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Oxidation of Organic Compounds in Water by Unactivated Peroxymonosulfate Yi Yang, Gourab Banerjee, Gary W. Brudvig, Jae-Hong Kim, and Joseph J. Pignatello Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.8b00735 • Publication Date (Web): 17 Apr 2018 Downloaded from http://pubs.acs.org on April 17, 2018

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Environmental Science & Technology

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Oxidation of Organic Compounds in Water by Unactivated Peroxymonosulfate

7 8 9 10

Yi Yang1, Gourab Banerjee2, Jae-Hong Kim3, Gary W. Brudvig2, Joseph J. Pignatello1*

11 12 13

Submitted to

14 15 16


17 18

1 Department of Environmental Sciences, The Connecticut Agricultural Experiment Station,

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123 Huntington St., P.O. Box 1106, New Haven, Connecticut 06504-1106, United States

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2 Department of Chemistry, Yale University, New Haven, Connecticut 06520, United States

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3 Department of Chemical and Environmental Engineering, Yale University, New Haven, CT

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06520-8286, United States

23 24

*Corresponding author contact information: [email protected]; tel. (1)-203-974-8518

1 ACS Paragon Plus Environment


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Abstract. Peroxymonosulfate (HSO5-, PMS) is an optional bulk oxidant in advanced oxidation

26

processes (AOPs) for treating wastewaters. Normally PMS is activated by the input of energy or

27

reducing agent to generate sulfate and/or hydroxyl radicals. This study shows that PMS without

28

explicit activation undergoes direct reaction with a variety of compounds, including antibiotics,

29

pharmaceuticals, phenolics, and commonly-used singlet oxygen (1O2) traps and quenchers,

30

specifically furfuryl alcohol (FFA), azide, and histidine. Reaction timeframes varied from

31

minutes to a few hours at pH 9. Using a test compound with intermediate reactivity (FFA), EPR

32

and scavenging experiments ruled out sulfate and hydroxyl. Although 1O2 was detected by EPR

33

and is produced stoichiometrically through PMS self-decomposition, 1O2 plays only a minor role

34

due to its efficient quenching by water, as confirmed by experiments manipulating 1O2 formation

35

rate (addition of H2O2) or lifetime (deuterium solvent isotope effect). Direct reactions with PMS

36

are highly pH- and ionic strength-sensitive and can be accelerated by (bi)carbonate, borate, and

37

pyrophosphate (although not phosphate) via non-radical pathways. The findings indicate that

38

direct reaction with PMS may contribute to degradation pathways and must be considered in

39

AOPs and other applications. They also signal caution to researchers when choosing buffers and

40

1

O2 traps/quenchers.

41 42


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TOC Art

44 45 46 47



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INTRODUCTION

49

A growing need exists for economical technologies to remove organic contaminants from

50

reclaimed waters such as industrial wastewater, oil and gas produced water, municipal

51

wastewater, landfill leachate, reverse osmosis brine, and contaminated groundwater. Advanced

52

oxidation processes (AOPs) are attractive for this purpose, since they target destruction of

53

pollutants rather than just concentrating them in another medium, as does adsorption and

54

membrane separation. In recent years, the option of using peroxymonosulfate (HSO5−, PMS) as

55

the bulk oxidant in AOPs has received considerable interest. In such AOPs, PMS is “activated”

56

to produce sulfate radical (SO4•−) and/or hydroxyl radical (•OH) which are highly reactive and

57

non-selective toward most organic compounds.1-3 Activation requires the input of energy in the

58

form of light (often in the presence of photocatalysts), heat, ultrasound, or low-valent transition

59

metal ions (eqs 1-2).4-8

60

hυ , ∆, or ultrasound HSO5- → SO4•- + •OH

(1)

61

Mn+ + HSO5- → Mn+1 + SO4•- + OH-

(2)

62

Activated PMS-based AOPs have been explored for use in various water reclamation

63

processes, for example decolorization and decontamination of industrial and municipal

64

wastewater,9-11 treatment of pulp and paper wastewater following other AOPs,10, 12 and treatment

65

of landfill leachates13

66

efficiently remove polychlorinated biphenyls, anisole, chloroethylenes, and chlorophenols in

67

aqueous or sediment systems16-19, implying promising applications for in-situ chemical oxidation.

68

While the fundamental chemistry of most activation processes is fairly well understood,

69

there is surprisingly little information on the chemistry of PMS itself in water—i.e., without an

70

activation step that converts PMS to SO4•−/•OH. Since activation of PMS in AOPs does not result

71

in instant degradation of organics, the issue of background reactions of unactivated PMS with

72

target organics and byproducts cannot be ignored. PMS is also used without activation in

73

sanitation applications where organic solutes may be present. For example, PMS is often added

74

as a chemosterilant and “shock oxidizer” to swimming pools, where it can encounter significant

75

concentrations of personal care products, sunscreens, parabens, urine, source water organic

76

matter, organic disinfectant compounds, and organic matter from saliva, skin, sweat, and human

77

excreta.20, 21 Reactions of PMS with organic compounds in this context is of obvious interest.

14 15

. Additionally, PMS-based technologies have been demonstrated to


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78

It has long been known that PMS self-decomposes spontaneously in water, especially at

79

slightly alkaline pH, to form singlet oxygen (1O2).22-24 Singlet oxygen has received attention for

80

its potential ability to remove organic contaminants in water25, 28, 30

26 27-29

or to inactivate

81

pathogens.

82

gives radical products including SO4•- , •OH, and superoxide (O2•-). Qi et al.31 reported the

83

decomposition of Acid Orange 7, phenol, and bisphenol A by PMS in alkaline solution and, on

84

the basis of radical and spin trapping experiments, suggested that 1O2 and O2•- were the principal

85

reactive species. Lou et al.32 attributed degradation of dyes in phosphate-buffered solutions to

86

SO4•- and •OH. Zhou et al.26 reported that benzoquinone activates PMS toward degradation of

87

the antibiotic sulfamethoxazole and postulated that intermediates of benzoquinone decompose to

88

1

89

synthesis in nonaqueous solvents, where it can oxidize alkenes to epoxides,33 aldehydes to

90

carboxylic acids,34 phosphines to phosphine oxides,35 and sulfides to sulfones.36 In water, it is

91

known to convert amino acids to various products.37-39

In addition, some researchers have proposed that PMS without explicit activation

O2, which then oxidizes sulfamethoxazole. PMS has also been used as a mild oxidant in organic

92

We recognize that there is a great deal of uncertainty about the aqueous-phase chemistry of

93

PMS without activation, whether as background in AOPs or in sanitation applications. Questions

94

remain about organic compound reactivity; the identity of the active oxidant(s) involved; and the

95

influence of solution composition including pH, ionic strength, and electrolyte composition. We

96

screened a number of compounds for their reactivity toward unactivated PMS including

97

pharmaceuticals of concern in drinking water and wastewaters, as well as phenolic compounds

98

that have been widely detected in water and soil. A detailed kinetic and product study was

99

undertaken on PMS self-decomposition considering that its product 1O2 is a potentially important

100

active oxidant in otherwise unactivated PMS systems. The involvement of 1O2 in water deserves

101

scrutiny because 1O2 is rapidly quenched by water.40 A test compound of intermediate reactivity,

102

furfuryl alcohol (FFA), was selected for detailed mechanistic investigation. We found that PMS

103

is the predominant oxidant and its reactions are highly sensitive to pH and electrolyte

104

composition. We discuss the implications of these findings in the context of PMS-based AOPs

105

and other applications.

106 107

Materials and Methods

108

Materials. The pharmaceuticals included the antibiotics sulfamethoxazole, trimethoprim and 5 ACS Paragon Plus Environment


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ampicillin; the anti-convulsant carbamazepine; and the H2-histamine receptor antagonists

110

cimetidine and ranitidine. 4-Chlorophenol and 2,4,6-trichlorophenol were chosen to represent

111

phenolic contaminants. The 1O2 scavengers included furfuryl alcohol (FFA), azide ion, and L-

112

histidine. The sources of these and other reagents are listed in Text S1, Supporting Information

113

(SI) Section.

114

Reactions of PMS with Organic Compounds. Reactions were carried out in 40 mL amber vials

115

mixed end-over-end at 60 rpm in the dark at 20 ± 0.5 °C. Except when otherwise mentioned,

116

solutions contained PMS at 1 mM and organic compounds at 50-100 µM, and were buffered with

117

50 mM phosphate. The pH remained within ± 0.1 unit of the target pH over the course of the

118

experiments. Collected samples were supplemented with 20 µL of 0.5 M sodium thiosulfate to

119

quench the remaining PMS before organics analysis.

120

Spontaneous Decomposition of PMS. Except where otherwise specified, solutions were

121

maintained at constant ionic strength of 0.5 M by adding NaClO4. Preliminary experiments ruled

122

out interference by transition metal ion impurities in PMS self-decomposition tests (Text S2 and

123

Figure S1). The yield of 1O2 in PMS decomposition was determined by calibration against the yield of

124 125

1

126

given in Text S3 and S4 of the SI. Photochemical experiments were performed in a Rayonet

127

reactor equipped with 16 visible-light bulbs (RPR-5750A) and a filter that cut off wavelengths

128

below ~400 nm.

129

Analytical Methods. The organic compounds were analyzed by HPLC with UV detection (Text

130

S5). The concentration of PMS was measured by the ABTS method.6 Briefly, 0.1 mL of sample

131

was added to a mixture of 0.5 mL of 2 mM 2,2’-Azino-bis(3-ethylbenzothiazoline-6-sulfonic

132

acid) diammonium salt (ABTS) solution, 1 mL pH 4 acetate buffer solution, and 20 µL 1.5 mM

133

iodide solution, and diluted to 10 ml with water. One mole of PMS produces two moles of

134

ABTS•+, which was detected at 415 nm with an absorption coefficient of 34000 M-1 cm-1. The

135

total peroxides concentration (sum of PMS and H2O2) was determined by addition of both iodide

136

and horseradish peroxidase, which were catalysts for PMS and H2O2, respectively. The

137

concentration of H2O2 was calculated by subtracting the PMS concentration, which was

138

measured only in the presence of iodide, from the total peroxides concentration.

139

O2 in the photolysis of methylene blue (MB) based on published methodology.41, 42 Details are

Electron paramagnetic resonance (EPR) spectrometry was carried out using a Bruker 6 ACS Paragon Plus Environment

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ELEXSYS E500 EPR instrument with an SHQ resonator and Oxford ESR900 continuous flow

141

cryostat. 5,5-Dimethyl-1-pyrrolidine N-oxide (DMPO) was the spin-trapping agent for •OH and

142

SO4•−, while TEMP was the spin-trapping agent for 1O2. As signal intensity is subject to loading

143

volume and instrument tune up time, signals were only used for identification purposes. The

144

solutions contained 1 mM PMS and 0.5 M NaClO4 at pH 9.4, and were allowed to react for 5

145

min prior to EPR measurement. Further details are given in Text S6.

146 147

RESULTS AND DISCUSSION

148

Degradation of Target Organic Compounds by Unactivated PMS

149

The pharmaceuticals, phenols and furfuryl alcohol (FFA) were individually reacted with

150

PMS without activation by metal ions or energy sources (Figure 1a and b). The reactions were

151

carried out at pH 9, approximately where the rate for PMS and FFA loss in their mixture reached

152

a maximum (Figure S4). The reactivity of the target compounds varied widely. Ranitidine and

153

cimetidine disappeared within 1 min, and ampicillin almost completely disappeared within 6 min.

154

Sulfamethoxazole and trimethoprim required a few hours to disappear. After 6 h the phenols

155

declined in concentration by 35% (2,4,6-trichlorophenol) or 65% (4-chlorophenol), while only 7%

156

of carbamazepine was removed. The half-life of FFA was about 143 min.

157

Figure 1.

158

A central question is the identity of the active oxidant(s) in chemical transformation of the

159

test compounds. From among the test compounds, FFA was selected for detailed study because it

160

showed intermediate reactivity in the screen (Figure 1), and in view of its use as a benchmark

161

1

162

order rate constants of the test compounds with reactive oxygen species are listed in Table S2.

163

We first undertook a detailed study of PMS self-decomposition in anticipation that its product,

164

1

165

reactive oxygen species through scavenging and EPR experiments. Third, we evaluated the

166

importance of 1O2 by taking into account the rates of its formation and decay via reaction with

167

the organics. Finally, we investigated solution conditions that affect unactivated PMS oxidations

168

of organic compounds.

O2 scavenger. FFA is also highly reactive towards •OH, and presumably SO4•−. Known second-

O2, may be an important active oxidant in the systems. Second, we tested for the presence of



169

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PMS Self-Decomposition and Generation of 1O2

170

PMS decomposes slowly in water to give sulfate and oxygen gas.22, 43, 44 Previous studies

171

reported widely varying rate constants (0.02−0.11 M-1 s-1)22, 43, 44 and neglected to consider the

172

effects of ionic strength and buffering agent, which are appreciable (see below). For these

173

reasons we re-examined the kinetics of PMS self-decomposition and quantified the yield of 1O2.

174

The PMS decomposition profiles at pH 9.4, 0.5 M ionic strength, and various initial PMS

175

concentrations were best fit by a second-order rate law, with the slope (rate constant) being

176

independent of the initial PMS concentration (Figure 2a). Fits of the profiles to a first-order rate

177

law were poorer and the derived first-order rate constant was not independent of initial PMS

178

concentration (Figure S5b). The mean observed second-order rate constant kdecomp is 0.013 ±

179

0.0003 M-1s-1 under the specified conditions. Although this value is lower than previously

180

reported, unlike previous studies22,

181

substances were added in this study, since we had ruled out interference by transition metal ions

182

(Text S2 and Figure S1). The kdecomp is pH-dependent between pH 6 and 12 with a maximum

183

between pH 9 and 9.4 (Figure 2b), close to the reported pKa,2 of 9.4 (ionic strength not

184

specified),22 where the product [HSO5-]⋅[SO52-] would be at maximum.

43, 44

, neither chelating agents (e.g, EDTA) nor other

185

The second-order rate law and pH dependence are broadly consistent with a previously-

186

proposed mechanism,22 in which the deprotonated species SO52- nucleophilically attacks the

187

distal O of the protonated species HSO5- to give a trioxide intermediate HSO6-,44 which

188

spontaneously decomposes to products (eqs 3 and 4).

189

kPMS HOOSO3- + -OOSO3-  → HOOOSO3- + SO42-

(3)

190

HOOOSO3- → SO42- + O2 + H+

(4)

191

Using NaClO4 or Na2SO4 to adjust ionic strength, kdecomp greatly increases with ionic

192

strength, and the increase is much steeper below 0.2 M (Figure 2c). On the other hand, pKa,2

193

decreases with ionic strength (NaClO4), from 9.78 ± 0.04 near 0 M, to 8.93 ± 0.05 at 0.5 M, to

194

8.46 ± 0.05 at 1 M (Text S7 and Figure S6). Given the decrease of pKa,2 with ionic strength,

195

kdecomp at constant pH of 9.4 is expected to decrease if the rate were dependent solely on reactant

196

concentrations. That the opposite is true means that the effects of ionic strength on pKa,2 are

197

overridden by charge shielding of the transition state (a tri-anion), resulting in net acceleration of 8 ACS Paragon Plus Environment

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PMS decomposition. Figure 2c also shows that adjustment of ionic strength with MgSO4 also

199

increases the rate, but less effectively than does NaClO4 or Na2SO4. A reasonable explanation is

200

complexation by Mg2+, which reduces the nucleophilicity of SO52- (discussion in Text S8).

201

Figure 2.

202

The yield of 1O2 from PMS self-decomposition was quantified by calibration against the

203

ADPA-MB method (Text S3 and S4; Figures S2 and S3 in SI).41, 42, 45 Using this method, the

204

cumulative 1O2 concentration as a function of PMS consumed was linear with a slope of 0.53

205

(Figure 2d), indicating that two moles of PMS generate one mole of 1O2.

206

To summarize at this point, bimolecular self-decomposition of PMS quantitatively forms

207

1

208

that PMS self-decomposition and reaction with organic compounds is accelerated by certain

209

buffer salts.

210

Scavenging Experiments and EPR Spectroscopic Studies

O2 with a rate constant that varies acutely with pH and ionic strength. Later, it will be shown

211

Scavenging experiments and EPR analyses were carried out to determine potential

212

involvement of SO4•−, •OH, and 1O2 in reactions of PMS with FFA. t-Butanol (0.1 M), which

213

efficiently scavenges •OH, had no effect on FFA and PMS loss rates (Figure 3a and b). Methanol

214

(0.1 M), which efficiently scavenges both SO4•- and •OH, slightly inhibited FFA loss, but also

215

slightly accelerated PMS loss. Thus, the slight inhibition of FFA degradation by methanol was

216

most likely due to reduction of the PMS concentration by direct reaction of PMS with methanol,

217

which has been reported.33 The alcohol scavenging experiments point away from involvement of

218

SO4•- and •OH in FFA transformation. Moreover, the wide variation in test compound

219

transformation rates (Figure 1) seems inconsistent with the relatively non-selective nature of

220

SO4•- and •OH.

221

The 1O2 scavenger, L-histidine added in 100-fold molar excess of FFA and 10-fold excess of

222

PMS completely suppressed FFA transformation (Figure 3a). However, PMS mixed alone with

223

L-histidine

224

transformation was due to depletion of PMS in solution. Sodium azide, used in the same molar

225

excess amounts, initially accelerated FFA transformation, but then FFA transformation halted

226

after about 30 min (Figure 3a). Like L-histidine, azide greatly accelerated PMS loss in the

was eliminated within 4 min (Figure 3b), suggesting that suppression of FFA



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227

absence of FFA, causing its complete disappearance by 30 min. It is apparent that L-histidine and

228

azide cannot be used as quenchers to validate 1O2 formation in PMS systems because they react

229

rapidly with PMS. This finding calls into question the conclusions of earlier studies where azide

230

was used to verify 1O2 involvement.26, 31

231

EPR spectra were obtained with DMPO as a trap for •OH and SO4•-, or with TEMP as a trap

232

for 1O2. Fenton-type reactions were used as an authentic source of •OH (Fe2+/H2O2) and SO4•-

233

(Fe2+/PMS) at pH 2−3.4 The spectra of Fe2+/H2O2 and Fe2+/PMS both show the signal assigned to

234

DMPO-OH (a(N) = a(H) = 14.9 G) (Figure 3c). The signal for DMPO-SO4− was not observed in

235

the spectrum of Fe2+/PMS because DMPO-SO4− rapidly hydrolyzes within the timeframe of

236

instrument tune up to give DMPO-OH (Scheme S1).46 No signals for either trapped •OH or SO4•-

237

were observed in systems containing PMS alone (Figure 3c), which, together with the

238

scavenging results, essentially rules out their involvement as active oxidants in transformation of

239

the test organic compounds in the unactivated PMS system. The reaction between TEMP and 1O2

240

generates TEMPO, which shows a characteristic signal of three-lines of equal intensity (a(N) =

241

16.9 G).27 The appearance of the TEMPO signal (Figure 3d) is clear evidence for 1O2 production,

242

but does not prove its participation in transformation of the test compounds. Superoxide, another

243

reactive oxygen species implicated in PMS systems,31 is generally not very reactive with organic

244

compounds in aqueous solution.3 Moreover, it would be difficult to rationalize its presence in the

245

absence of •OH or SO4•-.

246 247

Figure 3. 1

O2 Oxidation vs Direct Reaction by PMS

248

Contribution of 1O2. To evaluate the role of 1O2, we carried out PMS-FFA reactions with

249

mixtures at different pH between 6 and 12 and compared the rates of FFA transformation and

250

PMS decomposition, assuming PMS loss is due solely to its quantitative decomposition to 1O2.

251

The observed pseudo first-order rate constant for FFA degradation (kobs,FFA) and the observed

252

second-order rate constant for PMS decomposition (kobs,PMS) follow a closely similar pattern with

253

pH, each reaching a maximum at ~pH 9 (Figure S4c). While this at first appears to be consistent

254

with 1O2 as the active oxidant, note that the so-calculated value of kobs,PMS is three times greater

255

than the kdecomp for PMS self-decomposition in the absence of FFA and at the same ionic strength

256

and pH. Were self-decomposition of PMS rate-limiting to 1O2 production, this finding would 10 ACS Paragon Plus Environment

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257

mean that FFA somehow accelerates bimolecular PMS self-decomposition in their mixtures, a

258

difficult proposition to accept. Thus, other reactions are possible. The maximum role of 1O2 is estimated as follows. Under the conditions of Figure S4c, the

259 260

initial rate of 1O2 formation by PMS self-decomposition in the absence of FFA is given by eq 5,

261 262

R 1O

2 ,form

=

d [ 1 O 2 ]t = k decomp [PMS]2 dt

(5)

and is equal to 0.0092 µM s-1. The fraction of 1O2 reacting with FFA ( f O 1

263

f 1O

2 ,FFA

=

k1O k1 O

2

2 ,FFA

[FFA]

+ k1 O ,H O 2

2

[FFA] ,FFA

2 , FFA

) is calculated by,

(6)

(1.2 × 108 M-1s-1) is the FFA-1O2 rate constant, which is independent of pH from 3

264

where k 1 O

265

to 12;47, 49 and k O

266

Thus, only 1.5 µM of 1O2 reacted with FFA in the first hour, while 33 µM of FFA was removed,

267

corresponding to 4.6% of FFA reacting through the 1O2 pathway. In the same way, for other test

268

compounds with known 1O2 rate constants (Table S2), 1O2 contributed 12.5% to transformation

269

of 2,4,6-trichlorophenol, 3.6% to transformation of 4-chlorophenol, and 99% D2O were compared at

289

[D+] = [H+] = 1 x 10-9 M (i.e., pH = 9.0 in water; pHmeasured = 8.6 in D2O), 1 mM [PMS]0, and 0.1

290

mM [FFA]0 (Figure S8). In D2O, the pKa2D of PMS was determined spectrophotometrically to be

291

9.76 ± 0.07 compared with 9.13 in H2O (Figure S6). The pKaD of FFA is taken to be 10.17 from

292

a linear relationship reported for weak acids51 (pKaH + 0.62).

O2 quenching rate constant ( k 1 O

2 ,D2 O

=1.6 × 104 s-1);48 b) pre-equilibria (i.e., acid dissociation);

293

H D The observed solvent isotope effect for PMS self-decomposition, SIEdecomp = kdecomp kdecomp , is

294

obtained from fits of the decomposition curves in the absence of FFA (Figure S8a) to a second-

295

H D order rate law, and is found to be 4.6. Assuming eq 3 is rate-limiting, kdecomp is equal to k decomp

296

(k

297

H D and K a2D , it follows that kPMS kPMS = 1.07, a value consistent with a secondary kinetic isotope

298

effect in which H(D) is once-removed from the atom undergoing bond transformation. The rate

299

law for reaction of FFA with 1O2 is given by,

300

D k PMS )( K a2H K a2D ) , where kPMS is the elementary rate constant for eq 3. Thus, knowing K a2H

H PMS

−

d [FFA] = k 1 O ,FFA [ 1 O 2 ]ss [FFA] 2 dt

(7)

301

where [1O2]ss is the steady-state 1O2 concentration and k 1 O

302

rate constant. The observed isotope effect for FFA transformation in the presence of PMS,

303

H D SIEobs = kobs kobs , is obtained from pseudo first-order fit to the experimental decay curves in

304

Figure S8b, and assuming [PMS] changes little, to be 2.0. The predicted solvent isotope effect

305

SIE pred, 1 O is equal to 2

k 1HO

2 ,FFA

k 1DO

2 ,FFA

⋅

2 ,FFA

is the second-order elementary

[ 1 O2 ]ssH . It may be assumed that k 1HO ,FFA k 1DO ,FFA =1 because 1 D 2 2 [ O2 ]ss

306

H(D) is remote from the diene reaction center. The respective [1O2]ss are obtained by assuming

307

equal rates of formation (eq 5) and decay (eq 7 plus solvent quenching) and are given by,


Page 13 of 26

308


1

H(D) 2 ss

[ O ]

=

H(D) kdecomp [PMS]2

(8)

kH(D)2 O +k 1H(D) [FFA] O ,FFA 2

309

The SIE pred, 1 O is thus calculated to be 0.49. The 4-fold discrepancy between SIE pred, 1 O and 2

310

2

SIEobs argues against a dominant role of 1O2 in FFA transformation.

311

Contribution of direct reaction. An alternative to oxidation by 1O2 is direct reaction between

312

PMS and FFA. The presumed second-order rate constant for direct reaction (kPMS,FFA) at low

313

PMS conversion is given by the product of kobs,FFA (from Figure S4c) and the initial PMS

314

concentration (1 mM). The kPMS,FFA is plotted as a function of pH from 6 to 12 in Figure 4.

315

In reactions with alkenes the species HSO52- behaves an electrophile since the rate increases

316

with the number of alkyl substituents.34 In view of the known chemistry of PMS, it is proposed

317

that the distal O of PMS electrophilically attacks the diene system of FFA. The pH dependence

318

of kPMS,FFA can be rationalized in terms of this mechanism and by taking into account the

319

speciation of both reactants (pKa,2 of PMS is 9.13 at the ionic strength used, see Figures S4c and

320

S6; the pKa of FFA is 9.5550). The following reactions may be written,

321

HSO5- + FFA → product 1

k1

(9)

322

HSO5- + FFA- → product 2

k2

(10)

323

SO52- + FFA- → product 3

k3

(11)

324

The expression for kPMS,FFA is given by

325

k PMS,FFA = k1α HSO- βFFA + k 2 α HSO- (1 − βFFA ) + k3 1 − α HSO5

5

(

5

) (1 − β

FFA

)

(12)

326

where FFA- is the acid-dissociated form of FFA; α HSO is the fraction of total PMS as HSO5-;

327

βFFA is the fraction of neutral FFA; and k1, k2, and k3 are the respective second-order rate

328

constants of eqs 9-11. The experimental curve was deconvoluted by nonlinear, 3-parameter

329

regression and the results are shown in Figure 4. The best-fit rate constants (units of M-1 s-1)

330

follow the order:

331 332

− 5

k2 (0.564 ± 0.067) > k1 (0.059 ± 0.004) > k3 (0.012 ± 0.004).

The order in rate constants is consistent with the proposed electrophilic pathway, if the 13 ACS Paragon Plus Environment


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333

reasonable assumption is made that HSO5- is more electrophilic than SO52-, and that FFA− is

334

more nucleophilic than FFA due to sigma electron density donation from the dissociated alcohol

335

to the diene.

336

Figure 4.

337

We now consider whether direct reaction is consistent with the observed SIEobs of 2.0 from

338

Figure S8b. A predicted isotope effect, SIE pred,PMS , can be calculated from eq 12. Neglecting

339

kinetic isotope effects on k1, k2 and k3, this gives SIE pred,PMS = 1.27. A secondary isotope effect

340

may be expected for k1 and k2, whereas no isotope effect is expected for k3. To make SIE pred,PMS =

341

SIEobs requires an isotope effect on k1 and k2 (considered together) of approximately 1.57, a value

342

that is reasonable for a secondary deuterium kinetic isotope effect on reaction at the distal O. We

343

conclude, therefore, that SIEobs is consistent with direct reaction between PMS and FFA.

344

To summarize at this point, it is clear that direct reaction predominates over 1O2 oxidation

345

of FFA. The same can be inferred for the other test compounds on the basis of the above

346

calculations that incorporate water-quenching.

347 348

Effects of Electrolytes on Unactivated PMS Reactions.

349

This study also looked at specific effects of the background electrolyte on PMS reactions

350

with organics, as well as on PMS self-decomposition. Phosphate, bicarbonate, and borate were

351

examined because of their appearance in some wastewaters, their widespread use as pH buffers

352

in research, and their reported ability to activate H2O2 and/or PMS. Yang et al.52 report that

353

phosphate activates H2O2 towards self-decomposition and oxidation of methylene blue at pH 10.

354

Lou et al.

355

H2O2 and can accelerate oxidation of organosulfides by both H2O2 and peroxydisufate (S2O82-).53,

356

54

357

Pyrophosphate was selected because it was found to accelerate PMS decomposition in

358

connection with a different objective in our laboratory.

32

observed phosphate acceleration of PMS reactions. Bicarbonate complexes with

Borate complexes with H2O2 and can accelerate H2O2 oxidation of organosulfides.55,

56

359

The results of PMS-FFA reactions, initially at 1 mM PMS and 100 µM FFA in 100 mM

360

buffer salt at constant ionic strength of 1 M (made up with NaClO4), appear in Figure 5. 14 ACS Paragon Plus Environment

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361

Phosphate had no effect on either FFA or PMS loss compared to the 1 M NaClO4 control. This

362

finding contrasts with that of Lou et al.,32 who observed that phosphate accelerated degradation

363

of organic compounds in the presence of PMS. Bicarbonate and pyrophosphate significantly

364

accelerated both FFA and PMS loss. Borate strongly accelerated PMS loss, whereas FFA loss

365

was faster initially but slowed down after ~90 min. The slowdown can be attributed to PMS

366

depletion. The scavengers, t-BuOH and methanol, had little or no effect on PMS loss in any case

367

(Figure S9). t-BuOH had no effect on FFA loss in borate, but slightly inhibited FFA loss in

368

bicarbonate and pyrophosphate solution. Methanol had no effect on FFA loss in borate, a slight

369

effect in bicarbonate, and a strong effect in pyrophosphate solution. While EPR spectra of these

370

systems were positive for 1O2 (Figure S10), they were negative for •OH or SO4•- in all cases (data

371

not shown). The strong inhibition by methanol in the pyrophosphate case deserves further study.

372

Relative to the ionic strength control, PMS self-decomposition (Figure S11) was unaffected

373

by phosphate; accelerated by borate; accelerated by pyrophosphate up to 70 mM pyrophosphate

374

(where it began to decline); and accelerated by bicarbonate above 50 mM bicarbonate.

375

This is the first report to our knowledge of PMS activation by bicarbonate, borate, or

376

pyrophosphate. These oxoanions may nucleophilically attack the distal O of PMS to form the

377

corresponding peroxoanion:

378

HOOSO3- + XOnm- → SO42- + HOOXOn-1(m-1)-

(13)

379

Previous studies found that peroxymonocarbonates or peroxyborates form rapidly in small yield

380

in solutions of H2O2 with bicarbonate or borate. The distal O of HOOXOn-1(m-1)- can engage in a

381

two-electron electrophilic attack on an organic compound, leading to displacement of XOnm- and

382

formation of the oxidized organic product.53,

383

reactions remain to be established.

384 385

56

The mechanisms of the corresponding PMS

Figure 5. Environmental Significance.

386

Few studies have addressed the implications of unactivated PMS reactions on contaminant

387

degradation in environmental media. We show that many compounds are susceptible to direct

388

reaction with PMS within timeframes relevant to activated PMS-based AOPs. There is a dearth

389

of information on the kinetics and mechanisms of direct reaction between PMS and a wide range 15 ACS Paragon Plus Environment


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390

of pollutants. While PMS itself likely reacts electrophilically with FFA, it can behave as an

391

electrophilic or nucleophilic oxidant depending on pH and the target compound. In activated

392

PMS-based AOPs, direct reaction may help or hinder the desired goal of hazard elimination by

393

routing degradation through pathways that lead to less or more toxic byproducts. Direct reaction

394

with PMS may also be useful on its own for water purification. While activated PMS-based

395

AOPs are attractive in many ways, they do have some drawbacks. The energy input is costly and

396

much of it is wasted on the surrounding matrix. Reactions based on solid-phase catalysis are

397

often limited by diffusion or adsorption to the catalyst surface. Catalysts and transition metal

398

reductants must be separated from the treated water. Reclaimable waters can be high in salts that

399

may interfere. (Bi)carbonate scavenges SO4•− and •OH to give less reactive carbonate radicals.

400

Halide ions scavenge SO4•− and •OH to form reactive halogen species (X•, X2•−, X2, X3−,

401

HOX/OX−) that can incorporate halogen into some organic structures, including alkenes,

402

aromatic and heterocyclic aromatic rings, ketones, aldehydes, and amines.57 Halogenated

403

byproducts pose a serious concern in view of the inherent toxicity of many halogenated

404

compounds. While OH• reacts only with bromide and iodide at ordinary pH, SO4•− reacts

405

efficiently with bromide, iodide and chloride,58 making reactive halogen species especially

406

problematic in activated PMS-based AOPs because of the ever-presence of chloride in

407

wastewaters. By contrast, PMS reacts only slowly with chloride.59

408

PMS is a milder but more selective oxidizing agent than SO4•− and •OH. It may be possible

409

to take advantage of the selectivity principle in certain applications—for example, where highly

410

hazardous contaminants susceptible to direct PMS oxidation exist in the presence of a large

411

background concentration of more innocuous compounds that would consume a non-selective

412

species like SO4•− or •OH. The finding that pyrophosphate, bicarbonate, and borate accelerate

413

PMS decomposition raises questions about their use as buffers in mechanistic studies of PMS

414

systems.

415

Non-photochemical means of 1O2 generation have been explored for potential application in

416

water treatment.25, 26 Unlike photochemical methods that generate 1O2 continuously as long as

417

the light is on, non-photochemical methods rely on steady-state generation of 1O2 that declines

418

with consumption of the precursor oxidant. Our experience with PMS/FFA suggests that in many

419

such cases the steady-state 1O2 concentration may be insufficient to achieve significant oxidation 16 ACS Paragon Plus Environment

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420

of the target organic compound due to its quenching by water. This limitation seems to have

421

been overlooked in the literature. Lastly, investigators of PMS chemistry should take note that

422

traditional 1O2 quenchers, like azide ion and L-histidine, and furfuryl alcohol, react directly with

423

PMS. It is possible that this has led to misinterpretation of experimental observations in past

424

studies.

425 426

ACKNOWLEDGMENTS

427

The authors thank the Chinese International Postdoctoral Exchange Fellowship Program

428

(No. 20160074) for support for Y.Y. The EPR spectroscopy work was supported by the

429

Department of Energy, Office of Basic Energy Sciences, Division of Chemical Sciences, grant

430

DE-FG02-05ER15646 (G.W.B. and G.B.).

431 432

SUPPORTING INFORMATION

433

The Supporting Information is available free of charge on the ACS Publications website at

434

DOI: [to be inserted]. It gives additional details on materials and methods; rate constants of

435

relevant compounds with reactive oxygen species; supplementary EPR spectra; a section on the

436

determination of pKa; a section on the kinetics of PMS self-decomposition; and supplementary

437

sections and data on the influence of electrolytes, pH, ionic strength, added hydrogen peroxide,

438

solvent deuterium isotope, and scavengers on PMS self-decomposition and reactions of PMS

439

with FFA.

440



Figures

Figure 1. Degradation of organic contaminants in water by PMS alone. ([phosphate]0 = 50 mM, pH = 9, for FFA: [PMS]0 = 1 mM, [FFA]0 = 100 µM; for carbamazepine, 2,4,6-trichlorophenol, 4-chlorophenol, trimethoprim and sulfamethoxazole: [PMS]0 = 1 mM, [substrate]0 = 50 µM; for cimetidine, ranitidine and ampicillin: [PMS]0 = 0.5 mM, [substrate]0 = 100 µM).

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Figure 2. PMS self-decomposition. (a) Second-order kinetic plots at different PMS concentrations (pH = 9.4, and 0.5 M ionic strength controlled by NaClO4). (b) Effect of pH on the rate constant ([PMS]0 = 1 mM, and 0.5 M ionic strength). (c) Effect of ionic strength on the rate constant ([PMS]0 = 1 mM and pH = 9.4). (d) Stoichiometry of 1O2 in D2O/H2O (0.95/0.05) in the absence and presence of H2O2 ([ADPA]0 = 200 µM, pH = 9.4, and 0.15 M ionic strength; for PMS alone, [PMS]0 = 3 mM; for PMS with H2O2, [PMS]0 = 2 mM and [H2O2]0 = 2 mM).



Figure 3. Detection of reactive oxygen species in unactivated PMS. Effects of scavengers on: (a) FFA degradation in PMS-FFA mixtures; and (b) PMS decomposition in PMS alone ([PMS]0 = 1 mM, pH = 9.0, [phosphate]0 = 50 mM, and [FFA]0 = 100 µM). EPR spectra of: (c) DMPO adducts; and (d) TEMPO in PMS alone and PMS/H2O2. (PMS alone: [PMS]0 = 1 mM, pH = 9.4, and 0.5 M ionic strength controlled with NaClO4; PMS/H2O2: [PMS]0 = 1 mM, [H2O2]0 = 1 mM, pH = 9.4, and 0.5 M ionic strength; Fe(II)/H2O2: [H2O2]0 = 1 mM, [FeSO4]0 = 300 µM, pH = 3; Fe(II)/PMS: [PMS]0 = 1 mM, [FeSO4]0 = 300 µM, pH = 3).


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Figure 4. Variation with pH of second-order rate constant for reaction between PMS and FFA, and deconvolution of rate constants for individual reactions, eqs 9-11. Symbols represent measured data. Solid line represents model prediction according to eq 12. Dashed lines represent the contributions of individual eqs 9-11 to the model prediction.



Figure 5. Effect of buffer salts on (a) FFA transformation and (b) PMS decomposition in PMSFFA mixtures relative to the control at the same pH (9) and ionic strength (1 M adjusted with NaClO4), [PMS]0 = 1 mM, [buffer] = 100 mM, [FFA]0 = 100 µM. (Note that both FFA degradation and PMS decomposition are slower in 1 M than 0.15 M ionic strength at pH 9 shown in Figure S4. This is because ionic strength favors the anionic forms of PMS and FFA, which are inherently less reactive; see text.)


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