Recent Advances and New Horizons in Zeolite Science and Technology (Studies in Surface Science and Catalysis, Volume 102)

Studies in Surface Science and Catalysis 102 RECENT ADVANCES AND NEW HORIZONS IN ZEOLITE SCIENCE AND TECHNOLOGY

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Studies in Surface Science and Catalysis Advisory

Editors:

B. D e l m o n

a n d J.T, Y a t e s

V o l , 102

RECENT ADVANCES AND NEW HORIZONS IN ZEOLITE SCIENCE AND TECHNOLOGY Editors H. Chon Department of Chemistry, KAIST, Yusung-ku,Taejon, 305-701 Korea

S.I.Woo Department of Chemical Engineering, KAIST, Taejon, 305-701 Korea S.-E, Park Industrial Catalysis Research Laboratory, KRICT,RO. Box 107, Taejon, 305-606 Korea

1996 ELSEVIER Amsterdam

- - L a u s a n n e - - N e w Y o r k m O x f o r d --- S h a n n o n - - T o k y o

ELSEVIER SCIENCE B.V. Sara Burgerhartstraat 25 P.O. Box 211, 1000 AE Amsterdam, The Netherlands

ISBN 0-444-82499-5 91996 Elsevier Science B.V. All rights reserved. No part of this publication may be reproduced, stored in a retrieval system or transmitted in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, without the prior written permission of the publisher, Elsevier Science B.V., Copyright & Permissions Department, P.O. Box 521, 1000 AM Amsterdam, The Netherlands. Special regulations for readers in the U.S.A.- This publication has been registered with the Copyright Clearance Center Inc. (CCC), 222 Rosewood Drive, Danvers, MA 01923. Information can be obtained from the CCC about conditions under which photocopies of parts of this publication may be made in the U.S.A. All other copyright questions, including photocopying outside of the U.S.A., should be referred to the copyright owner, Elsevier Science B.V., unless otherwise specified. No responsibility is assumed by the publisher for any injury and/or damage to persons or property as a matter of products liability, negligence or otherwise, or from any use or operation of any methods, products, instructions or ideas contained in the material herein. This book is printed on acid-free paper. Printed in The Netherlands

Preface Ever increasing interest and continuous developments in the zeolite science and technology have been reflected in the overwhelming response from all of the world to the 1 lth International Zeolite Conference. This book was conceived as a handbook for the 11th IZC Pre-conference summer school on zeolites, held in 1996 at Taejon, Korea. Three-day school on "Introduction to Zeolite Science and Practice" preceeding the 8th and 9th Internaional Zeollite Conference helped to improve the interaction between newcomers and experienced scientists in the field.

At the 10th IZC

Summer School, the lectures were given under the theme, "Advanced Zeolite Science and Applications" at an advanced level rather than introductory courses. Extending the concept of the 10th IZC summer school, the 11th IZC Summer School has also been organized to help those who have already actively worked on zeolite science and technology to be exposed to the latest new developments and the new horizons of zeolite science and technology for the 21 st century. The content of this book entitled, "Recent Advances and New Horizons in Zeolite Scinence and Technology" intended to give an extensive review and analysis of the important new findings of last 10 years on the synthesis, characterization and applications of zeolite materials as well as the prediction of new R&D directions for the next decade. We would like to express our appreciation to the authors who have written excellent manuscripts of their lectures in such a limited time. We sincerely hope this summer school has contributed to the advancement of the zeolite science and technology. Hakze Chon, Seong Ihl Woo, and Sang-Eon Park KAIST/KRICT, Taejon, Korea, April 1996

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vii List of contributors

T. Bein

Department of Chemistry, Purdue University, West Lafayette, IN 47907, USA H. van Bekkum Waterman Institute of Chemical Technology, Delft University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands M.E. Dry Catalysis Research Unit, Department of Chemical Engineering, University of Cape Town, Rondebosch 7700, South Africa M.J. den Exter Waterman Institute of Chemical Technology, Delft University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands S. Feast

Department of Chemical Engineering, University of Twente, P.O. Box 217, 7500 AE Enschede, The Netherlands J. van de Graaf Waterman Institute of Chemical Technology, Delft University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands R.F. Howe Deptmant of Physical Chemistry, University of New South Wales, P.O. Box 1, Sydney 20, Australia J.C. Jansen Waterman Institute of Chemical Technology, Delft University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands S. Kaliaguine P

9

9

Departement de Genie Chimique, Universite Laval, Ste-Foy, Quebec, G IK 7P4, Canada

viii

F. Kapteijn Waterman Institute of Chemical Technology, Delft University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands

S. Lee Department of Chemistry, University of Michigan, Ann Arbor, MI, 48109-1055, USA J.A. Lercher Department of Chemical Technology, Catalytic Processes and Materials, University of Twente, P.O. Box 217, 7500 AE Enschede, The Netherlands J.M. Newsam Molecular Simulations inc., 9685 Scranton Road, San Diego, CA 92121, USA C.T. O'Connor Catalysis Research Unit, Department of Chemical Engineering, University of Cape Town, Rondebosch 7700, South Africa A. Sayari

Department of Chemical Engineering and CERPIC, Universite Laval, Ste-Foy, Quebec, G1 K 7P4, Canada Karl Serf Chemistry Deparment, University of Hawaii, 2545 The Mall, Honolulu, HI 968222275, USA E. Van Steen Catalysis Research Unit, Department of Chemical Engineering, University of Cape Town, Rondebosch, 7700, South Africa M. St~cker SINTEF Oslo, Department of Hydrocarbon Process Chemistry, P.O. Box 124 Blindern, N-0314 Oslo, Norway

ix S.L. Suib Department of Chemistry, Department of Chemical Engineering, and Institute of Materials Science, University of Connecticut, Storrs, CT 06269-3060, USA D. Venkataraman Department of Chemistry, Cornell University, Ithaca, NY 14853, USA

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xi

Contents

Preface List of contributors

vii

Chapter 1. Periodic Mesoporous Materials- Synthesis, Characterization and Potential Applications A. Sayari Introduction Synthesis and characterization of M41S and related materials Catalytic Applications of M41S and related materials Other potential applications Concluding remarks References

1 2 27 33 33 37

Chapter 2. Synthesis, Characterization, and Catalysis with Microporous Ferrierites, Octahedral Molecular Sieves, and Layered Materials S.L. Suib Overview n-Butene isomerization with boron substituted zeolites and

47

ferrierites Synthesis of octahedral molecular sieves and octahedral

52

layered materials Characterization of octahedral molecular sieves and octahedral

55

layered materials Catalytic activity of octahedral molecular sieves and octahedral

65

layered systems References

67 69

xii

Chapter 3. Organic Zeolites ? Stephen Lee and D. Venkataraman Introduction Dianin's compound Helical tubulates of 2,6-Dimethylbicyclo[3.3.1]nonaneexo-2,exo-6-diol [Ag(1,3,5-tris(3-ethynylbenzonitrile)benzene)CF3SO3]-2CsH 6 [Ag(1,3,5-tris(4-ethynylbenzonitrile)benzene)C F3SO3]-2C6H8 Further directions References and notes

75 76 78 79 81 84 92

Chapter 4. Spectroscopic Characterization of Zeolites R.F. Howe Introduction EPR spectroscopy Infrared spectroscopy Raman spectroscopy UV-Visible spectroscopy X-ray absorption spectroscopy Mass spectrometry of zeolites References

97 98 106 123 127 130 134 136

Chapter 5. Characterization of Zeolitic Materials by Solid-State NMR -State of the Art M. St~cker Introduction High resolution solid state NMR spectroscopy- experimental techniques including 2D NMR Recent highlights about the framework characterization Pore architecture investigated by NMR In-situ NMR studies with zeolitic materials

141

Diffusion of adsorbed molecules monitored by NMR Acidity of zeolitic materials Concluding remarks

179 181 184

143 157 172 176

xiii

References

185

Chapter 6. Application of Surface Science Techniques in the Field of Zeolitic Materials S. Kaliaguine Introduction Photoelectron spectroscopy (XPS) and other surface analysis XPS of zeolites Acidity in zeolites Basicity in zeolites Active centers in zeolitic oxidation catalysts Conclusion References

191 192 204 209 217 221 225 227

Chapter 7. Computational Approaches in Zeolite Structural Chemistry J.M. Newsam Roles of simulation Structural characterization Occluded or sorbed molecules and clusters Meso- and macro-structure Dynamical behavior Sorptive behavior Intrazeolite chemistry Some five year challenge areas Conclusion References

231 233 247 250 251 253 255 255 259 260

Chapter 8. What Can Be in the Channels and Cavities of Zeolites ? Karl Serf Zeolites and their frameworks The contents of neutral zeolite frameworks

267

The contents of charged zeolite frameworks

270

Molecular and ionic sieving

274

Sorption of atoms and molecules

276

269

xiv

A single substance may react within a zeolite

279

Two different substances may react within a zeolite Encapsulation

285 286

An empty zeolite is a crystalline solvent

289

An appeal for care in the preparation of samples for study

289

Final comments

291

References

291

Chapter 9. Conjugated and Conducting Nanostructures in Zeolites T. Bein Introduction

295

Polyacetylene and derivatives in zeolites

304

Heteroaromatic conducting polymers in zeolites

305 314

Carbon-based conducting materials in nanometer channels Conclusions References

317 318

Chapter 10. New Catalytic Applications of Zeolites for Petrochemicals C.T. O'Connor, E. Van Steen, and M.E. Dry Introduction Catalytic cracking Alkylation Aromatization of alkanes/alkenes

323 325 336 344

Skeletal isomerization of 1-butene Alkene oligomerization

353

Isomerization of long-chain alkanes

353

References

355

349

Chapter 11. Synthesis of Intermediates and Fine Chemicals using Molecular Sieve Catalysts S. Feast and J.A. Lercher Introduction

363

Chemical functionalities of molecular sieves

365

Physical aspects of molecular sieve catalysis

xv

for chemical synthesis Conclusion and outlook References

396 400 404

Chapter 12. Zeolite-based Membranes, Preparation, Performance and Prospects M.J. den Exter, J.C. Jansen, J. van de Graaf, F. Kapteijn, and H. van Bekkum Introduction Dynamics of zeolite pores and consequences for adsorption

413

and permeation Small pore zeolites Medium pore zeolites Large pore zeolite-based membranes (12 membered

417 421 428

oxygen ring) Permeation through silicalite-1 membranes: examples

432

and modelling Zeolite membranes in catalytic conversions

433 446

Prospects References

450 451

Keyword index

455

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H. Chon, S.I. Woo and S.-E. Park (Editor,s) Recent Advances and New Horizons in Zeolite Science and Technology Studies in Surface Science and Catalysis, Vol. 102 9 1996 Elsevier Science B.V. All fights reserved.

Periodic mesoporous materials: synthesis, characterization and potential applications Abdelhamid SAYARI Department of Chemical Engineering and CERPIC, Universit~ Laval, Ste-Foy, Qc, Canada G1K 7P4

1. INTRODUCTION The design of materials with precise pore structure is of tremendous technological importance [ 1,2]. In addition to zeolites and molecular sieves, subject of our meeting, there are many other materials where the shape, connectivity and size distribution of their pores determine their usefulness. Such materials include pillared clays [3], carbon molecular sieves (CMS) [4], sol-gel materials [5], porous polymers [6,7] and other organic solids [8-10], hollow tubes of polypeptides [11,12], carbon [13,14], polymers [15], metals [16] and inorganic oxides [17,18] as well as imprinted organic [19] and inorganic [19,20] materials. For some applications such as air separation by CMS, changes in the optimum effective micropore size by only 0.3 ]k dramatically deteriorate the separating ability of the CMS [21]. As far as crystalline porous materials are concerned, zeolitic molecular sieves (zeotypes) are of particular interest not only because of their important industrial applications as adsorbents [22], ion-exchangers [23] and catalysts [24,25] but also because of theft potential applications as hosts for a variety of technologically advanced materials such as semiconductor clusters, molecular wires with tailored electronic and optical properties, enzyme mimicking organic complexes, etc. [26-30]. Zeotypes were historically divided into four groups: small, medium, large and ultra-large pore molecular sieves, with pore openings comprised of 8, 10, 12 and more than 12 tetrahedral (T) atoms, respectively. Figure 1 shows the pore size of typical zeotype materials. Before the discovery of M41S in 1991, the only ultra-large pore zeotypes were metallophosphates. These are A1PO4-8 [31], VPI-5 [32], Cloverite [33] and JDF-20 [34] with 14, 18, 20 and 20-membered rings, respectively. All microporous zeotypes have been reviewed extensively [35,36]. This paper is concerned with the particular class of molecular sieves having periodic mesoporous structure with pore sizes in the range of 2 to 10 nm. They are comprised of the M41S mesoporous molecular sieves and solids with related structures. In the first part, the preparation methods and characterization techniques will be reviewed and discussed. Silicatebased materials and non-silicate materials will be dealt with separately. In the second part of this review particular emphasis will be put on potential applications reported in both the patent and the open literature. Early progress in this field has been presented in the previous Summer School by Casci [37]. Potential catalytic applications of M41S were also reviewed r~.e~.ntlv rqRl_

Figure 1. Pore size of molecular sieves. 2.

SYNTHESIS AND CHARACTERIZATION OF M41S AND RELATED MATERIALS

The term "periodic" will be used interchangeably with the word "crystalline", because as will be shown in a later section, these materials have actually amorphous walls despite their long-range periodic structure. Such materials are prepared very much like zeolites by hydrothermal treatment of a gel, typically at 80 to 120 ~ for 24 to 48 h. The main difference between the preparation of zeolites and M41S materials is that in the former case a single cationic (e.g. alkali, organic ammonium or phosphonium cations) or molecular (e.g. amines) species is used as "template" (or structure directing agent [39-41]), whereas in the latter case an array of surfactant molecules plays the role of template.

Strictly speaking, a template is a structure (usually organic) around which a material (often inorganic) nucleates and grows in a "skin-tight" fashion, so that upon the removal of the templating structure, its geometric and electronic characteristics are replicated in the (inorganic) material. In this sense, except for the possible example of ZSM-18 [39,42], there is hardly any examples of true templating effect in zeolite synthesis. Because of geometrical constraints and the small size of species used as templates, the host-guest interactions are rather weak so that the shape of the pore most often does not reflect that of the template. This among other problems makes the "rational design" of zeolite synthesis particularly difficult [39,43]. In the early 90's Mobil researchers had the brilliant idea to use supramolecular surfactant templates for the preparation of porous silicates. This led to the discovery of the so-called M41S mesostructured molecular sieves and to a tighter control on the morphology and pore size of the desired material through simple manipulations of the synthesis conditions. The preparation of mesoporous silicates requires three ingredients in the appropriate amounts: a solvent, a source of silica and a surfactant. In addition, other reagents such as acids, bases, salts, expander molecules and cosolvents may be used. The nature and relative amount of all these ingredients may vary greatly, thus offering a high degree of flexibility for the design and preparation of periodic mesostructured materials. The Mobil group identified three main silicate phases. They consist of a hexagonal phase referred to as MCM-41, a cubic phase (space group: Ia3d) known as MCM-48, and a non stable lameUar phase (MCM-50) which can be stabilized by post-synthesis treatment in the presence of tetraethyl orthosilicate (TEOS). Among other facts, the striking similarity between these structures and those of known liquid crystal phases led Mobil researchers and others to propose the so-called liquidcrystal templating (LCT) mechanisms in order to rationalize the formation of mesostructured materials.

2.1 Synthesis Mechanisms of Surfactant/Silieate Mesophases The most studied crystalline nanoporous silicates are those prepared in aqueous medium in the presence of long chain alkyltrimethylammonium hydroxide or halide under basic conditions. It is therefore practical to deal with this case ftrst, then consider other pH conditions, surfactants and synthesis approaches. At this stage it is essential to learn about surfactant aggregation in water. A surfactant is an amphiphile molecule with a hydrophilic head group and a hydrophobic tail. Figure 2 shows a schematic phase diagram of cetyltrimethylammonium bromide (C16T/VIABI")ill water [44]. Depending on the temperature and concentration, surfactants tend to self-organize into aggregates with different shapes. At very low concentration, the surfactant molecules are randomly dispersed in solution. As the concentration reaches a critical level referred to as CMC1, spherical micelles are formed. The outer surface of the micelle is comprised of the hydrophilic heads of surfactant molecules, while the tails of these molecules are directed toward the center of the micelle. There exists a second critical concentration CMC2 corresponding to the further aggregation of spherical into cylindrical or rod-like miceUes. As the concentration increases further, a higher level of aggregation into liquid crystals takes place. There are three main liquid crystalline phases with hexagonal (H1), cubic (V 1) and lameUar (L) structures (Figure 3). The H 1 phase is the result of a hexagonal packing of cylindrical miceUes, while the L phase corresponds to the formation of surfactant bilayers.

The cubic phase may be regarded as a bicontinuous structure as suggested by Luzzati et al. [45]. To describe the tendency of an amphiphile to aggregate into a particular morphology, Israelachvili et al. [46,47] introduced the packing parameter g = v/aolc with 1c < lmax = 1.54 + 1.26 n (A) and v = 27.4 + 26.9 n (A3). Here 1c, v and n stand for the critical length, the volume and the carbon number of the hydrophobic chain, and ao for the optimum surface area of the headgroup. The packing parameter determines whether the amphiphile will form spherical micelles (g < 1/3), cylindrical micelles (1/3 < g < 1/2), vesicles or bilayers (1/2 < g < 1) or inverted structures (g > 1). The LCT mechanism as postulated by Mobil scientists [48,49] is represented in Figure 4. Along Pathway (1), the template is supposed to self-organize into a liquid-crystal phase such as H 1 before being encapsulated by inorganic species which then condense into rigid walls. This mechanism is consistent with recent data obtained in the presence of high concentrations of surfactant [50]. In the second proposed pathway, the inorganic species participate in the ordering process of the surfactant-inorganic mesophase and influence its morphology. The LCT mechanism was ftrst proposed based on (i) the close similarity between the morphology of the surfactant-inorganic mesophases and liquid crystals, and (ii) the dependence of the pore size on the surfactant chain length and on the amount of auxiliary organics such as 1,3,5 trimethylbenzene (TMB) [48,49]. Further work by the same group [5153] supports the occurrence of Pathway (2). Two main reasons were evoked, (i) M41S materials are often prepared in the presence of surfactant concentrations well below that required for the formation of a liquid-crystal phase, and (ii) the hexagonal, culSic and lamellar M4IS structures may be formed by changing only the silica concentration indicating that no preexisting liquid-crystal phase is required. Several other workers reached similar conclusions, but often with slight variations [5459]. Based on XRD, thermogravimetric analysis, 29Si and in situ 14N NMR, Davis et aL [54,55] concluded that rod-like miceUes coated with 2 to 3 monolayers of silica form before they spontaneously self-organize into a hexagonal phase, with further silica condensation during calcination. Obviously, this mechanism may be operative only if rod-like micelles are formed in the synthesis medium. To this end two requirements should be met: (i) the surfactant carbon chain should be long enough so that the formation of rod-like micelles is possible, and (ii) the surfactant concentration should be equal to at least CMC2, the minimum concentration for the generation of such micelles. While agreeing that a liquid-crystal templating mechanism initiated by silica occurs, Cheng et al. [57,58] claimed that in the presence of TEOS as the silica source synthesis takes place only if the surfactant concentration is equal to CMC1 or higher. Based on XRD and 14N NMR data, Steel et al. [59] proposed a modified LCT mechanism. They postulated that the silicate source first dissolves into the reaction medium and promotes the formation of a surfactant hexagonal mesophase. Then the silicate forms parallel layers in between the rows of cylindrical micelles. The hexagonal organic-inorganic mesophase is then formed by puckering of the silicate layers. If the concentration of silicate is high, the layers will be thicker and the puckering does not take place, thus leading to a lamellar phase. One of the most important contributions toward the elucidation of synthesis pathways of M41S and related materials was made by Stucky and his colleagues [60-66] at both Santa Barbara University (USA) and Johannes Gutenberg-Universi~t, Mainz (Germany). They fully

Figure 4. Mobil group proposed formation pathways of M41S [49].

documented the fact that the presence of preorganized liquid crystal structures or even rodlike micelles prior to adding the inorganic precursor is not required as a "static" template for nucleation and growth of the inorganic phase. This conclusion stems from several experimental observations [61]: (i) hexagonal MCM-41 phases can be prepared using C16TMABr, C16TMAC1 or C16TMAOH at concentrations much below CMC2, (ii) hexagonal MCM-41 phases can also be made in the presence of short chain surfactants such as C12TMAOH and C12TMAC1 which have not been reported to form rod-like micelles in water, and (iii) MCM-41 and MCM-48 can be prepared at temperatures above 70 ~ where rod-like miceUes are unstable. Moreover, Firouzi et al. [64] argued that addition of inorganic species to a micellar assembly of organic molecules often leads to reorganization into new morphologies depending on the electrostatic and steric interactions between organic and inorganic species. The cooperative templating mechanism proposed by Stucky and coworkers is shown schematically in Figure 5. Prior to silicate addition, the surfactant is in a dynamic equilibrium between spherical or cylindrical micelles and single molecules. Upon addition of a silica source, the predominantly multicharged silicate species ion-exchange with the OHor Br- anions to form organic-inorganic ion pairs accompanied by dissociation of the organic miceUes and aggregation of the ion pairs into a new mesophase. The multidentate interaction controls the number of surfactant molecules that can bind to a given inorganic species and determines the interface packing density and ultimately the biphase morphology. The last step is the polymerization and condensation of the inorganic species. In addition to the arguments mentioned earlier in favor of the cooperative organization of inorganic-surfactant mesophases, further support was obtained by investigating a system with very low surfactant concentration where the effects of self-assembly are decoupled from the kinetics of silica condensation. Using in situ small-angle neutron scattering, Stucky et al. [64,66] found that a 1% CTAB aqueous solution exhibits an isotropic miceUar distribution. Addition of a silicate solution leads to the formation of an inorganic-organic hexagonal array. Similar conclusions were drawn from freeze-fracture electron microscopy studies [64,67]. Moreover, in the absence of polymerization, in situ 21-1NMR of labelled surfactant showed that this mesophase referred to as a silicatropic liquid crystal (SLC) exhibits a behavior very similar to a lyotropic liquid crystal (LLC) including a reversible first-order transformation between the lamellar and hexagonal phases. Huo et al. [60,61] also demonstrated that such a cooperative organization process is not limited to ion pairs formed between cationic surfactants (S§ and anionic inorganic species (I-) but can be easily generalized to include other pathways. In addition to the S+I- route described above, three other pathways were considered. Pathway S-l+ involves the cooperative organization of a cationic inorganic solution species and an anionic surfactant. The other two routes correspond to the assembly of surfactant and inorganic ions with similar charges, mediated by small ions with the opposite charge. These pathways are referred to as S+X-I + (X- = CI-, Br-) and S-M~- (M+ = Na § K+). Typical syntheses to illustrate each of these pathways have been reported [60,61]. Fyfe and Fu [68] proceeded in two steps. They ftrst obtained a precipitate by reacting the octamer Si8028t~ with C16TMABr. The degree of condensation was then adjusted by titrating the oxoanions by acid vapor treatment during which the following sequence of structural transformations took place: layered precipitate ---) cubic ---) lamellar ---) hexagonal

mesophase. These findings strongly indicate that before extensive polymerization of the silicate species, the organic-inorganic mesophase exhibits the behavior of a liquid-crystal.

Figure 5. Cooperative templating mechanism [64]. In another important development, Pinnavaia et al. [69-72] used non ionic surfactants such as primary amines and polyethylene oxide to prepare silicates with cylindrical nanopores

referred to as HMS and MSU-n, respectively. Contrary to the case of charged surfactants where the electrostatic interactions between inorganic and surfactant ions play a key role in determining the morphology of the mesophase, in the presence of neutral surfactants hydrogen bonding becomes a predominant factor. Figure 6 represents the S~ ~ pathway in the case of primary amine surfactants. The authors postulated that the Si(OC2Hs)4.x(OH)x species formed by hydrolysis of TEOS participate in H-bonding interactions with the lone pairs of surfactant amine headgroups. This new organic-inorganic complex may be considered as an amphiphile with a very bulky headgroup which increases the likelihood for the formation of rod-like micelles. These micelles self-organize into a hexagonal packing followed by condensation of silanol groups and silica walls formation. k,-./ S~

i~

CnH2n+INH 2

+

Si(OEt)4.x(OH)x

B .# qHSN"

/ ;""e

$i o H

H~H

/?%.

!t

H

o OEt

#

"'/

--

Figure 6. S~ ~ templating mechanism [70]. Parallel to the discovery of M41S materials, Yanagisawa et al. [73-75] from Waseda University and Inagaki et al. [76-79] from Toyota prepared nanoporous silicates designated as FSM- 16 using a completely different strategy. They treated a layered kanemite polysilicate in a 0.1 M aqueous solution of C16TMABr at 70 ~ followed by filtration, drying and calcination. Further work on this method was carried out by Vartuli et al. [51] and Chen et al. [80]. As shown schematically in Figure 7A, Inagaki et al. [76] proposed a two-step mechanism for the formation of the FSM-16 material. First, the surfactant cation exchanges with Na§ and penetrates in between the silicate sheets. Subsequently, the flexible silicate layers wind around the exchanged C16TMA+ cations. Further condensation of silanol groups between adjacent silicate layers leads to a highly ordered honeycomb structure. The intermediate formation of a lameUar silica-surfactant mesophase has been observed recently by in situ XRD [75]. As represented in Figure 7B, Chen et al. [80] proposed that once

C16TMA + exchanges with Na + forming a bilayer intercalated silicate, local rearrangement of silicate species accompanied by the formation of rod-like miceUes and silanol groups condensation leads to the nanoporous structure. Table 1 Suffactant used for the preparation of mesostructured materials Cationic CnH2n+I(CH3)3N+X -

X- = CI-, Br-, OH-

n = 8-22

CnH2n+l (C2H5)3N+X -

n = 12, 14, 16, 18

(CnH2n+I)2(CH3)2N+X -

n = 10-18

C16H33N(CH3)2CnH2n+ 1

n = 1-12

Gemini

+ + [CnH2n+ 1(CH3)2N-CsH2s-N(CH3) 2CmH2m+ 1]Br2

n = 16, s = 2-12, m = 1-16

Anionic C14H29COOH, C 17H35COOH C12H25OPO3H 2, C14H29OPO3K n = 12, 14, 16, 18

CnH2n+IOSO3Na C16H33SO3H, C12H25C6H4SO3Na Neutral CnH2n+INH2

n = 10-16

CnH2n+IN(CH3)2

n = 10-16

Cll_15(EO)n

C = alkyl

n = 9, 12, 15, 20, 30

CnPh(EO) m

Ph = phenyl

n = 8 or 12, m = 8, 10, 18

(PEO) 13(PPO)30(PEO) 13

10

Figure 7. Formation methcanisms of FSM-16 according to (A) Inagaki et al. [76], and (B) Chen et al. [80]. More recently Fukushima et al. [79] carried out the ion exchange at pH = 12 followed by pH adjustment to 8.5. Elemental analysis and 29Si MAS NMR showed that at pH = 12, single sheets folded, but without condensation between them (only Q3 Si species were present in the solidphase). In addition significant amounts of silica dissolved. Upon adjustment of the pH to 8.5, an increase of Q4/(Q3 + Q4) to 0.6 was observed suggesting that dissolved SiO2 condense into walls 2 to 3 layers thick. Galameau et al. [81] used a similar approach to design porous clay heterostructures. They intercalated layered fluorohectorite by C16TMA+ cations followed by treatment in a solution of neutral amine and TEOS, then drying and calcination. In brief, the authors proposed that the interlayer galleries of the intercalated clay are further swollen by the amine followed by insertion of TEOS, formation of rod-like micelles and silica polymerization. 2.2 Synthesis Conditions of M41S and Related Mesoporous Silicates Periodic nanoporous silicates have been prepared in a wide variety of conditions. Different sources of molecular, and non molecular silica have been used. This includes TEOS, TMOS, fumed, colloidal and precipitated silicas. Depending on the synthesis conditions, particularly on the nature of the silica source, crystallization may take place in seconds at subambient temperatures [82], or at room temperature [60,61,69,72,83]. However, in most cases the crystallization temperature was set in the 80 - 120 ~ range. Liu et al. [84,85] found that the use of small amounts of colloidal particles (silica or titania) promotes the formation of ordered structures by providing nucleation seeds. The pH conditions varied from extremely acidic [60,61], to neutral [69,72] to very basic [48,49]. Ryoo and Kim [86]

11 showed that alternating hydrothermal treatments (373 K, 24 h) and pH adjustments using acetic acid (pH = 11) leads to MCM-41 products with improved crystallinity and in higher yields due to an equilibrium shift. All the surfactants listed in Table 1 have been used for the preparation of nanostructured materials, mainly for silicates. The most used surfactants are alkyltrimethylammonium hydroxides or halides. Depending on the synthesis conditions, they give rise to hexagonal, cubic or lamellar structures. Because of their high packing factors, two-tailed surfactants favor the formation of lameUar structures. On the contrary, surfactants with bulky headgroups promote the formation of mesophases with high surface curvature. For example, the use of surfactants such as alkyltriethylammonium and cetylethylpiperidinium under acidic conditions afforded a cubic phase named SBA-1 [87] with space group Pm~n [60,61]. Using gemini surfactants CnH2n+IN+(CH3)2(CH2)sN+(CH3)3 with very large headgroups, under either basic or acidic conditions, Huo et al. [87] discovered a new mesophase (SBA-2) that has three-dimensional hexagonal (P63/mmc) symmetry and apparently no lyotropic surfactant or lipid liquid mesophase counterpart. The structure of this material is derived from hexagonal close packing of globular silicate-surfactant arrays. The cage-structured mesoporous network of SBA-2 may offer advantages over the unidimensional channels of MCM-41, particularly in host-guest applications. Sayari et al. [88,89] used a series of alkyl cetyldimethylammonuim bromides, (C16H33)(CnH2n+1)(CH3)2N+Br - (n = 1 to 12), under basic conditions. For n < 9, an interesting odd-even effect was observed. Hexagonal and lamellar phases were obtained in the presence of surfactants with odd and even n values, respectively. For n >__9, the surfactant behaved as a long chain two-tailed molecule and gave only lamellar phases. These findings were related to the organization of the occluded surfactant as evidenced by 13C NMR. Two other mesoporous silicates, HMS and MSU-n, were prepared at room temperature in the presence of neutral amines, and polyethylene oxide surfactants, respectively. They have parallel cylindrical channels, but are not identical to MCM-41 [69,70,72]. The main structural difference is that the pore system of HMS and MSU-n silicates is much less ordered than that of MCM-41. The use of cosolvents may have different effects as illustrated in the following examples. Anderson et al. [82] found that the use of methanol or formamide as cosolvents enabled them to control the crystallization temperature. Moreover, the use of methanol as cosolvent and formamide as solvent allowed the fine tuning of the pore size to within 1/~. Huo et aL [87] found that the use of tert-amyl alcohol as a polar cosolvent in the presence of surfactants with bulky headgroups such as C16H33N+(C2H5)3 or CnH2n+IN+(CH3)2(CH2)sN+(CH3)3 increases the volume of the hydrophobic core and therefore the packing factor, leading to a hexagonal (MCM-41) instead of a cubic (Pm~ n) or a three dimensional hexagonal (P63/mmc) phase. However, the most known effect of co-solvents is that of non polar trimethylbenzene (TMB) which dissolves into the hydrophobic part of the micelle and acts as a swelling agent [48,49,61]. One of the most important aspects of the liquid-crystal templating strategy for the manufacture of mesoporous inorganic materials is the ability to adjust the pore size (or the interlayer distance) between ca. 2 and 10 nm. This may be achieved by using surfactants with different chain lengths (Figure 8). This method used for all three M41S phases has been extended to SBA-2 and more recently to FSM silicates [90]. Very fine pore size tuning may be achieved by using variable amounts of methanol as cosolvent [82]. For pore sizes above

12

ca. 40 A, expander molecules such as TMB may be used (Figure 9). These molecules increase the size of the surfactant miceUes by dissolving into their hydrophobic region. Another less known method [91] consists of preparing a MCM-41 silicate at relatively low temperature, e.g. 70 ~ then heating it in its mother liquor at high temperature, e.g. 150 ~ (Figure 10). This restructuring occurs via silica dissolution, transport and redeposition. Moreover, the pore size may be narrowed by post-synthesis silylation [49].

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Figure 8. Pore size of MCM-41 as a function of carbon number of RTMABr. n. ref. [49]" A: unpublished work.

Figure 9. Variation of dlo0 spacing of MCM-41 as a function of TMA/SiO 2 ratio [61].

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Aging time, day

2.3 Characterization of Mesoporous Silicates Mesoporous silicate molecular sieves were characterized by an arsenal of techniques including XRD, SEM, TEM, adsorption measurements, 29Si NMR, IR, Raman and XANES. XRD patterns of all nanoporous phases are dominated by low angle peaks. Figure 11 shows typical patterns for MCM-41, MCM-48, MCM-50 and SBA-2 phases. HMS, MSU-n and some "MCM-41" exhibit only the 100 peak either because of too small scattering domain sizes [69] or because of poorly ordered pore system [55,72,80].

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Figure 11. Typical XRD patterns. (A) MCM-41, 03) MCM-48, (C) Lamellar, (D) SBA-2. TEM combined with electron diffraction has been a key technique for proper characterization of nanoporous materials [48-55,58,60,61,80,92-94] and elucidation of their structures [87,95]. An interesting application of TEM was reported by Chen et aL [80] who used the Fresnel method to determine the thickness of the FSM-16 walls. According to this method the imaged wall thickness decreases with decreasing focus and reaches a minimum which is the real thickness at zero defocus. Adsorption measurements indicate that mesoporous silicates have high porosity (0.7 -1.2 cm3/g) and very high surface area often exceeding 1000 m2/g. As shown in Figure 12 (upper part) N 2 adsorption-desorption isotherms of MCM-41 [49,96,97-99], MCM-48 [49] and FSM [90] materials are of type IV in the IUPAC classification and have a typical shape. They are usually reversible and exhibit a sharp step at P/Po in the range 0.25 to 0.5 depending on the average pore size. This step increase in N 2 adsorption corresponds to capillary condensation within uniform mesopores. The sharpness of this step reflects the uniformity of the pore size whereas its hight the pore volume. Llewellyn et al. [ 100] reported that hysteresis in the Ar

14 and N 2 adsorption-desorption isotherms does occur for materials with pore sizes exceeding 2 and 4 rim, respectively. The authors concluded that these pore sizes represent the limit between "secondary micropores" and "mesopores" for each adsorptive. Branton et al. [96,101] also found that 0 2 and Ar adsorption isotherms at 77 K exhibit significant hysteresis loops. Rathousky et al. [102] studied the adsorption of cyclopentane on MCM-41 at 243-333 K and found a hysteresis loop when the adsorption is carried out below the pore critical temperature. Nitrogen adsorption isotherms were also calculated using molecular simulation [103] and density functional theory [104, 105]. Figure 12 (lower part) shows that the SBA-2 N 2 adsorption isotherm exhibit a H2 hysteresis loop consistent with bottle-shaped pores [87]. EFFECTIVE PORE DIAMETER (nm) 1.0

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Figure 12. N 2 adsorption-desorption isotherm and pore size distribution for MCM-41 (upper) and SBA-2 (lower) silicates. Water adsorption on MCM-41 was studied by 1H NMR [94,106-108] and by gravimetry and FTIR [ 109]. Schmidt et al. [ 106,108] measured the freezing point of adsorbed water by

15 1H NMR and used it to determine the pore size distribution of mesoporous silicates. They also derived the self-diffusion coefficient of water in MCM-41 and MCM-48 [94,107]. LleweUyn et al. [109] found that MCM-41 exhibits a type V water adsorption isotherm indicating an initial repulsive character followed by capillary condensation at higher pressures. The wall thickness of hexagonally packed silicates (MCM-41 and HMS) was determined as the difference between the repeat distance ao between pore centers (measured from TEMs or calculated from XRD data with the formula ao = 2d10o/~/3) and the Horvath-Kawazoe [110] pore diameter. In most studies it was found to be in the range of 9 to 12 A. Molecular dynamics simulations show that there is an excellent agreement between experimental XRD patterns and those calculated based on _> 10/~ wall of amorphous silica models [ 111]. This rather thin wall combined with very high porosity is at the origin of the relatively low mechanical stability of these materials [112]. Tanev and Pinnavaia [69] reported that the walls of HMS silicates are always thicker than those of MCM-41 and connected this observation to differences in formation mechanisms. Coustel et al. [ 113] found that the wall thickness of MCM-41 silicates may be varied from as low as 5/~ to 17 A, mainly through control of the OH- / S i t 2 ratio, thus the silicate solubility. Magic angle spinning (MAS) 29Si NMR was also used extensively for the characterization of M41S silicates [49,114]. One of the earliest observations is that 29Si NMR spectra are broad and closely resemble those of amorphous silica, a strong indication that the material waUs are actually amorphous with a wide range of O-Si-O angles. This conclusion was further supported by FTIR [115], Raman [115] and SiK XANES [116] data. Deconvolution of 29Si NMR spectra gave essentially two peaks at -100 and -110 ppm corresponding to the so-called Q3 [Si(OSi)3OH] and Q4 [Si(OSi)4] environments. A small peak (Q2) may appear at ca. -90 ppm. The Q3/Q4 ratio reflects the degree of polymerization of silica and the density of OH groups. Vartuli et al. [53] found that in as-synthesized M41S silicates prepared by the electrostatic S+I- pathway, Q3 remains almost constant at ca. 50%, consistent with the N/Si ratio being practically constant in all three phases. On the contrary, Steel et al. [114] found that Q3/Q4 is the same (59%) for as-synthesized MCM-41 and MCM48, but much higher (96%) for the lamellar phase suggesting that the latter phase has thinner walls. Likewise, Hut et al. [60] found Q,/Q4 ratios of 55 and 120% for hexagonal and lamellar silicates, respectively regardless of whether they are prepared by the electrostatic S+I- or S+X-I+ pathway. 29Si NMR also shows that for thermally stable materials the Q3/Q4 ratio decreases upon calcination [87,114,115] as additional cross linking takes place. Sayari e t al. [96] found that Q4/(Q2 + Q3) increases upon incorporation of boron in the silicate framework indicating that the substitution occurs through interaction with hydroxyl groups.

2.4 Synthesis and Characterization of Modified Mesoporous Silicates 2.4.1. Mesoporous Aluminosilicates The synthesis of aluminum containing MCM-41 molecular sieves was reported in both the patent [117-119] and the open [48,49,92,113,115,120-129] literature. There are also some reports on A1-HMS [130] and A1-FSM- 16 [77]. MCM-41 aluminosilicates were prepared under hydrothermal conditions, typically at 70-150 ~ during 1 to 10 days. Various sources of silica and alumina were used. Janicke et al. [121] reported that aluminum isopropoxide is a much better precursor than Catapal B. In the presence of isopropoxide, they prepared samples with Si/A1 ratios down to 16, with aluminum being entirely in tetrahedral positions.

16 Luan et al. [ 125] tested a number of aluminum precursors and found that aluminum sulfate leads to total incorporation of A1 in tetrahedral sites up to very high loadings (Si/A1 = 2.5). On the contrary, other researchers prepared M-rich samples using sodium aluminate [122124,128]. These discrepancies may be explained, at least partly, on the basis that the preparation methods used by various groups were different. Fu et al. [131] prepared MCM-41 aluminosilicates using the two-step approach described earlier [68]. They first prepared well defined aluminosilicate polyanions AlxSis.x(OH)xO2o. 4 (0 < x < 4) oligomers to be used as precursors. These precursors were then precipitated with C16TMABr and treated with water vapor at 110 ~ for 3 days. This method not only affords MCM-41 aluminosilicates with variable Si/A1 ratios down to the lowest possible ratio of 1/1, but it offers additional flexibility in the design of new materials by using suitable building blocks. Incorporation of aluminum in MCM-41 silicates brings about a dramatic decrease in the intensity of the XRD [124,125] and a significant broadening of the pore distribution [120]. As mentioned before, the MCM-41 silicate walls are amorphous with a wide range of T-O-T bond angles. The presence of A1 in such a highly distorted environment in addition to the limited flexibility of the O-A1-O angle compared to O-Si-O may generate a more defective structure with a broader pore size distribution. 27A1 MAS NMR was widely used to distinguish between "framework" and extraframework aluminum. Several workers found a linear relationship between the absolute intensity of the 27A1 NMR peak and the A1 content in as-synthesized samples [ 120,124,125,132]. However, it was also reported that A1 containing MCM-41 silicates exhibit a strong tendency to dealumination during the removal of surfactant by calcination. Dealumination is mostly due to hydrolysis of framework aluminum by steam generated during the combustion of the surfactant. Corma et al. [120] found that direct calcination of Al-rich MCM-41 at 540 ~ in air gives rise to material with significant amounts of extraframework aluminum, smaller pores and lower BrCnsted acid site density compared to samples treated first in N 2 and then in air at the same temperature. The two-step activation procedure has two advantages: (i) the local temperature is lower, and (ii) much less water vapor is formed. Luan et al. [125,133] found that dealumination depends on the nature of counter cations. Calcination of as-synthesized samples does not generate extraframework A1 species, but it brings about a broadening of the 53 ppm 27A1 NMR peak due to decreased symmetry. However, proton exchanged samples were found to be prone to dealumination, presumably because the small H § cation cannot satisfy the framework charge balance efficiently. A series of A1 containing MCM-41 samples with Si/A1 ratios in the range 62 to 2.5 were synthesized in our laboratory using sodium aluminate [132]. As show in Table 2, NMR data indicated that for both as-synthesized and for calcined samples at least 90-95% of all A1 in all samples was located in tetrahedral positions. Table 2 also shows that the BET surface area decreases sharply as the Si/A1 ratio drops below 10. This is consistent with TEM observation made by Kloetstra et al. [126] who found that in samples with low bulk Si/A1 ratios, most of the aluminum was part of a separate dense phase displaying a tetrahedral environment. This indicates that the conventional one-pulse NMR technique does not discriminate between tetrahedral A1 in the MCM-41 framework and tetrahedral A1 in the so-called dense phase [126], and should be combined with adsorption and TEM measurements for proper characterization of the state of aluminum.

17 The thermal and hydrothermal stability of a sample with Si/A1 = 26 was investigated in our laboratory by heating batches for 3 h in dry air, or in pure water vapor in the temperature range 550 to 850 ~ In dry air, the extent of dealumination increases with the treatment temperature but the crystallinity and the pore structure are preserved. However, under hydrothermal conditions, extensive dealumination takes place even at 550 ~ and the structure collapses above 650 ~ Another investigation of the thermal and hydrothermal stability of ion-exchanged A1-MCM-41 (Si/A1 = 39) was carried out by Ryoo et al. [127] using dry or water vapor saturated 0 2 for 2 h at different temperatures. Under both sets of conditions, the stability was found to depend on the nature of the counter cations in the following order: y 3 + _ Ca2+ > Na+ _ as-prepared A1-MCM-41 > pure silica MCM-41. Table 2 Thermal and hydrothermal stability of AI-MCM-41 (Si/A1 = 26) (a) Temperature (b)

Thermal treatment

Hydrothermal treatment

oC

(Sa~/g)

pore vol. (c) (cc/g)

AI(T) (a) (%/

S (ma~/g)

pore vol. (e) (cc/g)

550 650 750 850

1465 1230 1252 1320

1.42 1.16 1.13 1.03

92 61 49 45

974 1005 231 170

0.78 0.75 0.30 0.22

(a) From ref. (132); (b) 3 h in dry air for thermal treatment and in water vapor for hydrothermal treatment; (c) pore volume according to the MP-method; (d) tetrahedral aluminum (%). As for the acidity of A1 containing MCM-41, ammonia TPD data indicate that it is comparable to that of amorphous silica-alumina, and much lower than the acidity of zeolites such as USY or H-mordenite [115,120,134]. This is consistent with Raman, FTIR and 29Si NMR data which indicate that despite their long range order, M41S mesoporous silicates and aluminosilicates exhibit essentially amorphous walls [48,49,115].

2.4.2. Mesoporous Titanosilieates Titanium modified mesoporous silicates were first mentioned in the patent literature [135]. In 1994, several research groups reported independently on the synthesis and characterization of Ti-MCM-41 [136-138] and Ti-HMS [139]. Corma et aL [136,140] prepared Ti-MCM-41 under hydrothermal conditions (413 K, 28 h) using fumed Aerosil silica, TMAOH (25 %), C16TMAOH/Br and Ti ethoxide. Similar preparations were also used by other workers [137,138,141]. Pinnavaia et al. [139,142] prepared Ti-HMS at room temperature using long chain primary amines instead of charged surfactants. More recently, several other groups reported on Ti-HMS [143] and Ti-MCM-41 [144,145]. In a detailed investigation Gontier and Tuel [143] studied the effect of several parameters on the synthesis of Ti-HMS at room temperature. This includes the following: (i) presence of isopropyl alcohol, (ii) synthesis time, (iii) length of the surfactant carbon chain, (iv) nature of the Ti source, (v) surfactant/SiO 2 ratio. Maschmeyer et al. [145] grafted Ti species on the pore walls

18 of MCM-41 silicate using titanocene solutions in chloroform. They obtained materials with very high density of accessible Ti sites. Bugshaw et aL [72] prepared Ti-MSU-1 in the presence of a non-ionic polyethylene oxide template. Kim et al. [146] used non aqueous formamide medium to control the Ti alkoxide hydrolysis. They were able to prepare mesoporous titanosilicate samples with very high Ti content. The XRD patterns of Ti containing samples consist of a main peak at 2 0 < 3 ~ corresponding to the 100 diffraction, sometimes accompanied by much weaker 110, 200 and 210 reflections in the 2 0 range of 4 to 7 ~ HMS based catalysts exhibit only the 100 diffraction peak because of excessive broadening of the hk0 reflections due to too small scattering domain sizes [69,71,139,142] or more likely to the presence of a poorly ordered pore system [55]. Indeed, TEM studies indicated that the pore structure of Ti-HMS is much less ordered than that of Ti-MCM-41 [147]. In addition, SEM [143] and TEM [147] show that Ti-HMS is comprised of spherical particles with 0.2-0.3 lam in diameter. N 2 adsorption isotherms obtained by Pinnavaia et al. [71,139,142] showed that in addition to the frameworkconfined mesoporosity due to the presence of parallel channels, HMS materials display a weU developed textural mesoporosity. However, other workers found that the N 2 adsorptiondesorption isotherm of Ti-HMS is reversible [143,147]. The pore distribution was broader for HMS as compared to MCM-41 materials. The Ti-MSU-1 material also exhibited only the 100 diffraction peak due to the occurrence of disordered, hexagonal-like packing channels [72]. These samples displayed reversible N 2 adsorption-desorption isotherms with no hysteresis. Several authors used diffuse reflectance UV-Visible spectroscopy to probe the local environment of Ti sites [136-138,140,141,143,147]. Sayad et al. [137,141,147] found that both Ti-HMS and Ti-MCM-41 exhibit an absorption band at 220 nm with no indication of a band at 330 nm characteristic of anatase [148]. There was however a weak shoulder at 270 nm, particularly for Ti-rich samples. Corma et al. [ 140] also found a band at 205-220 nm and a shoulder at 270 rim. The absence of anatase was also conf'mned by the absence of the characteristic 140 cm "1Raman band [138,149]. The band at 220 nm was attributed to isolated framework Ti species in interaction with water molecules [150]. The 270 nm band was assigned to partially condensed hexacoordinated Ti species belonging to a silicon rich amorphous phase [155]. FTIR spectra of Ti modified crystalline mesoporous materials (Figure 13) display a band at ca. 960 cm 1 [136,140,141,144,147]. A similar band was also found in FTIR spectra of Timodified zeolites such as TS-1 [156,157], Ti-13 [151-153], Ti-ZSM-48 [158,159] and others. It was attributed to the stretching mode of at least five different entities [ 160]: (i) SiO4 units bonded to a titanium ion (Si-O...Ti) [148], (ii) titanyl (Ti=O) groups [161], (iii) Si-O in SiO...H groups [162], (iv) T i - O in TiO4 tetrahedra [163], and (v) Si-O in SiOH...(OH)Ti defectives sites [164]. However, the presence of the 960 cm -1 band in FTIR spectra of TiMCM-41 and Ti-HMS samples may not be regarded as a firm proof for Ti incorporation because Ti-free silicates also exhibit a similar band. Nevertheless, the intensity ratio of the 960 to the 800 cm "1 band (due to the symmetric stretching vibrations of SiO4 [41]) was significantly higher for Ti containing samples than for pure silicates. The relative enhancement of the 960 cm "1 band may be attributed to Ti incorporation [133]. X-ray photoelectron spectroscopy (XPS) has also been used to gain insight into the local environment of Ti sites. Interestingly, tetrahedral Ti(IV) has a (2P3/2) binding energy (BE) more than 1 eV below that of octahedral Ti(IV) [151,165-169]. Using a BE of 103.3 eV for Si(2p) as internal reference, the BE of Ti(2P3/2) in both Ti-MCM-41 and Ti-HMS was found

19 to be 459.9 +_ 0.1 eV (Figure 14) compared to 458.6 eV for Ti in octahedral coordination as in TiO 2 [141,147]. Several groups [144,145,170,171] used X-ray absorption techniques to characterize Ti sites in Ti-MCM-41. Their data were also consistent with the occurrence of Ti species in tetrahedrally symmetrical environment.

459.9

B

&, 9

1200

t

| |

1000

I

I !

800

I

I I

"

600

Wavenumber, c m

i

400 41 ;4

"1

Figure 13. Infrared spectra of Ti-HMS samples [147]. A, B, C, D, E correspond to samples with Si/Ti ratios of .o, 145, 76, 39 and 22, respectively.

459

464

469

Binding energy, eV Figure 14. XPS spectrum of Ti2p in Ti-HMS, Si/Ti = 76 [141,147].

2.4.3. Mesoporous Vanadosilicates Periodic mesoporous vanadium modified silicates (V-MCM-41 and V-HMS) were synthesized, characterized and tested in our laboratory [93,141,172-174]. V-MCM-41 samples were prepared hydrothermally at 373 K using Cab-O-Sil fumed silica, vanadyl sulfate and dodecyltrimethylammonium bromide [172]. The preparation of V-HMS was carded out at room temperature in the presence of dodecylamine as template [173,174]. As in the case of V-free Ti containing samples, the XRD patterns of V-MCM-41 consisted of a main 100 peak and weak 110 and 200 peaks, while those of V-HMS exhibited only the 100 diffraction peak. N 2 BET adsorption isotherms with theft characteristic sharp step were fully reversible with no hysteresis loops [175]. TEM showed that V-MCM-41 particles have a highly ordered porous structure with some typical hexagonal crystals [38,93,173,176]. However, as other HMS materials, V-HMS consisted of globular particles of 0.1 to 0.4 lain in diameter with disordered channel packing [173,175]. Diffuse reflectance UV-visible spectra of V-HMS exhibited two bands in the range of 373-385 nm and 252-272 nm. Similar bands were observed at 384 and 265 nm for V-

20 silicalite-1 [ 177]. The charge transfer (CT) band at ca. 380 nm was attributed to V+5 with a short V=O bond and three longer V-O bonds, possibly in interaction with water vapor. The ca. 265 nm band was assigned to V +5 in a tetrahedral environment [175,177]. 51V NMR is one of the most suited techniques for characterizing V sites [141,172-174]. Typical 51V NMR V-M CM-41-60 V-HM S-60 spectra are displayed in Figure 15. Table 3 shows relevant NMR parameters for our samples and some closely related systems. As seen in Fig. 15, dehydrated samples exhibited signals with significant anisotropy of the chemical shift, the largest anisotropy being for V-HMS samples (Table 3). The isotropic chemical shift for dehydrated samples was in the range of-650 to -720 ppm. b Adsorption of water brought about a ~k downfield shift of the isotropic component by 100 to 150 ppm and affected the anisotropy significantly. In C the case of V-HMS, the hydration led to a large anisotropy increase, while it had -50O -1000 o -5oo - ;oo no effect on the parameter of asymmetry. The parameters of anisotropy for V-HMS Chemical shift, ppm were very close to those reported for (SiO)3V=O species in amorphous Figure 15. 51VNMR spectra of V-MCM-41-60 V/SiO 2 [178,179], indicating that V and V-HMS-60 (ref. 173). a: static NMR species in V-HMS are similar to those of dehydrated samples; b: static NMR of in V/SiO 2. Scheme 1 is a simplified hydrated samples; c: MAS NMR of hydrated representation of V site in V-HMS and samples, the rotation speed was 3.3 kHz its interaction with water. This proposal for V-MCM-41-60 and 11.2 kHz for is in full agreement with the UV-visible V-HMS-60, (*) side bands. data already discussed. In contrast, exposure of V-MCM-41 to water vapor led to almost complete disappearance of the anisotropy. Two nearly isotropic lines at-508 and -527 ppm could be resolved. The small value of anisotropy reflects the high degree of symmetry of the V species. The "symmetrization" of the V environment upon exposure to water vapor is believed to be due to interaction with OH groups as shown in Scheme 2. The active participation of hydroxyl groups and water in the stabilization of vanadium in the molecular sieve "framework" was further substantiated by 29Si MAS NMR data [173].

21 Contrary to V modified zeolites [180-184], as-synthesized mesoporous vanadosilicates did not exhibit any EPR signals at 77 K [173,174]. This was attributed to the presence of V 4+ in highly symmetrical lattice positions. Because of the electronic degeneracy and the associated very short relaxation times, lower temperatures maybe needed for the observation of such species. The presence of V +5 or clustered V+4 in the as-synthesized materials was excluded based on NMR data [173]. Recently, Tuel and Gontier [185,186] also reported on the preparation and characterization of V-HMS. At variance with our data, their as-synthesized samples exhibited an anisotropic hyperfine EPR signal at room temperature. The main conclusion of this study is that Si--O--V bonds form during calcination leading to isolated and tetrahedraUy coordinated vanadium centers which interact readily with water vapor. Morey et al. [187] prepared a series of V/MCM-48 samples by impregnation using dry hexane solutions of O=V(OiPr) 3. Their 51V NMR data for hydrated and dehydrated samples were almost identical to those reported for V-HMS [173,174], suggesting that in the absence of water, vanadium sites consist of pseudotetrahedral (SiO)3V=O grafted on the channel walls. Interactions of these species with water were also investigated by UV-visible.

2.4.4. Other Mesoporous Metallosilicates Boron was successfully incorporated in the "framework" of MCM-41 silicate in our laboratory [93,99,188]. In addition to XRD, N 2 adsorption and TEM, samples were thoroughly characterized by liB and 29Si MAS NMR. In as-synthesized materials, all boron was found to be in the four-coordinated state up to a B/Si ratio of 16%. 29Si MAS NMR data indicated that the attachment of the boron to the lattice takes place through interaction with structural hydroxyl groups. Careful calcination under dry conditions transformed a significant fraction of the boron into trigonally coordinated species with complete retention of boron in the framework up to B/Si = 8%. Calcination of samples with higher B loading generated extraframework species. Exposure of calcined samples to moist air at room temperature led to partial deboronation by hydrolysis. Boron-containing silicates were also studied briefly by other workers [130,189]. Several authors [129,130,190] reported on Fe modified MCM-41 and HMS. Yuan et al. [190] interpreted their FTIR and EPR data on the basis of Fe incorporation in the silicate "framework". Tuel and Gontier [130] found that the EPR spectrum of Fe-HMS is comprised of three signals one of which corresponds to iron-substituted framework sites. A number of literature reports dealt with Ga-modified mesoporous silicate molecular sieves [129,130,191,192]. Cheng et al. [191] synthesized Ga-MCM-41 with Si/Ga from 10 to 120. The quality of the samples was very sensitive to the pH of the gel mixture. No extraframework gallium was detected by 71Ga NMR in the as-synthesized samples. However, during calcination, the 4-coordinated gallium was partially expelled from the structure for samples with Si/Ga < 20. Galloaluminosilicate [192] as well as Ti [193] ans manganese [ 194] modified mesoporous silicates were also synthesized and characterized. 2.5 Synthesis and Characterization of Non-Silica Mesostructured Materials Huo et al. [60,61,195] first extended the LCT strategy to the synthesis of non-silica-based mesostructures, mainly metal oxides. Both positively and negatively charged surfactants and inorganic species were used. It was found that a suitable metal oxide should have the following characteristics: (i) depending on the formation mechanism, the ability to form polyanions or polycations allowing multidentate binding to the surfactant, (ii) the polyanions

22 Table 3 Parameters of 51V NMR spectra of V-MCM-41 and V-HMS (a'b) Samples

Si/V

Anisotropy A~5, ppm _ 10

Parameter of asymmetry, rl

~istat,

-650 -527 -508 -660 -527 -720 -600 -715 -590

V/MCM-41-D V/MCM-41-H

60 60

-320 --50

0.3 - 0

V/MCM-41-D V/MCM-41-H V/HMS-D V/HMS-H V/HMS-D V/HMS-H

145 145 60 60 124 124

-315 --50 -475 -640 -480 -640

0.3 - 0 0.15 0.15 0.15 0.15

Ref. (178) [(C6H11)7(Si7012)VO] 2 OV(OSiPh3) 3 WSiO2-D WSiO2-H

-398 -422 -487 -620

-

ppm +_. 10

0.05 0.05 0.05 0.13

~iiMAs, ppm _ 3 -665 -527

-530 -708 -580 -711 -576 -714 -736 -710 -609

(a) From ref. (173);

la33-ai I---iaee-ai I, I~ilX-~ii[, 15i=1/3(~i11+~i22+~33),Afi=~i33-~Si,rl=l~i22-~ill I/la33-ai (c) D: dehydrated samples, H: hydrated samples. (b)

O

O

Io

.2o

Io

7S2\oCe..~

c7S2 ",,or 0

I~'~H

Scheme 1

t o,,~ ~,o-

,-..

p.-"-~ o

v / ~,,,.. o~o.--,~

H20

. . . . . .

"--

Scheme 2

X.._ _o:r "o -[ ".~r ' o - " s , /~

I;

23 or polycations should be able to condense into rigid walls, (iii) a charge density matching between the surfactant and the metal oxide is necessary to control the formation of a particular phase. As shown in Table 4 most of the metal oxides have a strong tendency to form lamellar structures, except Sb, W, Pb. The formation pathway of these materials depends on the charge of the surfactant and that of the inorganic ion involved in the synthesis. Different mesostructured Sb and W oxides were synthesized at room temperature by controlling the pH of the system. Regardless of their structures, all the resulting materials collapsed upon calcination. Stein et al. [196,197] independently explored the applicability of the use of surfactants for the synthesis of channel structures with transition metal oxide frameworks. Vanadium, niobium, molybdenum and tungsten oxides were studied. In the case of tungsten, hydrothermal reaction of sodium metatungstate with C16TMAOH gave the salt [C16H33N(CH3)a]6(H2W1204o). Despite the apparent similarity of TEM micrographs and XRD patterns of this material to those of mesoporous silicates, the salt contained unconnected Keggin ions H2W120406-. These Keggin ions pack in a puckered layer arrangement and create roughly spherical cavities for the suffactant micelle counterions. As a result, attempts to remove the template cations and condense the inorganic portion of the structure invariably led to dense WO3. x phases. As in the case of lamellar MCM-50 [ 198], the authors prepared a stable salt-gel by reacting the surfactant niobotungstate salt with TEOS. During this treatment Nb-O-Si linkages were formed. Removal of the cationic surfactant by acidextraction resulted in porous structures with surface areas up to 265 mE/g. Following the pioneering work of Huo et al., Antonelli and Ying [ 199] prepared the first stable mesoporous transition metal oxide, TiO 2, using a modified sol-gel method. They showed that the key step to obtain a desired phase is to control the hydrolysis rate of the organometallic precursor, titanium tris-isopropoxide. Acetylacetone was added to the system to stabilize the titanium compound and to lower the hydrolysis rate. A hexagonal phase was obtained only in the presence of phosphate surfactant. After calcination at 627 K, hexagonally packed TiO 2 prepared in the presence of tetradecyl phosphate had a BET surface area of 200 m2/g and a pore distribution centered at 32 A. However, IR spectroscopy revealed a strong absorption band at 1087 cm-1, indicating the presence of phosphate ions, even after calcination. Luca et al. [200] prepared mesostructured vanadium oxide. In their synthesis, cetyltrimethylammonium vanadate was first crystallized from water solution. It was then dissolved in alcohol followed by titration with HC1 to pH = 2.2. A hexagonal phase was identified by XRD. However, the material was thermally unstable. Abe et al. [201] synthesized a hexagonal vanadophosphate using a similar templating method under hydrothermal conditions. Characterization of the resulting hexagonal phase by IR and XRD showed that the inorganic part was basically amorphous and was similar to glass V2Os-P205. Sayari et al. [202,203] extended the LCT technique to the synthesis of mesostructured zirconium oxide. The use of long chain quaternary ammonium salts or primary amines as templates led to the formation of hexagonal and lameUar ZrO 2 phases, respectively. Zr(SO4) 2 was used as zirconium source, which provided a highly acidic medium, pH < 1.5. Consistent with the synthesis conditions and EDX analysis data a S+X-I+ mechanism where the suffactant-inorganic interaction is mediated by sulfate anions was proposed. Unfortunately, both structures collapsed upon removal of the surfactant either by high temperature calcination or by solvent extraction [203]. However, the hexagonal form was successfully

24 Table 4 Typical Synthesis Results Using Different Inorganic Precursors and Surfactants inorganic precursor

surfactant

phase

XRD d spacing (A)*

Sb oxide Sb oxide

ClgH37(CHa)aNBr C 18H37(CH3)3NBr

cubic (Ia3d) hexagonal

42.9 46.0

Sb oxide

ClgH37(CHa)3NBr

lamellar

37.5

W oxide

CId-I3s(CH3)3N-Br

hexagonal

40.0

W oxide

C16Hs3(CH3)3NBr

lamellar

28.3

zinc phosphate

C.I-I2.§

lameUar

21.6(10), 23.5(12), 26.0(14), 28.2(16), 30.5(18), 32.5(20)

alumina

C 12H25C6I-I4SOaNa

lamellar

28.9

Pb 2+ Pb 2+

C 1d-IssSOsH C 1d'Iss SOsH

hexagonal lamellar

45.8 38.5

Fe 2+

C 16I-I33SOaH

lamellar

41.0

Mg 2+

C12H25OPO3H2

lamellar

31.0

Mn 2+

C 12H25OPOsH2

lameUar

28.6

F es*

C 12H25OPOaH2

lamellar

26.9

Co2+

C 12H25OPO3H2

lamellar

3 0.8

Ni 2+

C12H25OPOsH2

lamellar

31.1

Zn 2+ .Al 3+

C 12H25OPOsH2 C 12H25OPOaH2

lamellar lamellar

29.6 26.4

Gas+

C 12H25OPOsH2

lamellar

27.2

Fe 2§

CnI-I2n+lOSO3Na

lamellar

21.0(12), 23.0(14), 27.3(16), 30.3(18)

Fe s§

C~H2~§

lamellar

23.1(12), 26.0(14), 28.1(16), 28.1(18)

Co 2+

C,H2,+IOSOsNa

lamellar

20.9 and 39.7(12), 22.8(14), 41.5 and 27.4(16), 28.4(18)

Ni 2+

C,H2,+~OSOaNa

lamellar

31.8, 23.5 and 23.2(14), 43.5 and 27.5(16), 24.3(18)

Mn 2§

C~-I2,§

lamellar

23.3(14), 42.2 and 28.9(16), 24.3(18)

tin sulfide

Ca6H33(CH3)3NBr

lamellar

25.8

* Carbon numbers of the surfactant chains are shown in parentheses.

25 stabilized by post-synthesis treatment with potassium phosphate at room temperature followed by air calcination at 627 K. The final product contained significant amounts of phosphorous, and had a surface area exceeding 500 m2/g but with a broad pore size distribution [202]. Knowles and Hudson [204] also prepared a mesoporous, high surface area zirconium oxide in a basic medium. At pH = 11.4-11.7, zirconium species formed a gel. Through a scaffolding process followed by calcination at 723 K, they obtained zirconium oxide with surface area in the range 238-329 m2/g. The d-spacing of as-synthesized compounds did not change with the length of the surfactant used, which is inconsistent with the LCT approach. An alternative mechanism was proposed. The positive surfactant ion first exchanges with the ions in zirconium hydroxide gel. Subsequently, controlled heating and scaffolding condense the inorganic structure, and calcination leads to a mesoporous, large surface area material. Pinnavaia et al. [69,72] prepared a hexagonally packed alumina through the neutral templating approach. In the presence of a polymer surfactant, (PEO)13(PPO)30(PEO)13, an alumina with a d spacing of 63/~ and a surface area of 420 m2/g was obtained [72]. It was also mentioned that non-layered alumina can be synthesized using octyl or dodecyl amine as template and a neutral aluminum alkoxide precursor. Compared with metal oxides, less attention has been paid to the synthesis of mesostructured metal sulfides [61,205]. The only systematic work was reported by Anderson and Newcomer [205]. The liquid-crystal templating approach was applied to metal sulfides, such as Mo, W, Co, Fe, Zn, Ga, Sn and Sb sulfides. All of the products were lameUar and consisted of bilayers or interdigitated layers of surfactant molecules sandwiched between metal sulfide layers. As mentioned in the introduction, microporous aluminophosphates form a large and important family of molecular sieves. With the ultimate goal of preparing stable ultra-large aluminophosphates, two research groups attempted to use a supramolecular surfactant array as template. Oliver et al. [206-208] reported on the synthesis of lamellar aluminophosphates in non aqueous medium. A typical synthesis gel was: 14 TEG : 0.9 A1203 : 2.5 H20 : 1.8 P205 : 3.0 CloH21NH2 , where TEG denotes tetraethylene glycol. SEM showed that in some parts of the resulting lamellar materials, there exists micrometer-scale surface patterns, including bowl, honeycomb and quilted shapes with superimposed finer columnar, sphere, mesh, and pore-like structural features. The origin of these pattems was explained based on a proposed vesicle templating mechanism along with a "cellular" model (Figure 16). The multi-functional role of TEG was emphasized. It acted as a solvent for the surfactant, a polydentate ligand to Al3+, a co-surfactant to control the bilayer curvature, and a demixing agent to promote surfactant-TEG phase separation and patterning of vesicle bilayer, and ionchannels to facilitate the transport of (TEG)A11u ionophores through vesicle bilayers, to access reactive phosphate sites and permit mesolamellar aluminophosphate nucleation and growth. The micrometer scale patterns obtained mimic some natural microskeletons such as diatom and radolarian. Sayari et al. [209-212] followed another path to the synthesis of mesolameUar aluminophosphates. They used AI20 3, H3PO4 and primary or tertiary amines as surfactants in aqueous media. The surfactant was found to be protonated while acting as template. Effects of synthesis parameters such as A1/P, P/amine, P/H20 ratios were studied systemically by XRD and 27A1 and 31p NMR. The connectivity between A1 and P was found to be dependent on the synthesis parameters.

26

Figure 16. Model for vesicle and bilayer control of macroscopic morphology and surface patterning, and mesolamellar structure of the aluminophosphates [206]. TEM studies [210] indicated that in some samples apart from the simple lameUar packing, there were extended areas comprised of a unique hexagonal-like packing of alternating concentric dark and bright tings (Figure 17). As shown schematically in Figure 18, these circular fringes were interpreted as the edge projections of cylindrical layers of inorganic A1PO4 materials separated by cylindrical vesicles of surfactant, all wrapped around a single rodlike miceUe. This new mesophase has no equivalent among known surfactant liquid crystal phases.

Figure 17. TEM Image of as-synthesized AIPO4 [210].

Figure 18. Representation of coaxial cylindrical bilayer growth of A1PO4 [210].

27 3.

C A T A L Y T I C APPLICATIONS OF M41S AND RELATED MATERIALS

Most potential applications reported in the literature use MCM-41, HMS or FSM-16 type of materials. Patents dealing not only with the applications, but also with the synthesis of crystalline mesoporous materials are almost exclusively assigned to Mobil Oil Corporation. Extensive lists of these patents are provided as appendixes. Catalytic applications may be conveniently divided into three categories: (i) acid catalysis, (ii) liquid phase redox catalysis, and (iii) other applications. 3.1 Acid Catalysis Acid sites in mesoporous silicates can be generated either by isomorphous substitution of trivalent cations such as A1 or B for Si, or by adding an acidic component such as a heteropolyacid, an ultra stable Y (USY) or a A1 containing ZSM-5 zeolite. 3.1.1. Mesoporous Aluminosilieates A1-MCM-41 based materials were tested in a number of petroleum ref'ming processes. A NiMo impregnated Al-MCM-41 catalyst (12 wt% MoO 3, 3 wt% NiO) was tested for hydrocracking of vacuum gasoil, and found to be more efficient in hydrodesulfurization and hydrodenitrogenation than NiMo loaded on USY or on amorphous silica-alumina. The higher performance of NiMo/MCM-41 was attributed to its high and freely accessible surface area and also to the higher dispersion of catalytically active ingredients. In addition, despite its lower acidity, the NiMo/MCM-41 catalyst was also found to have higher activity in mild hydrocracking of gasoil than USY or amorphous silica-alumina based catalysts. Other investigations reported in the open literature include microactivity tests (MAT) [214], oligomerization of propene [129], cracking of cumene [134] and bifunctional hydroismerization and hydrocracking of n-hexadecane [215]. Potential catalytic applications of MCM-41 based catalysts in the petroleum refining industry were also reported in the patent literature (Appendix 2). The hydrogen form of A1MCM-41 mixed with A1203binder in a ratio of 65 to 35 wt% was used for the cracking of a straight run naphtha at 540 ~ and ca. 3 atm [216]. At the same conversion (43-45 %), the MCM-41 based catalyst produced more C3-C5 olefins (74 vs. 54 %) and much less light gas and linear hydrocarbons (11 vs. 29 %) than medium pore ZSM-5 zeolite. In addition, MCM-41 exhibited higher selectivity towards valuable isobutane and isopentanes which can be further upgraded via alkylation by olefins or via dehydrogenation into isoalkenes. A catalyst comprised of 35% Al-MCM-41 and 65 % silica-alumina-kaolin clay matrix was also tested in fluid catalytic cracking and found to be more active and more gasoline selective than a catalyst containing 35 % USY. It also displayed higher selectivity towards C5 olefins. Apelian et al. [218] compared an A1-MCM-41 based catalyst (5.8% Ni, 29.1% W on 65% alumina, 35% A1-MCM-41) with a fluorided NiW/A1203 for the hydrocracking of a heavy wax. At low conversion (< 50%), both catalysts had similar lube yields, but at higher conversion NiW/MCM-41 exhibited higher yields. Furthermore, the MCM-41 based catalyst gave lubes with higher viscosities, thus allowing operation at a much higher wax conversion while still meeting viscosity specifications. NiMo/H-MCM-41/A1203 catalysts were found to be effective for the removal of heavy metals (Ni, V, As, Fe) and asphaltenes from resid and shale oil under relatively mild conditions [219]. Combination of the mildly acidic A1-MCM-41 with a strongly acidic zeolite such as USY in the presence of an alumina binder in addition to nickel and tungsten leads to an enhanced overall hydrocracking activity of the catalyst and a decreased yield of light gas, compared to USY free catalyst [220].

28 Protonated A1-MCM-41 catalysts were also used for the cracking of olefmic feedstocks such as FCC gasoline. They gave high selectivities towards valuable isobutene and isoamylenes which can be further upgraded into high octane components via etherification with methanol [221]. Other patent applications dealt with oligomerization of propene [222], dealkylation of 1,3,5 tri-tert-butylbenzene [223], alkylation of naphthalene with long chain alpha-olefins [224], and alkylation of benzene with ethylene [225]. Isobutane alkylation with butene over H2SO4 or BF3 promoted H-MCM-41 was also investigated [226]. In all these processes, H-MCM-41 based catalysts exhibited promising performances, and in some cases may be considered as a viable alternative to currently used commercial catalysts. Armengol et al. [227] used protonated A1-MCM-41 molecular sieve for alkylation of bulky aromatic compounds such as 2,4-di-tert-butylphenol with a bulky alcohol (cinnamyl alcohol). This reaction did not occur in the presence of large pore HY zeolite indicating the importance of the mesoporous structure of the H-MCM-41 catalyst and the accessibility of active sims. Kloetstra et al. [228] obtained excellent results during the tetrahydropyranylation of alcohols and phenols over A1-MCM-41 (Scheme 3). Bulky alcohols including cholesterol, adamantan-l-ol and 2-naphthol were converted into the corresponding tetrahydropyranyl ethers in relatively short periods of time.

Scheme 3 Shinoda et al. [229] found that FSM-16 mesoporous aluminosilicate compares favorably with the currently used liquid acid BF3.OEt2 for the synthesis of meso-tetraarylporphyrins from the corresponding aromatic aldehydes and pyrrole. In addition, contrary to soluble catalysts and to K10 acid treated montmonrillonite, FSM-16 could be used repeatedly after regeneration by calcination at 500 ~ Kloetstra [230] used Na§ and Cs+ exchanged A1-MCM-41 catalysts to carry out a base catalyzed reaction, namely the Knoevenagel condensation of benzaldehyde with ethyl cyanoacetate (Scheme 4).

O

H + H2C\ /CNco2Et

-H20

= ~~CO2Et ~

CN

Scheme 4 In the presence of Na-MCM-41 as catalyst, and water as solvent, almost 100 % selectivity was achieved at 90% conversion. However, this reaction did not take place in ethanol. Instead, benzaldehyde and ethanol reacted over residual acid sites giving equilibrium concentration of diethyl acetal. In addition, both H-MCM-41 and Na-MCM-41 catalyzed the condensation of benzaldehyde with acetophenone to chalcone (Scheme 5a with R = H) and

29 other aldol condensations of bulky molecules. They were also found to catalyze intramolecular Michael addition of the a-unsaturated ketones to flavones (Scheme 5a and 5b with R = OH). The strongly basic Cs-MCM-41 was found to be an effective catalyst for the Michael addition of chalcone and diethyl malonate. o

o

R ~

o

-~~

(b)~R = OH

o

Scheme 5

3.1.2.

Heteropolyacid (HPA) Supported MCM-41 Silicate Catalysts

To enhance the acidity of A1-MCM-41 Kozhevnikov et al. [231] added 10 to 50 wt% phosphotungstic acid. The obtained catalysts were found to be more efficient than H2SO4 or bulk phosphotungstic acid in liquid phase alkylation of 4-tert-butylphenol by isobutene and styrene. Using four test reactions, Kresge et al. [232] compared phosphotungstic acid loaded MCM-41 to more conventional catalysts. The reactions were: (i) n-butane conversion, (i.i) nhexane conversion, (iii) alkylation of isobutane with 2-butene, and (iv) alkylation of benzene with 1-tetradecene. For the first reaction, at similar n-butane conversions, HPA/MCM-41 gave much higher isobutane selectivities than ZSM-5. It was reported that n-hexane conversion and the isomerization selectivity over HPA/MCM-41 are significantly higher than those obtained in the presence of the unsupported ammonium salt of phosphotungstic acid [233] or the supported salt on silica or alumina [234]. The higher dispersion of HPA on MCM-41 may be at the origin of these differences. However, HPA/MCM-41 displayed lower isobutane alkylation activity than unpromoted H-MCM-22 zeolite. In addition, the ratio of trimethylpentanes to dimethylhexanes was low indicating that the quality of the alkylate was poor. As for the alkylation of aromatics, both benzene and tetradecene conversions were higher for the supported than unsupported phosphotungstic acid.

3.2 Liquid Phase Redox Catalysis Since the discovery in the early 80's of the remarkable catalytic activity of Ti-modified silicalite-1 (TS-1) in the selective oxidation of organic substrates by dilute H20 2, the field of transition metal modified zeolites grew tremendously as shown in a number of recent reviews [156,235,236]. In addition to its hydrophobicity, the major role of the zeolite matrix is the stabilization of isolated redox centers. However, the limited accessibility of these sites precluded the use of large substrate molecules. The discovery of crystalline mesoporous silicate was immediately perceived as an ideal solution to these limitations.

30 Ti, V and Sn-modified mesoporous silicates were reported to be active in a number of liquid phase oxidation reactions. Ti-containing samples were used for the selective oxidation of large organic molecules in the presence of ten-butyl hydroperoxide (TBHP) or dilute H20 2 [71,136,137,139-141,147,186,237]. Typical data shown in Table 5 indicate that both TiMCM-41 and Ti-HMS are efficient catalysts for the epoxidation of bulky olefms such as ctterpineol and norbomene in the presence of TBHP or H20 2. Comparison with Ti-13 indicates that the accessibility of active sites plays a critical role in the liquid phase oxidation of organic molecules. Mesoporous titanosilicates also exhibited remarkable activity in the hydroxylation of 2,6-di-tert-butyl phenol (2,6 DTBP) [142,147] and the oxidation of cyclododecanol [147], naphthol [147] aniline [237] and chloroaniline [186]. However, they were disappointingly poor catalysts for the liquid phase oxidation of n-hexane and aliphatic primary amines, as well as the ammoximation of cyclohexanone [ 147,238]. Table 5 Oxidation of tx-Terpineol (a) and Norbomene (a'b) over Ti-MCM-41 and Ti-HMS Catalyst

Time

o~-terpineol

Time

Norbornene

Ref.

(h)

epoxide

others

(h)

epoxide

alcohol

TiA1-MCM-41 (c)

3 8

23.8 31.5

4.0 8.6

5 11

26.4 42.3

3.1 6.4

140 140

Ti-13(el)

3 8

4.1 7.6

2.5 5.8

5 11

10.3 18.3

6.6 12.8

140 140

Ti-MCM-41 (e)

-

-

-

2

5.6

10.2

147

Ti-HMS 0)

-

-

-

2

16.4

3.6

147

(a) Reaction conditions in Ref. 140: temperature (313 K), catalyst (100 mg per mmol of substrate), TBHP in CH2C12 to olefin ratio: 1.2; (b) reaction conditions in Ref. 147: temperature (335 K), catalyst (100 mg), norbomene (100 rag), 30 wt% H20 2 (2.36 g), acetonitfile (10 ml); (c) Si/A1 = 97, Si/Ti = 55; (d) Si/Ti = 60, Si/A1 = ca. 200; (e) Si/Ti --49; (j) Si/Ti =76. Corma et aL [239] took advantage of both redox and acidic properties of Ti/A1-MCM-41 (Si/Ti = 68 and Si/A1 = 196) aluminosilicate to carry out the multistep oxidation of linalool to cyclic furan and pyran hydroxy ethers (Scheme 6). The reaction was carried out in acetonitfile at 353 K using TBHP as oxidant. Conversions as high as 80 % were obtained. As shown in Scheme 6, it was postulated that the reaction takes place via epoxidation over Ti sites followed by acid catalyzed intramolecular opening of the epoxide ring by the 3-hydroxy group. Ti-13 zeolite gave somewhat lower conversions in addition to the preferential formation of furans over pyrans (ratio of ca. 1.5) due to shape selectivity. Ti-MCM-41 and gave furan to pyran ratios of ca. 0.9, comparable to those obtained by the epoxidase conversion of linalool.

31

.

TBHP

,..

3a

H

~ 2b

H

3b

Scheme 6 Sayari et al. [ 172,174] found that like their Ti containing analogs, V-HMS and V-MCM41 are efficient catalysts for the hydroxylation of bulky aromatic molecules such as naphthol and 2,6 DTBP. Gontier et al. [186,237] found that V-HMS has no activity in the oxidation of aromatic amines in the presence of H20 2. However, it oxidizes aniline selectively into nitrobenzene when TBHP and acetonitrile are used as oxidant and solvent, respectively. Notice that Ti-HMS is active even in the presence of H20 2 and gives a different product distribution, particularly azoxybenzene [237]. Das et al. [193] also reported that Sn-MCM-41 is effective in the liquid phase hydroxylation of phenol and naphthol. Geometric constraints and related factors including active site accessibility, steric effects of transition states and diffusion limitation of reactants and products play a crucial role in liquid phase catalyzed reactions [240]. Several examples are presented hereafter for illustration: (i) TS-1 is more active in the oxidation of linear vs. branched alcohols [241], and in the epoxidation of linear vs. cyclic olefins [153]; (ii) in the presence of TBHP, TS-1 has no activity [242], while Ti-g is less active than Ti-MCM-41 [243]; (iii) large TS-1 particles are less effective catalysts than smaller ones [244]. All the above mentioned observations stem from diffusion limitations and steric constraints due to reactants or transition-state intermediates. However, this is not the only factor that governs the activity of Ti sites in molecular sieves. Indeed when considering the epoxidation of a smaU linear olefin such as 1-hexene, the activity of Ti silicates decreases in the order TS-1 >> Ti-13 > Ti-MCM-41 [153,171,243]. Furthermore, the more demanding reactions such as the hydroxylation of benzene and the oxidation of linear alkane and aliphatic primary amines take place on TS-1 but not in the presence on Ti-13 or Ti-MCM-41. The reason for such a strong dependence of Ti catalytic properties on the nature of the silicate matrix is not well understood. Indeed, none of the spectroscopic techniques used so far including X-ray absorption [171,245] shows any difference in the local environment of Ti in various crystalline micro- and meso-porous silicates. Decreasing hydrophobicity from TS-1 to Ti-l~ and Ti-MCM-41 is a contributing factor in the parallel decrease of the intrinsic catalytic activity of Ti sites. However, as suggested earlier [137,141,147], subtle variations of some important properties such as bond angles or redox potential may be at the origin of this behavior.

32 3.3 Other Catalytic Applications Because of their extremely high surface areas MCM-41 and FSM-16 (alumino) silicates were used as supports for catalytically active materials. Bhore et al. [246] found that Ni supported HMCM-41 has higher propene oligomerization activity and improved selectivity towards trimers and tetramers than Ni on medium pore zeolites. Several authors used Pt/MCM-41 [215,247-249] and Pt/FSM-16 [250] catalysts prepared by impregnation or by ion-exchange. Del Rossi et al. [247] used 0.6% Pt/HMCM-41 for the conversion of n-hexane in the presence of H e. The isoparaff'm yield obtained at a given n-hexane conversion over Pt/H-MCM-41 was significantly higher than over 0.6% Pt on amorphous SiO2-A1203. In addition, the MCM-41 based catalyst gave much less cracked products. Inui et al. [250] used Pt supported FSM-16 mesoporous silicate downstream of H-Fe-silicate to convert in one step propene into high octane branched alkanes in the gasoline range. This catalyst showed much better hydrogenation activity than Pt on amorphous silica with a broad distribution of mesopores. 0.85% Pd supported on a 65% - 35% extruded mixture of HMCM-41 and A1203 showed higher benzene hydrogenation activity than other Pd containing catalysts such as Pd/USY, Pd/ZSM-5 and Pd/SiO2 [251]. However, the Pd dispersion was also higher for the MCM-41 based catalyst than the others. MCM-41 supported V205-TiO2catalysts (6.1% Ti and 2.5% V) were used for NO x selective reduction [252] in the presence of a gas stream containing 125 ppm NO, 125 ppm NH 3, 0.12% 0 2 in He. The MCM-41 based catalyst was found to be more active than V205TiO2/SiO2, particularly at high temperature. Several patents were also devoted to the oligomerization of 1-decene over Cr203 impregnated A1-MCM-41 catalysts [253-255]. The products obtained had much higher viscosities than those obtained over chromia on silica. Immobilization of transition metal complexes such as metaUoporphorin [256], Fe(II)penanthroline [257] and others [258] on the walls of MCM-41 silicates is an interesting recent development in this field. Chibwe et aL [256] ion exchanged A1-MCM-41 with meso-tetra(Imethylpyridinium) porphyrin COO/) complex, [Co(II)TMPyP]~ . At low loading, the turnover numbers of the supported complex in the oxidation of 2,6 DTBP by H20 2 were up to two orders of magnitude higher than for the homogeneous catalyst. This remarkable increase compared for example to the 5-fold activity enhancement obtained in the oxidation of methyl cyclohexane over Fe phtalocyanine loaded NaY [259] indicates that the improved accessibility of the active complex in MCM-41 materials plays a crucial role in the catalyst performance. Liu et aL [257] prepared Fe(II)-Phen on A1-MCM-41 (A1/Si = 20) using an alcohol solution of [Fe(Phen)3]C12 at room temperature. This catalyst was used repeatedly in benzene hydroxylation by 30% H202 without any significant loss of activity. Huber et al. [260] also encapsulated a tin-molybdenum complex, Me3SnMo(CO)a(rI-CsH5) in MCM-41 silicate. The material was characterized by EXAFS and FTIR-TPD measurements. The complex was found to attach strongly to the channel walls, and starts to decompose at ca. 200 ~ At 300 ~ sub-nanometer size Sn-Mo clusters were obtained which may have interesting catalytic properties. Helldng et al. [261] found also that A1-MCM-41 materials display a promising behavior as phase transfer catalysts. The rate of the two-phase reaction between potassium iodide and 1-bromopentane was enhanced significantly upon addition of A1-MCM-41.

33 4.

OTHER POTENTIAL APPLICATIONS

Periodic mesoporous materials may have important applications in the area of seperation of biological materials. The fabrication of composite and non composite membranes based on M41S silicates has been reported in the patent literature [262]. Another area with potential growth is the encapsulation of technologically advanced materials. Preliminary findings dealing with the following materials have been reported: election transfer photosensitizers [263] semiconductors [264] - polymer wires [265-268] conducting carbon wires [269] - sensing devices [270] - materials with non linear optic properties [271] - quantum sized clusters [271,272] -

-

-

5.

CONCLUDING REMARKS

M41S crystalline mesoporous materials have expanded the area of microporous zeolites and molecular sieves into the mesopore range. Their discovery has created new opportunities in several areas. First, by its simplicity and diversity the synthesis strategy aroused the interest of zeolite synthesis scientists in the rich chemistry of surfactant-inorganic systems. This effort has already led to several important findings, in particular (i) the design of new synthesis routes using cheap polymer surfactants, (ii) the discovery of new morphologies without lyotropic surfactant counterparts, and (iii) the synthesis of some stable non-silica based mesoporous materials. However many challenging tasks are yet to be accomplished. Pertinent examples include the development of efficient methods for the removal of templates from non-silica framework without altering the porous structure, and the the design and synthesis of three dimensional mesoporous cage-structured materials using cheap and readily available surfactants. In addition, crystalline mesoporous molecular sieves have promising properties particularly in catalysis. The main advantages of these materials are their extremely high surface areas and the great accessibility of their pore systems. Because of the amorphous nature of their pore walls, mesoporous aluminosilicates have much lower acidity than acid zeolites such as H-Mordenite, HY and ZSM-5. However, reactions that do not require very strong acid sites will benefit from the enhanced accessibility of acid sites and the high surface area of these materials. In addition, the acidity drawback may be mitigated by combination with strongly acidic ingredients such as heteropolyacids. More important than the limited strength of acid sites is the stability of mesoporous aluminosilicates in terms of both aluminum retention and structure preservation. Issues that are crucial in evaluating these materials for possible commercial applications include resistance to dealumination and mechanical stability under working conditions, as well as the long term activity maintenance. As for transition metals modified mesoporous molecular sieves, innovative applications in selective oxidation of bulky molecules relevant to pharmaceuticals and agrochemicals should be envisaged. In this context, strong collaboration between scientists working in (i) catalysis, (ii) the design and fabrication of crystalline mesoporous materials, and (iii) organic synthesis is essential to identify opportunities where heterogeneous liquid phase redox catalysis may be useful. In addition, despite many advantages of using solid catalysts for

34 liquid phase reactions, leaching of active ingredients may occur. It is worthwile to undertake an in-depth evaluation of this problem during repeated reaction - regeneration cycles. ACKNOWLEDGMENTS

The author is grateful to Y. Yang, P. Liu, C. Danumah and H. Michel and J. Desgagn, for their help with the manuscript. APPENDIXES

Appendix 1 Mobil US Patents on Synthesis of Mesoporous Materials Authors

Patent Number

Publication Date

Title

Beck et al. Calabro et al.

5334368 5308602

02.08.94 03.05.94

Beck et al. Kresge et al.

5304363 5300277

19.04.94 05.04.94

Beck et al. Kresge et al.

5264203 5250282

23.11.93 05.10.93

Beck et al.

5246689

21.09.93

Chu et al. Kresge et al. Kresge et al. McCullen et al.

5215737 5211934 5198203 5156829

01.06.93 18.05.93 30.03.93 20.10.92

Degnan et al.

5156828

20.10.92

Johnson et al.

5112589

12.05.92

Calabro et al.

5110572

05.05.92

Beck et al.

5108725

28.04.92

Chu et al.

5104515

14.04.92

Kresge et al.

5102643

07.04.92

Kresge et al. Beck

5098684 5057296

24.03.92 15.10.91

Synthesis of Mesoporous Oxide Synthesis of Crystalline Ultra-Large Pore Oxide Materials Porous Materials Synthesis of Mesoporous Crystalline Material Synthetic Mesoporous Crystalline Materials Use of Amphiphilic Compounds to Produce Novel Classes of Crystalline Oxide Materials Synthetic Porous Crystalline Material Its Synthesis and Use Synthesis of Mesoporous Aluminosilicate Synthesis of Mesoporous Aluminosilicate Synthetic Mesoporous Crystalline Material Method for Stabilizing Synthetic Mesoporous Crystalline Material Method for Manufacturing Synthetic Mesoporous Crystalline Material Method for Synthesizing Mesoporous Crystalline Material Using Acid Synthesis of Mesoporous Crystalline Material Using OrganometaUic Reactants Synthesis of Mesoporous Crystalline Material Method for Purifying Synthetic Mesoporous Crystalline Material Composition of Synthetic Porous Crystalline Material, Its Synthesis Synthetic Mesoporous Crystalline Material Method for Synthesizing Mesoporous Crystalline Material

35 Appendix 2 Mobil US Patents on Catalytic Applications of Mesoporous Authors

Patent Number

Publication Date

Title

Baker et al. A p e l i a n et al.

5468368 5451312

21.11.95 19.09.95

Beck et al.

5370785

06.12.94

Shih

5344553

06.09.94

Kresge et al.

5 3 2 4 8 8 1 28.06.94

Degnan et al.

5290744

01.03.94

Marler et al.

5288395

22.02.94

Degnan et al.

5281328

25.01.94

Marler et aL Pelrine et al. Borghard et al.

5277792 11.01.94 5270273 14.12.93 5264641 23.11.93

Apelian et al.

5264116

Bhore et al.

5 2 6 0 5 0 1 09.11.93

Aufdembrink

5258114

02.11.93

Del Rossi et al.

5256277

26.10.93

Le et al.

5232580

03.08.93

Apelian et al. Beck et al.

5227353 5200058

13.07.93 06.04.93

Kresge et al. Degnan et al. Degnan et aL

5196633 5191148 5191147

23.03.93 02.03.93 02.03.93

Lubricant Hydrocracking Process Catalyst and Process for Producing LowAromatics Distillates Hydrocarbon Conversion Process Employing a Porous Material Upgrading of a Hydrocarbon Feedstock Utilizing a Graded, Mesoporous Catalyst System Supported Heteropoly Acid Catalysts for Isoparaffin-Olefm Alkylation Reactions Hydrocracking Process Using Ultra-Large Pore Size Catalysts Production of High Viscosity Index Lubricants Hydrocracking with Ultra Large Pore Size Catalysts Production of Hydrocracked Lubricants Olefin Oligomerization Catalyst Aromatics Saturation with Catalysts Comprising Crystalline Ultra-Large Pore Oxide Materials Production of Lubricants by Hydrocracking and Hydroisomerization Catalytic Oligomerization Process Using Modified Mesoporous Crystalline Material Ultra Large Pore Cracking Catalyst and Process for Catalytic Cracking Paraffin Isomerization Process Utilizing a Catalyst Comprising a Mesoporous Crystalline Material Catalytic Process for Hydrocarbon Cracking Using Synthetic Mesoporous Crystalline Material Hydroprocessing Catalyst Composition Catalytic Conversion over Modified Synthetic Mesoporous Crystalline Material Catalytic Conversion Isoparaff'm/Olefin Alkylation Isoparaffin/Olef'm Alkylation

23.11.93

et al.

36 Appendix 2 (continued) Mobil US Patents on Catalytic Applications of Mesoporous Materials Authors

Patent Number

Publication Date

Title

L e et al.

5191144

02.03.93

Le Kresge et al.

5191134 02.03.93 5 1 8 3 5 6 1 02.02.93

Degnan et al.

5183557

02.02.93

Schipper et al.

5179054

12.01.93

Kresge et al. Beck et al.

5 1 7 4 8 8 8 29.12.92 5143707 01.09.92

Bhore et al.

5 1 3 4 2 4 3 28.07.92

L e et al.

5134242

28.07.92

L e et al.

5134241

28.07.92

Le Pelrine et al.

5118894 02.06.92 5 1 0 5 0 5 1 14.04.92

Olefin Upgrading by Selective Conversion with Synthetic Mesoporous Crystalline Material Aromatics Alkylation Process Demetallation of Hydrocarbon Feedstocks with a Synthetic Mesoporous Crystalline Material Hydrocracking Process Using Ultra-Large Pore Size Catalysts Layered Cracking Catalyst and Method of Manufacture and Use Thereof Catalytic Conversion Selective Catalytic Reduction (SCR) of Nitrogen Oxides Catalytic Oligomerization Process Using Synthetic Mesoporous CrystaI1ine Material Catalytic Olefin Upgrading Process Using Synthetic Mesoporous Crystalline Material Multistage Olefm Upgrading Process Using Synthetic Mesoporous Crystalline Material Production of Ethylbenzene Production of Olefin Oligomer Lubricants

37 Appendix 3 Mobil US Patents on Other than Catalytic Applications of Mesoporous Materials Authors

Patent Number

Publication Date

Herbst et al. Kresge et al. Olson et al.

5378440 03.01.95 5 3 6 6 9 4 5 22.11.94 5 3 6 4 7 9 7 15.11.94

Beck et al.

5348687

20.09.94

HeUfing et al.

5347060

13.09.94

Roth et al.

5238676

24.08.93

Beck et al.

5220101

15.06.93

Beck et al.

5145816

08.09.92

Withehurst

5143879

01.09.92

Title Method for Separation of Substances Supported Heteropoly Acid Catalysts Sensor Device Containing Mesoporous Crystalline Material M41S Materials Having Nonlinear Optical Properties Phase-Transfer Catalysis with OniumContaining Synthetic Mesoporous Crystalline Material Method for Modifying Synthetic Mesoporous Crystalline Materials Sorption Separation over Modified Synthetic Mesoporous Crystalline Material Method for Functionalizing Synthetic Mesoporous Crystalline Material Method to Recover Organic Templates from Freshly Synthesized Molecular Sieves

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H. Chon, S.I. Woo and S.-E. Park (Editors) Recent Advances and New Horizons in Zeolite Science and Technology Studies in Surface Science and Catalysis, Vol. 102 9 1996 Elsevier Science B.V. All fights reserved.

47

Synthesis, Characterization and Catalysis with Microporous Ferrierites, Octahedral Molecular Sieves, and Layered Materials Steven L. Suib U-60, Department of Chemistry, Department of Chemical Engineering, and Institute of Materials Science, University of Connecticut, Storrs, CT 06269-3060 This review concerns the synthesis, characterization, and catalytic activity of microporous ferrierite zeolites and octahedral molecular sieves (OMS) and octahedral layer (OL) complexes of mixed valent manganese oxides. The ferrierite zeolite materials along with borosilicate materials have been studied as catalysts for the isomerization of n-butenes to isobutylene, which is an important intermediate in the production of methyltertiarybutylether (MTBE). The OMS materials have tunnels on the order of 4.6 to 6.9 A. These materials have been used in the total oxidation of CO to CO2, decomposition of H202, dehydrogenation of C6H14, C6H14 oxidation, 1C4H 8 isomerization, and CH 4 oxidation. The manuscript will be divided into two major areas that describes zeolites and OMS/OL materials. Each of these two sections will include a discussion of synthesis, characterization, and catalytic activity. A..Overview. In the context of microporous and mesoporous materials, IUPAC has provided a variety of recommendations for nomenclature and characterization of porous materals, that can be found in the literature. 1 Microporosity should not be based on structural data but on adsorption data. Sorption by materials that show Type 1 isotherms is an indication of a microporous material. Pore size distributions less than 20 A are related to microporous materials like zeolites. Materials having pores between 20 A and 500 A are refered to as mesoporous materials. Materials that have pores larger than 500 A are refered to as macroporous. The measurement of pore size distributions is well established. However, the use of BET surface area measurements for zeolitic materials has been called into question due to potential multiple adsorption and nonconformity of monolayer adsorption implicite in the BET theory. The type of gas used, the method of data analysis, and even the use of the term surface area for a zeolitic material has been seriously questioned lately. On the other hand, most commercial manufacturers supply a surface area determined often by a three point or even a one point procedure that some researchers feel tells something about the material.

48 For microporous materials there is considerable debate on how to interpret surface area data due to the presence of micropores. Usually BET data are reported when one purchases a zeolite from a commercial vendor. The specific value of the BET measurement may give some indication of the relative surface area of the zeolite. Several experimental methods have been developed to collect BET data and many different equations have been used to model adsorption data. There are several discussions in the literature on various methods that have been used to analyze adsorption data by Broekhoff, et al.2 and Masthan et al.3 There is considerable debate on which method is the best for measuring BET data, however, there is good consensus that pore size distributions can be measured accurately with little problem in data analysis. Synthesis of novel micropomus zeolites such as intersecting 10- and 12-ring pore zeolites 4 has been the major focus of several research groups in the past few years. One of the more recent examples of catalytic applications that exploit micropores is the conversion of n-butenes to isobutylene over 10 ring zeolites like ferrierite as will be discussed later in this manuscript. Mesoporous molecular sieves materials 5-8 designated M41S (which include the MCM-41 class of materials) have made a further major impact on the area of synthesis of porous materials. A variety of open framework structures that are mesoporous have recently been reviewed by Thomas.9 Activated charcoal, MCM-41, mesoporous tungsten oxide, and substituted MCM-41 materials are mentioned. This article primarily emphasizes potential applications of such materials and possible mechanisms of reaction. The mesoporous sysems are compared briefly to microporous materials such as zeolites, ALPOs, MeALPOs and SAPOs. MCM-41 has been prepared by a liquid crystal templating mechanism where surfactant molecules are belived to act as templates by Beck et al.10 Surfactants such as C16H33(CH3)3NOH/CI in aqueous solution have been added to silica and alumina sources such as HiSil, tetramethyl ammonium silicate, and Catapal, respectively and then heated under autogeneous pressure to temperatures near 150oc for extended periods of time such as 48 h. In such procedures, products are cooled to room temperature, filtered to remove the spent Mother liquor, washed with distilled deionized water to clean the surface of the crystallites and then dried in air. The resultant solid materials show X-ray powder diffraction peaks indicative of mesoporous materials with d-spacings on the order of 40 A. At times other peaks at larger angles are observed and these materials have generally been indexed to hexagonal or cubic structure types for the MCM-41 materials. The attractive feature of these systems is that the size of the mesopores can be controlled by controlling the length of the hydrocarbon chain. Materials ranging from about 20 A to about 200 A have been prepared in this way. Siliceous MCM-41, aluminosilicate MCM-41, and mesitylene based materials have also been reported by Beck et al.10 N-Brand sodium silicate solutions were added to acidic solutions with the subsequent addition of surfactant and generation of a gel. Siliceous MCM-41 materials resulted by mixing these gels with water and heating the mixture to temperatures of 100oc for 6 days. Similar materials with different elemental compositions were prepared by using C12H25(CH3)3NOH/CI surfactant solutions with sodium aluminate solutions. Ultrasil silica,

49 tetramethylammonium silicate solution, or tetramethylammonium hydroxide solutions were added to the surfactant aluminate solution while stirring. Mixtures were then put into an autoclave at temperatures of 100oc for 1 day. Aluminosilicate MCM-41 materials resulted in such preparations. It is clear in the synthesis of these materials that aging of the gel is an important process and that rearrangements and bond transformation s can occur in the gel after it is formed. Specifically, mesitylene MCM-41 materials were prepared by adding the mesitylene during the last step of the synthesis,. Mesitylene gels as in the other MCM-41 preparations were then heated in an autoclave and allowed to crystallize. The specific conditions include temperatures of 105oc and quenching of the autoclaved gel after 4 h of thermal treatment. The calcination of such materials is important and involves treatment in nitrogen gas at elevated temperature and subsequent thermal treatment in air at elevated temperatures for 6 h. Calcination removes the surfactant "template" and it is clear that specific calcination treatments are critical in order to avoid degradation of the MCM-41 product and to completely remove the surfactant. 10 The mesoporosity of these materials has been established by BET measurements and gas adsorption experiments. As the chain length of the surfactant was increased from C 8 to C16, the amount of adsorbed benzene was increased, indicating that there was a relationship between the size of the surfactant and the amount of gas adsorbent taken up by the MCM-41 material. In terms of a comparison to zeolite materials, experiments were done at 60 torr pressure and at 25oc. The USY zeolite sample had an uptake that was about 4 times less than that of MCM-41. The above mentioned MCM-41 materials all show pore size distributions with broad bands centered around 40 A. The pore size distribution measurements are a true indication of the size of the pores and can be used to verify the existence of mesopores. Further evidence of mesopomsity comes from X-ray powder difraction experiments which were done to determine the crystallinity of these materials. The position of the (100) reflection was found to correlate with the amount of uptake by the different materials, or in therwords, with the mesoporosity of these systems. Pores of the MCM-41 materials were shown to form in a hexagonal shape by using high resolution transmission electron microscopy data. 10 The d-spacings of the (100) reflection of the MCM-41 materials made with mesitylene and the ratio of the number of moles of mesitylene to the number of moles of surfactant show a linear relationship as evidenced by X-ray powder diffraction data. As mentioned above, it is possible to produce both hexagonal and cubic MCM-41 structure types. In the case of the mesitylene systems, when the ratio of surfactant to Si was less than one, hexagonal MCM-41 was oberved. When the ratio of surfactant to Si was greater than one, cubic MCM-41 materials phase were observed. It is clear then that the structural type of MCM-41 can be controlled by proper choice and amounts of reactants. The original reaction mechanism for the growth of M41S and MCM-41 materials was proposed by Mobil researchers. 10 This proposed mechansim involved formation of rod-like structures of micelles and concomitant formation of a hexagonal array of rods, after which an inorganic species would encapsulate the rods and surround the surfactant species. Calcination of these composite materials led to the

5o formation of either hexagonal or cubic arrays of MCM-41 as shown by work of Beck et al.,10 and Kresge et a!.11 The Mobil researchers suggested that the inorganic encapsulating species might be able to control the formation of the micellular liquid crystalline phase during synthesis. The carbon chain length on the surfactant was critical for the control of the resultant structure of the micelles which in turn was important in the control of the mesoporosity and the pore size distributions of the M41S materials. A liquid crystal template model was proposed in order to explain the various stages of the crystallization process of MCM-41. Compositional and analytical data for these MCM-41 materials clearly show that isomorphous substitution of divalent, trivalent, tetravalent, and pentavalent cations can occur in these systems as reported by Kresge et al. 12 Several of the above-mentioned materials have been used in catalytic applications especially in the area of oxidation catalysis. For example, incorporation with titanium and vanadium species into MCM-41 has led to interesting partial oxidation catalysis as reported by several groups including Sankar et al., 13 Tanev et al., 14 Reddy et al., 15 Corma et a1.,16 and other work by Corma et a1.17 A variety of other elements have been suggested to be incorporated in a series of patents by Mobil researchers: (Pelrine et al., 1-8 Pelrine et al., 19 Bhore et al., 20 Le et al., 21 Le et al., 22 Le et al., 23 Bhore et a1.,24 Le et al., 25 Shih, 26 Kresge et a1.,27 Degnan et a1.,28 Del Rosssi et a1.,29 HeUring et al.,30 and further work by Kresge et al.31). Other novel structural materials that can lead to microporous and open framework systems are the vanadium phosphate class of materials which have 18.4 A elliptical tunnels32 which have been pursued by Haushalter, Zubieta and coworkers. A particularly novel structure consisting of inorganic double helices, 33 was also prepared by this group. Propped open structures have been made by the intercalation of organic/inorganic SnS2-cobaltacene layered intercalation compounds, 34 as reported by Nicoud which are reminiscent of intercalation of dichalcogenide species that are important in high energy battery applications. Novel microporous and mesoporous materials such as TiO2 based systems35, 36 are continuing to be reported that have interesting structural and catalytic properties. A specific area of interest in all of the above-mentioned systems is the precise control of the porosity of these materials, as well as control of their acidity. 37 Systematic control of the acidity of these materials and related materials (superacids) 38 is also an ongoing area of interest and is important in the potential use of these materials in acid catalyzed reactions. The above discussion is important as regards the synthesis of both microporous ferrierite materials as well as the mixed valent maganese oxide OMS and OL materials to be discussed below. All of the above materials in general are insulating materials. In addition, they are often synthesized with charge compensation in mind. For example, A13+ substitution for Si 4+ in zeolites leads to an inherent cation-exchange capacity. This is observed in zeolites, M41 S, and other materials like clays and pillared clays. Another way to approach the generation of microporous materials is to generate mixed valency of one element in an oxidic structure which should also lead to cation exchange capacity. Mixed valent materials 39 have been the focus of considerable attention due to their role in a variety of fundamental and applied research studies, including electron transfer, photoredox systems, biological materials, magnetic

.5] materials, semiconductors, and batteries. Several types of coordination complexes, clusters, enzymes, and extended structures have been studied in this regard. One particular system that tends to allow mixed valency in oxides is that of manganese. This is likely a result of the unusually high number of multiple oxidation states that are available for this element. Most mineral phases of manganese oxide show some degree of mixed valency. Birnessite is a layered mixed valent manganese oxide material that consists of MnO 6 octahedra, however, some of the manganese ions are reduced from Mn4+ to Mn3+. Birnessite is the most abundant manganese mineral. Birnesite shows cation exchange capacity as do many other mixed valent manganese oxide minerals such as todorokite, pyrolusite, ramsdellite, and others. Syntheses of mixed valent manganese nodules in agar gels 40, spinels 41 and layered materials 42 via sol-gel methods have recently been pursued. In addition, stabilization of 10 A manganite materials 43 via low temperature hydrothermal treatment may serve as standard materials for geothermometers. Considerable recent work concerns lithiation of manganese oxides (mostly layered structures) as potential battery materials. 44 Secondary nonaqueous rechargeable batteries are the goal of this area of research. Manganese containing malachite materials have also been shown to undergo solid state redox equilibria based on XPS data. 45 Review articles concerning mixed valent manganese oxides are available. 46 It is clear that the redox properties of the solution phase precursors, of the intermediates, and the final mixed valent products are very important during nucleation and growth of these materials. Organic oxidations 47 with manganese oxide have focused on dehydrogenations (loss of H) and incorporation of electronegative species (O). The activity of manganese oxides depend on particle sizes, solvents, methods of preparation, and other factors.47,48, 49 MnO2 is often used for oxidation of allylic and benzylic alcohols to aldehydes or ketones. In addition, amines can be converted to imines, thiols to disulfides, sulfides into sulfoxides, etc.47, 48 Manganese oxides have been reported to deposit on surfaces of bacteria and play a role in the decomposition of humic substances. 50 Natural manganese oxides systems have also been shown to decompose halogenated hydrocarbons. 51 Decomposition of H202 has also been catalyzed by MnO2 .52 In most of the above cases besides the decomposition of H202, the manganese oxide materials act as stoichiometric reagents in oxidation reactions of conversions of for example alcohols to ketones. A further report of the oxidation ability of manganese nodules is that of Nitta. 53 Several reactions were carried out with natural manganese oxide nodules including oxidative dehydrogenations of aikanes and cycloalkanes, reduction of NO, total oxidation of CO, and use in the gettering of metal and mixed metal ions. For example, nodules were found to have a tremendous capacity for adsorption of heavy metals and toxic metals like Pb 2+, and Hg2+. In addition, nodules have been used to sequester metals that are present in petroleum fractions that can contain metals like V and Ni. These metals can cause degradation of the fluid cracking catalysts even at levels as low as 1 ppm. A final report showing the potential of manganese oxide nodules was that of Weisz. 54 Manganese nodules dredged from the Pacific basin were tested in a variety

52 of catalytic oxidations including the total oxidation of methane, CO, and butane. The activity of the nodules was compared to commercial oxidation catalysts like Pt on AI20 3 and CuO. In all cases, the nodules were more active than the commercial catalysts. In addition, a good correlation was found between activity and large surface area. For example, as the surface area (ranging up to 250 m2/g) was increased going from one nodule to another the catalytic activity for total oxidation increased. This is a good example of the potential of the use of high surface area manganese oxide materials. B. n-Butene Isomerization with Boron Substituted Zeolites and Ferrierites. A major recent effort in our research program has involved studies of butene isomerization catalysts. Boron ZSM-5 and ZSM-11 catalysts have been synthesized and characterized55 and studied for isomerization of n-butenes to isobutylene.56,57 The study of these catalysts with a variety of bases such as NH 3, butene, ethylene have led to a better understanding of the amounts and types of acid sites in these materials. 58 A specific model explaining the role of different catalytic sites on B-ZSM-5 and B-ZSM-11 is shown in Figure 1. Quantitative Temperature Programmed Desorption (TPD) data have been used to identify 3 types (different stedc

~

Ethylene

~ i f i e g t i o n of Acid Sites on B-ZSM-5 + B-ZSM-II: Site 0, Adsorption of NH3; Site I, Adsorption of Nt][3 + 1-C4H8; Site H, Adsorption ofNH 3 + I-C4H $ + C2]B4.

Figure 1. Three Different Sites of B-Containing Zeolites. accessibilities) of sites: (1) those that only interact with NH 3 (2) those that interact with both NH 3 and ethylene (strong acid sites that lead to oligomerization) and (3) those that interact with butene (moderately weak acidic sites) that lead to isomerization. A systematic series of synthetic, characterization and butene isomerization catalysis studies of ferrierite and ferrierite-like materials such as ZSM-22 59 ZSM23, 60 and ZSM-35,61 was undertaken to study optimization of isobutylene product. Coke deposits 62 in the pores of these materials play a key role in isobutylene formation as does the overall acidity and structure of the pore system. Such shape selective effects have been probed with TPD methods. A comparison of catalytic conversions, selectivities, yields and other properties of some ferrierite and ZSM-22 samples is given in Table I.

.53 Table I Com0adson of Catalyl;i~ Properties of ~

Like Materials."

Catalyst*

Conversion**

Selectivity**

Yield**

Polymer"

Rate***

FER

60

40

24

12.5

161

FE R/TMOS

43

40

17

2.2

115

FER/AI20 3

43

82

36

4.0

241

ZSM-22

61

51

31

10

389

ZSM-22' 61 51 31 10 389 - 420oc, 20 cc/min, 220 mg catalyst. TOS = 4 h FER, 10 h ZSM-22. FER = Tosoh Ferrierite; TMOS = tetramethyl orthosilicate, source of Si; AI20 3 mixed with 40% Dispal powder, then extruded. ' = 100 cc/min. -in %. - mmol i-C4H8/min-g catalyst. The data of Table I suggest that addition of alumina binder can have a marked influence on selectivity with minimization of polymer (> C4) formation. Deactivation via coke formation depends on the exact pore size and shape of the zeolite. The time it takes for each catalyst type to reach steady state is quite variable, i.e., 4 h for FER and 10 h for ZSM-22. The FER/AI20 3 catalyst has the best yield of all materials we have studied. 61,62 Even though the structures are relatively similar, there are important differences in overall activity and deactivation. In the case of ferrierite we observed an excellent quantitative correlation between the different types of acid sites as determined by TPD and crystallographicaily different and accessible (to n-butenes) oxygen atoms. The 8 different types of oxygen sites of ferriedte are shown in Figure 2 and distributions are summarized in Table I1. Table II Distribution of 0 2- !ons in a Unit Cell of .FER. TvDe #

1 4

;~ 16

3 16

4 4

5 4

6 8

7 8

8 12

TPD data for NH 3, 1-C4H 8 and i-C4H 8 are given in Table III for FER.62 The amount of i-C4H 8 sorbed (using a similar procedure to that described above for quantifying 3 types of sites on B-ZSM-5, Fig. 2) is 30.8% whereas 1-C4 H 8 is accessible to 74.1% of the total acid " s~tes " . The 0 2 - "~ons o f type 1,2 and 4 are located in both 10-member and 8-member rings and have the largest space around them (XL). Types 3 and 5 are located inside the channels of the 8-member rings

54

i

(a)

(b)

The 8 Different O 2- Sites of FER along, a- [001], b- [010]. 9

I

I

Figure 2. Eight Different Oxygen Sites of Ferrierite. Table III Total Amount (%) Acid i . ~ ~;amole FER

NH3 100

~

i_nFER (4.4 [H+]/Unit Cell).

IIIII

1-C4H8 74.1

i-C4H 8 30.8

and type 6 is inside channels of 8-member rings. Types 3,5,6 grouped together are of about the same size (L) but with a smaller accessibility than XL. Types 7 and 8 (S) are all located on 5-member rings which limits space around them. Molecular modeling with Biosym software was used to study the steric constraints of these different sites. 62 Statistically , the % H+(XL), H+(L), and H+(S) are the same as the 0 2distributions (data from Table il). Note that the % H+(XL) [33.3%] closely matches the

.55 amount of i-C4H 8 sorbed (30.8%, Table III). The % of H+(XL) plus H+(L) [72.2%] also quite closely matches the amount of 1-C4H8 adsorption (74.1%, Table III). In summary, acid sites on FER have different size constraints from a structural point of view. Coke (predominantly aromatic in nature) formation is limited to < 11 wt. % of the micropore volume of FER. Coke formation modifies desirable polymerization (dimerization) reactions. Such blocking produces the pore shapes and limits access to more strongly acidic sites that catalyze less significant contributions for shape selectivity for skeletal isomerization of n-butene. TPD results suggest that adsorption of NH 3, 1-C4H 8 and i-C4 H8 is shape selective. 62 The technological importance of butene isomerization and similar experiments with ferrierite by researchers at Shell Amsterdam63,64 have recently been reported. 65 Results of our studies (and scaleup studies at Texaco) have led to the commercialization of n-butene isomerization catalysts by Texaco, Inc.66 C. Svnthesis Of Octahedral Molecular Sieves and Octahedral Lavered Materials. The primary building block in octahedral molecular sieves is an octahedral unit. As shown in Figure 3, there are several ways to link these primary building blocks together, such as vertex sharing, edge sharing, or face sharing as described in the treatise by Wells.67 MnO6 octahedra in many materials prefer to link at verteces and edges. The objective in the preparation of octahedral molecular sives is to be able to control the linking of primary building blocks in order to generate porous materials. The structure of synthetic todorokite is shown in Figure 4. It consists of edge and corner shared MnO6 units. Three MnO6 units are shared on each side generating a 1 dimensional pore which measures 6.9 A on each side. The synthesis of synthetic todorokite or OMS-1 was achieved by reacting Mg(MnO4) 2 with Mn2+ salts in NaOH aqueous solutions. The first step in the synthesis is the precipitation of synthetic birnessite or OL-I. This layered material is shown in Figure 2 and consists of edge and corner shared MnO6 units as in OMS-I. However, a layer structure is produced which has exchangeable cations between the layers as well as water molecules. The second step in the production of OMS-1 is the ion-exchange of OL-1 with Mg2+ cations to form a layered material that has the buserite structure. The final step in the synthesis of OMS-1 is the autoclave treatment of synthetic buserite at elevated temperature (near 150oc) and under autogenous pressure.68,69

56

PrimaryBuildingBlock : MnO6 ~

TiO6

%

ReO6

SecondaryBuildingBlock"

Edge Sharing Figure 3. Building Blocks of Octahedral Molecular Sieves. We now know from elemental analyses and ion-exchange studies that a key ingredient in the synthesis of OMS-1 is Mg2+ ions which must be introduced in the first step of the synthesis. On the basis of charge balance and analytical data, about 3 weight % Mg 2+ goes into the framework of OL-1 which leads to stabilization of the layered structure and concomitant generation of synthetic buserite and synthetic todorokite. If no Mg2+ is present during the initial nucleation of OL-1, it is possible to generate the layered structure, however, we have not been able to convert the non Mg2+ containing OL-1 species into OMS-1.70

5"/

8yntheti c Tod oro kite

OMS-1 2+ 2+ Mn 4+ .5012 4.47.4 55 H2 0 Mgl-l. /!nl.9 4.4.4 Figure 4. Structure of Syntheitc Todorokite, OMS-1. The incorporation of Mg 2+ into OL-1 is consistent with the finding of Mg 2+ in some natural birnessite materials. The Mg2+ isomorphous substitution in octahedral sites of OL-1 is also similar to the substitution of Mg 2+ ions in octahedral layers of certain smectite clays where Mg2+ substitutes for some AI3+, such as in montmorillonite which has an ideal composition of Exx[AI2.xMgx]{Si4}O10,~OH)2, where Ex stands for exchangeable cations, the [ ] symbols signify that AI ~H" and Mg2+ an octahedral layer, and the {} symbols signify that Si 4+ is in a tetrahedral coordination. Mg2+ ions can also substitute for AI3+ in the clay mineral chlorite, and is present in octahedral layers of the clay mineral vermiculite. 71 The conversion of a layered material like OL-1 into a 1-dimensional porous material like OMS-1 is also not surprising. It is well known that certain fluid cracking three dimensional zeolite catalysts can be grown from layered clays. In addition, it is noted that birnessite is the most common manganese mineral. Its abundance may be important in geological transformation of birnessite into todorokite.68, 69 Another 1 dimensional tunnel structure material that we have studied is synthetic cryptomelane or OMS-2. The structure and composition of cryptomelane are given in Figure 5. This structure is composed of a 2 x 2 edge and corner shared structure which generates a 4.6 ,~ 1 dimensional tunnel. Cryptomelane is a K+ form of the Ba 2+ mineral hollandite. The chemical formula indicates as is the case for OL-1 and OMS-1, that cryptomelane is a mixed valent manganese oxide with waters of hydration. Cryptomelane also has ion-exchange capacity just like OMS-1 and OL-1, although the restricted size of the tunnel does not allow very large cations to be incorporated, or they are only incorporated to a small extent. 72 During the synthesis of OMS-1 and OL-1, several other phases can be produced. In particular, psilomelane, pyrolusite and ramsdellite can be formed. Pyrolusite has a composition of MnO2 which has a simple tetragonal rutile structure

58 or the b-MnO2 structure. 67 Psilomelane is a 2 x 3 phase, and ramsdellite is a 1 x 2 MnO 6 structure having slit-like pores on the order of 2.3 A by 4.6 A. In addition to these phases, reduction of MnO2 to Mn203 which is known as bixbyite [actually (Fe,Mn)20 3 in nature] or reduction even further to MnO (which as the face centered cubic structure of NaCI)67 can occur during synthesis. Another phase that needs to be avoided in order to prepare microporous materials is Mn304 or hausmannite. It is important to try to avoid nucleation of these phases, or to be able to convert these materials into a desirable phase. Finally, when ion-exchange and isomorphous substitutions are carded out, at temperatures ranging from 400oc to 800oc, it is possible to produce dense spinel phases which are quite stable.

Synthetic Cryptomelane OMS-2

KMn8 O16. nH20 Figure 5. Structure and Composition of Synthetic Cryptomelane. The composition of OMS-1 is (Mg2+,0.98.1.35Mn2+1.89.1.94Mn4+4.38. 4.54012-4.47-4.55 H20. About 3 weight % Mg ~'+ is part of the framework of the resultant OMS-1 material. The average oxidation state of manganese in OMS-1 is about 3.5 as determined by thiosulfate titrations. The thermal stability of OMS-1 is up to 500oc for degradation in vacuum or in N2 and up to 600oc when the degradation is done in the presence of 02.68,69 Synthetic cryptomelane or OMS-2 has a composition of KMn8016.nH20. In this case, there is no substitution of the framework with K + or other ions such as the Mg2+ incorporation with OMS-I. The average oxidation state of OMS-2 is about 3.9. The framework is primarily composed of Mn 4+ ions, however, some Mn3+ ions are found. The thermal stability of OMS-2 is about 800oc for decomposition in vacuum or in N2 and up to about 900oc when decomposition is done in the presence of 0 2. In both OMS-1 and OMS-2, the presence of 02 leads to healing of the structure by O atoms. Defect sites are believed to be oxygen vacancies that are formed during thermal treatment. Synthetic birnessite or OL-1 has an an octahedral layer (OL) structure and a composition of [K,Na]4Mn14027-21H20). The average oxidation state of the

59 manganese is about 3.6 to 3.7, similar to that in OMS-I. This similarity in oxidation state may be an important factor in the transformation of OL-1 into OMS-1, and implies that the transformation may not be as dependent on redox chemistry as is the initial precipitation reaction of OL-1 or in the formation of OMS-2. The thermal properties of OL-1 are similar to those of OMS-1, however, in all cases, the presence of cations, the specific synthesis procedure, and the crystallinity of the resultant OMS1 and OL-1 materials can lead to large differences in thermal properties, especially in differential scanning calorimetry studies.68,69 The structure and composition of OL-1 are shown in Figure 6.

Figure 6. Structure and Composition of OL-1. Several methods have been used to produce different types of OL-1, OMS-1, and OMS-2 materials. The materials that are produced by various methods lead to vastly different materials, that have unique chemical and physical properties. Some of the properties that can be controlled are particle size, color, morphology, average manganese oxidation state, thermal stability, ion-exchange capacity, electrical conductivity, magnetic properties, crystallinity, defect density, desorption of oxygen, and catalytic properties. Table IV summarizes 16 different classes of OMS-1, OMS-2, OL-1, and amorphous manganese oxide (AMO) materials that we have prepared. These materials are separated into different classes because they show different crystalline, chemical and physical properties. For the case of OMS-1 these materials

60 are more crystalline and thermally stable than either natural or other reported synthetic materials. Table IV Summary of !4 Classes of OMS and OL Materigls. Method 1. Reflux

Structure OMS-2

2. Calcination

OMS-2

3. Sol-Gel

OMS-2

4. Sol-Gel

OL-1

5. Precipitation

OL-1

6. Hydrothermal

OMS-1

7. Ion-Exchange

M-OMS-1

Reactants

KMnO4, MnS04~?O MnC]2,02, NaOH KMnO4, KOH#ucrose Cyclodextrin, KMnO4,

Conditions I00~ 17h

Reference 72

400~

18h

?2

gel; 500~ 2h gel, 400oc, 2h

73 73

KOH

(runnel) 8. Ion-Exchange

10. Isomorphous Substitution

M-OMS-2 (tunnel) M-OL-1 (interlayer) [M]-OMS-2 (framework)

11. Isomorphous Substitution 12. Isomorphous Substitution 13. Complexation

[M]-OL-I (framework) [M]-OMS-I (framework) AMO

14. Sol-Gel

OMS-2

'i5. Crystal Growth in Gels

OL-I

9. Ion-Exchange

KMnO4, Mn 2+, NaOH #5,Mg 2+ exchange; #6 plus ion exchanl~e #1,2 plus ionexchange #5, ion exchange

room tanperature

68,69

autoclave 175o12

68,69

room temperature

68,69,74,75

room temperature

68,69,74,75

room Temperature

68,69,74,75

#1,2 plus Dopants added to initial sol

# 1,#2, added @room temperature #5, added @ room temperature #6, added @ room temperature 800(:

68,69,74,75

0~ gel, 500~

?3

#5, plus dopants

added to initial sol #6 plus Dopants added to initial sol KMnO4, oxalic acid KMnO4, ma]eic acid Sodium silicate gel containing

1V~z+, KMnO4

Room temperature, 30 days

68,69,74,75 68,69,74,75

76

??

above hardened gel.

Each of the materials in Table IV will be described below. Entry 1 involves the precipitation of OMS-2 from solutions that are refluxed. This synthesis results in high surface area synthetic cryptomelane materials. For method 1, the resultant 3.9 oxidation state of manganese is obtained by starting with a Mn 7+ reactant (KMnO4) and a Mn2+ reactant and approaching an intermediate oxidation state. The reflux method and a subsequent ion-exchange is shown in Figure 7.

6]

KlVlnO 4 +

Mn 100Oc//

M+

Mn 2+ + NaOH + 0 2

Figure 7. Synthesis of OMS-2 Via Reflux Methods. OMS-2 can also be made by oxidation of Mn 2+ as summarized in method 2. In this case, calcination to higher temperatures (400oC versus 100oc for method 1) is needed in order to obtain crystalline OMS-2 material. The particle size of the OMS-2 prepared by calcination is larger than that prepared by reflux methods. Another method for preparation of OMS-2 that we have recently reported involves the use of sol-gel techniques. In this case, sugars like sucrose can be reacted with KMnO4 to reduce the manganese with concomitant formation of a sol. Aging of the sol followed by calcination at elevated temperatures leads to crystalline OMS-2. Sol-gel derived OMS-2 materials show much greater thermal stability than materials made by either procedures 1 or 2, and they also are of much larger particle size. As particle size is increased the number of defects also decreases. The amount of oxygen desorption from OMS-2 prepared by sol-gel techniques is significantly lower than from materials prepared by reflux or calcination methods. While other sugars besides sucrose can be used to obtain OMS-2, sucrose gives the most crystalline and pure product. Synthetic birnessite or OL-1 can be prepared by sol-gel methods such as the reaction described in method 4. In this method cyclodextrin can be oxidized by KMnO4 with formation of a sol. Aging and calcination of the resultant gel at 400oc leads to crystalline OL-1. The particle size of the OL-1 made by method 4 is larger than with other methods, and to date we have not been able to convert the sol-gel derived OL-1 material into OMS-1 by any method. Method 5 describes the precipitation of OL-I. Such reactions of KMnO4 and Mn 2+ lead to small particle size OL-1 that can be converted into OMS-1. The crystallinity of this OL-1 is lower than the sol-gel derived materials and the thermal desorption of oxygen is marked for the precipitated materials as opposed to small amounts of oxygen desorption for the sol-gel materials. More oxygen defects are present in the precipitated OL-1 that the sol-gel materials (as is the case for OMS-2 systems, and the defects are again due to oxygen vacancies. This method along with subsequent ion-exchange of OL-1 is shown in Figure 8.

62

KMnO 4 +

Mn

2-1-

+

NaOH Ru

0L-1 M-I-

Figure 8. Synthetic Scheme for Precipitation of OL-1. Method 6 involves the hydrothermal alteration of OL-1 into OMS-I. The overall reaction scheme involves the precipitation of OL-1, followed by ion-exchange with Mg2+ cations, followed by treatment in an autoclave at 175oc for several days. The resultant OMS-1 material has very small particle size, high surface area (250 m2/g) and a considerable number of defect sites. It is esssential that Mg2+ ions are present during the precipitation of OL-1 in order to end up with crystalline and thermally stable OMS-1 materials. An overall scheme depicting the synthesis of OMS-1 is shown in Figure 9. Method 7 involves the ion-exchange of OMS-I. In this case, divalent ions such as Mn 2+, Co 2+, Cu2+, Ni2+, and Zn 2+ can be exchanged into tunnels sites of OMS1. Ion-exchange at temperatures of 80oc can lead to enhanced exchange of OMS-I. Good evidence for ion-exchange of monovalent ions such as in OMS-1 also exists. The ion-exchange forms of OMS-1 will be given the acronym M-OMS-I. Ion-exchange of OMS-2 can also occur due to exchange of K+ ions out of tunnel sites. Both monovalent and divalent cations have been exchanged into OMS2. The acronym M-OMS-2 will be used to signify ion-exchange of OMS-2. Ionexchange at temperatures greater than room temperature again lead to enhanced amounts of cation-exchange with respect to room temperature exchange. This process is described in method 8. Ion-exchange of OL-1 can also occur much the same way as with OMS-1 and OMS-2 as described in method 9. The ion-exchange of OL-1 is a critical step in the formation of OMS-I. The hydrated cationic complex is believed to act as a template around which the tunnel can be formed. There is a small variation in ionic radius that allows the formation of OMS-1, with ions ranging in size from Mn2+ to Zn2+ being ideal. In addition to substitution of tunnel sites via ion-exchange, it is possible to isomorphously replace cations in the framework of OMS-I. The general synthetic scheme is reported in method 10 and involves the doping of small amounts of cations into the precursor solution, before OL-1 is precipitated. Divalent cations like Mn2+, Co2+, Cu2+, Ni2+, and Zn 2+ can be incorporated into the framework in this manner.

63 Template cations are still needed in the second step in order to produce [M]-OMS-1 materials where [M] signifies incorporation of M into the framework. Isomorphous substitution of OMS-2 is described in method 10. In this case, the dopants are added to the solution prior to precipitation and reflux treatment.

o,.1

.g2

Mg.OL.1 175~ Autoclave

io.s 1 1 M+

Figure 9. Synthesis Scheme for Preparation of OMS-1. [M]-OMS-2 materials are significantly different than M-OMS-2 materials in terms of chemical and physical properties. A similar type of nomenclature is used for describing the isomorphous substitution of OL-I. For example, [M]-OL-1 would signify isomorphous substitution in the MnO 6 layers of OL-I. This preparation method is dsecribed in method 11 of Table IV. Method 13 describes the generation of amorphous manganese oxide (AMO) materials that are made by the complexation of oxalic acid with KMnO4. This reaction is done at 80oc and leads to an amorphous gray black powder. The chemical and physical properties of AMO are very different than all the other materials listed in Table IV. K + ions are incorporated into AMO due to reduction of Mn 4+ ions. Analytical data suggest that some unreacted oxalic acid is incorporated into the AMO powder. This suggests that some Mn 4+ ions are reduced to Mn3+ creating a mixed valent species, as is the case for OL-1, OMS-1, and OMS-2. The AMO materials has been found to be an outstanding photooxidation catalyst for the conversion of alcohols to ketones such as Isopropanol to acetone. Two key features of this system concern the ability to desorb oxygen at very low temperature and the synergistic effect of AMO with other solid substrates such as MgO. Oxygen can be desorbed from AMO at room temperature during photolysis. The oxygen can be detected by chromatographic techniques and it is clear from X-ray photoelectron spectroscopy (XPS) studies that there is an enhancement of oxygen at the surface of the AMO during photolysis. The amount of oxygen desorbed at low temperaures is markedly higher than crystalline OMS and OL materials.

64 Another interesting feature of the AMO photocatalysts is the effect of diluent substrates such as MgO or activated C. Addition of substrates causes an increase in the rate of photoassisted catalytic oxidation of isopropanol. A synergistic effect is clear; specific amounts of diluent lead to an increase. Too much or too little diluent leads to a decrease in rate. The exact explanation of this synergistic effect is not known, however, it may related to the ability of species such as OH or adsorbed hydrocarbons and intermediates to travel back and forth across the AMO/substrate interface. There does not seem to be a correlation of rate with the surface area, acid base character, particle size or other physicaVchemical properties of the substrate. Method 14 involves the sol-gel synthesis of OMS-2. In this case dicarboxylic acids like maleic acid are used to reduce KMnO4. Highly crystalline low surface area OMS-2 materials can be made in this manner. The dicarboxylic acids are oxidized to CO 2 as is the case for the sol-gel sugar preparations (Method 3). It is apparent from studies of sol-gel syntheses that strong acids react very rapidly and do not generate stable sols. For this reason, weak acidic material like dicarboxylic acids, sugars, cyclodextrins and similar materials need to be used. The final method mentioned in Table IV involves crystal growth in gels. In this case, sodium silicate gel is acidified with a weak acid like acetic acid and a sol is formed. Mn2+ ions can be dissolved in the sodium silicate reactant before addition of the sodium silicate. The order of addition of acid to sodium silicate is important because addition of sodium silicate to acid causes instantaneous gelation and air and CO 2 can be trapped in the gel. When acid is slowly added to the Mn2+, a sol forms that takes about 8 h to form a gel via syneresis. The evolved water during syneresis must be allowed to vaporize, so it is important that the container is not sealed, but only lightly covered to avoid contamination from the atmosphere and dust particles. Once the gel has set, KMnO4 solutions can be added to the top of the hardened gel. At this stage it is important to stopper the reaction tube, so that the KMnO4 solution does not change its concentration. After 1 week nucleation of crystallites occurs, and after about 1 month large crystals of OL-1 (150 m) are formed. The crystal growth in gel method slows down the crystallization process and can decrease the number of nucleation events with respect to similar reactions carded out in solution. There are some other methods that might be envisioned that could lead to OMS and OL materials. One obvious direction would be to use structure directors or templates that are similar to those used in zeolite synthesi such as tetraalkylammonium halides. Unfortunately, we have observed that such structure directors and templates react with KMnO4 and get oxidized to CO2. Another seemingly obvious route would be electrochemical syntheses. Some research has been done in this area, however, it is difficult to synthesize a sizeable amount of material such as with controlled potential electrolysis. In addition, some early work showed the generation of amorphous materials that after inital formation can be heated to form spinel phases without apparently going through the OMS/OL phases. Another obvious direction would be to mix several of the methods outlined in Table IV. In this case, at least for the preparation of OMS-1, so far we have only been able to convert OL-1 which has been made by precipitation methods. This may be due to the small article size and large number of defects in OL-1 synthesized by precipitation methods (method 5). One combination that does lead to new materials is

65 the isomorphous substition of either OL-1 (method 11) or OMS-2 (method 10) followed by ion-exchange of the isomorphously substituted OL-1 or OMS-2. In these materials, two types of metal ions can be incorporated into resultant OMS and OL materials, one in the tunnel sites and one in the framework sites. A variety of combinations of metals is possible. By converting [M]-OL-1 to [M]-OMS-1 (method 9) with ion-exchanging the [M]-OL-1 precursor, it is possible to incorporate two types of metal ions into OMS-I. Literature preparations of OMS-2 types of materials such as hollandite and cryptomelane materials are quite common. A variety of other methods have been used to make hollandite type materials such as thermal treatment of mixed oxides and high presssure syntheses, however, the resultant maetrials usually are of very large particle size and have porosities and catalytic actvities that are significantly lower than OMS and OL materials described in Table IV. Exceptions to this trend are the sol gel materials of Table IV (methods 3, 4, and 13) which lead to relatively stable and large particle size systems. The crystals grown in gels (method 15) also lead to relatively inert materials that have large particle sizes. OMS-2 prepared by reflux methods has higher surface areas, smaller particle sizes, more acid sites and greater defects (primarily oxygen vacancies) than materials reported by others.78,79, 80 For the sol-gel preparations, we are not aware of use of this method to prepare OMS-2. Both tunnel (M-OMS) and framework [M]OMS substitution is possible. We are also not aware of attempts to incorporate transition metals into the framework of such materials. Note that the different preparations of OL (#4,5) and OMS-2 (#1-3,14) of Table IV yield unique materials with different particle sizes and other physical properties. D. Characterization of OM$ and OL Material~. The compositional, electrochemical, structural, mixed valency, magnetic, thermal, acidity, and surface properties of OMS and OL materials have been investigated in detail by our group.68"70, 72"77 The semiconducting nature of such systems is particularly novel.68,69,75 The ease of electron transfer,75 may allow applications in energy and electron transfer as well as battery applications. Isomorphous substitution,74, 75 of OMS-1, OMS-2 and OL-1 has led to a variety of new materials where chemical and physical properties can be controlled. For the case of magnetic environments, spin glass behavior,74 has been observed in these systems. The extreme variety in chemical composition68-70.72-77 of materials such as those outlined in Table IV is difficult to achieve in similar systems such as clays and zeolites, because many more transition metals prefer to be in octahedral coordination and can be substituted for manganese than the number of substitutions for AI3+ and Si4+ in tetrahedral sites. Acidity of OMS and OL materials can be controlled by varying the composition and structure. This is not unexpected based on studies of zeolites and clays. Tremendous adsorptive capacities of OMS and OL materials are observed68, 69 on the same order of magnitude of natural manganese nodules. The OMS and OL materials rival zeolitic uptakes and can be as high as 20 g adsorbate per 100 g OMS/OL material. Similar observations53,81 have been made for the natural

66

manganese nodule counterparts of OMS and OL systems, although such natural materials are often mixtures, less pure, less crystalline, and less readily available. A summary of electrochemical, magnetic, and conductivity properties is given in Table V. Note that these properties can be controlled by proper choice of OMS/OL framework and composition. Some specific applications of these materials are as sensors to discriminate size and charge, as regchargeable nonaqueous secondary batteries, as new semiconducting materials, and in magnetic applications. The correlation of structure and composition with properties shown in Table V has not been well developed. For example, we do not know why K-OMS-2 is a spin glass material whereas other OMS-2 materials and OMS-1 and OL-1 materials are not spin glasses. The same lack of understanding holds for the conducting properties of these systems. OMS-2 systems have conductivities that are orders of magnitude greater than OL-1 and OMS-1 systems. This may be due to the incorporation of traps in both OL-1 and OMS-1 that are divalent cations that impede electron transfer, but this is not known with certainty. The mechanism of conduction is also not clear. In certain cases, electrical conductivity is apparent whereas in other materials, ionic conductivity appears to dominate. It is likely that both types of conductivity occur in some materials. Table V Physical ProDerties and ADolications of OMS and OL Materials. PROPERTY Electrochemistry Cyclic Voltammetry _

Conductivitv OMS-1 Ea = 0.35 eV OMS-2 Ea = 0.25 eV OL-1 E,~ = 0.39 eV Maanetism Mg-OMS-1 K-OMS-2 Cu-OL-1 v

COMMENT

REFERENCE

Charge Selective, Size Selective

76

5 x 1o-5 (w-c=)-I 1 x 10 -3 (W-cm)-] 1 x 10 -6 (W-cm) "]

82

Weak Magnetism Spin Glass Weak Ma~lnetism

28

General synthetic conditions for preparing OMS and OL materials are summarized in a proceedings manuscript. 83 The synthesis of small particles of OMS2 having the cryptomelane structure as well as a Rietveld refinement have been done to verify space group assignments. 72 We have recently shown that various inorganic cations can be used as true templates to form synthetic todorokite. 84 Characterization with a variety of techniques was necessary in order to understand the changes in structure and role of the template. The resulting materials can incorporate Cu 2+, Zn 2+, Ni 2+, Co 2+ and other divalent ions in tunnel sites of these materials. Organic reducing agents such as fumaric acids have been used to develop

67 sol-gel preparations of cryptomelane (OMS-2).85 In addition, simple sugars have been used with KMnO 4 to prepare synthetic birnessite materials (OL-1).86 These solgel routes have opened up many doors for the synthesis of OMS and OL materials and are much easier to carry out than our earlier preparations. Infrared spectroscopy methods have been used in these systems to study the role of the organic reducing agents. A variety of characterization methods have been used to study the optical, electronic, surface, bulk, thermal, morphological, and magnetic properties of OMS and OL systems. Cyclic voltammetry methods have be used 87 on these systems to study charge transfer since they conduct. In addition, both dc and ac electrical conductivity measurements have been made on OMS-1, OMS-2 and OL-1, and results suggest that both electrical and ionic conductivity exist. 82 Both dc and ac magnetic susceptibility experiments have shown that some of these materials are spin glasses.74 Surface properties76, 88 have been studied with both X-ray photoelectron spectroscopy and scanning Auger microscopy and results suggest that there is good lateral and bulk homogeneity even for doped samples. X-ray absorption studies in collaboration with Steve Wasserman and Katie Carrado at Argonne National Labs89, 90 have been done to study the average manganese oxidation states of these systems, and EXAFS studies have been used to confirm the octahedral geometry and to help understand differences in structural properties of crystalline OMS and OL materials as well as amorphous manganese oxide photocatalysts. E. Catalvtic Activitv of OMS and OL Svstems.

The two major catalytic applications of OMS and OL materials involve oxidations and photooxidations. Some amorphous manganese oxide (AMO) systems have been prepared that are outstanding photooxidation catalysts for degradation of CH3Br and conversion of isopropanol to acetone. 76 Catalytic data for several OMS and OL systems are summarized in Table VI. Table Vl Catalytic Data for OMS and OL Systems. Catalysis CO Oxidation H20 2 Decomposition C6H14 Dehydrogenation C6H14 Oxidation 1-C4H 8 Isomerization CH 4 Oxidation Photocatalvsis i-C3HTO-H Oxidation III

Product . . . . CO 2, 100% S, RT H 2, 02, 90% S to 1-hexene 60% Co 80% S to 1-C6H12OH c/t-2-C4H 8 C2,C3,C4,C5,C6

Reference 68,69 91

CH3(CO)CH 3, Room T, 100% S

76

I

II

I

II

II

91 91 68,69 68,69

II

I

68 One of the most important properties of these OMS and OL systems is their ability to lose and recover oxygen.92,93,94,95 Temperature programmed desorption, reduction, Oxidation, and studies of lattice oxygen mobility and structural stability studies have been done on a variety of systems. A review of these and similar open framework structures has recently been submitted.96 The major focus of our ongoing research is catalytic studies of OMS and OL materials. Catalytic activity for both total and selective oxidations is excellent.68,69,91 Some of the reactions under investigation are the liquid phase oxidative hydrogenation of cyclohexane, gas phase oxidative dehydrogenation of cyclohexane, oxidative dehydrogention of hexane to 1-hexene, total oxidation of CO, selective oxidation of humic acid, decomposition of peroxide, and similar reactions. We are optimizing conversions, selectivities, and yields and studying the fate of the OMS and OL materials after reaction. Mechanistic studies of interactions of organic reagents with OMS and OL surfaces are also underway. We have shown that [Ni2+]-OMS-2 and [Ni2+]-OMS-1 catalyze the selective conversion of hexane to 1-hexene. Stainless steel flow reactors of 1/4" diameter containing 0.5 g catalyst, charges of 7 g n-hexane in 2 h, 1 atm pressure and temperatures of 500oc are used in these experiments. Both gas chromatography (GC) and mass spectrometry (MS) analyses are done to monitor product distributions. Conversions as high as 60% and selectivities of 90% (to the terminal olefin) have been observed for the OMS-2 system. This may be a consequence of the better shape selectivity of [Ni2+]-OMS-2 (4.6 A tunnel) versus [Ni2+]-OMS-1 (6.9 A). The latter material is not as selective or active. Systems that do not contain Ni ~'+ are totally inactive. 91 There is precedence for dehydrogenation activity of these systems since manganese nodules have been reported to be excellent catalysts for dehydrogenation of cyclohexane. 53 Unfortunately, the gas phase reaction of hexane to hexene lead to degradation of the catalysts. It is clear that the manganese is reduced predominantly to Mn 2+ in the crystalline residue which has the MnO structure. The XRD peaks are very broad indicating a highly defect type structure which still contains MnO 6 octahedral units. Due to this structural change in the catalyst we have been focusing on oxidative dehydrogenations. In this case the hope is that oxygen can be used to regenerate the mixed valent catalyst when the organic substrates are selectively oxidized. Liquid phase oxidations have been shown to preserve the structure of OMS during reaction and produce yields and conversions that are on the same order of homogeneous catalysts. The synthesis of OMS and OL materials has led to unique opportunities to study fundamental chemical and physical phenomena such as mixed valency, electron transfer, spin glass behavior, conductivity, isomorphous substitution, catalytic oxidations, photocatalysis, new synthetic methods, structural analysis, phase stability, and other areas.68-70, 72-77 Semiconducting molecular sieves are rare and are excellent materials for high resolution spectroscopic characterization (especially imaging, surface studies, etc.) and for catalytic studies for correlation of activity/selectivity and redox or electron transfer capability. The potential of OMS materials as adsorbents, catalysts and in secondary battery applications has been featured in Catalytica ,Highliqhts_,97 Scientific American98 and Science Digest.99

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H. Chon, S.I. Woo and S.-E. Park (Editors) Recent Advances and New Horizons in Zeolite Science and Technology

Studies in Surface Science and Catalysis, Vol. 102 9 1996 Elsevier Science B.V. All rights reserved.

75

Organic Zeolites? Stephen Lee Department of Chemistry, University of Michigan, Ann Arbor, M148109-1055, USA. and D. Venkataraman Department of Chemistry, University of Illinois, Urbana, IL 61801, USA. January 11, 1996

Introduction Zeolites are a class of inclusion compounds, in which the occluded guests can be removed without the collapse of the host structure. The porous structure that results upon the removal of guests (generally water) has proved to be technologically important in the areas such as shape selective catalysis, sensors and ion-exchange membranes. ~ Parallel to the developments that are taking place in the synthetic procedures of these aluminosilicates, there has been an ongoing effort for the construction of porous solids using organic compounds. It is believed that by the use of organic compounds, a greater level of control on the nature and size of the cavities may be possible. Also, these organic based materials may be more easily processed into thin films or fibers. One of the difficulties in the realization of porous materials using organic molecules is the isotropic nature of intermolecular interactions. This isotropy generally leads to closest packing. Directional intermolecular interactions such as hydrogen and coordination bonds have been used to ovemde the tendency for closest packing. If a host network has void space, then these voids may be filled in two ways: a) by inclusion of guest moieties such as solvent and/or counterions (like zeolites) and b) by the interpenetration of topologically equivalent nets - a process of self-inclusion. Generally, either mechanism results in a dense structure. Also, in contrast to zeolites, typical molecular inclusion compounds irreversibly lose crystallinity, undergo a phase change or alter their morphology upon loss of their guests. It has, however, been shown that some coordination and hydrogen bonded networks can rapidly exchange inclusions or counterions while maintaining crystal integrity. 2 The problems of closest packing and the general instability of the host lattice raises the question of whether a porous structure with zeolitic properties constructed from organic compounds i.e. an organic zeolite, can be realized? More specifically, can a porous organic solid, if realized, function like a zeolite? In this review, we shall address these questions by taking a few examples of inclusion compounds where the host structure is proven to be stable to solvent loss. We will also evaluate a future direction for these compounds in the context of building hosts for the synthesis of chiral organic molecules. These examples may be considered as the stepping stones towards the goal of an organic zeolite. 3

76

Dianin's Compound One

of

the

well-studied

inclusion

compounds

is

4-p-h y dro x yphen y 1-2,3,4-

trimethylchroman. 1, commonly known as Dianin's compound. 4 First

CH 3

prepared in 1914 by the Russian chemist A. P. Dianin, this compound has attracted considerable attention due to its ability to tenaciously hold on to certain organic solvents. A cage structure was predicted in the adducts of 1 based on space group, unit cell

H3C

dimensions and packing considerations 5 and was confirmed by detailed X-ray single crystal structure studies on the chloroform and

OH

ethanol adducts of 1 after 15 years. 6 Irrespective of the guest, the structure consists columns of independent cages of ca. 6/~ x 11 ~, running along the c-axis (see Figure 1). Six molecules of 1 link together through hydrogen bonds to form a complex such that the oxygen atoms (of the phenol) constitute an hexagon. Alternate molecules in the complex point up or down. The cage is formed when two of these complexes stack in a way that one of the hexagons of the H-bonded oxygen atom constitutes the ceiling and the other hexagon constitutes the floor of the cage. 6 The stacking of these cages along the c-axis result in columns whose topology resembles the interior of an Allihn condenser.

Figure 1: A stereoviewshowing an independentcage of in the structure of Dianin's compound (1). Based on the occupancies and thermal factors of the guest molecules in these structures, it was speculated that some of the cages might be completely empty. 7 This speculation led to the conclusion that Dianin's compound might retain its cage structure upon complete removal of the

-2-

77

guest species. A definitive proof that the cages are only partially filled and that the host is stable came from CP/MAS 129Xe and 13C NMR of the xenon occluded guests.

Ripmeester and co-

workers crystallized 1 from dodecane under varying xenon pressures. ~la In the 129Xe NMR spectra of the occluded complexes, they found two peaks, ca. 18 ppm apart, which varied in intensity as a function of the pressure of xenon used during the crystallization. The high field peak was predominant under low xenon pressures while the low field peak was predominant at higher xenon pressures. As illustrated in Figure 2, depending on the concentration of xenon, the cages can be completely

empty

(A),

partially

filled

(B)

or

completely filled (C, the maximum occupancy of a cage being two xenon atoms).

From their intensity

behavior, the high field peak was assigned to partially filled configuration (B) and the low field peak was attributed to the completely filled configuration (C). The authors of this paper mention that "it is not really clear whether the presence of the second guest in the cage is the primary reason for the shift, or whether the second guest causes a change in the configuration of the methyl groups at the neck of the cage which in turn is sensed by the xenon atom." Since they also reported similar behavior in the ethanol-xenon adducts and 13C NMR showed chemical shift changes in the methyl carbon, it is tempting to assign the 18 ppm shift to the change in configuration of the methyl groups. This NMR study attests the earlier speculation that some of the cages might be partially occupied.

In addition, although the ref'mement details remain unpublished, MacNicol and co-

workers noted that their X-ray studies confirm that even the unsolvated 1 retains its cage structure. 8 Since Dianin's compound retains its structure after the loss of solvent, can it function like a zeolite? In other words, can the adsorption and desorbtion of the guests take place without the loss of crystallinity or change of phase? Barrer and Shanson demonstrated that 1 does behave like a

I I I I I I

~ds

zeolite by studying its sorption properties using a variety of gaseous guests. 9 The guest molecules used were argon, krypton, xenon, carbon dioxide, methane, ethane, propane, n-butane, iso-butane and neo-pentane.

They found that the sorption isotherms of 1, like

-3-

P

Po

A ty.pe I adsorption isotherm

78

zeolites, were of Type I (Brunauer's classification). 1~ More recently, the dynamics of these occluded guest molecules in the cages of 1 have been studied by NMR spectroscopy. 11 Unlike zeolites, the cages in Dianin's compound are not extended.

The cavities that

constitute the ceiling and the floor of cage are ca. 2.8 ]k in diameter. This indicates that the entry of the guest molecules into the guest-free Dianin's compound may occur through some reorganization.

However, there is no satisfactory explanation to why the porous solid state

structure of 1 is stable to solvent loss. Although the host lattice may not be as rigid as zeolites, it is important to note that the structure does retains its pores after the removal of guests and has sorption properties like zeolites.

Helical tubulates of 2,6-Dimethylbicyclo[3.3.1]nonane-exo-2,exo-6-diol

An interesting inclusion compound derived from 2,6-dimethylbicyclo[3.3.1]nonane-exo-

2,exo-6-diol, 2, was reported by the group led by Bishop and Dance. 12 The crystal structure of the alicyclic diol 2 can be

HO

OH

construed as packing of helices along the c-axis. The parallel canals that result from the helical tubules have an unobstructed triangular cross-sectional area of roughly 20 A2 which are

H3C

occupied by the guests (see Figure 3). ~ The helices bear striking

"

u

CH~

2

similarities to that in the structure of a-quartz. The resemblance to a-quartz becomes important in the design strategies, as discusses in the second part of this review, to use these porous organic solids for asymmetric synthesis. 13

Figure 3: Stereoviewof the helical tubulates along c-axis in the guest adducts of 2.

-4-

79

Several guest adducts of 2 with similar helical tubuland structure have also been reported by Bishop, Dance and co-workers. ~~ The most interesting result they was reported was that guestfree sample of 2, irrespective of how it was prepared, retained its open channel structure. The guest-free samples were prepared in three ways: a) by heating the inclusion compounds under reduced pressure b) by sublimation and c) by crystallization from mesitylene. ~4 The products that were obtained by the three methods were identical by IR spectroscopy and elemental analysis. The elemental composition, C11H2002, corresponded to the pure diol 2. 13C CP/IVIAS NMR studies showed that the carbon resonances observed for the guest-free product were almost the same as for the helical tubulate compounds. The powder X-ray diffraction pattern of the guest-free product was in complete agreement with the calculated pattern from the single crystal coordinates of adducts of 2 but without the solvent. It is rather intriguing why 2 does not crystallize in a closest packed arrangement in the absence of solvent. Even though crystal structures of various guest adducts of 2 are known, TM no studies about the exchange of these guests have been reported. The examples described above show that it is possible to realize inclusion compounds which does not collapse upon removal of solvent. The host lattice in these structures is primarily held by intermolecular hydrogen bonds. Since the stability of zeolites results from the stronger SiO-AI covalent bonds, it would be of interest to explore intermolecular bonds that are reversible and have similar strength to that of covalent bonds. In this regard coordination bonds have attracted wide attention for the construction of networks. ~5 A general strategy that is being followed is the construction of molecular analogs of prototypic minerals using organic ligands and metal ions. We will discuss two coordination networks constructed by this approach~6 from the group led by Lee and Moore which retain their pores upon partial or complete loss of the included solvent.

[Ag(1,3,5.tris(3-ethynylbenzonitrile)benzene)CF3SOa].2C6H6 The

structure

of

[Ag( 1,3,5-tris(3-

CN

ethynylbenzonitrile)benzene)CF3SO3].2C6H6, 4, is a 3-connected, two-dimensional coordination network (see Figure

4 ) . 17

II

In this structure, the coordination

geometry around silver is trigonal pyramidal, with three nitriles of the network in the basal plane and a

NC

triflate counterion bound to the apical position. The orientation of the nitrile groups on 3 together with the above-mentioned silver coordination geometry gives a [12]annulene-like segment as the simplest cyclic motif

-5-

CN

1,3,5-tris(3-ethynylbenzonitrile)benzene (3)

80

of the network. These sheets are stacked in an ..-ABCD..- sequence creating channel structures that run at an oblique angle to [101 ] i.e., the layer normal (Figure 5). Difference Fourier analysis locates sixteen molecules of benzene per unit cell in the channels, four of which are disordered.

Figure 4: Illustration of the [12]annulene unit in a sheet that constitutes the structure of 6. These sheets stack at an oblique angle to[101] to generate channels of c a . 8 ,~, in diameter.

The thermogravimetric analyses (TGA) of microcrystalline powders of this complex trace revealed two discrete mass losses at 110 ~ and 145 ~ corresponding to a mass percent of four (4.5%) and twelve benzene molecules (13.5%) respectively. It was rationalized that the initial mass loss corresponded to the loss of four disordered benzene units. The mass loss at 145 ~ was attributed to the loss of the ordered twelve benzenes. Differential scanning calorimetry (DSC), optical microscopy and X-ray powder analysis showed that there was no phase change associated with the first mass loss. At 145 ~

much like an inclusion compound behavior, a solid-to-solid phase transition

occurs concomitant with the loss of the remaining benzene molecules. This high temperature solid phase eventually undergoes a melting transition at 169 ~

Upon cooling the melt, a glassy

material is obtained. DSC cooling traces indicated heat capacity jumps that could be attributed to a glass transition. No chemical decomposition of the ligands occurs up to 200 ~ as verified by redissolving the glassy material and recording its NMR spectrum. The

unit

cell

parameters

of

a

single

ethynylbenzonitrile)benzene)CFaSOa].2C6H6 heated to 110 ~

crystal

of

[Ag(1,3,5-tris(3-

for 10 minutes and subsequently

cooled to room temperature remain unchanged within the standard deviation of the original crystal. ~s TGA experiments on crystals of similar or larger dimension confirm that a mass loss

81

equivalent to four benzene molecules occurs under these conditions.

These crystals remain

optically transparent and uniformly birefringent when viewed between crossed-polarizers. Under higher magnification, interesting surface changes as a function of temperature were noted in these crystals. It could be seen from the optical micrographs that surface defects appear around 110 ~

increasing to 124 ~

However, the macroscopic crystals remained intact. It is

likely that there is a surface reconstruction after the removal of the disordered benzene molecules.

Figure 5: A stereoview showing the channels which are at an oblique angle to (101) in the crystal structure of 4. Benzene molecules occupy these cavity and are omitted for clarity. Microcrystalline samples heated to 145 ~ under vacuum and subsequently cooled to room temperature re-absorb benzene vapor in an amount that corresponds to the mass percent in the original, unheated sample. Sorption saturation is achieved in

ca.

60 h at room temperature. X-ray

powder diffraction shows that the original solid phase was reformed. This behavior is analogous to that of classical inclusion compounds.

In contrast, samples of 4 heated to 110 ~

re-absorb

benzene in an amount equivalent to four molecules of benzene in less than 45 min. without ever undergoing a phase change.

[Ag(1,3,5-tris(4~thynylbenzonitrile)benzene)CF3SO3]-2CrH6 Two polymorphs exists for the complex of 5 with silver(I) triflate. We have previously reported the crystal structure one of polymorphs, 6, and it is homeotypic with LaPtSi (ThSi2-type) structure. 2b,~9 In contrast to the LaPtSi structure where each of the atoms are roughly comparable in size, the tritopic ligand 5 is significantly larger than either the Ag(D or CF3SO3" ions. As a result, a single LaPtSi-type net constructed with the dimensions of the tritopic ligand 5 and silver

82

(I) generates large void spaces. The triflate anion is of insufficient size to fully occupy these voids. The voids along [010] and [001] are filled by the six mutually independent, interpenetrating LaPtSi-type lattices that constitute the [Ag(5)CF3SO3]-2C6H6 structure (Figure 6). 20

The

interpenetration occurs in a way to accommodate the propensity of aromatic rings to lie in a slightly staggered coplanar arrangement with an offset angle (o0 37 ~ at an interplanar distance (d) of ca. 3.3 ]k. To accommodate the geometrical requirements 21 (o~ and d) for stacking of the aromatic rings, the nets are sheared to an angle (0) of 60 ~ (0=90 ~ for an ideal LaPtSi net and 0 ~ for CaCuP (A1B2-type).

However the interpenetration leaves the channels along [100] near the

maximal size possible for this network (see Figure 7). This cavity of 15 x 22/~, is filled by the solvent molecules.

The coordinates of 12 benzene molecules/unit cell were located in the

refinement.

Earlier, we had reported the exchange of

CN

benzene with benzene-d6 in the channels of 6 without the destruction of crystallinity.

We have extended this

study to other guests like toluene, m-xylene, undecane, benzyl

alcohol,

2,6-di-tert-butylphenol

and

II

(+_)-l-

phenylethanol. 22 Cell constants determined by X-ray powder diffraction after exchange were close to the original cell constants.

Optical micrographs showed

that there was no dissolution and reformation of the crystals during the period of exchange.

NC

CN

1.3.5-tris(4-ethynylbenzonitrile)benzene(5)

Guest-free samples of 6 can be prepared by heating the guest included complex in a TGA furnace room temperature to 200 ~ under nitrogen. When the samples were heated in an open pan, no phase changes were observed in the DSC.

The X-ray powder diffraction data can be

indexed as a two-dimensional rectangular unit cell with no sufficient information along the a-axis. However the lattice constants b and c had changed only by 5%-10% from the initial single crystal model. 2~ The OkI reflections calculated from the single crystal model (orthorhombic cell) were in agreement with the observed Okl reflections (rectangular cell). Based on this model, it can be concluded that the 15 A x 22/~, pores have been retained in the guest-free samples. The guest-free samples were exposed to vapors of benzyl alcohol, benzene, m-xylene, cyclooctane, undecane, (+_)1-phenylethanol. 24 Thermogravimetric analyses showed that cyclooctane and undecane were not absorbed by the guest-free sample of 6 in 60 h. However, benzyl alcohol, (+_)-l-phenylethanol, benzene and m-xylene were absorbed with the guest to ligand stoichiometry of ca. 3.5:1.0. The lattice cell constants calculated from X-ray powder diffraction were identical to those calculated from solution exchanged samples.

83

Figure 6: A single LaPtSi-type net from the crystal structure of 6. Five more nets interpenetrate to fillthe void space created in a single net.

Figure 7: A stereoview showing the 15 ,A,x 22 ,~ channels along the a-axis in the crystal strucure of 6. The solvent molecules and the counterion occupy these channels and have been removed for clarity.

Recently Yaghi and co-workers reported a coordination network based on trimesic acid and Co(II) which retained its porous structure upon the removal of solvent molecules and selectively absorbs aromatic guests. 25 Also, Fujita, Ogura and co-workers have demonstrated that cyanosilation of aldehydes can be performed in the microchannels of a two-dimensional square material composed of cadmium(II) and 4,4'-bipyridine. 2~ From the examples described above, it is quite apparent that organic zeolites are more of a reality than a fantasy. 26 What is the future that await these porous compounds with zeolitic behavior?

84

Further Directions One future direction for organic zeolite analog chemistry will involve the synthesis of chiral organic molecules. The reasons are two-fold. First molecules in crystals are generally trapped in just a few conformations. If the host crystal environment is chiral then this chirality can be imparted to the guest. Second, there are a large number of readily available chiral organic building blocks which can be used in the preparation of chiral host crystals. This is especially true for organic molecules and hence there is a clear advantage for organic as opposed to inorganic crystal hosts. We highlight here a few studies in which the synthesis of chiral molecules has been achieved through the use of organic crystals in the hopes that this will prove a useful incentive and review. The reported studies fall into two natural categories. In the one case one starts with racemic mixtures or optically inactive compounds, crystallize these materials into chiral crystals and finally by subsequent reactions, trap this chirality in the final chemical products. In the second category one forms host-guest inclusion compounds in which the host is already an optically resolved compound. This in turn leads to the formation of optically active guest molecules. In the first class of studies the sole chiral influence derives from the asymmetric environment of the molecule in the crystal. This implies that while the initial molecular sample is either fully racemized or just not optically active in solution, the crystal is optically active and thus belongs to one of the 65 chiral space groups (out of a total of 230 possible space groups). As the starting mixture is achiral in all likelihood the crystalline product will be a mixture of the two chiral forms. However by either growing just a few large crystals out of the sample and/or by seeding the sample with previously selected crystals of a given chirality one may still obtain reasonable quantities of chiral crystals. While the hand-picking of crystals may play an important role, it has been shown that this is not necessary in imparting a given chirality to a bulk sample. One of the clearest examples of this is in the study by Wilson and Pincock on a racemic mixture of binapthyl, 7, crystals. 27

(R)-(-)-I, l'-binaphthyl

(S)-(+)-1,1'-binaphthyl

85

These authors observed that binaphthyl crystallize in two polymorphs. The one is stable at lower temperature, is centrosymmetric and is not optically active. This polymorph melts at 145 ~

The

second polymorph is stable at higher temperature but is metastable at room temperature. It is optically active and melts at 158 ~

Wilson and Pincock show that as one cycles in temperature

between room temperature and 150 ~

a sample which is initially the optically inactive low

temperature polymorph transforms to an optically active solid. After three or four cycles one achieves the maximum optical resolution which corresponds to 56% ee. The crux of the Wilson and Pincock experiment is that at 150 ~

the reaction physically resembles a solid state reaction in

which the low temperature form is melting in the near presence of high temperature polymorph crystals. These chiral crystals are therefore nucleating sites for further chiral crystal growth. As at room temperature binaphthyl retains its chirality, the resultant samples can then be dissolved with retention of stereochemistry. The studies by Wilson and Pincock are however not the first study to show chirality from an initial achiral host. In earlier work Penzien and Schmidt showed that 4,4'-dimethylchalcene, 8,

Br2 (g)

v

A

single Crystal

Ar =

~/-~CH3

r

8

~"-Br

A•COAr ....Br

enantiomer ratio, 53:47

crystallizes in the chiral space group P212121. 28 Further this central ethylenic bond undergoes addition with Br2 to yield for a given single crystal with 6% optical yield. In subsequent papers Schmidt and other workers at the Weizman Institute further explored the ability of olefins to form chiral products. '9 Many of their studies centered on [2 + 2] photoreactions between adjacent olefins. In the absence of pores in a given solid it may be seen that such photochemically induced reactions, which are either intramolecular or between neighboring molecules, are ideal as one of the reactants does not have to break through the initial solid state host, thus destroying the originally advantageous chirality. In reactions such as the olefin plus bromine addition reaction one needs the chirality to be preserved on the surface of the crystal up to the moment of the transition state. If the reactant molecule loses its initial chiral arrangement the chirality is of course not transmitted to the product. ~ Schmidt and his coworkers who had already developed great expertise in solid state photochemistry observed that if one crystallized a disubstituted olefin, the olefins could orient themselves in one of two arrangements, 9 or 10. 31

86

In the arrangement shown in 9 a subsequent photoreaction would lead to an achiral product. By contrast, the arrangement 10 would lead to a chiral substituted cyclobutane.

After studying a

a..~

number of disubstituted olefm compounds the authors discovered that such olefins tend to form in the former

\ b _ \....~

and not the latter arrangement. They therefore proposed

\b

a

two modified versions of the [2 + 2] photocyclizationf

a

a try /9

\~

In one set of studies researchers prepared a mixed

\b

crystal with two different types of disubstituted olefins. This is illustrated in 11, involving two olef'ms, one with a thiophene and the other a pure benzyl derivative. As

a

11 shows there are two possible ways for these two olefins to interact. However as there was no symmetry element relating the upper to the lower reaction of 11, a

tw

/---

b

/9

10

disproportionate amount of one reaction could and did occur with respect to the other. This led to a final optically active product. 32

Ph ~-~Ar

Th

4A

hv

axis

r ~---

Ar

Ar Th hv

PhL_

Ar

In a second approach the authors considered disubstituted diolefins as is generically illustrated in 12. 33 These disubstituted olefms tend to arrange themselves in crystals in the form shown in 12. There is therefore good potential for chiral product formation. On this note, some authors have found that mutual interaction between the a, b and X substituents (see 12) can insure

87

the overall enforcement of this stereochemistry. 33b Especially ideal are a or b units which are esters and X units which are phenyl spacers. Conversely carboxylic acids and amides tend to hydrogen bond in a symmetric fashion thus leading to a achiral crystal type. This latter method has led to a number of successes. 7 One interesting application involves the formation of chiral polymers from an initially racemic mixture of starting materials. 34 In all these experiments, a key intermediate step is the formation of a chiral crystal from an achiral starting material. Rational procedures for such a synthesis are therefore desirable. In a series of beautiful experiments Addadi, van Mil and Lahav demonstrated one such rational approach. 35 In this work they first made use of an chiral a-substituent (see 12, in this example a was a chiral sec-butyl group, X was a phenyl link and b was an ethyl group). As the initial diolefin was chiral, the resulting product was also of necessity chiral.

They then studied the resultant

structure. They found that sec-butyl groups from neighboring groups were in close proximity to one another.

They therefore noted that were they to transfer the methyl group of the sec-butyl

group to the neighboring molecule they would then have a formal mixture of 3-propyl and 3-pentyl units with nearly the same quality of packing. These latter two molecules are not optically active. They then demonstrated that these latter molecules when grown together into a co-crystal forms a unit cell somewhat like the initial crystal (in the

sec-butyl

case unit cell dimension are a = 13.17 A,

b = 6.94, c = 5.25, ot = 103.1 ~ 13 = 95.5 ~ and "/= 90.1 ~ while in the latter case a = 13.53 ,~,, b = 6.90,4,, c = 5.28 ,~,, ot = 102 ~ [3 = 104 ~ and ), = 94~

Both crystals are chiral and are in the P1

space group. (Pure 3-propyl and 3-pentyl crystals by contrast form in P-1 with rather different at cell dimensions.) The optically active crystals upon exposure to UV light do indeed react to form an optically active product mixture.

Finally in the same work they show that a crystal in P21

symmetry with a = 3-pentyl b = methyl and X = phenyl link (see 12 for the location of the a, b and X substituents) leads to a resultant chiral cyclobutane product with ee up to 100%.

t2

x hv x

t2

88

A number of fine studies have shown that high ee's can be achieved. 36 It has generally proven feasible to determine the resultant chiral form from the chiral conformation of the reactant molecules in the crystal. Generally the optical product closest in geometry to the chiral reactant is the true final stereoisomer.

Examples of this, 13 and 14, illustrate schematically the chiral

conformation of the reactants and the chiral form of the products. Some recent studies include those in reference 36. hv H3

Ar HO Ar = - - ~ C I

13

14 in (-)-crystal

14 in (+)-crystal

H3 ~ ~ - ~ P h

P H3(~ ,ON~CH3

H3~"~O CH3

CH3 hv

H3C

CH3

Ph

Ph

CH3

i i

CH3

Mirror plane 14

A second methodology for chiral synthesis is to use a host-guest method in which the initially chiral host imparts chirality to the final reacted guest product 37 Useful hosts have proven to be 15 through 18. In all cases shown here the host is not only chiral, but has two optically active phenol or alcohol groups. These alcohol or phenol units are held in a fairly rigid manner to one another and therefore can readily hydrogen bond to many guests in a stereospecific manner. Two examples, both taken from the insightful work of F. Toda illustrate (Is)

89

the practicality of such host ligands. In the first example the host 16b was co-crystallized with 19 to make a 1:1 inclusion compound. 12 While the molecule 19 is by itself is not chiral, once linked by hydrogen bonds to the chiral host, a twist is introduced into the backbone of the 19 molecular structure. Based on the X-ray structure the inclusion compound can be seen to be like that shown in 20. Further as 20 shows, 19 undergoes a photoinduced cyclization. In the inclusion compound this photoreaction proceeds with 100%ee to only one of the two possible stereoisomers. CH3 CI

~

H3C---~

HO

HO

/CH3

(S,S)-(-)17

(s, S)-(-) 16 16a R= 16b R=p-C6H5 16c R=m-C6H5 Ph.

OH

P

-0

O

O

Ph Php/\OH

"0

NMe2 19

R= - (CH3)2or cyclohcxylor cyclopentyl 18

oX

Ph

Ar

\

H

i Ph

/O

\\

.." "0

0 "/

Ph

NMe2

.H

Ph,.

OH

hv \CH3

20

90

In a related host-guest based chiral synthesis, 38 the molecule 16b was used to make a 1:1 inclusion compound with 21. Just as in the case of 19, 21 is achiral, contains

o i

two oxygens on adjacent carbon atoms and undergoes a photorearrangement.

.,r

The photorearrangement is a cyclization as is shown in 22. A stereodiagram of \

a portion of the crystal structure is shown in 23.

Recalling the Woodward-

Hoffmann rules that such reactions occur in a disrotatory manner one sees that

i

.~/C~R

J 21

of the two possible dirotatory pathways, one of the two (see 22) causes the rearranging groups to push into the neighboring host molecule. The reaction therefore proceeds in only one of the two stereochemical pathways.

Ph

HO Ph C1 l

O

O

hv

O

O

&O" "'"H

(IS, 5R)-(-)-22 R=Et

(IR, 55")-(+)-22

22

23

91

While these chiral host-guest inclusion compounds have been demonstrated to produce excellent ee's it has proven difficult to predict in an apriori fashion the direction of enantiomeric preference.

CH3

For example in the cyclization CH3

reaction of 21, host 16b produces one enantiomer in 98% ee while 16c produces the opposite enamtioner with 95% ee. 39 It should be noted that 16b and 16c are rather similar and it is therefore difficult without further 4o work to ascribe the exact cause for the different product outcome.

I

CH3

24

Another example of such difficulties is given in reference 36. One of the reasons for this lack of predictability is that in every host-guest system a new crystal must be made. It is easy to imagine that it would be difficult to transfer one stereo-orientation from one crystal to the next. It may prove in the future that the use of chiral porous organic solids as hosts will obviate this latter concern.

Acknowledgement A part of our work described here (Compounds 3-6) was done under the guidance of Prof. Jeffrey S. Moore at the University of Illinois at Urbana-Champaign.

We thank the National

Science Foundation (Grant CHE-94-23121) for financial support. A portion of our research was carried out at the Center of Microanalysis of Matierals, University of Illinois, which is supported by the U. S. Department of Energy under Grant DEFG02-91-ER45439. S.L. Thanks the J. D. and C. T. MacArthur foundation (1993-97) and the A. P. Sloan Foundation (1993-95) for fellowships. S.L. and D.V. thank Prof. Jeffery S. Moore for helpful discussions. We also thank Messrs. G. B. Gardner and Y. -H. Kiang for their contributions.

92

References and Notes

1.

(a) Newsam, J. M. Science 1986, 231, 1093. (b) Breck, D. W. Zeolite Molecular Sieves; Robert E. Krieger Publishing Co.: Malabar, FL, USA, 1974.

2.

(a) Hoskins, B. F.; Robson, R. J. Am. Chem. Soc. 1990, 112, 1546. (b) Gardner, G. B.; Venkataraman, D.; Moore, J. S.; Lee, S. Nature 1995, 374, 792. (c) Fujita, M.; Kwon, Y. J.; Washizu, S.; Ogura, K. J. Am. Chem. Soc. 1994, 369, 727. (d) Endo, K.; Sawaki, T.; Koyanagi, M.; Kobayashi, K.; Masuda, H.; Aoyama, Y. J. Am. Chem. Soc. 1995, 117, 8341. (e) Wang, X.; Simard, M.; Wuest, J. D. J. Am. Chem. Soc. 1994, 116, 12119.

3.

As emphasized in the text, the importance of a zeolite derives from its function. It is not only important to realize a structure which retains its cavities upon removal of the guest but should also demonstrate reversible absorption and desorption of guests, selectivity in the intake of guest molecules and stability under conditions of a typical reaction.

.

For a review of Dianin's compound and related systems see: MacNicol. D. D. Inclusion Compounds; Atwood, J. L., Davies, J. E. D., MacNicol, D. D., Eds.; Oxford University

Press: Oxford, 1984; Vol. 2, Chapter 1, pp 1-46. 5.

Powell, H. M.; Wetters, B. D. P. Chem. Ind. 1955, 256.

6.

Since 1971, a number of crystal structures of various adducts of 1 have been reported in the literature.

The atomic coordinates of the chloroform adduct is used as the starting

point for other adducts. For the atomic coordinates of the chloroform adduct see: Flippen, J. L.; Karle, J.; Karle, I. L. J. Am. Chem. Soc. 1970, 92, 3749. 7.

Note that as the thermal motion and occupancy factors are correlated, it is difficult to differentiate between low occupancy and high disorder.

A conclusive proof about the

occupancy of the can be obtained from the chemical shift information in NMR spectroscopy. 8.

(a) MacNicol, D. D.; McKendrick, J. J.; Wilson, D. R. Chem. Soc. Rev. 1978, 7, 65-87. (b) MacNicol. D. D. Inclusion Compounds; Atwood, J. L., Davies, J. E. D., MacNicol, D. D., Eds.; Oxford University Press: Oxford, 1984; Vol. 2, Chapter 1, p 14.

9.

Barrer, R. M.; Shanson, V. H. J. Chem. Soc., Chem Commun. 1976, 333.

93 10.

Brunauer, S. The Adsorption of Gases and Vapors; Princeton University Press: Princeton, 1965. Type IV and type V behavior have also been observed in other porous compounds. See Allinson, A.; Barrer, R. M. J. Chem. Soc. A 1969, 1718.

11.

(a) Lee, F.; Gabe, E.; Tse, J. S.; Ripmeester, J. A. J. Am. Chem. Soc. 1988, 110, 6014. (b) Pang, L.; Lucken, E. A. C.; Bemardinelli, G. J. Am. Chem. Soc. 1990, 11, 8754.

12.

(a) Ung, A. T.; Bishop, R.; Craig, D. C.; Dance, I. G. Scudder, M. L. J. Chem. Soc., Chem. Commun. 1991, 1012. (b) Ung, A. T.; Gizachew, D.; Bishop, R.; Scudder, M. L.; Dance, I. G.; Craig, D. C. J. Am. Chem. Soc. 1995, 117, 8745.

13.

For a review of helical canal inclusion networks see: Bishop, R.; Dance, I. G. Top. in Curr. Chem. 1988, 149, 137.

14.

Mesitylene is considered as too bulky a solvent to be included in the cage.

15.

(a) Hoskins, B. F.; Robson, R. J. Am. Chem. Soc. 1990, 112, 1546. (b) Moore. J. S.; Lee,S. Chem. Ind. 1994, 556.

16.

Based on this approach, few inclusion compounds have been made. see References 2a-c and 24.

17.

Venkataraman, D.; Gardner, G. B.; Lee, S. Moore, J. S. J. Am. Chem. Soc. 1995, 117, 11600.

18.

We could not collect data after 25" in 20 for the single crystal.

Hence a complete

structural solution and refinement was not possible. 19.

The second polymorph, polymorph B, can be considered as homeotypic with CaCuP (A1B2-type) structure. The structure consists of undulating hexagonal sheets. Five such nets interpenetrate to file the void space in a single net. see Venkataraman, D.; Lee, S.; Moore, J. S.; Zhang, P.; Hirsch, K. A.; Gardner, G. B.; Covey, A. C.; Prentice, C. L. Submitted for publication to Chemistry of Materials.

20.

(a) Wang, X.; Simard, M.; Wuest, J. D. J. Am. Chem. Soc. 1994, 116, 12119. (b) Duchamp, D. J.; Marsh, R. E. Acta Crystallogr. 1969, B25, 5. (c) Ermer, O.; Lindenberg, L. Helv. Chim. Acta. 1991, 74, 825. (d) Ermer, O. J. Am. Chem. Soc. 1988, 110, 3747.

21.

Gavezzotti, A.; Desiraju, G. R. Acta Crystallogr. 1988, B44, 427.

22.

Asgaonkar, A.; Gardner, G. B.; Kiang, Y. -H.; Lee, S.; Moore, J. S; Venkataraman, D. Manuscript in preparation.

94 23.

The lattice constants for the original orthorhombic cell were a -- 11.625 A, b - 19.110 ,~, c = 38.856 A. The lattice constants for the rectangular cell of the guest-free sample were b = 22.76/~ and c = 36.48 ]k.

24.

The exposure temperature was 40 ~ for benzene, 65 ~ for benzyl alcohol and 60 ~ for other guest molecules.

25.

Yaghi, O. M.; Li, G.; Li, H. L. Nature 1995, 378, 703.

26.

For other important examples of networks which does not collapse upon guest removal see: (a) Cartraud, P.; Cointot, A.; Renaud, A. J. Chem. Soc., Faraday Trans. 1 1981, 77, 1561. (b) Allinson, S. A.; Barrer, R. M. J. Chem. Soc. A 1969, 1717.

27

Wilson, K. R.; Pincock, R. E. J. Am. Chem. Soc. 1975, 97, 1474.

28.

Penzien, K.; Schmidt, G.M.J. Angew. Chem. Int. Ed. Engl. 1969, 8, 608.

29.

a) Green, B. S.; Lahav, M.; Rabinovich, D. Acc. Chem. Res. 1979, 29, 187. b) Addadi, L.; Lahav, M. Pure Appl. Chem. 1979, 51, 1269.

30.

Wudl, F.; Lightner, D. A.; Cram, D.J.J. Am. Chem. Soc. 1967, 89, 4099.

31.

Green, B. S.; Lahav, M.; Schmidt, G.M.J. Mol. Crvst. Liq. Cryst. 1975, 29, 187.

32.

a) Elgavi, A.; Green, B. S.; Schmidt, G.M.J.J. Am. Chem. Soc. 1973, 95, 2058. b) Heller, E.; Schmidt, G. M.J. Isr. J. Chem. 1971, 9, 449.

33.

a) Green, B.S.; Lahav, M.; Schmidt J. Am. Chem. Soc. 1973, 95, 2058. b) Cookson, R. D.; Franklel, J. J.; Hudec, J. Chem. Soc., Chem Commun. 1965, 16. c) Adam, G. Tet. Lett. 1971, 2030. d) Addadi, L.; Gati, E.; Lahav, M.; Leiserowitz, L. Isr. J. Chem. 1976-

1977, 15, 116. 34.

a) Addadi, L.; Cohen, M.D.; Lahav, M. Mol. Cryst. 1976, 32, 137. b) Addadi, L.; Cohen, M.D.; Lahav, M. J. Chem. Soc., Chem. Commun.1973, 471.

35.

Addadi, L.; Lahav, M. J. Am. Chem. Soc. 1979, 101, 2152. b) Addadi, L.; van Mil, J.; Lahav, M. J. Am. Chem. Soc. 1982, 104, 3422.

36.

a) Evans, S. V.; Miguel, G.-G.; Omkaran, N.; Scheffer, J. R.; Trotter, J.; Wireko, F. J. Am. Chem. Soc. 1986, 108, 5648. b) Sekine, A.; Hori, K.; Ohashi, Y.; Yagi, M.; Toda, F. J. Am. Chem. Soc. 1989, 111,697. c) Roughton, A. L.; Muneer, M.; Demuth, M.; Klopp,

I. Krtiger, C. J. Am. Chem. Soc. 1993, 115, 2085. d) Kaupp, G.; Haak, M. Angew.

95 Chem. Int. Ed. Engl. 1993, 32, 694. e) Toda, F. Synlen. 1992, 303. f) Toda, F.; Tanaka,

K.; Stein, Z.; Goldberg, I. Acta Cryst. 1995, 351,856. g) Ramamurthy, V.; Venkatesan, K. Chem. Rev. 1987, 87, 433.

37.

For reviews see (a) Toda F. Acc. Chem. Res. 1995, 28, 480, (b) Toda F. Synlett. 1992, 303 and references therein.

38.

Kaftory, M.; Tanaka, K.; Toda, F. J. Org. Chem. 1988, 53, 4391.

39.

Tanaka, K.; Kakinoki, O.; Toda, F. J. Chem. Soc., Chem. Commun. 1992, 1053.

40.

Hashizume, D.; Vekusa, H.; Ohashi, Y.; Matsugawa, R.; Miyamoto, H.; Toda, F. Bull. Chem. Soc. Jpn. 1994, 67, 985.

This Page Intentionally Left Blank

H. Chon, S.I. Woo and S.-E. Park (Editors) Recent Advances and New Horizons in Zeolite Science and Technology Studies in Surface Science and Catalysis, Vol. 102 9 1996 Elsevier Science B.V. All rights reserved.

Spectroscopic

Characterisation

of

97

Zeolites

Russell F Howe Department of Physical Chemistry University of New South Wales Sydney, Australia 2052

1.INTRODUCTION This review is concerned primarily with optical spectroscopic methods for characterising zeolites and molecules adsorbed in zeolites. The electromagnetic spectrum spans the range from radiofrequencies to Xradiation. Spectroscopic techniques included in this range are, in order of increasing frequency, NMR, EPR, infrared, UV-VIS and Raman, XPS, XAS and Mossbauer spectroscopies. The NMR and XPS techniques are discussed elsewhere in Chapter 4 (B and C respectively). This review will give an overview of recent advances in EPR, infrared, Raman, UV-VIS and X-ray absorption spectroscopies as applied to zeolitic materials. Brief mention will also be made of new techniques recently reported for obtaining mass spectra of zeolites and adsorbed species. Each of the spectroscopic techniques has an extensive literature which cannot possibly be reviewed within the confines of a single book chapter. The reader seeking introduction to the principles of the methods and prior literature is referred to references [1-3] at the end of the chapter. It is well understood by experienced practitioners of zeolite science that no problem can be fully solved by applying a single spectroscopic or physical technique. That theme will be repeated here; complete characterisation of a zeolite system requires a combination of chemical, physical and spectroscopic experiments, and it is necessary to have a good appreciation of the advantages and limitations of each technique. Where disagreements exist in the literature, it is often because different techniques under different conditions are giving different answers.

98

2.EPR SPECTROSCOPY 2.1

Background Electron paramagnetic resonance (or electron spin resonance) was first applied to zeolites more than 30 years ago, and continues to be utilised. EPR can in principle observe any species containing one or more unpaired electrons; in the case of zeolites this means paramagnetic transition metal or rare earth ions, solid state defects which are paramagnetic, and inorganic or organic radicals. The method is highly sensitive, detecting as few as 101 2 spins in favourable circumstances.This sensitivity does impose the need for caution in interpreting weak signals, which can be caused by trace level impurities irrelevant to the chemistry being studied. EPR spectroscopy has not seen the dramatic changes in instrumentation and techniques that have characterized NMR spectroscopy over the past 20 years. Until recently, commercial EPR instruments used exclusively the continuous wave m i c r o w a v e techniques first developed 50 years ago. The advent of digital electronics and computer controlled spectrometers has made the experiment easier to do, more reliable, and arguably improving signal to noise ratios, but the principle of measurement remains that of measuring (in derivative mode) the absorption of microwave radiation at a fixed frequency as the magnetic field is continuously scanned. The physics of the continuous wave EPR experiment are summarised in the spin Hamiltonian j~

= gl3H

+

aS.I

+

D S1S2

All three terms in the Hamiltonian contain chemical information. The g tensor is determined by the ground and excited state wave functions of the atom or molecule containing the unpaired electron(s) (strictly speaking the excited state wave functions coupled to the ground state through spin orbit interactions). It is thus sensitive to factors influencing the ground and excited state wave functions, such as the crystal field in the case of transition metal ions, or the electrostatic fields in zeolite cavities for adsorbed radicals. The components of the g tensor determine the magnetic field at which resonance occurs for a given microwave frequency. Almost all zeolite studies have used polycrystalline powders rather than single crystals, which means that the observed spectra are a summation of spectra from individual crystallites at all possible orientations with respect to the magnetic field. If the principal components of the g tensor are sufficiently different, the first derivative powder spectrum will show three clearly

99 resolved features at the magnetic fields corresponding to the g t e n s o r components. More commonly, computer simulation of the powder spectrum is necessary to extract reliable parameters. The isotropic and anisotropic hyperfine coupling terms in a arise from interactions b e t w e e n electron and nuclear spins, and provide information about the nature of the orbital containing the unpaired electron and the extent to which it overlaps with orbitals on adjacent atoms. The anisotropic term can cause similar difficulties to the g tensor anisotropy in analysing spectra of polycrystalline powders; extracting coupling constants from spectra of transition metal ions or radicals in zeolites can be difficult or impossible without computer simulation. The third term in the spin Hamiltonian is the so-called zero-field splitting term which arises in systems containing more than one unpaired electron. This is often encountered in spectra of transition metal ions, and can complicate measurement and interpretation of spectra. Depending on the magnitude and sign of the zero-field splitting, signals may be broad and poorly resolved, not detected at all, or may show anomalous temperature dependence (an example is high spin Co 2+ with three unpaired electrons, which is considered below). Long range dipolar interactions between an unpaired electron and nuclear spins on adjacent atoms will not normally be resolved in conventional powder EPR spectra.The pulse technique of electron spin echo modulation (ESEM) is in favourable cases able to detect very weak hyperfine interactions not seen in CW EPR. The method measures modulation of the electron spin echo signal by dipolar hyperfine coupling in the time domain at fixed magnetic field. Until recently, pulsed EPR experiments required home built instrumentation, and in the zeolite field the method has been restricted to a small number of practitioners. Commercial pulsed EPR instruments are however now available, and the ESEM technique and variants thereof will undoubtedly grow in popularity.

2.2

EPR Spectroscopy of Adsorbed Radicals Organic and inorganic radical formation in zeolites can occur spontaneously, on adsorption of molecules into a suitably activated zeolite, or as the result of radiolysis of adsorbed species. Once a radical is formed, EPR spectroscopy can be used to follow its subsequent reactions. For example, Trifunar et al have recently described the use of variable temperature EPR to investigate reactions of olefin radical cations generated in ZSM-5 zeolites. [4] . This work shows clearly the facile rearrangement of radical cations produced by irradiation of

100 olefins adsorbed in zeolites at low temperatures. Figure 1 shows EPR spectra observed at 100K and 160K after irradiation of 1,4 cyclohexadiene adsorbed in NaZSM-5 at 77K. The initial spectrum recorded is that of the 1,3 cyclohexadiene radical cation, confirmed by the calculated spectrum shown below. The spectrum of the parent 1,4 cyclohexadiene radical cation could not be detected at all, even when irradiation was carried out at 4 K. The isomerization of the 1,4 cyclohexadiene radical cation to the 1,3 cyclohexadiene radical cation is known to be energetically favourable in the gas phase, but does not occur in low temperature inert gas matrices in the absence of photo excitation. It is remarkable that reaction occurs so readily in zeolite pores. Trifunac et al. suggest that the reaction may be driven by excess energy remaining after ionisation of the parent molecule, however radical : zeolite interactions clearly also play a role. I

b)

c)

d)

t 50,G

I

_A

F i g u r e 1. EPR spectra recorded at (a) 100K and (c) 160K after irradiation at 77K of 1,4-cyclohexadiene in NaZSM-5. Spectra (b) and (d) are calculated spectra for the 1,3-cyclohexadiene radical cation and the cyclohexadienyl radical respectively. (From r e f e r e n c e 4, with permission).

101 The spectrum obtained after warming to 160 K is that of the neutral cyclohexadienyl radical (confirmed by the computer simulation shown). Transformation of the radical cation to a neutral radical is commonly observed in condensed phases and occurs via ion 9 molecule reactions. A similar explanation is proposed for the adsorbed species in NaZSM-5" formally proton transfer from the cyclohexadiene radical cation to an adsorbed cyclohexadiene molecule will produce the cyclohexadienyl radical and a diamagnetic carbenium ion. In HZSM-5 the Bronsted acid protons can also contribute to the chemistry occurring. The conversion of 1,3 cyclohexadiene radical cation to the cyclohexadienyl radical occurs in a similar manner to that in NaZSM-5, but at higher temperatures a third signal due to the cyclohexenyl radical is detected. In this case carbenium ions generated by protonation of cyclohexadiene at Bronsted sites act as scavengers of free electrons produced during radiolysis, forming H addition type radicals. The chemistry of these low temperature reactions of radicals and radical cations is certainly highly relevant to the catalytic chemistry of hydrocarbons in zeolites and warrants further detailed study by EPR spectroscopy in parallel with other spectroscopic techniques. The recent review by Rhodes [5] should be consulted for information about other work in this area.

2.3

Metal Ion Exchanged Zeolites EPR spectroscopy has been used with good effect to follow the state of Cu in Cu exchanged MFI zeolites used as catalysts in the selective reduction of NOx with hydrocarbons in the presence of oxygen. These catalysts, first reported by Iwamoto et al. [6] and Held et al. [7], have since become widely studied because of their high activity and selectivity to nitrogen [8]. Questions addressed using EPR spectroscopy include the location of the Cu in the zeolite, the extent of clustering of Cu at high Cu exchange levels, the oxidation state of the Cu, and the extent to which all of these vary during catalyst pretreatment or use. Kucherov et al.[9] have shown that at low Cu levels, all of the Cu exchanged into the zeolite is detected as isolated Cu 2+ cations existing in two different coordination states: four coordinate square planar and five coordinate square pyramidal. Figure 2 shows a typical spectrum of Cu 2 + exchanged MFI , showing the two different Cu 2+ species.[10]. Both species are located in the main channels of the MFI structure [9]. At higher Cu exchange levels, the EPR spectra become poorly resolved due to the onset of Cu:Cu interactions. Some authors have shown that heating Cu exchanged MFI zeolites in vacuo causes spontaneous reduction of Cu 2+ to Cu + and loss of the EPR signals [11,12]. Shelef [8] argues that it is only coupled Cu species that can be spontaneously

102 reduced in vacuo. In the presence of oxygen (ie. under conditions similar to those of the working catalyst) the isolated Cu 2+ species are not reduced by hydrocarbons such as propene. Shelef et al. have shown very recently using EPR spectroscopy that steam treatment at 650 C or calcination at higher temperatures causes an irreversible transformation in the coordination state of the Cu 2+ , suggested to involve a rearrangement of the local topography of the isolated Cu 2 + sites [13]. There is disagreement as to whether this is caused by dealumination of the zeolite [14]; Sachtler [15] has suggested that a Cu aluminate phase is produced in hydrothermally treated zeolites. More in-situ EPR studies of the type conducted by Shelef et al. are needed before the remaining questions concerning the Cu MFI catalysts can be answered.

I'

!

i

2600

I

I

I

I ~

.,,

I

3000

t/

I

.

I

3400

[G]

Figure 2. EPR spectrum of Cu 2+ in CuZSM-5 NO reduction catalyst.[10] 2.4

Metal Substituted Zeolites The synthesis of zeolitic materials containing transition elements substituted into the framework has become of considerable interest. A crucial issue for such materials is the question of whether the material

103 as synthesised really does contain the transition element substituted into the framework, and whether or not it remains in the framework following activation or use in a catalytic reaction. The sensitivity of the spin Hamiltonian parameters for a paramagnetic transition metal ion to the crystal field environment means that EPR spectroscopy should be able to clearly distinguish between a metal ion in the framework with tetrahedral or distorted tetrahedral coordination and an extra framework species with octahedral or distorted octahedral coordination. The formation of extra framework metal oxide or hydroxide clusters will also be evident from the EPR spectra in that such species will contain strongly coupled spins. The recent EPR study by Kevan et a1.[16] of cobalt substituted ALPO-5 illustrates the power of the EPR technique to characterise metal substituted zeolites. CoAPO-5 as synthesised is blue in colour; the electronic spectrum is characteristic of tetrahedrally coordinated Co 2+, suggesting that Co 2+ has been incorporated into the A1PO4 lattice. The corresponding generation of Bronsted acid sites indicates that Co 2 + substitutes for A13+.The material does however change colour to yellowgreen on calcination at high temperature; this has been interpreted to mean that cobalt is partially oxidised to Co3+ under these conditions.[17] Figure 3 shows EPR spectra measured by Kevan et al. at 7 K of CoAPO-5 before and after calcination. The anisotropic signal observed is characteristic of high spin Co 2+ in tetrahedral coordination (and because of effective spin lattice relaxation in the spin = 3/2 system can only be detected below liquid nitrogen temperatures). At first sight, the decrease in signal intensity on calcination would seem to support the suggestion that Co 2+ is oxidised to Co 3+ on calcination. However, Kevan et al. showed that the apparent signal loss is in fact due to a change in the temperature dependence of the signal , as illustrated in Figure 4. When the spectrum is measured at 20K there is no difference in intensity between the as synthesised and calcined samples. The as synthesised material follows approximately the normal Curie law dependence of signal intensity on temperature, whereas the calcined material deviates from Curie law behaviour below 20K. This behaviour can be accounted for in terms of distortion of the tetrahedral symmetry of Co 2+. Such distortion can induce a negative zero field splitting between the Ms= + 1/2 and Ms= + 3/2 doublets. As the temperature decreases, the Co 2+ ions tend to populate the lower energy + 3 / 2 doublet, which is EPR silent, thus diminishing the signal intensity. The calculated curve in Figure 4 assumes a zero field splitting o f - 1 3 cm-1, and agrees well with the experimental data. The exact cause of the symmetry distortion on calcination is not yet clear, although Kevan et al.

104 suggest that it may involve coordination of molecular oxygen to the Co 2+ ions in the ALPO-5 lattice.

500 G S

,

"

-

slier ~ l ~ n s l i o n --_

--.-

g " 5.45

14 ,

12

~

8

~

4

m

2

.........

Curie Law Calculated

0 ',, X \

b.

C]

After calcination

0

As-synthesized

W

0

10

20

30

40

T,K

Figures 3 and 4. EPR spectra of Co2+ at 7K and temperature dependence of signal intensity from as synthesized and calcined CoAPO-5. Reproduced with permission from reference 16. A second example of the use of EPR spectroscopy to investigate lattice substitution in zeolites is the recent work of Schoonheydt et al. on VAPO-5 materials [18]. The EPR spectra of VAPO-5 containing low concentrations of vanadium show two overlapping signals which both have the g-tensor and 51V hyperfine parameters typical of distorted octahedral V 4+. An additional broad poorly resolved signal is detected in higher vanadium content samples. The broad signal is undoubtedly due to extra lattice vanadium oxide species in which V4+ ions are strongly magnetically coupled. The isolated octahedral V4+ species may

105

be anchored to the surface of the zeolite, but the absence of any spectroscopic signature of V4+ tetrahedrally coordinated to 4 oxide ions of the lattice suggests that isomorphous substitution of V 4+ into the lattice does not occur. 2.5

ESEM Studies of Metal Ions in Zeolites The electron spin echo modulation technique detects directly the coordination environment around a paramagnetic ion by observing the dipolar coupling to nuclei of weakly coordinated ligands. This technique has been used extensively by Kevan, for example, to examine transition metal ion exchanged and substituted zeolite materials [19].

I0

8

6 ,,i,,o w

r c: -ira

.=_ O ..g:

,., ILl

2

o

t

z

3

4

T,/~s

Figure 5. Three pulse ESEM recorded at 4 K of (a) MnAPO-5 and (b) MnAPO-11 containing adsorbed D20. Reproduced with permission from reference 20.

106 Figure 5 illustrates the method showing three pulse ESEM traces for MnAPO-5 and MnAPO-11 both containing adsorbed D20, from the work of Brouet, Chen and Kevan [20]. The modulation observed on the spin echo decay curves is caused by dipolar hyperfine coupling between the unpaired electrons on Mn 2+ and the nuclear spins of deuterium atoms in surrounding water molecules. The modulation patterns are distinctly different for the two different ALPOs. The pattern for MnAPO-5 could be fitted with a model of 4 deuteriums at a distance of 0.30nm, consistent with solvation of two D 2 0 molecules about a Mn 2+ cation in a non-framework position. For MnAPO-11, on the other hand, the modulation pattern required a model of two deuterium atoms at a closer distance of 0.24nm and two at a longer distance of 0.36nm. This model is consistent with coordination of two water molecules to a negatively charged site, where the deuterium atoms of each water molecule are not equidistant from the manganese, suggesting (albeit indirectly) that the Mn 2+ in MnAPO-11 is substituting isomorphously for aluminium in the ALPO lattice.

3. INFRARED SPECTROSCOPY 3.1

Background Infrared spectroscopy, like EPR, has been used for more than 30 years to study zeolites and adsorbed molecules. The introduction of Fourier transform techniques has made the conventional transmission measurement in the mid infrared region much easier and faster to perform , which has allowed new experiments to be performed which were either impossible or extremely difficult with wavelength scanning instruments. For example, in situ measurements of zeolite catalysts under reaction conditions are now routinely performed.Time resolved measurements on the scale of seconds or less are now feasible, and spectra can be measured routinely in the energy limited far infrared region. Another exciting new development is infrared microspectrocopy, which permits spectra to be recorded from zeolite single crystals.Some of these newer developments will be illustrated in the examples which follow. 3.1

Zeolite Lattice Modes The vibrational frequencies of the so-called lattice modes of aluminosilicate zeolites (stretching and bending modes of the T-O linkages, plus specific vibrations of discrete structural units) were first studied in detail by Flanigen more than 20 years ago [21]. The lattice modes are sensitive to both the composition of the lattice and the structure. For example, Jacobs et al. showed that the T-O stretching

107

frequencies of different zeolites correlated with the average electronegativities of the zeolite frameworks [22]. For zeolite Y, a linear correlation has been reported between the frequencies of selected lattice modes and the aluminium content of the zeolite [23]. The lattice mode frequencies are also sensitive to the aluminium content of the lattice in the case of MFI zeolites. Campbell et al. [24] have recently reported that the lattice mode frequencies can be used to monitor changes in the lattice aluminium content of HMFI catalysts after hydrothermal treatment or use in the conversion of methanol to hydrocarbons. Figure 6 shows plots from reference 24 of the frequencies for two of the lattice modes ( t h e structure sensitive 544 c m -1 band, due to deformation modes of the MFI lattice, and the 790 c m -1 band, a structure insensitive T-O-T stretching mode) versus the lattice aluminium content determined by solid state 27A1 and 2 9 S i NMR for fresh and treated MFI zeolites. The correlation in this case is not linear, but is still a useful empirical guide to the extent of lattice dealumination.

5,55

'

'

' 0

9 a , ~ ,]~:83H[-6 .,.~~.,~

T_ 5~o

~

parent

i 805 I

I

I

545

,

i

'

,

1~

coo

-

795

0

2

4-

6

lattice .~lnm~ntum Content/.~1 (uc) -1

Figure

8

790

0

2

4

6

8

Lattice Alumtnlum C o n t e n t / A1 (ue) -1

6. Dependence of lattice frequencies on lattice aluminium content for HMFI zeolites. Reproduced with permission from reference 24.

108 The sensitivity of lattice modes to structural changes is illustrated by the recent study of Mueller and Connor [25] on the effects of cyclohexane adsorption on the structure of MFI zeolites. The adsorption of molecules such as paraxylene and benzene into MFI zeolites causes a structural transition from monoclinic to orthorhombic symmetry, which has been characterized by X-ray powder diffraction and 29Si NMR spectroscopy [26]. Cyclohexane has a slightly larger kinetic diameter than benzene or paraxylene (0.60 nm compared with 0.585nm), but does not cause the same structural transition. Cyclohexane adsorption does however affect the zeolite lattice mode vibrational frequencies. Figure 7 shows spectra taken from reference 25 before and after (upper spectrum) adsorption of cyclohexane in a low aluminium MFI zeolite. The OTO bending mode at 547.5 cm -1 is shifted to higher frequency and a new band appears at 401.6 cm-1 on adsorption of cyclohexane. Mueller and Connor attribute these changes to a flexing of the MFI lattice, causing small changes in the OTO bond angles (the 401.6 cm-1 band is tentatively attributed to a pore mouth vibration shifted up in frequency from below the 350cm -1 cut off of the spectrometer used). The amount of cyclohexane adsorbed is only half that of benzene or paraxylene; it is suggested that cyclohexane adsorbs only in the linear channels of the MFI structure, perturbing the lattice modes but not inducing the structural transformation to orthorhombic symmetry.

~SM-S

(3OO)

45S.~ 454.7---.~ ~

\

o

401.6

600 I

I

560 I

I

520 I

I

480 I'

I

440 .I

I

400 380 I

Wave Number

F i g u r e 7. Infrared spectra of ZSM-5 before and after adsorption of cyclohexane . Reproduced with permission from reference 25.

109

3.3

Low Frequency Cation Vibrations The far infrared region of zeolite spectra (here defined as the frequency region below 300 cm ~ shows bands which are attributed to stretching vibrations of the zeolite cations vibrating relative to the zeolite lattice. The vibrational frequencies of the cations depend on their charge, mass and interaction with the zeolite. In the case of zeolite NaY, for example, 4 intense bands are observed which have been assigned to Na + cations in the 4 different cation sites SI, SI', SII and SIII [27]. Esemann and Forster [28] have recently used far infrared spectroscopy to follow the progress of solid state ion exchange into zeolites HY and NH4Y and to compare solid state ion exchange with conventional aqueous ion exchange. Figure 8 shows far infared spectra from reference 28 of zeolite NH4Y after solid state ion exchange by heating in the presence of solid NaC1, KC1 and RbC1 respectively. In all three cases the spectra obtained by solid state exchange are identical to those observed after aqueous ion exchange. The 4 bands due to Na + cations in sites SI (160 cm-1), SI' (106 cm -1), SII (188 cm -1) and SIII ( 90 cm -1) are all shifted to lower frequencies by the amounts expected for the change in reduced mass when Na + is replaced by K + or Rb +. In the case of Cs +, the solid state exchange produced a band due to Cs + cations occupying site I in the double six rings of the faujasite structure (at 86 cm -1) which was not formed after conventional aqueous ion exchange with Cs +. This can be explained by the inability of the large hydrated Cs + cation to enter the double six-rings.

KY .8NaY .6-

.4-

.2-

:3;0

2;0

200 1;0 Wavenumbers

1;0

F i g u r e 8. Far-infrared spectra of zeolites prepared by solid state ion exchange. Reproduced with permission from reference 28.

110 The far infrared experiments also showed that the effectiveness of solid state ion exchange depends on the starting form of the zeolite. N H 4 Y allows effectively complete ion exchange on reaction with metal halide salts, due to the initial formation of NH4C1. With HY, on the other hand, it is difficult to achieve exchange levels beyond 50% due to the immediate formation of HC1 which dealuminates the zeolite to a considerable extent.

3.4.

Infrared Spectra of Probe Molecules Infrared spectroscopy has been used for many years to probe acid sites in zeolites. Typically, strong bases such as ammonia or pyridine are adsorbed, and the relative or absolute intensities of bands due to Lewis acid adducts or protonated Bronsted acid adducts are measured. The basicity of ammonia or pyridine is however much stronger than that of most hydrocarbon reactants in zeolite catalysed reactions. Such probe molecules therefore detect all of the acid sites in a zeolite, including those weaker acid sites which do not participate in the catalytic reaction. Interest has recently grown in using much more weakly basic probe molecules which will be more sensitive to variations in acid strength. It is also important in studying smaller pore zeolites to use probe molecules which can easily access all of the available pore volume. Nitrogen and carbon monoxide are both candidates as small probe molecules which may interact only with strong acid sites in zeolites and which can be observed by infrared spectroscopy. As an illustration of this method, consider the recent work of Wakabayashi et al. on N2 adsorbed in H-mordenite [29], HY and HZSM-5 [30]. References to infrared spectra of adsorbed CO include [31-33]. Figure 9 shows spectra (from reference 29) of N 2 adsorbed in Hmordenite as a function of pressure at a temperature of l l0K. The two bands observed are both due to N-N stretching vibrations of molecularly adsorbed N 2 species; this was confirmed by recording spectra with 15N 2 and 15 N 14N mixtures. The higher frequency band at 2352 cm-1 saturated at low nitrogen pressures, whereas the 2335 cm -1 band increased in intensity with increasing nitrogen pressures up to 3 kPa. The two bands also showed a different temperature dependence; the 2335 cm -1 band decreased steadily with increasing temperature and was totally lost at 280K, whereas the 2352 cm -1 band was still present at this temperature. These observations indicate that the species responsible for the 2352 cm -1 band is more strongly adsorbed than that giving the 2335 cm -1 band. Wakabayashi et al. observed also

111 changes in the zeolite v(OH) bands during nitrogen adsorption at low temperatures, and found that the growth of the 2335 cm -1 band correlated strongly with loss of the 3616 cm -1 v(OH) band due to Bronsted acid sites and the appearance of a new v(OH) band at 3510 c m - 1 . The 2335 cm -1 band was thus assigned to N 2 hydrogen bonded to Bronsted acid sites, the hydrogen bonding causing a 106 cm -1 shift in the v(OH) frequency. There was no interaction of nitrogen with the silanol groups responsible for a v(OH) band at 3752 cm -1.

The 2352

c m -1 band was assigned to N 2 interacting with Lewis acid sites in the zeolite; this band was suppressed if the zeolite was pretreated with water vapour, but restored if the zeolite was outgassed at high temperature before exposing to nitrogen. 2335

s

2400

2300

Figure 9. FTIR spectra of nitrogen adsorbed in H-mordenite at 110K. Nitrogen pressures of (a) 0.3, (b) 0.5, (c) 1.6 and (d) 3.7 kPa. Reproduced with permission from reference 29. Similar results were found for nitrogen adsorbed in HZSM-5 and HY zeolites[30]. A 2352 cm -1 band was observed only under conditions (high temperature evacuation) where Lewis acid sites could be anticipated. The frequency of the N 2 interacting with Bronsted acid sites is identical in HZSM-5 to that in H-mordenite (2334 cm -1 c o m p a r e d with 2335 cm-1). In HY, the corresponding frequency is 2338 cm -1,

112 suggesting a stronger perturbation of the nitrogen molecule in this case (the gas phase frequency of N 2 is 2330 cm -1, measured by Raman spectroscopy). In summary, physisorbed nitrogen appears to offer several advantages as an infrared probe of acid sites in zeolites. It clearly d i s t i n g u i s h e s between Bronsted and Lewis acid sites without interference from gas phase species, it is small enough to probe sites in smaller pore zeolites, and its interaction with the zeolite is sufficiently weak and reversible to have negligible influence on the zeolite chemistry. It is not yet clear whether the method can probe variations in Bronsted acid strength. The infrared spectrum of adsorbed nitrogen can also be used to probe cation sites in zeolites. Zecchina et al [34] compared vibrational frequencies of CO and N2 adsorbed at low temperatures in mordenite containing different alkali metal cations. In both cases the vibrational frequencies could be correlated with (R x + Rm)-2, where R x is the cation radius and R m the radius of the adsorbed molecule, suggesting a simple electrostatic field explanation for the frequency shifts between different cations. The appearance of a band due to N 2 interacting with a particular zeolite cation will also mean that that particular cation is located in sites accessible to the N 2 molecule. In contrast to the enormous body of literature associated with acid catalysis by zeolites, base catalysis has until recently received little attention. Infrared spectroscopy of probe molecules can however be used to study basic sites in zeolites in an analogous manner to the experiments characterising acid sites. Barthomeuf [35] first suggested using pyrrole as an acid probe of basic sites, and showed that the (NH) frequency of adsorbed pyrrole correlated well with the negative charge on the oxygen atoms of different zeolites calculated from average electronegativities. More recently, Huang and Kaliaguine [36] have used this method to examine a series of X, Y, mordenite and ZSM-5 zeolites exchanged with different alkali metal cations. The infrared spectrum of adsorbed pyrrole is complex, and care must be exercised in identifying the v(NH) mode and eliminating contributions to the spectrum from other forms of adsorbed pyrrole besides those present at basic sites. The v(NH) frequency of the free pyrrole molecule is 3497 cm-1; adsorption in basic zeolites shifts this to lower frequency by up to 320 cm-1. In the series of zeolites LiX, NaX, KX, RbX and CsX, for example, the v(NH) frequency of adsorbed pyrrole decreases from 3295 cm -1 to 3175 cm -1. Furthermore, in zeolites containing two different alkali

113

metal cations, two different v(NH) bands can be resolved. Figure10 shows spectra of pyrrole adsorbed in respectively NaX and CsNaX containing 6, 29, and 60 Cs per unit cell. The 3280 cm -1 band associated with Na cations is gradually diminished as the Na content decreases, and the 3175 cm -1 band associated with Cs increases. Note that a band at 3375 cm -1 is present in all of the X zeolites studied and was attributed to a weak structure dependant basic site independent of the cations present.[36]

3175

c-

05

..Q 04

,::5

j

3375 I

q"l

(d)

I

3280

(c) (b) (a)

Figure 10. Infrared spectra of pyrrole adsorbed in (a) NaX, and CsNaX containing respectively (b) 6, (c) 29, and (d) 60 Cs per unit cell. Reproduced with permission from reference 36. Huang and Kaliaguine interpret spectra such as those in Figure 10 to mean that the strongly basic sites in alkali metal exchanged zeolites are framework oxygen atoms immediately adjacent to the alkali metal cations, acting as Lewis bases. Since the v(NH) frequency of pyrrole adsorbed on MgO surfaces is at 3320 cm -1, those zeolites for which the

114 corresponding frequency is below this value can be taken to have basic sites stronger than those of MgO. This is the case for all alkali metal exchanged X zeolites, CsY, RbY, Cs-mordenite and Rb-mordenite. Alkali metal exchanged ZSM-5 zeolites, on the other hand, show v ( N H ) frequencies for adsorbed pyrrole of 3370 cm -1, suggesting that these zeolites contain only weakly basic sites. Huang and Kaliaguine also attempted to improve the correlation established by Barthomeuf between the v(NH) frequency of adsorbed pyrrole and the partial charge on the zeolite oxygen atoms, by calculating a local electronegativity for a fragment of the zeolite consisting of a single 6 ring plus one or more cations. The partial charge on the oxygen atoms in the 6 ring is then determined by this local electronegativity r a t h e r than an electronegativity averaged over the entire zeolite lattice. Figure 11 compares the correlation between the v(NH) frequency and oxygen charges calculated from the local electronegativity (solid line) and the average electronegativity (dotted trace). The local description clearly gives a smoother correlation for the more basic zeolites.

3500

!

3400

!

Li

3450 I

i

I

i

I

Li

K, Nb, Cs ~ Na -

E

3350

-I-

z 3300

ZSM-5

Cs

3250

~,

mordenite

3200

3150 0.210

K ~ Li Rb ~ k, '~Na

Rbk~,

I

0.260

I

0.310

Rb !

0.360

Y

I

X

0.410 0.460

Cs

1

0.510

oxygen charge Figure 11. Correlation of (NH) frequency of adsorbed pyrrole with oxygen atom charge calculated from global (dotted line) or local (solid line) electronegativity. Reproduced with permission from reference 36.

115

3.5

Hydrogen Bonding versus Proton Transfer A key issue in the chemistry and catalysis of basic molecules reacting in acid zeolites is the extent to which proton transfer occurs from the Bronsted site to the basic molecule. For strongly basic molecules like a m m o n i a or pyridine, infrared spectroscopy clearly identifies the protonated adduct (NH4 + or PyH +) from its characteristic vibrational frequencies. For trimethylphosphine, also a strong base, both infrared and NMR evidence for complete proton transfer are convincing[37]. For molecules which are less strongly basic, the question is not so easily answered. Hydrogen bonding of adsorbed molecules to the Bronsted OH groups in acid zeolites causes a shift to lower frequency and increases in the half width and intensity of the v(OH) infrared band, and there have been many studies correlating the extent of the frequency shift with the basicity of the adsorbate or the strength of hydrogen bonding.Where the hydrogen bonding becomes very strong, the infrared spectrum becomes more complex and harder to interpret. Figure 12 shows, for example, difference spectra obtained when methanol is adsorbed in HZSM-5 zeolites containing different concentrations of Bronsted acid sites at 423K. The negative peak at 3610 cm -1 shows that the Bronsted acid sites are interacting with methanol. The intense broad bands at around 2800, 2400 and 1700 cm -1 correlate with the concentrations of Bronsted acid sites in the different zeolites, and were originally attributed to v(OH) and 8(HOH) vibrations of a protonated methanol species C H 3 O H 2 + ; ie. proton transfer was considered to occur from the Bronsted acid sites to adsorbed methanol [38,39]. Note that the spectra in Figure 12 also contain contributions from methanol reacting with silanol groups and A1OH groups associated with extra f r a m e w o r k aluminium to form respectively CH3OSi and CH3OA1 species [40]. Similar intense bands at approximately the same frequencies are observed when other molecules of comparable proton affinity are adsorbed in HZSM-5 e.g.water [41],dimethylether [38,40,41], acetone and various carboxylic acids [41]. Pelmenschikov et al. [42,43] pointed out that these bands are very similar to the so-called A,B,C triplet of OH bands characteristic of strong molecular hydrogen bonded complexes in liquid or solid phases. The most widely accepted explanation for the A,B,C triplet in hydrogen bonded systems is that due to Claydon and Sheppard [44], who suggested that the A,B,C triplet are in fact pseudobands caused by the superposition onto a very broad single (OH) band of two so-called Evans transmission windows caused by Fermi resonance between the (OH) mode and the first overtones of in-plane ( 2 6(OH) =__ 2600 cm -1) and out of plane ( 2 y(OH) ___- 1900 cm -1) bending

116 modes respectively. If this is origin of the A,B,C triplet for molecules adsorbed in HZSM-5, then proton transfer is not occurring from the Bronsted site to the adsorbed molecules.

o er

b

.t) I,..,

O


- 1 / 2 transition, which is only depending on the second order quadrupolar interaction. This interaction decreases with increasing magnetic field strength B o and can partly be reduced by MAS. However, complete removal of the second order quadrupolar interaction can be achieved by applying either DOR or DAS (see chapter 2.). Obeying Loewenstein's rule (forbidden A1-O-A1 linkages), zeolitic materials give rise to quite simple 27A1 NMR spectra, consisting of signals due to only one type of tetrahedral aluminium environment [A1 (OSi)4] besides evtl. octahedrally coordinated aluminium. No relationships between the chemical shifts and the Si, A1 ordering or Si/A1 ratio have been established, whereas relations between the 27A1 chemical shifts and mean Si-O-A1 bond angles exist, with shift values carefully corrected for quadrupolar shift contributions (11). The quantitative use of 27A1 NMR data allows determination of the relative proportions of lattice and extra-lattice A1 in zeolitic samples, providing that all aluminium is detected in the spectra (no signal intensity loss due to quadrupolar interactions, no "NMR invisible Ar'). The distinction between A1 species having the same chemical shifts and strongly overlapping lines but different quadrupolar couplings can be made via the two-dimensional quadrupole nutation NMR technique (13) (see chapter 2.)

163 No doubt, 27A1 MAS NMR is a valuable tool in probing the coordination, location and quantity of AI in zeolitic materials, but less useful than 29Si MAS NMR in detailed studies of the lattice structure (2). However, introduction of DAS and DOR allows much more detailed insight with respect to structural information available by use of 27A1 NMR spectroscopy. Since a number of excellent review papers exist, which summarize the structural information obtained by application of 27A1 NMR on zeolites, I would like to draw the readers attention for further details again to those reviews (1-16) and highlight only few results published recently dealing with framework characterization of zeolites or A1PO4 molecular sieves by using 27A1 NMR spectroscopy.

(a)

(c)

VPI-5

t,

60

I

40

,

!

i

,

,I

20 0 '20 ppm from AI(H20)~

I

-40

Fig. 16. 27A1 NMR spectra of dehydrated and partially rehydrated VPI-5. a) MAS spectrum of dehydrated VPI-5. b) DOR spectrum of dehydrated VPI-5. c) DOR spectrum after 2 days of rehydration, d) DOR spectrum after 23 days of rehydration (24). Structure of VPI-5 (right) (Reproduced by permission of Elsevier Science Publishers B.V., Amsterdam). VPI-5 has received great attention during the last years mainly due to the unusual structure and properties of this material with extra-large pores consisting of 18-membered tings. Wu et al. showed that 27A1DOR is capable of resolving discrete framework

164 aluminium sites in VPI-5, permitting quantitative investigation of site-specific adsorbate interactions (typically water interactions) with the framework. Figure 16 shows the 27A1 MAS NMR spectrum of the dehydrated VPI-5 and the DOR spectrum of the same sample (24). Two peaks unresolved in the MAS spectrum, at 33.3 ppm and at 35.9 ppm, are observed in the DOR spectrum. From the 1 : 2 intensity ratio of these two signals the highfield peak is assigned to A1 sites in the double four ring, whereas the low-field signal is assigned to A1 sites in the six-membered tings. During dehydration, 6-coordinated A1 sites are converted to 4-coordinated sites, consistent with the disappearance of the peak at -18.4 ppm. Furthermore, in hydrated VPI-5, 27A1 tetrahedral sites are altered by dehydration to yield different 27A1 tetrahedral environments (15). 31p NMR spectroscopy is considered as the most sensitive technique performing information on the local structure and structural modifications involving the tetrahedrally coordinated framework elements phosphorus and aluminium in A1PO4 molecular sieves. The 31p MAS NMR spectrum of VPI-5 is quite unusual and has been an item of discussion. The structure of VPI-5 consists of two crystallographically unique phosphorus sites, which are located in the 6-membered rings and in the atomic positions that belong to the two adjacent 4-membered tings and are in the atomic ratio of 2:1, respectively (see Figure 16). The 31p MAS NMR spectrum of dehydrated VPI-5 indeed consists of two lines at about-26/-27 and-31/-32 ppm in an area ratio that closely approximates 2:1 (1). However, the 31p MAS NMR spectrum of hydrated or as-synthesized VPI-5 reveals three lines of equal intensity at about -23, -27 and -33 ppm (1). McCusker et al. investigated the structure of hydrated VPI-5 by synchrotron X-ray, and concluded that two water molecules complete an octahedral coordination sphere around the framework aluminium atom between the fused 4-membered rings (38). Furthermore, all the water molecules are located in the 18- ring channels of VPI-5 forming a triple helix structure. According to a large number of investigations, the two low-field signals at-23 and-27 ppm are finally assigned to the P atoms in the 6-membered tings and the remaining high-field signal at -33 ppm represents the P-positions in the double-four tings (1). This assignment has been confirmed by Klinowski using 2D 31p NMR spin-diffusion spectra of hydrated VPI-5 (15), where only 31p_31p dipolar interactions govern the 31p spin diffusion process (see Figure 17). The stability of VPI-5 and the transformation to A1PO4-8, a 14-membered ring A1PO4 molecular sieve, have been extensively studied by XRD and solid-state NMR spectroscopy. Depending on the quality of the sample and the treatment conditions (evacuation, extremely slow heating rate, careful removal of water etc.), the structrue of VPI-5 can be preserved up to at least 400 ~ C, as followed by 31p MAS NMR (1). Under sealed conditions and slightly elevated temperatures (70 to 150 ~ the hydrated VPI-5 undergoes a reversible dehydration/rehydration effect, which results in a merging of the two low-field signals in the 31p MAS NMR spectrum (1) and splitting again after cooling to room temperature. Under more drastic conditions (higher temperatures, faster heating rates and unsealed conditions), VPI-5 transforms to A1PO4-8, monitored by recording of a complete different

165 31p MAS NMR spectrum (see Figure 18). The 31p MAS NMR spectrum of hydrated A1PO4-8 shows three signals at -21, -25 and -30 ppm, with an intensity ration of 1:2:6, which is not incompatible with the five independent crystallographic sites of A1PO4-8, if three of the sites are equivalent with respect to the NMR chemical shift (1).

Fig. 17.2D 31p NMR spin diffusion spectrum of hydrated VPI-5 (Reproduced by permission of Elsevier Science Publishers B.V., Amsterdam).

3.2

Mesoporous Materials

A new family of silicate/alurninosilicate mesoporous molecular sieves designated as M41S and kanernite were introduced a few years ago (39-43), and NMR investigations have been done to characterize those materials. MCM-41, one member of the M41S family (see Figure 19), exhibits a hexagonal arrangement of uniform mesopores whose dimensions may be engineered in the range of about 15 to greater than 100 A (39). The 29Si MAS NMR spectra of MCM-41 closely resemble those of amorphous silica, that means the spectra can be separated into three very broad signals at -89/-92 (assigned to Si (2OSi)(2OH) = Q2), -98/-102 (assigned to Si(3OSi)(OH) - Q3) and -108/-111 ppm (assigned to Si(4OSi) = Q4), with the two highfield signals dominating the spectrum (39, 44-50) (see Figure 20). MCM-48, consisting of a cubic arrangement of uniform mesopores, has essentially the same 29Si MAS NMR spectrum as amorphous silica or MCM-41, the only difference being the higher Q3/Q4 ratio (48-50)(see Figure 20).

166

-27.5

42.2 a

"22"611-33.3 ~

A

-30.3 33 1

b

40.9

b

2.9 -24.4 / \-14.4t, -29.8

36.9

c

#

-29.6 d

37.2

d

-24.8

~'

:~o''':2'd"36"-4'6 ppm

"-~6 '~-

al-P MAS NMRspectra

~so ~oo

so

o

ppm

-so -~oo

27-AI MAS NMR spectra

Fig. 18.31 p and 27A1 MAS NMR spectra of a) VPI-5 dried at 60 OC/ovemight, b) VPI-5 evacuated at 54 OC/overnight and calcined at 250 oC/ovemight, c) A1PO4-8 (transformed from VPI-5 by calcination at 400 OC/ovemight) and d) hydrated A1PO4-8. Asterisks denotes spinning side bands (Reproduced by permission of Elsevier Science Publishers B.V., Amsterdam).

167

Fig. 19. Structure of MCM-41 (Reproduced by permission of Verlag Chemie, Weinheim). -110 ppm -102 ppm I I

O/oQ4

Q3/Q4

36

61

0.59

51

47

1.1

56

41

1.4

O/oQ3

=

L

I

I

!

-50

-100

-150

Chemical shift (p.p.m.)

Fig. 20. 29Si MAS NMR spectra of pure siliceous, as-synthesized MCM-41, MCM-48 and MCM-50 (Reproduced by permission of Macmillan Magazines Ltd., London).

168 The relative number of incompletely condensed (Q2, Q3) and fully condensed (Q4) silicon atoms in MCM-41, MCM-48 and MCM-50 (which represents a lamellar structur) can be determined by 29Si MAS NMR, indicating the degree of the formed mesostructures. In other words, the Q3/Q4 ratio measures the extent of silanol condensation, which means lower values for more condensed frameworks (40, 48-50) (see Figure 20). 27A1 MAS NMR spectroscopy has been used to follow the formation of tetrahedrally and octahedrally coordinated aluminium during the synthesis of mesoporous materials, and the main focus has been concentrated on MCM-41. Tetrahedrally coordinated aluminium resonates at about 50 ppm, and these Al-sites are regarded as belonging to the lattice of MCM-41 (44-47). However, the detailed "fine" (or microscopic) structure of these mesoporous materials are not yet elucidated, and the common opinion of the nature of the wall material reflects an amorphous character. The absence of octahedrally coordinated aluminium on MCM-41 samples with decreasing Si/A1 ratios (increasing aluminium content) has been investigated by several authors. Davis and coworkers were able to prepare MCM-41 material containing only tetrahedrally coordinated aluminium with a Si/A1 ratio higher or equal to 29 (45), whereas Corma et al. decreased this ratio to 14 (47), even lower to 8.5 by Schmidt et al. (51) and finally as low as 2 by Clearfield et al. (52). The stability of aluminium in the framework of MCM-41 materials was monitored using 27A1 MAS NMR spectroscopy by Klinowski et al., concluding that the structural A1 in MCM-41 is thermally less stable than, for example, in zeolite Y, because the MCM-41 structure lacks strict crystallographic order at the atomic level and the very small protons cannot satisfy the framework charge balance as efficiently as the sodium cations (53). Corma et al. describe the dealumination of an as-synthesized MCM-41 (with aluminium only in the lattice tetrahedral positions) by direct calcination, leading to decreasing BrCnsted acidity but increasing Lewis acidity due to the formation of extra-framework aluminium as followed by 27A1 MAS NMR (54). A range of MCM-41 materials has been synthesized using different sources of aluminium and characterized in detail by (among others) MAS NMR spectroscopy. 27A1 MAS NMR clearly shows that, when Catapal alumina or sodium aluminate is used, virtually all A1 in the solid is six-coordinated. On the other hand, by using aluminium sulfate, MCM-41 can easily be prepared with all aluminium in four-coordination and a framework Si/Al-ratio as low as 10 (55). Syntheses using monomeric alumina precursers have been shown to yield stable MCM-41 samples, with controlled incorporation of aluminium into the framework being readily achieved as monitored by 27A1 MAS and DOR NMR (56). Japanese researchers characterized kanemite, a highly ordered mesoporous material (FSM-16) prepared from layered polysilicates (41-43): the 29Si MAS NMR spectrum showed a sharp signal at-97 ppm due to Q3 units, indicating a single laffered structure. The spectrum of the silicate-organic complex showed two si~nals due to a QOunit (-99 ppm) and a Q4 unit (-109 ppm). The appearance of the strong Q~ unit clearly indicates the formation of a three-dimensional SiO 2 network from the layered kanemite (41, 42, 57). Mesoporous materials derived from kanemite and MCM-41 samples were synthesized and characterized

169 by Davis and coworkers (58). Both preparations yield mesoporous materials with narrowpore-size distributions and somewhat similar physicochemical properties. However, due to a higher degree of condensation in the silicate walls of the material derived from kanemite (as followed by 29Si MAS NMR spectroscopy), these samples have higher thermal and hydrothermal stability than MCM-41 (58). The synthesis mechanism of the formation of MCM-41 material is still a matter of discussion, since large surfactant molecules are used as templates, and their interactions with the different species during synthesis are quite complex. Therefore, certain efforts have been concentrated on following the template interactions by 13C CP/MAS NMR spectroscopy. The signals from the template (cetyltrimethylammonium bromide/hydroxide=CTMA) located inside MCM-41 were sharper and had markedly different intensity distribution and chemical shifts than those from CFMABr, as reported by Corma et al. (47). Considering the linewidths, the authors suppose that aliphatic chains of CrMA cations are more ordered inside the MCM-41 than in CTMABr and this can be explained by the liquid-crystal arrangement of the template inside MCM-41 and by the high water content of the CTMABr powder (47). Beck and coworkers reinforced the hypothesis that M4IS materials are formed through a mechanism in which aggregates of cationic surfactant molecules in combination with anionic silicate species form a supramolecular structure, whereas microporous materials are formed by molecular organic species, as concluded from their 13C NMR spectra (among others) (59). Furthermore, deuterium NMR was applied for studies related to the micelle formation in connection with the preparation of nanocomposite materials, as outlined by Chmelka et al. (60). The authors concluded that the isotropic 2H NMR signal was consistent with the isotropically mobile spherical micelles in dynamic equilibrium with monomeric surfactant molecules in solution (60). 14N NMR spectra of M41S gels containing CTMA surfactant molecules were recorded by Davis et al. (46) and Anderson et al. (61) in order to collect further information about the liquid-crystal templating mechanism. Their results are consistent with the presence of organic micelles interacting with the silicate species yielding tubular silica encapsulations around the external surface of the micelles. Boron containing MCM-41 molecular sieves were prepared, and l lB NMR spectra confirm the incorporation of boron into the lattice and show the coexistence of several B sites in calcined samples. The stability of such systems is relatively low, and part of the boron can be removed from the lattice by hydrolysis with water vapor at room temperature (62). Substitution of vanadium into the framework of MCM-41 molecular sieves has been achieved, and 51V NMR spectra clearly indicate the presence of two tetrahedral vanadium species with different local environments with isotropic chemical shifts of about -500 ppm (63) (see Figure 21).The absence of an NMR signal with a chemical shift of about -300 ppm indicates that the prepared V-MCM-41 material was free of V20 5 (63). 129Xe NMR spectra were recorded by Ryoo et al. following the formation of A1-MCM-41 applying pH adjustment during synthesis (64). The 129Xe NMR spectrum of A1-MCM-41 prepared with pH adjustment contains a narrow single Lorentzian line appearing at about 71

170 ppm, whereas the line width obtained from A1-MCM-41 samples synthesized without pH adjustment increased due to heterogeneous NMR line broadening (64).

I

-300

I

-.400

I

-500

!

-600

I

-700

~5[ppm]

Fig. 21. 51V NMR spectra of calcined V-MCM-41. a) static NMR and b) MAS NMR (Reproduced by permission of The Royal Society of Chemistry, Cambridge). 3.3

Octahedral Molecular Sieves

Most microporous materials known until recently were zeolites or A1PO4/SAPOs, all of which contain tetrahedraUy coordinated metal atoms. In 1989, a family of microporous titanosilicates (generally denoted ETS) was discovered in which the metal atoms (Ti 4+) are octahedraUy coordinated (65-68). The structure of one prominent member of this family, ETS-10, has been elucidated by Anderson et al. using HREM, electron and X-ray diffraction as well as solid-state NMR spectroscopy (69). This structure comprises corner-sharing SiO 4 tetrahedra and TiO 6 octahedra linked through bridging oxygen atoms. The pore system contains 12-membered rings and displays a considerable degree of disorder. Many ordered variants of ETS-10 exist, some of which are chiral (69). The ETS-10 display characteristics indicating a wide-pore material. The 29Si MAS NMR spectrum of ETS-10 is shown in Figure 22. It shows four lines (one with a shoulder) with chemical shifts of-94.1, -95.8, -96.5 and -103.3 ppm. The intensity ratio of these resonances is 2 9 1 9 1 9 1. By comparing the chemical shifts with other titanosilicate minerals such as zorite and lorenzenite, the three low-field signals can all be assigned to Si (3 Si, 1 Ti) (that is, silicon connected through oxygen bridges to three silicon atoms and one titanium atom) and the resonance at-103.3 ppm can be assigned to Si (4 Si, 0 Ti). These assignments are consistent with the framework connectivity of ETS-10 shown in Figure 23. For the perfect C2/c structure there are eight crystallographic types of Si (3 Si, 1 Ti). From an NMR point of view (short range ordering!) these can be reduced to four distinct

171 types in a 1 91 91 91 ratio (denoted as A, B, C and D in Figure 23). The 29Si MAS NMR spectrum is able to resolve three crystallogaphic types of Si (3 Si, 1 Ti) with two sites overlapping. In conclusion, the structure of ETS-10 is built from sheets and consists of two polymorphs. Polymorph A results in a 12-ring pore system having a zig-zag arrangement with either P41 or P43 symmetry. Polymorph B belongs to the space group C2/c. Consequently, polymorph A has a screw axis and will, therefore, display chiral symmetry

(69).

Si(3Si, 1Ti)

1

Si(4Si, OTi)

l 1

S I 95

I 1 O0 Chemical

shift

,,

I 105

(p.p.m.)

Fig. 22. 29Si MAS NMR spectrum of ETS-10 (Reproduced by permission of Macmillan Magazines Ltd., London). I :




sample

sample

1

sample

2

4~,,,~. 1

45 A

45 A

_>, r

800 c

A

sample 2

27

sample

18A

.a

E

600

400

3

O3 rr

A

18A

sample 3

Z

200

27

0 Q.

0

0

0.2

0.4

0.6

0.8

1

0

P/Po

-20

-40

-60

-80

-1 O0

Temperature [ ~

Fig. 24. Nitrogen adsorption isotherms at 77 K (left) and 1H NMR signal intensities of pore water confined in different siliceous MCM-41 samples (see text) versus temperature (fight). Filled symbols denote adsorption (left) and cooling (right), open symbols denote desorption (left) and heating (right).

v

>pi

O3

Z

LLI

I-Z i

_.I
0) is expected whatever the chemical cause for the decrease in electronic charge of atom A. Thus AF~s may reflect not only a change in the oxidation state of atom A but also any change that affects its charge like a difference in chemical environment, coordination, nature of ligands, and crystallographic position in a lattice. The relationship between E x and the electronic charge qa of atom A is often correctly represented by a simple electrostatic model developed by SIEGBAHN et aL:

Eu~ = E ~ , kq^ * I3 qj j.^ r~

(10)

where E ~ is an energy reference and qj the so-called point charges on neighbouring atoms j. rA- are of course the distances from atom A to j. ~lf atom A is represented as a hollow sphere charged with the valence charge qA then the potential has the same value at each point within this sphere namely qA/rA where r A is the average radius of the valence shell. Thus any change in qA changes this potential by AqA/rA and the model predicts that all core levels will be changed by this amount. With a change in chemical environment for atom A it follows from equation (10) that AE, = k AqA + A V

(11)

V is the summation in the right hand term of equation (10) and is designated as the MAGDELUNG potential. In molecular solids the sum is limited to atoms bonded to A but in ionic solids the summation must be extended to infinity.

2.1.1.2

Spin-orbit splitting

Whenever an orbital does not have spherical symmetry (quantum number 1 ~ 0) any subshell splits into two levels with quantum numbers j = 1 _ Is i- This splitting is therefore increasing with the atomic number on a given subshell (constant n, l) and with a decrease in 1 at constant n. Assuming the two levels have the same photoionization cross-section, the ratio of peak areas in the doublet is given by the ratio of their degeneracies (2j + 1). The values are given in Table 1. Table 1 Parameters for spin-orbit doublets Subshell

1

j

area ratio

s

0

/2

--.-

p d f

1 2 3

1/2, 3/2 3/2, 5/2 5/2, 7/2

1/2 2/3 3/4

198 2.1.2

Final state effects

2.1.2.1 Chemical shift It is indeed surprising that approximation (9) and equation (11) as a rule predict

essentially correctly the observed chemical shifts. Indeed these equations neglect completely the AEF term in equation (8) which following equation (7) should be written as ~E~ = AE j

where AE~ is the difference in relaxation energy for level j. Thus in many cases it seems correct to assume: Z~Ej - 0

but there are cases when this assumption is not justified. For example is has been found [5] that the difficulties in interpretation of the RU3d5/2 lines in Ru Y catalysts [6] were due to the fact that owing to final state effects the RU3d 5R level of R u t 2 was not significantly different from the one of metallic ruthenium in spite of a difference of 4 in oxidation number. The common XPS features discussed below may also be regarded as final state effects, namely the shake up satellites, energy loss bands and multiplet splitting. 2.1.2.2

Shake up satellites

One internal relaxation process is a two electron effect which involves an unpaired valence electron being excited to an upper free level of the valence shell. The corresponding energy is borrowed to the photoelectron so that its kinetic energy is lowered and a satellite peak appears on the high binding energy side of the main peak. Shake up satellites are noticeably high in transition metals compounds with unpaired 3d electrons and in rare earth compounds with unpaired electrons in the 4f shell. The intensity and splitting of these satellites is often of critical diagnostic value. For example Fe(II) in such compounds as FeO, FeMoO4 and Fe(COOH) 2 can be easily identified from its prominent shake up satellite of the Fe2p 3/2 peak [7]. 2.1.2.3 Energy loss peaks The inelastic collision which gives rise to the photoelectrons appearing as a step on the high binding energy side of each main peak (see Figure 4) is obviously one energy loss process which must be considered as an external relaxation event. Another such process is the so-called plasmon loss in which the photoelectron loses energy to a collective oscillation of conduction electrons. There are bulk and surface plasmons with characteristic respective frequencies t ~ and o~s. The plasmon loss peak may be used in several analytical applications. For example carbonaceous materials with high polyaromatic content show a plasmon loss peak of their Cls photoelectrons at 291.2 eV. This peak is clearly resolved from the Cls peak [8].

199

2.1.2.4 Multiplet splitting Whenever the system has unpaired electrons in the valence levels, after ejection of a photoelectron from a core level, the electron remaining on this core level may have its spin parallel to that of the unpaired valence electrons. In that case an exchange interaction can occur yielding a lower energy than for the case of antiparaUel spin.

2.1.3 Auger parameters It may be seen from equation (3) that the chemical shift which affects binding energies will also be reflected in the kinetic energy of Auger electrons. Wagner found that the difference between the kinetic energies of Auger [EK(jkl)] and XPS [EK(i)] emissions was also characteristic of the chemical environment in addition to being independent of reference level and charging effects. Thus he defined the classical Auger parameter as:

(= = EK(]kl ) _ EK(i )

(12)

which was subsequently modified as (='= r * hv = EK(jkl) .Ea(i )

(~3)

The Auger parameter it' is often used in diagrams like the one shown in Figure 5 designated as chemical state plots. In these plots experimental values of E~:(jkl) are plotted as a function of EB(i) so that states at constant values of (x' appear as lines with slope +1. The chemical plot in Figure 5 is for silicon in a series of zeolite. It shows that parameter (x" is only distributed over a short range 1711.3 - 1711.8 eV whereas the Si2a binding energy ranges from 101.3 to 103.5 eV. Another important application of Auger parameters stems from the fact that the extraatomic relaxation energy may be expressed as:

E~

= (1 - ~1) q=

2ro

(14)

where q is the charge, eo the dielectric constant and r o the effective screening distance of the electron. Thus if we assume that the intra-atomic relaxation energy is not affected by the chemical environment, equation (7) yields &EB(i) = - &e(i) + A E ~ ( i )

(15)

Then assuming that changes in initial state orbital energies between the chemical environments is the same for all core orbitals, it may be shown from equations (14) and (15) that:

200 Z~EK(jkl ) = Ae(j) + 3 A E ~ ( j )

(16)

,

where AE~ a (j) is the relaxation energy change associated with the formation of the photohole by XPS, whereas:

a='=

2&E~a(j)

.

(17)

1611

/

/V/o

A > LU v

/

/ ~9

1610

.. ~

1711

tu

1710

~"

UJ

._~ (D

~

1609

.d Y

1608

Z ,,"/ / /OW / 104

/

i//,,

102 103 2P Binding Energy (EV)

E "

~ 1709

a)
103 although on a more narrow range. Stoch et al. [30] o. published also a set of XPS binding energies of r 102 alumino silicates. There reported trends for Si2p and 101 Ols binding energies agree essentially with the ones 76 displayed on Figure 8 but they show an opposite trend A ..13-'" ...... [2 for the Al2p values which would decrease with an >. 7s f increased Si/A1. As discussed in reference [29] these data are not allowing the fight conclusion because they 74 ~ contain results for both sodium and proton forms of zeolites. In addition the increasing trend is indeed 73 , , , ..... , imposed by the two extreme points, the upper value 2.0 4.0 6.0 40.0 Si/AI Ratio of Zeolites corresponding to mullite (which is not a zeolite) and Figure 8. Binding energy shifts in Na zeolites [26]. t

10-~

e~

206 the lower one to a high silica H-ZSM-5 for which the low A12pbinding energy might correspond to extraframework aluminium. Thus if we exclude this set of data all other literature data on Na-zeolites indicates that all binding energies of lattice atoms and counter ions are shifted in the same direction when the Si/A1 ratio is varied. This is a peculiar behavior if the chemical shift is to be explained by a change in the charges [AqA in equation (11)] because not all charges can increase at the same time. Therefore it is usually expected that if some elements display a positive chemical shift then some other should yield negative binding energy shifts. Barr [27,31] tried to explain this anomaly by introducing the concept of a "chemical mixture" of group clusters SiO2 and N a ~ O 2 which would play the roles of cation-anion units instead of the individual elements. Moreover introducing one cluster say SiO2 into the chemical mixture with the other (NaA102) changes the covalency/ionicity in the M-O bonds. The concept of "chemical mixture" is however not very well defined and its relation to binding energy shifts is not clear. Since the variations in AEB presented in o X zeolites Figure 8 cannot be explained by changes in the /x ZSM-5 zeolites AqA term in equation (11) the question must be 9 Y zeolites formulated in terms of wether the observed ~, 5 3 3 trends are associated to differences in the Magdelung potential AV or in f'mal state effects -= 5 3 2 AER. Both Okamoto et al. [28] and Huang et al. [26] suggested changes in the Magdelung potential as the reason, and the latter authors < 531 based their arguments on the data reported in Figures 9 and 10. Figure 9 gives Al2p, Si2p and 104.0 Ols binding ener~es for alkali exchanged X, Y >, 10,3.5and ZSM-5 zeolites. It is seen that changing c 103.0the electronegativity of the cation affects very --~ 1 0 2 . 5 little the binding energies in ZSM-5 but more significantly the ones in Y and X zeolites. In ~102.0Figure 10 some of these data are reported as a to 1 0 1 . 5 function of the charge calculated using the electronegativity equivalence method and >= electronegativity values listed by Sanderson 533 [32]. It is interesting to note that both Ols and Si2p binding energies vary linearly with the charge thus following equation (11). The two _c 5 3 2 lines corresponding to X and Y zeolites are parallel having thus the same value of k but 0 531 different values of the Magdelung potential. H ti Na b Cs v

(ID e,-

"O t-

122 o.

I

I

i

I

I

I

I

i

I

I

!

,,

lID

W

r m

''1

A

eILl

O)

"O ern

Figure 9. Variations in binding energies of zeolite framework elements with the extraframework cations [26].

207

532.5 532.0

Na .,•Li "

O-531.5

Na o

531.0 530.5 103.0

Rb

Li ~

~

Cs Cs

K

, .., , 0.32 0.34 0.36 0.38 0.40 Negative Charge on Oxygen Atom

. o X Zeolites 9YZeolites

~ co

.. LIo

~ e ~ C s 102.0 101.5

Between X and Y the average AV may thus be estimated as 0.3 eV for Ols and 0.5 eV for Si2p. In reference [29], Griinert et al. reported calculated distributions of Magdelung potentials at O, Si, Al and Na sites for models of NaX, NaY, Na-mordenite and Na-ZSM-5 zeolites. These models were constructed using the "Catalysis" software of Biosym Technologies (San Diego) with lattice energy minimization. The potentials used and their parameters are described in references [33,34]. The models still have arbitrary aspects for example in the population of the cationic sites in NaX and NaY zeolites. Nevertheless the results are reported in Figure 11 which also shows the average Magdelung potential values of each of the X, Y, MOR and ZSM-5 structures.

Ke .,~eLi

"/Na Rb R

Rb/,~

K Cs 9' ' / ! '' ! I 0.02 0.07 0.12 0.17 0.22 Positive Charge on Silicon Atom

Figure 10. Relationship between charge and binding energy [26]. Table 2 Auger parameters of Na-zeolites (eV) [29] zeolite

og(Si)=EB(Si2p) +EK(Si KLL)*

og(O)=EB(Ols) +EK(O I ~ )

og(Na)=EB(Nals) +EK('NaK L L )

NaA NaX NaY Na-mordenite Na-ZSM-5

1711.4 1711.4 1711.6 1711.8 1711.7

1039.4 1038.9 1039.5 1039.1 1039.5

2061.0 2060.8 2061.1 2060.7 2061.1

1460.3 1460.2 1460.3 1460.7/1458.3 1460.9/1458.1

Average

1711.6 __.0.2

1039 _.+0.3

2060.9 _.+0.2

1460.6 _ 0.3/ 1458.2 _* O. 1

.

.

.

.

.

c~'(A1)=EB(AI2p) +EK(A1KLL)* .

,

,

The Si KLL and A1 KLL Auger lines were excited using the bremsstrahlung of the A1 Ka and Mg Ka sources respectively. These results show that the distributions extend over a 4-6 V range with widely different distribution curves. For all four atoms the average values increase in the order X > Y > MOR > ZSM-5 which is the order of the (Si/A1) ratios. It is interesting to note that the differences in average VM values between the NaX and NaY zeolites are of the order of 0.5 V for oxygens and .1 V for silicons. These values are in fair agreement with the AV values estimated from the results in Figure 10 reported above. Moreover Griinert et al. show a convincing similarity between the shapes of the average

208 Magdelung potential as function of A1/Si and the corresponding variations in Ols, Si2p, A12p and Nals binding energies. All this suggests that the trends observed in Figure 8 for the variations in binding energies with the Si/A1 ratio are essentially associated with changes in average Magdelung potential rather than in local charge. This explains the observations that Ols, Si2p, A12p and Nals binding energies all vary in the same direction with the Si/A1 ratio. MOR a)

X

Y

I

!

~o~ o ,o

(~

MOR .SM-5

X

y ~$M-5

T ii

30 20 10

g 0 24

26

28

30

32

-54

34

-52

c)

X~l

-50

-48

VM ,V

VM ,V

MOR ~ (r"'~ZSM'5

50

X

."II

~

Na

y

MOR / ZSM-5

( r~

40

~oI

/J,J-),,' .~ ',, II

/~i 9,', 7i~ 1t-57 II I~,

I

t t l~- /..\ i| i ,~

i ~k/'J

!

i\

,oq ,,,',,//ZI\V/,.".....k -38

-37

-36 VM,V

-35

-34 -11

- 10

-9

-8

-7

-6

VM ,V

Figure 11. Distribution of Magdelung potentials Vu at the atomic of zeolite models [29]. ( e ) NaX, (O) NaY, (-) Na-mordenite, (v) Na-ZSM-5. The possibility that these changes may be associated to final state effects has been ruled out by Grtinert et al. [29] on the basis of a set of measurements of the Auger parameters o( reported in Table 2.

209

Na-A Na-X Na-Y Na-MOR Na-ZSM-5 1395

1390

1385 KE, eV

1380

In Table 2 two values are reported for tx'(A1) of the high silica Na-mordenite and NaZSM-5. These double values arise from the double component A1 KLL Auger lines observed with these two samples as shown in Figure 12. As none of the ct" values varies with the Si/A1 ratio it may be concluded from equation (17) that no change in extra-atomic relaxation energy AE~a is responsible for the variations of binding energies shown in Figure 8. The low kinetic energy Auger AI KLL component line in the Na-mordenite and NaZSM-5 samples have not being assigned. We will come back to this result when we discuss acidic sites in high silica zeolite.

Figure 12. AI KLL Auger lines of Na-zeolites [29]. All this discussion leaves very little doubt that the changes in binding energy' observed in Figure 8 are essentially dominated by changes in the Magdelung potential in other words by changes in the spatial distribution of the charges on their neighbouring atoms. These effects overcome the initial state charges effects. They explain for example why the Si2p line does discriminate Si atoms with different numbers of Al-atoms as second neighbours. The well resolved 29Si MAS NMR signals are indeed reflecting differences in NMR chemical shifts which correspond to different electron densities. They explain also why the Ols lines do not differentiate the oxygens adjacent to the various cations in a mixte counter-cations zeolite lattice. These oxygens have different Lewis base Strengths, as will be shown below, and therefore significantly different electronic charges. The narrow Ols lines observed with a partially exchanged zeolite should therefore not be reflecting uniform oxygen charge within the lattice, but variously compensated Ols binding energies due to induced differences in local values of the Magdelung potential.

4.

ACIDITY IN ZEOLITES

XPS of chemisorbed nitrogen containing basic molecules such as pyridine and ammonia have been employed to monitor the strength of acid sites in H and cationic forms of zeolites. One interesting problem is to follow the changes in concentration of the Br/3nsted and Lewis acidic centers with the temperature of thermal treatment. Two mechanisms have been proposed for the thermal degradation. Dehydroxylation occurs through a process first described by Uytterhoven et al. [35]

210

H

H

I

I

Two Br0nsted acid sites give rise to two Lewis acidic sites, namely the trigonal A1 and Si atoms so that assuming that all A1 is intra-lattice in the initial sample: n~ = B + L

(27)

where n~u, B and L are the numbers of aluminium atoms, Br0nsted and Lewis acid sites per unit cell respectively. However according to Ktihl [36] this process may be followed by framework dealumination which involves the formation of extra-lattice Lewis acid A10§ species:

+

AIO +

When this happens only one Lewis center is formed with the disappearance of two Br0nsted OH's. Then n,~ = B + 2 L

(28)

Kazansky [37] has shown that for zeolites with Si/Al < 5 thermal degradation occurs via the mechanism of Kiihl whereas the process is limited to the Uytterhoven step for higher silica zeolites. 4.1 XPS study of faujasites The early work of Defosse and Canesson [38] had shown that the Nls line of pyridine chemisorbed on NH4Y zeolites calcined at 300 and 400~ displayed a strong peak at 402.4 and a shoulder peak at 400.4 eV. These two peaks were assigned to pyridine respectively chemisorbed on Br0nsted and Lewis acid sites. Figure 13 features the Nls lines of pyridine chemisorbed on a progressively dehydroxylated HY zeolite sample. The precursor was a commercial NH4Y sample with a Si/A1 ratio of 2.33 and 1.2% residual sodium. The dehydroxylated samples were prepared by slowly heating the precursor to temperatures ranging from 300 to 700~ in air for 10 hours in a shallow bed of less than 3 mm in depth [39]. The pyridine desorption temperature was 25~ At a calcination temperature of 300~ two peaks are observed at binding energies identical to the ones reported by Defosse and Canesson. The same peak assignment was therefore adopted. The main peak at 402.4 eV is ascribed to the pyridium ion formed upon adsorption of a pyridine molecule over a protonic OH. The peak at

211 400.4 eV corresponds to pyridine chemisorbed on Lewis acid centers which in this case may be either a residual sodium counter-ion or the Lewis sites X A1 /\

or

/ *Si /\

formed in the Uytterhoven dehy-

droxylation step. Upon heating to 400~ the relative intensity of the two peaks is strongly changed and the Brt~nsted acid sites diminish whereas the Lewis acid sites increase in proportion. From 500 to 700~ one witnesses the appearance of a new Nls peak on the low binding energy side (see also Table 3) which must correspond to the inset of dealumination and formation of AIO+ type of Lewis acid sites. The change in binding energy of this peak (component 1 in Table 3b) with temperature indicates that the nature of the extraframework aluminium oxihydroxide/hydroxide phase which bears these Lewis acid sites is changing with the calcination temperature. The binding energy of this new peak was found to be very close from the main peak observed upon adsorption of pyridine on 7-A1203. This result is in line with the views expressed by Lunsford et al. [40] who concluded that A1203 is formed upon dehydroxylation of HY at 700~ The sites associated with the new Nls peak (component 1 in Table 3b) were found to be weaker acids as pyridine desorbs from them at temperatures lower than from the Lewis acid sites with Nls peak at 402.4 eV.

400

_>, _c

700 392.4

397.4

402.4

407.4

421.4

Eb/eV

Figure 13. Deconvoluted Nls XPS spectra of pyridine chemisorbed on HY zeolites. Sample calcination temperature, in ~ is indicated on the figure [39].

212 Table 3a XPS binding energies and fwmh (eV) for the HY samples of Figure 13 Calcination Temperature ~

(Si/A1) s

(N/AI) s

300 400 500 600 700

3.60 3.18 4.40 4.10 2.48

0.38 0.30 0.30 0.24 0.25

Si2~ 103.3 103.3 103.3 103.3 103.3

(2.4) (2.4) (2.4) (2.3) (2.4)

A]2p 74.8 75.0 74.8 74.8 74.7

(2.5) (2.6) (2.4) (2.4) (2.4)

Table 3 b XPS binding energies and fwmh (eV) for the HY samples of Figure 13 (continued) Calcination Temperature oC 300 400 500 600 700

N~* Oxs 532.3 532.5 532.5 532.3 532.2

(2.6) (2.7) (2.5) (2.5) (2.5)

1

2

3

----399.6 399.3 398.6

400.4 400.2 400.9 400.9 400.4

402.4 402.2 402.5 402.4 402.2

* In this early work the binding energy scale was referenced to Si2p = 103.3 eV ** fwmh = 2.4 eV for all three components It is seen from the data in Table 3 that the (Si/A1) s ratio is higher than the bulk value (2.33). This suggests an aluminium depleted surface region. The low values observed for the ratio (N/A1) s reflect the fact that only part of the Br0nsted acid sites are accessible to pyridine. The pyridine molecule kinetic diameter of 5.9 A does not allow it to enter the sodalite cages with 2.2 A openings. Thus only the acid sites protruding in the supercages can chemisorb pyridine. The number of these molecules is estimated to be 24 per unit cell [41] and since the Y zeolite with a Si/A1 ratio of 2.33 has 57 A1 atoms per unit cell, the maximum N/AI ratio is 0.42. This value is reasonably close to the 0.38 value obtained after calcination at 300~ In a recent paper, Guimon et al. [42] advocate the use of NH 3 as a probe molecule. Interestingly working with commercial NH4Y samples they also find two peaks at 402.7 and 401.2 eV (referencing to Cls = 284.6) when the samples are calcined at 400~ and a third one close to 399 eV when the calcination temperature is 700~ The N/A1 ratio never reaches the value of 1, even at a desorption temperature of 100~ in spite of the fact that ammonia is a stronger base than pyridine.

213 Comparing the high binding energies peaks obtained when adsorbing ammonia and pyridine, it could be concluded that the Br/Snsted acid sites located close to the six oxygen ring of the sodalite cage are as strong acids as the ones in the supercages. 4.2

r

E" i.. v

(D r

8

396.0 400.5 405.0 Binding energy (eV~

High silica zeolites Figure 14 gives the Nls lines obtained with five different samples of H-ZSM-5 zeolites [43]. For their preparation, the sodium precursors were ftrst calcined at 500~ in air for 10 hours. Then the calcined samples were converted to ammonium form by repeated ion exchange with 1 M ammonium nitrate solutions. The protonic form was obtained by air calcination at 500~ for 10 hours and the residual sodium content measured by atomic absorption was less than 0.02%. Sample B was produced by acid leaching sample A with a 0.1 N solution of HC1 at room temperature. Table 4 gives some of the bulk and surface properties of these samples. Sample A shows a 25% difference in (Si/A1) ratio between the bulk and the surface indicating that A1 is not uniformly distributed throughout the zeolite grains. Moreover sample A has a relatively low BET surface area and a low pyridine uptake (at 200~ indicated by its rather low value of (N/A1)B. This suggests the presence of extraframework amorphous and non-acidic AI oxidic phase. Upon acid leaching of sample A the obtained sample B shows increased values of (Si/A1) which indicate extraction of some of the aluminium specially in the surface region. The BET surface area however has not reached the values of samples C, D and E and the pyridine uptake is not significantly increased even though the surface (N/A1) ratio is substantially increased. This suggests only partial extraction of the A1 containing phase upon acid leaching and a preferential extraction in the external surface region of the zeolite crystals.

Figure 14. Deconvoluted Nls XPS spectra of pyridine chemisorbed on H-ZSM-5 zeolites. Sample designation as in Table 4 [43].

214 Table 4 Bulk and surface properties of H-ZSM-5 samples Sample

(Si/A1)B

(Si/A1)s

(N/A1)B

(N/A1)s

So* m2/g

A B C D E

19.9 24.4 39.0 45.9 65.0

14.3 21.4 37.2 47.4 36.6

0.52 0.53 0.85 0.91 0.92

0.37 0.62 0.92 1.07 0.83

362 392 432 422 445

* So: Nitrogen BET surface area Samples C and D are different from sample A. Their (Si/A1)B and (Si/A1)s values are equal within 5% suggesting a uniform spatial distribution of A1 atoms. Their BET surface areas are very close to the known values for ZSM-5 suggesting absence of pore blockage by extraneous material which is confirmed by their high pyridine uptake. The (N/M) ratios both in the bulk and on the surface are close to 1 indicating that every A1 atom is associated with one acidic center (able to adsorb one pyridine molecule). Thus these two samples are specially pure and uniform, and they may be used as model solids for further study. Sample E seems also quite free of extra-lattice phase as its BET surface area is high and its (N/AI) values are both reasonably close to 1. However it (Si/A1)s value is 30% below its bulk value indicating a non-uniform distribution of its A1 lattice atoms.

.O i"r0 0 r 11) u) e~

o

2000c 390 395 400 405 410 Binding energy (eV)

4.2.1 Add strength and Nls binding energies Sample C was used in the study reported in Figure 15 and Table 5. The pyridine desorption temperature was varied from 50 to 400~ These data show that the high binding energy peak corresponds to a site which adsorbs pyridine much more strongly than the other two. The almost constant value observed for (Si/A1)s indicates that the desorption process has not affected the lattice and the constant [N(3)/A1]s ratio indicates that at 400~ pyridine has not desorbed from the sites responsible for the high binding energy component of the Nls signal. The other two sites are obviously weaker acids. Thus once again it is found that the Nls binding energies are in the order of acid strengths. In this respect it is interesting to note that from the Nls binding energies of chemisorbed pyridine it was indeed found that the BrOnsted acid sites are stronger in A1-ZSM-22 and Fe-ZSM-22 [44,45] and weaker in Fe-ZSM-5 and BZSM-5 [46] than in AI-ZSM-5.

Figure 15. Deconvoluted Nxs XPS spectra of pyridine chemisorbed on sample C. Pyridine desorption temperatures, in ~ are indicated on the figure [43].

215 Table 5 Effect of pyridine desorption temperature on some properties of sample C Pyridine desorption T~

XPS

Relative intensity of Nls components %

[N(3)/A1] s

,

(Si/A1)s

(N/A1)s

1

2

3

36.7 33.3 37.2 38.4 39.6

1.10 1.19 0.93 0.56 0.55

19.5 21.0 24.4 9.5 7.7

46.5 48.0 47.6 26.0 29.5

34.0 31.0 28.0 64.5 62.8

50 140 200 300 400

.

.

.

.

.

.

.

.

0.374 0.369 0.260 0.361 0.345

.

4.2.2 Nls peak assignment The assignment of the three component peaks in Nls spectra of H-ZSM-5 samples in Figures 14 and 15 is different from the one discussed in section 4.1 for faujasites. This assignment was made after a careful IR study of pyridine chemisorbed on the same samples. Table 6 reports the relative intensities of the three Nls component peaks for the samples of Figure 14 and Table 4. Table 6 Relative intensities and B/I, ratio for samples A-E. (Calcination temperature 500~ desorption temperature 200~

pyridine

,

Sample

A B C D E

Nls relative intensity %

1

2

3

14.7 26.9 24.4 18.3 26.7

46.4 52.6 47.6 65.4 48.5

38.9 20.5 28.0 16.3 24.8

(B/L)rR

(B/L)sl

(B/L)s2

4.5 2.7 3.7 4.7 4.7

0.64 0.26 0.39 0.19 0.33

5.8 2.7 3.1 4.6 2.7

In Table 6 the Br6nsted to Lewis (B/L) ratio is first estimated from the ratio of absorbances of IR lines at 1545 and 1455 cm -1" (l~L)l R - As eL AL ~.

(29)

The ratio of extinction coefficients I~L/EB was given the value 1.5 as suggested by Rhee Then (B/L) ratios were calculated from the relative intensity ratios of the three component Nls peaks, under two different hypotheses. First (B/L)sl was calculated assuming that peak 2 corresponded to a Lewis acid site and then et al. [47] for high silica zeolites.

216 (B/L)s2 under the assumption that peak 2 was for a BrOnsted acid site. Obviously the first hypothesis is very far from being verified whereas the second hypothesis yields B/L estimates in good agreement with the IR data, specially for samples B, C and D. These three samples are the ones which were found to have surface properties more representative of the bulk in the discussion of Table 4. Thus from these data it may be safely concluded that the three peaks in Figures 14 and 15 are associated to one kind of Lewis acid site (binding energy 398 __. 0.3 eV) and two kinds of Br/3nsted acid sites, one being rather weak (BE = 400.0 _.+0.3 eV) and the other one relatively strong (BE = 401.8 -4-_0.3 eV). Several authors have discussed the existence of two kinds of Br/3nsted acid sites in high silica zeolites. For example Datka and Tuznik [48] measured the desorption of pyridine from H-ZSM-5 at various temperatures while monitoring the changes in intensity of the 3610 cm "1 and the 1545 and 1455 cm "1 bands. They found that a certain amount of pyridinium ion was decomposed after the disappearance of the 1455 cm "1 band and before the reappearance of the 3610 cm 1. They ascribed this amount to pyridine chemisorbed on weak BrSnsted acid sites. From the spectra in Figure 14 it is shown that the component 2 corresponding to weak Br~3nsted acid centers is dominant in all five spectra. It seems thus that XPS is specially sensitive to the presence of these sites compared to bulk techniques such as IR spectroscopy. This suggests that the weak Brt~nsted sites may be surface species. One possibility [49] is that these OH's would be adjacent to a surface silanol:

H

I

0

X /

OH

X/

A1

/X

Si

/

X

Their total concentration would therefore be essentially proportional to the external surface area of the crystals. Similar surface AI atoms may be responsible for the low kinetic energy Auger peaks observed in Figure 12 for high silica zeolites. In this case the sodium cation instead of the proton would be adjacent to the silanol group.

4.2.3 Dehydroxylation process It should be noted that whenever equation (27) applies and dehydroxylation happens solely through the Uytterhoven step, assuming the stoechiometry of pyridine chemisorption is one molecule adsorbed per either BrOnsted or Lewis acid site, then the N/A1 ratio must be equal to one. If however dealumination follows dehydroxylation and equation (28) applies, then the N/A1 is given by:

217 N/L = 1 * B/L 2+B/L

(30)

Table 7 reports some data obtained for the sample D of Table 4 and Figure 14, calcined at increasingly high temperature. Table 7 Effect of dehydroxylation on B/L and N/L ratios of sample D Calcination Temperature

(Si/A1)B

(Si/A1)s

(B/L)IR

(B/L) s

oC 500 675 800 950

(N/AL) s XPS

45.9 45.9 45.9 45.9

47.4 40.1 43.4 41.9

4.71 1.17 0.46 0.49

4.60 0.65 0.46 0.44 ,

1.07 1.04 1.18 0.87

eq. (30)* 0.85 0.62 0.59 0.59

,

* Calculated from equation (30) using the XPS (B/L) s values It is seen from the minor variation of the (Si/Al)s ratio that the migration of hl atoms is rather minor. Since, however, the fB/L)m ratio decreases dehydroxylation is significant. The rather good agreement between (B/L)s and (B/L) m confwms the peak assignment. It is seen that the XPS (N/A1)s average value is equal to 1.03 _+ 0.15 and that the predicted (N/A1)s values calculated by substituting the (B/L) s values in equation (30) are clearly farther from the XPS values. It is thus concluded that the dealumination process of Ktihl [36] leading to equation (28) did not happen in these samples in agreement with the conclusions of Kazansky [37].

5.

BASICITY IN ZEOLITES

In ion exchanged zeolites, the counter cation has electron pair acceptor (EPA) capacity and is therefore a Lewis acid. The lattice plays the role of the conjugated base and has electron pair donor (EPD) properties. Vinek et aL [50] have discussed the relationship between EPA-EPD strength and the catalytic activity of the cation-anion site as well as the mechanism of elimination reactions. They also proposed to measure the EPD strength (or Lewis base strength) of oxides including MgNaX and MgNaY zeolites by their Ols binding energy. As will be shown below this is not a very useful method essentially because it is not sensitive enough. We have recently studied systematically the Lewis basicity of zeolites using pyrrole as a probe molecule. Pyrrole has H-donor properties and adsorbs on basic zeolites forming NH---O bridges. The NH stretching vibration was shown by Barthomeuf [51-53] and others [54] to be a measure of the Lewis base strength of 02. in zeolites.

218 Lavalley has recently compared the suitability of various IR-probe molecules for the surface basicity of oxides [55] and found that pyrrole (PYH) which would decompose to pyrrolate ions (PY') over stronger bases is suitable for zeolites [56]. We have used pyrrole as an XPS probe molecule and monitored the Nls line of this molecule chemisorbed on alkali exchanged X and Y zeolites [26,57]. We also studied this adsorbed species on the same samples by IR [58] and microcalorimetry [59]. A similar study was also performed using chloroform as the probe molecule in IR [60] and XPS [61]. Figure 16 shows the Nls lines of pyrrole chemisorbed on NaX and partially exchanged LiX, KX, RbX and CsX [57]. All samples contain thus Na § counter-ions and it is interesting to note that the peak which dominates the NaX spectrum is present in all of these spectra. These results are also represented in Table 8. Table 8 Nls binding energies for pyrrole chemisorbed on alkali exchanged X zeolites [57] Sample

Cationic composition

Nls binding* energies, eV 1

Relative intensity %

2

3

1

2

3

LiX NaX KX

Li543Na31.1 399.7 Nass.4 399.8 K48_3Na37.1 399.8

400.3 --399.1

401.5 401.5 401.5

37.3 87.0 44.0

50.4 --43.7

12.3 13.0 12.3

RbX CsX-2 CsX-3

Rb37.TNa47.7 399.8 Cs2s.gNa56.6 399.7 Cs59.gNa23.5 399.8

398.7 398.3 398.3

401.6 401.5 401.6

46.3 51.3 45.0

37.1 41.0 51.6

16.6 7.7 3.4

They show that component No 2 decreases regularly in binding energy from Li-X to CsX. A third minor peak at 401.5 eV is believed to belong to secondary adsorption of pyrrole molecules. Our previous IR-study of the same samples had also allowed to detect a common band in the NH stretching vibration region at 3280 cm "l and a second band characteristic of the second cation and ranoino from 3295 cm 1 for LiX to 3175 cm 1 for CsX. A minor component at 3375 cm "1 ascribed to secondary adsorption was also present in all spectra. The results reported in Figure 17 are the measured distributions of differential heats of adsorption of pyrrole on the same solids. Again it is found that the only peak in the distribution for NaX (117 kJ/mol) is present in the curves of KX, RbX and CsX and shifted to 123 kJ/mol in LiX. The second component of these diagrams varies regularly with the cation from 110 kJ/mol on LiX to 146 kJ/mol for CsX. The correlation between the three kinds of data is shown in Figure 18 which shows that the cation characteristic Nls binding energy (component 2) is linearly correlated to the corresponding NH stretching vibration and heat of adsorption of pyrrole.

219

117 1

123

I_iX

NaX

126

0

KX

R L.

Figure 17. Distribution of differential heats of pyrrole adsorption on the samples of Figure 16 [59].

Figure 16. Deconvoluted N~s XPS spectra of pyrrole chemisorbed on a/LiNaX; b/NaX; c/KNaX; d/RbNaX: e/and f./CsNaX [57]. x

'-E 3350 o

o

~Xx~

~'~ o

x~ z

X

"i

o

145 I

to tO

O

E

c~r

,-

9

50 100 150 Differential Heat of Pyrrole Adsorption, Q (kJ/mol)

Binding Energy

A

i

RbX

125 ~

3250-

E O

Ne _q

= 3150 "r" 398 Z

399 4~) Nls Binding Energy (eV)

, 105 401

Figure 18. Relationship between Nls binding energies, NH stretching frequencies and Qmax of chemisorbed pyrrole [59].

220 It seems thus that whenever two counter-ions are present in the same sample the oxygens adjacent to the two cations have different basic strengths and therefore different electron charges. This means that the basic strength is not a collective property of the zeolite lattice but rather a local property. This is in contrast with the Oxs peaks of these zeolites which show almost no increase in the peak width (fwhm) between NaX (2.3 eV) at mixte CsNaX (2.6 eV). This small change does not allow to deconvolute two Oxs component peaks in partially exchanged CsX. Thus the use of the Ols binding energy is less appropriate for the characterization of basic strength. The reason it best seen from Figure 19 which shows the values of Ols, Sizp and A12p binding energies for ion exchanged X and Y zeolites as a function of the Nls binding energy of pyrrole chemisorbed on the oxygen adjacent to the main cation. It is seen from the curves that the Nls binding energy varies from 398.3 eV for CsX to 401.2 eV for LiY, a range of close to 3 eV. The variation ranges for Ols, Si2r, and Al2p for the same samples are respectively 1.4, 1.3 and 1.4 eV. The pyrrole Nls binding energy is therefore a much more sensitive probe of the changes in base strength than the Ols, Si2p and A12p binding energies of the lattice O, Si and A1 atoms. 532.50

O

K Rb,~O .... Cs .-'~

531.50 -

530.50 -

= CS

Ha

.... c, ......

K

Na

101.50

Li

w

F~b

Cs -~176 a~176176176176

10~50

o~176

.......

'U -o

Li

Cs LJ

Na

Rb

74.50

K

_~.....oo'"

o,O o

.~

~176176




~ Na

3 400 3300-

KLiNa Rb ~ ~ ~ . ~ C

"~ 3 200 o i

i

i

i

0.18 0.20 0.22 0.24 0.26 Negative Charge on Nitrogen Atom

Figure 20. Relationship between nitrogen electronic charge and Nls binding energy of chemisorbed pyrrole [26].

6.

.o

o X Zeolites 9Y Zeolites

.-s

399 398

-'E 3500

co 3100

"r" Z

O. 8

0.'~0

0.:22. 0.24

s

0.:26

Negative Charge on Nitrogen Atom

Figure 21. Relationship between nitrogen electronic charge and NH stretching frequency of chemisorbed pyrrole [26].

ACTIVE CENTERS IN ZEOLITIC OXIDATION CATALYSTS

The discovery in the early 80's of titanium silicalites [62-64] opened the new application perspective of zeolitic materials as oxidation catalysts. Several reactions of partial oxidation of organic reactants using dilute solutions of hydrogen peroxide could for the first time be performed selectively in very mild conditions. Other elements inserted in the lattice of silicalites have since been shown to have similarly interesting catalytic properties including, vanadium, zirconium, chromium and more recently tin and arsenic [65]. Titanium silicalites with both MFI (TS-1) and MEL (TS-2) structures have however been the object of more attention and they still seem to display unmatched properties. Indeed some of these reactions like the oxyfunctionalization of alkanes [66-69] by H20 2 are not activated by other Ti containing catalysts (with the exception of Ti-A1-Beta [70]). The same situation

222 applies to several other reactions such as aromatics and phenols hydroxylation [62,64b,7174], alcohols oxidation [67,75,76], ketones ammoximation [73,77,78], synthesis of hydrazine [79]. By contrast the epoxidation of olef'ms by H20 2 (and organic peroxides) which is catalyzed by TS-1 and TS-2 [80-83] is also catalyzed by other Ti containing solids including Ti-AI-Beta [84,85], Ti-MCM-41 [86] and mixed TiO2-SiO2 oxides such as the ones described in [87,88]. These exceptional catalytic properties of titanium silicalites have triggered much analytical work aiming at the characterization of the environment of the active titanium site in these lattices. From the very beginning of these efforts, it had been recognized that the Ti2p lines of TS-1 contain a very unusual component doublet with binding energies significantly higher than the one of Ti in TiO 2 [89]. Table 9 shows the Ti2p 3n binding energies of a series of calcined (500~ TS-2 catalysts before and after acid leaching at room temperature with IN HC1 solution followed by another calcination at 500~ [90]. Table 9 Ti2p 3/2 binding energies (eV) and percent peak area of the two component peaks for TS-2 catalysts before and after acid leaching. Acid leached

As prepared Sample

Ti2p 3/2a

Ti2p 3/2b

Ti2p 3/2a

TiO 2 1.6 TS-2 4.2 TS-2 6.4 TS-2 9.1 TS-2 TiSiG c

458.3 (100) --457.1 (71) 457.8 (97) 458.3 (100) ---

--459.8 (100) 459.8 (29) 460.5 (3) --459.8 (100)

--458.0 (19) 458.4 (76) 458.3 (81)

Ti2p 3/2b 460.2 459.9 460.2 460.1

(100) (81) (24) (19)

a/high coordination site; b/low coordination site; c/Ti-Si glass. binding energy scale referenced to Si2p = 103.4 eV. The sample designation indicates the bulk value of the Ti/Ti + Si ratio expressed in %. It is seen that the 1.6 TS-2 sample does not show any line at 458.3 eV, which is the binding energy value in TiO 2. Our peak appears instead at 459.8 eV which is precisely the value obtained for Ti in Titanium Silicon glasses where titanium is known to be in tetrahedral coordination. A similar value was reported by Mukhopadyay and Garofalfiai [91] for tetrahedral Ti in the same kind of glass. The high binding energy Ti2p 3r~ peak was therefore ascribed to tetrahedral titanium in TS-2. It is seen from the data in Table 9 that the percentage area of this low coordination titanium peak increases upon acid leaching. This corresponds to the extraction of extraframework octahedral Ti species. The small upward change in Ti2p 3/2 binding energy observed after leaching was later shown to be associated with the desorption of Na§ ions from the surface. It was also found that pyridine adsorption on TS-2 led to a 0.7- 0.8 eV decrease in the Tip 3/2 binding energies. The Nls line of

223 chemisorbed pyridine was 399.0 eV which according to the discussion in section 4.2.2, should be assigned to pyridine chemisorbed on Lewis acid simms. This result was confirmed by IR analysis and thermodesorption of pyridine which showed the presence of very weak Lewis acid simmsin titanium silicalites [92]. The assignment of the peak with Ti2p 3t2 binding energy close to 460 eV, to tetrahedral titanium was confirmed by the spectra reported in Figures 22A and B [93]. Figure 22A shows the Ti2p photolines of the four TS-2 samples in Table 9, and Figure 22B the XANES spectra for the same hydrated samples. Figure 22B also gives the XANES spectrum for a TiO2-SiO2 glass which is a standard for tetrahedral Ti and for rutile TiO 2 in which Ti is indeed in octahedral coordination. The XANES data confirm the presence of the two coordinations of Ti in TS-2 samples, with tetrahedral species dominating in the sample with lower Ti loading. The interpretation of XANES data is also discussed in [94]. '

1

'

I

~

i

i

'

5

f 8

'

I

'

2

450.5 459.8 469.1 Binding energy (eV) A

'

I ' I'. -1 2

i ' 11 14

Energy (eV) B

Figure 22A. Ti2p photolines for TS-2 samples with 1.6, 4.2, 6.4 and 9.1 Ti / Ti + Si %. Figure 22B. XANES pre-edge features at Ti K edge. Top: tetmhedral Ti in a TiO2-SiO2 glass, from top to bottom 1.6, 4.2, 6.4 and 9.1 TS-2; bottom: octahedral Ti in rutile [93]. In reference [92] it was also shown that the Ols lines of TS-2 samples with a main peak at 532.7 eV for the regular lattice oxygens show a minor component at 529.7 eV whenever the loading is increased and octahedral Ti species appear. Similar values (533.2 and 530.1 eV referenced to Cls = 285.0 eV) were reported by Grohman et al. [95]. These authors also confirmed our Ti2p peak assignment as well as the effect of acid leaching. The capacity to detect tetrahedral Ti species from their Ti2p 3/2 lines was used recently by Bonneviot et al. [96] to establish the substitutional limit for Ti insertion in silicalite-1. The results are shown in Figures 23A and B. Figure 23A shows a series of Ti2p photolines for anatase and TS-1 samples having Ti/Ti + Si of 1.5, 2.4 and 4.6% respectively. Here again the tetrahedral line with Ti2p 3/2 binding energy at 459.8 eV is clearly resolved from the one of octahedral titanium at 457.8 eV. In Figure 23B the fraction of tetrahedral Ti

224 calculated from the XPS surface area ratio is plotted against the bulk Ti/Si ratio. The curve is compared to the one of tetrahedral Ti ratio estimated for the same samples (as well as other TS-1 samples) from the linear fit of Ti-K-edge XANES spectra following the method described in reference [97]. It is seen from Figure 23B that both methods yield a substitutional limit not exceeding 1.5 - 1.7% for Ti/Si. ........

. . -

Anatas

, .

.

~

.

.

.

5 400

i= 1 380~o
~4A. For microporous crystals, measurements of uptake of molecules of differing sizes and shapes, and product distributions obtained in various test catalytic reactions reveal aspects of the micropore architecture. Direct ways of quantifying pore volumes and pore coordination environments are available and an automatic volume filling algorithm helps quantify the number of molecules of a defined type that can be accommodated within a given micropore volume. The pore volume and its degree of interconnectedness can then be visualized using contour meshes or solid voblme rendering [14-16]. As noted above, known pore architecture information can also be used as a structural constraint in the simulated annealing model development approach [36].

25. Non-frAmework cations The positions of non-framework cations in al,lminosilicate zeolites can control or fine-tune their sorptive and catalytic properties. Measurement, however, requires careful and usually protracted analyses of accurate single crystal or powder diffraction data. In cases for which extensive experimental data are available, statistical mechanics analyses can yield insight into relative site energies [53-55]; earlier analyses have also attempted to quantify the relative importance of short and long-range interactions in controlling site occupancy patterns [56]. Earlier atomistic simulations in this area [57-62] had mixed results. Recent developments in methods and interatomic potentials have allowed non-framework cation positions to be simulated based solely on a knowledge of the framework structure in zeolite systems for which validatory experimental data are available [113]. Illustrative results come from two systems chosen, firstly, because reasonable structural data are available and, secondly, because the S i - A1 distributions are known precisely; they both have Si:A] ratios of unity and hence strict S i - AI alternation. The procedure applied was originally

245

developed for probing the preferred binding sites of molecular sorbates [63] and takes as input a suitable framework model. For zeolite Li-A(BW), this is the unit cell and framework Si, A1 and O coordinates taken from a crystallographic refinement [64,65]. For zeolite A, we again use accurate crystal structure data [66], but reduce the published a = 24.61A supercell to an a = 12.305~k triclinic, P1, subcell by trimming the full supercell contents to the 0 < x < 0.5, 0 _KE-Y,~HF. Although the production of the 2-phenyl isomer is desirable in the production of LAB, it is much more important to avoid the formation of diphenyl-isomers and branched phenyl isomers. The production of these compounds can be suppressed by using HY as a catalyst [98,60]. The pore size of the zeolites can be modified by introducing cations in the pores or by the formation of coke inside the zeolite crystals, which decreases the diffus'lvity of the products, but also of the reactants, in the zeolite channels. Chen et al observed an increase in p-xylene selectivity with modified and coked H-ZSM-5 [85].

3.4

Effect of crystal size and external surface

As mentioned above, it is di~cu]t to separate the influence of the crystal size and of the external surface on the shape-selectivity of zeolites rigorously. Increasing the crystal size of the zeolite means at the same time reducing the influence of the external surface. For the industrial production it is desirable to operate with small crystal sizes.

342 Using small zeolite crystals (< 0.5~tm) in the methylation of toluene, Chen et al [85] observed at 500~ an equih'bfium mixture of the xylenes (i.e. 23% p-xylene). Increasing the crystal size to 31xm enhanced the para-selectivity to 46%. A further increase in the paraselectivity up to 97% could be obtained by modifying the catalyst with phosphoric acid ending with P-loading of 8.5%. The enhanced para-selectivity was explained by the increase in the diffusional pathway by pore plugging which would favour the outward diffusion of paraxylene. Also for the ethylation of toluene, modification with phosphorous or metal oxides of ZSM-5 was required to obtain high para-selectivities [75]. Kaeding et al [99] stated that the modification with P effectively blocks the external surface. Paparatto et aL [ 100] observed, in the ethylation of toluene and in the isomerisation of m-xylene at low contact times over H-ZSM-5, an excess pf p-ethyltoluene whereas amorphous silica-ahamina yielded an excess of the o-isomer. This indicates that the ethylation of toluene over H-ZSM-5 takes place inside the micro-pores. At higher contact times the product composition reaches the thermodynamic equih'brium distn'bution. With zeolite samples with large primary particles the o-isomer was always absent, whereas small crystals yielded the equilibrium distribution at high contact times. They explained their observations by the primary formation inside the pores of the p-isomer, which can isomerize on the external surface, The conm'bution of the external surface to the total active surface per gram of catalyst becomes larger if the size of the catalysts is reduced. The increase in crystal size also reduced the observed conversion. External acid sites can be eliminated by building an inert iso-structural silica shell around the zeolite by continuing the synthesis in an Al-free synthesis gel. This increases the crystal size and the effective diffusional pathway and is an effective method to enhance the para-selectivity in toluene alkylation but reduces the conversion over the catalyst [ 101]. The modification of zeolites by Chemical Vapour Deposition (CVD) does not only eliminate the external acid sites but also causes pore mouth narrowing. I-h'bino et al. [102] showed that the rate of adsorption of xylenes is decreased by CVD-treatment of H-ZSM-5 with tetramethoxysilane and they ascn"oed their enhanced para-selectivity in the methylation of toluene to pore mouth narrowing. Wang and Ay [103] showed that larger crystals needed less silica on their surface to obtain high para selectivity in toluene ethylation and therefore they regard the role of the active sites on the external surface to be very important. Matsuda et aL [104] studied the disproportionation of 2-methylnaphthalene over HZSM-5 and this zeolite post-treated with (NH4)2SiF 6 to eliminate the external acid sites. They observed that the bulkier isomers were formed over H-ZSM-5, whereas over the treated zeolite only the isomers 2,6 and 2,7-dimethylnaphthalene ws observed. This was explained by a disproportionation reaction in the pores and an isomerization reaction over the external surface. The same was observed in the isopropylation of biphenyl over Mordenite, where modification of the external surface with tn'butyl phosphonate increase the 4,4'diisopropylbiphenyl content in the fraction of dialkylbiphenyls and reduced the deactivation [105]. This was also ascn"oed to the activity of the external surface. Another method to eliminate the external acid sites is the selective poisoning technique in which a stronger base is added to the feed which is too large to enter the pores of the zeolite. It was observed that injection of small amounts of [3-naphthoquinoline during the toluene ethylation over H-ZSM-5 (crystal size 21am) increased the selectivity towards p-ethyltoluene but at the same time decreased the ethene conversion [106]. Regular injections of the base molecule are necessary because of the reversa'ble adsorption and the decomposition of the base under reaction conditions.

343 3.5

Effect of silica to alumina ratio and dealumination

The effect of silica to alumina ratios in the alkylation of aromatics is difficult to study separately because upon changing the Si/A1 ratio in the synthesis gel both the ratio in the zeolite and the crystal size are changed [97], as well as the crystallinity and morphology. If the Si/Al ratio is changed in a post-treatment step by dealumination, which can be done either by acid washing or by steaming, extra framework aluminium species are formed. These species can block pores and thereby modify the diffush~es of the reactants and products in the pores and can form a complex with remaining framework aluminium which may result in a modified acidity of the catalyst. Vinek and Lercher [107] synthesized ZSM-5 with Si/A1 ratios between 20 and 240, but because the pyridine TPD yielded a lower Si/Al ratio the existence of extra-framework aluminium species which were ascrl'bed to be weak acid sites. They obtained a linear correlation between the specific rate of toluene disproportionation and xylene isomerization and the number of strong Bronsted sites, which indicates that the reaction rate is primarily a function of the concentration of acid sites. This was also concluded by Nayak and Riekert [89] and observed for the ethylbenzene disproporfionation [108,109,110]. For toluene dispropordonation a linear relation~ip between the rate constant, assuming the rate to be first order with respect to toluene, and the Si/A1 ratio was obtained [ 111]. This indicates, that the turn-over-number (TON) remains constant and independent of the silica to alumina ratio. Sastre et al. [112] studied the isomerization of m-xylene over Ot~etite and observed monotonical increase in m-xylene conversion upon exchange of the K+-cations. This was ascn'bed to the increase of the concentration of the protons and the increase in accessa'bility of the pores, which resulted in a higher selectivity for the isomerisation reaction at the expense of the disproportionation reaction. Only a slight increase in the p-xylene in the fraction of orthoand para-xylene was observed. Over Beta a maximum activity for the xylene isomersation was observed and this was explained by either a possible existence of a synergistic effect between extra-framework aluminium and the framework Bronsted acid sites or a concentration effect [113]. The alkylation of toluene with methanol is also catalyzed by both strong and weak acid sites [107]. The ideal alkylation catalyst should have a high concentration of weak acid sites and a low concentration of strong Bronsted sites in order to minimize the side-reactions, viz. disproportionation. On the other hand it was observed that for the alkylation of benzene with linear alkenes over zeolite Y the rate increased linearly with the number of t~amework aluminium atoms which means that the turnover number remains constant [98]. However, the turnover number increased with increasing degree of ion-exchange which causes an increases in acid strength. The selectivity towards the desired 2-phenyl-alkane increases with increasing degree of ion-exchange showing that alkylation is a demanding reaction. For the alkylation of polyaromatics and biphenyls Mordenite with high silica to alumina ratios seems to be the preferred catalyst. Lee at al. [64] observed that dealumination of Mordenite by acid washing with 6 N HNO 3 modified the pore structure of Mordenite resulting in an increase in the total pore volume and especially an increase in the volume of pores with a diameter between 20 and 1000 A. In the isopropylation ofbiphenyl an increase in the yield of diisopropylbiphenyl was obtained which might be ascribed to either the enhanced diffusion of the reactants and products via the newly created meso-pores or the decrease in the rate of deactivation during the alkylation reaction.

344 The effect of dealumination of Mordenite by acid washing, leaching with EDTA and steaming has been studied systematically [114]. The selectivity to 2,6-diisopropylnaphthalene in the alkylation ofnapthalene was enhanced by the removal of external sites by leaching with EDTA. On the other hand after deep bed calcination the catalyst with a high external acidity showed a high conversion and a high selectivity. Stezming followed by mild acid washing to remove the extra-framework a~minhnn showed the lowest external activity and the highest selectivity for the formation of 2,6-diisopropylnaphthalene. Coke formation during the alkylation of biphenyl over Mordenite is reduced by using Mordenite with a high silica to aluminium ratio [ 110, 61], but also the nature of the coke is different. Mordenite with a high Si/A1 ratio produces a volatile coke (Td==,vtion= 200 - 340~ which are mainly biphenyl derivates whereas mordenite with a low Si/AI ratio yields hard coke which is burnt off at ca. 500~ The content of 4,4'-diisopropylbiphenyl in the fraction of encapsulated diisopropyldiphenyl isomers in the highly siliceous mordenite is over 80% which indicates the effectiveness of the pore system of Mordenite to produce selectively the desired isomer 4,4'- diisopropylbiphenyl.

4.

AROMATIZATION OF ALKANES/ALKENES

4.1

Introduction

Besides being a key high octane component of gasoline light aromatics are important raw materials for the production of a wide variety of petrochemicals. Benzene ranks third in volume and together with ethylene and propylene accounts for about 75% of the world's petrochemical production. At present catalytic reforming of hydrofined naphtha is the main source of BTX (benzene, toluene and xylene). The standard Pt/Re/A1203/CI catalyst is not very effective for converting C 6 alkanes to benzene, the yield being typically only about 10% as against 60% for methycyclopentane (MCP) and 90% for cyclohexane [58]. When the phasing out of octane-boosting lead from gasoline was started there was considerable interest in the production of additional BTX. To this end several zeolite based processes were developed, e.g., BP/UOP's Cyclar process using refinery C3/C 4 gases as feed, Chevron's Aromax process using C6 to C 8 alkanes as feed. The benzene could also be sold into the growing petrochemical market. However, since the allowable benzene content of gasoline is being lowered to below 1% the interest in these processes waned. To lower the benzene content in reformate it can be alkylated to toluene and xylenes, or it can be extracted and sold into the petrochemical market. The latter option could further depress the need in the short term for new sources of benzene. Only in instances where there is a shortage of aromatics but an ample supply of C 3 to C 8 alkanes and alkenes (as in a Fischer Tropsch complex) may processes such as Cyclar or Aromax be of interest. Nevertheless there is continuing research interest in aromatization using zeolite based catalysts such as Cra/HZSM-5 and Pt/KL.

4.2

Acidic Catalysts

The catalyst of choice remains acidic Ga-HZSM-5. The BP/UOP Cyclar process [115] used this catalyst in the 1000 bpd plant at Grangemouth, Scotland, which operated for about two years and was shut down in December 1991. With butane as feed a typical product spectrum was 65% BTX, 5% hydrogen and 30% fuel gas. UOP's continuous catalyst

345 regeneration process was used. IFffs Aroforming process also uses Ga-HZSM-5 in isothermal tubular reactors which operates on dual cycles [132]. Mitsubishi's Z-Forming process was tested in a 200 bpd unit which was commissioned in September 1991. The success of the HZSM-5 catalyst is no doubt linked to the low coke forming tendency of this particular zeolite. Other acidic zeolites such as HY are initially active but deactivate very rapidly due to coke deposition. Bradley and Kydd [116] investigated the performance of several pillar interlayered clay minerals and although the Ga pillared montmoriUonite was found to be the most effective it had a much lower activity than Cra/HZSM- 5. As is well know gallium addition greatly improves the performance of HZSM-5, eg, HZSM-5 at 550~ has a BTX selectivity of only about 12% while the addition of 5% Ga increases the BTX selectivity up to 70% [117]. The conversion of alkanes or alcohols to aromatics over HZSM-5 involves firstly the formation of alkenes which are then subsequently converted to aromatics, strong acid sites being involved in both steps [118,119]. Alkenes react much faster than alkanes [ 120] and this is in keeping with the deduction that the initial alkane dehydrogenation is a slow step in the overall process. The addition of Ga provides additional routes for dehydrogenation of alkanes, alkenes and naphthenes thus increasing both the overall reaction rate and also the selectivity to aromatics. Dehydrogenation via acid sites involves hydrogen transfer with the formation of low molecular mass alkane such as methane and ethane [120,121] which being inactive represent a loss of feedstock carbon. Dehydrogenation via Ga, however, produces hydrogen gas (which is a valuable by-product in refineries) and so results in a better feedstock carbon utilization. Addition of zinc to HZSM-5 has also been found to be very effective but Ga is preferred because of its higher stability [122]. ZnO is slowly lost through volafflisation at the high operating temperatures. More recently other metals active in dehydrogenation have been investigated as co-catalysts with HZSM-5. Ibm et al [123] claim that when feeding n-pentane to a Ni HZSM-5 catalyst the aromatic yield was equivalent to that obtained with Ga or Zn. They report an aromatic selectivity of 64% with Ni as against 71% for Zn and 66% for Ga. It should be noted, however, that when feeding propane the aromatic selectivity was only 25% which is a poor result. Ono et al [124] found that their Ag-HZSM-5 catalyst produced less methane and ethane than G-a- or Zn-HZSM-5 and concluded that Ag enhances C-H bond cleavage whereas Ga or Zn enhances both C-H and C-C cleavage. With butane and isobutane at 500~ the Ag catalyst gave a higher aromatic selectivity, namely 50 to 60% as against 30% for Ga. It should be noted, however, that here again the reported Ga results appear to be poor. With butene and methanol as feeds aromatic selectivities of 85 and 73% were obtained respectively with the Ag-ZSM5 catalyst. Shpiro et al [125] investigated the effect of adding both Pt and Ga to HZSM-5. They report that Pt promoted Ga reduction and its migration, resttlting in a more stable catalyst with a higher aromatic selectivity.

As gallium plays a key role its effective distn'aution in the zeolite is important. Although it is generally assumed that Ga 3+ is the active form, migration occurs more readily in the reduced Ga "~ state. The addition of Pt promotes Ga reduction by hydrogen spill over [125]. Hamid et al [126] found that the Ga, prepared by ion exchange, was, as one would expect, concentrated on the outer skin of the zeolite particles but with reduction/oxidation cycles the Ga migrated into the interior. They speculated that Ga + migrated as Ga20 vapour. In the regeneration cycle the Ga + is oxidised to the more active Ga 3+ state resulting in an improved performance. Further studies showed that after several reduction/oxidation cycles the performance reached a plateau [ 127]. It was deduced from pyridine infra red studies that the reduction/oxidation cycles resulted in a decreased H + concentration, due to exchange by Ga ions, and an increase in the Lewis acidity due to better Ga dispersion. In another study [117] it was found that H 2 pre-reduction markedly increased the aromatic selectivity of

346 physically mixed Ga20 3 / HZSM-5 but it decreased the aromatic selectivity of samples prepared by incipient wetness impregnation or by ion exchange. It appears therefore that H 2 reductions only improves matters when the Ga is poorly distn'buted in the initial state of the catalyst. If Cra dism~aution is important it could be reasoned that HGa silicate (MFI) would be a good catalysts since the Ga here is atomically dispersed by being incorporated in the ~amework. It has in fact again been reported recently that this zeolite is more effective than Ga/HZSM-5 [128]. Choudhary et al [129] found that the aromatic selectivity of Hgallosilicate increased with the degree of I-I* exchange while it decreased with mcreasing calcination temperature or increasing steam content during calcination. The latter two effects would be due to sintering (ie lower dispersion) of the extra-~amework Ga. Lukvanov, Gnep and Guinet have modelled the kinetics of propene [119] and of propane [120] aromatization over both HZSM-5 and Ga-HZSM-5 obtaining good fits with the experimental results. Propane is converted to propene along two main routes, protolytic craclcing of C-H bonds and dehydrogenation at the Ga sites. Protolytic crack~g of C-C bonds produce methane and ethane. The acid site reactions result in a CH4/H 2 ratio of 2.6 while the Ga sites give a 0.26 ratio which is in line with the observation that Ga HZSM-5 produces more H 2 than H-ZSM5. The propene then oligomerizes to higher alkenes (acid reaction). The oligomers are converted to dienes with both acid and Ga sites contn"outing and the dienes are converted to cyclic alkenes (acid sites). Cyclic alkenes are then converted to cyclic dialkenes and then to aromatics (H transfer at acid sites and de-hydrogenation at Ga sites). It was estimated that the Ga sites contn'bute about 90% to the diene formation and about 50% to the formation of aromatics. The formation of aromatics via H-transfer should result in the production of alkanes but the majority of the latter are again converted to alkenes. The only stable alkanes to emerge are the nonaromatizable methane and ethane. The product spectrum when feeding hexene or octene is very similar to that when feeding propene which is expected i~ as was found, the primary reaction is the craclcing of these higher alkenes to propene and butenes [117]. In general the percentage conversion of the feed and the aromatic selectivity follow the same trend. Likewise C2H4 selectivity follows the BTX selectivity [117,129]. When considering the breakdown of the aromatics it appears that at low temperature (350 ~ xylenes are the dominant aromatics, the benzene being low. As the temperature is raised to 550~ the benzene increases, toluene remains fairly constant and the xylenes decrease [117].

4.3

Platinum on Neutral Zeolites

Although platinum alone or on a variety of neutral supports selectively converts nhexane to benzene most of these catalysts deactivate rapidly due to coke formation. With the neutral zeolite KL as a support, however, much longer on-stream times are feast~ole and within a few years of Bernard's original publication [130] the Aromax process had been developed by Chevron [131]. Table 6 compares the aromatic selectivity obtained with Pt-Ba KL and Pt Re Sn / Al203 - C1 reforming catalysts [58]. Associated with the much higher aromatic selectivity is a lower amount of light gas production. Since the neutrality of the support was an important aspect the influence of doping Pt KL with the alkali series Li to Cs has been investigated by various workers. Hicks and coworkers [133] exchanged BaKL individually with Li to Cs and then added Pt by incipient wetness impregnation. They reported that the activity for aromatic formation increased markedly ~om Li to Cs but that the selectivity only increased slightly. Earlier studies [134]

347 had reported that both the conversion and selectivity increased markedly as Pt/KL was promoted with Li to Cs. Clearly the more basic the catalyst the better the performance. From this point of view it is interesting that promotion with various halogen compounds enhanced performance [135,136] despite the electronegativity of the halogens themselves. Tatsumi et al [ 137] investigated the effects of added KF, KC1, KBr and KI on the performance of Pt/KL. They found that KF and KC1 gave the highest benzene selectivities but that KBr and KI were actually inferior to the unpromoted Pt/KL. Table 6 Alkane Aromatization over Pt on Neutral and on Acidic Supports [58] % Aromatic Selectivity Alkane Feed Pt-BaKL Pt Re Sn/Ai203CI C6 87 25 07 82 45 Cs 80 60

The high selectivity of Pt/KL has been ascribed to the presence of very small Pt particles [138,139] and thus that sintering of these particles is one of the causes of deactivation [140]. It has been shown previously that Pt on zeolites KL, HL, HZSM-5 and silicalite were dispersed by treating with C12 in nitrogen or HC1 in air at 350~ [141]. With standard Pt / A120 3 reforming catalysts the practice of redistn'bution of the Pt (after air regeneration) by treatment with chlorine is well known. In the light of the foregoing it appears probable that the positive effect of halogen pro-treatment [135,136,137] is largely due to its resulting in finely dispersed Pt clusters. Iglesia and Baumgartner [ 142] have pointed out that selective terminal adsorption and dehydrocyclization of hexane to benzene are intrinsic properties of any clean Pt particles and that the role of KL zeolite is that the size of the channels inh~it the formation of coke in these channels thus keeping the Pt clusters there clean. This is in keeping with the opinion expressed previously by Tamm et al [131]. If it were a matter of pore size then one could, however, expect neutral silicalite also to be a satisl~ctory support. It is well known that operating metal catalysts in a hydrogen atmosphere inh~its coke fouling. Hicks et al [143] found that with 0.6% Pt/KBaL the catalyst deactivated due to coke fouling at hydrogen partial pressures below 6 atmospheres. The conversion of heptane increased with increasing hydrogen pressure up to 6 atmospheres. Pt/AI20 3 reforming catalysts commonly also contain Re which improves the catalysts' resistance to coke deposition [131] and this raises the question whether the effect of adding Ke to Pt/KL has been investigated. Pt/KL has been shown to be very sensitive to sulphur poisoning [131,144] and the effect has been ascribed to sulphur accelerating Pt sintering and subsequent blocldng of the zeolite channels rather than normal surface poisoning [144]. The sensitivity to sulphur obviously requires very thorough desulphurization of this feed and from this aspect a feedstock derived from the normal Co or Fe based Fischer Tropseh catalytic process could present an advantage. Fischer Tropsch products, however, contain other non-paraffmic substances such as alkenes, alcohols and carbonyls. The effect of these and other contaminants on the aromatization of hexane over Pt/KL is currently being investigated in the authors' research group. Zeolite supports other than KL have also been investigated. Pt/K Beta because of its higher acidity yielded more isomerized and cracked products than Pt/KL [145]. Ion exchanging with Cs reduced its acidity and improved the aromatic selectivity and Ba improved

348 the dispersion of the Pt which also increased its aromatic selectivity but despite these improvements the Pt Beta catalyst was still inferior to that of Pt/KL. It was, however, less sensitive to sulphur than Pt KL. Ruckenstein et al [146] studied the performance of composite catalysts, consisting of Pt/Ba-K1 with either Pt/beta or Pt/USY. Feeding mixtures of n-hexane, methylcyclopentane and methylcyclohexane the composite catalysts gave higher C7+ aromatics than expected from theindividual catalysts and feeds. An interesting observation was that for all the various individual catalysts (including Pt/Ba-KL) and for the composites n-hexane gave a lower benzene sdecdvity than did methylcyclopentane. This is contrary to the results of others [131]. In normal Pt/AI203-CI naphtha reforming the reaction network is complex because both metal and acidic sites to varying degrees catalyze ring closure, isomerization, dehydrogenation and cracking. Pt apparently is mainly responsible for hydrogenation / dehydrogenation with the acid sites accounting mainly for isomerization [147]. The conversion of n-hexane to benzene apparently goes via methylcyclopentane (MCP) and this is supported by the observation that at low conversions the major product when feeding nhexane is MCP [ 147]. With neutral Pt/KL the reaction network is simpler because the Pt sites mainly account for all the products. A commonly assumed reaction sequence is depicted in Figure 12. 1-6 Ring closure of chemisorbed n-hexane yields cyclohexane while 1-5 closure yields methylcyclopemane. Ring opening of the latter accounts for the two isoalkanes (2MP and 3MP).

Benzene

m

Cyclohexane Hexenes

~~

2MP

T IT MGP

n-Hexane

, P e n t a n e + OH 4

l

Butane

+ OH4

3MP

Figure 12. Reaction sequence for n-hexane conversion over Pt/KL.

It is of interest to compare the observed product concentrations with those predicted by thermodynamics. Table 7 lists several relevant equilibrium ratios. From the values ofthe cyclohexane / n-hexane and the benzene / cyclohexane ratios one would expect that the amoum of cyclohexane emerging from the reactor would be low which indeed it is. At about 40% n-hexane conversion at 450~ a typical exit molar ratio of benzene to cyclohexane is about 300: I, which although much lower than predicted is in line with the known fact that cyclohexane is very rapidly dehydrogenated over Pt at high temperatures. The exit benzene:MCP ratio is about 7 (at 40% n-hexane conversion at 450~ which is also lower than the predicted value. This nevertheless indicates that I-5 ring closure occurs at a reasonable rate compared to I-6 ring closure (the latter being followed by rapid benzene formation). At 350~ and at conversions below 3% the MPC concentration in fact exceeds that of benzene by a factor of 3, which again shows that I-5 ring closure is fairly rapid.

349 Table 7 Equilibrium Ratios of varius mixtures at 1 atm H2 Ratio 600K 700K 800K Cylohexaneln-Hexane 0.017 0.062 0.17 Benzene/Cylcohexane 34 1.8X10 4 2x106 MCPIn-Hexane 0.10 0.56 2.1 Benzene/MCP 5.9 2x103 1.6x105 3MPI2MP 0.51 0.54 0.56 MCP = Methylcyciopentane; MP = methylpentane

The observed 3MP/2MP ratio is about 0.8 which is not very different from the predicted value of about 0.5. The latter ratio does not change much with increasing hexane conversion when, as expected, the benzene level increases. This shows that the near equilibrium state between MCP, 2MP and 3MP persists at different conversion levels indicating that ring opening and closing occurs fairly rapidly. The main cracked products are methane and pentanes which is in line with the known hydrogenolysis activity of Pt. The main alkenes in the product mixtures are trans and cis 2-hexene in that order. While the reaction pathway depicted in figure z is supported by the fact that feeding n-hexane, MCP, 3MP or 2MP individually over Pt KL all result in high and similar aromatic selectivities [131], the actual mechanL~m on a molecular level is still a matter of dispute [139,142].

5.

SKEI~ETAL ISOMERIZATION OF 1-BUTENE

Isobutene is an important petrochemical starting material and best known for its use in the production of MTBE which is added to fuel as an octane-enhancer. It is also used as a monomer for the production of butyl rubber. Furthermore isobutene can be converted into isoprene which is an important monomer for elastomers by the modified Prins reaction with formaldehyde over, for example, H-ZSM-5 at 175 - 4000C [148]. Partial oxidation of isobutene yields methacrolein/methacrylic acid which upon esterification yields alkylacrylates, which are used e.g. for the production of polymers (plexiglass) and in water-soluble paint. Presently, the need for isobutene is covered by its production in the FCC-unit. However, with a strongly increasing demand for this raw material, especially for the m~nufacture of MTBE, alternative routes for the formation of isobutene need to be explored such as the acid catalyzed skeletal isomerization of linear butenes. Thermodynamically the skeletal isomerization of alkenes is favoured at low temperatures and the reciprocal temperature increases with increasing carbon number. The equih~rium concentration of isobutene in the fraction of butenes decreases from ca. 50 % at 200~ to 37% at 500~ [149]. Thus, the conversion of n-butenes into isobutene at these temperatures will be limited by thermodynamic constraints. The skeletal isomerization of the alkenes with more than 4 carbon atoms is a relatively facile reaction step, which is carried out at ca. 290~ over H-Ferrierite [150] or at 340~ over ZSM-5 [151]. This reaction proceeds via the skeletal rearrangement of a carbenium ion yielding a secondary carbenium ion. The singular reaction mechanism indicates that side product formation can be minimized. Even the skeletal isomerization of C 5- and C6-alkanes over Pt-Mordenite, which is thought to proceed

350 via a dehydrogenation step is a relatively facile process [ 151]. This is nowadays an important process for increasing octane numbers [ 153]. Contrary to the isomerization of longer chain alkenes, the formation of iso-butene from n-butene over acid catalysts is a difficult reaction, which proceeds e ~ e r via a mechanism involving oligomerisation, skeletal isomerisation of the oligomers and subsequent cracking or via an energetically unfavourable primary carbenium ion mechanism [ 154]. The first proposed mechanism imnlies the unavoidable formation of C5+-oligomers and C3.-cracked species as by-products m this process. Fluorinated alumina seems to be a promising catalyst for the skeletal isomerization of linear butenes [155,156], but the need to add fluorine to the feed stream together with the associated corrosion and environmental problems might prevent its industrial application [ 157]. A promising alternative to fluorinated alumina are zeolites. A number of zeolites have been studied for their activity and selectivity for n-butene isomerization [156-161]. The conversion of n-butenes over H-ZSM-5 can be observed at 377~ [157], but the iso-butene selectivity for this catalyst is rather low (14 %) and especially the selectivity for C 1-C3 products is quite high. At 500~ high conversions are obtained but the yield is then limited due to thermodynamic constraints [158,160]. The high activity and low selectivity of ZSM-5 has been ascribed to its strong acidity [158]. The acidity of zeolites can be reduced by the incorporation of boron in the zeolite framework [16_2,163] and therefore B-substituted ZSM-5, ZSM-11 and Beta were tested [158,164]. A13§ free boron zeolites are inactive, but these zeolites with low levels of A13+ ions which can be obtained by adding A120 3 binder to the A13+ free boron zeolite have weak acidity and are moderately active at 500 - 600~ and isobutene selectivities of up to 50 % have been reported. At these conditions the observed activity and selectivity of B/A1-ZSM-5, B/A1-ZSM-11 and B/A1-Beta were similar and therefore it was concluded that the pore structure did not play a decisive role in the conversion of n-butene into isobutene [164]. However, A1 which migrates into the pores not only modifies the acidity but also modifies the effective pore diameter. The importance of pore size has been frequently emphasized [150-161]. Figure 13 shows the performance of H-Ferrierite, H-Mordenite and SAPO-11 on the skeletal isomerization of n-butene. The widepore zeolite, mordenite, shows a low conversion and a low selectivity towards isobutene. The selectivity to isobutene obtained with the medium pore size SAPO-11 is higher but, due to the low conversion, a lower isobutene yield is obtained. With H-Ferrierite, a higher selectivity (> 80%) is obtained yielding a composition which is close to thermod~(namic equih'bfium The ~lectivity of zeolite Theta-1, which has narrower pores (5.5 x 4.4 A) than ZSM-5 (5.6 x 5.3 A and 5.5 x 5.1 A) for isobutene at 377 - 379~ was reported to be three times higher but also the conversion over Theta-1 at these conditions was significantly lower. Ferrierite (4.2 x 5.4 A and 3.5 x 4.8 A) showed at these temperatures over 90 % selectivity to isobutene [159]. A comparison of the activity and selectivity of SAPO-5 (ca. 8 A), SAPO-11 (6.7 x 4 A) and SAPO-34 (4.3 A) at 400~ showed that the selectivity to isobutene increased with decreasing pore size [159]. It was further noticed that the conversion and catalyst deactivation decreased with decreasing pore size. The large pore zeolite Mordenite itself is not selective for butene isomerization but Mg-Mordenite has been found to be more selective than H-Mordenite [150]. Zeolites with pores smaller than 4.2 A like Erionite are not useful because of the diffusional constraint of the product isobutene [150, 160]. Studies have shown that Ferderite is an attractive catalyst for n-butene isomerization, because it is both active and selective at relatively low temperatures of 350 - 400~ [150]. Initially small differences in butene conversion and isobutene selectivity for Ferriefite with

351 different Si/AI ratios (Si/A1 = 9 - 43) were observed but after some time on stream the differences were negligible indicating an influence of the acidity on the initial performance of the catalyst but not on its steady-state performance. The selectivity of this zeolite has been explained in terms of shape selectivity and substantiated with computer simulations [160] because the pore structure of the Ferrierite strongly inh~its the diffusion of the intermediate trimethylpentenes and therefore increases the probability of cracking which yields iso-butene. Most laboratory studies has been performed with a highly diluted feed at atmospheric pressure [158,157,159,163,164] because at a butene partial pressure of 1 bar much lower isobutene selectivities [161] and shorter catalyst life-times for medium and large pore zeolites were reported [150]. This is consistent with the postulated tmi-molecular isomerization reaction and multi-molecular oligomerization reaction [158] and the observed first order with respect to the partial pressure of n-butene for the formation of iso-butene and an order larger than one for the formation of by-products assuming low coverage for boron substituted zeolites [164]. For industrial operation, however, it is desirable to operate at higher partial pressure of n-butenes. Ferrierite can operate at higher butene partial pressures with a moderate catalyst life time [ 150,160]. With time on stream the activity of the zeolites for butene conversion decreases and the selectivity for isobutene increases [150,157,159,160,165]. Ferrierite initially produces cracking and oligome"nsation products [165] but with time on stream the rate of formation of these by-products decreases whereas the rate of formation of isobutene first increases before showing a slow decrease. The still active and selective catalyst can contain up to 8 - 10 wt.-% coke which is aromatic in character (H/C = 1) and which reduces the accessible pore volume dramatically [160, 165]. In the case of Theta-1 it was observed by using lower calcination temperatures (325~ instead of 500~ that the isobutene yield increased si~ificantly from 4.6 mol-% to 25.5 tool-% [157]. This has been ascribed to residual template inside the pores [161]. Figure 14 shows the effect of temperature and space velocity on n-butene isomerization over H-Ferrierite. It has been observed, that the isobutene selectivity increases with increasing space velocity and with increasing temperature [150,157]. The usual explanation for an observed increase in selectivity with increasing space velocity is the reduction of secondary conversion of the compound and if the space velocity is high enough its primary formation. Primary formation of isobutene is consistent with the observed different reaction orders for the isomerization and by-product formation [164]. The enhanced selectivity for the coked catalysts can then be explained with a lower diffusivit~ of the reactants into the zeolite and thus a lower butene concentration in the pores, which would favour the reaction with the lower reaction order, i.e. the skeletal isomerization. Assuming the primary formation of isobutene the observed increase in selectivity with temperature would indicate a higher activation energy for its formation in comparison to the oligomerisation. This is mechanistically consistent, because the oligomerisation will proceed via secondary carbenium ions whereas the direct skeletal isomerization of butenes involves an energetically unfavoured primary carbenium ion which might be stabilized by the zeolite structure.

352

40 ) 35 3O

~ F e r r i e r i t e

25 -o 20 SAPO-11

15

"---t

10

H-Mordenite 0

I

I

t

6

12

18

24

Tim 9on Stream [m in]

Figure 13. Yield ofiso-butene from the skeletal isomerization ofn-butene.

100 90

9 9

80

E 0

>=

8

X

70

~

100

400"C, 14 hr-1 425"C. 14 hr-1 400"C, l B h r - 1 425"c, 7 hr-1 Export. (400~, 14hr- 1) qlxxt. (425"C, 14hr-1)

90

7O

~~ _-; so

F ~ y. ( .;2-3~,,. 7"~ - i"/

60 50

o

40

~

4 o l~ /

1~400"C' 18 hr'l

30

m

30 ~-

IX425~:.7",-~

20

20

10

10

O/ 24

48

7"2

Time on Steam [hr]

96

120

0

F''~'~""r-'

)

I

i

t

24

48

72

96

Time on Steam [hr]

Figure 14. Effects of temperature and space velocity pm n-butene isomerization over Ferrierite.

120

353 6.

ALKENE OLIGOMERIZATION

The oligomerization of light alkenes into dimers, trimers, tetramers and higher oligomers represents an important reaction for the production of aromatic-free higher alkenes. Although the use of ZSM-5, with Si/Al ratios of approximately 30 - 40, as an oligomefization catalyst was patented in the early 1980s and has been extensively reviewed [166,167] currently only the Mossgas Refinery in South Africa is using this technology [168]. The process is able to produce, after hydrogenation, mainly low branched alkanes and scarcely any aromatics. The high selectivity to alkenes in the diesel mode of operation is attn~outed to restricted transition shape selectivity which both favours alkene formation and inh~its the formation of typical cyclic coke precursors. These factors together with a high reaction pressure (typically 5MPa) and moderate reaction temperature (200 - 220~ are conducive to the formation of diesel fractions. The typical feed composition is 81.7% alkenes, 15% alkanes, 1.5% aromatics and 1.8% oxygenates and the typical liquid fuel yields, based on alkenes, of 97% and, when operated in distillate mode, yields 78% distillate and 19% gasoline. The product diesel has a high cetane number (about 53). The gasoline has an RON of between 81 and 85 and a MON between 74 and 75. The oxygenates may cause premature catalyst deactivation possa'bly due to stronger and irreversa'ble adsorption on the acid sites [169]. Apart from ZSM-5, there have been reports of the oligomerization activity of other zeolites. As expected the extent of chain branching, usually undesirable in most applications, increases as the pore size increases due to the shape selective nature of the reaction in the zeolite. Oligomerization of higher alkenes represents an important route to the formation of synthetic lube oils [ 180]. After hydrogenation such oils have excellent properties due to their low volat'flity for their viscosity, high thermal and oxidation stability, very low pour point and exceptional low-temperature performance [170]. Such oils however are usually expensive due to the relatively high cost of the olefinic feed. Synthetic lubricant base stocks can be prepared in good yields by oligomefizing long-chain alkenes using catalysts containing large pore zeolites with high Si/Al ratios. Internal alkenes are less reactive than the corresponding alphaalkenes and conversions decrease as the chain-length of the feed alkene increases. In general, however, the zeolites thus far reported are not as good as clay catalysts or the curently used boron trifluoride or aluminittm chloride catalysts. 7.

ISOMERIZATION OF LONG-CHAIN ALKANES

The need for lubes and middle distillate fuels with greater performance, safety, and environmental advantages is increasing. This need has focused attention on highly paraffmic feedstocks due to their high oxidation stability, low volatility for a given viscosity, and high viscosity index (>130). Because highly paraffmic feeds tend to have high wax contents, however, the production of lubes and fuels from these feeds has been limited due to the large loss upon wax removal. An alternative approach is to change the molecular structure of the wax by isomerization, such that low pour point, high performance products can be prepared with a high yield. Since isomerization preserves paraffmicity rather than lowering it, the quality of the feedstocks are maximized. The usefulness of wax isomerization will depend greatly upon its feed flexl"oility, i.e. its ability to produce high yields of dewaxed oils from feeds which vary extensively in boiling range and in chemical composition, particularly wax content [ 171]. Recently anumber of patents have appeared descn~oing the use of zeolites for this isomerization process [ 172]. The catalysts used for dewaxing are usually bifimctional in nature with Pt being the hydrogenation-dehydrogenation component and a large pore typically 12MR zeolite provides the acidic component [1]. ZSM-5 is the catalyst used in Mobil's Distillate Dewaxing (MDDW)

354 or Lube Dewaxing (MLDW) processes. In this process the straight chain, waxy normal or slightly branched alkanes are able to emter the pores where they are selectively cracked and the light products are removed by distillation. Currently more than 70% of the catalytic dewaxing units in operation are based on the Mobil zeolite catalyst and process technology [7]. Table 8 Isomerization of n-octane over Pt catalysts at 1000 psig, 2.8 WHSV, 16 H2/HC, and 30% conv.[177] Catalyst Temperature i-C8sel. 2M-C7/3M-Cr [*c] (wt.%) SiO2-AI2Oz HY ZSM-5 (80 SIO2/AI203) ZSM-5 (650 SIO2/AI203) Na-Beta SAPO- 11

C3+Cs/C4 molar ratio

i-CJn-C4

DM-C6sel. (wt. %)

371 257 260

96.4 96.8 56.6

0.67 0.71 1.54

0.95 0.64 2.1

0.96 3.5 1.1

8.5 12 1.8

343

58.4

0.88

1.2

0.98

5.6

367 331

74.3 94.8

0.70 1.07

0.68 1.0

1.7 0.92

10 2.3

302826242220-

= 9

=. . . . . . . . . . .

- - - o - - R-Silica Alumina

""l,

.~ 1 8 16 .==, o 14=E 1 2 -

/

'1

:~.

1086

Pt-SAPO-11

a ........Pt-ZSM-5

...

. . . . o . . . . -o . . . . -o

o,

,! 2 0 0

2

4

6

8

10

12

Carbon Number

Figure 15. Molar distn"oution of cracked product l~om hexadecane at 1000 psig, 3.1 WHSV, 30 HJI-IC and 94% conversion [177].

14

355 Catalysts containing a hydrogenation component and an intermediate-pore silicoaluminophosphate (SAPO) molecular sieve have recently been found to have a high selectivity for wax isomerization [ 173,174,175]. A new process for dewaxing high alkane lubricating oils, called Isodewaxing [176], is being commercialized by Chevron using Pt/SAPO-11 catalyst. Table 8 shows the hexadecane isomerization selectivities of a number of Pt -loaded catalysts [ 177,178]. SAPO-11 has a low selectivity to dimethyl isomers such that fewer branches are required to obtain a given degree of pour point reduction. Since increased branching reduces the wide-temperature range fluidity of an oil the oil made using SAPO-11 catalyst has a lower sensit'wity of viscosity to temperature. Figure 15 shows that SAPO-11 had a more even distn'bution over the carbon numbers than is commonly associated with intermediate-pore sieves such as ZSM-5 [177]. The cracked also contained fewer isomers with methyl branches separated by less than two carbons than, for example, silica-alumina. Secondary hydrocracking is low and the cracked by-product is all liquid and at the same time hydrocracking of the long chain alkenes is inh~ited. These properties of SAPO-11 for this application appear to be associated with its moderate acid activity and the one-dimensional nature of its pores. Much of the catalysis appears to occur at or near to the sieve external surface. [178]. In a separate study of the relative activities of USY, mordenite, ZSM-5, Beta and SAPO-11, the SAPO-11 was found to be the only catalyst capable of isomerizing normal alkanes in the presence of iso-alkanes without large yield losses due to unwanted cracking [179]. Pt-H mordenite and ferrierite have also been used for this reaction. Recently bifimctional forms of Beta have been found to give better isomerization selectivities relative to hydrocracking and this may represent a superior and economically attractive dewaxing process [3,66]. 8.

ACKNOWLEDGEMENTS

The authors wish to thank all their colleagues both from academic and industrial research groups who kindly contributed much of the source material used in this paper. Their kind assistance and ready response to a request for information is much appreciated.

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363

Synthesis of Intermediates and Fine Chemicals using Molecular Sieve Catalysts Saskia Feast and Johannes A. Lercher University of Twente, Department of Chemical Engineering, P.O. Box 217, 7500 AE Enschede, The Netherlands

Abstract The main principles of using molecular sieve catalysts for intermediates and fine chemical synthesis are reviewed and critically discussed. Emphasis is placed on describing the role of the elementary steps. The role of the catalytic functions (acid-base, redox and host for catalytically active sites) and the role of the pore constraints in activity and selectivity are compared. Examples of successful applications are presented. 1. INTRODUCTION The use of molecular sieves as catalysts or catalyst components for synthesis of intermediates and fine chemical has increased impressively over the last two decades. A large number of reactions has been explored over a growing number of microporous materials. Also the level of understanding of the catalytic chemistry and the structure-activity relationships has greatly improved. Since the first review of Venuto and Landi~ in 1968 [1] and the one of Venuto in 1994 [2] the discovery of medium pore zeolites such as ZSM-5 [3] and of phosphate based molecular sieves [4] had the largest impact on the field. The reasons why they had such an impact, however, were quite different for the two materials. Medium pore materials (especially ZSM-5) have enabled a quantum leap in controlling the selectivity by subtly adjusting the pores size and the tortuosity and so modifying the space available for transition states and/or the diffusivities of reactants and products (i.e., inducing pronounced shape selectivity). At the same time they provide a very robust material that is straightforward to synthesize and withstands severe reaction conditions [5]. Phosphate based molecular sieves such as A1PO4-, SAPO-, MeAPO and MeAPSO, on the other hand, have considerably extended the range of lattice properties and of chemical elements incorporated into the framework [6]. The breadth and depth of these developments are well reflected in a large number of papers, patents and reviews written on the subject during the last two decades (see Fibre 1 and refs.[ 1-2,7,8,11 ]). The growth in the use of molecular sieves as catalysts as compared with macro- and mesoporous oxides was stimulated by several factors: (i) The high concentration of active sites (in comparison with oxides) results in very active catalysts. (ii) The defined pore structure allows to exclude reactants from being converted and/or products to be formed or transported out of the pores due to a too large size. (iii) The active site and the environment of that site can be designed on an atomic level for example by ion exchange [9] or chemical functionalization of the framework [ 10]. (iv) It is possible to tailor the chemical properties of molecular sieves better than those of conventional macro and mesoporous oxides. Most of these advantages relate to the fact that acid/base sites of dense or

364

macroporous oxides, sulfates, etc. depend on the way the bulk is terminated, but that for molecular sieves the entire pore surface and, thus, most acid and base sites are an integral part of the crystal structure. This allows on the one hand an unsurpassed subtlety and reproducibility in the design and modification of the acid/base sites of the molecular sieves. On the other hand it is necessary that the reactant molecules diffuse through the channels (with the pore diameter being of comparable to the size of the molecule) to reach the acid/base sites, where they may experience further severe steric constraints during reaction.

Figure 1. A summary of the number of reviews, articles andpatents published on the synthesis of fine chemicals and intermediates overmicroporous materials during the last 29 years. Thus, the rate and selectivity of catalyzed reactions over molecular sieve catalysts are influenced by factors that are affiliated with the specific interface chemistry (chemi'cal induced selectivity) and the constraints induced by the steric limitations (shape selectivity). This advantage, however, also induces drawbacks that can only be partly overcome by adjustment of the mesoscopic properties of the molecular sieve. The most prominent limitation concerns the size of the molecular sieve channels that does not allow large organic molecules to be converted. Such limitations can be partially overcome by creating a secondary meso/macro pore structure that improves the transport in the microporous materials. The implementaion of such secondary pore structure represents an important option for future improvement of zeolite based catalysts. Another limitation lies in the fact the combination of the presence of larger molecules and of strongly basic or acidic functional groups in the pores causes the desorption of products frequently to be rate limiting or even to be impossible without the help of a co-reactant (adsorption assisted desorption). This is especially important for reactions like condensation, oligomerization and nucleophilic substitution. Finally, the catalytic chemistry inside a zeolite can be seen to occur in a microscopically small tubular reactor in which the active sites are distributed over the whole reactor. It can be intuitively understood that it is difficult to forecome sequential reactions of the same type in such an environment. There is one practical aspect that limits the progress towards a molecular understanding of such processes. The large scale applications of the petroleum industry, permit

365 from an economic point of view to allocate resources for detailed investigations of the catalytic chemistry of a particular reaction. Compared to this the processes in fine chemical industry are usually small scale and allocating comparable resources for catalyst and/or process development is rarely possible. Thus, we have to develop knowledge on the generic reactivity of functional groups of the reacting molecules and their way of interaction with active sites in the molecular sieve catalysts in order to extrapolate knowledge from one process to another one. In this way it should be possible to successfully design new catalysts and/or processes. There is a large number of excellent review articles published on the use of molecular sieves for fine chemical synthesis [ 1,2,7,8,11,12,13,13,14,15,16,17,18,19,20,21,22 ]. Most of these address the problem from the side of the organic conversion and how a molecular sieve changes activity and selectivity for a particular reaction as compared to a macroporous oxide. In the light of what is said above, we have chosen for a more material oriented approach. In first instance the focus will be on outlining the chemical requirements for functional groups in molecular sieves to catalyze a particular reaction and how these groups can be manipulated to optimize activity and selectivity. Then, we will discuss how a change in the pore structure can selectively influence the transport and sorption of reactants and products in the pores and/or the formation of different transition state complexes. Finally, we will provide examples of zeolite catalyzed organic reactions that have been successfully implemented.

2. C H E M I C A L F U N C T I O N A L I T I E S OF M O L E C U L A R SIEVES Molecular sieve catalysts may provide three basic functions. They may act as solid acids and bases, provide sites or be the carrier of sites capable of undergoing valence changes in redox processes and be the host for metalorganic complexes or metal particles offering a unique steric and chemical environment [23]. It will be the purpose of this review to show, how these functionalities can be realized and manipulated within the molecular sieve lattice, at ion exchange positions or by species entrapped in the molecular sieve pores. Our focus will be to highlight the structure - activity relationships characteristic of a generic type'of reaction and to show how the complex interplay between structure and chemical composition can be utilized to obtain a catalyst with highly specialized properties. 2.1. M o l e c u l a r sieves as acids and bases

Molecular sieves consist of a three-dimensional network of metal-oxygen tetrahedra (and to a lesser extent also octahedra) that provides a regularly sized micropore structure in which the acid and base sites are structurally embedded [24,25,26](see Figure 2). Acid sites result from an imbalance between the metal oxygen stoichiometry and the formal charge on the cations. This is seen most clearly in the case ofzeolites which have three-dimensional networks of Si-O tetrahedra. Formally, there is a 4+ charge on the Si cation and a 2- charge on the oxygen anion. As every oxygen belongs to two of such tetrahedra, each of them appears neutral and the resulting lattice does not possess acidic properties. If Si is partially substituted by A1, the formal charge on the metal cation changes from 4+ to 3+. Thus, the tetrahedron which contains the aluminum cation must be negatively charged. This negative charge is balanced by a metal cation or a proton that constitute a Lewis- or a Br6nsted-acid site, respectively [2, 27,28] (see Figure 2). By convention, the bare, negatively charged, tetrahedron must then be seen as the corresponding base. Depending upon the charge on the metal catiow'proton and the oxygen, the acidic or basic properties of the molecular sieves

366

will be dominating and consequently it will be called a solid acid or base [2, 29]. Note that these acid or base properties are not a simple function of the chemical composition, but that also the structure (framework density) has a major impact.

Figure 2. Representation of the three-dimensional network of metaloxygen tetrahedra that provides a regularly sized micropore structure in which the acid and base sites are structurally embedded. The possibilities that emerge when also metal-oxygen tetrahedra with metal cations of formal charges different from 4+ or 3+ are used can be schematically seen in Figure 3. Depending on the combination of the metal cation in the frame work neutral lattices and lattices with cation or anion exchange sites are conceptually conceivable [30]. Up to now only cation exchange molecular sieves have been found and there is considerable doubt, if anion exchanging lattices can be formed because of the instabilities of the metal-oxygen bonds involved in building such a lattice. Intuitively, it can be well understood that only the theoretically highest concentration of the acid (and base) sites can be assessed by comparing the formal charges on the tetmhedra involved in building the molecular sieve framework. The strength of the acid (and base) sites will depend on a variety of factors such as the nature of the lattice cations, the overall chemical composition of the lattice, the crystal structure, etc. [31,32,33,34]. The type of cation incorporated will change the polarizability and the real charge of the lattice oxygens, thus, leading to a wide variety of chemical properties. These relationships between the chemical composition, the crystal structure and the acid/base properties of molecular sieves have been studied extensively and were thoroughly reviewed (see e.g. refs. 6,35,36,37,38,39).

367

0

9

1.4 0/~

/0

O~s1/O~si

i ",,

\

0

~ \

0

~, 0

0

0

Pure-Silica -1

O

M+

O~

Si O~"AI/O

O O

o

O o

Zeolite +1

O

~

+5

o/l

"'' O

O~

/O~p./O

AI

. ~

s~

O O

O O

AIPO4 -n Figure 3. Possibilities of tailoring the zeolite by replacing framework Si4+ or A13+with different metal cations, i.e. ps+, after ref. 7. However, it should be emphasized that the measurement and quantification of the acid/base strength of zeolites is complex and that it is difficult to directly compare the acid/base strength of a solid with that of a liquid. This results from the fact that the stabilization of carbocations and carbanions in a microporous solid differs from that in strongly polar acid and base solutions. For zeolites, it can be stated that the concentration of aluminum in the lattice is directly proportional to the concentration of acid sites and the polarity of the lattice and to a first approximation indirectly proportional to the strength of acid sites [40]. For a given chemical composition of the zeolite, the polarity of the lattice increases with decreasing framework density [41 ].

2.1.2. Acid catalyzed reactions Reactions involving carbon-carbon bond rearrangements The early use and success of molecular sieve catalysis was spurred by the dramatic improvement in activity selectivity for catalytic cracking of vacuum gas oil achieved by using the faujasite based catalysts in comparison to the previously used amorphous SiOJA120 3. These catalysts had a factor of about 103 - 104 higher catalytic activity than the amorphous SiO_-,/AI203catalysts [42]. Paraffin, C4 to C8 isomerization [43] was one of the ftrst successful non-petroleum processing applications using zeolite catalysts. The complexity of tailoring zeolite catalysts, however, is well illustrated by the fact that is only four years back that Shell has developed the first zeolite based process for isomerization ofn-butene to isobutene [44]. Traditionally, industrial isomerization processes involve the use of Br6nsted and Lewis acids, such as H2SO4 and AIC13that are uneconomical to recycle for most of the chemical applications. Replacement of such bulk chemicals with recyclable zeolites is a very attractive

368 option that is only limited by the significantly lower proton (acid site) concentration of molecular sieves in comparison to the above mentioned acids. For example, 1g of H 2 S O 4 contains 0.02 moles of protons whereas 1g of zeolite H'Y, with Si/AI = 5, contains 0.003 moles of protons. This is a rough approximation of the acidic protons available for catalysis, since it assumes 100% dissociation for both samples and that every proton is accessible in the zeolite. Note that lg of H_,SO4 occupies far less volume than the equivalent mass of zeolite, i.e., approx. 0.5 c m 3 compared to 4-6 c m 3. However, due to increasingly stringent environmental policies, the interest in solid acid catalysts remains quite high as can be seen from the number of examples in several reviews [9-22, 45]. The uniformity of the site strength that can be achieved with high silica zeolites such as ZSM5 has been shown by Mirth et al. [46] for m-xylene isomerization using in situ i.r. spectroscopy as means to characterize the concentration of sorbed reactants. The reaction rate was shown to be directly proportional to the concentration of acid sites covered by mxylene indicating that all acid sites of the ZSM-5 zeolite used were able to convert m-xylene to o- and p-xylene with the same activity per proton (see Fig.4). Note at this point that for practical reasons (limitations in the accessibility, diffusional limitations, etc.) not all sites may actually participate in the reaction. The aspects of the shape selectivity that also influence activity in a complex way will be discussed in a later section. It should be emphasized that such uniform behavior of acid sites is usually confined to high silica zeolites or, in general, to molecular sieves with a low density of acid sites. With materials of higher acid site concentrations (such as zeolites Y and X) sites with distinct differences in the acid strength have been observed [47]. These differences were attributed mainly to the existence of neighboring acid sites that produce a local situation not unlike that in the acid H2SO4 also having two protons per molecule with differing acid strength. TOF (molecules/site.s) m-x3'len~r

0.0012

0.0008

0.0004

0

v

0

'

10

20

I

~

30 40 Coverage (%)

!

50

J

t

60

i

70

Figure 4. The reaction rate for the isomerisation of m-xylene to o- and p- xylene is directly proportional to the concentration of acid sites, i.e. the coverage, [46]. The presence of neighboring acid sites, however, may be important when bimolecular reaction steps are involved in the reaction network as illustrated in the following two examples. Over a series of ZSM-5 materials Halik et al. [48] showed that the conversion of

369 ethanol to intermediate size hydrocarbons was a non-linear function of the acid site concentration with a much lower catalytic activity found below a certain concentration of acid sites. This behavior was explained by the necessity of a critical concentration of acid sites being required to maintain reasonable rates of the bimolecular reaction steps that are part of the complex transformation. It should not be interpreted that two protonated species have to react, but rather that a higher concentration of acid sites also results in a higher concentration of reactants in the pores, that favors the bimolecular reactions. Note that Lewis acid sites present in the zeolite may also play a significant role in enhancing the concentration of the reactant in the pores of the zeolite [49]. In a similar way, hydride transfer reactions in alkane/alkene transformations depend in a nonlinear fashion upon the varying concentration of acid sites. Post et al. [50] showed elegantly that the rates of these bimolecular reactions depend upon the square of the concentration of the acid sites, while the rates of the monomolecular reactions (protolytic cracking [51 ]) were linearly dependent on the proton concentration. This suggests that similar effects can also be expected in more complex organic transformations, where less thoroughly developed structure-activity relations exist. The role of Lewis acid sites in such conversions is less understood. Karge et al. [52] showed that La3§ ion exchanged zeolites that do not contain hydroxyl groups are catalytically inactive for ethylbenzene disproportionation suggesting that protons are indispensable for the carbon-carbon bond rearrangement reactions. On the other hand a number of reactions have been reported (the absence of hydroxyl groups is not certain in all those eases) that are well catalyzed by trivalent metal cation exchanged zeolites [53]. The role of the metal cation is in these instances more that of mediating the acid strength and modifying the adsorption strength than being the active site by itself. The skeletal isomerization of tetmhydrodicyclopentadiene into adamantane is an example of a very complex rearrangement that is commercially carried out over strong Lewis acids with a hydride transfer initiator. The reaction can be catalyzed by rare earth (La, Ce, Y, Nd, Yb) exchanged faujasites (Scheme 1) in a Hz/I-IC1atmosphere at 250~ Selectivities to adamantane of up to 50% have been reported, when a metal function, such as Pt, capable of catalyzing hydrogenation is added [54]. Initially acid catalyzed endo- to exo- isomefiz~on of tetmhydro-dicyclopentadiene takes place and then a series of 1,2 alkyl shifts invo'lv~,ag secondary and tertiary carbonium ions leads eventually to adamantane[55]. The possible mechanistic pathways of adamantane formation from tetmhydro-dicyclopentadiene are discussed in detail in ref. [56].

H2/HCI 250 ~ Scheme 1. Skeletal isomerisation of tetrahydrodicyclopentadiene to form adamantane. Zeolites do not only catalyze isomerizations of pure hydrocarbons. Also for molecules bearing a polar functional group, double bond and skeletal rearrangements can be performed without conversion of the functional group. Suitable zeolites should be rather apolar with a low concentration of acid sites, e.g., HZSM5. The interaction of polar functional groups with the pore walls of these rather apolar zeolites are weak [57]and hence polarization and

370 activation of these groups can be minimized. An example for a double bond relocation is the isomerization of2-ethyl propenal into trans-2-methyl-2-butenal over CeBZSM5 [58] (Scheme 2). An example for a skeletal isomerization is the allylic rearrangement of 1,4 diacetoxybutene over ZSM-5 while retaining the functional group intact (see Scheme 3). H3C~ CH2 H2C

CeB-ZSM-5 ~

H3C / C H ' ~ CH3

300 ~

CHO

CHO

Scheme 2. Isomerisation of 2-ethyl propenal to trans-2methyl-2-butenal. o

o

A.

%A.

C

,O..

CH2 Y

CH3 CH.

o

~

ZSM-5 Oc 300

.. "

H2C O

CH3

CH

HC/ll "O~T/CH3 CH2

O

Scheme 3. Skeletal isomerisation of 1,4-diacetoxybutene over ZSM-5.

lsomerization involving heteroatoms Molecular sieves are also well suited as catalysts for isomerization of molecules containing heteroatoms. The weaker strength of the bond between a carbon and a heteroatom compared to a carbon-carbon bond usually allows, and even necessitates, working at relatively low temperatures. Good examples are the isomerization ofhalogenated aromatic molecules such as chlorophenols, chlorothiophene, bromothiophene and iodothiophene over Z.SM-5 zeolites [59]. The optimum reaction temperature for the last three molecules gradually drops from 300~ to 100~ in parallel with the increasingly weaker carbon-halogen bond.

Nucleophilic substitutions In nucleophilic substitutions one can distinguish between two mechanisms, i.e., the two step nucleophilic substitution (SN~)and the one step process (Sin). In the latter route, the highly polar intermediate species or, in the limiting case, the carbocation is stabilized by the catalyst. Direct evidence for the presence of carbonium and carbenium ions in the molecular sieve pores is scarce. Experiments point to such species only in the presence of very strong acid sites provided relatively basic reactant molecules are used [60]. Even in such cases the interpretation of the experimental data does not seem to be unequivocal [61 ]. Most results suggest that the true cation exists only in the transition state resulting in a quite complex reaction coordinate. The course of the reaction is determined by the chemical nature of the leaving and the substituting group, the acid/base properties ofthe molecular sieve, the influence of co-reactants and the availability of space for the reaction to take place. The majority of the nucleophilic substitutions involve the replacement of an-OH group with an-NH,-S,-SH,

371 -OR or another functional group. One of the major problems is that in many cases the resulting product interacts more strongly with the molecular sieve than the reactant. This leads to the situation that many reactions are desorption controlled and need either a reactant to desorb (adsorption assisted desorption) or a gaseous/liquid cocatalyst that also facilitates the desorption of the products without participating in the reaction. Note that for liquid phase reactions the solvent can take over the role of the cocatalyst. Etherifications, conceptionally one of the simplest reactions to catalyze, occur over most zeolites. Molecular sieves, however, have too low acid site density to make them interesting for commercial applications. Usually, resins like amberlyst are used for that purpose [62]. On the other hand etherification is experimentally and theoretically well studied and understood. Using the example of dimethylether formation from methanol one clearly sees, how the reaction conditions influence the reaction mechanism, i.e., whether the reaction proceeds along a Sm or a Sm pathway [63,64,65]. Temperature programmed reaction studies of methanol conversion over HZSM5 suggest that three reaction routes to form dimethylether exist, i.e., via an alkoxonium cation and via t w o alkoxy pathways [65]. At low temperatures the reaction proceeds v/a an Eley-Rideal type mechanism_ In the transition state one methanol molecule forms a methoxonium ion, water leaves the molecule and simultaneously another weakly sorbed methanol binds to the methyl group forming protonated dimethylether (see Scheme 4). The protonated dimethylether donates immediately the proton back to the zeolite and desorbs. M.S. Response (Arb.U.)

V~

nol

,,

~

C H , O H + CH,OH: + ~

D M E + H,O + I-I*

i_~.'. CH~OH + SiOCH3A[ - - ~ DME + SiOHAI CHjOH + SiOCH, ~

!

I

400

500

I

A

600 700 Temperature (K)

....

DME + SiOH

|

800

,

|

900

Scheme 4. Temperature programmed desorption/reaction of methanol on HZSM-5 [65]. As the reaction temperature increases, part of the methanol molecules will be transformed into methoxy groups that replace the proton in bridging (SiOHAI) and terminal (SiOH) hydroxyl groups. These methoxy groups react with weakly associated methanol to form dimethylether under simultaneous restitution of the hydroxyl group. While the methoxy group is covalently bound to the zeolite lattice, its reactivity increases with the acid strength of the hydroxyl group it replaced [65,66,67]. Thus, methoxy groups at bridging hydroxyl

372

groups produce dimethylether at lower temperatures than methoxy groups at terminal hydroxyl groups [67]. Comparison of the chemistry over various zeolites indicates that formation and reactivity of a specific type of methoxy group is connected in a complex way with the polarizability of the lattice and the overall acid/base properties. Methoxy groups at bridging sites are more easily formed and consumed on FAU type materials than on MFI type materials [68]. Recent theoretical calculations by Blazowski and van Santen suggest that indeed the pathway to form DME via the methoxonium ion is energetically favored over the pathway via methoxy groups [69]. The data clearly agree with the observed strong temperature dependence of the reaction mechanism as reported in ref. [65]. Blazowski and van Santen used ab initio calculations to show that the SN_,reaction involves a complex transition state in which four reactions have to proceed in a synchronous manner, i.e., (i) formation of a methoxonium ion by proton donation from the zeolite, (ii) cleavage of water from the methoxonium ion and formation of a methylcarbenium ion, (iii) binding of the methyl carbenium ion to the second methanol molecule to form protonated dimethylether and (iv) donation of the proton back to the zeolite. Note that according to the calculations all must occur in a concerted manner, as the protonated species are only found to be stable in the transition state. If this proves to be true the transition states may be quite difficult to achieve, i.e., transition entropy must be quite low and, hence, also the reaction rates must be low. Stabilization of the methoxonium ion by the catalyst would lead to a less complex transition state and hence, one might expect the intrinsic rates of the reaction to be higher. The initial results of methanol sorption on organic resins and heteropoly acids indicate that such a situation may be attained with these materials [70]. Note that this would make molecular sieves only preferable, if special properties, such as pronounced shape selectivity, would be required. Amination of alcohols follows a mechanistic pathway similar to etherification [71 ]. Due to the basicity of the reactants one might expect that for this reaction, ammonia will be present in the molecular sieve pores in the form of an ammonium ion and the alkyl group of the alcohol would substitute for one of the protons of the ammonium ion. Indeed, if an alcohol is passed over the ammonium form of a zeolite, amines are readily formed [72]. These alkylamines, however, cannot desorb and remain chemisorbed in the zeolite pores under typical reaction conditions (T = 353 ~ The apparent reaction mechanism can be classified a/s S'm, ammonia being stabilized by the molecular sieve in the form of the ammonium ion. In order to obtain a successful reaction, the ammonium ion must protonate the alcohol in the transition state, thus, generating a H20 leaving group. In a simultaneous step the alkyl group must dock onto the lone electron pair of the nitrogen forming an alkylammonium ion. Even under more severe reaction conditions (temperatures higher than 353 ~ the alkylammonium ions are unable to desorb from the acid sites [73,74]. These alkylamines released into the gas phase stem actually from a further nucleophilic substitution in which the alkylgroup of the alkylammonium ion is scavenged by weakly adsorbed ammonia (see Scheme 5) This shows that the type of zeolite might not be as important as the reaction conditions and indeed several zeolites have been claimed to be suitable for amination of alcohols [75,76,77]. The acid strength of such zeolites should be as high as possible in order to assure that all the acid sites are covered by ammonia and amines, thus, preventing the formation of ethers and higher hydrocarbons from the alcohol over free acid sites [78]. The requirement for zeolites of high acid strength for the alkylation of ammonia by alcohols contrasts with the need for weakly acidic zeolites for the addition reactions between alkenes and ammonia [79]. In these reactions the alkene has to be activated by the Bronsted acid site of the zeolite and that is only possible when the acid sites are not fully blocked by

373 ammonium ions. In addition to the weak acid sites of the catalyst high reaction temperatures and high pressures of alkenes are necessary to achieve this. H HAl

H

H

0

+ CH. I~" .N CH.I ~CH 3 ~CH~ O"

0

/ A!\ OO

H CH3 +1 .N~ CH31 CH3

(ii) proton transfer 0

\ / \ / \ / /Si \

O

(i) methyl scavenging

CH

0

\ / \

/Si ~ OO

O-

// S i N O

3

0

/ \ / AI\

O

OO

/ OO

/ si \

o

o

Scheme 5. Proposed mechanism for the removal of methyl amines by scavenging with ammonia, after [72]. A somewhat more involved example is the transformation of oxygen containing heterocycles into nitrogen or sulphur containing heterocycles [80]. Again, not the structure of the molecular sieve has been found to be important (provided there is enough space within in the zeolite pores to accommodate the reactants and products), but rather the acid strength and nature of the acid site. Hatada et al. reported the transformation of y-butyrolactone into 2-pyrrolidinone over a series of metal exchanged Y zeolites (Scheme 6) [81 ]. For alkali metal and alkaline earth metal exchanged FAU a direct dependence of the yield of 2-pyrrolidinone upon the strength of the electrostatic field of the cation was found. This indicates that the strength of the coordination of y-butyrolactone to the metal cation is the most important parameter influencing the catalytic conversion. For transition metal (Co, Ni, Cu and Zn) exchanged zeolite Y a correlation between the cation field strength and the activity was not observed. It is speculated that this is due to the non-spherical nature of transition metal cation orbitals, especially of their partially filled d-orbitals, and thus the simple electrostatic model is no longer applicable [82]. The reaction proceeds via initial polarization of the carbonyl group of y-butyrolactone by the metal cation. In the next step, ammonia binds to the carbon atom of the polarised carbonyl group forming an acid amide, which then rapidly dehydrates under ring closure. The stronger the electrostatic field of the metal cation the stronger the interaction between the carbonyl group and the metal cation leading to a more polarised C=O bond, which is then more reactive towards ammonia. (0.~0

Me-Y

(N-vO

2-pyrolidinone

7-butyrolactone B r o ~ acid catalysed side-reaction

~ A

~

o NH,_

co-hydro xybutryo nit r it e Scheme 6. The transformation of g-butyrolacone into 2pyrrolinone over MeY.

374 In contrast to the previously discussed carbon-carbon bond rearrangements, these results clearly show that Lewis acid sites can also act as catalytically active sites for nucleophilic substitutions. Note that if catalysts without Bronsted acid sites are used (i.e., with zeolites exchanged with monovalent cations) the competitive side reaction leading to o-hydroxybutryonitrile via protonation of the acid amide can be completely suppresscd (Scheme 6). .

.

I)

1I)

.

.

.

.

.

.

.

.

.

.

.

CI

OH

.

.

.

.

.

.

.

.

NH, +

HCI

+

H20

NH2

Scheme 7. The reactions of chlorobenzene with ammonia (I) and phenol with ammonia (1I). . _ _

Nucleophilic substitution of an aromatic ring is difficult to achieve, as the ~-eleetrons will repel the electron density of the incoming molecule and it is difficult for the aromatic ring to accommodate the additional electrons. The substitution becomes more likely when a strongly electron withdrawing group is replaced by a more electron donating one. Examples of this case are the reaction of chlorobenzene with ammonia to form aniline and HC1 or the reaction of phenol with ammonia to give aniline and water (see Scheme 7) [83,84]. The most selective zeolite catalysts are based on mordenite and ZSM-5 and contain copper or cobalt ions [85]. The rate determining step seems to be the release of the electron withdrawing group [76,86]. A similar reaction mechanism seems also to be responsible for the formation of diphenyl from two phenol molecules [87]. The zeolite faciliates nucleophilic substitutionSof aromatics by electron withdrawal from the aromatic ring via coordination on the metal cations

[88].

Addition and elimination reactions of carbonyl compounds The polar nature of the carbonyl group allows for addition ofnucleophiles at the carbon atom. Molecular sieves catalyze these reactions by enhancing the polarity of the carbonyl group through interactions between the Bronsted or Lewis acid sites and the oxygen of the carbonyl group. If the nucleophile retains a proton, water can be easily eliminated and the overall reaction leads to the replacement of the oxygen by another nucleophile (see Scheme 8 and refs. [89,90]). The reactions involve the addition ofH20, ROH, RSH, HCN and HSO3 to the carbonyl group yielding the corresponding hydrates, (semi)acetals, cyanhydrines etc. Acetal and ketal formation from aldehydes, resp. ketones and alcohols occurs over mordenite and other acidic zeolites [91 ] slightly above ambient temperatures in the liquid phase. The reaction is not confined to simple alcohols, diols can also be converted (e.g., cyclohexanone reacts with ethylglycol to 1,4, dioxaspiro(4,5)decane [2]). Note that it is likely that desorption controls the rate of such reactions as the product molecules are larger than the reactants and have, hence, a higher adsorption constant.

375

R.

Null9

R~

R,

,g O H

R

R2 + H20 --H

RI

Scheme 8. General mechanism of nucleophilic addition with subsequent elimination of water. The reaction ofacetonyi acetone to dimethylfuran, catalyzed by HZSM-5, is an example of an intramolecular addition reaction involving two carbonyl groups, followed by a 13elimination of water (Scheme 9). The Bransted acid site of the zeolite protonates one of the carbonyl groups, while the oxygen atom of the second C-O group binds to the positively charged carbon atom of the protonated carbonyl group. The use of the rather hydrophopic zeolite HZSM-5 facilitates the elimination of water after the ring closure reaction. ~___O H3C~ . . .

OH CH,

n +

+ H3C

H3C CH2

"

CH 3

CH 3

H+

OH ~C~ H

_-- o

i IH CH 3

H3C~ C ' - - - C H ~

O ~/

I

+ H20

H3C

acetonylaceto ne

dimethylfuran

Scheme 9. The reaction of acetonyl acetone to give dimethylfiwan over HZSM-5. The products of the ketone or aldehyde conversion with ammonia and amines depends upon the availability of a proton at the nitrogen atom. If such a hydrogen is present, e.g., in the reaction of benzaldehyde with NH3, the addition of ammonia to the carbonyl group is followed by a rapid elimination of water. The so formed benzylidine imine subsequently dehydrogenates to form benzonitrile in the presence of transition metal ions, such as Co, Cr, Cu, Zn or Mn, (Scheme 10) [74]. If the acidic proton at the nitrogen is lacking, e.g., it, ~ e reaction ofdiethylamine with cyclohexanone, the formation ofa C=N bond is prevented during the dehydration step and instead a ring C=C bond is formed. The zeolites used in this case are large pore zeolites such as CaX or HMOR. Note that with these catalysts drying agents have to be added and it appears to be likely that large pore hydrophobic zeolites would be a better choice as catalyst. N

+ NH3

-'~

Scheme 10. Reaction ofbenzaldehyde with ammonia yields benzonitrile.

376 The acidic 10 and 12 membered ring zeolites (H-MOR, ZSM-5, ZSM-11) can also be used to catalyze the condensation ofalkenes with aldehydes to form unsaturated alcohols, acetals etc. (Prins reaction)J92]. Chang et a/.[93] showed that this reaction involves in the initial step the activation of the aldehyde by a Bronsted acid site to generate an electrophilic species. The condensation with, e.g., isobutene leads then to a primary alcohol with a positive charge at the tertiary carbon atom. Elimination of water and addition of further aldehyde molecules may lead to a broad variety of products. Some of these reactions can be effectively blocked by chosing zeolites with the appropriate pore size [94,95].

Reatv'angements of nitrogen containing compounds The most prominent examples of this type of reaction are the Fischer Indole synthesis, the Beckmann rearrangement and the benzylamine rearrangement. For all three reactions rather complex mechanisms have been proposed. On comparing the stmctt~e- activity relationships for these transformations, it becomes clear that generalisations are difficult and that a complex interplay between pore shape and size, the acid strength and the polarity of the zeolite lattice seems to control the activity and selectivity for a given reaction. An example of the Fischer Indole synthesis of substituted indoles involves the initial condensation ofa phenylhydrazine and 3-heptanone to form a phenylhydrazone (see Scheme 11). The phenylhydrazone undergoes (Br~nsted or Lewis) acid catalysed tautomerisation to give the enhydrazine tautomer which further rearranges and then eliminates ammonia to form the indole. Two products are possible, the bulky 2-ethyl-3-propyl-indole (Scheme 11) and the more linear 2-butyl-3-methyl-indole. In the homogeneous phase the selectvity towards one of the two products is controlled by the acid strength of the catalyst. The role of the zeolite in controlling the selectivity in the heterogeneously catalysed process is not unabiguously resolved. Van Bekkum et aL [96] showed that the 'linear' isomer was predominantly produced over most zeolites suggesting at first sight that the constraints in the molecular sieve pores favor the product with the smaller minimum kinetic diameter. However, since HNaX was more selective to the linear isomer than the isostmctural HNaY it must be concluded that the selectivity is not exclusively goverened by classical zeolite shape selectivity.

o

II

+ NHNH,

~

C

~

- H,O - r-

cat

~ - NH 3

"bulky" (III)

_

(i) (II)

Scheme 11.

~

"linear"

(iv) The Fischer Indoie reaction of phenylhydrazine (I) with 3-heptanone (II) giving two indole products.

Similarly, for the vapor phase Beckman rearrangement of, e.g., cyclohexanone oxime into caprolactam (Scheme 12) the zeolite structure was initially thought to be the most decisive factor for selectivity. Small pore zeolite HA (pore size 4fi,) produced caprolactam with only

377 4% selectivity at 14% conversion, whereas medium pore sized HZSM-5 (5.5A) gave 50% selectivity at 100% conversion [97] and large pore HY (7.5A) 89% selectivity at 82% conversion [98]. Recent studies however, suggest that the external silanol groups are the dominant catalytically active sites. Sato et al. [99] observed that the catalytic activity and selectivity to E-caprolactam increased in parallel with the Si/AI ratio of ZSM-5 and were directly proportional to the concentration of weak acid sites on the external surface of H-ZSM5. Interestingly, amorphous silica that also contained a high concentration of such SiOH groups also gave a very high initial conversion, but deactivated rapidly due to coking. Highly crystalline zeolite samples were shown to be more selective and more active indicating that the regularity and/or the low density of such weakly acidic silanols are essential for high lactam selectivity. o

C;

"OH

H

Scheme 12.The Beckmann rearrangement of cyclohexanone oxime into e-caprolactam. Typical examples for benzamine warmngement are the conversion of aniline and 1,3diaminobenzene with ammonia to 2-methylpyridine (Scheme 13) and a-amino-a' -picoline. Although several acidic oxides were found to be active, the best results were obtained with HZSM-5. It was found to be more active than SIOJA1203(48% conversion compared to 29%), but showed similar selectivities (83% over HZSM-5 and 98% over SiO2/Al203) [ 100]. The reaction seems to proceed via the addition of ammonia to a protonated aminobenzene (probably present in the form of an cyclic enamine). After an enamine-amine isomerization, the ring is opened via a reverse aldol-type reaction. Upon addition of the amino group to the imine double bond the ring closes again. After elimination of ammonia from the resulting aminal the final product is obtained [ 101 ]. Note that this potentially provides a new simple route for the production ofaminopyridines replacing the current complicated industrial process [102]. NH~ NH 3 ~

~N]N ~

(I)

CHs

(III)

NH~

NH,_

(If)

HsC

N

(IV)

CH3

Scheme 13. Benamine rearrangements of aaniline (I) and 1,3-diaminobenzene (II) to apicoline (III) and a-amino a'-picoline (IV).

378 Electrophilic substitution on the aromatic ring

In general terms, this type of reaction is characterized by the attack of a species with a positive partial charge, a positively charged species or a radical (i.e, species that are electron deficient) on an aromatic ring, preferably on the carbon atom with the highest negative charge. A broad variety of such electrophilic species has been reported to exist in the pores of molecular sieves (see Fig 5.after ref. [2]). The generation of such species can take place via several pathways from amongst which protonation, hydride abstraction and cleavage of polar groups are the most important ones. A general mechanism can be visualized as depicted in Scheme 14.

RH ArR

o R--

o

\

- - X (X=CI, O C R

ROH, ROR, AK)R

HNO 3, N204, NO 2

Cl 2, Br 2, HBr + 02, 12 + 02

.

D-

|

~

H+

Figure 5. Range of electrophilic agents employed in electrophilic aromatic substitutions over zeolite catalysts, after ref [2]. In the first step, coordinative bonding between the ~-electrons of the aromatic ring and the electrophile frequently occurs. Recent spectroscopic evidence for such an intermediate was reported for the methylation of toluene [ 103]. The aromatic ring should only be weakly held by the zeolite in order not to decrease the availability of the re-electrons. Then, a localized interaction with a carbon atom of the ring preceeds the actual substitution. In the presence of a substitutent on the ring, the carbon atom position at which the interaction with the electrophile occurs will depend upon the inductive effects induced by the ring substituent. For electron donating substituents the preferred carbon atoms to accept the electrophile are those in ortho and para position to the substituent group [ 104]. H

E

E

+

Scheme 14.

H +

General scheme for electrophilic aromatic substitution.

The overall reactivity of the aromatic ring will also depend upon the nature of the substituent. Electron donating properties of the substituents increase the availablity of nelectrons at the aromatic ring, while the electron withdrawal properties reduce it. In that respect alkyl-, hydroxyl-, alkoxy-, or amine groups increase the reactivity, while the presence of halogen or nitro groups will reduce it. The reactivity ofheterocycles also depends upon whether or not the ring is 3z-electron excessive. This results in pyrrole and thiophene being

379 more reactive than benzene, while pyridine is less reactive [92]. The examples that we have chosen to demonstrate the required design of the molecular sieve catalysts and the necessary adjustment of the reaction conditions are alkylation, acylation, nitration and chlorination. Friedel Crafts type alkylations of benzene by alkenes involve the initial formation of a lattice associated carbenium ion, formed by protonation of the sorbed olefin. The chemisorbed alkene is covalently bound to the zeolite in the form of an alkoxy group and the carbenium ion formed exists only in the transition state. As would be expected from conventional Friedel Crafts alkylation, the reaction rate over acidic molecular sieves also increases with the degree of substitution of the aromatic ring (tetramethyl > trimethyl > dimethyl > methyl > unsubstituted benzene). The spatial restrictions induced by the pore size and geometry frequently inhibit the formation of large multisubstituted products (see also the section on shape selectivity). For similar alkylation reactions modified faujasites need lower temperatures to catalyze the reaction with the same rate (under otherwise identical reaction conditions) than amorphous silica-alumina catalysts [ 105]. The difference is explained with the higher site strength and density in the zeolite catalysts. The fact that the original Friedel Crafts catalyst (promoted Lewis acid - AICI3-HC1) is reactive at yet lower temperatures than modified faujasites, suggests that a microporous material with higher acid strength could push the operating temperatures even lower. In general, a suitable catalyst should have high acid site strength and sorb the substituting molecule strongly. A good example of this is the alkylation of benzene with propene for which the reaction rate over divalent cation exchanged Y zeolites was found to decrease in the order Mg--Ca > Sr > Ba in accordance with the decreasing acid strength of the materials [ 106].

IR INTENSITY (2400 cm-I)

I0-

6

2 0

w I

0

.t

t

I

0.000.'3

0.001

0.00 i 5

J,

t

t

0.(}02

0.00"-5

I'OF [molec site s]

F i b r e 6. Reaction rate of methylation of toluene over HZSM-5 is directly proportional to the concentration of chemisorbed methanol [108]. The alkylation of toluene with methanol over HZSM-5 proceeds at low temperatures via a protonated methanol species in the transition state [ 107] and weakly coadsorbed toluene

as classically predicted for Friedel Crafts alkylation. The reaction rate is directly proportional to the concentration of the chemisorbed methanol (in the presence of excess toluene) as shown in Figure 6 [ 108]. Alkylation leads preferentially to ortho- and para- substituted products which rapidly isomerise in the zeolite pores. Specific reaction conditions and tailoring of the catalyst pore structure can be employed so that para- substituted products are preferentially

380 produced [ 109]. The reasons for this selectivity and the methods for optimizing the catalyst performance will be discussed in a later section. The catalysis appears to be completely controlled by the Br/Snsted acid sites with the role of the Lewis acid sites being marginal [ 110]. For the alkylation of the more active phenol both Bronsted and Lewis acid sites are claimed to participate in the catalytic activation [ 111 ]. The Bronsted acid sites activate the alkylating agent by protonation, whereas the Lewis acid sites can activate the alkylating agent and phenol by coordination, and/or phenol by deprotonation. If activation of the alkylating agent and the phenol occurs on the same Lewis acid site, the predominant product will be the ortho- substituted isomer.

OH

OH ~ CH3CO2H

~

O

~ / ~

Scheme 15. The accylation of phenol with acetic acid to yield 2-hydroxyacetophenone.

Acylation is currently carded out industrially with stoichiometric amounts of metal chlorides or mineral acids. Zeolites can replace liquid acids in this two step process consisting ofesterifcation and the Fries rearrangement. Several possible starting compounds for acylation, such as acid halides, carboxylic acids and acid anhydrides exist. The type of acid site (i.e., Br~nsted or Lewis) in the molecular sieve has to be adjusted for the acylating agent. A Lewis acid, such as La3+in the zeolite, will not activate a carboxylic acid to give an acylium ion, but will rather form a carboxylate anion. In contrast, Bronsted acidic hydroxyl groups will readily help to generate an acylium ions. When using an acid chloride as the acylating agent, Lewis acid sites are a better choice than Brrnsted acid sites, since they assist heterolytic dissociation by forming strong bonds to the halogen anion. An example of the need'for Bronsted acid sites is the acylation of phenol with acetic acid to yield 2-hydroxyacetophenone [112] with HZSM-5 as catalyst (Scheme 15). The latter situation is exemplified with the paradirected acylation of toluene with several aliphatic acid chlorides over ZnY [113]. Other examples are the formation of anthraquinone from benzene and phthalic anhydride or from phthalic anhydride alone over NaCe and NaZn exchanged FAU, respectively [ 114]. Acylation ofheterocyclics, such as thiophene, seems to require a lower acid strength of the catalyst which is best met by B-ZSM-5. The reaction ofthiophene with acetic anhydride to 2-acetylthiophene proceeds with 99% selectivity at 24% conversion and the conversion of pyrrole with acetic anhydride to 2-acetylpyrrole with 98% selectivity at 41% conversion [115]. Nitration of aromatic compounds requires very strong acid sites to stabilize the NO2§ cation which is an important intermediate in liquid phase nitration [ 104]. Several nitrating agents such as HNO3, NO_,and N,_O4have been successfully applied using mainly dealuminated mordenites faujasites and elevated pressures [ 116]. The stability of the zeolites is a major problem given the highly acidic reaction medium. A combination of high crystallinity and sufficient extra-zeolite surface area (the presence of extra-lattice material) was found to be beneficial for stabilizing the catalysts.

381 In the halogenation of aromatic molecules the role of the zeolite is to polarize the CI_, or Br2 molecule in order to enable it to attack the aromatic nucleus. The polarization is aided by an alkali or an alkali earth cation [ 117]. In most cases addition of CI,_ to benzene dominates over MFI and FAU type molecular sieves leading to chlorocyclohexane intermediates. A minor portion of the aromatic molecules, however, also reacts directly to chlorobenzenes via electrophilic substitution. Larger pore zeolites usually lead to a higher degree of chlorination which can be explained by the availability of the space in the zeolite pores[ 118].

2.1.2 Reactions catalyzed by basic sites In contrast to the situation found with acid catalyzed reactions, the role of the zeolite is less well defined for base catalyzed reactions. This results from the fact that all "basic" zeolites contain alkali cations that act as (weak) Lewis acid sites. Thus, most of the chemistry described in this chapter involves Lewis acid and base sites. It should be stressed that for all acid/base catalyzed reactions both sites are involved in the reaction sequence. In many of the acid catalyzed reactions the importance of the acid sites dominates so drastically that attention is paid only to the acidic function [ 119]. We speak, therefore, of base catalyzed reactions, if the strength of the base sites is high enough to stabilize anionic or polarized species with a marked negative charge and if these species are part of the catalytic cycle. Interpretation of catalytic results with respect to the role of acid and base sites remains, however, always ambiguous as the stabilizing effect of the metal cation (for zeolites usually an alkali metal cation) is difficult to assess. There is a second problem affiliated with defining the catalytically active site for base catalyzed reactions. Acid sites, irrespective of whether they are of Lewis or Brmasted nature, are always a minority species. The majority of the molecular sieve lattice is comprised of the more electronegative oxygen. Consequently, it is straightforward to characterize the minority species (indeed a large variety of methods have been developed in that respect [ 120], while characterizing the majority species, i.e., base sites, of the catalyst still poses a major problem. Thus, evidence on the location of the base sites in the molecular "sieve channels is ambiguous [ 121 ]. The main question in this respect is whether or not the base sites are localized (e.g., next to the alkali cation) or whether all oxygens of the molecular sieve lattice act as base sites [ 122]. In base catalyzed reactions relatively low rates are achievable compared to acid catalyzed reactions and in many cases minor traces of acidic protons may change the selectivity of a reaction dramatically [ 110]. In order to overcome this problem, catalysts are prepared with a slight excess of the alkali cations. Very strong basic sites have been created by supporting metallic sodium in the zeolite pores [ 123,124]. Recently, the method of using an excess of alkali metal cations has been expanded to load zeolites systematically with alkali metal oxides. This approach results in the zeolites being used more as a support than as base catalyst [125]. The oxidic nanophase particles in the zeolites are created by thermal decomposition of the corresponding alkali metal acetate, nitrate or hydroxide [ 125, 122]. In contrast to the situation found with solid acids, basic mesoporous oxides are excellent catalysts and the use of basic molecular sieves might be advantageous only if shape selectivity is needed for a particular reaction. In general, the action of the basic catalysts can be twofold. On the one hand the high electrostatic field in the pores and the polar lattice of basic molecular sieves facilitates

382 proton abstraction from functional groups of reactant molecules. Depending upon whether this leads to a stabilized carbanion or to a polarized functional group of the reacting molecule, the reaction occurs in a one or a two step process. On the other hand, the functional group is polarised by an electron pair donor/electron pair acceptor interaction with the alkali metal cation. The positive part of the dipole in the polar group and the rest of the molecule may interact with the basic oxygens close to the cation. Hydride transfer, which is frequently part of the catalytic sequence in base catalyzed reactions, is more a consequence of the close vicinity of the sorbed molecules than of being induced by the basic nature of the zeolite. The reactions that will be employed to exemplify these general principals are alcohol dehydrogenation, olefin isomerization, aldol condensation and Meerwein-Ponndorf-Verley reductions. Dehydrogenation of alcohols occurs over base zeolites in the presence and in the absence of oxygen [126]. Dehydration is the prevailing reaction over acid zeolites [ 127]. Higher reaction temperatures are required for dehydrogenation than for dehydration, due to the higher energy of activation for the former reaction [128]. The catalytic activity is related to the concentration and the type of alkali cation, i.e., with increasing size of the alkali cation and increasing level of exchange the rate/selectivity to dehydrogenated products increases [ 126]. Mechanistically, the reaction is thought to proceed via abstraction of a proton from the hydroxyl group of the alcohol by forming an alkoxylate. [3-Hydride abstraction produces hydrogen that desorbs, together with the ketone/aldehyde formed to close the catalytic cycle [129]. Poisoning experiments with pyridine (to block the acid sites) and with phenol (to block the base sites) indeed show that dehydration requires strong acid sites to be catalyzed whilst dehydrogenation requires strong basic sites [130,131]. Conceptually, one might expect that higher concentrations of aluminum in the zeolites and (for a given alkali metal ion) higher concentrations of alkali metal ions would generate a stronger basic zeolite [ 132, 122]. However, the chemical composition of the molecular sieve seems to influence the base strength and, hence, the catalytic activity in a more complex way. Davies et al. [125] reported that Cs exchanged Y type zeolites are an order of magnitude more active than the corresponding X type materials for the catalytic dehydrogenation of isopropanol. In that context it should be emphasized that using the selectivity to dehydration find dehydrogenation of alcohols to characterize the acid and base properties requires the comparison of results at one (arbitary) reference temperature. Since the apparent energies of activation for the two reactions are quite different, it is difficult to judge whether or not changes in the selectivity observed at varying reaction temperatures are induced by changes in the acid/base properties or by the different energies of activation. Basic zeolites are able to catalyze double bond isomerization of olefins [133]. Although this can also be achieved with acidic zeolites, the lower reactivity of basic zeolites towards hydrocarbons (i.e., the complete absence of skeletal isomerization) leads to higher yields [ 134]. A good example for this is the double bond isomerization of 1-octene over potassium loaded NaY. It is claimed that high yields can be achieved in that way and that the impregnation of the zeolite with an excess of alkali cations is important to obtain a good catalyst [ 135]. Aldol condensations are catalyzed by acid and basic zeolites (see Scheme 16). In the base catalyzed route the anionic species is generated by the interaction of the basic site with the hydrogen in a-position to the carbonyl group. The a-carbon atom (bearing a negative partial charge) then forms a new C-C bond with the carbon atom of the carbonyl group of another aldehyde molecule generating a larger 13-hydroxy carbonyl compound. Subsequent

383 dehydration leads to the formation of an a,[3-unsaturated aldehyde. Successful examples include the synthesis of crotonaldehyde from acetaldehyde over SAPO34 [136] and the conversion of acetone into diacetonaldehyde, mesityloxide and subsequent products over various alkali exchanged and alkali oxide loaded large pore zeolites [137]. The pore size of the zeolite influences the product distribution via suppression of the formation of the bulkier products. The condensation of acetone over NaX and NaL type zeolites is an example of this shape selectivity. As outlined above acetone is converted to diacetonalcohol and mesityloxide which may further react to isophorone. The product ratio of mesityloxide to the bulkier isophorone was 0.75 for zeolite X and 1.87 for zeolite L [138,139].

Scheme 16. Side chain alkylation of toluene with methanol over basic zeolites.

A special case of an aldol condensation is the side chain alkylation of alkylaromatics over basic zeolites such as alkali containing faujasites. The reaction requires the complete absence of protons in the zeolite, since these would catalyze ring alkylation with a much higher rate. The most well studied example is the side chain alkylation of toluene with methanol over a variety of alkali containing zeolites. Note that also alkenes can be used as alkylating agents for this reaction, but they require a higher base strength, i.e., the presence of metallic Na [ 137]. The role of the basic zeolite is twofold. It polarizes the methyl group of toluene which leads in the limiting form to a carbanion structure [140,141] and it catalyzes the conversion of methanol to formaldehyde [ 142]. The negatively charged carbon at the toluene carbanion forms a C-C bond with the positively charged carbon atom of chemisorbed formaldehyde forming an intermediate that rapidly eliminates water and ~el'cls styrene (see Scheme 16). The reaction rate seems to be determined by the availability of toluene (which is more readily stabilized in the faujasite pores by the larger alkali cations than methanol) and formaldehyde. Indeed, addition of an extra dehydrogenating function by the addition of ZnO to the zeolite leads to a drastic improvement in the activity [ 143]. The stability of carbanions follows the opposite sequence to that of carbonium ions, i.e., carbanions at primary carbon atoms are more stable than those at secondary or tertiary carbon atoms [144]. Thus, one would expect that it might be possible to convert methane and ethane with methanol. Unfortunately activation and/or proton abstraction from a]kanes seems not to be possible to a significant extent, as attempts to react methanol with methane or ethane have up till now failed. Presumably, one needs to couple such experiments with oxidative dehydrogenation [ 145] in order to achieve feasible conversions. A special case in which a strongly basic catalyst was used to produce 4-methyl thiazole in a simplified reaction sequence (replacing a five step synthesis with a two step synthesis) has been reported recently [ 146]. The catalysts (Cs loaded MFI and BEA) proved to be effective for the conversion of a ketone to an imine, more specifically acetone and rnethylamine into the corresponding imine. In the second step this imine is converted with SO 2 into 4-methyl thiazole (Scheme 17). Using Cs sulfate as the Cs source resulted in the

384 best catalyst and given the acidity and basicity of the reactants, one can speculate that sulfate species may also prevail in the rate determining step.

)=

O

+

~Nk,

CH3NH 9 -

§ SO2

-~

,,

?-

N

\

+

H20

Cszeolite ~ ~ - - ~ S

Scheme 17. Base catalysed conversion of acetone into an imine, which is further reacted to give 4-methyl thiazole.

The Meerwein-Ponndorf-Vedey reaction is conventionally seen as a base catalyzed reduction of a complex aldehyde by a secondary alcohol, e.g., isopropanol. The reaction is catalyzed by alkali metal exchanged zeolites and the product distribution is influenced by the strength of the base sites and/or by spatial constraints in the zeolite pores. An example is the reduction of citronellal (I) with isopropanol (Scheme 18) which gives 86% isopulegol and 14% citronellol at 87% conversion with Li or NaX as catalysts, while with Cs exchanged faujasites 99% citronellol is produced at 77% conversion [ 147]. This change in selectivity is attributed to the steric hindrance induced by the larger Cs § ions,.but the influence of the increasing base strength cannot be ruled out. As with other base catalyzed reactions the role of the catalyst in this example is also twofold, i.e., the basic oxygen helps to abstract a proton from the hydroxyl group of the alcohol, while the metal cation stabilizes the resulting alkoxy species and polarizes the carbonyl group of the aldehyde. If both molecules are adlineated, hydride transfer from the alkoxy group to the polarized aldehyde group takes place inverting the nature of the two reactants. The remaining steps are the reverse reactions of the activation.

Lix

. NaX OH

87%

CsX r92%

~ ]

OH

A