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Intermolecular Forces in Zeolite Adsorption and Catalysis
C H A P T E R
9
Intermolecular Forces in Zeolite Adsorption and Catalysis Raul F. Lobo Contents 239 240 242 247 252 254 259 259 260
1. Introduction 2. Basic Thermodynamic Concepts 3. Dispersion (London) Forces 4. Induced Dipole and Quadrupole–Charge Interactions 5. Dipole–Charge Interactions 6. Interactions of Adsorbed Molecules with Acid Sites 7. Hydrophobic Interactions in High-Silica Zeolites 8. Final Remarks References
Abstract The concepts of classical intermolecular forces are used to understand isosteric heats of adsorption of small molecules adsorbed within zeolite pores. This is used as an organizing principle to comprehend the complex and multifaceted set of chemical properties of zeolites in general. We correlate molecular properties to the magnitude of molecule– zeolite forces and these to measurable quantities such as heats of adsorption. Several key thermodynamic concepts are reviewed and then the most relevant types of intermolecular forces as applied to microporous materials are described. The interaction of small molecules with zeolite Brnsted acid sites is then described as a chemical interaction, rather than in terms of physical forces. We conclude with a discussion of hydrophobicity as applied to siliceous zeolites. Keywords:
Adsorption, Intermolecular forces, Calorimetry, Catalysis, Zeolites
1. Introduction As molecules travel down the pore structure of a zeolite, they encounter cations and other chemical groups that expose the molecules to important classes of forces. These forces, and more frequently, the difference between forces for different molecules, ultimately lead to the useful properties of zeolites as adsorbents Ordered Porous Solids DOI: 10.1016/B978-0-444-53189-6.00009-3
#
2009 Elsevier B.V. All rights reserved.
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and catalysts, and should be understood as a foundation to the analysis of material properties in molecular terms. In this chapter, we provide the basis for this understanding with emphasis on the origin and magnitude of the various forces relevant to zeolites. We intend this approach to be an organizing principle to comprehend the complex and multifaceted set of chemical properties of zeolites in general. We explicitly attempt to correlate molecular properties to the magnitude of the forces and these to measurable quantities such as heats of adsorption. To keep the length of this chapter reasonable, it is assumed throughout that we are working at the dilute limit, and that molecule–molecule forces are insignificant. We start with some basic thermodynamic concepts, and then proceed to describe the most relevant types of intermolecular forces as applied to microporous materials and finish with a discussion of hydrophobicity as applied to siliceous zeolites.
2. Basic Thermodynamic Conceptsa The pair potential w(r)—in vacuum—or potential of mean force—in a medium— describes the interaction potential between two molecules.1 The force between the molecules is given by F ¼ dw(r)/dr and shows that the work w(r) that can be done by the force F is also a measure of the available energy or free energy of this system of two molecules. The medium in the case of microporous materials is the micropore and is different from more classical media such as a liquid solvent. First, in a liquid, the solute molecules can perturb the local structure of the solvent; in a zeolite the framework atoms are essentially fixed in space and barely move in response to the presence of an adsorbate in the pores. Second, when a molecule is introduced into a liquid, work must be done to create a cavity to accommodate the guest molecule; in zeolites, the cavity (the pores) is already there and no energy is needed to form it. A zeolite is, however, also like a liquid in the sense that for two solute molecules their pair potential includes molecule–molecule interactions, and also solute–solvent interactions (or equivalently molecule–molecule and molecule– zeolite interactions). A zeolite, just like a solvent, can also change the properties of the adsorbed molecule (dipole moment and polarizability, geometry) with respect to the molecular values in vacuum. A molecule in medium has a cohesive energy mi in a medium i given by the sum of its interactions with the surroundings. Inside a siliceous zeolite pore, these interactions are mainly the interaction between the molecule and Si–O–Si groups that surround the pore. It is possible to relate the cohesive energy to the pair potential w(r). For instance, consider the pair potential given by a power law such as
wðrÞ ¼ C=r n wðrÞ ¼ 1
a
for r > s for r s
The description of intermolecular forces used here follows closely the treatment used by Israelachvilli 1.
ð1Þ
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Figure 9.1 Illustration of the pore size of a 10-ring zeolite (in this case zeolite ZSM-5).The effec˚ and the oxygen-to-oxygen distance is 8.1 A˚. tive pore opening is 5.5A
where s is the hard sphere diameter of the molecule and r is the distance between the centre of the molecule and another atom (or molecule). We can calculate mi by summing all the pair potentials over all space, and then make use of the fact that zeolites are crystalline and atoms are located at specific positions in space. Alternatively, we can estimate this value by considering only the first row of zeolite oxygen atoms that surround the molecule and placing the molecule in the centre of the pore. For a 10-ring zeolite (see Fig. 9.1), the pore opening is 5.5 A˚ (see the effective pore dimensions of the MFIb framework type in the Atlas of zeolite structure-types2 or the web3) and the distance between oxygen atoms at opposite ˚ (the effective pore diameter ends of the pore is 2rpore 5.5þ2 1.3 ¼ 8.1A ˚ ). The net energy reported in the Atlas plus twice the ionic oxygen radius of 1.3 A change by taking a molecule from the gas phase to the centre of a zeolite 10-ring pore is given then by
mizeo 10 wðrpore Þ
rpore ¼ 4:05A˚
ð2Þ
where w(r) describes, in particular, the pair potential between zeolite oxygen atoms and the adsorbed molecule. This value is a lower bound of the cohesive energy because the off-axis position may be more favourable than the centre, and because contributions from other zeolite oxygen and silicon atoms further away b
We will use the three-letter-code described in the Atlas to designate unambiguously the topology of the framework of the various zeolite materials that will be discussed in this chapter.
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have not been included. This function already tells us, for instance, that the cohesive energy of a molecule in a zeolite decreases as the effective pore size increases, since w(r) is a monotonically decreasing function of r at the relevant molecular distances. Consequently, the molar cohesive energy U of an adsorbed molecule can be calculated simply by multiplying by Avogadro’s number U ¼ NA mizeo . The cohesive energy of a molecule can have different values in different parts of a system (say, in the zeolite pores and in the gas phase). Then, at equilibrium, the concentrations of the molecules in the two regions is given by the Boltzmann distribution
X1 ¼ X2 exp mi1 mi2 =kB T
ð3Þ
where X1 and X2 are dimensionless concentrations (usually mole fractions or volume fractions) and kB is the Boltzmann constant. One can expect differences in the concentration of species i if the differences in chemical potential are substantially larger than the thermal energy (at 300 K kBT is 2.5 kJ/mol and at 600 K is 5 kJ/ mol). This equation can be restated as
mi1 þ kB T ln X1 ¼ mi2 þ kB T ln X2
ð4Þ
Here, we assume the phases mix ideally and the activity coefficients can then be assumed to be one. Also, this equation is valid only if the condition of equilibrium between the two states is fulfilled. This expression can be extended to a system with many states by
min þ kB T ln Xn ¼ m ¼ constant for all states n ¼ 1; 2; 3 . . . In such a system, molecules will flow between different states until these n equations are satisfied. The term m is also known as the chemical potential, the total free energy per molecule including its interactions with its medium and its thermal energy. The term kB ln Xn is the entropic contribution to the total chemical potential (again assuming ideal mixing).
3. Dispersion (London) Forces Dispersion forces are one of the three components to the total van der Waals force (in addition to induced dipole and orientation polarization forces).1 They are very important because they are always present, independently of the molecular and geometric properties of the molecules. Dispersion forces are, in general, complex to calculate precisely because their magnitude can be affected by the presence of other molecules, that is, it is a non-additive interaction. The origin of this force can be thought of to be a quantum-mechanical induced polarizability force, and it is described for two different atoms 1 and 2 by London’s expression
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wðrÞ ¼
3 a0;1 a0;2 I1 I2 2 6 2 ð4pe0 Þ r ðI1 þ I2 Þ
ð5Þ
where 0,i is the electronic polarizability of molecule i, e0 is the permittivity of free space (8.8541012 C2 J1 m1) and I is the ionization potential (in J). Table 9.1 summarizes the values of the electronic polarizability for small molecules and a few molecular groups, and Table 9.2 lists the ionization potentials of a selected group of molecules. Polarizability values span roughly two orders of magnitude for this group of molecules, and thus this force can vary greatly, depending on the identity of the molecules in question. Note that water molecules have a rather low polarizability similar to the polarizability of argon. The energetics of adsorption of non-polar molecules in high-silica zeolites is dominated by dispersion forces, and thus these forces dictate the value of cohesive energy gained by a molecule as it enters a zeolite from the gas phase. The isosteric heat of adsorption (Hst) in the limit of zero coverage, a quantity that can be obtained experimentally, is a good estimate of this change in cohesive energy. Fortunately, Savitz et al.4 have conducted a comprehensive study of the adsorption of hydrocarbons and other simple molecules on siliceous zeolites. These data will be used to further elucidate the role of dispersion forces in zeolite–adsorbate interactions. Table 9.3 shows the zeolites used in this study and their relevant characteristics. Three unidimensional zeolites of different pore sizes ZSM-22 ([SiO2]-TON, 10-ring), ZSM-12 ([SiO2]-MTW, 12-ring) and UTD-1 ([SiO2]-DON, 14-ring) were used in addition to multidimensional 10-ring ([SiO2]-MFI) and 12-ring Table 9.1 Electronic polarizabilitiesa 0 of atoms, molecules and molecular groupsb Atoms and molecules
He H2 H2O O2 N2 Ar
0.208 0.787 1.501 1.562 1.710 1.664
CO NH3 HCl CO2 CH4 C2H6 C3H8
1.953 2.103 2.515 2.507 2.448 4.226 5.921
C2H4 CH3OH Cl2 CHCl3 C6H6 CCl4
1.28 1.13 2.03
CH2 Si–O–Si Si–OH
1.84 1.39c 1.60
4.188 3.081 4.4 8.2 10.453 10.002
Molecular groups
C–O–H C–O–C C–NH2 a b c
˚ 3¼(4pe0) 1030 m3¼1.111040 C2 m2 J1. Polarizabilities given in units of (4pe0)A Data from NIST Chemistry Webook8 if available (experimental values). Otherwise from compilation in Israelachvilli. This value is probably an underestimation. See discussion below.
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Table 9.2 Ionization potentiala of selected atoms and molecules (in eV)b Atoms and molecules
He H2 H2O O2 N2 Ar a b
24.59 15.42 12.62 12.07 15.58 15.56
CO NH3 HCl CO2 CH4 C2H6
14.01 10.07 12.74 13.78 12.61 11.52
C2H4 CH3OH Cl2 CHCl3 C6H6 CCl4
10.51 10.84 11.48 11.37 9.24 11.47
Data from NIST Chemistry Webook.8 1 J ¼ 6.242 1018 eV.
Table 9.3 Properties of high-silica zeolites used for adsorption measurements of hydrocarbons in Savitz et al.4a
a b
Zeolite framework type
Vpore (cm3/g)b
Si/Al ratio
Pore type
Pore size (n-rings)
Pore/Cavity ˚) dimensions (A
TON MTW DON FER
0.069 0.097 0.106 0.105
52 140 1 22
1-D 1-D 1-D 2-D
10 12 14 8, 10
MFI
0.173
300
2-D
10, 10
FAU
0.245
1
3-D spherical
5.54.4 6.25.5 107.5 4.83.5, 5.44.2 5.35.6, 5.55.1 12.6
Reproduced with permission from Savitz et al.4 Vpore ¼ Pore volume.
([SiO2]-FAU) materials. Ferrierite (FER) has a two-dimensional pore system with 10- and 8-ring pores running perpendicular to each other. The isosteric heats of adsorption for methane, ethane and propane in these zeolites (Table 9.4) follow the trends expected from the pore size of the zeolites and the polarizability of the molecules. For all molecules, the heat of adsorption decreases as the size of the pore or cavity increases. For a given structure, the heat of adsorption increases as the size of the molecule increases. In fact, the increments in heat of adsorption follow quite closely the changes in electronic polarizability (a0) of the adsorbate. For example, for silicalite-1 ([SiO2]-MFI) the heats of adsorption increase in the ratios 1:1.5:2, while the polarizabilities increase in the ratios 1:1.7:2.4, a consequence of London’s formula. Also note that the cohesive energy of the solid hydrocarbons (sum of the latent heats of melting and vapourization, DHmþDHv) is substantially lower than the isosteric heat of adsorption for all zeolites, thus highlighting the strength of the van der Waals forces between the zeolite and the hydrocarbon.
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Table 9.4 Zero-coverage isosteric heats of adsorption of small hydrocarbons in high-silica zeolites at 25 Ca
a
Hst (kJ/mol)
Zeolite framework type
CH4
C2H 6
C3H 8
FER TON MFI MTW UTD FAU DHmþDHv
27.7 27.2 20.9 20.9 14.2 11.0 9.1
41.7 39.0 31.1 29.5 22.2 19.7 17.5
53.3 48.8 41.4 37.6 28.1 26.2 18.6
CH4 Hst (kJ/mol) from theory
22.4 19.7 14.2 –
4
Reproduced with permission from Savitz et al.
The cohesive energy is closely related to the change in internal energy per mole upon adsorption, a quantity that is called differential heat of adsorption (Ud). In turn, the differential and isosteric heats of adsorption are related by
Hst ¼ Ud þ ZRT
ð6Þ
where Z ¼ PV/RT is the compressibility factor in the bulk gas phase (for an ideal gas Z ¼ 1). On the basis of this information, for instance, we can attempt to estimate the isosteric heat of adsorption of methane in zeolite ZSM-22 (TON) for methane. CH4 ¼ 1, Si–O–Si ¼ 2, a0,1 / (4pe0) ¼ 2.448 1030 m, I1 ¼ 12.61 eV ¼ 2.021018J, a0,2 / (4pe0) ¼ 1.39 1030 m, I2 2 1018J (we use a typical value), rpore ¼ 3.825 1010 m. Using London’s equation we obtain w(r) ¼ 1.64 1021 J. Consequently, Ust ¼ NA 10 w(r) ¼ 9.86 kJ/mol and Hst ¼ 7.37 kJ/mol. Compared to experiment, this estimate is quite low, reflecting first that these estimates are a minimum value since they exclude interaction with many nearby atoms, but probably also reflect that the value of a0,2 for Si–O–Si groups is underestimated (on the basis of the London formula it is not possible to reach the observed isosteric heats otherwise). Alternatively, one can consider an average potential energy over a layer of oxygen atoms distributed uniformly on the wall of a cylindrical pore as reported by Savitz et al.4 based on the model of Everett and Powl,5 using the Lennard-Jones 126 potential, with parameters e1s and s12. The potential energy of this model is given by
( " # " #) 1 1 r 2k r 2k s 4 X 5 21 s12 10 X 12 ð7Þ wðrÞ ¼ pe1s ak bk 2 32 R R R R k k where ak and bk are defined by
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1=2
ak
1=2 bk
Gð4:5Þ Gð4:5 kÞGðk þ 1Þ Gð1:5Þ ¼ Gð1:5 kÞGðk þ 1Þ
¼
ð8Þ
The gas–solid energy parameter e1s is given by e1s ¼ 6=5pros e12 s12 where rOs is the number of oxygen atoms per unit of area on the pore surface, and the collision ˚ using values of 2.7 A ˚ for diameter (sum of molecular radii) is s12¼3.195 A ˚ the diameter of an oxygen atom and 3.691 A for a methane molecule. The results of the calculation (Fig. 9.2) show that the shape of the potential energy as a function of the distance from the centre of the pore is different for the three zeolites. For ZSM-22 (TON), the minimum is at the centre of the pore while the minima for ZSM-12 (MTW) and UTD-1 (DON) are closer to the wall. In fact, the shape of the curve for DON is quite similar to the potential for a flat wall indicating that the transition from a microporous environment to a mesoporous molecular environment (i.e. one in which interactions resemble a flat surface) starts to be observed even for 14-ring zeolites.
Figure 9.2 Potential energy of a methane molecule in 1-D zeolites ZSM-22, ZSM-12 and UTD-1 ([SiO2]-TON, MTWand DON) as a function of distance from the pore centre.The value of the well depth was obtained from an independent calculation E12/kB¼133.33. The average radii for the three ˚ , respectively, and the zeolites (framework-typesTON, MTWand DON) are 3.825, 4.275 and 5.725 A ˚ 2. average densities of oxygen atoms on the pore surfaces are 0.166, 0.179 and 0.184 O atoms/A
Intermolecular Forces in Zeolite Adsorption and Catalysis
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The desired results are predictions of the isosteric heats of adsorption, and these can be obtained from the Boltzmann weighted average energy ―
RR
w ¼ R R0 0
wðrÞewðrÞ=kB T r dr
ðewðrÞ=kB T 1Þr dr
ð9Þ
and by the following equation ―
Hst ¼ NA w þRT
ð10Þ
The comparison of this theory to experiment is very reasonable (Table 9.4). The average difference is 8%, larger than experimental error but small nevertheless, with the highest deviation for ZSM-22. The best agreement is found for UTD-1, a zeolite with pores that are quite uniform—that is, framework oxygen atoms are at the same distance from the pore centre as one moves along the pore direction. The internal surface of the pores of ZSM-22 and ZSM-12, on the contrary, is corrugated and the difference between theory and experiment may reflect this deviation from the model assumptions. The important conclusion is that simple models based on van der Waals interactions can capture semiquantitatively the isosteric heats of adsorption of simple non-polar molecules on zeolite microporores. For large molecules like dodecane and above, van der Waals forces often dominate the interaction molecule–zeolite over other forces, and can be used as a first estimate of the selectivity of adsorption respect to the gas phase. Finally, in Fig. 9.3 the isosteric heat of adsorption at zero coverage for several non-polar molecules is plotted versus the electronic polarizability for zeolite [SiO2]MFI. These data show that there is indeed a very good correlation between the heats of adsorption and the electronic polarizability. For larger molecules, this correlation is expected to break down as London’s model assumes the molecules can be represented by one point, and clearly this is already incorrect even for molecules as small as propane. Nevertheless, the correlation is excellent.
4. Induced Dipole and Quadrupole–Charge Interactions We next consider the interaction of molecules with cations coordinated to the zeolite pore walls. An electric field induces a dipole, uind, on all molecules (polar and non-polar) that is proportional to the polarizability and the strength of the electric field (uind ¼ a E). The size of this induced dipole can be substantial (of the order of 1 Debye), and thus this interaction can be quite important in determining the overall energetics of an adsorbed molecule inside a zeolite pore. For non-polar molecules, the polarizability is composed only of the electronic polarizability term such that a¼a0, but for polar molecules the polarizability contains another term denoted orientational polarizability
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Figure 9.3 Isosteric heats of adsorption on [SiO2]-MFI zeolites versus electronic polarizability. Note the close correlation between molecular properties such as electronic polarizability and observable macroscopic properties such as isosteric heats of adsorption. (Data from Dunne et al.).22,23
that arises from the effect of an external field on the average orientation of the molecule in space. The net ion-induced dipole interaction free energy is given by
ðzeÞ2 a ðzeÞ2 u2 wðrÞ ¼ ¼ a0 þ 3kB T 2ð4pe0 eÞ2 r 4 2ð4pe0 eÞ2 r 4
ð11Þ
where e is the charge of the electron, z is the valence of the cation, u is the permanent dipole of the molecule and e is the dielectric constant of the medium. First, note that this interaction depends on the negative fourth power of r, and thus it is a longer-range interaction than the dispersion forces. Second, inside the zeolite pores it is not obvious what one should use as the value of the dielectric constant e. The dielectric constant of quartz is 4.8, while the one of amorphous silica glass is 3.8, and one may think that a density-scaled dielectric constant could be used as a good approximation.c This is adequate within mesoscopic and larger length scales; however, in the limit of zero coverage, the length scale of interest are distances comparable to the pore diameter or smaller, so a more appropriate value of e inside the zeolite pore is e 1 (i.e. comparable to air or vacuum). At the same time, the weaker dependence on distance makes it doubly important that we understand the role of
c
To the best of our knowledge the dielectric constant of siliceous zeolites has not been measured.
Intermolecular Forces in Zeolite Adsorption and Catalysis
249
other adsorbed molecules can have on this force. If the pore is filled with water molecules (eH2 O ¼ 88 in the bulk), for instance, the magnitude of this interaction will decrease by a factor of 80–90.d A non-uniform distribution of charge within the molecule gives rise to a second type of interaction with an electric field even when a molecule has no permanent dipole. This distribution of charge is called a quadrupole , and the interaction between a quadrupole and the electric field generated by a cation can be very important in zeolites as we will see shortly. In general, the quadrupole of a molecule is a tensor quantity, and its interaction with a charge (a cation) is angle dependent.6 Frequently though, we will be interested in molecules (such as N2 and O2) that have a linear distribution of charge; and in such cases, the quadrupole can be thought of as a two equidirectional but opposed dipoles (þþ). The energy of interaction w(r) between a linear molecule with quadrupole moment and a point charge is given by
wðrÞ ¼
ze ð3 cos 2 y 1Þ 2r 3 ð4pe0 eÞ
ð12Þ
This interaction can be attractive or repulsive depending on molecular orientation ˚ and the interaction is a and the sign of the charge and quadrupole. For N2 Y¼1.39 D A maximum when the molecule axis is aligned along the cation-molecule direction (y ¼ 0; ð3 cos2 y 1Þ ¼ 2), and minimized with a different sign when the molecule is perpendicular to the cation-molecule direction (y ¼ p=2;ð3 cos2 y 1Þ ¼ 1). Since the angle average of hcos2 yi ¼ 1=3, the average quadrupole interaction can be neglected at high temperatures when molecules tumble isotropically in the zeolite pores (i.e. h3 cos 2 y 1i ¼ 0). Cations in a zeolite can be highly exposed to molecules in the pores, can be partially blocked in such a way that only small molecules have direct contact with the cation, or can be coordinated to the zeolite framework such that there is no—or only very weak—interaction with adsorbed molecules. Because of this distribution of cation coordination environments, adsorbed molecules will find a complex landscape of adsorption sites. Moreover, different zeolites will show different zerocoverage heats of adsorption precisely because the distribution of coordination environments is different in each case. Because of the negatively charged zeolite oxygen atoms and due to charge transfer processes to the cation from the framework oxygen atoms, the effective charge of the cation is substantially lower than the charge of an isolated cation (i.e. charge < e). To clarify this point, Table 9.5 shows the isosteric heat of adsorption of cation-containing zeolites compared to their all-silica analogues. We can indeed see that the presence of sodium cations in the pores substantially increases the heat of adsorption of the molecules. Figure 9.4 compares the sodium cation affinity of several of the molecules in the table to the
d
The local dielectric constant of a medium is dependent on the structure of the solvent and inside a micropore such structure is bound to be quite different from the bulk structure of the solvent. The effective dielectric constant will be different from 88, although it is not possible to evaluate the local dielectric constant without recourse to atomistic models of the solvent in the pores. It should be dependent on pore size and zeolite composition.
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Table 9.5 Isosteric heats of adsorption of small molecules on MFI-type and FAU-type zeolitesa Hst (kJ/mo) [SiO2]Adsorbate MFI
Ar CH4 SF6 O2 N2 C2H6 CO2 a b
15.8 20.9 34.4 16.3 17.6 31,1 27.2
jHþj [SiAlO2]MFI
35.2 20.7 33.3 38.0
jNaþj [SiAlO2]MFI
18.0 26.5 42.0 17b 24.1 38.0 50.0
[SiO2]FAU
11.0
19.7
jNaþj [SiAlO2]-FAU (NaX)
12.7 19.2 28.2 15.0 19.9 27.0 49.1
Data from Dunne and Savitz et al.7,22–24 This datum estimated from figure in ref. 7.
Figure 9.4 Gas-phase sodium cation affinity versus excess isosteric heat above the value of the siliceous zeolite.The data show excellent correlation between the two sets of numbers indicating that the interaction of small molecules with cations inside zeolite pores resemble energetically (and probably structurally) the molecule-cation adducts formed in the gas phase. Sodium cation affinity data from NIST Chemistry Webook and from Table 9.5.
Intermolecular Forces in Zeolite Adsorption and Catalysis
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excess isosteric heat of adsorption due to the presence of cations on the zeolite (the difference between the sodium-cation containing zeolite and its all-silica analogue). We can clearly observe a very good trend between the two quantities. Within some scatter, the trend is linear, (although notice that the excess heat for methane in NaZSM-5 and in Na-X (FAU) differ by more than 3 kJ/mol). This linear trend is an indication that the geometry of the adsorbed molecule on the cation, and type of interaction in the gas phase and in the zeolite pores are quite similar. However, the magnitude of the excess heat of adsorption is substantially lower than the DHr of the interaction of the molecule with the cation in the gas phase. This large difference is the consequence of the coordination of the negatively charged framework to the cation, that is, the effective weaker charge of the cation. To illustrate this further, let’s consider a specific example of a cation coordinated to an AlO4/2 tetrahedron on the surface of a high-silica zeolite. For simplicity, we will consider an idealized model of an exposed cation as depicted in Fig. 9.5. This
Figure 9.5 Schematic of the coordination of sodium cations to a model AlO4/2 tetrahedron. Distances are not to scale. Since this site has a threefold axis of symmetry, the Al^O bond distance is 1.73 A˚ and the angle between the O^Al^O atoms is 109.47 (the tetrahedral angle), all distances can be calculated using elementary geometry.
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figure shows the sodium cation coordinated to three of the oxygen atoms of the alumina tetrahedron. The interaction energy at contact between a methane molecule and the sodium cation is given by the formulae above. For a NaþCH4 distance of 2.79 A˚ and an effective charge of þ1 e, w(r)¼27.8 kJ/ mol, much larger than the excess isosteric heat of 8.2 kJ/mol. This is of course because the effective charge on the cation is smaller due to the effect of the negatively charged oxygen atoms. Using 8.2 kJ/mol as the benchmark, the effective charge can be back calculated to be 0.54 e. Using this effective charge, we can then calculate the total interaction for N2, a molecule with a ˚ between Naþ and substantial quadrupole moment. Using a distance of 3.17 A the centre of mass of N2, and assuming N2 is aligned along the molecule-cation axis, the induced dipole and quadrupole interactions are 3.4 and 6.6 kJ/mol, respectively. The total 10 kJ/mol is very much in agreement with the excess isosteric heat between silicalite ([SiO2]-MFI) and NaZSM-5 (|Naþ|[SiAlO2]MFI) of 9 kJ/mol (Table 9.5).
5. Dipole–Charge Interactions The interaction of polar molecules—molecules with permanent dipole moments—with cations on the surface of a zeolite pore is the last type of electrostatic interaction that will be reviewed here. To start, let’s recall that the dipole moment of a polar molecule u is defined as u¼q l where l is the distance between two charges þq and q. Small molecules have dipole moment of the order of 1 Debye ¼ 3.3361030 C m. The energy of interaction between an ion and a dipole is given by
wðrÞ ¼
ðzeÞucosy 4pe0 er 2
ð13Þ
where the symbols are as defined before, and y is the angle between the direction of the dipole moment u and the line joining the centre of mass of the molecule and the ion. This interaction can be attractive or repulsive depending on the angle y. It is maximum when the angle is 0 . In practice, the interactions between an ion and a polar molecule can be divided into strong and weak depending on whether the interaction at y ¼ 0 is large compared to kBT. In such a case, the molecule will tend to be oriented towards the ion and would be coordinated to it for long periods of time. This is expected for small and multivalent cations interacting with small molecules containing a sizable dipole moment. For this case the energy is given by
wðrÞ ¼
ðzeÞu 4pe0 er 2
ð14Þ
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Intermolecular Forces in Zeolite Adsorption and Catalysis
Table 9.6 Comparison of calculated electrostatic interactions of CO and water molecules with sodium cations in the gas phase to isosteric heat of adsorption in Na–ZSM-5 Interaction
CO kJ/mol
H2O kJ/mol
Dipole–charge Quadrupole–charge Induced dipole Total DHst Silicalite-1 DHst Na–ZSM-5 DDHst
–3.65 –32.0 –17.9 –53.5 17 337 16
–96.7 –2.9 –34.1 –134 30 9525 65
For weak interactions, one can assume that the molecule tumbles with isotropic motion in the vicinity of the ion and one requires the angle-averaged charge–dipole interaction given by1
wðrÞ
q2 u2 6ð4pe0 eÞ2 kB Tr 4
for kB T >
qu 4pe0 er 2
ð15Þ
For instance, the heat of adsorption of CO on NaZSM-5 (33 1 kJ/mol at 195 K) and Silicalite-1 (17 1 kJ/mol)7 can be compared to the enthalpy of reaction of CO and Naþ(to form CO Naþ) in the gas phase (DHr ¼ 52.7 kJ/mol).5,e The difference between the two zeolites reflects the additional contribution of the sodium cations to the heat of adsorption, that is, 16 kJ/mol or about 30% of the gas-phase value. This is comparable to what we observed when we estimated the effective charge of sodium cations using the nitrogen molecule. This heat of adsorption on sodium cations has two contributions: one from the dipole–charge interaction, and another from the induced dipole moment on CO by the electric field of the cation. Using the following data (rCO ¼ 0.2 nm, rNa ¼ 0.095 nm, q ¼ ze ¼ 1.6 1019 C, a0 ¼ 1.95 1.11 10–40 C2 m2 J1, T ¼ 300 K, kB ¼ 1.381023 J K1, dipole moment of CO ¼ 0.113.3361030 C m, ¼ 9.47 10–40 C m2), the magnitude of the various interactions can be determined (Table 9.6). Water is an example of a small molecule with a large dipole moment. Water, in its simplest form, can be treated as a spherical molecule of 0.14 nm in radius. It has a dipole moment u ¼ 1.85 D and it can adsorb into a sodium cation (0.095 nm in radius). Calculation of the maximum interaction energy for this complex in the gas phase gives a total of w(r) ¼ 134 kJ/mol, including dipole, polarizability and quadrupole interactions. The dipole-charge contribution is by far the most important one. (Note that this value is higher than the value of 95 7.9 kJ/mol that has been measured experimentally for the reaction Naþ þ H2O ¼ Na H2Oþ in the gas phase.)8 The heat of adsorption of water in NaZSM-5 and NaA zeolites has
e
Values of 32 7.9 kJ/mol have also been reported.
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been measured by Dubinin and co-workers9 giving values in the range of 95–105 kJ/mol. This energy of adsorption is dominated by the dipole–charge interaction (see Table 9.6). For the zeolites, the contribution of the sodium cation corresponds to 65 kJ/mol, somewhat higher than the expected value on the basis of the partial charge of the sodium cation on the zeolite pores. This is the case because the water molecule can coordinate by hydrogen bonding to the oxygen atoms of the zeolite, adding further to the heat of adsorption.
6. Interactions of Adsorbed Molecules with Acid Sites The interaction of molecules with Brnsted acid sites in zeolites deserves a separate discussion because the binding of protons to the zeolite framework is very different from the binding of cations like sodium and potassium. Likewise, the coordination of small molecules to acid sites can be quite different from the coordination to alkali-metal cations. Figure 9.6 illustrates the structure and geometry of the acid sites in zeolites as has been determined for high-silica zeolites.10 The key element of the site is the presence of an Al–OH–Si group. The bond between the hydrogen atom and the zeolite oxygen atom is predominantly covalent in nature, and there is only a small effective charge on the hydrogen atom. This site can coordinate or donate the proton to adsorbed molecules forming 1:1, 2:1 and
Figure 9.6 Schematic of the structure and geometry of the classical Brnsted acid site in highsilica zeolites. The structures on the right illustrate the coordination of weak and strong bases to the acid site. For strong bases, there is proton transfer and charge separation between the framework and the adsorbed molecule.
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sometimes higher adducts. The oxygen atom in the acid site is tri-coordinated, and since it should keep a bond order of 2, there is a weakening of the three bonds, but particularly the O–H bond; this allows for facile transfer of the proton to adsorbed molecules. The top section of Fig. 9.6 also highlights the presence of three additional oxygen atoms coordinated to the aluminum that have Lewis basic character and that can coordinate through hydrogen bonding to molecules that exhibit both hydrogen-acceptor and hydrogen-donating groups, such as water or ammonia. Usually, two out of the three oxygen atoms—but sometimes only one—are physically accessible to the adsorbed molecules due to the specific configuration of the framework atoms with respect to the zeolite pore (see Fig. 9.5). These geometric differences appear to have no major thermochemical consequence11 and the heat of adsorption of strong bases on high-silica zeolites is the same—within experimental error—regardless of the specific geometry of the site. Figure 9.6 also shows the various bond lengths of the acid site. Note that the ˚ vs 1.73 A ˚ ); Al–OH bond is substantially longer than the other three bonds (1.92 A and as a consequence the coordination environment of the aluminum atom is not tetrahedral (as is the case for framework aluminum coordinated to alkali-metal cations), but an intermediate structure between tetrahedral and planar trigonal. The long aluminum–oxygen bond reflects, of course, the fact that this bond is weaker than the other three. Because of the elongation and weakening of the O– H bond, there is a change in the infrared (IR) absorption frequency from 3749 cm1 (for isolated Si–OH groups) to 3614 cm1 (Al–OH–Si).12 This change in vibrational frequency is also an indication that the ability to donate the proton of a zeolite Brnsted acid site is much larger than the ability of an isolated silanol group. There are two extremes of interaction of a zeolite acid site ZOH with a base B. For weak bases the adduct forms a hydrogen bond to the hydroxyl group (ZOH B), while for strong bases there is complete proton transfer and charge separation (ZO HBþ). And there are, of course, intermediate cases that fall between these two cases. A series of calorimetric, computational and spectroscopic investigations have shown that the molecular property that correlates more closely to the observed thermochemical and spectroscopic measurements is the proton affinity (PA) of the base. If the PA is below 858 kJ/mol for the 1:1 adduct, there is no proton transfer and the complex is described more closely as ZOH B. If the PA is above 858 kJ/mol, there is proton transfer and the complex is better described as ZO HBþ. Ammonia is right at this limit with a PA ¼ 858 kJ/mol, and definitely forms the ammonium ion upon adsorption on acid zeolites.13 Figure 9.7 shows the IR spectra of 1:1 adducts on zeolite beta at 1:1 coverage for bases of different PA. The perturbations of the IR spectra increase as the PA increases (see Table 9.7) up to pyridine, where there is proton transfer as confirmed by the observation of the pyridinium ion signatures in the IR spectrum. As illustrated in Fig. 9.6, the adsorption of weak bases slightly perturbs, but not substantially changes the geometry of the acid site. This weak hydrogen bond interaction is mainly of electrostatic character. It is weak because the effective charge of the hydrogen atom in the ZOH group is small. This is evident from Table 9.5 where we can observe that the excess isosteric heat for HZSM-5 versus NaZSM-5 is small, regardless of the small effective radius of the hydrogen atom.
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Figure 9.7 Background subtracted infrared (IR) spectra of bases of different proton affinity (PA). Nitrogen and CO conducted at 100 K, C3H6 at 180 K and the rest at room temperature.The perturbation of the IR spectra of clean H-beta increases with PA. The spectra become increasingly more complex due to mixing and doubling of bending and stretching modes. Reproduced with permission from Paze et al.12
The abstraction of the proton by a strong base, in contrast, changes the geometry of the aluminum from a distorted tetrahedron to a symmetric tetrahedron where all ˚ . For strong bases, both the differential heat Al–O bonds are similar and about 1.75 A of adsorption and the IR spectra of the adducts are independent of the zeolite structure, indicating that the complexes formed by a molecule on all zeolites are very similar (provided there is space to form the complex). Molecules with hydrogen accepting and donating groups are somewhat different because they can form 1:1, 2:1, and higher order complexes with acid sites in the zeolite pores (see Fig. 9.8). At low coverage (based on the Brnsted acid site ratio), IR spectra indicate that the 1:1 complex dominates for water and methanol, but even before the B:ZOH site ratios approach 1:1, evidence for clusters of molecules
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Table 9.7 Comparison of adsorption enthalpies in H–ZSM-5 and silicalite with gas-phase proton affinities
Molecule
Proton affinity (kJ/mol)
H^ZSM-5 DHads (kJ/mol)
Silicalite DHads (kJ/mol)
N2 CO C2H4 H2O CH3CN CH3OH (C2H5)2O NH3 N(CH3)3 Pyridine
494 594 680 724 798 774 838 858 938 922
20.7 26
17.6 17
90 110 115 135 145 205 200
30 75 65 70
Figure 9.8 Coordination of 1:1and 2:1water^zeolite complexes on acid sites.
are observed. For water and methanol above low coverages, 2:1 and higher order adducts are clearly observed, and here proton transfer is occurring, as may have been guessed from PA data. For instance, the gas phase PA of H2O is 723 kJ/mol (weak), but the PA of (H2O)2 is 850 kJ/mol, similar to ammonia (strong). Calorimetric investigations indicate that there is dynamic equilibrium between the 1:1 and 2:1 adducts (again, for water and methanol) at room temperature, even before the total coverage reaches 1:1 B:ZOH site ratios.14,15 The interaction of alkenes with acid sides is briefly discussed next because of its particular importance in petrochemical catalytic processes.16 The precise nature of this interaction remains a topic of ongoing research,17 yet important insights have been gained that start to clarify how organic molecules bind to zeolites. Alkenes are different from the two cases discussed above because there can be proton transfer to the alkene, yet a neutral species could be the result of this chemisorption process. The main question is then whether alkenes form silil-alkoxide species
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Figure 9.9 Comparison of the structure alkoxide complexes and stable carbocations formed on zeolite acid sites. The left side shows the structure formed by alkenes reacting with zeolite Brnsted acid sites.The right side shows an example of a stable carbocation formed by protonation of dimethyl-cyclopentadiene.
upon reaction with an acid site, or is there charge separation to form carbenium ions (Fig. 9.9). The evidence indicates that it depends on the structure of the organic molecule.18 Ethene is a weak base (PA ¼ 680 kJ) and at low temperature it forms a hydrogen-bonded complex with acid sites in zeolites, but upon reaction it is predicted to form the neutral ethoxide group, covalently bonded to the zeolite framework. The complex formed by propene has been difficult to establish because it rapidly forms oligomeric species.19 The tert-butyl alkoxide (from isobutene, PA ¼ 802 kJ/mol) has been isolated in some zeolites, but there are indications that steric hindrances can shift the equilibrium towards the tert-butyl cation. In summary, the bulk of the evidence indicates that primary and secondary alkenes will form covalent (neutral) complexes with zeolite acid sites. Tertiary alkenes can form covalent or charged complexes, depending on the specific structure of the alkene and the zeolite. In the last decade, there has been accumulating evidence demonstrating that organic molecules with very high PAs (such as 1,3-dimethylcyclopentadiene, PA ¼ 878 kJ/mol) do indeed form stable ions inside the zeolite pores.10,17,20,21 These stable carbocations can (Fig. 9.9), in fact, be very important in the conversion of hydrocarbons such as the methanol to gasoline and methanol to olefins processes. In addition to 1,3-dimethylcyclopentadiene, various stable carbocations have been identified inside the zeolite pores using nuclear magnetic resonance (NMR) spectroscopy, UV/vis spectroscopy and other techniques. Summarizing, the interaction of small molecules with zeolite acid sites is fundamentally a chemical interaction that can be broadly described as Lewis-base assisted, Brnsted acid–base interaction. Because it is a chemical interaction, often chemical bonds are made and broken and the end result of this process depends on the specific structure of the adsorbate. All acid sites in high-silica zeolites appear to be structurally very similar and the complexes they form with small molecules are, in thermochemical and spectroscopic terms, similar to each other (within a zeolite structure and between different zeolite structures as well). The gas-phase PA of the adsorbed molecule is the molecular property that gives the most useful information towards predicting the nature of the adduct formed upon adsorption.
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7. Hydrophobic Interactions in High-Silica Zeolites Water is unique because of its ability to form three-dimensional networks of hydrogen bonds, and this specific interaction leads to a remarkably high enthalpy of vapourization (DHvap ¼ 43.0 kJ/mol at 298 K) as compared to molecules of similar size like methane (DHvap ¼ 8.5 J/mol at 99 K). This is despite the fact that water molecules have a rather low electronic polarizability, and thus have relatively weak dispersion forces with other molecules. For instance, water has a polarizability of 1.5 ˚ 3 as compared to methane that has a polarizability of 2.5 of 4pe0 A ˚ 3 (see 4pe0 A Table 9.1). Inside a purely siliceous zeolite, at low coverage, the only force that governs the interaction of water molecules with the zeolite framework is the dispersion or London force. Because of its low polarizability, this force is low in magnitude, and it is especially low compared to its high enthalpy of vapourization. For comparison, methane’s heat of adsorption in siliceous MFI (silicalite) is 20.9 kJ/mol, much higher than its heat of vapourization, yet the heat of adsorption of water is smaller than the heat of adsorption of methane. For this reason, high-silica zeolites can be used to selectively adsorb hydrocarbons even in the presence of high vapour pressures of water, a property that it is nearly unique among readily available industrial sorbents. This is very different from high-surface-area porous silicas that contain a surface covered with silanol groups and are hydrophilic (in the absence of surface treatments). In fact, purely siliceous zeolites (without defects such as internal silanol groups or tetrahedral vacancies) belong to a class of materials called hydrophobic solids. In these materials, water condensation occurs at a pressure above the saturation vapour pressure. For instance, the condensation of water in defect-free silicalite-1 is about 100 MPa at room temperature. For zeolites with larger pores, the condensation pressure is still high: for siliceous zeolite beta (with a three-dimensional-12-ring pore system) it is above 50 MPa at room temperature. This additional pressure can be thought to be the work that is needed to accommodate the massive disruption of the hydrogen bond network that is unavoidable upon the incorporation of water molecules into a microporous, hydrogen-bond-free environment such as the one presented in the pores of siliceous zeolites.
8. Final Remarks This chapter has presented an overview of the important classes of zeolite– molecule forces that determine the adsorption and some catalytic properties of zeolite materials. Our intention is to relate molecular properties of the adsorbates to measurable quantities such as heats of adsorption. The data reveal that qualitatively, this approach helps to understand differences in the adsorption of small molecules for multiple classes of zeolite structures and compositions. At the same
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time, often differences are observed between the predictions of the simple models used here and the actual measurements of heat of adsorption. These reveal the limitations of this simplistic approach, yet also point to aspects of zeolite–molecule interactions that are not well understood. In particular, further calorimetric measurements of hydrogen-bond accepting and donating groups are needed to clarify the separate contribution of OH bonds and electrostatic forces.
REFERENCES [1] Israelachvili, J. N., Intermolecular and Surface Forces. 2nd ed. Academic Press: London, 1991; p 450. [2] Baerlocher, C., Meier, W. M., Olson, D. H., Atlas of Zeolite Framework Types. 5th ed.; Elsevier: Amsterdam, 2001. [3] Baerlocher, C., McCusker, L. B. Database of Zeolite Structures. http://www.iza-structure.org/ databases/ [4] Savitz, S., Siperstein, F., Gorte, R. J., Myers, A. L. ‘‘Calorimetric study of adsorption of alkanes in high-silica zeolites’’, J. Phys. Chem. B 1998, 102, (35), 6865–6872. [5] Everett, D. H., Powl, J. C., ‘‘Adsorption in slit-like and cylindrical micropores in Henry’s Law region: Medol for microporosity of carbons’’, J. Chem. Soc. Faraday Trans. I 1976, 72, 619–636. [6] Papai, I., Goursot, A., Fajula, F., Plee, D., Weber, J., ‘‘Modeling of N-2, and O-2 Adsorption in Zeolites’’, J. Phys. Chem. 1995, 99, (34), 12925–12932. [7] Savitz, S., Myers, A. L., Gorte, R. J., ‘‘A calorimetric investigation of CO, N-2, and O-2 in alkali-exchanged MFT’’, Micropor. Mesopor. Mater. 2000, 37, (1-2), 33–40. [8] NIST Chemistry Web book. http://webbook.nist.gov/chemistry/ [9] Dubinin, M. M., Rakhmatkariev, G. U., Isirikyan, A. A., ‘‘Differetial heats of adsorption of water vapor on NaA zeolite at 300 and 450 K’’, Izest. Ak. Nauk SSSR, S. Khim. 1989, 12, 2877–2778. [10] Beck, L. W., Xu, T., Nicholas, J. B., Haw, J. F., ‘‘Kinetic Nmr And Density-Functional Study Of Benzene H/D Exchange In Zeolites, The Most Simple Aromatic-Substitution’’, J. Am. Chem. Soc. 1995, 117, 11594–11595. [11] Savitz, S., Myers, A. L., Gorte, R. J., White, D., ‘‘Does the cal-ad method distinguish differences in the acid sites of H-MFI?’’, J. Phys. Chem. B 1998, 120, (23), 5701–5703. [12] Paze, C., Bordiga, S., Lamberti, C., Salvalaggio, M., Zecchina, A., Bellussi, G., ‘‘Acidic properties of H-beta zeolite as probed by bases with proton affinity in the 118-204 kcal mol (-1) range: A FTIR investigation’’, J. Phys. Chem. B 1997, 101, (24), 4740–4751. [13] Zecchina, A., Marchese, L., Bordiga, S., Paze, C., Gianotti, E., ‘‘Vibrational spectroscopy of NH4+ ions in zeolitic materials: An IR study’’, J. Phys. Chem. B 1997, 101, (48), 10128–10135. [14] Haase, F., Sauer, J., ‘‘Interaction Of Methanol With Bronsted Acid Sites Of Zeolite Catalysts An Ab-Initio Study’’, J. Am. Chem. Soc. 1995, 117, (13), 3780–3789. [15] Haase, F., Sauer, J., ‘‘Ab initio molecular dynamics simulation of methanol interacting with acidic zeolites of different framework structure’’, Micropor. Mesopor. Mater. 2000, 35-6, 379–385. [16] Haw, J. F., Song, W. G., Marcus, D. M., Nicholas, J. B., ‘‘The mechanism of methanol to hydrocarbon catalysis’’, Acc. Chem. Res. 2003, 36, (5), 317–326. [17] Clark, L. A., Sierka, M., Sauer, J., ‘‘Stable mechanistically-relevant aromatic-based carbenium ions in zeolite catalysts’’, J. Am. Chem. Soc. 2003, 125, (8), 2136–2141. [18] Nicholas, J. B., Haw, J. F., ‘‘The prediction of persistent carbenium ions in zeolites’’, J. Am. Chem. Soc. 1998, 120, (45), 11804–11805. [19] Gorte, R. J., White, D., ‘‘Interactions of chemical species with acid sites in zeolites’’., Topics Catal. 1997, 4, (1-2), 57–69.
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[20] Boronat, M., Viruela, P. M., Corma, A., ‘‘Reaction intermediates in acid catalysis by zeolites: Prediction of the relative tendency to form alkoxides or carbocations as a function of hydrocarbon nature and active site structure’’, J. Am. Chem. Soc. 2004, 126, (10), 3300–3309. [21] Song, W. G., Nicholas, J. B., Haw, J. F.,‘‘Acid-base chemistry of a carbenium ion in a zeolite under equilibrium conditions: Verification of a theoretical explanation of carbenium ion stability’’, J. Am. Chem. Soc. 2001, 123, (1), 121–129. [22] Dunne, J. A., Mariwals, R., Rao, M., Sircar, S., Gorte, R. J., Myers, A. L., ‘‘Calorimetric heats of adsorption and adsorption isotherms .1. O-2, N-2, Ar, CO2, CH4, C2H6 and SF6 on silicalite’’, Langmuir 1996, 12, (24), 5888–5895. [23] Dunne, J. A., Rao, M., Sircar, S., Gorte, R. J., Myers, A. L., ‘‘Calorimetric heats of adsorption and adsorption isotherms .2. O-2, N-2, Ar, CO2, CH4, C2H6, and SF6 on NaX, H-ZSM-5, and Na-ZSM-5 zeolites’’, Langmuir 1996, 12, (24), 5896–5904. [24] Savitz, S., Myers, A. L., Gorte, R. J., ‘‘Calorimetric investigation of CO and N-2 for characterization of acidity in zeolite H-MFI’’, J. Phys. Chem. B 1999, 103, (18), 3687–3690. [25] Dubinin, M. M., Rakhmatkariev, G. U., Isirikyan, A. A., ‘‘Energy of adsorption of water vapor adsorption on high and pure-silica zeolites’’, Izest. Ak. Nauk USSR, S. Khim. 1989, 12, 2862–2864.
9
Intermolecular Forces in Zeolite Adsorption and Catalysis Raul F. Lobo Contents 239 240 242 247 252 254 259 259 260
1. Introduction 2. Basic Thermodynamic Concepts 3. Dispersion (London) Forces 4. Induced Dipole and Quadrupole–Charge Interactions 5. Dipole–Charge Interactions 6. Interactions of Adsorbed Molecules with Acid Sites 7. Hydrophobic Interactions in High-Silica Zeolites 8. Final Remarks References
Abstract The concepts of classical intermolecular forces are used to understand isosteric heats of adsorption of small molecules adsorbed within zeolite pores. This is used as an organizing principle to comprehend the complex and multifaceted set of chemical properties of zeolites in general. We correlate molecular properties to the magnitude of molecule– zeolite forces and these to measurable quantities such as heats of adsorption. Several key thermodynamic concepts are reviewed and then the most relevant types of intermolecular forces as applied to microporous materials are described. The interaction of small molecules with zeolite Brnsted acid sites is then described as a chemical interaction, rather than in terms of physical forces. We conclude with a discussion of hydrophobicity as applied to siliceous zeolites. Keywords:
Adsorption, Intermolecular forces, Calorimetry, Catalysis, Zeolites
1. Introduction As molecules travel down the pore structure of a zeolite, they encounter cations and other chemical groups that expose the molecules to important classes of forces. These forces, and more frequently, the difference between forces for different molecules, ultimately lead to the useful properties of zeolites as adsorbents Ordered Porous Solids DOI: 10.1016/B978-0-444-53189-6.00009-3
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2009 Elsevier B.V. All rights reserved.
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and catalysts, and should be understood as a foundation to the analysis of material properties in molecular terms. In this chapter, we provide the basis for this understanding with emphasis on the origin and magnitude of the various forces relevant to zeolites. We intend this approach to be an organizing principle to comprehend the complex and multifaceted set of chemical properties of zeolites in general. We explicitly attempt to correlate molecular properties to the magnitude of the forces and these to measurable quantities such as heats of adsorption. To keep the length of this chapter reasonable, it is assumed throughout that we are working at the dilute limit, and that molecule–molecule forces are insignificant. We start with some basic thermodynamic concepts, and then proceed to describe the most relevant types of intermolecular forces as applied to microporous materials and finish with a discussion of hydrophobicity as applied to siliceous zeolites.
2. Basic Thermodynamic Conceptsa The pair potential w(r)—in vacuum—or potential of mean force—in a medium— describes the interaction potential between two molecules.1 The force between the molecules is given by F ¼ dw(r)/dr and shows that the work w(r) that can be done by the force F is also a measure of the available energy or free energy of this system of two molecules. The medium in the case of microporous materials is the micropore and is different from more classical media such as a liquid solvent. First, in a liquid, the solute molecules can perturb the local structure of the solvent; in a zeolite the framework atoms are essentially fixed in space and barely move in response to the presence of an adsorbate in the pores. Second, when a molecule is introduced into a liquid, work must be done to create a cavity to accommodate the guest molecule; in zeolites, the cavity (the pores) is already there and no energy is needed to form it. A zeolite is, however, also like a liquid in the sense that for two solute molecules their pair potential includes molecule–molecule interactions, and also solute–solvent interactions (or equivalently molecule–molecule and molecule– zeolite interactions). A zeolite, just like a solvent, can also change the properties of the adsorbed molecule (dipole moment and polarizability, geometry) with respect to the molecular values in vacuum. A molecule in medium has a cohesive energy mi in a medium i given by the sum of its interactions with the surroundings. Inside a siliceous zeolite pore, these interactions are mainly the interaction between the molecule and Si–O–Si groups that surround the pore. It is possible to relate the cohesive energy to the pair potential w(r). For instance, consider the pair potential given by a power law such as
wðrÞ ¼ C=r n wðrÞ ¼ 1
a
for r > s for r s
The description of intermolecular forces used here follows closely the treatment used by Israelachvilli 1.
ð1Þ
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Figure 9.1 Illustration of the pore size of a 10-ring zeolite (in this case zeolite ZSM-5).The effec˚ and the oxygen-to-oxygen distance is 8.1 A˚. tive pore opening is 5.5A
where s is the hard sphere diameter of the molecule and r is the distance between the centre of the molecule and another atom (or molecule). We can calculate mi by summing all the pair potentials over all space, and then make use of the fact that zeolites are crystalline and atoms are located at specific positions in space. Alternatively, we can estimate this value by considering only the first row of zeolite oxygen atoms that surround the molecule and placing the molecule in the centre of the pore. For a 10-ring zeolite (see Fig. 9.1), the pore opening is 5.5 A˚ (see the effective pore dimensions of the MFIb framework type in the Atlas of zeolite structure-types2 or the web3) and the distance between oxygen atoms at opposite ˚ (the effective pore diameter ends of the pore is 2rpore 5.5þ2 1.3 ¼ 8.1A ˚ ). The net energy reported in the Atlas plus twice the ionic oxygen radius of 1.3 A change by taking a molecule from the gas phase to the centre of a zeolite 10-ring pore is given then by
mizeo 10 wðrpore Þ
rpore ¼ 4:05A˚
ð2Þ
where w(r) describes, in particular, the pair potential between zeolite oxygen atoms and the adsorbed molecule. This value is a lower bound of the cohesive energy because the off-axis position may be more favourable than the centre, and because contributions from other zeolite oxygen and silicon atoms further away b
We will use the three-letter-code described in the Atlas to designate unambiguously the topology of the framework of the various zeolite materials that will be discussed in this chapter.
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have not been included. This function already tells us, for instance, that the cohesive energy of a molecule in a zeolite decreases as the effective pore size increases, since w(r) is a monotonically decreasing function of r at the relevant molecular distances. Consequently, the molar cohesive energy U of an adsorbed molecule can be calculated simply by multiplying by Avogadro’s number U ¼ NA mizeo . The cohesive energy of a molecule can have different values in different parts of a system (say, in the zeolite pores and in the gas phase). Then, at equilibrium, the concentrations of the molecules in the two regions is given by the Boltzmann distribution
X1 ¼ X2 exp mi1 mi2 =kB T
ð3Þ
where X1 and X2 are dimensionless concentrations (usually mole fractions or volume fractions) and kB is the Boltzmann constant. One can expect differences in the concentration of species i if the differences in chemical potential are substantially larger than the thermal energy (at 300 K kBT is 2.5 kJ/mol and at 600 K is 5 kJ/ mol). This equation can be restated as
mi1 þ kB T ln X1 ¼ mi2 þ kB T ln X2
ð4Þ
Here, we assume the phases mix ideally and the activity coefficients can then be assumed to be one. Also, this equation is valid only if the condition of equilibrium between the two states is fulfilled. This expression can be extended to a system with many states by
min þ kB T ln Xn ¼ m ¼ constant for all states n ¼ 1; 2; 3 . . . In such a system, molecules will flow between different states until these n equations are satisfied. The term m is also known as the chemical potential, the total free energy per molecule including its interactions with its medium and its thermal energy. The term kB ln Xn is the entropic contribution to the total chemical potential (again assuming ideal mixing).
3. Dispersion (London) Forces Dispersion forces are one of the three components to the total van der Waals force (in addition to induced dipole and orientation polarization forces).1 They are very important because they are always present, independently of the molecular and geometric properties of the molecules. Dispersion forces are, in general, complex to calculate precisely because their magnitude can be affected by the presence of other molecules, that is, it is a non-additive interaction. The origin of this force can be thought of to be a quantum-mechanical induced polarizability force, and it is described for two different atoms 1 and 2 by London’s expression
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Intermolecular Forces in Zeolite Adsorption and Catalysis
wðrÞ ¼
3 a0;1 a0;2 I1 I2 2 6 2 ð4pe0 Þ r ðI1 þ I2 Þ
ð5Þ
where 0,i is the electronic polarizability of molecule i, e0 is the permittivity of free space (8.8541012 C2 J1 m1) and I is the ionization potential (in J). Table 9.1 summarizes the values of the electronic polarizability for small molecules and a few molecular groups, and Table 9.2 lists the ionization potentials of a selected group of molecules. Polarizability values span roughly two orders of magnitude for this group of molecules, and thus this force can vary greatly, depending on the identity of the molecules in question. Note that water molecules have a rather low polarizability similar to the polarizability of argon. The energetics of adsorption of non-polar molecules in high-silica zeolites is dominated by dispersion forces, and thus these forces dictate the value of cohesive energy gained by a molecule as it enters a zeolite from the gas phase. The isosteric heat of adsorption (Hst) in the limit of zero coverage, a quantity that can be obtained experimentally, is a good estimate of this change in cohesive energy. Fortunately, Savitz et al.4 have conducted a comprehensive study of the adsorption of hydrocarbons and other simple molecules on siliceous zeolites. These data will be used to further elucidate the role of dispersion forces in zeolite–adsorbate interactions. Table 9.3 shows the zeolites used in this study and their relevant characteristics. Three unidimensional zeolites of different pore sizes ZSM-22 ([SiO2]-TON, 10-ring), ZSM-12 ([SiO2]-MTW, 12-ring) and UTD-1 ([SiO2]-DON, 14-ring) were used in addition to multidimensional 10-ring ([SiO2]-MFI) and 12-ring Table 9.1 Electronic polarizabilitiesa 0 of atoms, molecules and molecular groupsb Atoms and molecules
He H2 H2O O2 N2 Ar
0.208 0.787 1.501 1.562 1.710 1.664
CO NH3 HCl CO2 CH4 C2H6 C3H8
1.953 2.103 2.515 2.507 2.448 4.226 5.921
C2H4 CH3OH Cl2 CHCl3 C6H6 CCl4
1.28 1.13 2.03
CH2 Si–O–Si Si–OH
1.84 1.39c 1.60
4.188 3.081 4.4 8.2 10.453 10.002
Molecular groups
C–O–H C–O–C C–NH2 a b c
˚ 3¼(4pe0) 1030 m3¼1.111040 C2 m2 J1. Polarizabilities given in units of (4pe0)A Data from NIST Chemistry Webook8 if available (experimental values). Otherwise from compilation in Israelachvilli. This value is probably an underestimation. See discussion below.
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Table 9.2 Ionization potentiala of selected atoms and molecules (in eV)b Atoms and molecules
He H2 H2O O2 N2 Ar a b
24.59 15.42 12.62 12.07 15.58 15.56
CO NH3 HCl CO2 CH4 C2H6
14.01 10.07 12.74 13.78 12.61 11.52
C2H4 CH3OH Cl2 CHCl3 C6H6 CCl4
10.51 10.84 11.48 11.37 9.24 11.47
Data from NIST Chemistry Webook.8 1 J ¼ 6.242 1018 eV.
Table 9.3 Properties of high-silica zeolites used for adsorption measurements of hydrocarbons in Savitz et al.4a
a b
Zeolite framework type
Vpore (cm3/g)b
Si/Al ratio
Pore type
Pore size (n-rings)
Pore/Cavity ˚) dimensions (A
TON MTW DON FER
0.069 0.097 0.106 0.105
52 140 1 22
1-D 1-D 1-D 2-D
10 12 14 8, 10
MFI
0.173
300
2-D
10, 10
FAU
0.245
1
3-D spherical
5.54.4 6.25.5 107.5 4.83.5, 5.44.2 5.35.6, 5.55.1 12.6
Reproduced with permission from Savitz et al.4 Vpore ¼ Pore volume.
([SiO2]-FAU) materials. Ferrierite (FER) has a two-dimensional pore system with 10- and 8-ring pores running perpendicular to each other. The isosteric heats of adsorption for methane, ethane and propane in these zeolites (Table 9.4) follow the trends expected from the pore size of the zeolites and the polarizability of the molecules. For all molecules, the heat of adsorption decreases as the size of the pore or cavity increases. For a given structure, the heat of adsorption increases as the size of the molecule increases. In fact, the increments in heat of adsorption follow quite closely the changes in electronic polarizability (a0) of the adsorbate. For example, for silicalite-1 ([SiO2]-MFI) the heats of adsorption increase in the ratios 1:1.5:2, while the polarizabilities increase in the ratios 1:1.7:2.4, a consequence of London’s formula. Also note that the cohesive energy of the solid hydrocarbons (sum of the latent heats of melting and vapourization, DHmþDHv) is substantially lower than the isosteric heat of adsorption for all zeolites, thus highlighting the strength of the van der Waals forces between the zeolite and the hydrocarbon.
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Table 9.4 Zero-coverage isosteric heats of adsorption of small hydrocarbons in high-silica zeolites at 25 Ca
a
Hst (kJ/mol)
Zeolite framework type
CH4
C2H 6
C3H 8
FER TON MFI MTW UTD FAU DHmþDHv
27.7 27.2 20.9 20.9 14.2 11.0 9.1
41.7 39.0 31.1 29.5 22.2 19.7 17.5
53.3 48.8 41.4 37.6 28.1 26.2 18.6
CH4 Hst (kJ/mol) from theory
22.4 19.7 14.2 –
4
Reproduced with permission from Savitz et al.
The cohesive energy is closely related to the change in internal energy per mole upon adsorption, a quantity that is called differential heat of adsorption (Ud). In turn, the differential and isosteric heats of adsorption are related by
Hst ¼ Ud þ ZRT
ð6Þ
where Z ¼ PV/RT is the compressibility factor in the bulk gas phase (for an ideal gas Z ¼ 1). On the basis of this information, for instance, we can attempt to estimate the isosteric heat of adsorption of methane in zeolite ZSM-22 (TON) for methane. CH4 ¼ 1, Si–O–Si ¼ 2, a0,1 / (4pe0) ¼ 2.448 1030 m, I1 ¼ 12.61 eV ¼ 2.021018J, a0,2 / (4pe0) ¼ 1.39 1030 m, I2 2 1018J (we use a typical value), rpore ¼ 3.825 1010 m. Using London’s equation we obtain w(r) ¼ 1.64 1021 J. Consequently, Ust ¼ NA 10 w(r) ¼ 9.86 kJ/mol and Hst ¼ 7.37 kJ/mol. Compared to experiment, this estimate is quite low, reflecting first that these estimates are a minimum value since they exclude interaction with many nearby atoms, but probably also reflect that the value of a0,2 for Si–O–Si groups is underestimated (on the basis of the London formula it is not possible to reach the observed isosteric heats otherwise). Alternatively, one can consider an average potential energy over a layer of oxygen atoms distributed uniformly on the wall of a cylindrical pore as reported by Savitz et al.4 based on the model of Everett and Powl,5 using the Lennard-Jones 126 potential, with parameters e1s and s12. The potential energy of this model is given by
( " # " #) 1 1 r 2k r 2k s 4 X 5 21 s12 10 X 12 ð7Þ wðrÞ ¼ pe1s ak bk 2 32 R R R R k k where ak and bk are defined by
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Raul F. Lobo
1=2
ak
1=2 bk
Gð4:5Þ Gð4:5 kÞGðk þ 1Þ Gð1:5Þ ¼ Gð1:5 kÞGðk þ 1Þ
¼
ð8Þ
The gas–solid energy parameter e1s is given by e1s ¼ 6=5pros e12 s12 where rOs is the number of oxygen atoms per unit of area on the pore surface, and the collision ˚ using values of 2.7 A ˚ for diameter (sum of molecular radii) is s12¼3.195 A ˚ the diameter of an oxygen atom and 3.691 A for a methane molecule. The results of the calculation (Fig. 9.2) show that the shape of the potential energy as a function of the distance from the centre of the pore is different for the three zeolites. For ZSM-22 (TON), the minimum is at the centre of the pore while the minima for ZSM-12 (MTW) and UTD-1 (DON) are closer to the wall. In fact, the shape of the curve for DON is quite similar to the potential for a flat wall indicating that the transition from a microporous environment to a mesoporous molecular environment (i.e. one in which interactions resemble a flat surface) starts to be observed even for 14-ring zeolites.
Figure 9.2 Potential energy of a methane molecule in 1-D zeolites ZSM-22, ZSM-12 and UTD-1 ([SiO2]-TON, MTWand DON) as a function of distance from the pore centre.The value of the well depth was obtained from an independent calculation E12/kB¼133.33. The average radii for the three ˚ , respectively, and the zeolites (framework-typesTON, MTWand DON) are 3.825, 4.275 and 5.725 A ˚ 2. average densities of oxygen atoms on the pore surfaces are 0.166, 0.179 and 0.184 O atoms/A
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247
The desired results are predictions of the isosteric heats of adsorption, and these can be obtained from the Boltzmann weighted average energy ―
RR
w ¼ R R0 0
wðrÞewðrÞ=kB T r dr
ðewðrÞ=kB T 1Þr dr
ð9Þ
and by the following equation ―
Hst ¼ NA w þRT
ð10Þ
The comparison of this theory to experiment is very reasonable (Table 9.4). The average difference is 8%, larger than experimental error but small nevertheless, with the highest deviation for ZSM-22. The best agreement is found for UTD-1, a zeolite with pores that are quite uniform—that is, framework oxygen atoms are at the same distance from the pore centre as one moves along the pore direction. The internal surface of the pores of ZSM-22 and ZSM-12, on the contrary, is corrugated and the difference between theory and experiment may reflect this deviation from the model assumptions. The important conclusion is that simple models based on van der Waals interactions can capture semiquantitatively the isosteric heats of adsorption of simple non-polar molecules on zeolite microporores. For large molecules like dodecane and above, van der Waals forces often dominate the interaction molecule–zeolite over other forces, and can be used as a first estimate of the selectivity of adsorption respect to the gas phase. Finally, in Fig. 9.3 the isosteric heat of adsorption at zero coverage for several non-polar molecules is plotted versus the electronic polarizability for zeolite [SiO2]MFI. These data show that there is indeed a very good correlation between the heats of adsorption and the electronic polarizability. For larger molecules, this correlation is expected to break down as London’s model assumes the molecules can be represented by one point, and clearly this is already incorrect even for molecules as small as propane. Nevertheless, the correlation is excellent.
4. Induced Dipole and Quadrupole–Charge Interactions We next consider the interaction of molecules with cations coordinated to the zeolite pore walls. An electric field induces a dipole, uind, on all molecules (polar and non-polar) that is proportional to the polarizability and the strength of the electric field (uind ¼ a E). The size of this induced dipole can be substantial (of the order of 1 Debye), and thus this interaction can be quite important in determining the overall energetics of an adsorbed molecule inside a zeolite pore. For non-polar molecules, the polarizability is composed only of the electronic polarizability term such that a¼a0, but for polar molecules the polarizability contains another term denoted orientational polarizability
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Figure 9.3 Isosteric heats of adsorption on [SiO2]-MFI zeolites versus electronic polarizability. Note the close correlation between molecular properties such as electronic polarizability and observable macroscopic properties such as isosteric heats of adsorption. (Data from Dunne et al.).22,23
that arises from the effect of an external field on the average orientation of the molecule in space. The net ion-induced dipole interaction free energy is given by
ðzeÞ2 a ðzeÞ2 u2 wðrÞ ¼ ¼ a0 þ 3kB T 2ð4pe0 eÞ2 r 4 2ð4pe0 eÞ2 r 4
ð11Þ
where e is the charge of the electron, z is the valence of the cation, u is the permanent dipole of the molecule and e is the dielectric constant of the medium. First, note that this interaction depends on the negative fourth power of r, and thus it is a longer-range interaction than the dispersion forces. Second, inside the zeolite pores it is not obvious what one should use as the value of the dielectric constant e. The dielectric constant of quartz is 4.8, while the one of amorphous silica glass is 3.8, and one may think that a density-scaled dielectric constant could be used as a good approximation.c This is adequate within mesoscopic and larger length scales; however, in the limit of zero coverage, the length scale of interest are distances comparable to the pore diameter or smaller, so a more appropriate value of e inside the zeolite pore is e 1 (i.e. comparable to air or vacuum). At the same time, the weaker dependence on distance makes it doubly important that we understand the role of
c
To the best of our knowledge the dielectric constant of siliceous zeolites has not been measured.
Intermolecular Forces in Zeolite Adsorption and Catalysis
249
other adsorbed molecules can have on this force. If the pore is filled with water molecules (eH2 O ¼ 88 in the bulk), for instance, the magnitude of this interaction will decrease by a factor of 80–90.d A non-uniform distribution of charge within the molecule gives rise to a second type of interaction with an electric field even when a molecule has no permanent dipole. This distribution of charge is called a quadrupole , and the interaction between a quadrupole and the electric field generated by a cation can be very important in zeolites as we will see shortly. In general, the quadrupole of a molecule is a tensor quantity, and its interaction with a charge (a cation) is angle dependent.6 Frequently though, we will be interested in molecules (such as N2 and O2) that have a linear distribution of charge; and in such cases, the quadrupole can be thought of as a two equidirectional but opposed dipoles (þþ). The energy of interaction w(r) between a linear molecule with quadrupole moment and a point charge is given by
wðrÞ ¼
ze ð3 cos 2 y 1Þ 2r 3 ð4pe0 eÞ
ð12Þ
This interaction can be attractive or repulsive depending on molecular orientation ˚ and the interaction is a and the sign of the charge and quadrupole. For N2 Y¼1.39 D A maximum when the molecule axis is aligned along the cation-molecule direction (y ¼ 0; ð3 cos2 y 1Þ ¼ 2), and minimized with a different sign when the molecule is perpendicular to the cation-molecule direction (y ¼ p=2;ð3 cos2 y 1Þ ¼ 1). Since the angle average of hcos2 yi ¼ 1=3, the average quadrupole interaction can be neglected at high temperatures when molecules tumble isotropically in the zeolite pores (i.e. h3 cos 2 y 1i ¼ 0). Cations in a zeolite can be highly exposed to molecules in the pores, can be partially blocked in such a way that only small molecules have direct contact with the cation, or can be coordinated to the zeolite framework such that there is no—or only very weak—interaction with adsorbed molecules. Because of this distribution of cation coordination environments, adsorbed molecules will find a complex landscape of adsorption sites. Moreover, different zeolites will show different zerocoverage heats of adsorption precisely because the distribution of coordination environments is different in each case. Because of the negatively charged zeolite oxygen atoms and due to charge transfer processes to the cation from the framework oxygen atoms, the effective charge of the cation is substantially lower than the charge of an isolated cation (i.e. charge < e). To clarify this point, Table 9.5 shows the isosteric heat of adsorption of cation-containing zeolites compared to their all-silica analogues. We can indeed see that the presence of sodium cations in the pores substantially increases the heat of adsorption of the molecules. Figure 9.4 compares the sodium cation affinity of several of the molecules in the table to the
d
The local dielectric constant of a medium is dependent on the structure of the solvent and inside a micropore such structure is bound to be quite different from the bulk structure of the solvent. The effective dielectric constant will be different from 88, although it is not possible to evaluate the local dielectric constant without recourse to atomistic models of the solvent in the pores. It should be dependent on pore size and zeolite composition.
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Table 9.5 Isosteric heats of adsorption of small molecules on MFI-type and FAU-type zeolitesa Hst (kJ/mo) [SiO2]Adsorbate MFI
Ar CH4 SF6 O2 N2 C2H6 CO2 a b
15.8 20.9 34.4 16.3 17.6 31,1 27.2
jHþj [SiAlO2]MFI
35.2 20.7 33.3 38.0
jNaþj [SiAlO2]MFI
18.0 26.5 42.0 17b 24.1 38.0 50.0
[SiO2]FAU
11.0
19.7
jNaþj [SiAlO2]-FAU (NaX)
12.7 19.2 28.2 15.0 19.9 27.0 49.1
Data from Dunne and Savitz et al.7,22–24 This datum estimated from figure in ref. 7.
Figure 9.4 Gas-phase sodium cation affinity versus excess isosteric heat above the value of the siliceous zeolite.The data show excellent correlation between the two sets of numbers indicating that the interaction of small molecules with cations inside zeolite pores resemble energetically (and probably structurally) the molecule-cation adducts formed in the gas phase. Sodium cation affinity data from NIST Chemistry Webook and from Table 9.5.
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251
excess isosteric heat of adsorption due to the presence of cations on the zeolite (the difference between the sodium-cation containing zeolite and its all-silica analogue). We can clearly observe a very good trend between the two quantities. Within some scatter, the trend is linear, (although notice that the excess heat for methane in NaZSM-5 and in Na-X (FAU) differ by more than 3 kJ/mol). This linear trend is an indication that the geometry of the adsorbed molecule on the cation, and type of interaction in the gas phase and in the zeolite pores are quite similar. However, the magnitude of the excess heat of adsorption is substantially lower than the DHr of the interaction of the molecule with the cation in the gas phase. This large difference is the consequence of the coordination of the negatively charged framework to the cation, that is, the effective weaker charge of the cation. To illustrate this further, let’s consider a specific example of a cation coordinated to an AlO4/2 tetrahedron on the surface of a high-silica zeolite. For simplicity, we will consider an idealized model of an exposed cation as depicted in Fig. 9.5. This
Figure 9.5 Schematic of the coordination of sodium cations to a model AlO4/2 tetrahedron. Distances are not to scale. Since this site has a threefold axis of symmetry, the Al^O bond distance is 1.73 A˚ and the angle between the O^Al^O atoms is 109.47 (the tetrahedral angle), all distances can be calculated using elementary geometry.
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figure shows the sodium cation coordinated to three of the oxygen atoms of the alumina tetrahedron. The interaction energy at contact between a methane molecule and the sodium cation is given by the formulae above. For a NaþCH4 distance of 2.79 A˚ and an effective charge of þ1 e, w(r)¼27.8 kJ/ mol, much larger than the excess isosteric heat of 8.2 kJ/mol. This is of course because the effective charge on the cation is smaller due to the effect of the negatively charged oxygen atoms. Using 8.2 kJ/mol as the benchmark, the effective charge can be back calculated to be 0.54 e. Using this effective charge, we can then calculate the total interaction for N2, a molecule with a ˚ between Naþ and substantial quadrupole moment. Using a distance of 3.17 A the centre of mass of N2, and assuming N2 is aligned along the molecule-cation axis, the induced dipole and quadrupole interactions are 3.4 and 6.6 kJ/mol, respectively. The total 10 kJ/mol is very much in agreement with the excess isosteric heat between silicalite ([SiO2]-MFI) and NaZSM-5 (|Naþ|[SiAlO2]MFI) of 9 kJ/mol (Table 9.5).
5. Dipole–Charge Interactions The interaction of polar molecules—molecules with permanent dipole moments—with cations on the surface of a zeolite pore is the last type of electrostatic interaction that will be reviewed here. To start, let’s recall that the dipole moment of a polar molecule u is defined as u¼q l where l is the distance between two charges þq and q. Small molecules have dipole moment of the order of 1 Debye ¼ 3.3361030 C m. The energy of interaction between an ion and a dipole is given by
wðrÞ ¼
ðzeÞucosy 4pe0 er 2
ð13Þ
where the symbols are as defined before, and y is the angle between the direction of the dipole moment u and the line joining the centre of mass of the molecule and the ion. This interaction can be attractive or repulsive depending on the angle y. It is maximum when the angle is 0 . In practice, the interactions between an ion and a polar molecule can be divided into strong and weak depending on whether the interaction at y ¼ 0 is large compared to kBT. In such a case, the molecule will tend to be oriented towards the ion and would be coordinated to it for long periods of time. This is expected for small and multivalent cations interacting with small molecules containing a sizable dipole moment. For this case the energy is given by
wðrÞ ¼
ðzeÞu 4pe0 er 2
ð14Þ
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Intermolecular Forces in Zeolite Adsorption and Catalysis
Table 9.6 Comparison of calculated electrostatic interactions of CO and water molecules with sodium cations in the gas phase to isosteric heat of adsorption in Na–ZSM-5 Interaction
CO kJ/mol
H2O kJ/mol
Dipole–charge Quadrupole–charge Induced dipole Total DHst Silicalite-1 DHst Na–ZSM-5 DDHst
–3.65 –32.0 –17.9 –53.5 17 337 16
–96.7 –2.9 –34.1 –134 30 9525 65
For weak interactions, one can assume that the molecule tumbles with isotropic motion in the vicinity of the ion and one requires the angle-averaged charge–dipole interaction given by1
wðrÞ
q2 u2 6ð4pe0 eÞ2 kB Tr 4
for kB T >
qu 4pe0 er 2
ð15Þ
For instance, the heat of adsorption of CO on NaZSM-5 (33 1 kJ/mol at 195 K) and Silicalite-1 (17 1 kJ/mol)7 can be compared to the enthalpy of reaction of CO and Naþ(to form CO Naþ) in the gas phase (DHr ¼ 52.7 kJ/mol).5,e The difference between the two zeolites reflects the additional contribution of the sodium cations to the heat of adsorption, that is, 16 kJ/mol or about 30% of the gas-phase value. This is comparable to what we observed when we estimated the effective charge of sodium cations using the nitrogen molecule. This heat of adsorption on sodium cations has two contributions: one from the dipole–charge interaction, and another from the induced dipole moment on CO by the electric field of the cation. Using the following data (rCO ¼ 0.2 nm, rNa ¼ 0.095 nm, q ¼ ze ¼ 1.6 1019 C, a0 ¼ 1.95 1.11 10–40 C2 m2 J1, T ¼ 300 K, kB ¼ 1.381023 J K1, dipole moment of CO ¼ 0.113.3361030 C m, ¼ 9.47 10–40 C m2), the magnitude of the various interactions can be determined (Table 9.6). Water is an example of a small molecule with a large dipole moment. Water, in its simplest form, can be treated as a spherical molecule of 0.14 nm in radius. It has a dipole moment u ¼ 1.85 D and it can adsorb into a sodium cation (0.095 nm in radius). Calculation of the maximum interaction energy for this complex in the gas phase gives a total of w(r) ¼ 134 kJ/mol, including dipole, polarizability and quadrupole interactions. The dipole-charge contribution is by far the most important one. (Note that this value is higher than the value of 95 7.9 kJ/mol that has been measured experimentally for the reaction Naþ þ H2O ¼ Na H2Oþ in the gas phase.)8 The heat of adsorption of water in NaZSM-5 and NaA zeolites has
e
Values of 32 7.9 kJ/mol have also been reported.
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Raul F. Lobo
been measured by Dubinin and co-workers9 giving values in the range of 95–105 kJ/mol. This energy of adsorption is dominated by the dipole–charge interaction (see Table 9.6). For the zeolites, the contribution of the sodium cation corresponds to 65 kJ/mol, somewhat higher than the expected value on the basis of the partial charge of the sodium cation on the zeolite pores. This is the case because the water molecule can coordinate by hydrogen bonding to the oxygen atoms of the zeolite, adding further to the heat of adsorption.
6. Interactions of Adsorbed Molecules with Acid Sites The interaction of molecules with Brnsted acid sites in zeolites deserves a separate discussion because the binding of protons to the zeolite framework is very different from the binding of cations like sodium and potassium. Likewise, the coordination of small molecules to acid sites can be quite different from the coordination to alkali-metal cations. Figure 9.6 illustrates the structure and geometry of the acid sites in zeolites as has been determined for high-silica zeolites.10 The key element of the site is the presence of an Al–OH–Si group. The bond between the hydrogen atom and the zeolite oxygen atom is predominantly covalent in nature, and there is only a small effective charge on the hydrogen atom. This site can coordinate or donate the proton to adsorbed molecules forming 1:1, 2:1 and
Figure 9.6 Schematic of the structure and geometry of the classical Brnsted acid site in highsilica zeolites. The structures on the right illustrate the coordination of weak and strong bases to the acid site. For strong bases, there is proton transfer and charge separation between the framework and the adsorbed molecule.
Intermolecular Forces in Zeolite Adsorption and Catalysis
255
sometimes higher adducts. The oxygen atom in the acid site is tri-coordinated, and since it should keep a bond order of 2, there is a weakening of the three bonds, but particularly the O–H bond; this allows for facile transfer of the proton to adsorbed molecules. The top section of Fig. 9.6 also highlights the presence of three additional oxygen atoms coordinated to the aluminum that have Lewis basic character and that can coordinate through hydrogen bonding to molecules that exhibit both hydrogen-acceptor and hydrogen-donating groups, such as water or ammonia. Usually, two out of the three oxygen atoms—but sometimes only one—are physically accessible to the adsorbed molecules due to the specific configuration of the framework atoms with respect to the zeolite pore (see Fig. 9.5). These geometric differences appear to have no major thermochemical consequence11 and the heat of adsorption of strong bases on high-silica zeolites is the same—within experimental error—regardless of the specific geometry of the site. Figure 9.6 also shows the various bond lengths of the acid site. Note that the ˚ vs 1.73 A ˚ ); Al–OH bond is substantially longer than the other three bonds (1.92 A and as a consequence the coordination environment of the aluminum atom is not tetrahedral (as is the case for framework aluminum coordinated to alkali-metal cations), but an intermediate structure between tetrahedral and planar trigonal. The long aluminum–oxygen bond reflects, of course, the fact that this bond is weaker than the other three. Because of the elongation and weakening of the O– H bond, there is a change in the infrared (IR) absorption frequency from 3749 cm1 (for isolated Si–OH groups) to 3614 cm1 (Al–OH–Si).12 This change in vibrational frequency is also an indication that the ability to donate the proton of a zeolite Brnsted acid site is much larger than the ability of an isolated silanol group. There are two extremes of interaction of a zeolite acid site ZOH with a base B. For weak bases the adduct forms a hydrogen bond to the hydroxyl group (ZOH B), while for strong bases there is complete proton transfer and charge separation (ZO HBþ). And there are, of course, intermediate cases that fall between these two cases. A series of calorimetric, computational and spectroscopic investigations have shown that the molecular property that correlates more closely to the observed thermochemical and spectroscopic measurements is the proton affinity (PA) of the base. If the PA is below 858 kJ/mol for the 1:1 adduct, there is no proton transfer and the complex is described more closely as ZOH B. If the PA is above 858 kJ/mol, there is proton transfer and the complex is better described as ZO HBþ. Ammonia is right at this limit with a PA ¼ 858 kJ/mol, and definitely forms the ammonium ion upon adsorption on acid zeolites.13 Figure 9.7 shows the IR spectra of 1:1 adducts on zeolite beta at 1:1 coverage for bases of different PA. The perturbations of the IR spectra increase as the PA increases (see Table 9.7) up to pyridine, where there is proton transfer as confirmed by the observation of the pyridinium ion signatures in the IR spectrum. As illustrated in Fig. 9.6, the adsorption of weak bases slightly perturbs, but not substantially changes the geometry of the acid site. This weak hydrogen bond interaction is mainly of electrostatic character. It is weak because the effective charge of the hydrogen atom in the ZOH group is small. This is evident from Table 9.5 where we can observe that the excess isosteric heat for HZSM-5 versus NaZSM-5 is small, regardless of the small effective radius of the hydrogen atom.
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Figure 9.7 Background subtracted infrared (IR) spectra of bases of different proton affinity (PA). Nitrogen and CO conducted at 100 K, C3H6 at 180 K and the rest at room temperature.The perturbation of the IR spectra of clean H-beta increases with PA. The spectra become increasingly more complex due to mixing and doubling of bending and stretching modes. Reproduced with permission from Paze et al.12
The abstraction of the proton by a strong base, in contrast, changes the geometry of the aluminum from a distorted tetrahedron to a symmetric tetrahedron where all ˚ . For strong bases, both the differential heat Al–O bonds are similar and about 1.75 A of adsorption and the IR spectra of the adducts are independent of the zeolite structure, indicating that the complexes formed by a molecule on all zeolites are very similar (provided there is space to form the complex). Molecules with hydrogen accepting and donating groups are somewhat different because they can form 1:1, 2:1, and higher order complexes with acid sites in the zeolite pores (see Fig. 9.8). At low coverage (based on the Brnsted acid site ratio), IR spectra indicate that the 1:1 complex dominates for water and methanol, but even before the B:ZOH site ratios approach 1:1, evidence for clusters of molecules
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Table 9.7 Comparison of adsorption enthalpies in H–ZSM-5 and silicalite with gas-phase proton affinities
Molecule
Proton affinity (kJ/mol)
H^ZSM-5 DHads (kJ/mol)
Silicalite DHads (kJ/mol)
N2 CO C2H4 H2O CH3CN CH3OH (C2H5)2O NH3 N(CH3)3 Pyridine
494 594 680 724 798 774 838 858 938 922
20.7 26
17.6 17
90 110 115 135 145 205 200
30 75 65 70
Figure 9.8 Coordination of 1:1and 2:1water^zeolite complexes on acid sites.
are observed. For water and methanol above low coverages, 2:1 and higher order adducts are clearly observed, and here proton transfer is occurring, as may have been guessed from PA data. For instance, the gas phase PA of H2O is 723 kJ/mol (weak), but the PA of (H2O)2 is 850 kJ/mol, similar to ammonia (strong). Calorimetric investigations indicate that there is dynamic equilibrium between the 1:1 and 2:1 adducts (again, for water and methanol) at room temperature, even before the total coverage reaches 1:1 B:ZOH site ratios.14,15 The interaction of alkenes with acid sides is briefly discussed next because of its particular importance in petrochemical catalytic processes.16 The precise nature of this interaction remains a topic of ongoing research,17 yet important insights have been gained that start to clarify how organic molecules bind to zeolites. Alkenes are different from the two cases discussed above because there can be proton transfer to the alkene, yet a neutral species could be the result of this chemisorption process. The main question is then whether alkenes form silil-alkoxide species
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Figure 9.9 Comparison of the structure alkoxide complexes and stable carbocations formed on zeolite acid sites. The left side shows the structure formed by alkenes reacting with zeolite Brnsted acid sites.The right side shows an example of a stable carbocation formed by protonation of dimethyl-cyclopentadiene.
upon reaction with an acid site, or is there charge separation to form carbenium ions (Fig. 9.9). The evidence indicates that it depends on the structure of the organic molecule.18 Ethene is a weak base (PA ¼ 680 kJ) and at low temperature it forms a hydrogen-bonded complex with acid sites in zeolites, but upon reaction it is predicted to form the neutral ethoxide group, covalently bonded to the zeolite framework. The complex formed by propene has been difficult to establish because it rapidly forms oligomeric species.19 The tert-butyl alkoxide (from isobutene, PA ¼ 802 kJ/mol) has been isolated in some zeolites, but there are indications that steric hindrances can shift the equilibrium towards the tert-butyl cation. In summary, the bulk of the evidence indicates that primary and secondary alkenes will form covalent (neutral) complexes with zeolite acid sites. Tertiary alkenes can form covalent or charged complexes, depending on the specific structure of the alkene and the zeolite. In the last decade, there has been accumulating evidence demonstrating that organic molecules with very high PAs (such as 1,3-dimethylcyclopentadiene, PA ¼ 878 kJ/mol) do indeed form stable ions inside the zeolite pores.10,17,20,21 These stable carbocations can (Fig. 9.9), in fact, be very important in the conversion of hydrocarbons such as the methanol to gasoline and methanol to olefins processes. In addition to 1,3-dimethylcyclopentadiene, various stable carbocations have been identified inside the zeolite pores using nuclear magnetic resonance (NMR) spectroscopy, UV/vis spectroscopy and other techniques. Summarizing, the interaction of small molecules with zeolite acid sites is fundamentally a chemical interaction that can be broadly described as Lewis-base assisted, Brnsted acid–base interaction. Because it is a chemical interaction, often chemical bonds are made and broken and the end result of this process depends on the specific structure of the adsorbate. All acid sites in high-silica zeolites appear to be structurally very similar and the complexes they form with small molecules are, in thermochemical and spectroscopic terms, similar to each other (within a zeolite structure and between different zeolite structures as well). The gas-phase PA of the adsorbed molecule is the molecular property that gives the most useful information towards predicting the nature of the adduct formed upon adsorption.
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7. Hydrophobic Interactions in High-Silica Zeolites Water is unique because of its ability to form three-dimensional networks of hydrogen bonds, and this specific interaction leads to a remarkably high enthalpy of vapourization (DHvap ¼ 43.0 kJ/mol at 298 K) as compared to molecules of similar size like methane (DHvap ¼ 8.5 J/mol at 99 K). This is despite the fact that water molecules have a rather low electronic polarizability, and thus have relatively weak dispersion forces with other molecules. For instance, water has a polarizability of 1.5 ˚ 3 as compared to methane that has a polarizability of 2.5 of 4pe0 A ˚ 3 (see 4pe0 A Table 9.1). Inside a purely siliceous zeolite, at low coverage, the only force that governs the interaction of water molecules with the zeolite framework is the dispersion or London force. Because of its low polarizability, this force is low in magnitude, and it is especially low compared to its high enthalpy of vapourization. For comparison, methane’s heat of adsorption in siliceous MFI (silicalite) is 20.9 kJ/mol, much higher than its heat of vapourization, yet the heat of adsorption of water is smaller than the heat of adsorption of methane. For this reason, high-silica zeolites can be used to selectively adsorb hydrocarbons even in the presence of high vapour pressures of water, a property that it is nearly unique among readily available industrial sorbents. This is very different from high-surface-area porous silicas that contain a surface covered with silanol groups and are hydrophilic (in the absence of surface treatments). In fact, purely siliceous zeolites (without defects such as internal silanol groups or tetrahedral vacancies) belong to a class of materials called hydrophobic solids. In these materials, water condensation occurs at a pressure above the saturation vapour pressure. For instance, the condensation of water in defect-free silicalite-1 is about 100 MPa at room temperature. For zeolites with larger pores, the condensation pressure is still high: for siliceous zeolite beta (with a three-dimensional-12-ring pore system) it is above 50 MPa at room temperature. This additional pressure can be thought to be the work that is needed to accommodate the massive disruption of the hydrogen bond network that is unavoidable upon the incorporation of water molecules into a microporous, hydrogen-bond-free environment such as the one presented in the pores of siliceous zeolites.
8. Final Remarks This chapter has presented an overview of the important classes of zeolite– molecule forces that determine the adsorption and some catalytic properties of zeolite materials. Our intention is to relate molecular properties of the adsorbates to measurable quantities such as heats of adsorption. The data reveal that qualitatively, this approach helps to understand differences in the adsorption of small molecules for multiple classes of zeolite structures and compositions. At the same
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time, often differences are observed between the predictions of the simple models used here and the actual measurements of heat of adsorption. These reveal the limitations of this simplistic approach, yet also point to aspects of zeolite–molecule interactions that are not well understood. In particular, further calorimetric measurements of hydrogen-bond accepting and donating groups are needed to clarify the separate contribution of OH bonds and electrostatic forces.
REFERENCES [1] Israelachvili, J. N., Intermolecular and Surface Forces. 2nd ed. Academic Press: London, 1991; p 450. [2] Baerlocher, C., Meier, W. M., Olson, D. H., Atlas of Zeolite Framework Types. 5th ed.; Elsevier: Amsterdam, 2001. [3] Baerlocher, C., McCusker, L. B. Database of Zeolite Structures. http://www.iza-structure.org/ databases/ [4] Savitz, S., Siperstein, F., Gorte, R. J., Myers, A. L. ‘‘Calorimetric study of adsorption of alkanes in high-silica zeolites’’, J. Phys. Chem. B 1998, 102, (35), 6865–6872. [5] Everett, D. H., Powl, J. C., ‘‘Adsorption in slit-like and cylindrical micropores in Henry’s Law region: Medol for microporosity of carbons’’, J. Chem. Soc. Faraday Trans. I 1976, 72, 619–636. [6] Papai, I., Goursot, A., Fajula, F., Plee, D., Weber, J., ‘‘Modeling of N-2, and O-2 Adsorption in Zeolites’’, J. Phys. Chem. 1995, 99, (34), 12925–12932. [7] Savitz, S., Myers, A. L., Gorte, R. J., ‘‘A calorimetric investigation of CO, N-2, and O-2 in alkali-exchanged MFT’’, Micropor. Mesopor. Mater. 2000, 37, (1-2), 33–40. [8] NIST Chemistry Web book. http://webbook.nist.gov/chemistry/ [9] Dubinin, M. M., Rakhmatkariev, G. U., Isirikyan, A. A., ‘‘Differetial heats of adsorption of water vapor on NaA zeolite at 300 and 450 K’’, Izest. Ak. Nauk SSSR, S. Khim. 1989, 12, 2877–2778. [10] Beck, L. W., Xu, T., Nicholas, J. B., Haw, J. F., ‘‘Kinetic Nmr And Density-Functional Study Of Benzene H/D Exchange In Zeolites, The Most Simple Aromatic-Substitution’’, J. Am. Chem. Soc. 1995, 117, 11594–11595. [11] Savitz, S., Myers, A. L., Gorte, R. J., White, D., ‘‘Does the cal-ad method distinguish differences in the acid sites of H-MFI?’’, J. Phys. Chem. B 1998, 120, (23), 5701–5703. [12] Paze, C., Bordiga, S., Lamberti, C., Salvalaggio, M., Zecchina, A., Bellussi, G., ‘‘Acidic properties of H-beta zeolite as probed by bases with proton affinity in the 118-204 kcal mol (-1) range: A FTIR investigation’’, J. Phys. Chem. B 1997, 101, (24), 4740–4751. [13] Zecchina, A., Marchese, L., Bordiga, S., Paze, C., Gianotti, E., ‘‘Vibrational spectroscopy of NH4+ ions in zeolitic materials: An IR study’’, J. Phys. Chem. B 1997, 101, (48), 10128–10135. [14] Haase, F., Sauer, J., ‘‘Interaction Of Methanol With Bronsted Acid Sites Of Zeolite Catalysts An Ab-Initio Study’’, J. Am. Chem. Soc. 1995, 117, (13), 3780–3789. [15] Haase, F., Sauer, J., ‘‘Ab initio molecular dynamics simulation of methanol interacting with acidic zeolites of different framework structure’’, Micropor. Mesopor. Mater. 2000, 35-6, 379–385. [16] Haw, J. F., Song, W. G., Marcus, D. M., Nicholas, J. B., ‘‘The mechanism of methanol to hydrocarbon catalysis’’, Acc. Chem. Res. 2003, 36, (5), 317–326. [17] Clark, L. A., Sierka, M., Sauer, J., ‘‘Stable mechanistically-relevant aromatic-based carbenium ions in zeolite catalysts’’, J. Am. Chem. Soc. 2003, 125, (8), 2136–2141. [18] Nicholas, J. B., Haw, J. F., ‘‘The prediction of persistent carbenium ions in zeolites’’, J. Am. Chem. Soc. 1998, 120, (45), 11804–11805. [19] Gorte, R. J., White, D., ‘‘Interactions of chemical species with acid sites in zeolites’’., Topics Catal. 1997, 4, (1-2), 57–69.
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[20] Boronat, M., Viruela, P. M., Corma, A., ‘‘Reaction intermediates in acid catalysis by zeolites: Prediction of the relative tendency to form alkoxides or carbocations as a function of hydrocarbon nature and active site structure’’, J. Am. Chem. Soc. 2004, 126, (10), 3300–3309. [21] Song, W. G., Nicholas, J. B., Haw, J. F.,‘‘Acid-base chemistry of a carbenium ion in a zeolite under equilibrium conditions: Verification of a theoretical explanation of carbenium ion stability’’, J. Am. Chem. Soc. 2001, 123, (1), 121–129. [22] Dunne, J. A., Mariwals, R., Rao, M., Sircar, S., Gorte, R. J., Myers, A. L., ‘‘Calorimetric heats of adsorption and adsorption isotherms .1. O-2, N-2, Ar, CO2, CH4, C2H6 and SF6 on silicalite’’, Langmuir 1996, 12, (24), 5888–5895. [23] Dunne, J. A., Rao, M., Sircar, S., Gorte, R. J., Myers, A. L., ‘‘Calorimetric heats of adsorption and adsorption isotherms .2. O-2, N-2, Ar, CO2, CH4, C2H6, and SF6 on NaX, H-ZSM-5, and Na-ZSM-5 zeolites’’, Langmuir 1996, 12, (24), 5896–5904. [24] Savitz, S., Myers, A. L., Gorte, R. J., ‘‘Calorimetric investigation of CO and N-2 for characterization of acidity in zeolite H-MFI’’, J. Phys. Chem. B 1999, 103, (18), 3687–3690. [25] Dubinin, M. M., Rakhmatkariev, G. U., Isirikyan, A. A., ‘‘Energy of adsorption of water vapor adsorption on high and pure-silica zeolites’’, Izest. Ak. Nauk USSR, S. Khim. 1989, 12, 2862–2864.