫[푂3]푑푡} (Eq. 1) 푃0 ∫[푂퐻•]푑푡 푅퐶푇 = (Eq. 2) ∫[푂3]푑푡

∆[푃]푂퐻• 푘푂퐻•[푂퐻•][푃] 푓푂퐻• = = (Eq. 3) ∆[푃]푡표푡푎푙 푘푂퐻•[푂퐻•][푃]+푘푂3[푂3][푃]

푘푂퐻•푅퐶푇 푓푂퐻• = (Eq. 4) 푘푂퐻•푅퐶푇+퐾푂3

During phase 2 of ozone decay, the proliferation of OH• decreases and ozone decay stabilizes. This is a consequence of the fastest ERM-based reactions moving to completion. Less OH• are generated from initiation and propagation reactions as a result. RCT during this second phase decreases and is generally considered to be constant throughout in natural waters (literature value -8 for drinking water: RCT ~10 ). The auto-catalytic decomposition begins to dominate as the main reason for ozone decay in natural waters, as expressed in Section 2.4. Wastewater sources, on the other hand, continue to be heavily influenced by the water’s constituents (Buffle et al. 2006b). Decay rates in wastewaters remain much higher than in natural waters, while in both sources the decay rate of ozone decreases over reaction time. It should also be noted that changes in the transferred ozone dose will yield different observed first order decay rate constants for the same water quality during phase 2 (Buffle et al. 2004, 2006b; a; Elovitz et al. 2008; Elovitz and Von Gunten 1999). This is an artifact of second order reactions affecting the pseudo-first order regression. However, these differences in decay rates are not currently well elucidated mathematically. It is during the second phase that the majority of bulk organics oxidation and pathogen inactivation occur, as well as bromate formation.

2.4 Reactions of ozone and hydroxyl radicals with organics

As mentioned in Section 2.3, O3 and OH• reactions with organic matter in wastewater sources heavily influence reaction kinetics. Upon application to the process water, ozone can take many

9 different reaction pathways. The simplest reaction (Reaction 1) directly oxidizes a NOM molecule with an oxygen molecule as a product. Another pathway results in the radicalization of both the

NOM and O3, as depicted in Reaction 2. The resulting ozonide radical goes through a series of reactions in the superoxide radical pathway, resulting in OH• initiation (Reactions 3-5). The radicalized NOM molecule can then either terminate (Reaction 6) or propagate more radicals through the same superoxide pathway (Reaction 7) (von Gunten 2003a; Staehelln and Hoigné 1985).

푂3 + 푁푂푀1 → 푁푂푀1푂푋 (R. 1) + − 푂3 + 푁푂푀2 → 푁푂푀2 • +푂3 • (R. 2) − − 9 −1 −1 푂3 + 푂2 •→ 푂3 • +푂2 푘 = 1.6 ∗ 10 푀 푠 (R. 3)

− + + 10 −1 −1 푂3 • +퐻 ↔ 퐻푂3 • (푝퐻 < ~8) 푘 = 5 ∗ 10 푀 푠 (R. 4a)

푘− = 3.3 ∗ 102푠−1 (R. 4b)

5 −1 퐻푂3 •→ 푂퐻 • +푂2 푘 = 1.4 ∗ 10 푠 (R. 5)

− 푁푂푀4 • +푂2 → 푁푂푀4˗푂2 •→ 푛표 푂2 • 푓표푟푚푎푡푖표푛 (R. 6)

+ − 푁푂푀3 • +푂2 → 푁푂푀3˗푂2 •→ 푁푂푀3 + 푂2 • (R. 7)

Propagation reactions mainly occur through the production of carbon centered radicals, as depicted in Reactions 8 and 9, with methanol as an example (Staehelln and Hoigné 1985).

− 8 −1 −1 퐻2퐶푂퐻 + 푂퐻 •→ 퐻2푂 + 퐶푂2 • 푘 = 7 ∗ 10 푀 푠 (R. 8)

− − 퐶푂2 • +푂2 → 퐶푂2 + 푂2 • 푘 = 푓푎푠푡 (R. 9)

OH•, the product of many of the reactions, in turn reacts with NOM as well. From initiation reactions, OH• reacts with NOM to produce a radical NOM moiety (Reactions 10 and 11). These reactions do not directly generate more OH•, but the NOM• can either follow Reaction 6 or 7. In the latter sense, a propagation reaction occurs where more OH• are indirectly generated.

Termination reactions, in this sense, could then be defined as interactions with NOM• moieties which do not produce the ozonide or superoxide radical species (von Gunten 2003a).

− 푂퐻 • +푁푂푀3 → 푁푂푀3 • +퐻2푂 표푟 푁푂푀3 • +푂퐻 (R. 10)

푂퐻 • +푁푂푀4 → 푁푂푀4 • +퐻2푂 (R. 11)

2.5 Hydroxyl Radical Measurement Methods

Measuring the concentration and exposure of OH• is critical to the understanding of organics oxidation as well as the formation of bromate. However, due to their extremely low concentrations -8 (RCT ≈ 10 ) and short half-lives, direct measurement is effectively impossible. Therefore, surrogate compounds can be used as a probe for OH• exposure. Ideally, these probe compounds 8 -1 -1 have a very fast reaction rate with OH• (kOH• > 10 M s ) but a very low reactivity with O3. Using Equation 1, the degradation of these probe compounds can be used to calculate OH• exposure.

Traditionally, para-chlorobenzoic acid (pCBA) has been used as a OH• probe compound in ozonated systems. pCBA is an ideal probe compound due to its low reactivity towards O3 (푘 푂3 ≤ 푝퐶퐵퐴 −1 −1 9 −1 −1 0.15 푀 푠 ) but a high reactivity towards OH (푘 푂퐻• = 5 × 10 푀 푠 ). However, method • ( ) 푝퐶퐵퐴 requirements such as high-pressure liquid chromatography (HPLC) and reagent cost prevent method usage for many municipalities.

In other advanced oxidation systems, such as those with ultra-violet irradiation and hydrogen peroxide, 1,4-dioxane has been used as a OH• probe compound. 1,4-Dioxane has similar reaction rates in ozone processes as pCBA: a very low reaction rate with O3, but a high reactivity with OH• (Chitra et al. 2012; Tian et al. 2014).

2.6 Absorption spectrum of ozone-treated process water

Work by Gerrity (2012) and others has shown that OH• exposure can be well-correlated to changes in UV absorbance at 254nm using a pCBA as a OH• probe compound. It has been predicted that this correlation may be applicable to many different water and wastewater municipalities as a means of continuous on-line estimation of OH• exposure. Through this estimation, trace organic oxidation can also be estimated with known reaction rates with both OH• and ozone. However, it should be noted that compounds with high reactivity with ozone (k>104M-1s-1) interfere with this method, as their oxidation and subsequent absorbance change is dependent on ozone exposure rather than OH• exposure (Buffle et al. 2006b; Gerrity et al. 2012).

2.7 Disinfection

Disinfection within the ozonation system is dependent on ozone exposure alone, and not exposure from OH•. While OH• reactions do result in cellular and viral inactivation, their relative concentrations within the treatment process are far too low to allow for any meaningful disinfection. Disinfection credits are granted to utilities for maintaining ozone residuals at a specified residence time. The product of these parameters is ozone CT, an ozone exposure value with units of mg/L*min. According to the USEPA’s Long Term 2 Enhanced Surface Water Treatment Rule (LT2SWTR), log removal value credits (LRV) can be achieved with the following formulas (US EPA 2006):

푉푖푟푢푠 퐿표푔 퐶푟푒푑푖푡 = 2.1744 ∗ (1.0726푡푒푚푝) ∗ 퐶푇 (Eq. 5) 퐺푖푎푟푑푖푎 퐿표푔 퐶푟푒푑푖푡 = 1.0380 ∗ (1.0741푡푒푚푝) ∗ 퐶푇 (Eq. 6) 퐶푟푦푝푡표푠푝표푟푖푑푖푢푚 퐿표푔 푐푟푒푑푖푡 = 0.0397 ∗ (1.097576푡푒푚푝) ∗ 퐶푇 (Eq. 7)

Of the three pathogens listed, attaining credits for viruses is the easiest. Giardia removal is the next easiest and approximately 1 Giardia LRV is granted for every 2 virus LRV. Cryptosporidium, on the other hand, is orders of magnitude lower with respect to virus and Giardia inactivation due to their cystic morphology. Additionally, temperature is an important factor in attaining pathogen inactivation. As temperature lowers, reactivity of ozone with pathogen decreases (US EPA 2006).

2.8 Chlorinated and brominated disinfection byproducts (Haloacetic Acids and Trihalomethanes)

Chlorination is often utilized as a disinfection step in water treatment and serves as a residual disinfectant within distribution systems to prevent pathogenic regrowth. However, free chlorine reacts with organic constituents (namely humic and fulvic acids) to form unwanted disinfection byproducts. Additionally, chlorine oxidizes bromide to HOBr/OBr-, which can then react with the dissolved organic matter (DNOM). Through these oxidation steps, chlorinated and brominated disinfection byproducts are formed. In the ozonation process, bromide is also oxidized to HOBr/OBr- and can form brominated disinfection byproducts. These organic disinfection byproducts may act as bromine sinks during ozonation, preventing the element from forming the carcinogen bromate (Buffle et al. 2006a). Additionally, these organic-based DPBs can often be subsequently destroyed due to ozone’s oxidative capabilities. It is for this reason that ozone is occasionally used as a preoxidation step to limit chlorine-generated disinfection byproducts.

The United States’ Environmental Protection Agency (USEPA) has drinking water maximum contaminant levels (MCL) for total trihalomethanes (TTHMs) (chloroform, bromodichloromethane, dibromochloromethane, and bromoform) and a group of 5 haloacetic acids (HAA5) (monochloroacetic acid, dichloroacetic acid, trichloroacetic acid, bromoacetic acid, and dibromoacetic acid). TTHMs and HAA5 are regulated to 80 and 60 µg/L, respectively (United States Environmental Protection Agency 2006).

2.9 NDMA Formation

N-Nitrosodimethylamine (NDMA), part of the larger group of nitrosamines, is classified as a B2 carcinogen by the USEPA. This compound is formed during many different steps of the drinking water treatment process, particularly chlorination, chloramination, and ozonation. Though unregulated at the federal level, the state of California has set a health advisory limit of 10ng/L (California State Water Resources Control Board 2018).

Organic precursors are responsible for the formation of NDMA. These precursors include quaternary and tertiary amines. Some notable examples of these compounds may be found in

13 coagulants and polymers used in water treatment, such as polyacrylamide and poly- diallyldimethylammonium chloride (polyDADMAC). These compounds have been shown to react with chloramines within the water to form NDMA though a complex pathway. Within ozonation systems, dimethylamine (DMA) and dimethylsulfamide (DMS) have been shown to be important precursors, but a multitude of other complex amines, hydrazines, semicarbazides, sulfamides, and tertiary amines with dimethylamine functional groups also contribute to NDMA formation within waters with a complex water matrix, such as wastewater effluents.

While ozone forms NDMA, exposure to OH• within the ozone system degrades NDMA directly (von Gunten 2003b; Marti et al. 2015; Pisarenko et al. 2012). This leads to a balance of both formation and destruction within the system; increasing OH• exposure during the ozone process can lower both effluent NDMA concentrations as well as NDMA formation potential. NDMA can also be removed through downstream processes such as biofiltration and UV irradiation.

2.10 Bromate Formation

Bromide, a naturally occurring anion, is non-toxic and does not have any drinking water regulation. However, its presence in ozonation treatment processes can be troublesome due to its oxidized counterpart: bromate. Bromate is a suspected human carcinogen, regulated to 10µg/L by the WHO, EPA and EU, following evidence of the formation of renal tumors in mice and rats (Kurokawa et al. 1986). According to Buffle 2004, bromate can be formed during the ozonation process through two primary reaction mechanisms: directly oxidized by ozone, and indirectly oxidized by radical reactions as illustrated in Figure 1 (Buffle et al. 2004).

Figure 1: Bromate formation pathways. Adapted from Buffle et al, 2004

The first step of bromide’s transformation to bromate can follow either the direct or indirect pathway. The direct pathway forms hypobromous acid/hypobromite (HOBr/OBr-) through a - - reaction between O3 and Br . The indirect pathway is a reaction of Br with OH• to form bromine radicals (Br•), which quickly react through a sequence of reactions with bromide to form HOBr/OBr-. Alternatively, Br• can be further oxidized to BrO• by ozone. This compound will

- - - typically disproportionate into HOBr/OBr and BrO2 ; HOBr/OBr can be re-oxidized by OH• back into BrO•. Based on the above reactions, the product OBr- and its conjugate acid HOBr can be - - considered key intermediates in this pathway. HOBr/OBr can react with ozone to form BrO2 , the - final step prior to forming bromate. However, while OBr reacts relatively quickly with O3 to form - -1 -1 BrO2 , HOBr reacts slow enough for the reaction to be discounted (k=0.01 M s ).

Radical driven reactions in the bromate formation pathway have very fast reaction rates (k>108 M- 1 -1 s ). Conversely, direct reactions with O3 are relatively slow. Though OH• exposures are many

-8 orders of magnitude lower (RCT ≈ 10 ), the high reaction rates make them a critical route in the bromate formation pathway, especially in wastewater sources.

2.11 Bromate Suppression

According to Von Gunten (2003), bromate formation becomes a concern in ozonation processes where bromide concentrations exceed approximately 50µg/L. If bromate is problematic at these higher bromide concentrations, a combination of both operational changes and chemical mitigation strategies can be implemented to help meet the MCL (von Gunten 2003a).

2.11.1 pH suppression

Suppression of pH in ozonation systems is often considered the best available treatment. This suppression technique works in two ways: by shifting the equilibrium of hypobromous acid and hypobromite, and by reducing the degradation of ozone to OH• through the hydroxide ion pathway. - As depicted in Reaction 12 below, the pKa of HOBr/OBr is 8.8. As pH decreases, the ratio of

HOBr to OBr- increases. This reduces bromate formation as HOBr has a much slower reaction rate with ozone than OBr- (100M-1s-1 versus 0.01M-1s-1) (Buffle et al. 2004).

퐻푂퐵푟 ↔ 푂퐵푟− + 퐻+ pKa: 8.8 (R. 12)

While this treatment option is often considered best practice, waters high in alkalinity may see this method as cost prohibitive. Therefore, other suppression techniques are often sought (Buffle et al. 2004).

2.11.2 Bromate Formation Mitigation: Free Ammonia Addition

Free ammonia addition has been used to successfully reduce bromate formation, oftentimes coupled with pH suppression to yield up to a 50% reduction in bromate formation. Ammonia addition works through a direct reaction with HOBr to form bromamine. This bromamine is an intermediate in the bromate formation pathway which does not directly go on to form bromate: rather, it can be further oxidized by ozone to form nitrate and a reduced Br- ion, albeit slowly (k=40M-1s-1) (Buffle et al. 2004).

Figure 2: Ammonia suppression of bromate: the formation of bromamine and its' degradation via ozone. Adapted from Buffle et al., 2004

7 -1 -1 Though the reaction from HOBr to NH2Br is very fast (7.5*10 M s ), the efficacy of this method is limited. It has been suggested in literature that free ammonia’s suppression of bromate has diminishing returns above about 0.3mg/L (Hofmann and Andrews 2007), and evidence of diminishing returns can be seen in µg/L concentration ranges (Pinkernell and Von Gunten 2001). In waters with high bromide concentrations, free ammonia addition may not be able to fully suppress bromate formation below the required 10µg/L MCL (Buffle et al. 2004; Hofmann and Andrews 2007).

2.11.3 Bromate Formation Mitigation: Chlorine-Ammonia Process

One limitation of free ammonia addition is the need for HOBr to react with ammonia prior to further oxidation by ozone to bromate. To mitigate this issue, the chlorine-ammonia process was devised to preoxidize bromide into HOBr prior to the ozone process, reducing or eliminating the probability of HOBr oxidation by ozone (Buffle et al. 2004).

This process works by first adding free chlorine (in the form of aqueous hypochlorite) to the process water, followed by a set residence time. This residence time can vary from a short 30 seconds to 5 minutes: the former of which aims to minimize disinfection byproduct formation and the latter which aims to have bromide and organic oxidation reactions move towards completion.

Following this reaction time, free ammonia is added to react with the HOBr to form NH2Br, as well as to form monochloramine with the remaining free chlorine species (Buffle et al. 2004).

Figure 3: The chlorine-ammonia process, adapted from Buffle et al., 2004

2.11.4 Bromate Formation Mitigation: Monochloramination

A byproduct of the chlorine-ammonia process, monochloramine was found to have excellent bromate suppression capabilities. Previous works have cited more than 70% reduction in bromate formation (Buffle et al. 2004). However, the exact mechanism through which monochloramine suppresses bromate formation is still unknown. It has been hypothesized that monochloramine acts as a OH• scavenger, thereby reducing the magnitude of the indirect bromate formation pathways. It has also been suggested that monochloramine may be forming an as of yet unidentified intermediate product, removing bromine from the main bromate formation pathway through the duration of the ozonation process (Buffle et al. 2004).

2.12 Effects of preoxidation on ozone processes

The chlorine-ammonia process also boasts the benefit of preoxidizing process water with free chlorine. This helps to reduce ozone demand by oxidizing organic moieties which would contribute to ozone decay. This reduced ozone demand allows for smaller ozone doses to achieve

18 the same disinfection and oxidation characteristics as a system without preoxidation but with higher ozone doses. Reducing ozone dose also allows for less bromine oxidation into bromate.

2.13 Ozone control strategies

Controlling ozone dose is a vital aspect of any ozonation process and is critical to ensure excess bromate is not formed from overdosing while adequate oxidation and disinfection still occur. The most simplistic form of ozone dose control is through monitoring applied or transferred ozone doses: the amount of molecular ozone delivered to the process water. This method, however, does not account for changes in water chemistry, such as between surface water and wastewater sources.

Therefore, a ratio between ozone and total organic carbon (O3:TOC) is often implemented to characterize ozone dose regardless of organic loadings within a system. It should be noted that - nitrite, which exerts an immediate and high demand of ozone (3.43mg O3:1mg NO2 ) is discounted from this ratio.

O3:TOC, while factoring in relative expected demand of a water matrix, does not account for the reactivity of the water’s components themselves. Therefore, an alternative to ozone dose is monitoring of residual ozone concentrations at a set time to determine proper ozone doses. This method accounts for any variation in instantaneous ozone demand and ozone decay rate and allows for a consistent residual to be maintained for proper disinfection credit.

3. Manuscript 1: Balancing bromate formation, organics oxidation, and pathogen inactivation: the impact of bromate suppression techniques on ozonation system performance in reuse waters

3.1 Abstract

The ozonation process is becoming a critical step in potable reuse systems that do not rely on membranes. By providing a barrier to pathogens, increasing biodegradable carbon for downstream biofiltration, and oxidizing contaminants of emerging concern, the ozonation process can also form the disinfection byproducts, bromate and N-nitrosodimethylamine (NDMA). Bromate formation mitigation techniques including free ammonia addition, monochloramination, and the chlorine- ammonia process have been used in the past with varying degrees of success. However, the impact of these suppression methods has not been studied on NDMA formation, disinfection, organic oxidation, and bromate formation simultaneously. This study found that both preformed monochloramination and the chlorine-ammonia process were effective in controlling bromate formation, whereas free ammonia addition was found to be ineffective. The addition of preformed monochloramine and the chlorine-ammonia process was able to reduce bromate formation by up to 80%. Additionally, the chlorine-ammonia process was able to increase the attainable disinfection credits for the system by reducing ozone demand through preoxidation, reducing the required ozone dose for treatment objectives by 50%. NDMA was found to plateau after a residual ozone concentration was detected. Additionally, the chlorine-ammonia process formed slightly less NDMA than other suppression techniques, suggesting the preoxidation of precursor compounds prior to ozonation and decreasing NDMA formation potential. In terms of bromate formation, organics oxidation, and pathogenic inactivation, the chlorine-ammonia process was found to be the optimal treatment technique if implemented properly.

3.2 Introduction

3.2.1 Background of Ozone

Ozone was introduced in drinking water treatment as a disinfectant in the 1970’s. This treatment process has also proven itself in the field of water reuse as a disinfection and oxidation step. While providing a formidable barrier for pathogens, ozonation also effectively converts refractory organic carbon to assimilable, ideal for systems with downstream biofiltration (von Gunten 2003a). Ozonation has multiple advantages when compared with other advanced oxidation processes for these applications including low cost, replicability, and limited production of disinfection by-products (DBPs). Bromate, however, is a key DBP unique to the ozonation system and has an EPA, EU and WHO MCL of 10µg/L (U.S. Environmental Protection Agency 1998). For this reason, it is imperative to balance ozone oxidation and disinfection with bromate formation for maximum system efficiency.

3.2.2 Reactions of Ozone in Water

Upon application to process water, aqueous ozone decays into hydroxyl radicals (OH•) through autocatalytic decay or through reactions with organic moieties. Direct reactions between ozone and organics greatly increase the bioavailable fraction of dissolved organic carbon. Some of these reactions, typically with electron-rich moieties (ERMs) which have fast reaction rates with ozone, generate relatively large amounts of OH• (Buffle et al. 2004; von Gunten 2003a). This process of generating radicals through ozone-DOM reactions (OH• initiation) in turn allow for a greater number of organics to be oxidized by OH•, some of which are ozone-refractory. These organics can either generate further radicals (OH• propagation) or end the radical chain reactions (OH• termination) (von Gunten and Buffle 2006).

Though OH• can oxidize more organic compounds than ozone due to its higher redox potential, the short-lived existence and low concentration of OH• makes its impact on disinfection and bulk organic oxidation minimal (von Gunten and Buffle 2006). Trace organic constituents (TOrCs), such as certain pharmaceuticals and personal care products (PPCPs) and contaminants of emerging concern (CECs), often react only slowly with ozone, or may even be completely refractory. In

21 order to oxidize ozone-resistant organics such as 1,4-dioxane, advanced oxidation through OH• is required. Unfortunately, both OH• and O3 exposures result in the conversion of bromide to bromate. Figure 4 summarizes the actions of ozone and OH• in process water (Buffle et al. 2004, 2006b).

Figure 4: Actions of ozone and hydroxyl radicals (OH•) upon application to process water.

3.2.3 Ozone Decomposition Kinetics

While aqueous ozone decay is often modeled through first order regression, this approximation does not hold true for the entire reaction. Ozone has been described to have “biphasic” kinetics in water: a fast-initial decay phase followed by pseudo-first order decay. During the initial decomposition of ozone, often called instantaneous ozone demand (IOD), moieties with fast reactions with ozone (k>108 M-1s-1) cause a chain reaction resulting in radical proliferation (Buffle et al. 2004). In ozonation processes with sources high in organics, such as treated wastewater effluent, this IOD can consume more than 50% of the transferred ozone dose. During IOD, the ratio of the exposure of OH• to ozone exposure (RCT) is much higher than later in the reaction, where first-order decay dominates. Ozonation in wastewater effluents is often considered an

22 inherent advanced oxidation process (AOP) due to these elevated OH• exposures during the initial phase (Buffle et al. 2006b).

During phase 2 of ozone decay, the OH• exposure decreases, and ozone decay stabilizes. This is a result of the most reactive moieties reacting to completion with ozone and less OH• is generated as a result (von Gunten 2003a; b). RCT during this second phase decreases, and is generally -8 considered to be constant throughout (RCT ~10 ) (Buffle et al. 2004). It should be noted, however, that changes in the transferred ozone dose will yield different first order decay rate constants for the same water quality during this phase. This is due to more reactions going to completion during the first phase of ozonation. However, the differences in decay constants between phase 1 and phase 2 are not currently well interpreted mathematically, with the function of ozone decay not truly fitting a pseudo-first order regression (Buffle et al. 2006b).

3.2.7 OH• Exposure Measurement Methods

A OH•, generated during the ozonation process, is an extremely short-lived molecule which exists at very low concentrations. Consequently, OH• exposures (the product of concentration and time) in ozonated water are extremely low (~10-10 M*s) (Buffle et al. 2004). These low exposures are due to the high reactivity of OH• with compounds in the water matrix. The measurement of OH• is vital to discern reaction kinetics as these species have major impacts on both organics oxidation and bromate formation.

Due to their short half-life and reactivity, coupled with their low concentrations with respect to other oxidants present in the water, direct measurement of OH• in aqueous samples is effectively impossible. Therefore, probe compounds have been used to estimate total OH• exposures during ozonation processes. These compounds must be carefully selected, such that ozone does not directly react with them resulting in interference. Traditionally, para-chlorobenzoic acid (pCBA) has been used as a OH• probe compound (Gerrity et al. 2012; von Gunten and Buffle 2006). This is performed by measuring the amount of pCBA oxidized during ozonation. The formula for calculating OH• with pCBA oxidation may be seen in equation 8 (Gerrity et al. 2012).

[푝퐶퐵퐴] ln( ) ∫[푂퐻 •]푑푡(푀푠) = [푝퐶퐵퐴0] (Eq. 8) −5∗109

Due to method requirements such as high-pressure liquid chromatography (HPLC) and high capital cost, many municipalities do not have the capability to measure OH• exposures in ozonation systems through use of pCBA. Therefore, alternative methods are sought to increase the availability of OH• exposures. It is hypothesized that 1,4-dioxane would be a desirable alternative OH• probe compound due to the use of a GC-MS/MS Triple Quadrupole rather than HPLC, with the added benefit of gleaning contaminant abatement data. Oftentimes, 1,4-dioxane is also monitored in treatment processes in order to ensure the maintenance of water quality goals.

3.2.5 Bromate Formation

According to Buffle et al. 2004, bromate can be formed during the ozonation process through two primary reaction pathways: directly oxidized by ozone, and indirectly oxidized by radical reactions. The first step of the pathway can either be driven to Br• by hydroxyl radicals, or directly to HOBr/OBr- by ozone. Following the radical-driven pathway, Br• can be further oxidized to

- - BrO• through ozone. This compound typically reacts to form two products: OBr and BrO2 . The product OBr- and it’s conjugate acid HOBr can be considered key intermediates in this pathway: - - OBr can react with ozone to form BrO2 , the final step prior to forming bromate, while HOBr reacts slow enough for the reaction to be discounted (k=0.01 M-1s-1). The bromate formation pathway is illustrated in Figure 5 below (Buffle et al. 2004).

Figure 5: The bromate formation pathway, adapted from Buffle et al., 2004.

A variety of water quality parameters have an impact on bromate formation. Bromide, from sources such as salt/brackish water and landfill leachate, is the main consideration for bromate formation potential. Waters with elevated pH (>8) increase bromate formation due to the generation of more OH• and the shift in equilibrium of HOBr/OBr-. Temperature fluctuations also have a considerable impact on bromate formation, but whether there is a positive or negative impact depends on other water quality parameters as well as process design. NOM also has an impact on bromate formation but is specific to NOM reaction pathways. NOM may act either as a OH• terminator or initiator/propagator, as well as a sink for bromide-bromate intermediates. In waters with elevated alkalinity, OH• exposures decrease due to carbonate and bicarbonate ions acting as OH• sinks. These factors which affect bromate formation are summarized in Table 1.

Table 1: Factors affecting bromate formation

- Factor BrO3 Formation Reason Increased Bromide Increase Increased Precursor Concentration Elevated pH Increase Shifts HOBr/OBr- equilibrium; Generates more OH• Increased Temperature Increase Higher reaction rates Decrease Lower ozone exposure due to higher decay rate Higher Organics Loading Increase May initiate more OH• Decrease Decrease ozone exposure and may act as OH• sink Increased Alkalinity Decrease Acts as a OH• sink Higher ozone dose Increase Higher ozone and OH• exposure

3.2.6 Bromate Suppression Methods

Along with tighter control of ozone doses, residuals, and exposures, chemical suppression of bromate is widely practiced. All forms of bromate suppression techniques can be categorized into four separate groups, consisting of: ozone exposure limitation, pH depression, intermediate formation and subsequent masking, and hydroxyl suppression. Some commonly-practiced techniques for bromate suppression in waters with elevated bromide concentrations include free ammonia addition, monochloramine addition, and the chlorine-ammonia process (Buffle et al. 2004).

As depicted in Figure 6, ammonia can react with HOBr leading to the formation of bromamines. This intermediate product removes HOBr from the bromate formation pathway and remains relatively stable. Bromamine is oxidized slowly by ozone, forming bromide and nitrate. Bromamines can also react with organic matter in the water forming brominated organics. While these organic compounds sequester bromine throughout the rest of the ozonation process, little is known about their toxicity (Buffle et al. 2004).

Figure 6: Suppression of bromate formation through free ammonia addition, adapted from Buffle et al., 2004.

According to Buffle et al., 2004, the premise of the chlorine-ammonia process is to add free chlorine followed by free ammonia to the water source prior to ozonation in order to sequester bromine as an intermediate product, bromamine, thus slowing the formation of bromate. The free chlorine preoxidizes bromide ions into hypobromous acid to allow for a reaction between free ammonia and hypobromous acid to form bromamine. In this process, the excess free chlorine residual also reacts with the free ammonia to form monochloramine. It was noted in previous studies that having a free chlorine residual during ozonation leads to an increased formation rate of bromate, therefore excess free ammonia is critical to this process (Buffle et al. 2004). The impacts of di- and tri- chloramines are currently unknown in the bromate formation scheme.

3.2.7 Hypothesis and Objectives

The objective of this study is to determine the most effective bromate suppression technique in reuse waters in terms of bromate formation mitigation while maximizing organics oxidation and disinfection. It is hypothesized that free ammonia addition will have a negligible effect on bromate mitigation, organics oxidation and disinfection, as the mechanism through which bromate

27 inhibition acts does not affect oxidation kinetics. Preformed monochloramine is predicted to suppress bromate formation, though may have negative impacts on organics oxidation and attainable disinfection credits due to OH• suppression and direct reactions with dissolved ozone. On the other hand, the chlorine-ammonia process should also adequately suppress bromate formation, as monochloramine is formed as a byproduct. In addition to bromate suppression, disinfection should increase with similar ozone doses due to preoxidation of organics through free chlorine contact time. This preoxidation may negatively impact OH• exposure through the removal of radical promoting and initiating organics.

Figure 7: The chlorine-ammonia process’s suppression of bromate. Adapted from Buffle et al., 2004.

3.3 Materials and Methods

Type 1 water with a resistivity above 18 MΩ-cm was used to prepare all working solutions and dilutions. All standards and solutions used were analytical grade. Aqueous ozone stock solution containing ~1mM was prepared by sparging a ~13% ozone/oxygen mixture through a 15L high purity water vessel cooled to 4°C. Potassium indigo trisulfonate reagent was used to determine the aqueous ozone concentrations of the ozone stock solution and in accordance with the Standard Methods of the Examination of Water and Wastewater. 1,4-dioxane (99.8% concentration) was sourced through Fisher Scientific and diluted to a working stock concentration of 1000mg/L.

3.3.2 Preformed Monochloramines

Monochloramines for preformed experiments were created using a series of peristaltic pumps, as depicted in Figure 8. Carrier water consisted of type 1 water, diluting the monochloramine mixture to a concentration less than 2000mgCl2/L (preferably less than 1000mg/L) to reduce heat from the exothermic reaction as well as to minimize di- and trichloramine formation due to instantaneous reactions occurring prior to complete mixing. A pulsation dampener was installed to reduce the effect of pump pulsation-induce concentration gradients. pH was determined to be >10, allowing for proper monochloramine formation. The monochloramine stock solution’s absorbance spectrum was analyzed for chloramine speciation, and di- and trichloramine concentrations were deemed negligible as noted in the Appendix (Czech et al. 1961).

Figure 8: Bench-scale Preformed Monochloramine System

3.3.3 Ozone Demand Free Water

Ozone demand free water (ODF) was produced in a manner similar to chlorine demand free water (Wang et al. 2018). This was performed in order to ensure no ozone demand existed in dilution waters which may lead to inaccuracies in testing. Type 1 high purity water was first subjected to high concentration gaseous ozone in oxygen (~13%) for more than eight hours at or below 4°C, resulting in a dissolved concentration greater than 50mg/L. Ozone-containing water was then allowed to decay at room temperature for more than 24 hours in an enclosed vessel to permit hazardous concentrations of ozone to decay. To ensure the absence of dissolved ozone during testing, the water was then boiled for a minimum of 10 minutes to cause ozone to decay rapidly and to reduce solubility, driving off the remaining ozone in the gaseous phase. This is a similar procedure to those used in chlorine-demand experiments, where sample containers and dilution waters were subjected to residual chlorine prior to testing (Wang et al. 2018).

3.3.4 Sample Water

Sample water used for testing was collected from treated wastewater effluent of Hampton Roads Sanitation District’s Nansemond treatment plant. This treatment plant operates a five-stage Bardenpho process with methanol addition for further denitrification. Following secondary clarification, process water is fed into the Sustainable Water Initiative for Tomorrow’s Research Center (SWIFT-RC) treatment system. The flocculation and sedimentation process utilized 25mg/L (as product) aluminum chlorohydrate (Kemira PAX-XL19, Delaware) as a coagulant and 0.75mg/L cationic polymer (Polydyne Clarifloc C6220, Riceboro, Georgia) for flocculation. All sample water for bench testing was collected as a batch immediately following solids settling. Sample water was also filtered with 0.45µm filters (GWV high capacity in-line groundwater sampling capsules, product number 12178, Pall Corporation, Ann Arbor, Michigan) in order to inhibit microbial activity during storage at 4°C. 500µg/L of bromide was added to the sample water to simulate a worst-case scenario.

Sample water was collected as a grab sample with water quality parameters summarized in Table 2. Only trichloroacetic acid was detected of the halogenated compounds measured.

Table 2: Sample water collected from flocculation/sedimentation process effluent Compound Value Unit

Temperature 20.2 °C pH 7.4 Nitrite-N 0.012 mg/L N Nitrate-N 2.52 mg/L N Ammonia-N 0.06 mg/L N Conductivity 1042 µS/cm Bromide 223 µg/L Bromate 0.182 µg/L Total Organic Carbon 7.03 mg/L NDMA 3.63 ng/L Trichloroacetic Acid 2.66 µg/L

Temperature was maintained at 20°C ±2°C for all experiments. pH was neither adjusted nor buffered but was constant at 7.2 ±0.3. Though phosphate and borate buffers were used in other literature (Buffle et al. 2004), buffer solutions were not used in this study due to the potential

31 impact into monochloramine interactions with brominated species during ozonation (Vikesland et al. 2001).

3.3.5 Analytical Methods

Ozone was measured by the volumetric and gravimetric indigo methods in accordance with Standard Method 4500B using a Horiba Aqualog-UV-800-C. Ozone stock solution concentration was determined using the volumetric method at a 180:1 dilution factor (1mL of ozone stock solution in 20mL of indigo II reagent, diluted to 200mL). Ozone residual measurements were performed using a modified version of the gravimetric method. Ozonated sample water was dispensed into culture tubes containing 5mL of indigo II reagent and 0.5mL malonic acid for a maximum total volume of approximately 25mL. Malonic acid was not found to be needed in prior testing as monochloramine did not interfere with the indigo reagent but was added to create a standard test procedure. A complete description of the experimental setup is included in Section 3.3.6.

UV spectral scans were performed using the same instrument with a wavelength range from 200nm to 800nm every 1nm. Absorbance specimens were set aside for more than one week at 4°C to allow for the decay of mono-, di-, and trichloramine, preventing incidental absorbance measurements. Bromate was measured using EPA method 300.1 with a Dionex 5000 plus with IC columns AS19 and AS24. Bromide was measured using EPA method 300.0 with a Dionex Integrion HPIC and column AS27. Total chlorine, monochloramine, free ammonia, total ammonia, and nitrite were measured using a Hach SL1000 portable parallel analyzer. pH was measured using an Orion Star Series pH meter, serial number B28014 with a Pinnacle GB Combo pH probe (476086). Temperature was measured using a Hannah Instruments HI 93510. N- nitrosodimethylamine and 1,4-dioxane concentrations were analyzed using EPA methods 521 and 522, respectively, utilizing an Agilent 7010B GC/MS Triple Quadrupole.

An interval sampler (Spectra/Chrom® IS-95 Interval Sampler, Spectrum Chromatography) was set to dispense 120mL/min of ozonated sample water into a 28 mL tube in 10 second intervals for a maximum of 9 seconds of dispensing time. Sample tubes contained 5mL of indigo solution and 0.5mL of malonic acid for free chlorine masking. For high concentration ozone samples (>1mg/L),

32 the solenoid valve feeding the interval sampler was manually diverted back to the sample container allowing for the gravimetric indigo method to be utilized.

3.3.6 Experimental Setup

All experiments were performed at 20°C with 1.25L borosilicate amber glass bottles. These were filled with 750mL unbuffered sample water and variable quantities of ozone-demand free water such that the total volume (sample, ODF water, and ozone stock solution) totaled 1L. This was done to ensure that all experiments were carried out under the same dilution as not to effect ozone decay rates and other kinetics. Monochloramine stock solution, if added, was spiked less than 10 seconds before the addition of ozone. The chlorine-ammonia process began with the addition of a dilute sodium hypochlorite solution to 750mL of sample water, followed by the addition of an ammonium sulfate solution after a one-minute contact time. Ozone stock solution was added using a peristaltic pump in volumetric dose mode into the continuously stirred sample water. Stirring rate was adjusted such that the maximum stir rate was achieved before vortexing. A schematic of the bench-scale setup is in Figure 9. It should be noted that results shown represent the diluted water (75% sample water, 25% ozone demand free HPW) and not the raw water.

Figure 9: Bench-scale ozonation apparatus used for experimentation.

3.4 Results and Discussion

3.4.1 Bromate Formation with Increasing Ozone Dose for Different Suppression Methods

Bromate is the key DBP formed during the ozonation process and an understanding of its formation is critical to maintaining drinking water standards, especially in high bromide source waters. Sample water tested in this study was spiked with 500µg/L additional bromide; therefore, high levels of bromate were anticipated, particularly at the highest ozone doses. Figure 10 shows bromate formation versus ozone dose in units of O3:TOC. Interestingly, the addition of free ammonia performed poorly when compared to the control sample as bromate concentrations increased. While the explanation for this result is not immediately apparent, increases in ozone decay rate and OH• proliferation may be responsible. This conclusion is discussed further in Section 3.4.5.

All doses of preformed monochloramine suppressed bromate formation when compared to the control sample. As monochloramine concentrations increased, bromate formation decreased. Similar effects were observed with the chlorine-ammonia process; however, 1mg/L of free chlorine followed by free ammonia did not suppress bromate below the control bromate level and no monochloramine residual was measured in this experiment. This may be indicating that the chlorine-ammonia’s oxidation of bromide into hypobromous acid, followed by the formation of bromamine from free ammonia addition did not have a major impact on suppressing bromate formation. Rather, it may be concluded that the main mechanism behind the chlorine-ammonia’s suppression of bromate is the formation of monochloramine, observed at higher chlorine doses where a monochloramine residual was observed. For example, the 5mg/L chlorine-ammonia (Cl2-

NH3) process sample performed similarly to the 3mg/L preformed monochloramine sample, in which the Cl2-NH3 sample had a monochloramine residual of 2.62mg/L.

(a) 200 (b) 25 Background 0.4mg/L Ammonia 0.8mg/L Ammonia 1mg/L MC 20 150 3mg/L MC 5mg/L MC 1mg/L Cl2-NH3 15 3mg/L Cl2-NH3 5mg/L Cl2-NH3 100

Bromate formed (ug/L) Bromate formed (ug/L)

50 5

0 0 0.0 0.5 1.0 1.5 2.0 2.5 0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4

Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

Figure 10: Impact of ozone dose (in terms of the ozone dose to TOC ratio) versus bromate formed. Figure (b) shows bromate concentrations near the maximum contaminant level of 10 µg/L. Free ammonia addition increased bromate formation, whereas both monochloramine (MC) and the chlorine-ammonia process (Cl2-NH3) process inhibited bromate formation.

Figure 10b shows the ozone dose which caused the bromate MCL to be exceeded, and is more helpful for operations than the doses depicted in Figure 10a. Free ammonia addition and 1mg/L

Cl2-NH3 process all formed similar bromate concentrations as the background sample, surpassing the MCL at an ozone dose of 0.4-0.5 mg O3:mg TOC. The addition of preformed monochloramine and the chlorine-ammonia process showed similar results to those discussed above, with 5mg/L monochloramine performing the best, maintaining a safe level of bromate up to an ozone dose of

0.95 mg O3:mg TOC.

3.4.2 Ozone Exposure Calculations

Relating bromate formation to ozone dose alone does not fully describe the system, as changing water quality (like with free chlorination associated with the Cl2-NH3 process) will change the decay rate of ozone. Ozone exposure is the product of ozone residual concentration and reaction time and can be used to describe ozone kinetics. However, as depicted in Figure 11, this product does not illustrate the full area under the decay curve. This approximation excludes concentration

35 changes before and after the selected time. In order to determine the true ozone exposure for ozone kinetics, the full area under the first order decay curve must be known.

Figure 11: Illustration depicting the first order decay of aqueous ozone and C*t used for calculating disinfection credits.

Integrating the first-order regression equation (Eq. 9) yields the ozone exposure (OE) at a given time (Eq. 10), where Co is the calculated initial concentration based on first order linearization and k is the first order decay rate. By integrating time from zero to infinity (Eq. 11), the theoretical total ozone exposure (TOE) may be determined.

−푘푡 퐶 = 퐶표푒 (Eq. 9)

푡 퐶 푒−푘푡 퐶 푒푡 퐶 푒0 퐶 푒푘푡 퐶 푂퐸 = ∫ 퐶 푒−푘푡푑푡 = [ 표 ] 푡 = 표 − 표 = 표 − 표 (Eq. 10) 0 표 푘 0 푘 푘 푘 푘

∞ 퐶 푒−푘푡 퐶 푒−∞ 퐶 푒0 퐶 푇푂퐸 = ∫ 퐶 푒−푘푡푑푡 = [ 표 ] ∞ = 표 − 표 = − 표 (Eq. 11) 0 표 푘 0 푘 푘 푘

Utilizing the TOE equation depicted in equation 10, the impact of bromate suppression techniques were compared to ozone doses in terms of O3:TOC. These results are illustrated in Figure 13

Ideally, ozone exposures should be maximized in order to increase organics degradation for increasing assimilable organic carbon (AOC) content and to maintain high levels of disinfection, as discussed in Section 3.4.5.

3.4.3 Impact of Suppression Techniques on Bromate Formation and Ozone Exposure

Ammonia addition was shown to have similar ozone exposures per ozone dose as the background sample except for the highest ozone dose. This is most likely due to the highest ammonia doses only receiving 2.08-2.17mg O3:mg TOC versus the 2.35 mg O3:mg TOC received by the background sample, thus the exposure of the background at ~2.15 mg O3:mg TOC was not captured in the trend and may have a similar exposure to the ammonia-spiked samples. However, it is also possible that there is another phenomenon occurring, as discussed in Section 3.4.5. Notably, the decay rate in the background sample with the highest ozone dose (k = -0.0027 s-1) was roughly half that of the ammonia-spiked samples (k = -0.0053 s-1), allowing for more ozone exposure.

Increasing preformed monochloramine concentrations reduced ozone exposures with respect to ozone dose. This may be most likely attributed to the direct reaction of monochloramine with ozone, as depicted in reaction 13 (Haag and Hoigne 1983). The highest monochloramine dose of 5mg/L was shown to reduce ozone exposure by roughly half at the highest ozone dose, though lower doses (less than 2 mg O3:mg TOC) showed similar TOE to the background sample

+ − − −1 −1 푁퐻2퐶푙 + 3푂3 → 2퐻 + 푁푂3 + 퐶푙 + 3푂2 • 푘 = 20 푀 푠 (R. 13)

The chlorine-ammonia process was the only technique shown to increase TOE with respect to ozone dose. Notably, ozone residuals were detected at the lowest dose of 0.25 mg O3:mg TOC, whereas for the other suppression mechanisms no ozone residual was measured. This is indicative of the IOD phase decreasing, caused by the oxidation of organics from free chlorine contact. Ozone exposures were roughly doubled up to 1.5:1 O3:TOC due to the addition of free chlorine, after which showed diminishing returns. At the highest ozone dose, the 3mg/L chlorine dose seemed to outperform the 5mg/L chlorine dose. Though the cause of this is unknown, it should have been expected that the highest chlorine dose resulted in the highest ozone exposure.

(a) 2500 (b) 2500 Background 0.4mg/L Ammonia 2000 0.8mg/L Ammonia 2000 1mg/L MC 3mg/L MC 1500 5mg/L MC 1500 1mg/L Cl2-NH3 3mg/L Cl2-NH3

1000 5mg/L Cl2-NH3 1000

Ozone Exposure (mg*s/L) OzoneExposure (mg*s/L) OzoneExposure 500 500

0 0 0.0 0.5 1.0 1.5 2.0 2.5 0.0 0.5 1.0 1.5 2.0 2.5

Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

2000 2000

1500 1500

1000 1000

Ozone Exposure (mg*s/L) OzoneExposure (mg*s/L) OzoneExposure 500 500

0 0 0.0 0.5 1.0 1.5 2.0 2.5 0.0 0.5 1.0 1.5 2.0 2.5 Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

Figure 12: (a) Impact of increasing ozone dose on ozone exposures under different bromate suppression techniques, including (b) free ammonia addition, (c) preformed monochloramine addition, and (d) the chlorine-ammonia process with a 1 minute free-chlorine contact time prior to ammonia addition.

All the different chlorine doses in the chlorine-ammonia process showed similar ozone exposures. If the reaction time of free chlorine had been increased from the 1-minute in the experiment, more organics oxidation would have likely occurred, and a larger difference in ozone exposures may have been illustrated. These elevated exposures can be correlated to higher organics oxidation and increased AOC (Wert et al. 2007).

It should be noted that influent samples water, when collected, had almost no background ammonia present (0.07 mg/L-N). With background ammonia present, the chlorine-ammonia process becomes problematic. With a small amount of background ammonia present (i.e. 0.1mg/L in a 5mg/L applied free chlorine dose), the free chlorine fully oxidizes the ammonia into nitrogen trichloride, and breakpoint chlorination occurs. Between a 7.1:1 and 5:1 chlorine to ammonia mass ratio, dichloramine and trichloramine may become the dominant chloramine species. Dichloramine is known to form NDMA when left with high contact times in water (West et al. 2016). However, with only 1-minute reaction time studied, the contact time of dichloramine would be limited and it is hypothesized that NDMA would not form significantly. Upon addition of free ammonia, dichloramine-monochloramine equilibrium shifts in favor of monochloramine (Czech et al. 1961). Additionally, NDMA formed during the contact with dichloramine may be oxidized during the downstream ozonation process (Marti et al. 2015), reducing NDMA concentrations as well as NDMA formation potentials, though this requires further study. To help mitigate this issue, treatment facilities could add free ammonia upstream of the chlorine addition point under this condition. By reducing the chlorine to ammonia mass ratio below 5:1, monochloramine becomes the dominant chloramine species (Czech et al. 1961). In this case, a standard monochloramine suppression of bromate would occur during the ozonation process, and excess ammonia would not need to be added downstream of chlorine addition. An added benefit to this process would be the oxidation of fast-reacting organic moieties via chlorine prior to the formation of monochloramine, though this would not be as pronounced as with a free chlorine contact.

250

200 Background 0.4mg/L Ammonia 150 0.8mg/L Ammonia 1mg/L MC

100 3mg/L MC

5mg/L MC Bromateformed(ug/L) 50 1mg/L Cl2-NH3 3mg/L Cl2-NH3 5mg/L Cl2-NH3 0 0 500 1000 1500 2000 2500 Ozone Exposure (mg*s/L)

Figure 13: Bromate formation as a function of ozone exposure when utilizing differenct bromate suppression techniques, where ozone exposure is represented as total ozone exposure (TOE) rather than a C*t value. Bromate formation is known to increase as ozone exposure increases (Buffle et al. 2004). However, the addition of chemicals for bromate suppression changes ozone decay rate and the subsequent ozone exposure drastically. As noted in Figure 13, ammonia addition dramatically increased bromate formation per unit of ozone exposure. As the ozone decay rate for these samples was nearly double that of the background sample, it may be reasoned that more OH• proliferation occurred. This is explained further in Section 3.4.5. Preformed monochloramine addition and the chlorine-ammonia process were able to suppress bromate at higher ozone exposures. When compared to bromate formed versus ozone dose (Figure 10), the chlorine-ammonia process outpaced monochloramine in regards to TOE. This is a result of oxidation from free chlorine prior to ozonation, reducing IOD and stabilizing decay rate (Buffle et al. 2004). Ozone exposure is responsible for increasing disinfection (Section 3.4.4) and the biodegradability of organic carbon. It can be concluded that the proper implementation of the chlorine-ammonia process will increase assimilable organic carbon (AOC). This will impact carbon removal from biofiltration positively, reducing the organic loading on downstream granular activated carbon processes (Wert et al. 2007).

3.4.4 Impact of Chemical Addition on Disinfection Credits

Disinfection credits are granted for ozonation systems through the LT2SWTR (US EPA 2006). Unlike the previous section’s analysis, these credits are attained through the simple product of concentration and time. Figure 14 depicts C*t values calculated with variable reaction time. Though the highest residual ozone concentrations are at the beginning of the reaction, the low time values result in a small C*t product. As time increases, C*t values increase until concentration values begin to drop off; therefore, there is a point of diminishing returns. By deriving the first order regression equation and setting the slope of the derivation to zero, the maximum C*t product may be determined. This derivation, depicted in Figure 14, shows that the time of maximum C*t (therefore disinfection credit) occurs at the absolute value of the inverse decay rate. In online processes, multiple measurement locations along the ozone contactor are critical to maximizing the disinfection credit. Within a batch-scale ozone process, the time of the reaction in which the concentration was measured is also vital for maximizing disinfection credit.

Figure 14: C*t values calculated along an ozone decay curve. Insert shows the derivation for the time at which maximum C*t occurs (1/k).

Utilizing the inverse decay rate function defined above, maximum C*t values were calculated from the first-order linearization of ozone decay curves measured in the experiments. All curves followed near-identical trends as those illustrated in ozone exposure versus ozone dose: interestingly, this is due to the maximum C*t value being equivalent to TOE divided by Euler’s number, “e.” The formulas used for calculating disinfection credits in Figure 15 use the maximum C*t value directly, and are only slightly affected by temperature, which was held constant at 20°C ± 1°C throughout the experiments.

Virus disinfection credits, as depicted in Figure 15a and b, are the easiest to achieve, and a standard goal is 5 LRV within ozonation systems. In Figure 15b and Figure 15c, similar virus and giardia inactivation credit was attained by the background, ammonia-spiked and monochloraminated samples: achieved between an ozone dose between 0.4 and 0.55 mg O3:mg TOC. Chlorine- ammonia process samples showed great improvements in the required ozone dose for disinfection credit, lowering the required dose to approximately 0.3 mg O3:mg TOC. The 5mg/L Cl2-NH3 process sample likely had inaccuracies in the lowest ozone doses: the 0.5 mg O3:mg TOC ozone dose data point appears to have an arbitrarily low virus inactivation value; with this data point removed, the 5mg/L Cl2-NH3 process curve follows close to the 3mg/L Cl2-NH3 process curve. Cryptosporidium removal is more difficult in ozonation systems, and removal credits granted are more than an order of magnitude lower than virus and giardia removal credits (Figure 15d). However, the same results can be gleaned from the trends: the chlorine-ammonia process outperformed other suppression mechanisms. For the selected 1-log Cryptosporidium removal, this credit was attained with a maximum of 0.75 mg O3:mg TOC.

1200 50 (a) Background (b) 0.4mg/L Ammonia 45 1000 0.8mg/L Ammonia 40 1mg/L MC 35 800 3mg/L MC 5mg/L MC 30 600 1mg/L Cl2-NH3 25 3mg/L Cl2-NH3 20 5mg/L Cl2-NH3 400

Virus Log Removal Value Removal Virus Log Value Removal Virus Log 10 200 5 0 0 0.0 0.5 1.0 1.5 2.0 2.5 0.0 0.2 0.4 0.6 0.8 1.0

Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

Giardia Log Removal Value Removal Giardia Log 5 1.0

Cryptosporidium Log Removal Removal Value Log Cryptosporidium 0.5 0 0.0 0.0 0.2 0.4 0.6 0.8 1.0 0.00 0.25 0.50 0.75 1.00 1.25 1.50

Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

Figure 15: Log removal values calculated from LT2SWTR for (a & b) virus, (c) Giardia, and (d) Cryptosporidium parvum as a function of applied ozone dose in terms of O3:TOC

As mentioned in Section 3.4.2, background influent ammonia concentrations will have a significant impact on the chlorine-ammonia process. With increasing ammonia, the free chlorine contact will become less efficient, especially below free-chlorine breakpoint. Therefore, the boost in disinfection from the chlorine-ammonia process relative to ozone dose will become less

43 impactful with increasing ammonia concentrations, as fewer organic constituents are oxidized prior to ozonation.

250 20 (a) Background (b) 0.4mg/L Ammonia 18 0.8mg/L Ammonia 200 1mg/L MC 16 3mg/L MC 14 5mg/L MC 150 1mg/L Cl2-NH3 12 3mg/L Cl2-NH3 5mg/L Cl2-NH3 10 100 8

Bromate formed (ug/L) Bromate formed (ug/L) Bromate formed 50 4 2 0 0 0 500 1000 0 25 50 75 100 125 150 175 200 225 250 Virus Log Removal Value Virus Log Removal Value

Figure 16: Bromate formation as a function of Virus Log Removal, calculated using values from the LT2SWTR.

Balancing disinfection credits with bromate formation is vital for system operations in full scale plants striving to use the ozonation system as a pathogen barrier. The use of bromate suppression chemicals can increase pathogen inactivation through ozone stabilization and preoxidation while still reducing bromate formation. As shown in Figure 16, ammonia addition had a negative impact on disinfection credit per bromate formed. This phenomenon is explained further in Section 3.4.5. Monochloramine showed some improvement in increasing disinfection credit attained per bromate formed, but this was most likely due to bromate suppression rather than increasing disinfection credits themselves. Figure 16b shows that, while maintaining the MCL of 10µg/L bromate, the addition of 1mg/L, 3mg/L, and 5mg/L preformed monochloramine increased LRVs to 30, 55, and 70, respectively. The chlorine-ammonia process outperformed monochloramine alone due to the added benefit of preoxidation from free chlorine, increasing the attainable LRV to 20, 70, and 100, respectively. The 1mg/L dose of the chlorine-ammonia process did not perform as well as 1mg/L preformed monochloramine as all free chlorine was consumed during the 1-minute contact time, leaving no chlorine species available to form a monochloramine residual.

3.4.5 Hydroxyl Radical Exposure

OH• exposure was measured using 1,4-dioxane as a probe compound. 1,4-dioxane’s reaction with ozone directly is nearly non-existent (0.32 M-1s-1) but reacts quickly with OH• (2.9 *109 M-1s-1) (Buxton et al. 1988). Through the removal of 1,4-dioxane in the batch experiments, the exposures of OH• could be calculated through equation 8. This is a critical component to understanding the oxidative capabilities of the ozonation system as many trace organic contaminants are only slowly oxidized by ozone, with others being completely non-reactive. However, OH• exposure is also responsible for bromide oxidation through the indirect bromate formation pathway. Therefore, balancing radical oxidation with bromate formation is critical and must be addressed by the bromate suppression technique’s mechanism.

Figure 17 depicts OH• exposure versus ozone dose with different bromate suppression techniques. As noted in (Lee and von Gunten 2016), OH• versus ozone dose trends linearly. Based on this relationship, it can be stated that, for a given water quality, a specific fraction of OH• per ozone dose will be generated. This fraction of OH• per ozone added may be defined by the slope of the linear trend in Figure 17. However, the addition of chemicals in the water will change this ratio. Interestingly, as depicted in Figure 17b, the addition of ammonia increased the degradation of 1,4- dioxane (and the resulting calculated OH• exposure) by nearly double. The exact mechanism for this change is unknown: ammonia itself does not act as a radical promotor (Pinkernell and Von Gunten 2001). Therefore, two potential explanations may be made:

(a) Ammonia itself is acting as a catalyst for the degradation of 1,4-dioxane. This has never been reported in literature and is a highly unlikely phenomenon. (b) Sulfate (added in conjunction with ammonia as Ammonium Sulfate) may be acting as a OH• initiator/promotor. The sulfate radical has a similar or even a higher redox potential than OH• (2.6 V versus 1.77-2.74V (Buxton et al. 1988; Deng and Ezyske 2011)) and has been implemented as an advanced oxidation process (Deng and Ezyske 2011). Sulfate radicals, which may be formed from complex reactions during initial ozone decay, naturally decay into OH•, which would go on to react with 1,4-dioxane. Increasing the ammonium sulfate dose did not increase radical exposure.

(a) 1.E-09 Background (b) 1.E-09 0.4mg/L Ammonia 9.E-10 0.8mg/L Ammonia 9.E-10 1mg/L MC 8.E-10 3mg/L MC 8.E-10 7.E-10 5mg/L MC 7.E-10 1mg/L Cl2-NH3 6.E-10 3mg/L Cl2-NH3 6.E-10 5mg/L Cl2-NH3 5.E-10 5.E-10 4.E-10 4.E-10 3.E-10 3.E-10

2.E-10 2.E-10

Hydroxyl Radical (M*s) Radical Exposure Hydroxyl (M*s) Radical Exposure Hydroxyl 1.E-10 1.E-10 0.E+00 0.E+00 0.0 0.5 1.0 1.5 2.0 2.5 0.0 0.5 1.0 1.5 2.0 2.5

Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

5.E-10 4.E-10

4.E-10 3.E-10 3.E-10 2.E-10

2.E-10 Hydroxyl Radical (M*s) Radical Exposure Hydroxyl Hydroxyl Radical (M*s) Radical Exposure Hydroxyl 1.E-10 1.E-10

0.E+00 0.E+00 0.0 0.5 1.0 1.5 2.0 2.5 0.0 0.5 1.0 1.5 2.0 2.5

Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

Figure 17: OH* exposures as a function of ozone dose for different bromate suppression techniques: (a) all methods, (b) free ammonia addition, (c) preformed monochloramine addition, and (d) the chlorine-ammonia process with 1 minute free chlorine contact time.

As the remainder of the bromate suppression techniques also used ammonium sulfate, it is important to consider the free-ammonia scenario as the controlled experiment when comparing suppression techniques. Figure 17c depicts monochloramine’s impact on OH• exposure. Monochloramine is known to act as a radical scavenger, though it appears that the lowest doses of monochloramine increased radical exposure from the background. However, a comparison of

46 monochloramine to the free-ammonia samples shows that monochloramine suppressed OH• exposure with increasing monochloramine concentrations, where the highest doses of monochloramine of 3mg/L and 5mg/L had similar OH• exposures to that of the unperturbed sample water.

The chlorine-ammonia process showed similar OH• suppression to that of monochloramine addition. This may be directly attributed to the monochloramine residual formed during this process.

Bromate Formed vs Hydroxyl Radical Exposure 250

Background 200 0.4mg/L Ammonia

150 0.8mg/L Ammonia 1mg/L MC 100 3mg/L MC

5mg/L MC Bromateformed(µg/L) 50 1mg/L Cl2-NH3 3mg/L Cl2-NH3 0 0.E+00 2.E-10 4.E-10 6.E-10 8.E-10 1.E-09 5mg/L Cl2-NH3 Hydroxyl Radical Exposure (M*s)

Figure 18: Bromate formation as a function of hydroxyl radical exposure Figure 18 illustrates that the suppression technique’s formations of bromate are linear with OH• exposure, except for the control sample. With increasing doses of preformed monochloramine and chlorine-ammonia process concentrations, the total OH• exposure and bromate decrease linearly. This is indicating that monochloramine is decreasing radical exposure through radical scavenging and is a key mechanism behind bromate suppression. The chlorine-ammonia process had more bromate formed per OH• exposure, likely due to increased ozone exposure. The 5mg/L Cl2-NH3 experiment approaches monochloramine trends, due to the higher impact of monochloramine residual versus preoxidation through free chlorine contact. All suppression techniques caused a significant deviation from the background sample water’s bromate formation. It can be theorized

47 that this shift from an exponential increase to a linear trend is related to the formation of intermediate products such as bromamine or bromochloramine due to the presence of ammonia in the sample water (Allard et al. 2018; Buffle et al. 2004; Gazda and Margerum 1994; Wajon and Morris 1982).

9.00E-10 9.00E-10 (a) 0.25 (b) 8.00E-10 0.5 8.00E-10

7.00E-10 1 7.00E-10 1.5 6.00E-10 6.00E-10 2.5 5.00E-10 5.00E-10

4.00E-10 4.00E-10

3.00E-10 3.00E-10

Hydroxyl Radical Exposure (Ms) Exposure Radical Hydroxyl 2.00E-10 Hydroxyl Radical Exposure (Ms) Exposure Radical Hydroxyl 2.00E-10

1.00E-10 1.00E-10

0.00E+00 0.00E+00 0 1 2 3 4 5 0 1 2 3 4 5 Monochloramine Concentration (mg/L) Monochloramine Concentration (mg/L)

Figure 19: Impact of monochloramine concentrations on hydroxyl radical exposures for various O3:TOC ratios in (a) preformed monochloramine addition and (b) the chlorine-ammonia process.

The chlorine-ammonia process inherently forms monochloramine upon free ammonia addition, therefore it is important to resolve the differences between the two suppression techniques. Figure 19 compares the hydroxyl radical exposures from preformed monochloramine (Figure 19a) to the monochloramine formed through the chlorine ammonia process (Figure 19b) for different ozone doses. As expected, OH• exposures increased with increasing ozone dose. The chlorine-ammonia process suppressed OH• exposure with lower residual concentrations of monochloramine. This is most likely attributable to the oxidation of promoting organics via free chlorine oxidation prior to ozonation, reducing IOD. Therefore, it can be stated that the chlorine-ammonia process suppresses radical exposure more than preformed monochloramine addition for the same quantity of chemicals.

7.E-07 7.E-07 (a) Background (b) 0.4mg/L Ammonia 6.E-07 0.8mg/L Ammonia 6.E-07 1mg/L MC 5.E-07 3mg/L MC 5.E-07 5mg/L MC 4.E-07 1mg/L Cl2-NH3 4.E-07

3mg/L Cl2-NH3

CT CT R R 5mg/L Cl2-NH3 3.E-07 3.E-07

2.E-07 2.E-07

1.E-07 1.E-07

0.E+00 0.E+00 0.0 0.5 1.0 1.5 2.0 2.5 0.0 0.5 1.0 1.5 2.0 2.5

Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

6.E-07 6.E-07

5.E-07 5.E-07

4.E-07 4.E-07

CT CT

R R 3.E-07 3.E-07

2.E-07 2.E-07

1.E-07 1.E-07

0.E+00 0.E+00 0.0 0.5 1.0 1.5 2.0 2.5 0.0 0.5 1.0 1.5 2.0 2.5

Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

RCT, the ratio of OH• exposure to ozone exposure, is an important value for determining the relative oxidative capability of an ozonation process. Higher RCT values indicate that an ozonation process is leaning towards an AOP: ozone is rapidly transforming into OH• and ozone is decaying at a high rate. Elevated RCT is indicative of significant trace contaminant oxidation, but also of a higher propensity of bromide oxidation to bromate through the indirect bromate formation pathway.

Higher RCT values are observed during initial ozone decay (von Gunten 2003a), as supported by

Figure 20. As expected, ammonia did not have a direct impact on RCT. Monochloramine, on the other hand, did have an impact: The addition of the lowest dose of 1mg/L increased RCT, whereas

3mg/L and 5mg/L decreased RCT. Interestingly, 3mg/L monochloramine decreased RCT more than 5mg/L; this was most likely caused by inaccuracies in the measurement of ozone exposure at the lowest ozone doses.

The chlorine-ammonia process had the largest impact on RCT. Due to the prechlorination step, many of the organic constituents responsible for OH• initiation and propagation are already oxidized. As depicted in Figure 20d, the highest doses of 3mg/L and 5mg/L resulted in the near- neutralization of elevated RCT during IOD for the ozone doses tested. In conclusion, the chlorine- ammonia process is able to reduce OH• exposure while simultaneously increasing TOE, as indicated by RCT. This will shift bromate formation from the fast-indirect pathway to the slower direct formation pathway through ozone, reducing overall bromate formation. However, this will also negatively impact trace organic contaminant oxidation.

3.4.6 Impact of dose and suppression techniques on organics oxidation: UV spectral scans

Work by Gerrity et al., 2012 and others have correlated the change in UV absorbance at 254nm to the oxidation of organic constituents (Gerrity et al. 2012). Figure 21 shows the change in UV254 for the ozone doses studied under the influence of bromate suppression techniques. As ozone doses increase, the change in UV254 absorbance also increases, but at a diminishing rate. Figure 21b depicts a decrease in the reduction of UV254 when free-ammonia is added; this is in contradiction to the OH• exposure results explained in Section 3.4.5. The use of preformed monochloramine and the chlorine-ammonia process decreased UV254 abatement with increasing doses.

(a) 0.07 (b) 0.07

0.06 0.06

0.05 0.05

0.04 0.04 Blank 0.03 0.4mg/L NH4 0.03 0.8mg/L NH4 1mg/L NH2Cl 0.02 3mg/L NH2Cl 0.02

5mg/L NH2Cl UV Absorbance 254nm at Absorbance (1/cm) UV 0.01 1mg/L NH4-Cl2 254nm at Absorbance (1/cm) UV 0.01 3mg/L NH4-Cl2 5mg/L NH4-Cl2 0 0 0 0.5 1 1.5 2 2.5 0 0.5 1 1.5 2 2.5

Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

0.06 0.06

0.05 0.05

0.04 0.04

0.03 0.03

0.02 0.02 UV Absorbance 254nm at Absorbance (1/cm) UV UV Absorbance 254nm at Absorbance (1/cm) UV 0.01 0.01

0 0 0 0.5 1 1.5 2 2.5 0 0.5 1 1.5 2 2.5

Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

Figure 21: Changes in UV absorbance at 254nm as compared to the background sample as a function of ozone dose in mg O3:mg TOC for a variety of chemical suppression techniques, including (a) all experiments, (b) free ammonia addition (NH4), (c) preformed monochloramine addition (NH2Cl), and (d) the chlorine-ammonia process (Cl2-NH4) with a one minute free chlorine contact time

Further work by Gerrity et al. utilizes the change in UV254 to generate a OH• exposure model. In practice, a model based on UV254 is more favorable than the utilization of probe compounds due to a significant cost decrease and the ability for online measurement. The development of this model began with a graphical comparison and linearization of the percent decrease in para- chlorobenzoic acid concentrations and the measured change in UV254 absorbance. Unlike in this work, 1,4-dioxane was used as the OH• probe compound rather than para-chlorobenzoic acid due

51 for reasons previously stated. The relationships for all ozone doses and bromate suppression techniques studied are illustrated in Figure 22. All of the tests performed fit a linear trend well, and the slopes, intercepts, and coefficients of determination (r2) are tabulated in Table 3. However, it is noted that the 3mg/L preformed monochloramine results differed rather significantly: the trend can be best be described as an exponential increase. This result is most likely attributed to unknown sampling errors. Additionally, the 3mg/L chlorine-ammonia process samples had a much greater percentage change in UV254 absorbance per percent removal of 1,4-dioxane.

Overall, the entire testing matrix’s linear trend fit well (r2 = 0.80). This suggests that the addition of the bromate suppression chemicals studied do not interfere significantly for this relationship, and that OH• exposure may be modeled for all conditions studied without perterbation.

100

70 Blank 0.4mg/L NH4 60

0.8mg/L NH4 dioxane Concentration - 50 1mg/L NH2Cl 3mg/L NH2Cl 40 5mg/L NH2Cl 30 1mg/L NH4-Cl2 20 3mg/L NH4-Cl2 5mg/L NH4-Cl2

10 PercentDecrease 1,4 in

0 0 20 40 60 80 100 Percent Decrease in UV254 Absorbance

Figure 22: Percent decreases in 1,4-dioxane concentrations versus the percent decreases in UV254 absorbance for different bromate suppression techniques, subjected to various ozone doses.

Table 3: Linear regression parameters for the percent decrease in 1,4-dioxane versus the percent decrease in UV254 absorbance

Test Slope Intercept R2 Background 1.92 -27.8 0.9974 0.4mg/L Ammonia 2.55 -39.2 0.9768 0.8mg/L Ammonia 2.59 -43.3 0.8632 1mg/L Monochloramine 2.02 -22.0 0.9467 3mg/L Monochloramine 2.02 -18.1 0.6731 5mg/L Monochloramine 2.81 -50.6 0.9087 1mg/L Cl2-NH3 2.32 -43.8 0.9588 3mg/L Cl2-NH3 3.40 -98.9 0.9713 5mg/L Cl2-NH3 2.48 -50.5 0.9880 All Tests 2.11 -31.7 0.8018

Equation 11 below relates the percent change in 1,4-dioxane (1,4D) to the percent change in UV254 absorbance, as depicted graphically in Figure 22. Equation 12 is the calculated OH• exposure based on 1,4-dioxane and its reaction rate with OH•, 2.9*109 M-1s-1. By substituting Equation 11 into Equation 12, Equation 13 may be derived. This equation is an alternative form of the one used by other literature, using relationships formed by 1,4-dioxane rather than pCBA. The influence of ozone oxidation on 1,4-dioxane degradation was discounted in these equations, as they were deemed negligible (<<1%).

1,4퐷 푈푉 (1 − ) ∗ 100(%) = 푆푙표푝푒 ∗ [(1 − ) ∗ 100(%)] + 퐼푛푡푒푟푐푒푝푡 (Eq. 11) 1,4퐷0 푈푉0

1,4퐷 ln( ) ∫[푂퐻 •]푑푡(푀푠) = 1,4퐷0 (Eq. 12) −2.9∗109

푈푉 퐼푛푡푒푟푐푒푝푡 ln(1−푆푙표푝푒∗[(1− )]− ) ∫[푂퐻 •]푑푡(푀푠) = 푈푉0 100 (Eq. 13) −2.9∗109

(a) (b) 7E-10 7E-10 y = 0.9753x + 3E-11

R² = 0.7281

dioxane dioxane dioxane

6E-10 6E-10

- -

5E-10 5E-10

4E-10 4E-10

(M*s) (M*s)

3E-10 3E-10

2E-10 2E-10

1E-10 1E-10

Measured OH˙ Exposure through 1,4 through Exposure OH˙ Measured 1,4 through Exposure OH˙ Measured

0 0 0 1E-10 2E-10 3E-10 4E-10 5E-10 6E-10 7E-10 0 1E-10 2E-10 3E-10 4E-10 5E-10 6E-10 7E-10 Modeled OH˙ Exposure through UV254 (M*s) Modeled OH˙ Exposure Through UV254 (M*s)

Figure 23: (a) Validation of the OH˙ exposure model using ∆UV254 when compared to measured OH˙ exposures through 1,4-dioxane. (b) The impact of bromate suppression technique on the accuracy of the OH˙ exposure model OH• exposures were calculated using the change in UV254 absorbance using Equation 13. These results were validated against measured OH• exposures using 1,4-dioxane as a probe compound in Figure 23. Results fit a linear trend near the 1:1 optimal line, and the coefficient of determination was determined to be 0.73 (Figure 23a). Figure 23b depicts the various bromate suppression techniques’ impacts on model accuracy. As predicted, the type of suppression technique did not significantly impact model results. However, free ammonia samples with the highest dose of ozone (2.5:1) were under-predicted by the model. This may be a relic of the phenomenon explained in Section 3.4.5. Additionally, the 3mg/L preformed monochloramine sample was over-predicted by the model, though the reason for this is not currently known. All other preformed monochloramine and chlorine-ammonia process samples were well described by the model.

3.4.7 Impact of dose and suppression techniques on NDMA formation 180 (b) 50 (a) Background 45 160 0.4mg/L Ammonia 140 0.8mg/L Ammonia 40 1mg/L MC 35 120 3mg/L MC 30 100 5mg/L MC 1mg/L Cl2-NH3 25 80 3mg/L Cl2-NH3 20 60

5mg/L Cl2-NH3 15 NDMA Concentration (ng/L) NDMA Concentration NDMA Concentration (ng/L) NDMA Concentration 40 10 20 5 0 0 0.0 0.5 1.0 1.5 2.0 2.5 0.0 0.5 1.0 1.5 2.0 2.5

Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

40 NDMA Concentration (ng/L) NDMA Concentration NDMA Concentration (ng/L) NDMA Concentration 40

20 20

0 0 0.0 0.5 1.0 1.5 2.0 2.5 0.0 0.5 1.0 1.5 2.0 2.5 Ozone Dose (mg O3:mg TOC) Ozone Dose (mg O3:mg TOC)

Figure 24: (a) Impact of various ozone doses on NDMA formation for the various bromate suppression techniques studied, including (b) free ammonia addition, (c) preformed monochloramine addition, and (d) the chlorine- ammonia process

Figure 24 shows NDMA formation during ozonation. All the results indicated that NDMA was formed during the ozonation process above the 10ng/L California State health advisory limit. The general trend of NDMA formation versus ozone dose shows a plateau once residual conditions are met. Free ammonia addition showed a decrease in NDMA formation; most likely caused by an

55 increase in OH• exposure as noted in section 3.4.5. 5mg/L preformed monochloramine and 3mg/L and 5mg/L of the chlorine-ammonia process formed 130-170ng/L of NDMA in process water without ozonation. When monochloramine is left in contact with process water for a significant contact time (in this case, 24 hours), mono- and dichloramine reacts with precursors in the water to form NDMA (West et al. 2016). However, sample water subjected to ozonation immediately reduced NDMA formation to 20-30ng/L. Therefore, it can be deduced that the addition of ozone reduces the amount of NDMA formation potential connected with these methods (von Gunten 2003b). Preformed monochloramine samples subjected to ozonation formed similar NDMA as the background, but the chlorine-ammonia process formed approximately 30-50% less with than background. This may be due to the free chlorination step reducing NDMA formation potential. Based on this figure, it can be stated that the addition of bromate suppression chemicals does not increase NDMA formation through ozonation but may in fact reduce NDMA formation with the use of the chlorine-ammonia process.

3.5 Conclusions

Water collected from the effluent of a 5-stage Bardenpho wastewater treatment plant was subjected to flocculation and sedimentation, followed by a wide range of ozone doses and bromate mitigation techniques, including free ammonia addition, preformed monochloramines, and the chlorine- ammonia process. Sample water was spiked with 500µg/L bromide to simulate the impact of landfill leachate on process water, and 400µg/L of 1,4-dioxane was added for use as a OH• probe compound. The following conclusions may be drawn from the experiments performed:

• Of the processes studied, unaltered sample water and sample water spiked with ammonium sulfate were found to not adequately suppress bromate below the MCL. However, both preformed monochloramine addition and the chlorine-ammonia process were able to reduce bromate formation to 10µg/L, with 5mg/L of preformed performing best, and

3mg/L preformed monochloramine and 5mg/L Cl2-NH3 performing second best. • Free ammonia addition increased OH• exposure while reducing ozone exposure, due to the potential of a sulfate radical forming and initiating a radical chain reaction. Preformed monochloramine addition did not decrease ozone exposure but did decrease OH• exposure. This provides additional evidence that monochloramine acts as a OH• scavenger. The chlorine-ammonia process increased TOE and decreased OH• exposure. Therefore, the chlorine-ammonia process is ideal in this water for maximizing ozone exposure, associated with increasing AOC for biofiltration and disinfection, while maintaining satisfactory bromate formation levels. • A OH• exposure model using 1,4-dioxane as a probe compound was calibrated based on prior work by Gerrity et al., 2012. This model can be implemented in full-scale processes of similar water quality to monitor OH• exposure and be periodically checked with 1,4- dioxane degradation rates. • NDMA was formed at a consistent rate for all bromate mitigation techniques as soon as a residual condition was met, with the chlorine-ammonia process forming slightly less NDMA. The chlorine-ammonia process and preformed monochloramine samples without ozone addition formed considerable NDMA; therefore, it can be concluded that ozonation, while forming NDMA itself, can also work as a treatment step following these processes

57 to limit NDMA formation. The 30-40ng/L NDMA formed through ozonation may be removed downstream through biofiltration, granular activated carbon, or UV disinfection.

References

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4. Engineering Significance

Water reuse is becoming increasingly vital to the future of humanity as population increases and fresh water supplies decrease. Central to the processes of water reuse is the ozonation process, for the purpose of disinfection, contaminant abatement, and increasing the biodegradability of organic constituents for downstream biofiltration. However, the ozonation process does have the potential to form concentrations of bromate exceeding the EPA’s maximum contaminant level. To mitigate bromate formation, a combination of ozone dose limitation and chemical suppression techniques are utilized, such as free ammonia addition, monochloramine addition, and the chlorine-ammonia process. In the reuse water studied, free ammonia addition was not found to decrease bromate formation significantly as compared to unaltered sample water subjected to ozonation. Preformed monochloramine and the chlorine-ammonia process, on the other hand, were more than able to reduce bromate formation below the MCL when coupled with ozone dose limitations. Preformed monochloramine seemingly outperformed the chlorine-ammonia process when comparing bromate formed versus ozone dose; however, other metrics for determining system efficiency must be considered when comparing these processes.

Ozone exposure can be directly attributed to increasing assimilable organic carbon and disinfection. When comparing preformed monochloramine addition to the chlorine-ammonia process in terms of bromate formed per unit of ozone exposure, the chlorine-ammonia process came out decidedly ahead. When observing attainable disinfection credit benchmarks, the chlorine-ammonia process required nearly half the ozone dose as the unperturbed sample.

Ozone exposure alone, however, is not responsible for some trace contaminant oxidation, and OH• exposure is required. One such trace contaminant is 1,4-dioxane, which was spiked to 400µg/L to serve as a OH• probe compound. Both preformed monochloramine addition and the chlorine- ammonia process were found to slightly decrease OH• exposure, most likely due to the monochloramine molecule acting as a OH• scavenger. There was no significant difference between preformed monochloramine’s suppression of OH• exposure and the chlorine-ammonia process’s, though dissimilarity was noted when comparing the resulting monochloramine residual of the chlorine-ammonia process. Utilizing information gleaned from RCT values, it was concluded that the chlorine-ammonia process decreased OH• exposure further through the preoxidation of

62 initiating and promoting organics prior to the ozonation process. The OH• suppression resulting from both preformed monochloramine addition and the chlorine-ammonia process are similar, though through different mechanisms.

N-Nitrosodimethylamine formation was also monitored. With the addition of ozone, comparable NDMA formations were observed for all bromate suppression techniques tested. However, the chlorine-ammonia process slightly outperformed the rest, but was still above the 10ng/L California Health Advisory limit. With downstream biofiltration and UV disinfection, the ~30ng/L NDMA formed should be removed post-ozonation in the treatment process.

Summarizing the information provided, it is recommended that free ammonia addition not be used for suppressing the formation of bromate, as no positive impact was observed. Both preformed monochloramine and the chlorine-ammonia process were able to adequately suppress bromate formation, but the chlorine-ammonia process boasts additional capabilities. When compared to preformed monochloramines, lower ozone doses may be used with the chlorine-ammonia process to achieve necessary disinfection credits, lowering capital cost of a larger ozone generator. It should be noted that the chlorine-ammonia process has a drawback in wastewater treatment plant effluent: when influent ammonia concentrations swing high, the chlorine-ammonia process becomes less efficient. Further study is required to determine the proper protocol for operating a chlorine-ammonia process in waters with variable ammonia.

Appendix A

Figure 25: Absorbance spectrum of the monochloramine stock solution used for 1mg/L preformed monochloramine testing.

Figure 26: Absorbance spectrum of the monochloramine stock solution used for 3mg/L preformed monochloramine testing.

Figure 27: Absorbance spectrum of the monochloramine stock solution used for 5mg/L preformed monochloramine testing. UV Spectral Scan Data

A common metric of water quality is the absorbance of ultraviolet light, particularly in oxidation systems such as ozonation. Organic constituents absorb ultraviolet light at different wavelengths within the 240-330nm range. Typically, water and wastewater treatment facilities utilize the absorbance at 254nm (UV254) as a point of reference for treatment goals. However, it is important to understand the impacts of an oxidation system on the full spectrum of absorbance changes in order to determine if a better reference point exists. Figure 28 illustrates the change in absorbance for the bromate suppression techniques studied at a variety of ozone doses. Naturally, as the ozone dose increased, the reduction of UV254 absorbance increased. The background sample in Figure 28a showed the maximum reduction in absorbance occurs at 240nm and minimally at 310nm; a decrease in 0.062cm-1 and 0.036cm-1, respectively.

(a) 0.07 (b) 0.07 (c) 0.07 0.06 0.06 0.06 0.05 0.05 0.05 0.04 0.04 0.04 0.03 0.03 0.03 0.02 0.02 0.02

0.01 0.01 0.01

(d) 0.07 (e) 0.07 (f) 0.07 0.06 0.06 0.06 0.05 0.05 0.05 0.04 0.04 0.04 0.03 0.03 0.03 0.02 0.02 0.02

0.01 0.01 0.01

(g) 0.07 (h) 0.07 (i) 0.07 0.06 0.06 0.06 0.05 0.05 0.05 0.04 0.04 0.04 0.03 0.03 0.03 0.02 0.02 0.02

0.01 0.01 0.01

Absorbance Absorbance Difference (1/cm) Absorbance Difference (1/cm) Absorbance Difference (1/cm) 0.00 0.00 0.00 240 250 260 270 280 290 300 310 240 250 260 270 280 290 300 310 240 250 260 270 280 290 300 310 Figure 28: Absorbance Wavelengthozonated (nm) water samples from 240nm to 330nm for theWavelength different (nm) chemical suppression techniques studied. (a)Wavelength no broma (nm)te suppression, (b) 0.4mg/L-N free ammonia, (c) 0.8mg/L-N free ammonia, (d) 1mg/L preformed monochloramine, (e) 3mg/L preformed monochloramine, (f) 5mg/L preformed monochloramine, (g) 1mg/L chlorine-ammonia process, (h) 3mg/L chlorine-ammonia process, and (i) 5mg/L chlorine-ammonia process.

In the samples spiked with ammonium sulfate (Figure 28a and b), the sample not subjected to ozonation did not vary from the unperturbed sample water significantly (±0.003cm-1). However, when subjected to ozone it becomes apparent that UV absorbance removal efficiency decreased at the lower wavelengths (240-250nm) for the ozone doses of 0.25 and 0.5mg O3:mg TOC. This is in contradiction to previous data collected through OH• exposure calculations: the addition of free ammonia in these results showed an increase in advanced oxidation, and an increase in UV reduction should have been observed near the 254nm wavelength (Gerrity et al. 2012).

The addition of 1mg/L, 3mg/L and 5mg/L preformed monochloramine (Figure 28d, e, and f, respectively) illustrated a reduction in UV absorbance with a peak at 280nm for non-ozonated sample water. This is most likely a relic from the testing protocol, as the monochloramine residual was not quenched via sodium thiosulfate to prevent interference, and the sample water’s monochloramine residual was allowed to fully decay to prevent positive inference from monochloramine itself. This residual monochloramine was therefore able to reduce UV absorbance through slow oxidation reactions over its decay. This phenomenon also impacted the 3mg/L and 5mg/L chlorine-ammonia process samples, as background monochloramine was present, though more reduction in absorbance was noted due to free chlorine oxidation (Figure 28g, h, and i).

For the various ozone doses studied, less reduction in UV absorbance was noted with the addition of preformed monochloramine. The addition of 3mg/L and 5mg/L preformed monochloramine showed a significant decrease in UV absorbance near the wavelengths of 240-260nm. This is indicative of a reduction in advanced oxidation through OH• exposure (Gerrity et al. 2012).

The apparent UV reduction in monochloramine samples at 280nm did not change from the background or free ammonia samples. These impacts were nearly identically mimicked by the chlorine-ammonia process, when observing their monochloramine residuals (1mg/L Cl2-NH3 had no monochloramine, 3mg/L Cl2-NH3 had 1mg/L monochloramine, and 5mg/L Cl2-NH3 had 3mg/L monochloramine). It can therefore be hypothesized that the reduction in OH• exposure does not affect the change in UV absorbance reduction at 280nm, and this measurement may be attributed to ozone oxidation alone.