Efficient Reductive Dechlorination of Monochloroacetic Acid by Sulfite

Jun 8, 2012 - School of Civil Engineering, Harbin Engineering University, 145 Nantong Street, Harbin 150001, P.R. China. •S Supporting Information...
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Efficient Reductive Dechlorination of Monochloroacetic Acid by Sulfite/UV Process Xuchun Li,† Jun Ma,*,†,‡ Guifang Liu,§ Jingyun Fang,† Siyang Yue,† Yinghong Guan,† Liwei Chen,† and Xiaowei Liu† †

State Key Laboratory of Urban Water Resource and Environment, Harbin Institute of Technology, Harbin 150090, P.R. China National Engineering Research Center of Urban Water Resources, Harbin 150090, P.R. China § School of Civil Engineering, Harbin Engineering University, 145 Nantong Street, Harbin 150001, P.R. China ‡

S Supporting Information *

ABSTRACT: Most halogenated organic compounds (HOCs) are toxic and persistent, and their efficient destruction is currently a challenge. Here, we proposed a sulfite/UV (253.7 nm) process to eliminate HOCs. Monochloroacetic acid (MCAA) was selected as the target compound and was degraded rapidly in the sulfite/UV process. The degradation kinetics were accelerated proportionally to the increased sulfite concentration, while the significant enhancement by increasing pH only occurred in a pH range of 6.0−8.7. The degradation proceeded via a reductive dechlorination mechanism induced by hydrated electron (eaq−), and complete dechlorination was readily achieved with almost all the chlorine atoms in MCAA released as chloride ions. Mass balance (C and Cl) studies showed that acetate, succinate, sulfoacetate, and chloride ions were the major products, and a degradation pathway was proposed. The dual roles of pH were not only to regulate the S(IV) species distribution but also to control the interconversion between eaq− and H•. Effective quantum efficiency (Φ) for the formation of eaq− in the process was determined to be 0.116 ± 0.002 mol/einstein. The present study may provide a promising alternative for complete dehalogenation of most HOCs and reductive detoxification of numerous toxicants.

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INTRODUCTION An increasing number of halogenated organic compounds (HOCs) are environmentally widespread and routinely found in water, air, sediments, and soils, and they are of great concern for their bioaccumulation, persistence, and toxicity.1−4 They occupy a large part of emerging persistent organic pollutants (POPs) and toxic compounds,3 e.g., polychlorinated biphenyls (PCBs), polybrominated diphenyl ethers (PBDEs) extensively used as brominated flame-retardants (BFRs), and disinfection byproduct (DBPs) commonly generated during the widely applied chlorine disinfection process.2,4 From a toxicological point of view, HOCs may enter the food chain from environmental media and will threaten the health of both humans and wildlife. Under oxidized conditions, HOCs likely possess low reactivity due to the electron-withdrawing C−X (X = F, Cl, Br, I) groups in the molecular structure.3,5 Cleavage of the C− X bond (dehalogenation), which likely causes the toxicity and persistence of HOCs, seems difficult because of its high bond energy, which follows the order of C−F > C−Cl > C−Br > C− I.5 Various technologies have been developed to destroy HOCs, such as biodegradation,6 photochemical treatment,7 zerovalent metals (ZVMs),5,8,9 electrochemical methods,10 and radiation.11 However, most current technologies indeed show distinct performance in removal of various HOCs, depending on the reactivity of HOCs with radicals (e.g., HO•),12 active sites on ZVMs,5,8,9 enzymes,6 etc. Significant accumulation of halogen© 2012 American Chemical Society

containing intermediates during the treatment of HOCs by most current technologies (e.g., AOPs and ZVMs)9,10,13,14 will probably cause the risk of toxicity.15 Furthermore, nanomaterials or noble metal catalysts are usually used to achieve high efficiency,5,10 resulting in a potential risk of ecotoxicity and increased cost. Hydrated electron (eaq−) related techniques may be alternative methods for degradation and detoxification of HOCs. As one of the most reactive species with the standard reduction potential of about −2.9 V,12,16 eaq− can react rapidly with most HOCs even at diffusion-controlled rates,12 leading to rapid breakage of the C−X bond and release of halide ion. For example, radiolysis could completely dechlorinate PCBs in industrial transformer oil through generating eaq−, but the application of radiolytic techniques may be limited because of the cost of equipment and operation.11 It has been reported that reactive species (e.g., eaq−, H•) can be generated at attractive quantum efficiencies during the photolysis of S(IV) (sulfite and hydrogen sulfite),17,18 as shown in the following reactions: SO32 − + hν → SO3•− + eaq − Received: Revised: Accepted: Published: 7342

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March 3, 2012 May 30, 2012 June 8, 2012 June 8, 2012 dx.doi.org/10.1021/es3008535 | Environ. Sci. Technol. 2012, 46, 7342−7349

Environmental Science & Technology HSO3− + hν → SO3•− + H•

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and organic acids26 were analyzed by an ion chromatograph (Dionex ICS-3000), equipped with a SP gradient pump, an anion self-regenerating suppressor (ASRS 300, 4 mm), a conductive detector, and an AS40 autosampler. A Dionex AS19 column (4 × 250 mm) and a Dionex AG19 guard column (4 × 50 mm) were used to separate the organic acids, NO2−, NO3−, Cl−, and S2O62‑, with 30 mM KOH as isocratic eluent at 1.0 mL min−1 and a suppressor current of 75 mA. A Dionex AS11-HC column (4 × 250 mm) and a Dionex AG11-HC guard column (4 × 250 mm) were selected for separation of SO32‑ and SO42‑ with 20 mM KOH as isocratic eluent at 1.5 mL min−1 and a suppressor current of 75 mA. The injection volume was 100 μL with a 500 μL sample loop, and the detection limits were about 50 nM for all anions. As the depletion of monochloroacetic acid (MCAA) in sulfite solution was negligible (less than 6% depletion in 72 h with 1.0 mM sulfite at 25 °C), MCAA samples were analyzed (always within 24 h) without quenching of residual SO32‑ in the samples unless otherwise noted. SO32‑ was stabilized by addition of sufficient formaldehyde to form hydroxymethanesulfonate,17 and samples were analyzed within 2 h. Low-level (less than 2 μM) MCAA was determined by GC/ECD (Aglient GC 6890N) according to EPA Method 552.2.27

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Ultraviolet (UV) irradiation is a recently established disinfection technology widely used in water treatment,19 and UV-assisted processes for pollution remediation have also received increasing attention. Additionally, S(IV) has also been widely used as dechlorination agents20 in chlorinated wastewater and as preservatives in the food industry. Considering the rapid reactions induced by eaq−,12 the sulfite/UV process was thus proposed as a method to remove HOCs. To the best of our knowledge, there has been no report on pollution remediation by the combination of S(IV) and UV. The objective of the present study was to investigate the dechlorination process in the sulfite/UV process and specifically to focus on (i) dechlorination efficiency, (ii) the roles of S(IV) and pH, (iii) the dechlorination mechanism, and (iv) the effective quantum efficiency of the process. Monochloroacetic acid (MCAA) was selected as the target compound for its comparative recalcitrance to most current abiotic treatment processes7,8,12,14 and relatively simple structure.

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MATERIALS AND METHODS Materials. All chemicals were obtained from commercial sources and are listed in the Supporting Information (Text S1). All reagents were used as received without further purification. Solutions were prepared in 18.2 MΩ cm Milli-Q water (Millipore). Stock solution of sodium sulfite was prepared freshly at a concentration of 1 M with deaerated (N2) Milli-Q water everyday. Experimental Procedure. Photolysis experiments were conducted in a sealed cylindrical borosilicate glass reactor (V = 800 mL), and a 10 W low-pressure mercury UV lamp (253.7 nm, ozone-free, type GPH212T5L, Heraeus Noblelight) was used as the UV light source (Supporting Information, Figure S1). The photon flux (I0, 253.7 nm, amount basis) entering the solution from the UV source was determined to be (2.98 ± 0.02) × 10−6 einstein s−1 by the KI/KIO3 method,21 and the average fluence rate (Is) was estimated to be about 1.33 × 10−8 einstein s−1 cm−2, or 6.27 mW cm−2. Since the UV light in the reactor was not a strictly collimated beam, the fluence rate distribution would depend on reactor geometry and physics (e.g., reflection, refraction). So the effective path length (L)22,23 was selected to characterize the photoreactor and was obtained to be 3.57 ± 0.02 cm from the photolysis kinetics of dilute H2O2.23,24 Temperature was controlled at 25 ± 0.5 °C. Phosphate buffer (10 mM, for pH 5.0−8.0) and borate buffer (10 mM, for pH 8.0−10.5) were used to control the pH, and negligible absorption of the buffers was observed at 253.7 nm. Oxygenfree conditions were achieved by purging with N2 gas. Samples were withdrawn at predetermined time intervals. All experiments were repeated independently at least two times, and average values along with one standard deviation (±SD) are presented. More details concerning experimental procedure are provided in the Supporting Information (Text S2). Analytical Methods. Solution pH was measured by a pH meter (UB-7, Denver Instrument), and the calibration using standard buffers at pH 4.00, 6.86, and 9.18 was made routinely before measurements at room temperature. The concentrations of H2O2 were determined by the I3− method.25 UV−vis absorbance and absorption spectrum were carried out on a Cary 300 UV−vis spectrometer (Varian). Inorganic anions17

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RESULTS AND DISCUSSION Degradation Efficiency of MCAA in the Sulfite/UV Process. Figure 1 shows the degradation of MCAA in the

Figure 1. Degradation of MCAA in the sulfite/UV process. Conditions: [S(IV)]0 = 1.0 mM (no addition for UV only), [MCAA]0 = 45 μM or 2 μM, pH = 9.2, 25 °C, I0/V = 3.73 × 10−6 einstein L−1 s−1, and oxygen-free. Error bars represent the standard deviation from at least duplicate experiments.

sulfite/UV process. Less than 1% of MCAA was degraded within 15 min (UV dose of about 5643 mJ cm−2) under the UV irradiation (I0/V) of 3.73 × 10−6 einstein L−1 s−1 (or average fluence rate of about 6.27 mW cm−2). The negligible direct photolysis of MCAA was interpreted by its weak UV absorption7 (Supporting Information, Text S3 and Figure S4b). Surprisingly, the sulfite/UV process showed a high efficiency in degradation of MCAA. About 45 μM of MCAA underwent almost complete degradation in 15 min (UV dose of about 5643 mJ cm−2), while complete decay of 2 μM of MCAA required only 3 min (UV dose of about 1129 mJ cm−2). Figure 1 also indicates that the degradation rate (r, μM min−1) of MCAA was dependent upon its concentration. However, r was maintained relatively constant for high concentrations of MCAA (C0 = 45 μM), and the r was 3.29 ± 0.04 μM min−1 (corresponding to UV fluence-based rate of about 8.75 nM cm2 mJ−1, 253.7 nm) with correlation coefficient (R2) higher than 7343

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degraded at a rate of about 3.05 μM min−1 (about 8.11 nM cm2 mJ−1) under the conditions of pH 9.2, S(IV) concentration of 1.0 mM, and the average fluence rate of about 6.27 mW cm−2. It should be noted that S(IV) was present primarily (more than 99% in molar ratio) in the form of SO32‑ at pH 9.2 (Supporting Information, Figure S4a). Compared with SO32‑, the contribution of HSO3− was ruled out for its low fraction (less than 1% of S(IV) in molar ratio at pH 9.2) and much weaker absorption17,33 (Supporting Information, Figure S4b) as well as less photoactivity.17,33 It is further supported by the similar linear relationship observed between r and the UV absorption by SO32‑ (Supporting Information, Text S4 and Figure S5). Therefore, SO32‑ intrinsically governed the degradation of MCAA in the process. In addition, photolysis of SO32‑ was rather slow with about 17% depletion in 60 min (Supporting Information, Figure S6). According to Figure 2b, no significantly negative impact of SO32‑ (in the range of 0.0−2.0 mM) on the degradation of MCAA was observed, and the linear dependence (Figure 2b) implies that a high amount of SO32‑ can be applied. Role of pH in the Degradation of MCAA by the Sulfite/UV Process. Solution pH determines the distribution ratio of SO32‑ (Supporting Information, Figure S4a), and it may thus affect the degradation efficiency of MCAA in the process. Figure 3a shows the effect of pH on the degradation of MCAA. The degradation was promoted obviously with increasing pH from 6.0 to 8.7 but only slightly enhanced by further elevation

0.999, under the conditions of pH 9.2, S(IV) concentration of 1.0 mM, and the average fluence rate of about 6.27 mW cm−2. The degradation appeared to be much faster than that by H2O2/UV,7 biodegradation,28 and the Fe0,8 and Pd/Fe bimetallic system,29 which usually took several hours or even days. Though some DBPs underwent base-catalyzed decomposition with sulfite,30 no significant degradation of MCAA (less than 6% within 72 h) was observed by 1.0 mM S(IV) at pH 9.2 and 25 °C. It should be noted that dissolved oxygen played complex roles17,31 in the process, and the degradation efficiency of MCAA was obtained under oxygen-free conditions. Role of Sulfite in the Degradation of MCAA by the Sulfite/UV Process. S(IV) here exists in the forms of sulfite (SO32‑), hydrogen sulfite (HSO3−), and sulfurous acid (H2O·SO2) with pKa1 = 1.76 and pKa2 = 7.20,32 and SO32‑ possesses the strongest UV absorption (Supporting Information, Text S3 and Figure S4). In order to study the role of S(IV), the effect of concentration of S(IV) on MCAA degradation was investigated. Figure 2a shows that the

Figure 2. (a) Effect of S(IV) concentration on the degradation of MCAA in the sulfite/UV process. (b) The degradation rates (r) as a function of S(IV) concentration. Conditions: [MCAA]0 = 50 μM, pH = 9.2, 25 °C, I0/V = 3.73 × 10−6 einstein L−1 s−1, and oxygen-free. The solid line indicates the best linear fit (b), and error bars represent the standard deviation from duplicate experiments (a and b). Note that more than 99% (molar ratio) of S(IV) is present in the form of SO32‑ at pH 9.2. Figure 3. (a) Effect of pH on the degradation of MCAA in the sulfite/ UV process. (b) The degradation rates (r) versus pH. Conditions: [MCAA]0 = 50 μM, or [NO2−]0 = 100 μM (b), 25 °C, [S(IV)]0 = 1.0 mM, I0/V = 3.73 × 10−6 einstein L−1 s−1, and oxygen-free. The solid line part b is the predicted pH-dependent r of MCAA or nitrite, which is simply based on a linear relationship (Figure 2b) and the pHdependent distribution of SO32‑ (Supporting Information, Figure S4a). Error bars represent the standard deviation from duplicate experiments.

degradation of MCAA was enhanced obviously by the increased concentration of S(IV). Figure 2b further presents the relationship between the degradation rates (r, R2 > 0.99) of MCAA and concentration of S(IV). The r showed a positive linear correlation with the concentration of S(IV), and gave r/ [S(IV)]0 of (3.05 ± 0.11) × 10−3 min−1 (Is of about 6.27 mW cm−2). The linear dependence indicates that MCAA can be 7344

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NO3− underwent negligible direct photolysis due to their weak absorption and low quantum yields at 253.7 nm.35 The results can also reveal which reactive species were dominant and responsible for the degradation. The nearly complete inhibition of MCAA degradation by N2O, NO2−, and NO3− strongly indicates that eaq− and/or H• caused the degradation. The H• reacts slowly with MCAA (3.6 × 106 M−1 s−1)12 but about 200 times faster with NO2− (eq 6), so the slight difference between degradation rates of MCAA and NO2− at pH 9.2 (Figure 3b) strongly indicates a minor contribution of H•. According to the rapid reaction eq 6, the observation that much NO2− was accumulated when NO3− was used as the competitor (Supporting Information, Figure S8) also confirms much less formation of H• than that of eaq− at pH 9.2. Furthermore, although the H• scavenging activity of NO2− (eq 6) is about 500 times higher than that of NO3− (eq 8), the same amount of NO2− led to less inhibition of MCAA degradation than NO3− (Figure 4), so it supports the hypothesis of little formation of H•. In fact, considering about 200 times faster reaction kinetics of H• with NO2− (eq 6) than that of H• with MCAA, only about 49% depletion of NO2− accompanied by 18% of MCAA in 15 min also clearly demonstrates the low level of H•. Consequently, H• played a negligible role in the degradation of MCAA at pH 9.2 due to its slow reaction with MCAA and little formation. Sulfite radical (SO3•‑) is also produced simultaneously in eqs 1 and 2, and it is a mild oxidant but poor reductant.36 Under the N2O-saturated conditions, the formation of SO3•‑ was approximately doubled (Supporting Information, Text S7), but it caused little degradation of MCAA (Figure 4), so it is rational to regard that SO3•‑ made no contribution to MCAA degradation. Therefore, it is highly probable that eaq− was the dominant reactive species that accounted for MCAA degradation (eq 9).

of pH. The utmost removal efficiency was likely obtained at pH 10.1, which was supported by the distribution of SO32‑ versus pH (Supporting Information, Figure S4a). However, the distribution of SO32‑ failed to fully interpret the degradation of MCAA at circumneutral pH. For example, according to the linear dependence (Figure 2b) and fraction of SO 3 2‑ (Supporting Information, Figure S4a), the degradation efficiency of MCAA at pH 7.1 was expected to reach about half of that obtained at pH 10.1, but the actual efficiency was much lower (ca. 22%). Further study on the relationship between the degradation rates (r, R2 > 0.99) of MCAA and pH was conducted. Figure 3b shows the dependence of r on pH, and r increased sharply in a narrow pH range from 6.0 to 8.7. The predicted curve, which was simply based on the contribution of SO32‑ according to the linear relationship in Figure 2b and the distribution of SO32‑ in Figure S4a (Supporting Information), overestimated the degradation efficiency of MCAA, especially under the near-neutral conditions. At pH 7.1, the predicted r of 1.45 μM min−1 was about two times higher than the actual r of 0.69 μM min−1. In contrast, there was a good agreement between the degradation kinetics of nitrite (Figure 3b and Figure S7, Supporting Information) and the predicted curve, suggesting significant and complex roles of pH in the system. Degradation Mechanism of MCAA in the Sulfite/UV Process. It was reported that some reactive species (eqs 1 and 2), e.g., eaq− and H•, could be produced during the photolysis of S(IV).17,18,33 To explore the degradation mechanism of MCAA in the sulfite/UV process, NO2−, NO3−, and N2O were selected as quenching compounds to investigate their respective effects on inhibition of MCAA degradation. As known, NO2−,12,34 NO3−,12,34 and N2O12 are good quenchers for eaq− (eqs 3−5), while NO2− and a high level of N2O are good for H• (eqs 6−8).12,34 Figure 4 demonstrates the negative effects of NO2−,

NO2− + eaq − → (NO2•)2 −

(k 3 = 4.1 × 109 M−1 s−1)12 (3)

NO3− + eaq − → (NO3•)2 −

(k4 = 9.7 × 109 M−1 s−1)12 (4)

−

•

N2O + eaq → N2 + HO + OH

−

(k5 = 9.1 × 109 M−1 s−1)12

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NO2− + H• → NO• + OH− (k6 = 7.1 × 108 M−1 s−1)12 H• + N2O → N2 + HO•

Figure 4. Inhibition effect of electron scavengers on the degradation of MCAA in the sulfite/UV process. Conditions: [MCAA]0 = 47 μM, [S(IV)]0 = 1.0 mM, [NO2−]0 = 92 μM, [NO3−]0 = 92 μM, N2O saturation (about 25 mM), pH = 9.2, 25 °C, I0/V = 3.73 × 10−6 einstein L−1 s−1, and oxygen-free. Error bars represent the standard deviation from duplicate experiments.

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(k 7 = 2.1 × 106 M−1 s−1)12 (7)

NO3− + H• → (NO3H•)−

(k 8 = 1.4 × 106 M−1 s−1)12 (8)

CH 2ClCOO− + eaq − → •CH 2COO− + Cl− (k 9 = 1.0 × 109 M−1 s−1)12

NO3−, and N2O on the degradation of MCAA. The presence of NO3− or N2O almost completely inhibited the degradation of MCAA, while the inhibition by NO2− was slightly less. Though equal amounts of NO2− and NO3− were applied, the greater inhibition by NO3− was observed, which was probably attributed to significant formation of NO2− (Supporting Information, Figure S8). It should be noted that NO2− and

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To further study the degradation pathway of MCAA in the sulfite/UV process, identification and quantitative determination of major products were also carried out. Figure 5 shows the formation of various products and the total masses of Cl and C during the degradation of MCAA. Acetic acid (HAc), succinic acid (SA), sulfoacetic acid (SAA), chloride ion (Cl−), and dithionate ion (S2O62‑) were concurrently generated during the 7345

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that of Cl− (via eq 9) suggests that most sulfur-containing products were identified and quantified. The minor variance (ca. 13%) was possibly caused by the discrepancies in formation kinetics17,33,37 of S2O62‑ and SO42‑. Consequently, SO3•‑ mostly underwent radical recombination reactions17,33,37 to produce SO32‑, SO42‑, S2O62‑, and SAA, further confirming its little contribution to the dechlorination of MCAA. Therefore, eaq− fully accounted for the degradation of MCAA in the sulfite/UV process. The transformation pathway of MCAA was proposed in Scheme 1. As one of the most reactive Scheme 1. Proposed Pathways for MCAA Degradation and Products Formation in the Sulfite/UV Process Figure 5. Formation of major products and mass balances during degradation of MCAA in the sulfite/UV process. Total carbon (TC) and total chlorine (TCl) were calculated as follows (molar concentration): TC = 2 × [MCAA] + 2 × [HAc] + 4 × [SA] + 2 × [SAA] and TCl = [MCAA] + [Cl−], respectively. Conditions: [MCAA]0 = 110 μM, [S(IV)]0 = 1.0 mM, pH = 9.2, 25 °C, I0/V = 3.73 × 10−6 einstein L−1 s−1, and oxygen-free. Error bars represent the standard deviation from duplicate experiments.

degradation of MCAA. The mass balances of carbon (TC, 97− 104%) and chlorine (TCl, 93−101%) demonstrate that almost all of the C- and Cl-containing products were identified and quantified. They also support that the role of H• was negligible in the degradation of MCAA, otherwise more other unidentified Cl-containing intermediates would be formed due to the lack of contribution to MCAA dechlorination by H• involved in hydrogen abstraction reactions.12 The mass balance of Cl, i.e., the concerted process between MCAA degradation and Cl− release, reveals that the degradation of MCAA was a reductive dechlorination process induced by eaq− and also indicates that the complete dechlorination could be readily achieved in the sulfite/UV process. Table 1 presents the formation rates of the identified products. The degradation of MCAA, release of Cl−, and the

species,12,16 eaq− is always engaged in reductive reactions, even at diffusion-controlled rates.12 During the reactions between HOCs and eaq−, the electron attachment and halide formation may proceed via two mechanisms:38 (i) one is the direct dissociation to generate halide for halogenated aliphatic compounds, (ii) while the other undergoes the formation of a stable negative ion of a molecule with subsequent activation and dissociation into halide for halogenated aromatic derivatives. So the dechlorination of MCAA by eaq− proceeded via the former mechanism to release Cl− directly. Simultaneously formed with Cl−, •CH2COO− was subsequently transformed into three different organic products, namely HAc, SA, and SAA. Without involvement in the dechlorination of MCAA,12 H• was steadily produced (eqs 10−12) during scavenging of eaq− by H+,12 H2O,12 and HSO3−.33 The decrease in pH led to the decreased fraction of SO32‑ and the increased formation of H• and consequently caused the significant influence of pH on MCAA degradation. The high formation of H• was responsible for the significant discrepancy between actual degradation kinetics of MCAA and the predicted curve (i.e., prediction for formation of eaq− in eq 1) and also the much faster degradation kinetics of nitrite than that of MCAA at circumneutral pH. According to Figure 3b, about 52% (calculated from difference in degradation rates of MCAA and nitrite) of eaq− was converted to H• at pH 7.1, while the total formation of H• was low under alkaline conditions (higher than 8.7) due to low level of H+ and contribution of eq 13. It is also backed up by the minor difference in degradation rates of MCAA and nitrite at pH 9.2 and the confirmation between actual degradation rates of MCAA and predicted curve at pH higher than 8.7.

Table 1. Rates (r) of MCAA Degradation and Major Products Formation in the Sulfite/UV Processa MCAA/products MCAA chloride ion (Cl−) acetic acid (HAc) succinic acid (SA) sulfoacetic acid (SAA) dithionate (S2O62‑) TCPsc TSPsd

r (μM min−1)b

R2

± ± ± ± ± ± ± ±

0.9982 0.9999 0.9982 0.9993 0.9977 0.9965 nae na

3.7560 3.5166 1.6060 0.5890 1.0778 0.3778 7.7236 3.0721

0.0685 0.0112 0.0239 0.0056 0.0181 0.0080 0.1064 0.0603

Conditions: [MCAA]0 = 110 μM, [S(IV)]0 = 1.0 mM, I0/V = 3.73 × 10−6 einstein L−1 s−1, pH = 9.2, 25 °C, and oxygen-free. bError is given as the standard deviation. cTotal carbon-containing products (μM C), including HAc, SA, and SAA. dTotal sulfur-containing products (μM S), including SO32‑, SO42‑, S2O62‑, and SAA; formation of SO32‑ and SO42‑ was calculated according to the kinetics in ref 17. ena, not available. a

eaq − + H 2O → H• + OH−

(k10 = 1.9 × 101 M−1 s−1)12 (10)

formation of carbon-containing products agreed with each other, indicating the similar mass balances throughout the course of the MCAA degradation as shown in Figure 5. By assuming that the formation of all the sulfur-containing products was attributed to SO3•‑ from eq 1,17,33,37 the coincidently slight deviation of their formation rates from

eaq − + H+ → H•

(k11 = 2.3 × 1010 M−1 s−1)12

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eaq − + HSO3− → H• + SO3•− (k12 = 2 × 107 M−1 s−1)33 7346

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halogenation degree (more halogen atoms substituted).12 So it is rational to expect efficient and complete dehalogenation of most HOCs (e.g., PCBs, PBDEs, and DBPs) and consequently great decrease in toxicity4 as well as increase in biodegradability42 in the sulfite/UV process without the requirement of expensive catalysts. Because fluorine (F) has no low-lying vacant d orbital to accept an electron, the reactivity of fluorinated chemicals (FCs) with e aq−, 43 especially of perfluorochemicals (PFCs), is much lower than that of other HOCs. The process may show finite efficiency in FCs elimination, though it is still an attractive candidate method for the destruction of FCs.44 Considering eaq− involved reaction kinetics,12 the process will also be efficient in reductive detoxification of many toxicants (e.g., bromate) and chemical denitrification of nitrate, nitrite, and NOx with the advantages of negligible formation of ammonium and little release of greenhouse gas N2O.34 A large amount of toxic SO2 in exhaust gas or air45,46 always needs further costly and complex desulfurization treatment,39 though the sulfite/UV process can make full use of SO2 to remove pollutants. However, the efficient and nontoxic removal of residual S(IV) is also necessary. Figure S12 (Supporting Information) shows that it could be removed rapidly by oxygen with an apparent quantum efficiency of about 117 under UV irradiation, and nontoxic sulfate was the final product. Furthermore, the chain photooxidation17,33 simultaneously depletes dissolved oxygen rapidly and will minimize the negative effects8,11,29 of oxygen on the dehalogenation treatment (Supporting Information, Text S9). The sulfite/UV process still faces some important challenges in practical applications. Considering the well-known reductive chemistry of S(IV), the efficiency will be reduced primarily due to consumption of S(IV) by some oxidants (e.g., ozone, chlorine). When applied in drinking water treatment, the sulfite/UV process should be followed by ozonation or chlorine disinfection process, and total removal of residual S(IV) is necessary to minimize the disinfectant demand and to ensure the disinfectant residual. The dehalogenation efficiency will also be influenced by the presence of electron scavengers (e.g., nitrate, nitrite), and thus it will be energy-intensive (additional UV dose required to reduce the scavengers) to apply the process. When the levels of nitrite and nitrite are high enough, their efficient removal (e.g., ZVMs,47,48 microbial denitrification,49 and adsorption50) may be required prior to dehalogenation treatment by the sulfite/UV process. Additionally, because the reduction efficiency of the process depends on photons absorbed by sulfite, it can also be affected due to light attenuation by light absorbers (e.g., natural organic matter) and light scatterers (e.g., particles). The efficiency of the sulfite/UV process is dependent upon UV fluence rate, hydraulics, water matrix, etc., so more detailed study must be performed prior to any recommendation for its practical applications. Investigations concerning the toxicity assessment, degradation mechanisms, and UV fluence-based kinetic modeling of some HOCs in the process are underway, and further research on approaches to optimize and enhance the process is also being carried out.

(k13 = 2.2 × 107 M−1 s−1)12 (13)

Effective Quantum Efficiency of the Sulfite/UV Process. The relatively constant degradation rate of MCAA (C0 = 45 μM) in Figure 1 suggests that MCAA consumed almost all the generated eaq− at pH 9.2. Under strong alkaline conditions, the influence of side reactions that compete for eaq− can be neglected,12 and it is consequently rational to employ the degradation kinetics of MCAA to characterize the production efficiency of eaq− and the utmost reductive capacity in the sulfite/UV process. The effective quantum efficiency (Φ, mol/einstein)35 of the sulfite/UV process is defined as the moles of eaq− formed (eq 1) divided by the moles of photons absorbed by sulfite and can be calculated using eq 14 (Supporting Information, Text S5) Φ=

r0V I0(1 − 10−εCL)

(14) −

−1

where r0 is the formation rate of eaq in eq 1 (M s ), V is solution volume (L), I0 is the photon flux entering the solution (einstein s−1), ε is the molar absorption coefficient of SO32‑ (M−1 cm−1), L is the effective path length (cm), and C is the concentration of SO32‑ (M). Assuming that all eaq− participated in eq 9 under strong alkaline conditions, Φ1 = 0.116 ± 0.002 mol/einstein (253.7 nm) was obtained from the degradation rate of MCAA at pH 10.1 (Figure 3b), which was consistent with Φ2 = 0.116 ± 0.002 mol/einstein from degradation kinetics of nitrite at pH 9.2 (Figure 3b and Supporting Information Text S6) and Φ3 = 0.118 ± 0.003 mol/einstein from the depletion kinetics of MCAA and nitrite (Supporting Information, Text S8). It should be noted that Φ2 and Φ3 were calculated with the assumptions of almost complete consumption12,34 of eaq− and H• by nitrite with or without MCAA. Though obtained from independent methods separately, Φ1, Φ2, and Φ3 could all intrinsically characterize the effective quantum efficiency of the process. Considering possible interference from intermediates during reduction of nitrite,34 it is more appropriate to select Φ1 as the effective quantum efficiency of the process. The Φ is indicative of diffusive eaq− that escapes from the “solvent cage”, freely diffuses, and can be freely available in bulk solution.18,35 The Φ (253.7 nm) here was consistent with 0.144 mol/einstein (253.7 nm) for hydrogen formation39 but differed considerably from 0.39 mol/einstein (253.7 nm) obtained from sulfite photodecomposition and products formation.17 Generally, higher excitation energy (shorter wavelength) leads to increased excess energy of photofragments, higher efficiency of direct ionization, and blocking of the alternative photoreaction routes and consequently increases prompt quantum yield, escape efficiency, and the effective quantum efficiency.18,35 So our results agreed with 0.108 mol/einstein at 248 nm,18 0.231 mol/einstein at 200 nm,18 and 0.391 mol/einstein at 193 nm18 for eaq− formation. Interestingly, the Φ of the sulfite/UV process for eaq− generation was much higher than that of the widely studied TiO2/UV process, that is, 0.057 mol/einstein for hole photogeneration40 and 0.024 mol/einstein for dichloroacetic acid (DCAA) depletion.41 Technical Implication. The eaq− involved reactions are likely controlled by the availabi1ity of a vacant orbital of the reagents into which the electron can transfer.12 For HOCs, the corresponding reaction kinetics likely follow the order of F ≪ Cl < Br < I substitution5,12 and also increase with higher

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ASSOCIATED CONTENT

S Supporting Information *

Figures S1−S13 and Texts S1−S9. This information is available free of charge via the Internet at http://pubs.acs.org. 7347

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AUTHOR INFORMATION

Corresponding Author

*Phone: +86-0451-86282292; fax: +86-0451-86283010; e-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors greatly thank Z. H. Liu for his assistance on the experiments and Q. D. Qin, J. Zhang, X. L. Shao, J. Zhao, J. J. Yang, P. C. Xie, and J. Zou for discussions. This research was supported by the Funds for Creative Research Groups of China (51121062), the Natural Science Foundation of China (50978067 and 51108117), 863 High Tech Scheme (2009AA06Z310), and the Science and Technology Ministry of China (2009ZX07424-005; 2009ZX07424-006). We thank the three anonymous reviewers for their critical comments and constructive suggestions that helped in improving the quality of this paper greatly.

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