Surface Properties of Calcinated Titanium Dioxide ... - ACS Publications

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Surface Properties of Calcinated Titanium Dioxide Probed by Solvatochromic Indicators: Relevance to Catalytic Applications Omar A. El Seoud,*,† Adham R. Ramadan,‡,§ Bruno M. Sato,† and Paulo A. R. Pires† Institute of Chemistry, UniVersity of Sa˜o Paulo, P.O. Box 26077, 05513-970, Sa˜o Paulo, S. P., Brazil, Department of Chemistry, and The Yousef Jameel Science and Technology Research Center, American UniVersity in Cairo, P.O. Box 74, 11835 New Cairo, Egypt ReceiVed: October 7, 2009; ReVised Manuscript ReceiVed: April 27, 2010

Titanium dioxide was obtained by hydrolysis of the corresponding ethoxide, followed by washing, drying, and calcination at 80, 160, 240, 320, 400, and 700 °C, respectively. The following surface properties of the solids obtained were determined as a function of the calcinations temperature: TCalcn; area by the BET method; Brønsted acidity by titration with sodium hydroxide; empirical polarity, ET(30); Lewis acidity, RSurf; Lewis basicity, βSurf; and dipolarity/polarizability π*Surf, by use of solvatochromic indicators. Except for βSurf whose value increased slightly, heating the samples resulted in a decrease of all of the above-mentioned surface properties, due to the decrease of surface hydroxyl groups. This conclusion has been corroborated by FTIR. Values of ET(30), RSurf, and π*Surf are higher than those of water and alcohols; the Brønsted and Lewis acidities of the samples correlate linearly. The advantages of using solvatochromic indicators to probe the surface properties and relevance of the results to the applications of TiO2 are discussed. Introduction The impetus for the continued interest in studying the surface properties of solids is that they control the adsorption/desorption of species of interest (reagents, activated complexes, and products) which, in turn, control many useful applications. The adsorption of substrates and the desorption of the products are important steps in heterogeneous catalytic processes and are often rate limiting.1 In some cases, such as supported phasetransfer catalysts, the compatibility of the surface with the reactants is critical.2,3 The solvent affects the activity of heterogeneous catalysts, especially when the former is dipolar. The reason is that solvent dipoles compete with the reactants for the (acidic or basic) active sites of the catalyst, suppressing its catalytic activity.4 Consequently, the correlation of surface properties with catalytic activity of solid surfaces is an important field of material science research.5,6 Since its production on a commercial scale, titanium dioxide (TiO2) has been widely used as a pigment;7 in sun screening lotions;8 paints;9 ointments and toothpaste,10 etc. An enormous effort has been devoted to research on TiO2, which has led to many promising applications in different areas, e.g., photovoltaics, photocatalysis, photoelectrochromics, and sensors.11,12 Another surge in interest is motivated by the potential applications of TiO2 in nanoscience and nanotechnology.13 Information on the surface properties of titania has been derived, inter alia, from titration of the (acid or base) sites present;14,15 spectroscopic data of adsorbed probes, e.g., NH3, pyridine, and its derivatives, NO, CO;16,17 or fluorescent probes,18 thermogravimetric analysis,19,20 and inverse gas chromatography.21-23 Titration gives quantitative information with regard * Corresponding author. Phone/Fax: +55-11-3091-3874. E-mail: [email protected]. † University of Sa˜o Paulo. ‡ Department of Chemistry, American University in Cairo. § The Yousef Jameel Science and Technology Research Center, American University in Cairo.

to acidic/basic sites; spectroscopic data, in particular FTIR, report on Brønsted and Lewis acidity/basicity, and probe the state of surface OH groups, and the strength of solid-substrate interactions. Inverse gas chromatography focuses on the contribution of the Lifshitz-van der Waals nonspecific (dispersion) interactions to the free energy of the surface. In summary, different techniques, each one providing specific information, are required for a complete description of the surface properties that are relevant to applications of this important material. An alternative is to extend to TiO2 the same approach that has been employed in order to discern medium effects on chemical phenomena (rate constants, equilibrium constants, spectroscopic shifts, etc.); see also literature on the determination of the empirical polarities of solid surfaces, e.g., silica, alumina, and titanium dioxide.24-29 A brief discussion of this approach for liquids (solvents and their mixtures) is in order. Medium effects are complex, and hence, cannot be analyzed in terms of a single property, e.g., the relative permittivity of the solvent, ε, or any of the Kirkwood dielectric functions. More realistically, we describe these effects in terms of specific and nonspecific solute-solvent interactions, including hydrogen bonding; iondipole and dipole-dipole interactions; dipole induced-dipole; dispersion; or London interactions. Examples of the quantification of medium effects are the solvation free energy relationships, e.g., the Taft-Kamlet-Abboud equation:30,31

SDP ) constant + aRSolv + bβSolv + s(π*Solv + dδ) + h(δH2) (1) where the solvent dependent phenomenon, SDP, is modeled as a linear combination of two hydrogen-bonding terms, in which the solvent is the hydrogen-bond donor (aRSolv), or the hydrogenbond acceptor (bβSolv), a dipolarity/polarizability term [s (π*Solv + dδ)], and a cavity term (h(δH2)), related to Hildebrand solubility parameter. The properties RSolv, βSolv, and π*Solv, are

10.1021/jp909619c  2010 American Chemical Society Published on Web 05/18/2010

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Figure 1. Molecular structures of the solvatochromic probes employed in the present work. These are cis-dicyano-bis(1,10)-phenanthroline iron(II), Fe(phen)2(CN)2; 4,4′-bis(N,N-dimethylamino)benzophenone, Michler’s Ketone; 2,6-dibromo-4-[(E)-2-(1-butylpyridinium-4-yl)ethenyl] phenolate, BuPMBr2; 3-(4-amino-3-methylphenyl)-7-phenylbenzo[1,2b:4,5b′]difuran-2,6-dione, ABF; and 2,6-diphenyl-4-(2,4,6-triphenyl- pyridinium-1yl)phenolate, RB.

known as solvatochromic parameters because they are determined by using solvatochromic substances (see discussion below); the subscript (Solv, for solvent) is employed so that they are not confused with other parameters, e.g., R and β of Brønsted catalysis equation.32-35 Note that RSolv and βSolv refer to partial solute-solvent electrons transfer; that is, they refer to Lewis acidity and basicity, respectively; both types of acidity (Brønsted and Lewis) can be distinguished by titration with sterically hindered amines, e.g., 2,6-di-tert-butylpyridine;16b titration with NaOH determines sites of different acid strength, e.g., of pKa ) 2.9 and 12.7 for titania.14 Solvatochromic dyes (hereafter designated as “probes”) are substances whose UV-vis spectra, absorption, or emission, are particularly sensitive to the properties of the medium. The probes depicted in Figure 1 exhibit this sensitivity because each possess a medium-sensitive intramolecular charge-transfer, e.g., from the phenolate oxygen to the positively charged nitrogen for BuPMBr2 and RB. The equation that converts the electronic transition within the probe into the corresponding intramolecular charge-transfer energy is:32

ET(probe) ) 28591.5/λmax

(2)

where λmax is the maximum wavelength of the CT band in nm and ET(probe) is an empirical medium polarity scale in kcal/mol. Values of ET(probe) are then rationalized in terms of specific and nonspecific probe-medium interactions, vide supra.34,35 For our purpose, therefore, ET(probe) is substituted for SDP in eq 1. In the present study, we have raised the following questions: What information can we gain by probing the surface of TiO2 by solvatochromic indicators? How does this information bear on the applications of titania? The solids employed in the above-mentioned studies were mostly commercial products. We decided that it is important to examine the effects of a systematic treatment of TiO2 (calcination) on its surface properties, as revealed by the probes shown in Figure 1. Additionally, we were interested in investigating how the solvatochromic properties correlate with those determined by independent techniques. Using the first four indicators of Figure 1 we have calculated the values of overall empirical polarity ET(probe)Surf; Lewis acidity (RSurf); Lewis basicity (βSurf), and dipolarity/polarizability (π*Surf); the subscript (Surf) refers to the solid surface. These properties showed systematic dependence on TCalcn; Lewis and (titration-based) Brønsted acidities linearly correlated; the results have been explained; our reasoning has been corroborated by studying the samples by FTIR, and by theoretical calculations on model systems.

Experimental Section Chemicals and Probes. Commercial TiO2 sample (P25) was obtained from Evonik Industries; the other chemicals were from Aldrich, aPACgão2 Química, and Merck. Dichloromethane was agitated, then distilled from CaH2, and kept over activated type 4A molecular sieves. Michler’s ketone was recrystallized from ethanol. The Ti(OC2H5)4 used for the preparation of TiO2 was obtained as a dispersion in ethanol; standardized NaOH solution was used for the determination of Brønsted acidity. The probes Fe(phen)2(CN)2 and BuPMBr2 were synthesized as given elsewhere,36,37 and ABF was a gift from Prof. S. Spange (Chemnitz, Germany). TiO2 Sample Preparation. TiO2 was obtained by hydrolysis of its dispersion in ethanol. In order to ensure complete precipitation, excess deionized water was added slowly, with continuous stirring. The hydrated TiO2 precipitate was filtered and thoroughly washed with water, then dried to constant weight. It was then calcinated in air for 2 h in a Thermolyne 48000 furnace at temperatures of 80, 160, 240, 320, 400, and 700 °C. After heating, the samples were cooled under reduced pressure and kept in tightly stoppered bottles. Determination of Brønsted Surface Acidity. Brønsted surface acidity was determined volumetrically by the adsorption of sodium hydroxide from solutions of different concentrations. Accurately weighed TiO2 samples, ca. 0.2 g, were shaken (Burell Wrist-Action shaker) with 10 mL of the base for 6 h, then left overnight. The residual base was back-titrated with standardized HCl. Brφnsted acidity was obtained from the difference between each titration and a “blank” run.14 Investigation of the Surface Properties of TiO2 by Use of Solvatochromic Probes. Samples of TiO2, 0.2 g each, were introduced into 2 mL volumetric tubes; the solids were dried for 4 h at room temperature, under reduced pressure, over P4O10. Each sample was quickly covered with 2 mL of the probe solution in CH2Cl2 (5 × 10-5 mol/L) and agitated for 1 h, at room temperature, at 60 rpm, by using a tube rotator (model 099A, Glas-Col, Terre Haute). The TiO2 suspensions were transferred (Pasteur pipet) to 2 mm path length UV-vis (stoppered) quartz cells, and the reflectance of the solid was recorded by using a Shimadzu UV-2550, UV-vis spectrophotometer, equipped with model IST-204A (double beam) integrating sphere reflectance attachment. The conditions were: Each spectrum was recorded at least twice, against P25 titania as a (white) reference; rate ) 140 nm/min; slit width ) 1.0 nm; sampling interval 0.5 nm. The values of λmax were calculated from the first derivative of the absorption spectra; the uncertainty

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in λmax was (1.0 nm, corresponding to an uncertainty of 0.2 kcal/mol in ET(BuPMBr2). FTIR. We have employed Bruker Vector 22 FTIR spectrophotometer, equipped with a DRIFTS diffuse reflectance attachment (EasiDiff, Pike technologies, Madison). TiO2 was weighed and dried as described above, then transferred to a N2flushed inflatable polyethylene chamber (AtmosBag, Aldrich) and thoroughly mixed with dry KBr, at a ratio of 1/10. The solid mixture was transferred to the trough of the abovementioned accessory and its reflectance recorded; 24 spectra were added at 1 cm-1 resolution. During measurement, the sample compartment of the spectrophotometer was continuously flushed with dry nitrogen. Nitrogen Adsorption Measurements. The surface areas of the TiO2 samples were determined by nitrogen adsorption at 77 K by using a Micromeretics ASAP 2020 equipment. Typically, the samples were degassed at 60 °C for 6 h prior to nitrogen adsorption. The values of the surface areas (SBET) were obtained in the usual manner, i.e. by employing the Brunauer, Emmett, and Teller, BET equation, and the assumption that each nitrogen molecule occupies an area of 0.162 nm2 at 77 K.38 Results and Discussion Note: Details of all calculations are given in the Calculations section. Choice of the Probes and the Experimental Protocol Employed. As shown in the Calculations section, the values of the solvatochromic parameters were obtained from the UV-vis spectra of adsorbed probes. We used BuPMBr2 because its pKa in water is conveniently low, 5.22, so that any surface-induced pKa increase, if it occurs, does not constitute a problem (the CT band disappears if the indicator is protonated at its phenolate oxygen). RB, however, is the most extensively studied solvatochromic probe; its empirical polarity scale, ET(30), is available for more than 400 solvents and solvent-mixtures.31 Therefore, we have employed the linear correlation between ET(BuPMBr2) and ET(30) (34 protic and aprotic solvents),37 in order to convert the former into the latter empirical polarity scale. The use of the other solvatochromic probes is standard. The relevant requirements in the experimental protocol are as follows: (i) The water contents of the TiO2 samples should remain unchanged during the experiment; (ii) The spectrum of the adsorbed probe reflects the corresponding property of the surface. We have verified that the drying protocol is adequate because the samples showed the same weight after drying for 2, 3, and 4 h, respectively. Additionally, probe adsorption experiments repeated over a six-month period showed the same λmax ( 2 nm. Probe adsorption from its solution in CH2Cl2 has proved to be experimentally more convenient than applying solution of the latter to TiO2, followed by solvent evaporation. When the former approach was employed, the TiO2 particles were evenly colored and the supernatant was colorless, as confirmed by recording the absorbance spectrum of the latter, after filtering off the suspended solid. Reducing the dye concentration in the stock solution to half its original value resulted in a decrease of reflectance, without changing the value of λmax. Thus, there is no dye “stacking” on the surface of TiO2, i.e., the spectrum of probe reflects a true surface property. Dependence of surface properties of TiO2 on TCalcn. Figure 2 shows typical spectra of the solvatochromic dyes, adsorbed on TiO2. The dependence of the spectroscopic data on TCalcn is shown in Table 1, and the calculated solvatochromic

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Figure 2. UV-vis absorption spectra of the adsorbed dyes: Fe(phen)2(CN)2, curve 1; ABF, curve 2; BuPMBr2, curve 3; and Michler’s Ketone, curve 4 for the original sample (A), and that calcinated at 400 °C (B).

parameters are listed in Table 2, along with the corresponding surface areas and Brønsted acidity; the main emphasis is on samples heated up to 400 °C. Some correlations of the parameters calculated are depicted in Figure 3; all correlations, e.g., with TCalcn, are listed below, where r ) correlation coefficient.

ET(30) ) 70.005 - 3.160 × 10-3TCalcn

(Brønsted

r ) 0.9809 (3)

acidity)10-3 × µmol/g ) 1.161 1.648 × 10-3TCalcn r ) 0.9656 (4)

RSurf ) 3.255 - 5.638 × 10-4TCalcn

r ) 0.9551

(5) βSurf ) -5.760 × 10-3 - 2.145 × 10-4TCalcn

π*Surf ) 1.128 - 1.966 × 10-4TCalcn

r) 0.9940 (6)

r ) 0.9634

(7) RSurf ) 2.878 + 0.328[(Brφnsted acidity)10-3 × µmol/g] r ) 0.9208 (8) Regarding these data, the following is relevant: (i) The surface of TiO2 consists of coordinatively unsaturated (Ti4+) cations and (O2-) anions, as is the case of metal oxide surfaces.14,16 In hydrated TiO2, dissociative adsorption of water occurs in order to reduce the coordinative unsaturation of surface sites; leading to the formation of surface OH groups. Brønsted acidity results from doubly coordinated hydroxyl groups, and those polarized by the cations.14 Calcination leads to the loss of water from the structure, as confirmed by the thermal behavior of the samples (results not shown), and by FTIR, vide infra. (ii) Calcination, therefore, reduces the amount of surface hydroxyl groups. This would result in a reduction of Brønsted acidity, and the formation of aprotic oxide (O2-) ions, i.e., Lewis basic sites, in agreement with the results shown in the third and sixth columns of Table 2. It is interesting that the changes in βSurf correlate linearly with both types of acidity, albeit with different slopes, βSurf ) -0.149 + 0.125 (Brønsted acidity; 10-3 × µmol/g); r ) 0.968, and βSurf ) -1.213 + 0.370 RSurf; r ) 0.976.

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TABLE 1: Dependence of the Spectroscopic Data of Adsorbed Probes on TCalcn of TiO2 TCalcn, °C

ν˜ max (Fe(phen)2CN2), 10-3 cm-1

ν˜ max (ABF), 10-3 cm-1

ν˜ max (Michler’s ketone), 10-3 cm-1

λmax (BuPMBr2), nm

original 80 160 240 320 400

19.36 19.39 19.40 19.42 19.45 19.48

19.24 19.32 19.41 19.49 19.57 19.65

24.68 24.73 24.79 24.81 24.82 24.84

418.8 420.3 421.6 422.0 423.1 424.1

TABLE 2: Dependence of the Surface Properties of TiO2 on Its Calcination Temperature, TCalcn TCalcn, °C

surface area, m2/g

Brønsted acidity; × 10-3 µmol/g

ET(30), kcal/mol

RSurf

βSurf

π*Surf

original 80 160 240 320 400 700 P25a

250 248 195 142 135 132 2 52

1.105 0.964 0.922 0.742 0.605 0.456 0.021

70.05 69.69 69.36 69.27 69.01 68.77 65.00 63.2(63.5)a

3.26 3.20 3.14 3.13 3.08 3.03

-0.01 -0.02 -0.04 -0.06 -0.07 -0.09

1.13 1.11 1.09 1.08 1.07 1.05

1.67

-0.41

a

Surface characteristics of this sample were taken from ref 28. Values of ET(30) of P25 were calculated as 63.2 kcal/mol (present work) and 63.5 kcal/mol.24

Figure 3. Correlations of some of the parameters calculated with TCalcn. For uniformity, we used reduced values on the Y axis, calculated from: reduced parameter ) (value at TCalcn - value at 25 °C)/ (value at 25 °C - value at 400 °C).

(iii) The empirical polarities and relative permittivities of protic solvents, e.g., alcohols are usually higher than those of typical aprotic solvents. For example, the values of ET(30) and ε are 55.4 kcal/mol, and 33 (methanol), and 42.2 kcal/mol, and 21.01 (acetone). Therefore, substitution of surface hydroxyl groups by the aprotic (O2-) anions is akin to reduction of the overall surface polarity, and its dipolar character, as shown in the corresponding columns of ET(30) and π*Surf of Table 2. (iv) Based on the preceding discussion, a decrease in Lewis acidity is expected. Equally important, however, is the satisfactory linear correlation between Brønsted and Lewis acidities. This argues against the suggestion that water converts Brønsted sites into Lewis sites,39-41 unless the percentage conversion in both types of sites is the same, or is proportional; this is unlikely. Thus the observed effect of water on the IR bands of adsorbed bases, e.g., ammonia (an increase in NH4+ band at the expense of that of NH3) is probably due to displacement of the strongly adsorbed basic molecules from Brønsted sites, rather than the conversion of Lewis sites into Brønsted sites.42 (v) The term π*Surf stands for surface dipolarity/polarizability. Catala´n et al. have introduced a solvation free-energy equation based on the use of a limited number of solvatochromic indicators, in comparison with those employed in the TaftKamlet-Abboud approach. They have suggested a procedure in order to separate solvent dipolarity from its polarizability, although both properties are linearly correlated for several

solvents, e.g., saturated and aromatic hydrocarbons.43 This separation has not been attempted for TiO2, because we have recently shown that the use of Catala´n’s parameters to describe the solvation of merocyanine probes in 34 protic and aprotic solvents has led to unexplained results. Whereas the susceptibility toward solvent acidity (regression coefficient (a) of eq 1) is similar for both approaches, the susceptibility toward solvent dipolarity/polarizability (regression coefficient (s) of eq 1) is inexplicably low when Catalan’s parameters are employed.44 The origin of this discrepancy is still unclear. (vi) The effect of heating on the surface hydroxyl groups is most readily detected by FTIR. We have calculated the areas of the peaks at ca. 3430 and 1630 cm-1 corresponding to the stretching and bending frequencies of OH groups in the structure,45 Figure 4. As shown in Figure 5 (parts A and B) both areas decrease linearly up to 700 °C. The slopes and intercepts of eqs 9 and 10 that describe this decrease are almost the same showing, convincingly, that both frequencies originate from the same group, namely, surface OH:

ratio of areas at 3430 cm-1 ) 1.039 - 0.001TCalcn

r ) 0.9876 (9)

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Figure 4. IR absorption spectra of the original TiO2 (A) and that calcinated at 400 °C (B).

ratio of areas at 1630 cm-1 ) 0.940 - 0.001TCalcn

r ) 0.9700 (10)

(vii) An interesting observation is that the empirical polarity, RSurf, and π*Surf of all samples, even that heated at 700 °C (only ET(30) was determined) are higher than those of water, and alcohols;32,33 a similar result has been observed before for different solids, e.g., P25, and alumosilicates.24 The high empirical polarity may be attributed to dipole-dipole interaction between the surface and the probe, coupled with hydrogenbonding between the surface hydroxyl groups and the phenolate oxygen of the adsorbed probe,46 see Figure 6. These interactions stabilize the (zwitterionic) ground state of BuPMBr2, relative to its (quinonoid) excited state, leading to an increase in ET(BuPMBr2). A similar rationale can be invoked to explain the higher values of the other solvatochromic properties. (viii) Figure 6 shows two probe-TiO2 interaction schemes. In Figure 6A, the interaction shown is that between the probe phenolate oxygen and the surface OH group; in Figure 6B, both ions of the probe are shown to interact with the OH groups. The representation depicted in Figure 6A is in line with previous literature, e.g., that on the interactions between the OH groups of solid surfaces (containing both Lewis and Brφnsted acids) and the oxygen atom of the CdO group of Michler’s ketone.24,29 It is also in agreement with the low βSurf values reported in Table 2, and with the adsorption of amino acids, e.g., glycine and proline on TiO2. In the latter, the zwitterions lie perpendicular

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Figure 6. Two schematic representations of the interactions of BuPMBr2 with TiO2. For simplicity, we represent the titania surface by the hatched area, and place the OH groups symmetrically, perpendicular to the surface.

to the surface, due to the interaction of the carboxylate group of the amino acid with the surface OH groups.47 In order to shed more light on the problem of probe orientation, we have carried out theoretical calculations on model systems, namely the solvation of the probe by methanol and diethyl ether. The former solvent can act as an acid, by interacting with the probe phenolate oxygen (probe-O- · · · H-OCH3), or as a base, (probe-N+ · · · O(H)CH3). On the other hand, diethyl ether can act only as a base, via (probe-N+ · · · O(C2H5)2). Therefore, comparing the energetic of solvation in these solvents should corroborate our decision about part (A) of Figure 6. The calculated Gibbs free energies of solvation of BuPMBr2 in methanol and diethyl ether were found to be -22.5 and -11.9 kcal/mol, respectively. As shown by the ionization potentials, IP, diethyl ether (IP ) 9.51 eV) is more basic than methanol (IP ) 10.85 eV),48 i.e., the interaction (probe-N+ · · · O(C2H5)2) is expected to be stronger than (probe-N+ · · · O(H)CH3). Nevertheless, probe solvation by methanol is energetically much more favorable than that by ether, most certainly because (probe-O- · · · H-OCH3) is the dominant solute-solvent interaction; this favors the representation shown in Figure 6A. Monte Carlo (MC) simulations corroborate the latter conclusion, as clearly shown in Figure 7. The parts of this figure refer to the radial distribution functions (Gr) of the relevant atoms of the solvents around (probe-O-) and (probe-N+). Namely, (probe-O- · · · O(H)CH3), Figure 7A; (probe-O- · · · H-OCH3), Figure 7B; (probe-N+ · · · O(H)CH3),

Figure 5. Ratios of peak areas relative to that of sample 1 (original). Part (A) is for ν˜ OH at ca. 3430 cm-1 (area of sample 1 ) 694.60 a.u.); part (B) is that for δOH at ca. 1630 cm-1 (area of sample 1 ) 69.73 a.u.).

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Figure 7. Radial distribution functions (Gr) of the phenolate oxygen and quaternary nitrogen of BuPMBr2 and the atoms of methanol or diethyl ether. The parts refer to (probe-O- · · · O(H)CH3), A; (probe-O- · · · H-OCH3), B; (probe-N+ · · · O(H)CH3), C; (probe-O- · · · O(C2H5)2), D; and (probe-N+ · · · O(C2H5)2), E.

TABLE 3: Equations Used to Calculate the Solvatochromic Parameters

a

equation

r

sd

ET(BuPMBr2) ) 21.547 + 0.667ET(30)a RSurf ) (ET(30) - 52.19 - 0.99π*)/5.12b βSurf ) 3.06 + 0.17ν˜ max(Fe(phen)2CN2)10-3 - 0.33ν˜ max(ABF)10-3c π*Surf ) 13.89 - 0.251ν˜ max(Fe(phen)2CN2)10-3 - 0.32ν˜ max(Michler’s ketone)10-3b

0.9727 0.964 0.96 0.57

1.0753 0.24 0.10 0.15

Reference 37. b Reference 24. c Reference 51.

Figure 7C; (probe-O- · · · O(C2H5)2), Figure 7D; and (probe-N+ · · · O(C2H5)2), Figure 7E. We discuss parts (A) and (B) together. The calculated average radial distance between the two oxygen atoms (Gr; probe-O- · · · O(H)CH3)) is 2.75 Å, whereas (Gr; O- · · · H-OCH3) is 1.85 Å. The latter distance is in good agreement with the value of 1.908 Å, calculated (ab initio) for the length of the hydrogen bond between methanol and diacetamide (CH3O-H · · · (OdCCH3)2NH).49 The difference between (probe-O- · · · O(H)CH3) and (probe-O- · · · H-OCH3), 0.9 Å, represents the length of the O-H bond of methanol; it is slightly smaller than the standard value recommended for MC simulations, 0.945 Å.50 This can be taken to indicate that the average (O- · · · H-O) angle of (probe-O- · · · H-OCH3) is 162°. With regard to Figure 7 C, the (probe-N+ · · · O(H)CH3) distance is 4.65 Å, which is much larger than the length of an N-O covalent bond, 1.1 to 1.2 Å;48 that is, this is a negligible interaction. The average radial distances calculated for parts (D) and (E) of Figure 7 are 4.65 and 4.55 Å; both are much larger that the respective O-O and N-O covalent bonds, 1.2 and 1.1-1.2 Å, respectively.48 In summary, the only strong solute-solvent interaction is (probe-O- · · · H-OCH3). By analogy, provided that solute-solvent interactions can be taken as a proper model for the probe-TiO2 counterpart, it is expected that the relevant probe-TiO2 interaction is (probe-O- · · · H-O-Ti), as depicted in Figure 6A. (ix) We now address the questions raised in the Introduction about the use of solvatochromic dyes to probe the surface of TiO2. First, we have calculated a set of important surface properties by employing a small number of probes; this is not feasible with the other techniques. For example, there is no obvious way to measure the overall surface polarity, except by use of solvatochromism. Although π*Surf has not been separated into its components, it is probably a more useful quantity than the dispersion term alone (determined by inverse gas chromatography) especially for reactions whose transition states are

dipolar, i.e., where dipole-dipole interactions are dominant. Knowledge of ET(30) values of both surface and solvent helps in choosing the latter, i.e., one that is dipolar enough to solvate the species of interest without excessive deactivation of the catalyst. The protic and dipolar nature of the titania samples should be important for reactions where there is large polarity difference between the reagents and transition states. Surface polarity and acidity decrease linearly as a function of increasing TCalcn; that is, the catalyst efficiency can be “fine-tuned” by heat treatment. In summary, the information obtained from the use of these probes allows a better control of the efficiency and selectivity of TiO2, and solid catalysts in general. Conclusions The use of solvatochromic indicators in order to investigate surface properties of solid surfaces is clearly advantageous. The probes are readily synthesized and the experimental part is straightforward, provided that certain precautions are taken into account. More importantly, these probes offer data that are difficult to be quantified by other techniques, and that are important to applications, e.g., in adsorption and heterogeneous catalysis. These include the empirical (total) polarity; Lewis acidity and basicity, and the dipolarity/polarizability. Except for basicity, all parameters calculated decrease as a function of increasing TCalcn. Theoretical calculations on model systems are very helpful in probing dye adsorption on titania. The present study indicates that catalytic efficiency and selectivity of TiO2 can be adjusted for the application sought by help of solvatochromic data. For our samples, Brφnsted and Lewis acidities correlated linearly. If this (linear) trend proves to be general, then solvatochromism, along with simple titration, can be employed as an expedient procedure for determining the two types of acid sites.

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Calculations Calculations of the Solvatochromic Parameters from the UV-Vis Data. We have employed Shimadzu UV-Probe, version 2.21 program for the transformation of the reflectance spectra into the corresponding absorption curves, and for calculation of the values of λmax. In the FTIR measurements, the reflectance from the solid mixture was transformed into absorbance (Opus version 3.1, Bruker), and the areas of the OH peaks were calculated by a commercial software (GRAMS/32 version 5.10, Galactic Industries). The equations in Table 3 have been employed in order to calculate the parameters, where r, sd, and ν˜ refer to the correlation coefficient, the standard deviation, and frequency of the intramolecular charge transfer band, respectively. Other equations are known to calculate some of these solvatochromic parameters.28 Although the choice can be dictated by the nature of the solid surface (e.g., (Fe(phen)2CN2)) is suitable to examine relatively acidic surfaces), and maybe a matter of personal preference, the results obtained should agree with chemistry. For example, our RSurf decrease as a function of increasing TCaln, in agreement with the results of (reliable) acid-base titration. Simulation of the Interactions BuPMBr2-TiO2: Ab Initio Calculations and Monte Carlo Simulations of the Solvation of BuPMBr2 by Methanol and Diethyl Ether. Ab Initio Calculations of the Energy of SolWation of BuPMBr2. The geometry of merocyanine in the gas phase was optimized by using B3LYP hybrid density functional with 6-311+G(d,p) basis set, as implemented in Gaussian 03 rev. D.01.52 Atomic charges and solvation free energy (for solvent continuum) were calculated by using the Merz-Kollman-Singh scheme of electrostatic potential-derived charges, and IEF-PCM solvation model, respectively. Monte Carlo Simulations. The structure of solvent molecules (methanol or diethyl ether) around the merocyanine probe has been calculated by using MC simulations, by employing standard procedures for the Metropolis sampling technique in the isothermal-isobaric ensemble, where the number of molecules N, the pressure p (1 atm), and the temperature T (298 K), are fixed. We have used the periodic boundary conditions and image method in a cubic box with one merocyanine molecule embedded in 900 molecules of the solvent. The merocyanine and the solvent molecules interact by the LennardJones and Coulomb potentials. The all-atom optimized parameters were employed for liquid simulations (OPLS/AA).53 The MC simulation was performed with the DICE program,54 and involved a thermalization stage of 15 000 MC steps followed by an averaging stage of 90 000 MC steps. In this averaging stage, the radial distribution functions (rdf) were calculated. The conformations of methanol and diethyl ether, Lennard-Jones, and Coulomb potentials were those published elsewhere.50,55 Figures SI-1 (Figure 1 of Supporting Information) and SI-2 show snapshots of one of the conformations calculated by MC simulations. They show the BuPMBr2 molecule surround by the first solvation shell of methanol and diethyl ether molecules, respectively. Figure SI-3 shows the molecular structure and atom numbering of BuPMBr2, whereas Table SI-1 shows the X, Y, and Z coordinates, charge (q), and Lennard-Jones ε and σ parameters of BuPMBr2 atoms employed in the ab initio calculations and MC simulations. Acknowledgment. We thank the AUC for research funds to A.R.R., FAPESP (State of Sa˜o Paulo Research Foundation) for financial support to O.A.E.S. and a predoctoral fellowship to

El Seoud et al. B.M.S., and the CNPq (National Council for Scientific and Technological Research) for a research productivity fellowship to O.A.E.S. We are indebted to Profs. S. Spange (Chemnitz), J. Ragai (AUC, Cairo), and K. Canuto (Institute of Physics, the University of Sa˜o Paulo, Sa˜o Paulo) for a sample of the ABF indicator, for relevant comments on the work, and for helpful discussions on the use of the DICE program, respectively. Supporting Information Available: Figures SI-1 and SI2, snapshots of BuPMBr2 solvated by methanol and by diethyl ether, respectively; Figure SI-3 the molecular structure and atom numbering of BuPMBr2; Table SI-1, the X, Y, and Z coordinates; charge; Lennard-Jones parameters of BuPMBr2 atoms employed in the theoretical calculations. This material is available free of charge via the Internet at http://pubs.acs.org. References and Notes (1) Tomoi, M.; Ford, W. T. J. Am. Chem. Soc. 1981, 103, 3821. (2) Tundo, P.; Venturello, P. J. Am. Chem. Soc. 1979, 101, 6606. (3) Regen, S. L.; Bolikal, D. J. Am. Chem. Soc. 1981, 103, 5248. (4) Zimmermann, Y.; Spange, S. J. Phys. Chem. B 2002, 106, 12524. (5) Corma, A. Chem. ReV. 1995, 95, 559. (6) Haw, J. B.; Nicholas, J. B.; Xu, T.; Beck, L. W.; Ferguson, D. B. Acc. Chem. Res. 1996, 29, 259. (7) Pfaff, G.; Reynders, P. Chem. ReV. 1999, 99, 1963. (8) Salvador, A.; Pascual-Marti, M. C.; Adell, J. R.; Requeni, A.; March, J. G. J. Pharm. Biomed. Anal. 2000, 22, 301. (9) Braun, J. H.; Baidins, A.; Marganski, R. E. Prog. Org. Coat. 1992, 20, 105. (10) Yuan, S. A.; Chen, W. H.; Hu, S. S. Mater. Sci. Eng., C 2005, 25, 479. (11) Hagfeldt, A.; Graetzel, M. Chem. ReV. 1995, 95, 49. (12) Linsebigler, A. L.; Lu, G.; Yates, J. T., Jr. Chem. ReV 1995, 95, 735. (13) Chen, X.; Mao, S. S. Chem. ReV. 2007, 107, 2891. (14) Boehm, H. P. Discus. Faraday Soc. 1971, 52, 264. (15) Benesi, H. A. J. Phys. Chem. 1957, 61, 970. (16) (a) Kno¨zinger, H. Nato ASI Series, Series C 1988, 231, 35–46. (b) Farcasiu, D.; Lezcano, M.; Lukinskas, P.; Waldeck, D. H. J. Phys. Chem. A 2000, 104, 5190. (17) Kno¨zinger, H. In Fundamental aspects of heterogeneous catalysis studied by particle beams; Brongersma, H. H., van Santen, R. A., Eds.; Plenum Press: New York, 1991; p 167. (18) Hungerford, G.; Pereira, M. R.; Ferreira, J. A.; Viseu, T. M. R.; Coelho, A. F.; Isabel, M.; Ferreira, C.; Suhling, K. J. Fluoresc. 2002, 12, 397. (19) Mueller, R.; Kammler, H. K.; Wegner, K.; Pratsinis, S. E. Langmuir 2003, 19, 160. (20) Chen, Y.; Dionysiou, D. D. J. Mol. Cat. A 2006, 244, 73. (21) Papirer, E.; Brendle´, E. J. Chim. Phys. 1998, 95, 122. (22) Sun, C.; Berg, J. C. J. Chromatogr. A 2002, 969, 59. (23) Burry, W. M.; Keller, D. S. J. Chromatogr. A 2002, 972, 241. (24) Spange, S.; Vilsmeier, E.; Zimmermann, Y. J. Phys. Chem. B 2000, 104, 6417. (25) Michels, J. J.; Dorsey, J. G. Langmuir 1990, 6, 414. (26) Macquarrie, D. J.; Tavener, S. J.; Gray, G. W.; Heath, P. A.; Rafelt, J. S.; Saulzet, S. I.; Hardy, J. J. E.; Clark, J. H.; Sutra, P.; Brunel, D.; Renzo, F.; Fajula, F. New J. Chem. 1999, 23, 725. (27) Khristenko, I. V.; Kholin, Y. V.; Mchedlov-Petrossyan, N. O.; Reichardt, C.; Zaitsev, V. N. Colloid J. 2006, 68, 511. (28) Spange, S.; Prause, S.; Vilsmeier, E.; Thiel, W. R. J. Phys. Chem. B 2005, 109, 7280. (29) Zimmermann, Y.; Anders, S.; Hofmann, K.; Spange, S. Langmuir 2002, 18, 9578. (30) Abraham, M. H.; Grellier, P. L.; Abboud, J. L. M.; Doherty, R. M.; Taft, R. W. Can. J. Chem. 1988, 66, 2673. (31) Laurence, C.; Nicolet, P.; Dalati, M. T.; Abboud, J. L. M.; Notario, R. J. Phys. Chem. 1994, 98, 5807. (32) Reichardt, C. SolVents and SolVent Effects in Organic Chemistry, 3rd ed.; Wiley-VHC: New York, 2003; pp 5, 329, and 389. (33) Reichardt, C. Pure Appl. Chem. 2004, 76, 1903. (34) Reichardt, C. Pure Appl. Chem. 2008, 80, 1415. (35) El Seoud, O. A. Pure Appl. Chem. 2009, 81, 697. (36) Schilt, A. A. J. Am. Chem. Soc. 1960, 82, 3000. (37) Martins, C. T.; Lima, M. S.; El Seoud, O. A. J. Org. Chem. 2006, 71, 9068. (38) Gregg, S. J.; Sing, K. S. W. Adsorption, Surface Area and Porosity; Academic Press: London, 1982.

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