Predicting the Contribution of Chloramines to Contaminant Decay

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Remediation and Control Technologies

Predicting the Contribution of Chloramines to Contaminant Decay during UV/Hydrogen Peroxide Advanced Oxidation Process (AOP) Treatment for Potable Reuse Zhong Zhang, Yi-Hsueh Chuang, Nan Huang, and William A. Mitch Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.8b06894 • Publication Date (Web): 19 Mar 2019 Downloaded from http://pubs.acs.org on March 21, 2019

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Environmental Science & Technology

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Predicting the Contribution of Chloramines to Contaminant Decay during UV/Hydrogen Peroxide Advanced Oxidation Process (AOP) Treatment for Potable Reuse

12 13 14 15 16 17

Zhong Zhang1, Yi-Hsueh Chuang1, Nan Huang2, and William A. Mitch1,*

18 19 20 21 22 23

1

Department of Civil and Environmental Engineering, Stanford University, 473 Via Ortega, Stanford, California 94305, United States

24 25 26 27

2

Environmental Simulation and Pollution Control State Key Joint Laboratory, State Environmental Protection Key Laboratory of Microorganism Application and Risk Control (SMARC), School of Environment, Tsinghua University, Beijing, 100084, PR China

28 29 30 31 32

*Contact Information: email: [email protected], Phone: 650-725-9298, Fax: 650-723-7058

33 34 35 36 37 38 39 1

ACS Paragon Plus Environment


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ABSTRACT

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Chloramines applied to control membrane biofouling in potable reuse trains pass through

43

reverse osmosis membranes, such that downstream UV/H2O2 advanced oxidation processes

44

(AOPs) are de facto UV/H2O2-chloramines AOPs.

45

which use inaccurate chloramine quantum yields and ignore the fate of •NH2, are unable to

46

simultaneously predict the loss of chloramines and contaminants, such as 1,4-dioxane. This

47

study determined quantum yields for NH2Cl (0.35) and NHCl2 (0.75). Incorporating these

48

quantum yields and the formation from •NH2 of the radical scavengers, NO and NO2-, was

49

important for simultaneously modeling the loss of chloramines, H2O2 and 1,4-dioxane in the

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UV/H2O2-chloramines AOP. Although radical production was higher from the UV/H2O2-

51

chloramines AOP than the UV/H2O2 AOP, the UV/H2O2 AOP was at least two-fold more

52

efficient for 1,4-dioxane degradation, because chloramines efficiently scavenged radicals. At

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low chloramine concentrations, the UV/chloramines AOP efficiency increased with increasing

54

chloramine concentration, as radical production increased relative to radical scavenging by the

55

dissolved organic carbon (DOC) in RO permeate. However, the efficiency leveled out at higher

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chloramine concentrations as radical scavenging by chloramines offset increased radical

57

production.

58

the UV/H2O2-chloramines AOP when the residual chloramines in RO permeate were ~50 M

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(3.3 mg/L as Cl2). Initial cost estimates indicate that the UV/chloramines AOP using the

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residual chloramines in RO permeate could be a cost-effective alternative to the current

61

UV/H2O2-chloramines AOP in some cases, because the savings in reagent costs offset the ~30-

Current models for UV/chloramine AOPs,

1,4-Dioxane degradation was ~30-50% lower for the UV/chloramines AOP than

2


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50% reduction in 1,4-dioxane degradation efficiency.

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3



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INTRODUCTION

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Purified municipal wastewater has been increasingly emphasized as a local, reliable source

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water for drinking water plants.1,2 Potable reuse facilities frequently employ Full Advanced

67

Treatment (FAT) trains, typically including microfiltration (MF), reverse osmosis (RO), and

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the UV/hydrogen peroxide advanced oxidation process (UV/H2O2 AOP), wherein H2O2 is

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photolyzed to produce hydroxyl radical (•OH; equation 1) to destroy contaminants passing

70

through RO membranes.2 The ability to remove 0.5-log of 1,4-dioxane, an ingredient in

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chlorinated solvents and personal care products,3,4 has served as a metric to validate AOP

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performance.5 1,4-Dioxane is poorly rejected by RO membranes, because of its neutral charge

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and low molecular weight.

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H2O2 + hν → 2 •OH

75

Chloramines are commonly applied upstream of MF to control biofouling6, and readily

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pass through RO membranes because of their low molecular weight and neutral charge.7 While

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monochloramine (NH2Cl) is the predominant chloramine species upstream of RO, two factors

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promote dichloramine (NHCl2) formation. First, when chloramines are formed by concentrated

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free chlorine injection into ammonia-containing wastewater, the high chlorine:ammonia ratio

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at the injection point promotes NHCl2 formation.8 Moreover, the low pH in RO permeate (~5.5)

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and rejection of NH4+ by RO promote the conversion of NH2Cl to NHCl2 (equation 2).9

82

Accordingly, many UV/H2O2 AOP systems are de facto combinations of UV/H2O2, UV/NH2Cl

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and UV/NHCl2 AOPs.

(1)

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NH2Cl + NH2Cl + H+ ↔ NHCl2 + NH4+

85

Studies have only recently attempted to characterize the formation of radicals from the

(2)

4


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UV/chloramines AOP.10,11 Monochloramine absorbs UV light more efficiently than H2O2

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(εNH2Cl,254 nm = 371 M-1 cm-1 vs. εH2O2,254 nm = 18.6 M-1 cm-1),10,13 but its photolysis produces the

88

amidogen (•NH2) and chlorine (•Cl) radicals (equation 3).14,15 Chlorine radical can degrade

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contaminants directly or form •OH via equations 4-6.10,16-18 Chuang et al.10 indicated that the

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UV/NH2Cl AOP efficiently produces radicals and exhibits comparable performance to the

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UV/H2O2 AOP regarding 1,4-dioxane degradation. Accordingly, using only the chloramines in

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the RO permeate for the UV/chloramines AOP could avoid H2O2 addition and the chlorine

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demand required to quench the residual H2O2 prior to leaving a chloramine residual for

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distribution. More recently, Patton et al.12 studied the impact of chloramines on the UV/H2O2

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AOP. Although chloramines alone can produce radicals upon photolysis, the inclusion of

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chloramines in the UV/H2O2 AOP decreased 1,4-dioxane removal efficiency due to scavenging

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of UV photons and •OH by NH2Cl (kNH2Cl, • OH = 1.02 × 109 M-1 s-1).10,12

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suggested quenching the chloramine residual upstream of the UV/H2O2 AOP.

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concentrations of 1,4-dioxane (5-250 µM) and chloramines (0.2-6 mM) applied in previous

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studies10-12 were much higher than those relevant to RO permeate (~0.05 µM (4 g/L) 1,4-

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dioxane and 50 µM (3.5 mg/L as Cl2) chloramines). Moreover, the experiments were conducted

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in deionized water, hindering an assessment of the benefits or drawbacks of chloramines under

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realistic conditions.

Thus, they However, the

104

NH2Cl + hν → •NH2 + •Cl

(3)

105

•Cl

+ OH－↔ ClOH•－

(4)

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•Cl

+ H2O ↔ ClOH•－ + H+

(5)

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ClOH•－ ↔ •OH + Cl－

(6) 5



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108

Kinetic modeling is essential for understanding the net impact of increased radical

109

production and photon and radical scavenging by chloramines on the performance of the

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UV/H2O2-chloramines AOP. There are critical data gaps for the accurate modeling of the

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UV/NH2Cl AOP, and models have not yet incorporated the potentially important contributions

112

from the UV/NHCl2 AOP (equation 7). For example, there is significant uncertainty regarding

113

the quantum yields for photolysis of NH2Cl (0.26 – 0.62)14,15,19,20 and NHCl2 (0.82, 1.8)15,19.

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These quantum yields typically were calculated from the observed loss of chloramines, and thus

115

include decay by both direct photolysis (equations 3 and 7) and reactions with daughter radicals.

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Distinguishing direct chloramine photolysis from chloramine degradation by daughter radicals

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is important for accurate modeling of contaminant degradation because chloramines and target

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contaminants compete for reaction with these radicals. Chuang et al.10 used kinetic modeling to

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separate the direct photolysis from radical reactions and thereby improved the estimate of the

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NH2Cl quantum yield (0.20). Moreover, the subsequent reactions of •NH2 have been neglected

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due to its low reactivity towards contaminants,21 but these reactions may be important for AOP

122

modeling. Previous models have not been able to account for nitrite and nitrate formation as

123

products of chloramine decay.10-12 These products likely form from •NH2 reaction with O2, and

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one previous study showed that the photodegradation rate of NH2Cl is two times faster in the

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absence of oxygen.19 These results suggest that •NH2 reactions could significantly affect the

126

UV/NH2Cl AOP.

127

NHCl2 + hν → •NHCl + •Cl

128

The first objective of this study was to incorporate •NH2 and •NHCl reactions into the

129

UV/chloramines model to understand their importance. The second objective was to combine

(7)

6


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the updated model with experimental data to separate chloramine direct photolysis from their

131

degradation by radical reactions in order to obtain accurate quantum yields for NH2Cl and

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NHCl2 photolysis. The improved model was validated by its ability to simultaneously predict

133

the loss of oxidants (i.e., H2O2, NH2Cl and NHCl2), 1,4-dioxane decay and nitrite and nitrate

134

formation under scenarios relevant to potable reuse for both single (e.g., NHCl2 only) and mixed

135

oxidant (e.g., H2O2, NH2Cl and NHCl2) AOPs. This validation was conducted using both

136

deionized water and authentic RO permeate from a potable reuse facility. The third objective

137

was to combine the model with experiments to understand the contribution of chloramines to

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contaminant decay during UV/H2O2-chloramines AOP treatment of RO permeate and whether,

139

and under what conditions, the UV/chloramines AOP can serve as a suitable replacement.

140 141

MATERIALS AND METHODS

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Materials. Stock NH2Cl solutions (20 mM) were prepared daily by titrating 40 mM sodium

143

hypochlorite into 48 mM ammonium chloride (a 1:1.2 molar ratio) in deionized water adjusted

144

to pH 8 with sodium hydroxide, as described previously.22 Stock NHCl2 solutions were prepared

145

by maintaining the pH of a 3.5 mM NH2Cl solution near 3.7 using phosphoric acid and

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equilibrating for 1 hour. Then the solution was placed on ice with no headspace overnight to

147

convert NH2Cl to NHCl2 (equation 2). The total absorbances (A) at 245 nm and 295 nm were

148

measured using a Cary 60 UV-visible spectrophotometer with a 1 cm quartz cuvette. The

149

concentrations (C) of NH2Cl and NHCl2 were calculated by combining the total absorbances

150

with their molar absorption coefficients at these wavelength (εNH2Cl,245 = 445 M-1 cm-1, εNHCl2,245

151

= 208 M-1 cm-1, εNH2Cl,295 = 14 M-1 cm-1, εNH2Cl,245 = 267 M-1 cm-1) to solve the matrix 7



152

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encompassed by equations 8 and 9.22

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A245 = εNH2Cl,245CNH2Cll + εNHCl2,245CNHCl2l

(8)

154

A295 = εNH2Cl,295CNH2Cll + εNHCl2,295CNHCl2l

(9)

155

Here l is the path length (1 cm). The NH2Cl and NHCl2 stock solutions contained < 1% of the

156

other species. A H2O2 stock solution in deionized water was standardized by UV absorbance at

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254 nm (ε245 nm = 18.6 M-1 cm-1).13 Sources for other reagents are provided in Text S1.

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Grab samples of RO permeate were collected from a potable reuse facility (Facility 1)

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during two sampling events and an additional sample was collected from another potable reuse

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facility (Facility 2). Samples were maintained on ice with no headspace until experiments were

161

initiated, and experiments were conducted within 24 hours to minimize the evolution of

162

chloramine species. The concentrations of NH2Cl and NHCl2 in RO permeate were calculated

163

as described above after measuring the total absorbances at 245 nm and 295 nm using a 10 cm

164

quartz cuvette. General water quality parameters are provided in Table S1.

165

Experimental procedures. The UV absorbance of diluted NH2Cl or NHCl2 stock solutions

166

was measured at 254 nm to determine their molar absorption coefficients as 371 M-1 cm-1 and

167

136 M-1 cm-1, respectively. UV irradiation was applied using a semi-collimated beam apparatus

168

containing three 15 W Philips low pressure mercury lamps emitting at 254 nm, as described

169

previously.10,23 Briefly, light from the UV lamps shone down through a shutter onto a 750 mL

170

crystallization dish, which was continuously mixed by a magnetic stir bar. Incident irradiance

171

(0.60 mW cm-2) was determined by iodide-iodate actinometry.24,25 Experiments using either

172

deionized water or RO permeate were buffered with 2 mM phosphate at pH 5.5. In the

173

experiments evaluating the effect of dissolved oxygen, the deionized water was purged with 8


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nitrogen gas for 30 min and the solution was dispensed into a cylindrical dish with a quartz

175

cover for illumination. The direct photolysis of a compound (C) at 254 nm can be described by

176

equation 10 under conditions of minimal light absorbance (𝜀Cl < 0.02)26:

177

𝑑𝐶

(10)

－ 𝑑𝑡 ≅ 2.303Φ𝜀𝐼0𝐶

178

where Φ is the photolysis quantum yield in mol/Einstein, 𝜀 is the molar absorption

179

coefficient (M-1 cm-1), I0 is the incident light intensity (mEin cm-2 s-1) and l is the light

180

pathlength (cm). Except where noted, experiments were conducted using a 0.9-cm solution

181

depth to maintain 𝜀Cl < 0.02.

182

For experiments with only chloramines, aliquots were analyzed for total residual chlorine

183

by the DPD method.27 For experiments with only H2O2, H2O2 was measured by its oxidation of

184

DPD catalyzed by peroxidase enzyme.28 When H2O2 and chloramines were both present, one

185

aliquot was measured for residual H2O2 by the titanium oxalate method,29,30 while another was

186

measured for residual chloramines by the DPD method after first quenching residual H2O2 using

187

catalase.31 1,4-Dioxane was extracted into MtBE and analyzed by GC-MS (Text S1). Chloride,

188

nitrite and nitrate were measured using a Dionex DX-500 ion chromatography system. Total

189

nitrogen was measured using a Shimadzu TOC-L analyzer with a nitrogen detector. Ammonia

190

was measured using HACH method 10023. Dissolved oxygen (DO) was measured using a YSI

191

ProODO optical probe.

192

Kinetic modeling. A chemical kinetics model combining 94 elementary reactions obtained

193

from the literature, measured in this study, or estimated by analogy with similar reactions (Text

194

S2 and Table S3) was implemented using Kintecus 4.55.32 This model was based on our

195

previous model,10 with additional reactions, as discussed below. 9



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196 197

RESULTS AND DISCUSSION

198

Poor predictive capability of current UV/chloramines models: Initial experiments treated

199

0.2 M 1,4-dioxane at pH 5.5 in buffered deionized water with UV in the presence of 3.3 or

200

6.5 mg/L as Cl2 of preformed NH2Cl (47 M or 92 M) or NHCl2 (25 M or 44 M). The

201

concentrations of chloramines and 1,4-dioxane and the pH were considered similar to

202

conditions encountered during AOP treatment of RO permeate, but without hydrogen peroxide.

203

The degradation of chloramines and 1,4-dioxane both followed first-order kinetics (Figure S1).

204

As discussed below, the first-order degradation of 1,4-dioxane occurred even though at least 60%

205

of the chloramines degraded during the experiments, and chloramines were the source of the

206

radicals involved in 1,4-dioxane degradation. Figure 1 provides the pseudo-first-order UV

207

fluence-based degradation rate constants (kobs) observed for chloramines and 1,4-dioxane.

208

Figure 1 also provides the kobs values that would be predicted by two kinetic models. Patton

209

et al. presented a series of elementary reactions to describe the UV/NH2Cl and UV/NHCl2

210

AOPs, and highlighted quantum yields for chloramine photolysis obtained from previous

211

literature (0.54 for NH2Cl and 0.82 for NHCl2).11,12 We implemented this series of equations

212

(the 21 reactions in Table S2) as a model in Kintecus 4.55, but this model overestimated the

213

degradation rates of NH2Cl, NHCl2 and 1,4-dioxane by ~70%, ~25%, and ~300%, respectively

214

(Figure 1).

215

for the quantum yields for chloramine photolysis. These values were determined by assuming

216

that all of the experimentally measured chloramine loss was associated with direct photolysis.

217

Since the radicals produced by chloramine photolysis also contribute to chloramine degradation,

One factor contributing to the discrepancy was the use of these literature values

10


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these quantum yields overestimate the innate quantum yields associated with direct photolysis

219

of chloramines. When implemented in a kinetic model including daughter radical reactions, the

220

chloramine degradation rate and radical production rate would be overestimated, resulting in

221

overestimation of the 1,4-dioxane degradation rate.

222

A second model for the UV/NH2Cl AOP by Chuang et al.10 supplemented the elementary

223

steps in the Patton et al.12 model with additional elementary steps for radical reactions obtained

224

from the literature (70 reactions in Table S3). Experimental kobs values for NH2Cl decay were

225

fit to equation 11, where the first term accounts for direct photolysis of NH2Cl. The second term

226

accounts for NH2Cl degradation by reactions with its daughter radicals (•R), as described within

227

the Kintecus model. An optimal quantum yield of 0.20 was obtained from fitting the data. The

228

ability of this model to predict the experimental kobs values for degradation of NH2Cl and 1,4-

229

dioxane was better (Figure 1), but still underpredicted NH2Cl loss by up to 40%. The

230

discrepancy suggested the need to refine the model. In particular, we considered the need to

231

consider the fate of the •NH2 or •NHCl radicals.

232

𝑑𝐶

― 𝑑𝑡 = 𝑘𝑜𝑏𝑠𝐶 = 2.303Φ𝐼0𝜀𝐶 + 𝑘 ∙ 𝑅[ ∙ 𝑅]𝐶

11


(11)


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233

Importance of the •NH2 radical: Chloramine photolysis produces •Cl and •NH2 or •NHCl

234

radicals (equations 3-7). The previous models focused on the fate of •Cl and its daughter radicals

235

(i.e., Cl2•- and •OH). The fate of •NH2 was not evaluated since •NH2 was considered significantly

236

less reactive with contaminants.21 Previous models also did not include reactions that could

237

explain the formation of nitrite and nitrate, which are significant products of NH2Cl and NHCl2

238

photolysis.12,15,19,33 Reactions of •NH2 and •NHCl with dissolved oxygen would form NH2O2•

239

or NHClO2•, respectively (e.g., equation 12).34,35 In our experiments, the kobs values for NH2Cl

240

and NHCl2 photodecomposition were 1.38 and 1.27 times faster in oxygen-free solutions ([O2]

241

< 0.1 mg/L) than in aerated solutions ([O2] ≈ 8 mg/L) (Figure S2), concurring with previous

242

results by De Laat et al.19 regarding the effect of O2. De Laat et al.19 suggested that •NH2

243

scavenging by O2 reduced the observed NH2Cl degradation rate by inhibiting the reaction

244

between •NH2 and NH2Cl. Although the reaction rate constant between •NH2 and NH2Cl is

245

unknown, this explanation seems unlikely given the low reactivity of •NH2.21

246

•NH

247

As an alternative, we incorporated into our kinetic model additional elementary reactions

248

from the literature describing the fate of NH2O2•.34,35 NH2O2• forms nitric oxide (•NO), which

249

reacts rapidly with •OH (1.0 x 1010 M-1 s-1)21 to form nitrite (equations 13 and 14). Nitrite is

250

also a potent •OH scavenger (1.2 x 1010 M-1 s-1),21 forming •NO2 (equation 15). Additional

251

reactions lead to the formation of nitrate (equations 16-20). NH2O2• has also been proposed to

252

decay by a poorly characterized pathway to form N2O (equation 21).15 N2O is not a potent •OH

253

scavenger.36 Besides explaining nitrite and nitrate formation, these reactions might account for

254

the O2 effect, because scavenging of •OH by •NO and NO2- would reduce the NH2Cl

255

degradation rate by preventing the rapid reaction of NH2Cl with •OH (1.02 x 109 M-1 s-1).10 An

256

analogous series of reactions likely occurs with •NHCl.15

2

(•NHCl) + O2 ↔ NH2O2• (NHClO2•)

12


(12)

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NH2O2• (NHClO2•) → •NO + H2O (HOCl)

(13)

258

•NO

(14)

259

NO2－ + •OH → •NO2 + OH-

(15)

260

2 •NO + O2 → 2 •NO2

(16)

261

•NO

(17)

262

N2O3 + H2O → 2 NO2－ + 2 H+

(18)

263

2 •NO2 ↔ N2O4

(19)

264

N2O4 + H2O → NO2－ + NO3－ + 2 H+

(20)

265

NH2O2• (NHClO2•) → transient species → N2O

(21)

266

We evaluated the importance of these reactions for modeling the UV/chloramines AOP.

267

Reaction rate constants for most of these elementary steps have been determined (Table S3).10,21

268

However, the rate constant for the reaction of •NH2 with O2 has been estimated at 1.2 × 108 M-1

269

s-1 or higher.34,35 The only other unknown rate constants are those for the reaction of •NHCl

270

with O2, and for the decay of NH2O2• and NHClO2• to •NO or N2O. We assumed that the rate

271

constants for the reactions of •NH2 and •NHCl with O2 were 1.2 × 108 M-1 s-1, and that the rate

272

constants for the formation of •NO were 1.0 × 108 s-1. These assumptions were based upon a

273

sensitivity analysis indicating no change in the predicted loss of NH2Cl and formation of •NO

274

when these rate constants were varied between 1.2 × 108 M-1 s-1 and 1.2 × 1010 M-1 s-1 and

275

between 1.0 × 103 M-1 s-1 and 1.0 × 109 M-1 s-1, respectively; •NH2 and •NHCl formation from

276

NH2Cl and NHCl2 photolysis were rate-limiting. What mattered were the relative rate constants

277

controlling the branching pathway for formation of •NO or N2O from NH2O2• or NHClO2• (e.g.,

278

equations 13 and 21). To estimate these branching ratios, 100 M of NH2Cl or NHCl2 was

+ •OH → NO2－ + H+

+ •NO2 ↔ N2O3

13



279

treated with 2160 mJ/cm2 UV fluence, sufficient to degrade at least 90% of the chloramines.

280

The solutions were analyzed for nitrite and nitrate. The solutions were also purged gently with

281

nitrogen for 30 min and analyzed for total nitrogen. The difference between total nitrogen

282

measured in the solution before photolysis and after exposure to 2160 mJ/cm2 and nitrogen

283

purging was assumed to represent the N2O that had volatilized from the solution (Figure S3).

284

The ratios of the calculated N2O to the sum of nitrite and nitrate during photolysis of each

285

chloramine species were used to estimate rate constants for the production of N2O from NH2O2•

286

or NHClO2• of 6.0 x 108 s-1 and 6.7 x 108 s-1, relative to the 1.0 × 108 s-1 assumed for the

287

formation of •NO.

288

Experiments were conducted to determine the kobs values for the decay of NH2Cl (10-50

289

M) and NHCl2 (4-50 M) in the presence and absence of acetate at 50-fold higher

290

concentration as a radical quencher. The enhanced kinetic model (94 reactions in Table S3) was

291

combined with the experimental kobs values within the framework of equation 11 to determine

292

optimized quantum yields of 0.35 and 0.75 for NH2Cl and NHCl2, respectively (Text S3). Using

293

the optimized quantum yields, the model was able to predict experimental NH2Cl and NHCl2

294

loss over the range of chloramine concentrations (4 – 50 M) with and without the quencher

295

(Figure S5).

296

Additional experiments evaluated the predictive capability of the model during treatment

297

by the UV/NH2Cl or UV/NHCl2 AOPs. Although the quantum yields were developed solely by

298

fitting chloramine loss, the improved model effectively predicted NH2Cl or NHCl2 loss and

299

provided decent estimates for nitrite and nitrate formation during treatment of 100 M of these

300

chloramines in the absence of 1,4-dioxane (Figures 2A and 2B). The molar yields of nitrate 14


Page 14 of 37

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301

(9%) and nitrite (2%) we observed during UV photolysis of 100 M NH2Cl (Figure 2A) were

302

comparable in magnitude to the yields of nitrate (3.2%) and nitrite (6.5%) observed by Watts

303

and Linden33 in unbuffered deionized water; the preferential formation of nitrate over nitrite

304

that we observed at pH 5.5 concurred with the preference for nitrate over nitrite as the pH

305

decreased from 8.5 to 6.5 observed by Li and Blatchley.15 Although the predicted nitrite was

306

not observed during NHCl2 treatment, nitrate was well-modeled. Note that we assumed that

307

•NHCl

308

account for the discrepancy regarding nitrite, further research is needed to evaluate the reaction

309

pathways associated with •NHCl. Regardless, nitrite formation is a minor pathway, such that

310

the predicted level of nitrite was low (60% loss of

366

chloramines (Figure S1); while chloramine loss reduces the radical production rate, the 17



Page 18 of 37

367

reduction in chloramine concentration simultaneously decreases a dominant source of radical

368

scavenging. To compare the efficacy of the different AOPs, this study therefore used 0.2 µM

369

1,4-dioxane to approach conditions relevant to potable reuse.

370

Production rate

2.303ΦεI0[H2O2] k2[1,4 - dioxane] + ∑𝑘𝑖𝑠𝑖

[•OH]ss = Scavenging rate ≅ k1[H2O2] +

(22)

371

We compared the degradation of 0.2 M 1,4-dioxane by the UV/H2O2, UV/chloramines

372

and UV/H2O2-chloramines AOPs in two RO permeate samples collected from a wastewater

373

reuse facility (Facility 1 in Figure 4). The background chloramine concentrations were 28 M

374

NH2Cl and 9.4 M NHCl (47 M Cl[+1] chloramines) during Event 1, and 43 M NH2Cl and

375

15 M NHCl (73 M Cl[+1] chloramines) during Event 2. The UV/H2O2 AOP, where

376

background chloramines in the RO permeate were quenched by a stoichiometric addition of

377

sulfite before adding 100 M H2O2, showed the highest 1,4-dioxane degradation rate during

378

both events. The kobs for 1,4-dioxane for the UV/H2O2-chloramines AOP, where 100 M H2O2

379

was added without quenching the chloramines, was ~50% lower than for the UV/H2O2 AOP

380

during both events. However, for the UV/chloramines AOP, using only the background

381

chloramines in the samples, the kobs were similar for the two events, despite the different

382

background chloramine concentrations, and only 35% (Event 1) and 30% (Event 2) lower than

383

for the UV/H2O2-chloramines AOP. Indeed, the addition of 27 M preformed NH2Cl to the

384

background chloramines during Event 2 (100 M Cl[+1] chloramines total) did not

385

significantly increase the kobs for the UV/chloramines AOP (Figure 4). These results suggest

386

that the performance of the UV/chloramines AOP could nearly equal that of the UV/H2O2-

387

chloramines AOP, the AOP relevant to current practice at reuse facilities.

388

The model accurately predicted 1,4-dioxane degradation in authentic RO permeate 18


Page 19 of 37


389

during treatment by the UV/chloramines (2%-11% error) or UV/H2O2-chloramines (7%-22%

390

error) AOPs when chloramine concentrations were ≥ 47 M (Figure 4). However, the model

391

overestimated the 1,4-dioxane degradation rate by ~90% for the UV/H2O2 AOP. We

392

hypothesized that the high predictive capability for the AOPs containing chloramines arose

393

from chloramines serving as the predominant radical scavengers, which is relatively easy to

394

model.

395

radical scavengers.

396

total carbonate species were around 0.5 mM, consisting of 0.42 mM H2CO3 and 0.08 mM

397

HCO3- at pH 5.6. We determined rate constants for •OH reaction with DOC in RO permeate to

398

be 2.49 (±0.11)×104 (mg-C/L)-1 s-1 (average ± range of experimental duplicates) for the sample

399

from Facility 1 Event 1 and 2.12 (±0.009)×104 (mg-C/L)-1 s-1 for the sample from Facility 1

400

Event 2 (Text S4). These rate constants are lower than the 3.33 × 104 (mg-C/L)-1 s-1 rate

401

constant we had determined previously for the DOC in RO concentrate,38 suggesting that RO

402

treatment favors passage of the less reactive portions of DOC. Figure 4 provides the kobs

403

predicted after considering •OH scavenging by DOC and carbonates,

404

total production of radicals, differentiated according to their fate. After incorporating

405

scavenging by DOC and carbonates, the model predicted loss of oxidants within 6-28% and

406

1,4-dioxane degradation within 2-31% under all conditions, including the UV/H2O2 AOP.

407

Attempts to model nitrite and nitrate formation were unsuccessful because the concentration of

408

nitrate present in the RO permeate samples (1.16 mg/L as N (83 M) in Facility 1 and 0.95

409

mg/L as N (68 M) in Facility 2 (Table S1)) was much higher than the 0.9 cm.

453

When αl > 0.02, the direct photolysis rate of a compound (C) can be described by equation 23:

454

𝑑𝐶

― 𝑑𝑡 =

εΦ𝐼0(1 ― 10 ―𝛼𝑙) 𝛼𝑧

𝑐 21


(23)


455

where z is the solution depth (cm), which equals the pathlength during illumination with a

456

collimated beam.

457

During Event 1 for Facility 1, the UV254 was 0.0149 cm-1, where chloramines accounted for

458

79% of the total absorbance. As the solution depth increased from 0.9 cm to 3.6 cm, αl increased

459

from 0.0134 to 0.0536. During UV/chloramines AOP treatment with just the background

460

chloramines, kobs for chloramine decay decreased by 14%, while the kobs for 1,4-dioxane decay

461

decreased by 18% as the depth increased from 0.9 cm to 3.6 cm (Figures 5 and S11). When 100

462

M H2O2 was added, the kobs for 1,4-dioxane decay also decreased by 21% as the depth

463

increased from 0.9 cm to 3.6 cm (Figures 5 and S11). As a result, the UV/chloramines AOP

464

remained only 24% less efficient than the UV/H2O2-chloramines AOP. The UV/H2O2 AOP was

465

less affected by solution depth due to the low molar absorption coefficient of H2O2 (18.6 M-1

466

cm-1).13

467

To evaluate the potential for performance variations between RO permeate matrices, a RO

468

permeate sample collected from a second reuse facility (Facility 2), containing 35 M NH2Cl

469

and 9 M NHCl2 (53 M Cl[+1] chloramines), was treated by different AOPs while controlling

470

the depth at 3.6 cm. Again, the UV/H2O2 AOP exhibited the highest kobs for 1,4-dioxane

471

degradation (0.0033 cm2/mJ; Figure S12), comparable to that observed at Facility 1 for a 0.9

472

cm depth (Figure 4). The kobs for the UV/H2O2-chloramines AOP was also similar to that

473

observed in the sample from Facility 1 Event 1 for 47 M chloramines and a 3.6 cm depth

474

(Figure 5). However, the kobs for the UV/chloramines AOP was half that of the UV/H2O2-

475

chloramines AOP for this facility (0.00056 cm2/mJ vs. 0.0011 cm2/mJ), and lower than the

476

0.00082 cm2/mJ value observed for 47 M chloramines alone at the 3.6 cm depth at Facility 1. 22


Page 22 of 37

Page 23 of 37


477

At 20 M chloramines, the kobs for the UV/chloramines AOP was also lower than at Facility 1

478

(0.00042 cm2/mJ vs. 0.00067 cm2/mJ). Given the similarity of the kobs for the UV/H2O2 AOP

479

between the two facilities, the differences observed for the UV/chloramines AOP suggest that

480

the RO permeate from the second facility contained scavengers of •Cl that effectively inhibited

481

its conversion to •OH, thereby reducing 1,4-dioxane degradation.

482

Implications: The high molar absorption coefficients of chloramines, coupled with their

483

elevated quantum yields proposed in some previous research14,15,19,20 suggested that the

484

UV/chloramines AOP could be highly effective for RO permeate treatment. We developed an

485

improved kinetic model that can simultaneously predict the loss of H2O2, chloramines and 1,4-

486

dioxane in RO permeate, and used the model to develop better estimates of the NH2Cl and

487

NHCl2 quantum yields. Combining experimental data with our improved model, we

488

demonstrated that the UV/H2O2 AOP with 100 M (3.4 mg/L) H2O2 is roughly two-fold more

489

efficient for 1,4-dioxane degradation than the UV/H2O2-chloramines AOP, the de facto AOP

490

currently employed at many potable reuse facilities. For the UV/chloramines AOP, the 1,4-

491

dioxane degradation rate increased as the background chloramine concentration in RO permeate

492

increased from 20 M (1.4 mg/L as Cl2) to 47 M (3.3 mg/L as Cl2), but leveled off at higher

493

chloramine concentrations as the increase in radical production rate was balanced by

494

scavenging from chloramines. However, since chloramine photolysis dominated radical

495

production, the 1,4-dioxane degradation rate with 47 M total chloramines was ~30-50% lower

496

than when 100 M H2O2 was added, depending on the facility. These results suggested the need

497

to compare the efficiency gain associated with H2O2 addition against the added reagent costs.

498

We conducted initial operating cost estimates for two potential alternatives to the current 23



499

UV/H2O2-chloramines AOP (Text S5). We targeted 0.5-log removal of 1,4-dioxane with a 3.6

500

cm light pathlength. For the UV/H2O2-chloramines reference AOP, this target requires ~1000

501

mJ/cm2 UV fluence for 100 M H2O2 and 50 M chloramines, leaving ~90 M H2O2 and 20

502

M chloramine residuals (Figure 3; Text S5). We also targeted a 35 M (2.5 mg/L as Cl2)

503

chloramine residual for final distribution. The first alternative would be to switch to the

504

UV/H2O2 AOP by injecting sodium bisulfite upstream of the AOP to quench the chloramine

505

residual prior to H2O2 addition, thereby reducing the UV fluence needed to degrade 1,4-dioxane

506

to 360 mJ/cm2 (Text S5). The cost estimates considered the cost of bisulfite and the additional

507

chlorine and ammonia to replace the 20 M chloramine residual quenched by bisulfite to reach

508

the 35 M chloramine residual target. The initial cost estimate (Text S5) indicates that these

509

chemical costs ($13.22/million liters (ML)) exceed the electricity savings ($4.06/ML), for a net

510

cost increase of $9.15/ML. It may be difficult to dose the bisulfite to match the residual

511

chloramines. Excess bisulfite would quench a portion of the H2O2 added downstream. Under-

512

dosing would leave a partial chloramine residual, potentially negating the benefits of

513

chloramine quenching.

514

due to the need for higher fluence to degrade N-nitrosamines.

Moreover, it may not be possible to reduce the fluence to 360 mJ/cm2

515

The second alternative would be to avoid H2O2 addition, relying solely on the background

516

chloramines in the RO permeate. To achieve the same level of 1,4-dioxane removal as the

517

UV/H2O2-chloramines AOP, this would increase the UV fluence to 1400 mJ/cm2 at the first

518

facility and 2050 mJ/cm2 at the second facility, increasing the electricity costs by $2.54/ML and

519

$6.67/ML, respectively. The cost estimates also considered the reagent savings by avoiding the

520

cost of H2O2 addition and chlorine addition to quench the residual H2O2 ($14.67/ML), but also 24


Page 24 of 37

Page 25 of 37


521

the costs of any sodium hypochlorite and ammonium sulfate needed to replace the additional

522

chloramines degraded at the higher fluence. Altogether there would be a net savings of

523

$12.13/ML at Facility 1, and $4.44/ML at Facility 2 (Text S5). However, for utilities with low

524

chloramine residuals in their RO permeates, additional chloramines should be added upstream

525

of the AOP to reach ~50 M. For a utility with a 20 M chloramine residual in the RO permeate,

526

the increased costs associated with boosting the chloramine residual would be $11.73/ML. Thus

527

the net savings at Facility 1 would be $0.41/ML, but there would be a net increase in costs of

528

$7.29/ML at Facility 2 (Text S5). This alternative also simplifies the treatment train by

529

removing one reagent (hydrogen peroxide) and the additional fluence could also increase

530

disinfection credits and removal of contaminants such as N-nitrosamines. The present bench-

531

scale study employed environmentally-relevant 1,4-dioxane concentrations in authentic RO

532

permeate. However, the bench-scale UV reactor required ~30 mins to achieve 1000 mJ/cm2

533

fluence, compared to ~30 seconds in a full-scale system. It is possible that the prolonged

534

irradiation time heightened the importance of concurrent dark reactions, suggesting the need

535

for additional pilot-scale testing.

536 537

25



538

Acknowledgements: This work was supported by funding from the Water Research

539

Foundation (Reuse-16-01) and the National Science Foundation Engineering Research Center

540

for Re-Inventing the Nation’s Urban Water Infrastructure (ReNUWIt, EEC-1028968).

541

Supporting Information Available: Additional materials and methods; basic water quality;

542

model descriptions; procedure to determine quantum yields; modeled predictions for radicals;

543

determination of hydroxyl radical reaction rate constants with NHCl2 and DOC; additional

544

results for effect of solution depth; initial cost estimates.

545 546

References

547

1.

548

through Reuse of Municipal Wastewater; The National Academies Press: Washington, DC.

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throughout the world. J Water Supply Res Technol. 2013, 62, 321-338.

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National Research Council, 2012. Water Reuse: Expanding the Nation’s Water Supply

Gerrity, D.; Pecson, B.; Trussell, R. S.; Trussell, R. R., Potable reuse treatment trains

Kwon, K. D.; Jo, W. K.; Lim, H. J.; Jeong, W. S., Volatile pollutants emitted from selected

Steinemann, A. C.; Gallagher, L. G.; Davis, A. L.; MacGregor, I. C., Chemical emissions

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Chloramine Speciation Causing Re-Formation of NDMA during Potable Reuse of Wastewater.

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10. Chuang, Y. H.; Chen, S.; Chinn, C. J.; Mitch, W. A., Comparing the UV/Monochloramine

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and UV/Free Chlorine Advanced Oxidation Processes (AOPs) to the UV/Hydrogen Peroxide

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AOP Under Scenarios Relevant to Potable Reuse. Environ Sci Technol 2017, 51, 13859-13868.

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11. Patton, S.; Li, W.; Couch, K. D.; Mezyk, S. P.; Ishida, K. P.; Liu, H. Z., Impact of the

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Ultraviolet Photolysis of Monochloramine on 1,4-Dioxane Removal: New Insights into Potable

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12. Patton, S.; Romano, M.; Naddeo, V.; Ishida, K. P.; Liu, H., Photolysis of Mono- and

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Dichloramines in UV/Hydrogen Peroxide: Effects on 1,4-Dioxane Removal and Relevance in

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13. Morgan, M. S.; Vantrieste, P. F.; Garlick, S. M.; Mahon, M. J.; Smith, A. L., Ultraviolet

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15. Li, J.; Blatchley, E. R., UV Photodegradation of Inorganic Chloramines. Environ Sci

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16. Nowell, L. H.; Hoigne, J., Photolysis of Aqueous Chlorine at Sunlight and Ultraviolet

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photolysis as an advanced oxidation process for drinking water treatment. Environ Sci-Wat Res

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18. McElroy, W. J., A Laser Photolysis Study of the Reaction of SO4- with Cl- and the

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Subsequent Decay of Cl2- in Aqueous Solution. J Phys Chem 1990, 94, 2435-2441.

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19. De Laat, J.; Boudiaf, N.; Dossier-Berne, F., Effect of dissolved oxygen on the

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photodecomposition of monochloramine and dichloramine in aqueous solution by UV

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irradiation at 253.7 nm. Water Res 2010, 44, 3261-3269.

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20. Cooper, W. J.; Jones, A. C.; Whitehead, R. F.; Zika, R. G., Sunlight-induced photochemical

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21. NDRL/NIST Solution Kinetics Database. http://kinetics.nist.gov/solution/

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22. Schreiber, I. M.; Mitch, W. A., Influence of the order of reagent addition on NDMA

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formation during chloramination. Environ Sci Technol 2005, 39, 3811-3818.

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23. McCurry, D. L.; Krasner, S. W.; Mitch, W. A., Control of nitrosamines during non-potable

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and de facto wastewater reuse with medium pressure ultraviolet light and preformed

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monochloramine. Environ Sci-Wat Res 2016, 2, 502-510.

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24. Bolton, J. R.; Stefan, M. I.; Shaw, P. S.; Lykke, K. R., Determination of the quantum yields

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25. Goldstein, S.; Rabani, J., The ferrioxalate and iodide-iodate actinometers in the UV region.

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26. Schwarzenbach, R.P.; Gschwend, P.M.; Imboden, D.M. Environmental Organic Chemistry,

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Concentrations of Hydrogen Peroxide by the Peroxidase Catalyzed Oxidation of N,N-Diethyl-

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38. Yang, Y.; Pignatello, J. J.; Ma, J.; Mitch, W. A., Effect of matrix components on UV/H2O2

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Goldstein, S.; Rabani, J. Mechanism of nitrite formation by nitrate photolysis in

645 646 30


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Page 31 of 37


647

31



648 649 650 651 652 653

Figure 1. Experimental and modeled pseudo-first order observed fluence-based rate constants (kobs) for the degradation of 0.2 µM 1,4-dioxane (14D) and different concentrations of chloramines during treatment using the UV/NH2Cl and UV/NHCl2 AOPs in 2 mM phosphate buffer at pH 5.5. Error bars represent the range of experimental duplicates. The elementary reactions, rate constants and quantum yields applied in the model were obtained from Patton et al.11, 12 (model 1 results) and Chuang et al.109 (model 2 results).

32


Page 32 of 37

Page 33 of 37

654 655 656 657


Figure 2. Experimental vs. modeled results of UV/chloramine for (a) 100 µM NH2Cl loss and NO2- and NO3- formation, (b) 100 µM NHCl2 loss and NO2- and NO3- formation, (c) decay of 46 µM NH2Cl and 0.2 µM 1,4-dioxane, (d) decay of 45 µM NHCl2 and 0.2 µM 1,4-dioxane in 2 mM phosphate buffer at pH 5.5. Error bars represent the range of experimental duplicates.

658 659 660

33



661 662

663 664 665 666

Figure 3. Experimental vs. modeled results for the UV/H2O2-chloramines AOP containing 28 µM NH2Cl, 9.4 µM NHCl2, 98 µM H2O2 and 0.2 µM 1,4-dioxane in authentic RO permeate (Facility 1, Event 1). Samples were buffered at pH 5.5 with 2 mM phosphate. Error bars represent the data range of experimental duplicates.

667

34


Page 34 of 37

Page 35 of 37


668

669 670 671 672 673 674 675 676 677

Figure 4. Experimental and modeled fluence-based pseudo-first order rate constants (kobs) for loss of 0.2 µM 1,4-dioxane during various AOP treatments in authentic RO permeate collected during two sampling events. The samples contained 28 µM NH2Cl and 9.4 µM NHCl2 (Facility 1, Event 1) and 42.6 µM NH2Cl and 15 µM NHCl2 (Facility 1, Event 2). H2O2 indicates the addition of 100 M hydrogen peroxide. Error bars represent the data range of experimental duplicates. Modeled kobs values are provided with and without consideration of radical scavenging by DOC. Using the model considering radical scavenging by DOC, the total radicals produced for each AOP after 1000 mJ/cm2 fluence were distinguished by their fates.

678

35



679 680 681 682 683 684

Figure 5. The impact of light pathlength on pseudo-first order observed fluence-based rate constants for the degradation of 0.2 µM 1,4-dioxane and the residual chloramines in a RO permeate sample (Facility 1, Event 1), containing 28 µM NH2Cl and 9.4 µM NHCl2. H2O2 indicates the addition of 100 M hydrogen peroxide. Error bars represent the data range of experimental duplicates.

685 686 687 688

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Predicting the Contribution of Chloramines to Contaminant Decay

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