H2O2 Advanced Oxidation Process

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Environ. Sci. Technol. 2007, 41, 1696-1703

Trichloroethene Degradation by UV/H2O2 Advanced Oxidation Process: Product Study and Kinetic Modeling K E L I , †,| M I H A E L A I . S T E F A N , * ,‡,⊥ A N D JOHN C. CRITTENDEN§ Department of Civil and Environmental Engineering, Michigan Technical University, Houghton, Michigan 49931, Department of Chemistry, The University of Western Ontario, London, Ontario, Canada N6A 5B7, and Department of Civil and Environmental Engineering, Arizona State University, P.O. Box 872803, Tempe, Arizona 85287-2803

The broadband UV irradiation of 1.1 mM trichloroethene (TCE) aqueous solution in the presence of 10.4 mM H2O2 resulted in formic, oxalic, dichloroacetic (DCA), and monochloroacetic (MCA) acids, as organic byproducts. The organic chlorine was converted completely to chloride ion as a final product. TCE and its degradation products were completely mineralized in 30 min, under a volumeaveraged UV-C irradiant power of 35.7 W/L from a 1 kW medium-pressure mercury vapor arc lamp. TCE degraded primarily through hydroxyl radical-induced reactions and only to a low extent through direct UV photolysis and chlorine atom-induced chain reactions. The experimental patterns of TCE, H2O2, and detected reaction products combined with the literature information on radical reactions in the aqueous phase were used to postulate a degradation mechanism and to develop a kinetic model to predict the TCE decay, formation and degradation of byproducts, and pH and oxygen profiles. The agreement between the model calculations and the experimental data is satisfactory.

Introduction Trichloroethene (TCE) is a volatile organic compound that has been proven to cause liver damage and kidney failure in humans and was assessed as carcinogenic to animals. TCE is produced worldwide in very large quantities due to its extensive use in the chemical industry. Leakage from storage tanks and, sometimes, improper disposal practices led to soil, groundwater, and surface water contamination with TCE. Consequently, trichloroethene is listed as a priority pollutant on the U.S. EPA Chemical Contaminant List and is strictly regulated in drinking water to a maximum contaminant level of 0.005 mg/L. The UV/H2O2 process is currently successfully applied as an effective treatment of trichloroethene in groundwater and * Corresponding author phone: (519)457-3400; fax: (519)457-3030; e-mail: [email protected]. † Michigan Technical University. ‡ The University of Western Ontario. § Arizona State University. | Current affiliation: Department of Civil and Environmental Engineering, Arizona State University, P.O. Box 872803, Tempe, AZ 85287-2803. ⊥ Current affiliation: Trojan Technologies, Inc., 3020 Gore Road, London, ON, Canada N5V 4T7. 1696

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reclaimed water. The treatment process relies primarily on the •OH radical-induced oxidation of TCE. Trichloroethene exhibits an absorption band in the wavelength range of 200250 nm [max(200 nm))7183 M-1 cm-1 (1)] and undergoes direct photolysis with an overall quantum yield of φ ) 0.354 (1). Extensive studies have been conducted on TCE degradation by the UV/H2O2 process (2-8). Most of the published literature focused on the kinetics of TCE decay in various water qualities and on the associated treatment cost, and only a few studies reported the intermediate products formed. Hirvonen et al. (9) examined the formation of byproducts during TCE degradation in the UV/H2O2 process applied to groundwater remediation and detected di- and monochloroacetic acids. Studies on γ radiation-induced degradation of TCE indicated also the formation of chlorinated acids as reaction intermediates (10, 11). Haloacetic acids (HAAs) are formed during the drinking water disinfection with chlorine and are commonly known as disinfection byproducts (DBPs), which could cause bladder, colon, and rectal cancer (12). According to the U.S. EPA regulations for drinking water the HAA total concentration must not exceed 60 µg/L. In this work, we investigated the TCE degradation mechanism in the presence of H2O2 under irradiation with a 1 kW medium-pressure Hg lamp. Based on the proposed reaction mechanism and the kinetic parameters available in the literature, we have developed a kinetic model to predict the dynamics of this system and found a satisfactory agreement between the calculated and experimental data.

Experimental Section Reagents. All chemical reagents were of analytical reagent grade and were used as received. Catalase from bovine liver (Sigma, 19 900 units/mg) was used to destroy any residual hydrogen peroxide in the sample vials prior to IC analysis. Apparatus. All experiments were carried out in a stainless steel 1 kW bench-scale Rayox reactor described elsewhere (13). The total reaction volume was 28 L, and the reactor was airtight. In the wavelength range of 200∼300 nm, the volumeaveraged photon flow entering the reactor was 7.8 × 10-6 einstein s-1 L-1, as determined by potassium persulfate actinometry (14). The physical path length of the reactor was 10.4 cm. Additional information on the experimental setup is available elsewhere (1). Analytical Methods. The analytical methods for TCE and other organic acids were reported previously (1). The molar absorption coefficients of TCE, H2O2, and organic acids are provided in the Supporting Information. Hydrogen peroxide was determined according to the triiodide method (15). The oxygen level was monitored in the holding reservoir with an ATI Orion Model 840 DO meter.

Results and Discussion Product Study. Trichloroethene (1.1 mM) in distilled water was irradiated in the presence of 10.4 mM H2O2. At these concentrations, the fractions of UV light absorbed by TCE and H2O2 within 200-300 nm range were 24.9% and 59.5%, respectively. Under such conditions, TCE degraded through both oxidative chain reactions initiated by the OH radicals and chlorine atoms, and direct UV photolysis due to its strong absorption band below 240 nm and large quantum yield [φ)0.354 (1)]. TCE decayed by 2 orders of magnitude in less than 10 min, and fewer byproducts with lower yields were observed than during the direct UV photolysis alone (1). Figure 1 shows the time profiles of TCE, H2O2, and degradation byproducts. 10.1021/es0607638 CCC: $37.00

 2007 American Chemical Society Published on Web 02/02/2007

The major byproducts were formic and oxalic acids. Mono- and dichloroacetic acids were detected at low levels. No chlorinated ethynes and aldehydes were measured, likely due to their low levels and high reactivity toward hydroxyl radical. Chloride ion was released as a final product from early stages of irradiation through the end of the degradation of chlorinated byproducts, and the organic carbon total mineralization was achieved within ∼30 min exposure time. The chlorine mass balance was calculated and indicated that about 6% chlorine was missing from the system, most likely due to some TCE volatility losses in the headspace zone of the reactor. As the organic and hydrochloric acids were formed, the solution pH dropped sharply from 5.9 (t)0 min) to 2.6 (t)5 min). Dissolved oxygen level decreased from ∼7.1 mg/L (t)0 min) to 1.7-2.0 mg/L at 4-5 min reaction time and is associated with the mineralization of approximately 70% of the organic carbon. After 4 min, the oxygen level increased continuously up to ∼23 mg/L through the photodecomposition of H2O2 and due to a lower chemical oxygen demand than at short exposure times [under oxygen-saturated conditions, maximum O2 level in aqueous solutions could be as high as 40 mg/L [1.25 × 10-3 M (16)]. The pH and oxygen patterns over the irradiation time are shown in Figure 2. Reaction Mechanism for Photolysis of H2O2. The H2O2 photolysis and subsequent thermal reactions of the generated radicals were extensively studied (17-21). The reactions relevant to this system and considered in the kinetic model are as follows:

H2O2 + hv f 2•OH H2O2 + •OH f H2O + HO2•

φOH ) 1.0 (17)

109 M-1 s-1 (27); and 2.6 × 109 M-1 s-1 (28). In our kinetic model, a rate constant of 2.4 × 109 M-1 s-1 (25) was considered. The geminal chlorohydrins eliminate HCl very rapidly with first-order rate constants k > 7 × 105 s-1 (27, 29); a similar reaction occurs from chlorohydrin radicals:

ClCH(OH)-C•Cl2 f OHC-C•Cl2 + H+ + Cl-

k ) 5.1 × 105 M-1 s-1 (30) (7)

The reactions between the carbon-centered radicals and O2 are diffusion-controlled processes (k>109 M-1 s-1) forming peroxyl radicals. The peroxyl radical generated in reaction 8 evolves primarily into oxyl radicals and molecular oxygen (10, 11).

OHC-C•Cl2 + O2 f OHC-C(Cl2)OO•

k ) 3 × 109 M-1 s-1 (31) (8)

2OHC-C(Cl2)OO• f 2OHC-C(Cl2)O• + O2

k ) 4 × 108 M-1 s-1 (31) (9)

The oxyl radicals decay through fast unimolecular fragmentation either by a C-C bond scission or by elimination of a chlorine atom with formation of molecular products, which undergo rapid hydrolysis.

OHC-C(Cl2)O• f COCl2 + •CHO

k ) 1 × 106 M-1 s-1 (31) (10a)

f OHC-C(O)Cl + Cl•

(1)

k ) 2.7 × 107 M-1 s-1 (22) (2)

H2O2 + HO2•/O2•- f •OH + H2O/OH- + O2

Phosgene hydrolysis occurs with a pseudo-first-order rate constant of 6 s-1 (32) - 9 s-1 (16), with formation of CO2 and hydrochloric acid.

k ) 3.6 M-1 s-1 (22) (3)

•

•

•-

COCl2 + H2O f CO2 + 2H+ + 2Cl-

-

OH + HO2 /O2 f H2O/OH + O2

k ) 1 × 1010 M-1 s-1 (23) (4)

HO2• + HO2•/O2•- f H2O2/HO2- + O2

k ) 1 × 106 M-1 s-1 (20) (5)

At 254 nm, the primary quantum yield (φ) of hydrogen peroxide photolysis in eq 1 is 0.5. The extended set of the H2O2 photolysis reactions is incorporated in a complex kinetic model, which was validated for UV/H2O2 applications in completely mixed batch-reactors (24). TCE Decay and Byproduct Formation. A simplified representation of the proposed mechanism for TCE degradation is given in Figure 3. The three major routes that describe the TCE degradation start with the following primary processes: OH radical addition to the double bond, direct photolysis, and chlorine atom addition to the double bond. TCE Degradation through OH Radical-Induced Reactions. The OH radical addition to the double bond of trichloroethene yields a geminal chlorohydrin carbon-centered radical:

ClHCdCCl2 + •OH f ClCH(OH)-C•Cl2

k ) 2.4 × 109 M-1 s-1 (25) (6)

There are various rate constant values reported in the literature for reaction 6: 2.4 × 109 M-1 s-1 (25); 2.9 × 109 M-1 s-1 (26); 3.3 × 109 M-1 s-1 (10); 4.0 × 109 M-1 s-1 and 4.3 ×

k ) 1 × 105 M-1 s-1 (fitted) (10b)

(11)

The unstable chloroglyoxal OHC-C(O)Cl hydrolyzes to glyoxylic acid (reaction 12), whereas the formyl radical •CHO, after hydration and reaction with O2, generates formic acid through a fast HO2• elimination reaction [k>1 × 106 s-1 (33)]. The formyl radical can also be oxidized to CO and HO2• (10) [k)4 × 105 M-1 s-1 (fitted)].

OHC-C(O)Cl + H2O f OHCCOOH + H+ + Cl-

(12)

O2

CHO + H2O 98 HCOOH + O2•- + H+

•

k ) 1 × 106 M-1 s-1 (33) (13)

Therefore, the •OH-induced degradation of trichloroethene results in formic and glyoxylic acids as intermediates and chloride ion and CO2 as final products. TCE Direct Photolysis. The fraction of incident light (200300 nm) absorbed by TCE at t ) 0 min is 24.9%, as compared to 42.6% for the same initial concentration and light source, in the absence of hydrogen peroxide (1); on the other hand, H2O2 absorbs a significantly larger fraction of light than TCE under the conditions used in this work (Fabs(H2O2))59.5%, t)0 min). Therefore, the contribution of the direct photolysis to the overall product yield of TCE degradation in the presence of H2O2 is expected to be significantly smaller than that of the OH radical pathway. The direct photolysis of TCE occurs with an overall quantum yield of φ ) 0.354 (1) and is represented by the VOL. 41, NO. 5, 2007 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 1. Time profiles for the decay of TCE, H2O2, and intermediates in UV/H2O2 treatment process (Solid lines represent the computer modeling fits.). following primary processes:

ClHCdCCl2 + hν f ClHCdC•Cl + Cl• φ ) 0.13 (1) (14) H2O, φ ) 0.1 (1)

ClHCdCCl2 + hν 98 k ) 3.2 × 105 s-1 (fitted)

ClHC(OH)-CHCl2 98

Cl2HC-CHO + H+ + Cl- (15)

The mesomeric form of oxyl radical formed in reactions 19 is a carbon-centered radical, which generates another oxyl radical via reaction with oxygen (reaction sequence 20), which decays through (a) chlorine elimination followed by hydrolysis to glyoxylic acid (reaction sequence 20) and (b) rapid intramolecular 1,2-H shift to a C-centered radical, which leads to oxalic acid (reaction 22) through hydrolysis of dichlorogyoxal formed in reaction sequence 21. The rate constant fitted for this reaction sequence was 1 × 106 s-1. O2

ClHCdCCl2 + hν f HCtCCl + Cl2

φ ) 0.032 (1) (16)

ClHCdCCl2 + hν f ClCtCCl + HCl φ ) 0.092 (1) (17) The geminal chlorohydrin formed in reaction 15 releases HCl to form dichloroacetaldehyde, a precursor of dichloroacetic acid (reaction 18). O2

Cl2HC-CH(OH)2 + •OH 9 8 9 -1 -1 k ) 2 × 10 M

s

(fitted)

k′ ) 1 × 107 M-1 s-1 (fitted)

Cl2HC-C(OH)2OO• 98 Cl2HC-COOH + HO2• (18) Monochloroacetic acid was detected at levels as low as 10-6 M. The 1,2-dichlorovinyl radical generated in reaction 14 is the major route to MCA formation, through the reaction with O2 with a rate constant k ∼ 4.3 × 109 M-1 s-1 (16), followed by chlorine atom elimination and rapid hydrolysis of chloroketene intermediate (reaction sequence 19). This pathway is supported by the experimental observation that MCA was not detected after 14 min of reaction when the fraction of absorbed UV light by TCE was negligible (less than 3%). Also, MCA was not reported in the •OH-induced degradation of TCE without UV irradiation (11). O2

ClHCdC•Cl 9 8 9 -1 -1 k ) 4.3 × 10 M

s

(16)

•

ClHCdC(Cl)OO 9 8 9 -1 -1 k ) 2 × 10 M

s

(fitted)

H2O

ClHCdC(Cl)O• 9 8 ClH2CCOOH + Cl• 6 -1 k ) 1.1 × 10 s

1698

9

(fitted)

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 41, NO. 5, 2007

ClHC•-C(O)Cl 98 +H2O,-Cl•

•

O(Cl)HC-C(O)Cl 98 OHCCOO- + Cl- + 2H+ (20) O2, -HO2•

•

O(Cl)HC-C(O)Cl f HO(Cl)C•-C(O)Cl 98 Cl(O)C-C(O)Cl (21) Cl(O)C-C(O)Cl + H2O f HOOCCOOH + 2H+ + 2Cl(22)

The quantum yield of dichloroethyne formation (reaction 17) was estimated as φ ) 0.092 (1). Hydrolysis and oxidation induced by the hydroxyl radical are the expected degradation pathways: H3O+

H2O

ClCtCCl 98 ClCH2C(O)Cl 98 ClCH2COOH

(23)

•OH

ClCtCCl 9 8 9 -1 -1 k ) 3 × 10 M •

s

(fitted)

-ΗCl

H2O

Cl(OH)CdC Cl 98 OdCdC•Cl 98 HOOC-HC•Cl (24) O2

HOOC-HC•Cl 98 HOOC-HC(Cl)O• f CO2•- + H+ + HC(O)Cl (25) Formyl chloride HC(O)Cl decomposes into HCl and CO with a rate constant as high as 104 s-1 (31). The photolytic path to monochloroethyne (reaction 16) occurs with a quantum yield of φ ) 0.032 (1), and its contribution to the TCE decay must be negligible. Therefore, the major degradation intermediates expected to originate from TCE photolysis are mono- and dichloroacetic acids; oxalic and glyoxylic acids are predicted mecha-

FIGURE 2. pH and O2 patterns (Solid line represents the computer modeling fits.).

FIGURE 3. Schematic representation of TCE degradation mechanism [/photolysis products (1)]. nistically, but their yields must be very low. Chloride ion is formed as a final product. TCE Degradation through Chlorine Atom-Induced Reactions. The chlorine atoms are formed through the C-Cl photolytic bond cleavage in TCE molecule (reaction 14) and C-Cl bond scission in chlorinated organic radicals. To our knowledge, the rate constant for the addition of chlorine atom to the TCE double bond (reaction 26) is unknown. The chlorine reactions with substituted olefins occur with rate constants lower by approximately one order of magnitude than those for the OH radical reactions: kCl,Cl2C)CCl2 ) 2.8 × 108 M-1 s-1 (30); kCl,H2C)CH2 ) 2.0 × 108 M-1 s-1; kCl,H2C)CHCH2Cl ) 3.0 × 108 M-1 s-1 (34). Therefore, we believe that a rate constant of k ) 1.9 × 108 M-1 s-1 for reaction 26 would be appropriate and used it in our kinetic modeling.

ClHCdCCl2 + Cl• f Cl2HC-C•Cl2

k ) 1.9 × 108 M-1 s-1 (fitted) (26)

The carbon-centered radical leads to dichloroacetic acid through the reaction sequence 27:

O2

Cl2HC-C•Cl2 98 H2O [k ) 1 × 105 s-1 (fitted)]

Cl2HC-CCl2O• 98

Cl2HCCOOH + H+ + Cl-+ Cl• (27)

H2O, O2

Cl2HC-CCl2O• 98 CO + CO2 + H+ + Cl-

k ) 9 × 105 s-1 (fitted) (28)

The last step in reaction sequence 27 was also suggested as a pathway to DCA in pulse radiolysis studies on TCE (11). The oxyl radical can also undergo C-C bond cleavage followed by hydrolysis of phosgene and oxidation of Cl2HC• radical to inorganic products (reaction 28). Therefore, the intermediate product resulted from the TCE degradation through chlorine atoms is dichloroacetic acid. The chlorine atoms are effectively scavenged by both H2O2 (reaction 29) and chloride ions (reaction 30) due to their large concentrations relative to that of TCE as well as to the large reaction rate constants. Bicarbonate ions resulted through mineralization processes react with the chlorine atoms, generating carbonate radical ions. Such competitive VOL. 41, NO. 5, 2007 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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reactions minimize the role of chlorine atoms on the overall TCE degradation rate. Dichlorine radical ion Cl2•- (reaction 30) reacts with both H2O2 (reaction 31) and TCE with rate constants much smaller than those of chlorine atoms. Therefore, slower degradation rates are expected for TCE once the chloride ion builds up in the solution. •

+

-

k ) 3 × 108 M-1 s-1 (assumed) (29) k ) 8.6 × 109 M-1 s-1 (35) (30)

Cl2•- + H2O2 f H+ + 2Cl- + HO2•

k ) 1.4 × 105 M-1 s-1 (36) (31)

The rate constant for reaction 29 is not known, but it is expected to be similar to that of bicarbonate-chlorine atom reaction [k)2.2 × 108 M-1 s-1 (30)]. Also, the rate constant for the TCE reaction is unknown, but based on the literature data for unsubstituted olefins k ) ∼5 × 106-5 × 107 M-1 s-1 (34), it is conceivable to assume it as ∼1 × 107 M-1 s-1. The chlorine atom and chlorine radical ion reactions discussed here were considered in the kinetic model. Although no attempts were made to distinguish experimentally among various pathways of TCE decay, based on the intermediate product yields and patterns and the competition kinetics for the free radicals generated during the course of irradiation, the kinetic model predicts the following order for the relative contributions of these pathways to the degradation rate of TCE, integrated over the first 10 min reaction time: •OH . UV photolysis > Cl2•- . Cl• (Kinetic Modeling section). Degradation of Organic Acids. The direct photolysis of detected organic acids is negligible, their degradation occurring through OH radical-induced reactions. Glyoxylic acid reacts with both H2O2 and the OH radical. Its degradation products are formic and oxalic acids, which are oxidized further to CO2, H2O, and formoyl radical.

OHCCOOH + H2O2 f HCOOH + CO2 + H2O

k ) 0.3 M-1 s-1 (37) (32) H2O, O2

OHCCOOH (OHCCOO-) + •OH 98 HOOCCOOH (HOOCCOO-) + HO2• (33) The rate constants for the OH radical reactions of protonated and unprotonated forms of glyoxylic acid were determined recently as k ) 5.9 × 108 M-1 s-1 and k ) 1.3 × 109 M-1 s-1, respectively (38).

HCOOH (HCOO-) + •OH f CO2•- + H+ + H2O

(34)

HOOCCOOH (HOOCCOO-, -OOCCOO-) + • OH f CO2 + CO2•- + H+ + H2O (35) The rate constants for the •OH reactions of protonated/ unprotonated forms of formic and oxalic acids are known (22) and used in the kinetic model. Both MCA and DCA degraded through H-abstraction reaction by •OH, followed by the peroxyl radical chemistry as shown below.

ClCH2COOH + •OH f ClC•HCOOH + H2O

k ) 4.3 × 107 (4 × 108) M-1 s-1 (25, 39) (36)

ClC•HCOOH + O2 f •Ο2(Cl)HCCOOH 1700

9

Cl2CHCOOH + •OH f Cl2C•COOH + H2O

k ) 2.75 × 107 (9.2 × 107) M-1 s-1 (39) (39)

Cl2C•COOH + O2 f •O2(Cl2)CCOOH

(40)

2•O2(Cl2)CCOOH f 2•Ο(Cl2)CCOOH + O2

(41)

•

Cl + H2O2 f H + Cl + HO2 Cl• + Cl- f Cl2•-

2•Ο2(Cl)HC(Cl)COOH f 2•Ο(Cl)CHCOOH + O2 (38)

(37)

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 41, NO. 5, 2007

The oxyl radicals formed in eqs 38 and 41 can decay through both C-C scission and elimination of a chlorine atom: •

Ο(Cl)CHCOOH f CO2•- + H+ + HC(O)Cl (42a) f OHCCOOH + Cl•

(42b)

Ο(Cl2)CCOOH f CO2•- + H+ + Cl2CdO

(43a)

f Cl(Ο)COOH + Cl• f HOOC-COOH

(43b)

•

From the photocatalytic degradation of TCE, Mao et al. (40) concluded that the Cl-elimination was not a significant route because the yields of glyoxylic and oxalic acids resulted from the degradation of MCA and DCA represented less than 6% and 3% of the total carbon degraded, respectively. Formoyl radical (CO2•-) is produced from various radical or molecular intermediates and decays through the reactions with O2 and H2O2:

CO2•- + O2 + H+ f CO2 + HO2•/O2•-

k ) 2 × 109 M-1 s-1 (34) (44)

CO2•- + H2O2 f CO2 + OH- + •OH

k ) 6.3 × 105 M-1 s-1 (41) (45)

Bimolecular decay of CO2•- does not occur given the strong competition from the reactions 44 and 45. Peroxyl radical chemistry is more complex than as depicted above. Cross-termination reactions among themselves or involving HO2• which reaches concentrations as high as 1 × 10-6 M are possible as well as other routes of decay of the tetroxide intermediates. The mechanism suggested above comprises the most relevant reactions in this system and are characteristic to the halogenated peroxyl radicals. Other reactions involving bicarbonate ion and the resulting carbonate radical ion were not described in the mechanism but were considered in the kinetic model.

Kinetic Modeling The model predicts the destruction of TCE and the formation and fate of MCA, DCA, and oxalic and formic acids, along with the pH and oxygen time profiles (Figures 1 and 2). The reactions and kinetic parameters used in the model are provided throughout the paper and summarized in Table S1 in the Supporting Information. Selection of the Species and Reactions Used in the Model. A critical screening of the reactions occurring in this system was performed based on the competition kinetics at given reaction times. Some basic considerations are given below. The ionized and nonionized forms of organic acids have different reactivity toward the OH radical, and the kinetic model accounts for that. The concentrations of the protonated/unprotonated forms were calculated according to the appropriate acid/base equilibrium constants. The H2O2/ HO2- and HCO3-/CO32-equilibria were ignored given the pH range observed during the irradiation and the corresponding pKa values [11.75 (42) and 10.75 (43), respectively].

The superoxide radical anion O2•- reactions were considered because the concentration of this radical is about 10 times larger than that of its conjugated acid HO2• at the beginning of irradiation and cannot be ignored within the first minute reaction time. The self-termination reaction of HO2• with regeneration of H2O2 and its reaction with •OH were included because the concentration of HO2• reaches levels as high as 1 × 10-6 M. Such reactions are also among the few ones that contribute to the O2 formation. The carbonate species not only affect the pH but are also scavengers of radical species, such as •OH and Cl•. The electron-transfer reactions with these radicals lead to carbonate radical ion CO3•-, which disappears primarily through the reaction with H2O2 [4.3 × 105 M-1 s-1 (34)]. The chlorine atom reactions were captured in the model. The steady-state concentration of Cl• was calculated as 10-1310-14 M over the first 5 min exposure time and decreased to 10-18 M by 30 min. The major scavengers of Cl• are H2O2 (reaction 29), Cl- (reaction 30), and HCO3- (reaction 46).

HCO3- + Cl• f CO3•- + H+ + Cl-

k ) 2.2 × 108 M-1 s-1 (30) (46)

The dichlorine radical ion Cl2•- is an electrophile, which reacts with organic and inorganic species through H-atom abstraction and addition to the double bond, and electron transfer, respectively (34, 36). H2O2, TCE, and HO2• were considered the major sinks for Cl2•-, based on their concentrations at given reaction times and rate constants: 1.4 × 105 M-1 s-1 (34), 1 × 107 M-1 s-1 (assumed), and 4.3 × 109 M-1 s-1 (34), respectively. Formic acid accumulates to approximately 4 × 10-4 M, but the rate constant for the Cl2•reaction is small [3 × 105 M-1 s-1 (36)], and thus, this reaction was disregarded in the model. The •OH concentration was estimated by the model as ∼4 × 10-12 M at ∼3.5 min reaction time and increased by 1 order of magnitude by the end of the experiment. The rate constant for the reaction between •OH and Cl- is pH-dependent, and in acid solution it cannot be ignored. It leads to rapid formation of Cl2•- through a complex mechanism, which can be simplified as follows: •

OH + Cl- f HOCl•k ) (4.3 ( 0.4) × 109 M-1 s-1 (44) (47) HOCl•- f •OH + Cl-

k ) (6.1 ( 0.8) × 109 s-1 (44) (48)

HOCl•- + H+ f Cl• + H2O

k ) (2.1 ( 0.7) × 1010 M-1 s-1 (44) (49)

Cl• + Cl- f Cl2•-

k ) 8.6 × 109 M-1 s-1 (35) (30)

Cl2•- f Cl• + Cl- k ) 5.3 × 104 s-1 (35)

(50)

In the model, a set of consecutive steps were simulated as a single reaction only when the rate-limiting step was clearly identified and the corresponding rate constant was known or could have been reasonably approximated; otherwise, the model used the individual reactions steps, with the rate constants taken from the literature (see the Supporting Information). Modeling Approach. The kinetic equations used in the model were derived using the governing equation in a completely mixed batch reactor (CMBR)

dCA ) rA, dt

CA|t)0 ) CA0

(51)

where CAo and CA are the concentrations of species A at time 0 and t, respectively, and rA is the overall kinetic rate expression for species A in the reaction system. The full set of kinetic equations was obtained by substituting the rate expressions for all main species in the TCE-UV/H2O2 system and is provided in the Supporting Information. Twelve kinetic parameters were evaluated by fitting the experimental data obtained for TCE, H2O2, MCA, DCA, formic and oxalic acids, and pH using a genetic algorithm developed by Carroll (45). A range of the possible minima and maxima for each rate constant was provided to the genetic algorithm. This range was determined based on the literature reported values for each type of reaction, such that the resulting 12 fitting parameters exhibit reasonable values. For example, the rate constant for O2 addition to the carbon centered radicals was set within the range of 5 × 108-1 × 1010 M-1 s-1 (31) and that for β-scission of alkoxyl radicals from 1 × 104 to 1 × 107 s-1 (31). The detailed analysis of the algorithm is available in the literature (46) and is briefly covered in the Supporting Information. TCE Decay Modeling. Figure 1 shows the comparison between the predicted and experimental patterns of TCE and its byproducts. The model accounts for the four photolysis pathways and the reactions with •OH, Cl•, and Cl2•- radicals. TCE decay exhibits two kinetic regimes. Over the first 2.5 min it can be described by both zero- and firstorder kinetics, due to the contribution of the direct photolysis to the decay rate. After the fraction of light absorbed by TCE decreased by ∼35%, a first-order kinetics is obeyed over ∼1.6 orders of magnitude decay of TCE concentration, with a rate constant of (9.4 ( 0.6) × 10-3 s-1. The contribution of each major pathway to TCE decay was calculated using the concentration profiles of •OH, Cl•, and Cl2•- and the time step TCE photolysis rate (Supporting Information). The calculation shows that over the 12 min exposure time where TCE decays by ∼2.5 orders of magnitude, the integrated relative contribution of each pathway to the overall TCE degradation is 87.4%, 8.5%, 3.7%, and 0.4%, for •OH, UV photolysis, Cl2•-, and Cl•, respectively. The actual time profiles of such contributions are given in Figure 4. Given the complexity of radical reactions in this system and the lack of kinetic data for TCE reactions with chlorine and dichlorine radicals, we believe that the model predicts relatively well the experimental data of TCE decay. The bigger discrepancies lie in the lower oxygen level zone (4-5 min), where the mechanism could be more complex than described above and, thus, not accounted for by the model. H2O2 Decay Modeling. The degradation of H2O2 follows zero-order kinetics, with an experimental rate constant of 5.47 × 10-6 M s-1, in good agreement with the model predicted data (5.73 × 10-6 M s-1). The model predicts well the H2O2 decay. Intermediates Modeling. The discrepancies between the model-calculated patterns of DCA, MCA, and formic acid and the experimental data are less than 10%, 2%, and 7%, respectively, and could be attributed to both experimental analytical errors and model imperfections. pH Modeling. Simulation of pH pattern was based on the charge balance as described elsewhere (1). In the pH calculations the model calculates the inorganic carbon balance at each reaction time assuming that all inorganic carbon exists either as H2CO3* or HCO3-. The concentration of bicarbonate is then calculated according to the pKa of H2CO3*. Carbonate ion was ignored due to the low experimental pH data in the system. Since almost all organic and inorganic acids were considered in the model, the predicted pH pattern agrees well with the experimental data. Negatively charged radicals, i.e., Cl2•-, CO2•-, and CO3•-, were also included in the charge balance. VOL. 41, NO. 5, 2007 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 4. Model predicted relative contributions of the various pathways to the overall decay of TCE (12 min integrated time). O2 Modeling. The O2 pattern was simulated by calculating the O2-consuming and generating rate for all oxygeninvolving reactions included in the model (Table S1 in the Supporting Information). For example, the overall decay of formic acid, DCA, MCA, and glyoxylic acid, is associated in the model with the O2 loss rate of 0.5 mol per each mol of compound, based on the stoichiometric O2 required for peroxyl radical formation (1 mol) (e.g., reaction 37) and the amount of O2 released through bimolecular termination of the corresponding peroxyl radical (0.5 mol) (e.g., reaction 38). However, due to the complexity of the peroxyl radical reactions, the actual O2 recovery from peroxyl radical termination could be less than 0.5 mol per mol peroxyl radical formed. The predicted data follows the experimental data for the first 10 min reaction time; after 10 min, the model does not predict the measured O2 data probably due to the following reasons or the their combination: not all peroxyl radical reactions were included in the model, yet they should be dominant primarily within the 10 min reaction time; the mass transfer between the atmosphere (reactor headspace) and the solution was not considered. To our knowledge, no other similar study on UV/H2O2 applications attempted to model the experimental O2 pattern, and we believe that the experimental and predicted O2 concentration time profiles in this work agree reasonably well.

Acknowledgments The authors are very grateful to the reviewers of this manuscript for their thoughtful comments. The modeling work was supported partially by the Center for Clean Industrial and Treatment Technologies (CenCITT) sponsored by the U.S. EPA at Michigan Technological University. Partial support was provided by the Arizona State University Foundation and the Richard Snell Presidential Chair Funds.

Supporting Information Available Calculation of UV photolysis rate, brief description of the genetic algorithm, full set of the reactions and kinetic equations used in the model, and the calculation of the contribution factors of different TCE decay pathways as well as the radical concentration time-profiles (0-30 min). This material is available free of charge via the Internet at http:// pubs.acs.org. 1702

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Received for review March 30, 2006. Revised manuscript received October 16, 2006. Accepted November 2, 2006. ES0607638

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