Photocatalytic Transformation of Organic Compounds in the Presence

Jan 21, 2000 - This first paper deals with the effect of fluoride on the photocatalytic transformation of phenol. Phenol photocatalytic ..... Equation...
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Photocatalytic Transformation of Organic Compounds in the Presence of Inorganic Anions. 1. Hydroxyl-Mediated and Direct Electron-Transfer Reactions of Phenol on a Titanium Dioxide-Fluoride System C. Minero, G. Mariella, V. Maurino, and E. Pelizzetti* Dipartimento di Chimica Analitica, Universita` di Torino, Via Pietro Giuria 5, 10125 Torino, Italy Received March 19, 1999. In Final Form: November 15, 1999 The effect of fluoride ions on the photocatalytic degradation of phenol in an aqueous suspension of TiO2 has been investigated. Fluoride ions displace surficial hydroxyl groups and coordinate surface-bound titanium atoms directly. For 0.01 M fluoride concentration and 0.10 g L-1 of TiO2 in the range pH 2-6, the degradation rate of phenol is up to 3 times that in the absence of fluoride ions. This behavior has been correlated with the computed surface speciation. The decrease in the degradation rate of phenol as a function of the substrate concentration observed in naked TiO2 at a high concentration of phenol (over 0.01 M) is largely diminished in the presence of fluoride ions. A photocatalytic model which takes into account the primary events and recombination reactions is able to account for these experimental results. The competition between OH-radical-mediated reaction versus direct electron transfer is discussed. Finally, under a helium atmosphere and in the presence of fluoride ions, phenol is slowly but significantly degraded, although total organic carbon does not decrease, suggesting the occurrence of a photocatalytically induced hydrolysis.

Introduction The presence of inorganic ions affects the mineralization process of organic compounds under photocatalytic conditions.1-3 Early works envisaged the inhibitory effect of chloride and bromide anions on the photocatalyzed degradation of halogenated aliphatics.4 Mineral acids such as HNO3, HCl, and HClO4 have been chosen among the scavengers to investigate charge-carrier recombination dynamics for titanium dioxide.5 Because there is almost unanimous consent that the photocatalytic reactions occur at the surface of the semiconductor particle,6 specific adsorption of ions may reasonably influence the system performance. Specific adsorption of ions can be treated in terms of surface coordination reactions at the oxide-water interface.7 The binding of anions can be related to the electrostatic interaction with the surface, depending on the relative values of pH and point of zero charge (PZC), and to exchange reactions with the surface hydroxyl groups. The surface occupation by anions may be competitive with adsorption of organic molecules. This effect is directly related to their coverage fraction. Whatever the interaction with the surface, there the ions are subjected to redox transformations after electron transfer with photogenerated charge carriers. This second effect could produce inhibition by competition of inorganic ions with the organics. Inhibition by strongly adsorbed anions (1) Abdullah, M.; Low, J. K. C.; Matthews, R. W. J. Phys. Chem. 1990, 94, 6820-6825. (2) Kormann, C.; Bahnemann, D. W.; Hoffmann, M. R. Environ. Sci. Technol. 1991, 25, 494-500. (3) D’Oliveira, J. C.; Guillard, C.; Maillard, C.; Pichat, P. J. Environ. Sci. Health 1993, A28, 941-956. (4) Hsiao, C. H.; Lee, C. L.; Ollis, D. F. J. Catal. 1983, 82, 418-423. (5) Martin, S. T.; Herrmann, H.; Hoffmann, M. R. J. Chem. Soc., Faraday Trans. 1994, 90, 3323-3330. (6) Minero, C.; Catozzo, F.; Pelizzetti, E. Langmuir 1992, 8, 481486. (7) Stumm, W.; Morgan, J. J. Aquatic Chemistry; Wiley: New York, 1996.

has been reported, while for weakly adsorbed anions such as nitrate and perchlorate, a negligible effect has been observed.1-3 The redox reactions of adsorbed ions form reactive species, which can further interact with the organics and their intermediates. This was reported for CHCl3, for which in the presence of chlorides under photocatalytic conditions CCl4 is formed.8 The effects of several common anions, such as halides (F-, Cl-, Br-, I-) and oxoanions (NO3-, ClO4-, SO42-, HCO3-, H2PO4-), on the photocatalytic processes involving some target organic molecules will be examined in this and some forthcoming papers. The organic substrates have been chosen by their aromatic/aliphatic structure and their solubility in water, to discriminate between the chemical reactivity and adsorption properties and to assess the possibility of degradation through oxidative or reductive pathways (also in the absence of oxygen). The rationale for this investigation was to get information on (i) the mechanism of inhibition as cited above, (ii) the effect of anions on the photocatalytic pathways of organic transformation, (iii) the extent of inhibition as a function of the nature and concentration of the anions, (iv) the extent of formation of some undesired intermediates (for example, the chloro and bromo derivatives), and (v) the possibility of minimizing the inhibition effect by proper selection of operative conditions. This first paper deals with the effect of fluoride on the photocatalytic transformation of phenol. Phenol photocatalytic degradation has been extensively investigated.9,10 As far as fluoride peculiarity is concerned, this anion shows strong adsorption on TiO2,11 and the estimated redox (8) Minero, C.; Maurino, V.; Calza, P.; Pelizzetti, E. New J. Chem. 1997, 21, 841-842. (9) Okamoto, K.; Yamamoto, Y.; Tanaka, H.; Tanaka, M.; Itaya, A. Bull. Chem. Soc. Jpn. 1985, 58, 2015-2022. (10) Augugliaro, V.; Palmisano, L.; Sclafani, A.; Minero, C.; Pelizzetti, E. Toxicol. Environ. Chem. 1988, 16, 89-109. (11) Bohem, H. P. J. Chem. Soc., Faraday Discuss. 1972, 66, 264275.

10.1021/la9903301 CCC: $19.00 © 2000 American Chemical Society Published on Web 01/21/2000

Photocatalytic Transformation of Organic Compounds

potential of the couple F•/F- is 3.6 V.12 This potential makes fluoride stable against oxidation by TiO2 valence holes, even in acidic media. Thus, for this anion the redox competition with the organics is not possible. Experimental Section Materials and Reagents. All degradation experiments were carried out using TiO2 Degussa P25 as the photocatalyst. To avoid possible interference from ions adsorbed on the photocatalyst, the TiO2 powder was irradiated and washed with bidistilled water until no signal due to chloride, sulfate, or sodium ions could be detected by ion chromatography. Phenol, catechol, quinol, and NaF (Aldrich) were used as received. HClO4, HNO3, and NaOH (reagent grade) were used to adjust the pH. Samples of TiO2 with adsorbed cathecol with or without fluoride ions have been prepared for reflectance spectra. The powders of TiO2 were prepared from P25 Degussa by irradiating for 24 h, washing, and drying at 80 °C for 2 h. This specimen was used as a reference. The TiO2 powder with adsorbed cathecol was prepared from a slurry containing 2 g L-1 of P25, by addition of 2 × 10-4 M cathecol in the presence of 40 mM NaCl/HCl at pH 3.4. The slurry changes color within 10 s after addition of cathecol. After 1 h of equilibration time in the dark, the TiO2-catecholate surface-modified titania has been recovered by centrifugation and dried at 80 °C for 2 h. The TiO2 powder for the specimen with cathecol in the presence of fluoride ions was prepared from a slurry containing 2 g L-1 of P25, by addition of 2 × 10-4 M cathecol in the presence of 40 mM NaF/20 mM HF. The measured pH was 3.4. After 1 h of equilibration time in the dark, the fluorinated TiO2 has been recovered by centrifugation and dried at 80 °C for 2 h. Irradiation Procedures. Aqueous solutions of the organic compounds and previously prepared suspensions of catalyst were mixed in the cells used for the irradiation experiments (final volume 5 mL). Typical final concentrations were in the range of 0.1-10 mM (from 3 to 1000 part per million, ppm) for the organics and 100 mg L-1 (ppm) of TiO2. The cells containing the reaction slurry were kept in the dark in a water bath at a temperature close to that of the lamp housing used for illumination until they reached the working temperature (about 50 °C). Irradiation was carried out in cylindrical Pyrex glass cells (4.0 cm diameter and 2.3 cm height) on 5 mL of the aqueous suspension containing the desired amount of the photocatalyst powder and substrate, using a 1500 W xenon lamp (Solarbox, CO. FO. MEGRA, Milan, Italy) equipped with a 340 nm cutoff filter. The irradiance spectrum and the cells were described elsewhere.13 Different Solarbox devices have been used. The total photon flux (340-400 nm) in the cell and the temperature during irradiation were (1.3 ( 0.2) × 10-5 einstein min-1 and 50 ( 5 °C, respectively, for all illuminating devices. The ratio of the measured rate of disappearance and this photon flux gives the photon efficiency. The samples were magnetically stirred during irradiation. For the experiments in the absence of oxygen, the cells were purged with helium for 1/2 h before addition of the organic substance and subsequent irradiation. Kinetics. After the established irradiation time and filtration through 0.45 µm cellulose acetate membranes (Millipore HA), the whole sample was analyzed by HPLC for phenol and other dihydroxybenzenes. On the filtrates TOC measurements and aldehyde quantitation were performed when required. The time evolution of substrate concentration was fitted by a low first-order exponential rate (until 1/e). The fit is normally excellent until 1/2e. The initial rate was calculated by ro ) kC0, where k is the exponential decay constant and C0 is the initial substrate concentration. Several kinetic rates have been averaged over three different experiments, performed either at different times or with different illuminating devices. The reproducibility is within (15%. Measurements. Phenol and dihydroxybenzenes were detected by HPLC using a Rheodyne 7125 injector, an RP C18 column (Lichrochart, Merck, 12.5 × 0.4 cm, 5 µm packing), a high-pressure (12) Stanbury, D. M. Adv. Inorg. Chem. 1989, 33, 69-138. (13) Piccinini, P.; Minero, C.; Vincenti, M.; Pelizzetti, E. J. Chem. Soc., Faraday Trans. 1997, 93, 1993-2000.

Langmuir, Vol. 16, No. 6, 2000 2633 two-pump gradient (Merck Hitachi L-6200), and a UV-vis detector (Merck Hitachi L-4200). Formaldehyde and other aldehydes were derivatized with (2,4-dinitrophenyl)hydrazine and detected at 360 nm. The chromatographic analysis was accomplished by HPLC as above. Total organic carbon (TOC) and inorganic carbon (IC) were measured on filtered suspensions using a Shimadzu TOC-5000 analyzer (catalytic oxidation on Pt at 680 °C). Calibration was achieved by injecting standards of potassium phthalate. Reflectance spectra on powders have been obtained with a near-IR-vis-UV Perkin-Elmer Lambda 19 spectrophotometer equipped with an integrating sphere. BaSO4 was used as the reference. Measurements of cathecol and phenol adsorption have been performed by analyzing the aqueous phase. On solutions containing 0.2 g L-1 of TiO2 P25 powder kept in the dark, 2 × 10-5 M cathecol or phenol, and no salt or 10 mM NaF/5 mM HF (measured pH 3.4) or 10 mM NaNO3/HNO3 (measured pH 3.4), the cathecol or phenol concentration in the aqueous phase was measured as a function of time by HPLC as above. Measurements of adsorption have been performed after equilibration.

Photocatalytic Process The primary events occurring at the catalyst surface in the presence of an anion X- that can undergo surface complexation are summarized in reactions 1-17. The primary photochemical act, following the near-UV light absorption by TiO2 (λ < 380 nm) is the generation of electron/hole pairs in the bulk of the semiconductor (eq 1). The charge carriers can either recombine according to eq 5 or migrate rapidly to the surface where they are ultimately trapped (the electron as a surface Ti(III) and the hole as a surface radical hydroxyl group; eq 3). In the presence of surface complexation by ligands other than OH- (eq 2), trapping by these species is possible (eq 4). If electron acceptors (Ox2) or electron donors (Red1, solvent) are present at the surface (adsorbed), interfacial electron transfers may occur according to reactions 8-11 and 15. Competitive with these are the interfacial recombination (eqs 6 and 7), secondary surface trapping (eq 13), and backreactions (eqs 16 and 17). hν

charge separationTiO2 98 e-CB + h+VB

(1)

K2

surface complexationtTiOH + X- [\] tTiX + OH- (2) k3

surface trapping h+VB + tTiOH 98 tTiO• + H+ k4

h+VB + TiX 98 tTiX•+

(3) (4)

k5

recombination e-CB + h+VB 98 heat

(5) k6

e-CB + tTiO• + H+ 98 tTiOH k7

e-CB + tTiX•+ 98 tTiX

(7) k8

interfacial charge transfer h+VB + Red1 98 Ox1 k9

tTiO• + Red1 98 Ox1′ k10

tTiX•+ + Red1 98 Ox1′′ k11

(6)

h+VB + H2O 98 H2O•+

(8) (9) (10) (11)

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Minero et al.

k12

H2O•+ 98 H+ + OH•

(12)

k13

OH• + tTiOH 98 tTiO• + H2O k14

OH• + Red 98 Ox1′′

(13) (14)

k15

e-CB + Ox2 98 Red2

(15)

backreactions Red2 + h+VB k16

(or tTiO• or tTiX•+) 98 Ox2 (16) k17

Ox1 (or Ox1′ or Ox1′′) + e-CB 98 Red1

(17)

Reactions 8-10, 14 and 15 can generate different (radical) species depending on the type of primary active species and the further chemical reactions (redox, reaction with solvent or oxygen, and dimerization). In aerated systems, oxygen acts as an efficient electron scavenger in reaction 15

O2(ads) + e- f O2-ads

(18)

In the presence of organic compounds, a series of oxidation/reduction processes, both initiated through eqs 8-15, and involving other species such as O2, H2O2, and O2-, may lead to complete degradation of the organic compound to CO2 and inorganic anions.14-16 Results and Discussion Fluoride Adsorption and Surface Speciation. Coagulation and adsorption experiments, as well as literature data on XPS measurements,17 indicate that fluoride ions strongly interact with titanium dioxide. A maximum value of adsorbed fluoride, 2.5 × 10-4 mol/g of TiO2 P25, was measured.18,19 An exchangeable surface site density was then calculated to be ca. 3-4 sites/nm2.19 As already noted in previous reports concerning bidentate ligands, such as salicylic acid and catechols,20,21 white TiO2 particles turn yellow upon immersion in a catechol solution at pH 3.5, whereas the solution remains colorless. The development of the yellow color is a clear indication of the formation of charge-transfer Ti(IV)catecholate surface complexes. The correspondence with aqueous Ti(IV)-catecholate complexes has already been reported.21 The band with the maximum absorbance at 420 nm shown in Figure 1 can be assigned to the intramolecular ligand-to-metal charge-transfer transition within the inner-sphere titanium(IV)-catecholate surface complexes. However, in the presence of 0.01 M NaF the (14) Pelizzetti, E.; Minero, C. Electrochim. Acta 1993, 38, 47-55. (15) Bahnemann, D. W.; Cunningham, J.; Fox, M. A.; Pelizzetti, E.; Pichat, P.; Serpone, N. In Aquatic and Surface Photochemistry; Helz, G. R., Zepp, R. G., Crosby, D. G., Eds.; Lewis Publ.: Boca Raton, FL, 1994; pp 261-316. (16) Hoffmann, M. R.; Martin, S. T.; Choi, W.; Bahnemann, D. W. Chem. Rev. 1995, 95, 69-96. (17) Sanchez, J.; Augustynski, J. J. Electroanal. Chem. 1979, 103, 423-426. (18) Torrents, A.; Stone, A. T. Environ. Sci. Technol. 1993, 27, 10601067. (19) Vasudevan, D.; Stone, A. T. Environ. Sci. Technol. 1996, 30, 1604-1613. (20) Dagan, G.; Tomkiewicz, M. J. Phys. Chem. 1993, 97, 1265112653. (21) Rodriguez, R.; Regazzoni, A. E.; Blesa, M. A. J. Colloid Interface Sci. 1996, 177, 122-131.

Figure 1. Diffuse reflectance spectrum of TiO2 P25 (A). The differential diffuse reflectance spectrum of TiO2-catecholate surface complexes and TiO2/F-catecholate surface complexes are reported as B and C, respectively.

yellow color that develops at the TiO2 surface is strongly depressed as shown in Figure 1, suggesting a strong competition for adsorption. In pure water 100 mg L-1 of TiO2 (P25) gives a milky suspension stable for days. Light-scattering measurements upon aqueous solutions of TiO2 (P25) pointed to the existence of 1-5-µm-sized aggregates of 30-nm-sized primary particles.22,23 The addition of 0.01 M NaCl at pH 3.5 (HCl) does not modify significantly the stability, whereas 0.01 M NaF at pH 3.0 (HF) leads the suspension to settle down within 1/2 h. Supposing a simple exchange equilibrium like eq 2, the two surface acid-base equilibria (eqs 19 and 20) and the

TiOH2+ / TiOH + H+

(19)

TiOH / TiO- + H+

(20)

number balance for surface sites Cs ) [tTiOH2+] + [t TiOH ] + [tTiO-] + [tTiF], the surface coverage in the dark can be calculated. Figure 2 shows the result of the computation using pK19 ) 3.9, pK20 ) 8.7,19 pK2 ) 6.2 (referring to eq 2 where X ) F),24 and 3 × 10-5 M total exchangeable surface sites (for 100 mg L-1 of TiO2 P25). Many different values (within 2 orders of magnitude) are reported for these constants.25-27 We have adopted those recently confirmed by Stone for acid-base equilibria of surface protonation constants determined at an ionic strength of 0.1 M using the technique of extrapolation to zero surface charge7 and the value of Bohem24 for hydroxyl-fluoride exchange (pK2). Changes in these parameters (or considering that the surface potential developed by adsorbed ions may shift the apparent constants by about 1 order of magnitude) obviously modify the surface speciation. It is worth mentioning that TitF remains the dominant species in the acidic pH range and also when other reported data or model calculations are used. (22) Bickley, R. I.; Gonzales-Carreno, T.; Lees, J. S.; Palmisano, L.; Tilley, R. J. D. J. Solid State Chem. 1991, 92, 178-193. (23) Chen, H. Y.; Zahraa, O.; Bouchy, M.; Thomas, F.; Bottero, J. Y. J. Photochem. Photobiol. A 1995, 85, 179-186. (24) Herrmann, M.; Kaluza, U.; Bohem, H. P. Z. Anorg. Chem. 1970, 372, 308-313. (25) Gruebel, K.; Davis, J.; Leckie, J. Environ. Sci. Technol. 1995, 29, 586-591. (26) Augustynski, J. Struct. Bonding (Berlin) 1988, 69, 1-61. (27) Martin, S. T.; Kesselman, J. M.; Park, D. S.; Lewis, N. S.; Hoffmann, M. R. Environ. Sci. Technol. 1996, 30, 2535-2542.

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Figure 2. Surface speciation of TiO2 particles as a function of pH and fluoride concentration, calculated with the equilibrium constants pK19 ) 3.9, pK20 ) 8.7, pK2 ) 6.2, and pKa ) 3.2 for HF/F- and Cs ) 3 × 10-5 M.

Figure 3. Photocatalytic degradation of phenol (2 × 10-4 M) with and without fluoride (0.01 M). TiO2 0.1 g L-1 and pH 5.0.

Photocatalytic Degradation of Phenol. An example of phenol degradation in the presence/absence of fluoride ions at pH 5.0 is reported in Figure 3. Figure 4A reports the rate observed at pH 3.6 as a function of the fluoride concentration. As the concentration of fluoride is increased, the degradation rate increases. For fluoride concentration higher than 0.01 M, the degradation rate starts to decrease. This is partially due to an ionic strength effect, as comparisons among experiments at 0.10 M fluoride and 0.01 M fluoride + 0.09 M NaClO4 suggest. The variation of the degradation rate due to the presence of fluoride anions could be attributed to different causes. The most relevant should be the following: (i) The diminution of tTiOH as [F-] is increased will decrease the ability to trap the holes as tTiO• (OH•ads) according to eq 3. Because reaction 4 is not allowed for X ) F, this will increase h+ availability and consequently affect either the direct electron transfer to the organic electron donors or the recombination of the charge carriers. In addition, the depletion of tTiOH will inhibit the formation of surficial peroxide. Fluoride anion is wellknown as the inhibitor of the complex formation between Ti(IV) and peroxide.28 (28) Stauff, J.; Huster, H.-J. Z. Phys. Chem. (Frankfurt) 1967, 55, 39-52.

Figure 4. Rate of phenol degradation in the presence of different concentrations of NaF at pH 3.6 (A) and 0.01 M NaF at different pHs (B), relative to that at [F-] ) 0. (C) Linearization of data according to eq 22 (see text). TiO2 100 mg L-1 and initial phenol concentration 2 × 10-4 M.

(ii) The displacement of -OH by F- changes the adsorption and the surface interactions. In particular, for organics that interact with -OH groups, the extent of this interaction is largely modified, as is reported above for cathecol. The adsorption of oxygen may be affected as well documented at the gas/solid interface.29,30 Then the rate of electron scavenging by O2 and of superoxide formation will be influenced. (iii) The shift of the potential of the valence and conduction bands is due to the F- substitution of -OH.31 (iv) The surface of titania, in the presence of 0.01 M Fin the range from pH 3.5 to 6.5, has a very low charge, as calculated with the constants used for Figure 2. This may be important particularly for charged substrates or intermediates and for the possibility of interfacial electron transfer. The low value of the surface charge in the range pH 3-6 causes coagulation of the particles. This could modify the reactive surface area32 and the scattering of incident photons. (29) Munuera, G.; Navio, J. A.; Rives-Arnaud, V. In Fourth International Conference on Photochemical Conversion and Storage of Solar Energy; Rabani, J., Ed.; The Hebrew University: Jerusalem, 1982; pp 141-143. (30) Munuera, G.; Gonzales-Elipe, A. R.; Rives-Arnaud, V.; Navio, J. A.; Malet, P.; Soria, J.; Conesa, J. C.; Sanz, J. In Studies in Surface Science and Catalysis; Che, M., Bond, G. C., Eds.; Elsevier: Amsterdam, The Netherlands, 1985; Vol. 21, pp 113-125. (31) Wang, C. M.; Mallouk, T. E. J. Phys. Chem. 1990, 94, 42764280. (32) Karmakar, S.; Greene, H. L. J. Catal. 1995, 151, 394-406.

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The influence of pH in the range 2.6-6.0 was also examined. As previously reported for phenol9,10 and chlorophenols33-35 in this pH range, the observed reaction rates vary by a factor ca. 2 on naked TiO2. Measurements in the presence of fluoride showed that below pH 6 fluoride ions have a positive effect on the degradation rate of phenol. Figure 4B shows the ratios of the observed rate constants at [F-] ) 0.01 M with respect to those at [F-] ) 0 at different pHs. The bell shape correlates with that reported in Figure 2 for tTiF species, although the maximum is slightly shifted toward higher pHs. Conversely, at pH 10.0 addition of up to 0.10 M fluoride has the effect of slightly decreasing the degradation rate, from 4.2 × 10-6 to 3.0 × 10-6 M min-1 at [F-] ) 0.01 M. All of these evidences suggest that the reaction rate is strongly dependent on the surface coverage of tTiF species and that, for a surface covered by fluorine, the redox process will follow routes different from those in the presence of tTiOH species. The nature of the predominant initial interfacial reaction of holes generated in the valence band of TiO2 and migrating to the titanium dioxide/aqueous solution microinterface carrying adsorbed solutes and solvent has been extensively discussed.36,37Surface-bound hydroxyl (reaction) is a highly probable site for trapped holes. Encounters with solvent (eq 11) would also be possible, and this possibility should increase on TiO2/F because of the decreased availability of bound hydroxyl. This path should be statistically more probable than the direct electron transfer with the organics. The radical H2O+ eventually formed, when very close to the organic molecule, leads to a charge transfer along the surface/solution monolayer and yields the organic cation radical. This process is dependent on the competition with eq 12. In the following the term electron transfer regarding the organics will be used irrespective if it involves directly h+ or it is mediated through H2O+. Whereas the •OH radical can hardly leave the naked TiO2 because it rapidly reacts with surficial -OH38 (eq 13), similar evidence is not available and probably is not occurring for TiO2/F. The possibility of reaction either via eqs 9 and 14 or via eq 8 controls the reaction pathways. As far as the mechanism is concerned, hydroxyl radical addition or direct electron abstraction is undistinguishable based on the detected intermediates for phenol. Scheme 1 shows the main initial events in a photocatalytic process involving phenol. The reversibility of the formation of the dihydroxy radical was assumed on what was reported for the reaction between benzene and the hydroxyl radical (in the gas phase).39 In TiO2 photocatalysis, the evidence in support of both mechanisms has been obtained,40-46 and two recent contributions gave further support to the existence of two different types of (33) Barbeni, M.; Pramauro, P.; Pelizzetti, E.; Borgarello, E.; Gratzel, M.; Serpone, N. New J. Chem. 1984, 8, 547-552. (34) D’Oliveira, J.-C.; Al-Sayyed, G.; Pichat, P. Environ. Sci. Technol. 1990, 24, 990-996. (35) Theurich, J.; Lindner, M.; Bahnemann, D. W. Langmuir 1996, 12, 6368-6376. (36) Cunningham, J.; Sedlak, P. In Photocatalytic Purification and Treatment of Water and Air; Ollis, D. F., Al-Ekabi, H., Eds.; Elsevier: Amsterdam, The Netherlands, 1993; pp 67-81. (37) Fox, M. A. In Photocatalytic Purification and Treatment of Water and Air; Ollis, D. F., Al-Ekabi, H., Eds.; Elsevier: Amsterdam, The Netherlands, 1993; pp 163-167. (38) Lawless, D.; Meisel, D.; Serpone, N. J. Phys. Chem. 1991, 95, 5166-5170. (39) Knipsel, R.; Koch, R.; Siese, M.; Zetzsch, C. Ber. Bunsen-Ges. Phys. Chem. 1990, 94, 1375-1379. (40) Percherancier, J. P.; Chapelon, R.; Pouyet, B. J. Photochem. Photobiol. A 1995, 87, 261-266.

Minero et al. Scheme 1

traps and to the possibility of both mechanisms.47,48 Comparison of the substrate degradation rate with the increase in the current density upon substrate addition, the extent of degradation due to a path via h+ and •OHmediated can be quantified. This ratio should depend on the nature of the organic substrate. A kinetic scheme taking into account these peculiarities has been proposed for rationalizing the quantum yield of the spin-trapping adduct formation in irradiated particulate TiO2. This allowed an estimate of the ratio of the rate constant of surface •OH formation over direct oxidation of the spintrap by the hole.49 The present data in the presence of fluoride suggest the possibility that the phenol transformation proceeds through the reaction with either the free •OH (eq 14) or direct electron transfer (eq 8). Literature and the actual results can be rationalized in light of the resonance between the adsorbed hydroxyl radical and lattice -•Oat the TiO2 surface. The surface-bound hydroxyl radical is long-lived and unreactive for direct electron transfer, whereas other traps may be in a thermally activated equilibrium with free holes, exhibiting a very high oxidation potential and direct electron-transfer capability. The displacement of surface hydroxyl groups by fluoride should decrease the formation of deeply trapped holes (tTiO•), favoring less deep surface trapping. Electron transfer from these sites, or direct hole transfer to phenol, is faster because the increase in the free energy of reaction increases the rate of the electron transfer according to the Marcus theory.50,51 The effect of fluoride to shift the band potentials to more positive values (as reported in acetonitrile31) should lead to the same experimental effects. The surface-bound hydroxyl radicals, although presenting a lower reduction potential with respect to free (41) Draper, R. B.; Fox, M. A. Langmuir 1990, 6, 1396-1401. (42) Mao, Y.; Schoneich, C.; Asmus, K. D. J. Phys. Chem. 1991, 95, 10080-10089. (43) Carraway, E. R.; Hoffmann, A. J.; Hoffmann, M. R. Environ. Sci. Technol. 1994, 28, 786-793. (44) Richard, C.; Boule, P. New J. Chem. 1994, 18, 547-552. (45) Sun, Y.; Pignatello, J. J. Environ. Sci. Technol. 1995, 29, 20652072. (46) Richard, C.; Bosquet, F.; Pilichowski, J.-F. J. Photochem. Photobiol. A 1997, 108, 45-49. (47) Kesselman, J. M.; Weres, O.; Lewis, N. S.; Hoffmann, M. R. J. Phys. Chem. B 1997, 101, 2637-2643. (48) Bahnemann, D. W.; Hilgendorff, M.; Memming, R. J. Phys. Chem. B 1997, 101, 4265-4275. (49) Grela, M. A.; Coronel, M. E. J.; Colussi, A. J. J. Phys. Chem. 1996, 100, 16940-16946. (50) Eberson, L. Electron-Transfer Reactions in Organic Chemistry; Springer: Berlin, 1987. (51) Ferry, J. L.; Glaze, W. H. J. Phys. Chem. B 1998, 102, 22392244.

Photocatalytic Transformation of Organic Compounds •OH,38

react similarly, abstracting hydrogen atoms and/ or adding to the aromatic ring or unsaturated bonds. Free •OH normally does not react by one-electron transfer because of the high solvent reorganization energy associated with transfer of charge.52 Conversely, holes cannot abstract hydrogen atoms or add to the aromatic ring but can carry out electron-transfer oxidation either from aromatic structures, generating cation radicals, or, preferentially, from carboxyl over other functional groups. On titania the preference for the carboxylate group may stem from its ability to coordinate directly with underlying Ti(IV), facilitating direct electron transfer. Data are not available for comparison with the Ti-F bond. The rate at which the organic disappears is given by eq 21, where four kinetic pathways are taken into account, namely, the reaction with surface-trapped holes, with subsurface holes when the surface is covered by fluorine or hydroxyl groups, and with OH free in solution. This relationship holds when the predominant surface species are tTiOH and tTiF, namely, from inspection of Figure 2, when pH g 4 for [F-] ) 0.01 M, and in the presence of at least 1 mM fluoride ions at pH 3.6.

-d[A]/dt ) k9{Red1,O}{tTiO•} + k8,O{h+}{Red1,O} + k8,F{h+}{Red1,F} + k14[Red1][•OH] (21) where { } denotes the surface concentrations, the subscripts F and O denote adsorption on fluorinated and hydroxylated surface sites, ki the specific rate constants for reactions 8, 9, and 14 are assumed independent of pH, and [ ] are bulk solution concentrations. Two different kinetic constants have been considered for the direct electron transfer, because there can be some effect of adsorbed fluoride on the valence band in an aqueous solution or on the electron-transfer specific rate constant, as expected for these reactions, if shifts are operating.50 Under steady illumination and constant {Red1} (i.e., initial rate extrapolation), the concentrations of {h+}, {t TiO•}, and [•OH] are stationary. This allows to one assume d{tTiO•}/dt ) 0 and d[•OH]/dt ) 0. The steady-state hypothesis can also be applied to d{h+}/dt and d{e-}/dt. This complicates a lot the solution of the kinetic system, as reported elsewhere.53 Thus, for the purpose of this paper, {h+} will be considered constant. From eqs 3, 11, and 12, it can be easily demonstrated that {tTiO•} ) k3{h+}{t TiOH}/(k9{Red1,O}) and [•OH] ) k11{h+}{H2O}F/(k14[Red1]). The considered adsorption sites are tTiF, and tTiOH. Although the relationship of the fractions RF ) {tTiF}/Cs and RO ) {tTiOH}/Cs with pH is analytical but involved, they can be numerically calculated effortlessly. The surface concentration of the organic substrate may be suitably related to the solution concentration by a Langmuir isotherm. For partition of phenol, which is scarcely adsorbed, and considering the low concentrations used in these experiments, simplified relationships of the form {Red1,O} ) KOads[Red1]ROCs and {Red1,F} ) KFads[Red1]RFCs can hold. Different adsorption constants KFads and KOads are considered, because the adsorption properties of t TiOH, and tTiF sites may be different. Similarly, {H2O}F ) RFCsKH2O, where KH2O is simply a proportionality constant between the number of surface sites and the number of adsorbed water molecules. Considering that in the absence of fluoride ions only the first two right-hand-side terms of eq 21 are pertinent, (52) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J. Phys. Chem. Ref. Data 1988, 17, 513-886. (53) Minero, C. Kinetic Analysis of Photoinduced Reactions at the Water Semiconductor Interface. Catal. Today 1999, 54, 205-216.

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because the •OH radical rapidly reacts with surficial -OH38 (eq 13), the previous hypotheses transform eq 21 into eq 22 where k′8,F ) k8,FKFads, k′8,O ) k8,OKOads, and k′11 ) k11KH2O

(

)

rF k′8,F[Red1] + k′11 ) β RO + RF rO k′8,O[Red1] + k3

(22)

and β ) {h+}F/{h+}O is the ratio of stationary state concentrations of holes in the presence and absence of fluoride ions. Equation 22 is strictly valid at fixed concentrations of substrate and electron-scavenging species, as well as of light intensity, because their variation strongly influences {h+} and {e-}, as reported elsewhere.53 Although β is expected to vary with fluoride concentration, it will be considered constant. The two hole concentrations may be different as the surface speciation changes, because at a constant production rate (constant adsorbed light flux) scavenging of holes is different in the two cases (presence and absence of fluoride ions). In addition, the different aggregation of primary particles in the presence of fluorides may change the actual adsorbed light flux. However, in the limit [F-] f 0, β ) 1. Equation 22 clearly shows that the ratio rF/rO is proportional to the surface speciation of the titania particle, as is evident by comparison of Figures 2 and 4. The uncertainty about the possible value of β (see above) and RF (see the discussion referring to Figure 2) precludes a confident estimation of the ratio P ) (k′8,F[Red1] + k′11)/ (k′8,O[Red1] + k3). Equation 22 can be linearized by plotting rF/(rORO) versus RF/RO. The fit of data of parts A and B of Figure 4 using the same parameters of Figure 2 (pK19 ) 3.9, pK20 ) 8.7, pK2 ) 6.2, and pKa ) 3.2 for HF/F- and Cs ) 3 × 10-5 M) is reported in Figure 4C. The best fit of all of the experimental points lying in the range for which eq 22 was derived, assuming intercept ) β ) 1, gives P ) 3.26 ( (0.37)95% confidence. This value leads us to conclude that when the surface is covered by fluorine, the kinetic pathways for reaction with subsurface holes and with OH free in solution are predominant with respect to the reaction, when the surface is hydroxylated, with surfacetrapped holes and direct electron transfer. The experiments of Figure 4 are silent about the relative role of water-mediated oxidation (k11; see eqs 11 and 12) and direct electron transfer. In general, the kinetic effect of surface fluoride complexation will depend on the relative weights of terms k3, k′8[Red1], and k′11, which are specific for a given catalyst and organic molecule. For the relative weight of k′8,O and k3, on naked TiO2 a value between 15 and 60% for the direct electron transfer (k′8,O) with respect to the overall reaction rate for 4-chlorophenol photocatalytic degradation47 and 8% for carbetamide were recently suggested.40 Among the obvious tentative assumptions on the relative weights of the terms present in the ratio P, the hypothesis that k′8,F is almost equal to k′8,O leads to the following view. Assuming from literature a ratio k3/k′8,O ) 2 (that is, an average 30% for the direct electron transfer (k′8,O) with respect to the overall reaction rate on hydroxylated titania), it follows from the found value of P that k11/k′8,F is about 10 (that is, the direct electron transfer. (k′8,F) is at maximum 10% with respect to the overall reaction rate on fluorinated titania). As a consequence, most of the reaction would proceed through the free OH radical pathway. The shapes of TOC variation with time at different fluoride concentrations are reported in Figure 5. It may be rationalized by recognizing that compounds without an aromatic ring or a carboxylate group seem to react

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Minero et al.

Figure 6. Photocatalytic degradation of quinol and catechol on 0.1 g L-1 of TiO2 in the absence or presence of [F-] ) 0.01 M.

Figure 5. TOC disappearance as a function of the irradiation time for phenol (2 × 10-4 M) in the presence of increasing concentration of fluoride. TiO2 0.1 g L-1; (A) pH 3.6, (B) pH 5.0.

preferentially through the tTiO• pathway (eq 9). Whereas the phenol degradation rate increases with added fluoride, a little variation in the initial TOC disappearance rate is observed. At longer times an increase in the TOC disappearance is observed at pH 3.6, according to the hypothesis that carboxyl-bearing groups are more efficiently oxidized by direct electron transfer with holes.45 At pH 5.0, which is closer to pHPZC, TOC decreases slightly slower than the primary compound over the entire time scale. Catechol and quinol are the main primary nonradical intermediates detected and quantified by chromatographic analysis in the phenol degradation. Their degradation rate increases in the presence of fluoride, as shown in Figure 6. In these experiments the catechol concentration (1 × 10-4 M) is exceeding that of available sites at the catalyst surface (ca. 3 × 10-5 M). As reported above, catechol is strongly adsorbed on naked TiO223,27,54 and largely displaced by fluoride.19 However, the presence of fluorides still increases its degradation rate, implying that reaction 14 or direct electron transfer (eq 8) prevails over the availability at the catalyst surface. However, because phenol and quinol react almost with the same rate as cathecol and their adsorption onto naked TiO2 is negligible, it follows that also adsorption must be considered from a different point of view. When adsorbed, catechol and related structures, such as salicylic acid,55,56 may act also as recombination centers (eqs 8 and 9 vs eq 17, see below). (54) Connor, P. A.; Dobson, K. D.; McQuillan, A. J. Langmuir 1995, 11, 4193-4195.

Figure 7. (a) Concentration of phenol and dihydroxybenzenes during the photocatalytic process on 0.1 g L-1 of TiO2 at pH 3.6. Initial phenol concentration 1.1 × 10-3 M. (b) As in a part with [F-] ) 0.01 M. The initial rate of formation is indicated by a dotted straight line.

The time evolution of the dihydroxybenzenes is shown for the degradation of phenol on naked TiO2 and TiO2/F in parts A and B of Figure 7, respectively. To obtain more reliable experimental data, the investigation at 1 × 10-3 M phenol was carried out. The time dependence strongly resembles that reported by Okamoto et al.9 under comparable conditions. It appears that, irrespective of the (55) Tunesi, S.; Anderson, M. A. Langmuir 1992, 8, 487-495. (56) Tunesi, S.; Anderson, M. A. J. Phys. Chem. 1991, 95, 33993405.

Photocatalytic Transformation of Organic Compounds

Langmuir, Vol. 16, No. 6, 2000 2639

•OH or h+ routes, phenoxyl and dihydroxycyclohexadienyl

radicals, or the cation radical, preferentially form o- and p-dihydroxybenzene. Scheme 1 depicts quinol and catechol as the primary stable intermediates. Their initial rate of formation (see the initial tangents reported in Figure 7) would be 1:2, respectively, as required for statistical reasons. In Figure 7B this expectation is reasonably fulfilled, whereas the initial rates in Figure 7A (naked TiO2) are comparable. Because the intermediates are measured in the aqueous phase, the ratio