J. Phys. Chem. 1991, 95, 3399-3405
3399
Influence of Chemlsorption on the Photodecomposition of Salicylic Acid and Related Compounds Using Suspended TiO, Ceramic Membranes S.Tunesi and M. Anderson* Water Chemistry Program, University of Wisconsin, Madison, Wisconsin 53706 (Received: March 20, 1990; In Final Form: October 29, 1990)
The adsorption and photodecomposition on semiconducting membranes of molecules with different functional groups were studied. Ceramic membranes, prepared by a sol-gel technique, were employed because this technique provides some control over surface area, porosity, and crystal form of these materials, properties which affect their photochemical behavior. Salicylic acid was chosen as a prototype molecule. 3-Chlorosalicylic acid, benzoic acid, phenol, and 4-chlorophenol were also investigated. The adsorption densities of salicylic and 3-chlorosalicylic acids decreased with both increasing solution pH and membrane firing temperature, in correlation with the number of positive adsorption sites on Ti02surface. No adsorption for benzoic acid, phenol, and 4-chlorophenol was observed. Photodecomposition rates were found to depend on the adsorption characteristics of the organic compound; for salicylic acid the degradation rate diminished with increasing pH. Benzoic acid, whose degradation products adsorb on the oxide surface, showed a similar trend. No dependence on pH was detected for phenols. Methanol was found to affect the degradation rate of salicylic acid only under conditions of high pH. It is proposed that, for chemisorbing organics, the initial step of the photodecompositionprocess is governed by two different mechanisms that depend on the adsorption behavior of the organic compound at a particular pH. For adsorbed salicylate, an orbital configuration of the chelate ring is proposed to interpret direct electron transfer from the organic molecule to the semiconductor.
Introduction Several mechanisms have been proposed to account for the initial steps of semiconductor-mediated photodegradation of aliphatic and aromatic organics. The heterogeneous reaction mechanisms proposed are similar to their homogeneous counIn aqueous solutions of benzene derivatives and terparts.'-" aliphatic hydrocarbons, these mechanisms can be summarized as (i) direct charge transfer from the semiconductor to the dissolved molecule4
@ %@ Ti0 +h+
and3 RCHOH
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-
-
Degradation Paths
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(1)
or (ii) generation of radicals from water decomposition which then attack the aromatic ring4 H
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and4 RCHZCH3
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Similar degradation pathways have been proposed to be effective for the heterogeneous photodegradation of environmentally damaging compounds,' when using titanium dioxide. In addition, the Ti02particulate semiconductors have proven useful for organic ( I ) Barbeni, M.; Minero. C.; Pellizzetti, E.; Borgarello, E.; Serpone, N. Chemosphere 1987, 16, 2225 and bibliography therein. (2) Wallinn. C. Acc. Chem. Res. 1975. 8. 125. (3) Kormain, C.; Banhemann, D. W.:Hoffman, M. R. Enuiron. Sci. Technol. 1988. 22. 798. (4) Hashimoto,'K.; Kawai, T.; Sakata, T. J . Phys. Chem. 1984,88,4083. ( 5 ) (a) Barkni, M.; Remauro, E.; Pellizzetti, E.; Borgarello, E.; Serpone, N. Chemosphere 1985, 14, 195. (b) Barbeni, M.; Premauro, E.; Pellizetti, E.; Borgarello, E.; Serpone, N.; Jamieson, M. A. Chemosphere 1986, 15, 1913. (c) Tunesi, S.; Andenon, M. A. Chemosphere 1987,16, 1447. (d) AI-Ekabi, H.; Serpone, N.; Pellizzetti, E.; Minero, C.; Fox, M. A.; Draper, R. B. Langmuir 1989,5,250. .(e) Pruden, A. L.; Ollis, D. F.Enuiron. Sci. Technol. 1983, 17, 628. (0 Ollis, D. F. Enuiron. Sei. Technol. 1985, 19 ( 6 ) , 480.
0022-365419112095-3399$02.50/0
synthesk6 In related semiconductor studies concerning the photosplitting of water using Ti02photocatalysts, emphasis has been placed mainly upon optimizing the energy requirements for this last p r o ~ e s s . ~ *In ~ ~all* of these systems, it is important to consider the mechanisms controlling the kinetics of these processes at the Ti02/aqueous solution i n t e r f a ~ e . ~ . ' ~ Ceramic membranes, composed of TiOz particulates, offer the same advantage as the suspended colloids in that the membranes have similar photoactivities, surface areas, the same crystalline phase, and the same high capacities for adsorbing organics. These ceramic substrates can be permanently fvted onto reactor surfaces, such that there is no attrition of the catalyst." The advantage offered by the sol-gel technique, employed in obtaining our photocatalytic substrate, residues in the fact that crystal morphology, surface area, and porosity can be controlled by varying initial reactant ratios and firing conditions. Through these variables it is possible to affect adsorption and diffusion, parameters that directly influence photodegradation rates. In this paper, we investigate the role played by surface phenomena in determining photodegradation kinetics both for organic molecules which exhibit pH-dependent adsorption behavior in aqueous TiOz suspensions and for organic solutes which do nor adsorb on the hydrous oxide surface. It is rarely possible to predict the interaction between an oxide surface and a given solute solely from their crystal field and chemical structures. Nor has it been possible to a priori predict bond strengths in these systems. Adsorption on Ti02 catalytic substrates has been investigated in the gas phase, but the presence of water a t the oxide surface strongly affects the bonding behavior of the compounds to be photooxidized. In fact, water and hydroxyl groups adsorb on Ti sites at the crystal surface; in chemisorption these ligands are exchanged with adsorbing solutes, whereas in physisorption there is no such direct interaction between the solid and the dissolved species. ( 6 ) Fox, M. A.; Chen, C. C. J . Am. Chem. Soc. 1981, 103,6757. ( 7 ) Tafalla, D.; Salvador, P. J . Electroanal. Chem. 1987, 237, 225. (8) Munera, G.; Rives-Arnau, V.; Saucedo, A. J . Chem. Soc., Faraday Trans. I 1979, 75, 736. ( 9 ) (a) AI-Ekabi, H.; de Mayo, P. J . Phys. Chem. 1985, 89, 5815. (b) AI-Ekabi, H.; Serpone, N. J . Phys. Chem. 1988, 92, 5726.
(IO) (a) Okamoto, K.; Yamamoto, Y.; Tanaka, H.; Itaya, A. Bull. Chem. Soc. Jpn. 1985, 58, 2023. (b) Okamoto, K.; Yamamoto, Y.; Tanaka, H.; Itaya, A. Bull. Chem. SOC.Jpn. 1985, 58, 2015.
( 1 1 ) Sabate, J.; Anderson, M. A.; Hill, C. G.; Kikkawa, H. J . Catal., submitted for publication.
0 1991 American Chemical Society
3400 The Journal of Physical Chemistry, Vol. 95, No. 8, 1991
The amount of adsorption of organic compounds on Ti02 ceramic membranes is studied here as a function of pH. By combining these adsorption results with spectroscopic data, it can be possible to determine whether chemical bonding affects photodegradation. Salicylate was chosen as a prototype molecule for its suitable stereochemical configuration; the other organic compounds investigated may be considered as derivatives of the salicylate structure. Methods Membrane Preparation. The Ti02 suspensions (sols) and membranes were synthesized, using molar ratios of 0.4:1:150 for H+/Ti/H20, according to the procedure outlined by Anderson et al.I2 After drying the titanium dioxide sol in a 60 OC oven, we obtained a thin transparent xerogel. Changes in surface area and crystallographic structure were obtained by firing this unsupported xerogel for 4 h at different temperatures from 325 to 550 f 10 'C, at a heating rate of 1 "C/min. This firing cycle produced Ti02membranes having a range of porosities and surface areas that were characterized by BET analysis using N2 In many experiments, these membranes were crushed with a mortar and pestle prior to use. Crushed membranes typically displayed a 10% increase in surface area. Crystallographic analysis was performed on crushed membranes by using powder X-ray diffraction. The isoelectric pH (pHIEP)of a crushed membrane was measured with a Pen Kem System 3000 microelectrophoresis instrument and found to be 6.0 f 0.2 for a sample fired at 400 OC. Adsorption Studies. In all adsorption studies (Le,, equilibrium and kinetic), we used crushed membranes. This procedure ensured rapid and complete equilibration, and it enabled the use of CIRFTIR spectroscopy as an in situ technique to identify adsorbed species. It is assumed that the behavior of a supported membrane, i.e., one coated on a supporting substrate prior to firing, would be comparable to the behavior of a crushed unsupported one. Stock suspensions of titanium dioxide crushed membranes were prepared by adding deionized Milli-Q water to the solids, to give a concentration of 13 g/L. This suspension was allowed to equilibrate 1 week, thereby assuring complete hydrolysis of the Ti02surface. After this equilibration period, the stock suspension was diluted with additional deionized water to 1.3 g/L. These diluted suspensions (50 mL) were used to establish adsorption isotherms. In order to ensure reproducibility and minimize errors associated with a possible solids concentration effect, we used the same stock suspension of crushed membranes for membranes fired at a specific temperature. Kinetic Adsorption Studies. Kinetics studies of the adsorption of salicylate and related compounds on Ti02were begun by adding 100 pL of the particular organic compound from a 20 mM stock solution to the Ti02suspension. This initial amount was chosen to avoid the total saturation of surface sites. All other conditions of this experiment were identical with the equilibrium adsorption isotherm studies discussed below. Reactions were quenched by filtration, using a Nuclepore apparatus with 0.05-pm polycarbonate filters. Equilibrium Isotherms. Adsorption isotherms were obtained at 23 f 2 OC. In order to study reproducibility, the second and fifth data points were replicates, and the results always agreed to within 5%. All equilibrium adsorption studies were performed with 50 mL of a 1.3 g TiO,/L suspension. The suspension was equilibrated overnight in a temperature-controlled shaker (New Brunswick). The appropriate additions of acid (HCl), or base (KOH), and inert electrolyte (KCI) were made to adjust pH and to buffer ionic strength, respectively. All organic adsorbates used in this study (salicylic acid, 3-chlorosalicylic acid, benzoic acid, phenol, and 4-chlorophenol) 'were obtained from Aldrich and used without any further purification. Adsorbates were added from 20 mM stock solutions. After this addition, when needed, the suspension pH was adjusted with HCI or KOH (0.1 M). The samples were kept in the dark for the entire length of the ex-
Tunesi and Anderson TABLE I: Effect of Firing Tempemtwe on the Surface Area and Crystalline Phase of Titania Membranes crvstalloarsphic form' firing temp, O C anatase rutile surface area, m2/g
60 350 315 400 425
am
88
W
m S
ww
S
W
450
S
W
415 500 525 550
S
W
W
s
76
66
S
S
' s = strong signal, m = medium signal, w = weak signal, ww = very weak, and am = amorphous.
periments in order to avoid photoexcitation of the Ti02. We selected two adsorption times, 5 and 20 h, at which to evaluate equilibration, pH variation, solute diffusion kinetics, and dark degradation. Analysis of the compounds was performed with a Varian DNM8O UV spectrophotometer (detection limits: salicylic acid and 3-chlorosalicylic acid, 0.7 X lod M; benzoic acid, 0.4 X lod M). Photodegradation Studies. Having determined the role of surface area on adsorption, we next used the membranes fired at 450 OC to perform photodegradation studies. This firing temperature was chosen because, at this temperature, anatase is the predominant crystalline phase. This phase has been shown previously to be the most favorable for photodecomposition of water and photodegradation of organic solutes in aqueous Ti02 suspension^.'^ A stock suspension containing 0.68 g/L of crushed Ti02 membrane was prepared in a single batch of approximately 20 L, and all comparisons were made using the same batch. In each experiment, 600 mL of a stock suspension was adjusted to the desired pH and ionic strength and equilibrated overnight. The organic solute was added from a 20 mM stock solution 1 h before the degradation experiments in order to obtain nearly complete adsorption on the Ti02surface, as was previously determined in the adsorption studies. All degradation rates were determined at three different pH values (pH = 4, 7,9.5), maintaining constant ionic strength M, KCI). Considering salicylate, the pH of the illuminated suspensions changed only when the experiments were conducted at highest pH value of 9.5,diminishing to a final value of pH = 8. The continuously stirred batch reactor consisted of a Pyrex vessel irradiated directly by the UV source, a X e H g 200-W lamp (Photon Technologies). A quartz cylinder, through which water was constantly recirculated, removed the heat produced by the light source. Temperature control (25 f 2 "C) was achieved by connecting the reactor to a temperature-controlled water bath. A Pyrex filter was used to exclude frequencies of light that would cause direct photodegradation of the dissolved compounds. The flux of photons entering the reactor was 3 W (0.9 X einstein/s) in the 300-400-nm wavelength range. To minimize errors associated with system variability, the irradiation was undertaken with an intensity of light corresponding to the semiconductor saturation zone. Under these conditions the variability between replicate experiments was approximately 4%. Suspension filtration was performed as described earlier. Salicylic and 3-chlorosalicylic acids were analyzed as before, while benzoic acid, phenol, and 4-chlorophenol were determined by using the UV spectrofluorimetric technique of Mathews" (detection limits: benzoic acid, 2 X 10" M; phenol, 0.5 X 10-6 M; khlorophenol, 2.5 X 10-6 M). No degradation was observed when the experiments were conducted in the dark or in the absence of the semiconductor. (13) (a) Borgarello. E.; Kiwi, J.; Pellizzetti, E.; Visca, M.; Gratzel, M. J.
Am. Chem. Soc. 1981, 103, 6324. (b) Rao, M. V.; Rajshwar, K.; Pai
(12) Anderson, M. A.; Gieselman, M.; Xu, Q.J . Membr. Sci. 1988, 39, 243.
Verneker, V. R.; DuBow, J. J . Phys. Chem. 1980,84, 1987. (c) Ward, M. D.; White, J. R.: Bard, A. J. J. Am. Chem. Soc. 1983, 105, 27. (14) Matthews, R. W. J . Phys. Chem. 1987, 91, 3328.
Photodecomposition of Salicylic Acid and Related Compounds n
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The Journal of Physical Chemistry, Vol. 95, No. 8,1991 3401
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Equilibrium Salicylate Concentration (pM) Figure 1. Adsorption isotherms for salicylic acid adsorption, T = 23 2 O C , I = IO? Firing temperature: 0 , 3 7 5 OC;A, 400 O C ; 0, 450 OC. pH = 5.7; filled symbols pH = 3.7.
Equilibrium Concentration (CCM) Figure 2. Adsorption isotherms normalized with rwpect to surface area. Firing temperature: 0,375 O C ; A, 400 O C ; 0 , 4 5 0 O C . pH = 3.7; filled symbols pH = 5.7.
Results Effect of Firing Temperature on Membranes. Table I summarizes the effects of firing temperature on surface area and crystalline phase of Ti02 crushed membranes. In the temperature range from 350 to 400 OC, the Ti02membrane changes from an amorphous material, with broad X-ray peaks centered at the position of crystalline anatase, to a pure anatase phase. The temperature of transition, from anatase to the more thermodynamically stable rutile, appeared to be in the range 475-500 OC. A surface area decrease is clearly observed with increasing tiring temperature. During the firing process, surface hydroxyl groups are lost, and the bulk is subject to crystallographic plane migration and subsequent shrinkage. These factors lead to surface area reduction and increasing pore size, as verified by BET measurements. These results show that, through the variation of the preparation procedures, it is possible to control the physicochemical characteristics that affect the behavior of the Ti02 membranes. Studies of Adsorption Kinetics. Membranes fired at 375,400, and 450 OC were selected for adsorption studies because these temperatures produced a change in surface area without a concomitant change in the crystalline form of the membrane, which remains anatase. The adsorption kinetics for salicylic acid, for both pH = 3.7 and 5.7,showed that, if the adsorption at 6 h is taken as an equilibrium condition, most of the adsorption (90%) occurs within the first hour, with very little increase over a 24-h period. Adsorption studies over a longer time frame may be compounded with problems due to biological degradation of the salicylic acid.Is As the adsorption of salicylic acid takes place, a bright yellow color is formed on the Ti02membrane, while no detectable color is observed in the aqueous phase even after 24 h. This observation is best explained by assuming that a charge-transfer complex is formed between singly charged salicylate ions and the titanium atoms at the oxide surface. This complex, which effectively removes salicylate from the solution, is very stable with respect to titania dissolution. Formation of a similar yellow complex, when salicylic acid and titanium(1V) solutions are mixed, has been used as a quantitative UV test for the determination of the metal.I6 Equilibrium Isotherms. In Figure 1, we compare the adsorption isotherms for the three different membranes at two pH values. Higher adsorption was obtained for experiments performed at pH = 3.7 than a t pH = 5.7. The higher adsorption, 71 pmol/g, was obtained at pH = 3.7, with a 375 O C membrane, while for the same membrane at the higher pH the adsorption was 31 pmol/g. One notices that there is not a well-defined maximum in the isotherms; furthermore, as firing temperature increases adsorption
decreases. Decreased adsorption at the higher pH is most likely due to the reduced surface charge on the hydrous oxide surface, which lowers the electrostatic attraction between the positively charged surface and the organic anion. It should be remarked, however, that salicylic acid also demonstrated a positive amount of adsorption when the analysis was performed at pH = 6.4 (results not shown) where both the oxide surface and the salicylate ion are negative, which further indicates the formation of covalent bonds. To test the system for surface heterogeneity, we attempted to model multiple-site adsorption using different isotherms. The lack of an adsorption maximum may mean that we have not saturated the surface sites, have several types of sites, or possibly have a continuum of sites. The presence of Ti02hydroxyl groups having several degrees of acidity has been postulated in the literature." At these sites, the titanium cations, whose charge distribution is affected by the ligands that form the coordination sphere, would show varying tendencies to bond with the adsorbate. A simple Langmuir equation, which assumes the presence of only one type of surface site, did not fit our data. When the method proposed by Spositol* for a two-site Langmuir model was used, the geometrical fitting was satisfactory. However, best results were consistently obtained when fitting the data with a Freundlich isotherm (see Figure 1 insert), which could account for a continuous distribution of adsorption sites. In fact, the fit to the Freundlich equation is interpreted as if the strength of the oxide-adsorbate bond varied with coverage. This variance might be taken as a weighted average of the effects of the different surface sites or as a change in steric or electrostatic factors as levels of adsorption change. This level of uncertainty in modeling has prompted us to perform in situ FTIR spectroscopy which will be reported in a later paper. The decrease in adsorption with increasing firing temperature corresponds directly to a decrease in surface area and therefore in the numbers of titanium atoms and hydroxyl groups present per unit area available for bonding. When the adsorption isotherms were normalized with respect to the surface area of the adsorbing membrane, the curves shown in Figure 2 were obtained. After this transformation, data for all membranes at a given pH fit on the same curve. At pH = 3.7, an average value of 0.84 pmol/m2 was obtained for the maximum observed adsorption, corresponding to 0.5 molecule/nm*. To define the role of the adsorbate functional groups on the adsorption of organic compounds on the TiOl membrane, we investigated the adsorption behavior of 3-chlorosalicylic acid, benzoic acid, phenol, and 4-chlorophenol. Figure 3 summarizes the results obtained a t pH = 3.7 and 5.7 for a membrane fired a t 450 O C . Salicylic and 3-chlorosalicylic acids have similar
(IS) (a) Yost. E. C. Ph.D. Dissertation, University of Wisconsin, 1987. (b) Yost, E.C.; Tejedor-Tejedor,M.I.; Anderson, M. A. Emiron. Sci. Technol. 1990, 21, 822. (16) (a) Muller, J. H.J . Am. Chem. Soc. 1911,33, 1506. (b) Goncalves, A. A.; de Salles Andradc, H.A.; Frwnius, Z . Anal. Chem. 1978,292, 299.
(17) Boehm, H.P.Discuss. Faraday Soc. 1971,52,264. (b) Yamanaka, T.; Tanabe, K.J. Phys. Chem. 1975, 79, 2409. ( 1 8 ) S p i t o . G . The Surfuce Chemistry of Soil; Oxford University Press: London, 1984; p 118.
Tunesi and Anderson
3402 The Journal of Physical Chemistry, Vol. 95, No. 8, 1991 I
50 I
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Equilibrium Concentration (pM) Figure 3. Adsorption isotherms on a 450 O C fired membrane. pH = 3.7: 0,salicylic acid. pH = 5.7: A, salicylic acid; filled symbols, 3-chlorosalicylic acid. -
100
01 0
I
I
I
I
I
10
20
30
40
50
Time (minutes) Figure 5. Comparison of salicylic acid degradation (initial concentration 50 pM) in aqueous TiO, suspensionsand suspensions with added methanol (0.01 M): 0,pH = 4; A, pH = 7.5; 0,pH = 9.5; filled symbols, with added methanol. I
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Time (minutes) Figure 4. Salicylic acid photodegradation on a Ti02 450 OC fired membrane: 0, pH = 4; A, pH = 7.5; 0, pH = 9.5. Initial concentration 40 pM;filled symbols, initial concentration 80 pM. adsorption behavior with 45 and 15 pmol/g adsorbed at high and low acidity, respectively. These results were confirmed by several replicate experiments, performed with different Ti02membranes under different acidity conditions. These data indicate that the o-chlorine substituent presence on the salicylic acid molecule does not sterically hinder the formation of the adsorption complex. Similar behavior has been observed by Yostl5 in comparing salicylate and 2,4-dihydroxybenzoate adsorption on goethite (aFeOOH). With the analytical technique used in this study, no adsorption could be detected for benzoic acid, phenol, and 4chlorophenol. Photodegradation Mechanisms. The membrane fired at 450 OC (crystalline anatase) was selected for photocatalytic degradation studies. (When normalized with respect to surface area, adsorption density did not vary with firing temperature.) The effect of pH on the degradation rate of salicylic acid is shown in Figure 4 for initial organic concentrations of 40 and 80 pM. A detailed analysis of Figure 4 illustrates several experimental results. At the lower concentration, which represents data corresponding to the initial part of the adsorption isotherm (an unsaturated surface), a marked difference in degradation rates can be observed with varying pH. On the other hand, at the higher concentration levels, differences in degradation rates with changing pH values are not as profound. These experiments also show that increasing salicylic acid concentration reduces photodegradation at low values of pH, while an opposite trend is observed a t high pH values. In order to investigate the effect of chemical competition on the degradation process, methanol was added to the suspension. This alcohol has been found to both scavenge holes and react with OH' radical^.'^-*^ Figure 5 illustrates the pH dependence of (19) Ulmann, M.;Augustynsky, J. Chem. Phys. Lcrr. 1987, 141, 154. (20) Matthews, R.J . Chem. SOC.,Faraday Trans. I 1984,80,457.
0
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Time (minutes) Figure 6. Benzoic acid photodegradation, on 450 "C fired membrane (initial concentration 60 pM): 0, pH = 4; 0 , pH = 7.5; A, pH = 9.5.
2
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'I 20
0
0
10
20
30
40
50
Time (minutes) Figure 7. Phenol photodegradation, on 450 OC fired membrane (initial concentration 48 pM): 0, pH = 4; 0 , pH = 7.5; A, pH = 9.5. salicylic acid degradation (initial concentration 60 p M ) in the presence of methanol (0.1 M). The alcohol appears to alter the photodegradation only at the two higher pH values (7,9.5) and does not substantially affect degradation a t pH = 4.0. Figure 6 shows the behavior of benzoic acid, which was not found to adsorb at the Ti02surface. The similarity of degradation rates, with respect to changes in pH, between benzoic acid and salicylic acid suggests that benzoic acid degradation kinetics are affected by the products of the photoreaction, which certainly include hydroxylated compounds able, when the hydrous oxide is protonated, to chemically bond to the TiOl membrane surfacem Lastly, Figure 7 shows the degradation behavior of phenol for an initial concentration of 48 pM, comparable to the lowest value shown for salicylic acid degradation. In this case, it is difficult
Photodecomposition of Salicylic Acid and Related Compounds
The Journal of Physical Chemislry, Vol. 95, NO. 8, 1991 3403
to identify a trend with respect to pH. The same was observed for 4-chlorophenol photodegradation (result not shown).
the 4-fold coordinated Ti atoms were present per nm2 of the surface which had a surface area of 25 m2/g. As implied by surface studies," surface area and particle size could significantly affect this value. In our study we calculate that at maximum salicylate adsorption, which occurred at pH = 3.7, coverage amounts to 0.5 molecule/nm2, comparing favorably with the value of 1.9 reported above. This lower value results also from the oxide charge density, which under these pH conditions is not fully charged to its maximum positive value. The lack of adsorption for benzoic acid ( 3 4 % of initial solution concentration, detected with spectrofluorimetry) and the total absence of adsorption for phenol and 4-chlorophenol can be explained by considering that these molecules do not have the stereochemical configuration which makes ring formation possible. Substitutions that lead to ring formation are favored over similar two-ligand substitutions. This is due to the entropic contribution given by the replacement of two leaving ligands (water or hydroxyl groups) by the one adsorbate molecule.26 Photodegradation Mechanisms. Several mechanisms have been proposed for the observed photodegradation reactions occurring at the Ti02 surface under UV irradiation. It has been suggested26 that the formation of surface bound species is the dominant feature of n-Ti02 semiconductors. In this mechanism, holes from the valence band react with surface bound hydroxyl groups (and/or water and hydroxyl ions), generating surface radicals:
Discussion Adsorprion Mechanisms. In studies of the Ti02 surface, Hadjiivanov2' showed that, for samples of anatase prepared by different procedures, it is difficult to assume which crystallographic planes are predominantly exposed. This is due to the fact that the preparation methods affect the crystal cleavage. The electronic environment of the Ti cation is affected by the number of bonds with oxygens and also by the coordination number of these oxygen ligands. As a result, several types of titanium cations can exist at the different crystallographic planes. Infrared studies of CO adsorption on anataseu have confirmed the existence of a heterogeneous array of adsorption sites. Structural heterogeneity at zero adsorbate coverage arises from the ability of the Ti02 conduction electrons to transmit electronic effects along a few unit cells of the crystal. When adsorbates are present, their electronic effects are distributed over a few rows of atoms producing site heterogeneity. With increasing adsorbate coverage, the interaction of the organic adsorbate with both the crystal lattice and other adsorbates generates further heterogeneity. The resulting heterogeneous surface accounts for the poor fit to the one-site Langmuir isotherm. Therefore, to depict the oxide surface for purposes of describing adsorption mechanisms, it seems more correct to consider the coordination sphere of the most numerous and chemically active atoms, rather than the average characteristics of the oxide surface. The most active of these sites are likely the titanium atoms present on the surface whose coordination sphere is incomplete. These are present predominantly at the (1 10) plane, (1 lO)X( 100)and (OlO)X( 100)edges, and several corners. These atoms are 4-fold coordinated to oxygen and have two unfilled orbitals which are exposed at the oxide surface. From infrared studies23it is recognized that this type of titanium cation has the highest degree of Lewis acidity of all surface sites because it can accept two "lone pairs" from electron donors to complete its octahedral coordination. This kind of surface species has been invoked to interpret the anatase surface reactivity toward catalytic decomposition of alcohols in the gas In an aqueous medium, Ti cations can bond to oxygen atoms of water molecules. If salicylate ions are present in solution, the ability of their oxygen atoms to form chelate structures leads to the replacement of the solvent by the organic ligand. Using "in situ" CIR-FTIR techniques, Tejedor-Tejedor et aI.'* identified surface complexes formed between goethite and salicylate that were chelates of this type:
Tiso- + h+VB
-
Tis@
(3) These radicals would be responsible for (i) formation of surface peroxo species
-
2(Tis@) Tis-O-O-Tis (4) (ii) surface recombination of h+VBand e-CB,and (iii) electron exchange between the semiconductor and the species in solution. According to this model, compounds that can adsorb at the T i s o radical centers would undergo photooxidation more efficiently than water and OH- ions, also reducing the number of recombination centers by diminishing the production of peroxo groups. Several authors, with different emphasis, have suggested that OH' radicals, derived from water or surface hydroxyls through photooxidation, reactions 5 and 6, are involved in the photooxidation of organic compounds:28
The hydroxyl radical formed in these reactions could generate H202,according to
0
20H' H202 (7) The formation of H202 on semiconductor surfaces has been the object of numerous investigation^.^-'^^ Additional proposed mechanisms include 4
-
+ O2 + 2e-,, H202 2H20 + 2 h + v ~ H202 + 2H+ 2H+
-+
.
-
We interpret our adsorption experiments using a similar picture for salicylate adsorption on Ti02, where the Ti cations are at the center of the surface bond. CIR-FTIR studies confirming this hypothesis are to be reported in a future paper. Carrisoza et al.23 calculated that, for the anatase used in their experiments, 1.9 of (21) (a) Hadjiivanov, K.; Klissunki, D.; Davidov. A. J. Chem. Soc., Faraday Trans. I 1988,8437. (b) Hadjiivanov, K.; Kliwunki, D.; Davidov, A. Proc. 6th In?. Symp. Heterogen. Catal., Sofia 1987 (Part I), 365. (22) Garrone, E.; Bolis, V.; Fubini. B.; Morterra, C. hngmuir 1989, 5,
__
892. -.
(23) Carrisoza. I.; Munuera, G.; Castanar, S.J . Catal. 1977, 49, 265. (24) Ramh, G.; Busca, G.; Lorenzelli, V. J . Chem. Soc., Faraday Trans. I 1987,83. 1591. ( 2 5 ) Anpo. M. Res. Chem. Intermed. 1989, 1 1 , 67.
(9)
Under circumstances where electron donors are present, holes are readily removed from the valence band, thereby decreasing the recombination rate. In this case, the formation of H202is postulated to occur through conduction band electrons, according to reaction 8. In the absence of electron donors, reaction 7 would be activated. According to reaction 9, hydrogen peroxide may (26) Myers, R. T. Inorg. Chem. 1978, 17, 952. (27) Augustinsky, J. In Chemistry of Solid Materials; Springer-Verlag: Berlin, 1989. (28) (a) Bahnemann, D. W.; Fisher, C. H.; Janata, E.; Henglein, A. J . Chem. SOC.,Faraday Trans. I 1987,83. 2559. (b) Howe, R. F.; Graetzel. M. J . Phys. Chem. 1987, 91, 3906. (c) Cunningham. J.; Srijaranai, S.J . Photochem. Photobiol., A 1988.43, 329. (d) Tennakone, K.; Punchihewa, S.;Wickremanayaka,S.; Tantrigoda, R. U. J . Photochem. Photobiol., A 1989, 46, 247. (29) Harbour, J. R.; Tromp, J.; Hair, M. L. Can. J . Chem. 1985,63,204.
3404 The Journal of Physical Chemistry, Vol. 95, No. 8, 1991
also derive from water oxidation. It should be noted that H 2 0 2can be also destroyed through the following mechanisms:
-
+ 2e-ce + 2H+ 2 H 2 0 HZ02 + 2h'VB 0 2 + 2H+
HzOz
+
(10)
(11)
On the basis of our experimental observations and from these briefly reviewed mechanisms, the decrease of salicylate degradation with increasing pH can be interpreted by assuming that two different reaction pathways take place under conditions of different acidityUwWe suggest that electron transfer and indirect hydroxyl radical attack are the predominant reaction mechanisms at low and high pH, respectively. At low pH, the positively charged titanium hydroxide offers a suitable surface for salicylate chemisorption, according to the chelate model proposed above. This bond has a relatively high covalent character, and the oxygen atoms of the salicylate ions, being relatively strong electron donors, are able to direct interact with valence band holes, similarly to the mechanism proposed for interfacial hydroxyl groups in eq 3. Increasing the pH reduces adsorption and gradually increases the electrostatic repulsion between the salicylate anion (pK, = 2.914) and the oxide surface (pHIEP= 6). The increased distance between the two reactants no longer allows direct charge transfer. With only limited salicylate adsorption occurring, most of the decomposition is probably mediated by OH' radicals formed at other semiconductor surface sites. The decrease in photodecomposition rates at pH 9.5 is probably due to Coulombic repulsion between the organic anions and the highly negative charged oxide surface. Decomposition would thus depend on diffusion of surface-generated OH' radicals to the low concentration of anion in the double layer, a slower process than direct charge transfer. At low pH, where salicylate ions are degraded by direct injection of electrons into the valence band, the lower concentration of holes at the surface would interfere with hydrogen peroxide photochemistry by precluding degradation (eq 11) and diminishing the formation of OH' radicals (eqs 5 and 6). When direct chemisorption of the organic electron donor does not occur, it has been shown that the TiOz surface can photodecompose H,O2 very efficiently, following eqs 10 and 1 1 .3 Thus, it appears that, at both high and low pH, the role of hydrogen peroxide in oxidizing the organic is insignifi~ant.~~ Photodegradation experiments conducted at different initial salicylate concentrations (Figure 4) also appear to confirm that different reaction mechanisms control the photodegradation under different conditions than those observed with varying pH. Increasing the initial salicylate concentration produces opposite trends. At low pH, the initial relative photodegradation rate decreases while it increases at high pH. At low pH, the higher salicylate surface coverage leads to displacement of OH- and water ligands at the titanium centers, thus reducing the rate of photoformation of OH' radicals. Slower degradation of nonadsorbing species would result. The decomposition of chemisorbing species would be affected as well at higher concentrations, because these species also compete at the catalyst surface with the initial reagent for the same active sites. At higher pH values, where less chemisorption occurs and indirect charge transfer predominates, diffusion through the double layer controls reactions with OH' radical^.^' Since the relative rate of diffusion-limited reactions increases at higher reagent concentrations, degradation is faster. Further evidence for these proposed photodegradation mechanisms is obtained by adding methanol to the system. For salicylate, at low pH, the lack of change in degradation rate would (30) (a) Peral, J.; Casado, J.; Domenech, J. J . Photochem. Phorobiol., A 1988.44.209. (b) Anpo, M.; Tomonari. M.; Fox, M.A. J. Phys. Chem. 1989, 83, 7300. (c) Cunningham, K. M.; Goldberg, M. D.; Weiner, E. R. Enuiron. Sci. Technol. 1988, 22, 1090. (d) Takagi, K.; Fujioka, T.; Sawaki, Y.; Iwamura, H. Chem. Le??.1985,913. ( e ) Kawaguchi, Emiron. Technol. Lett. 1984, 5, 471. (31) Mulvaney, P.; Swayambunathan, V.; Grieser, F.; Meisel, D. J . Phys. Chem. 1988.92, 6732.
Tunesi and Anderson be expected if direct charge injection is predominantly responsible for photodegradation. Since methanol does not strongly adsorb on TiOz in aqueous systems and does not compete for sites, it does not affect the photodegradation rate of a species degrading by direct charge injection. The slight reduction observed is likely due to an increased salicylate solubility caused by the methanol cosolvent effect. Only at higher pH, where the reaction is diffusion-limited and less salicylate is adsorbed, would methanol quench the degradation. As adsorption levels decrease, the sacrificial efficiency of methanol increases noticeably. Charge transfer at the TiOZ surface has been proposed to explain halide oxidation, although by two different mechanisms. In one, iodine oxidation is mediated by OH' radicals coordinated to surface Ti atoms3* In the other, an indirect two-step path for photoanodic oxidation of halide ions is assumed: the formation of an inner-sphere surface complex that can sustain a direct charge transfer between the excited Ti02 semiconductor and the halide anion. Such charge transfer can be explained by using the model of the Ti02 band structure proposed by Hoffman and G r a e t ~ e I . ~ ~ J ~ They suggest that the bonding orbitals in the TiO, valence band are predominantly O(2p) in character, while the edge of the nonbonding or slightly antibonding conduction band is primarily Ti(3d). When a Ti0,-salicylate chelate forms, it is most likely that the salicylate bonds to O(2p) surface orbitals of the Ti02. During irradiation, these O(2p) orbitals would become sources of holes and attract electrons from the organic ligand, leading to oxidation of salicylate. For benzoic acid the observed change in degradation rates with varying pH appears to be influenced by photooxidation products adsorbed at the surface. In fact, as mentioned in the Results section, compounds that can strongly adsorb on the semiconductor are very likely to be generated during benzoic acid degradation. These adsorbed products, in a fashion similar to salicylate ions, likely photodegrade rapidly at pH 4, thus reducing the competition for active sites and inducing the fast decomposition of the starting reagent. At higher pH values, where even the products do not adsorb, the degradation rates of both reagents and photooxidation products are decreased, and the pH effects become less evident. For phenol and 4-chlorophenol, unlike the former cases, different photodegradation mechanisms do not appear to occur. Since phenol and, most likely, its photodegradation products do not significantly adsorb on TiOz surfaces, the degradation kinetics do not appear to be i n f l u e n d by surface conditions. This result suggests that when the species does not modify the surface by binding to active surface sites, a steady-state concentration of oxidant OH' is reached, leading to similar degradation rates at all pH values.
Conclusions The adsorption mechanism outlined above suggests that the surface has a high number of heterogeneous sites, as supported by the good fit of the data to a Freundlich equation. With respect to chemisorption processes, the surface is not homogeneous. Instead, the local chemical environment for those titanium atoms that are chemically most active as well as abundant must be known in order to understand the overall behavior (e.g., adsorption mechanisms) of the system. The efficiency of the photoredox processes for the organic molecules studied depends upon their adsorption behavior. When direct charge transfer occurs at low pH, this mechanism is more efficient than free-radical attack. In this case, however, degradation rates decrease as the adsorbate concentrations increase, probably due to increased competition for active surface sites. At high pH values, for similar concentrations of organic compounds, the rates of degradation tend to be more similar among the compounds studied. This is probably due to the constant concentration reached by OH' radicals under these irradiation con(32) Gutierrez, C.; Salvador, P. J . Elecrrochem. Soc. 1986, 133, 924. (33) Hoffman, R. Angew. Chem., In?. Ed. Engl. 1987, 26, 846. (34) Grletzel, M.; Rotzinger, F. P. Chem. Phys. Let?. 1985, 118, 474.
J . Phys. Chem. 1991, 95, 3405-3409 ditions. For comparable initial concentrations of adsorbing and nonadsorbing compounds, the photodegradation rates are consistently reduced for the latter compounds, this effect becoming most evident when there is no adsorption of either the reactant or the reaction products. These results indicate that degradation rates for organic molecules should be reported with respect to the amount of the
3405
species adsorbed and to the pH20*30 a t which these degradation rates have been determined. These parameters would allow one to better interpret the interfacial relationship between semiconducting solids and adsorbates. Registry No. TiO,, 13463-67-7; salicyclic acid, 69-72-7; 3-chlorosalicyclic acid, 1829-32-9; benzoic acid, 65.85-0; phenol, 108-95-2; 4chlorophenol, 106-48-9; methanol, 67-56-1.
Dehydrobromination Reaction of Para-Substituted 2-Phenylethyl Derivatives in Functional Micelles Kazimiera A. Wilk Institute of Organic and Polymer Technology, Technical University of Wroclaw, 50370 Wroclaw. Poland (Received: June 25, 1990; In Final Form: November 16, 1990)
Dehydrobromination reactions of para-substituted 2-phenylethyl bromides, i.e., p-Y-C6H4CH2CH2Br(Y = NOz, CI,H, OCH,) have been examined in aqueous micelles of the functional surfactant N,N-dimethyl-N-(2-hydroxyethyl)-n-hexadecylammonium bromide (1) in the presence of sodium hydroxide. The kinetic experiments have been performed for a partially deprotonated nucleophilic head group. The variation of the overall first-order rate constant, k, with concentration of 1 can be fitted to equations describing the micellar effect on the bimolecular reaction in the cationic micelle as well as acid-base and ion-exchange equilibria.
Introduction The 1,2-elimination reaction is one of the most extensively studied because of the interesting possibilities for stereochemistry, orientation, substituent, and medium effects.l The possibility of variable transition states has presented an extraordinary mechanistic challenge,2 to which many investigators have responded. The 2-phenylethyl derivatives have played a key role in understanding the structure-reactivity correlations in bimolecular 1,2-elimination~.'*~More recently, they also figured in our studies addressed to the nonfunctional micelle-catalyzed classical E2 r e a c t i ~ n . ~ According to the phase separation approach? the overall reaction rate in a micellar medium comprises rates of two processes: the reaction occurring in the micellar pseudophase and that in the aqueous pseudophase. Rate enhancements of bimolecular reactions by nonfunctional micelles can be quantitatively described by evaluating the concentration of reactions in the micellar pseudophase and, furthermore, by calculating a second-order rate constant in that pseudophase.b8 The effect of nonfunctional cationic micelles on the elimination reaction of para-substituted 2-phenylethyl bromides promoted by hydroxide ion was found to be mainly due to the concentration of both reactants in the small (1) (a) Saunders, W. H., Jr.; Cockerill, A. F. Mechanisms of Eliminution Reoctions; Wiley: New York, 1973. (b) Fry, A. Chem. Soc. Reu. 1972, 163. (2) (a) Jencks, W. P. Chem. Reo. 1985, 85, 511. (b) Saunders, W. H., Jr. Acc. Chem. Res. 1976, 9, 19. (c) Ford, W. T. Acc. Chem. Res. 1973, 6, 410. (d) Bordwell, F. G. Acc. Chem. Res. 1972, 5, 374. (3) Kwart, H.; Wilk, K. A. J . Org. Chem. 1985, 50, 817. (4) Yano, Y.; Yoshiba, Y.; Kurashima, A.; Tamura, Y.; Tagaki, W. J . Chem. SOC.,Perkin Truns. 2 1979, 1128. (5) (a) Wilk, K.A.; Burczyk, B. J. Phys. Chem. 1989,93,8219. (b) Wilk, K. A. J . Phys. Chem. 1989. 93, 7432. (c) Wilk, K. A,; Burczyk, B. In Proceedings of the 81h Conference on Surface ond Colloid Chemistry, Liblice, Crechoslouakiu; 1989; pp 282-287. (6) (a) Yatsiminki, A. K.; Martinek, K.; Berezin, I. V. Teiruhedron 1971, 27,2855. (b) Martinek, K.; Yatsimirski, A. K.; Levashov, A. V.; Berezin, I. V. In Micellirulion, Solubilirurion, und Microemulsions; Mittal, K.L., Ed.; Plenum Press: New York, 1977; Vol. 2, p 489. (7) Bunton, C. A. In Surfucronrs in Solution; Mittal, K.L., Lindman, B., Eds.; Plenum Press: New York, 1984; Vol. 2, p 1093. (8) (a) Bunton, C. A. In Solurion Chemistry of Surfucrunrs; Mittal, K. L., Ed.; Plenum Press: New York, 1979; Vol. 2, p 519. (b) Bunton, C. A.; Savelli, G.Adu. Phys. Org. Chem. 1986, 22, 213.
0022-3654/91/2095-3405$02.50/0
volume of the Stern layer of the micelle.Sa Besides specific examples in the pesticide field9 and similar assemblies,'O a classical E2 reaction has not been quantitatively studied in functional micelles. The functional micelles and comicelles are effective nucleophilic or basic reagents." In many systems reaction involves attack of an oximate, hydroxamate, alkoxide, or thiolate moiety formed by deprotonation of a weak acid so that reaction rates are pH dependent.lS According to many kinetic findings, the sources of rate enhancements in functional and nonfunctional micelles are similar.8J6 As a continuation of our studies dealing with a quantitative treatment of micellar effects upon the E2 reaction in cationic micelles: we have examined the dehydrobromination reaction of para-substituted 2-phenylethyl bromides, p-Y-C6H,CH2CHzBr, where Y = NO2,CI, H, and OCH3promoted by aqueous solutions of N,N-dimethyl-N-(2-hydroxyethyl)-n-hexadecylammonium bromide (1) in the presence of sodium hydroxide. n-C16H33N+Me2CH2CH20H s 1
+
n-C16H33N+Me2CHzCH20- H+ la Micelles of this hydroxyethyl surfactant are effective nucleophiles at high pH in depho~phorylation,'~~~~~ deacylation,18 and nu191 la) Rezende. M. C.: Rubira. A. F.: Franco. C.: Nome. F. J. Chem. SOC.,Perkin Truns. 2 1983, 1075. (b) Nome, F.; Schwingel, E.' W.; Ionescu, L. G. J . Org. Chem. 1980, 45, 705. (IO) Nome, F.; Neves, A.; Ionescu, L. G. In Solution Behuuior of Surfactunrs; Mittal, K. L., Fendler, E. J., Eds.; Plenum Press: New York. 1982; Vol. 2. D 1157. (1 Ij'For general discussions see, refs 8b and 12-14. (12) Bunton, C. A. Curd Rev.-Sci. Eng. 1979, 20, I . (13) (a) Tonelatto, U. In Solurion Chemistry o/Sur/octunts; Mittal, K. L., Ed.; Plenum Press: New York, 1979; Vol. 2, p 541. (b) Fornasier, R.; Tonellato, U. J. Chem. Soc., Furaduy Truns. I 1980, 76, 1301. (c) Moss, R. A.; Lee, Y.-S.; Alwis, K. W. In Solution Eehuuior of Surfucrunrs; Mittal, K . L., Fendler, E. F., Eds.; Plenum Press: New York, 1982; Vol. 2, p 993. (14) Cordes, E. H. Pure Appl. Chem. 1978,50, 617. (15) (a) Bunton, C. A.; Ionescu, L. G. J . Am. Chem. Soc. 1973,95.2912. (b) Bunton, C. A.; Diaz, S. J. Am. Chem. SOC.1976, 98, 5663. (16) (a) Bunton, C. A,; Hamed, F.; Romsted, L. S . J . Phys. Chem. 1982. 86,2103. (b) Bunton, C. A.; Hamed, F.; Romsted, L. S. Telruhedron Let?. 1980, 21, 1217. (c) Brown, J. M.; Bunton, C. A.; Diaz, S.; Ihara, Y. J. Org. Chem. 1980, 45, 4169. .
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0 1991 American Chemical Society
