Article pubs.acs.org/est
The Roles of Reactive Species in Micropollutant Degradation in the UV/Free Chlorine System Jingyun Fang,*,†,‡,§,∥ Yun Fu,‡ and Chii Shang*,‡ †
School of Environmental Science and Engineering, Sun Yat-sen University, Guangzhou 510275, China Department of Civil and Environmental Engineering, The Hong Kong University of Science and Technology, Clear Water Bay, Kowloon, Hong Kong § Guangdong Provincial Key Laboratory of Environmental Pollution Control and Remediation Technology, Guangzhou 510275, China ∥ SYSU-HKUST Research Center for Innovative Environmental Technology (SHRCIET) School of Environmental Science and Engineering, Sun Yat-sen University, Guangzhou 510275, China ‡
S Supporting Information *
ABSTRACT: The UV/free chlorine process forms reactive species such as hydroxyl radicals (HO•), chlorine atoms (Cl•), Cl2•−, and O•−. The specific roles of these reactive species in aqueous micropollutant degradation in the UV/chlorine process under different conditions were investigated using a steady-state kinetic model. Benzoic acid (BA) was chosen as the model micropollutant. The steady-state kinetic model developed fitted the experimental data well. The results showed that HO• and Cl• contributed substantially to BA degradation, while the roles of the other reactive species such as Cl2•− and O•− were negligible. The overall degradation rate of BA decreased as the pH increased from 6 to 9. In particular, the relative contributions of HO• and Cl• to the degradation changed from 34.7% and 65.3% respectively at pH 6 to 37.9% and 62% respectively at pH 9 under the conditions evaluated. Their relative contributions also changed slightly with variations in chlorine dosage, BA concentration and chloride concentration. The scavenging effect of natural organic matter (NOM) on Cl• was relatively small compared to that on HO•, while bicarbonate preferentially reduced the contribution of Cl•. This study is the first to demonstrate the contributions of different reactive species to the micropollutant degradation in the UV/chlorine system under environmentally relevant conditions.
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(LP) mercury UV lampare 59 M−1cm−1, 66 M−1cm−1, ∼19 M−1cm−1, 47.5 M−1cm−1 and 12.5 M−1cm−1 respectively.11−13 The quantum yields of chlorine photolysis are at least ∼1.0 mol Es1− under UV irradiation from an LP or a medium pressure (MP) Hg lamp.5,7−9 The quantum yields of H2O2 and persulfate photolysis are 0.5 and 0.7 mol Es−1 under LP UV irradiation.12,14 The UV/chlorine process has been reported to be more effective than the UV/H2O2 process against some micropollutants such as trichloroethylene and certain pharmaceuticals under mildly acidic conditions.4,6,15 Furthermore, it can provide multiple disinfection barriers in drinking water treatment and residual protection in water distribution systems,16 without the need to quench residual oxidants such as H2O2 as in the UV/H2O2 process. In the UV/chlorine process, reactive species such as HO• and chlorine atoms (Cl•) are formed due to the UV photolysis of HOCl and OCl−:5
INTRODUCTION Chlorine or UV alone is used for disinfection and/or oxidation worldwide. When chlorine is used for preoxidation prior to UV disinfection, constituents in water are exposed to UV and chlorine simultaneously. UV disinfection is commonly used to treat chlorinated pool waters and to purify drinking water containing residual chlorine at household levels, both scenarios provide a UV/chlorine coexposure environment.1,2 The UV/chlorine coexposure derives from the UV photolysis of free chlorine and/or chloramines.2 The UV/chlorine process, particularly the UV photolysis of free chlorine, is being considered as an alternative advanced oxidation process (AOP) to the UV/hydrogen peroxide (H2O2) process for the destruction of emerging contaminants (e.g., desethylatrazine, sulfamethoxazole, carbamazepine, diclofenac, benzotriazole, tolyltriazole, iopamidole and 17αethinylestradiol), and other micropollutants such as nitrobenzene, p-chlorobenzoic acids, cyclohexanoic acid and trichloroethylene,3−10 for chlorine’s comparable UV absorbance and quantum yield to H2O2 and easy retrofitting of the process into the existing treatment train by adding a compact UV system after chlorine dosing. The molar absorption coefficients of HOCl, OCl−, H2O2, persulfate and peroxymonsulfate at 254 nmthe emission peak of a monochromatic low pressure © XXXX American Chemical Society
HOCl/OCl− → HO• /O•− + Cl•
(a)
Received: August 13, 2013 Revised: January 7, 2014 Accepted: January 8, 2014
A
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The Cl• so formed will react with the chloride ions that can always be found in a HOCl/OCl− solution to form Cl2•−:17
organic matter (NOM) and alkalinity may further complicate the system by affecting the formation of radicals and their interconversion and consumption. Their impacts are also considered in the model. The objective of this study was therefore to improve the understanding of the reactive species responsible for micropollutant degradation in the UV/chlorine process using a kinetic modeling approach. Simulated water matrices were employed to uncover the reactive species responsible for degradation and reveal the effects of changes in water constituents such as pH, chlorine dosage, micropollutant concentration, chloride concentration, NOM concentration, and alkalinity. Benzoic acid (BA) and nitrobenzene (NB) were selected as the model compounds for the reactive species, as they are reactive to HO• and/or Cl• but inert to either chlorination or UV photolysis. They also represent chemical pollutants that are frequently found in the environment.
Cl• + Cl− ↔ Cl 2•− k+ = 6.5 × 109M−1s−1, k − = 1.1 × 105s−1 (b)
where, k+ and k− represent the rate constants of forward and backward reactions, respectively. Thus, Cl2•− is also present and in equilibrium with Cl• in the UV/chlorine system. Both Cl • and Cl 2•− are strong oxidants with the standard reduction potentials of 2.4 and 2.0 V, respectively,18,19 which are comparable to that of HO•. HO•, Cl•, and Cl2•− react with organics by one-electron oxidation, H-abstraction, and additions to unsaturated C−C bonds.19−21 Cl• is a selective oxidant and its reactivity can be higher than that of HO• in some instances.22 For example, Cl• reacts more quickly with acetic acid, benzoic acid and phenol than does HO•.22 Cl2•− is, however, generally much less reactive than HO• or Cl•.22 HO• has been deemed as the major or only reactive species in the degradation of UV- and chlorine-persistent micropollutants, while the contributions of Cl• and Cl2•− are often neglected or underestimated in the UV/chlorine process.5,6,23 Cl• was found to play a negligible role in the oxidation of nitrobenzene and 1-chlorobutane, due to its low reactivity with the two organic compounds.5,23 It does, however, react rapidly with some pollutants such as phenol and ethylene.22 Therefore, its contribution and that of Cl2•− in the UV/chlorine process should be considered in line with their reactivity with micropollutants. The conversion of radicals further complicates the UV/ chlorine process. Cl• reacts with OH− to form ClOH•− with a rate constant of 1.8 × 1010 M−1s−1 (eq 16 in Table 1), and HO• reacts with Cl− to form ClOH•− with a rate constant of 4.3 × 109 M−1s−1 (eq 6 in Table 1). ClOH•− then decays and produces HO• and Cl• with rate constants of 6.1 × 109 s−1 (eq 7 in Table 1) and 23 s−1, respectively.17 The pH is known to affect the conversion.17 It influences the UV/chlorine process by affecting the HOCl/OCl− speciation, which determines the absorbance and the quantum yields of chlorine degradation and thus the radical production.7,9,24 HOCl and OCl− also exert different radical scavenging effects on Cl• and HO•.17 In addition, the production of HO• decreases at pH equal or close to the pKa of HO•/O•− (11.9).9,20 In this paper, we chose a mathematical modeling approach to improve our understanding of the roles of various reactive species in the UV/chlorine process for process development and optimization. The pseudosteady state approximation for the kinetic description of free radical species has been successfully used to predict the degradation of micropollutants in AOPs such as the UV/H2O2, UV/ozone and UV/ peroxymonosulfate processes.12,14,25 It should also be applicable to the establishment of a kinetic model for the UV/chlorine system. Only one study has reportedly dealt with the kinetic modeling of the UV/chlorine process.6 That study built a mathematical model of trichloroethylene degradation at various pHs using Matlab, based on the assumption that only HO• contributes to the degradation while the contributions of other reactive species such as Cl• to the degradation are negligible. To elucidate the roles of Cl• and Cl2•−, as well as that of HO•, in the UV/chlorine system, a more sophisticated model considering all the reactions involved is required. The variations in chloride concentration and pH and the presence of natural
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EXPERIMENTAL SECTION Chemicals. Solutions were prepared with reagent-grade chemicals and purified water (18.2 MΩ/cm) from a NANOpure system (Barnstead). Solutions were stored at 4 °C in the dark and were brought back to room temperature before use. A NaOCl stock solution (1350 mg/L active chlorine as Cl2) was prepared by bubbling chlorine gas into a 0.02 M NaOH solution. The chlorine gas was produced by reacting KMnO4 with NaCl catalyzed by concentrated H2SO4. The stock solution was periodically standardized by the DPD/FAS titration method.26 BA, NB, sodium sulfite and phosphoric acid of HPLC grade were purchased from Sigma-Aldrich. Methanol of HPLC grade was purchased from Merck KGaA. Stock solutions (10 mM) of target pollutants were prepared by adding pure compounds to the purified water and stirring overnight. Phosphate buffers were used to adjust pH and the pH values during reaction were monitored and maintained at the predetermined pH ± 0.1. Sodium sulfite was prepared weekly. A NOM stock solution was prepared by dissolving an aliquot of Suwannee River RO NOM isolate (Cat. No. 1R101N, International Humic Substance Society) in the purified water and then the solution was filtered through a 0.45 μM membrane. The stock solution (filtrate) was quantified and standardized by dissolved organic carbon (DOC) analysis with a total organic carbon (TOC) analyzer (TOC-VCPH, Shimadzu). UV Exposure. UV exposure was carried out using a benchscale UV irradiator, consisting of four LP Hg UV lamps (254 nm, G15T8, 10 W, Sankyo Denki) housed in a shuttered box, with a vertical tube extending from the bottom.21,27 A 250 mL sample with a water depth of 5 cm was placed in a glass reactor with an inner diameter of 8 cm which was placed directly beneath the collimated tube for UV treatment, which schematic diagram is shown in Supporting Information (SI) Figure S1. A quartz cap was used to reduce the volatilization of chlorine and the sample was stirred rapidly at ambient temperature (22 ± 2 °C). The distance between the solution and the lamps was about 30 cm. As the UV light in the reactor was not a standard collimated beam,28 the fluence rate distribution would depend on reactor geometry and physics (e.g., dispersion, reflection, refraction, etc.). The effective path length (L) was selected to characterize the photoreactor and was measured and calculated to be 4.9 ± 0.1 cm from the kinetics of the photolysis of dilute H2O2 (see Text S1 and Figure S2 in the SI for details).14 The photon flux (I0)29 entering the solution was determined to be B
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Table 1. Principal Reactions in the UV/Chlorine System no.
rate constant
ref.
1
HOCl/OCl− → HO• /O•− + Cl•
I r HO• = ΦHClOf HClO 0 L(1 − 10−A ) V I r O•‐ = ΦOCl−fClO− 0 L(1 − 10−A ) V I r Cl• = (ΦHClOfHClO + ΦOCl−fClO− ) 0 L(1 − 10−A ) V A = (εHOClC HOCl + εOCl−COCl− + εBA C BA + ...)L fHClO = εHOClC HOClL /A fClO− = εOCl−COCl−L /A
14,12
2
HOCl ⇔ OCl− + H+
pKa1 = 7.5
31
3
HO• ⇔ O•− + H+
pKa2 = 11.9
20
4 5 6 7 8 9 10 11 12 13
reaction
•
−
k1 = 6.5 × 109 M−1s−1
Cl •− 2
Cl + Cl →
k2 = 1.1 × 10 s
• − Cl •− 2 → Cl + Cl •
−
HO + Cl → ClOH ClOH
•−
ClOH
•−
ClOH
•−
ClOH
•−
−
→ HO + Cl •
−
+
•
−
Cl •− 2
+ Cl →
+ OH
•
HO + OCl → ClO + OH
17 −1 −1
k7 = 1.0 × 10 M s
−1 −1
−
34 35
−1 −1
k10= 5.5 × 10 M s
HO + HO → H 2O2
33
k9 = 8.8 × 109 M−1s−1 9
•
17
k8 = 2.0 × 10 M s 9
HO + HOCl → ClO + H 2O •
17
−1
5
−
•
−
21
k4 = 6.1 × 109 s−1 k6 = 2.1 × 1010 M−1s−1
+ H → Cl + H 2O
•
k3 = 4.3 × 10 M s k5 = 23 s
→ Cl + HO
17 32
−1 −1̀
9
•−
•
•
−1
5
−1 −1
20
14
Cl + HOCl → H + Cl + ClO
k11 = 3.0 × 10 M s
36
15
Cl• + OCl− → Cl− + ClO•
k12 = 8.2 × 109 M−1s−1
32
k13 = 1.8 × 1010 M−1s−1
17
16 17 18
•
+
•
−
Cl + OH → ClOH
Cl 2
•−
−
9
•
•‐
−
+ OH → ClOH
•−
−1 −1
k14 = 4.5 × 10 M s 7
−
+ Cl
−1 −1
33
k15 = 1.8 × 10 M s
20
k16 = 1.3 × 1010 M−1s−1
20
k17 = 5.9 × 109 M−1s−1
20
Cl + C6H5COO → product
k18 = 1.8 × 1010 M−1s−1
37
− Cl •− 2 + C6H5COO → product
k19 = 2.0 × 10 M s
•−
O
•
•
+ H 2O → HO + OH −
•−
6
−
19
HO + OH → O
20
HO• + C6H5COO− → product
+ H 2O
In the Presence of Benzoic Acid 21 22 23 24
•
•−
O
−
+ C6H5COO → product
C6H5COOH ⇔ C6H5COO + H In the Presence of HCO3
−1 −1
+
38
k20 = 4.0 × 10 M s
20
pKa3 = 4.2
39
7
−
−
−1 −1
6
−
HO• + HCO−3 → H 2O + CO•− 3
k21 = 8.5 × 106 M−1s−1
Cl• + HCO−3 → H 2O + CO•− 3
k22 = 2.2 × 10 M s
27
− − + •− Cl •− 2 + HCO3 → 2Cl + H + CO3
k23 = 8.0 × 10 M s
34
28
H 2CO*3 ⇔ H+ + HCO−3 ⇔ 2H+ + CO32 −
pKa,1 = 6.3, pKa,2 = 10.3
39
29
HO• + NOM→
k24 = 2.5 × 104 (mg/L)−1s−1
40
30
Cl• + NOM→
k25 = 1.3 × 104 (mg/L)−1s−1
a
k26 = 3.9 × 109 M−1s−1
20
k27 is negligible
5
k28 = 6.0 × 108 M−1s−1
20
k29 = 3.0 × 108 M−1s−1
41
25 26
−1 −1
8
−1 −1
7
20 22
In the Presence of NOM
In the Presence of NB 31 32
HO• + NB→ •
Cl + NB→ In the Presence of t-BuOH
33 34 35 a
HO• + (CH3)3 COH→ •
Cl + (CH3)3 COH→ Cl •− 2 + (CH3)3 COH→
k30 = 700 M
−1
s
−1
38
Estimated in this study with details shown in SI Text S3. C
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0.063 ± 0.003 μE s−1 in a separate experiment using iodide/ iodate chemical actinometry.30 The corresponding average UV fluence rate (EP0) was about 0.58 mW cm−2 (see SI Text S1 for details). Experimental Procedures. In order to simulate UV/chlorine coexposure, a 250 mL testing solution spiked with 5 μM BA and buffered at a certain pH (2 mM phosphate buffer), with or without NOM or bicarbonate, was dosed with the NaOCl stock solution to give an initial chlorine concentration of 10−100 μM and was simultaneously exposed to UV light. Samples were collected at different time intervals, quenched with freshly prepared sodium sulfite at a sulfite-to-chlorine molar ratio of around 1.5:1, and analyzed for remaining BA concentrations within one day. In parallel tests, exposure to either UV light or chlorination was conducted using a similar approach as the controls. Another test was conducted in the same manner with the coexistence of 5 μM BA and 5 μM NB and buffered at pH 6, 7.5, and 9. All tests were conducted at least twice. The error bars in all data plots represent the maximum and minimum of the experimental data of the duplicated test results. Analytical Method. Concentrations of benzoic acid were determined with a reverse phase liquid chromatograph system (VP series, Shimadzu) equipped with a Waters symmetry C18 column (4.6 mm ×150 mm, 5 μm particle size) and a UV detector set at 227 nm. An eluent of water (pH 2, adjusted using phosphoric acid) and methanol (55:45, v/v %) was used to separate BA and its products at a flow rate of 1.0 mL/min. The pH value was measured with a pH meter (420-A, ORION). Chlorine concentrations were determined by the DPD/FAS titration method.26 The quantum yield of free chlorine decomposition was determined by dividing moles of decomposed free chlorine by einsteins absorbed at 254 nm (see SI Text S2 for details).5,9 Model Equations. The model was established based on the hypothesis that the degradation of BA in the UV/chlorine process mainly depends on HO•, Cl•, O•−, and Cl2•− from the photolysis of chlorine, while the contribution of other radicals such as ClOH•− and ClO• in the system can be neglected due to their low reactivity with BA.22 The degradation rates of BA by UV alone and by chlorine alone are 8.3 × 10−6 s−1 and 6.7 × 10−6 M−1s−1 (obtained in this work), respectively, which contribute to only 2.6% and 1.9% of the degradation within the first 20 min of the UV/chlorine coexposure and thus can be neglected. Table 1 summarizes the possible reactions in the UV/chlorine system along with their rate constants obtained from the literature. Based on the steady-state assumption, the kinetics of BA degradation in the UV/chlorine system can be modeled as follows: −
OCl− and the conversion of HO• to O•−. The overall BA degradation can thus be attributed to a number of reactive species. The kinetic expressions of HO•, Cl•, Cl2•−, O•− and ClOH•− in the UV/chlorine system are shown in eqs 2−6. rHO•, rCl•, and rO•− in eqs 2, 3 and 5 are the formation rates of HO•, Cl• and O•− from the photolysis of chlorine as shown in eq 1 in Table 1. The derivation of the eqs 2−6 is shown in SI Text S3. d[HO•] = r HO• + k4[ClOH•−] + k15[O•−][H 2O] dt − k17[HO•][BA] − k 8[HO•][HOCl] − k 9[HO•][OCl−] − k3[HO•][Cl−] − k16[HO•][OH−] −
∑ kSi,HO [HO•][Si ] •
i
(2)
d[Cl•] = r Cl• + k 2[Cl 2•−] + k5[ClOH•−] dt + k6[ClOH•−][H+] − k18[Cl•][BA] − k1[Cl•][Cl−] − k11[Cl•][HOCl] − k12[Cl•][OCl−] − k13[Cl•][OH−] −
∑ k Si,Cl [Cl•][Si] •
(3)
i
d[Cl 2•−] = k1[Cl•][Cl−] + k 7[Cl−][ClOH•−] dt − k 2[Cl 2•−] − k19[Cl 2•−][BA] − k14[Cl 2•−][OH−] −
∑ kSi,Cl
[Cl 2•−][Si ]
•− 2
i
(4)
d[BA] = (k17[HO•]ss + k18[Cl•]ss + k19[Cl 2•−]ss + k 20[O•−]ss )[BA] dt s [BA] = k 0,BA
(1)
•
where [HO ]ss is defined as the steady-state concentration of HO•, [Cl•]ss the steady-state concentration of Cl•, [Cl2•−]ss the steady-state concentration of Cl2•−, [O•−]ss the steady-state concentration of O•− and ks0,BA the overall pseudo-first-order rate constant of BA degradation. It is assumed that HO• is formed from the photolysis of HOCl and the conversion of ClOH•− and O•− to HO•; Cl• from the photolysis of HOCl and OCl− and the conversion of ClOH•− and Cl2•− to Cl•; Cl2•− from the reaction between Cl• and chloride and the conversion of ClOH•− to Cl2•−; and O•− from the photolysis of
Figure 1. Time-dependent BA degradation in ultrapure water by chlorination alone, UV irradiation alone and the UV/chlorine combined process in the presence and absence of 0.5-mM t-butanol. Conditions: pH 6.0, [BA]0 = 5 μM, initial free chlorine dose = 70 μM. D
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d[O•−] = r O•− + k16[HO•][OH−] − k15[O•−][H 2O] dt − k 20[O•−][BA] −
∑ kSi,O
[O•−][Si ]
where [Si] is the specific concentration of any scavenger in water (i.e., NOM, alkalinity or other micropollutants) and kSi,HO•, kSi,Cl•, kSi,Cl2•−, kSi,O•−, and kSi,ClOH•− are its corresponding second-order rate constants when it is reacting with specific radicals. The rate constants of BA used in this study were those of benzoate in reaction with the radicals, and the small fraction of benzoic acid at pH ≥ 6 should not affect the kinetics. Based on the steady-state assumption, the net formation rates of HO•, Cl•, O•−, Cl2•−, and ClOH•− are all zero. The modeling steady-state concentrations of HO•, Cl•, Cl2•−, O•−, and ClOH•−, which are dependent on the operating conditions, can be derived as follows (see SI Text S3 for derivation details):
•−
(5)
i
d[ClOH•−] = k 3[HO•][Cl−] + k13[Cl•][OH−] dt + k14[Cl 2•−][OH−] − k4[ClOH•−] − k6[ClOH•−][H+] − k 7[ClOH•−][Cl−] −
∑ kSi,ClOH
[ClOH•−][Si ]
•−
(6)
i
ΦHOCl εHOCl
•
[HO ]ss =
10 pKa − pH[chlorine] 10
pKa − pH
k 3[Cl−] + k17[BA] + k 8
+1
(10 pK a − pH + 1)
[Cl ]ss =
−
[Cl 2•−]ss =
k11
[chlorine]10 pK a − pH 10 pK a − pH + 1
•
−
+ k12
(1 − 10−A ) VA
[chlorine]10 pKa − pH 10 pK a − pH + 1
(ΦHOCl εHOCl10 pK a − pH + ΦOCl−εOCl−)[chlorine] •
I 0L
I 0L
(1 − 10−A ) VA
[chlorine] 10 pKa − pH + 1
[O ]ss =
10 pK a − pH + 1
I 0L
(1 − 10−A ) VA
+ k16[OH−][HO•]ss
k 20[BA] + k15[H 2O] + ∑i k Si,O•−[Si ] (10)
[ClOH•−]ss =
[chlorine] 10 pK a − pH + 1
+ k16[OH−] + ∑i k Si,HO•[Si ]
(7)
+ k 2[Cl 2•−]ss + k5[ClOH•−]ss + k6[H+][ClOH•−]ss
(8)
the degradation in the UV/chlorine process are minor and can be ignored. The first-order rate constant of BA degradation by UV/ chlorine coexposure in the presence of 0.5 mM t-butanol decreased by about 50% (from 2.0 × 10−3 s−1 to 1.0 × 10−3 s−1). The 0.5-mM t-butanol can compete with BA for up to 91% of HO•. This result indicates the contribution of other reactive species besides HO• to the BA degradation, such as Cl• and Cl2•− present in the system. To elucidate the contribution of other radicals, both NB and BA were used as the probe compounds. The former is only reactive with HO•,5 and the latter can react with HO• and Cl• and Cl2•− in the system. The result shows that Cl• and Cl2•− contributed more to BA degradation than did HO• (see SI Text S4 and Figure S3 for details). Effect of Chlorine Dosage. The BA degradation rate increased with increasing chlorine dosage (SI Figure S4). Figure 2 shows the experimental and predicted values of overall pseudo-first-order rate constants of BA degradation at different chlorine dosages (0−100 μM) and pH 6. The experimental values increased from 8.8 × 10−4 s−1 to 2.1 × 10−3 s−1 as chlorine dosage increased from 10 μM to 100 μM. Rapid increases in the rate constants were observed at low chlorine dosages (
