Evidence for Discrepancy between the Surface Lewis Acid Site

J. Phys. Chem. B 2004, 108, 16499-16507

16499

Evidence for Discrepancy between the Surface Lewis Acid Site Strength and Infrared Spectra of Adsorbed Molecules: The Case of Boria-Silica A. Travert,*,† A. Vimont,† J.-C. Lavalley,† V. Montouillout,†,‡ M. Rodrı´guez Delgado,§ J. J. Cuart Pascual,§ and C. Otero Area´ n§ Laboratoire Catalyse et Spectrochimie, CNRS-ENSICAEN, UniVersite´ de Caen, 6 BouleVard du Mare´ chal Juin, 14050 Caen Cedex, France, and Departamento de Quı´mica, UniVersidad de las Islas Baleares, 07122 Palma de Mallorca, Spain ReceiVed: May 14, 2004; In Final Form: July 23, 2004

The acidity of amorphous B2O3-SiO2 has been investigated by infrared spectroscopy using the following three probe molecules presenting a wide range of basic strength: pyridine, acetonitrile, and carbon monoxide. The results are compared to those obtained on γ-Al2O3. No coordination of carbon monoxide is observed on B2O3-SiO2 even at low temperatures, whereas strongly coordinated CO species are formed on γ-Al2O3 under such conditions. Coordinated pyridine and acetonitrile show important infrared frequency shifts on both metal oxides, indicating strong charge transfer from the probe molecules to the surface Lewis acid centers. However, the thermal stability of coordinated species is much lower on B2O3-SiO2 than on γ-Al2O3, which suggests that there is no direct correlation between charge transfer and adsorption energy. Density functional theory (DFT) calculations on the interaction of these probe molecules with simple models representing Al3+ and B3+ Lewis acid sites adequately reproduce experimental observations. The main difference between Al3+ and B3+ results from the higher energy required to convert boron from a trigonal planar conformation to a tetrahedral conformation upon adsorption of the probe molecule. Despite strong charge transfer, this leads to a weaker adsorption of pyridine and acetonitrile on B3+ as compared to Al3+ Lewis acid sites. Carbon monoxide is not basic enough to compensate for the energy required for the conformational change of the B3+ Lewis acid center.

1. Introduction Infrared spectroscopy of adsorbed basic probe molecules is one of the most often used technique for characterizing the acidic properties of solid surfaces, particularly those of metal oxides.1 Provided that the molecular probe has been well chosen, its infrared spectrum shows absorption bands that are characteristic of its interaction with the surface (hydrogen bonding, protonation, or coordination) and allows the nature (Brønsted or Lewis acid) of the surface adsorption sites to be determined. The number of surface acid sites of each type can then be assessed by measuring the intensities of the IR bands of the adsorbed probe molecules, provided that their molar absorption coefficients are known. More problematic is the determination of the strength of surface acid sites. Infrared spectroscopy could allow this measurement to be performed, provided that correlations between the IR spectra and adsorption heats can be established. The Brønsted acid sites of metal oxides present the advantage of being directly observable by infrared spectroscopy through the ν(OH) absorption bands when H-bonded complexes are formed. By analogy with liquid solutions,2 empirical relationships between ν(OH) frequency shifts and the adsorption heats of H-bonded basic probe molecules could be established;3,4 the higher the ν(OH) shift, the higher the adsorption heat. When * Corresponding author. Phone: +33(0)2 31 45 28 23. Fax: +33(0)2 31 45 28 21. E-mail: [email protected]. † Universite ´ de Caen. ‡ Present address: CNRS-CRMHT, 1D Avenue de la Recherche Scientifique, 45071 Orleans Cedex 2, France. § Universidad de las Islas Baleares.

spectra in the ν(OH) range are too complex to be accurately analyzed, which is often the case for many metal oxides, frequency shifts of specific IR absorption bands of the adsorbed probe molecule can also give useful information on surface acid strength; referring to the infrared spectrum of the probe molecule in a liquid or gas phase, shifts of specific IR bands can thus be correlated with interaction energy.5-7 On the other hand, Lewis acid sites which are formally electron lone pair acceptors cannot be directly detected by infrared spectroscopy. Hence, the strength of these sites can only be assessed by the infrared spectra of molecules interacting with them. Charge transfer occurring from the adsorbed probe molecule leads to important modifications of its electron density, which in turn is reflected in frequency shifts of characteristic infrared absorption bands.8 Carbon monoxide,1,9-12 acetonitrile,1,13-16 and pyridine1,17-21 are the most frequently used molecules for probing surface Lewis acidity; their ν(CO), ν(CtN), and ν(CdC) (ν8a and ν19b ring vibrations) bands are very sensitive to the charge transfer to the adsorption site:

Despite their very different basicities, these probe molecules give rise to similar trends for the surface Lewis acidity of metal oxides showing a wide range of acid strength.1,16 In particular, the observed frequency shifts were directly correlated to the

10.1021/jp0479365 CCC: $27.50 © 2004 American Chemical Society Published on Web 09/23/2004

16500 J. Phys. Chem. B, Vol. 108, No. 42, 2004 Lewis acid site strength which is expected on the basis of the charge, coordinative unsaturation, and polarizing power of the cation.16 Moreover, the adsorbed molecules showing the higher frequency shifts are also those that are most stable toward thermal desorption16,19,21 which indicates that these species are more strongly adsorbed on the surface. Hence, all this experimental evidence tends to demonstrate a general relationship between the IR spectra of adsorbed probe molecules and the strength of the Lewis acid sites to which they are coordinated; the higher the shift, the stronger the site. This paper reports on the case of boria-silica mixed oxides for which the above relationship is clearly ruled out. In the first instance, the adsorption of probe molecules (pyridine, acetonitrile, and carbon monoxide) on B2O3-SiO2 as followed by IR spectroscopy is described. Then, the obtained results are compared to the case of γ-Al2O3, which has been often described in the literature10,22,23 and which is well-known to show strong Lewis acid sites. It will be shown that the adsorption of pyridine and acetonitrile on B2O3-SiO2 leads to coordinated species presenting higher IR frequency shifts than those observed on Al2O3, but a much lower thermal stability. Also, density functional theory (DFT) calculations on the interaction of the probe molecules with small clusters representing surface Al3+ and B3+ Lewis acid sites will be given and compared with the experimental results. 2. Experimental Section To prepare the boria-silica samples, boric acid (analytical grade) was dissolved in ethanol, ∼10 g in 25 mL. To aliquots of this solution, silicon tetramethoxide was added in the required amounts to have nominal compositions of 5, 10, and 20 wt % B2O3; the corresponding samples are hereafter denoted BS-5, BS-10, and BS-20, respectively. For each sample, the solution was kept under stirring at room temperature for 30 min, after which concentrated aqueous ammonia was added dropwise until no further precipitate was formed. The resulting gel was aged 72 h at room temperature, washed with ethanol, vacuum-dried in a desiccator, and calcined for 1 h at 573 K. The values of the specific surface area, calculated from the corresponding nitrogen sorption isotherms, were found to be 611, 508, and 408 m2 g-1 for BS-5, BS-10, and BS-20, respectively. γ-Al2O3 was obtained from Rhone-Poulenc; it had a specific surface area of 223 m2 g-1. Infrared transmission spectra were recorded on self-supporting wafers (2 cm2, 8-15 mg) which were placed into an infrared quartz cell (KBr windows) connected to a vacuum line. They were activated under vacuum (10-4 Pa) by slowly heating (1 K min-1) them up to 673 K and then outgassing them at this temperature for a period of 2 h. Additions of accurately known amounts of probe molecules into the cell were made by using a calibrated volume and measuring the pressure inside the IR cell. Spectra were recorded at room temperature at a resolution of 4 cm-1. The IR spectrometer was a Nicolet Nexus apparatus equipped with an extended KBr beam splitter and a mercury cadmium telluride (MCT) detector. CH3CN, CD3CN, and pyridine (Aldrich, 99+% grade) were dried on molecular sieves prior to use. The isotopic purities of CD3CN and CH3C15N were 99.95 and 95%, respectively. CO was provided by Air Liquide. 3. Computational Methods The Lewis acid sites of γ-Al2O3 and SiO2-B2O3 were modeled using a cluster approach. DFT calculations were carried

Travert et al. out with the Gamess package24 using the B3LYP functional25 and the 6-31+G** basis set.26 Recent periodic DFT simulations have shown that the strongest Lewis acid sites of dehydroxylated γ-Al2O3 surfaces are the trigonal planar AlIII sites of the (110) plane.27 Much less is known about the local structure of the Lewis acid centers of the SiO2-B2O3 systems. However, it will be shown that trigonal BO3 units, characterized by an IR band at 1380 cm-1, are directly involved in the coordination of pyridine and therefore are located at the surface. More problematic is the determination of the environment of these surface BO3 units of the SiO2-B2O3 samples. It has been shown, however, that the coordinative environment of oxygen has only a limited influence on B-O bond lengths,28 that is, the very local structure of the adsorption site. Thus, in the absence of detailed information, B3+ Lewis acid sites were modeled by a B(OH)3 cluster (denoted BIII hereafter) where hydrogen atoms ensure charge neutrality. The B-O bond length was 138 pm at the equilibrium geometry, which corresponds to the mean value for trigonal borates.28 For consistency, a similar Al(OH)3 model was used for Al3+ Lewis acid sites. To avoid internal hydrogen bonds, the OH bonds were forced to lie in planes perpendicular to the MO3 plane (M ) Al or B), with all other internal coordinates being fully relaxed. With these constraints, Al and B were found to lie slightly below the oxygen plane, with the M-O bonds forming angles of ∼86° with the C3 axis. Surface relaxation of metal oxides usually gives rise to similar rearrangements.29 Assuming that cation coordination essentially determines the adsorptive properties of Lewis acid centers, it is expected that these models, despite their small size, can qualitatiVely reproduce the main experimental observations reported in the present study. The adsorption of pyridine, acetonitrile, and carbon monoxide on these clusters was computed at the equilibrium geometry (full relaxation) as well as at various intermolecular distances. The vibrational frequencies of the complexes and monomers were computed at their equilibrium geometries by finite differences and scaled with respect to the experimental values in the gas phase. For all complexes, the total interaction energy was computed as

∆Etot ) Ec(AB) - E°(A) - E°(B) where Ec(AB) is the total energy of the complex and E°(A) and E°(B) are the energies of the isolated monomers at their equilibrium geometries. Basis set superposition errors (BSSEs) were checked at the equilibrium geometries of the complexes with the counterpoise (CP) method.30 To take into account the effect of the deformation of the Lewis center and of the probe molecule upon complexation, the total interaction energies were separated into inter- and intramolecular terms:31

∆Etot ) ∆Eint + ∆Edef(A) + ∆Edef(B) where ∆Edef(A) and ∆Edef(B) are the deformation energies of the acid center and the probe molecule, respectively:

∆Edef(A) ) Ec(A) - E°(A) and ∆Edef(B) ) Ec(B) - E°(B) where Ec refers to the energy of the isolated species at their geometry in the complex. ∆Eint is the interaction energy of the probe molecule with Lewis acid sites at their geometries in the complex:

Lewis Acid Site Strength and IR Spectra of B2O3-SiO2

J. Phys. Chem. B, Vol. 108, No. 42, 2004 16501

∆Eint ) Ec(AB) - Ec(A) - Ec(B) Finally, the amounts of charge transfer (∆q) from the probe molecules to the Lewis acid site were determined from Mulliken charges. It should be noted that the calculation of partial charges is always arbitrary. In particular, the absolute values of charge transfer (∆q) should not be given much meaning, and only relative values of ∆q should be compared. 4. Results and Discussion 4.1. IR Spectra of Activated Samples. Figure 1 shows the spectra of the BS-5, BS-10, and BS-20 samples after being outgassed at 673 K. In the ν(OH) range, the main band at 3745 cm-1 is assigned to the ν(OH) vibration of isolated silanol groups. A broad band with a low intensity, between 3400 and 3700 cm-1, characterizes H-bonded hydroxyl groups. This band was also observed in the spectrum of pure silica and assigned to vicinal and inner silanol groups.32 Another weak, well-defined band with a low wavenumber tail is detected at 3703 cm-1. This band, which was not observed in the IR spectrum of silica, is assigned to the ν(OH) vibration of B-OH groups.33 Its intensity increases with the boron content in the sample at the expense of the silanol groups. In the 1500-1200 cm-1 range, the strong band seen at 1380 cm-1 is assigned to the ν(B-O) vibration mode of the trigonal BO3 units, in agreement with a similar assignment for the IR spectra of SiO2-B2O3 glasses.34,35 The very high intensity of this latter band in BS-10 and BS-20 precludes a study of intensity variation versus boron content. 4.2. Pyridine Adsorption. Figure 2 shows that the adsorption of pyridine (133 Pa) at room temperature on the activated samples gives rise to bands at 1596 and 1445 cm-1, both on pure silica (spectrum a) and on B2O3-SiO2 samples (spectra b-d). These bands correspond to the ν8a and ν19b modes of H-bonded pyridine species. Accordingly, in the ν(OH) range (not shown), the silanol band (3745 cm-1) was found to be perturbed and shifted down to ∼3000-2800 cm-1, as also described in the literature.36 Simultaneously, a decrease in the intensity of the boranol band was observed, suggesting that these hydroxyl groups are also involved in pyridine adsorption. In no case were pyridinium species (characteristic band at 1545 cm-1) detected. All of the bands of H-bonded species were nearly completely removed by outgassing at room temperature (Figure 3A, spectrum a). Bands at 1626 and 1459 cm-1 are observed (Figure 2) only on B2O3-SiO2 samples and not on pure silica. These bands correspond to the ν8a and ν19b modes, respectively, of coordinated pyridine on Lewis acid sites (PyL).17 Interestingly, these bands are close to those (1629 and 1464 cm-1) observed for the pyridine-BF3 complex in benzene solution.38 Comparison of the spectra after pyridine adsorption at room temperature (Figure 2, spectra b-d) shows that the intensity of the bands at 1626 and 1459 cm-1 characterizing PyL species increases with the quantity of boron present in the sample. Quantitative measurements have been performed on BS-5 by adding at 373 K successive doses (10 µmol g-1) of pyridine on the activated sample. Under these conditions, only PyL species are formed. The variation of the integrated area of the ν19b band (1459 cm-1) versus the amount of pyridine dosed was found to be linear and led to an approximate value of 2.4 µmol cm-1 for the corresponding molar absorption coefficient. From this value, the amounts of coordinated pyridine on the surface of the different samples were estimated. They are approximately 130, 360, and 400 µmol g-1 for BS-5, BS-10, and BS-20, respectively. From theses values, it is deduced that the percentage of

Figure 1. Infrared spectra of boria-silica after activation at 673 K under vacuum (10-4 Pa): (a) BS-5; (b) BS-10; (c) BS-20.

Figure 2. Infrared spectra after the introduction of 133 Pa of pyridine into the IR cell: (a) silica; (b) BS-5; (c) BS-10; (d) BS-20. The inset shows the band at 1380 cm-1 before and after (line labeled c) pyridine adsorption on BS-5.

boron atoms involved in the formation of PyL species is ∼12% in BS-5 and BS-10 and only 7% in the case of BS-20. Lower dispersion of boria in the latter sample could account for this lower fraction of boron Lewis acid sites. Taking into account the specific surface areas, we find that the surface concentrations of Lewis acid sites of the BS samples are 0.1, 0.4, and 0.6 nm-2 for BS-5, BS-10, and BS-20, respectively. Interestingly, for the BS-10 and BS-20 samples, these values are close to those reported for the concentration of strong Lewis acid sites on alumina.37 The intensity of the IR absorption bands at 1627 and 1461 cm-1 (sample BS-20) was found to decrease slightly upon evacuation at room temperature, as shown in Figure 3A. In the case of the ν8a band at 1627 cm-1, this decrease could be partly due to the disappearance of a weaker band at 1622 cm-1 assigned to the (ν1 + ν6a) combination mode of H-bonded species, as observed in the spectrum of pure silica (Figure 2, spectrum a). Outgassing the samples at 373 K causes a further intensity decrease of the bands at 1627 and 1461 cm-1 due to coordinated species (Figure 3A, spectrum b). Both bands are hardly detectable after outgassing at 423 K (spectrum c) and completely disappear at 473 K (spectrum d). The intensity of the band at 1380 cm-1 in the spectrum of the activated BS-5 sample partly decreases after pyridine adsorption (inset of Figure 2). This band is still perturbed after outgassing at room temperature but is fully restored after the complete elimination of coordinated pyridine species, suggesting that it can be assigned to a surface vibration mode sensitive to

16502 J. Phys. Chem. B, Vol. 108, No. 42, 2004

Travert et al.

Figure 3. (A) Infrared spectra of BS-20 after thermodesorption of pyridine under vacuum at (a) room temperature, (b) 373 K, (c) 423 K, and (d) 473 K. (B) Infrared spectra recorded on Al2O3 after thermodesorption of pyridine under vacuum at (e) room temperature, (f) 373 K, (g) 423 K, (h) 473 K, and (i) 673 K.

the formation of coordinated species. Similar behavior has been reported by Scarano et al.39 in the case of the adsorption of several probe molecules on boralite materials. Bands in the 1500-1200 cm-1 range are usually attributed to the B-O stretching modes of trigonal planar BO3 units,35,39 which is the dominant conformation of boron in boron-oxygen compounds.40 This suggests that the surface sites of B2O3-SiO2 leading to the coordination of pyridine are trigonal planar BO3 Lewis acid sites. Figure 3B shows the infrared spectra obtained after pyridine adsorption on Al2O3. After dosing with pyridine (133 Pa) followed by outgassing at room temperature (spectrum e), two types of coordinated species characterized by ν8a bands at ∼1615 and ∼1625 cm-1 are detected, in agreement with previous studies.22 A band at 1449 cm-1, corresponding to the ν19b mode of both types of coordinated species, is also seen. From the intensity of this band, the surface concentration of total Lewis acid sites is estimated to be close to 0.9 nm-2. This surface density is comparable to that of the BS-20 sample. From desorption experiments at increasing temperatures (Figure 3B, spectra f-i), it clearly appears that coordinated species characterized by the band at 1625 cm-1 are more strongly adsorbed than those corresponding to the band at 1615 cm-1, since the intensity of the latter preferentially decreases as the temperature is increased. As previously reported,19,22 this clearly shows that Al2O3 coordinated species characterized by higher ν8a frequencies are more strongly held at the surface, since they are still present after outgassing above 473 K. In the case of B2O3-SiO2, the pyridine bands at 1627 and 1461 cm-1 (Figure 3) are clearly located at higher frequencies than those observed on alumina (1625 and 1455 cm-1) and even on AlF3.16 This suggests an important charge transfer upon the coordination of pyridine to the surface of B2O3-SiO2, presumably higher than that occurring on Al2O3. However, despite this strong charge transfer, a much easier thermodesorption of pyridine on B2O3-SiO2 is observed, since all species desorb after outgassing at 473 K, whereas, on alumina, coordinated pyridine species are still present after desorption at 673 K (Figure 3B, spectrum i). 4.3. Acetonitrile Adsorption. The spectra obtained after dosing with 665 Pa of CH3CN, CD3CN, and CH3C15N on activated SiO2 and on BS-20 are presented in Figure 4. In the 2100-2400 cm-1 range, ν(CtN) vibrations are observed as a doublet at 2265 and 2297 cm-1 for CH3CN due to the Fermi resonance between ν(CtN) and the (ν(CsC) + δs(CH3)) combination mode (Figure 4A). The very weak bands at 2214 and 2409 cm-1 correspond to those observed at 2202 and 2410 cm-1 in the spectrum of CH3CN in a liquid phase and are due

Figure 4. (A) CH3CN adsorption: (a) 665 Pa on SiO2; (b) 665 Pa on BS-20; (c) after the outgassing of BS-20 at room temperature. (B) CD3CN adsorption: (d) 665 Pa on SiO2; (e) 665 Pa on BS-20; (f) after the outgassing of BS-20 at room temperature. Inset: spectra j and k were recorded on BS-5 before and after dosing with 665 Pa of CD3CN. (C) CH3C15N adsorption: (g) 665 Pa on SiO2; (h) 665 Pa on BS-20; (i) after the outgassing of BS-20 at room temperature.

to combination modes.41 In the case of CD3CN, for which no resonance arises, ν(CtN) occurs as a single band at 2272 cm-1 (Figure 4B). Other weaker bands are observed at 2120 and 2218 cm-1 and are assigned to νs(CD3) and νas(CD3), respectively. All of these bands are observed on both SiO2 (Figure 4B, spectrum d) and BS-20 (Figure 4B, spectrum e), and consequently, they are assigned to the MeCN‚ ‚ ‚HO-Si complex (the corresponding ν(OH) of the perturbed silanol bands was found at 3420 cm-1). It is important to note in the spectra of BS-20 an additional weak band at 2349 cm-1, regardless of whether CH3CN or CD3CN is used. This band is reversible. It appears under an



Figure 5. CD3CN adsorption on Al2O3: (a) 133 Pa at room temperature and after outgassing at (b) room temperature, (c) 373 K, and (d) 473 K.

Figure 6. Infrared spectra of an activated BS-20 sample (spectrum a) and after the introduction of 200 and 400 µmol g-1 and 100 Pa of CO at 100 K (spectra b, c, and d, respectively).

equilibrium pressure of acetonitrile, its intensity increases with pressure, and it totally disappears upon outgassing at room temperature (Figure 4A, spectrum c, and Figure 4B, spectrum f). This indicates that this band corresponds to very weakly adsorbed acetonitrile species. Since this band is not detected on silica, it cannot be assigned to a combination or overtone mode due to H-bonded or physisorbed acetonitrile species. When CH3C15N is used (Figure 4C, spectrum h), this band is shifted to 2326 cm-1. This shows that it corresponds to the ν(CtN) band of another adsorption mode of acetonitrile. Acetonitrile is a weak base (pKa ) -10),42 but nonetheless, it is sensitive to Lewis acid strength; electron withdrawal from the nitrogen lone pair upon coordination leads to an increase of the ν(CN) frequency.43 Hence, the weak band at 2349 cm-1 likely corresponds to acetonitrile coordinated to boron Lewis acid sites; moreover, its wavenumber is close to that observed in the spectrum of the CH3CN‚ ‚ ‚BF3 complex in CH3CN solution38 (2351 cm-1). The fact that the observed ν(CN) frequency is the same for both CD3CN and CH3CN shows that the Fermi resonance is removed in the case of CH3CN. Such a phenomenon can only occur for a strong shift of the ν(CN) band, confirming an important charge transfer upon the coordination of acetonitrile to the surface of B2O3-SiO2. However, the very weak intensity of the ν(CN) band at 2349 cm-1 is very surprising, taking into account the number of strong Lewis acid sites detected by pyridine. A possible explanation would be a very low value of the molar absorption coefficient of the ν(CN) band of strongly coordinated species. Note that a high sensitivity of the ν(CN) band intensity of acetonitrile to its molecular environment has been reported.43 In the spectrum of the CH3CN‚ ‚ ‚BF3 complex, however, the intensity of the ν(CN) band is strong,38 suggesting that its weak intensity in the case of B2O3-SiO2 is due to the formation of a very small amount of coordinated species. This is in agreement with the very weak perturbation of the band at 1380 cm-1 observed on the BS-5 sample (Figure 4B, spectra j and k) which was strongly diminished upon pyridine adsorption (Figure 2, inset). Accordingly, the coverage of Lewis acid sites by acetonitrile is much lower than that obtained using pyridine as a probe molecule, thus confirming the weak adsorption energy of acetonitrile on boron Lewis acid sites. Figure 5 shows the infrared spectra obtained after CD3CN adsorption on Al2O3. The ν(CN) band at 2261 cm-1 characterizes H-bonded acetonitrile species. Coordinated species lead to the strong ν(CN) band at 2320-2330 cm-1.44 As previously observed for pyridine, outgassing at room temperature leads to

desorption of all the H-bonded species, whereas coordinated species are still present in a large amount (Figure 5, spectrum b). Outgassing at a higher temperature (Figure 5, spectra c and d) leads to further desorption of the coordinated species. It should be noted that, even at room temperature, part of CD3CN transforms at the surface,13,44 as indicated by the broad band at 2180 cm-1. Comparison between the results obtained for Al2O3 and B2O3-SiO2 indicates that coordinated species present a lower ν(CN) frequency on Al2O3 than on B2O3-SiO2 (2320 and 2349 cm-1, respectively), whereas their thermal stability is much higher on the former. In this respect, these results are in agreement with those obtained using pyridine as a probe molecule. As indicated by the ν(CN) frequencies, the coordination of acetonitrile species on B2O3-SiO2 gives rise to a higher charge transfer than that observed on Al2O3, whereas the corresponding adsorption energy is much weaker, as shown by the very low amount of coordinated species and the complete reversibility of adsorption at room temperature. 4.4. Carbon Monoxide Adsorption. Figure 6 shows the IR spectra of CO adsorbed at 100 K on BS-20 and also that of the blank sample. This low temperature was chosen because at room temperature no CO adsorption was found to occur. The IR spectra recorded after the adsorption of CO on BS-20 at 100 K (Figure 6, spectra b-d) show in the ν(CO) range two bands at 2157 and 2140 cm-1, which were also observed in the case of CO adsorbed on pure silica45 and which are assigned to CO interacting with silanol groups by hydrogen bonding and physisorption, respectively. In the ν(OH) range, the ν(OH) band of the silanol groups initially observed at 3749 cm-1 is shifted down to 3655 cm-1 by interaction with CO, in a similar way to that observed on silicalite.46 The adsorption of CO at 100 K does not significantly perturb the boranol band at 3703 cm-1. Moreover, no ν(CO) or perturbed ν(OH) bands different from those due to CO adsorption on pure silica are detected. In no case are ν(CO) bands corresponding to coordinated CO (expected above 2180 cm-1) observed. This result contrasts with Al2O3 for which CO is coordinated to Lewis acid sites and gives rise to bands in the 2200-2235 cm-1 range that are detected (with a low intensity) even at room temperature.22 No such interaction was observed on BS samples, where CO interacts only with silanol groups. The ν(B-O) band at 1380 cm-1 was not affected at all by CO adsorption. Such a behavior has been already observed in the case of borosilicalites,


Travert et al.

Figure 7. Evolution of the total interaction energy (∆Etot, ---), deformation energy of the Lewis acid center (∆Edef(A), ‚‚‚‚‚‚‚), deformation energy of the probe molecule (∆Edef(B), -‚-‚-), and interaction energies of the probe with the Lewis center at their geometry in the complex (∆Eint, - - -) for the six acid-base complexes as a function of the M-L distance (M ) B or Al; L ) C or N).

for which it has been reported that pyridine and NH3 give rise to seemingly strongly coordinated species,39,47 whereas CO does not interact. This has been explained in terms of the conversion of boron from a trigonal planar conformation to a tetrahedral conformation upon interaction with a strong base.47 However, the thermal stability of coordinated pyridine and ammonia has not been investigated for the above samples which precludes assessment of their Lewis acid strengths. 4.5. Summary of the Experimental Results. The main experimental findings of this study may be summarized as follows: (1) As shown by pyridine adsorption, B2O3-SiO2 samples present Lewis acid sites consisting of trigonal planar BO3 units, the amount of which is similar to that of strong Lewis acid sites of Al2O3. (2) Both acetonitrile and pyridine coordinated to B2O3-SiO2 show characteristic frequency shifts suggesting a larger charge transfer to the surface than on Al2O3, but a much lower adsorption energy. This is particularly clear in the case of acetonitrile for which (i) coordination is totally reversible at room temperature and (ii) the coverage of boron Lewis acid sites is very low. (3) No coordination of CO was observed on the B2O3-SiO2 Lewis acid sites, whereas CO is known to coordinate Lewis acid centers on alumina surfaces. Hence, regardless of the probe molecule used to characterize the surface, the thermal stability of coordinated species shows that B2O3-SiO2 presents surface Lewis acid sites that are much weaker than those of Al2O3. However, coordination of pyridine

and acetonitrile to these weaker sites gives rise to stronger charge transfer than that on Al2O3 Lewis acid sites, as indicated by infrared frequency shifts. This behavior apparently contradicts a direct correlation between the surface Lewis acid site strengths and IR frequency shifts of probe molecules, which has been established for a wide range of metal oxides, in particular, when tested with pyridine and acetonitrile.16 To gain more insight on this discrepancy, the following section compares the theoretical results obtained using simple BIII and AlIII site models with the experimental observations. 5. Computational Results and Comparison with Experiment To obtain a better understanding of the phenomena occurring during the adsorption process, the interaction and deformation energies for the various complexes have been computed as a function of the M-L distance (M ) B or Al; L ) C or N), with the other internal coordinates being relaxed. Variations of these energies are reported in Figure 7, and their values at the equilibrium geometries (full relaxation) are reported in Table 1 for the six Lewis acid-base complexes. Table 1 shows that the total interaction energies (∆Etot values) obtained for AlIII complexes compare well with those experimentally obtained on Al2O3, where microcalorimetric studies report adsorption energies for the strongest Lewis acid sites of ∼55 kJ mol-1 for CO 48 and ∼150 kJ mol-1 for pyridine.49,50 This indicates that, despite its small size, the AlIII model used adequately reproduces the adsorptive properties of the strongest



TABLE 1: Energy (kJ mol-1) and Amount of Charge Transfer for the BIII and AlIII Complexes complex

∆Etota

∆Edef(B)b

∆Edef(A)b

∆Eintc

∆qd

CO-BIII CD3CN-BIII Py-BIII CO-AlIII CD3CN-AlIII Py-AlIII

-2.1 -9.0 -83.3 -52.8 -102.6 -150.6

0.0 1.2 3.2 0.4 0.8 2.3

0.4 89.0 126.6 14.0 30.6 40.9

-2.5 -99.3 -213.1 -67.2 -134.0 -193.7

0.05 0.18 0.60 0.24 0.28 0.56

a ∆Etot: total interaction energy. b ∆Edef(B) and ∆Edef(A): deformation energies of the probe and the Lewis acid center, respectively. c ∆Eint: interaction energy of the probe and the Lewis acid center at their geometry in the complex. d ∆q: charge transfer from the probe to the Lewis acid center.

Lewis acid sites of γ-Al2O3. To the best of our knowledge, no adsorption energies have been reported yet for the Lewis acid sites of boron oxide or SiO2-B2O3 mixed oxides. For both Lewis acid sites, BIII and AlIII, the total interaction energy increases in the following order:

CO < CH3CN < C5H5N As could be expected, this ranking is the same as that of the basicity of the three molecules, and it is in agreement with our experimental results; the more basic the probe molecule, the stronger the interaction. On the other hand, for a given probe molecule, the total interaction energy is much higher on AlIII than on BIII, showing that AlIII is a stronger Lewis acid site than BIII. At the equilibrium geometry, the interaction energy between pyridine and BIII is lowered by ∼70 kJ mol-1 with respect to AlIII. The most striking differences are obtained for the weaker bases, CO and CD3CN. Whereas their interaction energy with AlIII is rather high (∼50 and ∼100 kJ mol-1, respectively), they are much lower with BIII (2 and 10 kJ mol-1, respectively (Figure 7 and Table 1). Also, it is striking that for the CO-BIII complex no minimum exists at short C-B distances; CO forms only a van der Waals complex with BIII. These computational results are in agreement with the foregoing experimental observations and confirm (i) the lower thermal stability of coordinated pyridine and acetonitrile species on SiO2-B2O3 as compared to Al2O3 and (ii) the total absence of coordinated CO species on SiO2-B2O3. Examination of the energy terms (Table 1 and Figure 7) shows that, for all complexes, the deformation energy of the probe molecule does not significantly contribute to the total interaction energy. Hence, the latter essentially results from the following two opposite contributions: (i) ∆Eint, the interaction energy of the probe with the Lewis acid center (at their geometry in the complex) which is stabilizing, and (ii) ∆Edef(A), the deformation energy of the Lewis acid center which is destabilizing. Figure 7 shows that, for all complexes, ∆Eint presents a deep minimum at short M-L distances (170-200 pm) that is characteristic of charge-transfer complexes. Moreover, Figure 8 shows that, for complexes at their equilibrium geometries, ∆Eint is roughly correlated with the amount of charge transfer derived from Mulliken populations (Table 1). This indicates that charge transfer from the probe to the Lewis center is likely to be responsible for the ∆Eint values of both the AlIII and BIII complexes at their equilibrium geometry. Examination of Table 1 and Figure 7 also shows that, for a given Lewis acid, ∆Eint becomes more negative when the basic strength of the probe molecule increases. As could be expected, stronger bases lead to a more important charge transfer and increase the stability of the surface complex (Table 1).

Figure 8. Correlation between ∆Eint and ∆q for the six complexes at their equilibrium geometries (solid symbols, BIII complexes; open symbols, AlIII complexes).

TABLE 2: Selected Geometric Parameters (distances in pm, angles in degrees) for the BIII and AlIII Complexesa CO-BIII CD3CN-BIII Py-BIII CO-AlIII CD3CN-AlIII Py-AlIII

r(M‚ ‚ ‚L)b

r(M-O)c

β(LMO)d

309.6 171.3 165.1 218.5 200.2 199.1

138.3 (+0.2) 143.7 (+5.6) 145.3 (+7.2) 171.6 (+2.5) 172.9 (+3.8) 173.7 (+4.6)

87.4 (+0.8) 99.9 (+13.3) 102.8 (+16.2) 93.5 (+7.1) 96.8 (+10.4) 98.3 (+11.7)

a The values in parentheses represent the variation with respect to the isolated Lewis acid center. b Intermolecular distances (M ) B, Al; L ) C, N). c Al-O and B-O distances. d Mean angle of M-O bonds with respect to the C3 axis.

On the other hand, the deformation energies of both Lewis centers monotonically increase when the M-L distance is reduced (Figure 7). Selected geometrical data for the complexes at their equilibrium geometry are reported in Table 2. For both types of Lewis acid centers, the M-L distance decreases when the basic strength increases, whereas the M-O bond length and the L-B-O angle increase. These geometrical changes are consistent with the classical picture of similar charge-transfer adducts (e.g., boron and aluminum halides51-53), changing their configuration from planar MX3 to tetrahedral LMX3 upon complexation. Table 1 and Figure 7 show that, in all cases, deformation energies are much lower for AlIII than for BIII complexes. Hence, the total interaction energies (∆Etot values) of AlIII complexes are close to the ∆Eint value at all M-L distances. Conversely, for BIII complexes, ∆Etot is much higher than ∆Eint, particularly at short M-L distances. Thus, although the amount of charge transfer (from the probe molecule to the Lewis acid center) in Py-BIII and MeCN-BIII is not much different from that in the corresponding AlIII complexes, the stability of the boron adducts is much lower. For the CO-BIII complex, no minimum exists at short C-B distances, although perturbed CO and BIII do interact, as shown by the minimum of ∆Eint (Figure 7). Because of the high deformation energy of BIII, CO does not coordinate on this Lewis acid site but forms only a van der Waals complex. Boron donnor-acceptor complexes have been widely studied, particularly for the case of boron halides.38 The energy involved in the formation of these complexes has been explained53 by similar considerations as those discussed above. Among the factors affecting stability, π-interaction between the formally empty pz orbital of boron and the filled pz orbitals of halogens is usually invoked. Similar interactions could partly explain the differences between BIII and AlIII Lewis acid sites. Figure 9 shows isodensity contours of the HOMO-5 molecular orbitals of BIII and AlIII. Whereas the pz orbitals of oxygen atoms are involved in both molecular orbitals, significant overlap with the


Travert et al.

Figure 9. Isodensity surfaces of the HOMO-5 orbitals of the BIII and AlIII Lewis acid sites. The contour level was made at 0.03 au.

TABLE 3: Comparison of Computed and Experimental Vibration Frequenciesa AlIII

Al2O3b

BIII

B2O3-SiO2

ν(CO)/cm-1 2227 (+82) 2230 (+87) 2144 (+ 1) c ν(CN)/cm-1 2340 (+77) 2320 (+57) 2362 (+99) 2349 (+86) ν8a(Py)/cm-1 1636 (+52) 1625 (+42) 1625 (+ 41) 1628 (+44) a The values in parentheses represent the frequency shifts with respect to the gas phase. b Experimental values corresponding to the strongest Lewis acid sites of Al2O3. c Not observed.

pz orbital of the cation is only observed in BIII which clearly indicates the π-character of B-O bonds, as usually found in boron oxides.28 This is not observed in the case of AlIII for which the Al pz orbital contribution is much less significant and slightly antibonding with respect to Al-O bonds. The π-character of B-O bonds, estimated at ∼20%,55 explains the high deformation energy of the BIII Lewis acid sites. As stated in the Introduction section, IR frequency shifts of coordinated probe molecules are largely due to charge transfer from the probe to the Lewis acid center. Table 3 reports the calculated frequencies for the ν(CO), ν(CN), and ν8a modes of the CO, CD3CN, and Py complexes, respectively. These values are compared with the highest frequencies observed for coordinated species on SiO2-B2O3 and Al2O3 samples. The computed frequencies are higher than those experimentally observed by ∼10-20 cm-1 for most of the complexes. These variations can arise from the simplicity of the computational model used. The magnitude of the shifts, however, is comparable with that found experimentally. In the case of CO, the computed frequency for CO-AlIII (2227 cm-1) is in agreement with the value observed on Al2O3 (∼2220 cm-1). For CO-BIII, the ν(CO) frequency is very close to the gas-phase value, as could be expected for a van der Waals complex. It should be noted that, if this type of adsorption occurred on the B3+ Lewis acid sites, it would not be distinguishable from the contribution of other physisorbed species. The magnitudes of the ν(CN) and ν8a frequency shifts of the MeCN and Py complexes are also comparable with the experimental values. As found experimentally, the ν(CN) frequency is higher on BIII than on AlIII. However, the reverse occurs for the ν8a mode of pyridine, which disagrees with experiment. This discrepancy could be due to the approximations used to compute the force constants (finite differences) and to the small size of our cluster models. In any case, however, the magnitude of the IR frequency shifts for these complexes confirms that an important charge transfer occurs from MeCN and Py to the AlIII and BIII Lewis acid centers. In view of the strong differences in stabilities of the MeCN-AlIII and MeCN-BIII complexes, for example, the computed IR frequency shifts confirm that there is no direct correlation between the adsorption energy of the probe molecule on one hand and its IR frequency shift, or the amount of charge transfer, on the other. This is in agreement with recent theoretical studies of donor-acceptor complexes in the gas phase. Although linear correlations between binding energy and charge transfer have been established for several complexes,54 significant discrep-

ancies may occur for some of them.51-53 Numerous studies have established general relationships between the IR spectra of adsorbed probe molecules and Lewis acid strength,1,16,21 demonstrating that such comparisons are valid for most metal oxides. In this respect, the case of B2O3-SiO2 Lewis acid sites is an exception. According to the calculations shown here, this discrepancy is essentially due to the high energy required to deform the BO3 unit from a nearly planar configuration to a tetrahedral configuration in the adsorption complex. 6. Conclusions This contribution reports on the characterization of the Lewis acidity of B2O3-SiO2 and γ-Al2O3 by infrared spectroscopy using carbon monoxide, acetonitrile, and pyridine as probe molecules. No coordination of carbon monoxide is observed on B2O3-SiO2 even at low temperatures, whereas coordinated CO species are formed on γ-Al2O3. Similarly, coordinated pyridine and acetonitrile show a much lower thermal stability on B2O3-SiO2 than on γ-Al2O3, indicating that B2O3-SiO2 presents much weaker Lewis acid sites than γ-Al2O3. On the other hand, coordinated pyridine and acetonitrile species show infrared frequency shifts that are larger on B2O3-SiO2 than on γ-Al2O3, suggesting that charge transfer from these probe molecules is more important on B3+ than on Al3+ Lewis acid sites. DFT calculations of the interaction of these probe molecules with simple models representing Al3+ and B3+ Lewis acid sites adequately reproduce these experimental observations. The weak strength of the B3+ Lewis acid sites is ascribed to the π-character of B-O bonds, which disfavors the conversion of boron from a trigonal planar conformation to a tetrahedral conformation upon adsorption of probe molecules and decreases the adsorption energy of pyridine and acetonitrile despite a strong charge transfer. For carbon monoxide, its basic strength is not large enough to compensate for the conformational change of the B3+ Lewis acid center. Finally, B2O3-SiO2 samples present the peculiarity to give rise to unstable coordinated species having a strong charge transfer to the surface. More work is needed to establish in which extent this unique behavior could affect the adsorption and activation of reactants in catalytic processes. Acknowledgment. Computational resources were provided by CRIHAN (Saint-Etienne du Rouvray, France), funded under the framework of the “Contrat de Plan Etat/Re´gion (CPER, fiche 15)”. References and Notes (1) (a) Kno¨zinger, H. In Elementary Reaction Steps in Heterogeneous Catalysis; Joyner, R. W., van Santen, R. A., Eds.; Kluwer: Dordrecht, The Netherlands, 1993; p 267. (b) Lercher, J. A.; Grundling, C.; Erder-Mirth, G. Catal. Today 1996, 27, 353. (c) Busca, G. Catal. Today 1998, 41, 191. (d) Rodrı´guez Delgado, M.; Morterra, C.; Cerrato, G.; Magnacca, G.; Otero Areań, C. Langmuir 2002, 18, 10255. (e) Payen, E.; Grimblot, J.; Lavalley, J. C.; Daturi, M.; Mauge´, F. In Handbook of Vibrational Spectroscopy; Chalmers, J. M., Griffiths, R., Eds.; Wiley: 2001. (2) Pimentel, G. C.; McClellan, A. L. The Hydrogen Bond; W. H. Freeman and Co.: San Francisco, CA, 1960. (3) Makarova, M. A.; Zholobenko, V. L.; Al-Ghefailli, K. M.; Thompson, N. E.; Dwyer, J. J. Chem. Soc., Faraday Trans. 1994, 90, 1047. (4) Davydov, V. Ya.; Kiselev, A. V.; Kuznetsov, B. V. Russ. J. Phys. Chem. 1970, 44, 1. (5) Horill, P.; Noller, H. Z. Phys. Chem. 1976, 100, 155. (6) Rouxhet, P. G.; Sempels, R. E. J. Chem. Soc., Faraday Trans. 1 1974, 70, 2021. (7) Cairon, O.; Chevreau, Th.; Lavalley, J.-C. J. Chem. Soc., Faraday Trans. 1998, 94, 3039. (8) Person, W. B. In Spectroscopy and Structure of Molecular Complexes; Yarwood, J., Ed.; Plenum Press: London, 1973.

Lewis Acid Site Strength and IR Spectra of B2O3-SiO2 (9) Otero Areań, C.; Rodrı´guez Delgado, M.; Montouillout, V.; Lavalley, J.-C.; Fernandez, C.; Cuart Pascual, J. J.; Parra, J. B. Microporous Mesoporous Mater. 2004, 67, 259. (10) Kno¨zinger, H.; Ratnasamy, P. Catal. ReV.sSci. Eng. 1978, 17, 31. (11) Kno¨zinger, H.; Huber, S. J. Chem. Soc., Faraday Trans. 1998, 94, 2047. (12) Zecchina, A.; Lamberti, C.; Bordiga, S. Catal. Today 1998, 41, 169. (13) Kno¨zinger, H.; Kietenbrink, H. J. Chem. Soc., Faraday Trans. 1 1975, 71, 2421. (14) Sempels, R. E.; Rouxhet, P. G. J. Colloid Interface Sci. 1976, 55, 263. (15) Scokart, P. O.; Declerk, F. D.; Sempels, R. E.; Rouxhet, P. G. J. Chem. Soc., Faraday Trans. 1 1979, 75, 271. (16) Busca, G. Phys. Chem. Chem. Phys. 1999, 1, 723. (17) Parry, E. P. J. Catal. 1963, 2, 371. (18) Ward, J. W. J. Catal. 1968, 10, 34. (19) Morterra, C.; Chiorino, A.; Ghiotti, G.; Garrone, E. J. Chem. Soc., Faraday Trans. 1 1979, 271, 75. (20) Connel, G.; Dumesic, J. A. J. Catal. 1987, 105, 285. (21) Morterra, C.; Cerrato, G. Langmuir 1990, 6, 1810. (22) Morterra, C.; Magnacca, G. Catal. Today 1996, 27, 497. (23) Yates, J. T.; Ballinger, T. H. Langmuir 1991, 7, 3041. (24) Schmidt, M. W.; Baldridge, K. K.; Boatz, J. A.; Elbert, S. T.; Gordon, M. S.; Jensen, J. H.; Koseki, S.; Matsunaga, N.; Nguyen, K. A.; Su, S. J.; Windus, T. L.; Dupuis, M.; Montgomery, J. A. J. Comput. Chem. 1993, 14, 1347. (25) (a) Becke, A. D. J. Chem. Phys. 1993, 98, 5648. (b) Stephens, P. J.; Devlin, F. J.; Chablowski, C. F.; Frisch, M. J. J. Phys. Chem. 1994, 98, 11623. (c) Hertwig, R. H.; Koch, W. Chem. Phys. Lett. 1997, 268, 345. (26) (a) Ditchfield, R.; Hehre, W. J.; Pople, J. A. J. Chem. Phys. 1971, 54, 724. (b) Dill, J. D.; Pople, J. A. J. Chem. Phys. 1975, 62, 2921. (c) Hehre, W. J.; Ditchfield, R.; Pople, J. A. J. Chem. Phys. 1972, 56, 2257. (d) Francl, M. M.; Pietro, W. J.; Hehre, W. J.; Binkley, J. S.; Gordon, M. S.; DeFrees, D. J.; Pople, J. A. J. Chem. Phys. 1982, 77, 3654. (e) Hariharan, P. C.; Pople, J. A. Theor. Chim. Acta 1973, 28, 213. (27) Digne, M.; Sautet, P.; Raybaud, P.; Euzen, P.; Toulhoat, H. J. Catal. 2002, 211, 1. (28) Takada, A.; Catlow, C. R. A.; Price, G. D.; Hayward, C. L. Phys. Chem. Miner. 1997, 24, 423. (29) Henrich, W. E.; Cox, P. A. The Surface Science of Metal Oxides; Cambridge University Press: Cambridge, U.K., 1994. (30) van Duijneveldt, F. B.; van de Rijdt, J. G. C. M.; van Lenthe, J. H. Chem. ReV. 1994, 94, 1878. (31) Morokuma, K.; Kitaura, K. In Molecular Interactions; Ratajczak, H., Orville-Thomas, W. J., Ed.; Wiley: 1993; p 21.

J. Phys. Chem. B, Vol. 108, No. 42, 2004 16507 (32) Burneau, A.; Gallas, J.-P. In The surface properties of silicas; Legrand, A. P., Ed.; Wiley: 1998. (33) Carteret, C.; Burneau, A. Phys. Chem. Chem. Phys. 2000, 2, 1747. (34) El-Egili, K. Physica B 2003, 325, 340. (35) Kamitsos, E. I.; Karakassides, M. A.; Chryssikos, G. D. J. Phys. Chem. 1987, 91, 1073. (36) Pichat, P.; Mathieu, M.; Imelik, B. Bull. Soc. Chim. Fr. 1969, 8, 2611. (37) Mohammed Saad, A. B.; Ivanov, V. A.; Lavalley, J.-C. Appl. Catal., A 1993, 94, 71. (38) Taillandier, M.; Taillandier, E. Spectrochim. Acta, Part A 1969, 25, 1807. (39) Scarano, D.; Zecchina, A.; Bordiga, S.; Geobaldo, F.; Spoto, G.; Petrini, G.; Leofanti, G.; Padovan, M.; Tozzola, G. J. Chem. Soc., Faraday Trans. 1993, 89, 4123. (40) Cotton, F. A.; Wilkinson, G.; Murillo, C. A.; Bochmann, M.; Grimes, R. N. AdVanced Inorganic Chemistry, 6th ed.; Wiley: 1999. (41) Taillandier, E. Ph.D. Thesis, Paris, 1970. (42) Collumeau, A. Bull. Soc. Chim. Fr. 1968, 12, 5087. (43) Purcell, K. F.; Drago, R. S. J. Am. Chem. Soc. 1966, 88, 919. (44) Escalona Platero, E.; Penãrroya Mentruit, M.; Morterra, C. Langmuir 1999, 15, 5079. (45) Storozhev, P. Yu.; Otero Areań, C.; Garrone, E.; Ugliengo, P.; Ermoshin, V. A.; Tsyganenko, A. A. Chem. Phys. Lett. 2003, 374, 439. (46) Zecchina, A.; Bordiga, S.; Spoto, G.; Marchese, L.; Petrini, G.; Leofanti, G.; Padovan, M. J. Phys. Chem. 1992, 96, 4991. (47) Datka, J.; Cichocki, A.; Piwowarska, Z. Stud. Surf. Sci. Catal. 1991, 65, 681. (48) Della Gatta, G.; Fubini, B.; Ghiotti, G.; Morterra, C. J. Catal. 1976, 43, 90. (49) Paukshtis, E. A.; Solanov, R. I.; Yurchenko, E. N. React. Kinet. Catal. Lett. 1982, 19, 105. (50) Clark, A.; Holm, V. C. F. J. Catal. 1963, 2, 21. (51) Jonas, V.; Frenking, G.; Reetz, M. T. J. Am. Chem. Soc. 1994, 116, 8741. (52) Timoshkin, A. Y.; Suvorov, A. V.; Bettinger, H. F.; Schaefer, H. F. J. Am. Chem. Soc. 1999, 121, 5687. (53) Rowsell, B. D.; Gillespie, R. J.; Heard, G. L. Inorg. Chem. 1999, 38, 4659. (54) Gurjanova, E. N.; Goldstein, I. P.; Romm, I. P. Donor-Acceptor Bond; Wiley: New York, 1975. (55) Pauling, L. The Nature of the Chemical Bond, 3rd ed.; Cornell University Press: Ithaca, NY, 1960.

Evidence for Discrepancy between the Surface Lewis Acid Site

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