Perfluorooctanoic Acid Degradation Using UVâPersulfate Process

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Perfluorooctanoic Acid Degradation Using UV/Persulfate Process: Modeling of the Degradation and Chlorate Formation Yajie Qian, Xin Guo, Yalei Zhang, Yue Peng, Peizhe Sun, ChingHua Huang, Junfeng Niu, Xuefei Zhou, and John Crittenden Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.5b03715 • Publication Date (Web): 21 Dec 2015 Downloaded from http://pubs.acs.org on December 22, 2015

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Environmental Science & Technology

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Perfluorooctanoic Acid Degradation Using UV/Persulfate Process:

2

Modeling of the Degradation and Chlorate Formation

3 4

Yajie Qian†,‡, Xin Guo‡, Yalei Zhang†, Yue Peng‡, Peizhe Sun‡, Ching-Hua Huang‡,

5

Junfeng Niu§, Xuefei Zhou†,*, and John C. Crittenden‡,*

6 †

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200092, China

8 9

‡

School of Civil and Environmental Engineering, Georgia Institute of Technology, Atlanta, Georgia 30332, United States

10 11 12

School of Environmental Science and Technology, Tongji University, Shanghai

§

State Key Laboratory of Water Environment Simulation, School of Environment, Beijing Normal University, Beijing 100875, China

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* Corresponding Authors. Phone: 404-894-5676; Fax: 404-894-7896.

15

Email: [email protected] (John C. Crittenden). Phone: 86-21-65980624;

16

Fax: 86-21-65989961; E-mail: [email protected] (Xuefei Zhou).

17 18

Manuscript submitted to

19


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August 1, 2015

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(Manuscript words count: 7535)

1

ACS Paragon Plus Environment


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pH Simulation and Chlorate Simulation:

23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 2


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ABSTRACT

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In this study, we investigated the destruction and byproduct formation of

43

perfluorooctanoic acid (PFOA) using ultraviolet light and persulfate (UV/PS). Additionally,

44

we developed a first-principles kinetic model to simulate both PFOA destruction, and

45

byproduct and chlorate (ClO3-) formation in ultrapure water (UW), surface water (SW) and

46

wastewater (WW). PFOA degradation was significantly suppressed in the presence of

47

chloride and carbonate species and did not occur until all the chloride was converted to ClO3-

48

in UW and for low DOC concentrations in SW. The model was able to simulate the PS decay,

49

pH changes, radical concentrations and ClO3- formation for UW and SW. However, our model

50

was unable to simulate PFOA degradation well in WW, possibly from PS activation by NOM,

51

which in turn produced sulfate radicals.

52 53 54 55 56 57 58 59

3



60

INTRODUCTION

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PFOA has been widely found in surface waters in China, the U.S., Japan, and Europe.1-4

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High concentrations near fluorochemical plants in the U.S. have been found. For example,

63

about 500 ng·L-1 of this compound has been detected in the Tennessee and Ohio Rivers.5, 6

64

Various concentrations of PFOA were observed in sewage treatment plants in different states

65

of the U.S. including Georgia.7-9 PFOA has a high thermal and chemical stability and is

66

reported to be persistent in the environment and to bioaccumulate in biota.10, 11 Granular

67

activated carbon, ion exchange and membrane separation have been used to remove PFOA

68

from water.12-14 However, the residuals that contain PFOA still need further treatment to

69

destroy this persistent compound. Due to the strong electronegative property of fluorine (i.e.,

70

E0 (F/F-) = 3.6 V)15 and numerous tightly bonded fluorine atoms to carbon (i.e., 116

71

kcal·mol-1),16 PFOA is recalcitrant to conventional chemical oxidation treatment schemes.

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Hydroxyl radical (HO·) are effective for degrading most persistent organic contaminants.17, 18

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(For example, Luo et al. provided a promising application of Cu2O/TiO2 in p-nitrophenol

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degradation.19) However, HO· is not effective in destroying PFOA. In fact, the electron

75

withdrawing functional group -COOH in the structure of PFOA has been reported to be

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unreactive with HO·.20, 21

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Recent studies have demonstrated that the UV/persulfate (UV/PS) process produced

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sulfate radical (SO4-·) and SO4-· were shown to degrade PFOA.21 Indeed, SO4-· is a very

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electrophilic radical and can be more effective than HO· in degrading organic compounds.21-23

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Several published works on PFOA degradation by SO4-· have focused on the removal

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efficiency of PFOA and different activation methods for producing SO4-·.21, 24 Few studies 4


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however have reported either the kinetics of PFOA degradation or the nature of the oxidation

83

products generated in the UV/PS process. Many important rate constants for SO4-· reactions

84

with PFOA and the oxidation byproducts are still not available in the literature. In this study,

85

we report the rate constants that we have determined by fitting our data.

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The UV/PS process is considered a potential water treatment option. For example, it has

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been shown to degrade of a range of contaminants including 2,4-dichlorophenol and butylated

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hydroxyanisole.25, 26 Empirical kinetic models for the degradation of atenolol and bisphenol A

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by the UV/PS process have been developed based on the experimental observations.27, 28

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Additionally, the SO4-· and HO· distributions in the presence of Cl-, Br- and HCO3- have been

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modeled using the computer program Kintecus.29 Although these studies focused on the

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contaminant degradation by UV/PS, limitations are evident. First, contaminant degradation

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only was investigated in these previous studies with no attempt to develop kinetic models of

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the degradation process. Some studies did use empirical models to describe contaminant

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destruction but did not provide sufficient details of the UV/PS process. Second, while kinetic

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models have been developed in some studies, these models were based on pseudo-steady state

97

approximations and an assumption of constant pH despite the fact that pH change always

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occurs in the UV/PS process because protons are generated during PS activation.30, 31 Many

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studies demonstrated that this pH change plays an important role in radical distributions and

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the efficacy of the UV/PS process.32,

33

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SO4-· formation, because SO4-· reacts with HO- to generate HO· under basic conditions,34 and

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HO· is incapable of degrading PFOA. In this study, we develop a first-principles kinetic

103

model which incorporates these effects.

Acidic conditions were more favorable for

5



104

Advanced oxidation processes (AOPs) are promising technologies that destroy organic

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compounds. However, it is recognized that undesired byproducts may be produced during the

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treatment process.35-37 Bromate and chlorate are known disinfection byproducts (DBPs) in

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drinking water treatment. Bromate is currently regulated in U.S. drinking waters with a

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maximum contaminant level (MCL) of 10 µg·L-1.38 Chlorate is not currently regulated in

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drinking water though the U.S. EPA has recommended a chlorate health reference level (HRL)

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of 210 µg·L-1.39 The formation of byproducts in ozone-mediated AOPs has been widely

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studied.35, 40 To date, research regarding the formation of byproducts in the UV/PS process is

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still quite limited. Most researchers have focused on the influence of Cl- on the performance

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of the UV/PS process. Their work primarily focused on Cl-based radicals generation and the

114

role of these radicals in contaminant degradation.41, 42 Only one study reported the conversion

115

of Cl-based radicals to ClO3-. In that study, ClO3- formation under different pH values was

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examined with acidic pHs found to be more favorable for ClO3- formation.43 Indeed, with the

117

decomposition of PS, protons are generated in the system and this promotes the formation of

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ClO3-. Fang et al. investigated the BrO3- formation in UV/PS system and also found that BrO3-

119

formation was enhanced under acidic conditions.44 Fang et al. used the empirical

120

pseudo-first-order kinetics to fit the experimental data; and, because of this, the detailed

121

reaction rates are unknown for the formation of BrO3-. To elucidate the mechanism of ClO3-

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formation and provide a quantitative understanding of how it is formed in the UV/PS process,

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a model considering all the elementary reactions is developed here.

124

The objective of this study was to investigate the kinetics of PFOA destruction and

125

byproduct formation in the UV/PS system and to develop a model to describe process 6


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performance. Various water matrices were examined in order to determine the effects on

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PFOA degradation of constituents present in surface and waste waters. Based on our

128

experimental results and kinetic parameters that available in the literature, a first-principles

129

kinetic model was developed. The destruction of PS, pH change and radical concentrations

130

were modeled to improve the understanding of PFOA degradation in UV/PS process. The

131

formation of ClO3- in the presence of Cl- was also investigated and modeled. The applicability

132

of our model was verified by examining the degradation of PFOA in surface water and

133

secondary wastewater effluent samples. Finally, our model was used to minimize the

134

electrical energy per order (EE/O), by examining the impact of PS dosage and UV intensity.

135 136

MATERIALS AND METHODS

137

Chemicals. Sources of chemicals and reagents are provided in Text S1.

138

Water Sampling and Characterization. Samples of surface water (SW) were collected

139

from the Lake Lanier in Atlanta. Samples of treated wastewater (WW) were obtained from

140

secondary effluent prior to disinfection from a local municipal wastewater treatment plant.

141

The characteristics of the water samples are shown in Table S2. Details of the sample

142

collection and characterization are summarized in Text S2.

143

Analytical methods. Analytical details are provided in Text S3.

144

Experimental procedures. The reactor is illustrated in Figure S13. The lamp output

145

spectra peaks at 254 nm as measured using a UVX radiometer (SR-1100, Spectral Evolution,

146

MA). The average light intensity in the reactors was measured by potassium

147

tris(oxalato)ferrate(III) photoreduction (Hatchard-Parker actinometry) to be 2.88 × 10-7 7



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Einstein·L-1·s-1.45 The detailed description of the experimental procedures are provided in Text

149

S4.

150

Modeling Approaches. We used literature reported rate constants when we could find

151

reported values. The remaining values were obtained by fitting experimental data. The genetic

152

algorithm (GA) was used to minimize the objective function (OF) and determine the rate

153

constants.46 Accordingly, a reasonable range of minima and maxima for each rate constant

154

were provided to the GA. For the reactions which cannot be found in literature, the range was

155

evaluated based on the analogy with similar reactions which were available in the literature.

156

The backward differentiation formula (BDF) method [i.e., Gear’s method]47 was used to solve

157

all the ordinary differential equations (ODEs). Our model had 93 rate constants, 71 were from

158

the literature and 22 were determined by fitting the model to the data. The detailed modeling

159

approaches were provided in Text S5 and Table S1. When bicarbonate was not added to the

160

reaction feed, we assumed that CO2 (H2CO3*), bicarbonate (HCO3-) and carbonate (CO32-)

161

concentrations were estimated assuming that they were in equilibrium with atmospheric CO2.

162

The calculated concentrations of HCO3- and CO32- were 6.6×10-8 mol·L-1 and 4.2×10-14

163

mol·L-1, respectively. It is important to point out that the reaction bottles were sealed with a

164

cap at this point to prevent any additional dissolution of the CO2. And the initial carbonate

165

concentrations were essentially negligible in this instance. Their impact on PFOA degradation

166

was estimated by a quenching coefficient QR. 48

QR = 167

kmCm k R CR = 1− kmCm + kRCR kmCm + k RCR

(1)

168

where k R is the second-order rate constant for destruction of PFOA with SO4-·, M-1·s-1; km

169

is the second-order rate constant for scavengers (bicarbonate and chloride) in the water matrix 8


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with SO4-·, M-1·s-1; C R is the concentration of PFOA, M; Cm is the concentration of the

171

scavengers in the water matrices, M. The details were provided in Text S6. The calculated QR

172

for the case when carbonate species were in equilibrium with the atmosphere was 0.9939,

173

which means the quenching of the reaction rate was negligible. The concentrations of the

174

protonated and unprotonated species were considered in the model using the predicted pH and

175

the pKa values. We assumed equilibrium for all the carbonate species for a given pH because

176

the reaction rate for proton addition or loss is nearly instantaneous. Superoxide radical

177

(O2-· and HO2·) was not considered in this model because O2-· and HO2· cannot degrade

178

PFOA. (The pKa value is 4.8 for HO2·,49 which means the concentration of O2-· is extremely

179

low at acidic pH.) Due to the acidic condition, the reaction between SO4-· and HO- was not

180

included. Instead, the reaction between SO4-· and H2O was added in the model. The detailed

181

kinetic equations were provided in Text S9.

182 183

RESULTS AND DISCUSSION

184

PFOA Degradation and Byproducts Formation. Figure S1 shows the degradation of

185

PFOA under direct UV photolysis, PS dark oxidation and UV/PS combined oxidation. No

186

degradation of PFOA was observed for both dark and photolysis experiments. However,

187

significant degradation of PFOA was observed for UV/PS process. For example, 85.6% of

188

PFOA was destroyed after 8 hours. Thus, all of our ensuing work focused only on the PFOA

189

destruction mechanisms for the UV/PS process. It should be noted that the destruction of

190

PFOA took a long time in this work, because a low light intensity was used in our

191

experiments. In practice, the light intensities that are used can be 2 orders of magnitude higher 9



192

and will cut the reaction time by 2 orders of magnitude.48 To determine the role of HO· in

193

PFOA degradation, we compared PFOA degradation using UV/H2O2 and UV/PS processes.

194

For the same UV light intensity, no removal of PFOA was observed with 15 mM and even

195

higher H2O2 dosages (Figure S2). Thus, we assumed that PFOA degradation involved only

196

SO4-· and our kinetic model only considered SO4-· for the destruction of PFOA.

197

The pseudo-first order rate constants kobs for the destruction of PFOA are listed in Table

198

S3. The kobs increased about 41% (i.e., from 0.18 h-1 to 0.30 h-1) as PS concentration increased

199

from 5 mM to 15 mM. However, the kobs (i.e., 0.30 h-1) leveled off with further increase of PS

200

dosage due to the scavenging of SO4-· or by other species. To further elucidate the destruction

201

efficiency of PFOA, the F- and TOC were measured to calculate the defluorination and

202

mineralization of PFOA. The defluorination and mineralization increased with the increasing

203

PS dosage from 15 to 20 mM, but then decreased for larger PS concentrations (i.e., 30 mM).

204

One explanation might be that the excess SO4-· reacted with the intermediates of PFOA in the

205

range of 15 to 20 mM PS, thus the defluorination and mineralization increased but the kobs for

206

PFOA remained constant. When the concentration of PS further increased to 30 mM, however,

207

the scavenging caused by PS became dominant in the destruction process and the kobs did not

208

increase. As we discussed below, our model considers the scavenging reactions.

209

The oxidation of PFOA by SO4-· produced five main products with molecular weights

210

(MW) of 363, 313, 263, 213, 163 and 113. All these products were formed with the sequential

211

loss of 50 Da from the molecular weight, and this loss agreed with the loss of a CF2 group

212

from the parent compound. We proved this using standards solutions of perfluoroheptanoic

213

acid (PFHpA), perfluorohexanoic acid (PFHxA), perfluoropentanoic acid (PFPeA), 10


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perfluorobutyric acid (PFBA) and pentafluoropropionic acid (PFPrA). These five byproducts

215

were the main intermediates of PFOA that we found and this was consistent with the previous

216

reports.21, 24

217

The mass balance for carbon and fluorine is discussed in Text S7 and the data are

218

presented in Table S5. The initial measured TOC concentration was 1.07 mM; and, after 8

219

hours it was 0.76 mM. Accordingly, 0.31 mM of the TOC was mineralized according to TOC

220

measurements. Based on PFOA and PFOA byproduct measurements (hereafter called the

221

theoretical TOC), the theoretical initial and final TOC values were 1.17 mM and 0.61 mM,

222

respectively. The initial measured TOC was 0.10 mM lower than the theoretical TOC and it is

223

likely that the heat activated PS process that is used to measure TOC does not completely

224

mineralize the PFOA. However, the TOC measurement after 8 hours was 0.15 mM higher.

225

This could be experimental error and it is worth noting that the measured could be more

226

accurate; because, according to our kinetic model, the second order rate constants between

227

PFOA and PS are much lower than that for the byproducts. Consequently, the TOC

228

measurements after 8 hours reaction time should be more accurate, because the heat activated

229

PS process that is used for the TOC measurement would be more complete. (See Table

230

S1. C7F15COOH and C6F13COOH have similar values but C5F11COOH is 2.7 times larger

231

and C2F5COOH is 360 times larger). Also, the TOC measurement after 8 hours maybe higher

232

than the theoretical value because the theoretical TOC does not include all the byproducts that

233

are present.

234

The degree of mineralization was 29% based on TOC measurements and 48% based on

235

PFOA species measurement. The degree of mineralization based on the initial theoretical 11



236

TOC and measured TOC after 8 hours was 35%. Based on the aforementioned discussion, the

237

degree of mineralization was greater than 35% but less than 48%. Our mass balance on

238

fluorine showed that there were 2.19 mM of F initially and 2.02 mM after 8 hours (includes

239

F- ion and F contained in PFOA and its products). So we could see 17% of the fluorine was

240

contained in byproducts that we did not measure. The fluoride ion concentration in solution

241

was 0.89 mM; consequently, 59% of the fluorine the in the PFOA was converted to fluoride

242

ion. As shown in Figure 1, the concentration of PFHpA gradually increased in the initial 4

243

hours, and then slightly decreased after 4 hours. The evolution of PFHeA exhibited a similar

244

trend. On the other hand, the concentration of PFPeA, PFBA and PFPrA increased over time.

245

These observations indicated that short-chain PFCAs were generated in a stepwise manner

246

from longer-chain PFCAs. SO4-· is a powerful electrophile, which could draw an electron

247

from carboxylate group to generate carboxylate radical.23 Then, the carboxylate radical

248

decarboxylated and finally generate perﬂuorinated alcohols or cause the loss of a COF2 group

249

to yield perﬂuorohexyl radical.24,

250

Figure S3.

50-52

The detailed degradation pathways are proposed in

251

Table S3 lists the final pH under different initial dosage of PS. With the increase of initial

252

PS concentration (i.e., from 5 mM to 30 mM), the final pH in the system decreased from 2.81

253

to 2.04. Under UV irradiation, PS decomposition has been shown to release H+ as shown in

254

Eq. 2 and 3.30 pH drop during the reaction was not only observed in UV/PS, but also reported

255

in thermally activated PS.30, 53 In a previous study, the formation of H+ was found to be 1.7

256

mol H+/mol PS decomposed.23 The pH drop was always observed to be almost the same in

257

UV/PS process when PFOA was present (Figure S4) or not present (Figure S5). The reaction 12


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of PFOA oxidation by SO4-· can be described as Eq. 4.54 The H+ released from the

259

decomposition of PFOA is negligible, because its concentrations is low as compared to PFOA.

260

Instead the PS generates a significant amount of hydrogen ion and the pKa of HSO4- is lower

261

than the final pH; consequently, it cannot buffer the solution pH and the pH drops

262

dramatically. Consequently, PS decomposition main contributor for the pH decrease during

263

the reaction. S2O82- + H2O

264 265

HSO4-

2HSO4- + 0.5 O2

H+ + SO42-

pKa = 1.99

14 SO 4- + C7F 15COOH + 14 H2O

8 CO2 + 15 F- + 14 SO 42- + 29 H +

(2) (3) (4)

266 267

Kinetic Modeling of PFOA Degradation and Products Transformation. The model

268

was able to simulate PFOA degradation, products evolution, along with PS decay and predict

269

pH changes. The sample deviation (SD) (Eq. 5) was used to measure the relative error

270

between experimental data and calculated data. If the difference between the model

271

predictions and the data follows a Guassian distribution, then the model prediction ± the SD

272

will contain 68% of all the data. This assumes that the SD is equal to the standard deviation

273

which will require a large number of data but SD is still informative. If SD is 0.05, then the

274

model describes the data very well because ± 5% of the concentration data values will be

275

within 68% of the model predictions. Hence, SD was used to guide our model development

276

and this included both which reactions to include and rate constant determination.

277

SD =

1 [(Cexp − Ccal ) / Cexp ]2 ∑ n −1

(5)

278

where n is the number of data points with the same dosage of PS; Cexp and Ccal are the

279

experimental and calculated concentrations of different species, respectively. 13



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280

In the UV/PS process, SO4-· and S2O8-· are two important radicals for PFOA degradation.

281

SO4-· is the reactive radical in the system and S2O8-· is formed by a reaction between PS and

282

SO4-·. To our best knowledge, no one has reported the scavenging effect of S2O8-·. Yu et al.55

283

reported

284

2( SO 4 + S 2O 8

285

constant but they did not report the rate constant of the backward reaction. Due to the high

286

concentration of SO42- presented in the system, both the forward and the backward reactions

287

were considered in our model. Also we considered the SO4-· reaction with H2O (reaction 4 in

288

Table S1) which forms HO·. In addition, we also accounted for HO· formation from

289

S2O8-· (S 2O 8 + H 2O

290

simulation results of PS decay and PFOA destruction. The SD decreased when these reactions

291

were included in the model (e.g., SDPS decreased from 8.22 to 0.50; and, SDPFOA decreased

292

from 2.14 to 0.16 for a [PS] = 15 mM). Since protons were generated by these reactions, the

293

SDpH decreased when these reactions were included in the model (e.g., SDpH decreased from

294

0.63 to 0.09 for a [PS] = 15 mM). In addition, we also proposed reactions of SO4-· with

295

scavengers S2O8-· and SO5-· and the details are shown in reactions 8 and 20 in Table S1. We

296

found that SO4-· reactions with scavengers had to be included in the model, otherwise,

297

SO4-· increased dramatically. This SO4-· accumulation resulted in a much lower model

298

predicted PFOA concentration time profile as compared to the experimental results. In fact,

299

the SDPFOA decreased from 0.74 to 0.16 for [PS] = 15 mM, when the scavenging was included.

300

Figure S6 and Figure S7 show the model fits of PS and PFOA, respectively. The SD values of

301

major species in the UV/PS system were all listed in Table S4. SDPFOA were all below or equal

that

the

forward

and

backward

reactions

between

SO4-·

and

S2O82-

S2O 8- + SO4 2- ) exist. They reported a value for the forward rate

HO + H+ + S2 O82-). This reaction has a significant impact on the

14


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302

to 0.27. The inset of Figure S6 presents the experimental data of pH and the calculated results.

303

The SD values for pH were also listed in Table S4. SDpH were all smaller or equal to 0.24. To

304

our best knowledge, no other similar modeling work on PS decay and pH change system has

305

been reported in the literature.

306

Also, we evaluated the scavenging and reaction of S2O8-· with water. The proposed

307

reaction between S2O8-· and H2O has a significant influence on the decay of PS. Without

308

considering this reaction, the decay of PS would be under estimated. The proposed reactions

309

between SO4-· with S2O8-· and SO5-· also influenced the model simulations of PFOA

310

degradation.

311

SO4-· concentration and an under estimation of PFOA concentration. This discussion

312

confirmed that these reactions needed to be included in our model.

Without

these

reactions,

there

would

be

an

over

estimation

of

313

By fitting our model to experimental data, we determined the rate constants of SO4-· with

314

target compounds. The rate constants for the elementary reactions of all these products have

315

not been previously reported. To our knowledge, there are more than 35 elementary reactions

316

for all the PFOA intermediates and their rate constants cannot be determined using the genetic

317

algorithm simultaneously. Hence, the reactions of SO4-· with intermediates were expressed as

318

lumped reactions in our work as reactions 21 - 26 in Table S1. Table 1 provides the fitted rate

319

constants between SO4-· and the byproducts and their SD values. The rate constants increased

320

with a decrease in chain-length, which is consistent with a previous study about PFCAs decay

321

by electrochemical oxidation.56 The SD values for different products were in a reasonable

322

range (i.e., from 0.15 to 0.46) except for PFPrA. Larger experimental analytical errors are

323

expected for the extremely low PFPrA concentrations and these errors might be responsible 15



324

for the larger discrepancy between model fits and experimental data. It should be noted that

325

we have used the genetic algorithm to determine the rate constants for reactions between

326

radicals and intermediates.49,

327

combination with the genetic algorithm gives a good fit of the data.

57

We have found that the objective function, Eq. 5 in

328

We compared our model to the model developed by Yang et al..29 Figure S8 shows the

329

simulation results of PS decay, PFOA degradation and pH variation obtained from Yang et

330

al.’s model and our model. Yang et al. used the computer program Kintecus to predict the

331

radical distributions in the UV/PS system. Their model used the pseudo-steady state

332

assumption and literature-reported reactions. The pH values in their model was kept constant.

333

Yang et al.’s model under predicted PS decay and the PFOA destruction was over estimated.

334 335

Effect of Chloride Ions. In this study, we investigated the impact of Cl- on PFOA

degradation and the formation of ClO3- during the UV/PS treatment process.

336

Effect of Cl- on PFOA Degradation in UW. Figure 2a presents the experimental profiles of

337

PFOA degradation and model fits for various Cl- concentrations. SDPFOA values listed in Table

338

S4 were all below or equal to 0.13. Figure 3a shows the decay of Cl- under different initial Cl-

339

dosages. In the presence of Cl-, there was no PFOA degradation. SO4-· reacts first with Cl-

340

because its rate constant is 4.7×108 M-1·s-1,58 which was three orders of magnitude larger than

341

the model fitted rate constant of 2.59×105 M-1·s-1 between PFOA and SO4-·. Accordingly, Cl-

342

must first produce ClO3- before PFOA could be degraded.

343

An evaluation of contaminant degradation under different concentrations of Cl- impact

344

(i.e., 0.5 mM, 1 mM and 2 mM) was conducted using the quenching analysis QR. Table S6

345

listed the calculated quenching rate caused by Cl- for a variety of second order rate constants 16


Page 16 of 36

Page 17 of 36


346

of contaminants with SO4-· (kR). In the presence of 0.5 mM Cl-, QR would be above 70% for

347

kR values above 3.66×109 M-1·s-1 for treating 150 µM target compound. As listed in Table S6,

348

for lower the target compound concentrations, the scavenger effect caused by Cl- is more

349

pronounced. For example, all kR values would be above 1.10×1010 M-1·s-1 for QR to be above

350

70% for a target compound concentration of 50 µM. For target compound concentrations of

351

10 µM and 1 µM, all kR values would be above 5.48×1010 M-1·s-1 and 5.48×1011 M-1·s-1 for QR

352

above 70%, respectively. These are very large rate constants, perhaps approaching the

353

diffusion limit. Thus, the competition of Cl- with target compound to react with SO4-· is

354

significant. And Cl- has great influence on the UV/PS system.

355

The model predicted SO4-· concentration in the presence of Cl- is shown in Figure 2b.

356

With 2 mM Cl-, the concentration of SO4-· was much lower than that without Cl-. Indeed,

357

according to our model predictions, the concentration of SO4-· was reduced by more than 3

358

orders in the presence of Cl-. The concentrations of SO4-· under different initial dosages of Cl-

359

were predicted and shown in Figure S9. With the increasing of Cl- concentration, it took

360

longer before SO4-· began to degrade PFOA. The concentration of SO4-· increased gradually

361

only after nearly complete Cl- conversion to ClO3-. The simulated radical concentrations also

362

confirmed that the UV/PS process was highly sensitive to Cl-. This a report that the reaction

363

between SO4-· and Cl- was rapid and dominated the reduction of SO4-·.58 To the authors’ best

364

knowledge, no similar studies have predicted these radical concentrations as a function of

365

time.

366

Formation of Chlorate. Figures 3a and 3b show the decay of Cl- and the formation of

367

ClO3- under different doses of Cl-. No other intermediates, such as hypochlorous 17



368

acid/hypochlorite (HClO/ClO-) or chlorite (ClO2-), were observed in our experiment. Based

369

on the experimental results, a mole balance analysis of the formation ratios of ClO3- under

370

different Cl- concentrations was conducted. The stoichiometric conversion of Cl- to ClO3-

371

were 1.04, 1.09, 1.05 and 1.12 corresponding to Cl- concentration of 0.5 mM, 1 mM, 2 mM

372

and 3 mM, respectively. The mole balance results indicated that within experimental error all

373

Cl- was converted to ClO3- in the system. A slight decrease of ClO3- concentration was

374

observed late in the conversion process. A possible reason for this decrease was that UV light

375

photolysis of ClO3- and this needed to be studied further.

376

On the basis of the experimental data and our kinetic model of UV/PS process, a model of

377

Cl- decay and ClO3- formation in the UV/PS system was developed. Our model included the

378

reactions of bicarbonate species and their byproduct (i.e., CO3-·), and Cl- and its byproducts

379

(i.e., ClO-, ClO2-, Cl· and Cl2-·). The Cl- byproducts are important but these reactions are not

380

available in the literatures. Fang et al. reported similar reactions of SO4-· with Br- produced

381

BrO3- from BrO- and exhibited pseudo-first-order kinetics.44 Similarly, the literature-reported

382

reactions between HO· with HClO and ClO2· have been reported (reactions 51 and 56 in

383

Table S1).59 It should be noted that it was HClO that predicted here not ClO-, since the pKa of

384

HClO/ClO- was 7.6.60 We deduced that the similar reactions of SO4-· with HClO and

385

ClO2· would occur. The reaction rate constants between SO4-· with HClO and ClO2· were

386

obtained through fitting against the experimental data using the genetic algorithm. When these

387

two reactions were included in the model, SDClO3- reduced from 0.68 to 0.30 (reactions 57 and

388

58 in Table S1). Our model was able to fit the concentration profiles of Cl- decay and ClO3-

389

formation at different Cl- initial dosages. The SDCl- and SDClO3- were all below 0.55. 18


Page 18 of 36

Page 19 of 36


390

Combining the literature reported reactions with the reactions proposed in our work, we could

391

conclude that the ClO3- was generated from the oxidation of ClO2·, and ClO2· was produced

392

from the oxidation of HClO. During the oxidation process, SO4-· was the dominant reactive

393

species and all the Cl- chain reactions were started by SO4-·.

394

Effect of Bicarbonate. PFOA degradation for various concentrations of HCO3- is

395

presented in Figure S10. SDPFOA values are listed in Table S4. With the increase of HCO3-

396

dosage, PFOA degradation decreased. This was caused by the competition of carbonate

397

species to react with SO4-· and reduces the concentration of SO4-· in the system (reactions 72

398

and 73 in Table S1). 61, 62

399

PFOA can be degraded by both CO3-· and SO4-· in the UV/PS process. However, the

400

second order rate constant between PFOA and CO3-· is much small than SO4-· and PFOA. 63

401

Consequently, our model does not include the reactions between PFOA and CO3-·. And

402

CO3-· which is formed by SO4-· will reduce PFOA destruction by lowering the

403

SO4-· concentration.

404

To further clarify the decay of PFOA in the presence of HCO3-, the SO4-· concentrations

405

was also modeled and presented in Figure 2b. Compared with the SO4-· reduction caused by

406

15 mM HCO3-, the reduction of SO4-· caused by 2 mM Cl- was higher. QR analysis was also

407

performed for Cl- and HCO3- to compare their scavenger effects on PFOA degradation. As

408

shown in Table S6 and Table S7, QR for 2 mM Cl- was 4.12 × 10-5 and this is much lower than

409

7.19 × 10-4 for 15 mM carbonate species. These results demonstrate that Cl- is a much better

410

scavenger than bicarbonate for SO4-·.

411

For the kinetic modeling, H2O2 is an important scavenger of ·CO3- (reaction 77 in Table 19



412

S1).59 However, the reaction between ·CO3- and S2O82- was unavailable in the literature. This

413

reaction had great impact on the model. For example, SDPFOA was lowered from 0.75 to 0.04

414

when this reaction was included. Therefore, the reaction between ·CO3- and S2O82- was

415

2included (S 2O 8 + CO3

CO 32- + S 2O8 - ).

416

Model Application in Real Waters. Surface water (SW) and wastewater (WW) samples

417

were used to evaluate the effectiveness of UV/PS process in real waters. The background

418

concentration of constituents of SW and WW are shown in Table S2, including dissolved

419

organic carbon (DOC), Cl-, NO3-, PO43- and alkalinity. The absorption coefficient of water

420

matrix at 254 nm for SW and WW were 0.011 cm-1 and 0.085 cm-1, respectively. They were

421

included in the model as Eq. S3 in Text S5. PFOA degradation, Cl- decay and ClO3- formation

422

in ultrapure water (UW), SW and WW samples shown in Figure 4. pH decrease were also

423

observed in real water samples. pH sharply decreased to 2.89 and 3.03 in SW and WW

424

samples, respectively. For PFOA degradation, there was no obvious difference observed

425

between UW and SW. However, the degradation of PFOA was very slow for the first 2 h in

426

WW sample. The Cl- concentrations were 0.12 mM and 1.51 mM in SW and WW,

427

respectively. As discussed in the previous section, the presence of Cl- decreases the PFOA

428

degradation. By looking at the Cl- decay, we could conclude that the 2 hours lag-phase of

429

PFOA degradation in WW was mainly caused by Cl- oxidation. It should be noted that the

430

relatively long degradation time and lag phase was also partly due to the low UV intensity.

431

Compared with Cl-, NO3-, PO43- and alkalinity did not have notable effect on PFOA

432

degradation. The formation of ClO3- was observed in both SW and WW samples. A mole

433

balance showed that the stoichiometric conversion from Cl- to ClO3- were 0.98 and 0.95 in 20


Page 20 of 36

Page 21 of 36


434

SW and WW, respectively, indicating almost all Cl- was converted to ClO3-. Above results

435

highlighted the effect of Cl- on UV/PS treatment process. First, Cl- will significantly reduce

436

the concentration of SO4-·. Second, in the presence of Cl-, the byproduct of ClO3- will be formed.

437

The model predictions of PFOA degradation, Cl- decay and ClO3- formation are shown in

438

Figure 4; pH predictions are shown in Figure S11; and, SD values are listed in Table S8. Our

439

model predicts the PFOA degradation in UW and SW samples very well. However, the model

440

predictions for PFOA are much higher than the data for WW sample. The DOC in WW

441

sample was 17.69 mg/L, which is much higher than that for UW (~0.0 mg·L-1) and SW (0.198

442

mg·L-1). NOM appears to be activating PS to produce SO4-· and this is consistent with

443

literature reports about PS activation from NOM, phenols and quinones.

444

pursue any further because the complexity of the radicals formation in the presence of NOM

445

is not understood well enough to include in the model. But the model predicts a lower

446

degradation rate and it might be useful for preliminary design. Because if the model shows

447

that it is feasible for WW when we do not include PS activation, then we would expect that

448

the actual destruction would be greater. Accordingly, it would make sense to conduct

449

experiments to determine the actual performance. As listed in Table S8, SDpH was below 0.20

450

in SW and WW sample. For the prediction of Cl-, all the SD values were below or equal to

451

0.50 except for the SDCl- (i.e., 0.82) in SW. The bigger discrepancy of SDCl- in SW could be

452

attributed to the analytical errors in the presence of a low concentration of Cl- in SW. No other

453

similar work on UV/PS applications has attempted to model the destruction of PFOA and the

454

formation of ClO3- under changed pH conditions. In general, the model predictions agreed

455

very well with experimental data. They are predictions in this case because all the parameters 21


64, 65

We did not


456

came from previous model fits. Our model provided an approach to predict radical

457

distributions (i.e., SO4-·, HO·) in UV/PS system which could provide a better understanding

458

of the contaminant degradation.

459

Electrical energy per order (EE/O) is a useful way to evaluate the energy and costs in

460

UV/PS process. 66 As discussed above, our model was in good agreement with experimental

461

data in different water samples. Thus, we calculated EE/O for UW, SW and WW using our

462

model and this is described in Text S8. Two operational parameters, UV intensity and PS

463

doses, were varied to determine the most energy efficient condition. Figure S12 shows the

464

calculated EE/O values for UW, SW and WW. For a 15 mM PS doses, the EE/O

465

corresponding to different UV intensities were compared. The optimal UV intensity for UW

466

and SW was around 6×10-7 Einstein·L-1·s-1. However, due to the scavenger of Cl- in WW, high

467

UV intensity was required and the predicted optimal UV intensity was around 1.5×10-6

468

Einstein·L-1·s-1. For the UV intensity of 2.88×10-7 Einstein·L-1·s-1, the most energy efficient

469

PS doses were between 6-8 mM for UW and SW. As the UV intensity applied in our work

470

was relatively low and the energy cost for WW is not efficient here.

471

Engineering Implications. We examined PFOA destruction using the UV/PS process.

472

PFOA cannot be degraded by HO·. We developed a model that simulates the pH decrease,

473

chloride and bicarbonate scavenging and these factors were not considered in previous models.

474

Our work provides a comprehensive understanding on UV/PS process and quantitative

475

insights into chloride and bicarbonate scavenging under changing pH conditions. Our model

476

may have important implications for the design and optimization of the operational

477

parameters, such as PS initial dosage and UV intensity. 22


Page 22 of 36

Page 23 of 36


478

We found that chloride was a significant scavenger and that PFOA could not be degraded

479

until all of the Cl- is converted to ClO3-. Consequently, the UV/PS process will be ineffective

480

in degrading compounds that can only be degraded by SO4-·, when chloride is present.

481

At this time, we can speculate about the efficacy of the UV/PS process for compounds

482

that could be degraded other radicals, such as HO·, Cl· which are produced in the process in

483

the presence of chloride.67, 68 As shown in Figure S9 (b), model predicted HO· concentrations

484

were approaching to zero until all the chloride was converted to chlorate. Consequently, the

485

presence of chloride would also severely affect the destruction of compounds that can also be

486

degraded by HO· in the UV/ PS process. However, this should be proven experimentally.

487 488

ASSOCIATED CONTENT

489

Supporting Information. Text S1-S9, Tables S1-S8 and Figures S1-S13. This material is

490

available free of charge via the Internet at http://pubs.acs.org.

491 492

ACKNOWLEDGMENTS

493

This work was supported by National Science Foundation Award (NO. 0854416), the

494

Ministry of Science and Technology of China (NO. 2012BAJ25B04). The authors also

495

appreciate support from the Brook Byers Institute for Sustainable Systems, Hightower Chair

496

and Georgia Research Alliance at Georgia Institute of Technology. Y.Q. gratefully

497

acknowledges financial support from the China Scholarship Council. Y.Q. also thanks Dr.

498

Zhongming Lu for help with the modeling effort.

499 23



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Table 1.

Rate Constants of the Intermediates Degradation by SO4-· and the SD Values for different Intermediatesa

a

Intermediates

k

SD values

PFOA

2.59×105

0.16

PFHpA

2.68×105

0.15

PFHeA

7.02×105

0.46

PFPeA

1.26×106

0.45

PFBA

1.05×107

0.23

PFPrA

9.31×107

4.43

Rate constants were determined by fitting the data; rate constant unit are M-1·s-1.

688 689 690 32


Page 33 of 36


691 692 693 694 695 696 697 698 699 700 701 702 703 704 705 706 707

Figure Captions: Figure 1. Concentration profiles of PFOA and intermediates in the treatment process. The dots show the experimental results, and the solid lines represent the computer modeling fits. Conditions: [PFOA] = 150 µM (62.11 mg·L-1), PS dosage = 15 mM, UV intensity = 2.88 ×10-7 Einstein·L-1·s-1, without pH adjustment. Figure 2. (a) The concentration profiles of PFOA under different concentrations of Cl-. The dots show the experimental results, and the solid lines represent the computer modeling fits. (b) The model predicted SO4-· distributions for 15 mM HCO3- and 2 mM Cl-, along with the control group without HCO3- and Cl-. Conditions: [PFOA] = 150 µM (62.11 mg·L-1), PS dosage = 15 mM, UV intensity = 2.88 ×10-7 Einstein·L-1·s-1, without pH adjustment.

708 709 710 711 712

Figure 3. The decay of Cl- and the formation of ClO3- under different concentrations of Cl-. The dots show the experimental results, and the solid lines represent the computer modeling fits. Conditions: [PFOA] = 150 µM (62.11 mg·L-1), PS dosage = 15 mM, UV intensity = 2.88 ×10-7 Einstein·L-1·s-1, without pH adjustment.

713 714 715 716 717 718

Figure 4. (a) Concentration profiles of PFOA decay in ultrapure water (UW), surface water (SW) and wastewater (WW) samples, (b) Cl- decay and ClO3- formation in UW, SW and WW samples. The dots show the experimental results. The solid lines represent the model results. Conditions: [PFOA] = 150 µM (62.11 mg·L-1), PS dosage = 15 mM, UV intensity = 2.88 ×10-7 Einstein·L-1·s-1, without pH adjustment.

719 720

33


PFOA and the Intermediates (mM)


Page 34 of 36

0.15 PFOA (C7F15COOH) PFHpA (C6F13COOH) PFHeA (C5F11COOH)

0.10

PFPeA (C4F9COOH) PFBA (C3F7COOH) PFPrA (C2F5COOH)

0.05

0.00 0

2

4

6

8

Time (h) 721 722 723 724 725 726 727 728 729 730 731 732 733 734 735 736 737 738 739 740 741 742 743 744 745

Figure 1. Concentration profiles of PFOA and intermediates in the treatment process. The dots show the experimental results, and the solid lines represent the computer modeling fits. Conditions: [PFOA] = 150 µM (62.11 mg·L-1), PS dosage = 15 mM, UV intensity = 2.88 ×10-7 Einstein·L-1·s-1, without pH adjustment.

34


Page 35 of 36


Concentration of PFOA (mM)

(a) 0.15

0.10

-

3 mM Cl 2 mM Cl 1 mM Cl 0.5 mM Cl blank (0 mM Cl )

0.05

0.00 0

2

4

6

8

6

8

Time (h) 746

Concentration of sulfate radical (mM)

(b) blank 15 mM carbonate species 2 mM Cl -10

4.0x10

-10

2.0x10

0.0 0

2

4

Time (h) 747 748 749 750 751 752 753 754 755

Figure 2. (a) The concentration profiles of PFOA under different concentrations of Cl-. The dots show the experimental results, and the solid lines represent the computer modeling fits. (b) The model predicted SO4-· distributions for 15 mM carbonate species and 2 mM Cl-, along with the control group without these scavengers. Conditions: [PFOA] = 150 µM (62.11 mg·L-1), PS dosage = 15 mM, UV intensity = 2.88 ×10-7 Einstein·L-1·s-1, without pH adjustment.

35



(b) 3

-

2

1

0

2

1

0 0

756 757 758 759 760 761 762

-

3 mM Cl 2 mM Cl 1 mM Cl 0.5 mM Cl

-

Concentration of ClO3 (mM)

3 mM Cl 2 mM Cl 1 mM Cl 0.5 mM Cl

-

Concentration of Cl (mM)

(a) 3

Page 36 of 36

2

4

6

8

0

2

4

6

8

Time (h) Figure 3. The decay of Cl- and the formation of ClO3- under different concentrations of Cl-. The dots show the experimental results, and the solid lines represent the computer modeling fits. Conditions: [PFOA] = 150 µM (62.11 mg·L-1), PS dosage = 15 mM, UV intensity = 2.88 ×10-7 Einstein·L-1·s-1, without pH adjustment.

Concentration of Cl and ClO3 (mM)

763

775 776 777 778 779 780

(b) 1.6

-

0.16

0.12

1.2 -

0.8

Cl (SW) ClO3 (SW)

0.4

Cl (WW) ClO3 (WW)

-

765 766 767 768 769 770 771 772 773 774

Concentration of PFOA (mM)

(a)

764

0.08 PFOA (SW) PFOA (WW) PFOA (UW)

0.04

0.00 0

2

4

6

8

Time (h)

-

0.0 0

2

4

6

8

Time (h)

Figure 4. (a) Concentration profiles of PFOA decay in ultrapure water (UW), surface water (SW) and wastewater (WW) samples, (b) Cl- decay and ClO3- formation in UW, SW and WW samples. The dots show the experimental results. The solid lines represent the modeling predictions. Conditions: [PFOA] = 150 µM (62.11 mg·L-1), PS dosage = 15 mM, UV intensity = 2.88 ×10-7 Einstein·L-1·s-1, without pH adjustment.

781 782 36


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