Adsorption and Photooxidation of Salicylic Acid on Titanium Dioxide: A

Ciudad Universitaria, 1428-Buenos Aires, Argentina. Received June 23, 1997. In Final Form: December 16, 1997. The adsorption and photooxidation of sal...
0 downloads 0 Views 255KB Size
Download PDF
868

Langmuir 1998, 14, 868-874

Adsorption and Photooxidation of Salicylic Acid on Titanium Dioxide: A Surface Complexation Description Alberto E. Regazzoni,*,† Pablo Mandelbaum,‡ Mayumi Matsuyoshi,† Sabina Schiller,‡ Sara A. Bilmes,*,‡ and Miguel A. Blesa†,‡ Unidad de Actividad Quı´mica, Comisio´ n Nacional de Energı´a Ato´ mica, Av. del Libertador 8250, 1429-Buenos Aires, Argentina, and INQUIMAE, Facultad de Ciencias Exactas y Naturales, Universidad de Buenos Aires, Pabello´ n II, Ciudad Universitaria, 1428-Buenos Aires, Argentina Received June 23, 1997. In Final Form: December 16, 1997 The adsorption and photooxidation of salicylic acid on dispersed TiO2 (Degussa P-25) particles was studied as a function of substrate concentration and pH. Salicylic acid chemisorbs at the particle interface, forming inner-sphere titanium(IV) salicylate surface complexes. The visible differential diffuse reflectance spectra of the surface complexes present a band, with maximum absorption at 420 nm, which is assigned to the internal ligand to metal charge-transfer transition. The surface excess of salicylic acid increases with decreasing pH and levels off around pKa1. At constant pH, the surface excess increases with the concentration of salicylic acid, the isotherm reflecting surface site heterogeneity. Photooxidation rates in air-saturated solutions, on the other hand, are independent of both pH and salicylic acid concentration, in the entire studied range. Chemisorption results are accounted for by a multisite surface complexation model in which two different surface titanium sites and three complexation modes are considered. The mismatch between salicylic acid surface excess values and photooxidation rates is interpreted in terms of the different reactivities of the titanium(IV) salicylate surface complexes and is attributed to the fastest hole capture by bidentate salicylate binding a single surface titanium ion. The advanced rationale illustrates the importance of the basic principles of coordination chemistry in the interpretation of apparent kinetic orders in photolyte concentration.

Introduction In pursuit of more efficient water treatment procedures, the study of light-induced oxidation reactions catalyzed by titanium dioxide has received growing attention during the past decade; ideally, only UV radiation and dissolved oxygen should be required to achieve mineralization of common pollutants. The progress has been notable, and photodegradation of dissolved organic contaminants on TiO2 slurries or films has become a promising procedure for water purification.1 The overall process involves the photogeneration of electron-hole pairs within the colloidal catalyst, followed by the competition between recombination and interfacial charge transfer to adequate acceptors. Despite the vast wealth of available information, the mechanism of photooxidation is a matter of current debate. One important issue refers to the recombination reaction, which, ultimately, determines the overall efficiency. Both classical2-4 and stochastic models5,6 have been advanced to account for electron-hole recombination kinetics. Formally, the concentration of holes capable of reacting with dissolved electron donors is one key for the final photooxidation rate law. Thus, the many assumptions involved in the description of the coupling between recombination and * To whom correspondence should be addressed. † Comisio ´ n Nacional de Energı´a Ato´mica. ‡ Universidad de Buenos Aires. (1) Legrini, O.; Oliveros, E.; Braun, A. M. Chem. Rev. 1993, 93, 671. (2) Turchi, C. S.; Ollis, D. F. J. Catal. 1990, 122, 178. (3) Kesselman, J. M.; Shreve, G. A.; Hoffmann, M. R.; Lewis, N. S. J. Phys. Chem. 1994, 98, 13385. (4) Grela, M. A.; Coronel, M. E. J.; Colussi, A. J. J. Phys. Chem. 1996, 100, 16940. (5) Rothenberger, G.; Moser, J.; Gra¨tzel, M.; Serpone, N.; Sharma, D. K. J. Am. Chem. Soc. 1985, 107, 8054. (6) Grela, M. A.; Colussi, A. J. J. Phys. Chem. 1996, 100, 18214.

interfacial charge transfer reflect themselves in the different rate expressions that can be found in the literature. Another important point of controversy refers to the mechanism of the oxidation reaction itself. Both oxidation via OH• radicals, formed upon hole trapping by adsorbed OH- or water molecules, and direct hole transfer to adsorbed reductants have been diagnosed.7-11 Whereas adsorption of the electron donor is an essential requisite for direct hole transfer to occur, oxidation via OH• radicals may however take place in the aqueous phase.2 The operation of the different mechanisms should certainly manifest in the kinetic law. Reported photooxidation rates, however, obey mostly a Langmuir-Hinshelwood dependence on photolyte concentration.12-16 Such a dependence has been usually rationalized in terms of substrate adsorption,17-19 but in many cases adsorption is far from being Langmuirian (see, e.g., refs 10 and 20). (7) Draper, R. B.; Fox, M. A. Langmuir 1990, 6, 1396. (8) Richard, C.; Boule, P. J. Photochem. Photobiol., A: Chem. 1994, 84, 151. (9) Vinodgopal, K.; Stafford, U.; Gray, K. A.; Kamat, P. V. J. Phys. Chem. 1994, 98, 6797. (10) Tunesi, S.; Anderson, M. A. J. Phys. Chem. 1991, 95, 3399. (11) Minero, C.; Catozzo, F.; Pelizzetti, E. Langmuir 1992, 8, 481. (12) Matthews, R. W. Water Res. 1986, 20, 569. (13) Matthews, R. W. J. Catal. 1988, 111, 264. (14) Matthews, R. W. J. Chem. Soc., Faraday Trans. 1 1989, 85, 1291. (15) Matthews, R. W. J. Phys. Chem. 1987, 91, 3328. (16) Sabate, J.; Anderson, M. A.; Kikkawa, H.; Edwards, M.; Hill, C. G., Jr. J. Catal. 1991, 127, 167. (17) Al-Ekabi, H.; Serpone, N. J. Phys. Chem. 1988, 92, 5726. (18) Minero, C.; Aliberti, C.; Pelizzetti, E.; Terzian, R.; Serpone, N. Langmuir 1991, 7, 928. (19) Terzian, R.; Serpone, N.; Minero, C.; Pelizzetti, E. J. Catal. 1991, 128, 352. (20) Cunningham, J.; Sedla´k, P. J. Photochem. Photobiol., A: Chem. 1994, 77, 255.

S0743-7463(97)00665-3 CCC: $15.00 © 1998 American Chemical Society Published on Web 01/30/1998

Salicylic Acid on Titanium Dioxide

The complexities underlying the interpretation of kinetic orders in reductant concentration are further illustrated by the striking mismatch between adsorption and initial photooxidation rates that has been recently reported.20,21 Noteworthily, most of the conflicting evidence appears to arise from the intrinsic nature of the unlike substrates and the widely different experimental conditions which were explored. Thus, published data must be analyzed within the context imposed by the characteristics of each photolyte. Aromatic and nonaromatic carboxylic and hydroxycarboxylic acids (e.g., salicylic and its chloro derivatives, oxalic, glycolic, etc.), as well as polyhydroxyaromatics (e.g., catechol and its chloro derivatives), constitute one important group of model photolytes. All share in common great affinity for aqueous Ti(IV) and a strong tendency to chemisorb.10,22-25 The most sound description of the chemisorption of this type of photolytes at the TiO2/aqueous solution interface is offered by the surface complexation approach.26,27 In this approach, chemisorption is described as a mass-law process, in which formation of coordinate bonds between adsorbing ligands and surface metal ions gives rise to inner-sphere surface complexes.28,29 Complexation of surface titanium ions by salicylate,30 oxalate,31 catechol,32 and 4-chlorocathechol33 has been demonstrated by in situ ATR-FTIR and CIRFTIR evidence. Surprisingly, the surface complexation approach has not yet been used to interpret kinetic orders in photolyte concentration. Aiming at assessing the role of chemisorption in the kinetics of heterogeneous photooxidation of organic acids, this work presents a detailed study of the influence of substrate concentration and pH on the adsorption and photooxidation kinetics of salicylic acid on dispersed TiO2 particles and offers a new rationale based on the multisite surface complexation description of chemisorption. Experimental Section Materials. Adsorption and photooxidation experiments were carried out using Degussa P-25, a microcrystalline titanium dioxide powder composed mainly of anatase (rutile amounts for ca. 10%). The solid, the same used in our previous study,26 was thoroughly purified by dialysis against water, vacuum-dried at room temperature, and stored in a desiccator in the dark. The BET surface area, determined from N2 adsorption data at 77 K, was 51.4 m2 g-1. The density of exchangeable OH groups, measured after equilibration in KOH solutions, was 1.8 sites nm-2. All reagents were analytical grade and used without further purification. Solutions were made up using deionized water obtained from a Milli-Q apparatus (conductivity ca. 0.1 µS cm-1). (21) Cunningham, J.; Al-Sayyed, G. J. Chem. Soc., Faraday Trans. 1990, 86, 3935. (22) Rupprecht, H. Cosmet. Toiletries 1976, 91, 30. (23) Moser, J.; Punchihewa, S.; Infelta, P. P.; Gra¨tzel, M. Langmuir 1991, 7, 3012. (24) Matsuyoshi, M.; Bilmes, A. S.; Blesa, M. A.; Regazzoni, A. E. Unpublished. (25) Ludwig, C.; Schindler, P. W. J. Colloid Interface Sci. 1995, 169, 284. (26) Rodrı´guez, R.; Blesa, M. A.; Regazzoni, A. E. J. Colloid Interface Sci. 1996, 177, 122. (27) Vasudevan, D.; Stone, A. T. Environ. Sci. Technol. 1996, 30, 1613. (28) Stumm, W. Chemistry of the Solid-Water Interface; WileyInterscience: New York, 1992. (29) Blesa, M. A.; Morando, P. J.; Regazzoni, A. E. Chemical Dissolution of Metal Oxides; CRC Press: Boca Raton, FL, 1994. (30) Tunesi, S.; Anderson, M. A. Langmuir 1992, 8, 487. (31) Hug, S. J.; Sulzberger, B. Langmuir 1994, 10, 3587. (32) Connor, P. A.; Dobson, K. D.; McQuillan, A. J. Langmuir 1995, 11, 4193. (33) Martin, S. T.; Kesselman, J. M.; Park, D. S.; Lewis, N. S.; Hoffmann, M. R. Environ. Sci. Technol. 1996, 30, 2535.

Langmuir, Vol. 14, No. 4, 1998 869 Adsorption Measurements. All adsorption experiments were performed in the dark at 298.0 ( 0.2 K using a thermostated borosilicate amber-glass vessel. Measurements were carried out both on oxygen-purged (under a O2-free N2 stream) and aerated suspensions. During the experiments, particles were kept in suspension by means of magnetic stirring; to break aggregates, suspensions were briefly ultrasonicated using a titanium probe. In a typical experiment, 0.25 g of TiO2 was suspended in 25 cm3 of freshly prepared salicylic acid solution of a given concentration, and the pH was adjusted by addition of NaOH or HCl. During the experiment, pH was periodically monitored and readjusted when necessary. After a preset equilibration time (1 h) the solid was filtered off using a 0.2-µm pore size polycarbonate membrane and the supernatant stored; a separate set of experiments showed that adsorption equilibrium (i.e., no further variation of the residual salicylic acid concentration) is attained within 15 min. The analytical concentration of salicylic acid in the filtrates was measured photometrically at 296 nm ( ) 3530 mol-1 dm3 cm-1) after pH adjustment to ca. 5.0, and the surface excess, Γ, was calculated according to

Γ ) (V/SSw)(Ci - Ceq)

(1)

where V is the suspension volume, SS is the specific surface area, w is the mass of solid, and Ci and Ceq are the initial and equilibrium analytical concentrations of salicylic acid, respectively. In selected cases, the UV-visible diffuse reflectance spectra of filtered solids were recorded after drying under vacuum in the dark. Photooxidation Experiments. Experiments were carried out in a photochemical batch reactor composed of a borosilicate glass rectangular cell placed in an aluminum holder, a 150 W XBO Osram Xe lamp fitted in a high-collection-efficiency housing, and an IR CuSO4 solution filter. The incident spectrum was therefore a continuum with cutoff edges at 300 and 700 nm; thus, direct photolysis of salicylic acid by low-wavelength UV radiation was avoided. Experiments were performed at 299.0 ( 0.3 K as follows: 0.025 g of TiO2 was suspended in 25 cm3 of freshly prepared salicylic acid solution of known concentration and pH and left standing in the reaction cell under a continuous flow of O2 for 1 h; magnetic stirring kept particles in suspension. Before starting the photochemical reaction, 10 cm3 of this suspension was pipetted off and the concentration of dissolved salicylic acid at zero reaction time (C0) measured after filtration. The reaction was then triggered by opening a shutter; O2 bubbling throughout the experiment maintained a constant dissolved oxygen concentration of 2.4 × 10-4 mol dm-3. After prescheduled time intervals, the shutter was closed, the cell removed, and the residual salicylic acid concentration determined as before; the spectra of the supernatant solutions did not show any interference from possible intermediates. pH remained constant (within (0.1 units) during the kinetic runs. Due care was taken to ensure constant geometry and light intensity during the course of the experiments; constancy of light intensity at the cell position was controlled using a beam splitter and a calibrated photodiode. The photon flux reaching the reaction volume (15 cm3), measured by trisoxalateferrate(III) actinometry,34 was 6.2 × 10-7 einstein s-1, taking 1.2 as the average quantum yield.

Results and Discussion Chemisorption of Salicylic Acid onto P-25 TiO2. As already noted by Tunesi and Anderson10 and by Dagan and Tomkiewicz,35 white P-25 TiO2 particles turn yellow upon immersion in salicylic acid solutions; it is noteworthy that supernatant solutions are colorless. The development of the yellow color is a clear indication of the formation of charge transfer titanium(IV) salicylate surface com(34) Kuhn, H. J.; Braslavsky, S. E.; Schmidt, R. Pure Appl. Chem. 1989, 61, 187. (35) Dagan, G.; Tomkiewicz, M. J. Phys. Chem. 1993, 97, 12651.

870 Langmuir, Vol. 14, No. 4, 1998

Figure 1. Differential diffuse reflectance spectrum of surface titanium(IV) salicylate complexes; in arbitrary units.

Figure 2. Salicylic acid surface excess plotted as a function of equilibrium pH: total salicylic acid concentration, 8.0 × 10-4 mol dm-3; surface-to-volume ratio, 514 m2 dm-3; T ) 298 K.

plexes. Indeed, the differential diffuse reflectance spectrum shown in Figure 1 closely resembles the UV-visible spectra of aqueous titanium(IV) salicylate complexes.36 The same correspondence between analogous surface and aqueous complexes has been observed for titanium(IV) catecholate and iron(III) thiocyanate species.26,37 The band in Figure 1, with maximum absorption at 420 nm, can be assigned to the intramolecular ligand-to-metal charge transfer (LMCT) transition within the inner-sphere titanium(IV) salicylate surface complexes. The overall affinity of salicylic acid for the surface titanium ions of P-25 TiO2 is strongly dependent on both pH and free ligand concentration. The pH dependence of the overall affinity is illustrated by Figure 2, which depicts the variation of the salicylic acid surface excess with equilibrium pH for experiments performed at a constant total salicylic acid concentration (Ci ) 8.0 × 10-4 mol dm-3) and at a fixed surface-to-volume ratio (SSw/V ) 514 m2 dm-3). In agreement with previous reports on salicylic acid adsorption onto powdered TiO2 calcined membranes10 and P-25 TiO2,22 Γ increases with decreasing pH and levels off in the range 2.5-3.1, that is, around the first pKa of salicylic acid (pKa1 ) 2.97).38 The observed trend is typical of complexing anion chemisorption and results, mainly, from the coupling of homogeneous and heterogeneous protolytic equilibria with ligand-exchange surface reactions (see below).28,29 The dependence of Γ on the equilibrium concentration of salicylic acid at pH 3.6 is presented in Figure 3, which includes data for oxygen-purged and air-saturated sys(36) Hultquist, A. E. Anal. Chem. 1964, 36, 149. (37) Regazzoni, A. E.; Blesa, M. A. Langmuir 1991, 7, 1652. (38) Martell, A. E.; Smith, R. M. Critical Stability Constants; Plenum: New York, 1976; Vol. 3.

Regazzoni et al.

Figure 3. Salicylic acid adsorption isotherm at pH 3.6: (b) O2-free solutions; (O) air-saturated systems; T ) 298 K.

tems. As expected, dissolved oxygen has no influence on salicylic acid chemisorption. The shape of the Γ versus Ceq plot is representative of high-affinity adsorption isotherms and cannot be described by the simple Langmuir model; in fact, a poor fit, which significantly underestimates Γ values at Ceq e 10-4 mol dm-3, was obtained when fitting the data to the Langmuir isotherm. High-affinity adsorption isotherms, such as that shown in Figure 3, are a consequence of surface site heterogeneity, with the more energetic adsorption sites saturating at the lower substrate concentrations. Similar results were reported by Tunesi and Anderson,10 who found a Freundlich isotherm to be the best representation for the adsorption of salicylic and 3-chlorosalicylic acids onto powdered TiO2 membranes. In seeming discrepancy, data reported by Cunningham and Al-Sayyed21 comply with a Langmuir adsorption isotherm, but they rather reflect the simultaneous influence of Ceq and pH on Γ (cf. Figures 2 and 3). Multisite Surface Complexation Model. In our previous description of surface complexation at the TiO2(anatase)/aqueous solution interface,26 the particle surface is represented as an array of two types of sites, that is, ∼VTiOH1/3- and >OH1/3+ (for clarity, cf. below, a different notation is used here). These are the sites formed upon dissociative water chemisorption onto the most stable, perfectly cleaved, (001) and (011) crystal faces of anatase; ∼VTiOH1/3- (A sites) are surface titanium ions, pentacoordinated to oxo anions, that complete their coordination sphere binding OH-, whereas >OH1/3+ (B sites) are surface O2- ions, dicoordinated to titanium ions, that bind protons (see Figure 4).39-41 Both A and B sites undergo surface protonation-deprotonation reactions, but, as shown earlier,26 only A sites are prone to surface complexation. Even though this description suffices to account for catechol chemisorption, from the point of view of surface complexation, this is a single-site model. In contrast, the high-affinity adsorption isotherm shown in Figure 3 requires, at least, two different chemisorption sites. In fact, our previous model overlooks the intervention of surface titanium ions with fewer coordination positions occupied by O2- (say, those at edges and corners), which, as suggested by Tunesi and Anderson,30 should display the largest affinity for the adsorbing ligands. In what follows, to account for the chemisorption behavior of salicylic acid, we adopt a multisite surface complexation (39) (a) Hiemstra, T.; van Riemsdijk, W. H.; Bolt, G. H. J. Colloid Interface Sci. 1989, 133, 91. (b) Hiemstra, T.; de Wit, J. C. M.; van Riemsdijk, W. H. J. Colloid Interface Sci. 1989, 133, 105. (40) Boehm, H. P. Discuss. Faraday Soc. 1971, 52, 264. (41) Dore´mieux-Morin, C.; Enriquez, M. A.; Sanz, J.; Fraissard, J. J. Colloid Interface Sci. 1983, 95, 502.

Salicylic Acid on Titanium Dioxide

Langmuir, Vol. 14, No. 4, 1998 871

(4)

Note that titanium ions in species I are heptacoordinated, whereas those in species III are hexa-coordinated. The corresponding equilibrium constants are given by

Kint I )

{I}[H+] {∼VTiOH1/3-}[H2L] Kint II )

Kint III )

Figure 4. (A) Illustration of the freshly, vacuum-cleaved, (001) anatase perfect surface; dashed lines denote dangling bonds on titanium ions. (B) Sketch representation of water chemisorption (hydroxylation) on the perfectly cleaved (001) anatase surface; the lower row depicts a CBABABC array of sites (see text).

exp(-Fψ0/RT)

{II} {∼VTiOH1/3-}2[H2L] {III} {∼IVTi(OH)2/3}H2L 2

(5)

(6)

(7)

where ψ0 is the surface potential, { } denotes surface concentration expressed in mol m-2, and H2L represents salicylic acid. The surface complexation mode depicted by eq 2 was found to be essential to account for catechol and 4-chlorocatechol chemisorption.26,33 However, if species I are formed, the surface titanium ions in species I must adopt the uncommon seven coordination. In view of the small size of Ti4+, this seems quite unlikely. Therefore, a more plausible description would involve the intervention of a B site adjacent to ∼VTiOH1/3- to yield species I′,

(8)

model in which the latter type of surface titanium ions are recognized. These sites, hereafter C sites, are assumed to be 4-fold coordinated to O2- and are denoted by ∼IVTi(OH)(OH2)1/3+ and/or ∼IVTi(OH)22/3- (Figure 4b). Since C sites represent defect positions, we further assume that they are only a small fraction of the total number of surface titanium ions, NA + NC. This sum equals NS, the number of dissociatively chemisorbed water molecules (Figure 4b), which is the measured density of exchangeable hydroxyl groups.26 The total number of B sites, NB ()NA + NC/2), is given by the electroneutrality condition of the dehydrated surface. According to refs 26 and 33, complexation of A surface titanium ions by salicylate gives rise to two different surface complexes, in which the dianion coordinates either one (species I) or two neighboring titanium ions (species II),

(2)

in which surface titanium ions retain their normal six coordination. Equations 2 and 8 are entirely equivalent; actually, only the surface concentration of ‘free’ >OH1/3+ is altered. As reaction 8 is driven by the affinity of salicylate toward surface titanium ions, the mass-law equation given by eq 5 also accounts for the dependence of {I′} on salicylic acid concentration and pH. Coupled with surface complexation, the heterogeneous protolytic equilibria involving the A, B, and C sites are described by the following mass-law expressions:

Kint A )

{∼VTiOH1/3-}[H+] {∼VTiOH22/3+}

exp(-Fψ0/RT) ) 4.17 × 10-6 mol dm-3 (9)

1/Kint B )

{>O2/3-}[H+] {>OH1/3+}

exp(-Fψ0/RT) ) 2.51 × 10-8 mol dm-3 (10)

Kint C ) (3)

Chemisorption onto C sites, on the other hand, yields species III.

{∼IVTi(OH)22/3-}[H+] {∼IVTi(OH)(OH2)1/3+}

exp(-Fψ0/RT) )

2.82 × 10-7 mol dm-3 (11) int int 26 Kint A and KB values were derived earlier, while KC was determined by fitting our previously reported σ0 versus pH and ζ versus pH data26 using the Stern-Grahame-

872 Langmuir, Vol. 14, No. 4, 1998

Regazzoni et al.

Table 1. Set of Equations Describing the TiO2(anatase)/ Solution Interface

Figure 5. Speciation of surface titanium(IV) salicylate complexes as a function of the equilibrium concentration of salicylic acid at pH 3.6; the dashed line is the sum {I′} + {II} + {III}; symbols are the experimental Γ values given in Figure 3.

Table 2. Site Densities and SGGC Model Parameters NS/nm-2 1.80a a

NC/nm-2

K(/mol-1 dm3

C1/F m-2

C2/F m-2

0.15

1.41a

1.40b

0.23c

Reference 26. b Reference 42. c Reference 43.

Gouy-Chapman model outlined in Table 1 and the parameters listed in Table 2; capacities are literature values.42,43 The best description of the chemisorption behavior of salicylic acid onto P-25 TiO2 is obtained with pKI′int ) 3.3, int pKint II ) - 2.9, and pKIII ) - 6.5. Figures 5 and 6, which compare calculated Γ values (i.e., {I′} + {II} + {III}) with experimental data points, illustrate the excellent performance of the proposed model. These figures also show the speciation of titanium salicylate surface complexes in the framework of our experimental conditions. The largest affinity of salicylate for the C titanium sites is revealed by the initially fast increase of {III} with Ceq (Figure 5) and its marked insensitivity toward pH (Figure 6). Despite the large Kint III value, species III dominate the chemisorption behavior only at low Ceq and relatively high pH, a clear consequence of the low NC/NA ratio. As expected, species I′ contribute only modestly to the overall surface excess. The present model, incidentally, improves our previous description of catechol chemisorption. Chemisorption Kinetics. Adsorption of salicylic acid onto P-25 TiO2 is a rather slow process that takes approximately 15 min to reach equilibrium. This time, which coincides with the equilibration time found earlier for catechol,26 is shorter than that reported by other authors (ca. 1 h);21,35 salicylic acid adsorption onto powdered TiO2 membranes10 and aerogels35 is even slower. In these latter cases, and also in ours, the time scale of the overall chemisorption reaction is longer than that needed to attain equilibrated adsorption on, say, iron(III) (42) Yates, D. E.; Levine, S.; Healy, T. W. J. Chem. Soc., Faraday Trans. 1 1974, 70, 1807. (43) Foissy, A.; M’Pandou, A.; Lamarche, J. M.; Jaffrezic-Renault, N. Colloids Surf. 1982, 5, 363.

Figure 6. Speciation of surface titanium(IV) salicylate complexes as a function of pH; the dashed line is the sum {I′} + {II} + {III}; symbols are the experimental Γ values given in Figure 2.

oxides and certainly much longer than the time required to equilibrate surface protolytic reactions (see, e.g. ref 29 and references therein). In some cases, surface complexation has been measured to be diffusion-controlled, whereas, in others, the slower rates revealed a kinetic barrier by the surface substitution reaction. In the case of salicylic acid chemisorption on P-25 TiO2, it seems reasonable to assume that ring closure (i.e., formation of species II), following an unfavorable fast monodentate equilibrium (i.e., formation of species I′), determines the observed overall sluggishness. Formation of species III, on the other hand, is likely to be much faster and, probably, close to diffusion control. Noteworthily, the yellow color (cf. Figure 1) develops as soon as salicylic acid gets in contact with TiO2. Heterogeneous Photooxidation Kinetics. Under the experimental conditions explored in the present work, photooxidation rates (R ) -dCt/dt) are constant. This is shown in Figure 7, which depicts the decay of salicylic acid concentration during three typical photooxidation runs. Linear, or nearly linear, Ct versus t plots are not rare in the literature. In fact, such a time dependency has already been observed during photooxidation of salicylic acid on powdered TiO2 membranes10 and P-25 TiO2;35 similar results were also reported for the photooxidation of 3-chloro-4-hydroxybenzoic acid on P-25 TiO2.21 Linear Ct versus t profiles indicate, in principle, zeroorder kinetics on photolyte concentration. Noteworthy, observed rates (Figure 7) are almost insensitive toward changes in C0, the initial concentration of salicylic acid. This zero-order dependence is better seen in Figure 8a,

Salicylic Acid on Titanium Dioxide

Langmuir, Vol. 14, No. 4, 1998 873

Reactive Surface Species. After the photogeneration of electron-hole pairs, those holes which survive recombination can react with salicylic acid through two possible reaction pathways: (i) hole capture by surface hydroxyls followed by OH• attack to the organic solute, for example, (12)

and/or (ii) direct, inner-sphere, electron injection by chemisorbed salicylate, for example,

Figure 7. Residual concentration of salicylic acid as a function of irradiation time: (b) pH 3.6; (O) pH 5.0; T ) 299 K; surfaceto-volume ratio, 51.4 m2 dm-3.

(13)

The overall photooxidation rate is thus given by

R ) ROH + RIS

Figure 8. Photooxidation rates of salicylic acid plotted as a function of the initial photolyte concentration (A) and as a function of salicylic acid surface excess (B): pH 3.6; T ) 299 K; surface-to-volume ratio, 51.4 m2 dm-3; part B includes data for other pH values.

which shows that, at pH 3.6, measured photooxidation rates oscillate randomly around an average value of 5.2 × 10-6 mol dm-3 min-1. Despite the rather large scattering of the data, the observed trend is unquestionable: within the concentration range 4.5 × 10-5 to 1.56 × 10-3 mol dm-3, R is independent of salicylic acid concentration. As shown in Figure 7, the photooxidation rate is also independent of pH. This was further confirmed by experiments carried out at a fixed total salicylic acid initial concentration (8 × 10-4 mol dm-3) in the pH range 2.05.0. They yielded R values that were within those presented in Figure 8a. In agreement with the earlier observation by Cunningham and Al-Sayyed,21 comparison of rate data (Figures 7 and 8a) with dark adsorption results (Figures 2 and 3) shows that, under the present conditions, photooxidation rates have no correspondence with Γ. This is readily seen in Figure 8b, where all measured R values are plotted as a function of calculated Γ values; note that the low surfaceto-volume ratio used in the photochemical runs precluded the meaningful experimental determination of Γ. The insensitivity of photooxidation rates toward substrate concentration (or adsorption density) was interpreted by Cunningham and Sedla´k20 as an indication of rate control by a step which either precedes or succeeds interfacial charge transfer. Here, we will forward a different rationale which applies to photolytes that chemisorb.

(14)

where the first rate term embodies the contributions due to all bound hydroxyls and RIS includes the contributions due to all titanium(IV) salicylate inner-sphere surface complexes. The contributions of the different charge-transfer reaction pathways to the photooxidation rate depend on both the energetic levels of the surface-coordinated species and their surface concentrations; note that surface complexes are the chemical counterparts of surface states. Since the depth of surface hole traps must be determined by the Lewis basicity of the ligands that complete the coordination sphere of the surface titanium ions, chemisorbed salicylate must be a deeper hole trap than bound OH-. Likewise, hole trapping by chemisorbed salicylate is expected to be kinetically less reversible. In agreement, the observed LMCT absorption band (Figure 1) favors the easier hole capture by chemisorbed salicylate. Therefore, as already suggested by Tunesi and Anderson,10 direct hole transfer must be the prevailing salicylic acid photooxidation path, that is,

R = RIS ) (SSw/V)(kI′{I′} + kII{II} + kIII{III})[h+] (15) Whenever chemisorption is null, for example, in alkaline media, photooxidation must proceed via the OH• path.10 In principle, the reactivities of the different surface complexes, that is, the corresponding k values, should also differ. Among the three titanium salicylate surface complexes, species III is expected to be the most reactive one, for its bidentate configuration should allow for the largest electron donation. In fact, comparison of photooxidation rates (Figure 8) with surface speciation diagrams (see Figure 5) shows a clear correspondence between R and {III}, only. Thus, eq 15 reduces to

R ) k′III{III}[h+]

(16)

where k′III embodies SSw/V. If, as suggested by Kesselman et al.,3 both oxygen reduction and salicylic acid photooxidation proceed at the same rate, eq 16 yields

R ) (k′III{III})1/2(ke(O2)Iabs/krec)1/2

(17)

where ke(O2) is the (oxygen dependent) cathodic rate

874 Langmuir, Vol. 14, No. 4, 1998

Regazzoni et al.

volume. At constant oxygen partial pressure and Iabs, that is, under our experimental conditions, eq 17 predicts a half-order dependence on {III}. Alternatively, if we assume that under such conditions the number of holes surviving recombination is a constant (cf. ref 18), the photooxidation rate is given by

R ) k′′III{III}

(18)

where k′′III is a function of Iabs and oxygen partial pressure. In either case, the noted independence of R with salicylic acid concentration and pH stems from the saturation of C titanium sites by chemisorbed salicylate, viz, the apparent zero kinetic order results from the constancy of {III} within the studied pH and concentration ranges (Figure 9a). The proposed model, which is based upon the basic concepts of coordination chemistry implicit in the surface complexation approach, accounts very well for the photooxidation kinetics of salicylic acid on TiO2. Notwithstanding, eq 18, a suitable modification of the LangmuirHinshelwood expression, also predicts first-order, or nearly first-order, kinetics for very diluted and/or near neutral solutions (Figure 9b), as has already been reported in the literature for salicylic13-15 and 3-chlorosalicylic acids.16

Figure 9. Surface concentration of species III (A) and predicted apparent kinetic order in photolyte concentration (B) as a function of pH and salicylic acid concentration.

constant, krec is the second-order recombination rate constant, and Iabs is the photon absorption rate per unit

Acknowledgment. This work was partially funded by Universidad de Buenos Aires (UBACyT X022 and X803), Fundacio´n Antorchas, Volkswagen Stiftung, and Gesellschaft fur Technische Zusammenarbeit. A.E.R., M.A.B., and S.A.B. are members of CONICET. UBA scholarships to S.S. and P.M. are gratefully acknowledged. LA970665N