Role of Quinone Intermediates as Electron Shuttles in Fenton and

Jul 30, 1997 - The light source was a L201 tunable 150W xenon illuminator from Photon Technology International (South Brunswick, NJ). The band-pass wa...
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Environ. Sci. Technol. 1997, 31, 2399-2406

Role of Quinone Intermediates as Electron Shuttles in Fenton and Photoassisted Fenton Oxidations of Aromatic Compounds RUZHONG CHEN AND JOSEPH J. PIGNATELLO* Department of Soil and Water, The Connecticut Agricultural Experimental Station, P.O. Box 1106, New Haven, Connecticut 06504

Fenton and related reactions are potentially useful oxidation processes for destroying toxic organic compounds in water. In these reactions, H2O2 is combined with Fe(II) or Fe(III) in the presence or absence of light to generate hydroxyl radicals (HO•). A relatively neglected area of research is the influence of organic species on the reactivity of iron, and hence on the rate or course of the reaction. This study examined the oxidation of phenol (2 mM) by Fenton systems in the dark (Fe3+/H2O2, Fe2+/H2O2, and mixed Fe3+/Fe2+/H2O2 systems) and under UV/visible light (Fe3+/ H2O2/hν). Reactions were conducted at an initial pH of 2.8, with H2O2 in excess and iron in catalytic concentrations. In all cases, the reactions display autocatalysis. In dark reactions, the lag phase decreases (a) with increasing total iron concentration, [Fe]T; (b) as initial [Fe2+] increases when [Fe]T is held constant; (c) in proportion to the amount of hydroquinone (1,4- or 1,2-) added; and (d) in proportion to the amount of quinones (1,4-benzoquinone or 5-hydroxy1,4-naphthoquinone [juglone]) added. The hydroquinones reduce Fe3+ to Fe2+. An important result is the finding that quinones serve as electron-transfer catalysts between dihydroxycyclohexadienyl radicalsthe HO• adduct of phenolsand Fe3+ by way of a semiquinone radical. This conclusion is supported by simulations with a kinetic model employing known and proposed steps. In irradiated solutions, the lag phase decreases with decreasing wavelength (480-300 nm; 10 nm band-pass). A cause of light initiation, especially above ∼410 nm where photolysis of Fe3+ and/or H2O2 is negligible, is direct photolysis of quinones to HO• and semiquinone radicals, which subsequently reduce Fe3+ and re-form the quinone. Thus, quinones play an important catalytic role in Fenton oxidation of aromatic compounds.

Introduction Hydroxyl (HO•) is a reactive, nonselective radical (1, 2) underlying the chemistry of many advanced oxidation processes (AOPs) for degrading toxic organic compounds in water. Recently there has been a renewed interest in the use of Fenton-type AOPs, which generate HO• from hydrogen peroxide and an iron compound (3-6). A number of variants of the Fenton reaction have been studied, including those using Fe2+ or Fe3+, in the presence or absence of complexing ligands and in the presence or absence of light (3, 5-9). The * Corresponding author telephone: (203) 789-7237; fax: (203) 7897232; e-mail: [email protected].

S0013-936X(96)01064-4 CCC: $14.00

 1997 American Chemical Society

core thermal and photochemical inorganic reactions involved in HO• production and decay are well established (eqs 1-13, Table 1) (10). Little is known, however, about the influence of the target compounds or their degradation products on the reactivity of iron. Organic compounds may act as ligands (L) or redox agents. Possible reactions include (i) photolysis of Fe(III)-L complexes (5, 11); (ii) formation of reactive highvalent iron-oxo or iron-peroxo complexes stablized by L (12); (iii) reduction of Fe3+ by neutral organic molecules (13, 14) or radicals (10); and (iv) oxidation of Fe2+ by radicals or carbocations (10). An understanding of these reactions is important for the successful application of Fenton-type remediation technologies. We have undertaken a study of the Fenton degradation of phenol as a model for aromatic pollutants. Previous studies show that HO• initially adds to the aromatic ring to give a substituted hydroxycyclohexadienyl radical (I, Scheme 1) (10). Depending on pH and availability of suitable oxidants or reductants, I may react further by: acid-catalyzed elimination of water to form a radical cation (path A); formation of a dioxygen radical adduct (path B); oxidation to a cyclohexadienyl cation (path C); orswhen paths B and C are unavailablescoupling and disproportionation (path D) (15, 16). The impetus for this study was 2-fold. First, it was reported by Hamilton et al. (17) and later by Litvintsev et al. (18) that catalytic amounts of 1,4-hydroquinone (1,4-HQ) or 1,2hydroquinone (catechol, 1,2-HQ)sboth intermediates in phenol degradationsgreatly increased the rate of the Fe3+/ H2O2 degradation of substituted benzenes. They proposed that the active oxidizing agent is a ternary HQ-Fe-H2O2 complex. However, since 1,2- and 1,4-HQ are known to reduce Fe3+ rapidly at low pH (13, 14), it seemed more likely to us that rate acceleration was due to cut off production of Fe2+, which then generates HO• through the normal Fenton pathway (eq 4). Nevertheless, the observations suggested at least the possibility of a novel role for organics in the pathway, one that might involve an alternative oxidant. Second, it has been shown in studies where HO• was generated by radiolysis (15) that quinone intermediates of phenol degradation are able to oxidize I (Scheme 1, quinone ) OX in path C). Since the reduction product of quinonesa semiquinone radicalsmay reduce Fe3+ to regenerate the quinone, this suggested that quinone intermediates may act as electron-transfer catalysts between I and iron ions.

Experimental Section Chemicals. Water was distilled and then further purified in a Barnstead Nanopure system. Chemicals were used as received except where noted. The following were from Aldrich: 1,4-benzoquinone (1,4-BQ; 98%), 1,2-HQ (99+%), 1,4-HQ (99+%), resorcinol (99+%), NaClO4 (99%), Fe(ClO4)2‚6H2O (98%), Fe(ClO4)3‚xH2O (11.8% Fe by weight), and FerroZine reagent [3-(2-pyridyl)-5,6-diphenyl-1,2,4-triazinep,p′-disulfonic acid, monosodium salt; 97%]. 1,4-BQ was resublimed under vacuum. Juglone (5-hydroxyl-1,4-naphthoquinone, 98+%) was from Lancaster. Hydrogen peroxide (30%), K2C2O4, FeCl3, and 1,10-phenanthroline were from Fisher Scientific. Methanol and acetonitrile (HPLC grade) were from J. T. Baker. Phenol was purchased from Sigma (99+%, no preservatives) and redistilled under reduced pressure of N2. Reproducible results were obtained with this material. The lag period in phenol degradation was found, however, to be sensitive to the source of phenol due to the presence of impurities with a quinone or hydroquinone structure. A sample from Sigma containing 0.1% H3PO2 as preservative to

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2.7 × 107 4.3 × 108

k7f ) k7rKa k7r ) 5 × 1010 1.0 × 104

1.2 × 106

8.5 × 105

1.5 × 108

1.0 × 107 9.7 × 107 7.3 × 109

2.7 × 107 4.3 × 108 pKa ) 4.8 k7r ) 5 × 1010 1.0 × 104

1.2 × 106

8.5 × 105

1.5 × 108

1.0 × 107 9.7 × 107 7.3 × 109

5

6

7

8

9

10

11

12

13

14

24e

23e

22

21

20

19

18

17e

16

15

no.

reaction

d

p: ∼109

k23r ) 1.2 × 10-3

usedc

1 × 109

k23f ) 4.4 × 104 k23r ) 1.2 × 10-3

k22f ) 4.4 × 102 k22r ) 1.1 × 103

1 × 105

1 × 109

1 × 109

1 × 109

3.7 × 109

k16a ) 1.5 × 109 k16b ) 1.0 × 109

5 × 108

k (M-1 s-1)

k22f ) 4.4 × 102 k22r/k23f ) 0.024

na

∼109

na

∼109

o: 3.7 × 109 p: 3.7 × 109

na

o: ∼108 p: ∼109

lit.b

a In light, reaction 26 in the text also occurs. b References: k (24); k (10); k , k , k c 3 4 5 6 14 (2); k7, k10, k13 (25); k8, k9, k11, k12 (26); k15, k17 (15); k18 (27); k20 (28); k22, k23 (13, 14); k24 (29). The rate constants used in computer simulations. k16a, k16b, k21, k19, and k23f are treated as variables in the simulations (see text). d Simulations which include HO2• as a product here give essentially similar results as long as k16a ) k16b ) 1.05 × 109 M-1 s-1. e Analagous reactions for 1,2-benzoquinone.

76

76

4

0.01

usedc

0.01-0.02

reaction

k (M-1 s-1)

3

2

1

no.

lit.b

TABLE 1. Reactions in Iron-Catalyzed Oxidation of Phenol by Hydrogen Peroxidea

SCHEME 1

suppress air oxidation showed a decreasing lag occurring over several months of use. Procedures. All experiments were conducted in aerated solution at 0.1 M ionic strength (NaClO4) and initial pH of 2.80 adjusted with HClO4. All solutions were made fresh daily. Stock solutions of Fe3+ or Fe2+ at 10-20 mM concentration were prepared in 0.05 or 0.1 M HClO4. Stock solutions of H2O2 were prepared by diluting 30% H2O2. The final reaction solution was usually made by mixing equal volumes of solutions A and B, each adjusted to pH 2.8: solution A contained iron at twice target concentrations; solution B contained H2O2 and phenol at twice target concentrations. Other chemicals, if needed, were added as quickly as possible to the mixture. For the juglone experiment, chemicals were added to pH-adjusted Fe3+ solutions in the order phenol, H2O2, and juglone. Dark experiments were carried out in Teflon-lined septum screw-cap 40 mL vials or 250 mL Erlenmeyer flasks in a water bath at 25 ( 0.1 °C in a room illuminated by dim red light. The solution was stirred with a Teflon-coated magnetic stir bar. The solution was sampled through the septum by syringe with a Teflon needle (Aldrich); typically, 1 mL was withdrawn and mixed with an equal volume of methanol, which quenched the reaction. The actual volume of liquid withdrawn was determined by weight. Photolyses were carried out in a 3-mL Teflon-lined septum screw-cap quartz cell. The light source was a L201 tunable 150W xenon illuminator from Photon Technology International (South Brunswick, NJ). The band-pass was set at 10 nm. The reaction solution was prepared as above, and 3 mL was transferred into the quartz cell containing a small Tefloncoated magnetic stir bar. Typically, 50 µL of sample was withdrawn by syringe and mixed with 350 µL of methanol. Analyses. Phenol, 1,4-BQ, and 1,4-HQ were analyzed by HPLC on a 5 µm, 25 cm × 5 mm Spherisorb ODS-2 C-18

column and detected with a Hewlett-Packard 1050 diode array UV/vis detector. The mobile phase was 50% acetonitrile/ 50% water at 1.0 mL/min for the analyses of phenol (monitored at 210 nm and/or 270 nm; retention time, tR, 4.7 min) and 1,4-BQ (monitored at 250 nm; tR, 4.2 min) in dark and photoassisted Fenton experiments. In the 1,4-BQ quantum yield experiments, the mobile phase was 30% acetonitrile/70%water with 0.08% trifluoroacetic acid at 1.0 mL/min for the analyses of 1,4-HQ (monitored at 220 nm and/or 290 nm; tR, 4.0 min) and 1,4-BQ (monitored at 250 nm; tR, 6.4 min). HPLC analysis was usually done within 3 h, but never longer than 6 h, after sampling. Ferrous ion was determined by a modified FerroZine method (19). A methanolic solution of 2 mM FerroZine reagent and 18 mM NH4F was mixed with an equal volume of reaction solution and the absorbance at 562 nm was measured. The strong complexation of Fe3+ by F- was used to suppress any further reduction of Fe3+ after the reaction. When the measurement is carried out within 2 min, the minimum time required for manipulation, a correction of only e3% in [Fe2+] is needed. Hydrogen peroxide concentration was measured by iodometry (20). The light intensity was determined by chemical actinometry (21). Kinetic Modeling. Reaction rates were simulated with GEAR (PC version 1.11, Project Seraphim, Department of Chemistry, University of Wisconsin, Madison, WI 53706), which finds a numerical solution to the set of differential equations defined by the users’ multistep reaction scheme (22).

Results and Discussion Dark Reactions. Figure 1 shows phenol loss (A), 1,4-BQ formation (B), and Fe2+ formation (C) in a series of experiments carried out at 2 mM phenol, 4 mM H2O2, and different

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FIGURE 1. Dark reactions. Effect of the ratio [Fe3+]i/[Fe2+]i (mM/ mM). Initial concentrations: 2 mM phenol; 0.05 mM [Fe]T; 4 mM H2O2. Lines represent simulations. initial Fe3+:Fe2+ ratios, where total iron concentration [Fe]T was kept constant at 0.05 mM. In all cases, the loss of phenol displays autocatalysis; that is, there is a lag period followed by a fast, first-order decomposition. In general, the pH remained constant throughout the lag phase and dropped slightly (to pH ∼2.7) during the fast reaction phase. Several important results of Figure 1 are to be noted. (i) The lag period declines as the initial concentration of ferrous ion ([Fe2+]i) increases but still exists even when [Fe2+]i ) [Fe]T. (ii) The rate of reaction after the lag period is nearly independent of the initial Fe3+:Fe2+ ratio. (iii) When Fe is present initially in the ferric state, [Fe2+] exists at undetectable levels throughout most of the lag period but then increases rapidly as phenol enters its reaction phase, reaching steady state at ∼0.6[Fe]T. (iv) When Fe is present initially in the ferrous state, [Fe2+] drops by more than an order of magnitude within one minute but then recovers to the same steady state concentration as when Fe was present initially in the ferric state. These results can be explained in the following way. When [Fe]T ) [Fe3+]i, the lag phase of phenol actually represents a slow rate-limiting reduction of Fe3+ by H2O2 (eq 3, Table 1) to Fe2+ which sustains the Fenton reaction (eq 4). Progression from the lag phase to the reaction phase is due to the buildup of [Fe2+] through additional, more efficient reactions. One such reaction is the reduction of Fe3+ by hydroquinones, 1,4and 1,2-HQ, which are major initial products of phenol degradation (Scheme 1), producing 1,4- and 1,2-BQ, respectively (13, 14). These reactions take place by two one-electron transfer steps with a semiquinone radical as intermediate (eqs 22 and 23). Reduction of Fe3+ by this route is much faster than reduction of Fe3+ by H2O2 and accounts for the transition from the lag phase to the reaction phase. In Figure 1B it can be seen that the maximum rate of 1,4-BQ formation coincides with that of phenol decomposition. 1,2-BQ was not monitored due to its instability in water (23) and lack of a reference standard. The final yield of 1,4-BQ is relatively low because it is only one of several products formed (see Scheme 1) and because it is consumed by reaction with HO•. Depletion of H2O2 toward the end of the reaction accounts for the leveling off in [1,4-BQ]. The early rapid decline in [Fe2+] observed when all Fe is present initially in the ferrous

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FIGURE 2. Dark reactions. (A, B) Effect of added hydroquinone (1,4HQ). (C) Effect of added catechol (1,2-HQ) on phenol loss. Initial concentrations: 2 mM phenol; 0.05 mM [Fe3+]i; 4 mM H2O2. Lines represent simulations. state is due to its consumption in the Fenton reaction (eq 4); the subsequent rapid recovery in [Fe2+] is due to production of reducing organic intermediates such as 1,2- and 1,4-HQ. This accounts for the decreasing length of the lag period with decreasing Fe3+:Fe2+ ratio. The involvement of reducing intermediates such as hydroquinones is confirmed by the results of Figure 2, which shows that amendment of Fe3+/ H2O2 systems with 1,4-HQ or 1,2-HQ greatly reduces the lag period in relation to the amount added. The rate of the fast phase becomes nearly independent of the amount of HQ added as the stoichiometry of the reduction (HQ:Fe ) 1:2) is approached and then exceeded. The phenol lag phase persists until reducing intermediates build up to sufficient levels. Inspection of the rate constants in Table 1 shows that HO• will preferentially react with phenol over H2O2 or Fe2+ until [phenol] declines below ∼0.005 mM. The initial product of HO• attack is dihydroxycyclohexadienyl radical (DHCD•, I, R ) OH) which reacts mainly by dehydration and peroxidation pathways in the presence of O2 (paths A and B, respectively; Scheme 1). Hydroxylated products are among several formed in these pathways. For benzene, path B gives only about 50% hydroxylated products (16). Path C leads exclusively to hydroxylation but requires a suitable oxidant. Subsequent experiments showed that quinones play a role as oxidants in path C, driving the reaction toward hydroxylation. Addition of 1,4-BQ reduces the lag period of Fe3+/ H2O2 reactions (Figure 3) and Fe2+/H2O2 reactions (Figure 4) in proportion to the amount of 1,4-BQ added. Furthermore, 0.05 mM 5-hydroxy-1,4-naphthoquinone (juglone) had the same effect as 0.05 mM 1,4-BQ in the Fe3+/H2O2 system (Figure 3). These results are consistent with the involvement of quinones in the oxidation of DHCD• in path C. One-electron oxidation of DHCD• by 1,4-BQ affords a semiquinone radical and dihydroxycyclohexadienyl cation (DHCD+) (eq 17, Table 1). Subsequently, DHCD+ loses a proton to give a dihydroxybenzene. The semiquinone radical and two of the three possible dihydroxybenzeness1,2- and 1,4-HQ but not the 1,3-isomer (resorcinol)sare capable of reducing Fe3+ (13, 14). This results in regeneration of the corresponding quinone.

FIGURE 5. Effect of substituting reaction 25 for reaction 17 on simulations of the data previously shown in Figures 3A and C. Dashed lines correspond to simulations with k25 ) 0.001k17. Solid lines correspond to simulations with k25 ) k17.

FIGURE 3. Dark reactions. Effect of added 1,4-benzoquinone (open symbols) or juglone (closed symbol) on reactions starting with Fe3+. Initial concentrations: 2 mM phenol; 0.05 mM [Fe3+]i; 4 mM H2O2. Lines represent simulations.

FIGURE 4. Dark reactions. Effect of added 1,4-benzoquinone on reactions starting with Fe2+. Initial concentrations: 2 mM phenol; 0.05 mM [Fe2+]i; 4 mM H2O2. Thus, the quinones serve as catalysts to shuttle electrons from DHCD• to Fe3+. Quinone should have no effect on a mechanism involving a ternary HQ-Fe-H2O2 complex as the active oxidant, as proposed by the Hamilton (17) and Litvintsev (18) groups. Kinetic Modeling. To test the proposed redox chain mechanism, we simulated phenol decomposition using a kinetic model that incorporates the reactions and their rate constants (when available) listed in Table 1 (2, 10, 13-15, 24-29) except the photolysis reactions 1 and 2. In several of the reactions, the rate constant depends on the substitution pattern (o, m, p) in the benzene ring. The m position is ordinarily the least reactive, and we have assumed this to be true in cases where it has not been established. Values of k14 (2), k18 (27), and k20 (28) are reported as composite values, since the isomer-specific values either are unknown, or else the isomer distribution could not be determined experimentally. For k15, an average of the o and p values was used (15). For reactions 22 and 23, the k22f and k23r and the ratio k22r/k23f are known for both o and p isomers (13, 14); we arbitrarily

used the p- isomer values but showed that the quality of fit is independent of whichever is used. For reaction 24, we used the rate constant for the p isomer, since that of the o isomer is unknown. For reaction 16 the branching ratio between hydroxylation (k16a) and ring-opening (k16b) was allowed to vary around 50:50, the approximate ratio reported for benzene (16). In summary, a total of 28 reactions were used to simulate 10 sets of data with five unknowns (k16a, k16b, k19, k21, k23f). Simulations of data on [phenol] and [Fe2+] are shown in Figures 1, 2, and 3 for systems of variable initial Fe3+:Fe2+ ratio, added hydroquinone, and added 1,4-BQ, respectively. The values of the five unknowns resulting from the simulations are in Table 1. While an exact fit of the data was not obtained, considering the complexity of the reaction mechanism the kinetic model does a good job of predicting the curve shape and the position of the curves relative to each other. The results of a sensitivity analysis are depicted in graphical forms in the Supporting Information of this paper. Briefly, this analysis shows that any combination of reasonable values for the unknowns predicts qualitatitively the general curve shape and position of the curves relative to one another. The simulations are sensitive to k16 () k16a + k16b), as well as the hydroxylation to ring-opening branching ratio, k16a:k16b. The best value of k16 is 25 times greater than that reported for anisole, while k16a:k16b (60:40) is close to the 50:50 ratio for benzene (16). The values of the other unknowns seem reasonable; in any case, the simulations are relatively insensitive to 1 order of magnitude variations in them. Several other important aspects of the simulations are to be noted. First, the model consistently overestimates the rate of the fast reaction phase. This is because the model assumes HO• reacts only with phenol. We attribute the overestimation to competition from degradation intermediates whose concentrations build up as phenol concentration declines. Second, we have evaluated the contribution of reaction 25 to oxidation of DHCD•.

Although Fe3+ has a greater reduction potential than 1,4-BQ [E° ) 0.77 V (1) and 0.31 V (30), respectively], previous results (15, 31), when taken together, show the order of reactivity toward cyclohexadienyl radicals to be Fe(CN)63- ∼ 1,4-BQ > Cu2+ > Fe3+. When reaction 25 is added to the kinetic model

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encompassing reactions 3-24, it has a negligible effect on the simulations as long as k25 e 0.1k17. When k25 is between 0.1k17 and k17, the effect of reaction 25 is merely to smooth out the transition from the lag period to the fast reaction period in the phenol curves. On the other hand, when reaction 25 is substituted for reaction 17, the kinetic model fails completely to simulate even the qualitative features of the experimental curves and fails to register the dependence on quinone concentration, regardless of the value assigned to k25 within the range 0.001k17 to k17. This is shown in Figures 5A and B, which are to be compared with Figures 3A and C, respectively. Simulations to be compared with the results in Figures 1 and 2 are in the Supporting Information. The failure of simulations based on reaction 25 lends further support for the hypothesized role of electron-transfer catalysis by quinones. Photoassisted Reactions. Figure 6A shows the reaction of 2 mM phenol with 4 mM H2O2 and 0.05 mM Fe3+ in the dark and irradiated at different wavelengths from 300 to 480 nm. The lag, which is present under all conditions, declines in light compared to dark and with decreasing wavelength. Figure 6B shows that irradiation at 300 nm of phenol solutions containing either H2O2 alone or Fe3+ alone gives no obvious lag in phenol decomposition and that the kobs values (Table 2) are 2-3 orders of magnitude smaller than when H2O2 and Fe3+ are present together. The observed first-order rate constant normalized to light intensity (kobs/Io) for the fast reaction phase in the Fe3+/H2O2/hν reaction is only slightly wavelength-dependent, declining about 3-fold as the wavelength increases from 300 to 480 nm (Table 2). The relative insensitivity of kobs/Io to wavelength suggests that phenol oxidation is controlled by the Fenton and other thermal reactions governing the steady state [Fe2+] and that light initiates the reaction. Initiation becomes more efficient as the wavelength decreases. Photolysis of H2O2 (eq 1, Table 1) is possible below its absorption edge of 350 nm. Photoreduction of Fe3+ (actually FeOH2+) (eq 2) is possible below its absorption edge at ∼410 nm, although it has a very low quantum yield above 350 nm (32). Nevertheless, it can be seen in Figure 6A that wavelengths well above 400 nm are also capable of initiating the reaction.

TABLE 2. Rate Constants for the Fast Reaction Phase of Phenol Oxidation by Photoassisted and Dark Fenton Reactions and by Direct Photolysis of H2O2 or Fe3+ wavelength I0 ×106 b kobs × 103 (nm) (Einstein/L s) (s-1)

reaction Fe3+/H

2O2/hν

a

(kobs/I0) × 10-2 (L/Einstein)

300 320 350 400 450 480

3.46 4.01 5.66 6.78 6.76 8.45

2.1 2.0 1.8 1.7 1.6 1.8

6.1 5.0 3.2 2.5 2.4 2.1

H2O2/hνb Fe3+/hνc

300 300

3.46 3.46

0.016 0.0017

0.045 0.0049

Fe3+/H2O2a

dark

1.5

a

For photoassisted and dark Fenton reactions, 4 mM [H2O2] and 0.05 mM [Fe3+] were used; kobs were calculated from Figure 6A assuming first-order kinetics. b For direct photolysis of 4 mM H2O2. kobs was calculated from corresponding curve in Figure 6B and divided by 2 to permit direct comparison with photoassisted Fenton reactions. c For direct photolysis of 0.05 mM Fe3+. kobs was calculated from corresponding curve in Figure 6B and divided by 10 to permit direct comparison with photoassisted Fenton reactions.

We suggest that initiation also takes place by direct photolysis of 1,2- and 1,4-BQ intermediates, resulting in products that can reduce Fe3+. Figure 1B, where [Fe3+]i ) [Fe]T, shows that during the lag phase in the dark, thermal reactions cause a gradual increase in [1,4-BQ] from initially very low levels. 1,4-BQ has three major absorption bands in the UV/visible region: λmax ) 248 nm (Log  ) 4.32), 298 nm (2.48), and 426 nm (1.33). According to Ononye et al. (33), photolysis of 1,4-BQ alone in water gives semiquinone radical and hydroxyl radical (eq 26),

with subsequent reactions producing 1,4-HQ and hydroxy-

FIGURE 6. Irradiated reactions. (A) Effect of wavelength on phenol loss under photo-Fenton conditions. The dark reaction is also shown. Dashed line is simulation of the dark reaction; dotted line is simulation of the photoassisted reaction including eq 26 (k26 ) ΦI0El ) 0.003 M-1 min-1, l ) 1 cm, the light pass length). Initial concentrations: 2 mM phenol; 0.05 mM [Fe3+]i; 4 mM H2O2. (B) Phenol loss during photolysis of 0.49 mM Fe3+ alone (O) or 8mM H2O2 alone (0) at 300 nm. Initial concentration: 2 mM phenol.

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TABLE 3. Quantum Yields for the Conversion of 2.48 mM 1,4-Benzoquinone as a Function of Wavelength in Irradiated Aqueous Solution (0.1 M NaClO4, pH 2.8)

SCHEME 2. Quinone (BQ) and Hydroquinone (HQ) Roles in Fenton Reactionsa

I0 × 106 E final observed wavelength (Einstein/ (L/mol photolysis [hydroquinone quantum (nm) L s) cm) time (min) (mM) yield, Φ 300 320 350 400 450 480

3.46 4.01 5.66 6.78 6.76 8.45

300 179 18 17 15 4

8 10 25 25 25 25

0.20 0.27 0.18 0.17 0.18 0.07

0.28 0.34 0.40 0.34 0.40 0.40

1,4-benzoquinone. A similar reaction sequence presumably occurs for 1,2-BQ. No complexation occurred between 1,4BQ and Fe3+ by UV/vis. We observed 1,4-HQ accumulation in irradiated solutions of 1,4-BQ alone at pH 2.8 and 0.1 M NaClO4; Table 3 lists the quantum yield for 1,4-BQ conversion (based on 1,4-HQ formation) as a function of wavelength from 300 to 480 nm. The quantum yield varies between 0.28 and 0.40 almost independently of λ . Quantum yields of 0.3 (34) and 0.47 (33) using polychromatic light have been reported. Note in Table 3 that significant amounts of 1,4-HQ are generated at all wavelengths. In photoassisted Fenton reactions, the products of quinone photolysis (eq 26) represent 2 equiv of HO• because semiquinone reduces Fe3+ to Fe2+ (eq 23). Moreover, the quinone is regenerated. Thus, quinones act as photocatalysts in the presence of Fe3+ and H2O2 and cause a reduction in the lag period even above wavelengths where Fe3+ and H2O2 are not photolyzed. Computer simulation incorporating reaction 26 into the kinetic scheme of Table 1 (but leaving out eqs 1-2 to simulate photolysis at >410 nm and taking  ) 17 L mol-1 cm-1) shows that the lag phase is indeed reduced compared to the dark reaction but is still longer than the experimental lag phases for the photoassisted reaction. This suggests that other reactions contribute to reduction of the lag period; possibilities include photolysis of a Fe(III)-peroxide complex (5) and photolysis of Fe(III) complexes of organic acids (11). Between 400 and 300 nm, photolyses of Fe3+ and H2O2 (eqs 1 and 2) play an increasingly dominant role. Experimental and simulation work have demonstrated that quinone intermediates shuttle electrons from the HO• radical adduct of the starting aromatic compound to Fe3+, thus facilitating the degradation of the starting aromatic compound. In light, quinones serve as photosensitizers via their conversion to semiquinone, which reduces Fe3+, and HO•. The role of quinone is summarized in Scheme 2. Phenol degradation can be rationalized according to this scheme without having to invoke the ternary complex (17, 18). In waste treatment applications it is desirable to use iron in catalytic concentration in order to hold down reagent costs and minimize iron oxide sludge production. For the reaction to be catalytic in iron, pathways must be available for the regeneration of Fe(II) from Fe(III). Such pathways are provided by reactions 22 and 23, as well as by reaction 2 in light. In terms of effective degradation, hydroxylation and peroxidation pathways (paths B and C, Scheme 1) are more productive with respect to mineralization than dehydration (path A). Therefore, the presence of co-oxidants, such as oxygen and quinones, plays a crucial role in steering oxidation into the desired pathways.

Acknowledgments This work was done under a grant from the National Science Foundation (BES 9414594). We thank Tina Arounsak and Susan Devlin for technical assistance.

Supporting Information Available Seven figures plus their captions showing the results of the simulation sensitivity analysis (9 pp) will appear following

a Legend: SQ, semiquinone; DHCD+ and DHCD•, dihydroxycyclohexadienyl cation and radical, respectively.

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Received for review December 23, 1996. Revised manuscript received March 27, 1997. Accepted April 8, 1997.X ES9610646

X

Abstract published in Advance ACS Abstracts, June 15, 1997.