Environ. Sci. Technol. 1996, 30, 3457-3463
Reduction of Perchloroalkanes by Ferrioxalate-Generated Carboxylate Radical Preceding Mineralization by the Photo-Fenton Reaction
(near-UV) than the above AOTs, excepting TiO2/UV, and is catalytic in iron (3-9). In the photo-Fenton system, acidic solutions of Fe(III) are photolyzed to generate Fe(II) and OH• (eq 1). With excess H2O2 present, the Fe(II) is reoxidized to Fe(III) as in eq 2.
PATRICK L. HUSTON† AND JOSEPH J. PIGNATELLO*
The photo-Fenton reaction has been studied in detail and has been found to successfully degrade pesticides, phenols, and halogenated hydrocarbons in water (3-5, 7, 10, 11). However, carbon tetrachloride (CCl4) and other saturated perhalogenated aliphatic compounds do not react rapidly with OH• (1, 2, 12) and are not efficiently degraded by standard AOTs (1, 13). For example, in one study (13) the half-life of CCl4 in TiO2/UV photocatalysis was nearly 6 h under conditions where most organic substrates were degraded with half-lives of 2-6 min. Even this slow loss of CCl4 is not due to the direct reaction with OH•. Instead, it is the electron in the conduction band of TiO2 that reacts with CCl4 and the slowness of the reaction is due to a competition with oxygen. Recently, it was reported that CCl3• radical was formed in the reaction of lignin peroxidase with CCl4 in the presence of H2O2, veratryl alcohol, and oxalate (C2O42-) (14). It was suggested that veratryl alcohol was enzymatically oxidized to the corresponding cation radical, which subsequently oxidized oxalate. Presumably, the CO2•- radical that formed was able to reduce CCl4 (eq 3).
Department of Soil and Water, The Connecticut Agricultural Experiment Station, P.O. Box 1106, New Haven, Connecticut 06504
Perhalogenated aliphatic compounds in water are difficult to degrade by advanced oxidation processes that rely on hydroxyl radical as the reactive oxidant. We have successfully degraded carbon tetrachloride and hexachloroethane in UV-illuminated (300-400 nm) acidic oxic or anoxic solutions containing Fe(III) and oxalate. Kinetic and product studies were carried out to elucidate the mechanism and role of O2. The initial step is a one-electron reduction of the perhaloalkane by CO2•-, which is generated by photolysis of iron(III) oxalate complexes. The products are lowerchlorinated compounds, several of which are unreactive toward CO2•-. In the presence of O2, the reaction is catalytic in Fe(III) due to the re-oxidation of Fe(II) by oxo-radical intermediates; however, O2 also retards the rate by scavenging CO2•-. After the perhalogenated substrate has been transformed, the system can be made oxidizing by the addition of H2O2, resulting in mineralization of the remaining organochlorine intermediates by a photo-Fenton reaction.
Introduction Advanced oxidation technologies (AOTs) such as TiO2/UV, O3/UV, O3/H2O2, O3/H2O2/UV, and Fenton reactions, Fe(II,III)/H2O2, have been studied in recent years to determine whether they would be practical for destroying organic contaminants in water (1). The common factor that these methods share is the generation of hydroxyl radical (OH•). The second-order rate constant for reaction of OH• with the organic substrate often approaches the diffusioncontrolled limit (2), making these AOTs extremely efficient once OH• has been generated. A recent AOT that has been proposed is the photo-assisted Fenton reaction which, while also generating OH• radicals, requires lower energy radiation * To whom correspondence should be addressed; telephone: (203)789-7237; fax: (203)789-7232; e-mail:
[email protected]. † Present address: Chemistry Department, University of Evansville, Evansville, IN 47722.
S0013-936X(96)00091-0 CCC: $12.00
1996 American Chemical Society
hν
Fe(OH)2+ 98 Fe2+ + OH•
(1)
Fe2+ + H2O2 f Fe3+ + OH- + OH•
(2)
CO2•- + CCl4 f CO2 + CCl3• + Cl-
(3)
This contradicts a pulse radiolysis study that found CO2•to be relatively unreactive with CCl4 (15). However, there must be a strong reducing agent present to react with CCl4 in the veratryl alcohol-oxalate system, and CO2•- is a likely candidate since its estimated reduction potential (E° ) ∼ -1.9 V, NHE) (16) is more negative than that of CCl4 (E° ) -0.229 V) (14) and comparable to the reduction potentials of R-hydroxyalkyl radicals, which are known to reduce CCl4 (17). One method of generating CO2•- is to photolyze ferric oxalate complexes, e.g., ferric trioxalate (eq 4) (18-20). The oxalyl radical (C2O4•-) produced in the photochemical step decomposes to carbon dioxide and CO2•- with an estimated first-order rate constant k5 of 2 × 106 s-1 (eq 5) (21), which is so fast as to preclude direct reaction of C2O4•- with millimolar concentrations of organic substrates. hν
Fe(C2O4)33- 98 Fe2+ + 2C2O42- + C2O4•-
(4)
C2O4•- f CO2 + CO2•-
(5)
Ferrioxalate, furthermore, is compatible with the photoFenton system, and others have used it as the source of Fe(II) in the photo-Fenton reaction (8, 22). It has recently been shown that reduction can be achieved in AOTs under certain conditions. Reduction of CCl4 can occur during TiO2 photocatalysis in the presence of sacrificial organic electron donors such as alcohols and
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carboxylic acids (23). The donor serves primarily to scavenge valence-band holes or adsorbed HO• on the surface, thus favoring reduction of CCl4 by conductionband electrons. Oxalate was found to be a poor donor presumably because it does not react rapidly with HO•. It has also been shown that alcohols may react with OH• to form R-hydroxyalkyl radicals R2C•_OH (R2C•_O-) that can reduce CCl4 due to its relatively weak C-Cl bond. Peyton (24) achieved reduction of CCl4 (as well as nitroaromatic compounds) during O3/UV or H2O2/UV treatment by including ethanol in the reaction mixture. The reaction was sensitive to O2, which scavenged R-hydroxyalkyl radicals. In this work, we investigated the decomposition of CCl4 and hexachloroethane (C2Cl6) during photolysis of aqueous solutions containing Fe(III) and oxalate. Kinetic and product studies were carried out, and the effect of oxygen on the reaction was explored. We show that the organic reduction products are subsequently mineralized when photo-Fenton conditions are resumed upon the addition of H2O2.
Experimental Section Water was distilled and deionized (Barnstead NANOPure). Perchloric acid (J. T. Baker), hydrogen peroxide (30%, Fisher), iron(III) perchlorate (GFS Chemicals), and organic compounds (J. T. Baker; Alrich) were used as received. The supplier’s elemental analysis for iron in ferric perchlorate was used in calculating its concentration in reaction mixtures. Reactions were carried out in a cylindrical 330-mL borosilicate double-walled reaction vessel with water circulated through the walls to maintain constant temperature. The vessel, which was acid-washed before each use, was filled to minimize the headspace and sealed with PTFE-lined screw caps. The photochemical reactor chamber (Rayonet RPR-200) was equipped with 16 14-W fluorescent black lamps, which emit in the range 300-400 nm. The lamps were turned on for 10 min prior to reaction to obtain constant output. With all the lights on, the intensity in the vessel was 1.2 × 1019 quanta s-1 L-1 as measured by ferrioxalate actinometry. In the dark, all reagents except H2O2 or Fe(III) (depending on the experiment) were added to the vessel, and the temperature was equilibrated to 25 °C. The pH was adjusted to 2.8 with HClO4. The reaction was initiated by addition of H2O2 or Fe(III), followed by insertion of the vessel into the chamber. Reactions under anoxic or oxygen-saturated conditions were carried out by sparging solutions with N2 or O2, respectively, prior to addition of the organohalide stock solution. Samples were withdrawn at timed intervals and immediately extracted into hexane for GC/ECD analysis. Product identities were confirmed by GC/MS (HP5890 GC interfaced to HP5988A quadrapole MS) after comparison with standards. Chloride was determined by use of a Corning ion-selective electrode or by ion chromatography using a Dionex IONPAC ASA4 column with post-column micromembrane suppression and conductivity detection. Oxalate was determined by ion chromatography on the same column. Iron(II) was determined by the absorbance of its phenanthroline complex (25).
Results When solutions of CCl4 were photolyzed under photoFenton conditions (7 × 10-6 M Fe(III) and 5.0 × 10-3 M
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FIGURE 1. Transformation of CCl4 in initially air-saturated aqueous solution under photo-Fenton conditions: 7 × 10-6 M Fe(III), 5.0 × 10-3 M H2O2, and 3.0 × 10-3 M C2O42- at pH 2.8 and 25.0 °C.
FIGURE 2. Degradation of CCl4 and formation of products under air-saturated excess ferrioxalate conditions: no added H2O2, 7 × 10-4 M Fe(III), and 3.0 × 10-3 M C2O42-.
H2O2, pH 2.8 and 300-400 nm light) only 5-10% loss of CCl4 was observed over a 2-h period. However, upon addition of oxalate, slow, but significant loss of CCl4 occurred. For example (Figure 1), in an initially airsaturated system containing 4.5 × 10-4 M CCl4, 7 × 10-6 M Fe(III), 3.0 × 10-3 M C2O42-, and 5.0 × 10-3 M H2O2, a 72% decline in [CCl4] occurred gradually over a 1-h period before the reaction stopped. Fe(III) appears to be a catalyst because loss of [CCl4] corresponded to ∼45 times the initial Fe(III) concentration. Initial experiments to determine the mechanism of this reaction were conducted without peroxide. In the first experiment, the reagents were present in excess with respect to CCl4 [7 × 10-4 M Fe(III) and 3 × 10-3 M oxalate; hereafter referred to as “excess ferrioxalate” conditions] in initially air-saturated solution. Figure 2 shows that rapid and complete loss of 5.0 × 10-5 M CCl4 occurred when photolyzed under these conditions. There appears to be
FIGURE 3. Degradation of CCl4 and formation of products under air-saturated catalytic conditions: 7 × 10-6 M Fe(III) and 3.0 × 10-3 M C2O42-.
FIGURE 4. Degradation of CCl4 under anoxic excess ferrioxalate conditions showing incomplete loss of CCl4 and formation of stable products: no added H2O2, 0.7 × 10-4 M Fe(III) and 3.0 × 10-3 M C2O42-.
an induction period since the reaction rate increased with time during the first few minutes of the reaction. Hexachloroethane appeared after 5 min and reached a maximum yield of 38% (based on [CCl4]i ) before disappearing in ∼40 min. Tetrachloroethene (C2Cl4) was formed over a 20-min period in 67% yield and remained as a stable product. In a separate, identical experiment, the sum of Cl- and organochlorine in detected products accounted for 97% of total Cl. A dark control and an oxalate-only photolyzed control (not shown) each showed less than 10% loss of CCl4 over 2 h. The reaction was repeated under the same conditions except that Fe(III) was present in catalytic amount (7 × 10-6 M; hereafter referred to as “catalytic” conditions) (Figure 3). Decomposition of 1.3 × 10-4 M CCl4 under airsaturated catalytic conditions was rapid and complete. Compared to excess ferrioxalate conditions, under catalytic conditions (a) the lag is more apparent; (b) C2Cl6 and C2Cl4 are formed in lower yields and are less persistent in the reaction mixture; and (c) chloroform (CHCl3) is a trace product. In a similar catalytic experiment, except that [CCl4]i was slightly lower (0.7 × 10-4 M), it was determined that decomposition of oxalate practically ceases after CCl4 and C2Cl6 are destroyed and that 9.4 mol of oxalate/mol of CCl4 is consumed. Chloride ion, combined with identified organochlorine products, accounted for 85% of total chloride at 40-min reaction time. Additional experiments were carried out in the absence of O2 to help elucidate the mechanism. Photolysis of CCl4 under anoxic excess ferrioxalate conditions is shown in Figure 4. About 92% of CCl4 was decomposed in 4 min without an induction period to produce C2Cl6 and smaller amounts of C2Cl4, whereupon the reaction stopped. Chloroform was detected in trace yields later in the reaction. The cause of the abrupt halt is exhaustion of the Fe(III) catalyst by its photoreduction to Fe(II). This is clear from the experiment in Figure 5 in which a supplement of iron(III) perchlorate added after CCl4 degradation leveled off resulted in a second rapid loss of ∼50% of the remaining CCl4. Also, stoichiometric formation of Fe(II) was confirmed in a separate experiment.
FIGURE 5. Influence of Fe on degradation of CCl4 under similar initial conditions as in Figure 4. After 15 min, 2.8 × 10-4 M additional Fe(III) was added resulting in additional loss of CCl4.
Hexachloroethane (28 µM) also degraded rapidly under anoxic excess ferrioxalate conditions (Figure 6) forming C2Cl4 quantitatively. These results indicate that the reduction of CCl4 leads to CCl3• radicals, which combine to give C2Cl6, and that C2Cl6 undergoes further reduction to give C2Cl4:
CCl4 + [e-] f CCl3• + Cl-
(6)
2CCl3• f CCl3CCl3
(7)
CCl3CCl3 + [e-] f CCl3CCl2• + Cl-
(8a)
CCl3CCl2• + [e-] f Cl2CdCCl2 + Cl-
(8b)
Figure 7 shows the effect of O2 concentration on the catalytic reaction. Under anoxic conditions, the reaction
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FIGURE 6. Degradation of C2Cl6 and formation of tetrachloroethene under anoxic excess ferrioxalate conditions: 0.7 × 10-4 M Fe(III) and 3.0 × 10-3 M C2O42-.
FIGURE 7. Oxygen effect on catalytic reaction: 7 × 10-6 M Fe(III) and 3.0 × 10-3 M C2O42-.
is slow and incomplete. This is due to exhaustion of the Fe(III) catalyst by reduction to Fe(II) as discussed above. Upon comparing the results under air-saturated and O2saturated conditions, it is clear that O2 is inhibitory. However, the reaction eventually becomes fast even under O2-saturated conditions. In both the air- and O2-saturated systems, complete loss of CCl4 occurred, and the amount of CCl4 reacted was more than 17 times the inital [Fe(III)]. Thus, in the presence of O2 the reaction is catalytic in Fe(III) despite the initial absence of hydrogen peroxide to re-oxidize the Fe(II). The nature of the species reducing the organohalide was deduced from experiments in which OH• was generated by direct photolysis of 5 × 10-3 M H2O2 in anoxic solution at pH 2.8. The absorption spectrum of H2O2 (absorption edge, ∼360 nm) partly overlaps the lamp emission spectrum. Special care was taken in these experiments to exclude iron by prewashing glassware with HNO3 and by avoiding any
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contact of reagents with metal surfaces. Under peroxideonly conditions, CCl4 was stable. However, when the reaction mixture contained 0.10 M HC2O4-, which is oxidized by OH• to C2O4•- (kOH ) 4.7 × 107 L mol-1 s-1) (26), decomposition of CCl4 occurred to give the same products as those in ferrioxalate photolyses (Figure 8a). These results indicate that oxalate is a source of radicals that can react with CCl4. In another experiment, the reaction mixture contained 0.10 M HCO2H, which produces CO2•- directly upon reaction with OH• (kOH ) 1.3 × 108 at pH 1 and kOH ) 3.2 × 109 at neutral or alkaline pH) (2). Decomposition of CCl4 occurred in this system also, again without induction, to give the same products as in the oxalate system plus very low and transient yields of pentachloroethane (CCl3CHCl2) and 1,1,2,2-tetrachloroethane (CHCl2CHCl2) (Figure 8b). (These results were unaffected by including 1 µM ethylenediaminetetraacetic acid, EDTA, as a further precaution against the effects of trace iron.) Taken together, the oxalate and formate results indicate that CO2•- itself reacts with CCl4. Competition studies were carried out to determine the reactivity toward other chlorinated alkanes of the reductant produced in the iron(III) oxalate systems. C2Cl6 was employed as a reference compound. Under anoxic excess ferrioxalate conditions, 3.5 × 10-7 M C2Cl6 reacts rapidly (80% loss in 10 min), while 2.72 × 10-4 M CHCl2CHCl2 present in the same reaction mixture was not degraded at all (Figure 9). This indicates that the rate constant for reduction of C2Cl6 is at least 4 orders of magnitude greater than that of CHCl2CHCl2. Similarly, under comparable conditions, CH2ClCH2Cl was not degraded at all in the presence of C2Cl6 (data not shown). Lastly, to effect complete mineralization of CCl4, we tested a sequential reduction-oxidation process in which reduction would take place in the presence of oxalate followed by oxidation by the addition of H2O2 to attain photo-Fenton conditions. The main byproduct of the reduction step is C2Cl4, which is a pollutant of concern, but is slow to react under those conditions. Reduction of CCl4 is inefficient in the presence of H2O2, presumably due to scavenging of the reductant by H2O2. Peroxide, however, may be added after CCl4 is depleted in order to degrade C2Cl4. Figure 10 shows degradation in an air-saturated catalytic system that was ammended with H2O2 after all CCl4 had disappeared (∼20 min) as compared with catalytic and excess ferrioxalate systems in which no H2O2 was added. In the H2O2-ammended case, C2Cl4 was completely removed, and the yield of Cl- was quantitative. In the unammended cases, C2Cl4 was persistent or degraded at a much slower rate.
Discussion Photolysis of Fe(C2O4)33- initially leads to formation of the oxalyl radical, C2O4•- as in eq 4. This radical can reduce a second Fe(III) (eq 9) or reduce O2 to generate the HO2• radical (eq 10), which is unreactive toward most organic compounds (26). Although C2O4•- may be capable of reducing CCl4 (eq 11), its rapid decomposition to CO2•- (eq 5) precludes this possibility. At typical CCl4 concentrations used (e.g., 1 × 10-4 M), k11 would have to be close to the diffusion-controlled limit (∼1 × 1010 M-1 s-1) to compete with the decomposition reaction (k5 ) 106 s-1) and even much higher than that (1012-1013 L mol-1s-1) to achieve CCl4 loss below the detection limit. Thus, if we accept the literature estimate for k5, C2O4•- cannot be the active
FIGURE 8. Direct photolysis of 5.0 × 10-3 M H2O2 leading to degradation of CCl4 and formation of halogenated products in the presence of (a) 0.10 M H2C2O4 or (b) 0.10 M HCO2H. The inset in b shows the appearance of CCl3CHCl2 and CHCl2CHCl2.
FIGURE 9. Competitive reaction of C2Cl6 (3.5 × 10-7 M) and 1,1,2,2tetrachloroethane (2.72 × 10-4 M) under anoxic excess ferrioxalate conditions: 7 × 10-4 M Fe(III) and 3.0 × 10-3 M C2O42-; photon intensity, 1.5 × 1018 quanta s-1 L-1.
FIGURE 10. Degradation of tetrachloroethene byproduct by sequential treatment (H2O2 added at arrow) compared with the usual catalytic or excess ferrioxalate systems in which no H2O2 was added. Conditions (air-saturated): 3.0 × 10-3 M C2O42-; 7 × 10-4 M Fe(III) (excess ferrioxalate); or 7 × 10-6 M Fe(III) (sequential and catalytic).
reductant, and reaction 11 is unimportant.
To test this hypothesis, we generated C2O4•- by reaction of OH• from H2O2 photolysis with oxalate as in eq 12. Rapid loss of CCl4 occurred. This is consistent with either C2O4•or CO2•- as the reductant. Confirmation that CO2•- itself reacts with CCl4 was obtained by observing rapid loss of CCl4 upon reaction of OH• with formate, which gives CO2•directly (eq 13). Although oxalate may eventually build up in formate solution by dimerization of CO2•- (27) (eq 14), the absence of a lag period rules this out. Thus, when oxalate is donor, CO2•- and not C2O4•- itself is the reactant. The CO2•- radical will also react with H2O2 yielding OH• (eq 15), but this process is relatively slow (k15 ) 7.3 × 105 L mol-1 s-1) (28). Loss of CCl4 is predicted to occur at identical rates in 0.1 M HC2O4- and 0.1 M HCO2H if decomposition of C2O4•to CO2•- (eq 5) is fast relative to other reactions of C2O4•-,
C2O4•- + Fe(C2O4)33- f 2CO2 + Fe2+ + 3C2O42-
(9)
C2O4•- + O2 + H+ f 2CO2 + HO2•
(10)
C2O4•- + CCl4 f 2CO2 + CCl3• + Cl-
(11)
HO• + HO2CCO2- f H2O + C2O4•-
(12)
HO• + HCO2H f H2O + CO2•- + H+
(13)
2CO2•- f C2O42-
(14)
CO2•- + H2O2 f CO2 + OH- OH•
(15)
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and if CCl4 does not react by any other route. However, the degradation of CCl4 in formate solution is ∼4 times faster, gives a much higher yield of CHCl3, and gives two additional products, CCl3CHCl2 and CHCl2CHCl2. This indicates that CCl4 is also consumed in a chain reaction propagated by the initially formed CCl3• radical, as in eqs 16 or 17 followed by eq 3.
CCl3• + HO2CCO2- f CHCl3 + CO2 + CO2•- (16) CCl3• + HCO2H f CHCl3 + H+ + CO2•-
(17)
In this chain, k17 > k16, in keeping with the weaker formyl C-H bond relative to the oxalyl O-H bond [e.g., kOH (formate) . kOH (oxalate), see above], resulting in longer chains in formate solution; this accounts for the faster rate and the higher yield of CHCl3. The trace products CCl3CHCl2 and CHCl2CHCl2 are reasonably formed by consecutive reduction of C2Cl6:
CO2•- + CCl3CCl3 f CCl3CCl2• + Cl- + CO2 •
+
(18a) •-
CCl3CCl2 + HCO2H f CCl3CHCl2 + H + CO2 (18b) CO2•- + CCl3CHCl2 f •CCl2CHCl2 + Cl- + CO2 (18c) •
CCl2CHCl2 + HCO2H f CHCl2CHCl2 + H+ + CO2•(18d)
The reactions in oxalate solution corresponding to those in eq 18b-d are less important due to the weaker H atom donor ability of oxalate as compared to formate. The rate of CO2•- reaction with the alkyl halide follows the order C2Cl6 . CHCl2CHCl2, CH2ClCH2Cl (by many orders of magnitude); Figure 8b (inset) shows that CCl3CHCl2 is reactive but its position in the order is unknown. It may also be inferred from Figure 8b that the order is CCl4 > CHCl3. These results are consistent with an electron transfer reaction. A Cl-atom abstraction mechanism is possible but less likely in our view. The unpaired electron of CO2•- appears to be centered on C [coupling through C occurs in some reactions of CO2•- (29-31)] and Cl abstraction by carbon radicals is known (32). However, the neighboring effects of Cl in the R or β positions with respect to the incipient organoradical are relatively small (33). At this time, we can offer no explanation for the apparent discrepancy between our results and the literature (15) regarding whether CO2•- reacts with CCl4. The effect of O2 must be considered because it is present in many waste streams. Dioxygen is expected to scavenge C2O4•- (via eq 10) and/or CO2•- (via eq 19) and thus inhibit the reaction. The value of k19 is 1 × 109 L mol-1 s-1 (34) while k10 is unknown. Scavenging of CO2•- is confirmed by the existence of an induction period that increases with increasing [O2]. Obviously, all or most O2 in solution must be consumed before reduction of organohalide can begin. Oxalate serves as a sacrificial electron donor that is consumed both in reduction of O2 and the organohalide; thus, its cost may be of concern for treating large volumes of aerated water with very low levels of contaminant.
CO2•- + O2 + H+ f CO2 + HO2•
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(19)
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Reactions 10 and 19 are useful, however, because the product HO2• itself or its dismutation product H2O2 (eq 21) can reoxidize Fe(II) (via eq 20 and eq 2, respectively). Degradation of perhaloalkane is catalytic in Fe(III) in the presence of O2. Figures 4 and 5 show that reduction of perchloroalkane stops when all of the Fe(III) is reduced. Thus, the reaction will continue as long as the Fe(II) is somehow reoxidized. Another possible oxidant of Fe(II) is the trichloromethyl peroxyl radical (eqs 22, 23), which is known to be oxidizing (17, 34-36). The high yield of free Cl- observed here agrees with the appearance of Cl- upon hydrolysis of CCl3OOH (eq 24) (37).
HO2• + Fe(II) + H+ f H2O2 + Fe(III)
(20)
2HO2• f H2O2
(21)
CCl3• + O2 f CCl3OO•
(22)
CCl3OO• + Fe(II) + H+ f CCl3OOH + Fe(III)
(23)
CCl3OOH + H2O f 3H+ + 3Cl- + CO2 + 1/2O2 (24) The formation of chlorine-containing byproducts such as C2Cl4 and CHCl3 that are relatively stable to further reduction is a potential drawback because further treatment would be required to remove them. However, compounds containing unsaturation or abstractable H normally react rapidly with OH• and, therefore, are susceptible to oxidation by existing AOTs. The photo-Fenton system is especially suitable in this regard because it can be put into effect simply by adding H2O2 to the reaction mixture. Moreover, ferrioxalate is advantagous as a photo-Fenton precursor because it strongly absorbs visible light up to 550 nm and has a high quantum yield for the generation of Fe(II) (22, 38, 39). The superior overlap with the solar spectrum and the efficiency of the photochemical processes leading to reducing radicals make this an ideal system for reductive dehalogenation of perhalogenated aliphatic compounds. It is shown here that complete mineralization of CCl4 and C2Cl6 in contaminated water can be achieved in the presence of O2 by a tandem reaction consisting of ferrioxalate photoreduction followed by photo-Fenton oxidation (i.e., Figure 10).
Acknowledgments Technical assistance was provided by Susan J. Devlin. Funding was provided by the USDA Water Quality Special Grants Program and NSF Bioengineering and Environmental Systems Program.
Literature Cited (1) (a) Legrini, O.; Oliveros, E.; Braun, A. M. Chem. Rev. 1993, 93, 671-698; (b) von Sonntag, C.; Mark, G.; Mertens, R.; Schuchmann, M. N.; Schuchmann, H.-P. J. Water Supply Res. Technol.Aqua 1993, 42, 201-211. (2) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J. Phys. Chem. Ref. Data 1988, 17, 513-886. (3) Lipczynska-Kochany, E. Environ. Technol. 1991, 12, 87-92. (4) Pignatello, J. J. Environ. Sci. Technol. 1992, 26, 944-951. (5) Sun, Y.; Pignatello, J. J. Environ. Sci. Technol. 1993, 27, 304-310. (6) Haag, W. R.; Yao, C. C. D. Environ. Sci. Technol. 1992, 26, 10051013. (7) Rupert, G.; Bauer, R.; Heisler, G. J. Photochem. Photobiol. 1993, 73, 75-78. (8) Zepp, R. G.; Faust, B. C.; Hoigne´, J. Environ. Sci. Technol. 1992, 26, 313-319.
(9) Sedla`k, P.; Lun ˜ ak, S.; Brodilova`, J.; Lederer, P. React. Kinet. Catal. Lett. 1989, 39, 249-253. (10) Pignatello, J. J.; Huang, L. Q. Water Res. 1993, 27, 1731-1736. (11) Pignatello, J. J.; Sun, Y. Water Res. 1995, 29, 1837-1844. (12) Guittonneau, S.; De Laat, J.; Dore, M.; Duguet, J. P.; Bonnel, C. Rev. Sci. Eau 1988, 1, 35-54. (13) Sabin, F.; Turk, T.; Vogler, A. J. Photochem. Photobiol. A: Chem. 1992, 63, 99-106. (14) Khindaria, A.; Grover, T. A.; Aust, S. D. Environ. Sci. Technol. 1995, 29, 719-725. (15) Ko¨ster, R.; Asmus, K.-D. Z. Naturforsch. 1971, 26b, 1104-1108. (16) (a) Schwarz, H. A.; Dodson, R. W. J. Phys. Chem. 1989, 93, 409414. (b) Koppenol, W. H.; Rush, J. D. J. Phys. Chem. 1987, 91, 4429-4430. (17) Packer, J. E.; Willson, R. L.; Bahnemann, D.; Asmus, K.-D. J. Chem. Soc. Perkin Trans. 2 1980, 296-299. (18) Demas, J. N.; Bowman, W. D.; Zalewski, E. F.; Velapoldi, R. A. J. Phys. Chem. 1981, 85, 2766-2771. (19) Parker, C. A.; Hatchard, C. G. J. Phys. Chem. 1959, 63, 22-26. (20) Parker, C. A. Trans. Faraday Soc. 1954, 50, 1213-1221. (21) Mulazzani, Q. G.; D’Angelantonio, M.; Venturi, M.; Hoffman, M. Z.; Rodgers, M. A. J. J. Phys. Chem. 1986, 90, 5347-5352. (22) Safarzedeh-Amiri, A. Photocatalytic Method for Treatment of Contaminated Water. U. S. Patent 5,266,214, 1993. (23) Choi, W.; Hoffman, M. R. Environ. Sci. Technol. 1995, 29, 16461654. (24) Peyton, G. R.; Bell, O. J.; Girin, E.; Lefaivre, M. H. Environ. Sci. Technol. 1995, 29, 1710-1712. (25) Greenberg, A. E.; Clesceri, L. S.; Eaton, A. D. Standard Methods for the Examination of Water and Wastewater; American Public Health Association: Washington, DC, 1992; pp 3-66-3-68. (26) Getoff, N.; Schwoerer, F.; Markovic, V. M.; Sehested, K.; Nielsen, S. O. J. Phys. Chem. 1971, 75, 749-755. (27) Neta, P.; Simic, M.; Hayon, E. J. Phys. Chem. 1969, 73, 42074213.
(28) Kishore, K; Moorthy, P. N.; Rao, K. N. Radiat. Phys. Chem. 1987, 29, 309-313. (29) Popp, J. L.; Kalyanaraman, B.; Kirk, T. K. Biochemistry 1990, 29, 10475-10480. (30) Goldstein, S.; Czapski, G.; Cohen, H.; Meyerstein, D. J. Am. Chem. Soc. 1988, 110, 3903-3907. (31) Farahani, M.; Surdhar, P. S.; Allen, S.; Armstrong, D. A.; Scho¨neich, C.; Mao, Y.; Asmus, K. D. J. Chem. Soc. Perkin Trans. 2 1991, 1687-1693. (32) Russell, G. A. In Free Radicals; Kochi, J. K., Ed.; John Wiley & Sons: New York, 1973; Vol. 1, Chapter 7. (33) Sutton, H. C.; Seddon, W. A.; Sopchyshyn, F. C. Can. J. Chem. 1978, 56, 1961-1964. (34) Packer, J. E.; Mahood, J. S.; Mora-Arellano, V. O.; Slater, T. F.; Willson, R. L.; Wolfenden, B. S. Biochem. Biophys. Res. Commun. 1981, 98, 901-906. (35) Mertens, R.; von Sonntag, C.; Lind. J.; Merenyi, G. Angew. Chem. Int. Ed. Engl. 1994, 33, 1259-1261. (36) Bonifacic, M.; Schoneich, C.; Asmus, K.-D. J. Chem. Soc., Chem. Commun. 1991, 1117-1119. (37) Mo¨nig, J.; Bahnemann, D.; Asmus, K.-D. Chem.-Biol. Interact. 1983, 45, 15-27. (38) Calvert, J. G.; Pitts J. N., Jr. Photochemistry; John Wiley & Sons: New York, 1966; pp 783-786. (39) Balzani, V.; Carassiti, V. Photochemistry of Coordination Compounds; Academic Press: New York, 1970; pp 167-174.
Received for review January 29, 1996. Revised manuscript received July 25, 1996. Accepted July 29, 1996.X ES960091T X
Abstract published in Advance ACS Abstracts, October 15, 1996.
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