J. Phys. Chem. B 2001, 105, 3917-3921
3917
Characterizing Zeolite Acidity by Spectroscopic and Catalytic Means: A Comparison† S. Kotrel,‡,§ J. H.Lunsford,§ and H. Kno1 zinger*,‡ Department Chemie, Physikalische Chemie, LMU Mu¨ nchen, Butenandtstrasse 5-13 (Haus E); 81377 Mu¨ nchen, Germany, and Texas A & M UniVersity, Department of Chemistry, P.O. Box 30012, College Station, Texas 77842-3012 ReceiVed: June 16, 2000; In Final Form: October 27, 2000
Adsorption of H2, N2, and CO on four different protonated zeolitessH-ZSM-5, H-β, H-Y, and dealuminated H-Ysat low temperatures was studied by transmission Fourier transform infrared spectroscopy. The introduction of the basic probe molecules caused a red-shift of the IR stretching bands of the zeolitic acidic OH groups. This perturbation, which is commonly interpreted as a hydrogen bonding between the acidic OH group and the adsorbate and often taken as a measure of the acidic strength, was then compared with intrinsic activities for the acid-catalyzed cracking of n-hexane previously published for the same zeolite samples. Catalytic and spectroscopic characterization of the acidity is consistent only within the same class of zeolites, e.g. comparison of differently pretreated faujasites. Spectroscopic and catalytic observations for different types of zeolites do not match perfectly, because additional effects, such as interactions of larger molecules with pore walls and the stabilization of transition states and intermediates, can influence the course of an acid-catalyzed reaction.
1. Introduction Despite extensive industrial and academic interest, the nature and strength of acid sites in protonated zeolites remain largely unresolved. A widely used characterization technique employs the extent of perturbation of the particular acidic hydroxyl group by basic probe molecules. The resulting interaction between probe molecule and hydroxyl group is typically monitored by IR spectroscopy1 or NMR.2 The strength of this interaction is often taken as a measure of the acid strength and, hence, the catalytic activity. For example, Sigl et al.3 demonstrated that the strength of an H-bonding interaction between the acidic group and weak bases, e.g., H2, N2, and CO, correlates very well with the catalytic activity of isomorphously substituted MFI zeolites for the acid-catalyzed disproportionation of ethylbenzene. However, Farcasiu et al.4 demonstrated that a correlation of the hydrogen bond donor ability with the acid strength and catalytic activity of acid catalysts is not cogent unless there is a close structural similarity of the participating acids and bases. The objective of this work was to evaluate the correlation of the hydrogen bond donor ability with the activity for the acid catalyzed n-hexane cracking of four types of zeolites, namely, H-ZSM-5, H-β, and nondealuminated and dealuminated faujasite. Therefore, the interaction of H2, N2, and CO with the acidic OH groups of the particular zeolites was studied by transmission Fourier transform IR spectroscopy (FTIR) and then compared with intrinsic activities for n-hexane cracking of the same zeolite samples that were published previously.5,6 2. Experimental Section 2.1. Materials. H-β (PQ-Corp. VALFOR CP-811Al-25), H-ZSM-5 (PQ-Corp. CBV 5020E), and dealuminated H-Y (Linde Y-84 C1523732) zeolites were transformed into their ammonium forms by ion exchange in order to remove possible †
Part of the special issue “John T. Yates, Jr. Festschrift”. LMU Mu¨nchen. § Texas A&M University. ‡
sodium impurities. Approximately 20 g of the zeolite were exchanged in 500 mL of a 1 M NH4NO3 solution 6 times for 8 h at 358 K. The solid was filtered off between the treatments. 2.2. FTIR Measurements. For infrared measurements, a thin self-supporting wafer (10-15 mg/cm2) of a sample was degassed in the preheating zone of a purpose-made FTIR cell by heating the catalyst under a constant nitrogen flow stepwise to 373, 473, 573, and 673 K for 1 h at each temperature. Once the pretreatment was completed, the sample was evacuated and the cell was cooled with liquid nitrogen. When the temperature of the sample compartment reached 85 K, the sample was lowered from the preheating zone into the sample compartment. For the adsorption experiment, the evacuated sample was exposed stepwise to increasing partial pressures of the particular probe molecule (H2, N2, and CO). The FTIR spectra were recorded in the range 950-5000 cm-1 using a Bruker IFS 66 equipped with a MCT detector. Spectra were collected over 128 scans with a resolution of 1 cm-1. 3. Results 3.1. OH Stretching Bands of the Acidic Hydroxyl Groups. Figure 1 presents the infrared spectra in the O-H stretching region of the four zeolites under investigation and assigns the most prominent bands to the hydroxyl groups of these materials following the widely accepted interpretation of zeolitic OH stretching bands.7-9 The observed bands can be assigned as indicated in the figure to • terminal silanol groups, • internal silanol groups due to structural defects, • OH groups located on extraframework aluminum or partially hydrolyzed framework aluminum species (EFAL), and • acidic bridged OH groups connected to the framework of the zeolite (SiOHAL). The IR spectra of the OH stretching region for H-ZSM-5 and H-β (Figure 1A) possess one band in the range 3600-3620 cm-1 (spectral range d) that can be assigned to acidic bridged
10.1021/jp002161v CCC: $20.00 © 2001 American Chemical Society Published on Web 01/31/2001
3918 J. Phys. Chem. B, Vol. 105, No. 18, 2001
Kotrel et al. TABLE 2: Observed Frequency Shifts ∆ν(OH) of the Acidic OH Groups and of Carbonyl Stretching Bands ∆ν(CO) for H-ZSM-5, H-β, H-Y, and Dealuminated H-Y IR frequency shift, cm-1 ∆ν(OH) H-ZSM-5 H-β H-Y dealuminated H-Y (HF/HF′)
Figure 1. Infrared spectra for H-ZSM-5, H-β (A), H-Y, and dealuminated H-Y (B).
TABLE 1: Band Positions of OH Stretching Bands ν(OH) for H-ZSM-5, H-β, H-Y, and Dealuminated H-Y ν(OH), cm-1 OH type
H-ZSM-5
H-β
SiOHAL EFAL type I term SiOH int SiOH EFAL type II
3618 3668 3747
3612 3672 3748 3739 3780
ν(OH), cm-1 OH type
H-Y
dealuminated H-Y
HF(/HF′) LF(/LF′) SiOH EFAL
3642 3547
3627/3601 3555/3531 3744 3670
hydroxyl groups.10 The IR spectra of protonated faujasites in the region of the OH stretching bands (Figure 1B) are dominated by two main components, the high frequency (HF) and the low frequency (LF) band (spectral ranges c and d, respectively). Both bands are associated with bridging hydroxyl groups and are considered acidic, but they differ in their particular position within the zeolite. While the LF component originates from hydroxyl groups in the β cages and hexagonal prisms, the HF band results from hydroxyl groups in the supercages of the zeolitic structure.11-13 Besides the original HF and LF hydroxyl groups, which are now shifted to somewhat lower wavenumbers, two additional zeolitic hydroxyl groups appear for dealuminated HY at lower frequencies than those of the original hydroxyl bands. These bands were previously assigned to bridging hydroxyl groups interacting with extraframework species14-16 and are denoted HF′ and LF′ in the discussion to follow.9 Table 1 summarizes the observed band positions. 3.2. Adsorption of H2, N2, and CO. As expected from previous reports,10,17,18 the acidic bridging hydroxyl groups of the four investigated zeolite samples interact most strongly with the weakly basic probe molecules. Table 2 summarizes the observed frequency shifts of the Brønsted acid site upon CO
-317 -307 -283 -344/-374
∆ν(CO) +32 +32 +31 +34/37
exposure. The red-shift of the hydroxyl groups as well as the blue-shift of the CO stretching frequency indicates a similar acid strength of H-ZSM-5 and H-β. In contrast to the findings for H-ZSM-5 and H-β, faujasites behave quite differently in the presence of CO depending on the pretreatment and the extent of dealumination. Generally, the LF/LF′ bands are not affected by the presence of CO at low pressures under the applied conditions. This demonstrates the restricted accessibility of the hydroxyl groups in the β cages and hexagonal prisms, even for small probe molecules such as CO.18 It is interesting to note that these hydroxyl groups do interact with the much larger pyridine molecule. The reasons for this apparent inconsistency are the significantly higher base strength of pyridine and the different adsorption temperatures, namely, 80 K for CO and 295 K for pyridine.19 The HF band of nondealuminated faujasite shows the weakest interaction with CO (smallest O-H stretching frequency shift) of all four zeolites. Upon dealumination the acidic character of the Brønsted acid sites drastically changes and the results indicate a very strong acidity. Moreover, the unsymmetric shape of the perturbed Brønsted acid site band with increasing CO pressure suggests the presence of a distribution of sites with different acid strength. Consistent with the frequency shifts induced by CO on the HF and HF′ bands (see Table 2), the HF′ band is perturbed at lower CO pressures than the HF band. This indicates that OH groups characterized by the HF′ band possess a higher acid strength than those characterized by the HF band. Based on the observed interaction with CO, the acid strength can be ranked in the following order: dealuminated H-Y > H-ZSM-5 ≈ H-β > H-Y. When N2 or H2 is adsorbed on the zeolite samples, the IR band of the acidic hydroxyl groups is also shifted to lower frequencies. As expected, the extent of the resulting red-shift corresponds to the proton affinity of the probe molecules. Figure 2, parts A and B, present difference spectra of an H-ZSM-5 and an H-β sample after adsorption of H2, N2, and CO, respectively. All three probe molecules cause a red-shift of the band of the acidic hydroxyl groups but to different extents. For the H-ZSM-5 sample, H2 induces the smallest red-shift of only 49 cm-1, followed by N2 with a red-shift of 121 cm-1. The vertical lines in Figure 2A are drawn through the maxima of the red-shifted bands of H-ZSM-5, but these vertical lines coincide almost perfectly with the maxima of the corresponding bands of the H-β sample. This demonstrates that the observed red-shifts for the H-β sample are almost identical to the observed red-shifts for H-ZSM-5. The results for H-ZSM-5 and H-β are summarized in Table 3. Hence, the spectroscopic characterization does not indicate a fundamental difference in the nature of the acidic hydroxyl groups as demonstrated by the comparable position of the O-H stretching band of the perturbed hydroxyl groups. Figure 3A depicts a linear correlation of the square root of the observed frequency shifts with the proton affinity PA of the probe molecules,20 which shows almost identical lines for H-ZSM-5 and H-β.
Characterizing Zeolite Acidity
J. Phys. Chem. B, Vol. 105, No. 18, 2001 3919
Figure 2. Difference spectra of H-ZSM-5 (A) and H-β (B) upon exposure of H2 (A, 5 kPa; B, 10.5 kPa), N2 (A, 0.006 kPa; B, 0.015 kPa), and CO (A, 0.003 kPa; B, 0.007 kPa).
Figure 4. Difference spectra of H-Y (A) and dealuminated H-Y (B) upon exposure of H2 (A, 18.5 kPa; B, 16.0 kPa), N2 (A, 0.09 kPa; B, 0.07 kPa), and CO (A, 0.02 kPa; B, 0.02 kPa).
TABLE 3: Comparison of the OH Frquency Shifts ∆ν(OH) Induced by H2, N2, and CO
As with the observation for H-ZSM-5 and H-β, the HF(OH) band is red-shifted to lower frequencies upon adsorption of the probe molecules. Again, the extent of the red-shift correlates with the proton affinity of the probe molecules with H2 displaying the smallest effect and CO the largest. Interestingly, the structural similarity of the two faujasite samples does not result in a similar strength of the interaction between the OH groups and the probe molecules. The vertical lines in Figure 4 indicate a clear offset toward lower frequencies of the perturbed HF(OH) band for the dealuminated sample as compared to the perturbed bands of the nondealuminated faujasite sample. A drastically increased red-shift of the HF(OH) band of dealuminated faujasite is clearly visible for all probe molecules. The enhanced acidic character of dealuminated HY is also confirmed by the correlation plot of the square root of the ∆νOH shift against the proton affinity of the probe molecules (Figure 3B). The results are also summarized in Table 3.
IR frequency shift, cm-1 ∆ν(OH) catalyst
ν(OH)init
CO
N2
H2
H-ZSM-5 H-β H-Y dealuminated H-Y
3618 3612 3641 3627
-317 -307 -283 -344
-121 -119 -110 -139
-49 -49 -45 -56
4. Discussion
Figure 3. Correlation of proton affinity PA of the probe molecules and frequency shift for (A) H-X-ZSM-5 (o), H-β (2), (B) H-Y (2), and dealuminated H-Y (o).
In contrast to the similarity of H-ZSM-5 and H-β, the results obtained for HY do not match those of the two previous materials, and a comparison of the nondealuminated and the dealuminated faujasite sample (Figure 3B) indicates a considerable variation in the acidic character of these two materials.
Kotrel et al.5,6 published catalytic activities for n-hexane cracking over the same catalyst samples. In this study, the activities for each individual catalyst were corrected for the amount of n-hexane adsorbed in the zeolitic pores and the concentration of catalytically active sites. These data were believed to reflect the true acidic strength of the samples. Therefore, one is justified in comparing the results derived from catalytic tests with the spectroscopic results that are presented in this study. Table 4 summarizes the intrinsic unimolecular cracking activities of the zeolites under investigation. The uniquely high intrinsic activity for H-ZSM-5 compared to H-β and dealuminated H-Y could indicate that H-ZSM-5 is the strongest proton donor, but a comparison with the spectroscopic results makes this conclusion questionable. The interaction with weak bases is not exceptionally strong for H-ZSM-5 and was found to be comparable with H-β. Catalytic cracking is far more susceptible to structural properties of the catalyst material than CO adsorption at low temperatures. Besides, in addition to the
3920 J. Phys. Chem. B, Vol. 105, No. 18, 2001 TABLE 4: Intrinsic Rate Constants for H-ZSM-5, H-β, H-Y, and Dealuminated H-Y Zeolites at 623 K and pHex ) 5.3 kPaa catalysts H-ZSM-5 H-β dealuminated H-Y H-Y
active site/FAL 0.75 0.8 0.15
kint, s-1 × 100 based on FAL
based on active site
111 148 51 64 6 40 no measurable activity
a Based on either the total amount of framework aluminum or the fraction of framework aluminum which is associated with strongly acidic sites.6
interaction between hydroxyl groups and n-hexane molecules, the zeolite walls can also interact with the reactant molecule. Several studies indicated that the adsorption energy greatly depends on the interaction between the oxygen framework (functioning as a Lewis base) and the n-hexane molecule.21-25 The intrinsic activities given in Table 4 account for this effect by correcting the apparent catalytic activities for the adsorptive influence on the n-hexane molecule, but it is difficult to evaluate the structural effect on transition states and intermediates of catalytic cracking. Additional structural effects might assist the process of n-hexane cracking over zeolites with pores comparable to the dimensions of transition states and intermediates. In its role as a host, the microenvironment in the channels of zeolites is in this sense comparable to that of solvents, because the guest molecules in zeolites are enveloped by the small pores.26 Zeolites may be considered as solid analogues of solvents.27 The higher the polarity of a zeolite is, the more likely neutral guest molecules are to become polarized or even ionized upon occlusion into the zeolite.28,29 Therefore, zeolites that are more polar should display a higher catalytic activity for hydrocarbon conversions, which require the separation of charges at some point along the reaction coordinate.30 This effect should generally be more pronounced for structures with smaller pore diameters, e.g., H-ZSM-5. The enhanced acidity of dealuminated faujasites compared to nondealuminated faujasites is observed by spectroscopic characterization and catalytic experiments. For discussing the acidity of faujasites, the H-β sample may be taken as a reference catalyst, since both β and faujasite are large pore zeolites. The nondealuminated faujasite sample was found to be less acidic than H-β according to the interaction with probe molecules and catalytic performance. Upon dealumination, spectroscopy and catalytic results indicated a dramatic increase in acid strength relative to the nondealuminated faujasite, but the comparison with H-β is more complex. Based on the interaction with probe molecules, the acid strength of the dealuminated faujasite sample seems to exceed the acidity of H-β, whereas the catalytic results point to a similar intrinsic activity for H-β and dealuminated H-Y. Clearly, the ability of interacting with small weak bases at liquid nitrogen temperature does not quite probe the same property as a catalytic test. The infrared results characterize the ability of the material to form a hydrogen bonded complex between the acid and a standard base. Because the hydrogen bond formation is a preliminary step of proton transfer, it is frequently supposed that the strength of hydrogen bonding is a measure of the acid strength.17,31,32 Farcasiu et al.4 argued that the hydrogen bond strength correlates with the acid strength only when the acids and bases involved in the comparison consist of very close structural relatives. The increased activity of dealuminated faujasites is partially due to a synergistic effect between the Brønsted site and some form of extraframework aluminum.14,16,33,34 This implies that the
Kotrel et al. structural origin of strong acidity in dealuminated faujasites is somewhat different from the structural origin of Brønsted acidity in the H-ZSM-5 and H-β samples. A comparison of acid strength exclusively on the basis of the ability to form hydrogen bonds might therefore not be justified. Assuming that the red-shift of the hydroxyl groups does provide a direct comparison of the proton donor ability, catalytic and spectroscopic observations would be consistent only if the true number of active sites for dealuminated faujasites were much smaller than the frameworkto-aluminum ratio of Table 4 suggests. This would make the intrinsic activity of dealuminated faujasites considerably larger than the intrinsic activity for H-β. Recently, Kung et al.35 have attempted to explain the enhancement in hydrocarbon cracking activity of a steamdealuminated H-Y zeolite by the increased rate of pore diffusion that results from the creation of voids within the structure. They argue that a change in acidity is not a primary factor, and they correctly point out the ambiguity of acidity measurements based on heats of adsorption of strongly basic probe molecules. In contrast to their views concerning the importance of pore diffusion, the results reported here clearly show that the large differences in activity for the H-Y zeolites pretreated differently may be related to differences in acidity, as determined by the interaction between the zeolitic protons and weak bases. Parenthetically, if diffusion limitations were responsible for the exceptionally small activity of the unsteamed H-Y zeolite, it is difficult to understand how the activity of the H-ZSM-5 zeolite, with its relatively small diameter pores, could be so large (Table 4). 5. Conclusions In conclusion, spectroscopic and catalytic observations do not match perfectly for the four zeolite materials that were compared. The discrepancies between spectroscopic and catalytic results draw attention to additional factors that would have been missed had the zeolitic acidity be monitored by only one of the two techniques. Comparing frequency shifts for hydroxyl groups upon CO adsorption and intrinsic cracking activity for H-ZSM-5 and H-β shows that there must be an additional effect influencing the course of n-hexane cracking over these two types of zeolites. Since the pore sizes of H-ZSM-5 and H-β are the most prominent differences between these two types of zeolites, the structural aspect is presumably responsible for the observed discrepancy between these two types of catalysts. The most valid comparison is between the two faujasite materials because structural factors do not play a role. Indeed, the much greater catalytic activity of the dealuminated H-Y zeolite compared with the normal H-Y zeolite is consistent with the stronger hydrogen bonding. The agreement does not extend, however, to the H-β zeolite, for which the hydrogen bonding is less than with the dealuminated zeolite, but the intrinsic catalytic activity is comparable. Since the two types of zeolites are structurally similar (i.e., they both have relatively large cavities), other secondary factors must be responsible for the less than expected intrinsic activity of the dealuminated H-Y zeolite. Acknowledgment. This work was financially supported by the Deutsche Forschungsgemeinschaft (SFB 338) and the Fonds der Chemischen Industrie. References and Notes (1) Kno¨zinger, H. In Handbook of Heterogeneous Catalysis; Ertl, G., Kno¨zinger, H., Weitkamp, J., Eds.; VCH-Verlag: Weinheim, 1997; Vol. 2, p 707.
Characterizing Zeolite Acidity (2) Pfeifer, H. In Handbook of Heterogeneous Catalysis; Ertl, G., Kno¨zinger, H., Weitkamp, J., Eds.; VCH-Verlag: Weinheim, 1997; Vol. 2, p 732. (3) Sigl, M.; Ernst, S.; Weitkamp, J.; Kno¨ezinger, H. Catal. Lett. 1997, 45, 27. (4) Farcasiu, D.; Hancu, D. Catal. Lett. 1998, 53, 3. (5) Kotrel, S.; Rosynek, M. P.; Lunsford, J. H. J. Phys. Chem. B. 1999, 103, 818. (6) Kotrel, S.; Rosynek, M. P.; Lunsford, J. H. J. Catal. 1999, 182, 278. (7) Sauer, J. J. Mol. Catal. 1998, 54, 312. (8) Hegde, S. G.; Kumar, R.; Bhat, R. N.; Ratnasamy, P. Zeolites 1989, 9, 231. (9) Cairon, O.; Chavreau, Th.; Lavalley, J. C. J. Chem. Soc., Faraday Trans. 1998, 94, 3039. (10) Corma, A.; Fornes, V.; Melo, F.; Perez-Pariente, J. ACS Symp. Ser. 1988, 375, 49. (11) Uytterhoeven, J. B.; Christner, L. G.; Hall, W. K. J. Phys. Chem. 1965, 69, 2117. (12) White, J. L.; Jelli, A. N.; Andre, J. J.; Fripiat, J. J. Trans. Faraday Soc. 1967, 63, 461. (13) Jacobs P. A.; Uytterhoeven, J. B. J. Chem. Soc., Faraday Trans. 1972, 69, 359. (14) Carvajal, R.; Chu, P. J.; Lunsford, J. H. J. Catal. 1990, 25, 123. (15) Makarova M. A.; Dwyer, J. J. Phys. Chem. 1993, 93, 6337. (16) Khabtou, S.; Chevreau T.; Lavalley, J. C. Micropor. Mater. 1994, 3, 133. (17) Zecchina, A.; Bordiga, S.; Spoto, G.; Scarano, D.; Petrini, G.; Leofanti, G.; Padovan, M.; Arean, C. O. J. Chem. Soc., Faraday Trans. 1992, 88, 2959.
J. Phys. Chem. B, Vol. 105, No. 18, 2001 3921 (18) Kno¨ezinger H.; Huber, S. J. Chem. Soc., Faraday Trans. 1998f, 94, 2047. (19) Daniell, W.; Topsøe, N. Y.; Kno¨zinger, H., submitted. (20) Pimentel, G. C.; McClellan, A. L. The Hydrogen Bond; W. H. Freeman: London, 1960. (21) Hopkins, P. D. J. Catal. 1973, 29, 112. (22) Weber, G.; Simonot-Grange, M. H. Zeolites 1994, 14, 433. (23) Derouane, E. G.; Nagy, J. B.; Fernandez, C.; Gabelica, Z.; Laurent, E.; Maljean, P. Appl. Catal. 1988, 40, L1. (24) Eder, F.; Lercher, J. A. Zeolites 1997, 18, 75. (25) Eder, F.; Stockenhuber, M.; Lercher, J. A. Stud. Surf. Sci. Catal. 1995, 97, 495. (26) Gates, B. C. Catalytic Chemistry; John Wiley & Sons: New York, 1992; 274. (27) Rabo, J. in: Zeolite Science and Technology; Ribeiro, F. A., Rodrigues, A. E., Rollmann, L. D., Nacchache, C., Martinus, Nijhoff, Eds.; The Hague: 1984; 291. (28) Dutta P. K.; Turbeville, W. J. Phys. Chem. 1991, 95, 4087. (29) Handreck, G. P. H.; Smith, T. D. J. Chem. Soc., Faraday Trans. 1 1988, 84, 1847. (30) van Santen R. A.; Kramer, G. J. Chem. ReV. 1995, 95, 637. (31) Farneth W. E.; Gorte, R. J. Chem. ReV. 1995, 95, 615. (32) Zecchina,; A. Geobaldo, F.; Spoto, G.; Bordiga, S.; Ricchiardi, G.; Buzzoni, R.; Petrini, G. J. Phys. Chem. 1996, 100, 16584. (33) Lo´nyi, F.; Lunsford, J. H. J. Catal. 1992, 136, 566. (34) Lunsford, J. H ACS Symp. Ser. 1991, 452, 1. (35) Kung, H. H.; Williams, B. A.; Babitz, S. M.; Miller, J. T.; Haag, W. O.; Snurr, R. Q. Top. Catal. 2000, 10, 59.
