Decomposition of Environmentally Persistent Trifluoroacetic Acid to

Table 1 lists the reaction yields under various combinations of the catalyst, light irradiation, and reaction atmosphere. In the absence of either lig...
0 downloads 0 Views 99KB Size
Download PDF
Environ. Sci. Technol. 2003, 37, 418-422

Decomposition of Environmentally Persistent Trifluoroacetic Acid to Fluoride Ions by a Homogeneous Photocatalyst in Water HISAO HORI,* YUKO TAKANO, KAZUHIDE KOIKE, KOJI TAKEUCHI, AND HISAHIRO EINAGA National Institute of Advanced Industrial Science and Technology (AIST), AIST Tsukuba West, 16-1 Onogawa, Tsukuba, Ibaraki 305-8569, Japan

Decomposition of trifluoroacetic acid (TFA) was achieved with a tungstic heteropolyacid photocatalyst H3PW12O40‚ 6H2O in order to develop a technique for measures against TFA stationary sources. This is the first example of C-F bond cleavage in an environmentally harmful perfluoromethylgroup-containing compound using a homogeneous photocatalyst. The catalytic reaction proceeds in water at room temperature under UV-visible light irradiation in the presence of oxygen. The system produces only F- ions and CO2; the (mole of formed F-)/(mole of decomposed TFA) and (mole of formed CO2)/(mole of decomposed TFA) ratios were 2.91 and 2.09, respectively. GC/MS measurements showed no trace of other species such as environmentally undesirable CF4, which is the most stable perfluorocarbon and has a very high global warming potential. When the (initial TFA)/(initial catalyst) molar ratio was 20:1, the turnover number of TFA decomposition reached 5.58 by 72 h of irradiation, accompanying with no catalyst degradation. The catalytic reaction mechanism can be explained by a redox reaction between the catalyst and TFA, involving a photo-Kolbe process.

Introduction Trifluoroacetic acid (TFA), which is harmful to aquatic organisms and may cause long-term adverse effects in the aquatic environment (1), has recently been detected in the environmental waters in the world (2-5). Hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs), alternatives for ozone-destroying chlorofluorocarbons (CFCs), are thought to be a source of TFA (6). However, the environmental levels of TFA cannot be explained by only the degradation of HCFCs and HFCs. Recently, it has been revealed that the commercial fluoropolymers such as poly(tetrafluoroethylene) (PTFE) produce TFA when heated and suggested that degradation of the polymers can be a significant source of TFA in rainwater around industrialized areas (7). The polymer manufacturing and processing plants may act as stationary sources of TFA. Although TFA has been shown to decompose under extreme anaerobic conditions (8), it does not degrade under normal environmental conditions (9-12). Hence, it is desirable to develop an artificial method for decomposing TFA to environmentally harmless species under mild condi* Corresponding author telephone: +81 298 61 8161; fax: +81 298 61 8258; e-mail: [email protected]. 418

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 2, 2003

tions. If the C-F bonds of TFA are cleaved to yield F- ions, the F- ions can readily combine with Ca2+ to form environmentally harmless CaF2. The heterogeneous photocatalyst TiO2 has been widely used for the decomposition of air and water pollutants (13-15). However, the reactivity of TiO2 toward TFA is estimated to be very low (16) because OH radicals in aqueous solution are poorly reactive with respect to TFA (17, 18). Water-soluble heteropolyacids are attractive candidates for photocatalysts for the decomposition of TFA because of their multielectron capabilities and high stability under highly acidic conditions (pH < 1) (19-21). Here we show the effective photocatalytic decomposition of TFA using a homogeneous system consisting of the tungstic heteropolyacid H3PW12O40‚6H2O (1), water, and oxygen. In this system, most of the fluorine and carbon atoms in the decomposed TFA molecules are transformed into F- and CO2, respectively. The overall decomposition reaction can be written as follows: 1, hν

CF3COOH + H2O + 1/2O2 98 2CO2 + 3H+ + 3F- (1)

Experimental Section Materials. Tungstic heteropolyacid 1 was obtained from Nippon Inorganic Chemical Co., Ltd. It was purified by ether extraction and recrystallized from water. Other reagents and solvents were of high purity commercially available from Kanto Chemical Co., Inc. and Wako Pure Chemical Industries Ltd. Photochemical Procedures. A cylindrical pressure resistant Inconel reactor (200-mL volume, 4.5 cm i.d.) equipped with a sapphire window on the top for introduction of light was used. The inner wall of the reactor was coated with PTFE. A small gold vessel (25 mL, 3.8 cm i.d.), which is stable to highly acidic solutions, was placed in the reactor and filled with 23 mL of an aqueous solution containing 1 (0.46 g, 1.54 × 10-4 mol; 6.70 mM) and a 1-20 molar excess of TFA. The pH of the solution was 0.8. After the reactor was purged and then pressurized to 0.55 MPa with oxygen gas, the solution was irradiated with UV-visible light from a high-pressure mercury lamp (500 W). For the light irradiation, a water filter and a liquid-type optical fiber were used. This combination filters out IR light and light with wavelengths below 260 nm. In the case of quantum yield measurements, a band-pass filter (313 nm) was also used to produce monochromatic light (light intensity ) 1.60 × 10-6 einstein min-1; 1 einstein ) 6.022 × 1023 photons). After irradiation, the pressure was released, and the gas was collected in a sampling bag (3 L volume) and subjected to GC/MS and GC. The liquid phase was also subjected to ion chromatography and ion-exclusion chromatography measurements to detect F- and TFA, respectively. Measurements. An ion-chromatograph system (Tosoh 8020) was used to measure the F- ion concentrations, where standard F- solutions containing 1 with the same concentrations as those of the reaction sample solutions were used. A Tosoh ion-exclusion chromatograph system consisting of a column (TSKgel OApak-A, 7.8 mm i.d., 30 cm), a pump (CCPS), and a conductivity detector (CM-8020) was used to measure the TFA concentrations, where the mobile phase was a mixture (80:20, v/v) of benzoic acid (2 mM) and MeOH. A Hewlett-Packard GC/MS system consisting of a gas chromatograph (HP5890) with a column (Chrompack, Poraplot Q, 0.32 mm i.d., 25 m length), a mass spectrometer (HP 5972A), and a workstation (HP G1034CJ) was used to identify the products in the gas phase. Carrier gas was He. 10.1021/es025783y CCC: $25.00

 2003 American Chemical Society Published on Web 12/11/2002

FIGURE 1. Wavelength distribution for the absorption of (A) TFA (67.4 mM in water), (B) catalyst 1 (6.70mM in water), and (C) the emission from the mercury lamp. The concentrations of 1 and TFA were the same as those in the following catalysis experiment (Figure 2). Path length for the absorption spectral measurements was 1.0 cm. Standard gases of CF3H, CF4, C2F6, and CO2 were used for qualitative and/or quantitative analysis. The concentrations of standard gases were 500-1000 ppm in Ar or N2. The CO2 amount was also measured by a Yanako G-3800 GC system using an active carbon column and a thermal conductivity detector.

Results and Discussion Photocatalysis. In the present reaction condition, the aqueous solution of 1 and TFA was irradiated with UVvisible light from a high-pressure mercury lamp through a water filter and a liquid-type optical fiber. Under this condition, the lamp with filters emits 260-600 nm of light (Figure 1C). On the other hand, TFA has no absorption above 240 nm (Figure 1A), whereas 1 has absorption from the UV region to 380 nm (Figure 1B). Hence, 1 is the only species that can absorb the light from the lamp. Figure 2 shows the irradiation time dependence of the photoreaction. In this case, a 10-fold molar excess of TFA based on 1 was used. As expected, the amount of TFA decreased, and F- and CO2 were found as products. After 96 h of irradiation, 33.3% of the initial TFA molecules were decomposed, corresponding to a turnover number [(mole of decomposed TFA)/(mole of initial 1)] of 3.36. At the same irradiation time, the yield of F- [(mole of formed F-)/(mole of initial TFA × 3)] was 32.4%. The (mole of formed F-)/ (mole of decomposed TFA) ratio was 2.91; that is, almost all fluorine atoms in the decomposed TFA molecules were converted to F- ions. Carbon dioxide accumulated in the gas phase. GC/MS analysis of the gas phase revealed no trace of any other gas product. The present system produces no environmentally undesirable gas species such as CF4, which is the most stable perfluorocarbon and has a global warming potential of at least 3900 times that of CO2 (22), and is often observed in the decomposition of perfluorocompounds by extremely high energy techniques such as electron beam irradiation (23). The yield of CO2 [(mole of formed CO2)/(mole of initial TFA × 2)] after 96 h of irradiation was 34.9%, and the (mole of formed CO2)/(mole of decomposed TFA) ratio was 2.09. These findings indicate that virtually all fluorine and carbon atoms in the decomposed TFA molecules were transformed into Fions and CO2.

FIGURE 2. Irradiation time dependence of the TFA decomposition: detected molar amounts of (A) TFA, (B) F-, and (C) CO2. An aqueous solution (23 mL) containing 1 (1.54 × 10-4 mol; 6.70 mM) and TFA (1.55 mmol; 67.4 mM) was irradiated with a mercury lamp under oxygen (0.55 MPa). The light from the mercury lamp was filtered to pass 260-600 nm of light, and the temperature of the solution was maintained at 25 °C.

TABLE 1. Reaction Yields under Several Combinations between Catalyst, Light Irradiation, and Reaction Atmospherea

no.

catalyst

light irradiation

1 2 3 4 5 6

presentb presentb none presentb presentb presentb

present none present present present present

atmosphere TFA (pressure, decomposition MPa) (%) O2 (0.55) O2 (0.55) O2 (0.55) Ar (0.10) air (0.10) O2 (1.10)

19.3 ( 1.7 0 0 1.3 15.6 20.1

Fyield (%) 17.6 ( 2.1 0 0 1.0 14.6 20.1

a Reaction time was 48 h, initial TFA amount was 1.55 mmol, and the reaction solution volume was 23 mL. b Amount of initial 1 was 1.54 × 10-4 mol.

The UV-visible absorption spectrum of the reaction solution after 96 h of irradiation was the same as that before irradiation, and the Raman and 31P NMR spectra after the reaction also showed no change. There was no sign of catalyst degradation after 96 h of irradiation, demonstrating that 1 is stable over long periods of irradiation. Table 1 lists the reaction yields under various combinations of the catalyst, light irradiation, and reaction atmosphere. In the absence of either light irradiation or 1, no reaction occurred (entries 2 and 3). When the reaction was carried out under argon instead of oxygen, only a small amount of F- (yield: 1.0%; entry 4, the corresponding value for an oxygen-atmosphere reaction was 17.6%; entry 1) was detected. Thus, it is clear that a combination of 1, oxygen, and light irradiation is required for effective TFA decomposition and that 1 works as a photocatalyst. When the reaction was carried out under air instead of oxygen, the reaction yields were decreased (entry 5). On the other hand, when the oxygen pressure was raised to 1.10 MPa, (entry 6), the yields were almost the same as those at 0.55 MPa. Thus, the 0.55 MPa of oxygen is enough to cause the catalytic reaction. The apparent quantum yield of TFA decomposition with 313 nm of irradiation under 0.55 MPa of oxygen was 0.008 when a 10-fold molar excess of TFA based on 1 was used. Therefore, about 1% of the photons absorbed by the solution are employed for the photocatalytic reaction. VOL. 37, NO. 2, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

419

FIGURE 3. Dependence of (A) F- formation, (B) CO2 formation, and (C) TFA decomposition on the initial amount of TFA. An aqueous solution (23 mL) containing 1 (1.54 × 10-4 mol; 6.70 mM) and a 1-20 molar excess of TFA was irradiated (260-600 nm) for 72 h under oxygen (0.55 MPa). The amounts of F- and CO2 at the maximum amount of initial TFA (3.08 mmol) were 2.32 (yield: 25.1%) and 1.49 mmol (24.2%), respectively.

SCHEME 1. TFA

Proposed Mechanism for the Decomposition of

The photocatalytic decomposition of TFA was greatly affected by the initial amount of TFA (Figure 3). For a constant amount of 1 and a reaction time of 72 h, the amounts of Fformed, CO2 formed, and TFA decomposed all increased with increasing initial amount of TFA. When the (initial TFA)/ (initial 1) molar ratio was 20:1 (initial TFA ) 3.08 mmol), the turnover number after 72 h of reaction was 5.58 and the amounts of F- and CO2 formed were 1.7-2.0 times those for a10:1 ratio. In accordance with the increase in reactivity with increasing initial amount of TFA, the apparent quantum yield at 313 nm of irradiation increased to 0.015 at the (initial TFA)/(initial 1) ratio of 20:1, which is 1.9 times that at the ratio of 10:1. Reaction Mechanism. Scheme 1 shows the proposed mechanism for the decomposition of TFA. In our case, the concentration of 1 is high (6.70 mM), so 1 is present as H+ and [PW12O40]3- in the reaction solution. Photoexcitation from the ground-state species [PW12O40]3- to the ligand-to-metal charge-transfer excited-state species [PW12O40]3-* is generally accepted as the initiation process of the photocatalysis by 1 (19-21). The relaxation process of [PW12O40]3-* is very fast. In the case of sodium and tetrapropylammonium salts of [PW12O40]3- in water, the reported emission lifetime is 1.5 ( 0.5 ns; however, the excited state is completely quenched by very small quantities (10-3 M) of alcohol because of the static mechanism based on the precomplexation of [PW12O40]3-and 420

9

FIGURE 4. UV-visible spectra of 1 (6.70 mM) in water: (A) without and (B) with TFA (67.4 mM). The concentrations of 1 and TFA were the same as those in the catalysis experiments shown in Figure 2.

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 2, 2003

FIGURE 5. UV-visible spectra of the sample solutions after 48 h of irradiation under (A) argon and (B) oxygen. The initial amounts of 1 and TFA were the same as those in Figure 2. The reaction solutions were transferred into quartz cells without dilution and then subjected to measurement. The spectra of the sample solutions prior to light irradiation were identical to spectrum B under both argon and oxygen atmospheres. alcohol (24). Thus, the enhancement of the catalytic reactivity with increasing initial amount of TFA observed in Figure 3 may indicate that the precomplexation of [PW12O40]3- with TFA is important to cause the subsequent reaction of [PW12O40]3-* with TFA. In fact, as shown in Figure 4, the UVvisible spectrum of 1 in water is red-shifted (∼6 nm) by TFA addition. This observation supports the precomplexation between [PW12O40]3- and TFA. To elucidate the reaction mechanism, we measured the UV-visible spectra of sample solutions after reaction under argon and oxygen. When the reaction was carried out under argon, the spectrum contained a broad absorption in the 500-1000-nm region (Figure 5A). This absorption reflects theappearanceoftheone-electron-reducedspecies[PW12O40]4-, identified by comparison with the spectrum of [PW12O40]4obtained electrochemically (25). The absorbance at 752 nm indicates that 5.1% of 1 converted into the reduced species. When the reaction was carried out under oxygen, catalytically producing F- ions and CO2, the spectrum was the same as that before the reaction, and no near-IR absorption was observed in the

FIGURE 6. GC/MS (total ion) chromatograms of the gas phase after 24 h of irradiation under (A) argon and (B) oxygen. The initial amounts of 1 and TFA were the same as those in Figure 2. The amounts of CO2 and CF3H measured in (A) were 2.0 × 10-5 and 1.3 × 10-6 mol, respectively. The control reaction (B) yielded 28.3 × 10-5 mol of CO2 and no trace of CF3H. 500-1000-nm region (Figure 5B). This observation clearly indicates that the catalytic reaction proceeds by a redox reaction between 1 and TFA. In the absence of oxygen, [PW12O40]4- is reoxidized to [PW12O40]3- very slowly. The presence of oxygen enhances reoxidation, leading to the catalytic decomposition of TFA to F- ions, a process similar to the generally accepted photocatalytic process by heteropolyacids (19). As for the fate of O2-, caused by this reoxidation process of the catalyst, it is in equilibrium with HO2 as shown in eq 2 with a pKa value of 4.8 (26). In our case, the reaction solution is highly acidic (pH 0.8); therefore, most of O2- should be transformed into HO2, and then, the HO2 radicals cause spontaneous disproportionation to give H2O2 and O2 (eq 3) (26):

O2- + H+ f HO2

(2)

HO2 + HO2 f H2O2 + O2

(3)

When the reaction was carried out under argon, GC/MS analysis of the gas phase revealed CF3H in addition to CO2 (Figure 6A). Trifluoromethane was not detected when a control reaction was carried out in the presence of oxygen (Figure 6B). These facts can be explained by assuming that decomposition of TFA proceeds through a photo-Kolbe mechanism to yield CO2 and CF3 radical. Under argon, the

CF3 radical preferentially reacts with H to produce CF3H. A similar photo-Kolbe mechanism was proposed in the decomposition of acetic acid to CH4 and CO2 using a polymolybdate in degassed water (27), where the reduced molybdic complex formed after irradiation combines with H+ from the acetic acid, and then the reduced complex reacts with CH3 radical to form CH4. Similarly, the H of CF3H in our case may come from H+ in the solution via the complexation with [PW12O40]4-. In the presence of a large amount of oxygen, CF3 predominantly reacts with O2 to form CF3O2, followed by disproportionation to CF3OF and CF2O (12). These species can cause hydrolysis, which results in F- ions. After we completed our study of the decomposition of TFA, we examined the possibility of reusing 1. Purification of heteropolyacids by ether extraction has been proposed in the literature (28) but has not been adopted for the recovery of catalysts from a reaction mixture. To accomplish this, the reaction solutions, which contained a total of 3.52 g (1.18 mmol) of 1, were collected in a flask and concentrated using a rotary evaporator. The concentrated solution (15 mL) was transferred in a separatory funnel, diethyl ether (40 mL) was added, and the mixture was shaken. After the mixture was allowed to stand, a new phase appeared at the bottom of the separatory funnel. We collected this phase and evaporated it to dryness. Spectroscopically (UV-visible, IR, 31P NMR) and ion chromatographically (no trace of F-), pure 1 was obtained with a recovery of 80%. We reused 1 for the photoreaction with TFA. The fresh and recovered 1 gave the TFA decomposition yields of 19.3 and 20.1%, respectively, where a 10-fold molar excess of TFA based on 1 was used and the reaction time was 48 h. Thus, even though our system is a homogeneous catalytic system, it was easy to separate catalyst 1 from the reaction mixture and reuse it. Our system can decompose TFA in higher concentrations such as 6.70-134 mM in water. Hence, it is suitable for the decomposition of TFA from stationary sources. Although stationary sources of TFA have not been well-clarified, fluoropolymer manufacturing, processing, and recycling plants, where the polymers are treated by heat, can act as the stationary sources because thermolysis of the polymers in air at 360-500 °C causes TFA in high yields (∼8%) (7). Furthermore, TFA is used as a solvent and/or reactant in many chemical plants, which are included in possible sources of TFA (29). Our system is also applicable to the decomposition of TFA from such plants. Of course, to treat wastewaters from the stationary sources, effects of other chemicals in the wastewaters should be taken into account. On this point, it has been shown that heteropolyacid photocatalysts including 1 can decompose chloroacetic acids (30). Recently, several other perfluorinated acids such as bioaccumulative perfluorooctane sulfonate have been detected in the environment (31, 32). Further application of this method to other perfluorinated acids is being investigated in our laboratory.

Literature Cited (1) Material Safety Data Sheet; Aldrich Chemical Co. Inc., Milwaukee, 2001. (2) Von Sydow, L. M.; Grimvall, A. B.; Bore´n, H. B.; Laniewski, K.; Nielsen, A. T. Environ. Sci. Technol. 2000, 34, 3115-3118. (3) Scott, B. F.; Mactavish, D.; Spencer, C.; Strachan, W. M. J.; Muir, D. C. G. Environ. Sci. Technol. 2000, 34, 4266-4272. (4) Berg, M.; Mu ¨ller, S. R.; Mu ¨hlemann, J.; Wiedmer, A.; Schwarzenbach, R. P. Environ. Sci. Technol. 2000, 34, 2675-2683. (5) Jordan, A.; Frank, H. Environ. Sci. Technol. 1999, 33, 522-527. (6) Wallington, T. J.; Hurley, M. D.; Fracheboud, J. M.; Orlando, J. J.; Tyndall, G. S.; Sehested, J.; Møgelberg, T. E.; Nielsen, O. J. J. Phys. Chem. 1996, 100, 18116-18122. (7) Ellis, D. A.; Mabury, S. A.; Martin, J. W.; Muir, D. C. G. Nature 2001, 412, 321-324. (8) Kim, B. R.; Suidan, M. T.; Wallington, T. J.; Du, X. Environ. Eng. Sci. 2000, 17, 337-342. VOL. 37, NO. 2, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

421

(9) Martin, J. W.; Franklin, J.; Hanson, M. L.; Solomon, K. R.; Mabury, S. A.; Ellis, D. A.; Scott, B. F.; Muir, D. C. G. Environ. Sci. Technol. 2000, 34, 274-281. (10) Ellis, D. A.; Hanson, M. L.; Sibley, P. K.; Shahid, T.; Fineberg, N. A.; Solomon, K. R.; Muir, D. C. G.; Mabury, S. A. Chemosphere 2001, 42, 309-318. (11) Boutonnet, J. C.; Bingham, P.; Calamari, D.; de Rooij, C.; Franklin, J.; Kawano, T.; Libre, J.-M.; McCulloch, A.; Malinverno, G.; Odom, J. M.; Rusch, G. M.; Smythe, K.; Sobolev, I.; Thompson, R.; Tiedje, J. M. Hum. Ecol. Risk Assess. 1999, 5, 59-124. (12) Pehkonen, S. O.; Siefert, R. L.; Hoffmann, M. R. Environ. Sci. Technol. 1995, 29, 1215-1222. (13) Hoffmann, M. R.; Martin, S. T.; Choi, W.; Bahnemann, D. W. Chem. Rev. 1995, 95, 69-96. (14) Alfano, O. M.; Bahnemann, D.; Cassano, A. E.; Dillert, R.; Goslich, R. Catal. Today 2000, 58, 199-230. (15) Fujishima, A.; Rao, T. N.; Tyrk,; D. A. J. Photochem. Photobiol. C: Photochem. Rev. 2000, 1, 1-21. (16) The´ron, P.; Pichat, P.; Guillard, C.; Pe´trier, C.; Chopin, T. Phys. Chem. Chem. Phys. 1999, 1, 4663-4668. (17) Mas, D.; Delprat, H.; Pichat P. In Photoelectrochemistry; Rajeshwar, K., Peter, L. M., Fujishima, A., Meissner, D., Tomkiewich, M. Eds.; The Electrochemical Society Inc.: Pennington, 1997; PV97-20; pp 289-299. (18) Maruthamuthu, P.; Padmaja, S.; Huie, R. E. Int. J. Chem. Kinet. 1995, 27, 605-612. (19) Papaconstantinou, E. Chem. Soc. Rev. 1989, 18, 1-31. (20) Okuhara, T.; Mizuno, N.; Misono, M. Adv. Catal. 1996, 41, 113252. (21) Hill, C. L.; Prosser-McCartha, C. M. In Photosensitation and Photocatalysis Using Inorganic and Organometallic Compounds;

422

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 2, 2003

(22)

(23) (24) (25) (26) (27) (28) (29) (30) (31) (32)

Kalyanasundaram. K., Gra¨tzel, M., Eds.; Kluwer Academic Publishers: Dordrecht, The Netherlands, 1993; pp 307-330. IPCC 2001: Climate Change 2001: The Scientific Basis. Contribution of Working Group I to the Third Assessment Report of the Intergovernmental Panel on Climate Change; Houghton J. T., Ding, Y., Griggs, D. J., Noguer, M., van der Linden, P. J., Dai, X., Maskwell, K., Johnson, C. A., Eds.; Cambridge University Press: Cambridge, 2001; pp 388-390. Pacansky, J.; Waltman, R. J. Phys Chem. 1991, 95, 1512-1518. Fox, M. A.; Cardona, R.; Gaillard, E. J. Am. Chem. Soc. 1987, 109, 6347-6354. Varga, G. M., Jr.; Papaconstantinou, E.; Pope, M. T. Inorg. Chem. 1970, 9, 662-667. Bielski, B. H. J.; Cabelli, D. E.; Arudi, R. L.; Ross, A. B. J. Phys. Chem. Ref. Data 1985, 14, 1041-1100. Yamase, T.; Kurozumi, T. J. Chem. Soc., Dalton Trans. 1983, 2205-2209. Pope, T. Heteropoly and Isopoly Oxometalates; Springer-Verlag: Berlin, 1983. Scott, B. F.; Spencer, C.; Marvin, C. H.; Mactavish, D. C.; Muir, D. C. G. Environ. Sci. Technol. 2002, 36, 1893-1898. Mylonas, A.; Hiskia, A.; Papaconstantinou, E. J. Mol. Catal. A: Chem. 1996, 114, 191-200. Giesy, J. P.; Kannan, K. Environ. Sci. Technol. 2001, 35, 13391342. Giesy, J. P.; Kannan, K. Environ. Sci. Technol. 2002, 36, 146A152A.

Received for review May 13, 2002. Revised manuscript received November 11, 2002. Accepted November 11, 2002. ES025783Y