J. Phys. Chem. B 2001, 105, 1935-1942
1935
ARTICLES An Examination of Confinement Effects in High-Silica Zeolites L. Yang, K. Trafford, O. Kresnawahjuesa, J. Sˇ epa, and R. J. Gorte* Department of Chemical Engineering, UniVersity of PennsylVania, Philadelphia, PennsylVania 19104
David White Department of Chemistry, UniVersity of PennsylVania, Philadelphia, PennsylVania 19104 ReceiVed: August 15, 2000; In Final Form: NoVember 5, 2000
The environment provided by the zeolite channels has been examined in high-silica zeolites having the MFI (ZSM-5), CHA (Chabazite), MOR (Mordenite), TON (ZSM-22), MTW (ZSM-12), FER (Ferrierite), and FAU (Faujasite) structures. Calorimetric measurements of CH4 and O2 at ∼210 K showed that structure affects the adsorption properties in a manner which depends on the pore dimensions. The differential heats for CH4 at low coverage were 28 kJ/mol in FER, 27 kJ/mol in TON, 25 kJ/mol in MOR, 21 kJ/mol in MFI, 20.5 kJ/mol in MTW, 19.5 kJ/mol in CHA, and 14 kJ/mol in FAU. However, calorimetric data for acetonitrile in the acidic forms of these zeolites gave heats that were independent of structure within experimental error, with differential heats for the 1:1 adsorption complexes all being 100 ( 10 kJ/mol. A possible explanation lies in the additional orientation-dependent interaction resulting from the hydrogen bonding. These topology-sensitive, orientational interactions are observable in the temperature-sensitive, methyl-proton, NMR line shapes, where it is found that the barriers to reorientation of 1:1 adsorption complexes are much higher in small-pore zeolites. The implications of these measurements to “confinement” effects in catalysis are discussed.
Introduction Acid-catalyzed reactions in zeolites are of great industrial importance.1 While zeolite structure is known to have a significant affect on many reactions, questions remain regarding what role structure plays. It is sometimes stated that certain zeolites are more acidic than others and differences in the SiO-Al bond angles could result in the proton affinity changing with structure.2 However, structure can influence the reactivity of sites in other ways as well. For example, in the case of alkane cracking, dispersion forces between the alkane molecules and the siliceous walls of the zeolites are likely the most important interactions for stabilizing adsorption in the cavities.3 Since the proton affinity of alkanes is so low that one would not predict, nor does one see experimentally,4 any preferential adsorption on the Brønsted sites, the isotherm simply describes physisorption of the alkane in the cavity. Furthermore, it is reasonable to assume that the reaction rate should be written in terms of the surface coverage, using a Langmuir isotherm:5-7
r ) kintθ ) kintKP/(1 + KP) ≈ kappP K, and therefore kapp, changes exponentially with the heat of adsorption, so that the apparent activation energy for reaction and the absolute rates can vary with structure. Since the effect of zeolite structure on the heat of adsorption can be large, rates can be dramatically different for two zeolite structures, even if the acid sites are otherwise identical. For example, the difference in the zero-coverage heats of adsorption for propane in FER and FAU is approximately 27 kJ/mol and the difference will be even larger for larger molecules in these two structures.8
Other “confinement” effects can also lead to structuredependent, catalytic properties. For example, the spatial constraints imposed by the solid structure are quite different from that found in solution. These spatial restrictions on the transition state can affect selectivity, as first demonstrated by Csicsery in the reaction of 1-methyl-2-ethylbenzene in H-MOR.9,10 In this case, bimolecular reactions are suppressed in favor of the monomolecular reaction due to the bulky nature of the transition state for the bimolecular reactions. Several relatively recent reviews make it clear that spatial constraints and sorption effects are important for a wide range of other reactions as well.11,12 In some of these, the spatial constraints require a much more detailed understanding of the site geometry than that required for the 1-methyl-2-ethylbenzene reaction. For example, SiCl4modified H-MOR can exhibit significantly increased selectivity toward production of dimethylamine, rather than trimethylamine, in the reaction of ammonia with methanol.13,14 This has been explained as the result from tailoring of the pore size through silation of the pores,13 but treatment with SiCl4 is also used for dealumination and could therefore remove certain acid sites. Other reactions for which site geometry has been used to explain reactivities or selectivities include cracking of branched alkanes,15 Friedel-Crafts alkylations, condensation of carboxylic acids, and halogenation of aromatics.12 We distinguish two situations where “confinement” can affect reactions. In the first, the reactant molecules interact primarily with the siliceous walls of the zeolite, so that “confinement” is unconstrained by the location of the acid sites. For example, it has been shown that differential heats of adsorption for n-hexane
10.1021/jp002964i CCC: $20.00 © 2001 American Chemical Society Published on Web 02/08/2001
1936 J. Phys. Chem. B, Vol. 105, No. 10, 2001 TABLE 1: Properties of Zeolites Used in This Study zeolite pore dimens36 Si/Al2 [H+],a µmol/g MOR MTW FER TON MFI CHA FAU a
8 (2.6 × 5.7) 12 (6.5 × 7.0) 12 (5.9 × 5.5) 8 (4.2 × 5.4) 10 (3.5 × 4.8) 10 (4.4 × 5.5) 10 (5.3 × 5.6) 8 (3.8 × 3.8) 12 (7.4 × 7.4)
sources and/or refs to characterization
20
750
Conteka, CBV20A
140 21
155 850
synthesized45 Zeolyst46
52 60 40 85
400 500 700 60
synthesized46 Chemie Uetikon41 S. Zones19 PQ47
See refs 43 and 44.
adsorption in MFI zeolites are the same, whether or not acid sites are present.4 Since cracking rates for n-hexane should depend on the concentration of n-hexane at acid sites, the confining interactions can be studied in the absence of acid sites. Therefore, we used calorimetric measurements with CH4 and O2 to probe the zeolite cavities, since these two molecules do not interact in a significant manner with residual hydroxyls or alkali metal cations.16 By performing measurements at ∼200 K, it is possible to approach complete pore filling of the cavity volume, so that an accurate picture of the dispersion interactions with the zeolite cavity can be described. The second situation exists when adsorption is localized at the Brønsted-acid sites and the orientation of the adsorption complex is constrained by the hydrogen bond to the zeolite. In this case, studies of physically adsorbed species cannot adequately describe the environment of the reactant molecules. To examine the specific topology surrounding the acid sites, we used both calorimetric and the 1H NMR studies of CH3CN adsorbed in the various zeolites. We have previously demonstrated that acetonitrile forms localized, 1:1 adsorption complexes with Brønsted-acid sites for loadings below one/site and temperatures below ∼370 K.17,18 In our earlier studies of acetonitrile in CHA and MFI,19,20 we have demonstrated that the NMR studies allow one to probe the barriers to molecular reorientations resulting from van der Waals interactions between the molecule and the siliceous walls of the zeolite. The barrier to reorientation depends strongly on the zeolite structure5 and is found to depend on the detailed nature of the cavity, not simply the pore opening, making NMR a powerful method of differentiating zeolite structures. Because NMR probes only that part of the cavity near the acid sites, while the calorimetric measurements probe the entire volume, the two measurements provide complementary information. Experimental Section The samples used in this study came from various sources and are listed in Table 1. Characterization of some of these samples has been described in other publications that are also referenced in the table. All samples were placed in the hydrogen form prior to performing adsorption measurements, and the XRD patterns for each sample agreed with the expected patterns. Simultaneous temperature-programmed-desorption and thermogravimetric-analysis (TPD-TGA) measurements with either n-propylamine or isopropylamine were used to measure the Brønsted-site densities.21,22 The TPD-TGA technique is based on the fact that the amine molecules protonated by the Brønsted sites react to propene and ammonia prior to desorption. The calorimeter used in this study for CH4 and O2 has been described elsewhere.16,23 The calibration constant for this instrument was defined so that the differential heat of adsorption for ethane in MFI at low loadings was 31 kJ/mol, the value
Yang et al. determined from isotherm measurements from a number of groups.8 The agreement between heats measured calorimetrically and values obtained from isotherms was checked for several other adsorbates and found to be excellent in all cases. For purposes of this paper, we define the differential heat of adsorption as the enthalpy change in going from the adsorbedphase to the gas phase. The uncertainty in our measurement of each point was ∼2%, corresponding to ∼0.5 kJ/mol; however, the uncertainty in the absolute values of the differential heats is slightly larger, ∼1 kJ/mol. A separate, high-temperature calorimeter, also described elsewhere,24,25 was used for the measurements with acetonitrile. These measurements were performed at 400 K in order to promote rapid diffusion to the acid sites. Because the high-temperature calorimeter was not designed to accurately measure the quantity of adsorbate remaining in the gas phase, heats could not be measured accurately at coverages above one molecule/site. For the NMR measurements, approximately 100 mg of zeolite was degassed and weighed at 700 K in a Cahn microbalance at 10-6 Torr. The samples were then transferred into specially designed glass sample tubes which were then attached to a vacuum manifold and again degassed at 700 K and 10-6 Torr.26 To obtain the deuterated form of the samples, the acidic form of each zeolite was repeatedly exposed to D2O vapor and evacuated, after which it was gradually heated to 750 K in vacuo until the pressure fell below 10-6 Torr. Using NMR as the analytical tool, it was found that the proton background of the deuterated samples was always less than 3% of the total signal when CH3CN was adsorbed. The degassed samples were exposed to controlled amounts of acetonitrile vapor using a calibrated volume that allowed the adsorbate coverage to be known to within 2%. To avoid bed effects upon adsorption, the zeolites were spread along the length of a long, evacuated, 1/2-in.-diameter tube during exposure to the gaseous adsorbate26,27 and then poured into a smaller tube without exposure to air. The smaller tube was sealed with a torch and inserted into a static NMR probe for spin counting measurements, which in all cases agreed with the measured dosing volumes. The 1H NMR spectra were obtained at a magnetic field of 3.5 T (150 MHz for 1H). All measurements were performed using a home-built spectrometer previously described,28 in conjunction with a Libra pulse programmer and software provided by Tecmag.29 For all experiments, a single-coil, doubleresonance variation of the static probe of Pines, Gibby, and Waugh30 was employed, in which the sample could be immersed in either liquid N2 or heated gases contained in a small dewar surrounding the coil. The pulse cycle for the methyl proton line shapes was the quadrupolar echo sequence, (90°)0-τ1-(90°)90°-τ1-acquire. At low temperatures, the delay time, τ1, was set at 74 µs in order to obtain a line shape characteristic of a free-induction decay. As demonstrated earlier in the study of matrix-isolated acetonitrile,31 the three methyl protons behave as a quasi-quadrupolar nucleus of spin 3/2, but due to dipolar coherences, the line shapes are dependent on the quadrupolar echo delay time.31,32 At room temperature, measurements with values of τ1 corresponding to the averaged dipolar couplings showed little change from those presented here, other than that these line shapes showed some spectral rounding. For all spectra, the repetition rate on signal averaging was greater than 2T1’s at all temperatures to avoid significant saturation effects. Since the quantity of adsorbed molecules in most samples was approximately 35 µmol, the number of scans required to obtain a reasonable
Confinement Effects in High-Silica Zeolites
Figure 1. Isosteric heats measured for CH4 (circles) and O2 (diamonds) on MFI. The open points were measured at 295 K, and the solid points, at 210 K.
J. Phys. Chem. B, Vol. 105, No. 10, 2001 1937
Figure 2. Isosteric heats measured for CH4 (circles) and O2 (diamonds) on MTW. The open points were measured at 295 K, and the solid points, at 210 K.
signal-to-noise was several hundred at 78 K and several thousand in the vicinity of room temperature. Results CH4 and O2 Adsorption. As discussed in the Introduction, the reactant molecules in some reactions interact primarily with the siliceous walls of the zeolite, so that “confinement” is unconstrained by the location of the acid sites and can be studied using adsorption of CH4 and O2, which do not interact with the Brønsted-acid sites. Calorimetric Measurements for CH4 and O2. Isosteric differential heats are shown in Figures 1-7 for adsorption measurements at ∼210 K for both O2 and CH4 on each of the seven zeolites. On most of the zeolites, room-temperature measurements were also performed for CH4, with results shown as open symbols in the figures. Because the results were virtually identical at both temperatures and because it was not possible to achieve high adsorbate loadings at the higher temperatures, we focus primarily on the low-temperature data. Figure 1 shows isosteric heats for CH4 and O2 on siliceous MFI. For both, the heats are essentially constant with coverage and close in value to what was reported in earlier calorimetric studies, 21 kJ/mol for CH48 and 16 kJ/mol for O2.16 The ratio of the heats of adsorption for O2 and CH4 is approximately 0.75, a value which was essentially the same for all the zeolites studied. The common value for this ratio provides further evidence that these molecules interact with the cavities primarily through dispersion forces. The data for O2 and CH4 at 210 K on MTW, shown in Figure 2, are very similar to that for MFI, showing values that are perhaps slightly lower, 15.5 kJ/mol for O2 and 20.5 kJ/mol for CH4. Again, the differential heats are constant with coverage. The value for CH4 also agrees very well with reports from previous measurements taken at room temperature.8 Figure 3 gives the analogous results for the high-silica MOR. While we do not report all the data here, measurements performed on two other high-silica MOR samples, including a sample that was synthesized in high-silica form and not steamed,33 were found to give essentially identical results, so that the conclusions are not sample dependent. The differential heats are not as constant for MOR as for MFI and MTW and appear to fall at a coverage of ∼0.5 mmol/g, close to 0.7 mmol/g which corresponds to one molecule in each of the 8-ring side cavities. Below 0.5 mmol/g, the differential heats are 25.5 kJ/ mol for CH4 and 20.5 kJ/mol for O2. The differential heats fall to 22.5 and 19 kJ/mol, respectively, at higher coverages.
Figure 3. Isosteric heats measured for CH4 (circles) and O2 (diamonds) on MOR. The open points were measured at 295 K, and the solid points, at 210 K.
Figure 4. Isosteric heats measured for CH4 (circles) and O2 (diamonds) on TON. The open points were measured at 295 K, and the solid points, at 210 K.
Data for TON and FER are shown in Figures 4 and 5, and the higher differential heats clearly reflect the smaller channel sizes for these zeolite structures. For CH4 on TON, the isosteric heats start at ∼27.5 kJ/mol and rise with coverage to almost 30 kJ/mol, probably due to attractive interactions between adsorbate molecules.34 The heats for O2 on TON are fairly constant at 22.5 kJ/mol. On FER, the isosteric heats were fairly constant at 29 kJ/mol for CH4 and 23 kJ/mol for O2. Again, for both TON and FER, the ratio of the isosteric heats for O2 and CH4
1938 J. Phys. Chem. B, Vol. 105, No. 10, 2001
Yang et al. TABLE 2: Calculated Differential Heats at Low Coverages
Figure 5. Isosteric heats measured for CH4 (circles) and O2 (diamonds) on FER. The open points were measured at 295 K, and the solid points, at 210 K.
Figure 6. Isosteric heats measured for CH4 (circles) and O2 (diamonds) on CHA. The open points were measured at 295 K, and the solid points, at 210 K.
Figure 7. Isosteric heats measured for CH4 (circles) and O2 (diamonds) on FAU. The open points were measured at 295 K, and the solid points, at 210 K.
is close to 0.75, indicating that van der Waal interactions dominate the adsorption process. Results for the two zeolites that have large cavities which must be accessed by smaller rings, CHA and FAU, are reported in Figures 6 and 7. On CHA, the isosteric heats are constant at ∼19 kJ/mol for CH4 and 14 kJ/mol for O2, while corresponding values on FAU are 14 and 11 kJ/mol, respectively. While higher heats have been reported for O2 on other FAU samples, it should be noted that most FAU samples have a much higher Al content
zeolite
calcd ∆H (kJ/mol)
zeolite
calcd ∆H (kJ/mol)
CHA FER MFI
20.5 26.5 25.5
MOR MTW TON
31.0 23.0 27.5
than the sample we used and that the polarizability of the lattice oxygen in zeolites is affected by the Si/Al ratio.35 The data we show here probably represents the high-silica limit for FAU. Simulation Results for CH4. Zero-coverage heats of adsorption for CH4 on the various structures were calculated using the molecular simulation package, Cerius2, Release 3.5, from Molecular Simulation, Inc (MSI). In all simulations, the zeolites were considered to have rigid frameworks with lattice vectors given in the literature.36 For each structure, a crystalline superlattice was created with a length of at least 2.6 nm, corresponding to a cutoff radius of 1.3 nm. Differential heats were determined using a grand canonical, Monte Carlo simulation at a pressure of 0.1 kPa and 195 K. The simulation consisted of 2-5 million steps in order to allow the system to stabilize. The consistent-valence force field (CVFF), which is basically a Lennard-Jones, 6-12 potential function for the purely siliceous zeolites, was used for the calculations, with parameters supplied by MSI. We also assumed that all lattice oxygen have the same polarizability, independent of structure, even though bond angles and bond lengths can affect this value.35 The calculated heats for each of the zeolites, with the exception of FAU, are shown in Table 2. We did not attempt to calculate heats for the FAU structure because of the difficulties in dealing with the sodalite cages in this structure, and we did not attempt to optimize the potential function we used to fit adsorption data. The parameters used in the potential function were fit to liquid properties, rather than adsorption properties, and represent methane in an all-atom model. Furthermore, the collision diameter of methane in the simulation package is 3.82 Å, which contrasts with the experimental value of 3.70 Å.37 The quality of the force field is important in any molecular simulation, but simulations in large pores are more “forgiving” than in small pores because the large collision diameter can result in an overlap of the probe molecule with the pore walls, giving incorrect results. Also, comparison of zeolite structures having intersecting channels (MFI) with straight, nonintersecting channels (TON, MTW) is problematic in the calculations. However, the trends exhibited in the calculations, which take into account only the geometry of the cavity on the adsorption, help to explain some of the experimental observations very well. The much higher heats for MOR compared to MTW are explained by the fact that the initial CH4 molecules adsorb in the 8-ring side pockets of the structure.38 Indeed, the calculations suggest that the difference should have been even higher than the experimental observations indicate. Also, CHA, which is a small-pore zeolite, has adsorption properties more similar to a large-pore zeolite, primarily because the cavity inside the pore structure is very large. For FER, the cavities which are separated from the main, 10-ring channels by the 8-ring openings must have a similar interaction with the methane as do the main channels, or else the calculated heats should not be similar to the values for TON. Again, this agrees well with the experimental observations. Adsorption of CH3CN. Calorimetric Studies. Figure 8 shows the calorimetric data for acetonitrile in each of the samples except FAU. (The site density in the FAU sample was too low for the results to be reliable.) The measurements were performed at 400 K in order to ensure enough mobility for the adsorbate
Confinement Effects in High-Silica Zeolites
J. Phys. Chem. B, Vol. 105, No. 10, 2001 1939
Figure 8. Differential heats of adsorption for acetonitrile measured at 400 K in the acidic forms of various zeolites: FER (2), TON (×), CHA (O), MTW (b), MFI (/), and MOR (0).
molecules to diffuse to the acid sites on the time scale of the experiment. Coverages are reported as molecules/site, using site densities given in Table 1. We observed a significant increase in the pressure above each sample at coverages above one per site, indicating that the molecules at lower coverages interact with the sites, while molecules in excess of one per site tend to remain in the gas phase at this temperature. Because this instrument was not designed to easily measure the amount remaining in the gas phase, experimental heats at higher coverages are unreliable. Two interesting observations are apparent from the data in Figure 8. First, there is a relatively strong interaction between acetonitrile and the acid sites in each of the zeolite structures, with each of the samples exhibiting a differential heat of 100 ( 10 kJ/mol. In the case of MFI, measurements on a siliceous sample showed a heat of adsorption of ∼60 kJ/mol,39 indicating that the hydrogen-bond strength must be at least 40 kJ/mol on this structure. Second, the differential heats on the various samples are essentially indistinguishable. Because the physical adsorption of acetonitrile in the siliceous forms of these zeolites will be quite different, this observation is quite surprising. There are two possible explanations for the heats of adsorption being the same on each of the samples. First, the hydrogenbond strength could decrease as the cavity dimension decreased, so that the sum of the energies for physical adsorption and for the hydrogen bond remains a constant. A more likely explanation is that a molecule that is bound to the acid site cannot orient itself to maximize its interactions with the cavity walls. For example, a molecule bound at a site in the channel intersection of H-ZSM-5 will not be able to lie flat on the siliceous walls of the 10-ring channels in the manner of a molecule physisorbed in a siliceous ZSM-5. Indeed, this explanation was used to explain why the heats of adsorption for a series of alkylamines in H-ZSM-5, up to n-butylamine, increased in a 1:1 manner with proton affinities of the amines.40,41 Those same studies suggested that, starting with n-butylamine, additional van der Waal interactions between the end of the n-butyl group and the siliceous walls could be observed. An important implication of this is that heats of adsorption cannot be considered to be a simple sum of the specific, hydrogen-bonding interaction and the physisorption energy, as some have suggested.6,7 1H NMR Results for CH CN. Although static parameters, like 3 heats of adsorption, provide a crude measure of the van der
Figure 9. 1H NMR results for CH3CN in the deuterated, acidic forms of TON and MTW as a function of temperature. The adsorbate coverage was less than one per site.
Waal interactions that confine and restrict the adsorption complex, a more meaningful approach for understanding reactions at finite temperatures involves an investigation of the molecular dynamics of the complex. CH3CN is an especially useful probe molecule because the molecular axis is rigid and, to a good approximation, the molecule remains bonded to the site on the NMR time scale at loadings less than one per Brønsted-acid site and temperatures below approximately 370 K.20,32 Therefore, the only motions possible are due to rotations of the CH3 group in a plane perpendicular to the C3 symmetry axis of the molecule and reorientation of the molecular axis with respect to the zeolite framework. We have previously shown that the NMR powder patterns of the localized, hydrogenbonded, adsorption complex are characteristic of free rotation, or tunneling, of the CH3 group above 78 K. Therefore, the 1H NMR studies of CH3CN essentially probe the excitation of librational modes (rocking-type motions of the molecular axis about the hydrogen bond) of the complex. These are very sensitive to zeolite topology.5 Since the librational frequencies are typically several wavenumbers, the experiments provide dynamically average information relating to the amplitudes that describe the statistics of the molecular reorientations. This results in changes in the line shape with increasing temperature which are entirely due to an increasing librational amplitude that tend to motionally narrow the dipolar splitting (the outer shoulders of the spectra) while leaving the (1/2 spin transitions (the center line) essentially unaffected. In Figures 9-11, we show the methyl proton, dipolar spectra for the stoichiometric, 1:1 adsorption complexes for CH3CN adsorbed in the deuterated zeolites as a function of temperature. In Figure 9, the spectra for CH3CN in TON (channel size of 0.44 × 0.55 nm) and MTW (channel size of 0.59 × 0.55 nm) are compared to show the effect of channel size in onedimensional structures. The effect of increasing temperature on the dipolar part of the spectrum is much more pronounced on MTW at a given temperature. At 295 K, the dipolar spectrum for MTW is almost completely motionally narrowed, while the motion is still highly anisotropic in TON at this temperature. This result is consistent with a larger barrier to librational
1940 J. Phys. Chem. B, Vol. 105, No. 10, 2001
Figure 10. 1H NMR results for CH3CN in the deuterated, acidic forms of FER and MOR as a function of temperature. The adsorbate coverage was less than one per site.
Yang et al. cavity, while the end is blocked with MOR. One again observes a larger barrier to reorientation with the smaller channels in FER; however, a comparison of the results for these two materials with TON and MTW in Figure 9 shows that the 8-membered-ring openings have a major effect in restricting the motion of the 1:1 complex. For example, the complex is essentially rigid at 295 K in MOR, while motional narrowing of the spectrum begins already at 190 K in MTW. Figure 11 shows a comparison of 1H NMR results for MFI and CHA. While MFI consists of intersecting 10-memberedring channels, the sequence of spectra most resemble the results for MTW, a larger pore zeolite, probably due to the relatively large volume available at channel intersections in MFI. Indeed, the results would suggest that the acetonitrile molecule is probably located at the channel intersections. The data for CHA show higher barriers for reorientation of the complex compared to that for MFI or MTW; however, the barriers for CHA appear to be smaller than for any of the other structures, even though the openings are 8-membered rings. Obviously, the relatively low barrier associated with CHA must be due to the large size of the cages. The 8-ring openings also do not form channels that might confine the acetonitrile molecule. Clearly, the overall topology, as well as the size of the cavity openings, plays an important role in confining molecules that are bound to the acid sites. Discussion
Figure 11. 1H NMR results for CH3CN in the deuterated, acidic forms of MFI and CHA as a function of temperature. The adsorbate coverage was less than one per site.
reorientation in the smaller channels of TON. To describe these barriers more quantitatively, it is necessary to compare the experiments with simulations, such as were carried out for CHA;19 however, computational limitations for the larger unitcells of TON and MTW make this unfeasible at the present time. In Figure 10, we show a comparison of 1H NMR results for CH3CN in FER and MOR. These two materials have onedimensional channels, 10-membered rings for FER and 12membered rings for MOR, with 8-membered-ring cavities intersecting. With FER, the 8-ring opening leads to a larger
It is clear from the data in this paper that the topology of the zeolite structure can have a significant effect on reactivities. As pointed out by Derouane,6 one should not be quick to ascribe differences in reactivity for two different zeolite structures to differences in acid strengths. There is much more required to understand reaction rates than simply understanding the protontransfer processes at the acid sites. However, on the basis of our results, it also seems clear that “confinement” cannot be defined simply in terms of an effective pore radius for the zeolite structure. The overall topology of the zeolite structure is important for understanding how the reaction occurs. In cases where the reactants and products interact weakly with the acid sites, such as with alkane cracking, it is important to characterize physisorption of the reactants and products in order to understand reactions.6,7 Here, the specific interaction with the acid sites is relatively weak and protonation of the reactant is almost certainly an activated process. In these cases, calorimetric measurements in siliceous materials can provide insights into the nature of these interactions. The effect of zeolite topology can be very large. For CH4, we have shown that heats of adsorption can double in going from the large-cavity zeolite, FAU, to the small-pore zeolite, FER. While this amounts to a difference of only 14 kJ/mol for CH4, one will get a proportional increase for larger alkanes,23 so that the heats of adsorption for a molecule like n-hexane will be most significant. Indeed, MOR is an interesting case to consider. The acidic form of this zeolite has been referred to as a “mild superacid” since it exhibits somewhat higher rates for catalytic cracking of alkanes than H-ZSM-5 or H-Y.42 On the basis of our present results, which show that the 8-ring side channels have a significant effect on the heats of adsorption of alkanes, one should at least consider the possibility that the acid “strength” of this zeolite is the same as the others, with only the confining interactions being different. In this latter picture, the MOR is simply more effective in concentrating reactant molecules in the vicinity of the acid sites. In the case of MOR, it is interesting that we do not observed two types of sites in the calorimetric measurements with CH4,
Confinement Effects in High-Silica Zeolites one corresponding to the 8-ring pockets and one to the 12-ring channels. One should expect some averaging of the energies due to the Boltzmann distribution at finite temperatures;23,25 however, the expected energy difference between the 8-ring and 12-ring channels should be at least 10 kJ/mol on the basis of the differences observed between MTW and TON, large enough so that one should expect to see sequential filling of the most energetic sites and then the weaker sites. Obviously, we do not observe a distribution of sites, even though our equipment was easily able to distinguish sites having different energies in ionexchanged zeolites.16,23 For reactants that interact more strongly with the acid sites, calorimetric measurements in purely siliceous zeolites do not provide the same kinds of insights. One can no longer model the interaction of the reactant and the zeolite as a simple sum of the specific, hydrogen-bonding interaction and a physical interaction. The localization of the adsorbate by the acid site profoundly affects the orientation and the molecular motions of the adsorbed molecule, proving that localization affects the ability of the adsorbate to optimize its interactions with the cavity walls. This is demonstrated in both the NMR and calorimetric results. In the NMR, the barriers to reorientation are strongly dependent on topology. In the calorimetry, the increase in differential heats of adsorption for acetonitrile, expected on the basis of the physisorption results, is not observed. It is worth discussing recent claims that the one-for-one increase in differential heats with proton affinities observed for a series of alkylamines seen for H-ZSM-5 and H-MOR40,41 is actually due to increases in the interactions of the alkyl group with the siliceous walls, with the specific interaction at the acid site remaining constant.6,7 This argument assumes that the alkylgroup interaction can be modeled by the physical adsorption heats, which we have shown here is probably not true when the functional group is strongly bonded to the site. Second, the argument that gas-phase proton affinity is not important for understanding the adsorption complex cannot explain why methylamine (PA ) 896 kJ/mol) has a heat of adsorption 40 kJ/mol higher than ammonia (PA ) 858 kJ/mol) in both H-ZSM-5 and H-MOR or why n-butylamine (PA ) 917 kJ/ mol) has a heat of adsorption that is 75 kJ/mol higher than ammonia.40,41 Physical interactions for the methyl and n-butyl groups cannot increase the adsorption energy by this much, even if these functional groups could optimize their interactions with the siliceous pores. That proton affinities are crucial for understanding the specific interactions with the acid site is also demonstrated by looking at the series pyridine (PA ) 922 kJ/ mol), 2-fluoropyridine (PA ) 886 kJ/mol), and 2-methylpyridine (PA ) 936 kJ/mol), where the measured differential heats in H-MOR are 200, 145, and 235 kJ/mol, respectively.41 Similarly, nonspecific interactions with the cavity walls cannot explain why ethanol (PA ) 796 kJ/mol, ∆Hads ) 130 kJ/mol in H-ZSM5) has a higher heat of adsorption than trifluoroethanol (PA ) 731 kJ/mol, ∆Hads < 90 kJ/mol) or why acetonitrile (PA ) 799 kJ/mol, ∆Hads ) 100 kJ/mol) has a higher heat of adsorption than trichloroacetonitrile (PA ) 760 kJ/mol, ∆Hads < 75 kJ/ mol).39 Nonspecific interactions clearly do play a role in stabilizing the adsorption complex, but proton affinity is certainly crucial for understanding the nature of the specific interactions with the acid site. An intriguing observation from the present study is that characterization of the dynamical properties of the adsorption complex may provide important information on the geometry that confines the adsorption complex. Large differences in the
J. Phys. Chem. B, Vol. 105, No. 10, 2001 1941 barriers to reorientation of the acetonitrile complex are observed for different zeolite structures, and the data clearly suggest that this is due to the size and structure of the cavities in the vicinity of the acid site. Performing these kinds of measurements with various types of probe molecules, such as a series of isotopically labeled methyl pyridines, could well allow one to map the cavity volume surrounding the acid sites. Characterization of the dynamical properties of adsorption complexes has not been explored extensively, and we feel this is a very rich topic for future investigations. Summary The size and structure of zeolite cavities significantly change heats of adsorption of physically adsorbed species, so that calorimetric measurements can provide useful insights into “confinement” effects for molecules that do not interact strongly with the acid sites. However, for molecules which form strong hydrogen bonds that orient the molecule with respect to the acid sites, heats of adsorption are not a simple sum of the specific, hydrogen-bond interaction and the heat of physisorption. Dynamical studies using NMR provide more useful information on the interactions which localize and orient the adsorbate at the Brønsted-acid site. Acknowledgment. The authors gratefully acknowledge financial support from the NSF, Grant CTS97-13023. We are grateful to F. Siperstein for helpful advice. References and Notes (1) Venuto, P. B. Microporous Mater. 1994, 2, 297. (2) Fripiat, J. G.; Berger-Andre, F.; Andre, J. M.; Derouane, E. G. Zeolites 1983, 3, 306 (3) Haag, W. O. Stud. Surf. Sci. Catal. 1994, 84, 1375. (4) Savitz, S.; Myers, A. L.; Gorte, R. J.; White, D. J. Am. Chem. Soc. 1998, 120, 5701. (5) Gorte, R. J.; White, D. Microporous Mesoporous Mater. 2000, 356, 447. (6) Derouane, E. G., J. Mol. Catal., A 1998, 134, 29. (7) Derouane, E. G.; Chang, C. D. Microporous Mesoporous Mater. 2000, 35-6, 425. (8) Savitz, S.; Siperstein, F.; Gorte, R. J.; Myers, A. L. J. Phys. Chem. B 1998, 102, 6865. (9) Csicsery, S. M. J. Catal. 1970, 19, 394. (10) Csicsery, S. M. J. Catal. 1971, 23, 124. (11) Weitkamp, J.; Ernst, S.; Puppe, L. In Catalysis and Zeolites: Fundamentals and Applications; Weitkamp, J., Puppe, L., Eds.; SpringerVerlag: Berlin, 1999; p 327. (12) Espeel, P.; Parton, R.; Toufar, H.; Martens, J.; Ho¨lderich, W.; Jacobs, P. In Catalysis and Zeolites: Fundamentals and Applications; Weitkamp, J., Puppe, L., Eds.; Springer-Verlag: Berlin, 1999; p 377. (13) Segawa, K.; Tachibana, H. J. Catal. 1991, 131, 482. (14) Veefkind, V. A.; Grundling, C.; Lercher, J. A. J. Mol. Catal., A 1998, 134, 111. (15) Haag, W. O.; Lago, R. M.; Weisz, P. B. Faraday Discuss. Chem. Soc. 1982, 72, 317. (16) Savitz, S.; Myers, A. L.; Gorte, R. J. Microporous Mesoporous Mater. 2000, 37, 33. (17) Biaglow, A. I.; Gorte, R. J.; White, D. J. Phys. Chem. 1993, 97, 7135. (18) Sˇepa, J.; Gorte, R. J.; White, David; Kassab, E.; Allavena, M. Chem. Phys. Lett. 1996, 262, 321. (19) Trout, B. L.; Suits, B. H.; Gorte, R. J.; White, D. J. Phys. Chem. B 2000, 104, 11734. (20) Suits, B. H.; Sˇ epa, J.; Gorte, R. J.; White, D. J. Phys. Chem. B 2000, 104, 5124. (21) Gorte, R. J. Catal. Lett. 1999, 62, 1. (22) Farneth, W. E.; Gorte, R. J. Chem. ReV. 1995, 95, 615. (23) Savitz, S.; Myers, A. L.; Gorte, R. J. J. Phys. Chem. B 1999, 103, 3687. (24) Parrillo, D. J.; Lee, C.; Gorte, R. J. Appl. Catal. A 1994, 110, 67. (25) Parrillo, D. J.; Gorte, R. J. Thermochim. Acta 1998, 312, 125. (26) Biaglow, A. I.; Gorte, R. J.; Kokotailo, G. T.; White, D. J. Catal. 1994, 148, 779. (27) Biaglow, A. I. Ph.D. Thesis, University of Pennsylvania, 1993.
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