Production of Sulfate Radical and Hydroxyl Radical by Reaction of ...

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May 19, 2015 - compounds at a high reaction rate constant (kO3,pH7 > 105. M. −1 s. −1. , e.g., diclofenac and carbamazepine).1,3 Hydroxyl radical (.
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Production of Sulfate Radical and Hydroxyl Radical by Reaction of Ozone with Peroxymonosulfate: A Novel Advanced Oxidation Process Yi Yang, Jin Jiang,* Xinglin Lu, Jun Ma,* and Yongze Liu State Key Laboratory of Urban Water Resource and Environment, Harbin Institute of Technology, Harbin, Heilongjiang 150090, China S Supporting Information *

ABSTRACT: In this work, simultaneous generation of hydroxyl radical (•OH) and sulfate radical (SO4•−) by the reaction of ozone (O3) with peroxymonosulfate (PMS; HSO5−) has been proposed and experimentally verified. We demonstrate that the reaction between the anion of PMS (i.e., SO52−) and O3 is primarily responsible for driving O3 consumption with a measured second order rate constant of (2.12 ± 0.03) × 104 M−1 s−1. The formation of both •OH and SO4•− from the reaction between SO52− and O3 is confirmed by chemical probes (i.e., nitrobenzene for •OH and atrazine for both •OH and SO4•−). The yields of •OH and SO4•− are determined to be 0.43 ± 0.1 and 0.45 ± 0.1 per mol of O3 consumption, respectively. An adduct, − O3SOO− + O3 → −O3SO5−, is assumed as the first step, which further decomposes into SO5•− and O3•−. The subsequent reaction of SO5•− with O3 is proposed to generate SO4•−, while O3•− converts to •OH. A definition of Rct,•OH and Rct,SO4•− (i.e., respective ratios of •OH and SO4•− exposures to O3 exposure) is adopted to quantify relative contributions of •OH and SO4•−. Increasing pH leads to increases in both values of Rct,•OH and Rct,SO4•− but does not significantly affect the ratio of Rct,SO4•− to Rct,•OH (i.e., Rct,SO4•−/Rct,•OH), which represents the relative formation of SO4•− to •OH. The presence of bicarbonate appreciably inhibits the degradation of probes and fairly decreases the relative contribution of •OH for their degradation, which may be attributed to the conversion of both •OH and SO4•− to the more selective carbonate radical (CO3•−). Humic acid promotes O3 consumption to generate •OH and thus leads to an increase in the Rct,•OH value in the O3/PMS process, while humic acid has negligible influence on the Rct,SO4•− value. This discrepancy is reasonably explained by the negligible effect of humic acid on SO4•− formation and a lower rate constant for the reaction of humic acid with SO4•− than with •OH. In addition, the efficacy of the O3/PMS process in real water is also confirmed.





OH proceeds via reactions 4 and 5, while O2•− undergoes electron transfer to O3 and forms •OH (reactions 6 and 7).

INTRODUCTION The application of ozone (O3) has been widely used to remove unwanted organic contaminants in both drinking water purification and wastewater reclamation.1,2 Ozone (O3) is a selective oxidant which reacts with electron-rich organic compounds at a high reaction rate constant (kO3,pH7 > 105 M−1 s−1, e.g., diclofenac and carbamazepine).1,3 Hydroxyl radical (•OH) generated from O3 consumption is a nonselective oxidant, which rapidly reacts with various organic compounds at nearly diffusion controlled rates,4 including many ozone-refractory substances. The reaction of O3 with H2O2, termed “peroxone process”, is one of the most common advanced oxidation processes to produce •OH for pollutant abatement. It was first investigated in an admirable work by Staehelin and Hoigne.5 Recently, the underlying mechanism has been revised on the basis of thermokinetic and quantumchemical data by proposing the formation of an adduct, HO5− (reaction 1), which subsequently decomposes in two parallel reactions (reactions 2 and 3).6,7 The conversion of O3•− to © 2015 American Chemical Society

HO2− + O3 → HO5−

HO5−

HO5− O3•− •−

O



+ HO2

→ 2O2 + OH

(1) •

(2)



(3)



↔ O + O2 •

+ H 2O ↔ OH + OH

(4) −

(5)

HO2• ↔ O2•− + H+

(6)

O2•− + O3 → O3•− + O2

(7)

Received: Revised: Accepted: Published: 7330

O3•−

December 31, 2014 May 14, 2015 May 19, 2015 May 19, 2015 DOI: 10.1021/es506362e Environ. Sci. Technol. 2015, 49, 7330−7339

Article

Environmental Science & Technology

by an ABTS method.23 More details are provided in the Supporting Information. Stock solutions of O3 were prepared by sparging O3 containing oxygen through 4 °C DI water and standardized spectrophotometrically by measuring absorbance at 260 nm (ε = 3200 M−1 cm−1).7 Experimental Procedures. Experiments were usually conducted in 10 mM phosphate buffer solutions spiked with a mixture of contaminants (ATZ and NB) and 40 μM tBuOH to mimic the •OH scavenging capacity of natural water (∼1 mg-C/L dissolved organic carbon (DOC), k•OH,DOC = 2.5 × 104 (mg-C/L)−1 s−1).24 All experiments were performed at 20 ± 1 °C, and solution pH was adjusted by perchloric acid and/or sodium hydroxide. Reactions were initiated by adding certain volumes of an O3 stock solution into buffered solutions containing PMS at desirable concentrations. Samples were periodically withdrawn and quenched by excess sodium thiosulfate. The experiments of humic acid effect were carried out without tBuOH. Measurements on the rate constant of O3 consumption by PMS in the presence of radical scavengers were performed at varying concentrations of PMS in excess following the procedure described above. A high concentration of tBuOH (10 mM) was added in order to competitively scavenge all • OH. Some experiments were carried out with filtrated waters from two drinking water treatment plants using groundwater (GW) and surface water (SW) as raw waters. These two water samples were filtered through the glass fiber filters (Whatman; the nominal pore size of 0.7 μm) and stored at 4 °C. SW was characterized by a low DOC and low alkalinity (DOC 2.3 mg/ L, alkalinity 0.2 mM, UV254 = 0.045, pH 7.2), while GW had a relative high DOC and high alkalinity (DOC 4.4 mg/L, alkalinity 3 mM, UV254 = 0.034, pH 7.6). To increase the buffer capacity, the water samples were buffered to pH 7 by adding 2 mM borate. Varying ionic strength by different buffer concentrations had no effect on the reaction of O 3 consumption with PMS herein (Supporting Information, Figure S3). These experiments were performed at 15 ± 1 °C to simulate drinking water treatment conditions. Analytical Methods. ATZ and NB were analyzed using a Waters 1525 HPLC with a Waters 2487 dual λ detector. Chromatographic separations were performed using a Waters symmetry C18 column (150 mm × 4.6 mm, 5 μm). The concentrations of ATZ and NB were quantified at λ = 230 and 260 nm, respectively, with an eluent of 0.1% acetic acid and methanol with ratio of 40:60 (v/v) at a flow rate of 1 mL/min. Dissolved O3 in reaction solutions was determined by the Indigo method.25 Electron paramagnetic resonance (EPR) experiments were performed on a Bruker A200 spectrometer with DMPO as a spin-trapping agent (Supporting Information, Text S1). Formaldehyde was determined as its 2,4-dinitrophenylhydrazone by HPLC.26 0.2 mL of 2,4-dinitrophenylhydrazine (9 mM in acetonitrile) and 0.1 mL of perchloric acid (1 M in acetonitrile) were mixed with 1.7 mL of sample. After standing for 45 min in the dark, the resulting solution was analyzed by HPLC at λ = 350 nm with an eluent of 0.1% acetic acid and methanol with ratio of 40:60 (v/v) at a flow rate of 1 mL/min. Characterization of the O3/PMS Process. The rates of NB and ATZ degradation can be expressed using second order kinetics by eqs 8 and 9, respectively.

In recent years, sulfate radical-based oxidation processes have received much attention for their efficient destruction of organic contaminants, such as pesticides, perfluorocarboxylic acids, and cyanotoxin.8−10 Sulfate radical (SO4•−) is a strong oxidant with a high redox potential (2.5−3.1 V)11 and reacts with many organic compounds at nearly diffusion controlled rates, which are comparable to •OH.11 SO4•− reacts with organic compounds primarily by a one-electron transfer mechanism and promotes the decarboxylation of carboxylic acids.12,13 For instance, it is reported that both aliphatic and aromatic acids undergo more efficient mineralization by SO4•− than by •OH.14,15 Additionally, SO4•− is less influenced by competing constituents relative to •OH, such as alkalinity and natural organic matter in real water,8,11,16 implying that SO4•− is more favorable to destruct high reactive organic contaminants. As another important sector of inorganic peroxides, peroxymonosulfate (PMS) is often used to produce SO4•− in many studies.17−19 The activation of PMS can be achieved by UV, transition metals, or metal oxides,17−20 similar to the cases involving H2O2.6,7 Anipsitakis and Dionysiou17 investigated the interaction of PMS with different transition metals and concluded that Co(II) and Ru(III) were the best metal catalysts for PMS activation, where sulfate radical (SO4•−) was found to be the primary radical species during these processes. Guan et al.18 reported that the photolysis of PMS, through the cleavage of the O−O bond by UV, generated both •OH and SO4•−, which effectively destructed benzoic acid. Interestingly, a recent study reported an enhancement of ozonation on pchlorobenzoic acid degradation by PMS, indicating that PMS might have a similar effect as H2O2 to promote the production of •OH in ozonation.21 However, the fundamental aspects (e.g., reaction kinetics and yields of radicals) of the reaction between O3 and PMS have not been explored so far. In this study, we seek to investigate the reaction of O3 with PMS and elucidate possible pathways contributing to the formation of radicals (i.e., •OH and SO4•−). The first objective is to verify the formation of •OH and SO4•− in O3/PMS by using atrazine and nitrobenzene as chemical probes. The second objective is to determine the kinetics of the reaction of O3 with PMS and quantify the yields of •OH and SO4•−. The third objective is to evaluate the effects of factors (i.e., pH, bicarbonate, and humic acid) on the efficiency of contaminant degradations in O3/PMS.



MATERIALS AND METHODS Materials. Peroxymonosulfate (PMS, available as Oxone), sodium perdisulfate (PDS), atrazine (ATZ), nitrobenzene (NB), benzoic acid (BA), p-chlorobenzoic acid (pCBA), 2,4dinitrophenylhydrazine, 2,2′-azino-bis(3-ethylbenzothiazoline6-sulfonic acid) diammonium salt (ABTS), and 5,5-dimethyl1-pyrrolidine N-oxide (DMPO) were purchased from SigmaAldrich. tert-Butanol (tBuOH) was of guaranteed reagent grade and purchased from Sinopharm Chemical Reagent Co. Ltd., China. A commercial humic acid (Sigma-Aldrich) was used as a model natural organic matter (NOM) and was purified following the procedure described in the literature.22 All other reagents were of reagent grade quality or higher and were used without further purification. All solutions were prepared using deionized (DI) water (18.2 MΩ/cm) from a Milli-Q purification system (Millipore, Billerica, MA). Stock solutions of PMS were freshly prepared by dissolving weighed amounts of Oxone in DI water and standardized spectrophotometrically 7331

DOI: 10.1021/es506362e Environ. Sci. Technol. 2015, 49, 7330−7339

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Environmental Science & Technology −



d[NB] = k •OH,NB[NB][•OH] dt

f•OH,ATZ =

d[ATZ] = k O3,ATZ[ATZ][O3] + k •OH,ATZ[ATZ][•OH] dt + k SO4•−,ATZ[ATZ][SO4•−]

=

∫ [•OH]dt

⎛ [ATZ] ⎞ ln⎜ ⎟ = −(k O3,ATZ ⎝ [ATZ]0 ⎠

∫ [•OH]dt + k SO

•−

4

fSO •−,ATZ = 4

fO ,ATZ = 3

Rct ,SO4•− =

∫ [SO4•−]dt /∫ [O3]dt

(11)

(12)

∫ [O3]dt

RESULTS AND DISCUSSION Degradation of Contaminants by O3/PMS. The degradation of ATZ by ozonation in the presence of PMS, PDS, or H2O2 was compared in Figure 1. O3/PMS achieved 81% removal of ATZ in 10 min, while O3 alone only achieved 27% removal in 20 min. The first order rate constant of O3 consumption also increased by a factor of 2 with the addition of PMS (Supporting Information, Figure S7). No degradation of PMS, ATZ, and NB was observed in the control without O3 (Supporting Information, Figure S8). Although O3 consumption was faster in O3/H2O2, less ATZ was removed compared with that in O3/PMS. On the contrary, the effects of PDS on the O3 consumption and ATZ degradation were negligible in the investigated time scale. The discrepancy between PMS and PDS implied that the neighboring structures of peroxy group (O−O) determine the reactivity of peroxides toward O3. The species of PMS in the pH range of current interest exist as its monoanion form (HSO5−) and dianion form (SO52−). PDS only exists in a form of S2O82−, in which the peroxy group (O− O) bridges sulfonate group (Supporting Information, Figure S9). It is well-known that sulfonate group does not react with O3.7 The reaction of O3 with HO2− and peroxyl radicals are assumed to form an adduct as the first step.7,33 Thus, we

(14)

⎛ [ATZ] ⎞ ln⎜ ⎟ = −(k O3,ATZ + k •OH,ATZRct , •OH ⎝ [ATZ]0 ⎠ + k SO4•−,ATZRct ,SO4•−)

∫ [O3]dt

k O3,ATZ + k •OH,ATZRct ,•OH + k SO4•−,ATZRct ,SO4•−



(13)

Substitution of eqs 12 and 13 into eqs 10 and 11 yields ⎛ [NB] ⎞ ln⎜ ⎟ = −k •OH,NBRct , •OH ⎝ [NB]0 ⎠

k O3,ATZ

Determination of Rate Constants for the Reactions of ATZ, NB, tBuOH, and Humic Acid with •OH and SO4•−. Since the reported second order rate constants for the reactions of ATZ and NB with •OH and SO4•− varied appreciably among previous studies (k•OH,ATZ = (2.5−3.0) × 109 M−1 s−1,29,30 k•OH,NB = (3.9−4.7) × 109 M−1 s−1,4 and kSO4•−,ATZ = (2.6−3.5) × 109 M−1 s−1 8,31,32), which affected the model calculations significantly, we reinvestigated the second order rate constants for their reactions with •OH and SO4•− by competition kinetics.7 For tBuOH and humic acid, the second order rate constants were evaluated with varying concentration ratios of BA to tBuOH or humic acid. More details about the competition kinetics are provided in the Supporting Information. The second order rate constants determined from this study were k•OH,ATZ = (2.6 ± 0.1) × 109 M−1 s−1, k•OH,NB = (3.9 ± 0.1) × 109 M−1 s−1, kSO4•−,ATZ =(2.6 ± 0.1) × 109 M−1 s−1, and kSO4•−,tBuOH = (7.6 ± 0.9) × 105 M−1 s−1, respectively. The latter one was also reported in the range of (4.0−9.1) × 105 M−1 s−1 in the literature.11 The measured rate constants for the reaction of humic acid with •OH and SO4•− were (2.5 ± 0.4) × 104 (mg-C/L)−1 s−1 and (5.1 ± 0.5) × 103 (mg-C/L)−1 s−1, respectively, consistent with previous studies as well.8,24



∫ [•OH]dt /∫ [O3]dt

k O3,ATZ + k •OH,ATZRct ,•OH + k SO4•−,ATZRct ,SO4•−

(18)

(10)

∫ [SO4•−]dt )

k SO4•−,ATZRct ,SO4•− (17)

where ∫ [O3]dt (i.e., O3 exposure), ∫ [•OH]dt (i.e., •OH exposure), and ∫ [SO4•−]dt (i.e., SO4•− exposure) represent the time-integrated concentrations of O3, •OH, and SO4•−, respectively. A term Rct,•OH is defined as the ratio of •OH exposure to O3 exposure (eq 12),28 while a term Rct,SO4•− is defined to describe the ratio of SO4•− exposure to O3 exposure (eq 13). Rct ,•OH =

k •OH,ATZRct ,•OH k O3,ATZ + k •OH,ATZRct ,•OH + k SO4•−,ATZRct ,SO4•−

(16)

∫ [O3]dt + k OH,ATZ

,ATZ

k O3,ATZ ∫ [O3]dt + k •OH,ATZ ∫ [•OH]dt + k SO4•−,ATZ ∫ [SO4•−]dt

(9)

where k•OH,NB is the second order rate constant for the reaction of •OH with NB and kO3,ATZ, k•OH,ATZ, and kSO4•−,ATZ are the second order rate constants for reactions of ATZ with O3, •OH, and SO4•−, respectively. The reactions of NB with O3 and SO4•− are negligible because of their very low reaction rate constants (kO3,NB = 0.09 M−1 s−1 27 and kSO4•−,NB < 106 M−1 s−1 11). Rearranging and integrating eqs 8 and 9 obtains ⎛ [NB] ⎞ ln⎜ ⎟ = −k •OH,NB ⎝ [NB]0 ⎠

k •OH,ATZ ∫ [•OH]dt

(8)

(15)

Therefore, the values of Rct,•OH and Rct,SO4•− can be calculated by plotting the logarithm of contaminant degradations vs O3 exposure. The fraction of ATZ reacting with •OH and SO4•− can be calculated by using Rct,•OH (eq 16) and Rct,SO4•− (eq 17). The fraction of ATZ oxidized by O3 is expressed by eq 18. Such calculations enable estimating their relative contributions of O3, • OH, and SO4•− on ATZ degradation. 7332

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concentration completely scavenges •OH but partially scavenges SO4•−. NB was selected as a probe of •OH (k•OH,NB = (3.9 ± 0.1) × 109 M−1 s−1), because of its very low reaction rate constant with SO4•− (kSO4•−,NB < 106 M−1 s−1).11 In contrast, due to the medium rate constants for the reaction of ATZ with • OH (k•OH,ATZ = (2.6 ± 0.1) × 109 M−1 s−1) and SO4•− (kSO4•−,ATZ = (2.6 ± 0.1) × 109 M−1 s−1), ATZ was used to confirm the formation of SO4•− in the presence of 2 mM tBuOH, which scavenged 13% of SO4•−. The results of quenching experiments (Figure 1B) showed that the degradation of NB was completely inhibited by 2 mM tBuOH, while the corresponding elimination of ATZ was 62%. This indicated that both •OH and SO4•− were responsible for ATZ degradation, confirming the simultaneous generation of •OH and SO4•− in O3/PMS. EPR spectra further verified that both •OH and SO4•− were generated in O3/PMS (Supporting Information, Figure S10). Kinetics of the Reaction of O3 with PMS. Assuming that the peroxy group of SO52− was responsible for driving an acceleration of O3 consumption, the rate of O3 consumption should be significantly promoted with increasing pH. To this end, we determined the rate constant of O3 consumption in the presence of 10 mM tBuOH with varying concentrations of PMS in excess. •OH was scavenged by tBuOH, and SO4•− was assumed to be unreactive toward O3 based on a previous study.33 O3 consumption followed pseudo-first order kinetics with rate constant, kobs,O3, obtained from the slope of the logarithm of residual O3 vs time. The rate constant for the reaction of O3 consumption with PMS, kobs,O3,PMS, was calculated by eq 25:

Figure 1. (A) Degradation of ATZ in O3/H2O2, O3/PMS, and O3/ PDS. [O3]0 = 1 mg/L, [ATZ]0 = 1 μM, [NB]0 = 1 μM, [tBuOH]0 = 40 μM, [H2O2]0 = 10 μM, [PMS]0 = 10 μM, [PDS]0 = 20 μM. (B) Inhibition effect of radical scavenger on the degradations of ATZ and NB in O3/PMS. [O3]0 = 1 mg/L, [ATZ]0 = 1 μM, [NB]0 = 1 μM, [PMS]0 = 10 μM. All solutions contained 10 mM phosphate buffer at pH 8.

kobs,O3,PMS = kobs,O3 − k O3,control

where kO3,control was the rate constant of O3 consumption without PMS in control experiments. As shown in Figure 2A, kobs,O3,PMS exhibited a linear relationship with PMS concentration at a given pH, indicating second order kinetics of the reaction between O3 and PMS. The rate of O3 consumption can be described by following equations:

assume that the peroxy group of SO52− can react with O3 to form an adduct (−O3SO5−) (reaction 19). This adduct may decompose in two parallel pathways (reactions 20 and 21). O3•− converts to •OH according to reactions 4 and 5. The subsequent reaction of SO5•− with O3 is proposed to generate SO4•− and O2 (reaction 22).33 SO5•− also decays bimolecularly (reactions 23 and 24).33 Therefore, it is likely that the O3/PMS process can simultaneously produce •OH and SO4•−, consistent with the speculation of a previous study.21 −

O3SOO− + O3 → −O3SO5−

(19)



O3SO5− → SO5•− + O3•−

(20)



O3SO5− → SO4 2 − + 2O2 k = 1.6 × 105 M−1s−1

(22)

2SO5•−

→ 2SO4

•−

+ O2

2SO5•− → S2O82 − + O2

8

−1 −1

k = 2.1 × 10 M s

(23)

k = 2.2 × 108 M−1s−1

(24)

d[O3] = −kobs,O3,PMS[O3] dt

(26)

⎛ d[O3] ⎞ ⎜ ⎟ = −k O3,PMS[PMS]T [O3] ⎝ dt ⎠ pH

(27)

where kO3,PMS is the second order rate constant for the reaction at a given pH and obtained from the slopes of lines in Figure 2A and [PMS]T is the total concentration of PMS. Meanwhile, it was found that kO3,PMS increased by 1 order of magnitude per pH unit increase (Figure 2A inset). This observation suggested that PMS reacted with O3 in its deprotonated form (SO52−), similar to the case of H2O2. Figure 2B depicts an increase in kobs,O3,PMS with SO52− concentration, which was calculated on the basis of pKa2 of PMS. This indicated that kobs,O3,PMS was first order in SO52− concentration (Supporting Information, Text S5), expressed as eq 28:

(21)

SO5•− + O3 → SO4•− + 2O2

(25)

To identify radical species produced in O3/PMS, quenching experiments were performed with chemical probes, ATZ and NB. The rate constant for the reaction of tBuOH with •OH (k•OH,tBuOH = 6.0 × 108 M−1 s−1)4 is about 3 orders of magnitude greater than with SO4•− (kSO4•−,tBuOH = (7.6 ± 0.9) × 105 M−1 s−1), indicating that tBuOH at a relatively high

d[O3] = −k O3,SO52−[SO52 −][O3] dt 7333

(28)

DOI: 10.1021/es506362e Environ. Sci. Technol. 2015, 49, 7330−7339

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Environmental Science & Technology

rate of ATZ degradation is given by eq 29, while the rate of SO4•− degradation is given by eq 30. −

d[ATZ] = k SO4•−,ATZ[ATZ][SO4•−] dt



d[SO4•−] = (k SO4•−,ATZ[ATZ] + k SO4•−,PMS[PMS] dt

(29)

+ k SO4•−,tBuOH[tBuOH])[SO4•−]

(30)

The concentrations of PMS (i.e., 1 mM) and tBuOH (i.e., 100 mM) in large excess can be considered to be constant as the reactions progress. According to the mathematical derivation in a previous study,38 substitution [SO4•−] from eq 29 yields k SO4•−,ATZ[ATZ] + k SO4•−,PMS[PMS] + k SO4•−,tBuOH[tBuOH] k SO4•−,ATZ[ATZ]

d[ATZ]

= d[SO4•−]

(31)

Upon integrating eq 31 with the limits of starting and end concentrations (at infinite time in theory), the yield of SO4•− (ΦSO4•−) is obtained. k SO4•−,PMS[PMS] + k SO4•−,tBuOH[tBuOH]

[ATZ]0 − [ATZ]∞ +

k SO4•−,ATZ

⎛ [ATZ]0 ⎞ × ln⎜ ⎟ = [SO4•−]0 = [O3]0 ΦSO4•− ⎝ [ATZ]∞ ⎠

(32)

The term [ATZ]0 − [ATZ]∞ + ((kSO4•−,PMS[PMS] + kSO4•−,tBuOH[tBuOH])/(kSO4•−,ATZ)) ln([ATZ]0/[ATZ]∞), defined as Aexp, can be experimentally obtained by monitoring the decay of ATZ in a trace amount (1 μM) as a function of O3 concentration. A plot of Aexp vs O3 concentration gave a straight line in Figure 3A, suggesting that the yield of SO4•− was 0.45 ± 0.01. Determination the •OH Yield by tBuOH Assay. The tBuOH assay is developed to determine the •OH yield by quantification of formaldehyde formation, where the •OH yield is twice the formaldehyde yield.7 However, little is known about the reaction of SO4•− with tBuOH. According to the reaction mechanisms of SO4•− with organic compounds,13 it is inferred that the reaction of SO4•− with tBuOH can also form a tBuOHderived radical (reaction 33), which will react with O2, giving rise to its corresponding peroxyl radical (reaction 34).39 The decay of this peroxyl radical yields formaldehyde in further reactions.39

Figure 2. (A) Measured pseudo-first order rate constants of O3 consumption vs PMS concentration at different pH values. Inset indicates logarithm of second order rate constant kobs,O3,PMS vs pH. (B) Measured pseudo-first order rate constants of O3 consumption vs SO52− concentration. [O3]0 = 1 mg/L and [tBuOH]0 = 10 mM. All solutions contained 10 mM phosphate buffer.

where kO3SO52− was (2.12 ± 0.03) × 104 M−1 s−1, obtained from the slope of the line in Figure 2B. The rate constant for the reaction of O3 with SO52− is lower than with HO2− ((2.12 ± 0.03) × 104 vs (9.6 ± 2) × 106 M−1 s−1).34 This may be ascribed to the strongly electronegative sulfonate group in SO52−, which decreases its reactivity toward electrophilic O3 compared to HO2−. pKa2 of PMS (9.4)35 is lower than H2O2 (11.8).36 This implies that SO52− concentration is about 2 orders of magnitude higher than HO2− at the same concentration of PMS and H2O2 for a given pH. Hence, the lower rate constant for the reaction of O3 with PMS can be compensated by the higher concentration of SO52− species, thereby exhibiting a comparable O3 consumption rate to that in O3/H2O2. The calculated activation energy of O3 consumption initiated by PMS was (75.7 ± 4.6) kJ mol−1, while that by H2O2 determined in this study was (66.8 ± 4.0) kJ mol−1 (Supporting Information, Figure S11). The latter one was comparable with the value of 73.5 kJ mol−1 reported in the literature.37 The •OH and SO4•− Yields from the Reaction of O3 with PMS. Determination the SO4•− Yield by Competition Kinetics. A competition kinetic method by using one •OH scavenger (e.g., tBuOH) and one reference compound (e.g., ATZ) has been suggested to quantify the •OH yield in the reaction of O3 with H2O2.38 This method is adopted for the quantification of the SO4•− yield in O3/PMS by monitoring ATZ degradation, when tBuOH is in large excess to completely scavenge •OH but partially scavenge SO4•−. In this case, the

SO4•− + (CH3)3 COH → SO4•− + H+ + •CH 2C(CH3)2 OH (33) •



HOC(CH3)2 CH 2 + O2 → HOC(CH3)2 CH 2OO

(34)

Although the rate constant for the reaction of tBuOH with SO4•− is very low, its contribution on formaldehyde formation cannot be negligible. The formaldehyde yield of SO4•− in UV/ PDS was determined to be 0.53 ± 0.006 (Supporting Information, Figure S12), comparable to the formaldehyde yield of •OH (i.e., 0.49 ± 0.003) in UV/H2O2. Therefore, 1 M tBuOH was used to completely scavenge •OH and SO4•− in O 3 /PMS. A slope of formaldehyde formation vs O 3 concentration is 0.44 (Figure 3B). Because both yields of • OH and SO4•− are twice the formaldehyde yield, the sum of • OH and SO4•− yields was 0.88. Subtracting the contribution of 7334

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Figure 3. (A) Competition kinetics of tBuOH with ATZ for SO4•−. Consumption of total organic compounds by SO4•− (Aexp) as a function of O3 concentration. [PMS]0 = 1 mM, [ATZ]0 = 1 μM, [tBuOH]0 = 100 mM. (B) The formation of formaldehyde as a function of O3 concentration by the tBuOH assay. [PMS]0 = 1 mM and [tBuOH]0 = 1 M. The solid lines are fitted to experimental data. All solutions contained 2 mM phosphate buffer at pH 7.

Figure 4. Logarithm of ATZ (A) and NB (B) degradations vs O3 exposure at different pH values. All solutions contained 10 mM phosphate buffer, [O3]0 = 1 mg/L, [ATZ]0 = 1 μM, [NB]0 = 1 μM, [tBuOH]0 = 40 μM, [PMS]0 = 10 μM.

the values during the second phase), increasing pH from 7 to 8.5 resulted in increases in kobs from (1.8 ± 0.1) × 10−3 s−1 to (1.4 ± 0.1) × 10−2 s−1, Rct,SO4•− from (1.4 ± 0.2) × 10−8 to (2.3 ± 0.3) × 10−7, and Rct,•OH from (1.2 ± 0.1) × 10−8 to (1.6 ± 0.2) × 10−7. Although higher pH promoted the formation rates of •OH and SO4•− that resulted in a faster rate of ATZ and NB degradations, it had no substantial effect on the Rct,SO4•−/Rct,•OH value (i.e., 1.18−1.49). The reaction of SO52− with O3 is more important than OH−. Additionally, when pH < 9, the conversion of SO4•− to •OH is not significantly affected by pH. Both factors lead to a similar value of Rct,SO4•−/Rct,•OH at different pH values. Therefore, the fractions of ATZ reacted with •OH or SO4•− (calculated by eqs 14 and 15) were less affected by pH. We also noticed that the fraction of ATZ oxidized by O3 ( f O3) was less than 10% at pH 7 and decreased to 0.6% at pH 8.5. This was ascribed to a decrease in O3 exposure by the fast consumption of O3 with increasing pH. Bicarbonate. Bicarbonate is an important scavenger of •OH and SO4•− to generate CO3•− in natural water. Due to the lower reactivity of CO3•− relative to •OH and SO4•−,41 the conversion of •OH and SO4•− to more selective CO3•− will reduce the degradation efficiency of refractory contaminants. The eliminations of ATZ and NB decreased from 81% to 54% and from 73% to 33%, respectively, with increasing bicarbonate concentration from 0 to 5 mM (Supporting Information, Figure S14). However, CO3•−, as a weak oxidant, can selectively react with some contaminants.42,43 We assumed that CO3•− is essentially unreactive toward NB, considering that SO4•− is

SO4•− determined from the competition kinetics, the •OH yield of 0.43 ± 0.01 was obtained, suggesting that the yield of • OH was almost equal to SO4•− in O3/PMS. The formation of an adduct, −O3SO5−, in the first step (reaction 19) has been assumed. The experimental observation that the total efficiency of SO4•− and •OH formations is ∼90% can be rationalized if − O3SO5− decays predominantly to form SO5•− and O3•− by reaction 20. The subsequent reactions of SO5•− (reaction 22) and O3•− (reactions 4 and 5) are proposed to generate SO4•− and •OH equally. The bimolecular reactions 23 and 24 can be neglected in O3/PMS. Effect of Water Quality Parameters on the O3/PMS Process. Solution pH. The rates of ATZ and NB degradation were promoted with increasing pH (Supporting Information, Figure S13). By plotting the left-hand side of eqs 14 and 15 vs the O3 exposure, the values of Rct,•OH and Rct,SO4•− can be extracted from the slopes of the plots in Figure 4A,B, respectively. Due to the trace impurity of DOC in reaction solution, O3 consumption was observed to be kinetically divided into an initial and a second phase.40 Linear relationships between the logarithm of contaminant degradations (i.e., ATZ and NB) and O3 exposure were observed during the second phase of O3 consumption (t > 20 s), indicating that the Rct,•OH and Rct,SO4•− values remained constant. The influence of pH on the O3 consumption, •OH formation, and SO4•− formation could be evaluated by the values of kobs, Rct,•OH, and Rct,SO4•−, respectively. As shown in Table 1 (only including 7335

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Table 1. O3 Consumption Rate Constant (kobs), Rct,•OH and Rct,SO4•− Values, and f•OH,ATZ and f SO4•−,ATZ Values in Various Conditions pHa kobs (s−1) 7 (1.8 7.5 (3.4 8 (6.0 8.5 (1.4 bicarbonatea

± ± ± ±

0.1) 0.1) 0.5) 0.1)

× × × ×

Rct,•OH 10−3 10−3 10−3 10−2

(1.2 (3.2 (8.5 (1.6

± ± ± ±

0.1) 0.3) 0.9) 0.2)

Rct,SO4•−

× × × ×

10−8 10−8 10−8 10−7

(1.4 (3.7 (9.5 (2.3

kobs (s−1) 0 mM 0.5 mM 1 mM 3 mM 5 mM humic acidb

(6.0 (5.7 (5.0 (4.8 (4.6

± ± ± ± ±

0.5) 0.5) 0.4) 0.3) 0.3)

× × × × ×

SW O3 O3/H2O2 O3/PMS GW O3 O3/H2O2 O3/PMS

0.2) 0.3) 1.2) 0.3)

× × × ×

Rct,SO4•−/Rct,•OH

f O3,ATZ (%)

f•OH,ATZ (%)

f SO4•−,ATZ (%)

1.19 1.18 1.11 1.49

8.4 3.2 1.3 0.6

41.9 44.5 46.7 40.0

49.7 52.3 52.0 59.4

10−8 10−8 10−8 10−7

Rct,•OH 10−3 10−3 10−3 10−3 10−3

(8.5 (7.9 (5.6 (3.4 (2.2

kobs (s−1) pH 7, O3 0.1 mg-C/L 0.2 mg-C/L 0.3 mg-C/L 0.5 mg-C/L pH 7, O3/PMS 0.1 mg-C/L 0.2 mg-C/L 0.3 mg-C/L 0.5 mg-C/L pH 8, O3 0.1 mg-C/L 0.2 mg-C/L 0.3 mg-C/L pH 8, O3/PMS 0.1 mg-C/L 0.2 mg-C/L 0.3 mg-C/L water matricesc

± ± ± ±

± ± ± ± ±

0.9) 0.8) 0.5) 0.3) 0.2)

× × × × ×

Rct,SO4•−

± ± ± ±

0.2) 0.4) 0.1) 0.2)

× × × ×

10−3 10−3 10−2 10−2

(1.1 (1.7 (2.6 (5.1

± ± ± ±

0.1) 0.1) 0.2) 0.4)

× × × ×

10−7 10−7 10−7 10−7

(4.1 (7.8 (1.2 (2.2

± ± ± ±

0.3) 0.5) 0.1) 0.2)

× × × ×

10−3 10−3 10−2 10−2

(2.0 (2.9 (3.9 (6.3

± ± ± ±

0.1) 0.1) 0.2) 0.4)

× × × ×

10−7 10−7 10−7 10−7

(5.5 ± 0.5) × 10−3 (9.2 ± 0.8) × 10−3 (1.4 ± 0.2) × 10−2

(1.8 ± 0.1) × 10−7 (3.6 ± 0.4) × 10−7 (5.9 ± 0.7) × 10−7

(1.5 ± 0.1) × 10−2 (1.9 ± 0.2) × 10−2 (2.4 ± 0.3) × 10−2

(6.6 ± 0.7) × 10−7 (7.5 ± 0.7) × 10−7 (8.1 ± 0.9) × 10−7

f•OH,ATZ (%)

fΣR•,ATZ (%)d

1.3 1.4 1.6 2.5 3.5

46.7 44.6 39.0 36.8 33.7

52.0 54.0 59.4 60.7 62.8

10−8 10−8 10−8 10−8 10−8

Rct,•OH

(3.1 (6.3 (1.1 (1.9

f O3,ATZ (%)

Rct,SO4•−/Rct,•OH

(4.6 ± 0.9) × 10−8 (4.3 ± 1.0) × 10−8 (4.8 ± 1.6) × 10−8 N.A.e

(3.5 ± 0.7) × 10−7 (3.5 ± 0.7) × 10−7 (3.7 ± 0.9) × 10−7

f O3,ATZ (%)

f•OH,ATZ (%)

4.2 1.7 1.1 0.7

95.8 98.3 98.9 99.3

1.0 0.7 0.5 0.5

80.5 86.5 88.6 99.5

1.3 0.6 0.4

98.7 99.4 99.6

0.2 0.2 0.2

65.2 68.0 68.5

0.23 0.15 0.12 N.A.e

0.53 0.47 0.46

f SO4•−,ATZ (%)

18.5 12.8 10.9 N.A.e

34.6 31.8 31.3

kobs (s−1)

Rct,•OH

f O3,ATZ (%)

f•OH,ATZ (%)

fΣR•,ATZ (%)d

(1.3 ± 0.1) × 10−3 (7.3 ± 0.5) × 10−3 (3.7 ± 0.2) × 10−3

(1.7 ± 0.1) × 10−8 (7.5 ± 0.6) × 10−8 (3.7 ± 0.4) × 10−8

14.6 3.0 3.3

85.4 97.0 79.8

16.9

(3.7 ± 0.3) × 10−3 (2.0 ± 0.2) × 10−2 (1.2 ± 0.1) × 10−2

(7.1 ± 0.5) × 10−9 (4.0 ± 0.3) × 10−8 (1.4 ± 0.1) × 10−8

24.4 5.4 8.2

75.6 94.6 50.4

41.4

Solutions contained 10 mM phosphate buffer at pH 8 with [O3]0 = 1 mg/L, [ATZ]0 = 1 μM, [NB]0 = 1 μM, [PMS]0 = 10 μM, and [tBuOH]0 = 40 μM. bSolutions contained 3 mM phosphate buffer at pH 8 with [O3]0 = 1 mg/L, [ATZ]0 = 1 μM, [NB]0 = 1 μM, and [PMS]0 = 10 μM. cThe water samples were buffered by 2 mM borate at pH 7. [O3]0 = 2 mg/L, [ATZ]0 = 1 μM, [NB]0 = 1 μM, [PMS]0 = 20 μM, [H2O2]0 = 20 μM. dfΣR•,ATZ included the fraction of ATZ oxidized by SO4•− and CO3•−. eN.A. means “not available”. Due to the very low value of Rct,SO4•−/Rct,•OH, the Rct,SO4•− value could not be accurately determined in the presence of 0.5 mg-C/L humic acid. a

inert to NB. In this case, NB was still considered as an •OH probe. The second order rate constant for the reaction of CO3•− with ATZ is 3.7 × 106 M−1 s−1,41 suggesting that the oxidation involving SO4•− and CO3•− could not be differentiated by ATZ. Therefore, the effect of bicarbonate on ATZ degradation was evaluated by the fraction of ATZ oxidized by both SO4•− and CO3•− radicals, fΣR•,ATZ, rather than individual SO4•−, f SO4•−,ATZ. The value of fΣR•,ATZ can be calculated by eq 35:

f∑ R•,ATZ = 1 − f•OH,ATZ − fO ,ATZ 3

(35)

Table 1 shows the calculated fractions of ATZ in the reactions with O3, •OH, and other radicals (i.e., SO4•−.and CO3•−). Varying bicarbonate concentrations from 0 to 5 mM increased fΣR•,ATZ by 21% but reduced f•OH,ATZ by 28%. The second order rate constants of bicarbonate with •OH and SO4•− are (0.85−1) × 107 M−1 s−1 4 and (0.28−0.91) × 107 M−1 s−1,44,45 respectively, suggesting comparable scavenging 7336

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Environmental Science & Technology effects of bicarbonate on •OH vs SO4•−. Therefore, the increase in fΣR•, ATZ with bicarbonate concentration might be ascribed to the reaction between CO3•− and ATZ. A previous study also demonstrated that CO3•− may contribute to the degradation of atrazine in UV/S2O82−.46 Humic Acid. NOM in water treatment is reactive toward O3 and •OH. Specific moieties in NOM (e.g., phenols and amines) directly react with O3 to generate •OH.40 The reaction of NOM with •OH can accelerate O3 consumption via a propagation step.47,48 The presence of NOM also competes with organic contaminants for oxidants and thus decreases their eliminations in water matrices. Humic acid, an important constituent of NOM, was selected as a surrogate to evaluate the effect of NOM on organic contaminant destructions in O3/ PMS.47 We chose easily accessible Sigma-Aldrich humic acid in this study. Previous studies found a strong correlation between O3 consumption rate and UV absorbance at 254 nm (UV254).49,50 Specific ultraviolet absorbance (SUVA) of humic acid determined at 254 nm is 0.108 (mg-C/L)−1 cm−1, which is about 2−8 times higher than NOM isolated from natural water sources.49 Therefore, a lower concentration of humic acid (i.e., 0.1−0.5 mg-C/L) was used to maintain the UV254 of a solution at 0.0108−0.054 cm−1, comparable to most drinking water sources (0.01−0.1 cm−1, mostly around 0.03 cm−1).50 As shown in Table 1, an increase of humic acid concentration accelerated O3 consumption and •OH formation as indicated by the increases of kobs and Rct,•OH in the absence of PMS, while the presence of PMS further increased the values of kobs and Rct,•OH. Figure 5 depicts linear relationships between kobs and

and (5.1 ± 0.5) × 103 (mg-C/L)−1 s−1, respectively, suggesting that the effect of humic acid was more pronounced in the •OHbased process than in the SO4•−-based process. This was generally consistent with the finding (see Table 1) that humic acid concentration negligibly affected the Rct,SO4•− value. Due to the increase in the Rct,•OH value, the Rct,SO4•−/Rct,•OH value was reduced with increasing humic acid concentration. For instance, the Rct,SO4•−/Rct,•OH value at pH 7 decreased by a factor of 2 when humic acid concentration varied from 0.1 to 0.3 mg-C/L. After complete consumption of O3, the residual PMS was 7.3 μM and 5.0 μM on average at pH 7 and 8, respectively (Supporting Information, Table S1). Its impact on the following water treatment processes needs further studies. Comparison of O3/PMS and O3/H2O2 in Authentic Water Matrices. To evaluate the efficiency of O3/PMS compared with O3/H2O2, the degradation of ATZ was investigated in two authentic water matrices (Figure 6). An increase of DOC

Figure 6. Degradation of ATZ in O3/PMS and O3/H2O2 in authentic water matrices. [O3]0 = 2 mg/L, [ATZ]0 = 1 μM, [NB]0 = 1 μM, [PMS]0 = 20 μM, [H2O2]0 = 20 μM, T = 15 °C.

concentration accelerated O3 consumption as shown in Table 1. Due to the higher concentration of carbonates scavenging • OH, the Rct,•OH value in GW was lower than in SW. This also resulted in a lower elimination of ATZ in GW. The rate of ATZ degradation in O3/H2O2 was faster than in O3/PMS. No enhancement of ATZ removal observed in O3/PMS might be ascribed to the presence of Cl− (∼0.15 mM). Cl− reacts rapidly with SO4•− to form Cl-containing radicals (3.0 × 108 M−1 s−1),51,52 which might decrease the efficiency of ATZ degradation. In contrast, Cl− had no significant effect on • OH-based AOPs at neutral pH.42 This present work demonstrates the combination of O3 and PMS as a novel advanced oxidation process to simultaneously produce •OH and SO4•−. In O3/PMS, SO4•− is favorable toward degradation of reactive organic contaminants, while • OH contributes to the destruction of SO4•−-refractory organics. SO4•− is less reactive toward dissolved organic matter (DOM), indicating that O3/PMS is more suitable in drinking water production. High DOM in wastewater will compete with PMS for O3 to produce •OH. Furthermore, the O3/PMS process is attractive because of good stability and easy handling of PMS in solid form. However, due to the high rate constants for the reactions of SO4•− with halides, further investigations will be conducted to evaluate their effects on organic contaminant degradations as well as halogenated byproduct formations in the O3/PMS process.

Figure 5. Effect of humic acid on O3 consumption rate. All solutions contained 3 mM phosphate buffer, [O3]0 = 1 mg/L, [ATZ]0 = 1 μM, [NB]0 = 1 μM, [PMS]0 = 10 μM.

humic acid concentration for both pH 7 and pH 8. The slope of the line at pH 8 was comparable to that at pH 7, suggesting that the increase in kobs by humic acid was not significantly affected by pH in the investigated range. However, as discussed above, increasing pH by one unit increased the rate constant for the reaction of O3 with PMS by 1 order of magnitude. Thus, the Rct,SO4•− value on average increased from (4.6 ± 0.3) × 10−8 at pH 7 to (3.6 ± 0.1) × 10−7 at pH 8, resulting in higher values of Rct,SO4•−/Rct,•OH at pH 8 relative to pH 7. Additionally, humic acid is also a scavenger of •OH and SO4•−. The eliminations of ATZ and NB (Supporting Information, Figure S15) decreased from 82% to 64% and from 85% to 75%, respectively, with increasing humic acid concentration from 0.1 to 0.5 mg-C/L at pH 7. The determined rate constants for the reactions of humic acid with •OH and SO4•− were (2.5 ± 0.4) × 104 (mg-C/L)−1 s−1 7337

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ASSOCIATED CONTENT

S Supporting Information *

Detailed descriptions of materials and methods as well as supporting figures and tables. The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/es506362e.



AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]; tel: +86 451 86283010; fax: +86 451 86283010. *E-mail: [email protected]; tel: +86 451 86283010; fax: +86 451 86283010. Notes

The authors declare no competing financial interest.



ACKNOWLEDGMENTS This research was financially supported by the National Science & Technology Pillar Program of China (No. 2012BAC05B02), the National Natural Science Foundation of China (Nos. 51178134 and 51378141), the Funds of State Key Laboratory of Urban Water Resource and Environment (HIT, 2013TS04), the Foundation for the Author of National Excellent Doctoral Dissertation of China (201346), and the Fundamental Research Funds for the Central Universities (AUGA5710056314).



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DOI: 10.1021/es506362e Environ. Sci. Technol. 2015, 49, 7330−7339

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DOI: 10.1021/es506362e Environ. Sci. Technol. 2015, 49, 7330−7339