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Remediation and Control Technologies
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
Environmental Science & Technology
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
20
2 Department of Chemistry, Yale University, New Haven, Connecticut 06520, United States
21
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
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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.
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INTRODUCTION
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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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It has long been known that PMS self-decomposes spontaneously in water, especially at
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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
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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
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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
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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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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,
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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
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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).
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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.
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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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REFERENCES 1. Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B., Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals. J. Phys. Chem. Ref. Data 1988, 17, (2), 513-886. 2. Neta, P.; Huie, R. E.; Ross, A. B., Rate constants for reactions of inorganic radicals in aqueous solution. J. Phys. Chem. Ref. Data 1988, 17, (3), 1027-1284. 3. NIST, NDRL/NIST Solution Kinetics Database on the Web. http://kinetics.nist.gov/solution/ 2017, (Accessed September, 2017). 4. Anipsitakis, G. P.; Dionysiou, D. D., Radical generation by the interaction of transition metals with common oxidants. Environ. Sci. Technol. 2004, 38, (13), 3705-3712. 5. Guan, Y.-H.; Ma, J.; Ren, Y.-M.; Liu, Y.-L.; Xiao, J.-Y.; Lin, L.-q.; Zhang, C., Efficient degradation of atrazine by magnetic porous copper ferrite catalyzed peroxymonosulfate oxidation via the formation of hydroxyl and sulfate radicals. Water Res. 2013, 47, (14), 54315438. 6. Yang, Y.; Jiang, J.; Lu, X.; Ma, J.; Liu, Y., Production of Sulfate Radical and Hydroxyl Radical by Reaction of Ozone with Peroxymonosulfate: A Novel Advanced Oxidation Process. Environ. Sci. Technol. 2015, 49, (12), 7330-7339. 7. Guan, Y.-H.; Ma, J.; Li, X.-C.; Fang, J.-Y.; Chen, L.-W., Influence of pH on the formation of sulfate and hydroxyl radicals in the UV/peroxymonosulfate system. Environ. Sci. Technol. 2011, 45, (21), 9308-9314. 8. Su, S.; Guo, W.; Yi, C.; Leng, Y.; Ma, Z., Degradation of amoxicillin in aqueous solution using sulphate radicals under ultrasound irradiation. Ultrason. Sonochem. 2012, 19, (3), 469-474. 9. Ghanbari, F.; Moradi, M.; Manshouri, M., Textile wastewater decolorization by zero valent iron activated peroxymonosulfate: compared with zero valent copper. Journal of Environmental Chemical Engineering 2014, 2, (3), 1846-1851. 10. Jaafarzadeh, N.; Omidinasab, M.; Ghanbari, F., Combined electrocoagulation and UV-based sulfate radical oxidation processes for treatment of pulp and paper wastewater. Process Saf. Environ. Prot. 2016, 102, 462-472. 11. Ghanbari, F.; Moradi, M., Application of peroxymonosulfate and its activation methods for degradation of environmental organic pollutants: Review. Chem. Eng. J. 2017, 310, (Part 1), 41-62. 12. Jaafarzadeh, N.; Ghanbari, F.; Ahmadi, M.; Omidinasab, M., Efficient integrated processes for pulp and paper wastewater treatment and phytotoxicity reduction: Permanganate, electro-Fenton and Co3O4/UV/peroxymonosulfate. Chem. Eng. J. 2017, 308, 142-150. 13. Zhang, W.; Yang, S.; Niu, R.; Shao, X.; Shan, L.; Yang, X.; Wang, P. In Microwave-assisted COD removal from landfill leachate by hydrogen peroxide, peroxymonosulfate and persulfate, Bioinformatics and Biomedical Engineering (iCBBE), 2010 4th International Conference on, 2010; IEEE: 2010; pp 1-4. 14. Rivas, F. J.; Beltrán, F. J.; Carvalho, F.; Alvarez, P. M., Oxone-promoted wet air oxidation of landfill leachates. Ind. Eng. Chem. Res. 2005, 44, (4), 749-758. 15. Sun, J.; Li, X.; Feng, J.; Tian, X., Oxone/Co 2+ oxidation as an advanced oxidation process: comparison with traditional Fenton oxidation for treatment of landfill leachate. Water Res. 2009, 43, (17), 4363-4369. 16. Rastogi, A.; Al-Abed, S. R.; Dionysiou, D. D., Sulfate radical-based ferrous– peroxymonosulfate oxidative system for PCBs degradation in aqueous and sediment systems. Applied Catalysis B: Environmental 2009, 85, (3), 171-179.
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Environmental Science & Technology
17. Yu, M.; Teel, A. L.; Watts, R. J., Activation of Peroxymonosulfate by Subsurface Minerals. J. Contam. Hydrol. 2016, 191, 33-43. 18. Watts, R. J.; Yu, M.; Teel, A. L., Reactive oxygen species and associated reactivity of peroxymonosulfate activated by soluble iron species. J. Contam. Hydrol. 2017, 205, 70-77. 19. Devi, P.; Das, U.; Dalai, A. K., In-situ chemical oxidation: Principle and applications of peroxide and persulfate treatments in wastewater systems. Sci. Total Environ. 2016, 571, 643-657. 20. Teo, T. L.; Coleman, H. M.; Khan, S. J., Chemical contaminants in swimming pools: occurrence, implications and control. Environ. Int. 2015, 76, 16-31. 21. DuPont DuPont Oxone Monopersulfate Compound. http://www2.dupont.com/Oxone/en_US/index.html 22. Ball, D. L.; Edwards, J. O., The kinetics and mechanism of the decomposition of Caro's acid. I. J. Am. Chem. Soc. 1956, 78, (6), 1125-1129. 23. Yuan, Z.; Ni, Y.; Van Heiningen, A. R. P., Kinetics of peracetic acid decomposition: Part I: Spontaneous decomposition at typical pulp bleaching conditions. The Canadian Journal of Chemical Engineering 1997, 75, (1), 37-41. 24. Yuan, Z.; Ni, Y.; Van Heiningen, A. R. P., Kinetics of the peracetic acid decomposition: Part II: pH effect and alkaline hydrolysis. The Canadian Journal of Chemical Engineering 1997, 75, (1), 42-47. 25. Bokare, A. D.; Choi, W., Singlet-Oxygen Generation in Alkaline Periodate Solution. Environ. Sci. Technol. 2015, 49, (24), 14392-14400. 26. Zhou, Y.; Jiang, J.; Gao, Y.; Ma, J.; Pang, S.-Y.; Li, J.; Lu, X.-T.; Yuan, L.-P., Activation of Peroxymonosulfate by Benzoquinone: A Novel Nonradical Oxidation Process. Environ. Sci. Technol. 2015, 49, (21), 12941-12950. 27. Lee, J.; Hong, S.; Mackeyev, Y.; Lee, C.; Chung, E.; Wilson, L. J.; Kim, J.-H.; Alvarez, P. J. J., Photosensitized Oxidation of Emerging Organic Pollutants by Tetrakis C60 Aminofullerene-Derivatized Silica under Visible Light Irradiation. Environ. Sci. Technol. 2011, 45, (24), 10598-10604. 28. Moor, K. J.; Valle, D. C.; Li, C.; Kim, J.-H., Improving the Visible Light Photoactivity of Supported Fullerene Photocatalysts through the Use of [C70] Fullerene. Environ. Sci. Technol. 2015, 49, (10), 6190-6197. 29. Tratnyek, P. G.; Hoigne, J., Oxidation of substituted phenols in the environment: a QSAR analysis of rate constants for reaction with singlet oxygen. Environ. Sci. Technol. 1991, 25, (9), 1596-1604. 30. Kohn, T.; Nelson, K. L., Sunlight-Mediated Inactivation of MS2 Coliphage via Exogenous Singlet Oxygen Produced by Sensitizers in Natural Waters. Environ. Sci. Technol. 2007, 41, (1), 192-197. 31. Qi, C.; Liu, X.; Ma, J.; Lin, C.; Li, X.; Zhang, H., Activation of peroxymonosulfate by base: Implications for the degradation of organic pollutants. Chemosphere 2016, 151, 280-288. 32. Lou, X.; Wu, L.; Guo, Y.; Chen, C.; Wang, Z.; Xiao, D.; Fang, C.; Liu, J.; Zhao, J.; Lu, S., Peroxymonosulfate activation by phosphate anion for organics degradation in water. Chemosphere 2014, 117, 582-585. 33. Zhu, W.; Ford, W. T., Oxidation of alkenes with aqueous potassium peroxymonosulfate and no organic solvent. J. Org. Chem 1991, 56, (25), 7022-7026.
ACS Paragon Plus Environment
Page 24 of 26
Page 25 of 26
Environmental Science & Technology
34. Travis, B. R.; Sivakumar, M.; Hollist, G. O.; Borhan, B., Facile oxidation of aldehydes to acids and esters with oxone. Org. Lett. 2003, 5, (7), 1031-1034. 35. Woźniak, L. A.; Koziołkiewicz, M.; Kobylańska, A.; Stec, W. J., Potassium peroxymonosulfate (oxone) — An efficient oxidizing agent for phosphothio compounds. Bioorg. Med. Chem. Lett. 1998, 8, (19), 2641-2646. 36. Bunton, C. A.; Foroudian, H. J.; Kumar, A., Sulfide oxidation and oxidative hydrolysis of thioesters by peroxymonosulfate ion. Journal of the Chemical Society, Perkin Transactions 2 1995, (1), 33-39. 37. Ramachandram, M. S.; Vivekanandam, T. S.; Raj, R. P. M. M., Kinetics and mechanism of the oxidation of amino acids by peroxomonosulphate. Part 2. Effect of formaldehyde. Journal of the Chemical Society, Perkin Transactions 2 1984, (8), 1345-1349. 38. Kannan, R. S.; Ramachandran, M., Studies on the autocatalyzed oxidation of amino acids by peroxomonosulfate. Int. J. Chem. Kinet. 2003, 35, (10), 475-483. 39. Chesney, A. R.; Booth, C. J.; Lietz, C. B.; Li, L.; Pedersen, J. A., Peroxymonosulfate Rapidly Inactivates the Disease-Associated Prion Protein. Environ. Sci. Technol. 2016, 50, (13), 7095-7105. 40. Haag, W. R.; Gassman, E., Singlet oxygen in surface waters—Part I: Furfuryl alcohol as a trapping agent. Chemosphere 1984, 13, (5-6), 631-640. 41. Evans, D. F.; Upton, M. W., Studies on singlet oxygen in aqueous solution. Part 1. Formation of singlet oxygen from hydrogen peroxide with two-electron oxidants. J. Chem. Soc., Dalton Trans. 1985, (6), 1141-1145. 42. Miyamoto, S.; Martinez, G. R.; Martins, A. P. B.; Medeiros, M. H. G.; Di Mascio, P., Direct Evidence of Singlet Molecular Oxygen [O2 (1∆g)] Production in the Reaction of Linoleic Acid Hydroperoxide with Peroxynitrite. J. Am. Chem. Soc. 2003, 125, (15), 4510-4517. 43. Goodman, J.; Robson, P., 534. Decomposition of inorganic peroxyacids in aqueous alkali. Journal of the Chemical Society (Resumed) 1963, 2871-2875. 44. Koubek, E.; Levey, G.; Edwards, J. O., An Isotope Study of the Decomposition of Caro's Acid. Inorg. Chem. 1964, 3, (9), 1331-1332. 45. Lindig, B. A.; Rodgers, M. A. J.; Schaap, A. P., Determination of the lifetime of singlet oxygen in water-d2 using 9,10-anthracenedipropionic acid, a water-soluble probe. J. Am. Chem. Soc. 1980, 102, (17), 5590-5593. 46. Davies, M. J.; Gilbert, B. C.; Jonathan K. Stell, t. l.; Whitwood, A. C., Nucleophilic Substitution Reactions of Spin Adducts. Implications for the Correct Identification of Reaction Intermediates by EPR/Spin Trapping. Journal of Chemical Society, Perkin Transactions 2 1992, 3, (23), 333-335. 47. Appiani, E.; Ossola, R.; Latch, D. E.; Erickson, P. R.; McNeill, K., Aqueous singlet oxygen reaction kinetics of furfuryl alcohol: effect of temperature, pH, and salt content. Environmental Science: Processes & Impacts 2017, 19, (4), 507-516. 48. Cabrerizo, F. M.; Thomas, A. H.; Lorente, C.; Dántola, M. L.; Petroselli, G.; Erra‐Balsells, R.; Capparelli, A. L., Generation of Reactive Oxygen Species during the Photolysis of 6‐(Hydroxymethyl) pterin in Alkaline Aqueous Solutions. Helv. Chim. Acta 2004, 87, (2), 349-365. 49. Scully, F. E.; Hoigné, J., Rate constants for reactions of singlet oxygen with phenols and other compounds in water. Chemosphere 1987, 16, (4), 681-694. 50. Martin, T. J., Phosphine Catalysis using Allenoates with pro-Nucleophiles or Arylidenes; Development of an Asymetric Phosphine Catalyst; and Allenes as π-Ligands in
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Environmental Science & Technology
Copper-Mediated Cross-Coupling. 2014. 51. Robinson, R.; Paabo, M.; Bates, R. G., Deuterium isotope effect on the dissociation of weak acids in water and deuterium oxide. National Bureau of Standards 1969. 52. Yang, B.; Pignatello, J. J.; Qu, D.; Xing, B., Activation of Hydrogen Peroxide and Solid Peroxide Reagents by Phosphate Ion in Alkaline Solution. Environmental Engineering Science 2016, 33, (3), 193-199. 53. Richardson, D. E.; Yao, H.; Frank, K. M.; Bennett, D. A., Equilibria, Kinetics, and Mechanism in the Bicarbonate Activation of Hydrogen Peroxide: Oxidation of Sulfides by Peroxymonocarbonate. J. Am. Chem. Soc. 2000, 122, (8), 1729-1739. 54. Jiang, M.; Lu, J.; Ji, Y.; Kong, D., Bicarbonate-activated persulfate oxidation of acetaminophen. Water Res. 2017, 116, 324-331. 55. Roy, A.; Reddy, K.; Mohanta, P. K.; Ila, H.; Junjappat, H., Hydrogen peroxide/boric acid: An efficient system for oxidation of aromatic aldehydes and ketones to phenols. Synth. Commun. 1999, 29, (21), 3781-3791. 56. Davies, D. M.; Deary, M. E.; Quill, K.; Smith, R. A., Borate‐Catalyzed Reactions of Hydrogen Peroxide: Kinetics and Mechanism of the Oxidation of Organic Sulfides by Peroxoborates. Chemistry-A European Journal 2005, 11, (12), 3552-3558. 57. Yang, Y.; Pignatello, J. J., Participation of the Halogens in Photochemical Reactions in Natural and Treated Waters. Molecules 2017, 22, (10), 1684. 58. Das, T. N., Reactivity and role of SO5•- radical in aqueous medium chain oxidation of sulfite to sulfate and atmospheric sulfuric acid generation. J. Phys. Chem. A 2001, 105, (40), 9142-9155. 59. Fortnum, D. H.; Battaglia, C. J.; Cohen, S. R.; Edwards, J. O., The Kinetics of the Oxidation of Halide Ions by Monosubstituted Peroxides. J. Am. Chem. Soc. 1960, 82, (4), 778782.
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