Persulfate

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Bromate Formation from Bromide Oxidation by the UV/Persulfate Process Jing-Yun Fang and Chii Shang* Department of Civil and Environmental Engineering, The Hong Kong University of Science and Technology, Clear Water Bay, Kowloon, Hong Kong S Supporting Information *

ABSTRACT: Bromate formation from bromide oxidation by the UV/persulfate process was investigated, along with changes in pH, persulfate dosages, and bromide concentrations in ultrapure water and in bromide-spiked real water. In general, the bromate formation increased with increasing persulfate dosage and bromide concentration. The bromate formation was initiated and primarily driven by sulfate radicals (SO4•−) and involved the formation of hypobromous acid/hypobromite (HOBr/ OBr−) as an intermediate and bromate as the final product. Under the test conditions, the rate of the first step driven by SO4•− is slower than that of the second step. Direct UV photolysis of HOBr/OBr− to form bromate and the photolysis of bromate are insignificant. The bromate formation was similar for pH 4−7 but decreased over 90% with increasing pH from 7 to above 9. Less bromate was formed in the real water sample than in ultrapure water, which was primarily attributable to the presence of natural organic matter that reacts with bromine atoms, HOBr/OBr− and SO4•−. The extent of bromate formation and degradation of micropollutants are nevertheless coupled processes unless intermediate bromine species are consumed by NOM in real water.

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transition metals.10−13 It has been demonstrated effective in oxidative destruction of endocrine disrupting compounds, herbicides, phenols, perfluorinated compounds, and chlorinated compounds in water treatment and groundwater remediation fields.10,11,14−17 SO4•− has also been considered for disinfection of Escherichia coli in swimming pools.18 Its reported promising features include the persistent nature of its precursor compounds (i.e., persulfate and peroxymonosulfate) compared to peroxide,10 its versatile activating means,10−13 its wider operating pH range compared to that of the Fenton process,19 and its less influence by other competing constituents, such as alkalinity and chloride in real water.20 However, being a powerful oxidant similar to HO•, SO4•− may inherently possess the same drawback of HO•: being able to oxidize bromide to bromate. However, nothing is known about the bromate formation potential by SO4•−. In our recent study on the degradation of dibromoacetamide by UV/persulfate in ultrapure water, we discovered that up to 100% of the bromine atoms in dibromoacetamide were converted to bromate. The dibromoacetamide was exposed to UV (254 nm) light at an intensity of 2.19 μEinstein L−1·s−1 and an initial persulfate dosage of 500 μM at pH 6 for 20 min (Supporting Information Figure S1). In the UV/persulfate process, the UV light activated persulfate to generate SO4•− with a quantum yield of 1.4:19

INTRODUCTION Bromate (BrO3−) is classified as a B2, probable human carcinogen. U.S. EPA and European Commission standards and WHO guidelines for drinking water regulate bromate at a maximum contaminant level (MCL) of 10 μg/L.1,2 Bromate can be formed excessively over its drinking water standards in bromide (Br−)-containing water subjected to ozonation or certain hydroxyl radical (HO•)-based advanced oxidation processes (AOPs) under typical water treatment conditions.3,4 The bromate formation from ozonation of bromide and in HO•-based AOPs is well understood, which involves oxidation of Br− to HOBr/OBr− and then to BrO3− by ozone or HO•.3,5 Nevertheless, the existence of peroxide or electrons in some AOPs decreases the net production of bromate by chemical reduction of HOBr/OBr− or bromate to bromide, respectively.5,6 Bromate is also formed after simultaneous UV and chlorine treatment of bromide-containing water, where Br− is oxidized to HOBr/OBr− by chlorine, followed by oxidation of HOBr/OBr− to BrO3−. The second step of the stepwise oxidation is likely associated with UV photolysis, halogen radical attack, or HO• attack but the actual mechanism is unclear.7 In recent years, sulfate radicals (SO4•−) have been proposed and evaluated for micropollutant destruction, owing to their high standard reduction potential (2.5−3.1 V),8 which is comparable to that of HO• at acidic pH (2.4−2.7 V) and higher than that of HO• at alkaline pH (1.9−2.0 V).9 In general, SO4•− is more selective than HO• and reacts rapidly with organic substrates by hydrogen abstraction or addition reactions.8 SO4•− can be generated from the activation of persulfate (S2O82−) or peroxymonosulfate by UV, heat, bases or © 2012 American Chemical Society

Received: Revised: Accepted: Published: 8976

February 17, 2012 July 22, 2012 July 26, 2012 July 26, 2012 dx.doi.org/10.1021/es300658u | Environ. Sci. Technol. 2012, 46, 8976−8983

Environmental Science & Technology

Article

bromide convert bromide to HOBr/OBr− in seconds at pH 6. A NOM stock solution was prepared by dissolving an aliquot of Suwannee River RO NOM isolate (Cat. No. 1R101N, International Humic Substance Society) into ultrapure water and then the solution was filtered through a 0.45-μm membrane. Phosphate and borate buffers were used to adjust the pH from 4 to 10 and the pH values during reactions were monitored and maintained at the predetermined pH ± 0.1. The treated real water sample without final disinfection for simulating bromate formation in a more realistic water matrix was collected from the Shatin Water Treatment Works in Hong Kong, where the source water from a river goes through coagulation, sedimentation and filtration processes. The water sample was further filtered through a 0.45-μm membrane before use. Its characteristics are shown in Supporting Information Table S1. Analytical Methods. Concentrations of persulfate were determined by reacting persulfate with Fe(NH4)2(SO4)2 at acidic pH (1-M H2SO4) to form ferric ions, which further reacted with NH4SCN to form colored Fe(SCN)2+ to be read with a UV−vis spectrophotometer (Lambda25, PerkinElmer) at a wavelength of 450 nm.26 HOBr/OBr− was determined by DPD colorimetric method,24 with additions of 2-mL 5% EDTA only when persulfate was used to overcome its interference.18 Dissolved organic matter (DOC) was quantified with a total organic carbon (TOC) analyzer (TOC−VCPH, Shimadzu) and pH was measured with a pH meter (420-A, ORION). Bromide and bromate were analyzed, without further sample pretreatment, using a reagent-free ion chromatography system (ICS3000, Dionex) coupled with a conductivity detector. A highcapacity hydroxide-selective analytical column (AS19, 4 × 250 mm, Dionex) and its respective guard column (AG19, 4 × 50 mm, Dionex) were used for separation. The detection limits for bromate and bromide were 0.75 μg/L and 1.03 μg/L, respectively, with an injection volume of 250 μL. For the chromatographic method, calibration curves were established before and after a series of 30−50 samples. Concentrations of p-chlorobenzoic acid (pCBA) were determined by HPLC (VP series, Shimadzu) with a C18 column (4.6 mm × 150 mm, Waters), using a methanol-phosphate buffer (pH 2.0, 10 mM) (45:55, v/v) eluent at 1 mL/min and a detection wavelength of 233 nm. Experimental Procedures. The photochemical experiments were performed in a 750-mL cylindrical borosilicate glass reactor at ambient temperature (22 ± 2 °C). The reactor was operated in batch mode with rapid mixing provided in the bottom of the reactor. The height and inner diameter of the reactor were 35 and 6.5 cm, respectively. A low-pressure mercury UV lamp (254 nm, GPH 135T5 L/4, 10 W, Heraeus) with a quartz sleeve (outer diameter of 25 mm) was placed in the centerline of the cylindrical reactor axially along the length of the reactor. The incident light intensity of the UV lamp measured by iodide-iodate chemical actinometry27 was 2.19 μEinstein L−1·s−1. The optical path length at 254 nm was determined to be 5.2 cm by measuring the photolysis rate of H2O2.28 In all cases, a water sample that contained a specific makeup of bromide buffered (phosphate at 2 mM, or borate at 5 mM) at a given pH was added to the vessel reactor. Persulfate of a predetermined concentration was added to the reactor and immediately subjected to UV irradiation. An orthogonal matrix experimental design was employed under the baseline condition of 20-μM Br−, pH 7.0 (phosphate, 2 mM), 200-

hv

S2 O82 − → 2SO4•− Φ = 1.4

(1)

In this activation process, one mole of persulfate produces two moles of SO4•− without the production of HO•. Interconversion of SO4•− to HO• is possible at alkaline pH:8 SO4•− + OH− → HO• + SO4 2 − k = 6.5 × 107 M−1 s−1 (2)

SO4•− + H 2O → HO• + SO4 2 − k < 60 M−1 s−1

(3)

However, the reaction (eq 3) is slow and can be neglected at pH 6, suggesting that the bromate formation from the HO• pathway should not be significant. Therefore, we suspected that the degradation of dibromoacetamide by UV/persulfate led to the release of Br−, which was then oxidized to BrO3− by SO4•−. To our best knowledge, bromate formation from bromide oxidation in SO4•−-based AOPs has never been reported in the literature. Nevertheless, it is known that the reaction between Br− and SO4•− forms bromine atoms (Br•) with a rate constant of 3.5 × 109 M−1 s−1 (eq 4),8 which is higher than that of the reaction between Br− and HO• to form Br• through eq 5:9,21−23 Br − + SO4•− → Br• + SO4 2 − k = 3.5 × 109 M−1 s−1 (4) −

•

•

9

Br + HO → ... → Br k ≈ 1.1 × 10 M

−1 −1

s

(5)

It should be noted that eq 5 consists of several forward and backward reactions, some of which are pH dependent.21−23 It is also known that Br• can go through a series of reactions with Br− and H2O to form HOBr/OBr−:3 Br − + Br• → ... → HOBr/OBr −

(6)

To produce bromate, we therefore propose that HOBr/OBr−, as an intermediate, can be oxidized to BrO3− by SO4•−, in addition to the oxidation of HOBr/OBr− to BrO3− by UV photolysis.7 The bromate formation initiated by the oxidation of bromide by SO4•− is thus expected to involve stepwise conversion of Br− to HOBr/OBr− and then to BrO3−. The formation pathways deserve investigation. The objectives of this study are (1) to investigate the timedependent bromate formation from bromide oxidation in the UV (at 254 nm)/persulfate process, along with changes in pH, persulfate dosages, and bromide concentrations; (2) to investigate the possible pathways contributing to the formation; and (3) to reveal the significance of bromate formation in real water. We selected the UV/persulfate process for its simplicity and its efficient production of SO4•−.

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EXPERIMENTAL SECTION Chemicals. Chemical solutions were prepared from reagentgrade chemicals and ultrapure water produced by the Barnstead NANOpure Diamond ultrapure water system. Sodium persulfate, sodium bromide, sodium bromate, sodium chloride, and sodium nitrite were purchased from Sigma-Aldrich. All chemicals were used as received without further purification. A free chlorine (HOCl/OCl−) stock solution (3000 mg/L as Cl2) was prepared from diluting a 4% sodium hypochlorite (NaOCl, Sigma-Aldrich) solution and periodically standardized by DPD/ FAS titration.24 Bromide-free HOBr/OBr− solutions were prepared freshly by adding the free chlorine in stoichiometry to bromide-containing solutions for over 2 min at pH around 6 according to ref 25. Using the rate constant reported in this reference, the acid-assisted reactions of HOCl/OCl− and 8977

dx.doi.org/10.1021/es300658u | Environ. Sci. Technol. 2012, 46, 8976−8983


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μM persulfate. The parameters including persulfate dosage (20−500 μM), Br− concentration (0.1−50.0 μM), and pH (4− 10, 2-mM phosphate buffers for pHs 4−8, and 5-mM borate buffers for pHs 9 and 10) were varied one parameter at a time from the baseline condition. Samples (1.5 mL) were withdrawn at predetermined time intervals, quenched using excessive sodium nitrite at a nitrite to persulfate dosage ratio of around 3:1, and then subjected to bromate analysis. The formation of bromate was tracked for up to 20 min. The UV intensity, reaction times and persulfate dosages used in the current study were lower than or comparable to those used in the literature for oxidative destruction of micropollutants.10,14 20 μM (1.6 mg/L) of Br− was used to that ensure the measured concentrations of the intermediates and the products were always statistically significant. Some of the tests were conducted with tracking of the evolution of Br− to HOBr/OBr− and then to BrO3−. Another set of additional tests were conducted to evaluate bromate formation from the oxidation of HOBr/OBr− to BrO3− by UV/ persulfate in the same manner as described above except that HOBr/OBr−, not Br−, was added as the starting compound. Control tests of the formation of bromate and HOBr/OBr− from UV alone (at the same UV intensity) and persulfate alone (at the initial persulfate dosage of 200 μM) were also carried out using synthetic solutions containing 20-μM Br− or 20-μM HOBr/OBr− and buffered at pH 7. One data set (shown with error bars in each figure) was duplicated for quality control. The error bars in all data plots represent the maximum and minimum of the experimental data of the duplicated test results.

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RESULTS AND DISCUSSION Bromate Formation at Different Persulfate and Bromide Concentrations. Figure 1a shows the bromate formation in ultrapure water by UV/persulfate at an initial Br− concentration of 20 μM and different persulfate dosages (20− 500 μM) under the UV intensity of 2.19 μEinstein L−1s−1 at pH 7. The conversion of bromide to bromate by UV/persulfate in the 20-min time frame was dependent on the persulfate dosage, and rates of conversion of about 5.8%, 19.6%, 45.7%, 92.2%, and 100% were recorded at the persulfate dosages of 20, 50, 100, 200, and ≥300 μM, respectively. On the other hand, no bromate was found from UV alone and persulfate alone in the control tests. The kinetics of the bromate formation displayed a lag-phase (shown in the inset of Figure 1), especially at the low persulfate dosage of 20 μM, where the persulfate was not in excess. Otherwise, the bromate formation generally followed pseudozero-order kinetics, as both bromide and persulfate were in excess. Figure 1b shows the relationship between the pseudozero-order rate constants (k) of bromate formation and the persulfate dosages of 50−500 μM. The rate constants increased linearly with increasing persulfate dosage. The dominant reactive species in UV/persulfate is SO4•− and its formation rate depends on the persulfate dosage: the higher the persulfate dosage, the higher the SO4•− formation rate10 and the bromate formation rate. However, the regression line in Figure 1b does not go through the origin. This finding, together with the observed lag-phase at the persulfate dosage of 20 μM, suggests that the bromate formation involves multiple steps, which require sufficient SO4•− exposure to convert bromide to bromate in stepwise manner. Insufficient SO4•− exposure converts bromide to intermediates (HOBr/OBr−, as discussed later) but not bromate.

Figure 1. (a) Time-dependent bromate formation in the UV/ persulfate process and (b) the corresponding pseudo-zero-order rate constants (k) as a function of persulfate dosage. The inset in panel a enlarges the first 10-min results at persulfate dosages of 20, 50, and 100 μM. The solid line in panel b represents the best linear fit. Conditions: pH 7.0 (2 mM phosphate buffer), [bromide]0 = 20 μM, and UV intensity = 2.19 μEinstein L−1·s−1.

Figures 2a and 2b show the bromate formation in ultrapure water by UV/persulfate at initial Br− concentrations of 0.1−5.0 μM and 10−50 μM, respectively, the persulfate dosage of 200 μM, and pH 7. In general, bromate formation increased with increasing bromide concentration. It was also found that, at low initial bromide concentrations up to 10 μM, the formation gradually reached a plateau when all of the dosed bromide has been converted to bromate. However, increasing the initial bromide concentration from 10 μM to 50 μM did not lead to much of an increase in the bromate formation rate, suggesting that the SO4•− formation rate limited the bromate formation rate. The evolution of HOBr/OBr− and bromide during Br− oxidation by UV/persulfate was tracked to examine their involvement in the bromate formation. HOBr/OBr− has been well documented as a requisite intermediate in bromate formation during bromide oxidation by ozone and HO•.3,4 Figure 3a displays the resulting concentrations of HOBr/OBr−, 8978



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Figure 3. (a) Evolution of bromine species during the oxidation of Br− in the UV/persulfate process. Solid lines show simulated results. (b) Comparison of bromate formation between the oxidation of Br− and HOBr/OBr− by UV/persulfate and UV processes. Conditions: pH 7.0 (2 mM phosphate buffer), [Br−]0 = 20 μM, [HOBr/OBr−]0 = 20 μM, [persulfate]0 = 200 μM, and UV intensity = 2.19 μEinstein L−1·s−1.

Figure 2. Time-dependent bromate formation in the UV/persulfate process at initial bromide concentrations of (a) 0.1−5.0 μM and (b) 10−50 μM. Conditions: pH 7.0 (2 mM phosphate buffer), [persulfate]0 = 200 μM, and UV intensity = 2.19 μEinstein L−1·s−1.

bromide, and bromate and the Br mass balance during the conversion. The concentration of bromide decreased steadily, while the concentration of HOBr/OBr− increased rapidly to the maximum within 2−3 min and then decreased gradually. Bromate was then formed as the final product. The fairly constant Br mass balance throughout the course of the conversion indicates that the primary intermediate during the oxidation of Br− to bromate was indeed HOBr/OBr−, while the concentrations of other possible intermediate species such as bromite and oxybromine radicals accounted for less than 5% of the total Br. The formation of HOBr/OBr− was primarily attributable to SO4•−, because 20-min control tests showed that persulfate alone (at the same concentration) and UV alone (at the same intensity) did not produce detectable concentrations of HOBr/OBr− (not shown). Figure 3b compares the bromate formation from the oxidation of Br− and HOBr/OBr− at initial concentrations of 20 μM by UV/persulfate and UV processes. After 20 min of UV irradiation in the presence of persulfate, almost all HOBr/ OBr− and Br− were converted to bromate of the same quantity. Nevertheless, the initial bromate formation rate was much

higher during the oxidation of HOBr/OBr− than that during the oxidation of Br−. The UV photolysis (at the same UV intensity) of HOBr/OBr− also formed bromate but at a slower rate, indicating that UV photolysis of HOBr/OBr− contributed partially to the bromate formation in the UV/persulfate process. A supplementary test using heat to activate persulfate verified that bromate was also formed by the oxidation of HOBr/OBr− by SO4•− alone (not shown). In addition, the low bromate yield from the UV photolysis of HOBr/OBr− after 20 min indicats that some HOBr/OBr− was converted back to Br−, a common occurrence in this type of reactions.6 On the other hand, the direct UV photolysis of bromide in the absence of persulfate did not form any bromate under the irradiation condition tested. All these results indicate that the bromate formation involves stepwise conversion of bromide to bromate with the following reaction pathways as shown in Scheme 1. SO4•− is needed to convert bromide to HOBr/OBr− and, once HOBr/OBr− becomes available, both the SO4•− attack and UV photolysis 8979



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Scheme 1. Proposed Pathways of Bromate Formation from Bromide Oxidation by UV/Persulfate

will convert the HOBr/OBr− finally to bromate. Some HOBr/ OBr− and bromate can be converted back to bromide by UV photolysis, which can then be oxidized to HOBr/OBr− again by SO4•−. According to Scheme 1, a mathematical model consisting of the sequential oxidation of bromide to form bromine and then to bromate by SO4•− (eqs 7 and 8), the UV photolysis of HOBr/OBr− to form bromide and bromate (eqs 9 and 10), and the UV photolysis of bromate to form HOBr/OBr− and bromide (eqs 11 and 12)29 was constructed to simulate the evolution of Br−, HOBr/OBr−, and bromate during Br− oxidation by UV/persulfate, by assuming the reactions 7 and 8 could be described by pseudo-first-order kinetics. SO4•−

Br − ⎯⎯⎯⎯⎯→ HOBr/OBr − k1 = ? SO4•−

HOBr/OBr − ⎯⎯⎯⎯⎯→ BrO3− k 2 = ?

Figure 4. Formation of bromate at pH 4−10 in the UV/persulfate process. Conditions: pHs 4−8 (2-mM phosphate buffers), pHs 9 and 10 (5-mM borate buffers), [bromide]0 = 20 μM, [persulfate]0 = 200 μM, and UV intensity = 2.19 μEinstein L−1·s−1.

proposed and tested to explain the pH-dependent trend at pH 7−10: (1) pH may affect the radical exposure. As shown in Supporting Information Figure S4, the depletion rates of persulfate under the UV irradiation at pH 7 and 10 were quite similar, suggesting the formation rates of SO4•− were similar at pH 7−10. In addition, SO4•− transforms to HO• at alkaline pH. However, using the rate constant available in the literature (eq 2), the radical interconversion was calculated to be lower than 10% at pH ≤ 10 under the testing conditions. These results suggest that the change of radical exposure at pH 7−10 is not the major reason attributable to the dramatic decrease in bromate formation at alkaline pH. (2) pH may affect the first step of the bromate formation, that is, the oxidation of bromide to HOBr/OBr−. Tests were conducted to monitor and compare the evolution of HOBr/OBr− at pH 7 and 10. As shown in Supporting Information Figure S5, the peak concentration of accumulated HOBr/OBr− at pH 7 was over 10 times higher than that at pH 10. This pH dependency can be explained by the series reactions from Br• to form HOBr/ OBr−.21−23

(7) (8)

hv

HOBr/OBr − → BrO3− k 3 = 2.50 × 10−4 s−1

(9)

hv

HOBr/OBr− → Br− k4 = 1.12 × 10−3 s−1

(10)

hv

BrO3− → HOBr/OBr− k5 = 1.83 × 10−5 s−1

(11)

hv

BrO3− → Br − k6 = 1.00 × 10−4 s−1

(12)

Where k1 and k2 were fitted with the experimental data in Figure 3a, and k3−k6 were obtained from the experimental results of UV photolysis of HOBr/OBr− and UV photolysis of BrO3− (Supporting Information Figures S2 and S3, respectively) with the help of Matlab-Simulink 7.10 (MathWorks). As shown in Supporting Information Figure S2, the UV photolysis of HOBr/OBr− follows a first-order kinetics, with the firstorder rate constants of 2.50 × 10−4 (k3) and 1.12 × 10−3 s−1 (k4) for formation of BrO3− and Br−, respectively. The UV photolysis of bromate (Supporting Information Figure S3) also follows the first-order kinetics, with the first-order rate constants of 1.83 × 10−5 (k5) and 1.00 × 10−4 s−1 (k6) for formation of HOBr/OBr− and Br−, respectively. Using MatlabSimulink 7.10, the values of k1 and k2 are found to be 2.45 × 10−3 and 6.5 × 10−3 s−1, respectively. Figure 3a shows the simulated results, which agree reasonably well with the experimental data, suggesting the model covers all the major reactions and the assumptions above are reasonable. The much larger k2, compared to k3, obtained from the model indicates that the bromate formation by UV/persulfate is primarily driven by SO4•− and the bromate formation from UV photolysis of HOBr/OBr− can be neglected. Effect of pH. Figure 4 shows the bromate formation by UV/persulfate at different pH values, with an initial Br− concentration of 20 μM and a persulfate dosage of 200 μM. Rather unexpectedly, the bromate formation were similar for pH 4−7 but decreased over 30% and 90% with increasing pH from 7 to 8 and above 9, respectively. Three aspects were

Br• + Br − ↔ Br2− k+ = 1.0 × 1010 M−1 s−1, k − = 1.0 × 105 s−1

(13)

Br• + OH− ↔ BrOH− k+ = 1.3 × 1010 M−1 s−1, k − = 4 × 106 s−1 −

−

(14) −

−

8

−1 −1

BrOH + Br → Br2 + OH k = 2.0 × 10 M s

(15)

Br• can react fast with Br− to form Br2− (eq 13), which is an intermediate to form HOBr/OBr−. Br• also react fast with OH− to form BrOH− (eq 14). This is a reversible reaction and higher pH shifts the equilibrium toward to BrOH−. BrOH− also reacts with Br− to form Br2− (eq 15), however, at a rate 50-folds slower than that of the reaction between Br• and Br−. Thus, less Br• can be oxidized to HOBr/OBr− at higher pH. This aspect provides a plausible reason that explains the dramatic decrease in bromate formation at alkaline pH. (3) pH may affect the second step of the bromate formation, that is, the oxidation of HOBr/OBr− to bromate. Tests of the UV/persulfate treatment 8980



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starting with HOBr/OBr− solutions were conducted. As shown in Supporting Information Figure S6, the depletion rates of HOBr/OBr− and the initial bromate formation rates, up to 5−7 min, were similar at pH 7 and 10, suggesting the oxidation of HOBr/OBr− to bromate by UV/persulfate was not influenced by pH. Although it is known that UV photolysis of HOBr/ OBr− is pH dependent, due to the lower UV absorbance of OBr− at 254 nm, compared to that of HOBr (as shown in Supporting Information Figure S7), as discussed earlier, the contribution of UV photolyis of HOBr/OBr− to bromate is minor in the UV/persulfate process. Supporting Information Figure S6 also shows that, at pH 10, the bromate concentration remained quite constant after 10 min, due to the exhaustion of HOBr/OBr−. Because the produced bromide from the photolysis of HOBr/OBr− cannot be oxidized back at pH 10, as indicated in the aspect (2), there is no more HOBr/OBr− available for converting to bromate. At pH 7, nevertheless, the continuous oxidation of the recycled bromide to HOBr/OBr− is rapid (shown in Figure 3a), providing more HOBr/OBr− to form bromate by both SO4•− and UV photolysis. On the basis of the discussion above, we suggest that the pH dependency of the bromate formation is related to the slower transformation rate of Br• to HOBr/OBr− at alkaline pH. Bromate Formation in the Real Water Sample. Experiments were performed in the real water sample spiked with 20-μM bromide to examine the significance of bromate formation during UV/persulfate treatment under a more realistic condition. The bromate formation in the bromidespiked ultrapure water and in the bromide-spiked real water sample is compared in Figure 5a. The bromate formation in the real water sample was suppressed at the first 10 and 5 min at persulfate dosages of 200 and 500 μM, respectively. Also, after 20 min of reaction, the bromate concentrations in the real water sample were much lower than those in ultrapure water at the corresponding persulfate dosages. Nevertheless, 20% and 65% of the spiked Br− was converted to BrO3− after 20 min at persulfate dosages of 200 μM and 500 μM, respectively. As shown in Supporting Information Table S1, the concentrations of natural organic matter (NOM), alkalinity and chloride in the real water sample are 1.22 mg/L, 0.3 mM, and 286 μM, respectively. Chloride and alkalinity at these concentrations affected the bromate formation little in ultrapure water (Supporting Information Figures S8). The significant reduced and retarded bromate formation in the real water sample should be primarily attributable to the presence of NOM of 1.22 mg/L. NOM can absorb the UV light at 254 nm so it is an inner filter to reduce the efficiency of persulfate photolysis in producing SO4•−. However, the filtering effect is not significant (around 9.0% calculated from the comparison of the first-order degradation rate constants of dibromoacetonitrile by the direct UV photolysis in ultrapure water and the real water sample). On the other hand, NOM could react with and deplete the intermediates, such as Br• and HOBr/OBr−, to prevent their further oxidation to bromate.4,30 Little HOBr/OBr− was detected in the first few minutes in the real water sample (not shown), suggesting that NOM readily consumed HOBr/ OBr− and/or other intermediates such as Br• at the initial phase. NOM also reacted with and consumed SO4•−. The UVfluence first-order degradation rate of 20-μM pCBA (kSO4•− pCBA = 3.6 × 108 M−1 s−1), decreased 11% in 1-mg/L Suwannee river NOM spiked water compared to NOM-free

Figure 5. Comparison of bromate formation in the UV/persulfate process (a) in ultrapure water and in treated water sample without final disinfection and (b) as a function of degradation rate (%) of pCBA in ultrapure water spiked with different concentrations of NOM. Conditions: pH 7.2 ± 0.1, [bromide]0 = 20 μM, [persulfate]0 = 200 μM or 500 μM, [pCBA]0 = 20 μM and UV intensity = 2.19 μEinstein L−1·s−1.

water (not shown), which is attributed to NOM acting as a scavenger of SO4•−. These two types of consumption shall be the major reasons attributing to the large decrease in the bromate formation. The impact of NOM on bromate formation was further evaluated using pCBA as the probe micropollutant in ultrapure water spiked with different concentrations of Suwannee River NOM. Figure 5b displays bromate formation versus percentages of pCBA degradation at NOM concentrations of 0−1.0 mg/L. The monotonic increases in bromate formation with increasing pCBA degradation suggest that the extent of bromate formation and degradation of micropollutants are coupled processes. However, the pCBA degradation-based bromate formation decreased significantly with increasing NOM concentrations from 0 to 1.0 mg/L, suggesting the importance of NOM in consuming the intermediate bromine species, particularly at higher NOM concentrations. Nevertheless, the percentages of conversion from bromide to 8981



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Water Res. 2003, 37 (7), 1469−1487, DOI: 10.1016/S0043-1354(02) 00458-X. (5) Buffle, M. O.; Galli, S.; von Gunten, U. Enhanced bromate control during ozonation: The chlorine-ammonia process. Environ. Sci. Technol. 2004, 38 (19), 5187−5195, DOI: 10.1021/es0352146. (6) Zhang, X.; Zhang, T.; Ng, J.; Pan, J. H.; Sun, D. D. Transformation of bromine species in TiO2 photocatalytic system. Environ. Sci. Technol. 2009, 44 (1), 439−444, DOI: 10.1021/ es902592w. (7) Zhao, Q.; Shang, C.; Zhang, X. Halate formation from UV/ chlorine process in the presence of bromide. Presented at the AWWA Water Quality Technology Conference (WQTC), Seattle, Washington, USA, November, 2009. (8) 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 (page 1146−1149) DOI: 10.1063/1.555808. (9) 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 (•OH/•O−) in aqueoussolution. J. Phys. Chem. Ref. Data 1988, 17 (2), 513−886, DOI: 10.1063/1.555805. (10) Lau, T. K.; Chu, W.; Graham, N. J. D. The aqueous degradation of butylated hydroxyanisole by UV/S2O82‑: Study of reaction mechanisms via dimerization and mineralization. Environ. Sci. Technol. 2007, 41 (2), 613−619, DOI: 10.1021/es061395a. (11) Waldemer, R. H.; Tratnyek, P. G.; Johnson, R. L.; Nurmi, J. T. Oxidation of chlorinated ethenes by heat-activated persulfate: Kinetics and products. Environ. Sci. Technol. 2007, 41 (3), 1010−1015, DOI: 10.1021/es062237m. (12) Furman, O. S.; Teel, A. L.; Watts, R. J. Mechanism of base activation of persulfate. Environ. Sci. Technol. 2010, 44 (16), 6423− 6428, DOI: 10.1021/es1013714. (13) Anipsitakis, G. P.; Dionysiou, D. D. UV/transition metalmediated activation of common oxidants in water. Abstr. Pap. Am. Chem. Soc. 2004, 228, U613−U613. (14) 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, DOI: 10.1021/es2017363. (15) Chan, T. W.; Graham, N. J. D.; Chu, W. Degradation of iopromide by combined UV irradiation and peroxydisulfate. J. Hazard. Mater. 2007, 181 (1−3), 508−513, DOI: 10.1016/j.jhazmat.2010.05.043. (16) Hori, H.; Yamamoto, A.; Hayakawa, E.; Taniyasu, S.; Yamashita, N.; Kutsuna, S. Efficient decomposition of environmentally persistent perfluorocarboxylic acids by use of persulfate as a photochemical oxidant. Environ. Sci. Technol. 2005, 39 (7), 2383−2388, DOI: 10.1021/es0484754. (17) Anipsitakis, G. P.; Dionysiou, D. D. Transition metal/UV-based advanced oxidation technologies for water decontamination. Appl. Catal. B-Environ. 2004, 54 (3), 155−163, DOI: 10.1016/j.apcatb.2004.05.025. (18) Anipsitakis, G. P.; Tufano, T. P.; Dionysiou, D. D. Chemical and microbial decontamination of pool water using activated potassium peroxymonosulfate. Water Res. 2008, 42 (12), 2899−2910, DOI: 10.1016/j.watres.2008.03.002. (19) 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, DOI: 10.1021/es035121o. (20) Liang, C. J.; Wang, Z. S.; Mohanty, N. Influences of carbonate and chloride ions on persulfate oxidation of trichloroethylene at 20 °C. Sci. Total Environ. 2006, 370 (2−3), 271−277, DOI: 10.1016/ j.scitotenv.2006.08.028. (21) Zehavi, D.; Rabani, J. Oxidation of aqueous bromide ions by hydroxyl radicals. A pulse radiolytic investigation. J. Phys. Chem. 1972, 76 (3), 312−319, DOI: 10.1021/j100647a006. (22) Kläning, U. K.; Wolff, T. Laser flash photolysis of HClO, ClO−, HBrO and BrO− in aqueous solution, Reactions of Cl- and Br- atoms. Ber. Bunsen-Ges. Phys. Chem. 1985, 89, 243−245.

bromate at the same target of micropollutant degradation are expected to be significant if water contains lower quantities of NOM. Engineering Implications. Bromide is widely present in water worldwide. In the U.S., its concentration ranges from 0 to 2.3 mg/L, with an average of 100 μg/L.31 At UV doses and persulfate dosages relevant to those used for micropollutant degradation,10,14 the large percentage of conversion from bromide to bromate by the UV/persulfate process obtained from this study suggests a potential problem of excessive production of bromate when SO4•−-based AOPs are used to treat bromide-containing drinking water, groundwater, and wastewater. NOM in real water can reduce the bromate formation by reacting with the intermediates (i.e., HOBr/OBr− and Br•) and SO4•−. The former shall lead to the formation of bromine-containing byproducts of potential health concern. The latter reduces the quantity of SO4•− available for micropollutant destruction. Therefore, based on the same SO4•− exposure that is required for micropollutant destruction, the bromide to bromate conversion by SO4•− is expected to be still significant, unless there also exists abundant NOM to consume intermediate bromine species. For controlling the bromate formation, as demonstrated in Figure 4, elevating pH to above 8 may be an option for degradation of micropollutants that is not negatively impacted by increasing pH. Examples of the micropollutants are butylated hydroxyanisole (an endocrine disruptor chemical) and phenol.10,32

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ASSOCIATED CONTENT

S Supporting Information *

Table S1 and Figures S1−S8. This information is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Phone: (852)2358 7885. Fax: (852)2358 1534. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported in part by Hong Kong’s Research Grants Council under grant number 618312. We gratefully acknowledge Prof. Urs von Gunten at Eawag, Swiss Federal Institute of Aquatic Science and Technology for his suggestion on the modeling work, and Tat-Ming Sze, Yan-Na Zhou and Yun Fu at The Hong Kong University of Science and Technology for their help on the experimental work. We also appreciate the valuable comments from the anonymous reviewers.

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REFERENCES

(1) WHO. Guidelines for Drinking-Water Quality, 4th ed.; World Health Organization: Geneva, Switzerland, 2011. (2) U.S. EPA. National Primary Drinking Water Regulations: Stage 2 Disinfectants and Disinfection Byproducts Rule; United States Environmental Protection Agency: Washington, DC, 2006. (3) von Gunten, U.; Oliveras, Y. Advanced oxidation of bromidecontaining waters: Bromate formation mechanisms. Environ. Sci. Technol. 1998, 32 (1), 63−70, DOI: 10.1021/es970477j. (4) von Gunten, U. Ozonation of drinking water: Part II. Disinfection and by-product formation in presence of bromide, iodide or chlorine. 8982



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(23) Mamou, A.; Rabani, J.; Behar, D. Oxidation of aqueous Br− by hydroxyl radicals, studies by pulse radiolysis. J. Phys. Chem. 1977, 81 (15), 1447−1448, DOI: 10.1021/j100530a003. (24) APHA-AWWA-WEF. Standard Methods for the Examination of Water and Wastewater, 20th ed.; American Public Health Association/ American Water Works Association/Water Environment Federation: Washington DC, USA, 1998. (25) Kumar, K.; Margerum, D. W. Kinetics and mechanism of general-acid-assisted oxidation of bromide by hypochlorite and hypochlorous acid. Inorg. Chem. 1987, 26 (16), 2706−2711, DOI: 10.1021/ic00263a030. (26) Huang, K.-C.; Couttenye, R. A.; Hoag, G. E. Kinetics of heatassisted persulfate oxidation of methyl tert-butyl ether (MTBE). Chemosphere 2002, 49 (4), 413−420, DOI: 10.1016/S0045-6535(02) 00330-2. (27) Rahn, R. O.; Stefan, M. I.; Bolton, J. R.; Goren, E.; Shaw, P. S.; Lykke, K. R. Quantum yield of the iodide-iodate chemical actinometer: Dependence on wavelength and concentration. Photochem. Photobiol. 2003, 78 (2), 146−152, DOI: 10.1562/0031-8655. (28) Baxendale, J. H.; Wilson, J. A. The photolysis of hydrogen peroxide at high light intensities. Trans. Faraday Soc. 1957, 53, 344− 356, DOI: 10.1039/TF9575300344. (29) Phillip, N. H.; Gürten, E.; Diyamandoğlu, V. Transformation of bromine species during decomposition of bromate under UV light from low pressure mercury vapor lamps. Ozone: Sci. Eng. 2006, 28 (4), 217−228, DOI: 10.1080/01919510600705384. (30) Westerhoff, P.; Song, R. G.; Amy, G.; Minear, R. NOM’s role in bromine and bromate formation during ozonation. J. Am. Water Work Assoc. 1998, 90 (2), 82−94. (31) Amy, G. L.; Siddiqui, M.; Zhai, W.; Debroux, J.; Odem, W. Nation-wide survey of bromide ion concentrations in drinking water sources. Proc., 1993 AWWA Annu. Conf. 1993, 1−19. (32) Lin, Y.-T.; Liang, C.; Chen, J.-H. Feasibility study of ultraviolet activated persulfate oxidation of phenol. Chemosphere 2011, 82 (8), 1168−1172, DOI: 10.1016/j.chemosphere.2010.12.027.

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