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156 Chapter 3

Figure 3.59. The molecular orbital energies for the BeH2 molecule. 2δ is the di erence in energy of the bonding and antibonding σ-orbital pairs; δ is the di erence in energy of the two 2pz , and 2s, based bonding and antibonding orbitals, respectively.

One constructs hybridized atomic orbitals on Be by the combination of atomic orbitals along the symmetry axis of a molecule. In our example for Be:

They are illustrated in Fig. 3.60., where ϕ+ is an orbital oriented to the right and ϕ− is an orbital directed to the left. The energies of those orbitals are:

Figure 3.60. As can be seen, ϕ+ is oriented to the right and ϕ− to the left. The hybridized orbitals ϕ+ and ϕ− are orthogonal.

When the ideal hybridization condition is satisfied, β+− = 0. The hybridized Beϕ+ orbital has the following overlap energies with hydrogen atoms a and b, located respectively right and left from Be in BeH2:

A similar relation then holds between hybridized orbital ϕ− and the hydrogen atomic orbital on the atom at the left hand position of Be.

Figure 3.61. Ideal hybridization molecular orbital interaction scheme in BeH2.

The ϕ+ orbital interacts with the hydrogen atom to the right of the Be atom, the ϕ− orbital interacts with the hydrogen atom at the left of the Be atom. The di erence between the resulting degenerate bonding and antibonding orbitals (see Fig. 3.61) becomes in the

As long as ∆δ 1 (see Fig. 3.59), the hybridization model can be considered a satisfactory description of the chemical bond.

The condition for the approximate validity of this description is that the overlap energies of the 2s and 2p atomic orbitals with their neighboring atoms are approximately the same and these overlap energies are large compared with the Be atomic orbital energy di erence of 2s and 2pz .

158 Chapter 3

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CHAPTER 4

Shape-Selective Microporous Catalysts, the Zeolites

4.1 Zeolite Catalysis, an Introduction

4.1.1 Zeolite Structural Features

Zeolites are crystalline aluminosilicates which assemble into well-defined three-dimensio- nal structures comprised of microporous channels that interconnnect cavities which approach molecular scale dimensions ranging from 2 to 12 ˚A. Mesoporous silica materials with pore sizes over 40 nm have also been made. They are relatively easy to synthesize and o er outstanding control over the pore size as well as its three-dimensional pore architecture, which makes them ideal for gas separations and shape-selective catalysis. Zeolites are perhaps the most widely recognized catalytic material. They are found in nature in over 50 distinct mineral forms. In addition, over 140 man-made zeolites have also been synthesized, with a total of 165 di erent framework structures[1] .

In addition to their microporous environment, zeolites o er a range of other important chemical properties. They are highly acidic materials. One can tune both the number of acid sites and the strength of these sites and thus control their overall acidity. They have high thermal stabilities. Metal ions can be readily exchanged, impregnated or added into the synthesis to open up many other types of reactions.

The primary building units of the zeolite are tetrahedral (MO4) structures typically comprised of a silicon or aluminum atom that sits at the center of the tetrahedron and four oxygen atoms which sit at the vertices of the tetrahedron. These oxygen atoms interconnect tetrahedra forming an interconnected network of three-dimensional channels. Each tetrahedron is bound to the others through the oxygen atoms at their vertices. The tetrahedra can subsequently assemble and form di erent secondary building units (16 SBUs are currently known) or chain-like structures which can go on to form 6, 8, 10 and 12 rings. These rings can be considered to assemble into two-dimensional structures that go on to form three-dimensional cavities or cages bounded by well-defined 6-, 8- , 10and 12-ring apertures. The cages are connected via di erent SBUs to comprise the tertiary zeolite framework. The porous network can form either straight or zig-zag pore structures depending on how the tertiary structure assembles. Two well-known, catalytically important zeolite structures mordenite and chabazite are shown in Fig. 4.1.

Zeolites have wide industrial uses covering a range of di erent commercial processes including: catalytic cracking of gas oil to gasoline, hydrocracking gas oil to kerosene, dewaxing middle distillates to lubricants, benzene alkylation to styrene, toluene disproportionation to xylenes, xylene isomerization, methanol-to-gasoline conversion, methanol-to- alkenes, halogenations and nitrations of arenes, isomerization of di erent hydrocarbons, hydration and dehydrations of alcohols and acids, hydrogenation and dehydrogenations of hydrocarbons, hydroformylation, and oxidation, to name just a few.

In this chapter, we focus solely on the catalysis of zeolitic systems. Zeolite synthesis is discussed in Chapter 8 and mesoporous systems are discussed in sections dealing with biomineralization in Chapter 9.

The purely silica framework (SiO2) alone is charge neutral. The substitution of Si4+, however, with Fe3+, B3+ or Al3+ will impart a negative charge on the framework oxygen atoms. Protons attached to the framework or positively charged cations located in the cavities or zeolite channels subsequently act to maintain overall charge neutrality.

This framework charge is compensated by positively charged cations positioned in

Molecular Heterogeneous Catalysis. Rutger Anthony van Santen and Matthew Neurock

Copyright © 2006 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim

ISBN: 3-527-29662-X

162 Chapter 4

Figure 4.1a. (Lefthand side) Ball-and-stick representation of mordenite. (Righthand side) Space filling model of chabazite. Note the di erence in size of oxygen and silicon atoms. Ethane is adsorbed in two of the cavities.

the zeolite channels and cavities. Ammonium ions can be used to compensate for the negative charge on the framework lattice. Ammonium ions, however, thermally decompose at higher temperatures, leaving behind a proton which binds to the framework oxygen. This subsequently leads to high Brønsted acidity at the framework oxygen sites. The concentration, type and size of the cations contained in the micropores ultimately control the zeolite’s catalytic activity. For example, Zn2+ or Fe2+ can act as Lewis acid or redox centers when used as charge-compensating cations of negatively charged zeolite lattices. Another class of zeolites contain reducible cations such as Co2+ (as in AlPO4), or Fe3+ and Ti4+ in the center of the lattice tetrahedra. As such, these materials are useful for oxidation catalysis.

The AlPO4 polymorphs are quite similar to the zeolites. The framework for the standard AlPO4 is charge neutral. The substitution of Al3+ with Zn2+, Co2+ or Ni2+, on the other hand, can lead to a negatively charged framework. Much of what we have described for the zeolites holds for the AlPO4 polymorph systems also.

In this section, we discuss the general aspects of chemical bonding in zeolites and the zeolite O–H bond. Brønsted and Lewis acid catalysis by zeolites is presented in Section 4.2. Section 4.3 covers redox catalysis by zeolites. The final three sections describe the catalytic cycle and the role of adsorption and di usion on catalytic performance. An important question that arises in each of these sections is the relation between the micropore structure of the zeolite and its activity and selectivity.

A fundamental understanding of the nature of chemical bonds within the zeolite and between the zeolite and adsorbate molecules is necessary in order to provide a more comprehensive picture of the molecular aspects that control zeolite catalysis. We describe both the structure and the electronic features that control bonding. We start by first describing the purely siliceous systems and then compare these with the results from cation-exchanged systems, exploring proton-induced acidity and the addition of metal cations for Lewis acidity and redox chemistry. The chemical bonding in the siliceous zeolitic polymorphs of silica is largely covalent. The electrostatic contribution to the chemical bond energy is only about 10%[2].

Figure 4.2. The structure of zeolites;(a) Brønsted acidic site. (b) HZSM-5 zeolite with Si/Al = 45; (1) 1H MAS NMR spectrum of OH groups; (2) DRIFT spectrum, adapted from Kazansky et al.[5a].

The dielectric constant of these materials is not sensitive to the micropore size. The siliceous part of the zeolite framework is hydrophobic and creates an apolar environment for adsorbed reactants. The chemical reactivity within the zeolite more closely resembles the reactivity in a vacuum than in a solution. Interestingly, the hydrophobic part of enzymes, as discussed in Chapter 7, has a dielectric constant comparable to that of the zeolite framework suggesting that zeolites, in principle, may be able to carry out similar reaction processes to those in the enzyme.

As one would expect, the local structure and bonding within zeolite are similar to those of silica. The potential energy surface of Si–O–Si angle bond deformation is rather flat. For instance, a 10◦ change in angle changes the energy only by a few kJ/mol. The energy needed to deform the tetrahedral configuration around Si is substantially greater since the rehybridization energy of the covalent SiO tetrahedral bond is quite large. In addition, a significant amount of energy is required to stretch the Si–O bond.

Structural changes of zeolites can fairly easily be accommodated by small changes in the Si–O–Si bond angles. This is the reason for the minor di erences in the heats of formation of di erent siliceous polymorphs of the zeolites. The di erence in energy between high-density and low-density zeolites usually does not exceed 15 kJ/mol per site. This is another indicator that the long-range electrostatic interactions result in only minor contributions to the chemical bonding in these systems.

In contrast to the purely siliceous framework bonds, the interaction between the negatively charged framework and the charge-compensating cations, such as Na+ and Zn2+, is primarily electrostatic and therefore quite strong[3]. Protons are covalently bound to the oxygen atoms that bridge the lattice Si and Al atoms (Fig. 4.2a). The charge on the proton is essentially zero. The proton–oxygen bond is therefore quite strong and predominantly covalent. Bonding between the proton and the oxygen changes the hybridization of the oxygen atom to which it is bond.

In classical chemical bonding theory, the hybridization on oxygen would change from approximately sp, when the Si–O–Al angle is nearly linear, to approximately sp2, when oxygen is three coordinated. This is in line with the decrease in the average Si–O–Si angle

164 Chapter 4

of 144◦ in quartz to a smaller Si–O–Al angle of about 120◦ in the protonated system. This leads to the formation of a strong O–H bond. The energy for the heterolytic cleavage of the O–H bond is about 1250 kJ/mol typically.

The bond between the zeolite and the proton can be considered an internal silanol (SiOH) group which interacts with a framework Al3+ cation through the oxygen atom. The SiO–H bond weakens when the silanol oxygen atom connects with the Al atom. The change in hybridization of the silanol oxygen from approximately sp to sp2 implies a decrease in the s-character of the chemical bond. The corresponding weakening of the OH bonds arises from the fact that the 2s states are lower in energy than the 2p states. The OH bond of the [Si–OH–Al] unit is therefore weakened compared with that of the silanol. The increased polarizability of the zeolitic proton compared with that for silanol is nicely illustrated by the greater infrared adsorption intensities of the zeolitic proton compared with that for a free surface silanol group.

For instance, the results from NMR studies on HZSM-5, as seen in Fig. 4.2b(1), show a dominance of surface silanol groups compared with acidic protons. The results shown in Fig. 4.2b(2), however, show a much larger infrared adsorption intensity, which should be assigned to the acidic protons the low-frequency peak in Fig. 4.2b(2). Normalizing these intensities per proton results in an 8-fold increase in the infrared intensity of the zeolitic proton compared with that of silanol. Since the infrared intensity is proportional to the polarizability, one concludes that there is a larger polarizability of the zeolitic proton. Hence it is easier to induce a positive charge on the zeolitic bonded proton than that on silanol when it interacts with a Lewis basic molecule. The H–O bond energy, which is approximately 1250 kJ/mol, is rather strong.

When the oxygen–proton bond is broken within a catalytic reaction, the metal–oxygen bond gains more s character, the Si–O and Al–O bond energies therefore increase and the irrespective distances decrease. Consequently, the e ective volume of the [Al–O–Si] unit is smaller than that of the [Si–OH–Al] unit. The resulting stress on the neighboring atoms is partially reduced by changes in their bond distances and angles. This results in

small di erences in local acidity in a zeolite due to di erences in the local compressibility of the structure[4] .

The di erences in proton bond cleavage energies in zeolites are also related to the framework composition. The OH bond energies at Al/Si ratios that are greater than 0.1 are 10–40 kJ/mol higher than those at Al/Si ratios that are smaller than 0.1. The di erences in energy relate to di erences in the local lattice-relaxation energies and also electronic relaxation e ects when the proton–oxygen bond is cleaved. At high Al concentrations, the e ective negative charge excess increases and, thus, the proton interaction energy increases. This increased OH bond energy corresponds to a weaker acid site for catalysis.

The strong electrostatic interaction with micropore cations will induce significant local structural changes. This local lattice relaxation in which the Si–O–Si and Si–O–Al angles are altered at small energy cost controls significantly the di erences in the relative energies of cations adsorbed in di erent exchange locations [5b].

For non-di usion-limited reactions carried out in low Al/Si ratio systems, the overall rate for a proton-catalyzed reaction increases linearly with the proton concentration, as illustrated schematically in Fig. 4.4[6]. The rate, when normalized against the framework proton concentration, however, is a constant. When the lattice Al/Si exceeds 10%, the proton–zeolite interaction energy increases. This increase in the proton–zeolite interaction decreases the intrinsic Brønsted acidity of the zeolite. At this concentration, the tetrahedra containing Al start to share a silicon tetrahedron. This increases the e ective negative

Figure 4.3. Dependence of the rate of a zeolite-catalyzed reaction on Al/Si ratio (schematic).

charge on the oxygen atoms. As a consequence, the zeolitic protons can no longer be considered to be independent. The e ect of having a negative charge on two or more tetrahedra containing Al prevents them from being in nearest-neighbor tetrahedra. This is the so-called L¨owenstein rule.

4.2 Activation of Reactant Molecules

4.2.1 Proton-Activated Reactivity

When a molecule adsorbs on the siliceous part of the micropore of a zeolite, the main interaction it experiences is a dispersive van der Waals-type interaction. This is due to the dominant interaction with the large polarizable oxygen atoms that make up the zeolite framework. For example, the interaction between a hydrocarbon CH3 or CH2 group and the siliceous framework typically results in an interaction energy that is on the order of 5–10 kJ/mol. These electrostatic interactions are small.

When an organic molecule approaches the zeolitic proton, in addition to the van der Waals dispersion forces, there is a weak additional interaction which is on the order of 5 kJ/mol. The large interaction energy of hydrocarbons with the siliceous zeolite channel [e.g. hexane in silicalite (ZSM-5), 60 kJ/mol] is due to the fact that there are multiple contacts between the hydrocarbon and the zeolite channel which are additive in nature. This helps to illustrates the fact that the siliceous zeolite channel is quite hydrophobic[7].

The interaction of a zeolitic proton with a polar adsorbate is much stronger than with an apolar hydrocarbon. The heat of adsorption of CH3OH, for example, to the zeolitic proton is on the order of 80 kJ/mol. This is largely due to the strong interaction between

the OH group on methanol and the zeolite proton. Its bonding features are well understood and sketched in Fig. 4.4[8].

Figure 4.4. Coordination of CH3 OH to zeolitic proton. Solid lines indicate covalent bonds whereas dotted lines represent hydrogen bonds.