Comparison of Photodegradation Performance of 1,1,1

May 7, 2012 - The chemical oxidation of 1,1,1-trichloroethane (TCA), a widely detected groundwater pollutant, by UV/H2O2 and UV/S2O82– processes was...
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Comparison of Photodegradation Performance of 1,1,1Trichloroethane in Aqueous Solution with the Addition of H2O2 or S2O82− Oxidants Xiaogang Gu,† Shuguang Lu,*,† Zhaofu Qiu,† Qian Sui,† Zhouwei Miao,† Kuangfei Lin,† Yongdi Liu,† and Qishi Luo‡ †

State Environmental Protection Key Laboratory of Environmental Risk Assessment and Control on Chemical Process, Shanghai Key Laboratory of Functional Materials Chemistry, East China University of Science and Technology, Shanghai 200237, China ‡ Shanghai Academy of Environmental Science, Shanghai 200233, China ABSTRACT: The chemical oxidation of 1,1,1-trichloroethane (TCA), a widely detected groundwater pollutant, by UV/H2O2 and UV/S2O82− processes was investigated. The effects of various factors were evaluated, including peroxide/TCA molar ratio, solution pH, Cl− and HCO3− anions, and humic acid (HA). The results showed that TCA oxidation fit to a pseudo-first-order kinetic model. The optimum H2O2/TCA molar ratio was 5:1, with TCA removal of 54.2% in 60 min. In the UV/S2O82− process, higher molar ratios (from 1/1 to 10/1) resulted in higher TCA oxidation rates, and TCA could be completely removed after 60 min with a S2O82−/TCA molar ratio of 3/1. In addition, acidic conditions were favorable for TCA removal in the UV/S2O82− process, while maximum TCA removal was observed at pH 6 in the UV/H2O2 process. Both Cl− and HCO3− anions adversely affected TCA oxidation performance, and higher concentration of HA resulted in a lag phase for TCA oxidation in both processes. Several reaction intermediates, including 1,1,1,2-tetrachloroethane, carbon tetrachloride, chloroform, tetrachloroethylene, 1,1-dichloroethylene, and tri- and dichloroacetic acids, were first identified during TCA oxidation by S2O82− chemistry, while only monochloroacetic acid was detected in the UV/H2O2 process. The results indicated that the UV/S2O82− process was much more effective than the UV/H2O2 process, but the latter was more environmentally friendly because fewer toxic intermediates were produced.

1. INTRODUCTION 1,1,1-trichloroethane (TCA), a chlorinated solvent, has been used widely for several decades as one of the major chemical solvents in adhesives, aerosols, textile processing, extraction solvents, industrial solvent blends, and for metal degreasing. This substance has been reported to be a contaminant in at least 823 of the 1662 National Priorities List sites identified by the U.S. Environmental Protection Agency (USEPA).1 TCA in groundwater is susceptible to abiotic and biotic transformations; the biotransformation of TCA can produce even more toxic intermediates, such as 1,l-dichloroethane, 1,1dichloroethylene (1,1-DCE), and vinyl chloride through the action of microorganisms, as well as chloroethane, which was observed to be the terminal product of TCA dechlorination through biotransformation.2 As TCA has the potential to cause liver, nervous system, and circulatory system problems from long-term exposure, the maximum contaminant level of TCA in drinking water has been set at 0.2 mg L−1.3 Ex-situ treatment with a UV peroxide system is an efficient and effective technology for contaminated groundwater remediation.4 The UV/H2O2 process is also a technique used in the treatment of chlorinated solvent contamination in groundwater based on the generation of the highly reactive hydroxyl radical with a standard redox potential of 2.70 V. Trichloroethene (TCE) is proved to be completely degraded by UV/H2O2 processes.5 However, chlorinated alkanes are more resistant to UV/H2O2 processes than alkenes. Beltran et al.6 demonstrated that the rate constant of TCE with •OH was © 2012 American Chemical Society

approximate 65 times higher than that of TCA, and approximately 80% of TCA removal (including volatilization) could be achieved when TCA and H2O2 were applied at concentrations of 100 μg L−1 and 10−2 M, respectively, but the products and mechanism during TCA oxidation were not mentioned. However, the potential of alternative advanced oxidation processes (AOPs) to improve the removal performance is always of interest, and persulfate (S2O82−) shows promise as a candidate for such a purpose. Being one of the strongest oxidizing agents known, persulfate (S2O82−) has emerged to have great potential for TCA oxidation through the activation by means of heat, ultraviolet irradiation, chelated or unchelated transition metals, hydrogen peroxide, and alkaline pH to generate sulfate radicals (SO4−•) and other reactive species.7,8 Currently, synthetic iron and manganese oxides, as well as some specific organic compounds, are confirmed to be able to activate persulfate to generate oxidants and reductants.9,10 The SO4−•, a stronger oxidant than S2O82−, has a higher redox potential (E0 ≈ 2.6 V) to decompose many organic contaminants, including chlorinated aliphatics and BTEX compounds (benzene, toluene, ethyl benzene, and xylenes).11,12 TCA has been reported to be readily oxidized by thermally activated S2O82− oxidation from 30−60 °C.11,13,14 Received: Revised: Accepted: Published: 7196

November 29, 2011 May 4, 2012 May 7, 2012 May 7, 2012 dx.doi.org/10.1021/ie202769d | Ind. Eng. Chem. Res. 2012, 51, 7196−7204

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Block et al.15 observed significant degradation of TCA with hydrogen peroxide or alkaline activation of S2O82−. Moreover, photolysis of S2O82− has gained much attention due to its ability to decompose a wide range of organic contaminants. For instance, butylated hydroxyanisole was observed to be minimized 80−100% at acidic, neutral, and basic pH ranges by UV/S2O82− and the decay mechanism was dependent on solution pH.16 Perfluorinated compounds also appear amenable to photodegradation in the presence of S2O82−.17−19 Lin et al.20 demonstrated that phenol was successfully degraded at various pH conditions in UV/S2O82− processes. Salari et al.21 showed that the efficiency of C.I. basic yellow 2 (BY 2) decolorization in UV/S2O82− processes was sensitive to operational parameters, such as initial concentrations of S2O82−, BY 2, light intensity, flow rate, and solution pH. However, no known study has yet focused on TCA oxidation by UV/S2O82− processes. It is reported that the photooxidation performance of carbofuran (CBF) by UV/H2O2 and UV/S2O82− processes was diverse, but the intermediates during CBF oxidation were not involved in their study,22 and there is no report on direct comparison of these two photochemical processes on chlorinated contaminant degradation. Therefore, the objective of this work is to investigate and compare the effect of UV/ H2O2 and UV/S2O82− processes toward TCA oxidation in aqueous solution. The effects of solution conditions, such as the molar ratios of H2O2/TCA or S2O82−/TCA, solution pH, Cl− and HCO3− anions, and dissolved organic matter (DOM), such as humic acid (HA), on TCA removal performance are investigated. The intermediates during TCA oxidation in both processes were determined by gas chromatography/mass spectrometry (GC/MS), and a possible TCA oxidation pathway is proposed. The purpose of the investigation is to elucidate the mechanisms of AOPs based on •OH and SO4−• for TCA removal and therefore provide technical support for TCA-contaminated groundwater remediation.

ground, and the sampling plot equipped with a ground glass stopper as well to avoid volatilization of TCA. The aqueous solution was loaded and filled into the reactor, and the temperature was kept constant at 20 °C during all experiments with a cooling water jacket using a thermostat circulating water bath (SCIENTZ SDC-6, Zhejiang, China). A magnetic stirrer was located at the base of the reactor to ensure the solution remained homogeneous. Aqueous samples were taken at desired time intervals and analyzed immediately. The initial pH in all experiments was unadjusted except in the tests for investigating the influence of pH. All experiments were conducted in duplicate, and the mean values were reported. 2.3. Analytical Methods. The concentration of TCA was quantified by a gas chromatograph (Agilent 7890A, Palo Alto, CA) equipped with an electron capture detector, an autosampler (Agilent 7693), and a DB−VRX column (60-m length, 320-μm i.d., 1.4-μm thickness). The temperatures of the injector and detector were 240 and 260 °C, respectively, and the oven temperature was isothermal at 75 °C. The amount of injected sample was 1 μL with a split ratio of 20:1. Aqueous samples were analyzed after extraction with hexane. The volatile organic intermediates formed in TCA oxidation experiments were identified by the EPA SW-846 Method 5030B and 8260B using an automatic purge and trap (Tekmar Atomx, Mason, OH) coupled to a GC/MS (Agilent 7890/5975) and with the same DB-VRX column as before. The carboxylic acid intermediates were identified with the EPA Method 552.3 using a GC/MS (Shimadzu GC/MS-QP 2010, Kyoto, Japan) equipped with an HP-5 column (30-m length, 250-μm i.d., 0.25-μm thickness). 30 mL samples were acidified with concentrated sulphuric acid (pH < 0.5) and extracted with 3 mL MTBE after the addition of 13.5 g Na2SO4. The ether fraction was methylated with acidic methanol. The concentration of S2O82− was determined using a spectrophotometric method with potassium iodide.23 H2O2 was also analyzed using a spectrophotometric method using titanium−sulfuric acid solution.24 The pH was measured with a pH meter (MettlerToledo DELTA 320, Greifensee, Switzerland). The concentration of chloride ions was detected by anion chromatograph (Dionex ICS-I000, Sunnyvale, CA).

2. MATERIALS AND METHODS 2.1. Materials. The following reagents were purchased from Shanghai Jingchun Reagent Co. Ltd. (Shanghai, China) and used without further purification: TCA (99%), sodium persulfate (Na2S2O8, 98.0%), sodium chloride (99.5%), sodium bicarbonate (NaHCO3, 99.5%), sodium phosphate dibasic dodecahydrate (Na2HPO4·12H2O, 99.0%), sodium dihydrogen phosphate dihydrate (NaH2 PO4 ·2H 2O, 99.0%), sodium phosphate (Na3PO4, 99.0%), potassium iodide (KI, 99.0%), sodium thiosulfate (Na2S2O3, 99.0%), titanium oxysulfatesulfuric acid hydrate (TiOSO4·xH2SO4·xH2O, 93%), ammonium sulfate ((NH4)2SO4, 99.0%), hydrogen peroxide (H2O2, 30%), humic acid (fulvic acid >90%, as fulvic acid is one of the main fractions of dissolved organic matter with high solubility), hexane (C6H14, 97%), and methyl tert-butyl ether (MTBE, 99.9%). Ultrapure water from a Milli-Q water process (Classic DI, ELGA) was used for preparing aqueous solutions. 2.2. Photoreaction Experiments. Irradiation experiments were conducted in a 1 L cylindrical glass reactor (an inner diameter of 7.0 cm and a height of 25 cm) with a quartz tube in the center of the reactor. A 10 W, low-pressure mercury vapor lamp (Guangdong, China), emitting a monochromatic wavelength at 254 nm, was placed in the quartz tube with a photon flux of 2.1 × 10−5 Einstein cm−2 s−1. The TCA concentration in all tests was prepared in an aqueous solution at 20 mg L−1 (0.15 mM), except for the intermediate investigation, where 100 mg L−1 TCA was used. The joint of the reactor and quartz tube was

3. RESULTS AND DISCUSSION 3.1. Performance of UV/H2O2 and UV/S2O82− Processes in TCA Oxidation. Preliminary experiments revealed that there was no observable TCA removal when UV irradiation was applied in the absence of H2O2 or S2O82−, nor was there TCA removal after addition of H2O2 or S2O82− without UV radiation (data not shown). However, TCA removal did occur when H2O2 or S2O82− at various oxidant/ TCA molar ratios was applied together with UV irradiation (Figure 1). In addition, the TCA oxidation in the above two processes followed the pseudo-first-order kinetic model. The apparent reaction rate constants (k) calculated under various experimental conditions are listed in Table 1, which shows that the k values in UV/H2O2 processes increase from 0.0082 min−1 to 0.012 min−1 when the molar ratio increases from 1/1 to 5/1, but further increments result in a retardation of the oxidation rate because at higher H2O2 dosage, hydroperoxyl radical (HO2•), which is less reactive than •OH is formed (eq 1, k = 1.2−4.5 × 107 M−1 s−1), and HO2• can further undergo a chain termination reaction to scavenge hydroxyl radical (eq 2, k = 1 × 1010 M−1 s−1). Moreover, excess of hydroxyl radicals also dimerize to H2O2, according to eq 3 (k = 4.2−5.3 × 109 M−1 7197

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light, and sulfate radicals would therefore be formed more readily than hydroxyl radicals. It should be noted that H2O2/ TCA and S2O82−/TCA molar ratios in subsequent experiments were set at 5/1 and 3/1, respectively. 3.2. Photodegradation of TCA at Different Solution pH in UV/H2O2 and UV/S2O82− Processes. The TCA photodegradation performances at different solution pH in UV/H2O2 and UV/S2O82− processes are shown in Figure 2. In pH-unadjusted tests, the solution pH dropped from 6.2 and 5.6 to 3.5 and 3.1 in UV/H2O2 and UV/S2O82− processes, respectively. The pH reduction was partially due to the production of protons and acid byproduct.16,29 In subsequent tests investigating the influence of pH, the solution was varied from pH 3 to pH 11 and was phosphate (0.1 M) buffered, except at pH 3. The pH varied within 0.2 units during experiments in buffered solutions. The results indicated that the rate of TCA oxidation in both processes was pH-dependent. However, it should be noted that the best results were obtained in the pH-unadjusted test in both UV/H2O2 and UV/S2O82− processes because the buffered solution prevented the decrease in solution pH and limited TCA removal. In the UV/H2O2 process, the maximum rate of TCA oxidation occurred at pH 6 with an apparent reaction rate constant of 0.0093 min−1. Based on the results in Table 1, the efficiency of TCA removal varied by pH, in the following order: pH 6 > pH 8 ≥ pH 3 ≥ pH 7 > pH 11. Since the pKa of hydrogen peroxide is 11.6530,31 so that the dissociation of H2O2 did not significantly occur in the pH range 3−8. At pH 11, a considerable amount of H2O2 (approximate 20%) would dissociate in ionic form. In addition, the self-decomposition of H2O2 (eq 4) is found to be strongly dependent on pH, and the first-order self-decomposition rate constant of H2O2 tended to increase with increasing solution pH, especially at high pH.32 Therefore, TCA removal efficiency decreased significantly at pH 11.

Figure 1. Oxidation of TCA under various peroxide/TCA molar ratios. (a) UV/H2O2; (b) UV/S2O82−. (TCA = 20 mg L−1; 20 °C).

2H 2O2 → H 2O + O2

s −1 ), which reduces the overall oxidative capacity in solution.22,25 •

OH + H 2O2 → H 2O + HO2•

Unlike in the UV/H2O2 process, the oxidation rate in the UV/S2O82− process was highest at pH 3 and decreased with an increase in initial solution pH (Figure 2b), indicating that acidic conditions are propitious to TCA oxidation. However, alkaline activation of S2O82− has been reported to effectively generate some reactive species, which have potential to decompose oxidant and reductant probe compounds.8,33 Yet, an excess molar ratio of base to persulfate was required in the alkaline activation system. In our present study, the quantity of the base was less compared with the S2O82− dosage at pH 11; hence, we speculated that the base used simply consumed S2O82− as an inhibitor instead of functioning as an activator. House34 ever reported that the observed rate constant of decomposition of peroxydisulfate increased with decreasing pH at various temperatures, and under acidic conditions, the breakdown of S2O82− into sulfate free radicals can be further acidic-catalyzed according to eqs 5 and 6. Acidic pH:

(1)

•

OH + HO2• → O2 + H 2O

(2)

2•OH → H 2O2

(4)

(3)

UV/S2O82−

Such retardation was not observed in process, however, and the TCA oxidation rate increased from 0.014 min−1 to 0.23 min−1 when the molar ratio of S2O82−/TCA increased from 1/1 to 10/1, demonstrating that the UV/S2O82− process was more effective for TCA removal than the UV/ H2O2 process. A study developed by Neta at al.26 indicated that SO4−• was more effective in many aromatic compounds oxidation than •OH, with higher selectivity. The results could also be attributed to the presence of both sulfate and hydroxyl radicals produced by the photolytic activation of S2O82−.27 Alternatively, S2O82− is suggested to be more easily cleaved than H2O2 because of the distance of the O−O bond.28,29 Anipsitakis and Dionysiou28 also demonstrated that S2O82− absorbs more strongly at 254 nm than H2O2 does and the degradation efficiency of 2,4-dichlorophenol is in the order UV/K2S2O8 > UV/H2O2. Yang et al.29 found that the degradation rates of azo dye Acid Orange 7 were in the order of UV/Na2S2O8 > UV/H2O2 as well. Therefore, we concluded that S2O82− could be more easily activated by UV

S2 O82 − + H+ → HS2 O8−

(5)

HS2 O8− → SO4 −• + SO4 2 − + H+

(6)

Finally, it has to be noted the possible effects of phosphate buffers used in the two processes. In the UV/H2O2 process, • OH are considered to be the major reactive species, the rates 7198

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Table 1. Results of TCA Oxidation Performance under Various Conditions pseudo-first-order rate constant, k (min−1) operational condit.

UV/H2O2

H2O2/TCA = 1 H2O2/TCA = 5 H2O2/TCA = 10 H2O2/TCA = 100 H2O2/TCA = 200 S2O82−/TCA = 1 S2O82−/TCA = 3 S2O82−/TCA = 5 S2O82−/TCA = 10 pH = 3 (unbuffer)a pH = 6 (buffer) pH = 7 (buffer) pH = 8 (buffer) pH = 11 (buffer)b Cl− = 1 mM Cl− = 100 mM HCO3− = 1 mM HCO3− = 100 mM HA = 1 mg L−1 HA = 10 mg L−1

0.0082 0.0119 0.0108 0.0108 0.0062

UV/S2O8

correlation coefficient, R2

2−

UV/H2O2

UV/S2O82−

0.9945 0.9936 0.9906 0.9894 0.9917 0.0143 0.0559 0.1446 0.2257 0.0375 0.0272 0.0140 0.0106 0.0038 0.0280 0.0014 0.0025 0.0010 0.0528 0.0258

0.0075 0.0093 0.0066 0.0077 0.0043 0.0092 0.0019 0.0075 0.0008 0.0125 0.0120

0.9858 0.9602 0.9821 0.9627 0.9883 0.9768 0.6781 0.9588 0.9000 0.9902 0.9931

pH (initial/final) UV/H2O2

UV/S2O82−

6.15/3.81 6.15/3.48 6.18/3.67 6.10/3.74 6.04/4.01 0.9977 0.9775 0.9555 0.9880 0.9759 0.9794 0.9595 0.9384 0.9635 0.9658 0.8303 0.9893 0.8778 0.9670 0.9850

3.01/2.85 6.01/5.96 7.10/7.00 8.00/7.99 11.01/10.89 6.21/3.45 6.60/4.23 6.80/6.74 8.55/8.35 /3.71 /3.75

5.80/3.50 5.55/3.13 5.40/2.89 5.14/2.56 2.70/2.44 6.00/5.89 7.00/6.96 7.95/7.88 10.99/10.97 6.01/3.20 6.33/3.29 6.63/5.57 8.50/7.42 /3.09 /3.07

a Solution was unbuffered and initial pH was adjusted by 0.1 M sulfuric acid. bSolution was buffered with 0.1 M Na2HPO4·12H2O and 0.1 M Na3PO4.

and extensive inhibition occurred at the higher concentrations. As hydroxyl radical scavengers, the rate constants for Cl− and HCO3− reactions with •OH were 4.3 × 109 and 8.5 × 106 M−1 s−1, respectively.40,41 The suppressive effect of Cl− on TCA removal in the UV/S2O82− process was similar to the UV/H2O2 process, but a more significantly inhibitive effect was caused by HCO3− (Figure 3b), as only 26.7% and 10.1% of TCA removal was achieved at HCO3− concentrations of 1 and 100 mM, respectively. The possible chemical scavenging mechanisms of Cl− and HCO3− in the UV/S2O82− process are shown in eqs 10−13. The rate constants for eqs 10 and 12 are 4.7 × 108 and 2.8 × 106 M−1 s−1.42,43 The adverse effect of HCO3− on TCA oxidation rate was possibly due to elevation of the initial solution pH and the solution buffering capacity as bicarbonate was added, thus suppressing the oxidation rate, as discussed in section 3.2.

of which reacting with the various phosphate species are shown in eqs 7−9.35 •

OH + PO34− → product

•

OH + HPO24 − → product

k 7 < 1 × 107 M−1 s−1

(7)

k 8 = 1.5 × 105 M−1 s−1 (8)

•

OH + H 2PO−4 → OH− + H 2PO4• k 9 = 2.0 × 104 M−1 s−1

(9) •

On the other hand, the rate constants of OH with TCA are reported to be 2 × 107 to 1 × 108 M−1 s−1,6,36 much higher than the rates in eqs 7−9. Whereas, the rate constants of SO4−• with HPO42− and H2PO4− are 1.2 × 106 and 7× 104 M−1 s−1.37 Maruthamuthu and Neta37 found that SO4−• may react with phosphate anions to propagate a chain reaction and form various phosphate radicals, the oxidation potentials of which were all less than SO4−•. Therefore, it can be deduced that •OH and SO4−• reacted slower with phosphate anions and the impact of buffered phosphate on TCA oxidation performance in this study is minimal. The results are also consistent with the conclusions of some other researchers.35,38,39 For instance, Liang and Bruell39 have demonstrated that phosphate-buffered system may not have significant effect on heat-activated persulfate oxidation of TCE. 3.3. Effects of Solution Matrix on TCA Oxidation in UV/H2O2 and UV/S2O82− Processes. Chloride and bicarbonate ions are ubiquitous in natural waters, necessitating a study of the influence of Cl− and HCO3− on TCA removal in UV/H2O2 and UV/S2O82− processes (Figure 3a,b). The results of this study demonstrated that both Cl− and HCO3− had significant scavenging effects at the anion concentrations tested (1 and 100 mM). In the UV/H2O2 process, a much more inhibitive effect was observed for HCO3− compared with Cl−,

SO4 −• + Cl− → SO4 2 − + Cl•

(10)

Cl• + Cl− → +Cl 2−•

(11)

SO4 −• + HCO3− → SO4 2 − + HCO3•

(12)

HCO3• ↔ H+ + CO3−•

(13)

Natural organic matter (NOM) is a complex heterogeneous mixture of organic compounds present in surface or groundwater. Westerhoff et al.44 demonstrated that rate constants for the reactions between •OH and seven DOMs fell within the range 1−5 × 108 M−1 s−1. In this study, the influence of HA (1 and 10 mg L−1) on TCA removal was investigated, and the results are shown in Figure 3c. Table 1 also shows that, in the UV/H2O2 process, the rate constants are 0.0119 min−1 (from 0−120 min), 0.0125 min−1 (from 0−120 min), and 0.0120 min−1 (from 20−120 min) in the presence of 0, 1, and 10 mg L−1 HA, respectively. A slight promotion of TCA oxidation was 7199

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Figure 2. Effect of initial solution pH on TCA oxidation performance. (a) UV/H2O2: molar ratio of H2O2/TCA = 5/1. (b) UV/S2O82−: molar ratio of S2O82−/TCA = 3/1. (TCA = 20 mg L−1; 20 °C; “control” represents the tests of TCA volatilization without UV and oxidants).

observed with 1 mg L−1 HA and a lag phase of TCA oxidation was observed with 10 mg L−1 HA. In the UV/S2O82− process, more than 90% of the TCA removal was achieved after 120 min at the higher HA concentration. However, a marked lag phase was observed during the initial 40 min, which indicated that HA competed with TCA for reactive radicals, with a rate constant of 0.026 min−1 from 30−120 min. 3.4. Performance of the Oxidant Decomposition in UV/H2O2 and UV/S2O82− Processes. Since the oxidant demand is an important parameter in determining the efficacy and cost of the contaminated sites remediation, the performances of the oxidant decomposition are evaluated under various conditions investigated before (Figure 4). Both H2O2 and S2O82− decomposition in UV processes followed the pseudofirst-order kinetic model and the rate constants under selected conditions are listed in Table 2. In the UV/H2O2 process, the rate constants were less different under various conditions (around 0.02 min−1), indicating that the performance of H2O2 decomposition was not significantly influenced by other factors. However, Chu32 demonstrated that the first-order constants of H2O2 in UV process marginally increased from pH 2.5 to 7.0 and shot up from pH 7.0 to 10.5. In the UV/S2O82− process, when 0.1 M HCO3− or 10 mg L− HA was added, or the solution was adjusted to pH 6, there were no obvious

Figure 3. Effect of Cl− and HCO3− anions and HA on TCA oxidation performance. (a) UV/H2O2: molar ratio of H2O2/TCA = 5/1; (b) UV/S2O82−: molar ratio of S2O82−/TCA = 3/1. (TCA = 20 mg L−1; Cl− = HCO3− = 1 mM and 100 mM; 20 °C); (c) TCA = 20 mg L−1; molar ratio of H2O2/TCA = 5/1; molar ratio of S2O82−/TCA = 3/1; 20 °C.

differences in S2O82− decomposition rate constants. However, S2O82− decomposition rate decreased visibly when 0.1 M Cl− was added. Even so, UV/S2O82− process is less oxidantconsuming and more effective with higher oxidant utilization than heat- and alkaline-activated S2O82− processes. For instance, in our previous study,14 TCA could be completely 7200

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expected that the oxidation products in both processes would be diverse. To the best of our knowledge so far, little has been reported on the intermediates and oxidation pathways of TCA by either •OH or SO4−• in aqueous solution, except acetic acid being identified from •OH-induced degradation of TCA in oxygen-containing aqueous solutions reported by Lal et al.46 The UV irradiation of 1,1,2-trichloroethane in the presence of H2O2 resulted in monochloroacetic and dichloroacetic acids (MCAA and DCAA) as organic products.47 Waldemer et al.48 identified only trans-DCE and cis-DCE as the major intermediates of cis-DCE and trans-DCE oxidation by heatactivated S2O82−, whereas no organic intermediates were detected during TCE and tetrachloroethylene (PCE) oxidation. In this study, MCAA was the only chlorinated compound identified in the UV/H2O2 process as a major intermediate, and no volatile intermediates were detected, while DCAA and trichloroacetic acid (TCAA) were confirmed by derivatization in the UV/S2O82− process. In addition, the identification of volatile intermediates by GC/MS revealed that 1,1-DCE, 1,1,1,2-tetrachloroethane, PCE, chloroform and carbon tetrachloride (CT) were formed in the UV/S2O82− process. Figure 5

Figure 4. Photodecomposition of H2O2 (a, molar ratio of H2O2/TCA = 5/1) and S2O82− (b, molar ratio of S2O82−/TCA = 3/1) under selected conditions. (TCA = 20 mg L−1; 20 °C).

Table 2. Results of H2O2 and S2O82− Photodecomposition Performance under Selected Conditions pseudo-first-order rate constant, k (min−1) operational conditions

H2O2

only H2O2 only S2O82− pH = 6 (buffer) Cl− = 100 mM HCO3− = 100 mM HA = 10 mg L−1

0.0223 0.0177 0.0166 0.0183 0.0198

S2O82−

Figure 5. Evolution of volatile organic intermediates in UV/S2O82− process. (TCA = 100 mg L−1; molar ratio of S2O82−/TCA = 3/1; 20 °C; “control” represents the tests of TCA volatilization without UV and oxidants).

correlation coefficient, R2 UV/H2O2

UV/S2O82−

0.9949 0.0188 0.0201 0.0103 0.0185 0.0203

0.9972 0.9977 0.9897 0.9738

0.9953 0.9941 0.9924 0.9971 0.9925

shows the fate of the above volatile intermediates during TCA oxidation, all of which were also degraded by S2O82− chemistry under the experimental conditions, except for trace CT (